Broadening Electrochemical Horizons - Principles and Illustration of Voltammetric and Related Techniques

0
.B
School of Chemistry, Monash University, Victoria, Australia
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Dedication
To Tunde, Stephen, and Andrew
1998 I was privileged to be present at the University of Oxford as the
inshelwood Lecturer in the Physical and Theoretical Chemistry Laboratory
and as a Christensen Fellow at St. Catherine's College. Financial support for
the lectureship was kindly provided by ICI.
The Hinshelwood lecture series, named after Cyril Norman Hinshelwood
(1897-1967), 1956 Nobel Laureate in Chemistry (with N.N. Semenov), was
entitled 'Broadening Electrochemical Horizons'. The presentation provided an
opportunity to develop an integrated series of lectures aimed at illustrating the
principles and applications of voltammetric methodology to a variety of problems of a fundamental and applied significance. In view of the fact that the
audience for the lectures had a wide range of levels of expertise in electrochemistry, the standard chosen for the presentation was equivalent to that appropriate
for a senior level undergraduate or postgraduate course rather than for specialists
in the field of electrochemistry.
The highly active Oxford University electrochemical environment was
invaluable to the lecture presentation because examples of studies undertaken
in the author's research progamme at Monash University in Australia could
be supplemented by extensive reference to material kindly made available by
ofessors Richard Compton, Allen Hill and Fraser Armstrong and colleagues
their Oxford University research groups.
At the completion of the lecture series, it was suggested that a written version
of the material could be developed as a guide for researchers and educationalists who wish to learn how the principles of voltammetry can be systematically
ited to solve problems relevant to a wide area of scientific disciplines. The
version of the Hinshelwood lectures, written in the years 1999-2001, is
fore expressly aimed at achieving this goal. I hope that the readers of this
will find the account of studies presented of assistance in providing them
a systematic approach to unravelling the mysteries that frequently accomstudies of processes that take place at or near to an electrode surface during
urse of voltammetric experiments. It is, of course, just these 'curiosities'
that have fascinated this author throughout his scientific career.
nash University
A. M. B.
I would like to specifically acknowledge the contributions of the following
people and organizations:
The Physical and Theoretical Chemistry Laboratory, St. Catherine's
College, and ICI for their generosity in hosting the 1998 Hinshelwood
Lectureship.
The School of Chemistry at Monash University for generously allowing
me leave at the University of Oxford in 1998 in order to present the
Hinshelwood Lectures.
Professor Richard Compton for extensive editorial advice during both the
Hinshelwood Lectures and the preparation of this book.
Professors Richard Compton, Allen Hill, Fraser Armstrong, and members
of their research groups at Oxford University and my colleagues at Monash
University for their friendship and for generous provision of material referred
to in this book. At the time the Hinshelwood Lectures were presented, material provided by colleagues often was unpublished. However, during the
time the book version was developed, much of this work has now appeared
in the open literature under the authorship of the relevant colleagues so that
due recognition can now be given via citation of these recently published
articles.
Glenda Oliver who typed the text without ever complaining, despite being
confronted with numerous drafts and then amendments to what she had
been advised was the final version.
My wife, Tunde, and children, Stephen and Andrew, who displayed significant patience and understanding while the book was written. I have
dedicated this book to them as an expression of my appreciation of their
tolerance in allowing me to spend countless hours hidden away in my study
during evenings and at weekends during times which otherwise would have
been devoted to fimily activities.
The fundamentals of electrochemistry
1 Introductory remarks
Redox reactions, electrochemical cells, and standard potentials
2
3 Thermodynamics versus kinetics
Calculation of reaction volumes and entropies from the
4
dependence of the reversible potential on temperature and pressure
5 Voltammetry and kinetics
Application of the principles of electrochemistry to fuel cells,
6
photovoltaic cells, and the lead-acid battery
6.1 Fuel cells
6.2 Dye-sensitized photovoltaic cells
6.3 Lead-acid battery
References
rinciples of voltammetry, electrolysis, spectroelectrochemistry,
and other techniques employed in studies involving solution
ase and surface-based electrode processes
An overview
The electrochemical cell used for voltammetric experiments
The electrodes used in voltammetric experiments
3.1 Working electrodes
3.2 Reference electrodes
3.3 Counter /auxiliary electrodes
The two major classes of voltammetry
4.1 Transient voltammetry
4.2 Steady-state voltammetry
Evaluation of electrode reaction mechanisms
Factors contributing to the nature of the current-potential curve
obtained in voltammetric experiments
6.1 Faradaic and non-Faradaic currents
Understanding the basic features of an electrode process
when the redox active species are soluble in the solution phase
7.1 Mass transport
7.2 Electron transfer
7.3 Homogeneous chemical kinetics
7.4 Electrochemical and chemical reversibility
Cyclic voltammetry under transient conditions when the
redox active species are soluble in the solution phase
8.1 Theory of cyclic voltammetry
xii
Contents
9
10
11
12
13
14
15
36
17
18
Hydrodynamic voltammetry
9.1 Rotating-disc electrode voltammetry
9.2 Channel electrodes
9.3 Wall-jet electrodes
9.4 A survey of the use of the theory of hydrodynamic
voltammetly
Voltammetric studies at microelectrodes when the redox active
species are soluble in the solution phase
10.1 Principles of the theory of microelectrode voltammetry
Semi-integration and semi-differentiation
(convolution voltammetry)
11.1 Some valuable properties of the semi-integral
11.2 Measurement of uncompensated resistance by
semi-integration
11.3 Semi-differentiation
General features associated with the modelling of
voltammetric experiments
12.1 Information required to solve voltanimetric theory
12.2 Methods used for solving voltammetric theory
A sunimary of the theoretical principles of voltammetry
13.1 Application of Faraday's law
13.2 A general approach to understanding a voltamnletric problem
Comparison of voltammetric techniques when the redox active
species are soluble in the solution phase
14.1 A quantitative comparison of the kinetic discrimination of
homogeneous reactions at common electrode geometries under
voltammetric steady-state conditions
14.2 A comparison of the homogeneous kinetic discrimination of
steady-state and transient experiments
Bulk electrolysis
15.1 Theory of bulk electrolysis
15.2 Cells for bulk electrolysis
Spectroelectrochemistry
16.1 ESR spectroelectrochen~istry
16.2 IR spectroelectrochemistry
16.3 UV-visible spectroelectrochemistry
16.4 NMR spectroelectrochemistry
16.5 Combining mass spectrometry and electrochemistry
Voltammetry at variable pressure and temperature
Voltammetric studies on solids attached to electrode surfaces in the
form of thin films
18.1 General aspects
18.2 Electron transfer in ideal redox active thin films attached
to electrode surfaces
68
69
71
74
75
78
79
83
84
85
87
88
88
90
91
91
92
95
96
98
100
102
104
107
108
113
116
117
117
123
128
128
131
Contents
...
xi11
18.3 Chemical reactions coupled to ideal thin-film
electron-transfer process
18.4 Nuances associated with adsorption
19 Techniques for obtaining molecular level
information on reactions associated with the
voltammetry of surface-attached species
19.1 The use of scanning probe microscopies in electrochemistry
19.2 The electrochemical quartz crystal microbalance
References
llustrating the basics of voltammetry for solution-soluble redox
active species involving reversible electron transfer and reversible
coupled chemical reactions: the reduction of electrochemically
rich polyoxometalate compounds
Introduction
Structural features of polyoxometalates
Coupled electron- and proton-transfer reactions associated with
~ ~ ] ~ the reduction of a- [ P ~ w ~ ~ oand
comparison of simulated and experimental cyclic voltammograms
obtained in aqueous media as a function of pH
3.1 Reduction of ~ - [ P ~ w ~ ~ o ~ ~ ] ~ 3.2 Reduction o f a - [ ~ ~ ~ ~ ~ 0 ~ ~ ] ~ 3.3 Discussion of results obtained from the simulation of the
reduction of a - [ ~ 2 ~ 1 ~ 0and
~ ~
a ]- ~[ - ~ 2 ~ ~ ~ 0 ~ ~ ] ~ Studies of the electrochemical reduction of ~ - [ s ~ M o ~ ~ino ~ ~ ] ~ aprotic and protic media3
~ ~ ] ~acetonitrile
4.1 Voltammetry of C X - [ S ~ M O ~ ~inOaprotic
media
4.2 Spectroelectrochemistry in acetonitrile
4.3 Electrochemical synthesis of one- and two-electron reduced
forms of [ s ~ M ~ ~ ~ o ~ ~ ] ~ 4.4 A systematic approach to chemical synthesis of a two-electron
reduced form of [s2Mo180 6 ~ 1 ~ 4.5 Voltammetry of
in acidic (95/5)
acetonitrile/water media
using
O a~ ~ ] ~
4.6 Photoelectrochernical studies of [ s ~ M O ~ ~
hydrodynamic channel electrode
Use of voltammetric techniques to identifjr the products formed
when [ ~ ~ ~ o reacts
~ with
~ 0 Ph3P
~ and
~ ]Bu3P
~ in- (95/5)
CH3CN/H20
5.1 Reaction of [ s ~ M ~ ~ ~with
o ~Ph3P
~ ]in~(95/5)
CH3CN/H20
5.2 Reaction of [ S ~ M O
$30621 4- with Ph3P under irradiative
conditions
xiv
Contents
with
o "Bu3P
~ ~ ] in~ (95/5)
5.3 Reaction of [ s ~ M ~ ~ ~
CH3CN/H2 0 : an explanation of differences relative to
reaction with PPh3
6 An overview of results obtained by application of
voltammetric, simulation, and spectroelectrochemical techniques to
polyoxometalate reduction studies
6.1 Cyclic voltammetry
6.2 Rotated-disc electrode voltammetry
6.3 Channel-electrode voltammetry
6.4 Microdisc-electrode voltammetry
6.5 Spectroelectrochemistry
6.6 Bulk electrolysis
6.7 Combinations of techniques
References
4
Electrode processes that illustrate the influence of irreversible
homogeneous reactions and the competition between reactions
that occur in the solution phase and on the electrode surface:
fundamental studies, photovoltaic dye-sensitizers, stripping
voltammetry and glucose biosensors
1
Introduction
2 Elucidation of the homogeneous reaction pathways that accompany
the electrochemical oxidation of cis, m e r - ~ n( ~ ($
0 -dpm)
) ~ (q2-dpm)
Br(dpm = Ph2PCH2PPh2)in dichloromethane
2.1 Voltammetric studies in dichloromethane
2.2 Bulk electrolysis and spectroelectrochemical experiments
2.3 Simulation of the voltammetry
2.4 Conclusions derived from electrochemical studies on
cis, m e u - M n ( ~ 0(ql-dpm)
)~
(q2-dpm)~r
3 Electrochemical studies on the [v(co)~]-'~ process in
aqueous media
3.1 Voltammetric oxidation of [V(CO)6]- in
acetone solutions containing water
3.2 Voltammetric, EQCM, and chronocoulometric studies on
the oxidation of p ( C 0 ) 6 ] - in water
3.3 Conclusions derived from voltammetric studies on
[V(C0)6]- in aqueous media
4 Voltammetric studies on the oxidation of the highly surface-active
polypyridyl ruthenium photovoltaic sensitizer cis-R~(II)(dcbpy)~
(NCS)2(dcbpy = 2, 2'-bipyridine-4, 4'-dicarboxylic acid
4.1 Reference studies on model mass-transport-controlled
processes
4.2 Electrochemical studies on cis-R~(dcbpy)~(NCS)~
in acetone
in tetrahydrofuran,
4.3 Voltammetry of cis-R~(dcbpy)~(NCS)~
acetonitrile, and dimethylformamide
248
248
b
250
251
253
260
261
262
262
264
268
269
272
280
Contents
4.4
Conclusions related to the voltammetry of surface-active
cis-R~(dcbpy)~(NCS)~
Stripping voltammetry
5.1 Anodic stripping voltammetry with thin-film mercury
electrodes
5.2 Theoly for a reversible process
5.3 Comparison of experimental results and theory
5.4 Mechanism associated with the adsorptive stripping
voltammetry of cobalt (and nickel) dimethylglyoxime
complexes at mercury electrodes
Glucose biosensors
6.1 The ferrocene-based glucose sensor
6.2 Optimization of the performance of a solution-phase
electrochemical glucose biosensor
6.3 Fabrication of a glucose bioelectrochemical sensor employing
glucose oxidase immobilized onto an electrode surface
6.4 Glucose analysis of whole blood with a commercially available
glucose bioelectrochemical sensor
References
llustration of the principles of voltammetry at
solid-electrode-solvent (electrolyte) inteflaces when redox active
microparticles are adhered to an electrode sudace
Introduction
Strategies to detect factors that may be important in the voltammetry
of redox active microparticles adhered to an electrode surface
Mechanistic aspects of the electron and ion-transport processes across
the electrode-solid-solvent (electrolyte)interface when arrays of
non-conducting microparticles are attached to an electrode
3.1 The oxidation of decamethylferrocene
3.2 Electrochemistry of microparticles of trans-Cr(CO)2(dpe)2,
trans-[Cr (CO), (dpe)2][XI salts, and ~ i s - C r ( C 0(dpe)2,
)~
(dpe = bidentate Ph2PCH2CH2PPh2,X - = anion) attached to
an electrode surface
3.3 Overview of factors that influence the voltammetry of
decarnethylferrocene and trans-Cr (CO)2(dpe)2 attached to an
electrode surface
3.4 Problems with a theoretical description of the
voltammetry of non-conducting microcrystals
Voltammetry of T C N Q adhered to an electrode surface: detection of
solid-state transformation, redistribution, and dissolution processes by
application of X-ray diffraction, electron scanning microscopy,
atomic force microscopy and electron spin resonance techniques
xvi
Contents
4.1 Solution-phase voltammetry of T C N Q
4.2 Voltammetric studies on microcrystals of T C N Q adhered to
, ~ b ' , and CS+
electrode surfaces in contact with ~ a +K+,
containing electrolytes
4.3 Electrochemically driven transformation of microcrystalline
T C N Q to tetraalkylammonium [TCNQ-] salts
4.4 Dissolution of solid T C N Q and [TCNQ-] salts
from electrode surfaces
4.5 Comparison of electrochemical data with microcrystals and
other forms of surface-confined T C N Q
4.6 Conclusions related to the electrochemistry of T C N Q
adhered to electrode surfaces
Voltammetric studies on systems where coupled electron and
5
ion transport within an adhered microparticle are rate determining
6
Voltammetric studies on adhered microparticles where
'thin-film' behaviour is exhibited
7 An overview of the techniques used in electrochemical studies of
microparticles adhered to electrode surfaces
References
6
Use o f metalloprotein voltammetry to illustrate the nuances o f
electrochemistry related to blocked electrodes, chemically
modified electrodes, electrode functionality, and microscopic
aspects o f electrode behaviour
1
Introduction
2 Structural features of metalloproteins that may give rise to
features that are different to those encountered in the
voltammetry of small molecules
3
Studies on protein-surface attachment to a
gold electrode by in situ scanning probe microscopy
The influence of surface attachment of
4
metalloproteins on voltammetric studies
4.1 General features of voltammetry of metalloproteins at
bare (unmodified) gold electrodes
4.2 The transient nature of the voltammetry of cytochrome c at
'bare' gold electrodes: an explanation based on a
self-blocking mechanism
Voltammetry of metalloproteins at chemically modified
5
gold electrodes
Voltammetry of metalloproteins at naturally and deliberately
6
functionalized carbon electrodes
6.1 Cytochrome c
6.2 Plasto cyanin
6.3 Ferredoxin
Contents
6.4
General conclusions concerning the voltammetry of
metalloproteins at carbon electrodes
Quantitative use of a microscopic model to explain the unusual
features of metalloprotein voltarnmetry at carbon electrodes
7.1 Cytochrome c voltammetry at carbon macrodisc electrodes
7.2 Cytochrome c voltammetry at carbon microdisc electrodes
7.3 Conclusions derived from modelling the voltarnmetry of
cytochrome c at carbon electrodes
Evidence that chemical modification of the
electrode surface can alter the reversible potential
8.1 The thermodynamic effects of chemical modification of
graphite electrodes on rubredoxin electrochemistry
8.2 Thermodynamic effects of chemical modification of
graphite electrodes on ferredoxin electrochemistry
8.3 Conclusions concerning the dependence of the reversible
potential on the presence of a surface modifier
Long-range electron-transfer effects encountered in cytochrome c
voltammetry at long-chain alkane thiolate modified electrodes
Voltammetry of metalloproteins in surfactant environments
Conclusions related to the voltammetry of metalloproteins
References
d
ndex
mistry is regarded as a mature scientific discipline, having a
ed two-hundred-year-old history (Table 1.1). Electrochemistry
iants such as Faraday and Nernst are well known to all students of chemistry.
urthermore, important applied devices based on electrochemical technology
ave been in widespread use by the general public throughout the twentieth
century. An obvious example of an electrochemical 'product' that has been
well known for generations is the lead-acid battery used in motor vehicles
(Section 6.3). Examples of other 'products' frequently encountered in everyday
life and derived from electrochemical technology include: fuel cells (Section 6.1)
and photovoltaic cells (Section 6.2 and Chapter 5); objects made from metals
such as aluminium, copper (Section 5) or zinc which are produced by electrorefining techniques; a hand-held sensor for monitoring glucose in diabetes
atients (Chapter 4); pacemakers and bionic ears. Corrosion also represents an
important example of an electrochemical process that has a significant impact
on modern society.
Given the industrial significance of electrochemical technology for the last
two centuries and the widespread teaching of the basics of electrochemistry in
undergraduate university and senior school science courses, it could be logically
assumed that most chemists have an excellent understanding of the theoretical and experimental aspects of the subject. However, informal surveys of
student attitudes by this author and other colleagues who teach the subject,
indicate that electrochemistry is regarded as a subject which is very difficult to
master from a theoretical point of view, and that experimentally it is seen to
be a 'fiddly' technique. With respect to the theory, it seems that the sign conventions associated with electrode potentials have caused misery for generations
of students and, from the experimental point of view, cries of 'my electrode is
blocked/broken', frequently reverberate around school and university student
laboratories.
During the course of presenting the Hinshelwood Lectures at The University
of Oxford in 1998, the author had the opportunity of assessing the status
of electrochemistry at that institution. The exchange of views between an
examiner and a science candidate during the course of a natural science viva
Table 1.1 A chronology of some important events in the history of electrochemistry over the period 1791-1960"
L. Galvani (1737-98) published his results on electrical experiments with
frogs (De Bononiensi Scientiarum et Artium Institute atque Academia
Commentarii VII (1791) 363).
A. Volta (1745-1827) communicated the construction of his electrochemical
pile (voltaic battery) in a letter to Sir Joseph Banks, published in Philos.
Trans. 90 (1800) 403.
W. Nicholson (1753-1815) and A. Carlisle (1768-1840) used Volta's pile to
demonstrate electrolytic conduction (electrolysis) (Nicholson's Journal 4
(1800) 179).
W. Cruikshanks (1745-1800) published the first qualitative analysis (copper)
performed with the aid of electrolysis (Nicholson's Journal 4 (1800) 187).
H. Davy (1778-1829) published his theory of electrolysis (Philos. Trans. 97
(1807) 1).
J.J. Berzelius (1779-1848) published his electrochemical theory and
developed the concept of the electrochemical series (Journalfur Chemie und
Physik 6 (1812) 119, Essai sur la Thkorie des Proportions Chimiques et sur
l'ln$uence Chimique de ~'~lectn'citk,
Paris, 1819).
G.S. Ohm (1787-1854) published his law (Schwei~er'sJournal 46 (1826) 137).
M. Faraday (179 1-1 867) published numerous observations based on data
obtained from electrochemical experiments, including the famous law, and
introduced modern electrochemical nomenclature (e.g. ion, anion, cation,
electrolyte, electrode) (Philos. Trans. 1832-34).
W.R. Grove (1811-96) described the first practical fuel cell (Philos. Mag. 14
(1839) 127).
E. Becquerel (1820-90) reported the photovoltaic effect (Compt. Rend. Acad.
Sci. (Paris) 9 (1839) 561).
R.L.G. Plant6 (1834-89) demonstrated the operation of the lead-acid battery
(Compt. Rend. Acad. Sci. (Pavis) 50 (1860) 640).
W. Gibbs (1822-1 908) published the first electrogravimetric analysis
(Fresenius' 2.Anal. Chem. 3 (1864) 334).
W. Nernst (1864-1941) published his fundamental equation which relates the
potential to ion activities in his dissertation (Die elektromotorische
Wirksamkeit der Jonen, Leipxig, 1889).
F.G. Cottrell (1877-1948) published the equation which is now known as
the Cottrell Equation (Z. phys. Chem. 42 (1903) 385).
B. Ku?era (1874-1 92 1) introduced the dropping mercury electrode for
electrocapillary studies (Ann. Phys. 11 (1903) 529).
J. Heyrovsk9 (1890-1967) published his first results with the dropping
mercury electrode method using a technique which he soon calls
polarography (Chem. Listy 16 (1922) 256, Philos. Mag. J. Sci. 45 (1923)
303).
D. Ilkovic (1907-80) published the equation which is now known as the
Ilkovi? Equation (Coll. Czech. Chem. Commun. 6 (1934) 498).
J.E.B. Randles (1912-98) (Trans. Faraday Soc. 44 (1948) 327) and A. Sevzik
(Coll. Czech. Chem. Comm. 13 (1948) 349) independently described the
technique of cyclic voltammetry.
G.C. Barker (1915-2000) and I.L. Jenkins introduced square-wave
techniques (Analyst 77 (1952) 685).
G.C. ~ a r k e and
r A.W. Gardner introduced pulse techniques (Fresenius' Z.
Anal. Chem. 173 (1960) 79).
"Adapted from information provided by courtesy of F. Scholz, C. Russell, and The Royal Society
of Chemistry.
Introductory remarks
(oral examination) late in the nineteenth cent1~ r is
y reproduced below:
Natural Science viva, c. 1890
Examiner:
Candidate:
What is Electricity?
Oh, Sir, I'm sure I have learnt what it is-I'm sure I did know-but
I've forgotten.
Examiner: How very unfortunate. Only two persons have ever known what Electricity is, the Author of Nature and yourself. Now one of the two has
forgotten.
Source: Falconer Madan, Oxfovd Outside the Guide-Books, 1923.
Quoted in: The Oxford Book of Oxfovd, Jan Morns, OUP, 1978.
Undoubtedly, many equivalent exchanges have occurred during the course of
oral examinations of electrochemistry-based theses at the University of Oxford
throughout the course of the twentieth century. In 1991, Professor P.W. Atkins
of Lincoln College, University of Oxford, one of the most distinguished chemical educators of our times, summarized his views on the subject by stating [I],
I think we ought to expand our view of what electrochemistry is about. It is not just about
electrode potentials and electrolysis, it is half of modern inorganic chemistry.
'k
and also,
Electrochemistry-in the broadest sense-will
be one of the great achievements in
chemistry in the next millennium, and one should prepare our people for it.
Thus, whilst electrochemistry may be seen as a 'tricky' discipline to master,
nevertheless according to Atkins, it will be of considerable significance in this
twenty-first century. Concurring with Atkins' thesis that education in electrochemistry will be even more important in the future than in the past, the author
titled the 1998 Hinshelwood Lectures, Broadening Electrochemical Horizons. The
aim of the Hinshelwood Lectures, and this book, which was generated from
the lecture series, is to demonstrate via presentation of a systematic account of
the subject, that the present commonly accepted limitations and correspondingly
conservative image of the subject, are unwarranted. As in many technologydriven subjects, there have been, in fact, numerous innovations achieved in
recent times which now make it possible to apply electrochemical methodology to solving problems in almost all branches of science. Developments in
electrochemistry in the twenty-first century may well be restricted more by
limited imagination, than by fundamental scientific impediments.
Intrinsic to many of the themes associated with the 'broadening of electrochemical horizons' concept, are recent innovations that have become possible via advances in materials science. Electrode sizes used, now range from the
nanometre (molecular dimensions) to the metre (electrorefining)size scales. The
materials from which electrodes are now constructed, and the electrode configurations and media in which they can be used, are now very wide-ranging. Up
4
Thefundamentals of electrochemistry
until the middle of the twentieth century, electrodes used in both fundamental studies and in applied devices were usually made from highly conducting
metals such as platinum, lead, silver, gold or mercury, or graphitic forms of
carbon. Now electrodes may be manufactured also from semiconducting and
even poorly conducting materials, many forms of carbon, including glassy carbon (GC) (a high-temperature form used widely in analytical applications of
voltammetry) and doped forms of diamond, and from an infinite range of composite materials. Furthermore, the electrodes may be modified extensively to
achieve characteristics associated with molecular recognition by addition, for
example, of conducting polymer coatings or enzymes, to the electrode suriace.
Advances in membrane technology also now enable a wide range of chemical
separation and speciation features to be coupled with the well-established operational aspects of an electrode or electrochemical cell. Thus, the generation ofthe
'smart' electrode system is now emerging in the area of sensor technology. Naturally, advances in instrumentation (especially computer technology) are also
having a significant impact in almost all aspects of electrochemical experiments,
although in the last decade, the rapid expansion of new electrode materials and
innovations in electrochemical cell design, arguably have had an even greater
impact on the subject than instrumentation advances aided by computer technology. However, in reality, it is of course the combination of advances in
materials science, electronics, computing, mathematics, physics, chemistry, and
the biological sciences, that has enabled the marvellous electrochemical packages
to be constructed. These now allow electrochemical techniques to be applied
routinely in liquid, solid, and gas phases, under conditions where restrictions
associated with high resistance, capacitance, size, or the need to achieve charge
neutralization were thought to be major impediments to progress.
At the start of this millennium, it needs to be recognized that the situation has been reached, where almost any problem involving reduction or
oxidation of a chemical moiety may be addressed, or even exploited, by an
appropriate electrochemical technique. Of course, as is presently the case,
there may be superior spectroscopic or other techniques to solve a given problem, and electrodes will still foul up under many circumstances, so that an
electrochemical solution to a problem will always be employed only if the
right sets of circumstances apply. However, via both the Hinshelwood Lecture Series and publication of this book, the author wishes to convey the
message that recently the electrochemical horizons have been broadened so
significantly, that one may predict confidently that techniques of electrochemistry will continue to be highly significant in both the fundamental and applied
senses in the twenty-first century. Atkins, in the quotations cited earlier, may
have exaggerated slightly the level of brightness of the electrochemical future,
but the general thrust of his highly positive remarks, in all probability, will be
realized.
In presenting the Hinshelwood Broadening Electrochemical Horizons Lecture
Series on a broad-based subject with a very long history, the need to focus the
content on only one or two aspects of the subject was essential. At the same time,
the more general theme of Broadening Electrochemical Horizons needed to emerge
Introductory remavkr
5
as a general conclusion. T o achieve both these objectives, this Hinshelwood
lecturer chose to use the widely used electrochemical technique of voltammetry
to illustrate the theme of the chosen topic, using in the main, examples taken
from recent research undertaken in the author's or colleague's laboratories.
In general terms, voltammetric techniques are associated with the measurement and interpretation of current-potential-time (I-E-t) relationships. In the
sense that current flows in the course of an experiment of this kind, voltammetry must be defined as a dynamic rather than an equilibrium technique of
electrochemistry. Simple consideration ofthe operational principles and outputs
of batteries, photovoltaic cells, and the glucose monitor mentioned previously,
reveals that current must flow through the devices when they are being used
for their intended purposes. Consequently, a detailed understanding of voltammetric techniques and methodology enables the behaviour of the devices to be
understood at both the applied and fundamental levels. In contrast, the wellknown pH and ion-selective electrode electrochemical sensors operate under
equilibrium conditions (zero current flow). Usually, the principles of these
so-called potentiometric devices may be understood by using the well-known
ernst and related equilibrium or thermodynamic equations [2].
Of course, a dynamic voltammetric technique may operate under conditions
where the response is almost equivalent to the equilibrium response given by the
ernst or other thermodynamic equations. An equilibrium-type response can
occur in the special case where zero current flow conditions prevail or where
all electron-transfer and coupled chemical reactions are so fast, within the timescale of the measurement, that the voltammetric response is indistinguishable
from that calculated for the reversible or equilibrium situation. However, irrespective of whether or not a voltammetric device is operating under apparently
equilibrium conditions, a fundamental understanding of the theory and practice of voltammetry requires a detailed knowledge of both the heterogeneous
kinetics, which define the electron-transfer step taking place at the electrodesolution, electrode-redox active solid, or relevant multi-phase interface, and the
homogeneous or heterogeneous reactions that may be coupled to the electrontransfer step. Naturally, the interfacial region where these dynamic reactions
occur may contain both resistance and capacitance terms, which are also likely to
exert a profound influence on the I-E-t behaviour of the voltammetric response.
That is, the theoretical description of voltammetry represents a solution to an
inherently complex problem, because a wide range of kinetically controlled heterogeneous and homogeneous reactions need to be described in an interfacial
region of space where resistance and capacitance are likely to be present.
T o conclude these introductory remarks, it may be noted that, as is always
the case with any branch of science, there are both 'good news' and 'bad news'
aspects of the subject that need to be given when presenting an overview of a
subject. The 'good and bad news' features of electrochemistry are summarized
in Table 1.2.
6
Thefundamentals of electvochemistry
Table 1.2 Dynamic electrochemistry
The good news
Integral to many problems of fundamental and applied significance:
e batteries
e photovoltaic cells
e biologically important electron-transfer reactions
e glucose sensors
electrochemical synthesis of metals (Zn, Cu, Al)
The bad news
Without great care, years of research, and 'good luck' they break, become
Electrodes:
fouled, commonly stop working
Resistance is always present and has a deleterious effect on an
Ohm's law:
electrochemical process when current flows
Capacitance: Associated with undesirable background currents or time constants when
time-dependent phenomena are operative
Theory:
Often considered complex because:
e historically, different sign conventions have been used in different
countries and at different times in the reporting of standard electrode
potentials and oxidation and reduction currentsa
e second-order differential equations need to be solvedb
e knowledge of surfaces and interfaces required
e knowledge of homogeneous and heterogeneous kinetics required, and
e knowledge of thermodynamics required.
'Use of International Union of Pure and Applied Chemistry rules is strongly recommended to
avoid this difficulty (e.g. (i) standard electrode potentials being reported as reduction reactions (see
Table 1.3);(ii) designation of oxidation currents as positive and reduction currents as negative) when
reporting voltarnrnetric data.
b~nalytical
solutions are rare, but numerical methods are now readily used (Chapter 2) to solve the
relevant equations.
Despite the fact that a kinetic theoretical description is required to explain
electrochemical devices or experiments involving current flow, the majority
of textbooks that focus on dynamic aspects of the subject still commence with
a description of thermodynamic relationships that exist when species are oxidized and reduced in an electrochemical cell (Fig. 1.1) containing two electrodes
referred to as either the negative or positive electrodes, or the anode (where
oxidation occurs) and the cathode (where reduction occurs).' The reason for
' ~ o t ethat when a cell is in an equilibrium state, it is inappropriate to apply the names
'anode' or 'cathode' to either electrode because no net chemistry is occurring. That is,
the cell is neither galvanic nor electrolytic and it is an equilibrium cell. Additionally,
the definition of which electrode is the anode and which is the cathode changes, for
example, when a lead-acid battery is charging or discharging. T o minimize these kinds of
R e d o x reactions, electvochemical cells
Galvanic cell
(b)
7
Electrolytic cell
Fig. 1.1 Schematic diagrams of typical (a) galvanic and (b) electrolytic cells.
introducing the subject of electrochemistry in this manner is that the fundamental thermodynamic relationships, relating chemical and electrochemical energy
(but not of course the mechanistic details), readily emerge from a relatively simple equilibrium treatment of the subject. Furthermore, and importantly from
a didactic viewpoint, electrochemical equilibrium relationships associated with
electrochemical cells can be readily understood, since only the potential difference between two half-cell reactions need to be considered under a range of
conditions, to illustrate the Nernst-type relationships associated with potential
difference, the nature of the redox couples in each half-cell, and the influence of
concentration (activity) of species present in the half-cell reactions. However,
most importantly, since these equilibrium-type principles must also represent a
limiting case of the dynamic response associated with generally non-equilibrium
voltammetric techniques, which are of paramount interest in this book, a brief
review of relationships between classical chemical redox reactions and reactions
that take place in a conventional electrochemical cell (Fig. 1.1) is also presented in this introductory chapter. Extensive details of these classical accounts
of equilibrium electrochemistry are available in references [3-51, while the,
inherently, far more complex kinetic description required to explain the details
of voltammetry is introduced in Chapter 2.
In order to qualitatively understand the nature of a process involving the
transfer of electrons, it is useful to consider the visually obvious features associated with a redox reaction such as the so-called 'copper nail' experiment
in which a nail, made of iron, is placed in a blue aqueous solution of copper(I1) sulphate. In this experiment, details of the course of the reaction may be
monitored conveniently by periodically removing the nail from the copper(I1)
sulphate solution and noting changes that have occurred, in both the nail and
the solution, as a function of time. An obvious visual conclusion reached with
respect to the initially pure iron nail is, that it has become progressively coated
with metallic copper, while close inspection of the solution phase would reveal
ambiguities, the terms 'anode' and 'cathode' are avoided wherever possible in this book.
Fortunately, the terms 'working' and 'reference' (or 'counter') electrodes can be introduced
in Chapter 2 onwards to avoid any confusion related to the function of the electrodes used
in a voltammetric experiment.
8
Thefundamentals of electuochemisty,
a concomitant decrease in intensity of the blue colour, associated with a lowering in the concentration of the blue copper(I1) ions in solution. Quantitative
analysis of the composition of the solution phase by atomic absorption spectrometry, would confirm that the copper concentration in the solution phase
has indeed decreased and, in fact, has been replaced by iron now dissolved in
the solution phase. Quantitative analysis of the chemical composition of the nail
would reveal that the percentage of metallic iron decreases as the percentage of
copper increases. That is, the overall reaction that occurs is2
Of specific interest to an electrochemist is the fact that the overall reaction
written in eqn (1. l ) may be treated as the summation of two reactions, each
involving an electron-transfer process because Fe(meta1) has been converted or
oxidized to Fe2+(solution),and c u 2 +(solution)converted or reduced to Cu(meta1).
Based on this kind of formalism, eqns (1.2a) and (1.2b) can be used to represent
the two half-reactions that on summation give eqn ( 1 . 2 ~which
)~
(as required)
is identical to eqn (1.I):
f ~e(so1id)--+ i Fe2+(solution)+ e-
$ ~u~+(solution)
+ e- ---t i solid)
(oxidation) (1.2a)
(reduction) (1.2b)
f Fe(so1id)+ f cu2+(solution)--+ $ Fe2+(solution) + i solid) (overall redox
reaction)
(1.2~)
It should be noted that representing electrode reactions as one-electron
processes as above, with respect to eqns (1.2a-c), simplifies calculations of thermodynamic parameters that may be derived from data obtained from electrochemical experiments. This practice is adopted therefore at all times in this book.
In an analogous manner, the chemically spontaneous reaction
+ 2~e~+(solution)
+ Iz(solution)+ 2ce3+(solution)
21- (solution)
(1-3)
21n the early stages of this Introductory Chapter, half-cell and overall reactions are written
(e.g. eqns (1.1)-(1.9)and Fig. 1.l) as irreversible processes (+) in order to highlight the
dominant direction of the reaction. In principle, thermodynamic treatments require that
all processes in a reaction scheme need to be considered as being reversible (
) rather
than irreversible (+). The concept of treating processes as being inherently reversible will
be adopted whenever quantitative calculations of electrochemicalreactions are undertaken.
R e d o x reactions, electrochemical cells
9
may be written in the formally equivalent electrochemical format shown in
eqns (1.4a) and (1.4b) to give the overall redox reaction (eqn 1 . 4 ~ ) :
I- (solution) --+- 1 Iz(solution)
+ e-
+
~e~+(s01ution)e- +ce3+(solution)
I- (solution)
(oxidation)
(1.4a)
(reduction)
(1.4b)
+ ~e~+(solution)
+f (solution) + ce3+(solution) (overall redox
12
reaction) ( 1 . 4 ~ )
- (solution)
+ 2ce4+(solution)+I2(solution) + 2ce3+(solution)
(1.4d)
At this stage it should be noted that a spontaneous reaction is one which occurs
without any input of energy, in contrast to the reverse reaction, which has to
e driven by input of energy, which may, of course, be electrochemical energy
redox reactions as two half-reactions enables the relationship
nt electrochemical half-cell reactions to be readily identified.
mica1 cell, the reduction reaction occurs at one electrode (the
electrode that accepts electrons) and the oxidation reaction at the other electrode
(the electrode that donates electrons), and an electrochemical cell may be seen
to consist of two half-cell reactions, each of which, effectively, may be said to
have its own potential or energy.
Conventional schematic forms of representation of electrochemical cells, with
examples of half-cell reactions, are shown in Fig. 1.1. Figure 1.1(a) represents
a galvanic cell and Fig. 1. l (b) an electrolytic cell. In the galvanic cell shown
in Fig. 1.1(a), the two half-cells are separated by a salt bridge or membrane
to prevent c u 2 + ions present in one half-cell coming into contact with the
metallic zinc which would lead to the occurrence of a spontaneous chemical
redox reaction. Thus, prior to connecting the two half-cells, a chemical form
of energy is stored, but may be made available for conversion into electrical
ergy. When the electrochemical cell is connected as in Fig. 1.1(a), current
w is associated with the occurrence of the two half-cell reactions
+
~n(so1id)+ f zn2+(solution) e;cu2+ (solution)
+ e- -+ fcu(so1id)
(1.5a)
(1.5b)
continues until the equilibrium position for the reaction
Zn(so1id)
+ c u 2 +(solution)
z n 2 +(solution)
+ Cu(so1id)
(1.6)
is reached. The magnitude of the current flowing in a galvanic cell at any given
time will be determined by the rates of numerous processes, which include the
eterogeneous and homogeneous reactions that occur in each half-cell, as well
10
Thefundamentals of electrochemistry
as the rate at which species present in the solution phase can be transported to
and from the electrodes by diffusion, migration or convection. Normally, at the
start of the experiment when the circuit is connected, the current will be at its
maximum value, since the forward reactions in eqns (1.5a) and (1.5b) will be
at their maximum values and the backward reaction rates will be negligible. In
contrast, at equilibrium, no net reaction occurs, since the forward and reverse
reaction rates are equal by definition, so that no net current flows. Clearly, a
relationship must exist between the magnitude of the current as a function of
time and rates of heterogeneous reactions at the electrode-solution interface,
and mass transport of species towards or away from electrodes. Voltammetry, the
subject ofmajor interest in this book, is all about these kinetic factors, because in
this technique it is the application of a suitable potential which drives a specific
half-cell reaction at a finite rate and, in turn, it is the occurrence of the reaction
that gives rise to a kinetically controlled current flow which is a function of the
net rate of the numerous processes that occur in a half-cell. The other half-cell in
a voltammetric experiment is a reference electrode system, which is assumed to
operate under equilibrium conditions and hence independently of any current
that may flow through this part of the circuit (Chapter 2). Figure 1.l(b) is an
electrolytic cell in which energy from a battery or a current source must be
supplied to drive the overall reaction
+
+ 2Cu(solid)
2 ~ ~ ~ + ( s o l u t i o n2H20(liquid)
)
+ O2(gas) + 4H+(solution) (I.7)
since the reverse direction of this reaction
is the spontaneously favoured one. It should be noted that a salt bridge to
separate the two half-cells is not required in this particular cell configuration
because the product at the copper electrode (metallic copper) adheres to the
electrode and the product at the platinum electrode (oxygen) bubbles out of
the solution, so that direct contact of the products which would give rise to
a spontaneous reaction (eqn 1.8) does not occur. The overall reaction given
in eqn (1.7 or 1.9c,d) may be written as the two half-cell reactions shown in
eqns (1.9a) and (1.9b).
f c u 2 +(solution) + e- + fcu(solid)
Cu
Pt
1H20(1iquid)
-+
2
+
0 2 (gas)
+ H+(solution)+ e-
2~u~+(solution)2H20(liquid)+ 2Cu(solid)
(1.9a)
(1.9b)
+ O 2(gas) + 4~+(solution)
(1.9d)
R e d o x reactions, electrochemical cells
11
While current flows, an electrochemical cell, by definition, is a kinetically
controlled device. In the chemical sense, the equilibrium position of the overall
reaction may be calculated from the difference in Gibbs free energy of products
and reactants, which is given the symbol A GO, when the reaction is carried out
under standard conditions [2-51. In the electrochemical sense, the equilibrium
osition (no net current flow) of the overall reaction may be calculated from the
difference in potentials of the two half-cell reactions. Therefore, under standard
conditions of X ° C , and with unit activity of all species involved in the reaction,
it follows that [2-51:
where K is the equilibrium constant for the overall reaction, R is the universal
gas constant, F is the Faraday constant, T is the temperature (in Kelvin), AE0
is the difference in the standard potentials of the two half-cell reactions and the
number of electrons associated with each charge-transfer step is unity.3
Table 1.3 gives a list of standard potentials (E0 values) for selected half-cell
reactions which, for both convenience and ease of calculation of the thermodynamics, are all written as reversible one-electron processes. By convention, the
value of the standard hydrogen electrode (SHE) process solution) e2(gas))is given a value of exactly zero.
It is of historical interest to note the relationship of these quantitative data to
e qualitative development of the Electrochemical Series (Table 1.4) developed
by Berzelius [6] (Table 1.1) almost 200 years ago. This form of the series was
based on qualitative observations made during the course of electrolysis experiments on numerous metal salts and compounds. Generally speaking, standard
tentials of half-cell reactions involving elements towards the top of Berzelius7
ctrochemical Series have large positive values (e.g. oxygen, 1.23 V), while
those towards the bottom have large negative values (e.g. potassium, -2.93 V).
The historical background to the development of Berzelius7 Electrochemical
eries is intriguingly related, in a recent publication by Russell [6], to the work
of the pioneers of electrochemistry (see Table 1.1).
At the beginning of the nineteenth century, the results of two enormously
significant discoveries were published. The possibility of a continuous electric
current was demonstrated by the Italian physicist Volta, who took alternating
discs of copper and zinc and sandwiched them between damp cardboard to
create the first wet battery. Within months, Nicholson and ~arllsle,in London,
found that decomposition of water occurred, when in contact with both poles
of what, in those days, was termed a pile [7]. They had discovered what Michael
araday later called 'electrolysis'. Shortly afterwards, Humphrey Davy used this
technique to decompose a number of solutions and fused materials, and by 1807
he had isolated sodium and potassium for the first time [6].
+ +
+
3~alf-cellreactions and equations derived from them are written from now on as reversible
) to signifj. that they have thermodynamic significance even when the
processes (
reaction for either the forward or backward direction may be heavily favoured. Also see
footnote 2 in this chapter.
+
12
Thefundamentals of electrochemistry
R e d o x reactions, elect~ochemicalcells
13
Table 1.4 Berzelius' Electrochemical Seriesa
O ~gen
Y
Sulfur
Nitrogen
Muriatic radical
Fluoric radical
Phosphorus
Selenium
Arsenic
Molybdenum
Chromium
Tungsten
Boron
Carbon
Antimony
Tellurium
Tantalum
Titanium
Silicon
Osmium
Hydrogen
Gold
Iridium
Rhodium
Platinum
Palladium
(
+ 1.23V)
(0.00 V)
( + I S 2 V)
(+1.19V)
Mercury
Silver
Copper
Nickel
Cobalt
Bismuth
Tin
Zirconium
Lead
Cerium
Uranium
Iron
Cadmium
Zinc
Manganese
Aluminium
Yttrium
Beryllium
Magnesium
Calcium
Strontium
Barium
Sodium
Potassium
'Adapted from reference [6]; order is from oxygen at top to potassium at bottom of series. Potentials in parenthesis are standard
potentials for half-cell reactions selected from the present day Electrochemical Series (Table 1.2 and reference [2]). With the exception
of oxygen and hydrogen, the half-cell reactions selected involve
reduction of the metal-ion (highest oxidation states commonly
found) to the metal.
erzelius, after learning of Davy's results, undertook new and enlightening
experiments with Volta's pile. In 1808 he discovered the use of liquid mercury
as an electrode material. Davy, in turn, promptly isolated magnesium, calcium,
strontium, and barium by using mercury as an electrode and distilling off the
from the resultant amalgams [6].
ver valuable these discoveries of elements may have been, they were
n importance by the theoretical concepts that subsequently emerged.
avy deduced that electrolysis involved the reversal of chemical combination. It
dawned on him that what kept elements combined in salts and other electrolytes
was electrical attraction, and that electrolysis neutralized the positive charge on
the metal and the negative charge on the other part of the electrolyte. He called
this an electrochemical theory [6].
14
Thefundamentals of electvochemistvy
Berzelius pursued the concepts of Davy much further and, in so doing, generated a philosophy of chemistry. For Berzelius, every compound was polar
and held together by electrostatic attraction. He then proceeded to develop the
concept that what determined an element's polarity was its position in a vast
electrochemical series (Table 1.4) ranging from the most electronegative of elements (oxygen) to the most electropositive (potassium). Thus, a very positive
metal would possess a larger polarity ('charge') than one further up the series.
There was much more to his theory than this, but essentially he was proposing
that inorganic compounds were held together by electrochemical forces and
could be pulled apart by electrolysis. The great strength of his electrochemical
theory became apparent when he combined his concepts with those of Dalton's
atomism in his essay on 'The Theory of Chemical Proportions and the Chemical Influence of ~lectricity',which was published in several languages over the
period 1814-19 [6].
Detailed treatment of electrochemical half-cells, and the relevant thermodynamics that apply to equilibrium reactions derived by combining two half-ceh,
are available in very readable form in the book by Compton and Sanders [2].
Obviously most textbooks on physical chemistry also treat this subject in considerable detail. In essence, electrochemical half-cell reactions are expressed by
convention as reversible reduction processes, the simplest being
Ox
+ e- 6 Red
(1.11)
where Ox is the oxidized form of a species, Red is the reduced form and
the number of electrons transferred in the charge-transfer step is assumed to
be unity again. Solely for reasons of convenience, the charges on Ox and
Red in eqn (1.11) have been omitted, although of course it must always be
remembered that charge balance is an important prerequisite in electrochemistry. The equation relating the equilibrium potential to the species involved
in an electron-transfer reaction was first derived by Nernst for a metal cationmetal electrode reaction of the kind given in eqns (1.5a) and (1.5b). The more
general form of the equation is now known as the Nernst equation, which for
the reaction given in eqn (1.1I), is:
where E is the potential of the electrode, EO is the standard electrode potential
(measured relative to the SHE), R (8.31451J K-' mol-') is the universal gas
constant, T is the temperature (in Kelvin), F (96484.6 C mol-') is the Faraday
constant, and ai is the activity of species i. The activity for species dissolved in
solution is commonly related to concentration by the expression
where yi is the activity coefficient of species i and [i] is the concentration of
species i, whilst for a metal or other pure solid substance it is unity.
R e d o x reactions, electrochemical cells
15
The tendency for the reaction to occur relative to the SHE is given by
hen concentrations are used instead of activities, as is normally the case, the
ernst equation becomes
ere EO
f is the formal potential4 (measured relative to the SHE) whose value
now depends on the activity coefficient and hence on the medium.
Calculations based on standard potentials, such as those in Table 1.3, and
s (1.10)-(1.15), are heavily emphasized in physical chemistry text books.
wever, such calculations enable only a very limited range of practical electrochemical problems to be addressed because most reactions do not occur
under standard conditions and they are kinetically rather than thermodynamically controlled. However, several features of some of the reactions considered
in eqns (1.1)-(1.9) may be understood via equilibrium-type calculations.
Knowledge of the standard electrode potentials of the half-cell reactions
ows the thermodynamically favoured direction of the cell reaction to be estab~ ~all~ cells when written in
lished. By convention, the standard potential, E : for
standard notation (see Figs 1.1 and 1.2) is calculated as:
Therefore, from eqn (1.10) it follows that
Thus, if A E0 is greater than zero, K will be greater than one, which means
that the forward direction of the overall reaction, when written as the summation of the cell reaction at the right-hand electrode (a reduction process) and the
cell reaction at the left-hand electrode (an oxidation process) is the spontaneous
ection of the cell reaction. In contrast, if AE' is negative, then K will be less
an unity for the cell reaction, so that the reverse reaction will be the spontaneous one, or energy will be required to drive the reaction in the forward
irection of the cell reaction. Thus, the standard electrode potentials of the
metal/metal-ion half-cells considered in eqns (1.5) ( c u 2 + / c u and z n 2 + / z n )
'Also known as the conditional potential.
16
The fundamentals of electrochemistry
(a)
7 7
1
EXTERNAL
CIRCUIT
1
I
Electron
Q
flow
( current
Z n 1 z n 2 +(aq) (aZ,z+) 11 c u 2 + (aq) (aC,2+) 1 C U @
flow
7 7
1
EXTERNAL
CIRCUIT
(b)
Current
I Cu 1 Cu2+(aq) (ac,,z+) 11 zn2+ (aq) (azn2+)1 Z n I Electron
@
Q
Fig. 1.2 Standard form of representation of an electrochemical cell. Irrespective of whether the
cell is written as in form (a) or (b), via convention E:,,~ = E,'& The value (sign) of
AGO = - F E : ~ ~=~ -RT In K calculated using this convention is used to determine the direction
-Eft.
of the spontaneous reaction and hence the direction of current flow in a galvanic cell. See text for
further details.
may be treated as follows:
The EO data provided for the Mn+/M couples in Table 1.3 gives the
thermodynamics of the cell reaction
Pt (solid) I Hz(gas)(P = 1 atm) ( solution) (a = 1)(1 solution) (a = 1) I M(so1id)
(1.18)
under standard conditions. Therefore the reversible chemical reaction for this
cell is
Thus, in the case of copper, the standard electrode potential is 0.34V for the
reversible reaction
Scu2'
(solution)
+- H~(gas) + S solid) + H+(solution)
and the standard free energy is
(1.20)
Redox reactions, electrochernical cells
17
(Table 1.3) gives an EL^^ value of -0.76 V for the reaction
+ iH2( p s )
i~n(so1id)+ H+(solution)
1zn2+(solution)
2
(1.22)
so that
A Go = 0.76 FJ mol-'
or AGO = -0.76 F J mol-' for the reverse reaction, which means that the
of metallic Zn with aqueous acid (H', a = 1) is strongly favoured in
the thermodynamic sense.
The data obtained in eqns (1.21) and (1.23) enable attention to be given to
the galvanic cell presented in Fig. 1.1. The overall reaction
$ ~n(so1id)+ 2 c u 2 +(solution) & zn2+(solution) +
K'
solid)
(1.24)
is obtained by subtraction of eqn (1.22) from eqn (1.20). Thus, for the reaction
in eqn (1.24)
which implies that this reaction is, as expected, strongly favoured in the forward
irection in the thermodynamic sense. In fact the equilibrium constant K' for
the cell reaction given in eqn (1.24), after noting that the activities of Cu(so1id)
and Zn(so1id) are unity by convention, may be calculated as follows:
K' =
(aZn2+)
'I2
= exp 1.1OFIRT = x
(acu") lI2
4 x 1018 (at 25°C)
This means that for the reaction written in the more usual form as
the equilibrium constant is given by
quivalent use of relevant data in Table 1.3 enables confirmation to be made
that reactions given in eqns ( 1 . 2 ~ (1.4c),
)~
and ( 1 . 9 ~ are
) spontaneous in the
thermodynamic sense, and their equilibrium constants can be calculated.
In principle, C~+(solution)could be included in the cell reaction given in
rather than ~e~+(solution)
could have been the prodig. 1.1 and ~e~+(solution)
uct of the 'copper nail' experiment. Why then have the reactions given not
included these metal-ion species in these different oxidation states? It is common chemical knowledge that both monovalent Cu(1) and divalent Cu(I1) ions,
18
Thefundamentals of electvochemistry
solution) and cu2+(solution)respectively, may exist under some conditions
in aqueous solution. The disproportionation reaction
may be considered in terms of the two half-cell reactions
and
solution) + e 6 ~u+(solution)
(1.29)
which, according to Table 1.3, have EO values of 0.52 and 0.16V respectively.
Thus, for the reaction
CU'
(solution)
+ $ H2(gas) Tf Cu(so1id)+ H+(solution)
AGO
(1.30)
= -0.52F J mol-'
and for the reaction
+ 1 Hz(gas) +
cu2+(solution)
AGO
CU+ (solution)
+ H+(solution)
(1-31)
= -0.16FJ mol-'
The two reactions (eqns 1.30 and 1.31) may be subtracted to give the disproportionation reaction in eqn (1.28), for which
The equilibrium constant for this disproportionation reaction is calculated as
K=-
acu2+
= exp 0.36FIRT = 1.2 x lo6 at 25°C
(~cU+)~
(1.33)
Thus, thermodynamically the disproportionation reaction is highly favoured
at 25°C. Indeed, the kinetics of the reaction are also very fast, so that uncomplexed copper(1) ions, rapidly formed when a copper(1) salt is dissolved in water,
disproportionate to give metallic copper and copper(I1) ions. In fact, it is only
when Cu(1) ions are stabilized by strong complexation, for example by chloride,
that the copper(1) oxidation state is thermodynamically stable in water.
Thermodynamics versus kinetics
19
general, it may be shown [2] for the disproportionation reaction
(a
+ b)MX+(solution)+ a~(~+~)+(solution)
+ bM("-')+ (solution)
(1.34)
EL~+/M(~+~)+
(1.35)
that
-
which leads to the conclusion that A G O will be negative, and therefore the
disproportionation reaction favourable, if
In the case of Cu+(solution),
thus disproportionation of CU' (solution) is favourable. In contrast
so the reaction
is not favoured. Thus, the addition of metallic iron to a solution of ~e~+(solution)
to the formation of Fe2+(solution).These two pieces of information explain
in the 'copper nail' experiment, the iron-containing product formed
reaction of metallic iron nails is Fe2+(solution)and not Fe3+(solution), and,
why no Cu+(solution)species are included in the reaction, as written in
.
in Table 1.3 also may be used to confirm that the spontaneous
eqn ( 1 . 2 ~ )Data
reaction direction written in eqn ( 1 . 4 ~is) correct. The direction of eqn 1.8 as
+ 1
spontaneous reaction under standard conditions, which include a ~ =
-- 0), may also be confirmed via calculations based on data contained in
e 1.3. However, at other pH values, different reaction pathways to that in
(1.8) may be favoured [2], as the value of A G, and hence the equilibrium
ion for the range of possible reaction pathways, is strongly dependent on
a ~ (pH).
+
The dependence of the equilibrium position of a reaction on pH,
at least one of the half-cell reactions contains a term for a ~ +is, discussed
ail in reference [2].
odynamics versus kinetics
n Section 2, the use of electrode potentials to predict the position of chemical
equilibria of redox reactions that may be written in terms of half-cell reactions
20
Thefundamentals of electrochemistry
has been illustrated. However, the predictions are subject to kinetic limitations [2]. That is, even if a reaction is thermodynamically feasible, the question
still arises as to whether the reaction proceeds at a reasonable rate. Consider the
hypothetical reaction
~ OH+
(complexation or precipitation).
which neglects interaction of M ~ with
Data in Table 1.3 advise that
which implies that for the reaction
:Nlg2+(solution)
+ H2(gas)
iMg(solid)
+ H+(solution)
the standard free energy change at 25°C is
Likewise, since (Table 1.3)
for the reaction
H 2 0(liquid) 4 OH- (solution)
+ H+(solution)
(1.45)
Thus, for the reaction of interest, when magnesium is in contact with water
(eqn 1.40)
This large negative A GO value implies, from a thermodynamic perspective,
that when magnesium metal is dipped in water, evolution of hydrogen gas
is expected to occur. However, in practice, little or no reaction is observed,
since a thin film of magnesium oxide, present on the metal surface, prevents
the reaction taking place at a significant rate. That is, the oxide layer passivates
the metal. Similar reasoning explains the well-known, and of course industrially
and commercially important, lack of reaction of 'thermodynamically reactive9
aluminium metal with water.
Reaction volumes and entropy calculations
21
ulation of reaction volumes and entropies from
bviously, most chemical redox and electrochemical cell reactions do not occur
under standard conditions of temperature (25°C) and pressure (1atm). Thus,
in most practical situations, the reversible potential dependence of AE,O on
these parameters needs to be established, and in fact may be used to calculate
the reaction volume (A V) and entropy (AS) of an electrochemical cell under
equilibrium conditions.
The use of the fundamental thermodynamic equations [8,9] enables the
required equilibrium relationships to be derived, which in turn enable A V and
A S to be calculated. Thermodynamics deals with the most probable microstate
of a system. At equilibrium, the system must have a uniform temperature and
ressure, a constant number of phases and a constant chemical potential of these
hases. For a single-phase half-cell of a redox system in which the number of
moles is constant, the free energy (G) is given by the equation
n eqn (1.48), H is the enthalpy, U is the internal energy, T is the temperature,
S is the entropy, P is the pressure and V is the volume for a system with only a
single phase, as may be encountered in a half-cell reaction of an electrochemical
cell involving only solution-soluble species. Thus, the change in free energy
(a G) for a closed system is given by the equation
For more complex systems involving two or more phases, where chemical
species move across boundaries, as is the case in solid-state electrochemical
piAni and yA (in a differentiated form) need to be
reactions, the terms
added to the above relationships, where pi and ni are the chemical potential
and number of moles of species i respectively, y is the surface tension and A the
surface area. Reactions involving deviations from equilibrium are also more
complex.
Equation (1.49) provides access to two very important relationships that may
be exploited via electrochemical measurement of the reversible potential in
order to define completely many solution phase redox systems in a thermodynamic sense. These additional relationships involve the change in free energy
22
Thefundamentals of electrochemistry
with respect to temperature or pressure as given in eqns (1S O ) and (1.51)
which, respectively, define the entropy and the volume for a half-cell reaction.
It has been established previously that A GO = -FA EO at standard conditions or A G = -FAE; when other conditions prevail. Since the reaction
volume (the difference in molar volumes of reactants and products in a chemical process) or A V is related to the pressure dependence of the Gibb's free
energy (A G) through eqn (1.5I), it follows that this parameter may be determined from the dependence of the formal potential of the electrochemical cell
reaction (equivalent to the chemical redox reaction) on pressure, and use of the
relationship
Analogously, the reaction entropy (the difference in entropy of reactants and
products in a chemical redox process) may be determined from temperature
dependence of the formal potential and use of the relationship
It follows from the above discussion that the temperature and pressure dependence of the reversible potential of the reference electrode (second half-cell
reaction) normally needs to be known in order to calculate A V and AS for a cell
reaction. The thermodynamic treatment of multi-phase redox systems, as noted
above, is more complex, and requires the introduction of additional relationships
associated with the interface, some ofwhich are considered in Chapter 4. However, in this introductory chapter, it is necessary to emphasize that measurements
of the reversible potential of electrochemical cells lead to more information than
measurements of A GO (or A G) and equilibrium constants of chemical reactions, which are the thermodynamic parameters emphasized in the majority of
textbooks on Physical Chemistry.
It was emphasized in Section 1 that the thermodynamic treatment of chemical
and electrochemical cell reactions, strictly applies only to equilibrium conditions
when zero net current is flowing. Thus, none of the equations presented to date,
Voltammetry and kinetics
23
contain terms that can account for the (time-dependent) current which flows
under dynamic conditions, when reactions are occurring spontaneously, as in
a galvanic cell, or else when chemical reactions are being electrochemically
driven by superimposition of an externally applied voltage or current onto the
electrochemical cell. Consequently, in order to understand the fundamentals of
the voltammetric techniques where an electron-transfer reaction is driven by
application of potential and, indeed, how the devices based on electrochemical
technology such as batteries, photovoltaic cells or glucose monitors operate, a
knowledge of the time dependence of the processes which contribute to the
experimentally measurable I-E-t relationship is required.
The importance of the dynamics of the various reactions coupled to the
electron-transfer step, is revealed by considering the large-scale production of
metals such as Al, Cu, and Zn by electrorefining technology. The half-cell
reactions of direct importance that occur at the cathode in these economically
important metal deposition processes may be written simply as:
. 1.3 Electrodeposition of copper at large-sized electrodes (cathodes) used for electrorefining of
copper. Photograph provided by courtesy of Copper Refineries, Townsville, Queensland, Australia.
24
Thefundamentals of electvochemistvy
with, of course, another half-cell reaction occurring at the anode. In essence,
these metal deposition reduction processes obey Faraday's law where, for a
100 per cent efficient electrolysis reaction, the product of I and t (current x time),
assuming constant current for the duration of the electrolysis, represents the
number of coulombs, Q, required to generate Q / n F moles of metal which
is plated onto the cathode. However, to obtain a metal deposit of very high
purity in the form of a smooth copper sheet attached to a giant-sized electrode
(see Fig. 1.3), via a reaction approaching 100 per cent efficiency, has required
decades of intensive research. The purity and nature of the metal deposit formed
on the electrode suriace, and the efficiency of the process, are determined by
the composition and design of both the cathode and anode, the composition
and purity of the electrolyte, the magnitude of voltage and/or current applied
to the cell, the distribution of potential and current, the absolute value and
constancy of temperature of the electrolyte, the hydrodynamics of the entire
electrolytic plant, the elimination of undesirable side reactions and, dare one
say it, possibly even 'the phase of the moon'.
An intriguing insight into the difference between the macroscopic and microscopic worlds of electrochemistry, is gained by comparison of the results of the
deposition of copper onto a giant electrode of the kind shown in Fig. 1.3, and
onto arrays of very small carbon microdisc-electrodes (Fig. 1.4). The microdisc
array electrode contains 7-pm diameter carbon fibres randomly distributed
Fig. 1.4 Electrodeposition of copper at a random assembly of carbon-fibre microdisc electrodes
M RAM^^ electrode). Photograph provided by courtesy of Stephen Fletcher, CSIRO, Division of
Minerals, Clayton, Victoria, Australia.
Application ofprinciples ofelectvochemistry
25
within an epoxy resin. At this electrode, mass transport (see Chapter 2) is much
enhanced at the edges of each carbon-fibre microdisc electrode. Furthermore,
copper is only deposited after a nucleation and growth phase of the reaction
occurs, so that the copper deposited onto each very small electrode is now
subject to the random statistics associated with stochastic processes, to give the
range of fascinating crystalline forms of metallic copper observed at different
carbon-fibre electrodes, as shown in Fig. 1.4. The absence of a copper deposit
on a particular carbon fibre in Fig. 1.4 could be a statistically expected result
or simply indicate that this fibre is not electrically connected or has 'failed7
during the course of the experiment. Clearly, in the electrorefined form of
copper (Fig. 1.3), the remarkably smooth deposit represents the result of the
careful averaging of many events. This commercially important 'averaged' result
represents the triumph of the combined skills of materials scientists, physicists,
chemical engineers, chemists, and electrochemists, who have, over many years,
learned how to control the dynamics of the numerous events that occur during the course of electrodeposition of a metal onto an electrode surface. The
example of copper deposition into electrode surfaces therefore gives another
opportunity to highlight the fact that the standard thermodynamic treatment of
a metal deposition reaction given in most physical chemistry textbooks, reveals
no more than the average amount of energy required to deposit copper onto
an electrode surface at 25°C. In reality, even this most basic of electrochemical
metal deposition processes represents the result of a series of complex dynamic
reactions that occur at a finite rate over a finite period of time, and which
encompass a wide range of homogeneous and heterogeneous processes. Consequently, usually a kinetic rather than thermodynamic description of all the
events taking place, is required to achieve any significant level of understanding
of a metal deposition, or indeed any dynamic electrochemical process, as always
will be encountered in voltammetry.
n the very first paragraph of this book, it was noted how frequently 'products7
erived from electrochemical technology are used in modern society. T o conclude this introductory chapter, it is now appropriate to highlight very briefly
how the general principles of electrochemistry presented above, apply to the
operation of fuel cells, photovoltaic cells and the lead-acid battery. These, and
all other electrochemical devices, are characterized by the performances of the
electrodes and the electrolyte. Since thermodynamics have to be favourable
by definition, it is the kinetics that usually determines the commercial viability. Thus, most research in the area of electrochemical technology is aimed
at speeding up desirable processes that occur in the electrochemical cell and
slowing down the rates of unwanted side reactions. Importantly, voltammetric techniques considered in the remainder of this book play a powerful role
Thefundamentals ofelectrochemistry
26
in quantifying the rates of the individual half-cell reactions, and hence the
performance of the overall cell reaction.
6.1 Fuel cells
Almost immediately following Volta7sdiscovery of the battery, Nicholson and
Carlisle used this new device to 'split' water into oxygen and hydrogen gases
(Table 1.1). That is, oxygen gas was produced at the anode of the battery and
hydrogen at the cathode. In today's terminology the following two half-cell
reactions can be written to describe the oxidation and reduction processes:
12
~0
2 (liquid)
solution)
-+
0 2 (gas)
+ H+(solution) + e- (oxidation)
+ e- -+ i ~ ~ ( ~ a s )
12 ~ 2 0 ( l i q u i d +
)
Hz(gas)
(reduction)
+ o2(gas)
or
2H20(liquid)-+ 2H2(gas)
+ 02(gas)
9
(1.57a)
(1.57b)
(1.57~)
(overall redox
reaction)
(1.57d)
Data in Table 1.3 reveal that under standard conditions of 1 atm of H2(gas)and
1 atm of Oa(gas)at 25OC
AGO
= 1 . 2 3~~m o l - '
(1.58)
which implies that a large amount of energy is required to drive this reaction.
However, for the reverse reaction
as) + 0 2 (gas) ---t H20(liquid)
(1.59)
A GO = -1 . 2 3J~mol-'
(1.60)
which advises that a great deal of energy will be released when this reaction
takes place. That is, large amounts of electricity should be available by reacting
these common gases in an electrochemical cell format. However, in the absence
of a catalyst, the reaction in eqn (1.59) is extremely slow, again highlighting
the need to understand the kinetics as well as the thermodynamics of reactions
when considering the practical usefulness of an electrochemical cell.
Fuel cells require the aid of 'catalytic7electrodes to exploit thermodynamically
favourable reactions such as that in eqn (1.59) and, like batteries, they transform
chemical energy into electricity. However, unlike batteries, fuel cells do not
store electrical energy. Rather, they convert energy from chemical reactions
directly into electrical energy.
The history and key features of the fuel cell have been reviewed recently by
Holper [I 01, Fuller [111 and Mobius [12]. William Grove produced the first
fuel cell over a hundred and fifty years ago (Table 1.1). He knew, on the basis
Application ofprinciples of electrochemistry
27
ofthe work of Nicholson and Carlisle, that sending an electric current through
water splits the water into its component parts ofhydrogen and oxygen. Thus, in
essence Grove demonstrated how to reverse the reaction and combine hydrogen
and oxygen to produce electricity and water in the first practical fuel cell.
Grove's and other simple fuel cells, as well as most electrochemically based
devices, consist of two electrodes separated by an ionic conductor (a salt or
acid solution) which acts as the electrolyte. In the hydrogen-oxygen fuel cell,
hydrogen is pumped to the anode, and oxygen to the cathode; a wire carries
the electrical current out of the fuel cell and ions carry the electrical current
through the electrolyte [lo]. This cycle continues, as long as the hydrogen and
oxygen fuel is supplied, with hydrogen and oxygen being turned into water
while generating electricity (Fig. 1.5).
Each hydrogen-oxygen fuel cell generates up to 1.23V under standard conitions (see eqn 1.58). However, most fuel cells operate at temperatures higher
than 25'C and pressures higher than 1atm so that the cell thermodynamics
usually must be derived on the basis of the equations and discussion presented
in Section 4. Individual cells can be wired together to produce greater voltages
or higher currents. The American space shuttle has 96 individual cells arranged
in three stacks. When hydrogen and oxygen are pumped into the shuttle's fuel
cells, they generate a 28-V power source or electricity supply as well as heat
and water. The heat is put to good use, vaporizing the liquid fuels before they
reach the fuel cells. Water flows into storage containers for drinking and other
cells are based on simple principles. However, the chemical reactions
d usually occur very slowly and unless special materials are used to construct the cells, very little current is produced. Consequently, much of the
research associated with commercial fuel cells is focused on the development
of suitable electrode materials and electrolytes to ensure that the rates of all
rocesses are sufficiently rapid so that power is available 'instantly'.
. 1.5 Diagrammatic representation of the hydrogen-oxygen fuel cell. Adapted from: Chem. in
Aust. October 1998, 21.
28
Thefundamentals of electvochemistry
Photovoltaic effects were first observed more than a hundred and fifty years
ago when Becquerel (Table 1.1) detected a photovoltage when sunlight was
allowed to shine on one of two electrodes he had placed in an electrolyte
solution. Figure 1.6 contains schematic representations of photovoltaic cells
based on the use of thin-film dye-sensitized cells introduced by Tsubomura
et al. [13] and then developed extensively by Gratzel to give efficiencies in
excess of 10 per cent [14]. In this regenerative Gratzel-type photoelectrochemical cell, a dye, usually based on compounds of the type Ru(bpy),(NCS),
(bpy = bipyridine; NCS- = N-bonded thiocyanate) absorbs light. As shown
in Fig. 1.6, the dye coats a nanostructured wide-band-gap semiconductor, such
as T ~ o ~which
,
permits efficient charge transfer. Again, -as for the fuel cell and
all other electrochemically based devices, two electrodes and an electrolyte are
present. In the ~ r a t z e l - t y ~photovoltaic
e
cell (Fig. 1.6 and Scheme 1. i ) , the
ruthenium dye, after excitation by light, takes on a standard potential that is
considerably different from the ground state or dark potential. The interaction
with light requires the use of a semi-conductor electrode rather than the simple metal or carbon-based electrodes previously considered. Furthermore, the
electrolyte used in a dye-sensitized photovoltaic cell commonly contains the
non-photoactive components that constitute a 'dark' half-cell reaction.
excitation:
-
e injection:
dye regeneration:
RU"
hv
(dcbpy,) ( N C S ) ~
F=+
T2
[RU"'
(dcbpy,) (NCS)~]' + e(dcbpy2)(NCS), + El
(dcbpyo-) (dcbpy)(NCS)~]'
[Ru"'(dcbpy2) (NCS)~]+ El- +
RU"'
electrolyte regeneration: El
+ e-
+
[RU"' (dcbpyo-) (dcbpy)(NCS)~]'
Pt
A
RU"
El-
Scheme 1.1 Schematic representation of reactions that occur in a dye-sensitized photoelectrochemical cell which demonstrates that both the ground and excited state redox potentials of the
c i s - [ ~ u ( d c b ~ ~ , ) ( ~couple
~ ~ ) ~are
] +important.
/~
El and E l are the oxidized and reduced forms of
the electrolyte (e.g. commonly ;1 and I- respectively). Adapted from reference [I 51.
A common sensitizer used in these ruthenium-titania Gratzel-type photovoltaic systems is cis-Ru(dcbpy),(NCS), (dcbpy = 2,2'-bipyridine-4,4'dicarboxylic acid). This sensitizer provides excellent absorption in the visible
spectrum, a high electron injection rate, high turnover rates and high stability in
photoelectrochemical cells. Scheme 1.1 summarizes the relevant reactions associated with a photovoltaic cell based on the use of the cis-Ru(dcbpy),(NCS),
dye, titania, and platinum electrodes, and acetonitrile containing I- and I,
which act as both the electrolyte and constitute the chemical components in the
f13(solution) + e'dark9 half-cell reaction.
$1- (solution)
(1.61)
Application ofprinciples of electvochemistvy
29
Ru(I1) Dye-sensitized solar cells
Porous T i 0 2 /
RU@PY)~(NCS)~
\ Clear / '
conducting
oxide
'
Electrolyte
containing I-/I<
(b) Light
111
-Clear conducting
oxide with electrocatalyst
I Glass
-Nanostructured T i 0 2
film
-Clear conducting
1 oxide
7 ,
Light
. 1.6 Schematic diagrams of dye-sensitized photoelectrochemical cells. (a) Provided by courtesy
of Leone Spiccia, Monash University, Victoria, Australia (bpy = bipyridine); (b) adapted from:
Electrochem. Soc. Inte?fdce6(3) (1997) 34 (I = ruthenium dye in ground (dark) state, I* = ruthenium
dye in photoexcited state).
As for fuel cells, extensive research is still taking place to improve the electrodes and electrolyte performance. Photon to electric current yields now being
achieved are as high as 33 per cent [16]. The field of high-efficiency solar cells
has been reviewed recently by Licht [17], who advises that progress in this area
of electrochemical technology over the last decade has been quite remarkable,
as indeed has been the case with fuel cells.
.3 Lead-acid battevy
The lead-acid battery was developed from the pioneering work of Planti in
1860 (Table 1.1). Figure 1.7(a) illustrates the main features of a three-cell 6-V
lead-acid car battery. Again, as in other electrochemical devices, negative and
30
Thefundamentals of electrochemistry
(a)
Positive electrode
Cell with electrolyte
-,
-
E' 1-0.356 V SHE
I
r Cell connector
I
r p b
EO
=+1.685
I
V SHE
I
Fig. 1.7 (a) A 6-V lead-acid storage battery is composed of three cells connected in series. Each
cell produces about 2 V; (b) reaction scheme for a lead-acid battery shown for the discharge mode
as a galvanic cell. Adapted from: J.E. Brady and J.R. Holum, Fundamentals of Chemistry, 3rd edn,
John Wiley, New York, 1988, p. 726.
positive electrodes and electrolyte are readily identified. The active component
of the positive electrode is lead dioxide, P b 0 2 , that of the negative electrode is
elemental lead and that of the electrolyte is aqueous sulphuric acid.
The half-cell (charge-discharge) reactions of the lead-acid battery for each
cell are given below.
At the positive electrode, EO = 1.685 V for the reaction
discharge
f P ~ (solid)
O ~+ 1H ~ S (aqueous)
O~
+ H' (aqueous) + e- F=+i p b ~ 0 (solid)
4
+ HzO(liquid)
charge
(l.62a)
At the negative electrode, EO = -0.356 V for the reaction:
discharge
Pb(so1id)
+ ; ~ ~ ~ ~ ~ ( a q u+
e o u s~) P ~ (solid)
S O+~H+(aqueous) + echarge
(l.62b)
verall, Ecell= 2.041 V for the reaction:
Lpbo2
(solid)
2
+ H 2 S 0 4(aqueous) + Pb
discharge
+
PbS04(solid) H20(liquid)
charge
(1.62~)
+ 2H2S04(aqueous)+ Pb +2PbS04(solid)+ 2H2O(liquid)
discharge
b 0 2(solid)
charge
(l.62d)
Thus, the three cells in the 6-V battery in Fig. 1.7(a) each contribute about
2 V. Obviously, in the common 12-V lead-acid battery used in automobiles
there are six 2-V cells. Under discharge conditions, the lead-acid battery acts as
a galvanic cell as shown in Fig. 1.7(b). In the charge mode, energy is required
to drive the reverse reaction, so the cell operates in a manner equivalent to an
electrolytic cell. Note that this implies that the assignment of the anode and
cathode can change depending on whether a lead-acid battery is being charged
or discharged, so as noted previously representation of the electrodes as positive
or negative is preferable terminology. Further details on the lead-acid battery
are available in references [3,4].
[I] P.W. Atkins, Chem. Aust., April 1991, 128.
[2] R.G. Compton and G.H.W. Sanders, Oxford Chemistry Primers, Electrode Potentials,
Oxford University Press, Oxford, 1996.
D.B.
Hibbert, Introduction to Electrochemistry, Macmillan, London, 1993.
[3]
[4] K.B . Oldham and J. C. Myland, Fundamentals of Electrochemical Science, Academic
Press, San Diego, 1994.
[5] P.H. Rieger, Electrochemistry, 2nd edn, Chapman and Hall, New York, 1994.
161 C.A. Russell, Chem. Br., September 1998, 36.
[7] S. Trasatti, J. Electroanal. Chem. 460 (1999) 1.
[8] See for example, Chemical Thermodynamics with Special Reference to Inorganic Chemistry,
Macdonald, London 1971, pp. 1-21 1.
[9] R. Van Eldik, T. Asano, and WJ. Le Noble, Chem. Rev. 89 (1989) 549.
32
[lo]
[111
[12]
1131
[14]
1151
[16]
[17]
Thefundamentals of electrochemistry
P. Holper, Chem. Aust., October 1998, 21.
T.F. Fuller, Electrochem. Soc. Integace 6(3) (1997) 26.
H-H. Mobius, J. Solid State Electrochern. 1 (1997) 2.
H. Tsubomura, M. Matsumura, Y. Nomura, and T. Amamiya, Nature (London) 261
(1976) 402.
M.K. Nazeeruddin, A. Kay, I. Rodicio, R. Humphry-Baker, E. Miiller, P. Liska, et
al., J. A m . Chem. Soc. 115 (1993) 6382.
R. Argazzi, C.A. Bignozzi, T.A. Heimer, F.N. Castellano, and G.J. Meyer, Inoz.
Chem. 33 (1994) 5741.
U. Bach, D. Lupo, P. Comte, J.E. Moser, F. Weissortel, J. Salbeck, et al., Nature
(London) 395 (1998) 583.
S. Licht, Electrochern. Soc. Interface 6(3) (1997) 34.
is,
ince Faraday's pioneering studies (Table 1.I), it has been known that the
transfer of electrons between a redox active species, dissolved in a solution phase,
and solid metal, or carbon, or liquid mercury electrodes, results in interesting
reaction pathways. However, voltammetric techniques, which involve the measurement and interpretation of I-E-t curves have only become popular since
the 1940s, when instrumentation required to conveniently conduct such experiments became readily available. The early studies invariably used a constant or
linear sweep of the direct current (DC) potential and measurement of D C current at the potential of interest. In the last 50 years a wide range of techniques
have emerged (Table 1.I). Thus, Randles and Sevtik reported the first cyclic
voltammetric studies in 1948 [l ,2], while the 1950s and 1960s saw the widespread use of hydrodynamic techniques such as rotating-disc electrode (RDE)
voltammetry [3] and alternating current (AC) pulse and square-wave voltammetry [4-61. The 1970s and early 1980s then witnessed the wide-spread use
of microelectrodes [6,7], initially under slow scan rate near steady-state conditions and subsequently, when advances in instrumentation occurred, under
extremely fast scan rate, transient D C voltammetric techniques [7,8].
Recent voltammetric studies in which the redox active components are soluble in the solution phase have reported the use of a wide range of waveforms
(time domains) and also combinations of macroelectrodes, microelectrodes, and
different forms of mass transport [9]. At the same time studies with solids attached
to electrode surfaces [lo] have emerged as a very important area of both pure
and applied research and are now almost as common as solution phase studies.
Thus, by the beginning of this millennium, an extensive range of voltammetric
techniques have been made available and can be applied to redox active species
in solution, solid and gas phases using conventional conducting electrodes, semiconducting electrodes and chemically modified electrodes. In this chapter only
'parts of Sections 2.1 to 2.13 of this Chapter have been adapted with permission from Adv.
Phys. Oy. Chem. 32 (1999) 1.
34
Principles of voltammetry
the relatively simple theoretical principles and techniques applicable to conventional studies will be considered, in which the redox active compound of
interest is soluble in the solution phase or else attached to an electrode surface
as an ideal thin film. Chapters 3 and 4 will provide a detailed discussion of
the application of the concepts contained in Chapter 2, while Chapter 5 will
deal with complexities associated with oxidation or reduction of solids adhered
to electrode surfaces in thick film or microcrystalline formats. The concluding
chapter describes the use of an integrated approach to problem solving under
specialized conditions encountered with metalloprotein voltammetry.
cell used for
ol
Figure 2.1 shows a typical experimental arrangement used for standard voltammetric experiments. This electrochemical cell has the following features:
(1) Three electrodes (working, reference and auxiliary) are present in close
proximity, with the working electrode being centrally placed. It is crucial
that the tip of the reference electrode is near to the surface of the working
electrode in order to minimize contributions of the IR, drop (I = current,
Ru = uncompensated resistance) to the applied potential. This situation
may be efficiently achieved using a Luggin capillary (Fig. 2.1).
(2) The cell contains an inlet and outlet for an inert gas which displaces electroactive oxygen from the solution. Typical inert gases include high purity
nitrogen and argon. Nitrogen is commonly used due to its low cost and
Connections to
potentiostat
1-
Water in from
thermostatic bath
Working
electrode
Fig. 2.1 Schematic diagram of a typical format for an electrochemical cell used in voltammetric
Chem. 32 (1999) 1. Copyright, Academic Press.
studies. Reproduced by courtesy: Adv. Phys. 0%.
Electvochemical cellfor voltammetvic experiments
35
ready availability in a highly pure form. However, for very air-sensitive
systems, argon is preferred, as it is heavier than air.
(3) The solution volume is typically in the range of 5-20 mL. However, specially designed cells in the pL range can be built, as can very large volume
cells containing manv litres of solution.
(4) The concentration of the electroactive species of interest when employing
DC voltammetric techniques is typically in the range 0.1-5 rnM. The upper
concentration limit is chosen to minimize IR, drop and possibly to minimize
the contribution of migration current from the electroactive species (see
later). At concentrations below 0.1 rnM, background current terms may
start becoming significantly relative to the Faradaic current of interest. In
analytical studies, trace concentrations down to
M, or even lower, may
be determined via use of AC, pulse or square-wave methods that enhance
the Faradaic-to-background current ratio [4,11].
(5) A high concentration of supporting electrolyte (>0.1 M) is usually added
to the solution to minimize the solution resistance, reduce transport of
electroactive ions by migration and to establish a well defined double layer.
Thus the supporting electrolyte has to be of high purity, should dissociate
substantially in the solvent of interest, consist of anions and cations that are
hard to oxidize and reduce (in order to provide as wide a potential window
as possible), and introduce no undesirable reactivity problems.2
(6) In the absence of constant temperature conditions within the laboratory, the
electrochemical cell may need to be thermostatted as voltammetric responses
can show significant temperature dependence.
U
J
The electrochemical cell in Fig. 2.1 is coupled to a three-electrode potentiostatted form of instrumentation [5,6]. If a two-electrode (working and
reference) system were to be used, the current would have to flow through
reference electrode, thus provoking instability in the reference potential.
thermore, in a two-electrode system, the IR, drop could be substantial. In
contrast, in the three-electrode potentiostatted system, essentially all the current
is forced to flow through the counter electrode, thereby minimizing problems
with the reference electrode. Additionally, much of the IR, potential loss is
ated by the potentiostatted circuitry [5,6,12], which drives the potential
the working and counter electrode to a value which compensates the
of the IR, drop. However, the use of a potentiostat does not remove
all of the IR, drop, since uncompensated resistance remains due to solution
resistance between the tip of the reference and working electrodes, and from
sistance inherently present in the working electrode and electronic circuitry.
bviously conditions where a very low value of I may occur, such as with a low
concentration of redox active species or when very small microelectrodes are
used [6,7], would represent the situations where a two-electrode voltammetric
2 ~ o m special
e
features encountered in voltammetric studies undertaken in the presence
of dilute supporting electrolyte or even the absence of added supporting electrolyte are
elucidated in reference 7.
36
Principles of voltammetry
electrochemical cell could be used, provided of course that the resistance, or
more importantly the IR, drop, is not excessive.
3
3.1
The electro
Working electrodes
This is the electrode at which the reaction of interest takes place, e.g. the simple
one-electron oxidation/reduction processes given in eqns (2.1) and (2.2) when
species X and X+ or Y and Y- are soluble in the solution (electrolyte) phase:
+ solution) + eY(so1ution) + e- + Y - (solution)
X(so1ution)
(oxidation)
(2.1)
(reduction)
(2.2)
However, it should be noted that the electrode processes are frequently
designated in this book via use of the formalism
or
Red
Ox
+ e-
(2.4)
where the charges and phases associated with the chemical species have been
omitted for convenience of presentation of mathematical relationships. In
eqn (2.4), the symbols Red and Ox (see Chapter 1) represent the reduced
and oxidized forms of the half-cell redox reaction respectively. In Chapter 1,
reversible half-cell reactions were written as reduction processes when defining the thermodynamics and relationships associated with the electrochemical
series. However, in voltammetry, by convention, oxidation currents are positive and reduction currents negative, so the half-cell reactions, frequently, will
be written in this chapter as the oxidation reactions in eqn (2.3) or (2.4) rather
than as a reduction reaction, for the purpose of a more convenient form of
presentation of the I-E-t curves derived from theory. A more detailed account
of the conventions is available in reference [13].
Working electrodes are fabricated from electrically conducting materials.
Common examples include liquid mercury, solid platinum or gold and some
forms of carbon (e.g. glassy carbon (GC) or graphite). Mercury electrodes [4]
are used in the form of dropping electrodes in which the surface is continuously
renewed, as a stationary hanging mercury drop electrode or as a thin film plated
onto a carbon or metal substrate. Historically, mercury electrodes have been
widely used for studying reduction processes (the positive potential range being
limited by the ease of oxidation of mercury). However, environmental concerns related to the toxicity of mercury and its compounds are beginning to limit
the use of mercury electrodes. Consequently, carbon, gold, and solid platinum
electrodes are now most commonly used as working electrodes, although there
is an almost infinite number of new generation electrodes being introduced
as advances in materials science occur. Typically, the solid electrode materials
Electrodes in voltammetric experiments
Side view
37
Top view
Wire connection
to potentiostat
Electrode material
typically smooth on p m scale
-
Diameter typically
mm+pm
Solder connection of wire
to back of electrode material
Conducting material
Electrode
(i) flush to insulating surface
(ii) well sealed in insulating mantle
(iii) polished
. 2.2 Schematic diagram of a typical disc working electrode employed in voltammetric studies.
Reproduced by courtesy: Adv.Phys. Org. Chem. 32 (1999) 1. Copyright, Academic Press.
are sealed into a non-conducting support (e.g. glass or Teflon) to form a disc
electrode (Fig. 2.2). Macrodisc electrodes have radii in the mm range, whereas
microdisc electrodes have radii in the nm to pm range (Section 10).
The quality of the voltammogram obtained depends on a number offactors:
(1) the integrity of the surface as measured by its cleanliness, smoothness, and
reproducibility of preparation. Failure to ensure that these ideal conditions
are operative may result in high, non-reproducible background currents as
well as distorted (with respect to theoretical expectations) Faradaic currents.
Usually electrodes are polished via systematic methods to achieve a high
quality surface using alumina or diamond paste as the polishing material [I41;
(2) the integrity of the seal between the electrode material and the inert, nonconducting, supporting mantle and the ability to ensure that electrode
and mantle surfaces remain flush at all times (see Fig. 2.2). A poor seal
or a recessed/protruding electrode may result in noisy, non-reproducible
voltammograms;
(3) the employment of a vibration-free environment and choice of conditions
where natural convection is minimized. This ensures that mass transport
(Section 7.1) is well defined;
(4) the ability to minimize problems associated with resistance. As noted above,
an important IRu drop term is associated with the uncompensated resistance arising from the solution between the working electrode surface and
the reference electrode. Thus, in order to minimize the undesirable I R u
potential drop term, small working to reference electrode separations and
adequately conducting solutions are generally employed. For microampere
38
Principles ofvoltammetvy
cell currents, the IR, term is in the mV range for typical organic solvent
electrolyte combinations e.g. CH3CN (0.1 M Bu4NPF6).T o achieve this
desired outcome, high resistance working, reference and auxiliary electrodes and highly resistive salt bridges and very low porosity, and hence
highly resistive, reference electrode frits, should be avoided. Any region
of high resistance in an electrochemical cell is likely to be 'bad news' with
result to IR, drop, noise, or R, C time constant ( C = capacitance).
3.2 Reference electrodes
The reference electrode ideally provides a fixed reference potential against
which the potential of the working electrode is measured. The most commonly used reference electrode in voltammetric studies, undertaken in aqueous
media, is the silver/silver chloride (Ag/AgCl) electrode depicted in Fig. 2.3.
The half-cell reactions utilized in this electrode are given in eqn (2.5)
Prior to environmental concerns being raised over mercury and mercury
compounds, the Calomel reference electrode (eqn 2.6)
2 Hg2C12(solid)
+ e- - Hg(1iquid)+ Cl- (solution)
A
(2.6)
was also very commonly used. The potentials of reversible reference electrodes
at 25°C and other standard conditions of activity etc. are defined thermodynamically via their Standard Electvode Potentials in conjunction with the
T o potentiostat
Ag wire
of (21- salt
riate solver
Salt
bridge
AgCl coating
on wire
Low porosity plug
Solution containing the same
solvent/supporting electrolyte
combination as in main cell
Fig. 2.3 Schematic diagram of a typical Ag/AgCl reference electrode and salt bridge. Reproduced
by courtesy: Adv.Phys. 0%.Chem. 32 (1999) 1. Copyright, Academic Press.
Electrodes in voltammetric experiments
39
ernst equation (see references [13,15,16], Chapter 1 and Table 1.3). The
source of aqueous solution-phase chloride ions for both the Ag/AgCl and
calomel electrodes is usually NaCl or KC1, in concentrations typically ranging from 1.0 M to saturated (in the case of the Calomel reference electrode,
en a saturated solution of KC1 is used, the special term of Saturated Calomel
ctrode (SCE) is used). Thus, reference electrodes are usually used under conditions of non-unity activity. However, for both these reference electrode types,
the potentials are frequently accurately known [5,15,16] relative to the SHE.
The SHE (eqn 2.7) is based upon the use of a high surface area 'platinum black'
in contact with hydrogen gas (1 atm) and hydrogen ions (unit activity) and, as
noted in Chapter 1 (Table 1.3), is defined at 25°C to have a potential of 0.000 V
any other forms of reference electrode are also used:
(1) Ag/Ag+(silver salt) electrodes [e.g. ~ ~ / A g + ( 0 . 0M
0 1AgNO,)] are commonly used for experiments involving conditions where AgCl is soluble.
For example, when acetonitrile is used as the solvent, AgCl is soluble and
a non-aqueous Ag/Ag+ (0.001 M AgNO,, CH3CN) reference electrode is
ideal provided care is taken to avoid photoreduction of the Ag+ ion.
(2) Ag or Pt wire quasi-reference electrodes3are frequently used for experiments
where the addition of deliberately added supporting electrolyte needs to be
avoided, or in solvents where no established reference electrode couple
exists.
(3) In voltammetric studies in organic solvents, it is now common practice to
~ H ~ ) ~ ,
measure the reversible potential for the oxidation of F ~ ( ~ ~ - cusually
called ferrocene (Fc), to the ferricinium cation (Fc'), versus the reference
electrode, or quasi-reference electrode3 actually used, and to subsequently
correct potentials to the FC/FC+ scale [l7,18]. In suitable cases, Fc may be
added to the solution being studied to enable an in situ form of calibration
to be achieved. This can be done when the potential of the FC/FC+ process
does not overlap with the process being studied and also where Fc and Fc+
are both chemically inert with respect to the oxidized and reduced forms of
compounds associated with the process of interest. More commonly, ex situ
reference electrode calibration is made by measurement of the reversible
potential versus the reference electrode used, and employing calculations
obtained from a voltammogram for the oxidation of a 5 x 1op4to 1 x 1on3M
ferrocene solution in the solvent (electrolyte) of interest. The reference electrode calibration procedure is analogous to the in situ or ex situ use of Me4Si
as a standard to calibrate 'H or 13cchemical shifts in N M R experiments.
he half-cell reaction giving rise to the potentials of these quasi-reference electrodes
may be unknown, but under the conditions of the experiment an acceptably reproducible
reference potential may still be achieved.
40
Pn'nciples of voltammetry
In order to minimize the risk of contamination from the salt present (e.g.
KC1 or NaCl), the reference electrode may be separated from the electrolytic
solution and working electrode by a salt bridge (Fig. 2.3). Care must be taken
to avoid precipitation of insoluble salts at salt bridge interfaces, for example,
K+ and C10,. Physical separation of the salt-bridge solution from that in the
reference electrode and that in the electrochemical cell is achieved by a low
porosity device such as a glass sinter, porous vycor (thirsty glass) or a membrane
(see Fig. 2.3).
Any junction (e.g. those associated with a salt bridge) may give rise to a liquid
junction potential (El,,). Such potentials, in combination with the ohmic potential drop term (IR,) may alter the potential applied between the working and
reference electrode (Eapp),SO that the effective cell potential (Ecen)is given [5] by
Careful choice of electrolyte in the salt bridge can be used to minimize Ehp.
For example, the liquid junction potentials at the reference/salt bridge and salt
bridge/working solution interfaces will have similar magnitudes but opposite
polarities and will, therefore, cancel each other out if a high concentration
of a salt whose constituent ions have similar ionic mobilities is present in the
salt bridge [5]. The liquid junction potential term is complicated if an aqueous
reference electrode is used in conjunction with an organic solvent test solution,
as a knowledge of the free energies of transfer of the ions between the aqueous
and organic solvents is required if corrections are to be made to compensate for
this term. Use of the Fc/Fc+ reference scale is a significant help in minimizing
problems with liquid junction potentials for measurements in organic ~ o l v e n t s . ~
3.3
Counter/auxiliavy electrodes
The terms counter or auxiliary electrode are used to describe the third electrode present in a typical potentiostatted voltammetric experiment (Fig. 2.1).
The counter electrode usually consists of a piece of platinum (wire or gauze)
or carbon (disc or rod) of large surface area placed directly in the test solution.
Since current flows through the counter electrode, it must have a sufficiently
large surface area relative to the working electrode to prevent limitation of the
current flowing in the total circuit. The current measured in a voltammetric
experiment flows between the working and counter electrode. Thus, if a reduction reaction is being studied at the working electrode a balancing oxidation
process (frequently electrolysis of the solvent medium) must occur at the counter
electrode or vice versa if an oxidation process is being considered.
In order to prevent extensive contamination of the test solution with products formed at the counter electrode or reaction of these products with the
electroactive solution of interest, a salt bridge should be used in conjunction
4~otentials
reported versus F C / F ~ +should be reproducible in any laboratory, irrespective
of the reference or quasi-reference electrode employed in the actual experiment.
Major classes of voltammetry
41
with the counter electrode in large-scale, long-time domain, bulk electrolysis
(Section 15) and equivalent classes of experiments.
asses of volta
.1 Tvansient voltammetry
In transient voltammetric experiments, a potential perturbation is applied to the
working electrode and the resulting current response associated with the reduction of oxidation reaction of interest is measured as a function of potential (time).
ransient techniques include D C linear sweep, D C cyclic, square-wave, pulsed,
and AC voltammetries (also see Section 8, and references [4-6,l2,l9,2O]). In
the first two cases, which are the voltammetric techniques considered most
extensively in this book, the D C potential of the working electrode is scanned
in a linear (or computer-generated staircase) fashion (Fig. 2.4) and the current
is monitored as a function of potential (time). The important temporal aspect
arises from the rate at which the potential E is ramped, dE/dt, known as the
scan rate, V . When the D C potential is swept in only one direction (Fig. 2.4(a)),
the technique is known as linear sweep (analogue instrument) or staircase (digital
instrument) voltammetry. If the potential is swept in one direction and then
reversed, and then this sequence is repeated, using a repetitive triangular potential excitation (Fig. 2.4), this technique is known as cyclic voltammetry. In most
alternating, square-wave, and pulse methods, a periodic waveform is superimposed onto the relevant D C waveform and the important temporal component
is usually the frequency of the superimposed signal. Figure 2.5 gives examples
of the waveforms used in square-wave and differential pulse voltammetry.
-1
CYCLE 1
F,
CYCLE 2 -1
0.2
0
5
10
15
20
Time (s)
. 2.4 Example of a typical waveform used in a D C cyclic voltammetric experiment. (a) Positive
potential scan from 0.25 to 0.75 V versus SCE (this is the waveform used in D C linear sweep voltammetry). (b) Direction of scan reversed at the switching potential of 0.75 V versus SCE. (c) Negative
potential scan from 0.75 V to 0.25 V versus SCE. (d) Termination of the first cycle. Reproduced by
courtesy: Laboratory Techniques in Electroanalytical Chemistry, Marcel Dekker, New York, 1984, p. 87.
42
Principles of voltammetry
(b)
//
Sample
width
\\
E A
I
I
I
I
I
I
I
I
I
I
/
/
/
II
1P
Pulse
period
Pulse
width
-
I
I
I
f
I
I
II _ - _ - _
I
I
II
II
II
II
A
,
/
//
/
!
t
t
Fig. 2.5 An example of a waveform used in (a) square-wave voltammetry in which a symmetrical
wave train (total amplitude 2Es,) is added to a staircase (step height AE) ramped voltage with a
period o f t . The current response is sampled at the end of both the forward and reverse half cycles (at
1 and 2). When a dropping mercury electrode is used a delay time Td is employed to allow the drop
to grow to a predetermined size. Reproduced by courtesy: Laboratory Techniques in Electroanalytical
Chemistry, Marcel Dekker, New York, 1984, p. 157; (b) Differential pulse voltammetry in which a
pulse is superimposed onto a staircase voltage and the difference in current before and after the pulse
is sampled.
4.2 Steady-state voltammetry
In the steady-state form of voltammetry, the concentration distributions of
each species in the electrode reaction mechanism are assumed to be temporally
invariant at each applied potential. Theoretically it takes an infinite time to reach
steady-state after the potential is changed. Thus, in a practical sense, steady-state
voltammetric experiments are conducted under conditions that approach sufficiently close to the true steady-state that the experimental uncertainty of the
steady-state value of the parameter being probed (e.g. current) is greater than
that associated with not fully reaching steady-state. The effective time-scale of
a near steady-state process is determined by the rate at which material reaches
the electrode surface. This time-scale may be varied in a number of ways:
(1) Altering the convective rate of transport, for example, by changing the
rotation frequency of a RDE (Section 9.1). Experiments in which the
Nature of the current-potential curve
43
convective rate of transport can be altered are known as hydrodynamic
techniques (see Section 9).
(2) Decreasing the size of the electrode so that the ratio of radial to linear diffusion of material to the electrode surface is enhanced. Voltammetric studies
at microelectrodes frequently exploit this method of altering the nature of the
diffusion characteristics to obtain a steady-state response (see Section 10).
ation of electrode reaction mechanisms
he kinetics of voltammetrically relevant reactions may be examined by varying
the critical time parameter of the experiment and monitoring its effect on some
voltammetric feature associated with the process being investigated. The general
cedure for obtaining quantitative kinetic data related to an electrode reaction
hanism using a voltammetric technique is schematically shown in Fig. 2.6.
he basic concept is that the experimental voltammetric data are collected
echanism for the electrode reaction mechanism is postulated. The promechanism may be theoretically simulated by solving the appropriate
matical problem. Satisfactory agreement between experiment and theory
is used to provide a quantitative description for a particular mechanism, but
with most kinetic studies, the identity of proposed reaction intermediates,
eally, should be confirmed by an independent technique, for example, an
ex situ spectroscopic or in situ spectroelectrochemical technique (Section 15). It
is inherently dangerous to assume the structure or even the identity of a reaction
roduct or intermediate solely on the basis of a voltammetric response.
Voltammograms are usually displayed as graphical representations of the current
endence (of the electrode reaction of interest) as a function of the potential
ifference applied between the working and reference electrodes. Three main
types of voltammetric wave shape are encountered, viz. sigmoidal or peakshaped; the latter may be symmetrical or asymmetrical. Figure 2.7(a) shows
the asymmetric peak-shaped response obtained under transient conditions of
C cyclic voltammetry when eqn (2.9) applies at a macrodisc electrode under
stationary solution conditions.
This class of transient response is characterized by oxidation and/or reduction
eak potentials (EpO"and/or E:~) and peak currents (1; and/or y d ) .
Figure 2.7(b) shows the sigmoidal-shaped response encountered in steadystate hydrodynamic and microelectrode voltammetry for the same class of
44
Principles of voltammetry
Electrode potential
defined manner
between electrode
Fig. 2.6 Schematic diagram illustrating the process of determining the mechanism of an electrode
process using voltammetric and spectroelectrochemical techniques. Reproduced by courtesy: Adv.
Phys. 0%.
Chem. 32 (1999) 1. Copyright, Academic Press.
electrode process. Steady-state voltammograms are usually characterized by a
and a half-wave potential
Figure 2.8 illustrates a
limiting current (Ilim)
combination of symmetrical and asymmetrical peak-shaped curves obtained in
square-wave voltammetry using the waveform depicted in Fig. 2.5(a) and the
mechanism in eqn (2.9). The current magnitudes in both transient and steadystate voltammetric techniques are usually proportional to the concentration of
the electroactive species and the electrode geometry as well as to the kinetics
of the heterogeneous and homogeneous reactions associated with the mechanism relevant to the electrode process being considered. Voltammetric studies
Nature of the cuvvent-potential curve
45
Oxidative
current
Oxidation
Reduction
Current
Current
Reductive
current
-
Applied potential (vs reference)
-
Applied potential (vs referent-e)
. 2.7 Commonly encountered voltammetric wave shapes: (a) asymmetric peak-shaped response
(e.g. cyclic voltammetry) and (b) sigmoidal-shaped response (e.g. steady-state hydrodynamic voltammetry). Reproduced by courtesy: Adv. Phys. 0%.Chem. 32 (1999) 1. Copyright, Academic
Press.
15
. 2.8 Square-wave voltammograms obtained in dimensionless current (q)form of presentation
when the waveform depicted in Fig. 2.5(a) is employed. (a) forward half-cycle; (b) reverse half-cycle;
(c) square-wave voltammogram obtained from the net current. Adapted from: Laboratory Techniques
in Electr~anal~tical
Chemistry, Marcel Dekker, New York, 1984, p. 158. Also see Anal. Chem. 53
(1981) 695 for further details.
on redox active solids (see Section 17 and also Chapter 5) give a vast range
of shapes because the mechanisms are inherently more complex and variable than is the case when the electroactive species are soluble in the solution
phase.
Principles of voltammetry
46
Cb)
f
Without unconipensated
resistance
Current
Current
Without uncompensated
resistance
With uncompensated
resistance
-
With uncompensated
resistance
,
'L'
Potential
Potential
v
_____)
Fig. 2.9 Effect of uncompensated resistance on (a) cyclic, and (b) steady-state voltammograms.
Reproduced by courtesy: Adv. Phys. 0%.Chem. 32 (1999) 1. Copyright, Academic Press.
Current
With double-layer
capacitance
/
Fig. 2.10 Effect of double-layer capacitance on a cyclic voltammetric response. Reproduced by
courtesy: Adv. Phys. Oy. Chem. 32 (1999) 1. Copyright, Academic Press.
The D C voltammograms displayed in Fig. 2.7 represent an ideal reversible
one-electron transfer process (eqn 2.9) in the absence of IRu drop or background current, although in real experiments the presence of contributions
from both these terms are unavoidable. Figure 2.9 shows the effect of uncompensated resistance for both transient and steady-state voltammograms, whilst
Fig. 2.10 shows the influence of double-layer capacitance which, usually, is
the origin of the background current present in a cyclic voltammetric curve.
However, it should be noted that with steady-state (time-independent) voltammetric techniques only a very low (ideally zero) capacitive background current
is expected. Both IRu drop and background current terms introduce distortions
which must be taken into account when comparisons are made with theory.
Unfortunately, many theoretical treatments of voltammetry only encompass the
Faradaic current and neglect IR, drop and background current contributions.
Cyclic voltammograms can be presented in an alternative format to that
shown in Fig. 2.7(a) by using a time, rather than potential, axis (see Fig. 2.1 1).
The equivalent parameters to scan rate (time) in steady-state voltammetric
Nature of the current-potential cuwe
47
2.11 Current-potential (a), and current-time (b) formats for presentation of a cyclic voltam~nogram.Reproduced by courtesy: Adv. Phys. 0%.Chern. 32 (1999) 1. Copyright, Academic
Press.
techniques are related to a hydrodynamic parameter (e.g. flow rate, rotation
speed) or a geometric parameter (e.g. electrode radius in microdisc electrode
voltammetry).
Faradaic and non- Faradaic currents
t is implied in the above discussion that the measured current contains a
contribution from both Faradaic and non-Faradaic (background) terms.
6.1.1 The Faradaic current
e Faradaic component is associated with the transfer of electrons resulting
from oxidation/reduction ofthe electroactive species ofinterest. The magnitude
of the Faradaic current is a function of many parameters and depends on the
exact nature of the voltammetric experiment and the mechanism associated with
e electrode process (see Section 8 for example).
6.1.2 The capacitive background cuwent
As a result of the layer of oppositely charged supporting electrolyte ions being
adjacent to the electrode surface [4-61, there is in effect a capacitive arrangement
in an electrochemical cell, which charges and discharges whenever the electrode
otential is changed. As a result, under conditions of cyclic or linear-sweep
oltammetry, a capacitive charging current, Ic, is generated which is directly
roportional to the scan rate, since
where C' is the capacitance of the double layer per unit area (A) and is usually
in the tens of pF ~ m range,
- ~ Q is the charge, E the potential, t the time and
w the scan rate. C' is assumed to be independent of electrode potential in the
above expression, which represents a considerable simplification. Under steadystate or more strictly, the near steady-state conditions which apply in practice at
0 so capacitive currents are very small with this form
slow scan rate, dE/dt
of voltammetry.
48
Principles of voltalnmetry
It will emerge from consideration of the theory for the Faradaic current
(IF)that IF oc v1I2 for a reversible process (eqn 2.9) under transient conditions at a macrodisc electrode with mass transport governed by linear diffusion
(Section 8.1) while from eqn (2.lo), Ic oc u . The different dependencies of IF
and Ic on scan rate means that transient voltammetric experiments will be limited by the scan rate which can be used because ultimately at a sufficiently high
scan rate, information associated with the Faradaic current will be swamped
by the presence of an unacceptably large capacitive background current. Some
of the problems associated with the capacitive current may be minimized by
using electrodes of smaller area (and hence lower overall capacitances), but still
0) near
retaining transient conditions, or by employing slow scan rate (dE/dt
steady-state conditions.
6.1.3 Other background currents
Background currents may also, of course, be Faradaic in nature and arise from
reduction of traces of oxygen in the solution or the presence of adventitious
redox active impurities on the electrode surface or in solution (e.g. oxygen
containing functional group on a graphite electrode or water in a non-aqueous
solvent). Since these Faradaic background currents are not capacitive in nature
they, of course, are not described by eqn (2.10).
Even a simple electrode process, in which the electroactive species are only
present in the solution phase (eqn 2.9), is quite complex because the overall
reaction consists of a considerable number of steps.
e
e
e
Mass transport of material to and from the electrode sudace.
Heterogeneous electron transfer between solid or liquid (mercury) electrodes and the solution soluble electroactive species.
Homogeneous chemical reactions coupled to the electron-transfer process.
7.1 Mass transport
In a voltammetric experiment in which the redox active species is soluble in
the solution phase, knowledge of the amount of electroactive material reaching
the working electrode, and the ability to alter the rate at which material reaches
5~raditionally,
the unit oflength used in electrochemistryhas been the centimetre. Accordingly, units for electrode area, the diffusion coefficient, and concentration, for example,
will be: cm2, cm2 s l , and mol cm-3 respectively. Use of the metre as a unit of length is
becoming more common in electrochemistry inwhich case the units for these parameters
would be m2, m2 s-l, and mol m-3 respectively.
Basic features of a n electrode process
Diffusion
due to concentration gradients
Convection
due to forced movement of solution
49
Migration
due to electrical fields
Movement of
species A
. 2.12 Different modes of mass transport of electroactive material to the electrode surface.
Chem. 32 (1999) 1. Copyright, Academic Press.
Reproduced by courtesy: Adv.Phys. 0%.
electrode are crucial for the determination of the reaction mechanism. There
three major pathways or modes of mass transport by which electroactive
material in solution may reach an electrode (Fig. 2.12). Thus, the total current
may be considered to consist of a linear contribution from three sources
where Id is the current associated with the diffusion process, I, is the current
associated with the migration process and Ic is the current associated with the
convection process.
iffusion involves the movement of species in solution due to a concentration
ent and is governed by Fick's two laws.
r one-dimensional diffusion of species A [5,6,9]
where Jd is the diffusional flux,6 DA is the diffusion coefficient of species A,
[A] is the concentration of A, and x is the distance from the electrode surface.
"umber
of moles of material diffusing through a unit of area in one second.
50
Principles of voltammetry
For diffusion in more than one dimension, the general expression [9]
may be used, where V is the Laplace operator and the flux density at the
electrode surface, JS, is related to the current, I, by the expression:
It follows from Fick's laws that the magnitude of the voltammetric current is
dependent upon the value of D, when the electroactive species is soluble in
the solution phase. The value of D, in turn, is a function of the solvent, the
molecular weight, and molecular dimensions of the electroactive species, the
temperature, and the electrolyte and at 25°C is usually in the range
cm2s-'
cm2 s-* for compounds of molecular weight below 2000 Da [21].
to 5 x
At macroelectrodes, under transient voltammetric conditions, semi-infinite
linear or one-dimensional diffusion is appropriate. For microelectrode geometries and steady-state conditions, the nature of diffusion is more complex, as
significant diffusion may occur in more than one dimension.
7.1.2 Migration
Migration results from motion of charged species, which occurs in the presence
of a potential gradient. Thus, charged electroactive species and the electrolyte
ions may migrate and contribute to the migration current. The total migrative
flux Vm)is related to the sum of the migration fluxes for each charged species.
For all the charged species (i) present, the migration flux due to a potential
gradient (aElax) is given by the Nernst-Einstein relationship (see reference [5],
for example):
-xiF
aE
Jm=
oi[i]all species i
RT
ax
or more generally for multi-dimensional migrational transport:
Jm=
-xiF
oi[i](BE)
RT
all species i
-
where z is the charge on the electroactive species, i. Obviously, migration may
be neglected for a neutral species. For charged species, the magnitude and sign
of the migration current is determined by the charge on the ion, xi. Thus,
if an anion is being reduced it will diffuse towards the (negatively charged)
working electrode but migrate to the (positively charged) counter electrode
in a potentiostatted three-electrode cell or the reference electrode in a twoelectrode cell. This migration of an anion, away from the working electrode,
effectively decreases the magnitude of the measured current (i.e. the current'in
the absence of electrolyte will be smaller than in the presence of electrolyte). Via
Basicfeaturesofanelectrodeprocess
51
With excess
supporting electrolyte
T
C
---------
Current
With no
electrolyte
f - g
I
I
I
I
I
Potential
0
. 2.13 Steady-state voltammograms obtained in the presence and absence of electrolyte for
oxidation of A2+ to A3+. Reproduced by courtesy: Adv. Phys. 0%.Chem. 32 (1999) 1. Copyright,
Academic Press. See J. Electroanal. Clzem. 337 (1992) 91 for details.
analogous logic it follows that when a cation is oxidized it will diffuse towards,
but migrate away, from the working electrode (positively charged), so again in
this situation the current measured will decrease from the diffusion-only value.
Conversely, if an anion is oxidized or a cation reduced, the current will increase
relative to the diffusion-only value. Figure 2.13 shows an idealized example of
a voltammogram obtained under microelectrode steady-state conditions, with
without added supporting electrolyte for the one-electron oxidation of
species A2+ (which has a counter anion X-) to A3+ [22]. Clearly, under these
circumstances, the measured current is expected to be suppressed by the presence of migration of the positively charged species A2+ away from the working
electrode (assumed to be positively charged). However, in most voltammetric
experiments a large excess of supporting electrolyte is utilized. Under these
conditions almost all of the migrational transport is associated with the supporting electrolyte and there will be a negligible contribution of migration of the
electroactive species to the overall appearance of the voltammogram 1131.
7 1.3 Convection
Convection current results from the movement of solution as a whole caused
by mechanical forces. The flux due to convection (7,) with a solution velocity,
v,, in one direction is given by:
or more generally, for a velocity vector, v:
Working electrodes which have material reaching them via a form of forced
convection are known as hydrodynamic electrodes. There are a wide range of
hydrodynamic electrodes, the common ones being rotating disc electrodes or '
Es [23], in which the electrode rotates at a fixed frequency and 'sucks up'
52
Principler of voltammetry
material to its sudace, and channel electrodes [24], over which the electroactive
species flows at a fixed volume flow rate (Section 9).
Each of the three mass transport terms defined above can be combined to
give a general mass transport equation describing the temporal variation of each
species in the electrode reaction mechanism. Thus, for species, A:
Note that opposite signs apply for the convection and diffusion terms in
eqn (2.19) because the concentration gradients resulting from each of these
processes are in opposite directions. Since migration of the electroactive species
of interest can be suppressed deliberately by the use of an excess of inert supporting electrolyte, generally, only diffusion and convection have been of interest to
the mechanistic electrochemist in the past. However, the addition of supporting
electrolyte is sometimes problematical and endeavours to undertake experiments without added electrolyte have recently become much more common
(see reference [7]).
Diffusion of material to the electrode may be controlled by altering the concentration of the electroactive species in solution and by changing the electrode
size from macro-dimensions where, in effect, diffusion is one-dimensional to
micro-dimensions, which may support two- or three-dimensional diffusion
(Section 10). However, while many electrochemical experiments are conducted
under diffusion-only conditions, when the redox active species is soluble in
solution, the rate of mass transport may be most easily changed by adjusting the
convective transport element.
7.2 Electron tranrfer
The voltammogram for a simple oxidative electron-transfer process when the
reduced (A) and oxidized (B) forms of the electroactive species are soluble in
the solution phase,
~/O,kO,cx
A
+B + ne-
(2.20)
can be described (charges on species A and B are omitted for simplicity) in terms
of the three parameters:
(1) EfO,the reversible formal potential ( V versus reference electrode); the use of
Eo, the standard electrode potential, is avoided because most voltammetric
experiments are not conducted under conditions of unit activity and standard
temperature and pressure (see Chapter 1).
(2) ko, the heterogeneous charge-transfer rate constant (cm s-l) measured at
EfO;the value of ko determines how far the peak potential (Ep)or half-wave
potential (Ell2)is removed from E~O.For fast reactions (e.g. ko 2 1.0 cm s-l)
Ell2 E~O.Table 2.1 contains examples of ko values for different systems.
le 2.1 Typical values of ko for one-electron oxidation or reduction processes
Charge-transfer process
Solvent/electrolyte/electrode
Temp.
(K)
Reference
k0
(cm s-l)
----
Oxidation of ferrocene
Reduction of anthracene
Reduction of benzoquinone
Reduction of co(H20)2'
Reduction of CO(NH~);+
Reduction of RU(NH~);+
Reduction of cyclo-octatetraene
Reduction of E U ~ +
Acetonitrile/O. 1 M Bu4NC104/
platinum electrode
Acetonitrile/O.6 M Et4NC104/
gold electrode
Acetonitrile/O.6 M Et4NC104/
platinum electrode
5.6 M HC104 (aq.)/
platinum electrode
0.1 M NaC104 (aq.)/
mercury electrode
0.5 M K2S04 (aq.)/
GC electrode
Dimethyl formamide/O. 1 M
Pt4NC104/mercury electrode
0.3 M NaC104 (aq.)/
GC electrode
"K.M. Kadish, J.Q. Ding, and T. Malinski, Anal. Chem. 56 (1984) 1741.
b ~ Bond,
. ~ .T.L.E. Henderson, D.R. Mann, T.F. Mann, W. Thormann, and C.G. Zoski, Anal. Chem. 60 (1988) 1878.
7.0. Howell and R.M. Wightman, Anal. Chem. 56 (1984) 524.
'J.M. Hale, Reactions ofMolecules at Electrodes (ed. N.S. Hush), John Wiley, Bath, 1971. Chapter 4.
'F. Marken, J.C. Eklund, and R.G. Compton, J. Electroanal. Chem. 395 (1995) 335.
f N.S. Hush, Electrochim. Acta 13 (1968) 1005.
54
Pn'nciples of voltarnrnetvy
Reaction coordinate for oxidation reaction
Fig. 2.14 Reaction coordinate diagrams for simple heterogeneous electron-transfer processes at an
electrode held at a potential of E: for a range of values of a . Reproduced by courtesy: Adv. Phys.
0%.Chem. 32 (1999) 1. Copyright, Academic Press.
(3) a, the charge-transfer coefficient; it is a dimensionless parameter and it
can be thought of in terms of the reaction coordinate diagram shown in
Fig. 2.14. a predominantly affects the shape and not the position of the
voltammetric response and, typically, it has a value of 0.5.
The relationship between these parameters and the current may be represented by the Butler-Volmer equation:
and [B],,o represent the electrode
where A is the electrode area and [A],'
surface concentrations of A and B. The full theoretical description of a voltammogram therefore is obtained by combining eqn (2.21) with the appropriate
mass transport equation (i.e. the appropriate version of eqn (2.19) which leads
In the case of fast electronto the required knowledge of [A],,o and [B],,'.
transfer kinetics, the theoretical expression derived in this manner, as expected,
becomes equivalent to that obtained by combining the Nernst equation with
the mass transport equation.
An alternative theory to the Butler-Volmer theory for electron transfer results
from considerations associated with the Marcus-Hush theory [25,26]. However,
in most cases over the potential region of interest in voltammetric studies when
species are soluble in the solution phase, use of the simpler Butler-Volmer
equation is usually adequate. In contrast, when species are confined to the
surface and a slow rate of electron transfer occurs, use of more sophisticated
relationships may usually be required (Section 17).
7.3 Homogeneous chemical kinetics
It is common for homogeneous chemical reactions to accompany the electrontransfer step. Thus, an electrochemical reaction mechanism may consist of a
combination of heterogeneous electron transfer and homogeneous chemical
reaction steps, each with their own individual rate constants. If the product, B,
Basic features of an electrode process
55
of eqn (2.20) undergoes a first-order, solution-phase chemical reaction with a
pate constant, kl,
~ k ' c
(2.22)
then a complete description of the electrode process when the initial chargetransfer step is reversible is given by the reaction scheme:
This frequently observed reaction scheme is commonly designated as an E C
or Ere,Cirrevmechanism7 [27,28]. Thus, homogeneous kinetic terms may be
combined with the expressions for diffusion and convection (i.e. an appropriately modified version of eqn (2.19)) to give the temporal variation of the
concentration of a species in an electrode reaction mechanism. In order to model
the voltammetric response associated with an ErevCirrev
mechanism, a knowledge of, E:, a , ko and kl is required, or else they may be deduced from a
theoretical-experimental comparison, and the set of concentration-time equations for species A, B and C must be solved, subject to the constraints of the
utler-Volrner equation and the experimental design.
Another example of a common electrode reaction mechanism encountered
in voltammetric studies is the ECE (e.g. ErevCirrevErev)
mechanism:
) ~ kl , the voltammogram
epending on the relative values of (E:)~, . ( ~ f oand
sociated with this ECE mechanism consists of two resolved one-electron
transfer processes (Fig. 2.15 (a)) or a single overall two-electron transfer process
2.15(b)) or intermediate situations. Obviously, it is possible to construct
trode reaction mechanisms with an infinite number of combinations of E
C steps [5,6,9,27,28], with different levels of reversibility being associated
with both the E and C components.
Electrochemical and chemical reversibility
A term that should be clearly defined and one that is often used haphazardly
in voltammetry is Reversibility. One must make a clear distinction between
Electrochemical Reversibility and Chemical Reversibility.
7~ubscripts
'rev' and 'irrev' stand for reversible and irreversible respectively, see Section 7.4.
56
Principles ofvoltammetry
f
One electron
.'. '.
\
f
ECE response
Current
One electron
response
Potential
Potential )
-
Current
T
F
/\
ECE response
/ One electron response
,//*
_--___-_-____
one
,,tl
,
I
--
I
Fig. 2.15 Peak-shaped (cyclic voltammetry) and sigmoidal-shaped (steady-state) voltammograms
< ( E ~ O ) ~ and (b)
> ( E : ) ~Reproduced
.
by
associated with an oxidative ECE process (a) (E~O)~
Chem., 32 (1999). Copyright, Academic Press.
courtesy: Adv. Phys. 0%.
(E~O)~
7.4.1 Electrochemical reversibility
The extent of electrochemical reversibility of a process is related to the
heterogeneous kinetics of electron transfer at the electrode surface. For a facile
electron-transfer reaction, equilibrium is achieved rapidly and the system is
defined as being electrochemically reversible. Effectively, this means that both the
forward and reverse electron-transfer reaction steps are rapid. Under conditions
of electrochemical reversibility, the Nernst equation
applies for the process described by eqn (2.20). This Nernst relationship arises
as a direct consequence of the fact that the electron-transfer kinetics for the
forward and reverse processes are so facile that equilibrium is attained at
each potential applied on the time-scale of the particular experiment. Thus,
an electron-transfer reaction under conditions of cyclic voltammetry may be
termed electrochemically reversible at a scan rate of 50 rnV s-' , but irreversible
at 1000 V s-'. The term is therefore a practical, rather than absolute, one and
is dependent upon the time-scale of the electrochemical measurement, which
means that a process may be reversible under conditions of cyclic voltammetry
(e.g. slow scan rate), but irreversible under some conditions of hydrodynamic
voltammetry (e.g. fast rotation rate), etc.
Cyclic voltammetry under transient conditions
57
7.4.2 Chemical reversibility
The level of chemical reversibility is associated with the stability of the species
involved in the electron-transfer step. Therefore, if species B irreversibly reacts,
as it is formed from the one-electron transfer process, to form species C,
as in eqn (2.23), then the overall process (A +- C) would be described as
eing chemically irreversible. However, if the chemical step associated with
the chemical reaction step of B was sufficiently fast, in both the forward and
backward directions, so that equilibrium is effectively maintained on the timescale of the voltammetric experiment, then the whole process would be termed
chemically reversible. Thus, concepts related to electrochemical and chemical
ersibility can be demonstrated by considering the E C mechanism described
eqns (2.26a) and (2.26b), and a range of scenarios.
(1) At extremely fast scan rates the electron-transfer step (E) is electrochemically
irreversible (Eirrev)
and the C step (Cirrev)
is outrun. Under these conditions,
the process would be described as an electrochemically irreversible process.
(2) At moderate scan rates, the C step (Cine,) is outrun, but the E step (Ere,)
is now reversible. This would be described as an electrochemically reversible
process.
(3) At slow scan rates, the E step (Ere,) remains reversible and the C step
(Cirrev)
is irreversible. This would be termed as an electrochemically reversible
and chemically irreversible process.
(4) At very slow scan rates the E (Ere,) and C steps (C,,,) are now reversible.
This would be described as an electrochemically and chemically reversible process.
Chapter 3 will contain detailed discussion of electrode processes where both
e heterogeneous charge transfer and homogeneous chemical steps are
reversible, while Chapter 4 will focus attention on what occurs when
irreversibility or other nuances arise in the description of the electrode processes.
Cyclic voltammetry is undoubtedly the voltammetric technique most widely
used by non-specialists in the subject of electrochemistry who are interested in
understanding the qualitative aspects of the mechanism of a Faradaic electrode
rocess. The experimental design usually consists of an electrochemical cell
containing the three electrodes described in Section 2 with both the working
electrode and solution being stationary. In cyclic voltammetric experiments,
58
Principles of voltarnrnetry
Fig. 2.16 Potential-time profile used in a typical cyclic voltammetric experiment. Reproduced by
courtesy: Adv. Phys. 0%.
Chem. 32 (1999) 1. Copyright, Academic Press.
the potential at the working electrode is usually swept at a constant scan rate
(v) from an initial potential value of El to a second potential E2 (Fig. 2.16(a),
also see Fig. 2.4). O n reaching E2, the direction of the sweep is reversed and,
when the potential returns to El, the scan may be halted and again reversed
or allowed to continue to a third potential E3 (not shown). Additionally, it is
common to cycle through the potential window of interest a number of times
to examine the stability of products formed via heterogeneous electron-transfer
reactions as well as to detect any new electroactive products formed as a result
of their decomposition. Figure 2.4 contains an example where two cycles of
the potential are employed.
A typical potential-time profile for a single cycle experiment is illustrated in
Fig. 2.16. The scan rate is represented by the magnitude of the slope of the
potential-time plot so that by this convention v is always said to be positive.
The expressions given in eqns (2.27) and (2.28) describe the potential (E(t)),
applied at the working electrode, as a function of time:
+ vt
Reverse Sweep: E(t) = -El + 2E2 - vt
Forward Sweep:
E(t) = El
(2.27)
(2.28)
These equations assume that the initial scan direction is positive, as normally
will be the case when studying an oxidation process. In eqns (2.27) and (2.28) it
is assumed also that the scan rate is the same in both the initial and reverse sweep
directions which need not always be the case (the scan rate may be increased
in the reverse scan in order to outrun homogeneous chemical steps associated
with species formed by heterogeneous electron transfer in the forward scan).
The scan rate may range from a few millivolts per second [5,6] to a million volts
per second [29-321. The lower scan rate limit is restricted in value by the effects
of natural convection which arises from the build-up of density gradients in the
solution resulting from such factors as inadequate thermostatting or mechanical
vibration. Natural convection adds to the rate of mass transport of material
to the electrode surface and thereby causes the experiment to deviate from
the diffusion-only regime. The upper scan rate limit is restricted by capacitive
charging (Section 6.1.2) since at very fast scan rates the capacitive current will
mask the current associated with the Faradaic process.
In cyclic voltammetric experiments and in the presence of excess electrolyte,
ideally, the sole form of mass transport to the electrode surface is diffusion, and
Cyclic voltammetry under transient conditions
59
. 2.17 Reaction coordinate system in a cyclic voltammetric experiment at a macrodisc electrode
for the process A + B e- and mass transport by linear diffusion. Reproduced by courtesy: Adv.
Phys. 0%.Chem. 32 (1999) 1. Copyright, Academic Press.
+
in the case of large macrodisc (millirnetre dimensions) electrodes, the diffusion
of material to the electrode occurs in the single dimension perpendicular to the
electrode surface. Figure 2.17 shows the reaction coordinate system that applies
to the process A + B -I- e under these macrodisc electrode conditions. As
will be discussed in Section 10 the situation is more complex for electrodes of
smaller (e.g. microdisc) dimensions.
Theory of cyclic voltarnrnetry
8.1.1 A reversible process
onsider a simple reversible one-electron oxidation process such as the oxidation of ferrocene (species A) to the ferricenium cation (species B) in acetonitrile
(0.1 M Bu4NC104 [33,34]). Initially only A is present in solution, the reaction
coordinate system in Fig. 2.17 applies and transient conditions are assumed. At
the usual macrodisc electrode (radius in mm range), material reaches the electrode by linear diffusion which is perpendicular to its surface (x-direction), and
the concentrations of A and B may be obtained as a function of time by solving
ick's second law of diffusion as applied to species A and B:
Additionally, the problem is subject to a number ofboundary conditions which
are defined in Table 2.2, and which are crucial to finding a solution to these
differential equations. It should also be noted that the time variation of the
electrode potential is given by eqns (2.27) and (2.28). Details of the solution of
eqns (2.29) and (2.30) are available in references [5,6].
The theoretical cyclic voltammogram shown in Fig. 2.18 for a reversible oneelectron oxidation process is obtained by scanning from an initial potential (El)
60
Principles of voltammetry
Table 2.2 Boundary conditions for the reversible one-electron oxidation
process A
B e under conditions of cyclic voltammetry
+
Time
coordinate
Spatial
coordinate
Species A
Species B
Reason for boundary condition
t=O
t>O
x>O
[A] = [Ale
[A] = [Ale
[B] = O
[B] = O
t>O
x=O
D A ( ~ [ A ] / ~ ~=
)X=O
(electrode --(a
[B]/a t)x=o
surface)
x=O
In([Blx=o/[A]x=o) =
nF/RT(E - E~O)
Initially, only A is in solution.
At large distances from the
electrode, the concentrations
of A and B tend to their
original values.
The rate at which A diffuses to
the electrode must equal the
rate at which B diffuses away.
The Nernst equation for an
electrochemically reversible
system applies at the site of
electron transfer, i.e. the
electrode surface.
t>O
X+OO
E~O
which is considerably less positive than
to a value E2 which considerably
more positive than EfO, and then scanning back to El. The forward and reverse
peaks constitute a voltammetric wave and the entire current-potential curve
represents what is termed a cyclic voltammogram.
The shape of the reversible cyclic voltammogram shown in Fig. 2.18 arises
for the following reasons. O n scanning the potential from El to more positive
values, the concentration of A at the electrode surface ([A],,o) decreases progressively as A is converted into B; this results in an increased concentration
gradient of A at the surface of the electrode, and thus the diffusional flux of
A to the electrode increases. The flux of material A to the electrode surface is
directly related to the current by the expression:
As the potential approaches E;, [A],,o decreases even further. Thus, the flux of
A to the electrode continues to increase, causing the current to rise. However,
eventually [A],=o reaches zero and the flux of A cannot change any further.
Under the conditions of this cyclic voltammetric experiment, once [A],=o = 0
the Nernst diffusion layer (the distance from the electrode at which concentration changes in A are associated solely with the electrolysis mechanism and
the resulting diffusion is the only form of mass transport) begins to relax further
into the solution as the diffusion process tries to equalize the concentrations of
A and B throughout the solution. Consequently, at very positive potentials, the
Cyclic voltammetry under transient conditions
61
oxidative
current
reduction
oxidation
reductive
current
I
I/ '
E ; ~
EpOX
E(vs arb. reference electrode)
Typical cyclic voltammogram obtained for a reversible one-electron oxidation process at
25°C. Reproduced by courtesy: Adv.Phys. 0%.
Chem. 32 (1999)l. Copyright, Academic Press.
ux of A drops, and thus the current is seen to decrease at potentials more posin the peak oxidative potential (EpO")to give the characteristic asymmetric
ssociated with a cyclic voltammogram. O n reversing the scan, initially
a high concentration of B at the electrode surface which decreases as B
ed back to A. However, at potentials sufficiently negative compared to
=o will return to zero and the same asymmetric peak-shaped response
is observed on the reverse sweep as for the forward sweep.
Consideration of the above discussion and examination of Fig. 2.18 shows
at there are a number of important parameters in cyclic voltammetry.
) The peak potentials, EpOX and E F ~ ,for the oxidation and reduction component of the experiment respectively; for a reversible process at 25"C, the
peaks will be separated by 56/nmV (where n is the number of electrons
transferred; in the example given in eqn (2.19), n = 1) and E; and
will
be independent of scan rate. The peak-to-peak separation is usually termed
the AEp value for the process.
(2) The midpoint, Em, or half-wave potential, Ell2, is related to the peak
potentials by the expression
~r~
62
Principles of voltammetry
If A and B have equal diffusion coefficients, Ell2 is identical to the formal
reversible potential (EfO).Otherwise, Ellz is related to E: by the expression:
E~O
For most redox couples, Ell2 only differs from
by a few millivolts.
respectively (note the base lines from which
(3) The peak currents, IpOxand
they are measured in Fig. 2.18); for a reversible process at 25'C, the value
of the peak current, in ampere is given by the Randles-SevCik expression
rd
where [Ale or [A],=, is the concentration of A in the bulk solution
(other symbols have already been defined) so that the magnitude of the
~ unity.
l
Equation 4.107 in Chapter 4 provides the general
ratio 1 l ~ / $ e is
expression for the peak current.
The increase in current with scan rate may be explained by the fact that
as the scan rate increases, less time is available for the Nernst diffusion layer
to relax into the bulk solution phase by diffusion. Consequently, as the scan
rate increases, the rate of change of concentration of A at the electrode surface
increases resulting in a greater flux of A to the electrode surface which in turn
gives rise to a larger observed current.
8.1.2 A n irreversible process
When the electron-transfer kinetics are slow (rate-determining) relative to mass
transport, the electrode process is no longer in equilibrium and therefore does
not obey the Nernst equation. As a result of the departure from equilibrium, the
kinetics associated with the rate of electron transfer at the electrode-solution
interface (rate constant k;' for the forward direction of the electron-transfer
process and kit for the backward reaction) have to be considered when discussing the voltammetry of non-irreversible systems. For the fully irreversible
process, kEt can be neglected so that the solution of the equation for the fully
irreversible process is achieved by replacement of the Nernst thermodynamic
boundary condition (Table 2.2) by a kinetic boundary condition which gives
for an irreversible oxidation process:
Thus, for a completely irreversible electron-transfer process, the rate limiting
step over a wide range of potentials is the electron-transfer step rather than diffusion. k;' is related to the electrode potential and the standard rate constant, kO,
Cyclic voltammetry under transient conditions
El
E (vs arb. reference electrode)
63
E2
. 2.19 Typical cyclic voltammogram obtained for an electrochemically irreversible one-electron
oxidation process. Reproduced by courtesy: Adu. Phys. 0%.Chem. 32 (1999) 1. Copyright,
Academic Press.
utler-Volmer theory (eqn 2.36).
kFt = k" exp (1 - a ) ( E - E;)~F/RT
(2.36)
Use of the Butler-Volmer equation (only the first term in eqn (2.21) for a completely irreversible process) and Fick's laws of diffusion enables the voltammetric
response of an electrochemically irreversible process to be calculated.
A typical voltammogram associated with an electrochemically irreversible
one-electron oxidation charge-transfer process is shown in Fig. 2.19. A number of differences, when compared to the electrochemicallyreversible case, may
be noted.
(1) There is no reverse peak in an irreversible cyclic voltammogram because
the reverse electron-transfer process does not occur at a measurable rate.
(2) The peak current, in amperes, at 25°C for an irreversible oxidation process
is given by the expression [4-61
I,"" = (2.99 x
Note that in eqn (2.37), n, refers to the number of electrons transferred in
the E steps before the rate-limiting electron-transfer step whereas n represents the total number of electrons transferred. Comparison of eqns (2.34)
and (2.37) under equivalent conditions reveals that the peak current for an
irreversible process is lower than the equivalent value for a reversible one.
This feature emerges because the kinetics of the electron transfer are relatively slow in the irreversible case, so that during the course of the potential
scan, diffusion has more time to relax the concentration gradient of A at
the electrode surface. Consequently, when [A],o = 0, the flux of A to
Principles of voltammetry
the electrode surface for an irreversible process is lower than for the equivalent reversible case, resulting in the occurrence of a decrease in the value
of the peak current. The sluggish electron-transfer kinetics also broadens
the voltammogram which results in the peak potential for the oxidation
process being shifted to a significantly more positive potential compared to
the formal potential (E:) for the electron-transfer process.
The peak potential is a function of scan rate, unlike the case for a reversible
process when the peak potentials are independent of scan rate. As the scan
rate increases, oxidation peak potential shift to more positive potentials.
8.1.3 A quasi-reversible process
Obviously, if both reversible and irreversible categories of process may exist,
then there must be an intermediate case in which the kinetics of both the
forward and reverse electron-transfer processes need to be considered. That is,
both kFt and kEt (forward and backward electron-transfer rate constants) must be
considered in solving the theory, which implies that both terms in the ButlerVolmer equation (eqn 2.21) are required. Such systems are described as being
quasi-reversible and, as would be expected for this category of electrode process,
the scan rate can have a considerable effect on the nature of the observed cyclic
voltammogram. At sufficiently slow scan rates, quasi-reversible processes appear
to be fully reversible. However, as the scan rate is increased, the kinetics of the
electron transfer are not fast enough to maintain (Nernstian) equilibrium. In the
scan rate region, when the process is quasi-reversible, the following observations
can be made.
The separation of the forward and reverse peaks (AEp)is larger than the value
of 56/n mV associated with a reversible process at 25°C. Importantly, AEp
increases with increasing scan rate and the value of the standard rate constant
for the electron-transfer process, ko, may be calculated from the separation
of the peaks in a quasi-reversible process [5,6], provided voltammograms
are corrected for solution resistance and background current effects.
The peaks become broader as the scan rate increases, and the peak current
is below the value expected for a reversible electron-transfer process.
The magnitude of ratio of the peak currents 1 1 ~ / 1 ~ ~ 1 is equal to unity for
a quasi-reversible system when a = 0.5.
It can be seen that the relative rates of electron transfer and the potential
scan rate may crucially determine whether voltammograms are observed to be
reversible, irreversible or quasi-reversible. Matsuda and Ayabe [35] proposed
the following (ko,V) regimes in order to define whether an electron-transfer
process will be observed to be reversible, quasi-reversible or irreversible:
k" 2 0.3 v ' I 2cm s-'
Quasi-reversible: 0.3 v'I2 > k" > (2 x
v'I2 cm s-'
Irreversible: k" 5 (2 x 1o - ~ )v 'I2 crn s-'
Reversible:
(2.38)
(2.39)
(2.40)
Cyclic voltammetry under transient conditions
65
It cannot be emphasized too strongly, that considerable care should always
be taken when interpreting the results of cyclic voltammetric experiments to
ensure that the effects of the double-layer capacitance and uncompensated
solution resistance are considered. Peak currents should be corrected for the
baseline capacitive charging current (for example by recording a background
voltarnrnogram obtained in the absence of the electroactive species and subtracting this response from voltammograrns containing the redox active component).
ince the charging current is proportional to scan rate, such background subtractions of this current are particularly crucial at fast scan rates. Uncompensated
solution resistance (R,) causes the peak-to-peak separation (AEp)to increase and
the peaks to broaden in a cyclic voltammogram which may make a reversible
system appear quasi-reversible. In the presence of uncompensated resistance, the
value of the applied potential will differ from the actual value by IR,. Modern
tentiostats are capable of compensating for the majority of the effects of this
tion resistance, by feeding back an additional potential, equal to IR,, to the
ied potential. The alternative and perhaps preferred approach when experimental data are to be compared to theory is to measure R, and include this
term in the theoretical simulation of the experiment.
.I .4 Chemical reactions coupled to the electron-transferprocess
ach process observed in a cyclic voltammogram is indicative of an electrode
reaction associated with material initially present in the bulk solution or generated at the electrode surface. Consequently, by varying the scan rate of the
cyclic voltammetric experiment, new waves may appear as a result of oxidation
or reduction of products formed at the relevant scan rate (time-scale). Therere, as well as providing an indication of the reversibility of an electron-transfer
ocess as described above, the presence of homogeneous chemical reactions
associated with the electron-transfer process may be detected by varying the scan
. The example below shows the power of cyclic voltammetry to interrogate
mechanisms of electrode reactions, even when they consist of a combination of heterogeneous (electron-transfer) and homogeneous (solution-phase)
chemical steps.
Examination of the cyclic voltammograms (Fig. 2.20) obtained [36] for the
oxidation o f f a c - M n ( ~ 0 (q2-dpm)
)~
~1 VacO),where dpm is the bidentate phoshine ligand Ph2PCH2PPh2species, in acetonitrile (0.1 M Bu4NC104)at slow
scan rates ( i 5 0 0 mV s-l) reveals the presence of a partially reversible oneelectron oxidation process having an oxidative peak potential of about 1.48 V
(versus Ag/AgCl) and reductive peak potential off 1.41 V. However, the peak
currents for the oxidation and reduction processes are not equal. The voltammogram is not fully reversible in the chemical sense because on the time-scale of
this experiment (scan rate = 500 mV s-') some of t h e f a c - M n ( ~ ~ ) , ( q ~ - d p m ) ~ l
cation Vat+) has isomerized to the mer cationic form (mer+) by the time the
potential is swept through the reversible potential for thefacO/faci couple. Confirmation of this conclusion is found by continuing the reverse (reductive) part
of the potential sweep through the facO/fac+ couple and noting that a new
+
66
Principles of voltammetry
Fig. 2.20 Cyclic voltammogram obtained at a platinum macrodisc electrode at a scan rate of
-+
f-
500 mV s-I for the ECE mechanism (eqns 2.41a-c) which applies to the oxidation of 1 rnM
f ~ c - M n ( C 0 )(r72-dprn)~1
~
in acetonitrile (0.1 M Bu4NC104)at 20°C. Adapted from: Inorg. Chem.
16 (1977) 155.
reduction feature is observed at a potential of fO.95 V, which can be shown to
correspond to the reduction of the mer- [Mn(CO), (q2-dpm)~ 1 1 ' (me?+) cation
to the m e r - ~ n ( ~ ~ ) , ( q ~ - (meuo)
d ~ m uncharged
)~l
species [36]. If the second
cycle of the potential is examined (Fig. 2.20), a new oxidative feature is observed
at a potential of +1 .O1 V corresponding to the oxidation of the mero to the meu+
cationic form. Thus, the first part of the mechanism at scan rates <SO0 mV s-'
-+
is described by the following EC scheme:
However, since the mer+ species may participate in a reversible one-electron
= +0.95V) on this voltammetreduction process (E,O" = +1.01 V,
ric time-scale, a third reaction needs to be added to the mechanism to give
~r~
-+
t
an ECE reaction sequence8 (eqns 2.4 1a-c) .
+- -+
he ECE notation is introduced to indicate that the two charge-transfer steps are oxidation and reduction processes respectively (or vice versa). This distinguishes the reaction
scheme from an ECE process where both charge-transfer steps are oxidation (or both
reduction).
Cyclic voltammetry under transient conditions
67
~t very high scan rates, the facO/fac+
couple appears to be reversible,
P
concomitantly, the mer+/merO couple disappears. Thus,
with I,oX = -Ped;
at very short time domains, the fac+ -+ mer+ isornerization step is outrun and the oxidation process becomes a single one-electron charge-transfer
process.
The kinetics of thefac' + mer+ isomerization and the rates of the electrontransfer steps can, in principle, all be determined quantitatively from the scan rate
dependence of the cyclic voltammetry and comparison with a voltammogram
simulated according to the proposed mechanism. The isomerization rate constants for this system have been determined by alternative voltammetric methods
[37].However, a quantitative cyclic voltammetric approach (Fig. 2.21) has been
presented recently for a related isomerization mechanism (see reaction sequence
in eqns (2.42a-c)) that occurs during the course of voltammetric oxidation of
Theory
-0.4
-0.6
0.0
-0.2
0.4
0.2
Potential (V vs FC/FC')
.
,
-0.6
.
,
.
,
.
,
.
,
-0.4
-0.2
0.0
0.2
Potential (V vs Fc/Fcf)
.
,
.
0.4
+ +
2.21 Comparison of experiment and simulation (according to the E C E mechanism in
0 -dpm)
) ~ ( ~ ~ - d p min) dichloromethane
~r
at 20°C
eqns (2.42a-c)) for oxidation of cis, m e r - ~ n ( ~(yl
using a 1-mm diameter platinum disc electrode (a) scan rate, 100 mV s-' (kl = 2.9 s-l), (b) scan-rate,
1000 m~ s-' (kl = 3.4 s-l). Reproduced by courtesy: Inoug. Ckem. 38 (1999) 2005. Copyright,
American Chemical Society.
68
Principles of voltammetry
cis, mer-M~(co),
(11'-dpm) ( q 2 - d p m ) ~in
r dichloromethane [38].
-+
E
(Efo)l
d
[cis, mer] ,
- [cis, mu]+
+ e- + [translo
(2.42a)
(E32
t
E
+ e-
[trans]'
(2.42~)
At 20°C, the voltammetric response was simulated using (E;), = 0.16 V versus
Fc/Fc+; (E:), = -0.40 V versus FC/FC+ and kl = 3.1 f 0.3 s-'. (Also see
Section 2, Chapter 4).
As discussed in Section 7.1, solution-soluble material may reach the electrode
surface by diffusion, migration or convection. In cyclic voltammetry at a stationary electrode, and assuming that migration can be neglected in the presence
of excess supporting electrolyte, diffusion is the sole form of mass transport.
However, additional material may be transported to the electrode by convection. Techniques where convection is a dominant form of mass transport, are
described under the heading 'Hydrodynamic Voltammetry'.
Hydrodynamic voltammetric methods have major advantages associated with
being steady-state techniques (Section 4.2). As a consequence, it is easy to
measure these classes of voltammograms as a function of the relevant convective parameter (flow rate, electrode angular velocity, etc.) in the absence
of significant problems arising from background capacitive charging current
(Section 6.1.2).
The potential profile associated with hydrodynamic techniques usually takes
the form of a linear or staircase sweep over the potential range in which the
oxidation or reduction processes of interest occur. As for cyclic voltammetry,
the gradient of the ramp represents the scan rate. However, the scan rate used
must be sufficiently slow to ensure that steady-state (within experimental error)
is attained at every potential during the course of the voltammetric scan. The
upper value of the scan rate that may be used under the steady-state regime is
therefore restricted by the rate of convective mass transport of material to the
electrode surface. Clearly, the faster the rate of convective mass transport, the
faster the scan rate that may be used to generate data that are consistent with
the existence of steady-state conditions.
With hydrodynamic voltammetry, it is the time parameter associated with
the rate of convection that is critical in the examination of the heterogeneous
and homogeneous kinetics associated with electrode reaction mechanisms. This
term plays a role analogous to that of the scan rate in cyclic voltammetric experi m e n t ~The
. ~ importance of this time parameter can be seen by examining the
' ~ i m eis implicit in hydrodynamic techniques whereas it is explicit in transient cyclic
voltammetry.
Hydrodynamic voltammetry
Slow convective transport
Fast convective transport
f
A
k
B---+
69
f Intermediates drawn
C
to electrode
\ Electrode
2.22 Schematic representation of an ECE reaction mechanism at a hydrodynamic electrode.
~ e ~ r o d u c by
e d courtesy: Adv.Phys. 0%.Chem. 32 (1999) 1. Copyright, Academic Press.
articular ECE mechanistic sequence given in eqns (2.24a-c). This mechanism
probed in hydrodynamic voltammetry by examining the effective number
of electrons transferred, NeE, as a function of mass transport of material to the
ode. NeE, which will vary between one and two for this form of ECE
anism, gives an indication of the competition between the loss of the
intermediates into the bulk solution and the second heterogeneous electrontransfer step. For rapid rates of convective mass transport, NeKtends to a value
of one, because the intermediates B and C are swept away from the electrode
into bulk solution before the second E step can occur (see Figs 2.15 and 2.22).
n contrast, at very low rates of mass transport, NeAtends to two, as B and C
remain in the vicinity of the electrode for sufficient time to allow C to undergo
an electron-transfer process at the electrode sudace (Fig. 2.22). Thus, in an
EGE process, the homogeneous kinetic process competes with mass transport
aterial to and from the electrode.
order to fully probe the kinetics of the C step in an ECE process of the kind
escribed above, the voltammetric response must be measured over a sufficiently
e range of mass transport rates so that NeEvaries between one and two. For
icularly rapid processes, this requirement implies that very fast rates of mass
transport are required in order to avoid Neg being equal to two at all transport
rates. Conversely for slow reactions, low rates of mass transport will be required
achieve significant deviations from Neff equalling one. Consequently, it can
appreciated that it is a study of the competition between the rates of mass
transport and chemical kinetics that leads to the quantitative determination of
electrode reaction mechanisms in hydrodynamic voltammetry. Importantly,
for each hydrodynamic technique, there is one assessable convective transport
arameter that directly relates to the kinetic time-scale.
.1 Rotating-disc electrode voltammetry
he rotating-disc or RDE consists of a disc electrode, made from a suitable
working electrode material [23,39,40], surrounded by a non-conducting sheath.
The complete electrode assembly is constructed so that the sheath and electrode
70
Principles of voltammetry
I
PTFE sheath
Fig. 2.23 Schematic representation of a rotating-disc electrode. Reproduced by courtesy: Adv.
Phys. 0%. Chem. 32 (1999) 1. Copyright, Academic Press.
Flow profile from below
r=O
Flow profile from side
Fig. 2.24 Convective flow profile associated with a rotating-disc electrode. Reproduced by
courtesy: Adv. Phys. Org. Chem. 32 (1999) 1. Copyright, Academic Press.
are flush (Fig. 2.23). Typically, the disc electrode faces downwards into solution
and is rotated around an axis perpendicular to and through the centre of the disc.
Under these conditions, a well-defined flow pattern distribution is established,
as illustrated in Fig. 2.24; in effect, solution is sucked towards the electrode and
then flung outwards.
The following experimental conditions must be met in order to ensure compliance with the well-established theory for this technique under steady-state
conditions:
(1) The electrode rotates in a single plane perpendicular to the axis of rotation.
(2) The frequency of rotation is stable with respect to the time required to
conduct a voltammetric scan.
(3) The electrode rotation frequency is sufficiently low to prevent localized
turbulent flow. This typically means that the rotation rate is in the range of
4-50 HZ.
Hydrodynamic voltammetry
71
The crucial parameter which controls the time-scale over which electrode
reactions are examined at a RDE is the electrode angular velocity, o(rad s-I),
\which is related to the rotation frequency, f (Hz), by o = 2nf. The related
rotating ring disc electrode is discussed in Section 4.4.1 of Chapter 5.
Channel electrodes
Figure 2.25 shows a schematic diagram of a channel electrode, which consists
of an electrode embedded in the wall of a rectangular duct through which
solution is made to flow under well-defined laminar steady-state conditions
[4O,4l].
In channel electrode voltammetry, the flow is treated as being twoimensional in the x-y plane. O n entering the channel, the solution velocity
profile is essentially plug flow. However, the effect of friction at the walls causes
retardation of the solution flow in the x-direction (Fig. 2.26).
After a distance 1, (the entry length) from the entrance, the hydrodynamic
layers from each wall merge, and the flow regime established is laminar in
form [42,43] in which separate layers (laminae) of solution have characteristic
velocities reaching a maximum (Vo)at the centre of the channel. The parabolic
ape of the ultimate velocity profile is therefore given by the following velocity
. 2.25 Schematic representation of a channel electrode. Reproduced by courtesy: Adv. Phys.
0%.
Chem. 32 (1999) 1. Copyright, Academic Press.
Growing boundary layers
Inviscid
flow
A
Entry length
(1,)
Boundary layers merge
/
Fully developed laminar
flow over electrode surface
2.26 Convective flow profile associated with a channel electrode. Reproduced by courtesy:
Adv.Phys. 0%.Chem. 32 (1999) 1. Copyright, Academic Press.
72
Principles of voltamrnetry
components
where V,, Vy,
and V, are the velocities in the x,y, and r directions respectively,
Vo is the velocity of the solution in the centre of the channel, h is half the cell
height and d is the width of the cell.
The velocity of the. solution in the centre of the channel is related to the
solution flow rate, Vf, by the expression:
It is this flow rate, Vf, or centre line velocity, Vo, parameter that is critical in
determining the time-domain over which chemical processes can be monitored.
Generally in order to develop laminar flow over the electrode, 1, should be
sufficiently long so that ideally
where Re is the Reynolds number which in terms of the kinematic viscosity v k
is defined as:
Laminar flow is generally achieved for Re < 2000.
Laminar flow is obtained either by a gravity-fed flow system [24] as shown
in Fig. 2.27 or by a pressurized flow system. In the gravity-fed flow system,
the inlet reservoir is at a greater height than the outlet and the solution may
flow through a variety of capillaries of differing diameter. Thus, the flow rate is
determined by the difference in the height of the solution in the inlet reservoir
and the level of the outlet tip as well as the diameter of the capillary through
which the solution flows. For typical (Fig. 2.25) cell heights (2h) of 0.04 cm
and cell widths (d) of 0.6 cm, flow-rates in the range IO-"-IO-' cm3 s-' are
readily attainable using this apparatus. A pressurized flow system [44,45] has
been recently developed which is designed to force solution through a flow
cell so as to induce much higher velocity gradients at the electrode surface.
Flow is achieved by applying a large pressure at the inlet reservoir end of the
system, while maintaining the outlet at atmospheric pressure, otherwise the
design is similar to the gravity-fed system. Values of Vo as large as 75 m s-I
can be obtained using this apparatus (h = 0.01 cm, d = 0.2 cm), allowing the
determination of first-order homogeneous kinetic parameters as high as lo5 s-'
from steady-state measurements [44].
A major advantage of the channel electrode technique, in addition to enabling
very fast reactions to be studied, is its ability to be utilized in conjunction with
other techniques (e.g. photochemical or spectroscopic). This advantage has
Hydrodynamic voltammetry
73
Counter
electrode
Flow
I
Channel
unit
Capillaries
Reference
II
7 Schematic representation of a gravity-fed channel electrode flow system. Reproduced by
courtesy: Adv. Phys. Og. Chem. 32 (1999) 1. Copyright, Academic Press.
been utilized effectively in the investigation of the photoelectrochemistry of
(co),(q2-dpm)~1 in acetonitrile [46]. As discussed in Section 8.1-4, in
ence of irradiation, this complex undergoes a one-electron oxidation
process forming the fac- [Mn(C0)3(q2-dpm)Cl]+ cation which then isomerizes
l ] + Upon irradiation of the channel
to the u n e v - [ ~ n ( ~ O ) , ( q ~ - d p m ) ~cation.
ectrode surface by 385-nm light, a new oxidative feature is seen at a less
positive potential than obtained for the parent oxidation wave (Fig. 2.28). The
half-wave potential for this new photoproduct corresponds to that expected for
ation of m e r - ~ n ( ~ O ) , ( q ~ - d ~ mConsequently,
)Cl.
it is logically postulated
the electrode reaction mechanism in the presence of light is CE in nature,
the C step involving the photo-isomerization of the fac species
light
C
fac +mer
E
mer ;
i
mer+
(2.47a)
+ e-
(2.47b)
A value of 0.07 s-' (light intensity = 40 mW cmF2)was obtained for the firstorder rate constant associated with the C step by examining the flow rate
ependence of the limiting current associated with the mer oxidation wave [46].
74
Principles of voltammetry
E/ V (vs SCE)
Fig. 2 -28 Channel electrode voltammogram for oxidation of 1.42 r n f a~c - ~ n(co)?
( ~ ~ - d c~1 m )
in acetonitrile (0.1M Bu4NC104)while being irradiated with light at 390 nm. Solution flow rate is
1o - ~cm3 s-' . Adapted from: J. Phyi. Chem. 97 (1993) 1601.
(4
Reference
electrode Elytrode
,
Electrode
\
Solution in
Fig. 2.29 (a) Schematic representation of a wall jet electrode, and (b) convective flow profile
Chem. 32 (1999) 1.
associated with a wall jet electrode. Reproduced by courtesy: Adv. Phys. 0%.
Copyright, Academic Press.
9.3
Wall-jet electrodes
In the wall-jet electrode configuration, a high, fixed-velocity jet of fluid is
fired through a nozzle of diameter, a, directly towards the middle of a disc
electrode (radius = rl), whose centre coincides with that of the nozzle as
shown in Fig. 2.29. The solution thus impinges upon the electrode surface and
is circulated outwards towards the extremities of the electrode surface, but the
recirculated solution can never reach the electrode a second time.
Hydrodynamic voltammetry
75
El 12
E (vs arb. reference electrode)
Voltammogram obtained for a simple reversible one-electron transfer process at a hydrodynamic electrode. Reproduced by courtesy: Adv. Phys. Og. Chem. 32 (1999) 1. Copyright,
Academic Press.
A survey of the use of the theory of hydrodynamic voltammetvy
n all forms of hydrodynamic voltammetry, electroactive material reaches the
electrode via diffusion and conve~tion.'~
In the cases of the RDE and channel electrode configurations under steady-state conditions, solutions to the
mass transport equations are combined with the Nernst equation to obtain
the reversible response shown in Fig. 2.30. A sigmoidal-shaped voltammogram
is obtained in contrast to the peak-shaped voltammetric response obtained in
cyclic voltammetry.
An overview of the theory for hydrodynamic voltammetry is available in
nce [9] where it is shown that there are two critical parameters that are
itatively important in this and other forms of steady-state voltammetry:
(1) The limiting or mass-transport-limited current, Jim At sufficiently positive
potentials where the mass-transport process associated with an oxidation process
is rate-determining, the current for a process of the kind A -+ B n e reaches a
fixed limiting current value (hi,) that is determined solely by the mass transport of
material to the electrode surface. Thus, the level of reversibility of the chargetransfer process is not relevant to the value of I&. Under these conditions,
material is continuously replenished at the electrode surface by convection, in
contrast to the situation in a cyclic voltammetric experiment where depletion
occurs and a peak-shaped response is observed. Table 2.3 gives the analytically
rived expressions for the limiting currents obtained for a process of the kind
-+ B ne- at the three electrode types discussed in this section.
(2) The halfwave potential, Elp For a reversible redox couple in which
the oxidized and reduced species have very similar diffusion coefficients, the
+
+
l0Nligration is assumed to be negligible if the usual excess supporting electrolyte is present.
Principles of voltarnrnetr),
Table 2.3 Expressions for the limiting current obtained for the process
A -+B ne- at a range of hydrodynamic electrodes
+
Electrode
Convective flow parameter
Expression for limiting current, hi,
RDE
ChE
Angular rotational velocity, w
Volume flow rate, Vf
0 . 6 2 n ~ ~ v,D
'l6 ~ [I ~ ~
WJE
Volume flow rate, Vf
1 . 5 9 n ~ ~ ~ / ~ v ; ~ ~ ~ [ ~ L] ~~ a: -~ ~ ~/ ~ 1 1 3 / ~
]
~
w
~
F = Faraday constant; A = electrode area; D = diffusion coefficient; w angular rotational
velocity; [Ale = bulk concentration of A; vk = kinematic viscosity; Vf = volume flow
rate; w =width of channel electrode (ChE); x, =length of electrode in the direction of
flow; h = channel flow-cell half-height; d = channel flow width; a = nozzle diameter of
wall jet electrode; rl = radius of wall jet electrode.
n
- slope:
~Kv
intercept: E l I 2
1
-40
Potential (mV)
~
~
0.925n~~~/~[A]~w(~~x~/h~d)'/~
8
-20
1
I
0
I
20
I
,
40
Potential (mV)
Fig. 2.31 'Log plot7form of analysis of a steady-state voltammogram for a reversible process at 25OC.
half-wave potential approximates to the formal electrode potential E ~ O . For a
reversible electron-transfer process, Ellz will not vary with the rate of convective mass transport to the electrode surface. Additionally, for a reversible
( ~I )~/ ,I ] ,the so-called 'log
electron-transfer process a plot of E versus 1 0 g ~ ~ [ plot', will have a slope of 2.303RTlnF or 59/nmV at 25°C. Figure 2.31 contains an example of a 'log plot9analysis of a reversible steady-state voltammogram
at 25°C. It follows, from analysis of the equation for a reversible process (see
Fig. 2.31), that the difference in the quartile potentials (E314- Ell4)is given by
the very simple expression 2.303RTlnF log 9 or 56/n mV at 25°C. Measurement of E3I4therefore provides an alternative form of (two points) data
analysis to the 'log plot9 (many data points) analysis for the determination of the
reversibility or otherwise of a hydrodynamic voltammogram.
As for all voltammetric techniques, sluggish electron-transfer kinetics require
the application of an additional potential (overpotential) to drive the electrontransfer process at the same rate as for the equivalent reversible process. Thus,
the observed voltammogram is broadened relative to that found for a reversible
Hydrodynamic voltammety,
77
E (vs arb. reference electrode)
Comparison of voltammograms obtained for a reversible and irreversible one-electron
oxidation process at a hydrodynamic electrode having equal half-wave potentials. Reproduced by
courtesy: Adv.Phys. 0%.Chem. 32 (1999) 1. Copyright, Academic Press.
process (see Fig. 2.32 where an irreversible and a reversible process with idenave potentials are shown for comparison). However, note that, as
, the limiting current is identical to that observed for the reversible
ed a potential sufficiently positive (oxidation) or negative (reducspect to the half-wave potential is applied. It should be noted that,
as in the case of homogeneous kinetics, competition exists between heteroectron transfer and transport of material to and from the electrode
nce, as the rate of convective mass transport increases, an initially
ectron-transfer process may become quasi-reversible and, finally, at
very high rates of mass transport, irreversible. Thus, a knowledge of the theoretical and experimental dependence of Ellz and/or hi, on the convective
time-scale enables a quantitative account of an electrode process to be given.
is is entirely analogous to the dependence of the cyclic voltammetric response
on scan rate (Section 7.4.1).
A number of studies have been conducted using fast rates of convective mass
transport in order to probe the kinetics of heterogeneous processes [45,47,48].
or example [45] the heterogeneous kinetics for the reduction of benzoquinone
at platinum electrodes in acetonitrile solutions was probed by examining the
variation of the mass transport in a channel electrode as a function of the flow
rate (over the range 10-~-3.5 cm3 s-l) of solution over the electrode surface.
A value of 0.30 cm s-' was obtained for the standard heterogeneous rate constant, which is in good agreement with the value obtained using the transient
cyclic voltammetric technique [B].
As discussed in Section 6, solution resistance and capacitive current can
lay a significant role in voltammetry. Due to the typically low scan rates
and steady-state conditions employed, capacitive charging presents a relatively
small problem in hydrodynamic techniques. However, solution resistance affects
the appearance of the voltammogram in much the same way as a decrease
in the rate of heterogeneous electron transfer, as is the case in the transient
78
Principles of voltammetvy
cyclic voltammetric technique. That is, uncompensated resistance broadens
the voltammetric response because the additional IR, potential term is still
present in hydrodynamic voltammetry. Thus, great care has to be taken to
ensure that resistance artefacts are accounted for when assessing the kinetics of heterogeneous processes using hydrodynamic and, indeed, all other
voltammetric techniques. However, note that the limiting current value
measured in steady-state techniques, unlike the peak current in transient
techniques, is unaffected by the solution resistance or slow electron-transfer
kinetics.
As noted earlier, varying the rate of convective mass transport of material to the electrode surface allows the elucidation of reaction mechanisms via
monitoring of the dependence of mass transport on a particular experimental parameter (e.g. N& for an ECE mechanism). As usual, comparison of the
experimental result with theory derived for a particular reaction mechanism,
provides quantitative detail concerning the kinetics.
Microelectrodes (also referred to as ultramicroelectrodes) are, as the name
implies, tiny electrodes which possess at least one dimension that is sufficiently
small so that the mass transport regime is a function of its size under steady-state
conditions [49]. In practice, the upper limit of this small dimension is approximately 20 pm because with larger sizes, natural convection is likely to occur
under very slow scan rate conditions required to achieve a steady-state voltammogram (Section 8). This effect is likely to interfere with the interpretation of
results, which are assumed to be made under diffusion-only conditions. The
lower limit is in the nm range, which approaches that associated with molecular
dimensions.
Attractive features of microelectrodes operating under steady-state conditions, relative to conventionally sized macrodisc electrodes and transient
conditions, include increased current density, reduced charging currents, and
reduced ohmic drop [7,49]. The latter feature permits experiments to be conducted in highly resistive media, particularly, for example, in non-polar solvents
[7] or solutions containing an absence or near absence of added supporting electrolyte [7]. Further aspects of the exploitation of these advantages of
microelectrode voltammetry under near steady-state conditions is contained in
reference 7.
Microelectrodes exist in a variety of geometries, the most important of which
is the microdisc electrode. Microbands, cylinders and rings are other possibilities
while the microsphere or hemisphere is often used to aid theoretical development since the rate of mass transport is invariant over the electrode surface.
Some of the different possibilities'for microelectrode designs are illustrated in
Fig. 2.33.
Voltammetric studies at microelectrodes
79
Schematic diagram showing microdisc (a), band (b), hemi-cylinder (c) and ring (d) electrodes. Reproduced by courtesy: Adv. Phys. Org. Chem. 32 (1999) 1. Copyright, Academic Press.
A schematic diagram of a rnicroband electrode. The arrows represent the directions of
Fi
diffusion to the electrode. Reproduced by courtesy: Adv.Phys. 0%.Chem. 32 (1999) 1. Copyright,
Academic Press.
1
rinciples of the theory ofrnicroelectrode voltarnrnetry
ionally, macrodisc electrodes operating under transient, linear,
n-only conditions are characterized by one-dimensional mass transport
h diffusion takes place normal to the electrode surface (Section 8.1).
In contrast, with microelectrodes, such mass transport is found only at short
times after electrolysis is initiated when the diffusion layer is small compared to
the shortest dimension of the microelectrode. At longer times, the rate of mass
transport varies locally over the electrode surface with the edges of microdisc
and line electrodes, for example, receiving a greater current density owing to
ossibility of convergent or radial diffusion, as illustrated in Fig. 2.34 for the
ular case of a microband electrode. It is this additional mass transport by
radial diffusion which leads to the significantly enhanced mass transport under
steady-state conditions.
or the case of a microdisc electrode of radius re, radial (convergent) diffusion
to a steady-state limiting current given [7,49] by eqn (2.48)
is equation shows that the limiting current scales with the electrode radius,
rather than area, thereby reflecting the non-uniformity of the current density.
Princiyles of voltammetry
The transition between linear (shorter times) and radial (longer times) diffusion is revealed if the equation relating to the current following a potential step
at a microhemisphere electrode is considered:
where A is the electrode area (A = 2 n r 3 , and t is the time. At short times,
the second term dominates and the mass transport is Cottrellian [5,6] so it
depends on t-'l2 as expected from linear diffusion, while at longer times,
when radial diffusion is dominant, the current tends to the steady-state (timeindependent) value predicted by the first term. However, note that microsphere
and microhemisphere electrodes are atypical microelectrodes in that their very
high symmetry dictates that the current density is uniform over the entire electrode surface. This is not the case for the microdisc, microband or microring
electrodes shown in Fig. 2.33 as edge mass transport dominates.
Implicit in the above discussion is the concept that current-voltage curves
measured at microelectrodes at sufficiently slow scan rates of potential are
characterized by a mass-transport-limited current plateau rather than a peak
current as in linear sweep voltammetry at a planar macrodisc electrode.
Figure 2.35(a) shows a typical microelectrode voltammogram for an electrochemically reversible system under near steady-state conditions. Of course as
the scan rate is increased, the voltammetric behaviour at a microelectrode converges to that expected for linear diffusion (Fig. 2.35 (b)), until at sufficiently fast
scan rates a characteristic transient-type cyclic voltammetric response is obtained
when the mass transport is predominantly governed by linear diffusion.
The increased rate of mass transport associated with the shrinking of the
electrode size means that electrode processes which appear electrochemically
reversible at slow scan rates with large electrodes may show quasi-reversible
or irreversible electrode kinetics when examined under steady-state conditions
at a microelectrode. However, the use of very fast scan rates may also enable
departures from reversibility to be detected with transient mode microelectrode methods, provided adequate correction for charging current and resistance
effects can be made under the latter conditions. The use of microelectrodes,
therefore, represents an opportunity to apply two very powerful approaches to
the determination of fast heterogeneous electrode kinetics, and rate constants in
excess of 1 cm s-' have been reported via use of both the slow scan rate (radial
diffusion) steady-state and very fast scan rate (linear diffusion) microelectrode
voltammetric approaches [49].
The above discussion implies that chemical information may be extracted
from microelectrode experiments either via steady-state measurements obtained
under slow scan rate conditions, or via transient, often cyclic, voltammetric
approaches utilizing fast scan rate techniques. In the former approach, measurements of the mass-transport-limited current are made as a function of the
electrode size, which is the electrode radius for the case of a microdisc electrode. This aspect of quantitative evaluation of steady-state microelectrode data
may be illustrated by reference to the ECE mechanism (eqn sequence 2.24a-c)
Voltammetric studies at microelectrodes
81
Potential
.35 Diagrams showing current-voltage curves measured at a microdisc electrode at scan rates
corresponding to the limits of (a) radial diffusion and (b) linear diffusion. Reproduced by courtesy:
Adv. Phys. Org. Chem. 32 (1999) 1. Copyright, Academic Press.
k is the first order rate constant for the C step. Kinetic and mechanistic
ation for this mechanism may be gleaned in an analogous manner to the
, channel, or wall-jet electrode steady-state techniques, through examination of the effective number of electrons transferred, N g , but in the case
of steady-state microelectrode voltammetry by studying NeKas a function of
microdisc radius. Figure 2.36 shows that Nen varies between 1 and 2 as the
value of the electrode radius increases. The former limit, NeE= 1, corresponds
to the case of fast mass transport (small radius) where B is lost to bulk solution
before it can be transformed into C while the latter, NeE= 2, case is reached
for slow mass transport (large radius) where B is nearly completely transformed
into C near the electrode surface. The parallel of steady-state microelectrode
voltarnrnetry to hydrodynamic voltamrnetry, where rotation rate, rather than
electrode radius, is the time-dependent variable (Section 9) is obvious.
The dimensionless parameter
ictates [49] whether the kinetics of the homogeneous first-order chemical
reaction in the ECE mechanism are 'fast' or 'slow'. The term r : / ~gives an
approximate measure of the time taken to move out ofthe diffusion layer of the
82
Principles of voltarnrnetry
Fig. 2.36 A working curve showing the relationship between Neffand the dimensionless parameter
( k r : / ~ )for an ECE mechanism. Reproduced by courtesy: Adv. Phys. 0%.Chern. 32 (1999) 1.
Copyright, Academic Press.
microdisc, while the term ( k - l ) corresponds to the time taken for appreciable
amounts of B to transform into C.
Examination of the theoretically generated working curves for an E C type
mechanism, analogous to the curves given in Fig. 2.36 for the E C E mechanism, suggest that for microdiscs of radii 1-10 pm, lifetimes of B in the range
0.1-100 ms should be amenable to quantitative study under steady-state conditions. A detailed summary of steady-state microelectrode voltammetry, and
its application to studies on homogeneous chemistry, is given in the review by
Montenegro [49].
As noted above, the dynamic range of microelectrode voltammetric experiments may be extended by the use of very fast scan rate techniques to give
transient conditions where the mass transport approximates that seen for conventional cyclic voltammetry. However, caution in interpretation of the results
is advisable under the theoretically 'tricky', intermediate conditions, where significant contributions from both linear and radial diffusion apply. For very fast
scan rate experiments, the major advantage in using a microelectrode resides
in its intrinsically small area which leads to a correspondingly reduced capacitance of the electrode-solution interface associated with its double layer. As
noted in Section 6.1.2, it is the current associated with the capacitance which
gives an upper limit to the scan rates accessible, if a Faradaic signal (arising from electron transfer between the electrode and solution-phase species)
is not to be masked. The capacitance scales directly with the electrode area
(eqn 2.10) so that there is a clear advantage in employing smaller-sized microelectrodes for both transient and steady-state measurements. Clearly, it also
follows that the R,Cr cell time constant is lowered by the use of micro rather
than macrodisc electrodes, which also leads to improved voltammetric performance, as does the decreased IR, (Ohmic) potential drop obtained with the
smaller electrode size.
Semi-integration and semi-d%ferentiation
83
u
0
-1
E/ V (vs SCE)
-37 Voltammograms (with background correction) for the reduction of 10rnM
2,6-diphenylpyrylium perchlorate in acetonitrile at a 5-pm radius platinum disc electrode using
scan rates of: (a) 250 (b) 150 and (c) 75 kV s-'. Adapted from J. Electroanal. Chenz. 324 (1992) 33.
n practice, useful measurements can be made with fast scan rate cyclic
voltammetry under transient conditions at microdisc electrodes to probe
lifetimes of unstable species which approach the nanosecond regime. This
requires the use of potential scan rates in the kV s-' or greater range. An elegant
example of a very fast scan rate study concerns the electro-reduction of the
2,6-diphenylpyrylium cation in acetonitrile solution [50] using a 5-pm radius
platinum microdisc electrode. Electrochemically, the cation is reversibly converted into the radical via a one-electron reductive charge-transfer process.
The resulting radical can subsequently dimerize in an irreversible second order
homogeneous reaction to form a species which displays no electroactivity
within the potential range examined. Figure 2.37 displays cyclic voltammograms measured at scan rates between 75 and 250 kV s-'. At the fastest scan rate
(250 kV s-') examined, re-oxidation of the radical is observed on the return
scan. This chemical reversibility is progressively lost as the sweep time becomes
comparable with the time taken for the radical to dimerize. Interpretation of
eak current data in terms of an EC2 mechanism, where C2 is ;he notation
for a second-order reaction, permitted the dimerization rate constant to
be evaluated as 2.5 x lo9M-' s-' which corresponds to a half-life of 20-50 ns
for the cation radical under the conditions studied.
It was shown in Section 10 that microdisc electrode voltammetry enables steadystate, time-independent I-E curves to be obtained at slow scan rates, while
transient voltammograms can be observed at fast scan rates. That is, the 'time
endence' can be introduced by increasing the scan rate. The converse opern of removing the 'time dependence' of a transient voltammogram may be
hematically by use of semi-integration, which is a form of fractional
84
Principles of voltammetry
11.1 Some valuable properties of the semi-integral
When applied to a voltammetric experiment, semi-integration with respect
to time, produces a quantity M(t) which has properties that are intermediate
between those of the charge Q(t), obtained by the usual form of integration,
and I(t) itself: Via this mathematical operation, and for a diffusion (linear)
controlled process at a macrodisc electrode, the time (t-'I2) dependence is
effectively removed from the transient voltammogram so that a sigmoidalshaped voltammogram, having many of the characteristics of a steady-state
response, is obtained. Figure 2.38 shows the application of semi-integration
to a reversible process, while Fig. 2.39 represents an example of applying the
same mathematical processes to a quasi-reversible cyclic voltammogram.
A valuable property of the semi-integral is that M(t) can be a unique function
of the applied potential E(t) and, hence, be totally independent of the history
by which the electrode arrives at E(t). Therefore, even the actual voltammetric
technique employed to obtain the semi-integral is irrelevant as are, for example,
the time or scan rate dependence of the cyclic voltammetric experimental data
from which the semi-integral is derived. This valuable property is achieved
when the following conditions are satisfied:
(1) The voltammogram starts from a potential where the current flowing is
effectively zero.
(2) The electrode is planar, or effectively so.
(3) Mass transport is effectively semi-infinite.
Fig. 2.38 A reversible cyclic voltammogram obtained at a macrodisc electrode with mass transport
by linear diffusion presented in the conventional form (a), and after semi-integration (b). Note the
complete overlap of forward and reverse scans in the serni-integral form of presentation. Adapted
from: Anal. Chem. 72 (2000) 3492.
Semi-integvation and semi-d@eventiation
r
+
Theoretical cyclic voltammogram for a quasi-reversible electrode process A + B neFi
emi-integral as a function of potential. In this dimensionless form of presentation
an
k0/2/=
= 0.145, a = nFv/RT, D is the common diffusion coefficient (DA = DB) and a
is assumed to be 0.5. In the semi-integral, M ( t ) and M(t),,, are the semi-integral and maximum
value of the semi-integral at time t. (at) is defined in Anal. Chem. 37 (1965) 1351 and J. Phys.
(1972) 1160. Adapted from: Anal. Chem. 45 (1973) 1298.
x
(4) Convection and migration are absent, so that transport is solely diffusive.
(5) The electrode reaction is of the kind A F+ B ne-.
(6) The electron process generates a solution-soluble product.
+
e electron transfer is uncomplicated by preceding, succeeding, or
adsorption reactions.
n-Faradaic current is absent, or has been corrected for.
compensated resistance is absent.
t is this 'uniqueness' of M ( t )that has led to powedul methods of 'global analysis' of reversible or quasi-reversible cyclic voltammograms and correction for
uncompensated resistance. The theoretical aspects of the use of semi-integration
the useful properties of the transformed data for calculation of E:, D, ko, and
cx or the potential dependence of a,associated with reversible or quasi-reversible
rocesses are available in references [52-601.
Measurement of uncompensated resistance by semi-integration1'
e simplification of data analysis achieved by semi-integration is readily made
evident by analysis of the effect of uncompensated resistance on a reversible
cyclic voltammogram. The one-to-one relation between M ( t ) and E(t) in the
"kidaped with permission from Anal. Chem. 72 (2000) 3492. Copyright, American
Chemical Society.
86
Principles of voltammetry
final condition listed above in Section 11.1 holds, even in the presence of
uncompensated resistance (R,), when E(t) is interpreted as the true potential
Etr,,(t) experienced by the electrode, rather than as the potential Eapp(t)that is
applied to the cell. The relationship between these two potentials is
Figure 2.40 shows the effect of uncompensated resistance on the cyclic voltammogram (Fig. 2.40(a)) and its semi-integral (Fig. 2.4O(b)). Figure 2.40 is scaled
exactly as Fig. 2.38, so that the differences between the two arise solely from the
presence of resistance. The abscissa in Fig. 2.40 is Eapp(t),though it is labeled
with an unsubscripted E ( t ) ,as is common practice. Notice that in the presence
of uncompensated R,, the two branches of the semi-integrated voltammogram
no longer overlap, the backward branch lying above the forward branch.
Consider the two points labeled m2 and ml in Fig. 2.41. ml is an arbitrary
point on the forward branch of the M(t) versus E(t) curve, while m2 lies on the
backward branch at the same ordinate level. The two points thus correspond
to equal semi-integral values, M (tz) = M (tl). By the uniqueness property, this
means that the true electrode potentials must have been equal at the two instants,
tl and t2, to which the points correspond, so that Et,,(t2) = E,,,(tl). The gap
between the two semi-integral branches arises from the resistance present, as
comparison of Figs 2.38 and 2.40 attests. The equality of the true potentials, at
the instants t2 and tl, can be combined with the general eqn (2.51), with the
Fig. 2.40 A reversible cyclic voltammogram obtained at a macrodisc electrode in the presence of
uncompensated resistance with mass transport by linear diffusion presented in the conventional form
(a), and after semi-integration (b). Note the absence of complete overlap of forward and reverse
scans in the semi-integral form of presentation, unlike the situation prevailing in Fig. 2.38 where no
uncompensated resistance is present. Adapted from: Anal. Chern. 72 (2000) 3492.
Semi-integration and semi-dtferentiation
87
1 Use of semi-integration to calculate the uncompensated resistance assocated with a cyclic
voltammetric experiment. Adapted from: Anal. Chern. 72 (2000) 3492.
subscript 'app' omitted, to give
e terms I(tl)and I(t2)are the values ofthe currents at the points labeled il and
i2 on the cyclic voltammogram itself. Hence, measurements of the appropriate
rdinates of the four points ml , mz, il and i2, permits R, to be calculated.
ince ml was selected arbitrarily, it follows that an unlimited number of simialculations may be made, ideally resulting in the same R, value. Obviously,
the precision of the measurement will be greatest where the separation between
the branches of the semi-integral is widest, that is, in the vicinity of Ell2. O n the
s of precision, use of the mid-section of the semi-integral between (Ell4)
(E314)of the wave height is recommended. Clearly, once the value of Ru
is calculated, theoretical analysis of cyclic voltammograms can be undertaken
inclusion of the value as shown in Fig. 2.42 for oxidation of ferrocene in
igh resistance solvent dichloromethane [60].
viously, it is also possible to semi-differentiate a voltammogram to generate
a symmetric peak-shaped curve from an asymmetric reversible cyclic voltammogram or indeed use any other form of operator from the field of fractional
or conventional (first derivative etc.) calculus. Figure 2.43 contains an example
of the asymmetrical responses obtained both in normal (Fig. 2.43(a)) and first
88
Principles of voltammetry
Fig. 2.42 Voltammograms at 25OC for reversible oxidation of 0.7 rnM ferrocene at a 0.5 mm
radius platinum disc electrode in dichloromethane (0.1 M Bu4NPF6)using a scan rate of 1V s-' ;
(a) cyclic voltammogram; (b) semi-integral of voltarnmetric data; (c) calculation of R, using data in
(b) and eqn (2.53); (d) comparison of cyclic voltammogram simulated with R, = 0 ( . . . ), with
R, = 2500 ohm as calculated from data in curve (c) (- - - -) and the experimental curve (-).
Adapted from: Anal. Chem. 72 (2000) 3492.
derivative (Fig. 2.4307)) cyclic voltammetry which may be compared with the
symmetrical semi-differential voltarnmogram (Fig. 2.43(c)).Additional details of
the mathematics and other advantageous features of semi-integration and semidifferentiation (often called convolution voltammetry) are available in references
[4-6,13,52-601.
12 General features associate
voltammetric experiments
e modelling of
The difficult part of deducing an unknown mechanism from a voltammetric
experiment is quantitatively extracting the chemical information from the I-E-t
curve. To do this, a model must be constructed to predict the current for a given
set of conditions and a postulated mechanism. An electrochemical model is
concerned with the concentration distributions of chemical species (and possibly
the potential distribution). If the concentration distributions of all the chemical
species dissolved in the solution phase can be simulated, the current flowing
at the 'working electrode' may be calculated by integrating the concentration
gradient zones at the electrode surface to gwe the total flux.
12.1 Inforrmatio required to solve voltammetric theory
Three pieces of information are required to define the experimental system:
Kinetics The rate of electron transfer and its potential dependence can be
described by the Butler-Volmer or Marcus-Hush relationships, with the former
Modelling of voltammetric experiments
89
Comparison of (a) normal, (b) first derivative and (c) semi-differential forms of presentaFi
tion of a cyclic voltammogram for a reversible one-electron oxidation process. Adapted from P.H.
nieger, Electrochemistry, 2nd edn, Chapman and Hall, New York, 1994.
being adequate on most occasions for standard voltammetric studies when the
electroactive species is soluble in solution (Section 7.2). An electron transfer
often initiates a cascade of homogeneous chemical reactions by producing a
reactive product which could be a radical, radical cation, radical anion, or inorganic compound in an unusual oxidation state. Each step in the mechanism can
be described mathematically by a rate equation, to form another part of the
electrochemical model. The rate law for the overall sequence is probed by the
voltammetric experiment.
transport Each of the three mass transport components may be described
matically, as discussed in Section 7. The effect of all three modes of mass
rt may be summed giving the partial differential equation (PDE)
re CA is the normalized concentration of species A (CA = [A]/[AIo).
tion (2.54) is equivalent to eqn (2.19) with [A] being replaced by CA.
quation (2.54) describes how the normalized concentration of species A at
a given point varies with time due to diffusion, convection, and migration,
relative to the bulk concentration, [Ale. This mass transport problem is the
second major component of the electrochemical model and it depends only on
the electrode geometry, the nature of which defines the Laplace (V2)operator.
Assuming a background 'supporting' electrolyte is used in excess to eliminate
migration effects from the experiment, only a diffusion-convective equation is
necessary to describe the mass transport.
90
Principles of voltammetry
Experiment The experimental technique controls how the mass transport and
rate law are combined to form the overall material balance equation. Thus,
migration effects may be eliminated by addition of supporting electrolyte,
steady-state measurements eliminate the need to solve the equation in a timedependent manner and the addition of excess substrate can reduce the kinetics
associated with a second-order reaction to pseudo-first-order. The material balance equations (one for each species), with a given set of boundary conditions
and parameters (electrode type, cell dimensions, flow rate, rate constants etc.)
define the I-E-t surface traversed by the voltammetric technique.
If all three pieces of information are known, the concentration distributions
of the species throughout the course of the voltammetric experiment may be
described mathematically by a set of simultaneous PDEs. The way these equations are perturbed during the course of the voltammetric experiment, and the
boundary conditions required to solve them, may also be deduced from these
three pieces of information.
12.2 Methods used for soloing voltammetvic theory
The final component of the theoretical model is the development of a suitable mathematical method to solve the relevant system of simultaneous PDEs,
often as a function of time as the concentration distributions evolve during the
experiment. The difficulty of solving these systems depends on the complexity
of the material balance equations and as to how they are linked to each other
by the kinetic terms.
For a simple electron-transfer reaction (e.g. eqns (2.1) and (2.2)), it is often
possible to solve the diffusion-only limiting current relationship analytically
under steady-state conditions. For example, the analytically deduced limiting
current steady-state microdisc equation for a spherical electrode is given by
eqn (2.55)
Ilim
= 4nnFD[AIoy,
(2.55)
For hydrodynamic electrodes, in order to solve the diffusion-convective
equation analytically for the steady-state limiting current, it is necessary to use
a first-order approximation of the convection function(s) (such as the Lkvcque
approximation for the channel electrode). Approximate expressions for these
hydrodynamic steady-state mass-transport-limited currents, as noted previously,
are contained in Table 2.3.
For planar or spherical electrodes, and under conditions where the mass
transport is a diffusion function in only one dimension, it is possible under some
circumstances to solve, analytically, the diffusion equation for the reversible
process as a function of time to give, for example, the peak current expression
for a linear sweep voltammogram (eqn 2.34). It is also possible to solve the
material balance equations for the spherical electrode at steady-state for a few
first-order mechanisms [61]. In order to tackle problems involving secondorder homogeneous reactions or kinetics, or more complex mechanisms, or to
solve time-dependent equations or model geometries with more complex mass
Summary of principles of voltammetry
91
transport, it is necessary to resort to numerical methods, a vast number ofwhich
are available.
ver the past few years, several commercial electrochemical simulation packages have appeared. ' ~ h e s epackages are currently only capable of simdating
mass transport in one spatial dimension and are therefore restricted to modelling
vo1tammeG-y at large planar and rotating-disc macroelectrodes, or spherical &
hemispherical electrodes of any size (macro or micro). Speiser [62] has undertaken a thorough assessment of these simulation packages in his recent review.
Use of a commercially available package to simulate the cyclic voltammograms
for a complex reaction mechanism is given in Chapter 3.
or many mechanisms, the steady-state Ell? or NeKvalue is a function of
just one or two dimensionless parameters. If simulations are used to generate
the 'working curve' to a sufficiently high resolution, the experimental response
may be interpolated for intermediate values without the need for further simion. A free data analysis service has been set up [63] via the World Wide
b (http://physchem.ox.ac.uk:8000/wwwda.htrd) based on this method and
currently supports spherical, microdisc, rotating-disc, channel, and channel
microband electrode voltammetry for a range of mechanisms.
receding discussion in this chapter has been aimed at providing the reader
book with a broad understanding of the experimental and theoretical considerations related to the case of the techniques of voltammetry when
the electroactive species are soluble in the solution phase. It is now possible,
appropriate, to briefly summarize the general features that are inherent
11 voltammetric techniques so that general principles are emphasized and
techniques of voltammetry discussed above employ a form of electrolysis,
t an appropriate form of Faraday's law is coupled with the mass transport
ons to give the theoretical description of the voltammogram.
-1 Application ofFavadayJs law
Consider an electrode-solution interface as in Fig. 2.44 and then apply Faraday9s
aw to the electron flow (current) in the circuit for the electrode process
A
B ne-. At the electrode-solution interface [13]
+
92
Principles of voltammetry
electrode
solvent
(electrolyte)
electrode reaction: A+B+ne-
Fig. 2.44 Transport processes occurring at the electrode-solution (electrolyte) interface for the
reaction A + B
ne-. Provided by courtesy: K.B. Oldham, Trent University, Peterborough,
Canada.
+
-Ji
where
and J; are the flux densities (mol cm2s-') of A and B respectively
at the electrode surface and other symbols are as defined previously.
It follows from the equalities in eqn (2.57) that the flux density, and, therefore,
the mass transport ofA and B, are intimately related to current by Faraday's Law;
this is why it is essential to understand the nature of transport processes or how
species move to and from electrode surfaces in electrochemical systems.
13.2 A general approach to undevstanding a voltammetvic pvoblem
In trying to understand any voltammetric problem, the following matters related
to mass transport need to be taken into account:
(1) Which species present at the electrode-solution interface are transported?
(2) What kind of transport mechanism is important in the problem being
considered?
(3) What is the relative contribution of the different kinds of transport in the
problem being considered?
(4) How is the voltammetric current controlled by the transport?
(5) How is the size of the electrode related to the transport?
(6) How are the transport, and heterogeneous and homogeneous chemical
reactions combined?
13.2.1 identifying the species transported
Let us consider some typical electrode processes [13] and identifj. the chemical
species that must be transported by use of bold type and those that do not by
use of italics.
Example 1
The electrodeposition of metallic copper onto a solid electrode
Summary ofprinciples of voltammetry
93
Example 2 The electrodeposition of metallic copper into a mercury electrode
to form an amalgam.
Unlike example 1, metallic copper can now diffuse into the mercury electrode
and is therefore transported into the electronic conductor.
In this case both the c u 2 + and Cu+ ions are soluble in solution and are
Example 4
Ag(so1id)
+ C1- (solution) ---+
AgCl(so1id) + e-
In all these examples, and generally for all reactions, including those involving
attached to electrodes (see later), at least one ion must be transported in
to balance the charge with that of e-. Thus, the total number of species
orted is always 2 1.
The transport mechanisms
tion (ions only) occurs in response to a gradient of potential.
sion (ions and molecules) occurs in response to a gradient of concentration.
(3) Convection (ions and molecules) occurs in response to a gradient ofpressure.
.3 A comparison of transport mechanisms
(1) The migration current, which is applicable to charged species only, and
is generally not large in magnitude, can be modelled mathematically with
relative ease via the Nernst-Einstein equation and can be avoided by the
resence of excess supporting electrolyte.
iffusion, like the migration current, is generally small in magnitude, is
mathematically easy to model via use of Fick's Laws, and is always present.
(3) Convection may give rise to large currents, is mathematically difficult to
model, but can be avoided by eliminating solution motion.
or the RDE the general mass transport equation in the absence of
homogeneous chemical kinetics is given by
94
Princlplesofvoltammetry
where r, x and 6' are cylindrical polar coordinates (see Fig. 2.23), and v,, v,, and
v~ are the respective components of the solution velocity.
For the channel electrode, the mass transport equation also contains convection and diffusion in the absence of homogeneous chemical kinetics and is given
(see Fig. 2.25 and Section 9.2 for an explanation of symbols) by
Other voltammetric techniques have mass transport equations described above
or else they are available in references [5,6] or [9].
13.2.4 Facton that control the current in a voltammetric experiment
(1) Transport control by migration, diffusion, and convection.
(2) Control by the chemistry of the redox reaction which may be thermodynamic (reversible process obeying the Nernst equation on the voltammetric
time-scale within experimental error) or kinetic (heterogeneous or homogeneous rate laws).
(3) Mixed transport and chemical control (usually the case).
13.2.5 The (I-E) voltammetric cuwe
In any form of voltammetric experiment, the dependence of I (t) on E (t) needs
to be known, with I(t) being measured and E(t) being imposed (sometimes
vice versa). For example, in the simple potential step experiment shown in
Fig. 2.45(a), where A is converted to B, current control will occur solely by
diffusion when the rate of the A -+ B e- heterogeneous electron-transfer step
for the oxidation reaction is inherently so fast that it is not rate-determining.
This will be true at potentials in the limiting current region, where I(t) is
measured as in Fig. 2.45(b) (macrodisc electrode, transient case). Furthermore,
if A is not a charged species, it cannot be transported by migration, so that mass
+
Fig. 2.45 A potential step experiment (a) which causes the reaction A -+B
diffusion-controlled manner (b).
+ ne- to occur in a
Comparison of voltammetric techniques
95
transport occurs solely by diffusion. Under these conditions I (t) depends on the
electrode size and has a characteristic mass transport time A I D (A, electrode
area; D, diffusion coefficient). For a typical value of D
cm2s-l) and for a
rnacrodisc electrode having an area of 10 mm2
rV
A/D(macroelectrode)
10-I
cm2
--
lop5 cm2 s-l
% 20 min
which is a long experimental timescale. In contrast, for a microdisc electrode
of area 10 pm2
&ich is a relatively short time-scale for the mass transport process.
An analogous result based on consideration of time-scales also occurs
r conditions of cyclic voltammetry. Thus, under both the potential step
itions described above and when the current is changing rapidly with
potential (cyclic voltammetry) the current is often diffusion-controlled at a
macrodisc electrode, but under control of both the electron transfer, and
transport by diffusion at a microdisc electrode. For a macrodisc electrode,
the purely diffusive component of the experiment is given by the CottreIl
( n whereas for a microdisc electrode,
equation, I(t) = FAD^''^ [ ~ ] ~ /t)'I2,
the time-independent steady-state equation, Ili, = 4nFreDA[AIo,applies. If
a combination of electron transfer and diffusion control occur under linear
potential sweep conditions, the resultant voltammetric response is constructed
from the two terms that may control the experiment (Cottrell and steady-state
ressions), with the measured value of the current at a given potential being
roximately given by whichever of the possible current terms is smaller at
that potential, that is, whichever of the steps in the process is rate-determining.
n any quantitative evaluation of the mechanism and kinetics of an electrode
process, the experimental technique and electrode geometry must be selected
match the kinetic time-scales. The time domain over which a first-order
mogeneous chemical process occurs, and which needs to be matched by
the voltammetric experiment, for example, would be l l k , where k is a firster rate constant. As concluded in the preceding qualitative discussion, this
matching of electrochemical and chemical time domains is achieved by varying the rate of mass transport via convection, electrode size/shape, or potential
scan rate. The access to solution to theoretical models and solutions for all techniques, made possible by numerical mathematical methods (Section 12), enables
a quantitative comparison of the kinetic discrimination of various voltammetric
techniques to be developed.
96
Principles of voltammetry
14.1 A quantitative comparison of the kinetic discn'mination of
homogeneous reactions at common electrode geometries under
vokammetric steady-state conditions
Each of the electrode geometries and voltammetric techniques has a characteristic time-scale parameter (t,) which is a function of the time required to reach a
steady-state [9,61]. For example, for an EC or ECE mechanism, dimensionless
homogeneous rate constants (K) for a first-order reaction and dimensionless time
(r)both may be formulated [9] using a characteristic time-scale, as in eqn (2.62)
K = kt,;
where k is the first-order homogeneous rate constant and t is the time. The
dimensionless homogeneous rate constant, K , is the unique parameter on which
the steady-state NeEor Ell2-values depend. Thus, the so-called 'working curve',
deduced from numerical modelling of the process, generates a plot of Ell, or
NeEversus log K ,in order to completely define the steady-state behaviour for a
particular mechanism at a specific electrode geometry and voltammetric technique. For example, the 'working curve' consisting of a plot of Neg versus
loglo (kr;/D), shown in Fig. 2.36, represents the case for a microelectrode of
radius re under steady-state conditions for a species undergoing an ECE mechanism where the initial material in bulk solution has a diffusion coefficient D and
a first-order rate constant k. Thus, the term ~ $ Z / Dused in this plot is dimensionless, as required; here r;/D is the characteristic time-scale parameter. The equivalent parameters for other steady-state techniques are available in reference [9].
The recent availability of working curves for a range of common mechanisms at a number of electrode geometries [9,61], allows a broad quantitative
comparison of the kinetic discrimination of common electrode geometries to
be made for both first- and second-order homogeneous processes. For example,
Table 2.4 shows the approximate range of time-scales and rate constants (for
ECE and ECzE reactions) that may be measured under conditions of steadystate voltammetry at various electrode geometries [9,61]. The range of rate
constants given in this table has been calculated from the values of the relefrom
vant dimensionless rate constant which give values of 1.1 and 1.9 for
the 'working curve' for each geometry and are considered to represent suitable
thresholds between which a kinetic process is 'voltammetrically visible' [64].
The calculations undertaken to obtain the results contained in Table 2.4 are
based on the following (typical) assumptions related to the range of experimental
parameters available with the relevant techniques:
(1) All species present in the bulk solution have diffusion coefficients of 1 x
loH5cm2s-'. The concentration of the electroactive species present in the
bulk solution is 1 x lov6mol ~ m - The
~ . kinematic viscosity of the solvent
is in the range 1 x lop3-1 x
cm2 s-'.
(2) Microdisc electrodes of radii 0.6-70 pm are used for steady-state measurements without problems associated with natural convection. The upper limit
le 2.4 A comparison of the kinetic time-scales accessible with steady-state voltammetry using
common electrode configurationsa
Electrode
configuration
Range of accessible
time-scales (t,)
Range of
log [dimensionless
rate constant (K)]
Range of rate constants (k)
which can be measured
ECE
EC2E
ECE(S-')
E C 2 E (mol-' cm3 s-')
6 x lo3-2 x 1012
1 x lo4-6x 1012
6 x lo4-2 x 10''
4 x lo1-8 x 1012
Hemispherical
Microdisc
Rotating-disc
Wall-j et
400 ps-5 s
400 ps-5 s
0.2-1 0 s
1 ms-100 s
3.82
3.93
2.47
3.03
4.31
4.47
3.71
3.93
2 x 10-~-2 x lo5
6 x 1 0 - ~ - 7 x lo5
9 x lop2-1 x lo3
2 x l0-~-2 x lo5
Channel
Conventional
Microband
Fast flow
0.1-1 0 s
3 ms-0.5 s
10 ps-10 ms
2.45
2.75
2.45
3.43
3.47
3.43
1 x lop2-5 x lo2 1 x lo4-4 x lo9
2 x 10-'-1 x lo6 3 x lo5-4 x 10"
2 0 ~ 1 0 ~ - 4 ~ 1 10 ~~ 1 0 ~ - 3 ~ 1 0 ~ ~
aAdapted from Adv. Phys. 0%.Chem. 32 (1999) 1.
Principles of voltammetry
has been chosen as 70pm, above which natural convection becomes significant. Dimensionless rate constants for spherical and microdisc electrodes
were interpolated from the working curves of Alden and Compton [61].
Hemispherical electrodes are experimentally realized using hanging mercury
drops for macroelectrodes and mercury-coated microdisc electrodes of radii
(0.6-70 pm) for microhemispheres.
For the RDE, the operating range of rotation frequency is assumed to
lie between 1 and 50 Hz and a typical electrode radius is 0.25 cm. Again,
dimensionless rate constants were interpolated from working curves.
In the case of the wall jet electrode, experimentally accessible flow-rates are
assumed to lie in the range 1 x 1o-~-I cm3 s-' , with a typical jet diameter
0.3 mm impinging on an electrode of radius 0.1-1 cm. Dimensionless rate
constants were interpolated from the working curves simulated using the
Backwards Implicit method [65]. MacPherson et al. [47,48] have recently
miniaturized the wall jet electrode so that this technique can be applied to
a uniformly accessible microelectrode system to give a micro-jet electrode.
cm3 s-' through
Flow-rates are reported in the range 2 x lop3-5 x
a nozzle of radius 30-60 pm at distances varying from tens to hundreds of
microns from the microdisc electrode.
For the channel electrode, the following typical parameters were used:
d = 0.6 cm; 2h = 0.06 cm; w = 0.4 cm; x, = 0.1-0.4 cm; Vf =
1 x 10-~-0.3 cm3 s-' . In the case where a microband electrode was used, the
smallest microband which could be fabricated reliably (x, = 1pm) was used.
Dimensions used for the fast-flow-cell were: d = 0.2 cm; 2h = 0.01 cm;
w = 0.15 cm. This can accommodate band electrodes of xe = 1-1 00 pm,
and flow-rates of 1 x 10-~-2.5 cm3 s-l. Dimensionless rate constants for the
channel electrode method were interpolated from the surfaces of Alden and
Compton [63].
The results of calculations of the kinetic discrimination for a range of steadystate techniques considered in this book and using the above-listed experimental
parameters and approximations are summarized graphically in Figs 2.46-2.48.
The overall rate constant 'window' (Fig. 2.36) of each geometry is the product
of the range of kinetic visibility at a particular geometry (Fig. 2.35) and the
range of time-scales that can be accessed (Fig. 2.34). It is clear from Fig. 2.34
that the hydrodynamic electrodes have a narrower kinetic 'window' (i.e. less
inherent kinetic discrimination) than diffusion-only systems, but that convection
allows shorter time-scales to be accessed.
14.2 A comparison of the homogeneous kinetic discrimination of
steady-state and transient experiments
Extremely high scan rates of potential can be used to accelerate the linear
sweep or cyclic voltammetry time-scale beyond that of steady-state methods.
The time-scale accessible by cyclic voltammetry is given by the expression
Comparison of voltammetric techniques
99
~ ' 1 s tflow channel
Microband channel
Channel
Micro-jet
Wall-jet
Rotating-disc
Microdisc
Hemisphere
-2
-3
1
0
-1
2
log (Tirnescale (s))
2.46 Time-scales accessible by steady-state voltammetry at common electrode geometries.
Adapted from Adv. Phys. 0%.
Chern. 32 (1999) 1.
Fast flow channel
Microband channel
Channel
Wall-jet
Rotating-disc
Microdisc
Hemisphere
-2
-1
0
1
2
3
4
5
6
7
log (Tirnescale (s))
Fig. 2.47 Range of rate constants that can be measured by steady-state voltammetry with common
electrode configurations for an ECE process. Reproduced by courtesy: Adv.Phys. 0%.
Chern. 32
(1999) 1 . Copyright, Academic Press.
100
Principles of voltammetry
Fast flow channel 1
Microband channel
Channel
Wall-jet
Rotating-disc
Microdisc
Hemisphere
log (Timescale (s))
Fig. 2.48 Range of rate constants that can be measured by steady-state voltammetry with common
electrode configurations for an EC2E process. Reproduced by courtesy: Adv. Phys. 0%.Chern. 32
(1999) 1. Copyright, Academic Press.
RT/ F v [5]. Typically, scan rates in the range of 10-1 0,000 V s-I are used in
cyclic voltammetry as the upper limit of the scan rate used when macrodisc electrodes are employed which correspond to timescales in the p s range. However,
it should be noted that when microdisc electrodes are used in the transient
regime, scan rates as fast as 1 x lo6 V s-' have been used, which achieve a
time-scale in the region of 10 ns [29-3 11.
The price paid for use of very fast scan rates in cyclic voltammetric experiments is precision. In this, and other, transient methods at short time domains,
the current changes so rapidly with time, that a small error in the measurement of the time at which the current is sampled, introduces a large error in
the calculated rate constant. Of course, analogous problems arise with precision
when steady-state methods are pushed to their limits (e.g. turbulent flow conditions are associated with the use of extremely high flow-rates in hydrodynamic
voltammetry). Transient methods, as previously noted, are also complicated
when short time domain measurements are undertaken in the presence of large
capacitive currents. Some of the practical difficulties related to the use of the
different voltammetric techniques are outlined in Table 2.5.
Electrochemical synthesis and subsequent isolation of gram or greater amounts
of a pure sample of product B, via the electrode process A -+ B ne-, requires
+
Parameter
Cyclic voltammetry
(transient)
Precision
Time errors; capacitive currents
become large at high scan-rates
Reduced IR, drop; significant
errors in electrode dimensions
Modelling
Time-dependent;
high-dimensional (I-E-t)
working surfaces
Two-dimensional diffusion
requires use of sophisticated
numerical methods of analysis
Use with hybrid
methods, e.g.
spectroscopic
detection
A relatively large cell is required
to maintain bulk concentration
which may present practical
difficulties
Reduced cell size is compatible
with most spectroscopic
apparatus. May pick up electrical
noise from other equipment
Electrode: practical
considerations
Static solution means that
passivating films may deposit on
the electrode. However,
electrode is easily removed and
cleaned. Electrolysis products
may build up in the bulk
solution. Natural convection
may become significant at low
scan-rates
Apparatus
considerations
Ohmic drop compensation
necessary at high scan-rates
Difficult to fabricate desired
electrode dimensions accurately.
Electrode elevation and recession
is much more significant.
Lithographically fabricated
microbands are very fragile and
cannot be cleaned mechanically.
Natural convection may interfere
with diffusion-only responses of
larger electrodes (e.g. >25 pm)
Small currents require high
amplification and shielding to
eliminate noise
"Adapted from Adv. Phys. Org. Chem. 32 (1999) 1.
icroelectrodes (steady-state)
ydrodynamic
Flow profile is perturbed by
electrodes which are not flush or
smooth
Generally steady-state. Apart from
RDE, models are
two-dimensional due to
convection
Spectroscopic and photochemical
methods are easily incorporated
into the small, transparent
channel flow cell. R D E is less
versatile
If electrode becomes passivated
and mechanical cleaning is
necessary, the cell must be
disassembled (introducing cell
height or electrode-jet distance
variation). For ChEs by-products
formed at the counter electrode
cannot reach the working
electrode
Flow-regulation apparatus must be
calibrated. Cell must ideally be
designed to ensure laminar flow
in order to simplift modelling.
R D E rotation must have stable
frequency and be axially
symmetric
102
Principles
ofvoltammetry
exhaustive or bulk electrolysis of electroactive material A at a large-size working electrode, although, of course, exhaustive electrolysis of A is frequently
achieved on the microscale (mg to pg) level in mechanistic studies, at smallersized electrodes. Thus, bulk electrolysis can be a useful large-scale synthetic
tool and, indeed, is widely used for commercial production of metals such
as Cu, Zn, and A1 (Chapter I), and for synthesis of some important organic
compounds. Bulk electrolysis experiments, when the data are analysed in a
coulometric form, also enable the n-value in a voltammetric electrode process to
be determined, provided no additional reactions occur on the longer time-scale
(typically minutes to tens of minutes) associated with such experiments. The
bulk electrolysis method may be applied in a constant potential (potentiostatic)
or constant current (galvanostatic) modes.
15.1
Theory of bulk electrolysis
In a controlled potential electrolysis (CPE) experiment, the working electrode
is held at a constant potential until the solution is exhaustively electrolysed.
The electrolysis potential is usually chosen to be slightly more positive than
the Ell2 value of a species being oxidized or slightly more negative than the
Ellz value of a species being reduced, and hence just into the limiting-current
region of a hydrodynamic voltammogram. A galvanostatic experiment, as the
name implies, is one in which the current is maintained at a constant value for
the desired period of time. In a coulometry experiment, usually undertaken
during the course of bulk electrolysis, the current passed as a function of time is
integrated by means of an electronic integrator or coulometer to give the charge.
Faraday's law of electrolysis requires that the quantity of electricity passed or
charge (coulombs) is directly proportional to the amount of chemical reaction
(equivalents) that has taken place at an electrode. Thus [66],
where I is the current at time t and Q is the charge or number of coulombs passed
during the course of the electrolysis experiment. Alternatively, this equation
may be expressed as
N = Q/nF
(2.64)
where N is the number of moles of electroactive material oxidized or reduced,
F is Faraday's constant and n is the equivalents per mole, that is, the number of
electrons transferred per molecule or ion that has undergone electrolysis. It can
be seen from eqn (2.64) that coulometric monitoring of the bulk electrolysis
experiment involving the reaction A -t Bfne- can be used to calculate n, ifthe
number of moles ofA initially present in the bulk solution is known. Conversely
the equation may be used to calculate the number of moles of species A intially
present, and hence the concentration of A in solution (provided the solution
Bulk electvolysis
103
volume and n are known). It also follows from eqn (2.64) that
in any experiment in which the current is kept constant for the duration of the
bulk electrolysis.
hen a constant potential exhaustive electrolysis is carried out under conditions where a constant level of convection is maintained for the duration
of the experiment, and the background or non-Faradaic current is relatively
insignificant, the Faradaic current, I, for a simple A -+ B ne- process decays
exponentially with time, t , according to the relation [66]
+
I = Ii exp (-pt)
(2.66)
where Ii is the initial Faradaic current and p is the mass-transport-controlled
rate constant for the reactant species undergoing electrolysis. Typical current
charge versus time plots for this simple mechanism are shown in Fig. 2.49.
For efficient electrolysis, a high value of p or very efficient mass transport
by convection is desirable. When the potential is held at a value corresponding
to the limiting current regime of a hydrodynamic voltammogram, p can be
where k,,, is the mass transport parameter for the reactant species, A is the
area of the electrode and V is the volume of the solution. The mass transport
parameter depends on the reactant diffusion coefficient, the solution viscosity
and the relative velocities of the solution and electrode surface, that is, the rate
of stirring or flow rate of the solution or rotation rate of the electrode. T o have
a large value ofp, the solution or electrode must be rapidly stirred and the cell
must have a high A/V ratio.
rovided there is only a negligibly small background current present, and unimass transport conditions operate throughout the course of the electrolysis
t, the current is proportional to the concentration of electroactive
the process A -+ B ne-. Thus, when the current has decayed
to a value of approximately 0.1 per cent of its initial value, the electrolysis is
99.9 per cent complete. Due to the exponential relationship of the decay of
the current with time, the rate of electrolysis decays with time as illustrated
+
(a) Faradaic current and (b) charge versus time plots obtained for CPE experiments under
constant convection conditions.
104
Pvinciples of voltammetvy
in Fig. 2.49. If mechanistic nuances are present, as is the case with ECE and
related reactions, the current-time and current-charge curves may, of course,
become much more complex but in principle may be analysed to enable the
rates of homogeneous reactions coupled with the electron-transfer reaction to
be calculated [66].
15.2
Cellsfor bulk electrolysis
Two different cells routinely used for CPE in the author's laboratories [67] are
described in detail in this book: (a) a large-volume cell (100mL) for electrosynthesis of gram quantities of material; (b) a small-volume cell (5 mL)
for electrolysis of relatively small amounts of material (up to 200 mg); this is
predominantly used for coulometry with voltammetric and/or spectroscopic
monitoring of the products formed.
15.2.1 Large-volume electvolysis cellfov electrosynthesis
A schematic diagram of a large-volume cell used for electrosynthesis is given
in Fig. 2.50. The large-volume GC cup serves as the working electrode and
also holds the solution to be electrolysed. The reference electrode is contained
Fig. 2.50 Schematic diagram of CPE cell and associated experimental arrangement used for
electrochemical synthesis experiments: (a) nitrogen gas bubbler (b) reference electrode (c) spring
clamp (d) Teflon o-ring (e) platinum basket auxiliary electrode within inner solution compartment
(f) 1.0-1.7 pm porosity glass frit (g) magnetic stirrer bar (h) glass outer housing (i) GC cup working
electrode forming outer solution compartment, and 6) glass supports for counter electrode compartment. Reproduced by courtesy: J. Chem. Sol. Perkin Trans. 2 (1995) 1365. Copyright, Royal
Society of Chemistry.
Bulk electrolysis
105
in the solution which is in contact with the working electrode. The auxiliary
electrode consists of a platinum gauze basket which is arranged symmetrically
inside the working electrode. Separating the working and auxiliary electrodes
is a glass cylinder with a porous glass frit in the base within which the platinum
auxiliaty electrode sits. The separation of these solutions is required to prevent
product generated during the course of the electrolysis at the auxiliary electrode
reacting with the products of interest that are formed at the working electrode.
The design is similar to that described by Moinet and Peltier [68], and Fry [69].
p he important features of the cell are given below.
( 3 ) There is a porous glass frit, separating the working and auxiliary electrode
compartments, which has a surface area of approximately 20 cm2. The
large frit surface area enables adequate current to be passed between these
two electrodes which in turn allows fast electrolysis times to be achieved.
owever, the use of a large frit also increases the amount of leakage between
the two compartments. In order to test the level of mixing of solutions that
occurs between the two compartments, fluorescein dye may be added to the
working electrode compartment, and the change in colour of the auxiliary
electrode compartment can be monitored as a function of time by UVspectroscopy. An experiment of this kind shows that after six hours
electrolysis conditions, 10 per cent mixing of the solutions occurs
using a 1.O-1.7 pm porosity glass frit (porosity no. 5) and the cell
epicted in Fig. 2.50. In contrast, a 10-20 pm porosity glass frit (porosity
o. 4) allows a much larger level of mixing of nearly 40 per cent to occur
over a three-hour period, making the use of the larger porosity glass frit
unacceptable with this particular cell design. The higher porosity frit does,
of course, permit a slightly higher current to flow, thereby allowing shorter
electrolysis times, but generally the disadvantage of a large increase in solution mixing between the two compartments more than offsets the advantage
gained by a faster electrolysis time.
(2) The working and auxiliary electrodes are arranged as symmetrically as posh respect to each other to assist the attainment of an even potential
ent distribution over the surface of the working electrode [70]. The
volumes of the working and auxiliary electrode compartments in the cell
are approximately 100 mL and 65 mL, respectively. The IR, drop which is
always present in any bulk electrolysis experiment, should be distributed as
uniformly as possible over the electrode surface. This requires that the frit
separating the two compartments, ideally, should also be arranged as symmetrically as possible. For practical reasons the highest possible symmetry
condition is not achieved, mainly because of leakage problems and the difficulty in obtaining a cylindrical frit of optimal diameter and pore size. The
asymmetry present in the cell in Fig. 2.50 results in a higher current density
and, therefore, more electrolysis occurring at the base rather than at the
sides of the GC working electrode cup, and also in a lower overall current
ow than the maximum which is theoretically possible. For this particular
cell, conditions are set so that current levels at the start of the electrolysis are
106
Principles ofvoltammetvy
in the 0.2-1 A range, the exact value depending on the concentration of the
species being electrolysed. In order to limit the amount of mixing between
the working and auxiliary electrode compartments, exhaustive electrolysis
experiments must be completed in less than two hours. This means that for
species with molecular weights of the order of 1000 Da, gram quantities of
electrolysed product may be obtained.
(3) The reference electrode, for example, A ~ / A ~(0.05
+ M AgNO,; CH3CN)
for electrolysis in acetonitrile solvent, is positioned as close to the surface of
the working electrode as possible to ensure a potential control which is as
uniform as possible [7I].
(4) The working and auxiliary electrode compartments are stirred using a
magnetic stirrer bar to increase the value o f p (see eqn 2.64).
(5) The working and auxiliary electrode compartments are easily degassed.
(6) Large currents of up to 1 A can be obtained with this cell configuration
when a standard 0.1 M concentration of supporting electrolyte is present.
However, that rate of electrolysis may slow down prematurely in some
situations unless an adequate concentration excess (at least 1.5 times) of
supporting electrolyte to compound being electrolysed is used. This premature rate decrease is highly significant when ion pairing occurs between
charged products of the electrolysis and the electrolyte, thus decreasing the
effective electrolyte concentration and lowering the overall conductivity of
the solution. The impact of precipitation of insoluble products containing
electrolyte ions may also need to be considered, for similar reasons.
(7) The cell can be easily assembled and dismantled for cleaning.
(8) The potential required for the electrolysis may be determined by taking
an aliquot of solution from the working electrode compartment, diluting
to approximately 1 mM, and recording a cyclic voltammogram, using, for
example, a 1-mm diameter macrodisc working electrode made from the
same material as the bulk electrolysis working electrode along with exactly
the same reference electrode as that used in the large-scale electrolysis cell.
The electrolysis experiment may also be monitored ex ritu by periodically
taking samples hom the working electrode compartment and recording
voltammograms or spectra (electronic, infrared (IR), NMR, etc.) in the
usual manner.
15.2.2 Small volume cell for coulometry and in situ voltammetnc or
spectroelectvochemical measurements
For coulometry measurements used to determine n-values associated with an
electron-transfer process, a smaller solution volume cell, which retains most of
the design features of the cell described in Section 15.2.1, can be used. In the
smaller volume cell, the 5-mL volume inner compartment contains a GC cylindrical working electrode (area approximately 28 cm2) as the working electrode
and a reference electrode separated by a 1.0-1.7 pm sintered glass frit from a
platinum mesh cylindrical auxiliary electrode. Solutions are again purged with
Spectroelectrochernistry
107
1 Cell for bulk electrolysis at a platinum gauze electrode. The working electrode compartment 13 sits inside the auxiliary electrode compartment A. (a) Pt mesh basket working electrode
(b) clip to pressure seal cell and lid (c) Pt mesh basket working electrode (d) porosity 5 glass frit
(e) magnetic stirrer bar, and (f) reference electrode. Reproduced by courtesy: Anal. Chem. 67
(1995) 1691. Copyright, American Chemical Society.
an inert gas prior to and during electrolysis and are stirred with a magnetic
stirrer bar. In this cell, sufficient space is available so that the standard elecarrangement used for voltammetric experiments may be inserted into the
ing electrode compartment and hence cyclic, rotated-disc or microelece voltammograms may be recorded during the course of the electrolysis.
tu spectroscopic (infrared, electronic, etc.) monitoring of products of elecis is also possible using detection aided by fibre optics transmission of
(Section 16). Of course equivalent designs with platinum or other metal
(Fig. 2.51), rather than GC working electrodes and porous vycor or other
materials (e.g. membranes) rather than sintered glass frits can be made using the
principles described in Section 15.
ee
ist
Comparisons of voltammetric experiment and theory frequently provide significant clues to the mechanisms of an electrode process, but only rarely can the
identity of intermediates and products be deduced solely from voltammetric
ata. Consequently, the characterization of intermediates and final products
proposed in a mechanism, whenever possible, should be confirmed by spectroopic identification (Fig. 2.6). Ex situ spectroscopic measurements made after
lk electrolysis experiments can obviously be used to identify stable products
tion 15). However, because in situ spectroelectrochemical measurements
y to much shorter time domains, they provide a powerful method of ideng the nature and even structures of intermediates involved in reaction
hanisms. The nature of the spectroelectrochemical technique should be
108
Principles of voltammetry
appropriate to the problem [72]. Thus, for example, electron spin resonance
(ESR-also called electron paramagnetic resonance (EPR)) spectroelectrochemical techniques are likely to be ideal for identifying radicals that are formed
by one-electron oxidation or reduction of diamagnetic organic species.
16.1
E S R spectroelectvochemistvy
Undertaking an electrolysis experiment on a flowing solution which moves
over a channel electrode and then passing the electrolysed solution into an ESR,
Raman, Infrared, UV-visible, or other kind of spectrometer (including an electrospray (ES) mass spectrometer), represents an ideal method of detecting and
identifying moderately stable species generated upstream at the electrode surface.
Figure 2.52 depicts an in situ channel-electrode cell design that has been used
in the ESR form of spectroelectrochemistry [73]. Alternatively, designs can be
achieved in which the entire electrochemical cell [74-761 may be placed inside
the cavity of the ESR spectrometer (Fig. 2.53) and when a microelectrode
working electrode is employed (Fig. 2.54), simultaneous ESR-voltammetric
experiments may be undertaken in situ using electrolysis in very small volumes
of solution [74].
Data illustrating responses obtained from channel electrodes and in situ small
volume microelectrode techniques are contained in Fig. 2.55 for the reduction
at a platinum electrode of dipropylpyridine-2,4-dicarboxylateto its anion radical [77]. Inspection of Fig. 2.55 reveals that the channel flow-cell experiments
(Fig. 2.52) display a markedly greater sensitivity as shown by the much enhanced
Fig. 2.52 Channel flow-ESR spectroelectrochemical cell. Reproduced by courtesy: J. Electroanal.
Chem. 144 (1983) 87. Copyright, Elsevier.
Spectroelectrochemistry
109
Electrolytic
cell
Lock-in
I
Resonator
Electromagnet
80 H signal
. 2.53 A highly sensitive microelectrode-ESR detection system obtained by placing the electrochemical cell within the ESR cavity. Reproduced by courtesy: Chem. Lett. (1995) 919. Copyright,
Japanese Chemical Society.
signal-to-noise ratio and the better developed hyperfine splitting patterns
observable. The significant improvement in sensitivity is largely attributable
to the relative electrode sizes. In the small volume cell (Fig. 2.54),spins are
injected through a 70-pm diameter microdisc electrode. In contrast, channel
electrodes typically have approximate dimensions of 4000 pm x 4000 pm, so
that the difference in electrode areas is approximately 4 x lo3 although this
relative advantage may be partly offset by the effect of convection in sweeping
radicals out of the ESR cavity once they are formed.
The effect of convection on the intensity S , of the ESR signals generis demonstrated in
ated by reduction of dipr~p~lpyridine-2,4-dicarboxylate
ig. 2.56, which shows signal/flow rate (Vi)/current (I data analysed according
Pn'nciples of voltammetry
Cu rods
Soldered connections
Nylon support
B7 joint (quartz socket,
pyrex cone)
o-ring
average filling level
(approx. 0.25 ml)
counter electrode
(250 pm Pt)
Sensitive region of cavity
Reference electrode
(70 pm Pt, coated with
Ag/AgCl)
1
\ Working electrode
(70 pm Pt)
Fig. 2.54 Small-volume microelectrode cell for simultaneous voltammetric-ESR experiments
inside the ESR cavity. Dimensions are given in mm. Adapted from: J. Electroanal. Chem. 404
(1996) 303 and J. Electrochem. Soc. 142 (1995) 863.
to eqn (2.68) which is known to apply to a stable radical in the laminar channel
flow cell [77]).
Specifically, a plot ofloglo{S/I}against loglo{Vf}is seen to be linear with a slope
of (-213) as predicted by eqn (2.68). This result provides clear evidence that the
radical is stable on the time-scale of the channel electrode ESR experiment (that
is, there is no time for decay before the radical leaves the ESR cavity). In practice,
given the range of solution flow rates normally employed in channel electrode
ESR work, this means that the anion radical has a lifetime in excess of 5-10 s.
Clearly, mechanistic information is accessible via channel electrode measurements in which the ESR signal is monitored as a function of the flow rate and
electrode current. Radical instability is revealed by curvature in plots of the type
Spectroelectrochemistry
(a)
200
111
0
I3470
3480
3490
3500
Field (G)
. 2.55 ESR spectra resulting from the reductive electrolysis of dipropylpyridine-2,4-dicarboxylate in acetonitrile solution; (a) 15 rnM (b) 5 mM using the small volume cell (Fig. 2.54), and
(c) 1.5 mM using the channel flow cell (Fig. 2.52) and a flow rate of 4 x 10-'cm3 s-l. Reproduced
by courtesy: J. Electroanal. Chem. 404 (1996) 303. Copyright, Elsevier.
displayed in Fig. 2.56 and as shown schematically in Fig. 2.57. Such plots can be
analysed to provide information concerning the mechanism of the radical decay
and the associated kinetics. Various EC, ECE and disproportionation processes
have been quantitatively studied in this manner [73,78-841. The small volume
cell does not readily permit the direct interrogation of electrode reaction mechanisms in an analogous way through the ESR signal. However, because of the
short time constant of the cell, fast scan cyclic voltammetry is viable, and may
1 12
Principles of voltarnrnetry
Fig. 2.56 Variation of the current normalized ESR signal intensity ( S I I ) with solution flow rate
(Vf) for reduction of dipropylpyridine-2,4-dicarboxylatein a channel flow cell. The line shown has
a slope in agreement with eqn (2.68). Reproduced by courtesy: J.Electroanal. Chem. 404 (1996) 303.
Copyright, Elsevier.
1%
S
A
'--,
I
Stable radical
R-
1% Vf
Fig. 2.57 Schematic behaviour of ESR signal-flow-rate-current data for stable and unstable radicals.
Reproduced by courtesy: J. Electroanal. Chem. 404 (1996) 303. Copyright, Elsevier.
be used to interrogate the electrode process in the conventional manner [73,77]
without recourse to the ESR signal. Cyclic or linear sweep voltammetry at
channel electrodes is restricted by the much larger time constant of the electrode, although the relevant theory which takes into account the influence of
the convective flow has been derived [85]. Instead, hydrodynamic voltammograms are normally recorded under steady-state conditions, and the variation
of the half-wave potential, and/or the transport-limited current, with solution
flow rate is used to derive kinetic and mechanistic information [86]. It is noteworthy that the demands of high sensitivity, which dictate as large an electrode
area as possible, are in direct conflict with those essential for undistorted cyclic
voltammetric studies. Finally, it can be noted that the channel flow cell is not
readily amenable to low-temperature work and that it is prone to consume relatively large quantities of electrolyte. Both these deficiencies are remedied in
the design considered in reference [74].
n summary, the advantages of the channel electrode flow through cell,
introduced by Coles and Compton [73], are as follows:
the deduction of kinetic and mechanistic parameters is readily accomplished
using the ESR signal and its dependence on flow rate and electrode current.
he advantages of the microelectrode, stationary solution, in situ cell designed
by Fiedler, Koppenol and Bond [74] are as follows:
only a small volume of solution is required (0.2 mL);
easy to work at variable temperature;
the electrode has a short cell time constant permitting, for example, fast scan
rate cyclic voltammetry;
the cell is of all-glass construction and is therefore usable in almost all solvents.
eferences [72,76,77,87,88] provide useful overviews of the subject of ESR
ectroelectrochernistry.
learly infrared spectroelectrochemical monitoring is likely to be excellent for
electrochemical studies on compounds where intermediates and/or final products exhibit strong infrared bands. Figure 2.58 illustrates what effectively acts
as a very small volume thin-layer cell design in order to achieve very efficient
electrolysis at a platinum macrodisc electrode [89,90]. With this design, infrared
iation only has to pass through a very thin layer of solution and after being
rbed by the species of interest can be reflected off the platinum electrode
and then detected by the I R spectrometer. The method is based on the use of
a modified infrared reflection-absorption spectroscopy (IRRAS) cell which is
mounted on a specular reflectance accessory located in the sample of a Fourier
Transform Infrared (FTIR) spectrometer [89,90].
Transition metal carbonyl complexes have intense infrared bands in the
2000 cm-' region and so studies on this class of compounds gain enormously
from the use of the infrared spectroelectrochemistry approach. Figure 2.59(a)
shows that the oxidation of Cr (CO), (C6Me6), under conditions of cyclic
voltammetry at a platinum macrodisc electrode, is chemically reversible
(scan rate = 100 mV s-l) at room temperature in dichloromethane. Monitoring
the course of this reaction by the IRRAS spectroelectrochemical method gives
the result shown in Fig. 2.59(b). The shift in the carbonyl infrared bands to
higher wave numbers is exactly as expected if [Cr(CO), ( c 6 ~ e 6 ) ]is' generated
via the electrode process Cr(CO), (C6Me6)+ [Cr (CO), (C6Me6)] e- .
In the experiment in Fig. 2.59(b), application of a constant applied potential
of 1000 mV versus Ag/AgCl is adequate to oxidize Cr (CO), (C6Me6)and FTIR
+
+
1 14
Principles of voltammetry
to micrometer
cooling gas in
-
---c
cooling gas out
platinum wire
B14 sept
platinum disc
working electrode
\ alkali halide/quartr
window
Fig. 2.58 Infrared spectroelectrochemical experiments based on the use of a modified infrared
reflection-absorption spectroscopy (IRRAS) cell [87,88]. A flow of cold nitrogen gas is maintained
for low-temperature experiments. Adapted from J. Chem. Soc. Dalton T m s . (1996) 2945 and
provided by courtesy of D.G. Humphrey, Monash University, Victoria, Australia.
spectra can be recorded rapidly to monitor the course ofthe electrolysis. Spectral
runs 0, 4, 8, and 9 are shown in Fig. 2.59(b) and because the differential
adsorbance is used to subtract the contribution from the solvent, the initial
spectrum (rc.0, in Fig. 2.59(b)) is inverted relative to the usual display format
used in infrared measurements. Consequently, the infrared spectrum of the
unoxidized form of the compound C T ( C O )(C6Me6)
~
lies below zero and decays
towards zero as the extent (time) of electrolysis increases, while that for the
lies above zero and increases with time.
product [Cr(CO)3(C6Me6)]+
Spectroelectvochemistry
2200
2100
2000
1900
115
1800
Wavenumber (cm-1)
2.59 (a) Cyclic voltammogram (scan rate 100 mV s-l, platinum macrodisc electrode) for
the oxidation of C S ( C O ) ~ ( C ~ in
M dichloromethane
~~)
(0.2 M Bu4NPF6) at room temperature.
(b) Infrared measurements obtained during the course of oxidation electrolysis using the cell described
in Fig. 2.58. Provided by courtesy of D.G. Humphrey, Monash University, Victoria, Australia.
Clearly, far more complex reaction schemes can be unravelled by infrared
troelectrochemistry in the IRRAS mode, particularly when the low temture mode of operation is employed in this thin-layer electrolysis cell.
tensive use of this method is described in references 191-931.
As an alternative to use of a thin-layer electrolysis method, systems based on
e use of fibre optic cable for transmission of infrared radiation (Fig. 2.60) may
e used in an in situ dip probe made for monitoring the course of conventional
ulk electrolysis experiments [94]. Unlike thin-layer cell methods, this spectroctrochemical technique requires no cell design compromises which diminish
e accuracy of the voltammetric data that may also be obtained during the
course of the electrolysis.
In the design described by Shaw and Geiger [94] a bifurcated 1-m fibre-optic
alcogenide cable mounted on the exit port of a FTIR spectrometer carries
radiation into the electrolysis cell and returns it to a HgCdTe detector.
he probe terminates in a tip into which the analyte solution is allowed to
ow freely through entry holes (Fig. 2.60). Infrared light enters the solution
\
screw threads
I------8
mm------1
Fig. 2.60 Infrared spectroelectrochemistry based on the use of sampling probe head at the end of
the fibre-optic cable. The portion shown schematically is immersed in the electrolysis solution, and
solution enters the sample chamber through holes in the stainless steel probe wall. Reproduced by
courtesy: Organometallics 15 (1996) 13. Copyright, American Chemical Society.
from the cable, bounces off a gold reflecting mirror attached to the end of an
adjustable screw, and is returned to the detector by the cable. The optical path
length may be changed (c.1-5 mm) by adjusting the screw. The oxidation of
1.0 rnM acetylferrocene ( E ~ O = 0.27 V versus FC/FC+) was studied using this
infrared spectroelectrochemical method in CH2C12(0.1 M NBu4PF6)at 273 K
by monitoring the position and intensity of the organic carbonyl band as the
reversible electrolysis to the cation proceeded at Eappl
= 0.53 V (v,, 1666 cm-'
for acetylferrocene, v,, 1700 cm-' for the cation) again as difference spectra
(Fig. 2.61 (a)). Progress of the conversion to the oxidized cation was also followed by voltammetry and coulometry [94]. Steady-state voltammograms at a
25-pm diameter Pt disc electrode (Fig. 2.61(b)) were recorded at approximately
the same time as the infrared spectra in Fig. 2.61 (a). Moreover, the absorbance of
the infrared bands was directly proportional (1700 cm-', product) or inversely
proportional (1666 cm-' , reactant) to the amount of charge passed during the
course of electrolysis as shown in Fig. 2.61 (c).
16.3
UV-visible spectroelectrochemistry
UV-visible spectroelectrochemical monitoring of the course of electrolysis
has been in widespread use for many years [72]. Thus, optically transparent
gauze minigrid gold or platinum (Fig. 2.62) or optically transparent tin or
iridium oxide semiconductor electrodes have been commonly used for in situ
UV-visible spectroelectrochemical measurements over very wide temperature
Spectroelectrochemistry
117
ranges. Alternatively, flowing solution channel electrode techniques can be
used conveniently (Fig. 2.63) with detection by UV-visible spectrophotometry in a manner akin to their ESR and I R spectroelectrochemical analogues.
The result obtained in a UV-visible spectroelectrochemical experiment is
shown in Fig. 2.64 for the reversible one-electron reduction of l-bromo9,lO-anthracenedione to its anion radical [95]. The series of isosbestic points,
expected for a simple reaction of this kind, is readily observed, as are parallel
absorbence and current transients in a potential step experiment (Fig. 2.65).
Further examples of the use of UV-visible spectroelectrochemical techniques
are provided in Chapter 3.
6.4 N M R spectroelectvochemistry
Coupling of solution-phase electrochemistry and N M R spectroscopy has
proved difficult even though the great value of N M R as a structural tool is
well recognized [96]. The desirability of spinning the N M R tube while simultaneously undertaking an electrolysis experiment represents a far more difficult
technical problem than that encountered in the ESR, IR, and UV-visible in situ
spectroelectrochemicaltechniques described in Sections 16.1-1 6.3. A solutionphase N M R spectroelectrochemical technique has been described by Prenzler
et al. [96] using the cell described in Fig. 2.66, which allows solution electrolysis
to occur in the receiving coil region of a superconducting, high-field N M R
spectrometer. Although the idea of coupling electrochemical generation with
M R detection is not new [97-1011, this cell is unique in that: the N M R
spectra may be accumulated as the electrolytic current is flowing; the electrolysis products are generated directly at the receiver coil; the outer tube spins
keeping the inner tube strictly centred and the sample mixed as electrolysis
occurs; and finally, the unit is designed to work in a commercial probe-that
is, no custom built probe is necessary. Figure 2.67 illustrates 'H N M R spectroelectrochemical results obtained for the two-electron, two-proton reduction
of p-quinone to (di)hydroquinone in acidified DzO. Both 0 = C 6 H 4 = 0 and
0- C6H4-OD are diamagnetic with a single, one line 'H N M R signal
in the aromatic region, and even though the two lines are quite close (10 Hz
apart with 300 MHz N M R instrument) excellent resolution is achieved under
electrolysis conditions.
6.5 Combining mass spectrometry and electrochemistry
An almost unlimited range of in situ and ex situ spectroelectrochemical techniques have been reported in the literature [72] and have become an essential
tool of the electrochemist's trade. Obviously, products obtained during or after
completion of an electrolysis experiment may also be examined by mass spectrometry, and masdcharge ratios and isotopic composition of charged products
determined.
Wavenumber
400
,
Volt vs Fc/Fci
0
1000
2000
Charge (mC)
o acetylferrocene (1666 cm-l)
3000
r acetylferrocenium (1700 cm-l)
Spectvoelectvochemistvy
119
he advantages of mass spectrometry (MS) as an analytical technique have
long been recognized. The high sensitivity and selectivity are particularly useful
fOl-the identification of trace levels of analyte in complex gaseous and solid
mixtures, but, until quite recently, a limitation of the technique was the inability
to routinely determine non-volatile species present in solution, as is required for
identification ofproducts formed by electrolysis in the solution phase. However,
following its initial development in the late 1970s [I021, thermospray MS has
been widely used for the identification of solution-based species [lO3,lO4]. Not
risingly therefore, the online coupling of electrochemistry and thermospray
has been reported for monitoring of electrochemically-generated species in
aqueous solution [I 041.
lectrospray MS (ESMS) has rapidly risen in popularity over the past fifteen
years [105-1071, and is clearly now the preferred MS technique for detection
of charged species present in solution. It is the 'softest' method for transfer of
ions from solution to the gas phase and has the advantage of minimal decomposition and fragmentation of ions. Thus it is ideally suited for the determination
of charged solution species which are frequently generated in electrochemical experiments. The ability to identify the charged products of electrolysis
reactions in solution by ESMS is very desirable, in both off-line and on-line
modes. However, a basic incompatibility exists between the two techniques,
and this problem has to be solved. Electrolysis reactions are normally carried
out with high concentrations of supporting electrolyte (typically about 0.1 M)
ith an approximately 100-fold excess of electrolyte over electroactive
und, whereas in ESMS, high concentrations of ions are very undesirusing problems of peak suppression [I081 and fouling of the interface
ates and lenses. This problem has limited the coupling of electrochemistry and
S. Dupont and co-workers describe the ESMS of products produced with
electrolysis of fullerenes [I091. Using 10 rnM Bu4NPF6 as electrolyte
pM fullerene, optimized conditions gave ESMS signals with the loss of
signal intensity relative to that expected under other conditions being attributed
the high concentration of supporting electrolyte. Zhou and Van Berkel also
ve reported the use of electrochemical cells online with ESMS [I 101. In this
study, the problems of high electrolyte concentrations were minimized by use
of an uncommon type of electrolyte, that is, one incorporating a small metal-ion
and a relatively volatile anion (for example, 20 rnM lithium triflate).
. 2.61 Infrared monitoring of oxidative bulk electrolysis of l.0mM acetylferrocene in
CH2C12(0.1 M NBu4PF6) at 273 K. An increasing amount of the acetylferrocenium cation is produced in the sequence A-F. (a) Difference infrared spectra; (b) Steady-state voltammograms (scan
rate 5 mV s-I at a 250-pm diameter Pt disc electrode) of solutions A-F like those in (a). Scan A
is that of 1.0 mM acetylferrocene and scan F that of nominally 1.0 mM acetylferrocenium cation.
(c) Absorbance changes observed in (a) plotted against the total amount of charge passed. Passage of
1 Faraday mol-I required 3300 mC in this experiment. Reproduced by courtesy: O~anovnetallics15
(1996) 13. Copyright, American Chemical Society.
Principles of voltammetry
Flow in
Flow out
Platinum mesh
Platinum wire
Low-temperature
melting glass sheet
\
s
+Silica plate
Fig. 2.62 Schematic diagram of an optically transparent platinum-mesh-electrode-thin-layer cell
design. The assembled cell can be placed in the sample holder of a UV-visible spectrophotometer
and electronic spectra recorded during the course of electrolysis. Diagram provided by courtesy of
G.A. Heath, Australian National University, Canberra, Australia.
Downstream Pt
counter electrode
Au minignd. Length of
grid was dictated by
Ag paint connection
to Cu wire
I
t-7 mm-i
t-9mm -I
-I 12.5mm 4
Upstream Ag wire
quasi-reference
electrode
Fig. 2.63 Schematic representation of a flow-through cell used for UV-visible spectroelectrochemistry. Reproduced by courtesy: Anal. Chem. 61 (1989) 1787. Copyright, American Chemical
Society.
Spectvoelectrochemistry
300
400
500
Wavelength (nm)
12 1
600
Development of the thin-layer spectrum with time when the cell depicted in Fig. 2.63
is used after solution flow is stopped for the reduction of 1.1 x ~ o - ~ M B in
A acetonitrile
(0.1 M Et4NC104) at 25°C: (-)
BA at t = 0, (- - -) BA and BA0- at t < 130s,
BAe- at t = 130s. The time elapsed to reach exhaustive electrolysis was 130 s.
and (.
(BA = 1-bromo-9,1 O-anthracenedione). Reproduced by courtesy: Anal. Chem. 61 (1989) 1787.
Copyright, American Chemical Society.
a
)
0
60
120
180
Time (s)
240
300
2.65 Parallel current and absorption transients using the cell depicted in Fig. 2.63 with a
stationary solution: (a) Absorption against time at 565 nm; (b) current against time. At (i) the potential
was stepped from 0 to -0.9 V versus a Ag wire pseudo-reference electrode, and at (ii) the potential
was stepped back to 0 V. The initial step (i) corresponded to stepping from a zero current to what
would be a limiting current under flowing solution conditions. Reproduced by courtesy: Anal.
Chem. 61 (1989) 1787. Copyright, American Chemical Society.
122
Principles of voltammetry
Reference
Working
electrode --------
J
- Counter
electrode
10-mrn
NMR -+
Tube
Receiver
coils
Fig. 2.66 Schematic diagram of an N M R spectroelectrochemical cell. Reproduced by courtesy:
Electrochem. Comm. 2 (2000) 5 16. Copyright, Elsevier.
Fig. 2.67 N M R spectroelectrochernical two-electron, two-proton reduction of diamagnetic
p-quinone to diamagnetic hydroquinone using an applied potential of -0.5 V versus A /AgC1 in
aqueous 0.1 M quinone (0.1 M KC1, 0.3 M DC1) inside a 300 MHz N M R spectrometer. H spectra
recorded for 130 min at 10-min intervals. Reproduced by courtesy: Electrochem. Comm. 2 (2000) 5 16.
Copyright, Elsevier.
F
More recently, work in the author's laboratory [I 111 has shown that off-line
(electrolysis in cell described in Fig. 2.5 1, but using very dilute electrolyte),
and, online ESMS electrochemical methods, again with low concentrations of
supporting electrolytes, can be used (Fig. 2.68). For online electrolysis with
ESMS detection of product, a syringe pump was used to transport the electrolysis solution through the flow cell, the outlet of which was connected
directly to the fused silica capillary of the ESMS interface. A schematic diagram is given in Fig. 2.68. The electrochemical flow cell consisted of two
10-cm lengths of Pt microtubing (internal diameter 100 prn) connected by
Voltammetry at variable pressure and temperature
123
Diagram of online electrochemical cell where products of electrolysis can be fed directly
into an ES mass spectrometer. (a) Syringe (b) syringe needle (c) Teflon tubing (d) Pt microtubule
(,) metal connector (0fused silica capillary of spectrometer (g) syringe pump and (h) constant
~oltagepower supply. Reproduced by courtesy: Anal. Chem. 67 (1995) 1699. Copyright, American
chemical Society.
3 cm of Teflon tubing. The microtube closest to the mass spectrometer funcas the working electrode and was connected to the positive terminal
of a high-voltage power supply. A constant potential was applied across the
two microtubes to cause electrolysis to occur. This very simple flow cell has
) ~acetonitrile
been shown to afford efficient electrolysis of 0.2 rnM N i ( E t 2 d t ~in
Bu4NPF6).Prior to electrolysis, the positive-ion mass spectrum in
Fig. 2.69(a) was recorded, and clearly verifies the presence of the Bu4Nf electe cation. After applying 200V across the two Pt microtubules, with the
closest to the mass spectrometer as the working (positive) electrode, and
after allowing time for the electrolysed solution to reach the mass spectrometer
r, the spectrum shown in Fig. 2.69(b) was obtained. Clearly, the prod(Et2dtc)3]+,is readily detected via its m / z value (masdcharge anion) and
the comparison of experimental and predicted isotopic patterns, after generation
by oxidative electrolysis. Thus, the ESMS data support previous proposals that
the electrochemical oxidation of Ni(Et2dtc)2occurs [I 121 via the overall reac-+ 2 [ ~ i ( d t c ) ~ ] +Ni2+ 4e-. The advantage of the online
i(Et2dt~)2
e is obvious: the flow line can be introduced to the mass spectrometer
modification of the instrument; for air-sensitive species, the solution
can be degassed prior to being introduced into the syringe; the required sample
volume is small. In contrast, in the off-line method, larger volumes of solution
are required, products must be stable in longer time domains, and prevention
with adventitious oxygen is difficult.
sly, gaseous products formed during the course of electrolysis may
erized by mass spectrometry [I131 and, of course, products of
bulk electrolysis may be identified with assistance from the relevant chromatographic separation techniques such as gas or liquid chromatography, or
electrophoresis, etc.
+
+
he equations relating to the thermodynamics (AE' or E: values) of an electrochemical cell, operating under equilibrium conditions, have been described
124
Princivles of voltammetry
1
[Ni ( ~ t ~ d t c ) ~ ] 502
'
495
500 505 510 515
m/z
Fig. 2.69 Positive-ion ES mass spectra, obtained using the online electrochemical-ES mass spectrometer flow cell shown in Fig. 2.68 for an acetonitrile solution containing equimolar (0.02 mM)
= diethyldithiocarbamate) and Bu4NPF6 (a) before electrolysis and (b) during
Ni(Et2dt~)2(Et2dt~
oxidative electrolysis. Reproduced by courtesy: Anal. Chern. 67 (1995) 1699. Copyright, American
Chemical Society.
as a function of pressure and temperature in Section 4 of Chapter 1. Voltammetric experiments may also be conducted at variable temperature and pressure,
and it, therefore, follows that the dependence of the kinetics of a reaction on
these parameters can be used to deduce the activation parameters [volumes
(A V*), energies (E,), enthalpies (AH*), and entropies (AS*) of activation]
of heterogeneous and homogeneous reactions associated with electron-transfer
reactions, in addition to the thermodynamically important reaction volumes
(A V), enthalpies (AH) and entropies (AS).
~fo,kO,u
Consider again a voltammetric process of the kind, A
B e-, where
E,O and ko (and in principle a) can be measured as a function of pressure or
temperature. Figures 2.70(a) and (b) illustrate the voltammetry observed for
the reduction of some iron complexes [I 141 as a function of pressure, and the
shift in the position of cyclic voltammograms with pressure can be noted via
data contained in Fig. 2.70(c). If the process remains electrochemically and
++
Voltammetry at variable pressure and temperature
125
chemically reversible as the pressure is varied at constant temperature, then the
position of the cyclic voltammograms essentially moves along the potential axis.
small changes due to the dependence of the diffusion coefficient on pressure are
likely. Under these conditions, the value of the reversible potential will reflect
the dependence of the thermodynamics of the overall process (combination
of working electrode and reference electrode half-cell reactions, Chapter 1)
essure. In contrast, if the temperature is varied at constant potential, and
ocess remains reversible at all temperatures studied, then the value of
the reversible potential may also vary with temperature, but in this situation a
change in shape of the voltammogram is predicted to occur as a result of the
nF term present in the theory.
the process is kinetically controlled by the heterogeneous charge-transfer
step (kO value) rather than being reversible, then k0 will change as either the
pressure or temperature is varied, assuming E, or A V* and AS* are not zero.
[Fe (phen) (cN)~]I-/'-
.+:
Fig. 2.70 Continued
126
Principles of voltammetry
$-@(,.I
a,)
(n1Y
[Fe (cN)~]~-'
0.0
25.0
50.0
75.0
100.0
125.0
Pressure (MPa)
Fig. 2.70 Dependence of voltammograms and reversible potentials on pressure. (a) Voltammograms ofthe [Fe(phen)(CN)4]1-/2-couple (phen = phenanthroline) in 1 M K N 0 3 at 0.1 Mpa (-),
36.2 Mpa (. . . ), 70.3 Mpa (- - -), and 105.2Mpa (- . -.) at a gold disc electrode using a Ag/AgN03
(0.01 M); 1.0 M K N 0 3 reference cell. Scan rate = 0.057 V s-'; T = 298 K. (b) Voltammograms of
the [ F ~ ( O X ) ~ couple
] ~ - / ~(ox = oxalate) in 0.10M NH4HC204at 0.1 Mpa (-1, 35.7Mpa (. . .),
69.7 Mpa (- - -), and 104.5 Mpa (- - -). Conditions as for (a), but using a Ag/AgN03 (0.01 M);
0.1 M K N 0 3 reference cell. (c) Variation of formal potential with pressure for [ F ~ ( c N ) ~ ] ~ - / ~ ]
(empty circles), and
(filled circles), [Fe(bpy)(cN)~]'-I2- (empty squares), [ ~ e ( b p y( )c~N ) ~ '+I0
Fe(bpy):+I2+ (filled squares) couples (bpy = 2, 2'-bipyridyl) relative to the formal potential at
0.1 Mpa. Ionic strength = 1.OM; T = 298 K. Reproduced by courtesy: Inorg. Chem. 33 (1994)
6180. Copyright, American Chemical Society.
The value of ko would, of course, be expected to decrease with temperature
as the rates of almost all reactions slow down as the temperature decreases.
An analogous situation with respect to temperature dependence applies, if a
rate-determining homogeneous chemical step accompanies a reversible chargetransfer process. Thus, consider the EC reaction
Voltammetry at variable pressure and temperature
c.)
A
0
0.4
0.8
1.2
Volt vs Ag/Ag C1
127
1.5
0
0.4
0.8
1.2
1.6
Volt vs Ag/Ag C1
Cyclic voltammograms for oxidation of [Cr(CO)51] in acetone at a platinum electrode: (a) T = 20°C; scan rate = 200mV s-'; electrolyte = 0.1 M Et4NC104 (b) T = -70°C;
scan rate = 200 mV s-l; electrolyte = saturated Et4NC104.Reproduced by courtesy: Inorg. Chem.
13 (1974) 602. Copyright, American Chemical Society.
As can be deduced from the discussion in Section 12, simulation of each of
the voltammetric responses obtained under the relevant conditions will enable
kl [ f ( T ) ]to be calculated as a function of pressure,f (P),or temper, respectively. The direction of change of kl with pressure variation
tive or negative. However, as the temperature is decreased, the maghe homogeneous reaction kl is expected to decrease. In Fig. 2.71,
on of [ C r ( C 0 ) 5 1 ]in acetone is shown under conditions of cyclic
voltammetry [I 151 at a scan rate of 200 mV s-' at both 20°C and -70°C, using
on-isothermal cell (reference electrode temperature maintained at 20°C). As
temperature is lowered it may be easily observed that the reaction scheme,
products
which the second process is chemically irreversible converts at 20°C to the
ly reversible scheme,
hat is, at -70°C, the value of kl decreases to the point where it becomes too
slow to be rate-determining on the voltammetric time-scale.
Quantitative measurements of the relevant rate constant as a function of
pressure [I 141 will give the volume of activation for the appropriate process via
use of the relationship
where j stands for the relevant heterogeneous (ko) or homogeneous ( k l ) rate
constant. The AS* value is analogously given from measurement of kj as a
function of temperature and use of the equation
The values of A V* and AS* are extremely valuable in assigning the nature of
a mechanism (inner sphere, outer sphere, associative, dissociative, twist, etc .) ,
and complement the measurement of thermodynamically significant A V and
AS values as described in Section 4 of Chapter 1.
Voltammetric measurements at very high pressures are, of course, potentially,
highly dangerous and require the use of carefully designed cells that can be operated safely. Figure 2.72 gives an example of a variable pressure electrochemical
cell that has been used to study a range of systems by voltammetry. A recent
review by Swaddle and Tregloan [I 161 details the theoretical and experimental
aspects of voltammetry at high pressure.
Variable temperature studies are more readily implemented than variable pressure ones. Commonly, a non-isothermal electrochemical cell arrangement is
employed in variable temperature voltammetry in which the reference electrode compartment remains at ambient temperature, while the temperature
of the working electrode compartment is varied [I 17,1181. If thermodynamic
calculations are being undertaken, then, of course, the relevant junction potentials and reference electrode behaviour as a function of temperature need to be
extremely well characterized. The use of a non-isothermal cell provides some
simplification with respect to the measurement of the temperature-dependent
terms that are of voltammetric significance. Details of experimental arrangements and data that can be obtained use of variable temperature measurements
for a range of reactions are available in references [I 17-1 2 I].
18.1
General aspects
Either explicitly or implicitly, voltammetric studies referred to above apply to
B ne- process are soluble in the
a process where both species in the A
solution phase. However, if both A and B in thick solid or thin-film formats
+ +
Voltammetric studies o n solids
129
Vertical section
Pt auxiliary
electrode
Electrical connection
through to high-pressure seal
and detection
Kel-F threaded electrode mount
Teflon seal
Disc working electrode
Ag reference electrode
and compartmens
Vycor junction and piston
Kel-F cell body
Section at AA'
Auxiliary
electrode
Reference
electrode
Disc
workmg
electrode
. 2.72 Schematic representation of the high-pressure electrochemical cell. Reproduced by
courtesy: Inorg. Chem. 33 (1994) 6180. Copyright, American Chemical Society.
are attached to an electrode surface, which is in contact with a solvent (electrolyte) solution phase, then an electron-transfer reaction could occur within
e solid phase to give a diffusionless system, at least with respect to the solution
(electrolyte) phase in which the electrode is placed. However, in these circumstances, charge transport across the redox active solid(s)-electrode-solution
(electrolyte) interface must still occur to achieve charge neutralization so that
both ionic and electron transport processes must occur within the attached solid
phase(s) as part of the mechanism of the electrode process.
In the absence of a solution phase, mass transport process being a ratedetermining step, voltammetry at stationary or rotated macro- or microdisc
electrodes can be inherently similar, since it is the charge neutralization or
electron-transfer step within the solid that is likely to be rate-determining. This
is about the only simplification, with respect to the theory, that is achieved
when studying the voltammetry of a surface confined process, as almost every
other aspect of the problem becomes inherently far more complex than when
the redox components are soluble in the solution phase. For example, the solid
130
Princijdes of voltammetry
may be adsorbed or attached to the electrode by chemical or physical process to
give thin films of monolayer or sub-monolayer coverage or thick films where
each layer has a separate level of activity. Alternatively, the solid may be attached
to the electrode as arrays of microcrystals or exhibit completely non-uniform
or random forms of electrode coverage and, of course, several phases may exist
simultaneously on the surface. Numerous permutations and combinations of
electrode (metal, chemically modified, conducting polymer) and attached solid
have been reported and all have nuances which need to be considered when
generating a theoretical description of the voltammogram. Finally, of course,
chemical transformations may occur that are related to the EE, EC, ECc,t,I,,ic
classes observed when the redox active species are solely present in solution.
Clearly, the C step when the redox active species is attached to the electrode can
be a heterogeneous solid-solution-phase reaction rather than a homogeneous
reaction.
Since it is not practical to consider every aspect of the voltammetry of surfaceattached species, only the simplest case of the redox transformation of a thin
film (monolayer or sub-monolayer) of non-interacting surface confined material will be considered in Section 18. In Chapter 5, the significantly more
complex situation when microcrystals (thick film) are adhered to an electrode
surface will be considered. Treatment of this moderately simple thin-film case
provides a basis for comparison with studies where reactants and products of
an electrode process are soluble in the solution phase as in previous sections
of this chapter. From a fundamental viewpoint, metalloprotein-films of the
kind prepared by Armstrong and colleagues [I221 come as close to fulfilling
the conditions required for so-called ideal thin-layer voltammetry. The beauty
of electron-transfer metalloproteins is that they are designed to accommodate
very fast rates of electron transport over long distances, and the interactions
between molecules present in thin films of these molecules are relatively small.
Figure 2.73 provides an idealized representation of a film of redox active protein
ox + e-A
kred
red
red
kox
Cell solution
ox + e-
redox
centre
[EOflsurf
Fig. 2.73 Idealized diagram of a layer of redox active protein adsorbed onto an electrode surface. All centres are assumed to have the same
value and exhibit identical electron-transfer
characteristics. Provided by courtesy: F.A. Armstrong and J. Hirst, University of Oxford, England.
(EF)~~~
Voltammetric studies on solids
131
Cyclic voltammogram obtained for reduction of a thin film of the 'blue copper' proomonas aeruginosa Azurin, at a pyrolytic graphite electrode using a scan rate of 500V s-'
(pH = 8.0, T = O°C). The potential axis is V versus SCE. Baseline subtracted peaks (not to
le) are shown in the centre of the voltammogram. Reproduced by courtesy: Chem. Soc. Rev.
(1997) 169. Copyright, Royal Society of Chemistry.
adsorbed as a monolayer onto an electrode surface, while Fig. 2.74 gives an
erimentally obtained cyclic voltammogram for a thin film of azurin attached
to a carbon electrode. In this ideal situation, a homogeneous electroactive
olayer of a metalloprotein such as azurin is formed to give a coverage
tween 10-l1 and 10-l2 rnol a n A 2 . Thus, only minuscule quantities of
sample are required to obtain data related to interfacial electron-transfer kinetics, information on chemical reactions coupled to the charge-transfer process,
including biologically important catalytic reactions, and reversible potentials of
biologically important processes [122,123]. However, it needs to be recognized
that kinetic and thermodynamic parameters obtained from voltammograms of
surface-attached species need not be the same, even with respect to units, as for
the solution-phase electroactive species.
-2 Electron transfer i n ideal redox active thin$lms attached to
electrode su faces
et us assume that species A (now labelled 'red' when in film form) can be
oxidized to species B (now labelled 'OX' when in film form) while attached to an
electrode surface placed in a solution (electrolyte) medium without dissolution
of red or ox. Let us also assume that the entire thin-film layer consisting of redox
~~,
active species red and ox has the same single reversible potential [ E ~ O ] ~ and
that for a given potential E, all values of the charge-transfer rate constants for
oxidation k,,, and reduction kred are equal. That is, red and ox are considered
132
Principles of voltammetry
to be non-interacting redox active centres. By analogy with the solution phase
case, if k,, and kred are both fast enough to maintain equilibrium as the potential
is scanned, then the potential of the electrode with surface-attached material
will be governed by the Nernst equation in its surface attached rather than
solution-phase form. Thus,
is used to denote the fraction of the surface coverage
where the symbol rred
of the reduced form of the compound, and r,, the fraction of the oxidized
form of the compound. This is equivalent to stating that the activities of red
and ox are proportional to their mole fractions in the ideal thin film, assumed
to be present on the electrode surface, unlike the case of a pure solid where the
activity is usually assumed to be unity. Other symbols in eqn (2.74) have their
usual meaning. For this reversible process and with the above assumptions, the
on the electrode surface will be governed by the
distribution of I?,, and rXd
Nernst equation to give the fractional coverage shown in Fig. 2.75 as a function
of potential.
This thin-layer model effectively assumes that the solid is adsorbed or otherwise attached to an electrode to give a planar uniform surface and that attachment
to the surface of ox and red is analogous to that which occurs in the Langmuir
model of gas adsorption when all sites are identical and behaviour akin to a
monolayer is observed. In this circumstance, and when no interaction between
redox centres occurs, it is reasonable to assume that the oxidation and reduction components of the experiment vary uniformly with potential so that the
charge-transfer coefficient (Section 7 -2) is 0.5.
Fig. 2.75 The variation of fractional coverage of the oxidized species with respect to potential, for a reversible process when the ideal thin-layer model given by eqn (2.74) applies for the
process red + ox e- at 0°C. The potential axis is V relative to [ ~ f o ] , , ~ . Provided by courtesy:
F.A. Armstrong and J. Hirst, University of Oxford, England.
+
Voltammetric studies on solids
133
18.2.1 ideal thin-filmvoltammetry and reversible electron transfer
For a diffusionless reversible (Nernstian) reaction involving the simultaneous
transfer of n electrons at the same potential, the reaction schemes
red
+ ox + ne-
ox
+ ne-
or
(2.75)
red
+
+
nlay now be used to distinguish them from the A + B new or B new + A
reactions given for the solution-phase. In eqn (2.75), charges again are omitted
for simplicity as is the ion that must be transported across the thin-film-solvent
(electrolyte) interface, even though this ion transport rather than electron transfer may be rate-determining in a kinetically controlled process. The role of the
ion transport required for charge neutralization will become clearer in Chapter
5 when voltammetric studies on microcrystals attached to an electrode surface
are considered.
or a cyclic voltammetric experiment, assuming the scan rate is constant, the
change of potential is linear in time and, therefore, the equilibrium current (i.e.
ed for suitably low scan rate/high electron-transfer rates) is simply given
rate of change of the redox state of the redox centers shown in Fig. 2.73.
Thus (Section 13.1)
ence except for the sign of the current12,a symmetrical result is obtained for
oxidation and reduction components of the experiment. Thus, combining the
=
ernst eqn (2.74), and the relationship for the total surface coverage rtot
yielding, via use of eqn (2.79) and substitution of the expression for the scan
rate, v = dE/dt.
1 2 ~ hpositive
e
sign required by convention for the oxidation current and the negative sign
for reduction current can be achieved in several ways. Equations (2.77) to (2.80) represent
one of the formats used in the literature to obtain this outcome.
134
Principles of voltammetry
Normalized
-
Fig. 2.76 A reversible voltammogram for a surface confined ideal one-electron thin-film reaction
at 0°C. Note that the peak separation is zero, the current has been normalized to the peak current
and the potential axis is V relative to [ E ~ O ] , , ~ . Provided by courtesy: F.A. Armstrong and J. Hirst,
University of Oxford, England.
The peak heights [IpIsud,for oxidation and reduction can be calculated from
the expression (dI/dE,,,) to give
=
[P
' IS""
n2F~
VAL,,
which has a positive value for oxidation and a negative one for reduction.13
Furthermore, for this reversible process, the peak current for both the oxidation and reduction components of a cyclic voltammogram occurs at the potential
[Ep],,d where E = [E~O],,,~.A convenient form of presentation of the resulting
process may be obtained by plotting the normalized current value (I/[IplSud)
versus potential E to give the voltammogram shown in Fig. 2.76. Note that the
areas under the peaks are related to the total number of electrons transferred (or
number of adsorbed molecules) as in eqn (2.82)
Area under peak
v
= nFAr,,,
where v is the scan rate in V s-' and the left-hand side of eqn (2.79) gives
the number of Coulombs (normalized with respect to scan rate). The peak
131n deriving the thin film theory, the scan rate is assumed to be positive for oxidation and
negative for reduction. However, ifthe scan rate is only regarded as having a magnitude and
hence always positive, formats required as in solution soluble theory presented in Section 8,
then different formats are required for presentation of all equations containing the scan rate
term.
Voltammetricstudies on solids
135
%tidthat half-height width, W'p, is also an important experimental parameter
in thin-film voltammetry and for a reversible process
or about 9 0 / n rnV at 25OC.
is the case with solution-phase voltammetry, when equilibrium cannot be
maintained, the thin-film voltammetric process is no longer reversible. This situation will arise either when the scan rate (rate of change ofpotential) is increased
or when the electron-transfer rate is too slow for equilibrium to be maintained.
Under this non-equilibrium condition (within relevant experimental domain)
the process is now termed irreversible or quasi-reversible in the intermediate
regions when the response lies between the reversible and irreversible. Since
the rate of change of potential is now faster than the rate of adjustment of the
redox centres, the current lags behind the potential, and as in the solution-phase
case, enhanced separation of the reductive and oxidative peaks is observed. An
of a non-reversible process is given in Fig. 2.77.
the dependence of the peak separation on scan rate, and also from the
wave shape, it is possible to derive the magnitude and the potential dependence
of the thin-film charge-transfer rate constants. However, for a non-reversible
process, an analytical theoretical solution is not available and a 'finite difference' simulation is required, using a model of the potential dependence of the
electron-transfer rate constants, rather than the Nernst equation [122].
Normalized
current
'T
-
Potential
-77 An example of a non-reversible voltammogram for a surface confined ideal one electron
thin-film reaction at 0°C. Note that the oxidation and reduction peaks have separated, the peak
height has decreased and the peak width has increased relative to the reversible process in Fig. 2.76.
The current axis is normalized with respect to the reversible process. The potential axis is V relative
to [ E ~ O ] , , ~ . Provided by courtesy: F.A. Armstrong and J. Hirst, University of Oxford, England.
136
Principles ofvoltammetry
18.2.2 Thin-film voltammetry and non-reversible electron transfer using
Butler- Volmer theory
The simplest model used to describe electron transfer between a thin film of
adsorbed redox centre and an electrode is based on Butler-Volmer Theory,
and hence the behaviour is phenomenologically closely related to that used
when the compound undergoes oxidation or reduction in the solution phase.
This model, as noted in Section 7.2, is based on a transition-state approach to
the electron-transfer process and the equations for the electron-transfer rates
therefore takes on a simple Arrhenius-like form (eqns 2.84 and 2.85)
ko, = ko exp
(
-anF
kred = ko exp KT ( E - [':lSud))
The important kinetic parameter in this equation is ko, the rate constant at E =
[ E ~ ] , , ~It.is important to note that the units of ko, k,, and kred are now s-' rather
than cm s-' as is the case for the heterogeneous charge-transfer rate constant that
is relevant when electron transfer takes place between a solution-soluble redox
active species and an electrode. The symbol ko is used for surface rate constant
to distinguish from kO used for heterogeneous charge-transfer reactions, while
the analogous parameters to k,, and kred in thin-film voltammetry are kfet and kEt
(see Section 8.1.2) in solution-phase reactions. The parameter a (Section 7.2)
again represents the degree of conversion between the initial and final states in
= 1 . For a 'symmetrical'
the transition state, and it is assumed that (a,, ared)
electron transfer in the oxidative and reductive directions, a is equal to 0.5.
The dependence of the rate constants k,, and kred on potential, as predicted by
the Butler-Volmer theory is given in Fig. 2.78.
+
P
Potential
Fig. 2.78 The exponential dependence of the Butler-Volmer rate constants on potential for a
surface confined ideal one-electron thin-film reaction. The rate constant at [~fo],,& is ko( = 600 s-l)
and cx = 0.5. The potential axis is V relative to [E:],,~. Provided by courtesy: F.A. Armstrong and
J. Hirst, University of Oxford, England.
Voltammetricstudies on solids
137
The dependence of voltammograms for a surface confined ideal one-electron thin-film
reaction on scan rate as modelled by Butler-Volmer theory. As the scan rate is increased the eaks
separate and broaden as predicted for irreversible electron transfer. In this example ko is 600 s-' and
the scan rates are 1 V s-' (lower scan rates are superimposed), 10 V s-' , 100V s-I ,1000 V s-' and
10 000 V s-' and the potential axis is V relative to
Provided by courtesy: F.A. Armstrong
and J. Hirst, University of Oxford, England.
[~fo],,~.
igure 2.79 contains a set of voltammograms predicted when the Butlerer thin-film relationship, combined with theory contained in Section
17.2.3, is used for a range of scan rates, with the same ko value, with the current
normalized to the value predicted for a reversible surface confined process. As
expected, increasing the scan rate has the same effect as decreasing ko. At the
lowest scan rates considered, the Nernstian response is observed with the noralized peak height, K , being, of course, unity. As the scan rate is increased,
peaks separate, but also decrease in height and broaden. With use of Butlerlmer theory, the peaks reach an 'irreversible' limit (Fig. 2.80) (e.g. at O°C,
p = 116 mV and K = 0.74) beyond which no further broadening occurs.
This dependence on the reversibility can be more easily visualized by plotting
the relevant parameters as a function of loglo (v). As also, shown in Fig. 2.80,
with increasing scan rate, peak positions separate symmetrically from [ E O ] , , ~ ,
eak widths at half-height increase and peak heights decrease. Finally, this concept may be generalized by normalizing with respect to ko as shown in Fig. 2.81,
to provide results that hold for all values of scan rate and ko.
18.2.3 Application of the Butler- Volmer and Marcus theories to
thin-film voltammetry
As noted above, metalloproteins spontaneously attached to electrodes exhibit
almost ideal thin-film voltammetric behaviour (Fig. 2.74) and non-reversibility
may be approximated by Butler-Volmer theory [122]. In contrast, while selfassembled monolayer (SAM) structures of the kind shown in Fig. 2.82 may also
138
Principles of voltammetry
Fig. 2.80 Variation of peak position (a, Epeak),
normalized peak height ( b , ~ )and
, peak half-height
width (c, W1/2)
predicted by the Butler-Volmer theory with loglo(scan rate). Note the attainment
of an 'irreversible limit' in peak height and peak half-height width and that the peaks appear to
separate linearly with log(v) at high scan rate. The example shown is for ko = 600 s-l. Provided by
courtesy: F.A. Armstrong and J. Hirst, University of Oxford, England.
exhibit close to ideal thin-film behaviour [122], usually, for these systems, the
rate of electron transfer is much slower than for metalloproteins and under these
circumstances it has been reported that Butler-Volmer theory is not adequate
and the more complex Marcus theory needs to be adopted. Several groups have
studied the electrochemical kinetics of SAM structures consisting of ferrocenes
Voltammetric studies on solids
139
. 2.81 Peak positions (a), normalized peak heights (b), and peak half-height widths (c) as a
function of log lo(norma1ized scan rate,v/ko), which provides a unique parameter applicable to all
data sets. Other symbols are defined in caption to Fig. 2.80. Provided by courtesy: F.A. Armstrong
and J. Hirst, University of Oxford, England.
linked to a gold electrode surface through alkane thiols of varying chain length.
Typically, the electrode is a gold mirror surface prepared by sputtering on silica
wafers. Notably, Chidsey [I241 was able to detect the Marcus theory prediction
of the variation of electron-transfer rate with free energy and temperature, and
his conclusions have been supported and extended in studies by Murray and
co-workers [I 251 and by Weber and Creager [I 261.
140
Principles of voltammetry
X is redox group, e.g. fenocene
or functionality interacting with protein
X
X
X
X
X
X
Distance fixed by
number (n) of
(CH2),,in spacer
S
S
S
S
S
S
Au electrode
Fig. 2.82 Cartoon showing the structure of electrode surfaces modified with a monolayer of functionalized alkanethiolate. The group X may be redox active or a functionality such as carboxylate
that is capable of interacting with the surface of a redox active protein. Reproduced by courtesy:
Chenz. Soc. Rev. 26 (1997) 169. Copyright, Royal Society of Chemistry.
The voltammetry of osmium complexes adsorbed on platinum electrodes
[127,128] also supports the use of the Marcus model. Monolayer coverage
of [O~(bipy)~Cl(X)]
(bpy = 2,2'-bipyridyl; X = 4,4'-bipyridyl or analogues
with two or three CH2 groups spacing the pyridyl rings) is achieved at platinum microelectrodes utilizing the pendant pyridyl-N atoms from ligands X as
anchors. The relatively ideal response obtained with this system has allowed
detailed investigation of the effects of electrolyte, solvent, temperature and
electron-transfer distance to be undertaken. At a sufficiently large overpotential,
electrochemical rates were observed to become independent of driving force,
indicating the need to use Marcus rather than Butler-Volmer theory. However,
in this study, and in the work on metalloproteins [122], some features of the
voltammetry are not predicted by either the Butler-Volmer or Marcus theories.
For example, a finite peak separation always remains at the lowest scan rates and
the rate constant is approximately, but not completely, independent of ionic
strength, giving rise to the possibility that ion transport/binding effects cannot
be neglected, as has been assumed.
The work on the SAMs and osmium complexes, as noted above, supports the
use of the Marcus rather than the Butler-Volmer theory. The Marcus theory
takes into account the dependence of the rate of electron transfer on distance
of the redox active centre to the electrode, and the re-organization energy, A,
defined as the energy required to change the nuclear coordination from the
equilibrium position of the reactant and product, without allowing electron
transfer to occur. Marcus theory also requires [5] that: (i) the electron-transfer
reaction can occur both from and to all Fermi levels in the electrode, not just
between two defined levels; (ii) the rate of the electron-transfer reaction is
affected by the occupancy of the Fermi levels involved in the reaction; (iii)
the
.
free energy or driving force is a direct result of potential differences.
A direct comparison of theoretical linear sweep and cyclic voltammograms,
predicted on the basis of the Butler-Volmer and Marcus models, has been
provided in references [122-1 24,1291. Adopting the simple formalism used in
,
Voltammetric studies on solids
141
[I251 provides the following brief comparison of the two models14:
For a simple, reversible one-electron charge-transfer process between an
electrode and a surfice-attached redox active species
red
+ox + eko*
the dependence of kox and kKd (the forward and reverse potential-dependent
half-reaction-rate constants) on the overpotential, q, as expressed by the ButlerVolmer relations is:
k,, = ko exp (2i:T
These are the same relations given in eqns (2.84) and (2.85); here q = E [~fo],,,~,
kg is the Boltzmann constant ( F I R = elkB), e is the charge on the
electron, the charge-transfer coefficient, a , is assumed to be equal to 0.5, and
a one-electron charge-transfer process is considered. The analogous Marcus
relations [I251
to the Butler-Volmer equations when q/h << 1, with h being the
ization energy. Thus, the Marcus relations applied to electrode reactions
that, as 11 approaches h, the rate constants do not continue to increase
nentially with r,~(as they always do in the Butler-Volmer formulation) but
mize at q = fh and actually decrease at larger q (classicalMarcus 'inverted'
region) [I24,13O,l3 11. Examples of differences resulting from the two forms of
relations are summarized in Fig. 2.83, while the impact of the difference on
ersible voltammograms is shown in Fig. 2.84.
owever, use of eqns (2.89) and (2.90) are not rigorously correct since,
when electron transfers occur at a metal interface it is necessary to account for
the energy distribution of electrons about the Fermi level in the metal which,
eference [I291 is a review using conventional electrochemical symbols, which can be
consulted for a more detailed comparison of the different electron-transfer models.
142
Principles of voltammetry
Fig. 2.83 Comparison of electrochemical rate constants predicted by Marcus and Butler-Volmer
theories. (a) and (b) Marcus theory rate constants for small (0.2 eV) and larger (0.4 eV) reorganization energies respectively, the same plateau level (2 x lo5 s-') is achieved in each case; (c) the
Butler-Volmer dependence with ko the same as for (b). Reproduced by courtesy: Chem. Soc. Rev.
26 (1997) 169. Copyright, Royal Society of Chemistry.
Normalized 0.8
current
I
I
Fig. 2.84 A comparison of the waveshapes premcted by Marcus theory (-) and Butler-Volmer
theory (--) for an irreversible reaction with equal ko values. The Marcus reorganization energy
is 0.2 eV and the Butler-Volmer ol is 0.5. Potential axis is V vs [E;],,~. Reproduced by courtesy:
Chem. Soc. Rev. 26(1997) 169. Copyright, Royal Society of Chemistry.
following Chidsey [I 241, leads to
OX = P P ~ T
B
" exp{-(x - e(h - ~ ) / k ~ T ) ~ ( k ~ ~ / 4 h ) }
1
+ exp (x)
dx
(2.91)
Voltammetric studies on solids
143
x is the electron energy relative to the Fermi level, p is the distance-
ent electronic coupling between the electrode and redox sites, and p is
ity of electronic states in the metal electrode. p is anticipated [I321 to
pelld exponentially on distance, d, according to the relation
P = Fo exP
(-m)
(2.93)
is the coupling at zero edge-to-edge reactant separation, and the decay
,f3 depends on the details of the structure through which tunnelling
occurs and which has been reported, for example, to be -1.07 per CH2 group
for ferrocene alkanethiol monolayers on gold electrodes [125]. The relations in
eqns (2.92) and (2.93) are referred to as the heterogeneous Marcus equations to
distinguish them from the Marcus equations given by eqns (2.90) and (2.91).
Assuming values of two of the three parameters, ko, A, or p p , automatically
fixes the value of the third. Figure 2.85 gives the dependence of k,, on q , as
calculated from eqn (2.92) at 273 K, for a fixed value of ko(l.Os-l) and a series
of values of h (and a corresponding series of p p values). At small values of ( E d), these curves are the same as the exponentially increasing reaction rates
ted by the Butler-Volmer relationship (eqn 2.88) for the same ko value,
but at larger overpotentials and at smaller values ofh, eqn (2.92)predicts reaction
-85 Calculated log(k,,) versus (E- [E:],,~)at 273 K and ko = 1.0 s-l, based on heterogeneous
s kinetics (-), top-to-bottom h = 1.00,0.9,0.8,0.7,0.6,0.5,0.4,0.3,0.2, and 0.1 eV. (- - -)
Butler-Volrner calculation for ko = 1.0 s-* and a = 0.5. Reproduced by courtesy: Anal. Chem. 66
(1994) 3 173. Copyright, American Chemical Society.
144
Principles of voltammetry
rates that increase less than exponentially with (E - [E:],,~). At sufficiently large
q, and/or small A, the heterogeneous electron-transfer rate constants fold over
to become, at q > A, essentially independent of overpotential. However, the
heterogeneous electron-transfer rates do not decrease at q > A, as predicted
in homogeneous solutions in the classical Marcus 'inverted' region, because of
the continuum of electronic states in the metal electrode [124,125,133]. The
limiting rate constants (at large q) in Fig. 2.85 are determined by the value of
,xp, which in the series of calculations presented at constant ko and decreasing
A, decrease in accord with A.
Calculation of voltammograms in which the reaction rate is controlled by
either the Butler-Volmer eqns (2.87) and (2.88) or different forms of the Marcus
theory (eqns 2.89 and 2.90 or 2.91 and 2.92) are made using principles analogous
to those for the reversible case (Section 17.2.1). The difference is that the current
for first-order reaction of a diffusionless electroactive species is now given by
where the reaction rate constants k,, and kred are given by eqns (2.84) and (2.85),
or (2.87) and (2.88)' or (2.89) and (2.90) and To, and rred
are the instantaneous
surface coverages of the oxidized and reduced forms of the redox surface confined species, respectively. The applied potential, relative to E:, is q = q* v t
where q* is the initial potential, v is the scan rate (V s-l), and t is time. The
mass balance relationship in eqn (2.95) again applies so that
+
and
where r,,, is total surface coverage and the initial surface coverages
are given by the Nernst equation (see Section 17.2.1).
To calculate voltammograms, q can be changed in small increments of dq
( < l mV, sufficiently small so that calculated currents are independent of the
chosen dq) thereby allowing the reaction to proceed during time intervals of
dt = dqlv at rates according to eqns (2.84), (2.85), or (2.87)' (2.88), or (2.89),
(2.90). The instantaneous values of r,, and I-',d at each potential are then
calculated from their initial values and the oxidative or reductive charges passed
since initiation of the potential sweep. Finally, the current is calculated from
eqn (2.94).
It has already been shown (Fig. 2.84) that cyclic voltammograms calculated
with the simplest form of the Marcus model may differ substantially from
those predicted on the basis of classical Butler-Volmer kinetics. Figure 2.86
shows a comparison of Butler-Volmer voltammograms calculated for several
values of log[v/ko] as the applied q is swept from -0.2 to 1.O V. (Increasing v is completely equivalent to decreasing ko in these calculations as shown
previously.)
Clearly, the traditionally used Butler-Volmer formalism does not take into
account the nature of the electronic states involved. Further, at least two significant major assumptions that cause limitations in the application of the theory
Voltammetric studies on solids
145
Calculated, normalized voltammetric waves based on (bottom, a = 0.5) Butler-Volmer
kinetics and (top, h = 0.85 eV) heterogeneous Marcus kinetics, at 273 K, for left-to-right values of
log[v/ko] = -1,0,1,2, and 3. Reproduced by courtesy: Anal. Chern. 66 (1994) 3173. Copyright,
American Chemical Society.
are made. First, it is assumed that the reaction surface is linear, so that potential
energy varies linearly rather than parabolically along the reaction coordinate,
and, second, that all electrode energy levels, apart from the Fermi level are
ignored. As a result of these assumptions, the variation of electrochemical
rate constants with overpotential is incorrect: Butler-Volmer theory predicts
an ever (exponentially) increasing electrochemical rate constant, whereas in
fact (and correctly predicted by Marcus theory) the rate reaches a constant
level (independent of applied potential) at high overpotential, as shown in
Fig. 2.83. This 'plateau' is the electrochemical equivalent of the inverted region
described in homogeneous solution-phase electron-transfer reactions at large
driving forces [I241.
The differences in the theories translate into important aspects of voltammeus for example, Butler-Volmer theory predicts that even a very sluggish
electron-transfer reaction should exhibit a relatively sharp waveform (since the
electron-transfer rate increases exponentially), the irreversible limiting value for
Wl/2(width at half-peak height) being 6 2 . 5 1 mV
~ ~ at 25"C, which is 125 mV
0.5. Broadening could, of course, with this model, be attributed to
lity, an obvious example being kinetic dispersion in which different
orientations of molecules attached to the electrode surface display a spectrum of
rate constants [134]. However, Marcus theory predicts broadening even for an
ideal, homogeneous array, since the rate of a process having a small reorganization energy quickly ceases to respond to an increase in driving force. This is
146
Principles of voltammetry
illustrated in Fig. 2.84, which compares wave shapes expected, in this situation,
for ~utler-Volker and Marcus models having ;he &me value of ko. Essentially, the heterogeneous Marcus model predicts electrochemical rate constants
by summing over all the individual rate constants for each Fermi level in the
electrode, where the separate rate constants are the standard Marcus long-range
electron-transfer rate constants specific to two energy levels. The actual rate
of electron transfer from or to a specific level is influenced by the probability
of occupancy of that level as predicted by the Fermi-Dirac distribution 11241.
In principle, the cyclic voltammetric behaviour of systems having poor electronic coupling and low reorganization energy, such as metalloproteins, should
be better accounted for by the Marcus model than by the Butler-Volmer model,
particularly under conditions offast scan-rates where the peaks occur at higher qvalues and hence lie closer to the plateau region. However, because of the rather
fast values of ko for metalloproteins, data, usually, can be matched fairly well with
the Butler-Volmer theory [122]. In the case of the Chidsey type experiments
[124-1261, systems of the kind given in Fig. 2.82 are employed and because of
the large electron-transfer distances, ko values are much smaller than for metalloproteins and use of the Marcus model is always required. Figure 2.74 contains an
example of a voltammogram of surface-confined metalloprotein azurin which
shows that even at a scan rate of 500 V s-l, the peak-to-peak separation is not
large. In contrast, the separations in peak potential for the thiol system attached
to a gold electrode (Fig. 2.87) show a very significant scan rate dependence with
peak-to-peak separations being above 1 V at a scan rate of 1000 V s-' . A number
of recent papers in the area of thin-film voltammetry have considered doublelayer and other effects [ I D ,135-1371. Thus, the overview given above still does
not necessarily incorporate all the nuances. Further, only the ideal monolayer
coverage is considered so that the probability of real systems exactly matching theory presented in Section 18 is still problematical, and, not surprisingly,
several aspects of the theory [I 29,135-1401 are being debated in the literature.
2
18.3 Chemical reactions coupled to ideal thinjilm electron-transfer process
As is the case when solution-phase electrochemistry is being considered, chemical steps may accompany charge-transfer processes when solids are attached
to electrode surfaces. The majority of studies on adsorbed proteins have been
carried out using cyclic voltammetry which is the most 'visual' of the dynamic
methods, revealing the potential and time perspectives in an interactively simple
manner. The scope for studying EC and more complex catalytic mechanisms
associated with surface-attached films is illustrated in Fig. 2.88. Three situations
are considered in this figure, each of which is simplified by being reversible in
terms of interfacial electron transfer, but they differ in how the E step is coupled
by C steps to further chemical processes.
Figure 2.88(a) shows the voltammogram expected for a simple, reversible
electron-transfer reaction for a thin-film surface-confined species as described
in Section 18.2. Thus, as explained previously, oxidative and reductive components of the cyclic voltammogram are symmetrical, and currents reach a
Voltammetric studies on solids
I
I
I
-0.5
0.0
0.5
1.0
+-----,----0.0
0.5
-0.5
E(V) vs Ag wire
-0.5
0.0
0.5
E(V) vs Ag wire
1.0
E(V) vs Ag wire
E(V) vs Ag wire
-0.5
0.0
0.5
E(V) vs Ag wire
1.0
2.8'7 Effect of sweep rate on voltammetry of mixed Fc - C O N H - Cl5 - SH/HO - Clh- SH
~ )aqueous 1.0 M HCIOl at room temperature.
thin-film monolayer (r = 9.5 x lo-" mol ~ m - in
Currents have been normalized by the factor vQF/RT, where v is the sweep rate and Q the charge
under the voltammetric wave. Sweep rates were (a) 0.1 (b) 1.0 (c) 10.0 (d) 100.0, and (e) 1000V s-'.
Reproduced by courtesy: Anal. Chern. 66 (1994) 3164. Copyright, American Chemical Society.
aximum value at the reversible potential [E~O],,~,
thereafter decreasing to zero
as all of the finite number of redox centres are transformed from one redox
level to the other. The separation between peaks (AE,) is zero, the half-height
eak width Wl12 is 3.53RTlnF and the peak current I, = PZ~F~VAI',,,/~RT.
bviously, on the basis of these relationships, a two-electron (n = 2) reaction
gives rise to a much more prominent signal than a one-electron (n = 1)process
since Ip is proportional to n2 and Wl12 varies as l l n . Integration of the I-E
curve gives the number of electrons transferred and hence direct calculation
of the electrode coverage (peak area = nFvAr,,,). For a protein of molecular mass of 100 000 Da, maximum (monolayer) coverage is in the region of
3 x 10-l2 mol cm-2 [122].
148
Pvinciples of volturnmetry
.
.:.
..
..
'"'.
E (V)
-1
eox +-!+
red
+
OX*. ............. red*
product
substrate
Fig. 2.88 (a) The ideal response expected from a thin-film monolayer of adsorbed electroactive
species when the electron transfer is reversible. Current values are normalized to n 2 ~ 2 ~ ~ l ' / ~
(b) A response obtained when the electron transfer is followed by chemical conversion to a more
stable electroinactive species. (c) Conversion of a one-electron reversible wave to a one-electron
reversible catalytic wave on addition of substrate. Potential scale is V vs [~J],,fi. Reproduced by
courtesy: Chem. Soc. Rev. 26 (1997) 169. Copyright, Royal Society of Chemistry.
An illustration of what happens when the electron transfer is coupled to a
first-order chemical reaction is shown in Fig. 2.88(b). The example given for
an electron transfer, coupled chemical reaction, is the thin-film version of the
well-known square-scheme model [I 41,1421, where the electrochemical and
chemical reactions are assumed to occur separately. The kinetics of the chemical
transformations, as in all electrochemical mechanistic studies, are investigated by
varying the experimental time-scale, for example by using a range of scan-rates
under conditons of cyclic voltammetry. The situation depicted in Fig. 2.88(b)
is one in which the time required for interconversion of reduced red (isostructural with ox) to the thermodynamically more stable species red* (different
structure to red) is comparable with the experimental time-scale, and where
red* is electroinactive within the experimentally examined potential range, all
species being surface confined. Obviously, under these conditions the reductive
and oxidative components of cyclic voltammograms are no longer symmetrical
and the reduction peak potential now contains a contribution from the thermodynamics associated with interconversion of red and red* instead of solely
red
reflecting the properties of the elementary electron exchange ox ereaction. Hence, the reductive wave is shifted to a less negative potential and
sharpened as the position of equilibrium for the primary electron-transfer process
is modified by removal of red. In the example given, the re-oxidation of red*
is effectively gated because before any oxidation current can be observed, it
must convert back to red. However, as the scan-rate is lowered, further time is
+ +
Voltammetric studies on solids
149
available for red* to be in equilibrium with red. If a slow enough scan-rate can
be achieved so that equilibrium between red* and red is achieved, the response
will revert to being symmetrical and the reduction potential under these circumstances will reflect the thermodynamic distribution of species. Other complex
heterogeneous-homogeneous processes that have been considered include dissolution/crystallization and more complex surface bound ECE type reactions
in which the electron-transfer step may be connected with the solid or solution
hase reactions [I 43-1 471.
Figure 2.88(c) shows the voltammogram expected when reversible electron
transfer is coupled to a catalytic scheme involving reaction of red with a substrate present in solution. Thus, ox is electrochemically transformed to red (in
this case it is reduced) and then restored (re-oxidized) to the initial state by
the substrate, whose mass transport in the solution phase can be controlled
hydrodynamically by, for example, rotating the electrode. Hence, electron
ort is no longer confined to the adsorbed film, and the balance produced
en electrochemical and catalytic redox transformation results in a steadystate voltammetric response. It is important to note that rotating the electrode
not affect the voltammetry of completely surface confined reactions. How, the catalytic redox scheme considered in Fig. 2.88(c) does involve a
solution-phase component, so the rotation rate is important. A wide range
talloprotein thin-film substrate catalytic schemes have been investigated
and the relevant kinetic parameters determined from the dependence of
limiting currents on rotation rate and substrate concentration [148].
8.4 Nuances associated with adsorption
n the above discussion, thin films of adsorbed or surface-attached material
are strongly and irreversibly attached to the electrode surface. However, if an
electrode is placed in a solution containing redox active material, a wide range of
possibilities other than monolayer or sub-monolayer coverage of red and ox may
ccur. Generally, these alternative situations give rise to complex voltammetry.
or example, if adsorption of the oxidized or reduced species is weak, or if
only either the oxidized or reduced form of the compound are adsorbed, then
mixture of surface confined and solution-phase voltammetric responses may
e observed which will be related to the free energies of adsorption and the
nature of the adsorption isotherm(s). Figures 2.89-2.9 1 provide examples of
voltammograms observed when different levels ofproduct (eqn 2.96) or reactant
(eqn 2.97) adsorption are present, so that both surface confined and diffusioncontrolled components are simultaneously present.
-ne
-
A(so1ution)6B (solution)
OX (adsorbed)
(2.96)
+ne-
+A(so1ution)+B (solution)
-ne
red(adsorbed)
Sne-
(2.97)
150
Principles of voltammetry
Fig. 2.89 Examples of the influence of adsorption in (a, b) linear sweep and (c) cyclic voltarnrnetry
for the process given in eqn (2.96). (a) Effect of variation of concentration for a reaction in which
the product is adsorbed. Relative concentrations C : B :A are 1 : 4 : 16. (b) Variation of voltammograms with product strongly adsorbed as a function of free energy of adsorption. Free energies of
adsorption are 29.4, 24.8, 20.2, and 15.6 kcal mol-', respectively, in curves A to D. (c) Influence
of scan-rate on cyclic voltammograms with product weakly adsorbed. Relative scan-rates A : B : C
are 4 x lo4 : 2.5 x lo3 : 1. Adapted from: Anal. Chern. 39 (1967) 1514. Copyright, American
Chemical Society.
An electrode may be modified, as in Fig. 2.82, by adsorption of a long chain
of alkanethiolate or other related species not having an electroactive group.
If this occurs, and an electroactive species is in solution, then the diffusioncontrolled response can be modified, as, effectively, electron transfer to and
from the electrode surface has to occur over a rather long distance and also the
double-layer region of the new interface at the chemically modified electrode
is significantly different from that at the bare unmodified electrode. Miller
et al. [149,150] have considered this situation and voltammograms of the kind
shown in Fig. 2.92 are observed for reduction of cytochrome c when long
and short chain alkane thiolates are used to modify the electrode and reduction
occurs according to the schematic diagram in Fig. 2.93. Of course, the chemical
modifier may also act as an insulator which will prevent any Faradaic current
being observed, although if pin holes or other imperfections are present at
such an electrode surface, then a small amount of material may still reach the
Voltanzmetric studies o n solids
151
. 2.90 Theoretical cyclic voltammograms for electrode processes involving adsorption: (a) reactant weakly adsorbed; (b) product weakly adsorbed; (c) reactant strongly adsorbed; (d) product
strongly adsorbed. Dashed lines indicate behaviour for uncomplicated Nernstian charge transfer.
Reproduced by courtesy: Anal. Chem. 39 (1967) 1514. Copyright, American Chemical Society.
I
I
I
I
-0.3
-0.2
-0.1
0.0
Volt vs SCE
ig. 2.91 Cyclic voltammograms for reduction of methylene blue in buffered aqueous solution
M; (c) 0.40 x lop4 M. Reproduced
(pH 6.5; u = 44.5 mV s-l): (a)1.00 x lon4M; (b) 0.70 x
by courtesy: Anal. Chem. 39 (1967) 1527. Copyright, American Chemical Society.
Principles of voltammetry
Potential vs SCE (V)
Fig. 2.92 Cyclic voltammograms ofyeast cytochrome c at a modified gold electrode. (a) With a thin
SAM, HO(CH2)3SHmodified electrode, the voltammetry reveals a well-formed almost reversible
response. (b) When the SAM is thicker [HO(CH2)11SH] the rate of the electron transfer process is
much slower, as evidenced by the appearance of a large shift in potential of approximately -0.6 V.
The flattened shape of the reduction peak is predicted by Marcus theory. Experimental conditons:
v = 0.5~s-lA
; = 0.13 cm2; T = 0°C; 1.1mM cytochrome c, 2 m M phosphate buffer (pH 7.1)
with 1.O M KC1. Reproduced by courtesy: J.Am. Chem. Soc. 118 (1996) 7857. Copyright, American
Chemical Society.
bare electrode surface by radial diffusion to give voltammograms of the kind
described in references [I 5 1,1521. Alternatively, if cytochrome c is adsorbed or
attached [I531 onto the modified electrode surface to give the voltammogram
shown in Fig. 2.94, then the distance dependence is given as in Fig. 2.95.
19 Techniques
information on r
voltammetry of s
s associated with the
-attached species
Clearly, the theory applicable to the eletrochemistry of surface-attached species
is almost invariably complex. In the discussion above, no consideration was
given to the possibility that the rate of a process in a film is limited by the incorporation or expulsion of ions from the surface-attached layer as must occur to
achieve charge neutralization. Additionally, significant problems with correction for uncompensated resistance, capacitance current and the modified double
layer have not been addressed, nor have problems that arise when more than
a monolayer coverage of electroactive film is present, when more than a single phase is attached to the surface, when microcrystals rather than films are
attached to the electrode surface, or when electrocrystallization is coupled with
electron transfer. Thus, it is probably not surprising that comparisons between
experimental data and inherently approximate theory are often relatively poor
when electrochemical data related to surface-attached species are scrutinized
over very wide time domains.
Obtaining moleculav level information
153
3 Cartoon of adsorbed cytochrome c on a COOH-terminated modified gold electrode.
A crystalline region for a SAM HS(CH2)lsCOOH film exhibiting a 30" alkyl chain tilt angle is
depicted. Alkanethiol molecules are represented as end-capped cylinders with a C O O H terminus (black) and a thiolate (grey) attachment to the gold electrode. Polypeptide line diagrams of
cytochrome c molecule are shown in an electrostatically favoured orientation. For clarity, the
haem group has been blackened. Reproduced by courtesy: Electmchern. SOC.liztegace 6(4) (1997)
40. Copyright, The Electrochemical Society.
The determination of what occurs in a chemical or physical sense on an
electrode surface when electron-transfer reactions occur with surface-attached
species, requires the use of spectroelectrochemical (ESR, IR, UV-Visible,
etc.) techniques in surface suitable formats, surface analysis techniques [I541
(e.g. electron microprobe) and microscopy (e.g. electron scanning microscopy,
atomic force microscopy (AFM)).That is, a battery of surface science techniques
are available to address the 'surface' problems that are present. A difficulty in
relying solely on data obtained from voltammetrically based experiments is that
since the electrochemical response usually reflects only the average response of
numerous processes, the details of the chemical or physical changes that occur
154
Principles of voltammetry
Fig. 2.94 Diffusionless cyclic voltammograms of yeast cytochrome c monolayers prepared by
adsorption onto COOH-terminated gold electrode modifiers of different thickness. Conditions: pH
7.0 phosphate buffer of ionic strength = 50 mM; sweep rate is 100 mV s-'. (a) HS(CH2)loCOOH
(b) HS(CH2)5COOH.Reproduced by courtesy: Electrochem. Soc. Inte6ace 6(4) (1997) 40. Copyright,
The Electrochemical Society.
Fig. 2.95 Dependence of the electron-transfer rate on distance for horse cytochrome c adsorbed on
COOH-terminated alkanethiol gold electrode. The logarithm of the standard electron-transfer rate
constant is plotted versus the number (n) of rnethylenes in the SAM HS(CH2),COOH alkyl chain.
Reproduced by courtesy: Electrochem. Soc. Inte6ace 6(4) (1997) 40. Copyright, The Electrochemical
Society.
on the surface are unlikely to be unravelled by voltammetric studies alone. As
well as utilizing knowledge gained from solid state spectroelectrochemistry, it
can be noted that a mass change invariably occurs when a surface confined
solid undergoes a redox reaction, and therefore the ability to 'weigh' the mass
Obtaining molecular level information
155
change, via use of the electrochemical quartz crystal microbalance (EQCM),
can be a very powerful tool to apply as an aid to understanding the nuances of
solid state electrochemical studies.
The surface spectroscopic, surface elemental analysis and electron microscopy
techniques are well-established methods used in many branches of surface science and do not need special discussion in this book devoted to electrochemistry.
In Chapter 5, important knowledge readily gained in electrochemical studies
of surface-attached solids, by employing these surface science techniques in
conjunction with voltammetry, will be demonstrated. In this chapter, only
brief details are provided on the use of scanning probe microscopy techniques,
which enable changes at atomic resolution of species on surfaces to be observed
at the same time as voltammetric experiments are undertaken, and the in situ
M method that enables the changes in mass taking place on an electrode
e during the course of a voltammetric experiment to be measured. Both
these techniques have been responsible for significant advances in the understanding of additional processes that need to be taken into account relative to
those present in the idealized thin-layer model in Fig. 2.73. Indeed, as will
e seen later, it is amazing that when electrode surfaces containing attached
solids are examined in molecular or atomic level detail, as is now possible, that
models based on Butler-Volmer or Marcus theoretical concepts provide even
a reasonable description of the voltammetry!
-1 The use of scanning probe microscopies in electrochemistry
e ability to actually 'observe7what happens on an electrode surface during
e of a voltammetric experiment has become possible via application
family of scanned probe microscopies (SPMs) that were developed in
s (see references [155,156], for example). Prior to the advent of these
s, only methods such as electron microscopy and X-ray diffraction
were available to obtain resolutions approaching molecular dimensions, and
ese methods usually had to be applied in an ex situ mode.
The new class of microscopes that have become an invaluable tool of the electrochemist7strade are typified by the scanning tunnelling microscope (STM) for
which Binnig and Rohrer [157,158] received a Nobel Prize in 1986. These
techniques examine surfaces at very close range with a probe that may be
t a single atom across and may detect features on electrode surfaces at sizes
roaching, or even reaching, molecular dimensions.
Figure 2.96 represents a schematic diagram of the basic components of many
forms of SPM instrumentation. In an STM, the tungsten probe or tip is ground
so fine that it may consist of only a single atom [155,156]. Piezoelectric controls
manoeuvre the tip within a few nanometres ofthe surface of a conducting sample
2.97(a)). At these short distances there is overlap of the electron cloud of
toms of the probe tip, and of the nearest atom of the sample being probed.
n a small voltage is applied to the probe tip, electrons tunnel across the
gap, generating a minuscule tunneling current. X and Y piezoelectric controls
move the probe back and forth across the sample surface in a raster pattern. The
156
Principles of voltammetry
Feedback electronics
Light or current detection
Computer system
Fig. 2.96 Schematic diagram of the basic components of a Scanning Probe Microscope system
which consists of an X,Y piezoelectric scanner, a sensor to monitor movement ofthe probe, feedback
circuits for controlling the Z-piezo and a computer system on which results can be displayed and
analysed. Reproduced by courtesy: Coord. Chem. Rev. 200-202 (2000) 41 1. Copyright, Elsevier.
probe is moved up and down as it tracks the topography of the surface using a
feedback mechanism which senses variations in the tunneling current and varies
the voltage applied to a third, piezoelectric control, Z. The Z control moves the
probe vertically to stabilize the current and maintain a constant gap between
the probe tip and the surface. The image produced by this STM technique
is not the true topography, but a surface of constant tunneling probability,
although often the two are closely related. Figure 2.97(b) shows an STM image
of highly oriented pyrolytic graphite, while Fig. 2.98 illustrates what happens
to this graphite surface during cycling of the potential [159]. The ability to
'see' an electrode surface at molecular levels of resolution is a great aid to the
interpretation of electrochemical phenomena [160].
The STM microscope can be used only to create images of conducting materials. In contrast the atomic force microscope (AFM), shown in Fig. 2.99, does
not require a conducting material to create an image. In this case, the probe tip
is an atomically sharp diamond mounted on a strip of metal foil which is moved
over the surface. An AFM records contours of 'force' rather than tunnelling
current. The 'force' is the repulsion generated by the overlap of electron clouds
of the tip, with the electron clouds of the surface atom. However, caution in
the interpretation of data is always required in SPM measurements because the
tip may damage the surface and artefacts may occur in the measurement [I611.
The AFM can image a wider range of materials than the STM and can be used
under water or other solvents in an in situ mode with voltammetric experiments.
Figure 2.100 shows AFM images of Cao (buckyball) obtained before and during the course of reduction of the solid at a Cao-GC-acetonitrile (electrolyte)
interface [I 621.
Other forms of SPM have also been developed which include [155,156]
the laser force microscope (LFM), the magnetic force microscope (MFM), the
electrostatic force microscope (EFM), the scanning thermal microscope (detects
Obtaining molecular level information
157
z Piezo
ig. 2.97 (a) Schematic view of an STM. The tip, shown as a rounded cone, is mounted on a
piezoelectric X, Y, Z scanner. A scan (dashed line) of the tip over the sample can reveal contours
of the surface down to the atomic level. (b) An STM image showing carbon atoms in a sample of
highly oriented pyrolytic graphite. This is a line scan image displayed as observed by a viewer 45'
above the surface. An STM image is made up of a series of line scans, each displayed in Y from the
previous one, and displays the path the tip followed over the surface.
surface-temperature variations as small as ten-thousandth of a degree), the scanning ion-conductance microscope (SICM) and the scanning electrochemical
microscope (SECM) [I 631.
Since the SECM technique is based on detection with a microelectrode of
the kind used in voltammetric studies (Section 10) and, therefore, is of direct
interest to electrochernists in a fundamental sense, it is discussed in more detail.
The technique of SECM was developed by the Bard group [163,164] as a novel
Fig. 2.98 Constant-current STM images of the oxidation of highly oriented pyrolytic graphite
(HOPG) in 0.1 M H S 0 4 at 0.05V versus a Ag quasi-reference electrode. KIP = -50 mV,
l;,nn,l= -7 nA. Scan speed = 260 A s-l. (a) HOPG topography before oxidation. (b) HOPG
topography after 20 potential cycles. (c) HOPG topography after further oxidation of the surface.
Reproduced by courtesy: J. Phys. Chem. 92 (1988) 5563. Copyright, American Chemical Society.
Obtainin, rnoleculav level infoormation
159
Laser
Optical
deflection
sensor
El-
Cantilever
and tip
Image
Piezo
scanner
2.99 An AFM scans a sample with a shard of diamond mounted on a thin metal arm. The
electron cloud of the diamond tip (which may end in a single atom) presses against the clouds of
individual atoms in the sample, generating a repulsive force that varies with the surface relief. The
force deflects the tip, whose movements are monitored by a laser beam reflected from the top of the
arm to a photodiode sensor. A feedback mechanism responds to the changes in the beam's path by
activating a piezoelectric control, which adjusts the sample's height so that the deflection of the arm
remains constant. The sample's movements are translated into a surface profile. Unlike the STM,
the AFM can readily image electrical insulators. Provided by courtesy: R.G. Compton, University
of Oxford, England.
variant on the use of voltammetric microelectrodes. Exact positioning of the
icroelectrode and measurement of features of an electrochemical process in
e spatial region behind the tip of the electrode surface and the surface, allows
resolution of surface features down to approximately 30 pm. Figure 2.101 shows
a schematic diagram of the typical apparatus used in the SECM method.
n a typical configuration [I651 the 10-50-pm diameter microelectrode is
formed from a glass sheath with a narrow central pore. A platinum wire is then
sealed into the pore with epoxy resin, and the end is polished flat. The position of the electrode can be controlled at two levels. First, a three-dimensional
microstage allows the microelectrode to be positioned manually in increments
of a few pm. The electrode tip is moved into an initial position using the
icrostage and with the aid of an optical microscope the area of the substrate
to be investigated is selected. The piezoelectric translator then allows fine positioning of the electrode with sub-micrometer resolution. The electrode may
thus be accurately positioned over the substrate.
The SECM may be operated in a variety of modes. In the 'collection' mode
B n e is used to monitor the current, for
and when the solution process A
example, [ F ~ ( c N ) , ]+
~ [ F ~ ( c N ) , ] ~ - e-, the substrate is held at a potential
where the species A in solution will be oxidized to B. The electrode tip is
held at potential Et, which causes the species B to be reduced back to A. The
tip current is then monitored as a function of time at constant height above
the substrate. Variations of this technique are usually related to whether E, is
+ +
+
Obtaining molecular level information
161
held at a constant value while E, is stepped, as in a potential step experiment,
or whether the substrate is held at constant potential and E, is cycled. In the
'AC generation/collection7 mode, an alternating signal is generated and applied
to the substrate and a lock-in-amplifier is used to analyse the tip response with
reference to the substrate signal. In the 'feedback' mode, a potential sufficient
to drive a redox reaction is applied to the tip of the electrode and the current
is measured as a function of tip position. As the tip approaches an insulating
substrate, the diffusion layer around the tip is obstructed and the current response
is diminished. At conductive substrates, the redox reaction may be reversed, as
the electrogenerated species will be in excess relative to the bulk solution to
which most of the conductive domain will be exposed. This 'reflected9material
will then react at the electrode again, and the 'feedback' between the tip and the
conductive substrate will enhance the current observed. This 'feedback7mode is
used to image the conductivity of surfaces. Variations in the current or potential
at the tip of the microelectrode, during the scanning of the surface, produce the
images. Figure 2.102 provides a summary of the basic principles of SECM, while
Fig. 2.103 represents a conventional height image of a portion of a composite
el-F/Au surface which obviously contains a mixture of conducting and nonconducting substrates [I 661. The constant-current image of the same Kel-F/Au
surface is shown in the upper half of Fig. 2.103. The constant-current image
shows more detail than the constant-height image and provides the topography
with higher accuracy.
Scanning photochemical techniques [I 67-1691 use, as a source of contrast,
the thermal effect associated with a focussed laser beam on the current or potential variations associated with an electrochemical process. This technique is
essentially a temperature jump experiment and it is complementary to the SECM
method in the sense that an electrochemical process is inherently associated with
the ability to obtain a high-resolution image. For irreversible processes, the photothermal response is due to thermally induced changes in the rate of electron
transfer, with contrast arising from local variations in dark current density, activation energy, and temperature change due to spatial variations in the absorption
coefficient of light. For reversible processes, the signal arises from the thermally
induced shift of the standard electrode potential, with contrast arising from
variations in the local diffusion-limited current or temperature change. Blurring and shadowing effects arise from the relaxation of the local temperature in
the electrode and diffusion field in the electrolyte as the scanning spot moves on.
ig. 2.100 In situ AFM images obtained at 0 V versus Ag/AgCl in acetonitrile (0.1 M Bu4NC104)
showing the changes in the morphology that occur when CbOmechanically adhered as a crystalline
~
solid to a G C electrode is subjected to redox cycling experiments (scan rate 0.1 V s-l). C 6 microcrystals after (a,b) 10 potential cycles between 0 and -1.OV (c,d) after two further cycles between
0 and -1.2 V, (e,f) after eight further cycles between 0 and -1.2 V (g) after three additional cycles
between 0 and -1.6 V. Scale (a) 1.5 pm (b) 0.9 ym (c) 2.2 pm (d) 0.6 pm (e) 3.0 pm (f) 1.4 pm, and
(g) 0.2 ym. Reproduced by courtesy: J. Phys. Chem. B 103 (1999) 5643. Copyright, American
Chemical Society.
162
Principles of voltammetuy
(2)
Bipotentiostat
I
Piezoelectric
translator
Electrode
insulated
up to tip
Reference
electrode
I
Fig. 2.101 Schematic diagram of the apparatus used in a SECM. (a) Instrumentation. (b) Expanded
version of microelectrode. Provided by courtesy: N. Stevens, Monash University, Victoria, Australia.
Insulating substrate
Conductive substrate
Fig. 2.102 Basic principles of SECM: (a) When the microelectrode (ME) is far from the substrate (d
the tip-to-substrate distance is much larger than electrode radius re), diffusion of A leads to a limiting
steady-state current IT,, = 4nFDr,[Alo (Section 10.1). (b) When the ME is near an insulating
substrate, hindered diffusion of A leads to IT < IT,,. (c) When the ME is near a conductive
substrate, positive feedback of A to the tip leads to IT > IT,oo.Reproduced by courtesy: Physical
Electrochemistry (ed. I. Rubenstein), Marcel Dekker, New York, 1995, p. 210.
19.2 T h e electrochemical quartz crystal microbalance
Mass changes can occur at an electrode surface under a wide range of conditions.
For example, the deposition of metallic lead from aqueous solution occurs as in
eqn (2.98).
ig. 2.103 (a) SECM surface plot of the microelectrode tip current recorded during a
constant-height scan. Scan size is 100 pm x 100 pm. Vertical axis is relative tip position in pm. (b)
SECM surface plot of the Z-piezo positioner voltage recorded during a constant-current imaging
scan. Scan size is 100 pm x 200 pm. Vertical axis is relative tip position in micrometres obtained
from the piezo voltage. The tip position was modulated at a frequency of 160 Hz with a 100-nm
peak-to-peak modulation amplitude. The negative and positive current reference levels were set at
500 and 780 pA, respectively, and IT,oo
was 80 PA, implying maintenance of about 2-ym tip-substrate
spacing over both conductive and insulating regions. The images were made with a 2 pm diameter
Pt tip, using a 2.1 rnM solution of [ R U ( N H ~ ) ~ in
] ~a+p H 4.0 buffer as the mediator species, and
a tip raster scan of 10 pm s-l. See Fig. 2.102 for definitions of symbols. Reproduced by courtesy:
Anal. Chern. 65 (1993) 1373. Copyright, American Chemical Society.
a
164
Principles of voltammetry
For this reaction, Faraday7slaw states that the charge consumed (Q) is related
to the mass deposited by the expression
n F (mass)
= molecular weight
mass =
(2.99)
(molecular weight) Q
nF
Typically, in a voltammetric experiment, the charge, Q, is only a few hundred microcoulombs (e.g. 10 pA for 10 s). Since F is 96 485 C mol-l, masses of
only up to about 5 pg are deposited under conventional voltammetric conditions. However, a metal deposition reaction involves much larger mass changes
than encountered with most other forms of reaction involving interaction of
a solid with an electrode. Consequently, the mass change that occurs during
the course of a voltammetric experiment is oiten less than a microgram, so that
very sensitive mass detection is needed. The quartz crystal (QC) microbalance
can be readily used to measure mass charges in the nanogram region. Hence,
when electrochemistry is combined with this mass sensor, an EQCM is constructed. In the EQCM experiment, a metal electrode (e.g. gold) is deposited
onto a Q C microbalance. Then if, say, Pb is electrochemically deposited as in
eqn (2.98), the interface looks like that in Fig. 2.1O4(a), whilst a voltammogram
and EQCM data for deposition and stripping of lead are shown in Fig. 2.104(b)
and (c) respectively.
The Q C microbalance itself is a piezoelectric device in which a crystal of
alpha quartz deforms when placed in an electric field (oscillates in an alternating
electric field). In order to fabricate a mass sensor, a large, single, alpha Q C is cut
into a wafer at a certain angle with respect to the crystalline axes to optimize
the piezoelectric activity. The application of an alternating electric field to the
crystal causes the crystal to oscillate. If the alternating electric field is applied
at the correct frequency, the crystal will oscillate at a resonant frequency, fo,
which depends on the thickness of the crystal.
where tq is the thickness of the crystal, p, is the density of quartz (2.648 g ~ m - ~ ) ,
and p, is the shear modulus (2.947 x 10" g cm-' s-').
19.2.1 The Sauerbvey equation
For thin, uniform, and rigid films attached to a QC, the following relationship
exists
Arn
relating the observed change in frequency to the change in mass at the surface
of the crystal. The term in brackets is a constant for a certain thickness of crystal
. 2.104 (a) Schematic diagram of electrode interface in an EQCM experiment in which metallic
lead is deposited onto a gold-coated quartz surface. (b) Cyclic voltammograms showing deposition
and stripping of lead in aqueous 0.1 M acetic acid/O.l M sodium acetate. (c) EQCM data obtained
simultaneously with cyclic voltammograms shown in (b). v = 100 mV sP1 and [Pb2+] = 8 mM.
Provided by courtesy: P.J. Mahon and G.A. Snook, Monash University, Victoria, Australia.
166
Principles of voltammetry
and the equation is simplified to
where Cf is the sensitivity factor for the crystal. Thus, a mass increase produces
a decrease in the observed frequency. For a 10-MHz crystal, Cf = 2.26 x
lo8 Hz cm2g-l. Equation (2.103) indicates that an observed frequency change
of 1 Hz corresponds to a mass change of 4.4 ng cmP2.In principle, the need for
calibration of each crystal is usually unnecessary because the 'sensitivity factor7
is only dependent upon well-known physical constants. In practice, calibration
is usually undertaken via the use of a well-defined metal deposition reaction
[170,171], because the frequency depends on additional factors, and a more
rigorous equation is
where the terms on the right-hand side of the equation refer to the effect of mass,
pressure, viscosity, and roughness, respectively. In electrochemical experiments
it is generally assumed that all factors, except mass, are constant. It is also assumed
that species are rigidly attached to the electrode surface.
In an EQCM experiment, the Q C may be mounted between two spring
clips and electrical contact is made to the connector holding the clips and
the crystal. A schematic diagram of an EQCM instrument of this design is
shown in Fig. 2.105. In addition to the standard components required for an
electrochemical experiment, there are two additional elements. The Oscillator
is a circuit which enables the crystal to oscillate without interfering with the
electrochemical experiment. The Frequency Counter measures the frequency
developed in the Oscillator (in some instruments a frequency-to-voltage converter is used). In a common instrument design, the working electrode is wired
so that it is true ground and the voltage between the two metal surfaces of the
crystal also varies relative to this ground point. The oscillator circuit develops a
voltage of approximately 0.4 V, peak-to-peak, across the crystal [17 1,1721.
A typical Q C with deposited metal electrode arrangement is shown in
Fig. 2.106. The deposited metal can be any metal that adheres strongly to the
,
Current in
nanoamps
Potentiostat
I
Counter
electrode
Quartz
45
Reference
electrode
nanograms
\Gold
working electrode
Fig. 2.105 A block diagram of an EQCM. Provided by courtesy: P.J. Mahon and G.A. Snook,
Monash University, Victoria, Australia.
Obtaining molecular level information
167
Typically the crystal has a
diameter of 13 mm.
T o resonate at 10 MHz the
crystal is cut so that it is
0.17 mm thick
Electrical
Electrical contact
Metal is deposited on both
sides of the crystal so that
the alternating electric field
can be applied across the
crystal.
-106 The gold-coated quartz electrode surface used in an E Q C M experiment. Provided by
courtesy: P.J. Mahon, Monash University, Victoria, Australia.
quartz surface. Typically Au is used, but other metals such as Ag, Al, Cu, Zn,
, and Pt have also been employed. It is also possible to deposit one metal
of another. Usually vapour deposition techniques are used to deposit
the metal onto the quartz, but any technique can be used as long as the crystemperature stays below 573°C. Above this temperature, the piezoelectrical
activity is destroyed.
ri&nally it &as thought that the application of the Q C microbalance in
solution-phase electrochemical studies would fail due to heavy damping of the
crystal by viscous liquids. The first electrochemical application involved the
three stage ex situ process of (i) measuring the frequency of the dry crystal,
(ii) deposition of a small quantity of metal in solution, and (iii) measuring the
frequency of the dried crystal. The first in situ application occurred in 1981
[170]. In this pioneering study, use of Faraday's law and the Sauerbrey equation
enabled a simple relationship between the change in frequency (Afm) and the
amount of charge (Q) passed to be obtained.
Af = -
(molecular weight) Cf Q
nF
he charge can be obtained by integrating the current (i.e. Q = J,fI dt).
Alternatively, the current can be related to frequency in the following way
I=-
d(Af,)/dE nFv
(molecular weight) Cf
ere v is the scan rate (V s-l).
19.2.2 Applications of the electrochemical quartz crystal microbalance
umerous applications of the EQCM method have been published in the last
ecade and have been reviewed in references [172-1751. In Chapter 5 the
168
Pn'ncipler of voltammetry
use of this method will be demonstrated with microcrystals adhered to electrode surfaces. In this chapter, application of the EQCM to the problem of
metal deposition has already been referred to in Section 18.2 (see Fig. 2.104).
To conclude this brief overview of the EQCM method, applications to studies on adsorption and desorption processes and conducting polymers will be
considered.
As noted above, a mass change will necessarily accompany an electron-transfer
process at an electrode surface only when either the oxidized or reduced species
is attached to the electrode surface, as occurs in the ~b~+(solution) 2e- +
Pb(rneta1) deposition process. However, a mass change also occurs when either
the oxidized or reduced forms of solution-soluble species become adsorbed
onto an electrode surface (see eqns 2.92 and 2.93). For example, consider the
bromide salt of the ferrocene derivative C H ~ ( C H ~ ) , ( C H ~ ) ~ N + CcontainH~FC
ing a long alkyl ammonium chain of length x attached at one position. The
unoxidized ferrocene (Fc) compounds are strongly adsorbed to gold surfaces
and studies on the interactions with this electrode surface have been made
with an EQCM [176,177]. When x = 12, and a dilute 22 yM solution of
CH3(CH2)12
( C H 3 ) 2 ~ + C H 2isFused,
~ oxidation of Fc(adsorbed) to solution)
leads to an EQCM frequency increase which corresponds to a mass decrease
at the electrode surface (Fig. 2.107) which is consistent with dissolution of the
oxidized FC+ form of the compound. Adsorption occurs when the potential
returns to less positive potentials, where Fc(adsorbed) is reformed by reduction
of solution). Under these conditions, eqn (2.104) applies.
+
However, if the concentration is increased to 500 yM (Fig. 2.108), then not all
the oxidized ferrocene dissolves and there are current contributions arising from
formation of both Fc+(adsorbed)and solution), because the reaction occurs as
in eqn (2.108).
This situation is akin to that described in Section 17.4. In this case, the molecular weight of the material cannot be calculated by simple use of the Sauerbrey
equation because an anion X from solution electrolyte enters the monolayer to
compensate for the increase in charge during the Fc(adsorbed) -+ Fc+(adsorbed)
component of the oxidation process, and so the frequency increase on oxidation corresponds to the difference of mass loss from dissolution and mass
Obtaining moleculav level infoormation
.080
I
,080
.240
.400
Potential (V vs Ag/AgCl)
I
I
I
I
I
.240
.400
Potential (V vs Ag/AgCl)
.560
I
.560
,720
I
I
I
.720
2.107 Simultaneous recording at a scan rate of 50mVs-' of (a) a cyclic voltammogram, and (b) EQCM frequency response for oxidation of low coilcentration (22 pM) of
CH3(CH2)12(CH3)2~f
CH2Fc in aqueous 1.0M H3P04. Reproduced by courtesy: Langmuir 5
(1989) 671. Copyright, American Chemical Society.
increase from incorporation of X and of course, vice versa, for the reduction
component of the cyclic voltammogram.
The fabrication and use of polymer modified electrodes has attracted much
attention in recent years [178,179]. Redox conversion in polymer film modified
electrodes is associated with the simultaneous presence of electron transfer and
exchange of electrolyte ions to maintain charge neutrality. The ionic switching
component of the process is elegantly monitored by the EQCM method. For
example, the oxidation of the ferrocene sites in poly-(vinylferrocene) in contact
Pvinciples of voltammetvy
.080
,240
.400
E (V vs SCE)
.560
.720
.080
.240
.400
E (V vs SCE)
.560
.720
I
I
Fig. 2.108 Simultaneous recording at a scan rate of 50 mV s-' of (a) a cyclic voltammogram and
(b) EQCM data for oxidation of a high concentration (0.5 rnM) of C H ~ ( C H ~ ) ~ ~ ( C H ~ ) ~ N + C H
in aqueous 0.2 M Li2S04 (pH 3). Adapted from: Chemically Modijed Electrode Sufaces in Science and
Industry, Gordon and Breach, New York, 1988, p. 377.
with solvent (electrolyte) creates a highly charged polymer coating. The frequency in the ECQM experiment is observed to decrease during the oxidation
process and subsequently increases to the original frequency during the reduction cycle (Fig. 2.109). The mass change corresponds to an influx of anions into
the polymer during the oxidation process, in order to maintain electroneutrality
throughout the polymer. During the reduction of the ferrocenium groups, the
anions are expelled into the solution. In the case shown in Fig. 2.109, the mass
change corresponded directly to the uptake of the electrolyte PF, anion [180].
However, use of the Sauerbrey equation requires that the solid should be
attached in an acoustically rigid manner. Clearly this assumption cannot always
be met in redox polymer electrodes, since such polymers usually have elastic
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9 Simultaneous measurement at a gold electrode (scan rate = 10 mV s-l) of (a) cyclic
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Principles of voltammetry
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uctio
most redox active compounds exhibit only a small number of
tric processes within the potential range available in the solvent (elecedium chosen for the study. However, polyoxometalates [I] represent
an electrochemist's dream class of compound because it is not uncommon
to have multiple charge-transfer processes taking place during the course of
a single voltammogram [2,3] and several examples, where more than twenty
electrons [4,5] can be added during the course of electrochemical reduction of
lyoxometalate, have been established.
o unravel the vast array of kinetic, thermodynamic, and synthetic information available from a voltammetric study of polyoxometalate compounds
res the application of the majority of procedures described in Chapter 2.
, the use of transient and steady-state studies at numerous classes of eleche use of spectroelectrochemical techniques, electrochemical synthesis,
ulation techniques, as well as photoelectrochemica1and electroanalytical
ions, abound in the polyoxometalate electrochemical literature.
n this chapter, detailed descriptions of the very rich redox chemistry of sevoxometalate compounds are presented. In particular, the contributions
ctrochemist's arsenal of techniques described in Chapter 2 are highthe establishment of the detailed picture now emerging on the redox
of this important class of compound. However, studies reported are
confined to cases where both the oxidized and reduced forms of the polyoxometalate species are soluble in the solvent(electrolyte) of interest. Examples
of voltammetric studies where solids are attached to electrode surfaces will be
emphasized in Chapter 5.
olyoxometalates form a large class of inorganic cluster-like mineral comounds, known for well over a century, that have become of increasing
178
Illustvating basics ofvoltammetvy
interest in recent years due to recognition of their extraordinary topological
and electronic versatility, and due to their current significance in disciplines
as diverse as oxidative catalysis [6-81, biology, medicine, and materials science
[I ,5,9-1 I]. They are formed principally from combinations of oxide ions and
transition metal cations in their do electronic configurations, held together only
by metal-oxygen bonds [I ,5,11,12]. Based on restrictions placed on acceptable
ionic radii ratios, only a limited number of metal-ions form these complexes
(v', ~ b ' , ~ a ' , MO", and w'' ) with the latter two forming by far the largest
number [I ,2,5,ll,121. The principal units that make up most polyoxometalates
Atomdesignations: M = e 7 O = O , P = @ , S i = O , W = O a n d X = @
Fig. 3.1 'Ball and stick' drawings of representative structural families of polyoxometalates:
(a) hexametalate structure, [M6OI9lX- (the charge, x, depends on M) ; (b) Keggin structure,
[XW12040]X-(X depends on the heteroatom, X); (c) Dawson structure, [X2wlX062]x-(X depends
on X); (d) Pope-Jeannin-Reyssler structure, [MPsW300110]x-(x depends on the central metal-ion,
M). Reproduced by courtesy: Chern. Rev.98 (1998) 327. Copyright, American Chemical Society.
Coupled electron- and proton-transfer reactions
179
6 octahedra. The 'polyhedral notation', commonly used by investigaare
tors who study polyoxometalate and related materials, including metal oxides,
the dominance of these octahedral units [9]. Deviations of the M 0 6 units
ure octahedral symmetry can be substantial, with the metal always being
ed towards the terminal or doubly bonded oxygen atoms. The M 0 6 units
in polyoxometalates can be linked together by a single oxygen atom, termed a
nkage. Octahedra joined in this manner are said to be 'corner sharing'.
present in the same polyoxometalate molecule as the corner-sharing
a are the edge-sharing octahedra. In the latter, two M 0 6 units are
gether by two oxygen atoms, termed a di-p-0x0 linkage.
Two large sub-categories of polyoxometalates exist: the 'isopoly' and the
'heteropoly' compounds. The former are constituted of only metal and oxygen
atoms, while the latter contain one or more p-, d-, or f-block 'heteroatoms'
at well-defined geometrical sites in the molecule, in addition to the metal and
oxygen atoms [2]. Up to 75 per cent of the elements in the periodic table
can function as heteroatoms in polyoxometalates. The heteroatoms in the heteropoly compounds can reside in either buried (not solvent accessible) or surface
ent accessible) positions in the polyoxometalate structure [2].
all and stick' forms of structural representation are given for a range of
polyoxometalate complexes in Fig. 3.1. Given the fact that the transition metals
sent in polyoxometalate complexes are generally in high oxidation states
oV1,W" etc.), and that there are non-equivalent environments for the
within the structures, it is not surprising that this class of compound
e extensively reduced under voltammetric conditions to give a rich array
-transfer reactions
aps the most widely used technique used to elucidate voltammetric mechas is cyclic voltammetry at macrodisc electrodes (see Section 8 in Chapter 2).
ost all recent voltammetric studies on polyoxometalate systems contain
les of the use of this technique, and if data are reported in aqueous
then the pH dependence is likely to have been examined. However, only rarely are voltammograms simulated according to the postulated
mechanisms reported, as recommended in Chapter 2. In this section simulationerimental comparisons for the initial processes observed in the reduction,
in aqueous media, of the ~ - [ P ~ w ~ ~and
o ~a ~ ]- ~ -[ ~ ~ systems
~ ~are~
presented [13] in order to demonstrate the value of undertaking a systematic
interactive experiment-theory approach to achieve a realistic comparison of
' ~ d a p e dwith permission from Anal. Chem. 71 (1999) 3650. Copyright, American
Chemical Society. Additional details may be obtained by consulting this reference.
0
~
180
Illustvating basics of voltarnrnetvy
experimental and theoretical voltarnrnograms under circumstances where the
mechanism is inherently complex. Limitations that may be encountered in
simulation-experimental comparisons also are identified.
3.1 Reduction ofa -[P, wI8
06#Six one-electron processes are observed for the reduction of a-[P2W18062]6(Structure C in Fig. 3.1) in aqueous electrolyte media at pH values above six
[14]. The processes under these conditions are summarized2in eqns (3.1)-(3.6).
However, in the presence of 6 M acid, there are three, apparently, two-electron
reduction events [I 4,151 instead of six, single one-electron processes. Typically,
the processes obtained in strongly acidic conditions have been written as in
eqns ( 3.7)-(3.9), although, as will be seen from subsequent discussion, protonation steps accompany electron transfer in acid media and more informative
descriptions of the coupling of electron- and proton-transfer reactions can be
developed via simulation-experiment comparisons.
3.1.1 Initial considerations
The cyclic voltammetric behaviour at a 3-mm diameter glassy carbon (GC)
macrodisc electrode is presented for the first four one-electron charge-transfer
processes (Scheme 3.1) over the pH range 5.2-1 -0 in aqueous 0.5 M NaC1
solution (Fig. 3.2).
Over the pH range 5.2-1 .O, the wave shapes and peak positions of the first two
couples 0 / 1 and 1/2 do not change [15], whereas processes 21'3 and 3/
lesce (Fig. 3.2). Thus, at some pH values, very complex cyclic voltammograms
2~somericnotation as in ~ ~ ! - [ P ~ w ~and
~ o~ ~~ ~! ]- ~
[H
- ~ W ~is ~commonly
O ~ ~ ] ~omitted
in
subsequent discussion for convenience and also because the isomeric forms of reduced
species are usually unknown.
Coupled electron- and poton-transfer reactions
I
-0.75
I
I
I
-0.5
-0.25
0
Potential (V vs Ag/AgCl)
I
0.25
181
I
0.5
Potential (V vs Ag/AgCl)
. 3.2 Cyclic voltammograms obtained at a scan rate of 100 mV s-I for reduction of 1 rnM
[ P ~ w ~ ~ oat~a GC
~ ] macrodisc
~ electrode in aqueous 0.5 M NaC1; (-) experiment; (- - . ) simulation. The poorer fitting at the switching potential is caused by the influence of additional processes
at more negative potentials. Reproduced by courtesy: Anal. Chem. 71 (1999) 3650. Copyright,
American Chemical Society.
are observed, but it will emerge that this does not correspond, in this case, to the
onset of either electrochemical or chemical irreversibility. The simulation will
e shown to be critical to gain an understanding of the pH-dependent factors
that govern the coalescence of these processes.
Illustrating basics of voltammetry
Scheme 3.1 Summary of the first four basic charge-transfer processes associated with reduction of
[ P ~ w ~ ~
in o
aqueous
~ ~ ]0.5
~MNaCI.
Experimental E;,? values are given in parentheses in rnV
versus Ag/AgCl.
3.1.2 Stepwise approach to the simulation
Even the simulation of the four uncomplicated one-electron charge-transfer
processes at pH 5.2 (Fig. 3.2(a)) requires the input of a large number of parameters, some of which are derived from the considerations in points (1) to (11)
below:
(1) Experimental conditions The potassium form of the polyoxometalate salt
was dissolved in an aqueous solution of 0.5 M NaCl at a known concentration
(1 or 5 mM). The pH value was adjusted with aqueous NaOH or HCl.
(2) E;,2-values for simple one-electron steps These were estimated from cyclic
voltammetric data as the average value of reduction and oxidation peak potentials [ ( E ; ~ E y ) /2] (Section 8.1 in Chapter 2) in solutions of sufficiently high
pH (Table 3.1, Scheme 3.1).
(3) Uncompensated resistance arisingfvom I R drop between the working and reference electrodes A value of 200 ohm was measured by impedance techniques
and found to be suitable for use in simulations over the scan rate range of
50-1 000 mV s-* .
Farad cm-2 was assumed,
(4) Double-layer capacitance A value of 2 x
on the basis of agreement between experiment and theory of the background electrolyte. Parameters (3) and (4) were optimized values deduced from
measurements at higher pH values where protonation effects were minimal.
(5) Temperature The simulation temperature was set to 298.2 K. Small differences with experimental data obtained at temperatures between 20°C and
23°C are not critical for this system.
(6) Charge-transfercoeficient a Assumed to be 0.5.
(7) Area of the working electrode This parameter was determined to be
0.0893 cm2 by application of the Levich equation (Table 2.3) to rotating-disc
voltammograms at the GC electrode obtained for oxidation of solutions of
1.0 mM ferrocene in CH3CN(0.1M Bu4NC104).
Each process
(8) Standard heterogeneous charge-trans& rate constants at E ; , ~ ( ~ O )
was considered to be essentially diffusion controlled in the scan rate range v =
50-1000 mV s-'. Consequently, the ko values were assumed to be fast and were
set at 1cm s-' for the purposes of the simulation. This implies that the processes
were assumed to be reversible on the cyclic voltammetric time-scale.
+
Table 3.1 Parameters used in the simulation of the cyclic voltammetry of 1 rnM solutions of polyoxometalate in aqueous
0.5 M NaCl
Couples
[ P ~ w ~ ~ o(Scheme
~ ~ ] ~ -3.4)
2/3
3/4
[H~wI~o~c$- (Scheme 3.5)
3h/4h
Of/1'
11/2'
lth/2'h
-490
-555
-755
-585
65
72
101
Protonic equil.
K~~
(M-I)
50
Dispropn. equil. 2 3 t,2 + 4
Equil. const
9.1 x lo-4
Diffusion coeff. Dan;,,
Kllh
2.8 x 103
2 1' t, 0'
4 x lo-4
Danion
K2'h
2.4 x lo6
1' 1' h t, 0'
lo6 cm2 s-I
2.0
33/2(mV)
AEp (mv)
2.5
+ 2'
+
0.35
DH+
19
+ 2'h
K2'h2
3.0 x 103
2 l'h t,0'
0.38
+ 2'h2
184
Illustvating basics of voltammetvy
(9) Dfusion coefficients of [P2W180 6 2 ] x ( x= 6-10) These were assumed to
be independent of x. A value of Danion= 2.0 x loh6cm2s-' was estimated for
[ P ~ w ~ ~
viaouse~ of~ hydrodynamic
] ~
voltammetry and the Levich equation
at a RDE (Section 9.1 in Chapter 2). Allowing variation of the single variable Danion,generated the simulated voltammogram in Fig. 3.2(a) with Danion
optimizing at 2.5 x 1o - ~cm2 s-' .
As noted above, the first and second processes remain unaltered as the pH of
the solution is lowered from the initial value of 5.2 (Fig. 3.2). In contrast, the
fourth wave appears to collapse and merge with the third wave (Fig. 3.2(b) and
(c)).At pH 3.0, and when the solution concentrations of H+ and [P2W18062]6are equal, only a single broad wave is detected, with an apparent peak-to-peak
separation, AE,, of 115 mV (Fig. 3.2(d)). As the pH value is further lowered,
this process gradually sharpens and, below pH 2.20, begins to move to more
positive potentials (Fig. 3.2(e)). By the time a pH-value of 1.0 (Fig. 3.2(f))
has been reached, AE, is 44 mV, and hence consistent with two unresolved
one-electron processes (28.5 mV is expected for a reversible couple when two
electrons are transferred simultaneously and reversibly at the same potential at
298 K).
In order to simulate this complex, pH-dependent, voltammetric behaviour,
the following stepwise approach is recornmended. Initially, a protonic equilibrium involving 4, presumably the most basic species (charge lo-), was included
into Scheme 3.1 to give Scheme 3.2.
Scheme 3.2
Subsequently, additional parameters associated with other added equilibrium
reactions need to be introduced as the proton concentration increases.
(10) The apparent dzjiusion coeficient of^+ In principle, the diffusion coefficients of four polyoxo anions need to be defined, as well as that of the proton.
In (9) it was decided to make the diffusion coefficient of all the polyoxo anions
equal to that measured for [ P ~ w ~ ~ o ~Further,
~ ] ~ - significant
.
ion-pair formation [16], that is likely to result from the combination of a high sodium cation
concentration and negatively charged polyoxo anions, was neglected. These
simplifications necessitate the use of an optimal apparent value of DH+ rather
than a measured value. In pure water at 25°C [17], DH+ = 9.3 x
cm2 s-* .
Its value in aqueous NaCl (0.5 M) is not known but is 7.5 x
cm2 s-* in
0.54 M NaCl at 25°C [18]. In contrast, a value of 1.0 x lop5cm2 s-' was found
Coupled electron- and proton-transfer reactions
185
give excellent agreement between experimental and simulated voltammograms in the pH range 1-6. This value was optimized in the intermediate pH
range, as the contribution from protonation was insignificant at higher values
and the electrochemistry was not sensitive to DH+at lower values. Clearly,
both approximations introduced above play a role in the apparent value of DH+
used to optimize the simulation. That is, DH+is subject to accumulation of
errors and hence may contain considerable uncertainty with respect to systematic error. Thus, while no great precision can be attached to the apparent value
H+, 1.0 x
cm2s-l, and indeed it is lower than that predicted from the
in literature [17,18], the important aspect is that this value is much larger
than that derived for the anions, a feature critical for the creation of a high pH
gradient close to the electrode in unbuffered media. This feature is essential in
order to achieve agreement between simulated and experimental data. Ideally,
input of independently measured diffusion coefficients and ion pairing equilibrium constants for all species and equilibria would improve the accuracy of
the simulation. Since such measurements could not be achieved, an optimal,
rent value of DH+was used over the entire pH range to minimize the
ut parameters. Errors in estimation of the area of the electrode
to the uncertainties in the diffusion coefficients.
mogeneous rate constants The rate constant for protonation, kqf (foreaction) of 4, was assumed to be extremely fast and arbitrarily set to
, which corresponds to an essentially diffusion-controlled value.
The back rate constant k4b was allowed to vary and therefore to determine the
of the equilibrium constant: K4h = k4f / k4b.
e parameters designated above, when used in the initial simulation, prod an acceptable fit to experimental voltammograms in the pH range 6.0-3.0
3.3; K4h = 5 X lo4M-l) but not at pH 1.0 where the simulated value
was too large, and the simulated peak potentials too negative, relative
to experiment (Fig. 3.4). Reducing K4h exacerbated these effects, and allowing
it to increase caused the wave to become too 'irreversible'. The latter effect is
traced to the constraint of assuming protonation to be
Thus, increasing Kqhreduces kqb. Consequently, less 4 s converted to 4 during
turn sweep leading to an unacceptably low cur
e inconsistencies between simulations based on Scheme 3.2, particularly at
be overcome by introducing an alternative route for the convervia 3h, the protonated form of 3 (Scheme 3.3). In the simulation
, the Eil2 value for the [ P ~ w ~ ~ o ~couple
~ ] ~in-Scheme
~ ~ ~ 3.2,
) remains fixed at the experimentally determined value, and K3h and Kqh
are allowed to vary. The Nernst equation ensures that these three quantities
define
) = E(3/4) - (2.303 RTInF) log (K3h/Kqh)
/4h) is the Ei12 value for the [ ~ ~ 2 ~ 1 8 0 6 2 ] ~couple.
-/~-
Illustrating basics of voltammetry
186
-0.50
-0.25
-0.00
Potential (V vs Ag/AgCl)
0.25
0.50
Fig. 3.3 Cyclic voltammograms obtained at a scan rate of 100 mV s-' for reduction of
~ oa GC
~ ~ macrodisc
] ~ electrode in aqueous 0.5 M NaC1 (pH 3); (-) experi1 ~ M [ P ~ w ~ at
ment; . .) simulation; employing Scheme 3.2 with K4h = 5 x lo4 M-' . Reproduced by courtesy:
Anal. Chem. 71 (1999) 3650. Copyright, American Chemical Society.
( 0
At pH 3.00, the experiment shows that the two waves have merged to a broad,
single wave (Fig. 3.2(d)).For this to occur, E(3h/4h) 5 E(2/3) and as, experimentally, E (213) - E(3/4) = 180 m v , eqn (3.3) requires that K3h/K4h <
This constraint allows a systematic search to be undertaken for the estimation
of K3h and K4h.
Values which provided acceptable simulation at individual pH values could
be defined but were not transferable (e.g. pH 5 3.0 : K3h = lo2M-l, K& =
I O ~ M - '(Fig. 3.2(a-d)); pH = 1.0 : K3h = IOM-', K4h = I O ~ M - ' ) .Consequently, the next step undertaken in order to seek good agreement between
simulated and experimental cyclic voltammograms was to include protonation
of 4h (Scheme 3.4). This seems desirable because:
(a) the fact that K4h2 depends upon [H+12 means that inclusion of this equilibrium reaction need only affect voltammetric behaviour at the lower pH
values where discrepancies presently exist with use of Scheme 3.3;
Coupled electron- and proton-transfer reactions
187
Potential (V vs Ag/AgCl)
Cyclic voltammograms obtained at a scan rate of 100mVs-' for reduction of 1 rnM
[ P ~ w ~ ~
at aoGC
~ macrodisc
~ ] ~ - electrode in aqueous 0.5 M NaCl (pH 1); (-1 experiment; (. . - .)
simulation; employing Scheme 3.2 with &, = 5 x lo4MM-'.Reproduced by courtesy: Anal. Chem.
71 (1999) 3650. Copyright, American Chemical Society.
Scheme 3.3
(b) anionic charge considerations and the expected increasing basicity, predicted
to occur as extent of reduction increases, suggest that if equilibrium 3 tz 3h
is important, then 4h tt 4h2 also should be included.
ystematic searching leads to a set of parameters (Table 3.1) which fitted the
cyclic voltammetric behaviour optimally in the pH range 6.0-1.0 (Fig. 3.2).
188
Illustrating basics of voltammetry
Scheme 3.4
E(3h/4h) is calculated to be -0.480 mV, some 10 mV more positive than
E(2/3), as required. Thus, a single set of three equilibrium constants can be
used to fit this complex voltammetric behaviour over a range of five orders of
magnitude in proton concentration.
Disproportionation reactions are known to be important in the chemistry
of polyoxometalates [5]. Relevant equilibrium constants related to disproportionation of reduced forms of the [ P ~ w ~ ~ osystem
~ ~ ] can
~ - be derived from
Table 3.1 (see Scheme 3.4 and eqns (3.11) and (3.12)):
As expected for a reaction scheme in which all electron-transfer steps and acidbase reactions are reversible, incorporation of these cross reactions had no impact
on the simulated cyclic voltammetric behaviour over the pH range 6.0-1.0.
3.1.3 Additional verification o f t h e acceptability of calculated parameters
The simulation parameter values established above have all been derived from
1mM solutions of polyoxometalate anion. A crucial test of their validity, is
their ability to be used to correctly predict voltammetric behaviour at different
polyoxometalate concentrations. Predictions for 5 mM solutions were surveyed
and the data suggested that significant differences in wave shapes were expected
at pH 3.45. A comparison of experimental behaviour and simulation at this
o ~ ~ ]in~ Fig.
3.5. Peak
pH for 1 and 5 mM solutions of ~ ~ ! - [ P ~ w ~ is~ shown
positions still match closely but current intensities reveal discrepancies. The
value
current magnitude may be fitted almost exactly by changing the D,,,
from 2.5 x
cm2 s-' (derived
cm2 s-l to the experimental value of 2.0 x
from rotated-disc voltammetry at 1mM concentration). A rationale for this
change is that an increase in anion concentration and a concomitant increase
Coupled electron- and proton-tranrfer reactions
I
-0.75
I
I
I
I
I
-0.5
-0.25
0
0.25
0.5
189
Potential (V vs Ag/AgCl)
. 3.5 Cyclic voltammograms for reduction of [P2 W18 0 ~ ~ (a)1 5~mM;
~ (b)
; 1mM. Other
conditions are as in Fig. 3.2. Reproduced by courtesy: Anal. Chern. 71 (1999) 3650. Copyright,
American Chemical Society.
ion pairing for these highly charged species is expected to decrease Danion.
owever, it can be noted that changing the electrode area slightly, assuming a
small amount of blockage has occurred, would also lead to a lowered current
agnitude and good fitof the data. The value of DH+was not altered in the
even though
simulation, when charging the concentration of a-[p2wI8
is species is expected to be a function of the ionic strength and environment.
A word of caution is needed to conclude this Section. With enough variables
available, it is of course possible to simulate the proverbial camel, so that a sysatic and chemically 'reasonable' strategy needs to be adopted when varying
parameters that may be input into a simulation. Thus, it is considered that
simulation of the [ P ~ w reduction
~ ~ sequence
~ ~ ~represents
] ~ the limit to
ich it is 'reasonable' to push the extent of credibility of a simulation of a comx problem if the data obtained from the simulation are to be quantitatively
lievable. Clearly, the more the number of available, independently measured
ameters, the better. Use of sensitivity analysis and smart statistical methods
calculation of the significance of each parameter [19] will assist in assessing
e reliability. Major weaknesses in the above simulation, as it presently stands,
es with uncertainties in the diffusion coefficients of D,;,,, and DH+and, to
lesser extent, with the electrode area. Setting the protonation reaction (e.g.
190
Illustrating basics of voltammetry
forward reaction k4f) at a value expected for the diffusion-controlled rate of
10'' M-' s-l is also somewhat arbitrary. Finally, it is never possible to guarantee that a unique solution to a mechanism has been achieved. In all senses, then,
ultimately, an improved simulation may emerge in the voltammetry than that
presented above, despite the relatively impressive results in Section 3.1.
3.2 Reduction ofa -[H2W I 2od016The initial definition of the reduction chemistry of ( . u - [ ~ 2 ~ 1 2 0 4 0by
] ~Pope
and Varga [20] was followed by examination of its extended redox chemistry which apparently involves addition of up to 32 electrons [20-241. At pH
above 4.0, the initial reduction processes O r / l ' and 1'/2' for this metatungstate
anion each involve transfer of one electron (Scheme 3.5) and as expected their
reversible potentials measured under these conditions are independent of p
Experimental E;,2 values in Scheme 3.5 are given in mV versus Ag/AgCl.
The metatungstate ion features two protons trapped in an internal cavity. Extra
protons present on reduced species are probably bound externally. However,
the two classes of protons are not differentiated in this reaction Scheme.
The voltammetric behaviour for reduction of 1mM solution of [H2Wl2O4()]
in the pH range 6.9-2.0 in aqueous 0.5 M NaCl is shown in Fig. 3.6. As the
pH value is decreased below 4.0, process l f / 2 ' diminishes in current intensity whereas the current intensity of O f / appears to increase. At pH 3.4,
a single wave (AEp = 87 mV at v = 100 mV s-l) is observed. This wave
gradually sharpens and begins to move to more positive potentials as the pH
is decreased below pH 3.2. At pH 1.1, A Ep = 55 mV, consistent with a
chemically reversible, apparently two-electron process constituted from two
non-resolved one-electron charge-transfer processes.
The approach to the simulation presented below follows that for the
[ P ~ w ~ system
~ o ~and
~ is] based
~ on Scheme 3.5. The simulations are presented
Scheme 3.5
Coupled electron- and pvoton-transfer reactions
-1.0
-0.8
-0.6
-0.4
-0.2
Potential (V vs Ag/AgCl)
0.0 -1.0
-0.8
-0.6
-0.4
-0.2
191
0.0
Potential (V vs Ag/AgCl)
. 3.6 Cyclic voltammograms obtained at a scan rate of 100mV s-' for reduction of 1 rnM
W12040]at a GC macrodisc electrode in aqueous 0.5 M NaC1; (-) experiment; (. . . ..) simulation. Reproduced by courtesy: Anal. Chem. 71 (1999) 3650. Copyright, American Chemical
Society.
in Fig. 3.6 and derived parameters listed in Table 3.1. Specific considerations
included:
(1) At v = 100 mV s-', processes 0'/1' and '/2' exhibit peak-to-peak separations, AE,, of 70 and 100 mV respectively, versus 57 mV expected for a
reversible one-electron couple at 298 K). Satisfactory simulation (Fig. 3.6(a))
of these processes was achieved by using values of 0.50 and 0.55 for a and
0.06 and 0.003 crn s-I for ko, respectively. These electron charge-transfer
192
Illustvating basics of voltammetry
steps may therefore be termed quasi-reversible, rather than reversible, as
was the case for the [ P ~ w ~ ~ osystem;
~ ~ ] ~ (2) The experimental value of D,,,
2.0 x low6cm2s-', was employed while
cm2s-' at intermediate pH values;
DH+was optimized at 1.9 x
(3) Relevant disproportionation reactions (eqns 3.13 and 3.14) were considered
(Table 3.1):
+ l'h --+0 + 2'h
1 h' + l'h + 0 + 2'
1'
K = 1.3 x
(3.13)
K = 0.35
(3.14)
Inclusion of disproportionation of lh' (eqn 3.14; kf = 100 M-' s-') significantly improved the fitting below pH 4;
(4) Predictions, based on parameters derived at 1mM anion concentration, of
behaviour for solutions of 5 m M concentration indicated that coalescence
of the one-electron processes O'/lf and 1\12' should occur at about one pH
unit lower at the higher [Hzw12
040]6- concentration. Simulations fitted
experiment well for pH values above 3 but increasingly poorly below pH 3.
Additional complexity could be introduced into the simulation to provide
an apparently better fit with experiment. However, reality may soon be left
behind and limitations in present levels of understanding of this system have
been reached.
3.3 Discussion of results obtainedfrom the simulation ofthe reduction of
w-[p2w ~ ~ oand~a!-[~
& ]w~ ~- ~ o ~ ~ ] ~ The comparisons of experiment and theory made in Figs 3.2 and 3.6 indicate
that the behaviours of both [ P ~ w ~ ~ oand
~ ~[ ]H~ -~ w ~ in
~ aqueous
~ ~ ~NaCl
] ~ (0.5 M) in the pH range 6.0-1.0 can be accounted for, in the main, by related
Schemes 3.4 and 3.5 respectively. Importantly, it is apparent that while the
observed voltammograms differ significantly in shape and apparent complexity
as the pH varies, this need not imply the onset of chemical irreversibility.
Definition of three protonation constants for each system, combined with
experimental values of E;,, for the two simple one-electron couples involved
(Table 3. I), provides a cdmplete thermodynamic description. The assumption
of fast protonation rates in all cases appears to be justified. In addition, the protocol available in the simulation [25] successfully handled the fact that the simple
one-electron couples O f / l ' (AE, = 72 mV; cx = 0.50; ko = 0.06 cm s-') and
lf/2' (AEp = 101 mV; cx = 0.55; ko = 0.003 cm s-') can be quasi-reversible
for reduction of [ H ~ w ~ ~ rather
~ ~ ~
than
] ~reversible,
- ,
as deduced for process
2/3 and 314
' in the [ P ~ w ~ ~ osystem.
~ ~ ] ~Minor
contributions from uncompensated resistance (200 Ohm) and capacitance (2 x
Farad ~ m - ~which
),
also give rise to non-ideality, are also accommodated by the simulation.
The assumption that the initial forms of the anions ( [ P ~ w ~ ~ o (2),
~~]~Scheme 3.4; [ H ~ w ~ (Of),
~ ~Scheme
~ ~ 3.5)
] ~ are
- not protonated is justified by
Coupled electron- and proton- tranrfer reactions
193
the results and in the case of 0'by the fact that fully oxidized polyoxometalates
are characteristically strong acids (pK,, 0-2) [26,27]. Schemes 3.4 and 3.5 are
also consistent with the observed two-electron-two-proton [20,26] (eqns 3.15
3.
and 3.16) Nernstian dependence of E;,, for the coalesced waves at pH i
A significant weakness may be the need to assume that the diffusion coefficient Danionis independent of the redox and protonation state of the anion
and that Dani,, and DH+are independent of pH in aqueous 0.5 M NaCl. Thus,
optimization of these parameters within chemically sensible ranges was necessary to simulate successfully the current magnitudes at anion concentrations of
1 and 5 mM (Fig. 3.5).
The detailed reasons for the significant differences in wave shapes as a function
of pH and concentration of polyoxometalate can be discerned from examination of plots of species concentration versus distance from the electrode [25].
Figure 3.7 presents simulated surface concentration data for reduction of 1 and
~ ] ~ - of -0.74 V, which corre5 mM solutions of ( ~ - [ P ~ w ~ ~ato a~ potential
sponds to a value just beyond the peak potential for the most negative process
considered in this study. For the 5 mM concentration case (Fig. 3.7(a)), this
potential corresponds to the reduction of 3 to 4 (also compare the morphology
of the inset in Fig. 3.7(a) with that of Fig. 3.2(a) for 1 mM concentration at pH
5.20). Note that at the 5 mM concentration, conversion of 4 to 4h to achieve
their equilibrium values has expended virtually all the available protons, which
leaves only a very low residual concentration of H+ at the electrode surface
(Fig. 3.7(a)). Hence, non-protonated 4 is the dominant species at the electrode
surface. In contrast, for the 1 mM bulk concentration (Fig. 3.7(b)), the surface concentration of H+ no longer falls to anywhere near zero, even though
[H+]/[anion]
1 : 3. This is because the diffusion coefficient DH+ 4D,,,,,
in the simulation, so that H+ diffuses to the surface faster than the anion, thereby
enabling a significant surface proton concentration to be maintained during the
entire forward sweep of the voltammogram. Consequently, under these conditions 4h, rather than 4, is now the dominant polyoxometalate species present
at the electrode surface and the change in voltammetric shape relative to that
when the bulk anion concentration is 5 mM reflects this feature. It is apparent that the differences in wave shapes as a function of anion concentration are
with any effect of ion pairing
driven by the ratios [H+]/[anion] and DH+/ Damon,
being neglected in this discussion. Thus, the fact that DH+is always significantly
larger than any other species present, is a significant feature in voltammetry of
electrode processes involving acid-base equilibria coupled to electron-transfer
processes. In many simulations it can be assumed that the diffusion coefficient of
all species participating in the electrode process are equal, but not when protons
participate in the redox process.
194
Illustrating basics of voltammetry
100
Distance (prn)
200
Fig. 3.7 Simulated speciation as a function of distance from the electrode surface for reduction of
[ P ~ w ~ ~ o(a)~ 5~mM;
] ~ (b)
- 1mM in aqueous 0.5 M NaCl at pH 3.45. Insets show cyclic voltammograms obtained at a scan rate of 100 mV s-l terminated at a potential of -0.74 V versus Ag/AgC1,
the sampling point for the calculation of the concentrations of each species considered in the simulation. Reproduced by courtesy: Anal. Chem. 71 (1999) 3650. Copyright, American Chemical
Society.
[p2w18
0 ~ ~ 1and [ HwI2
~ 0 4 ~ 1 formally carry the same negative charge.
However, the EiI2 values of the first two simple one-electron couples for the
latter species are more than 600 mV more negative (Schemes 3.1 and 3.5), a
situation probably related to the higher surface charge density on the smaller
6- o ~ ~ ]
[ H ~ w ~ ion.
~ ~In ~fact,
~ the
] ~
first- two reduction processes for [ P ~ w ~ ~
are insensitive to pH in the range 6.0-1.0 (Fig. 3.2), in stark contrast to
Electrochemical reduction ofa-[S2Mo180621~-
195
the equivalent situation in the [ H ~ w system
~ ~ (Fig.
~ 3.6).
~ ~ In ] addi~
tion, the derived protonation constants for [H2w ~ ~ o ~and
~ [] ~H - ~ w 8- ~ ~ ~ ~
(Scheme 3.5; Table 3.1) are two orders of magnitude larger than those for
180621~- and [ P ~ w ~ lo-,
~ o despite
~ ~ I the lower anionic charges. However, the pattern of relative values of the constants within a given system
(eqns 3.17 and 3.18), does reflect expectations based on the anionic charge:
K3h
K4h2 and
K3h/K4h
(3.17)
n summary, a systematic approach to the simulation of the cyclic voltammetry for reduction of the a-[p2w ~ ~ o ~and
~ a] ~- - [ ~ ~ ~ polyoxomet~ ~ 0 ~ ~ ] ~
alates in aqueous solution, based upon an established program [25], provides
a reasonably detailed understanding of the coupled electron-proton transfer
chemistry. Further refinement of the model to incorporate ion-pair formation
and phenomena related to media effects, and their influence on diffusion coefficients is probably required to afford a complete description. However, without
access to independent experimentally derived measurements of some of these
parameters, and their dependence on the medium, it is not clear that unique
ulations will be obtained if the complexity of the simulation is enhanced for
sole purpose of providing greater agreement with experiment. The reader
is therefore again cautioned against over-interpretation of the quantitative significance of a simulation involving the input of more than a few unknown
parameters. Thus, as noted in Chapter 2, while advances in the use of more
isticated forms of statistics can be expected to improve the level of confie in the results of simulations in the future, every effort should always be
e to evaluate any input parameters that can be independently measured, so
he number of parameters that have to be deduced from the simulation are
ept to the absolute minimum.
he successful simulation of experimental cyclic voltammograms obtained for
o ~a ~ -] ~[ - ~ ~ ~ has
~ helped
~ 0 to~quantify
~ ] sev~ reduction of ( ~ - [ P ~ w ~ ~and
eral aspects of the proton-transfer reactions that can accompany the electrontransfer processes of polyoxometalate compounds in aqueous media. In essence,
as reduction proceeds, the anions acquire an increasingly negative charge, which
increases the basicity of the reduced form, and encourages protonation. In turn,
protonation lowers the negative charge, encouraging further reduction, so that
very high levels of polyoxometalate reduction can occur under acidic conditions. However, cyclic voltammetry represents only the application of one
3 ~ a r tof
s Section 4 have been adapted with permission from Inoy. Chem. 32 (1993) 2416;
36 (1997) 2826; and 36 (1997) 4227. Copyright, American Chemical Society. Additional
details may be obtained by consulting these references.
196
Illustrating basics of voltammetry
technique of electrochemistry and, in its own right, does not lead to unequivocal identification of structures or properties of highly reduced species. In the case
studies have been undertaken in both protic and aprotic
of
media, in which many of the voltammetric, spectroelectrochemical, and electrochemical synthetic techniques described in Chapter 2 have been employed.
In this section, a review of data available on the reduction of a-[S2Mo18062]4is provided to highlight the advantages and limitations of the various electrochemical techniques and, also, to further illustrate how data obtained'from the
use of a wide range of techniques may be used in an interactive manner to
understand inherently complex redox chemistry.
~ ~ (Fig.
o ~3.8)
~ has
] ~approximately D3hpoint
The idealized a- [ s ~ M ~ anion
symmetry [28-321. Two hexagonal belts of, alternately, point- and edge-sharing
M o o 6 octahedra are each capped by an edge-shared Mo3Ol3 unit. The belts
are linked by six bridging 0x0 ligands located on the horizontal plane of symmetry. Since there are different Mo(V1) environments in this polyoxometalate
species, the location of electrons in reduced states is of considerable interest. In
principle, ~ - [ s ~ M ~ ~might
~ o be
~ ~expected
] ~ - to accept a total of at least 18
electrons, formally reducing each Mo(V1) (do electronic configuration) centre
to the Mo(V) (dl electronic configuration) state. Structural considerations noted
above, imply that reduction would occur over a range of potentials and not as a
single 18-electron process. Even more extensive reduction to generate, formally,
Mo(IV) (d2 electronic configuration) centres could be expected on the basis that
to
the meta-tungstate ion, a - [ ~ ~ ~ considered
~ ~ 0 in~ Section
~ ] 3~is reported
-
Fig. 3 -8 Polyhedral form of structural representation of ~ - [ s ~ M0o6I2 ~1 ~ ~ .
Electrochemical reduction Ofa -[S2Mo18o62I4-
197
accept at least 24 electrons, together with charge-compensating protons [20-241,
which formally, therefore, corresponds to reduction of all 12 W(V1) sites to the
~ ~ more
] ~ - , elec(IV) oxidation state. Apparently, with a- [ H ~ M O ~ ~ Oeven
trons can be accepted into the polyoxometalate framework [24,32], so that very
extensive addition of at least 36 electrons (do to d2 for each of the 18 Mo
atoms), and probably more, is predicted to be available in the case of reduction
of ~ ~ - [ s ~ M o ~
Electrochemical
~ o ~ ~ ] ~ .studies in protic and aprotic media presented below reveal the details of the extensive redox chemistry associated with
- [ ~ ~ ~ o ~ ~ 0 ~ ~ l ~ - .
4.1
Voltammetry of a - [ S ~ M Oob214in aprotic acetonitvile media4
deally, the fundamental nature of the electron-transfer properties of the
[ s ~ M ~ ~ ~system,
o ~ ~free
] ~from
- complicating, coupled protonation reactions,
should be revealed by voltammetric studies in aprotic solvents. The problem
with this concept is that traces of adventitious water and, hence, protons invariably present in aprotic media can cause complications, as also can interactions
with the high concentration of electrolyte anions (ion-pair formation). All these
facets of polyoxometalate redox chemistry studies are revealed in the voltammetry of [ s ~ M ~ in
~ nominally
~ o ~ ~aprotic
] ~ CH3CN. Figure 3.9 shows the
reduction of [ ~ ~ ~ ounder
~ ~conditions
0 ~ ~
of cyclic
] ~ voltammetry
at a 1-mm
diameter GC macrodisc electrode in 'dry' CH3CN (0.2 M Bu4NC104)over
the potential range of about +1.0-3.OV versus Fc/Fc+ at scan rates of 0.05,
0.5, and 10 V s-l (ion pairing ofpolyoxometalate anions with the Bu4N+ cation
appears to be small) [16]. Clearly, an extensive and complex reductive chemistry
is available in this solvent. An initial goal is to identify the primary electrontransfer processes and establish their reversible potentials; this may be achieved
in CH3CN (0.2M Bu4NC104)by a combination of fast scan rate and lowtemperature conditions which enable protonation reactions from adventitious
water or other impurities to be minimized.
4.1.1 Dependence of cyclic voltammetry on scan rate,
switching potential, and temperature
Initially, voltammograms, obtained only in the potential region where the
first four charge-transfer processes occur, are considered. These processes are
summarized in eqns (3.19)-(3.22) and designated I, 11, I11 and IV respectively.
'AS in Section 3.1 the a isomeric notation is now commonly omitted for convenience and
also because the isomeric forms of the reduced species are again usually unknown.
198
Illustrating basics of voltammetry
1
1
1
1
1
1
1
-3.50 -3.00 -2.50 -2.00 -1.50 -1.00 -0.50
1
1
1
0 +0.50 +1.00
Potential (V)
Fig. 3.9 Cyclic voltammograms for reduction of 2.0 rnM [ s ~ M o I ~ o ~ in
~ ] ~CH3CN
(0.2 M Bu4NC104)at a 1.O-mm diameter GC macrodisc electrode at 22OC; potentials versus FC/FC+;
scan rate (V s-I): (a) 0.05; (b) 0.50; (c) 10. Reproduced by courtesy: Inog. Chem. 36 (1997) 2826.
Copyright, American Chemical Society.
Experiments as a function of scan rate at GC electrodes for these initial four
processes involved switching the potential between the various processes. Scanning the potential regions of only process I or just processes I and I1 produced
almost ideal reversible one-electron reduction responses (after allowance for
Ohmic drop) for scan-rates in the range 0.02-1 0 V s-' (Fig. 3.10).
Thus, the voltammetric characteristics were initially indistinguishable from
the known reversible one-electron ferrocene (FC/FC+)oxidation process under
the same conditions. However, when the potential is switched after processes
I11 or IV, small deviations from ideality emerge in the form of a dependence
Electrochemical reduction ofa -[S2Mo180621~-
I
I
I
I
I
-1.00
-0.50
0
0.5
1O
.O
199
Potential (V)
ig. 3.10 Cyclic voltammograms for reduction of 2.0 rnM [ s ~ M ~ ~ ~when
o ~ the
~ ]potential
~ is
switched after process 11; other conditions as for Fig. 3.9. Reproduced by courtesy: Inorg. Chem. 36
(1997) 2826. Copyright, American Chemical Society.
of
on scan rate, and the appearance of minor processes (Fig. 3.11). In
particular, a new chemically reversible process is present at potentials just after
process IV and new processes are evident between I1 and 111. However, at a
fast scan rate of 10 V s-' , four, almost ideal, reversible, one-electron reduction
processes are observed (Fig. 3.1 1). Deliberate addition of water amplifies the
relative magnitude of the minor processes, implying that the three- and fourelectron reduced anions [ ~ ~ ~ o ~ ~ the
0 ~presumed
~ ] ~ -products
. ~ - of
, processes
11 and IV, are extremely basic and readily protonated by adventitious water
present in acetonitrile.
200
Illustrating basics of voltammetry
I
I
I
I
I
-2.00 -1.05 -1.00 -1.50
0
Potential (V)
I
I
0.50 1.00
Fig. 3.11 Cyclic voltammograms for reduction of 2.0 mM a - [ ~ ~ ~ when
o ~the~potential
0 ~ ~
is switched after process IV; other conditions as for Fig. 3.9. Reproduced by courtesy: Inorg. Chern.
36 (1997) 2826. Copyright, American Chemical Society.
Processes observed at more negative potentials than reduction waves I to
can now be considered. Clearly the complexity increases as more highly redu
species are generated (Fig. 3.9). The effect of scan rate variation on process
is illustrated in Fig. 3.12. At slow scan rates, V has a peak current considerably
larger than that expected for a one-electron reduction process (Fig. 3.12(a)).
the scan rate increases, the peak height of V approaches that of processes I(Fig. 3.12(b) and (c)).This feature is consistent with protonation accompanying
processes V and VI, thereby causing them to coalesce into a single two-electron
reduction wave (see analogous reactions in Section 3). However, at fast scan
rates, insufficient time is available for complete protonation to take place, so
that processes V and VI are resolved.
]
~
Electrochemical reduction of~x-[s~n/ilo~~
o62I4-
I
I
I
I
-2.00 -1.05 -1.00 -1.50
I
0
Potential (V)
I
I
0.50
1.00
. 3.12 Cyclic voltammograms for reduction of 2.0 rnM [ ~ ~ ~when
o the~ potential
~
0is
switched after process V; other conditions as for Fig. 3.9. Reproduced by courtesy: Inoq. Chem. 36
(1 997) 2826. Copyright, American Chemical Society.
Even more complex voltammograms are seen upon switching the potential
after process VI (not shown). Like V, process VI is only detected as a oneelectron reduction process at very fast scan rates when, presumably, the influence
of following chemical reactions is minimized. Very dry acetonitrile is required
to achieve even this fast scan rate result, and deliberate addition of water enormously increases the complexity of the cyclic voltammetry at negative potentials.
While processes V and VI observed in 'dry' acetonitrile (0.2 M Bu4NC104)
are not completely reversible in the chemical sense even at very fast scan rates,
simulations show that under this circumstance the reversible reduction potentials
~
~
]
202
Illustuating basics of voltammetry
Table 3.2 Cyclic voltammetric dataa for
reduction of 2.0 rnM [ s ~ M ~ ~ in~ o ~ ~ ] ~
CH3CN(0.2M Bu4NC104)
I
I1
I11
IV
v
v
VII
VIII
v = 10 V s-' ; 1 rnrn diameter GC macrodisc working electrode.
v v~ FC/FC+.
CAE1I2= difference in Ell2 between consecutive electrode
processes.
a
+
measured as (Erd E y ) / 2 will not be significantly shifted from their true values.
Thus, it can be concluded that E;,, values for I and 11 differ by -0.24V and
that similar differences also apply for the pairs 111, IV and V, VI. Based on a
separation of about -0.6 V between individual pairs (Table 3.2), the first process
for the next reversible pair (process V ) is predicted to occur at about -2.50 V
versus FC/FC+.However, switching
potential just prior to the solvent limit
of -3.0 V versus FC/FC+ in acetonitrile at 20°C shows very complex behaviour
(Fig. 3.9) and, hence, this experiment provides no direct evidence for the
anticipated additional pair of one-electron charge-transfer processes. At the
lowered temperature of -30°C, on the other hand, direct
found for the predicted seventh and eighth processes, VII an
close to the predicted potentials. Lowering the temperature t
o ~water,
~ ~ ~ ~ ]
to slow the rate of reaction of reduced forms of o l - [ ~ ~ ~ with
and possibly also lower the activity of water.
Data presented so far imply that in the absence of significant concentrations
s of protons, or ion pairing with the electrolyte, processes
as detected at a GC electrode, also correspond to simple
extension of the electron-transfer series (see eqns (3.19)-(3.26)).
Electrochemical reduction ofa -[S2 Mo180 ~ ~ 1 ~203
-
~ h u s eight
,
one-electron reversible charge-transfer processes appear to exist for
reduction of [ S 2 ~ 0 1 8 0 6 2 in
] ~aprotic media at potentials up to -3.0 V versus
c+, and these are summarized in the overall eqn (3.27).
resumably, even more one-electron processes would be detected if a more
negative potential range could be accessed in a highly pure (water-free) aprotic
solvent medium.
4.1.2
Steady-state voltammetry
e the transient cyclic voltammetric data provide some of the basic
ation, complementary studies using steady-state conditions are always
recommended to both confirm findings made by the transient method as well
as to provide additional details. For the reduction of [ ~ ~ ~ o it turned
~ ~ 0
out that the need to use slow scan rates to obtain the steady-state condition
t electrode 'fouling' in the negative potential region where
I are expected. Nevertheless, useful additional data associated
V can be obtained.
i
I
-1.50
-1.00
-0.50
0
Potential (V)
0.50
f
I
1
I
I
I
1.00 -3.00 -2.50 -2.00 -1.50 -1.00-0.50
1
0
I
I
0.50 1.00
Potential (V)
3.13 Steady-state voltammograms obtained at a scan rate of 0 . 0 0 5 V ~ - for
~ reduction of
2.0rnM [ s ~ M ~ ~ ~ino CH3CN
~ ~ ] ~(0.1
- M Bu4NC104) at 20°C; (a) 50pm diameter platinum
microdisc electrode; (b) 2.8 mm diameter GC RDE; rotation rate of 500 rpm. Reproduced by
courtesy: Inorg. Chern. 36 (1997) 2826. Copyright, American Chemical Society.
~
~
]
204
Illustrating basics of voltammetry
Microdisc Electrode Figure 3.13(a) shows a near steady-state voltammogram
obtained at a platinum microdisc electrode (50 pm diameter) over the potential
range encompassing waves -1V. While Ell2 values obtained
and I1 are the same as thos found by cyclic voltammetry (T
technique-dependent variations in Ell2 are evident for processe
vary stea
addition, the magnitude of the limiting currents (Ilim)
> IV (Fig. 3,.13)implying that the steady-state voltammetry does
not in fact solely consist of simple mass-transport-controlled, reversible processes
under these near steadyitions. The current contribution from minor
3.13) accounts for the majority of the
processes seen between
I. However, the magnitude of Ilim
for
'missing' current for process
process IV is too low to be consistent with simple mass-transport-controlled
conditions.
For electrochemically reversible processes, plots of E versus ln[(I& - I)/I]
or the so-called 'log' plots should be linear with intercept E;,2 and slope RT/nF
(Section 9.4 in Chapter 2) under steady-state conditions at a microdisc elecments with the slope of the 'log
trode. Processes I and
exhibit slight variations (apparent
plot' giving a value of n
n values calculated from the slopes of 'log' plots are 0.83 and 0.87, respectively). The fact that the limiting current of IV is considerably smaller than
that expected (ideally the limiting current values for all processes should be
identical) may be due to electrode blockage. Complete 'electrode blockage',
perhaps due to precipitation or adsorption of [ ~ ~ ~ o (or~ protonated
~ 0 ~ ~
forms) or double-layer effects [33] (some repulsion of anions may occur at the
negatively charged electrode), takes place at the plati
suriace soon after scanning the potential past process
scan rate employed in steady-state rnicrodisc electrode experiments combined
with the small surface area always enhances the risk of electrode fouling or
blockage relative to that encountered with fast scan rates and transient conditions at macrodisc electrodes. The very short time-scale of the steady-state
microdisc electrode and the very large negative charge and negative potential enhances the probability of a double-layer electrode kinetic effect [33].
Clearly, additional studies are required to identify the exact cause of this
problem.
Hydrodynamic Rotating-Disc Electrode A second series of steady-state voltammograms was recorded under hydrodynamic conditions using a GC rotatingmacrodisc electrode (2.8 mm diameter). It may be anticipated that the use of a
macrosized electrode and rotation will minimize electrode fouling or blockage
compared to that noted above under microdisc steady-state conditions. The
rotated-disc voltammogram in Fig. 3.13(b) shows that, in addition to
~ ~ omo~ ~ ] ~
process V is now also well defined and for the [ s ~ M ~ system,
ily interpreted voltammograms are obtained an with the platinum microdisc
and use of a rotated GC electechnique. Scanning the potential beyond
trode enables further processes to be observed which, on the basis of cyclic
]
Electrochemical reduction of a -IS2Mols o62I4-
205
voltammetric data, are attributable to multi-electron reduction of the mixture
of potonated and non-protonated, or adsorbed, or precipitated species now
present at, or on, the working electrode surface.
The first four processes obtained at the rotated electrode exhibit El12 values
are essentially independent of the rotation rate. Analysis of linear 'log
ersus ln[(Ii,,, - I)/I] for each of the first three processes at a rotation
00 rpm provide estimates of n-values of 0.99, 1.01, and 1.02, respecection 9.4 in Chapter 2). The apparent n-values obtained from the
' for the fourth process is 0.91, suggesting that a minor departure from
emins for processes at negati e potentials, even when the electrode is
he ideal behaviour of I and allow their total limiting current to be
culate the diffusion coeffici D of [ S 2 ~ o 1 8 0 6 2 ]in
4 -CH,CN using
equation (Chapter 2, Table 2.3). The value of 6.4 x 10-6cm2s-' at
e compared with values of 6.2 lop6cm2 s-' estimated from microelectrode voltammetry and use of process and the equation Ihm
= 4nFDr,[Alo
ection 10.1 in Chapter 2).
The injuence of electrolyte and solvent
electron-transfer steps observed in the reduction of [ ~ ~ ~ o ~ ~ 0 ~
l o 4 is used as electrolyte [34,35], mimic those found on addition of
d may be accounted for by assuming that ~ i acts
+ as a moderately
is acid [36] (or forms strong ion pairs [16]). For example, in (95/5)
/ H 2 0 (0.1 M NBu4C104), the voltammetric behaviour obtained on
of ~ i can
+ be simulated [34] according to Scheme 3.6:
Scheme 3.6
206
Illustrating basics of voltammetry
Fig. 3.14 Solvent dependence in CH3CN and DMF of cyclic voltammograms for reduction
of 1 mM [ s ~ M ~ ~ ~ato a~G~C ]electrode
~ (area 0.071 cm2) using a scan rate of 1 0 0 r n V ~ - ~ .
Reproduced by courtesy: Chern. Soc. Dalton Tmns. (2001) 1076. Copyright, Royal Society of
Chemistry.
The reversible potentials E1-E7 (V versus FC+/FC) in this mixed solvent
media are estimated to be 0.12, -0.11, -0.73, -0.91, -0.56, -0.67, and
-0.60, respectively, and the equilibrium constants Kl - Kg (M-l) to be lo1,
5 x lo2, lo6, lo3, and 8 x lo3, respectively.
In view of these results in Scheme 3.6, the role of ~ i ' , frequently contained
in the electrolyte, cannot be assumed to be thermodynamically innocent in
studies of polyoxometalates [21], nor can the solvent (e.g. CH3CN and
in Scheme 3.6), nor, presumably, even the tetraalkylammonium electrolyte
cations, as ion pairs are always likely to be significant when highly negative
polyoxometalate species are generated at electrode surfaces. Figure 3.14 shows
a comparison of the first four reductions steps of [ s ~ M ~ in
~ acetonitrile
~ Q ~ ~ ] ~
and dimethylformamide [37]. Clearly, a very large solvent effect is present in
the electrochemistry of polyoxometalates [38], so that interpretation of the
thermodynamics at a fundamental level can be very complicated.
4.1.4 Conclusions derivedjom voltammetric studies in acetonitrile
The data obtained from voltammetric studies indicate that, after avoidance of
any influence from the presence of trace water or acid, the choice of a suitable electrolyte with minimal ion pairing (e.g. Bu4NC1Q4), and the use of
transient conditions at a GC macrodisc electrode at low temperature, eight
reversible one-electron reductions of [ s ~ M o ~ ~ Q occurring
~ ~ ] ~ - , as four adjacent pairs, are accessible in CH3CN. The difference in potential within each
pair is remarkably constant, 0.26 z t 0.02 V, as are the separations between
adjacent pairs, 0.60 f 0.06 V (Table 3.2). The (Ei12
EfO) thermodynamic
208
Illustrating basics ofvoltammetry
shorter voltammetric time-scale. The observation of well-defined isosbestic
points (Fig. 3.15) in electronic spectra generated during the course of these
low-temperature reductive electrolysis experiments supports the notion that
no decomposition occurs at -45"C, as does the fact that oxidation experiments
after reduction, led to the quantitative reappearance ofthe initial [S2M018062]4electronic spectrum.
It can be seen from data contained in Table 3.3 that sequential oneelectron reduction to form [ s 2 ~ o 1
062]5-,
8
[ s ~ M 0o6 2~1 6~- , [ S 2 ~ 0 1 8 ~ 6 2 1 7 and [ ~ ~ ~ o generates
~ ~ 0a band
~ ~ in ]the~ red
- region of the visible spectrum that is absent in [ ~ ~ ~ o ~Additionally,
~ 0 ~ ~ it ]can
~ be
- noted
.
that this
band shifts to higher wave numbers on sequential reduction and the extinction coefficient increases in magnitude. [ ~ ~ ~ o 1has
~ each
0 ~molybdenum
~ ] ~ atom in a do configuration and the only transitions that occur are in the UV or
close to this region (c.30 800 cm-l). However, on sequential reduction, electrons occupy the formerly empty d orbitals of the molybdenum atoms in the
cluster, permitting d-d transitions to occur in the visible region of the spectrum [5]. The relatively high values for the measured extinction coefficients
suggest that the transitions occurring are spin-allowed intra-valence transitions
between separate molybdenum nuclei [5,36].
4.2.2
Electron paramagnetic resonance spectra of veduced species
Reduced forms of [ s ~ M o ~ ~ omay
~ ~be
] ~paramagnetic
or diamagnetic. If they
are paramagnetic, then EPR spectra may be observed, although there are, of
course, circumstances when paramagnetic species do not give rise to EPR signals. Thus, lack of detection of an EPR signal cannot be taken as evidence of
diamagnetism. As expected, no EPR signal was observed for [ ~ ~ ~ o ~
which is known to be diamagnetic on the basis of NMR and magnetic measurements. The EPR spectra (if any) of species, formed after well-defined and
known levels of reduction, may be obtained by measurements on reduced solutions obtained by coulometric titration using large-scale bulk electrolysis cells,
of the kind described in Section 15.2.1 in Chapter 2, and controlled potential
conditions, at appropriate potentials for each level of reduction. The respective
values with applied potentials (versus ~ c / F c + were
)
$0.12 V ( 0.02 V),
-0.12V (-0.28V), -0.78V (-0.88V) and -1.05 V (- l.lOV), for n = 1-4,
respectively. The extent of bulk electrochemical reduction at a platinum gauze
basket electrode was monitored by measuring both the charge passed (Faraday9s
Law) and the position of zero current in RDE voltammograms. That is, voltammetric confirmation that sequential one-electron reductions had occurred, as
predicted theoretically on the basis of the coulometric calculations (Section 15
in Chapter 2) was obtained by measurement of the sign and magnitude of the
limiting current for the four voltammetric one-electron waves in the RDE
steady-state voltammogram. Application of both of these methods for assigning the level of reduction, therefore, provided conclusive evidence that the
species formed had the appropriate number of electrons added and therefore that
+
~
210
Illustrating basics of voltammetry
Table 3.3 Summary of absorption maxima," in the visible region
(5000-25 000 cm-') for a variety of reduced species of [ s ~ M o ~ ~ o ~ ~ I ~ Reduced species
Medium
T ("C)
Absorption
a [ ~ 2 ~ ~ 1 8 ~has
6 2no
] 4absorption
maximum in the visible region (5000-25 000 cm-l).
b ~ ereferences
e
[5] and [36]for probable assignment of relevant intra-valence transitions that give
rise to the absorption bands.
'In CH3CN (0.1 M Bu4NC104).
d ~ a t obtained
a
by OTTLE experiments at a platinum gauze electrode at -45OC. Well-defined
isosbestic points obtained during the course of electrolysis to each redox level confirm the absence
of decomposition as do oxidation experiments after reduction, which quantitatively regenerate the
initial electronic spectrum of [ s ~ M ~ ~ ~ o ~ ~ ] ~ - .
eIn (9515)CH3CN/H20 (0.2 M HC104).
*Data obtained at 22°C by sequential (2eP, 2H') reduction to form [ H ~ S ~ M O ~ ~
and
O~~]~[ H ~ s ~ M ~ ~demonstrate
~ o ~ ~ ] that
~ - large shifis are observed in absorption band maxima when
protonation occurs.
a unique reduced polyoxometalate species was formed whose EPR properties
could be selectively probed.
Figure 3.16 shows EPR spectra recorded on frozen acetonitrile glasses (77 K)
formed after addition of a specified number of electron equivalents (n) to a
solution of 4 x
M [ s ~ M ~ ~
in acetonitrile
~ o ~ ~ (0.1
] ~M Bu4NC104)at
22°C. At the n = 1 stage of this coulometric titration, the intense, anisotropic
EPR spectrum of the frozen solution (Fig. 3.16(a)) was identical to that obtained
from an authentic sample of synthesized [Hex4NI5[S2M018062]dissolved in
acetonitrile (0.1 M Bu4NC104)solution and then frozen to form a glass under
the same conditions. The EPR signal associated with the n = 1 reduced species
decayed progressively upon further reduction to the n = 2 level (Fig. 3.16(b)),
indicating that the two-electron reduced anion, [ ~ ~ ~ o ~is EPR
~ 0 silent
~ ~ ] ~ or close to EPR silent at 77K. Diamagnetism would need to be established
by magnetic measurements. Upon further reduction to the n = 3 stage and
freezing of the solution, a new EPR signal was detected which was assigned to
the paramagnetic three-electron reduced species [ s ~ M ~ (Fig.
~ ~ 3.16(c)).
o ~ ~ ] ~
Electrochemical reduction of a -[S2 Adol80 ~ ~ 1 ~21
-1
DPPH internal
reference
. 3.16 Ex situ frozen glass EPR spectra (77 K) detected at different stages of a coulometric
titration of [ s ~ M ~ ~ ~in oacetonitrile
~ ~ ] ~ (0.1
- M Bu4NC104).Provided by courtesy: T. Vu, Monash
University, Victoria, Australia.
spectrum is consistent with the presence of a single unpaired electron
ion of [ s ~ M ~ to
~ the
~ no =~4 level,
~ ] led
~ to a progressive decrease
in the magnitude of the EPR signal assigned to [ s ~ M ~ ~ ~ oHowever,
~ ~ ] ~ - .
ty of the species present after the bulk four-electron reduction is
as the voltammograms of these solutions are more complicated than
edicted, if the only species present was the four-electron reduced species,
M O ~ ~ O ~In~all
] ~probability,
-.
protonation has occurred to give a mixture
ted and non-protonated 4e- reduced products. Thus, coulometric
titration experiments establish that frozen acetonitrile glasses of the pure oneand three-electron reduced species give intense EPR signals, whereas the pure
two- and four-electron reduced forms are EPR silent or EPR insensitive at
77 K. As for the n = 1 reduced species, no hyperfine coupling to the95797~o
i in the framework is observed for the n = 3 reduced species in glass phase
experiments at 77K. Apparently, the odd electrons are not localized at
a single site at this temperature on the EPR time-scale, but rather they are
elocalized over all or part of the polyoxometalate structure.
.3 Electvochemical synthesis of one- and two-electron
reduced fovms o f [ & ~ o ~ ~ ~ ~ ~ ] ~ The voltammetric and spectroelectrochemical data imply that large-scale electrochemical synthesis of at least solid [ s ~ M ~ and
~ ~[ so ~~M~~ ]~ ~~
should
o ~ ~ ] ~ -
212
Illustrating basics of voltammetry
be possible via electrochemical reduction in aprotic media using large GC or
platinum basket bulk electrolysis cells described in Section 15.2 in Chapter 2.
However, results also imply that prevention of protonation reactions is likely
to be difficult under the long time-scale conditions associated with largescale electrochemical synthesis. Thus, attention in this section is focussed on
the electrochemical synthesis and isolation of the non-protonated one- and
two-electron reduced species.
4.3.1 Synthesis of [S2Mo 8 0621560 mL of acetonitrile solution containing 2.4 rnM [ B u ~ N ] ~ [ S ~ M
andO ~ ~ ~
0.24 M Bu4NC104 was reduced at 400 mV versus Ag/AgCl. After an hour,
the solution had changed colour from transparent yellow to a deep transparent
green. Upon completion of electrolysis, the green solution was placed in a
refrigerator at 4°C. After several weeks, green crystals of [Bu4N15[S2M~18062]
were found to be present.
Synthesis of the green [ B U ~ N ] ~ [ S ~ M salt
O ~ via
~ Oreduction
~~]
of 2 rnM
[Bu4N14[S2M018062]
in dichloromethane (0.1 M Bu4NC104)is even more efficient as analytically pure product precipitates progressively during the course of
electrolysis.
4.3.2 Synthesis of[Pr4N]6[SzMo
and [ B U ~ N ] ~ [ H S Z18M0 O
s2]
A 2.4 rnM solution of [ H ~ x ~ N ] ~ [ Sin ~60M
mL~ CH3CN
~ ~ ~ ~ (0.24
~ ]
Bu4NC104) was reduced at -0.45 V versus FC/FC+. Pr4NBr (0.25 g) was
added to the deep-blue solution which was stored at -5°C. After one week,
analytically pure blue crystals -of [Pr4N]6[S2M018062]
were isolated, washed
with ethanol, and dried under vacuum. If no Pr4NBr is added and the
solution obtained after a two-electron reductive electrolysis is stored at 4"C,
then the protonated (2e-, H') product [Bu4N] [HS2Mo180 6 2 ] crystallizes.
Thus, even with the two-electron level of reduction, protonation readily
occurs. An X-ray structural determination of this protonated polyoxometalate anion [39] confirmed that the a-structural form is retained after addition
of two electrons. Salts of non-protonated species were not suitable for X-ray
~ ~ ]not
~ - ,
structural analysis. Salts of [ s ~ M ~ ~and~ [ os ~~ M~ ~] ~~~ ocould
be synthesized by bulk electrolysis. Rather, mixtures of species of different protonation levels were formed, again as predicted on the basis of the
voltammetry.
4.4 A systematic approach to chemical synthesis6 of
a two-electron vedcrced f o m of [SzMo 8 od214-
Figure 3.17 shows voltammograms at a rotated GC disc electrode over a potential range that encompasses the first two reduction waves of [ s ~ M ~ ~ ~ o ~ ~
(E;,, = 165 and -86 mV versus A g / ~ g + )and the one-electron oxidation
6 ~ d a p t e from
d
reference [39] and S. Juraja, MSc thesis, La Trobe University, 1999
Electvochemical reduction ofa -[S2M0i80621~-
-400
-200
0
200
213
400
Potential (mV vs Ag/Ag+)
3.17 Glassy carbon RDE voltammograms for 2rnM (a) decamethylferrocene;
and (b) [ s ~ M ~ ~ ~ino acetonitrile
~ ~ ] ~ - (0.1 M Bu4NPF6). The reference electrode is
A ~ / A ~ +M AgNO,; 0.1 M Bu4NPF6).
decarnethylferrocene ~ e ( q ~jMe5)2
-C
(EiI2 = -400 mV versus Ag/Ag')
in acetonitrile (0.1 M Bu4NPF6). (The E;12-values were determined by
clic voltammetry.) O n the basis of the Ei12 values (m
and callations derived from theory in Chapter 1, it can be predicted that
reaction of 2 mol of ~ e ( q ~ - C ~ M
with
e ~1 )mol
~ of [ s ~ M ~ ~ ~in Q
acetoni~ ~ ] ~ trile will produce 2 mol of [ ~ e ( q ~ - ~ ~ Mand
e ~1)mol
~ ] of
+ [ s ~ M ~ ~ ~aso ~ ~ ] ~
E~O
+
eh5-~5Me5)2 [ ~ 2 ~ 0 1 8 ~ 6 +
2 ] ~2Fe(q5-c~MQ)~]'
-k [ ~ 2 ~ 0 1 8 0 6 2 ] ~ (3.29)
that the reaction will be essentially quantitative.
Figure 3.18 shows RDE voltammograms when [ s ~ M ~ ~
is titrated
~ o ~ ~ ] ~
M acetonitrile
e~)~
(0.1 M Bu4NPF6).These
into a 2 rnM solution of ~ e ( q ~ - C ~ in
ata confirm, via noting the position of zero current, that addition of 1rnM
2 ~ 0 1 0621
8
4- is required to fully oxidize all of the 2 rnM decarnethylferrocene.
That is, the redox reaction between [ s ~ M ~ and
~ ~decamethylferrocene
o ~ ~ ] ~
occurs in the expected 1 : 2 rnol ratio (eqn 3.29). Analogous titration experiments show that when 1 rnM [ s ~ M ~ ~ ~is oadded
~ ~ ]to~a -1 mM solution
of two-electron reduced [ s ~ M ~ ~ ~ othat
~ ~[ ]s ~~ - M
, ~ ~ ~becomes
o ~ ~ ]oxi~ ized by one electron, and the [ s ~ M ~ ~ ~species
o ~ ~
reduced
] ~ - by one electron.
us, the reaction between the two-electron reduced [ s ~ M ~ ~ ~and
o ~ ~ ] ~ 2 ~ 0 1 8 0 6 2 occurs
]4
in the 1 : 1rnol ratio predicted on the basis of reversible
otentials as in eqn (3.30).
214
Illustvating basics of voltammetry
I
-700
I
I
I
-500
I
-300
I
I
-100
I
I
100
I
I
300
Potential (mV vs A ~ / A ~ + )
I
I
500
Potential (mV vs A ~ / A ~ + )
Fig. 3.18 Titration of a 2mM decamethylferrocene solution with [ s ~ M o ~ ~ oas~monitored
~]~by a GC RDE. Approximate mM ratios of decamethylferrocene: [ s ~ M o I ~ o ~are
~ ]given
~ - in
parentheses. The reference electrode is as in Fig. 3.17. Adapted from reference 39.
This is the reaction that produces the changes in the voltammograms in Fig. 3.18
when the ratio of decarnethylferrocene to [ ~ ~ ~ ois increased
~ ~ 0 from
~ 2~ : 1] ~ to 2 : 2. The experiment shown in Fig. 3.18 also confirms that [ ~ ~ ~ o ~ ~
is stable in 'dry' acetonitrile, since no evidence of processes associated with
formation of a protonated polyoxomolybdate is observed.
Rotated-disc voltammograms were also used to monitor the course of
titration of 2 rnM decamethylferrocene with the one-electron reduced polyoxomolybdate, [ s ~ M ~ ~ ~in oacetonitrile
~ ~ ] ~ -(0.1 M Bu4NPF6).It was expected
that since [ ~ ~ ~ ois already
~ ~ one-electron
0 ~ ~ ] reduced,
~ the stoichiometry of
Electrochemical reduction ofa -[S2Mo180621~-
215
the reaction between decamethylferrocene and [
~
~
~would
o be~ 1 : 1.
~
The occurrence of the reaction in eqn (3.31) was confirmed.
0
~
n CH3CN solution, it is evident that a very clean reaction occurs between
q5-C5Me5)2and [ s 2 ~ o 01682 1 4 - without intervention of protons from
adventitious water. However, isolated solid from this reaction mixture, always
had the formulation of a solvated [ ~ e ( q ~ - ~
[HS2Mo18062]
~ ~ e ~ ) compound.
~ ] ~
That is, as in the case of direct electrochemical synthesis from CH3CN
Bu4NC104) described in Section 4.3.2, the (2e-, H f ) form of the
d polyoxomolybdate was isolated.
n a typical chemical synthesis of [ F ~ ( ~ ~ - C ~ M ~ ~ ) ~
which
][HS~MO~
contains the two-electron reduced polyoxometalate and a one-electron
oxidized decamethylferrocene, 20 mL of an ether (Et20) solution (0.6 mol
e(q5-c5Me5)2) was added dropwise to 30 rnL of an acetonitrile
solution containing 0.2rnrnol of [ H ~ X ~ N ] [ S ~ M O The
~ ~ Oblue-green
~~].
itate was filtered and recrystallized from N,Nt-dimethylformamide
F). Elemental analysis of the solid was consistent with the formuion [ ~ e ( q ~ - ~ 5 M e5 [&H] S ~ M O ~( D
~M
OF
~ )~
(~~] t 2 0 ) ~The
.
5 : 1 ratio of
e(q5-C5~e5)2]+
to [ H S ~ M O I ~ O and
~ ~ ]the
~ - presence
,
of solvent molecules
of crystallization were confirmed by voltammetric and spectroscopic examination of solutions (obtained by dissolution of the solid in DMF) respectively.
lioltammetvy o f [ ~ ~ ~ o in~acidic
~ 0(9.54)
~ ~ ] ~ acetonitvile/watev media
Voltammetric generation of highly reduced forms of [ s ~ M ~ ~ ~in oacetoni~ ~ ] ~ trile has been shown to be complicated by the presence of adventitious sources
of protons. Clearly, an obvious method of deliberately seeking to understand the
nature of the interaction of the proton with reduced forms of [ s ~ M ~ ~ ~is o ~ ~ ] ~ to repeat the voltammetric measurements described above in aprotic acetonitrile in the presence of known concentrations of protons. A medium in which
the concentration of H+ can be deliberately varied is a (95/5, v/v) acetonitrile/water solvent mixture to which aqueous perchloric acid, of known H+
concentration, can be added.
'
1. Cyclic voltammetry Voltammograms at a 1-mm diameter glassy-carbon working
electrode in (95/5)CH3CN/H20 (0.02 M HC104) revealed the presence of at
least seven chemically reversible reduction waves in the narrow potential range
of 0.3 to -1.0 V versus FC/FC+ (Fig. 3.19). The processes observed in this
medium are labelled I, 11, etc., to distinguish them from the one-electron waves
216
illustrating basics of voltammetry
t
-1 .OO
I
-0.50
0
Potential (V vs FC/FC+)
0.50
Fig. 3.19 Cyclic voltammograms obtained with a 1-rnrn diameter GC disc electrode for reduction
of 1.0 rnM [ s ~ M ~ ~ ~in o(95/5)
~ ~C
]H
~ 3-C N / H 2 0 (0.1 M Bu4NC104;0.02 M HC1O4) at 20°C;
scan rate (VS-l): (a) 0.05; (b) 0.50; (c) 10. Reproduced by courtesy: Inorg. Chem. 36 (1997) 2826.
Copyright, American Chemical Society.
observed in CH3CN which have been labelled I, 11, etc., in earlier discussion.
Processes I-IV are particularly well defined in the presence of 0.02 M HC104
(Fig. 3.20). The reversible potentials in this medium are 0.19, 0.08, -0.12, and
-0.43 V versus Fc/Fc+.
Electrochemical reduction o f ~ ! - [ S ~ M0621~o~~
-0.50
0
217
0.50
Potential (V vs FC/FC+)
-20 Cyclic voltammograms obtained with a 1-mm diameter GC disc electrode using a scan
in o(95/5)
~ ~ ] ~ 0.5 V s-' for the first four reduction processes observed with 1.0 rnM [ s ~ M ~ ~ ~
CM3CN/H20 (0.1 M Bu4NC104;0.02 M HC104) at 20°C. Reproduced by courtesy: Inorg. Chem.
(1997) 2826. Copyright, American Chemical Society.
witching the potential after the first process, I, gave I ," /I;~(
= 1.0, as
uired for a chemically reversible process, and AEp = 0.038 v at a scan rate
of 50 mV s-l, consistent with an essentially diffusion-controlled overall twoelectron p r o ~ e s s .Simulations
~
presented in Section 3.1 on the reduction of
suggest that process I consists of two unresolved one-electron
cesses I and 11) accompanied by proton-transfer reactions, rather than
eous transfer of two electrons. Ell2 for process I is 0.19 V versus
FC+ and is independent of scan rate over the range 0.05-10 V s-' . This value
ositive than those of0.10 and -0.14 V versus FC/FC+for processes I and
3.2), highlighting the postive shift in potential that has occurred due to
nce of protonation reactions which accompany the electron-transfer
ilar conclusions can be drawn about processes 11-IV, which implies
re also overall two-electron, proton coupled processes, although Eliz
depend slightly upon scan rate, varying by 0.02 V in the range
0.05-10 V s-'. Thus, eight electrons have now already been added in the first
four processes, which is the total that could be added to [
~
~
~overo
entire potential range available in 'dry' acetonitrile.
rocesses V, VI, and VII (Fig. 3.19) overlap but extrapolation of information
available from earlier studies again indicate that they are derived from groups of
reversible multi-electron charge-transfer steps coupled to proton-transfer reactions. However, the reversible potentials of the simple one-electron processes,
from which these overall multi-electron processes are derived, must be more
7 ~ h nuances
e
associated with reversible two-electron EE and ECE processes are summarized in references [40-431.
~
~
2 18
Illustvating basics of voltammetry
Table 3.4 Steady-state voltammetric data for 2 rnM [ S 2 ~ 0 1 8 ~ 6 2in] 4 9515 C H 3 C N / H 2 0 (0.02 M HC104)'
Process Platinum microdisc electrodeb
GC rotating-disc electrodeb
" v = 5 m ~ s - l ;T = 20°C.
b~lectrode
diameter = 50 pm.
'Electrode diameter = 2.8 mm.
d~stimated
from plots of E versus h [ ( I L i r n- I ) / I ] .
'Estimated from ILim values.
f Estimated from data at 500 rpm.
negative than the acetonitrile solvent limit of about -3 V versus FC/FC+and
hence unobservable in that solvent (Table 3.2). Thus, protonation of extensively
reduced [
~
~
~musto be ~
the cause
~
0of potential
~
~ shifts
] well
~ in excess of
a volt. Switching at potentials more negative than the seventh wave reveals a
further group of reduction processes between -1.0 and - 1.8 V versus Fc/Fc+.
The total number of electrons that can be added to [ s ~ M ~ ~ ~ oas~ ~ ] ~
predicted in Section 4, must therefore be exceptionally large.
2. Steady-state voltammetry Figure 3.21(a) shows a near steady-state voltammogram obtained at a 50-pm diameter platinum microdisc electrode over the
potential range encompassing waves I-IV. Parameters Ell2, (Ell4 - E314)and
hi, are listed in Table 3.4. (Ell4 - E314)is expected to be 0.028 V for a simple
reversible two-electron process. The observed values for processes I11 and IV are
larger, implying that they do not involve simultaneous transfer of two electrons
at the same potential.
'Log' plots of E versus log[& - I)/I] for waves I and I1 provide n-values
of 2.1 and 1.9 (h0.I), as expected for reversible two-electron charge-transfer
processes. The intercepts provide Ell2 values (Table 3.4) similar to those found
in cyclic voltammetry. The same form of analysis of waves I11 and IV gave
apparent n-values of 1.6 and 1.5 electrons, respectively, based on the assumption that they are simple reversible processes. Clearly, they are not. However,
essentially equivalent limiting currents for each of the four waves confirm that
all these processes involve the overall transfer of two electrons.
Figure 3.21(b) shows a hydrodynamic steady-state voltammogram at a GC
RDE (diameter = 2.8mm) at a rotation rate of 500rpm. Well-defined
sigmoidal-shaped steady-state processes are seen out to -1 -5V versus FC/FC+
under these hydrodynamic conditions. Limiting currents for processes I-IV
Electrochemical reduction O ~ O ~ - [ S o62I4~ A ~ O ~ ~219
obtained at a rotation rate of 500 rpm are summarized in Table 3.4. Similar conclusions to those drawn from the above mentioned rnicroelectrode steady-state
data can be made about processes I-IV via the hydrodynamic RDE method.
he minimization of 'surface blockage7again present at very negative potentials with microdisc electrodes, but not under hydrodynamic conditions, means
that information on the number of electrons associated with processes V-VII
may be gained from data obtained with the RDE technique. Thus, since a
limiting current of -21 =t1 pA is associated with the transfer of one electron
at a rotation rate of 500 rpm in the absence of acid, the total current associated
with processes V, VI, and VII (Fig. 3.21(b)) of -210 f 4pA is equivalent to
the total transfer of ten electrons for these three processes. This conclusion is
ed for rotation rates in the range 500-3000 rpm. The combined cyclic
ady-state voltammetric data indicate that waves V-VII are chemically
reversible and involve a total of 10 electrons in the ratios 4 : 4 : 2. Consely, a total of 18 electrons has been added to the [ s ~ M o ~ ~ oanion
~~]~potential where process VII is observed to produce a (protonated) product, which is stable on the time-scale of these voltammetric experiments. The
agnitude of the limiting current of wave VIII indicates that at least eight furer electrons can be transferred in a series of unresolved processes, to give a
total of 26 added electrons in the presence of acid within the available potential
range.
I
I
I
I
I
I
I
I
I
-1.50
0
0.50
-1.50
-1.00
-0.50
0
0.50
1.00
Potential (V vs FC/FC+)
Potential (V vs FC/FC+)
ig. 3 -21 Steady-state voltammograms for reduction of 2.0 rnM [ ~ 2 ~ o l s 0 6 2 ]in
~ - (95/5)
C H 3 C N / H 2 0 (0.1M Bu4NC104; 0.02 M HC104) at 20°C; scan rate, 0.005 V s-l. (a) 50-pm
diameter platinum microdisc electrode; (b) 2.8-mm diameter GC RDE with a rotation rate of
500 rpm. Reproduced by courtesy: Inorg. Chern. 36 (1997) 2826. Copyright, American Chemical
Society.
220
Illustrating basics of voltammetry
4.5.2 Voltammetry as afunction of acid concentration in (95/5/ CH3 CN/H2 0
1. Simulation of the first two processes Simulations analogous to those presented
for the [ P ~ w ~ ~ oand
~ ~[ ]H~ -~ w ~ systems
~ ~ ~(Section
~ ] ~3)-at a range of acid
concentrations should enable an understanding of how at least the initial oneelectron processes converge into two-electron processes. As shown in Section 3,
the interesting regon with respect to acid concentration in unbuffered media
actually occurs when the concentrations of acid and polyoxometalate anion are
into
~ a 1 m~ solution of [ s ~ M ~ ~ ~in Q ~ ~
similar. Upon titration of H C ~ O
(95/5) C H 3 C N / H 2 0 , the first two one-electron waves convert progressively
to the first two-electron wave (Fig. 3.22). The processes responsible for the
one-electron waves observed prior to addition of acid have been summarized
in eqns (3.19) and (3.20).
The voltammograms of Fig. 3.23 were simulated assuming Scheme 3.7
applies, which incorporates eqns (3.19) and (3.20), and protonation equilibria
between the one- and two-electron reduced species in an analogous fashion to
reaction Schemes developed previously for the simulation of the [P2Wls062]6and [ H ~ w ~ ~6- osystems.
~ ~ ]
Simulation of Scheme 3.7 evolved from the following considerations:
(1) [ s ~ M ~ ~ ~was
o ~assumed
~ ] ~ -to be unprotonated. Experimentally, electronic spectra and electrospray mass spectra of solutions of [ s ~ M ~ ~ ~ o ~
in acetonitrile are independent of [Hf] up to 100 mM.
(2) Initial estimates of the two E;,, values of the two processes in eqns (3.18)
and (3.19) were made voltammetrically (Table 3.5).
cm2 s-' was used for the diffusion coefficient of
(3) A value of 6.4 x
[ s ~ M o ~ ~ The
o ~ same
~ ] ~value
.
was assumed for the reduced anions. The
diffusion coefficient for H+ is unknown in C H 3 C N / H 2 0 solutions. While
cm2s-', limitations
best fits were obtained with a DH+value of 1.8 x
arising from input of diffusion coefficients noted in Section 3 need to be
kept in mind.
(4) The forward rate constants kf of the three protonation equilibria were
assumed to be fast (10" M-' s-'). The backward rate constants kb were
allowed to vary.
(5) The electron-transfer processes were assumed to be reversible, so that values
of ko > 1 cm s-' were used in the simulations.
(6) Other input parameters such as electrode area, uncompensated resistance,
and capacitance were obtained directly by measurement as described in
Section 3.
Comparisons of experiment and theory are made in Figure 3.23 for [H'] =
0,0.3, and 0.7 mM. Agreement is excellent for all H+ concentrations examined
in the range 0-0.7 mM and for scan rates of 20, 50, 100, and 500 mV s-' . This
success validates the assumptions that the heterogeneous and homogeneous rates
of the redox and protonation processes are very fast relative to the voltammetric
Electrochemical reduction of a -[S2M O8 0621~~
22 1
Potential (V vs FC/FC+)
ig. 3.22 Cyclic voltammograms obtained with a 3-mm diameter GC electrode using a
scan rate of 0.5V s-I for reduction of [ s ~ M ~ ~ ~ino the
~ ~presence
] ~ - and absence of acid.
(a) (-)CH3CN (0.1M Bu4NC104)/(. . . .) (95/5) C H 3 C N / H 2 0 (3 M HC104); (b) (95/5)
CH3CN/H20 with indicated amounts of acid. Reproduced by courtesy: Inorg. Chern. 36 (1997)
4227. Copyright, American Chemical Society.
time-scale and that the redox processes are diffusion controlled under this range
of conditions. It is clear that while the observed waves differ in relative current
intensity as [H+] varies (Fig. 3.23), this does not imply irreversibility. Parameter ElI2V (versus Fc/Fc+) and K = kf/kb (M-') values used to simulate
voltammograms are listed in Table 3.5.
The simulations confirm that two one-electron reversible reduction processes
occur in (95/5) C H 3 C N / H 2 0 in the absence of acid. However, in the
Potential (V vs FC/FC+)
Fig. 3.23 Simulation of cyclic voltammograms for reduction of [ S ~ M O ~ ~ O
in ~(95/5)
~]'
CH3CN/H20.[HC104],mM: (a) 0; (b) 0.3; (c) 0.7. Experimental conditions are as in Fig. 3.22(b).
Reproduced by courtesy: Inorg. Chem. 36 (1997) 4227. Copyright, American Chemical Society.
Electrochemical reduction o f ~ ~ - [ s ~06214Mo~~
223
Scheme 3.7
Table 3.5 Summary of parameters used in the
simulation of cyclic voltammograms obtained
for reduction of [ s ~ M ~ ~ in~ (95/5)
o ~ ~ ] ~
C H 3 C N / H 2 0 (0.1 M Bu4NC104)in the presence
of HC104
Charge transfer reactions
E1/2 (V)
presence of acid, the (one-electron) reduction product [ s ~ M o ~ ~proo ~ ~ ] ~
tonates rapidly and reversibly to form an equilibrium with [ s ~ M ~ ~ ~and
o ~ ~ ] ~ [ H S ~ M O ~ ~ (Table
O ~ ~ ]3.5).
~-
The more basic (two-electron) reduction product [ ~ ~ ~ o forms
~ ~an 0
equilibrium favouring [ H ~ s ~ M ~ ~over
~ o[ ~H~S] ~~ M
- ~ ~ in
~ O
moderately
~ ~ ] ~ concentrated acid media.
~
~
]
224
Illustrating basics ofvoltammetry
The (one-electron) reduction product is essentially stable to disproportionation
in the absence of added acid, as noted in Section 4.1.4.
However, the Ell2 value for reduction of [ S ~ M O ~ ~ isOmuch
~ ~ ]more
~ - negative than that for its protonated form [ H S ~ M ~ ~ ~
(-0.13
O ~ ~
versus
] ~ -0.35 V;
Table 3.5 and so, in the presence of acid, the [ S ~ M O ~ ~ isOunstable
~ ~ ] ~to- the
disproportionation reactions in eqns (3.36) and (3.37).
Clearly, the final state of protonation of the two-electron reduction product will
depend on [H+] as quantified by the influence of the equilibria in eqns (3.33)
and (3.34).
In principle, the coalescence of all the (one-electron) reduction processes
into (two-electron) processes could be simulated. However, as the extent
of reduction increases the influence of redox cross reactions becomes very
complex and so only simulation of the initial processes has been attempted.
Nevertheless, qualitatively, certain conclusions may be reached. In (95/5)
CH3CN/H20(0.02MHC104), voltammograms exhibit four two-electron,
chemically reversible steps. The separations in Ell2 values are -0.1 1, -0.20,
and -0.3 1V, that is every two electrons (with accompanying protons) transferred to the [ s ~ M ~ ~ ~anion
o ~causes
~ ] ~a -0.1
V stabilization of the molecular
framework with respect to further reduction. In the absence of acid, the separation between adjacent pairs of one-electron reductions is constant at about
-0.6V (Table 3.2). The high basicity of the unprotonated reduced species is
arent and it can e concluded that process I is derived from processes
I and IV, process I11 from V and VI and process 1V from
VII and VIII. It therefore follows that processes V-VIII must arise from simple
electron-transfer processes whose reversible potentials are more negative than
the acetonitrile solvent limit.
2. Nernstian behaviour in hiqh acid concentrations For high acid concentrations,
it can be assumed that, during the course of the voltammetry, the concentration
of Hf at the electrode surface remains constant at its bulk solution value. This
simplifying feature means that an analytical theoretical solution is available to
The
interpret the dependence of Ei12 on [H+] when [H'] >>
overall equation for an electrode process for the reduction of a (ne-, y ~ + )
Electrochemical reduction of a -[s2M0i8 0621~-
225
reduced species is thermodynamically equivalent to
e Nernst equation relevant to eqn (3.38) indicates that a plot of Ell2 versus
ill have a slope of (x - y)RT/nF [33]. As discussed above, the
0 ~ to~be ]zero.
~ umber of protons y associated with [ ~ ~ ~ ocan~ be~assumed
r the first reduction process observed in the presence of excess
ows that the slope derived from use of the Nernst relationship will
ing steady-state voltammetry at the Pt microelectrode, Ell2 values for
two-electron processes were estimated as a function of [HC104]in
300 mM. Figure 3.24 shows a plot of E;/, for the first process versus - ln[H+] for reduction of 0.4 mM [ s ~ M o ~ ~ oin~(95/5)
~ ] ~ -C H 3 C N / H 2 0
containing 10 mMBu4NC104.There is a linear dependence, which gives x = 1
m the calculated value of slopes for [HC104] < 100rnM, and x = 2 for
C104] > 120 mM, consistent with eqns (3.39) and (3.40).
(V vs Ag/AgCl) on [HC104] for the initial two-electron reduction
. 3.24 Dependence of
(10 rnM Bu4NC104).Reproduced by courtesy:
process of [ s ~ M ~ ~ ~ino(95/5)
~ ~ CH3CN/H,0
] ~ Inog. Chem. 36 (1997) 4227. Copyright, American Chemical Society.
226
Illustrating basics ofvoltammetry
A dependence on [H+] corresponding to eqns (3.40)-(3.43) is observed for
processes I1 and 111.
All the species listed in eqns (3.39)-(3.44) are in equilibrium with other
protonated forms, although the form of analysis employed only identifies the
major species present at a given proton concentration range. At higher ionic
strengths (0.1 MBu4NC104),only eqns (3.40), (3.42), and (3.44) apply, so that
medium effects are again detected.
- acidic media
4.5.3 Spectroelectvochemical studies on [SZMo 0 6 ~ 1 ~ in
Electronic spectra may be obtained for the species predicted to exist in
Section 4.5.2 (eqns 3.40, 3.42, and 3.44) by undertaking spectroelectrochemical experiments in an OTTLE cell during the course of the reduction of
in (95/5) C H 3 C N / H 2 0 (0.2 M HC1o4). The potentials
applied in an OTTLE cell in order to bring about these reductions were
0.27V versus FC/FC+ (Ell2 = 0.31V) and 0.15V versus FC/FC+ (Ell2 =
0.22V). It can be seen from Table 3.3 and Fig. 3.25 that the two-electron
reduced, two-proton [ H ~ s ~ M ~ ~and
~ osix-electron
~ ~ ] ~ - reduced, six-proton
[ H ~ s ~ M O ~ species
~ O ~ ~have
] ~ three
absorption bands in the red-visible region,
again probably arising from suspected intervalence transitions, whereas the fourelectron, four-proton reduced [ H ~ s ~ M O ~ ~complex
O ~ ~ ]exhibits
~two bands
in this region. The data in Table 3.3 demonstrate that large wavelength shifts
are observed in the electronic spectra when protonation occurs. Typically, large
shifts in Ell2 are associated with large shifts in ,A
since both parameters are
associated with an electronic transition.
4.5.4 Directed electrochemical syntheses of veducedforms of
[SzMo 0 6 ~ 1 ~ in
- acidic media
Voltammetric studies predict that one-electron and two-electron reduced
species, [ s 2 ~ o 10 ~8~ 1 ' - and [ s 2 ~ o 1 8 0 66-,2 ] will be accessible synthetically in the absence of acid and indeed this is shown to be the case in
Section 4.3. Equations (3.39) and (3.40) indicate that the (2e-, 1 H+)- and
and [ H ~ s ~ M ~ ~ ~might
o ~ ~ ] ~ (2e-, 2 H+)-reduced species,
be synthetically accessible after a two-electron reduction at low and high [H+],
respectively. In contrast, the (le-, 1 H+)-reduced anion [ H S ~ M ~ ~ ~is O ~ ~
known to disproportionate (eqns 3.36 and 3.37) and hence will not be isolatable. Analogous arguments apply to the predicted, possible synthesis of the
Electrochemical reduction ~ ~ C Y - [ &o62I4M O ~ ~ 227
(4e-, 2 H') [ H ~ s ~ M ~
and~(4e-,
~ o4 H+)
~ ~ [ ]H~~ s ~ M O after
~ ~ aOfour~ ~ ] ~
electron reduction in high- and low-acid concentrations respectively (eqns 3.41
and 3.42). Definition of the optimum conditions for isolation of pure materials
from the solution phase was guided by the above considerations. However, relative solubilities and precipitation rates of the different salts formed on addition
of R ~ N +cations can, of course, determine the identity of the isolated solid.
Controlled potential electrolysis, coupled with coulometry, and monitoring
with steady-state voltammetry, confirmed that each of the first two two-electron
rocesses, observed in (95/5) C H 3 C N / H 2 0 , with [HC104] = 1 and 100mM
is chemically reversible: reoxidation leads to quantitative regeneration of the
original species. Further, salts of the two- and four-electron reduction products
may be obtained. The (2e-, 1 H') and (4eP, 2 H') products, [ H S ~ M ~ ~ ~ O ~ ~ ]
5000
10000
15000
25000
20000
30000
35000
Wavenumber (cm-l)
I
5000
10000
I
I
I
I
15000
20000
25000
Wavenumber (cm-l)
Fig. 3.25 Continued.
I
I
30000
I
35000
228
Illustrating basics ofvoltammetry
I
I
5000
10000
15000
25000
20000
30000
35000
Wavenumber (cml)
+
+
Fig. 3.25 Spectroelectrochemical experiments for (a) [ ~ ~ ~ o ~ 2H+
~ 0 2e~ ~ -+] ~ [ ~ 2 ~ 2 ~ 0 1 8 0 6 2 ] ~(b)
-; [ H Z S ~ M O ~ ~ O
2H+
~ ~ ] ~
2 e -+ [ H ~ S ~ M O I ~ O ~
and
~]~-;
~ ~ O ~processes
~ ] ~ for the Hex4N+
(c) [ ~ 4 ~ 2 ~ 0 1 8 0 6 2 ] ~2- ~ ' 2e- -+ [ H ~ S ~ M O reduction
salt (5 x lo-' M) in (95/5) C H 3 C N / H 2 0 containing 0.2 M HC104 and 0.1 M Hex4NC104 in an
OTTLE cell at 22°C. "Detector change. Provided by courtesy: T. Vu, Monash University, Victoria,
Australia.
+
+
+
+
and
were obtained when [HC104] = 1 mM while the (2eP9
2 H+) and (4eP, 4 H+) products, [ H ~ s ~ M ~ ~and
~ o[ ~H~ ~
] ~s - ~ M ~ ~
were obtained when [HC104] = 100rnM. A six-electron reduced form was
also produced efficiently in solution and reoxidation produced the oxidized form
quantitatively. However, upon standing, the six-electron reduced solution oxidized spontaneously (back to the four-electron level), probably by reduction
of H+ to H2, preventing isolation of salts. Similar observations of spontaneous
oxidation back to the four-electron level hold for even more highly reduced
solutions.
1. Isolation of [ P Y ~ N ] ~ [ H S ~ 0M6O2 ] and related salts A solution consisting of 5 mM [Hex4NI4[S2M~18062]
1 & HC1o4 and 0.2 M Bu4NC104 in
(95/5) M e C N / H 2 0 was reduced by exhaustive bulk electrolysis at 0.15 V versus Fc/Fc+. Voltammetric monitoring of the dark blue-green solution with
the Pt microelectrode confirmed a two-electron reduction (i.e. the point of
zero current was at the plateau between the first and second waves). Pr4NBr
was added. Storage at 4°C produced needle-like, dark blue-green crystals after
1h. These were isolated as the solvated salt. E~,+N+
and B U ~ N +
salts can be
synthesized in a similar manner, and synthesis of [Bu4NI5[HS2M018062]
salt as
an analytically pure non-solvated salt can also be achieved by isolation after
reduction in 'dry9 CH3CN (Section 4.3.2). Chemical synthesis of solvated
[ H S ~ M O ~ which
~ O ~ ~also
] includes the
[ ~ e (-qc5Me5)2j5
~
anion is described in Section 4.4.
2. Isolation of [ f i x 4 N ] 4 [ H 2 S z M o 1 8 0 6 2 ] and velated salts A solution of
5.0 mM [Hex4NI4[S2M~18062]
and 10 mM HC104 in (95/5) C H 3 C N / H 2 0
?
229
Electrochemical vedttction ofa -[S2Mo180621~-
was reduced by exhaustive bulk electrolysis at 0.14 V versus Fc/Fc+.
Voltammetric monitoring confirmed a two-electron reduction. Hex4NBr
was added. Storage at 4°C for four weeks produced blue crystals of
~ x ~ N ] ~ [ H ~ Pr4N+
S ~ Mand
~ Bu4N+
~ ~ ~ salts
~ ~ can
] . be synthesized in a
similar manner.
3. Isolation of [Pr4N]6[fiSzhfo180~ This salt was obtained in the form
of blue crystals using a procedure equivalent to that for [PrqNI5[HS2M018062]
except that the electrolysis potential was set at 0 V versus Fc/Fc+. Voltammetric
onitoring of the blue solution confirmed a four-electron reduction. A Bu4Nf
salt can be synthesized in a similar manner.
~ salt was obtained in the
4. Isolation of [ H ~ X ~ N ] ~ [ H ~ S ~ M OThis
form of deep blue crystals using a procedure equivalent to that for
ex4NI4[H2S2MoI8
0 6 2 ]except that the electrolysis potential was set at 0.09 V
versus Fc/Fc+. Voltammetric monitoring of the deep blue solution confirmed a
four-electron reduction. Pr4N+ and Bu4N+ salts can be synthesized in a similar
manner.
5.5 Voltammetric chavacterization of salts
and stoichiometry of all isolated salts were determined by elemental
s, and by monitoring the point of zero current and wave shapes of steadystate voltammograms of the salts dissolved in the media from which they were
isolated. The voltammetric technique allows, for example, the H e x 4 ~ salts
+
of [ s ~ M ~ ~ ~ o
the~ (2e7
~ ] ~2-H+)
, product [ H ~ s ~ M ~ ~ ~and
o ~the~ (dew,
] ~ - ,
4- to be distinguished. They differ in compo+) product
sition by only 2 and 4 hydrogen atoms in a molar mass of 4204 Da. The
I
-0.50
I
0
Potential (V vs FC/FC')
I
I
I
I
0.50
-0.50
0
Potential (V vs FC/FC+)
0.50
ig. 3.26 Near steady-state voltarnmograrns of 1rnM solutions of ~ e x ~ Nsalts
+ of polyoxornolybdates at a 100-prn diameter Pt microelectrode in (95/5) C H 3 C N / H 2 0 (0.2 M Bu4NC104:
0.1 M HC104): (a) [ H ~ s ~ M ~ ~ ~(b)
o ~[ H
~ ]~ ~s -~; M ~ ~ ~Reproduced
o ~ ~ ] ~ -by. courtesy: Inorg.
Chem. 36 (1997) 4227. Copyright, American Chemical Society.
230
Illustrating basics of voltammetry
(a)
I
I
I
-0.50
0
Potential (V vs FC/FC+)
0.50
Fig. 3.27 Near steady-state voltammograms at a 100-pm diameter Pt microelectrode of 2.3 mM
[(Pr4N)I6[S2M018062](a) in CH3CN (0.23 M Bu4NC104);(b) after addition of 3.8 rnM HC104.
Reproduced by courtesy: Inorg. Chem. 36 (1997) 4227. Copyright, American Chemical Society.
positions of the zero current in Fig. 3.26 confirm that [ H ~ x ~ N ] ~ [ H ~ S ~ M
is two-electron reduced and that [ H ~ x ~ N ] ~ [ H ~
is four-electron
S ~ M ~ ~ ~ ~ ~ ~
reduced. As noted in Section 4.5.1, the addition of acid to the one-electron
reduced anion [s2Mol80 ~ ~ 1causes disproportionation as evidenced by monitoring of the position of zero current in steady-state voltammograms. The
oxidation level of the two-electron reduced anion [ ~ ~ ~ ois confirmed
~ ~ 0 ~ ~ ]
by its steady-state voltammogram (Fig. 3.27(a)). Addition of acid leads to the
(2e-, 2 H') behaviour characteristic of [ H ~ s ~ M ~ ~in~acidified
o ~ ~ CH3CN
] ~ (Fig. 3.27(b)). The position of zero current is at the plateau between the first
two two-electron waves (Fig. 3.27(b)) and shows that [ s ~ M ~ has
~ ~been
o ~ ~ ]
protonated but does not disproportionate in the presence of acid.
The studies on the voltammetry of [ ~ ~ ~ oreveal
~ ~that0directed
~ ~ elec] ~ trochemical synthesis is possible for inherently complex systems, if adequate
voltammetric data are available.
4.6 Photoelectrochemical studies of [ S ~ M1O8 ~ 6 2 / 4 - using a
hydrodynamic channel electrode8
Since their discovery more than 100 years ago, many examples ofphotochemical
reactions involving polyoxometalate compounds have been identified. In the
case of [ ~ ~ ~ o ~exposure
~ 0 ~of ~
the]solid
~ -to ,sunlight for a period of time,
causes a colour change from yellow to green to blue to occur progressively.
Since [ s ~ M ~ is
~ green
~ o and
~ ~[ s] ~~ M ~ ~ ~is oblue,
~ ~these
] ~ -observations
are consistent with photoactivity involving reduction of the [ s ~ M ~ ~ ~ando ~ ~ ] ~
presumably oxidation of water, since photoactivity is not detected when light
is shone on samples contained in a vacuum.
' ~ d a ~ t ewith
d
permission from Inovg. Chem. 34 (1995) 3378. Copyright, American
Chemical Society.
Electrochemical reduction of a -[&Mol O62I4-
23 1
In order to study the photoelectrochemistry of
(and
M ~ ~ which
~ ois also
~ photoactive)
~ ] ~ a channel electrode (Section 9.2 in
apter 2) made of optical quality synthetic silica was employed with acetonitrile (0.1 M Bu4NC104)and solvent flow rates in the range 10-~-10-'cm3 s-' .
se of this flow-type cell makes it easy to shine light of the required wavelength
onto the electrode surface. Since the solution is flowing, heat is rapidly dissipated
so that thermal effects from the light source can be neglected. Finally, as noted
in Section 9.2 in Chapter 2, the hydrodynamics of the channel electrode may be
readily modelled so that the quantitative effect of the contribution of the light
to the reduction process can be calculated. The working electrodes were fabricated from thin platinum foil (thickness 0.025 mm), of dimension 4 x 4 mm, and
were periodically irradiated with a He-Cd laser light source having a maximum
irradiant power output of 40 f 5 mW cmW2.A silver wire pseudo-reference or
saturated calomel reference electrode was located upstream and a platinum foil
counter electrode downstream with respect to the channel electrode.
.6.1 Photoelectrochemical experiments in the
presence of toluene and tetrahydrojuran
As expected, on the basis of studies described in Section 4.1, hydrodynamic
channel electrode voltammograms at platinum electrodes for reduction of
~
~
~ exhibit
o four~ reversible
~
reduction
0
~ processes
~
over
]
the
~ potential range +0.7 to -1.3 V versus SCE in acetonitrile, which correspond
to the formation of [ s ~ M o ~ ~ o ~[ ~s ]~~ M
- , ~ ~ ~ o[S2M01
~ ~ ] 8~~ -6 2, ] 7 -and
,
[ s ~ M ~ ~ ~ oThe
~ ~Ell2
] ~ -values
.
are -I-0.49, +0.19, -0.46, and -1.20 V
versus SCE.
0621 5-, and
As will emerge, only the yellow [ s ~ M o ~ ~ o ~
green
~ ] ~[s2Mo18
-,
blue [ s ~ M o ~ ~areo relevant
~ ~ ] ~to the photochemistry so that the processes
relating to the formation of the 7- and 8- anions are not considered further
in Section 4.6.
From the voltammetric data, it can be deduced that if the potential of a
channel electrode is held at a value which is more positive than the first reduction
step, zero current (relative to the background) should be observed. However,
if [ s ~ M o ~ ~ oor~ ~
[ ]s ~~ - M ~ ~ ~areo generated
~ ~ ] ~ -photochemically, then an
oxidation current will be detected at these positive potentials.
Use of the platinum channel electrode, acetonitrile (0.1 M Bu4NC104)con~ irradiation
~ ] ~ - , with laser light of wavelength
taining 1op3M [ s ~ M ~ ~ ~ oand
325 nm, introduced no new features into the hydrodynamic voltammetry
over the potential range +0.9 to +0.2 V versus SCE. However, on irradiation of
in the presence of 0.6M toluene or tetrahydrofuran
(THF), steady-state oxidative currents were observed in this potential range.
These mass-transport limited currents are consistent with photochemical generation of reduced forms of [
~
~
~which
o are
~ then
~ detected
0
~ by ~their ]
transport-limited oxidation. Equations (3.45) and (3.46) summarize likely
photochemically induced processes, where ED = the electron donor
~
232
Illustrating basics of voltammetry
(toluene or THF):
The 'photocurrent' is much larger in the presence of THF than toluene, suggesting that the kinetics of the reaction of the photoexcited state molecule(s)
with THF must be faster than with toluene.
In initial experiments, electrode passivation at the platinum electrode (see also
Section 4.1.2) was substantial in the presence of light and toluene or THF and
the photocurrent decayed steadily with time. However, this problem was minimized by the use of a 'sacrificial electrode' which was irradiated upstream of the
channel detector electrode. The sacrificial electrode consisted of a platinum foil
(4 x 5 mm) located in the irradiation zone, immediately upstream of the working
electrode but insulated from it. Its purpose was to act as an adsorptive surface for
any surface-active minority species, either initially present or photogenerated in
solution, thus, preventing them from reaching and contaminating the working
electrode surface. The latter electrode was screened from the light source using
o ~ ~upstream
] ~ masking tape. In this way, [ s ~ M ~ ~ ~oro[ ~s ~ ]M~ ~- ~ ~formed
were swept to the working electrode for voltammetric interrogation while the
majority of the passivatingspecies were removed.
Using the above protocol, solutions containing known concentrations
of [
~
~
~ ando toluene
~
~or THF
0
were
~
~flowed
] through
~
the channel
electrode and the working electrode held at f0.8V versus SCE, correspond0621
0 682 1 6 - to
ing to transport-limited oxidation of [ s 2 ~ o 1
8 5- and [ s 2 ~ o 1
[ s ~ M ~ ~ ~ oLaser
~ ~ ]light
~ - .of wavelength 325 nm and intensity of 40 f
5 mW cmW2was periodically used to irradiate the area immediately upstream
of the working electrode. The resulting total steady-state transport-limited current (Fig. 3.28) was measured under irradiative conditions (A = 325 nm) as
a function of flow rate of solution over the electrode. After each measurement, the potential was switched to reducing potentials of -2.00 V versus SCE
for a period of time (about 1min) in order to clean the electrode. The photocurrents at 0.8 V versus SCE were measured over a wide range of flow rates
(6 x low4cm3 s-'-5 x 1oe2 cm3 s-') and concentrations (0.05-0.2 rnM), and at
two light intensities (Io) of 40 and 18 mW ~ m - In
~ .the absence of light there
was negligible oxidative current.
[ s ~ M o ~ ~ oabsorbs
~ ~ ] ~light
- strongly at 325 nm9 (E = 3.2 x l o 4 M-' cm-I).
It was also considered likely that a second electron transfer could occur between
photoexcited ( [ S ~ M O ~ ~ O ~
and
~ ]the
~ - electron
)*
donor, since [ s ~ M ~ ~ ~ o ~ ~
' ~ o t ethat the units of wavelength employed in Fig 3.15 are cm-' rather than nm.
Electrochemical reduction ofa-[S2Mol8o62I4-
233
LIGHT OFF
(b)
I
LIGHT OFF
~~=8.10~10-~cm~s-~
LIGHT ON
---+
time
1 minute
LIGHT ON
I
Itime
I
1 minute
ocurrent obtained in a channel electrode in acetonitrile (0.1M Bu4NC104) for
0 ~ ~ 0in ~
the~
presence
1 ~ -of (a) toluene; and (b) THF. Reproduced by courtesy: Inorg. Chem.
also exhibits strong absorption at 325 nm (E = 3.3 x lo4M-' cm-I). A theoretical CECEC model observed for photocurrents at potentials corresponding
to the transport-limited oxidation of [ ~ ~ ~ o and
~ [S2Mo
~ 0 8 ~0 ~ ~
]- ~
~ 1 ~
was
therefore developed as in eqns (3.47-3.51).
-
Since the donor concentration was in considerable excess, it was postulated
at photoreduction reactions occur to form [ s ~ M ~ ~ ~
and
o [~s ~~ ]M~ ~- ~ ~ o ~ ~ ] ~
(eqns 3.47 and 3.49) via a pseudo-first-order process which is followed by
their detection by electrochemical oxidation (eqns 3.48 and 3.50). The final
step postulated, involves decomposition of the donor (ED) cations, [toluene]+
and [THF]+. It was assumed that these cations decompose instantaneously on
the time-scale of the experiments and that this rapid decomposition therefore
recludes the occurrence of any back-electron-transfer reaction. Possible protonation of the polyoxometalate [ s ~ M o ~ by
~ oH+,
~ formed
~ ] ~ by decomposition
of ED', was also neglected because these voltammetrically distinguishable forms
234
Illustvating basics ofvoltammetvy
of the polyoxometalate were not detected under these conditions. The genera0 68
2 1 8 - or even more highly reduced species
tion of [ s 2 ~ o 01682 1 7 - and [ s 2 ~ o 1
was also not included in the model as the much higher energes required for their
formation made this a less likely event than the generation of [ S 2 ~ o 1 8 0 6 2 ] 5 and [ s ~ M o ~ ~ o ~NO
~ ]account
~ - . was taken, in the modelling, of the disproportionation step shown in eqn (3.28), since this reaction occurs to less than 1
per cent.
The experimental results were compared with the theoretical model for
CECE process derived from the solution of the relevant convective-diffusion
equations. [ s ~ M ~ ~ ~[ os ~~ ~M~ 5-,
~ -and
~ , ~[ so ~~M~~ ~~ ~were
o ~assumed
~ ] ~ to have equal diffusion coefficients of magnitude equal to that deduced for
cm2 s-' from rotating-platinum-disc
[ s ~ M ~ which
~ ~ was
o ~5.0~f]0.5~x
electrode voltammetry.
Additional modelling work was carried out to check if alternative mechanisms could give agreement with the experimental data for the photocurrent
variation with flow rate observed. The alternative mechanisms postulated all
had reversible chemical steps in which ED+ reacted with [ ~ ~ ~ o or~ ~
[ s ~ M ~ ~ ~toore-form
~ ~ ] [~ s- ~ M ~ ~ ~oro[ ~s ~ M
] ~~ - ~ ~ respectively,
o ~ ~ ] ~ - so
,
that the step for the decomposition of the donor cation became kinetically
significant.
The experimental fits for the CECE mechanism for both the toluene and
THF cases are shown in Fig. 3.29. The high correlation between theory and
experiment is indicative of the fact that the CECE mechanism provides an
accurate representation of the reaction. The rate obtained for steps kl and k2
are summarized in Table 3.6. The best fit values of kl and k2 were found to
be independent of [ s ~ M ~ ~ ~concentration,
o ~ ~ ] ~ - in the range 0.05-0.2 &,
(studies using higher concentrations were precluded due to electrode passivation, notwithstanding the sacrificial electrode approach), and as expected
theoretically, the rate constants were found to have a linear dependence on
light intensity.
A detailed knowledge of the voltammetry of [ ~ ~ ~ o has
~ enabled
~ 0 ~
directed synthesis of reduced forms of this polyoxomolybdate. This knowledge
also enables the course of redox reactions with other compounds to be understood. In particular, when acid-base reactions accompany electron-transfer
reactions, nuances such as why different types of products are formed when
[
~
~
~reacts
o with
~ Ph3P
~ and
0 Bu3P
~ can
~ be] established
~
by voltammetric
monitoring of the reaction pathway.
1°~daptedwith permission from Inorg. Chem. 37 (1998) 2366. Copyright, American
Chemical Society.
~
Use ofvoltammetric techniques
235
(a) 0.40
-
2. 0.30
w
4 2
2
9 0.20
;
-c0
0.10
0
0.010
0.020
0.030
0.040
Volume flow rate (cm3 sf1)
(b)
0.40
2* 0.30
w
5
g 0.20
;
0
2 0.10
0
0.01
0.02
0.03
0.04
0.05
Volume flow rate (cm3 S-l)
. 3.29 CECE mechanism, experiment and theory comparison for photocurrents obtained in a
o ~ (a)
~ ]toluene
~ - and (b) THF
channel electrode for (a) 0.2 mM; and (b) 0.1 mM [ s ~ M ~ ~ ~when
are the electron donors and the light intensity is (40 f5) m W crnp2. Reproduced by courtesy: Inorg.
Chern. 34 (1995) 3378. Copyright, American Chemical Society.
Table 3.6 Kinetic data obtained from photoelectro~ o ~ ~ ] ~
chemical measurements on [ s ~ M ~ in~ acetonitrile
(0.1 MBu4NC104)in the presence of toluene and THF
using a platinum channel electrode. Experimental details
are available in Inorg. Chem. 34 (1995) 3378
Rate constant
(s-l)
Value of k
Value of k
(Io= 55 m w ~ m - ~ ) (Io= 25 r n w
kl (toluene)
k2(toluene)
kl (THF)
k2 (THF)
0.0085 (*0.0005)
0.0047 ( f 0.0002)
0.0410 (A0.0050)
0.0370 ( f 0.0030)
0.0041 (f0.0005)
0.0020 (f0.0003)
0.0190 (f0.0060)
0.0150 ( f 0.0050)
236
Illustrating basics of voltammetry
5.1 Reaction of [&MOl 8 0 6 2 1 4 - with Ph3P in (95/5) C H 3C N / H 20
Figure 3.30 shows cyclic voltammograms obtained over the potential range
2]
-0.30 to +1.10V versus FC/FC+ for 1mM [ ( H e ~ ) ~ N ] ~ [ S ~ M o l 8in0 6(95/5)
C H 3 C N / H 2 0 (0.1 M Bu4NC104)in the absence and presence of 1 rnM Ph,p
at a GC macrodisc-electrode. This mixed solvent medium is the same as that
in which the simulations presented in Section 4.5.2 were undertaken, so that
quantitative details of the [ s ~ M o ~ ~redox
o ~ chemistry
~ ] ~
are available. In the
absence of Ph3P, and as seen previously, two reversible one-electron waves are
observed over the potential range considered in Fig. 3.30 with reversible El,,
values of 0.12 and -0.13 V versus Fc/Fc+. Clearly, in the presence of Ph3P, the
nature of the cyclic voltammograms and the
values change slowly with
time. Ph3P itself is irreversibly oxidized at f0.70 V under the conditions shown
in Fig. 3.30, in agreement with other studies in different media [44-481.
In the presence of Ph3P, the most positive process associated with the polyoxomolybdate system increases with respect to current magnitude and shifts to a
more positive potential, whereas the current for the second polyoxomolybdate
reduction wave decreases, as does that for the Ph3P oxidation wave (Fig. 3.30).
Steady-state voltammetry at a microdisc-electrode (Fig. 3.31, note position of
zero current) reveals that the [ s ~ M o ~ ~ osystem
~ ~ ] is~ extensively
reduced after
1 h of reaction time. After 48 h, only two processes can be detected (Fig. 3.30);
the new oxidation wave and the remaining reduction wave (Fig. 3.31) corresponding to reversible two-electron charge-transfer processes. Further, the
voltammetric feature associated with oxidation of Ph3P has completely vanished
(Fig. 3.30).
-0.4
-0.2
0.0
0.2
0.4
0.6
0.8
1.0
1.2
Potential (V vs FC+/FC)
Fig. 3.30 Voltammograms of 1mM [ ~ ~ ~ o in ~(95/5)
~ C
0 H ~3 C N
~ / H] 2 0~(0.1
- M Bu4NC104)
before and after addition of Ph3P (1 mM): (a) cyclic voltammograms at a G C disc electrode
(diameter = 3.0 mm; v = 100 mV s-l. (-, no Ph3P; - - - , 1 mM Ph3P, 1 h; . . . , 1mM Ph3P, 48 h.)
Reproduced by courtesy: Inorg. Chem. 347(1998) 2366. Copyright, American Chemical Society.
Use Ofvoltammetn'c techniques
237
2.>-After addition of Ph3P
Potential (V vs FC/FC+)
. 3.31 Steady-state voltammograms at a GC microdisc electrode (diameter = 10pm;
v = 10 mV s-') for the experiment described in Fig. 3.30. Reproduced by courtesy: Inorg. Chem.
(1998) 2366. Copyright, American Chemical Society.
-0.4
-0.2
-0.0
0.2
0.4
Potential (V) vs FC+/FC
0.6
ig. 3.32 Simulation of cyclic voltammograms of 1 m M [ S ~ M O ~ ~in O (9515)
~ ~ ] ~
CH3CN/H20; (0.1 M Bu4NC104)in the presence of differing concentrations of acid using the
parameters and mechanistic scheme described in the text. (-) [H'] = 0 mM; (- - -) [H'] = 0.3 mM;
( . . . ) [H+] = 0.7 mM).
The changes in voltammetric peak heights and potentials of the polyoxomolybdate processes as a function of time of reaction with Ph3P suggests that
an increase in proton concentration may be occurring during the course of the
reaction. Figure 3.32 shows simulated cyclic voltammograms using the parameters derived in Section 4.5.2 and the experimental conditions shown in Fig. 3.30
in the absence and presence of varying proton concentrations. The theoretical
voltammograms obtained in the presence of varying amounts of acid (Fig. 3.32)
clearly resemble, very closely, those obtained experimentally (Fig. 3.30) if loss
238
Illustrating basics ofvoltammetvy
of one mole of Ph3P is accompanied by generation of 2 mol of H+. To test
was examined experhenthis hypothesis, the voltammetry of
tally in the presence of 2 mol eq. HC104 and shown to be virtually identical
with that of the reaction solution after 48 h. The only difference is that the
first two-electron wave of the completely reacted solution is an oxidation rather
than a reduction process. O n the basis of the simulated and experimental data,
it is apparent that the final product in this reaction is the two-electron reduced
species, [ H S ~ M O ~ ~ Ogenerated
~ ~ ] ~ - ,by the initial one-electron reduction of
[ s ~ M ~ ~ ~too[ ~s ~ ~] ~
M - ~ by
~ Ph3P
~ o followed
~ ~ ] by
~ its disproportionation
in the presence of the protons released during the Ph3P oxidation process.
and Ph3P by 3 1 NMR
~
Monitoring the reaction between
spectroscopy in (95/5) C H 3 C N / H 2 0 both in the presence and absence of
electrolyte, showed that Ph3P (6 = -4.4ppm) is oxidized to P h 3 P 0 (6 =
30.1 ppm). However, the rate of reaction was found to be much faster in
the absence of electrolyte, presumably because of ion pairing (relatively weak)
of the very negatively charged polyoxomolybdate species with BU~N' that
occurs when 0.1 M Bu4NC104is present. 1 7 0 N M R experiments on solutions
enriched in H ~ (50
~ O
per cent 1 7 0 ) confirmed that the source of the oxygen
Electron
atom in the product P h 3 P 0 was solvent water and not
impact mass spectra on solids obtained from this 170-enriched reaction solution exhibited peaks at m / z = 278 and 279 in the intensity ratios expected for
~ per
O cent l 7 0 ) .
ph3p160and ph3p170produced from the presence of H ~ (50
The combination of voltammetric and N M R data therefore indicate that
[ S 2 ~ o l s ~ 6 2 undergoes
]4an overall two-electron reduction reaction to form
[ H S ~ M O ~ ~ while
O ~ ~ oxidizing
]~Ph3P to Ph3P0. O n the basis of the above
results, the reaction scheme is given in eqns (3.52) and (3.53) with the overall
reaction being represented by eqn (3.54).
Overall
Ph3P undergoes an irreversible one-electron oxidation at platinum electrodes in
acetonitrile solution and the initial product of this reaction is believed to be the
phosphonium radical cation (Ph3P+)[44-481. Hence, it is postulated that in the
C H 3 C N / H 2 0medium, Ph3P initially transfers one electron to [ s ~ M ~ ~ ~ o ~ ~ ]
Use of ooltammetn'c techniques
239
The Ph3P+ cation radical is proposed to react rapidly with water present in the
mixed solvent used, possibly forming a phosphine oxide precursor (such as the
radical Ph3POH) and protons [49].
This generation of protons will lead to protonation of [ s ~ M o ~ ~ o ~ ~ ] ~ and subsequent disproportionation to [ s ~ M ~ ~ ~and
o ~[ H
~ ]S ~ M
- O ~5-~ O ~ ~ ]
ection 4.5.2). The phosphine oxide precursor could react rapidly with
2 ~ 0 1 8 ~ 6 2 to
] 4form
the oxide and
(eqn 3.57), which
would again disproportionate.
n the basis of the above data, it can now be assumed that after addiof Ph3P to the solution of [ s ~ M ~ ~ ~ othe
~ ~
oxidative
] ~ - , component
of the current detected in voltammograms obtained under steady-state conns is predominantly due to the overall two-electron oxidation of the
final product
back to [ s ~ M ~ ~ ~ Under
o ~ ~ such
] ~ -con.
ditions, an estimate of the concentration of [ H S ~ M ~ ~ ~formed
O ~ ~may
] ~ be obtained from the equation (Section 10.1 in Chapter 2) for the lirniting current at a microdisc electrode, Ili, = 4nFDr,[Alo, where in this case
it is assumed that n=2, D = 6.4 x
cm2s-l for
and
[Ale is the concentration of [ H S ~ M O ~ ~ O ~ ~ ] ~ Other
-(~O
~
symbols
have
their usual significance. Knowledge of the initial concentrations of Ph3P and
~ 0 ~ ~in 0solution
~ ~and1 the
~ concentration
of the reaction product,
s ~ M ~ ~ ~ as
o a~function
~ ] ~ -of, time (as determined from rnicrodisc voltamtry) enables a second-order kinetic analysis to be developed which may then
utilized to estimate the rate constant, kl = 0.06 k 0 . 0 4 ~ - ' s - ' , for the
reaction between Ph3P and [ s ~ M ~ ~ ~(eqn
o ~3.54)
~ ] under
~ - non-irradiative
conditions.
-2 Reaction of [ S ~ M Ooh214with Ph3P under irradiative conditions
The photochemical reaction between [ ~ ~ ~ o and
~ Ph3P
~ 0 was
~ exam~ ] ~
ined in the presence of filtered light (300-400 nm, 10 mW ~ m - ~ )Irra.
o ~ in
~ (9515)
] ~
diation of an equimolar solution of [ s ~ M ~ ~and~ Ph3P
C H 3 C N / H 2 0 (0.1 M Bu4NC104)with this light source resulted in a colour
change from yellow to blue-green, and an overall voltammetric change identical
-
240
illustuating basics of voltammetry
to that shown in Fig. 3.30 for equivalent non-irradiated solutions. 3 1 NMR
~
confirmed P h 3 P 0 as the sole phosphorus-containing product of the photoreaction. It is therefore concluded that the overall reaction in the presence of
light is identical to that observed in its absence, the light merely accelerating the
rate-determining electron-transfer reaction between Ph3P and [S2M018062]4-.
The effect of irradiation on the kinetics of the reaction was quantified by channel electrode experiments (Section 4.6) conducted as rapidly
as possible ( < l h) in order to minimize the effects of the 'dark' reaction
(see above). An equimolar solution of Ph3P and [ ~ 2 ~ o l s 0 6 2in] ~(95/5)
C H 3 C N / H 2 0 (0.1 M Bu4NC104) was irradiated by light filtered down to
300-400 nm which represents the wavelength corresponding to the lowest
~ 0 ~ The
~ plat] ~ energy electronic absorption band of [ ~ ~ ~ o in ~acetonitrile.
inum channel electrode was held at a potential of +0.30 V, which will oxidize
photo-reduced forms of [ ~ ~ ~ o ~thus,
~ 0permitting
~ ~ ] rapid
~ - detection
,
of
such species. Neither Ph3P nor P h 3 P 0 undergo redox processes at +0.30 V in
acetonitrile.
Light (300-400 nm; intensity -10 mW ~ m - periodically
~ )
irradiated the platinum working electrode sudace, providing a phototransient response, similar to
o toluene
~ ~ ] ~or-T
that observed on irradiation of solutions of [ s ~ M ~ ~ ~and
(Section 4.6.1). The resultant photocurrent generated at the working electrode
was measured as a function of the flow rate (Vf = l0-~-10-' cm3 s-') of the
reaction solution over the electrode sudace.
It was postulated above that the rate limiting step in the dark reaction
~ one-electron-transfer
~ ] ~ process given
between Ph3P and [ s ~ M ~ ~ ~is othe
in eqn (3.55). It is now proposed that this step may be photo-accelerated by a
positive shift in redox potential in the photoexcited state:
5
[ s ~ M o ~ ~ o ~ ~ (] [~S- ~ M O ~ ~ O ~ ~ ] ~ - * )
accelerated rate
([s~Mo~~o~~
f ]Ph3P
~ - * )-----+
Product(s)
(3.58)
(3.59)
The final product, [ H S ~ M O ~ ~ Ois ~again
~ ] ~detected
-,
at the electrode surface by
its two-electron oxidation back to the original [ ~ ~ ~ oanion.
~ ~Therefore,
0 ~ ~
the reaction mechanism measured photoelectrochemically is postulated to be a
photo C2E reaction, where C2 represents the second-order reaction between
Ph3P and [S2Mo18
0G214-*
(eqn 3.60) and E represents the electro-oxidation of
o ~3.61)
~ ]: ~ back to [ s ~ M ~ ~ ~(eqn
the reduced species
]
Use of voltammetric techniques
24 1
t is assumed that water is in a large concentration excess and plays no part in
the rate determining step. kr reflects the second-order rate constant for eqn 3.60
in the presence of light and is analogous to the rate constant kl obtained in the
absence of light. A theoretical model for a C2Eprocess can be derived by solving
the relevant convective-diffusion equations, of each species, to the channel
electrode surface [50]. While considerable uncertainty exists because of the
contribution from the dark reaction, k; values obtained by comparison of theory
and experiment are in the range 0 . 6 - 3 . 0 ~ - l s-I for a variety of equimolar
concentrations of Ph3P and [S2Mo18062]4(0.4-1.0 mM), and water contents
(0-5 per cent v/v added H 2 0 ) . Comparison of these values with that for the
equivalent dark reaction in the presence of 5 per cent water (0.06 M-'s-')
shows that the redox reaction associated with eqn (3.55) when modified to give
eqn (3.59) is accelerated by up to two orders of magnitude in the presence of
light. This substantial acceleration in rate is again consistent with a significant
0 ~ ~ ] ~
positive shift in the excited state redox potential of the [ ~ ~ ~ o ~ ~and
couples over those of the ground state.
ed with Ph3P can be compared with the studies (Section 4.6)
acts
] " )as* a powerful oxidizing agent
ted ( [ S ~ M O ~ ~ O ~ ~
oluene and THF. Whereas toluene and THF have formal
reversible potentials greater than 1.1 V versus Fc/Fc+) and require photoexcitation of the polyoxo-cluster to undergo oxidation, the peak potential (and,
hence, the formal reversible potential) of Ph3P is lower and it is oxidized slowly
by [ s ~ M ~ ~ ~even
o ~in~the
] ~
absence
of light.
+
Reaction of [ S ~ M O
ob214~
with "Bu3P in (95/5)
of dgerences relative to reaction with PPh3
3 C N / H 20:an explanation
A colour change from yellow to green occurs upon mixing equimolar solutions
of [ s ~ M ~ and
~ ~"Bu3P
o ~in ~
(95/5)
] ~C H 3 C N / H 2 0 . However, the basic features of cyclic voltammograms obtained at a macrodisc-electrode are virtually
independent of time in the sense that two reversible one-electron processes are
served irrespective of the colour of the solution. This behaviour is very difent from that obtained in the presence of Ph3P where changes in wave height
otential occur. However, while the first two one-electron charge-transfer
~ ~ o the
~ position
~ ] ~ of zero current
sses associated with [ s ~ M ~ remain,
ed in steady-state voltammograms at a microdisc-electrode has shihed
and the initial one-electron reduction process has now become a one-electron
oxidation process (Fig. 3.33). The one-electron reduction, one-electron oxiation end point achieved with a [ ~ ~ ~ o to~nBu3P
~ 0molar
~ ~ratio
] of
~ 1 : 1.5. The voltammetric oxidation of "Bu3Poccurs [48] at about +0.5 V versus ~ c / F c + ,and the peak height for this process is observed to decrease with
time, using cyclic voltammetry, with 1.5 mM being required to convert 1 rnM
[ s ~ M ~ ~ ~to o1 rnM
~ ~ [] s~~- M ~ ~ ~ This
o ~ ~
result
] ~ is- consistent
.
with chemical
242
Illustrating basics ofvoltammetr,
After addition of 'Bu3p .
1n o "Bu3P
s2
-
V
2
_____-------
0
-1
2
-
,
0 -2 -
-4
-0.4
-0.2
0.0
0.2
Potential (V vs FC/FC+)
0.4
Fig. 3.33 Steady-state voltammograms of 1 mM [ S ~ M O ~ in
~ O(9515)
~ ~ ]CH3CN/H20;
~
(0.1M Bu4NC104)before and after addition of 1.5rnM nBu3Pat a 10-pm diameter GC microdisc
electrode. Reproduced by courtesy: Inez. Chem. 347(1998) 2366. Copyright, American Chemical
Society.
production of one-electron reduced [ S ~ M O ~5-,~ O
rather
~ ~ than
]
a protonated
form of the two-electron reduced species,
3 1 N
~ M R spectra for a 5 mM equimolar reaction solution revealed the
generation of a mixture of "Bu3P0 (6 = 57.6ppm) and "BU~PH' (6 =
13.7 ppm; J(31 P
H) = 480 Hz). Equivalent experiments in C H C~ N / ~ H
0~
showed that the source of hydrogen in "Bu3P2H+ was water, via observation of a triplet ( J ( 3 1 ~- 2 ~ =
) 74Hz) in place of the doublet observed
in ' ~ 2 0solutions. In summary, the nature of both the polyoxometalate
anion and phosphorus-containing products are different to those formed
when [ s ~ M o ~ ~ ois~reacted
~ ] ~ - with Ph3P. This difference may be rationalized by proposing that, [ ~ ~ ~ reacts
o again
~ with
~ "Bu3P
0 ~ in ~
an initial
] ~
one-electron-transfer step (compare with eqn 3.55).
-'
The "Bu3P+ cation radical, also reacts with water in a manner analogous to that
of Ph3P (compare with eqn 3.56)
However, "Bu3P is a much stronger base than Ph3P ( " B u 3 p ~ + , p K a=
8.43; ph3PH+,pKa = 2.73 [51]) so it reacts with the protons generated in
A n ovewiew of results
243
q n (3.63) to form "BU~PH+:
is leaves [ S ~ M O
062]
I~
as the finally observed polyoxomolybdate
s. The removal of protons by "Bu3P prevents the disproportionation of
1 8 0 6 2 1 5 - from occurring and, as a result, [ s ~ M ~ ~ ~
is othe~ product
~ ] ~ ofthe redox reaction between "Bu3P and [ ~ ~ ~ o ~The
~ overall
0 ~ reaction
~ ] ~ - .
in this case is
n summary, [S2Mo18
0 ~ ~ has
1 been shown to oxidize phosphine molecules,
oxidation processes being greatly accelerated by irradiation with light in the
300-400 nm range (corresponds to an absorption band of the polyoxometalate
ies). It is proposed that in (95/5) C H 3 C N / H 2 0 solutions, [ s ~ M ~ ~ ~ o ~ ~ ] ~
rgoes an initial one-electron-transfer process with both aryl and alkyl phosphines to form [ ~ ~ ~ o ~Concurrently,
~ 0 ~ ~ the
] phosphine
~ - .
is oxidized in the
presence of water to its phosphine oxide with generation of protons. If the prodisproportionation of [ s ~ M ~ then
~ ~theo final
~ polyoxometalate
~ ] ~
ected is [ H S ~ M O ~ ~InOcontrast,
~ ~ ] ~the
. more basic "Bu3Premoves
solution, forming "BU~PH+.
In this situation, the final polyoxroduct is [ s ~ M o ~ ~
~ ~ ] is stable to disproportionation
5-, o which
solutions and the phosphine products are a mixture of "Bu3P0
. The advantages of voltammetric monitoring of the course of
reactions can be appreciated by considering the data presented in Section 5.
ovewiew of results o
metric, simulation,
ues to polyoxometalate reduction studies
A detailed understanding of the redox chemistry of polyoxometalates in the
presence and absence of light has required the combined use of voltammetric,
ulation, spectroelectrochemical, and conventional spectroscopic techniques.
portantly, data presented in Chapter 3 should convince the reader that
once the nuances of the voltammetry have been understood, then this class
of electrochemical techniques can be employed as a powerful analytical tool for
monitoring reaction pathways. T o conclude Chapter 3, a summary of the application of different techniques, and the knowledge gained from the studies, is
resented, with respect to unravelling the redox chemistry of polyoxometalate
complexes.
244
6.1
Illustrating basics of voltammetry
Cyclic voltammetry
(1) Rapidly enabled the broad picture related to the number of reduction
processes as well as the extent of their chemical and electrochemical
reversibility to be established over a wide time domain.
(2) Enabled the reversible potential of some processes to be measured.
(3) Provided a quantitative understanding of the acid-base equilibria accompanying electron-transfer processes by comparison of simulated and experimental cyclic voltammograms.
6.2 Rotated-disc electrode voltammetry
(1) Frequently gave an excellent estimate of the relative numbers of electrons
involved in the numerous reduction processes by measurement of limiting
current values.
(2) Enabled the extent of reduction, and sometimes the identity of products, to be readily determined as a function of time by measurement of
voltammograms relative to the position of zero current during the course
of bulk-electrolysis experiments as well as during the course of chemical
reaction of [ s ~ M o ~ ~ owith
~ ~F
] ~~ -( V ~ - C ~Ph3P,
M ~ and
~ ) "Bu3P.
~ ,
(3) Enabled n-values to be determined by analysis of waveshapes for reversible
processes.
(4) Enabled the reversible potential of many processes to be established.
(5) Enabled diffusion coefficients to be measured, using the Levich Equation.
6.3
Channel-electvode voltammetry
(1) Enabled photochemical effects to be measured in the absence of significant
thermal effects, since heat is dissipated by solution flow.
(2) Enabled a quantitative understanding of photochemical reactions to be
achieved by simulations of channel electrode voltammetry.
6.4 Microdisc-electrode voltammetry
(1) Frequently provided equivalent data to the RDE experiments, but with
minimal ohmic drop so that accurate values of reversible potentials could be
established. Electrode blockage problems were more severe with the smallsurface-area microdisc-electrodes, than with macrodisc-electrodes used in
rotated-disc and cyclic voltammetry.
(2) Enabled diffusion coefficients to be determined from measurement of ILim
and use of a very simple equation.
References
245
Spectroelectrochemistry
(1) The technique of UV-visible spectroelectrochemistryhas demonstrated that
[ s ~ M ~ ~[S2M018062]~-,
~ ~ ~ ~ I [ ~~2 ~
- 0,1 8 0 6 2 1 and
~ - , even [ ~ 2 ~ o ~ ~ 0 ~ ~
can be generated quantitatively by electrolysis in a thin-layer OTTLE cell
in acetonitrile at low temperature.
(2) UV-visible spectroelectrochernical experiments enabled electronic spectra
of reduced non-protonated and protonated forms of the polyoxometalate
species to be obtained.
(3) EPR experiments after bulk electrolysis demonstrated that one- and
three-electron reduced forms of [ s ~ M ~ ~ ~are
o paramagnetic,
~ ~ ] ~ while
two- and four-electron reduced species are EPR silent.
Bulk electrolysis
(1) Enabled n-values to be determined via coulometry.
(2) Enabled directed electrochemical synthesis of reduced forms of the polyoxometalates to be achieved when used in combination with the results of
voltammetric experiments.
(3) Provided bulk quantities of material that could then be characterized by
spectroscopic techniques.
-7 Combinations of techniques
The knowledge gained by application of voltammetric, chemical, theoretical,
simulation, NMR, mass spectrometric, EPR, electronic spectra, X-ray diffraction, and other methods of measurement of the physicochernical properties of
2 ~ 0 1 8 ~ 6 2and
] 4 -its reduced forms, has been used to establish qualitative
and quantitative details of the redox chemistry of this polyoxometalate system.
The use of a wide range of physicochemical and other techniques in combination with the voltammetric methods supports the protocol recommended in
Chapter 2.
eferences
M.T. Pope, Comprehensive Coordination Chemistry, (ed. G. Wilkinson, R.D. Gillard,
and J.A. McClevert~),
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M.T. Pope and A. Muller, Angew. Chem., Int. Ed. Engl. 30 (1991) 34.
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M.T. Pope and B.W. Dale, Quarterly Rev. Chern. Soc. 22 (1968) 527.
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E. Papaconstantinou, Chem. Soc. Rev. 18 (1989) 1.
M.T. Pope and A. Miiller (ed.), Polyoxometalates: From Platonic Solids to Anti-Retrovira1
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H.T. Evans, Jr., Perspectives in Structural Chem. 4 (1971) 1.
P.D. Prenzler, C. Boskovic, A.M. Bond, and A.G. Wedd, Anal. Chem. 71 (1999)
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M.T. Pope and E. Papaconstantinou, Inorg. Chem. 6 (1967) 1147.
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J.P. Launay, Compt. Rend. 269C (1969) 971.
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J.P. Launay, Inorg. Nucl. Chem. 38 (1976) 807.
M. Rudolph, D.P. Reddy, and S.W. Feldberg, Anal. Chem. 66 (1994) 58912.
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e processes that illustrate the
uence of irreversi
ogeneous
reactions an
etition between
s that occur in the so ution phas
e electro e surface: fu
tovoltaic dye-sensitizers,
se biosensors
1 Introduction
In Chapter 3, the solution-phase voltammetric theory presented in Chapter 2
was applied to the electrochemically and chemically reversible electrode
processes associated with some polyoxometalate compounds. However, examples of this relatively straightforward reversible class of processes are not too
common and the influence of complexities that may arise in a voltamrnetric study are now illustrated in the examples of voltammetric studies in this
chapter.
Initially, the results of a fundamental study on the electrochemical oxidation
of the compound cis, r n e r - ~ n ( ~ o( )g, 1 - ~ h 2 ~ ~ (~g22~- ~~ hh22) ~ ~ ~ 2 ~ ~
in an organic solvent is presented to give an example of an electrode process
where a wide range of homogeneous processes accompany the charge-transfer
process. This is followed by a second fundamental study of the [v(co),]/O system in water, where one-half of the redox couple, [V(CO),]-, is soluble in the
solvent, whereas the V(CO), half is insoluble, so that both surface and solutionphase reactions are evident in the voltammetry. The chapter then concludes
with examples of electrode processes that are used in significant applications
of electrochemistry and also where careful control of different combinations
of homogeneous and heterogeneous processes is essential to the success of the
application.
The initial example relevant to applied applications of voltammetry is the
one-electron oxidation of Ru(2,2'-bipyridine-4,4'-dicarboxylicacid),(NCS),,
which is widely used in dye-sensitized photovoltaic cells (Section 6.2 in
Chapter 1). Since this compound has been designed specifically to be attached
Introduction
249
to a surface in the photovoltaic cell application, it is not surprising that adsorption of material accompanies this oxidation process. However, it will be shown
that despite the propensity of this system to interact with the surface, it is still
possible to obtain conditions that enable the reversible potentials, required in
the thermodynamic description of photovoltaic cells, to be calculated. Almost
since the invention of the electrochemical method over two hundred years
ago, this technique has been widely used in analytical chemistry (Chapter 1,
Table 1.1).In the case of voltammetry and related I-E-t-type techniques, the
most widely used analytical applications at the start of the twenty-first century are probably associated with trace analysis by stripping voltammetry and
monitoring of glucose in blood by diabetic patients using an electrochemical
biosensor. These two techniques of stripping voltammetry and electrochemiiosensors also utilize an inherently complex sequence of reaction pathways
that are described in the concluding section of this chapter. All the studies in
this chapter illustrate a wide range of the features that need to be addressed in
studies of complex electrochemical processes involving I-E-t principles, irrespective of whether they are being applied to problems associated with fuel
photovoltaic cells, or electroanalysis, or to fundamental studies associated
reactions involving inorganic, organometallic, organic, or biologically
rtant molecules.
4.1 IR and 3 1 ~NMR spectroscopic data obtained for the
associated with the electrochemical oxidation of cis, mern ( ~ 0(q~l-dpm)
) ~
( q ~ ~ - d ~ rinn dichloromethanea
)~r
c i s , r n c r - [ ~ n ( ~ ~ ) ~ ( ~ ' - d1866
~m)
1936
50.3
35.0
0.3
-28.1
53.2
49.3
19.0
-31.8
(v2-dpm)~rl
c i s f a c - [ ~ n ( ~ ~ ) ~ ( ~ ' - d ~ r1878
n)
1945
(v2-dpm)~rI
* c i s , r n e r - [ ~ n ( ~ ~ ) ~ ( ~ ' - d1950
~m)
(02-dpm)~r]+
* cisfac- [M~(co), (v l-dpm)
1960
(02-dpm)~r]+
(v '-dpm)
tvans-[Mn(~O)~
(v2-dprn)W
trans-[Mn(~~)~(~~-dprn)
(02-dpm)~r]+
trans-[Mn(C0)2(112-dpm)2]+
2022
203 1
1891
1971
1916
35.3
"Data taken from reference [lo].
"educed by reference to v ( C 0 ) data obtained for the isostructural neutral analogues.
250
Electrode processes
2 Elucidation o f the homogeneous reaction pathways
that accompany the electrochemical oxidation o f
cis,mer-Mn(CO),(ql-dpm)( q 2 - d p m ) ~(dpm
r
= PhzPC
in dichloromethanel
The electrochemical oxidation of stable, so-called 18-electron2 transition metal
organometallic complexes containing combinations of carbon monoxide, phosphine and other ligands has received much attention in the recent literature
11-31. Upon one-electron oxidation, a number of reaction routes for the
resulting 17-electron product may be possible [3]:
(1) If the resulting 17-electron species is stable, a chemically reversible oneelectron oxidation process is observed. This occurs in the case of oxidation of mer-c~(co), (r$dprn) (r12-dpm)[4,5] (dprn = Ph2PCH2PPh2)
and cis, rne~-Mn(CO)~
(Ph2P(CH2)2P(Ph)(CH2)2PPh2)Br [6] which are
reversibly oxidized to the isostructural 17-electron cations on the voltammetric time-scale.
(2) The product may isomerize. A classic example involves the oxidation of
fac-~n(~0),(~~-d~m)~1[7,8]
(see Section 8.1.4 in Chapter 2). In this case,
the initially generated 17-electron cationic fat' isomer, isomerizes to the
17-electron mer' cationic form. In turn, this species may be reduced to the
neutral 18-electron mer form at less positive potentials than those associated
with thefac/'ac+ couple. Overall, this sequence of reductions gives the wellknown square reaction scheme [3,7,8], summarized in eqn (4.1), which
describes a sequence of reactions, such as isomerization and cross redox
reactions that may accompany the electron-transfer step.
(3) The oxidized species may react to form a new product [3,8]. Oxidation of
mer- or ~ ~ c - M o ( c o (.rll-dpe)
),
(.r12-dpe)(dpe = PhzP(CH2)2PPh2results in
the formation of the 17-electron cationic species which may disproportionate into the starting material and dicationic [MO(CO),(11'-dpe) (q2-dpe)]2+[9].
~ d a ~ t ewith
d permission from Inorg. Chem. 38 (1999) 2005. Copyright, American
Chemical Society.
2 ~ h 18-electron
e
rule refers to a highly stable bonding arrangement in organometallic
chemistry. Details of the concept are contained in many Inorganic Chemistry textbooks.
See for example, J.E. Huheey, Inorganic Chemistry, 3rd edn, Harper International SI edition,
Cambridge, 1983, pp. 590-5. In contrast, 17-electron oxidized or 19-electron species
generated electrochemically (or chemically) tend to be far more reactive.
1
1I
i
1
!
[
!
i
I
D
Elucidation of honzogeneous reaction pathways
7
2
n
7
CO -
I-'
25 1
'
P
Structural representation of organometallic compounds associated with Section 2 where
dpm = Ph2CHzPPh2 and X = Br. Adapted from Inorg. Chem. 38 (1999) 2005. Copyright,
American Chemical Society.
his 16-electron product may then undergo internal nucleophilic attack by
the pendant phosphorus of the ql-dpe ligand to form the seven coordinate
(~~-dpe)~]~+.
18-electron species [MO(co),
In the study described in Section 2, the electrochemical oxidation
of cis, mer-[Mn(C0), (ql-dpm) (q2-dpm)Br] (Fig. 4.1 (a)) in dichloromethane
by voltammetric, bulk electrolysis, and spectroelectrochemical techniques is described [lo]. In principle, on oxidation of the 18-electron
cis, m e r - [ ~ n ( C 0(ql
) ~-dpm) (q2-dpm)Br], the resulting isostructural cis, mer+
cation may isomerize to give the 17-electron cis,fac+, and transf cationic versions of compounds whose structures are contained in Fig. 4.1 (b) and (c), or eject
d and/or undergo internal attack by the pendant phosphorus atom assowith the q1-dpm ligand to give the compound illustrated in Fig. 4.1 (d).
stem also is considered in Chapter 2, Section 8.1.4.
Voltarnmetvicstudies in dichloromethane
Figure 4.2(a) depicts a cyclic voltammogram obtained at a scan rate of
200 mV s-' and at 20°C for oxidation of cis, mer-[M~(co), (ql-dpm) (q2-dpm)
r] in dichloromethane (0.1 M Bu4NPF6). Over the potential range
-0.60 to 0.40 v3versus FC/FC+ and in the first cycle, a chemically irreversible
3 ~ potentials
1
in Section 4.2 are versus the FC'"(see Chapter 2, Section 2.3.2).
couple, where Fc denotes ferrocene
252
Electrode processes
(a)
-2nd scan
t--,
!A
'.
Potential (V vs FC"')
Fig. 4.2 Cyclic voltammograms obtained at a scan rate of 200 mV s-' for the oxidation of 1.0 mM
cis, m e ~ - [ M n ( C 0(rll-dpm)
)~
(r12-dprn)~r]
in dichloromethane (0.1 M Bu4NPF6)at a 1-mm diameter
Pt disc electrode; (a) 20°C; (b) -40°C. Reproduced by courtesy: hog. Chem. 38 (1999) 2005.
Copyright, American Chemical Society.
response is observed (process 1) with an oxidative peak potential (EF) of
f0.20 V. In addition, a reductive feature (process 2') is evident at E:~, -0.44 V
with its corresponding oxidation peak (process 2) being detected in second
and subsequent cycles at E F , -0.36V. As the scan rate increases, a reduction
peak associated with process 1' is detected (Fig. 4.3). Concurrently, the current
magnitude of process 2' decreases relative to that of process 1. The scan rate
dependence suggests that the product of process 1 seen at slow scan rates reacts
to form a new species that in turn undergoes a chemically reversible redox couple at less positive potentials (i.e. to generate processes 2 and 2' on second and
subsequent cycles). If the temperature is decreased to -40°C, redox couple 1
becomes chemically reversible (see Fig. 4.2(b)), while processes 2 and 2' disappear, indicating that the rate of the homogeneous chemical reaction following
the initial charge-transfer process is slowed upon lowering the temperature.
Process 1 was further studied using rotating-disc electrode which showed
that this oxidation exhibits Levich behaviour in the lirniting current region
(Section 9.1 in Chapter 2) over the rotation speed range of 500-3000 rpm.
Analysis of this lirniting current-mass-transport behaviour (Section 9.4 in Chapter 2 and Table 2.3) gives a diffusion coefficient of 7 x 10-~cm' s-' and suggests
that process 1 involves a one-electron charge-transfer process at all rotation
rates exarnined. Fast scan rate voltammetry using a 25-pm diameter platinum
microdisc electrode (Section 10.1 in Chapter 2) achieves chemically reversible
voltammograms which enables the reversible half-wave potential (Ei12) to be
Elucidation of homogeneous reaction pathways
8
1OOmV s-l
500mVs-1
+..*
;
;
253
;
'. %! ,
?.
. .
.
",
'
Potential (V vs ~c"')
-3 Cyclic voltammograms obtained at 20°C for the oxidation of 0.92mM
cis,mer-[~n(~0)2(q~-d~m)(q~-d~m)~r]
in dichloromethane (0.1 M Bu4NPF6) at a 1-mm
diameter Pt disc electrode using scan rates in the range 100-2000 rnV s-I . Reproduced by courtesy:
h o q . Chem. 38 (1999) 2005. Copyright, American Chemical Society.
E~O
determined for process 1 (equivalent to the reversible formal potential
assuming equal diffusion coefficients for reactant and product) from the average of
the oxidation and reduction peak potentials (Section 8.1 in Chapter 2). In the
scan rate range of 2000-5000 mV s-', the E;,,-value measured in this manner is
ependent of scan rate, as expected when the homogeneous reactions followthe one-electron transfer process are out-run. Thus, an E;,2 value for the
reversible potential of 0.16 V is obtained at 20°C. The same reversible potential
was obtained under near steady-state conditions at platinum or GC microdisc
electrodes. These microdisc data also imply that homogeneous reactions may
also be out-run under these short time domain near steady-state conditions
ection 10.1 in Chapter 2).
B u l k electrolysis and spectvoelectvochemical expeviments
ulk oxidative electrolyses experiments on solutions of cis,mer-[Mn(CO),
l-dpm) (72-dpm)~r]
in dichloromethane (0.1 M Bu4NPF6)were conducted
2OoC and -40°C using both a conventional bulk electrolysis cell with a
platinum basket working electrode and a thin-layer IRRAS cell (Section 16.2
in Chapter 2) contained within the sample compartment of an I R spectrometer enabling in situ reflective I R spectroelectrochemical experiments to be
conducted. Unless otherwise stated, the potential of the platinum working elece was held at f0.25 V which is sufficiently positive to achieve oxidation
ess 1. At both 20°C and -40°C, process 1 was determined by couloetry (exhaustive electrolysis conditions at a platinum basket electrode) to be
a one-electron process, as was also concluded to be the case in voltammetric
time-scale. The product(s) formed was examined by ex situ EPR measurements
and found to contain six lines of equal intensity indicative of a Mn(I1) product
( 5 5 ~ In =
, 512).
254
Electrode processes
In situ IR spectral monitoring the course of electrolysis experiments
at -40°C results in the two cis, m e r - [ ~ n ( c ~(ql-dpm)
),
(q2-dpm)~r]carbonyl bands at 1936 and 1866cm-' (see Fig. 4.4(a)) decreasing in
intensity while unresolved pairs of bands initially grow at approximately
2031, 2022, 1960, and 1950 cm-'. These new carborlyl bands are
attributed to the generation of two cationic cis isomeric forms of oxi) ~-dpm) (q2-dpm)~r]since the difference in fiedized cis, m e v - [ M n ( ~ 0\ql
quency of about 80 cm- relative to cis, m e v - [ M n ( ~ 0($-dpm)
)~
(q2-dprn)~d
is consistent with this interpretation. The two cis cation isomers possible
(a)
0.06 1
I
I
2100
I
1
2000
1900
Wavenumber ( c w 1 )
1
J
1800
Fig. 4.4 In situ I R spectra obtained in the carbonyl stretching range during electrolysis experiments in an IRRAS cell commencing with 1.0 mM cii,rner-[Mn(~0)~
(ql-dpm) ( q 2 - d p m ) ~ rin
]
dichloromethane (0.1 M Bu4NPF6) at -40°C. (a) Change in I R absorbance, relative to that
of the initial I R spectrum of cis, m e r - [ M n ( ~ o(ql
) ~-dpm) (q2-dpm)Br] during oxidative electrolysis at +0.25V versus Fco/+ (the t,j, arrows signify that the I R intensity initially increases
and then decreases); (b) change in I R absorbance relative to that of the I R spectrum of
trans-[Mn(~~)~(q~-d~m)(q~-dpm)~r]+
formed in (a) during reductive electrolysis at -0.55 V versus Fco/+. Reproduced by courtesy: Inorg. Chem. 38 (1999) 2005. Copyright, American Chemical
Society.
Elucidation of homogeneous veaction pathways
255
dpm)
(1950 and 2022 cm-') and
are the c i r , m e v - [ ~ n ( ~ 0 ) ~ ( q ~ - (q2-dpm)Br]+
i i s , - f a c - [ ~ n ( ~(ql
~ -dpm)
)2
(q2-dpm)Brlt (1960 and 203 1 cm-') species (see
Figs 4.1(a) and (b)). After initially increasing with time (seconds) the intensity
the overlapping v (CO) bands from compounds cis, mer- [Mn ( ~ (ql
0 -dpm)
) ~
(112-dpm)Br]+and cis,fac- [ M ~ ( c o ) (q
~ -dpm) (q2-dpm)Br] eventually decay.
The intense single carbonyl band observed to grow at 1971 cm-' is assigned
to the formation of the trans- [Mn(C0), (q -dpm) (q2-dpm)Br]+ isomeric form
ig. 4.1(c)).An additional minor product with carbonyl bands at 1878 and
1945 cm-l detected as an intermediate also disappeared upon complete electrolysis. Since cis, m e r - [ M n ( ~ 0(0'
) ~-dpm) (q2-dpm)Br] has I R bands at 1866
and 1936 cm-l, the transient species with a similar I R spectrum is attributed to
the generation of a small amount of cis,fac-[Mn(C0), (ql-dpm) (l12-dpm)~r].
1f only partial rather than exhaustive electrolysis is undertaken at -40°C,
then the I R spectrum shown in Fig. 4.5 is observed. This spectrum implies
that catalytic isomeric conversion of cis, mer- [Mn(Co), (ql-dpm)( q 2 - d p m ) ~ r ]
to cis,fac-[~n(C~)~(q'-dpm)(q~-dprn)Br] occurs in good yields in the presence of only trace amounts of the cis, m e r - [ ~ n ( ~ ~ ) ~ ( q ' - d(q2-dprn)Br]+
pm)
catlon.
pm)
redox couple has
resumably the c i s , f a c - [ ~ n ( ~ ~ ) ~ ( q ' - d(r12-dpm)~r]+/0
a similar reversible potential to that of the cis, mer-[M~(co), (0' -dpm) (q2-dpm)
~ r ] + / process
'
so that cis,fac-[Mn(C~)~
(0 -dpm) (q2-dpm)Br] formed in a bulk
trolysis experiment is oxidized to cis,fac-[M~(co), (ql-dprn) (r,~~-dprn)Br]+,
ch like the cis,mer cation form is also isomerized to the trans cationic isomer (see above), as shown by the detection of the carbonyl band at 1971 cm-'
4 . 4 ) . However, after only a small amount of electrolysis, all species required
for catalytic isomeric conversion of cis, rner-[M~(co), (0' -dpm) (l12-dpm)~r]
+
Wavenumber (cml)
. 4.5 In situ absolute I R spectrum obtained in the carbonyl stretching range during the
early stages of oxidative electrolysis at +0.25V versus FC'~' in an IRRAS cell of l.OrnM
cis, m e r - [ M n ( ~(ql
~ )-dpm)
~
(q2-dpm)~r]in dichloromethane (0.1 M Bu4NPF6) at -40°C. Note
( q ' (q2-dpm)~r]
at I878 and 1945 cm-' are marked with
that the bands for cis,f a c - [ M n ( ~ ~ ) ~-dpm)
an asterisk. Reproduced by courtesy: Inorg. Chem. 38 (1999) 2005. Copyright, American Chemical
Society.
256
Electrode processes
to cis,fac-[M~(co), (ql-dpm) (q2-dpm)Br] are present in bulk solution. The
reactions summarizing the catalytic generation of cis,fac- [Mn(CO), (q -dpm)
(q2-dpm)Br] after formation of small amounts of cis, mer-[Mn(CO), (ql-dpm)
(q2-dpm)Br]+ from cis, m e r - [ M n ( ~ 0(ql
) ~-dpm) (q2-dpm)~r]in bulk solution at early stages in the low temperature oxidation in the thin-layer IR
spectroelectrochemical cell are
'
-+ cis,fac-[~n(CO)~(q'-dpm)
(q2-dpm)~r]+
(4.2)
--+cis,fac - [ ~ n ( ~ ~ ) ~ ( q ' - (q2-dpm)Br]
dpm)
+ cis, mer- [M~(co), (q'-dpm) (02-dpm)~r]+
(4.3)
Under the conditions of electrolysis at -40°C in a conventional cell,
only reductive process 2' was observed in the cyclic and steady-state
microdisc voltammetry after exhaustive electrolysis, and there was no spectroscopic or voltammetric evidence of liberated Br-, C O , or dpm. From
steady-state microdisc electrode voltammetry, the magnitude of the reductive current associated with process 2' is the same as the oxidative current associated with process 1 before bulk electrolysis (compare Fig. 4.6(a)
and (b)). Thus, all bulk electrolysis data are consistent with the formation
of essentially a 100 per cent yield of trans-[M~(co), (ql-dpm) (q2-dpm)~r]+
1
Process 2'
Process 2 /
Potential (V vs FC"')
Fig. 4.6 Microdisc near steady-state voltammograms obtained at -4OoC with a 25-pm
diameter Pt macrodisc electrode (scan rate 10 mV s-l) from an initial 1.0 rnM solution of
cis,mer-[~n(~o),(q'-d~m)(q~-dpm)~r]
in dichloromethane (0.1 M BurNPF6) (a) before electrolysis, (b) after electrolysis at 4-0.25 V versus FC'/+, and (c) afier reduction of the solution formed in
(b) at -0.55 V versus FC'/+. Reproduced by courtesy: Inorg. Chem. 38 (1999) 2005. Copyright,
American Chemical Society.
Elucidation of homogeneous reaction pathways
257
m)
and c i s , f a c - [ ~ n ( C 0 ) ~ ( q ' - d ~ ~ )
with c i s , m e r - [ ~ n ( ~ ~ ) , ( q ' - d p(q2-dpm)Br]+
($-dpm)Br]+ cations, and cis,fac-[M~(co), (ql-dpm) (q2-dpm)Br] species
being generated at earlier stages of the experiment.
n re-reduction in the IRRAS cell at -0.55 V and -40°C, the IR signal at
1971 crn-l for trans-[M~(co), (ql-dpm) (q2-dpm)Br]+ decays with timeand is
replaced by a single carbonyl band at 1891 cm-' (Fig. 4.4(b))which is assigned to
the formation of the neutral 18-electron t r a n s - [ ~ n ( ~ ~ ) , ( q ~ (q2-dpm)Br]
-d~m)
species (Fig. 4.1 (c)).
e 3 1 N
~ M R spectrum of the product formed by the bulk oxidativetive electrolysis sequence obtained in a conventional cell was recorded
at -40°C with Bu4NC1O4being used instead of Bu4NPF6as the supporting
electrolyte. Four lines of equal intensity were observed as expected for formation
of trans- [M~(co), (q -dpm) (q2-dpm)13r](Table 4.1). The 31P N M R signal at
-3 1.8 ppm (trans-[M~(co), (ql-dpm)(q2-dpm)Br]) is consistent with retention of the pendant phosphorus. In addition, steady-state microdisc voltammetry
after this electrolysis sequence produced only one wave with an identical halfe potential and limiting current (now for process 2) to that observed prior to
ction. However, importantly the position of zero current indicates that the
species in solution is now in the reduced form (Fig. 4.6(c)). When this reduced
solution was warmed to 20°C, steady-state voltammograms indicate that concomitantly the limiting current for process 2 decreases and that for process 1
grows until process 2 finally disappears. In addition the v ( C 0 ) infra-red band at
1891 cm-' , assigned to trans-[M~(co), (ql-dpm)(q2-dpm)Br], decreases and
two new bands grow at 1936 and 1866cm-l as expected for formation of
cis,mer-[Mn(~~)~(~'-d~m)
(q2-dpm)Br],which must therefore be the thermoically favoured 18-electron form of [M~(co), (ql-dpm) (0,-dpm)Br].
en bulk electrolysis experiments were conducted at 20°C in the
conventional cell, the pink colour associated with the formation of
trans- [M~(co), (v' -dpm) ($-dpm)Br]+ gradually disappeared when the oxiized solution was left to stand (no applied potential) and the solution became
ale yellow. Cyclic voltammetric monitoring of this experiment towards comletion of the electrolysis reveals (compare Fig. 4.7(a) and (b)) that process
1 has almost vanished and that a response detected is associated with redox
uple 2. However, a previously undetected reversible process (couple 3) is
o observed with a half-wave potential of 0.37V. Further monitoring by
cyclic voltammetry after electrolysis is completed shows that couple 3 conhues to grow until 30min post-bulk electrolysis, redox couple 2 is not
etected while couple 3 is extremely well defined (Fig. 4.7(c)). The 3 1 ~
R spectrum from bulk electrolysis in dichloromethane (0.1 M Bu4NC104)
obtained at this time shows a single resonance having a chemical shift of
35.3 ppm. In situ I R spectroelectrochemical oxidation experiments in the
RAS cell at room temperature lead to a decrease in the intensity of the
cis,mer-[M~(co),(~'-dpm) (q2-dpm)Br]carbonyl bands at 1937 and 1866 cm-l
and concomitant growth of a single sharp band at 1971 cm-' due to formation of trans-[M~(co), (ql-dpm) (q2-dpm)Br]+,whereas after 30 min standing,
'
25 8
Electrode processes
Potential (V vs FC"')
Fig. 4.7 Cyclic voltammograms obtained at a 1-mm diameter GC macrodisc electrode with a
scan rate of 16OmVs-' during the course of monitoring electrolysis ex eriments at 20°C in a
conventional cell (a) before electrolysis of cis,mer-[~n(~~)~(q'-d~m)(q
-dpm)Br] (1.0 mM) in
dichloromethane (0.1 M Bu4NPF6), (b) after completion of bulk electrolysis at +0.25 V versus
FC'/+, and (c) 30 rnin after completion of bulk electrolysis. Reproduced by courtesy: Inorg. Chem.
38 (1999) 2005. Copyright, American Chemical Society.
f
the I R spectrum of the solution oxidized in a conventional cell, exhibits a
single carbonyl stretch at 1916 cm-'. The I R and NMR data of this latter
species are similar to that reported [ l l ] for the tvans-[~n(~~),(q~-d~m)~l+
cation (1916 cm-' and 33 ppm respectively) which suggests that it is this com~ -dpm) (q2-dpm)~r]'.
pound which is formed slowly from trans- [ M ~ ( c o )(q
This assignment is supported by the positive-ion electrospray mass spectrum
obtained from the bulk electrolysed solution which produced a strong signal at
cation.
) ~ ] ' Further879 (m/e), as expected for the t r a n s - [ ~ n ( ~ ~ ) , ( u ~ - d ~ m
more, the simulated and experimental mass spectra for this formulation of the
cation are in excellent agreement.
O n the basis of all the above evidence, it is therefore postulated that redox
couple 3 corresponds to the reaction
Analysis of the rate of increase in intensity of the carbonyl IR band at
1916 cm-' and decrease at 1971 cm-' , using the IRRAS cell, by a first-order
Elucidation of homogeneous reaction pathways
259
law gave a rate constant of 1.6 x
s-l at 20°C for the conversion of
( ~ o ) , ( v ' - d p m (q2-dpm)Brlt
)
to t r a n s - [ ~ n ( ~ ~ ) , ( q ~ - d ~ r n ) ~ ] + .
the evidences obtained imply that process 1, at short voltamdomains, must correspond to the oxidation of the 18-electron
( ~ (ql
0 -dpm)
) ~ (q2-dpm)Br] to the isostructural 17-electron
and
nganese(II) species cis,mer-[~n(~~),(q~-d~rn)(q~-d~m)~r]+
hat process 1' is the corresponding reduction process, so that couple 1 is
the reaction
However, c i s , m e r - [ ~ n ( C(ql
~ )-dpm)
~
(q2-dpm)Br]+ apparently is unstable
on longer time-scales so that under slow scan rate cyclic voltammetric conditions, cis,mer-[~n(~0)~(q~-dpm)(~~-dprn)~r]+
isomerizes to
(CO),(ql-dpm)(q2-dpm)Br]+. Thus, if processes 2 and 2' are associthe redox couple
then it follows that the slow scan rate voltammetric oxidation of
cis,mer- [M~(co), (ql-dpm) (q2-dpm)Br] occurs in dichloromethane via an
C scheme
is EC reaction scheme assumes that the catalytic process which gencis,fac- [Mn(C0), (ql-dpm) (q2-dpm)Br] under bulk electrolysis conons is too slow to be significant on the voltammetric time-scale.
Additionally, the reaction cis,mer-[M~(co), (ql-dpm) (q2-dpm)Br]+ -+
cis,fac-[~n(~~),(q~-dprn)(q~-d~m)~r]+may occur more rapidly than
reaction (4.7) and hence not be detected under these conditions
(not rate determining) or else it may not be detectable because of
lack of resolution of the cis,mer-[M~(co)~(O'-dpm) (r12-dpm)~r]+lo
and
cis,fac-[M~(co), (ql-dpm) (q2-dpm)Br]+/Oprocesses.
Since the 18-electron trans- [M~(co), ( ~ ~ - d p m ) ~compound
]+
is the
final stable product of bulk electrolysis, it can be concluded that the
260
Electrode processes
'
17-electron trans- [M~(co), (11 -dpm) (11,-dpm)Br]+ slowly releases bromide to
give trans-[M~(co), (r12-dp111)2]2+.
However, this trans di-cation is a strong
oxidant and can oxidize bromide to bromine, SO that steps (4.5-4.7) may be
followed by
---+ t r a n s - [ ~ n ( ~ O ) , ( ~ ~ - d p r n )+z Br]~+
k2 (slow)
ks (fast)
----+
trans-[~n(~~),(~~-f
d ~,$rm
2 )~]+
where the reaction in eqn (4.8) is the rate-determining step and the value of k2
s-l at 20°C as noted above.
associated with this reaction is 1.6 x
Thus, at 20" C the overall oxidation of cis,mer-[Mn(cO), (11' -dpm)
( r , ~ ~ - d ~ munder
) ~ r ] the time-scale of bulk electrolysis is represented by the
reaction
2.3
Simulation of the voltammetry
The cyclic voltammetric response obtained at 20°C was simulated using the
DigiSim software package [12] according to a form of the square reaction scheme
(eqn 4.1), which in this particular case can be described by eqns (4.5)-(4.7) in
combination with the slow bromide expulsion step described by eqn (4.8).
The reversible potentials for the redox processes 111' (0.16 V) and 212'
(-0.40 V) were obtained as described above. Double-layer capacitance values of
25 pF cm-2 and uncompensated resistances of 2000 C2 were used in the simulations, which are typical of a dichloromethane electrolyte system. The isomerization process (eqn 4.7) was assumed to be completely irreversible. Thus, solely for
the purposes of the simulation, a very high (lo1'), but chemically insignificant
value of the equilibrium constant was used to model eqn (4.7). The rate constant
(k2) used for the debromination of trans-[M~(co), (ql-dpm) (11,-dpm)Br]'
(eqn 4.8) was the value obtained from IR experiments, although it eventuated
that this rate constant is too slow to be significant at the scan rates employed
in voltammetric studies. The diffusion coefficients of all the species in the
cm2 s-' (the value deterelectrode-reaction mechanism were set to 7 x
m)
using microdisc
mined for compound c i s , m e r - [ M n ( ~ ~ ) ~ ( q ' - d p(q2-dpm)Br]
and rotating-disc voltammetry).
As can be seen in Fig. 4.8 excellent agreement is obtained between experiment and theory at scan rates of 100 and 1000 m~ s-' for the proposed
Elucidation of homogeneous reaction pathways
26 1
Theory
-0.6
-0.4
-0.2
0.0
0.2
Potential (V vs Fc'/O)
-0.6
-0.4
-0.2
0.0
0.2
Potential (V vs Fc+/O)
0.4
Comparison of experimental and simulated (according to the mechanism described by
eqns (4.4)-(4.7) and the parameters given in the text) voltammograms obtained at a 1-mm diameter
Pt disc electrode for the oxidation of 0.92mM c i s , r n e r - [ ~ n ( ~ ~ ) ~ ( ~ ~ - d ~ r n ) ( ~ ~at- d20°C
~rn)~r]
in dichloromethane (0.1 M Bu4NPF6). (a) Scan rate, 100 mV s-I (kl = 2.9 s-I), (b) scan rate,
1000m~ s-I (kl = 3.4 sP1). Reproduced by courtesy: Inorg. Chem. 38 (1999) 2005. Copyright,
American Chemical Society.
anism in dichloromethane when the heterogeneous charge-transfer rate
nts for electron transfer in steps (4.5) and (4.6) were both set at 0.06 s-'
en a kl value of (3.1 dz 0.3) s-' is used. Similarly good fits were obtained
for other scan rates in the range of 100-20001nVs-~ and concentrations of
cis,mev-[~n(~~)~(y~-dpm)(y~-d~m)~r]
(0.2-1.0 mM).
Conclusions derivedfvom electvochemical studies on
cis, m e r - M n ( C O ) z ( q*-dpm)(q2-dpm)
BY
e electrochemistry of cis,mer- [M~(co), (yl -dpm) (y2-dpm)~r]in organic
solvents, even though inherently complex, can be completely explained by a
combination of heterogeneous electron-transfer reactions and coupled solutionphase homogeneous chemical reactions. While there are many examples of
complex reaction schemes involving solely solution-phase processes coupled to
262
Electrode processes
electron transfer, frequently even greater complexity is associated with reactant or product interaction with the electrode su*ace. In the remainder of this
chapter, examples of nuances introduced by different forms of surface activity will be considered, as will strategies for minimizing the extent of surface
interaction in cases where this is an undesirable phenomenon.
3 Electrochemical studies on the [v(co),]-lo
aqueous media4
process in
The vanadium hexacarbon~lanion [V(CO),]- is another example of a redox
active organometallic system. By analogy with data presented in Section 2,
the stable 18-electron [ V ( C O ) , ] would be expected to be oxidized to
the neutral but inherently more reactive 17-electron V(CO),. In the study
described below, the reaction of water with V(CO), electrochemically generated from oxidation of the sodium diglyme (diglyme = 2-methoxyethyl ether,
, is
{ C H 3 0 C H 2 C H 2 J 2 0stabilized
)
[V(CO),]- salt, [ N a ( d i g l ~ m e[V(CO)6]
)~
considered as are nuances associated with the fact that V(CO), is only sparingly
soluble in water.
3.1 Voltammetric oxidation of [ V (CO)6]- in
acetone solutions containing water
In the absence of oxygen and water, the voltammetry obtained from dissolution
of [Na(diglyme)2[V(CO)6]
in acetone (0.1 M Bu4NPF6)solutions is extremely
well defined (Fig. 4.9(a)). Under conditions of cyclic voltammetry (scan rates
between 10 and 1000 mV s-'), a reversible one-electron process (eqn 4.11) is
observed at platinum, gold, and GC electrodes with the reversible half-wave
potential for the [v(co),]'/process being -0.35V versus FC+/FC in pure
acetone. Further details of the voltammetry in organic solvents are available in
reference [I 31.
(4.11)
[ ~ ( c o ) ~ ] v(co), e
+
+
Figure 4.9 shows the voltammetry of [V(CO),]- in acetone as a function of added water concentration. Clearly, the addition of water introduces
chemical irreversibility into the [v(co),]'/
couple, implying that dissolved
V(CO), rapidly reacts with water. Figure 4.10 shows the scan rate dependence of voltammograms obtained from a solution of [V(CO),]- in acetone
with 10 per cent added water. The couple is chemically reversible only at
scan rates >I000 mV s-'. Since the rate of reaction of V(CO), with water
also is dependent on the water concentration (Fig. 4.9), the system was simulated [12] as a pseudo-first-order EC reaction, according to the reactions given
4 ~ d a p t e with
d
permission from J. Phys. Chem. B 102 (1998) 1229. Copyright, American
Chemical Society.
Electvochemical studies o n the [ v ( c o ) ~ ] - ~ O
-0.6
-0.4
process
263
-0.2
Potential (V vs FC+/FC)
.9 Cyclic voltammograms obtained at 20°C at a scan rate of 1 0 0 m ~ s -and
~ with a Pt
macrodisc electrode showing the effect of addition of water (a) 0 per cent, (b) 3 per cent,
(c) 5 per cent, and (d) 10 per cent (v/v) to an acetone (0.1M Et4NPF6)solution containing 2 mM
[Na(diglyme)2][V(C0)6].Reproduced by courtesy: J. Phys. Chem. B 102 (1998) 1229. Copyright,
American Chemical Society.
in eqns (4.12) and (4.13):
e value of the pseudo-first-order rate constant, k[H20] was calculated at a
range of water concentrations by comparison of experimental and simulated
data. The second-order rate constant, k' was then calculated from a knowledge
of the known water concentrations to give a constant value of 0.3 M-' s-' for
the water concentration range 3-10 per cent. According to this model, the
chemical reversibility of cyclic voltammograms should be dependent on both
the scan rate and the concentration of water. Comparison of Fig. 4.9(d) with
Figs 4.9(b) and 4.10 shows that decreasing the water concentration by a factor of
1013 is equivalent to an increase in the scan rate by (1013)~
or 10, as required
if the postulated mechanism is correct.
264
Electrode processes
-0.6
-0.4
-0.2
Potential (V vs FC+/FC)
Fig. 4.10 Cyclic voltammograms obtained at 20°C with a Pt macrodisc electrode showing the effect
ofscan rate on an acetone (0.1 M Et4NPF6)solution containing 2 rnM [Na(diglyme)z][V(CO)6]an:
10 per cent (v/v) added water. (a) 100 mV sV1, (b) 200 mV sV1,(c) 500 mV sP1, (d) 1000 mV s- .
Reproduced by courtesy: J. Phys. Chem. B 102 (1998) 1229. Copyright, American Chemical
Society.
The overall reaction of V(CO), with a coordinating solvent such as water
(on the synthetic time-scale) has been described as a disproportionation reaction
[I41 where two-thirds of the vanadium is retained as [V(CO),]-, and carbon
monoxide gas is released
However, the pseudo-first-order kinetics observed on the voltammetric timescale indicate that the initial step in the decomposition ofV(CO), in the presence
of water probably occurs via a rate-determining substitution reaction
with subsequent reactions leading to the overall stoichiometry shown in
eqn (4.14) where V(I1) is the [ v ( H ~ o ) ~ ] ~
di-cation,
+
when the coordinating
solvent is water.
3.2 Voltammetric, EQCIZI, and chronocoulornetric studies on
the oxidation of[V(CO)& in water
Figure 4.11 shows multiple cyclic voltammetric scans obtained for oxidation of
2 rnM[Na(digl~me)~]
[V(CO),](solution) in water (0.05 M CsC104) at Pt, Au,
and basal-plane pyrolytic graphite macrodisc electrode^.^ In each case, a welldefined oxidation process, with a shape associated with diffusion control is
seen when the potential is scanned in the positive direction. A significantly
5 ~ these
n
studies in purely aqueous media, the reference electrode potential was calibrated
versus that of [ F ~ ( c N ) ~ ] ~ -process,
/~whereas in studies in acetone, the FC+/' scale was
used (see Figs. 4.9 and 4.10).
Electrochemical studies on the [v(co)~]-/'
process
265
Potential (V) vs [F~(cN),]~-'~-(1M KC1)
.I1 Ten voltammetric cycles obtained at 20°C for oxidation of [V(C0)6]- from an aqueous
[V(CO)6] at (a) platinum, (b) gold,
(0.05 M CsC104) solution containing 2 rnM [Na(digl~me)~]
and (c) pyrolytic graphite macrodisc electrodes. Reproduced by courtesy: J. Phys. Chem. B 102
(1998) 1229. Copyright, American Chemical Society.
larger reduction response (non-diffusion-controlled) is observed on the reverse
or negative potential scan. The symmetrical nature of the reduction response
indicates that the product of oxidation is insoluble in water and adsorbs or
precipitates onto the electrode surface. The overall process is consistent with
eqn (4.16)
[ v ( c o ) , ] (solution) 4 V ( C 0 ) 6(surface) -I- e-
(4.16)
Apparently, surface confinement prevents the oxidized V(CO)6(surface)species
from reacting with the water on the voltammetric time-scale so that in
urely aqueous media the voltammetry is described by eqn (4.16) rather than
eqns (4.12) and (4.13).
Confirmatory evidence for a surface-deposition process was obtained by
simultaneous cyclic voltammetry and electrochemical quartz crystal microbalance (EQCM) experiments (Section 19.2 in Chapter 2). A cyclic voltammogram (one cycle only) was recorded for 2 mM[Na(diglyme)2][V(CO),] in water
266
Electvode processes
(0.1 M NaC1) at a scan rate of 100 mV s-' over a potential range6 from -0.83 to
-0.23 V versus [ F ~ ( c N ) ~ ] ~ using
- / ~ - a Q C gold electrode. Upon oxidation, the
electrode mass increased as expected if deposition of the insoluble neutral species
V(CO),(solid) occurred. The mass deposited in this experiment was (28 f3 ng).
Assuming that the Sauerbrey equation (Section 19.2.1 in Chapter 2) is valid,
upon reduction, the mass decreased by almost the same amount (26 f 3 ng),
corresponding to reduction of V(CO),(solid) to a soluble [ V ( C O ) 6 ]salt. Using
the crystallographic data [15] for V(CO),, the radius of the molecule was calculated to be 0.39 nm, and assuming a packing factor of 0.91 for a hexagonally
close-packed flat layer, the mass increase corresponds to a final surface coverage of approximately two layers on the electrode. This result suggests that
V(CO), is not simply adsorbed onto the electrode surface, but is deposited as
solid V(CO), under conditions where the concentration of the neutral molecule exceeds the solubility limit in water and precipitates onto the electrode
surface.
[V(CO)6]in
Chronocoulometric experiments on 2.25 rnM [Na(digl~rne)~]
water (0.1 MNaC1) at Pt, Au, and GC macrodisc electrodes provide addii~n)
tional evidence for deposition of V(CO), when [ V ( C O ) 6 ] ( ~ ~ l ~ist oxidized
in aqueous media. These double-step chronocoulometric experiments consist
of an initial potential step from -0.76 to -0.26 V versus [ F ~ ( c N ) ~ ] ~ -and
/ ~ -a
reverse step from -0.26 to -0.76 V versus [ F ~ ( c N ) ~ ] ~ - 'Figure
~ - . 4.12 shows
an Anson plot [16] that confirms that the oxidation of [V(CO),]- to V(CO), at
an Au electrode is an uncomplicated diffusion-controlled process (upper straight
line). However, during the reduction of V(CO), (lower curve), there is an
initially rapid increase in charge as a function of time, corresponding to the
reduction of the majority of the surface-confined V(CO),, and dissolution of
I
0
10
Square root of time (rns')
20
Fig. 4.12 Anson plot obtained from double-potential step chronocoulometric data (gold electrode, step width = 500 ms derived from an aqueous (0.1 M NaC1) solution containing 2 mM
[Na(digl~me)~]
[V(co)(j].See reference [13] for further details. Reproduced by courtesy: J. Phys.
Chem. B 102 (1998) 1229. Copyright, American Chemical Society.
6 ~ e n.
e 5.
Electrochemical studies on the [v(co),]-lo
process in aqueous media
267
the resultant [V(CO),]-. The plot then becomes horizontal when the reduction
is completed.
alculation of the surface coverage from the chronocoulometric data [16] and
using the same packing factor assumptions as before showed that the quantity
.fV(CO), deposited is much greater than a monolayer and also dependent on
the step time. At a 500 ms step time, the chronocoulometric data corresponds
to 5.7 f0.2 (Pt), 5.5 f0.2 (Au) and 3 . 0 k 0 . 2 (GC) layers ofV(CO),, assuming
the deposition is in the form of uniform flat layers. At 1500 ms step time, deposition on Pt was 9.6 f 0.4 layers, on Au was 7.8 k 0.8, but on GC was highly
variable. Both the simultaneous cyclic voltammetry/EQCM and chronocoulometric experiments are consistent with precipitation of solid V(CO), onto the
e large apparent surface coverage suggests that V(CO), may be precipitated
as arrays of microcrystals (microparticles) on the electrode surface, rather than
as a thin film. Electron micrographs of V(CO), deposited from aqueous solution after oxidation of [Na(diglyme)2][V(CO),] showed the presence of only a
partially covered surface and the attached solid consists of lumps of approximate
diameter of 0.1 pm separated by distances of up to 1pm. Thus, it is apparent
that the solid is not in fact deposited as a uniform film (also see Chapter 5), and
hence, calculated surface coverage represents only average values of solid per
unit electrode area, rather than the extent of coverage in terms of monolayers
as is frequently assumed when using theoretical models of the kind described
above.
Conclusions devivedjom voltammetvic studies on
[ I / ( C O ) & in aqueous media
e solution-phase voltammetry of [ V ( C O ) , ] in organic solvents exhibits an
V(CO),
extremely well-defined, reversible, one-electron [V(CO),]e- process with an E;,,- value of -0.35 V versus FC+/FC.Addition of water
to the organic solvent causes the response to become chemically irreversible,
ut in this mixed solvent media all components remain soluble. Simulations of
e solution-phase voltammetry in acetone/water mixtures are consistent with
H 2 0 +-V(CO),(H20) C O being rate deterrnine reaction V(CO),
ing. The sodium diglyme stabilized [V(CO),]- salt ([Na(digl~me)~]
[V(CO),]),
is slightly soluble, but does not react with water. In contrast, solid V(CO),
is insoluble in water, but reacts so slowly that chemically reversible voltametry of the [V(CO),] - (solution)/V(CO), (solid) system is observed in aqueous
media. Consequently, oxidation of [Na(diglyrne)2][V(CO),] dissolved in water
(electrolyte)medium gives insoluble V(CO),, which precipitates onto the electrode surface. Cyclic voltammetry, EQCM, chronocoulometric, and electron
microscopy studies show that V(CO), is precipitated as lumps of solid rather
than as uniform layers. It is therefore concluded that the [V(CO)~]~'-process
provides an example of a reaction where competition between solution-phase
homogeneous reactions and surface-based processes occur.
+
- I
+
+
+
268
Electrode processes
Voltammetric studies on the oxidation of the
highly surface-active polypyridyl ruthenium
sitizer cis-Ru(11)(dcbpy)z(NCS)z(dcbpy =
photovoltaic
2,2'-bipyridin ,4'-dicarboxylic acid7
The electron transfer and photochemical properties of ruthenium polypyridyl
compounds have been studied extensively for many years (see references [l7-26]
for example) with voltammetric techniques being used to characterize their
redox properties and to determine their reversible potentials. In terms of development of photovoltaic cells based on ruthenium sensitizers (Section 6 in
Chapter 1) it needs to be noted that after substitution of the bipyridine ligands
in the 4,4' position with carboxylate groups, it is possible to attach ruthenium polypyridyl compounds to Ti02 semiconductor surfaces via ester linkages.
The combination of the light absorptive properties of the chemically attached
ruthenium sensitizers and the charge separation properties of the semiconductor electrode enable photo electrochemical cells to be constructed [27-321 with
energy conversion efficiencies of greater than 10 per cent (also see Section 6 in
Chapter 1).
The most common sensitizer used [27,32-351 in ruthenium-titania photovoltaic systems is cis-Ru(d~bpy)~(NCS)~(dcbpy= 2,2'-bipyridine-4,4'dicarboxylic acid). This sensitizer (Fig. 4.13) provides excellent absorption of
light in the visible region of the spectrum, a high electron injection rate, high
turnover rates and high stability in photoelectrochemical cells. O n the basis of
electrochemical studies on simple ruthenium bipyridyl complexes [17], it would
be expected that conventional cyclic voltammetric techniques at macrodisc electrodes could be used in a straightforward manner to measure the reversible
potential of the metal-based ground state [ ~ u ( d c b ~ y ) ~ ( ~
redox
~ S cou) ~ ] + ~ ~
ple, which is an essential component in the thermodyn.amics of the photovoltaic
cell that are summarized by eqns (4.17)-(4.20) .8
+
Dye regeneration: [ ~ u ( d c b ~(NCS)~]+
y)~
El-
+
-+ R ~ ( d c b p y ) ~ ( N C SEl
)~
Electrolyte regeneration: El
+ e-
El-
Excitation: ~ u ( d c b ~( yN) ~C S )[Ru(dcbpye-)
~ ~
(dcbpy)(NCS)2]*
(4.18)
(4.19)
+
e- Injection: [ ~ u ( d c b ~ y '(dcbpy)
-)
( N C S ) ~ ][ *~ ~u ( d c b ~ y ) ~ ( ~ ~e-S ) ~ ] +
(4.20)
7 ~ d a p t e dwith permission from J. Electrochem. SOL. 146 (1999) 648. Copysight, The
Electrochemical Society.
8 ~ hreation
e
scheme summarized by eqns (4.17)-(4.20) provides a schematic representation
of reactions that occur in a photoelectrochemical cell and demonstrate that both the ground
and excited state redox potentials of the c i i - [ ~ u ( d c b p ~ ) ~ ( ~couple
~ ~ ) ~are] +
important.
'~
El and El- are the oxidized and reduced forms of the electrolyte (commonly I; and I-,
respectively).
Fig. 4.13 Structural representation of cis-R~(dcbpy)~
(NCS)2.
However, since cis-Ru(dcbpy)2(NCS)2 (Fig. 4.13) has been designed to be
to electrode surfaces, the likelihood of surface activity being associated
with the Ru(II)/Ru(III)oxidation process is substantial. Importantly, if surfacebased reactions are coupled with the charge-transfer process then the application
of techniques such as cyclic voltammetry and diffusion-controlled theory may
not readily provide the values of the reversible potential (see Section 18.4 in
ion it is demonstrated that the combined use of a range of electrode
ectrochernical techniques (cyclic, rotating disc, and microelectrode
), and chemical modification of the electrode surface is required
to ensure that the effects of surface and other reactions present with the photovoltaic sensitizer are minimized so that the required reversible potential of the
mass-transport-controlled solution-phase process is correctly measured. Again,
ection 3, the extent ofthe surface activity will be shown to be conveniently
monitored by the EQCM method.
ejerence studies o n model mass-transport-controlled processes
The subtleties of the c i s - [ ~ u ( d c b ~ ~ ) ~ electrode
( ~ ~ ~ ) ~process
] + / ~associated
interactions with the surface are conveniently demonstrated by comparwith data obtained with model systems that are mass-transport controlled
~ ] ~ -in/ ~ (Section 7 in Chapter 2). For example, the [ s ~ M ~ ~ ~ o ~process
acetonitrile is essentially ideally behaved (Chapter 3) and classically the oxidation of Ferrocene (Fc) to the Ferricenium cation (Fc+) in many organic solvents
4.21) is assumed to be reversible and mass-transport controlled.
e close to ideal nature of this process is one of the reasons why the voltamtric oxidation of ferrocene is used as a reference potential scale (Section 3.2
PY)2]2+complex (bpy = 2,2'-bipyridyl) of which the photovoltaic sensitizer cis-Ru(dcbpy), (NCS)2 is a derivative, also exhibits ideal
Electrode processes
Table 4.2 Cyclic voltammetric data at 20°C for the oxidation
of 0.5 rnM [ ~ u ( b ~ y )in~ acetonitrile
]~+
(0.1 M Bu4NPF6)at
a 2-mm diameter GC disc electrodea
aData taken from reference [22].Symbols are defined in Chapter 2.
b~otentialsare reported versus the FC/FC' redox couple.
'The dependence of AEp on scan rate implies that the process is quasireversible.
d ~ lvalues
p
calculated as ( E r E F ~ ) / ~ .
+
mass-transport-controlled voltammetry under conditions of cyclic voltammetry
at platinum or GC macrodisc electrodes. [ ~ u ( b ~ ~exhibits
) ~ ] ~one
+ chemically
reversible metal-based oxidation (Fig. 4.14(a), eqn 4.22) and three chemically
reversible ligand-based reductions (Fig. 4.14(a), eqn 4.23).
Since the Ru(II)/Ru(III) oxidation process is of interest in this study, only
data relevant to eqn (4.22) are presented.9 The reversible E;,, value for the
Ru(II)/Ru(III)process in acetonitrile, derived from the average of the oxidation
and reduction peak potentials, is located at (888 4 4) mV versus FC/FC+ and is
independent ofscan rate (Table 4.2). The E;,, value in acetone was (882415) mV
versus FC/FC+.The peak current for the oxidation process in both, acetonitrile
and acetone, was found to be proportional to the square root of scan rate,
~ , 1 : 1.00( & 0.05). The diffusion
and the ratio of peak currents, 1 ~ / $ ewas
coefficient for [ ~ u ( b ~ y )in
~ ]acetone,
~+
calculated from the Randles-Sevzik
equation (eqn 2.34) was (2.2 & 0.5) x lov5cm2s-' .
Under steady-state conditions at a 6 pm radius GC microdisc electrode, one
oxidative and three reductive processes also were observed in acetonitrile and
'!See Chapter 2 for definition of the symbols and equations used to calculate parameters
reported in this section.
Voltammetric studies on polypyridyl ruthenium
271
2.0
(a)
1.o
-3
-
0.0
U
6
5:: -1.0
-2.0
-3.0
Potential (V vs FC/FC+)
0.5
0.6
0.7
0.8
0.9
1.0
1.1
Potential (V vs FC/FC+)
.14 Voltammetry of 1 mM [ ~ u ( b ~ ~ ) at, ]20°C
~ + (a) cyclic voltammetry with a scan rate
mV s-' at a 1 mm diameter Pt electrode over the potential range encompassing the single
oxidation and three reduction processes in acetonitrile (0.1 M Bu4NPF6),(b) RDE voltammogram
of the oxidation rocess in acetone (0.1 M Bu4NPF6),using a rotation rate of 1000 min-' , and a scan
rate of 10 mV s-', with a 3 mm diameter Pt disc electrode. Reproduced by courtesy: J. Electroihem.
46 (1999) 648. Copyright, The Electrochemical Society.
acetone. The Eil, value of the oxidation process in acetone was found to be
0.89(&0.01)V versus FC/FC+.
he RDE experiments on [ ~ u ( b ~ ~in) acetone
~ ] ~ + also produced wellefined oxidative (Fig. 4.13(b)) and three reductive processes (not shown).
e average value of E314 - Ell4 of 6 0 ( f 3) mV compares well with the
eoretical expected value of 56 mV (for T = 20°C) for a reversible oneelectron process (Section 9.4 in Chapter 2). Ell2 was found to be O.W(f 0.02) V
versus Fc/Fc+ for the oxidation process at the RDE. In summary, the
272
Electrode processes
[ R ~ ( ~ P 2+/3+
Y ) J process like that for FC/FC+ is essentially an ideal chemically
reversible mass-transport-controlled one-electron oxidation process.
4.2 Electrochemical studies o n c i s - R u (dcbpy)2 (NCS)2i n acetone
Comparison of the structures of cis-R~(dcbpy)~(NCS)z
and [Ru(bpy)J2+ suggests that cis-R~(dcbpy),(NCS)~
would also exhibit one reversible or close to
reversible metal-based oxidation process. Irreversible processes due to oxidation of the thiocyanate ligand also could be expected [36]. Simple inspection
of cyclic voltammograms for oxidation of the photovoltaic dye-sensitizer in
acetone (Fig. 4.15) shows that two oxidation processes are present, but that
neither corresponds to the expected simple diffusion-controlled process. The
even more complex series of reduction processes for Ru(dcbp~),(NCS)~
and
related photovoltaic sensitizers are described in references [23-26,371.
4.2.1 Cyclic voltammetry at macrodisc electrodes
O n the basis of analysis of cyclic voltammograms of the kind as shown in
Fig. 4.15 and more extensive data contained in reference [22], the mechanism for the oxidation of cis-R~(dcbpy),(NCS)~
may be postulated to involve
the expected solution-phase one-electron charge-transfer process coupled with
surface-based processes that occur prior to and after oxidation as well as oxidation of surface-modified species at positive potentials. Equations (4.24)-(4.28)
Potential (V vs F C / F C + )
-0.2
0.0
0.2
0.4
0.6
0.8
Potential (V vs F C / F C + )
1.0
Fig. 4.15 Cyclic voltammograms of cis-R~(dcbpy)~(NCS)~
at 2OoC in acetone (0.1 M Bu4NPF6)
obtained at different electrode materials diameter, (d), concentrations (c) and scan rates (v).
(a) c = 0.5 mM, v = 50 mV s-', GC electrode, d = 2rnrn; (b) c = 0.5 mM, v = 2000 mV s-',
GC electrode, d = 2 mm; (c) c = 0.2 mM, v = 50 m~ s-', Pt electrode, d = 1 mm; (d)
c = 0.5 mM, v = 1000 mV s-', GC electrode, d = 2 mm; (e) i = 0.5 mM, v = 100mV S-l,
GC electrode, d = 2 mm. Reproduced by courtesy: J. Electrockem Soc. 146 (1999) 648. Copyright,
The Electrochemical Society.
Voltammetric studies on polypyn'dyl ruthenium
273
the basic features of the postulated mechanism
Dsol
A
Dsurf
and Dsudrepresent the dye molecule, cis-R~(dcbpy),(NCS)~
in the
and surface-confined states respectively, and Xsud or Ysudrepresent
surface-attached forms of the dye, whose identities are unknown.
As theoretically predicted (Section 18.4 in Chapter 2), and as shown in
Fig. 4.15, when surface-based phenomena are coupled to the electron-transfer
the shape and nature of voltammograms are strongly affected by the nature
e working electrode, the scan rate, the switching potential, the extent of
potential cycling and the concentration. In this particular case, the extent
of material attached to the surface is related to the current magnitude of
the processes at positive potentials, and is less on platinum than on GC
or gold electrodes [22]. The scan rate, the electrode material and concenratio of the first
influence both the
tration of R~(dcbpy),(NCS)~
oxidation process and the prominence of the second oxidation peak (compare Figs 4.15(a)-(c)). The scan rate relationship is expected, since the primary
solution-phase oxidation process (eqn 4.26) becomes closer to diffusion cond at fast scan rates when it would be predicted that the extent or formation
was never unity
and hence YSudis minimized. The ratio of I,""/IF~
as required for a completely diffusion-controlled process (Section 8.1.1 in
Chapter 2), but at fast scan rates and when the potential is switched prior to
the onset of the second phase, it approaches 2 : 1 (see e.g. the cyclic voltammogram obtained at moderately fast scan rates of 1000 mV s-' in Fig. 4.15(d)).
A ratio greater than unity is expected (Section 18.4 in Chapter 2) when reactant
adsorption occurs [38]. Also noteworthy is the significant decrease in the peak
heights detected in the first and subsequent cycles of the potential (Fig. 4.15 (e)).
Apparently, during the course of sweeping the potential, surface-attached material is removed from the surface. Thus, when the potential is returned to the
initial value prior to commencing the second cycle, less time is available for
the reactant to become attracted to the electrode surface than is the case with
the initial cycle, where the electrode is held at the initial potential value for a
significantly longer period of time.
1,""/yd
274
Electrode processes
Table 4.3 Cyclic voltammetric data obtained for the oxidation of
c i ~ - R u ( d c b p ~ ) ~ ( N in
C Sacetone
)~
(0.1 M Bu4NPF6)at 1-mm diameter
platinum and 2-mm diameter GC macrodisc electrodes as a function of
scan rate (v) and concentrationa
aData taken from reference [22]. Symbols are defined in Chapter 2. Potentials are reported
versus Fc/Fc+. E 1 p values calculated as (EpO" E F ~ ) / ~ .
b ~ odetected.
t
+
For a simple diffusion-controlled process, the reversible Eil, potential is
E r d )121
expected to be approximated as the average value of E,O" and EFd[(E;
as is the case for the ideal FC'/+ and [ ~ u ( b ~ y ) ,13+
] ~processes
+
considered above.
The data in Table 4.3 actually indicate that the potential of the first oxidation
process calculated in this manner is remarkably insensitive to the experimental conditions, unlike other features of the voltammetry. The (E,O" EFd)/2
+
+
Voltammetric studies on polypyvidyl ruthenium
275
listed in Table 4.3 were all calculated from the third cycle of the
potential, but no significant difference was found in values obtained from other
cycles. The peak-to-peak separations are clearly larger than those obtained
for oxidation of [ ~ u ( b ~ ~ ) (see
, ] ~Table
+
4.2) at the same scan rate, and
part of this feature may be attributed to the surface effects.'' However, the
calculated (E,O" Erd)/2 value of (0.41 f 0.01) V versus FC/FC+ for the
c b p y ) 2 ( ~ ~ ~ ) 2 ]couple,
0 i + was found to be independent of concentration (0.05-0.5 M ) , scan rate (50-2000 mV s-l), and electrode (platinum, GC,
old) at the 10 mV uncertainty level, which implies that this is a good
approximation of the reversible potential (Ei12 value) for the solution-phase
diffusion-controlled process. Therefore, it is concluded that the influence of the
surface-based reactions does not appear to be significant in the thermodynamic
sense, with respect to the first oxidation process.
+
Microelectrode voltammetry
ar steady-state voltammetry of cis-R~(dcbpy)~(NCS)~
(0.05-0.5 M)
tinum microdisc electrode in acetone shows no evidence of a second
n process. In contrast, at a GC microdisc electrode there is a small
xidation response, detected as a slightly rising limiting current region
(a)),and which is enhanced as the concentration is increased. As noted
10 of Chapter 2, radial, rather than linear, diffusion is dominant under
-state conditiok, so it would appear that the greater flux of material away
microelectrode, that is associated with radial relative to linear diffusion,
has removed the majority of the oxidized material away from the electrode
ce before significant interaction occurs with the surface. The E;12 value
ned from microdisc-electrode experiments was found to be (0.40 f0.01) V
Potential (V vs FC/FC+)
Potential (V vs FC/FC+)
-16 Steady-state voltammograms of 1.1 mM cis-Ru(dcbpy)2(NCS)2 in acetone (0.1 M
Bu4NPF6)at 20°C, (a) GC microdisc electrode, d = 1 2 p n , c = 1.1mM, v = 10mVs-I, (b) GC
RDE, d = 3 mm, w = 500,1000,1500,2000,2500, and 3000 min-I , v = 10 mV s-I . Reproduced
by courtesy: J. Electrochem. Soc. 146 (1999) 648. Copyright, The Electrochemical Society.
10
he [ ~ ~ ( b 2+/3f
~ y ) and
~ ] [ ~ u ( d c b ~(NCS)~]'/+
y)~
processes are also probably both
quasi-reversible rather than reversible with respect to the electron-transfer process (Section 8.1.3 in Chapter 2) with the electron-transfer rate being faster for the [ ~ u ( b ~ y ) ~ ] ~ + ' ~ +
process.
276
Electrode processes
versus Fc/Fc+, which is in good agreement with the value obtained under
conditions of transient cyclic voltammetry at macrodisc electrodes. Assuming
that the limiting current is diffusion controlled under these near steady-state
conditions and use of eqn (2.48), gives a calculated diffusion coefficient of
(9.5 0.5) x 1o - cm2
~ s-' , which is slightly smaller than the diffusion coefficient
calculated for [ ~ u ( b ~ ~in)acetone
~ ] ~ +(Section 4.1).
+
4.2.3 Rotating-disc electrode voltammetry
The RDE voltammetry of cis-Ru(d~bpy)~(NCS)~
in acetone is consistent
with the suggestion that the second oxidation process is due to oxidation of
surface-attached material. At the lower rotation rates and at a concentration of
1.1 x lov3M, the second process is of almost the same size as the first, but as
the rotation rate is increased, the relative size of the second process decreases,
until at sufficiently high rotation rates it eventually disappears (Fig. 4.16(b)).
Again, the second process is less evident in platinum than GC or gold electrodes. Presumably, at high rotation rates, products are swept away from the
electrode suriace so rapidly that not enough time is available for interaction with
the electrode surface. This feature of the RDE experiment results in the elimination of the second peak as is the case when products are removed by radial
diffusion at a microdisc electrode. As expected for a mass-transport-controlled
process, a plot of limiting current is dependent on the square root of the rotation
frequency at a platinum electrode over the range of 500-3000 min-' and passes
through the origin (Section 9.4 in Chapter 2). The value of (E314- EIl41for the
first oxidation process is (66 f 3) mV at a rotation frequency of 3000 min-',
and (68 rt 3) mV at a rotation frequency of 500 min-' relative to a value of
56 mV expected for a reversible process (Section 9.4 in Chapter 2). Furthermore, the Ell2 value is independent of rotation rate and electrode material
and is located at (0.409 4~ 0.002) V versus Fc/Fc+, which is in agreement with
values calculated from cyclic voltammetric and microdisc-electrode voltammograms. Consequently, both the microdisc steady-state and rotating-disc
techniques give almost ideal voltammetric behaviour for a chemically reversible
mass-transport-controlled one-electron oxidation process.
4.2.4 Studies with an electrochemical quartz crystal microbalance
T o confirm the presence ofthe surface-based processes, studies were carried out
on gold electrodes using the EQCM technique, where small mass changes on
surfaces occurring during voltammetric experiments can be detected, assuming
that the Sauerbrey equation is valid (Section 19.2 in Chapter 2).
Figure 4.17 shows the EQCM response at the open-circuit potential for an
acetone solution (0.1 mM Bu4NPF6) before and after the solution is spiked
with 0.6 mM cis-Ru(dcbpy)2(NCS)Z.Apparently, a large amount of material
becomes attached to the gold electrode under open-circuit conditions at a stationary electrode. Under the conditions of Fig. 4.17, a rapid mass increase of
410 ng occurred onto the surface of the 5 mm gold electrode, which, assuming
Voltammetn'c studies on polypyridyl ruthenium
277
u
200
300
Time (s)
-17 EQCM experiments at 20°C in an acetone (0.1 M Bu4NPF6)solution which is spiked
with 0.6 mM ci~-Ru(dcbpy)2(NCS)~
at the time indicated. Open-circuit potential, 5-mm diameter
gold Q C electrode. Reproduced by courtesy: J. Electrochem. Soc. 146 (1999) 648. Copyright, The
~lectrochernicalSociety.
a flat surface and that the material is adsorbed in an unreacted form, corresponds
to a surface coverage of l-' = 3 x lop9mol ~ m - The
~ . surface coverage was
calculated assuming a flat electrode surface and that the cis-R~(dcbpy)~(NCS)~
molecules were spheres having a diameter of 1 4 k This diameter was estimated from X-ray structural data [39,40]. If cubic packing for the molecule is
assumed, a surface coverage corresponding to 30 monolayers is calculated at the
open-circuit potential.
e EQCM experiments during a cyclic voltammetric experiment in acetone
on a gold electrode give the results shown in Fig. 4.18(a). Results from this
experiment show that holding the potential at the initial value of -450 mV
versus Fc/Fc+ for 120 s and then scanning until a potential of about 200 mV
versus Fc/Fc+ is reached leads to only a small decrease of material initially
attached to the gold electrode. However, at about 200 mV versus FC/FC+which
corresponds to the onset of the oxidation process (eqn 4.26), a large amount of
material is rapidly removed from the electrode surface. The amount lost of about
350 ng is almost equal to the amount originally adhered onto the surface under
open-circuit conditions. When the potential is scanned to values slightly beyond
the half-wave potential, a very abrupt mass increase occurs, but as more positive
potentials are reached, the mass of material attached to the electrode decreases
to a value which is similar to that at the start of the experiment. Incorporation
F,, the electrolyte anion, into solid attached to the electrode surface may
occur after oxidation in order to achieve charge neutralization. Figure 4.18(b)
shows that the mass fluctuations became much smaller on repetitive cycling of
the potential until an almost constant mass situation is reached, with respect to
the potential cycle number. The EQCM data are consistent with the mechanism
proposed in eqns (4.24)-(4.28).
The identities of the species attached to the electrode surface and associated
with the second oxidation process at more positive potentials are unknown.
278
Electrode processes
(4
Potential (V vs FC/FC')
Time (s)
Potential (V vs FC/FC+)
I
0
I
I
200
.
I
'
I
400
600
Time (s)
'
I
800
Fig. 4.18 Simultaneous voltammetric-EQCM
experiments
cis-Ru(d~bpy)~
(NCS)2 in acetone (0.1 M Bu4NPF6). 5-mm diameter Q C gold electrode,
v = 20 mV s-' , vertical dashed line indicates start of potential scan. (a) One scan, (b) successive
) mass, (- - -) current. Reproduced by courtesy: J. Electrochem. Soc. 146 (1999) 648.
scans. (
Copyright, The Electrochemical Society.
Bron and Hoke [41] investigated the adsorption of free thiocyanate onto gold
electrodes. They found a strong affinity of thiocyanate towards gold electrodes
over a very wide potential range. The thiocyanate ions are adsorbed mainly
through the sulfur atom, although an observed potential dependence was related
to whether the adsorption occurred via the nitrogen or the sulfur. In case of
cis-R~(dcbpy)~
(NCS)2,the thiocyanate ligands are coordinated via the nitrogen,
which means that the sulfur atom is available for surface interactions. Alternatively, electrode-surface attachment via the carboxylate group could occur as
is oostulated at semiconductor titania electrodes [34,42,43].
. O n 'soft' metal
electrodes, sulfur-metal interactions probably are more dommant.
I
Voltammetric studies on polypyn'dyl ruthenium
I
-0.2
0.0
I
I
I
I
I
I
0.2
0.4
0.6
Potential (V vs FC/FC+)
I
I
0.8
I
279
I
1.0
9 Effect of addition of SS-bpy at 20°C on the voltammetry of 1.5mM
cis-R~(d~bpy)2(NCS)2
in acetone (0.1M Bu4NPF6) at a 1-mm diameter pt macrodisc electrode,
scan rate of 100 mV s-l. (- - -) without electrode modifier, (-)
after addition of 2 rnM SS-bpy.
~eproducedby courtesy: J. Electrochem. Soi 146 (1999) 648. Copyright, The Electrochemical
Society.
Free S C N added as Bu4NSCN was found to be irreversibly oxidized at
a potential of 0.18V, 0.12V, and 0.23V versus Fc/Fc+ under conditions of
cyclic voltammetry (scan rate 100mVs-') in acetone (0.1 M Bu4NPF6) at
platinum, GC, and gold macrodisc electrodes, respectively. Thus, the second
oxidation process is not associable with oxidation of SCN- released after oxidation of cis-Ru(dcb~y)~(NCS)~,
although it could be attributable to oxidation
of coordinated thiocyanate associated with surface-attached form of the dye.
owever, the process is not associated with oxidation of the thiocyanate ligand
coordinated to solution-soluble [ R ~ ( d c b p ~ ) ~ ( ~generated
C S ) ~ l ~via oxidation of cis-Ru(dcb~y)~(NCS)~,
since the second process is absent when the
u(d~bpy)~(~~
process
~ ) ~ is
] ~reversible
/'
and diffusion controlled.
Cyclic voltammetry in the presence ofsurfdce-activeSS-bpy
postulated mechanism is correct, then adding an electroinactive surfactant
solution may suppress surface attachment of cis-R~(dcbpy)~(NCS)~
and
thereby minimize the magnitude of the second oxidation process. Figure 4.19
s the effect on a voltammogram at a gold macrodisc electrode of adding the
electrode surface-active substance, 4,4'-bipyridyl disulfide (SS-bpy) [44]
in acetone. At equal or higher
to a 1.5 mM solution of cis-R~(dcbpy),(NCS)~
centrations of SS-bpy, the processes at positive potentials are eliminated,
ch implies that a rather high SS-bpy electrode modifier concentration is
necessary to compete for the adsorption sites."
l l ~ ~ - b palso
y modifies the response in an analogous manner at platinum electrodes, but
presumably because SS-bpy does not adsorb strongly on GC, the oxidation processes at
positive potentials remained when equivalent experiments were undertaken at this electrode
surface.
280
Electrode processes
)~
4.3 Voltammetvy of c i s - R ~ ( d c b p y ) ~ ( N CinS tetvahydvofcrvan,
acetonitvile, and dimethyljovmamide
Voltammetry in tetrahydrofuran was similar to that in acetone. However, due
to the limited positive potential range available in this solvent, detection of the
second oxidation process was difficult. Under conditions of cyclic voltammetry
at macrodisc electrodes in tetrahydrofuran, chemical reversibility was evident
at lower scan rates and higher cis-Ru(dcb~y)~(NCS)~
concentrations than in
acetone, presumably due to reduced surface effects. Determining the reversible
half-wave potential of 0.32 V versus FC/FC+ by steady-state voltammetry at
microdisc electrodes was straightforward, with a reversible, one one-electron
process being observed without evidence of adsorption. Half-wave potentials
determined by the RDE method gave slightly more positive values, which is
attributed to the presence of IR, (ohmic) drop resulting from the high resistance
of tetrahydrofuran.
The dye is not very soluble in acetonitrile, but even at low concentrations in
the loe5 M range evidence of surface activity is found under condition of cyclic
voltammetry at a macrodisc electrode (Fig. 4.20). Steady-state voltammograms
exhibit less evidence of surface interaction and give an E;,, value of 0.45 V
versus Fc/Fc+ in acetonitrile.
In dimethylformamide, transient cyclic voltammograms for oxidation of
cis-R~(dcbpy)~(NCS)~
require the use of very fast scan rates at platinum
microdisc electrodes (c.200 V s-l) before the onset of chemical reversibility
becomes evident (Fig. 4.21(a)). Increasing the scan rate beyond 200Vs-',
improved the chemical reversibility, but as in acetone, a peak current ratio
(1~11:~) of unity was never achieved in this solvent.
0.0
0.5
Potential (V vs FC/FC+)
Fig. 4.20 Background subtracted cyclic voltammogram obtained at 20°C with a scan rate of
M cis-R~(dcbpy)~(NCS)~
in acetonitrile (0.1 M Bu4NPF6), using a
5 0 0 m V ~ -for
~ 3.5 x
2-mm diameter Pt macrodisc electrode. Reproduced by courtesy: J. Electrochem. Soc. 146 (1999)
648. Copyright, The Electrochemical Society.
Voltammetric studies on polypyvidyl ruthenium
-0.5
0.0
0.5
Potential (V vs FC/FC+)
281
1.O
.21 Fast scan cyclic voltammetry of cis-R~(dcbpy)~(NCS)~
in DMF (0.1 M Bu4NPF6)
(a) 2.5mM cis-Ru(d~bpy)~(NCS)~;
(b) 4.0 mM c i i - R ~ ( d c b p y ) ~ ( N C with
S ) ~ 30 mM SS-bpy. 70-pm
diameter Pt microdisc electrode, background subtracted, v = 1000 V s-l. Reproduced by courtesy:
J. Electrochem. Soc. 146 (1999) 648. Copyright, The Electrochemical Society.
Addition of SS-bpy simplified the voltammograms obtained in dimethylformamide. In the presence of a sufficient concentration of this surface modifier,
a close to chemically and electrochemicallyreversible response was observed at a
scan rate at 100 V s-' (Fig. 4.21 (b)). Thus, provided a 10-fold excess of SS-bpy
was present, the (I:/I,'ed) ratio was close to unity over the scan rate range
of 100-4000 V s-'. From these fast scan rate cyclic voltammograms obtained
in the presence of a 10-fold concentration
from 2.5 mM cis-R~(dcbpy)~(NCS)~
excess of SS-bpy, the reversible half-wave potential for the oxidation process in
dimethylformamide was found to be 0.39 V versus Fc/Fc+.
L
L
.4 Conclusions related to the voltammetry ofsuface-active
u (d~bpy)( N C S )
reversible potential of the c i s - [ ~ u ( d c b ~ ~ redox
) ~ ( ~couple
~ ~ )may
~ ] ~
etermined in a variety of solvents using different electrochemical techniques
and electrode materials provided care is taken to minimize the interaction of
is system with the electrode surface. Reactant and product interaction with
e electrode surface cause departures from the mass-transport-controlled oxidation process for this compound under commonly used conditions of cyclic
voltammetry at a macrodisc electrode. Surface effects are stronger on GC and
gold than on platinum electrodes. Steady-state microdisc electrode and RDE
experiments provide close to chemically and electrochemically reversible masstransport-controlled voltammograms in acetone and low concentrations of the
dye and fast scan rates minimize the influence of the surface-based effects under
282
Electrode processes
the transient conditions of cyclic voltammetry. In dimethylformamide, scan
rates greater than 100 V s-I are needed to observe a significant level of chemical
reversibility, whereas only moderate scan rates were required in other solvents.
Addition of the electroinactive surfactant, SS-bpy, may also minimize surfacebased effects. The studies presented on the [ R U ( ~ C ~ ~ ~ ) ~ ( N Cprocess
S)~]~+'~+
have been included in this chapter to further emphasize the fact that use of a wide
range of techniques is generally needed to unravel mechanistic complexities
frequently associated with electrode processes.
5
Stripping voltammetry
Stripping voltammetry constitutes a class of techniques where electroactive
material is deliberately accumulated from the solution phase onto a solid electrode or into a liquid mercury electrode. After this pre-concentration stage,
the material is stripped back into the solution phase. Very low concentrations of electroactive compounds may be determined via this method [45-501.
Essentially, all stripping techniques possess three main steps viz. deposition,
equilibration, and stripping.
1. Deposition ov accumulation step The deposition step usually involves the
electrolytic or adsorptive deposition of a chemical species onto an electrode
surface at a constant D C potential. When metal ions are determined by anodic
stripping voltammetry [45-471 at a hanging mercury drop electrode, a sufficiently negative potential is applied to the working electrode to cause the metal
ion of interest to be reduced to the metal, which, in many cases, forms an
amalgam with the mercury electrode (e.g. Cd, Pb, T1 and Sb). In adsorptive
stripping voltammetry [45,46], a metal complex is accumulated at the electrode
surface by adsorption (e.g. Ni, Co, Sb, Ge). In stripping analysis, the deposition
step is usually facilitated by convective transport of the analyte to the surface of
the working electrode. This can be achieved by rotation of the electrode, by
stirring the solution or by flowing the solution over the electrode (Section 9 in
Chapter 2).
2. Equilibration step When the deposition step occurs under convective conditions, a quiet time usually follows this step in order to enable the electrode to
return to a quiescent state. This period is usually in the range of 10-30 s and is
called the equilibration step.
3. Stripping step In anodic stripping voltammetry, the stripping step is
achieved by the application of a voltage applied in the direction of positive
potential which, therefore, causes the metal or metal in the amalgam to be oxidized back to the solution-soluble metal-ion state. During the potential scan,
the accumulated metal is stripped from the surface, yielding a peak height for
each analyte, which is proportional to concentration. Ideally, the peak current
is linearly proportional to the concentration of the analyte in the bulk solution
and to the deposition time. The different stages associated with D C stripping
voltammetry in the linear potential sweep form are described in Fig. 4.22. In
Stripping voltammetry
283
3
g
0
Pi
Initial or
deposition potential
*-
4
Deposition
time
Quiet
time
2 Schematic diagram of the waveform used in DC linear potential sweep stripping
voltammetry.
adsorptive stripping voltammetry, the stripping step generally involves reduction
of the adsorbed metal complex by applying a negative direction potential scan.
g voltammetry represents one of the most widely used voltammetques for the determination of trace metal concentrations. However,
related to the analytical methodology [45-511 are not the focus of
. Rather, only the nuances of the electrode processes used in two
voltammetric methods are considered, to highlight how the same
established in other contexts also are valid in this important field of
ical endeavour. As conveyed by the description of the technique
ous paragraph, stripping voltammetry represents a combination
hase processes and processes where material to be determined is
then stripped from an electrode suliace. The first example to be
in Section 5.1 represents the electrochemical accumulation of metamalgams with thin films or droplets of mercury plated onto an
ce in a technique referred to as thin-film anodic stripping voltamsecond example, accumulation of material occurs by adsorption
lex onto a hanging mercury drop electrode which is followed
tated by electrochemical reduction in a technique referred to
ripping voltammetry. There are almost an infinite number of
trochemical stripping techniques [511, but the major features
electrode processes utilized in this field is conveyed by consideration of
two very widely used forms of stripping voltammetry.
.I Anodic stripping voltammetry with thin-jilm mercury electrodes12
n the anodic stripping voltammetric method for determining a metal ion Mn+
in aqueous solution, reduction at a mercury electrode for a carefully timed
12~daptedwith permission from Anal. Chem. 69 (1997) 2673. Copyright, American
Chemical Society.
284
Electrode processes
Table 4.4 Assignment of subscripts
used in the theory presented for
anodic stripping voltammetry
-
Subscript
Solute
Solvent
period, Atred,produces a metal that amalgamates
If the volume, V, of the mercury electrode is small, the average concentration of
the metal in the mercury phase soon exceeds manyfold the bulk concentration cp
of metal ions in the aqueous phase.13 It is this initial pre-concentration or plating
stage (see Fig. 4.22) that is responsible for the extreme sensitivity of stripping
voltammetry. As noted above, to foster a rapid transfer of metal into the mercury
phase, convection is sometimes applied during the pre-concentration stage of
stripping voltammetry, either by stirring the solution or rotating the electrode
[52]. However, convection is not needed when the electrode is a mercuryplated microdisc [53], or as in an array of such discs, as considered below,
because the efficiency of convergent diffusive transport to small inlaid discs [54]
is great enough to rival convective transport (also see Section 14 in Chapter 2).
The second of the three stages in traditional stripping voltammetry, the equilibration stage (see Fig. 4.22), occurs after the period, Atred,of reduction. This
stage, of duration Atwait,allows any convection to subside. If the thickness, 1,
of the mercury layer is small (compared with ( & ~ t , , ~ , ) ' / ~where
,
D2 is the
diffusion coefficient of the metal in mercury), as is the case with thin-film
mercury electrodes, the pre-equilibrium or inactive stage14 also allows total
homogenization of the amalgam to a uniform concentration c;.
In the stripping stage of anodic stripping voltammetry, the metal is electrochemically removed from the amalgam
often by applying a positive-going potential ramp of scan rate v to the electrode
(Fig. 4.22). Thus,
(4.31)
E = Einit v t
+
1 3 ~ o tthat
e the use of numerical subscripts to designate various solute species used in this
discussion is in accord with Table 4.4.
'"here
was no deliberate equilibration stage in the experiments considered in Section 5.1,
but the initial portion of the anodic scan provides the required brief period of inactivity.
Striping voltammetry
285
e potential, El,,, is the initial potential, and t represents time. As a
a time-dependent current, I, flows which, when plotted against E or t,
of M in the
lays a peak. The peak height I, is proportional to the amount VC,~
lgam phase at the beginning of the stripping phase and hence to cp. When
*-film mercury electrodes are used, but not necessarily otherwise [55-571,
is total, and as an alternative to measuring and interpreting the
the total content of metal M can be found by applying Faraday's
grating the voltammogram
00
I dt = QL = nFVc2b
notes the limiting voltammetric charge.
e recognized that, conceptually, carrying out chemical analysis via
dic stripping voltammetry involves two distinct steps:
1) determining the amount VC: of amalgamated metal from the voltammo2) relating VC; to the analyte concentration, c:, in the aqueous solution.
In practical chemical analysis, steps (1) and (2) are seldom disentangled; instead,
calibration or standard addition experiments are employed and the overall
proportionality IpO( C: is assumed and exploited.
u ' preparation of mercury thin-film electrodes
'in situ technique' of preparing a thin mercury film for anodic
mmetry, and the method used in data presented below, is available
ng liquid mercury
~ ~ : ( a q u e o u s ) f 2e-
-+ mHg(1iquid) m = 1 or 2
(4.33)
the pre-concentration phase. This is achieved by adding a soluble
ic (m = 2) or mercurous (m = 1) salt to the analyte solution to achieve
a concentration ck that is much greater than c,b. If the reduction potential is
action (4.29) proceeds where mass transport by diffusion is solely
ning (Section 7.1.1 in Chapter 2), then the deposited amount
is expected to be proportional to the bulk concentration c,b of the
Mn+ ion, to the charge n+ on the cation, and to some power p of its diffusion
coefficient. Accordingly
vc,"O( n@cp
(4.34)
VG
with no other M-specific terms being expected to enter this relationship. As
shown in Section 10.1 in Chapter 2, the power p equals uniq for microdisc
electrodes employed in experimental studies presented below, whereas p = 112
(Section 8.1 in Chapter 2) when co-deposition occurs by planar diffusion onto
286
Electrode processes
a macrodisc electrode, while p = 213 for hydrodynamic systems under laminar
flow conditions (Chapter 2, Table 2.3). Of course, similar proportionalities
apply to the
ion co-depositing according to reaction (4.33), but the lefthand term in eqn (4.34) is not then appropriate because mercury serves as the
solvent (Table 4.4). Instead, if the reasonable assumption that the atomic volume
of mercury in a dilute amalgam differs negligbly from that in the pure liquid
state is made, then the proportionalities
~~2
may be expected, where P H and
~ MHg are the density and atomic mass of
mercury respectively (PHg/hlHg= 67471 mol m-3 at 25°C). Note that the
concentration c,b
of amalgamated metal depends on the ratio (c,b/c,b) of aqueous
concentrations and is unaffected by plating duration.15
A complication that attends the metal-mercury co-deposition technique is
that, because most metals M of interest are considerably less noble than mercury,
the deposition of mercury continues unabated during the stripping stage. Thus
(as is apparent in figures shown later) the baseline of the stripping peak is not
zero, but occurs as a negative (i.e. reduction) current. When the electrode is a
single inlaid microdisc of radius re and H ~ (n~= +
2) is employed, this baseline
current is given (Section 10.1 in Chapter 2) by the relationship
Of course this constant baseline current must be subtracted from the measured
total current, I,,,,, to evaluate, I, the stripping current. For example the I term
in eqn (4.32) should be replaced by I,,,, - Ibase.
5.1.2 Comparison of macro- and microdisc mercury thinjilm electrodes
Individual or, even better, arrays of inlaid discs of carbon (Fig. 4.23) provide convenient substrates for co-deposition of thin layers of mercury suitable
for use in anodic stripping voltammetry. The size of a disc is characterized
by its radius r, or sudace area A = m e 2 . As described in Chapter 2, large
inlaid discs (macroelectrodes) may behave differently from small discs (microelectrodes). The distinction arises because diffusion to a 'large' disc electrode
is primarily linear (planar), whereas diffusion to a 'small' disc rapidly becomes
radial. Again, as noted in Chapter 2, the classification into 'large' or 'small'
electrodes is based on whether the disc radius is larger or smaller than the
'distance scale of the experiment'. The experimental distance scale during the
pre-concentration stage of anodic stripping voltammetry is (Dl A tred)'I2which
is much larger than the 3.5 pm radius of the inlaid discs considered experimentally below, so that these discs are clearly 'small7 during this stage. However,
1 5 ~ hconcentration
e
ratio applicable to experimental data reported below was chosen so
that the amalgam concentration never exceeds 0.3 atom per cent.
Stripping voltammetry
Current
collector
Carbon
fibres
287
Epoxy resin
. 4.23 The random assembly of carbon microdisc electrode or RAM@concept used in anodic
stripping voltammetric studies: (a) section view, (b) plan view. Provided by courtesy of Stephen
Fletcher, CSIRO Division of Minerals, Clayton, Victoria, Australia.
the situation is much more equivocal in the case of the stripping stage of the
experiments. Then, the experimental distance scale during the transit of the
stripping peak can be considered to be ( 2 . 9 4 ~ ~ ~ ~ l n n For
~ v the
) " values
~.
T = 298K, Dl = 8 x 1 0 - ' ~ m ~ s - ln, = 2, and v = 0.5Vs-', which correspond closely to the stripping conditions employed in experiments described
below, the experimental distance scale evaluates to 4 pm. Coincidentally, this
is very close to the 3.5 pm radius of experimentally used microdisc electrodes.
n a typical voltammetric experiment, hybrid behaviour midway between that
expected for macro- and microelectrodes [58] could therefore be expected (also
see Section 10 in Chapter 2). However, very atypical circumstances apply in
stripping from a thin film. During radial (quasi-spherical) diffusion to a small
inlaid disc, the current density is markedly non-uniform, most of the current
t
strippingfrom a
flowing through the perimetric region of the disc. ~ u during
thin mercury layer, the uniform distribution of the metal M solute will enforce
an almost uniform current density, as is the rule at 'large' discs. Enhanced current
density at the electrode edge is impossible in thin-film stripping. Accordingly,
the theoretical treatment can be based on linear (planar) diffusion.
Treatment of the mercury film as 'thin', as is the case in the present consideration, means that the thickness, 1, of the mercury layer is small in comparison
with the experimental distance scale associated with stripping metal M out of
the film. Thus
where the 6-pm length given in relationship (4.37) is based on D2 = 1.5 x
m2 ssl in addition to the parameter values cited above. The mercury volume deposited on a disc of radius r = 3.5 pm in experiments to be considered
below was calculated on the basis of expression (4.35) to be about 1.3x 10-l7 m3,
288
Electrode processes
which leads to a calculated average thickness 1 = v / n r e 2 = 3 x
m, so that
inequality (4.37) is amply satisfied. In the interest of realism, V I A instead of 1
is used as a measure of the mercury thickness, because [56,57] electrodeposited
mercury does not form a uniform thin layer, some bare carbon remaining long
after a uniform laver might be exoected.
Even after pre-concentration, the Faradaic currents derived from an electrode
of 3.5 pm radius in micromolar to nanomolar or even lower analyte concentration are very small and not easily measured with accuracy. To remedy this,
arrays of identical inlaid disc electrodes (Fig. 4.23) rather than a single disc can
be employed [59]. In analysing the results of array experiments (see below),
all the discs are regarded as functioning identically and independently, without
mutual interference by overlap of depletion zones. Consequently, the symbols
A, V , and I are used indiscriminately, to refer either to the area of a single
disc, the volume of mercury deposited on it, and the stripping current from
that disc, or to the total area, total mercury volume, and total stripping current
of all discs.
J
I
U
5.2 Theory for a reversible process
As in solution phase and studies with thin films described in Chapter 2, electrode
processes associated with stripping voltammetry may be classified as reversible,
quasi-reversible, or irreversible. Only the reversible case is considered in this
book. The theory of thin-layer anodic stripping voltammetry for a reversible
reaction was first developed by De Vries and Van Dalen [60] some thirty five
years ago. They adopted a model based on principles similar to those described
above, derived an appropriate integral equation, solved it numerically, and
reported expressions for the peak current
for the peak potential in terms of the half-wave potential
and for the width of the peak at half-height
The development of an equation that can be used to describe the shape of the
entire stripping voltammogram and which necessarily leads to these relationships
is taken from reference [61] and utilizes a range of concepts introduced in
Chapter 2.
The thinness of the mercury film implies that the concentration c2 of the
metal in the amalgam is effectively uniform at any instant during the stripping.
I
I
Stripping voltammetry
289
araday's law applied to the reaction in eqn (4.30) relates the amount of metal
removed from the amalgam at any time t to the integral of the stripping current
up to that instant:
It
nF(~c;-Vc~)=
Idt=Q
(4.41)
and when this result is combined with eqn (4.32), the expression
Idt
VC2 =
is derived and relates the amount of metal at time t to the difference between
the limiting charge and the stripping charge up to time t.
With respect to the solution phase, there is a time-dependent concentration c; of Mn+ ions at the electrode surface. This may be related to the bulk
concentration of this ion by the equation [62]
when transport is solely by linear diffusion. Here, M is the semi-integral of the
current I over the time interval from zero to t (Section 11 in Chapter 2).
When stripping voltammetry is being used for trace analysis, with massive preconcentration, cp is very small compared with
and is therefore generally
insignificant in comparison with ci . Accordingly,
ci
is a valid replacement for eqn (4.43).
In terms of its formal potential E ~ O , the Nernst equation for eqn (4.29)
(assumed reversible under the linear-scan stripping regime) is
After substitution from eqns (4.42) and (4.45), eqn (4.46) may be rearranged to
which shows that a graph of the logarithm of M I ( & - Q) versus potential should give a straight line of slope n F I R T and an intercept potential of
290
Electvode processes
Fig. 4.24 'Log plot' analysis of mercury thin-film stripping voltammograms based on eqn (4.47).
Reproduced by courtesy: Anal. Chem. 69 (1997) 2673. Copyright, American Chemical Society.
v2]
E: - (RT/2nF) ln ( A ~ D ~ / or EO
f - (RTInF) l n { ~ ~ : V}
' ~ /as illustrated
in Fig. 4.24. Such a linear construction to give a 'log-plot' is analogous to the
analysis of reversible steady-state or other forms of sigmoidal-shaped voltammograms by graphing the logarithm of (Ili, - I ) / I versus potential (Section 9.4
in Chapter 2).
Note that neither the bulk amalgam concentration c,b nor the scan rate v
appears in eqn (4.47), so that plotted data derived from experiments at several
scan rates or analyte concentrations should overlie each other.
It is convenient to define a reference potential for each stripping voltammogram by
(4.48)
It will transpire that the peak of the stripping voltammogram generally lies
within 1mV of this new parameter. The Nernst relationship [63] may be
reformulated as
in terms of this reference potential. Note that, unlike the intercept potential
shown in Fig. 4.24, E* does depend on the scan rate V .
Before constructing an equation for the stripping voltammogram shown in
Fig. 4.25, it is advantageous to adopt dimensionless variables. In linear sweep
anodic stripping voltammetry there is a linear relationship, namely eqn (4.31),
between potential E and time t . Both of these variables will be replaced by a
dimensionless counterpart
L
Stripping voltammetry
291
. 4.25 Predicted features of a thin-film anodic stripping voltammogram. The ordinate is the nor~ V~C ~ Vand the abscissa is an undirnensionalized potential scale
malized stripping current L = R T I / F~
f = ( n F / R T ) ( E- E*), where E* is defined in eqn (4.48).The peak occurs at ( = -0.05545. The
area beneath the curve is unity. Reproduced by courtesy: Anal. Chem. 69 (1997) 2673. Copyright,
American Chemical Society.
where
nF
co = -ERT
-
(4.51)
lnlt
e 6 variable is an undimensionalized potential scale with its zero at the refnce potential. The Eo term is a negative constant, equal to the value of 6 at
the instant when the stripping scan started. The current I can also be replaced
y a normalized dimensionless counterpart
ubstituting this definition into eqns (4.42) and (4.45), and simultaneously
replacing the time variable in those equations by 6, gives
and
respectively. Finally, eqns (4.53), (4.54), and (4.50) are inserted into the Nernst
relationship (4.49). The simple result
292
Electrode processes
emerges. In dimensionless form, this is the equation that describes the shape of
the stripping peak.
The next step in order to obtain an analytical expression to describe the
stripping voltammogram in Fig. 4.25 is to replace the lower limit toin eqn (4.55)
by -w.
This is equivalent to asserting that there would be no important change in
the voltammogram if the scan had started at any potential more negative than
the actual initial potential; experimental initial potentials are, of course, always
selected on such a basis. The valuable property [63] that semi-integration of
exp{bx} with respect to x, with a lower limit of -w, gives (114) exp{bx) for
any constant positive value of b, and the known limit
now permit the summation
to be recognized.
The coordinates of the peak in the Fig. 4.25 curve are
(c,,
ip) = (-0.0555,
0.29697)
(4.59)
while the coordinates of the half-peak points are
(cp12,
lpI2) = (-
1.6416,O.14848) and (1.2990,O.14848)
(4.60)
That the area under this peaked curve in Fig. 4.25 is exactly unity is guaranteed
by the eqn (4.57) limit as well as by purely chemical arguments.
O n returning to experimental variables, the equation of the stripping
voltammogram is seen to be
peak potential (Ep),and peak half-width ( Wl12)being
with the peak current (Ip),
Stripping voltammetry
293
and
RT
Wl12 = 2.9407nF
Observe that I, and E, depend on V , but Wl12does not. Furthermore, these values are in excellent agreement with those of De Vries and Van Dalen, reported
in eqns (4.38)-(4.40). Notice that the appearance of different diffusion coefficient~in eqns (4.63) and (4.39) arises because the Dutch workers chose to report
[ho]the peak potential relative to the polarographic half-wave potential,16equal
(RTI2nF) ln{D1/ D2}, instead of relative to the formal potential E ~ O as
to, E!
is the case in eqn (4.63).
For a divalent analyte ion at 25"C, the voltammetric peak is predicted to
lie 0.71 mV negative of the reference potential E* and the peak width at halft should be 37.78 mV. Observe the absence of diffusion coefficients from
eqn (4.62), unlike the case with solution-phase voltammetry, and note the
v re dictions that the peak current is proportional to the square of the electron
er and to scan rate v. Analysts tend to prefer measurements of peak height
eak area, but there is much to be said in favour of calculating C$ from QLvia
.32), rather than via the proportionality in eqn (4.62), because the charge
is more fundamentally related to the amalgam content and because integration
nishes effects of random noise. Obedience of experimental results to the
cted relationship
+
between peak area and peak height provides a powerful test of the applicability
or otherwise of the theoretical model to an experimental system being studied.
henever an electrode process involves interaction with the surface, the
itional inherent complexity relative to completely solution-phase studalways leads to the likelihood of poorer agreement between theory and
ment (see Chapter 5). In the case of thin-film stripping voltammetry,
a comparison of data obtained in reference [61] with the theory presented
above gives an idea of the level of agreement that is achievable in this area
of voltammetry.
Comparison of experimental results and theory
e thin mercury films for experimental studies were prepared by in situ deposition at - 1.100 V versus Ag/AgCl for 100 s onto an assembly of carbon discs
randomly inlaid into the end face of a 10-mm diameter cylinder of epoxy resin
olarography is the technique ofvoltammetry at a dropping mercury electrode developed
by the Nobel prize winner Heyrovsky. The theory of this technique is described in the
book Princbles of Polarography by J . Heyrovsky and J. Kuta, Academic Press, New York,
1966. Also see Table 1.1, of Chapter 1 and references [46,47] of this chapter for more
details on the history and practice of the polarographic method.
294
Electrode processes
using a RAM@electrode (Fig. 4.23), the carbon discs being the sectioned ends of
parallel 7-pm diameter pyrolytic graphite fibres. Based on calibration tests, there
were believed to be 653 separate inlaid carbon discs, electrically commoned, in
the electrode used. Thus the total electrode area was A = 2.5 x lov8m2.
At the end of the 100 s of plating, the electrode potential was immediately
ramped positively from the initial potential of - 1.100 V versus Ag/AgCl at a
rate of either 0.2, 0.4, 0.6, 0.8, or 1 V s-' until a final potential of 0 V versus
Ag/AgCl was attained. At the lowest scan rate, there would have been an
additional 3 or 4 s of plating time early in the stripping stage, but no correction
was made for the consequential enhancement of c:.
All experiments were conducted at 21°C in aqueous solutions containing 100 mM K N 0 3 (pH 1.7). Analyte solutions were 0.38 mM H ~ ~ + ,
1.OO pM pb2+, and 1.OO pM cd2+ and tp = 100 s (data obtained under other
conditions are available in reference [61]). A typical set of voltammograms
obtained under these conditions for the processes in eqns (4.66) and (4.67) is
shown in Fig. 4.26.
Fig. 4.26 Typical stripping voltammograms of Cd (first peak) and Pb (second peak) from thin
amalgam films on an array of microdisc electrodes. Scan rates: (lower curves) 0.4, (upper curves)
0.8 V s-l. The films were prepared by plating for 100 s from a solution containing 1.OO pM each
of Cd2+ and pb2+ with 0.38 m~ H ~ ~ Reproduced
+ .
by courtesy: Anal. Chew. 69 (1997) 2673.
Copyright, American Chemical Society.
Table 4.5 Width of experimental stripping
peaks at half-heighta from thin amalgam films
Cd (amalgam)
Pb (amalgam)
0.20
0.40
0.60
0.80
1.OO
46.8
47.2
47.2
47.7
47.8
43.7
43.3
41.4
39.3
41.7
Ave f SD
47.3 r f 0.4
41.9 f 1.7
'The theoretically predicted width is 37.3 mV; data
taken from reference [61]; 0.38 m~ H~"+, 1.OO pm pb2+,
1.00 pm cd2+.
For each metal at each scan rate, the stripping current I versus potential E
data was analysed in the following ways:
peak half-width was measured and compared with the theoretical value
7.3 mV given by eqn (4.64) for n = 2, T = 294 K. The experimental
are assembled in Table 4.5.
current data were semi-integrated with respect to time, so generating
current data were integrated with respect to time, producing a file of
a1 entries in the latter file, a value of & was identified. These
ta are presented in Table 4.6, together with $ values from item 6. Also
~stedin this tabulation are values of the quantity QLv/Ipwhich, according
eqn (4.65) should adopt the value 42.7 mV at 21°C when n = 2.
weighted linear regression corresponding to a log plot
rating three distinct pieces of information: (a) the slope,
to be compared with the predicted value nF/RT which equals 7 8 . 9 ~ - I
for n = 2, T = 294K; (b) the intercept potential; and (c) the correlation
coefficient, which provides a measure of the linearity of the ln{M/(QL- Q)}
ersus E relationship. Examples of experimental 'log plots7 are shown in
ig. 4.27 and the entire regression results are compiled in Table 4.7.
(6) The peak current and peak potential were identified. The former values,
after baseline correction, are reported in Table 4.6. The Epdata are listed
in Table 4.8 and compared with the reference potential E*, which is calculated from the intercept potential, described in (5b), by the addition of
the quantity
296
Electrode processes
Table 4.6 Limiting charge and peak current stripping dataafrom
thin amalgam films
tn
(3
100
100
100
100
100
v
(vs-')
0.20
0.40
0.60
0.80
1.00
Ave f SD
C d (amal)
Pb (amal)
Q
Ip
Ip
(PA)
(&vIIP)
(mV)
Q
(nC)
(nc)
(PA)
(QdIp)
(mV)
117
111
127
144
140
0.455
0.854
1.476
2.199
2.657
51.4
52.0
51.6
52.4
52.7
206
201
191
198
204
0.85
1.671
2.511
3.601
4.394
48.5
48.1
45.6
44.0
46.4
52.0 f 0.5
46.5 f 1.8
"Theoretical value of Qcv/Ip according to the model should equal 42.7 mV; data
.OO, pm pb2+, 1.00 pm c d 2 + .
taken from reference [61]; 0.38 mM H ~ ~1+
Fig. 4.27 'Log plots' for the Cd and Pb stripping peaks from thin amalgam films with scan rates of
0.40 V s-' (o) and 0.80 V s-' (A). The lines are the linear regression lines, determined from the points
shown. Other conditions as for Fig. 4.26. Reproduced by courtesy: Anal, Chern. 69 (1997) 2673.
Copyright, American Chemical Society.
Since the theoretical difference between these potentials is < l mV, close
agreement is expected.
(7) The molar amount of analyte metal in the amalgam, VC;, was measured by
three methods; (a) from Ip, via eqn (4.62); (b) from &, via eqn (4.32); and
(c) from the equation
Stripping voltammetry
297
le 4.7 Results of 'log plot' linear regression analysisa of anodic
stripping voltammetric data from thin amalgam films (eqn 4.47)
v
(V sf1)
Cd (amal)
slope
intercept
- 9
(V)
Pb (amal)
con. coef
slope
- 9
intercept
(V)
corr. coef
'Theoretically predicted slope is 7 8 . 9 ~ - l ; data taken from reference [61];
0.38 mM H ~ ~1.O
+O,pm pb2+, 1.OO pm ~ d " .
Table 4.8 Measured peak potentials, Ep,
compared with reference potentials, E*'
obtained from stripping voltammograms of
thin-film amalgams
'Calculated from the intercept potentials of the 'log
plots' (eqn 4.47); data taken from reference [61];
0.38 mM H ~ 1.O~O pm
+ pb2+,
~ 1-00prn cd2+.
which arises by combination of eqns (4.32), (4.64), and (4.65). The results
of these alternative determinations are assembled in Table 4.9. Reasonable
agreement is achieved. However, the correct value of vC,~ is unknown and
therefore the accuracy of the determinations of the amount of metal by the
rival methods cannot be assessed.
Table 4.5 shows that the stripping peaks are significantly wider than theory
redicts, the width at half-height exceeding prediction by (27 f 1) per cent for
cadmium (amalgam) and (12 f 5) per cent for lead (amalgam).It is commonly
found that voltammetric stripping peaks from thin amalgams of various varieties
exceed in width the value that theory predicts, though the excess peak width
for the cadmium case suggests that the model is imperfectly obeyed.
298
Electrode processes
Table 4.9 Comparison of three methods of calculating the amount of
amalgamated metal (pmol)" associated with stripping voltammetry from
thin amalgam films
v (V S - I )
Cd(ama1)
Pb (amal)
from Ip
from &
from
eqn (4.66)
from Ip
from C&
from
eqn (4.66)
0.20
0.40
0.60
0.80
1-00
0.503
0.472
0.544
0.608
0.587
0.606
0.575
0.658
0.746
0.725
0.632
0.598
0.689
0.778
0.754
0.940
0.924
0.925
0.995
0.972
1.067
1.042
0.990
1.026
1.057
1.100
1.073
1.028
1.050
1.087
Av f SD
0.543f
0.057
0.662f
0.074
0.690f
0.077
0.951 f
0.031
1.036&
0.030
1.068f
0.029
'Data taken from reference [61]; 0.38 mM H ~ ~1.+
OO,ym pb2+, 1.OO ym cd2+.
Fig. 4.28 Dependence of peak current on scan rate for the stripping of cadmium and lead from
thin-film amalgams. Reproduced by courtesy: Anal. Chem. 69 (1997) 2673. Copyright, American
Chemical Society.
Figure 4.28 demonstrates the expected linear dependence of peak height Ip
on scan rate v, for both amalgams. However, if eqn (4.62) is to be believed,
this diagram implies that the concentration of cadmium in the amalgam is only
60 per cent of the lead concentration. Table 4.6 reveals that both peak areas &
are essentially independent of v, as expected, and again the implication is that,
in the amalgam, the cadmium concentration is markedly lower than the lead
concentration, this time by a factor of 0.64. Values of (&v/I,), which average
(52.0 f.0.5) mV for cadmium amalgam and (46.5 & 1.9)mV for lead amalgam,
Stripping voltammetry
299
+
also are listed in this tabulation. (Here and elsewhere a number preceded by
is a standard deviation, SD.) Departure from the theoretical value of 42.7 mV is
modest for lead, suggesting that the model is acceptable for this metal, but the
cent disagreement for cadmium suggests that this process has additional
xity. Nevertheless note that, as for peak widths, the reproducibility of the
results is markedly better than for Pb.
linearity of the 'log plots' is excellent, as evidenced by Fig. 4.27 and
orrelation coefficients in Table 4.7. As is commonplace in other varieties of reversible 'log plot7analysis [64,65], the slopes of the ln{M/(QL - Q)}
versus E plots, which average (67.1 f0.8) V-' for cadmium amalgam and
(72.7 rt 1.4) V-' for lead amalgam, are smaller than the predicted Nernstian
of 78.9 V-'. The intercept potentials are predicted to be independent of
rate, but Fig. 4.27 suggests that this prediction is not realized in practice.
owever, the data in Table 4.7 reveal no discernable trend with scan rate in the
intercept potential. Rather, it appears that the potential scale is not reproducible
experiment to experiment to better than a few millivolts.
n contrast to the intercept potential, the potential of the stripping peak is
icted to depend on the scan rate, shifting positively with increasing v. There
convincing evidence of such a trend in Table 4.8. However, the theoretical
for the five-fold range of scan rates encompassed by these experiments is
10 mV and this small trend has probably become 'buried7 in the scatter.
n specific potentials are compared within a single experiment, the scatter is
ess intrusive. Recall that theory predicts that the (Ep- E*)potential difference
d be only -0.7 mV. The measured differences are (+4.0 0.9) mV for
+
. 4.29 Comparison of the theoretical shape of a linear-scan stripping voltammogram (full line)
and the experimental current-voltage curve for stripping of lead from thin amalgam films at a sweep
rate of 0.800V s-l. Other conditions as for Fig. 4.26. Reproduced by courtesy: Anal. Chem. 69
(1997) 2673. Copyright, American Chemical Society.
300
Electrode processes
Cd(ama1gam)and (-2.0 i3.7) mV for Pb(ama1garn).Once more, the lead data fir,
the model better than is the case with cadmium, but is less reproducible.
Figure 4.29 contains a comparison of the theoretical shape of a stripping
voltammogram, as predicted by eqn (4.61), with an experimentally obtained
lead amalgam curve. The curves have not been explicitly fitted at any point.
Instead, the VC;parameter, needed for eqn (4.61), was calculated via eqn (4.32)
from the experimentally measured value of &. Similarly, the needed E* reference potential was calculated by adding the quantity in eqn (4.66) to the
measured intercept potential of a graph ofln{M/(& - Q)}versus E (Table 4.8).
For the stripping of cadmium amalgam, agreement is significantly worse. The
stripping peak for Cd(amalgam)is lower and wider than expected on the basis of
the theory for a reversible process. If de-amalgamation reaction is not therrnodynamically reversible, it can be predicted that an irreversible or quasi-reversible
stripping peak would have a lower peak and a broader half-peak. Thus, it is
tempting to attribute the non-conformity of the cadmium peak to irreversible
stripping behaviour. However, this cannot be the whole story, because the
stripped charge for cadmium is only 64 per cent of the corresponding QL for
lead stripped from amalgam under similar conditions. The magnitude of the
stripped charge is not affected by the degree of reversibility of the electrode
reaction.
There are at least four conceivable explanations for the lower recovery of cadmium than lead. First, despite their equal concentrations in the bulk aqueous
solution, less cadmium might have been plated than lead. T o some extent, this
explanation is credible, because cd2+(aqueous)has a smaller diffusion coefficient
[66] than pb2+(aqueous).
However, DCp+/DPb2+has a value of about 0.83 and so,
in light of the discussion surrounding eqn (4.34), this cannot explain more than
a minor portion of the observed magnitude of the effect. A second possibility
is that an intermetallic compound (examples of which are well known [45] in
the case of some pairs of metals, such as Cu/Zn) might be formed within the
amalgam between Cd and Pb, and stripped with the lead. In that eventuality,
some of the & that has been attributed to lead should be ascribed to cadmium.
However, experimental lead stripping peaks were found to be independent of
whether or not the cadmium was co-deposited, lending no support to the possibility of intermetallic interference. Third, as postulated by Batley and Florence
[55], metals may plate on islands of bare carbon, from which it may not readily
be stripped. For the carbon-fibre-array electrode used in these studies, no cadmium peak was observed when mercury was absent, either because cadmium
was never plated onto the carbon substrate or because the plated cadmium could
not be stripped, whereas lead stripping peaks were observed in the absence of
mercury. A fourth possibility is that the solubility of cadmium in mercury may
have been exceeded in the surface layers during plating, resulting in a smaller
cadmium content than the more soluble lead. However, if this were the case,
the problem would be alleviated by using lower concentrations of c d 2 + in the
plating solution. There is no evidence of such an effect as graphs of I, and QI,
are linearly related to the aqueous concentration for both c d 2 + and pb2+.
Stripping vo~tammetry
30 1
1n summary, there are several ways in which the model of linear scan stripping
of &&-mercury films appears to be deficient. However, agreement between
the theoretical predictions of the model and the experimental results for lead
is regarded as being satisfactory and consistent with the level expected when
surface-basedprocesses form an integral component of an electrode process.
chanism associated with the adsoytive stripping voltammetry of
cobalt (and nickel) dimethylglyoxime complexes at mercury electrode^'^
5.
More than fifty years ago, Stromberg and Zelyanskaya [67] observed a large
increase in the polarographic18limiting current for the reduction of cobalt ions
ropping mercury electrode in an ammonia buffer solution when dimethylime was added to the solution. Furthermore, the unusually large adsorptive
ing current observed for reduction of cobalt in the presence of dimethylglyoxime at hanging mercury drop and mercury thin-film electrodes is now
widely utilized for the determination of trace concentrations of this element by
the adsorptive stripping voltammetric methods19 [45,46,68].
ecently, two papers have been published [69,70] which suggest that ligand
etal-based reduction occurs in an overall ten-electron reduction process
cobalt dimethylglyoxime complexes are reduced at mercury electrodes
and that it is this feature that gves rise to enhanced currents relative to a twoelectron reduction of cobalt ion to the metal that occurs in most other media.
Prior to that, considerable controversy existed concerning the nature of the
mechanism associated with reduction of the cobalt dimethylglyoxime complex.
ven though a consensus is now being reached [69,70] on many of the details,
it is useful to review the wide range of schemes proposed for this very important
analvtical mechanism over the Dast fiftv or so vears to illustrate how difficult
a task it is to even qualitatively establish the overall reaction process, let alone
quantitative details when a series of solution phase and surface processes are
esent. For convenience in writing the various reaction schemes, the symbol
o(II)(dmgH), will be used to denote the complexes (Fig. 4.30(a)), M = Co(I1)
formed by coordination of two deprotonated molecules of dimethylglyoxime
dmgH,). Brief mention of the closely related nickel analogue referred to as
i(dmgH,), (Fig. 4.30(a), M = Ni(I1)) also will be made.
J
L
J
J
l7Lidapted with permission from Anal. Chem. 70 (1998) 1312. Copyright, American
Chemical Society.
18See footnote 16 where the term polarography or voltammetry at a dropping mercury
electrode is defined.
191n the cobalt adsorptive stripping method at the hanging mercury drop electrode,
dimethylglyoxime is present in a significant concentration excess over cobalt and a
C ~ ( d r n g Hcomplex
)~
is formed in the solution phase. At the initial potential adsorption
of the cobalt dimethylglyoxime complex then occurs on the electrode surface during the
accumulation stage. Finally, the complex is reduced during the stripping component of the
experiment.
302
Electrode processes
Fig. 4.30 Structural representation of: (a) cobalt and nickel dimethylglyoxime complexes
M(drngH)2; (b) cobalt and nickel dimethylglyoximate analogues, M(CsdoH)2, where M is Co(11)
or Ni(1I).
Many of the various mechanisms involved in explaining the voltammetry
associated with the reduction of cobalt in the presence of dimethylglyoxime
have been summarized in reference [70]. According to the exclusively adsorptive
mechanism favoured at one stage, the reaction scheme may be summarized as
follows:
c o 2 +(solution)
+ 2dmgH2(solution) --+[Co(11)(dmgH)2](solution) + 2 ~ + ( s o l u t i o n )
(4.70)
[CO(11)(dmgH)2](solution) --+ [CO(11)(dmgH)2] (adsorbed)
[Co(II)(dmgH)2](adsorbed) -I- 2e- -k 2 ~ +
+ Co (0)
(metal)
(4.71)
+ [dmgH2](desorbed)
(4.72)
According to this scheme, when the mercury electrode reaches the required
potential, the cobalt dimethylgloxime complex is reduced in the adsorbed state
and dmgH2 is released after reduction of the absorbed complex.
In other studies, a slightly different mechanism was proposed in which a
reactive cobalt(1) intermediate is generated.
[C0 (1I)(dmgH)2] (solution) --+ [Co (II) (dmgH)2] (adsorbed)
[Co(II)(dmgH)2](adsorbed)+e--?\[Co(I)(dmgH)2]-(ad~~rbed)
[Co(I)(dmgH)2]- (adsorbed) f H+
(4.73)
(4.74)
+ e- --+ CO(O)(metal) + [dmgH2](desorbed)
1-(desorbed)
(4.75)
303
Stripping voltammetry
Compound X (identity unknown) was said to be produced via a chemical
decomposition reaction.
In other mechanisms proposed, reduction of both the central Co(I1) ion and
the ligand [dmgH]- occurs as in eqn (4.76)
[Co(II)(dmgH)2](adsorbed)
+ x e + yH+ --+ Co (0)(metal) + [dmglred (4.76)
where: x = (10-18) electrons and [dmglredis the product(s) of reduction of
dimethylglyoxime.
According to this mechanism, the coordinated ligand may be partially reduced
to 2,3-di(hydroxy1amino)butane
H ~ c / ~N-OH
\
or totally to 2,3-diaminobutane
+ 8 ~ +' Be-
I
+yC\
N-OH
I
+ 2H20
H ~ c / ~ ~ L
NH2
(4.78)
A range of catalytic schemes were also postulated. The cyclic catalytic
reduction mechanism may be summarized as follows
co2+(solution)
+ 2dmgH2(solution) --+ [Co(dmgH)2](solution) + 2~+(solution)
(4.79)
[CO(11)( d ~ x g H )(solution)
~]
-+ [CO(11)(dmgH)2](adsorbed)
[Co(II)
(adsorbed)
(4.80)
+ xe- + YH+(solution) --+ co2+(solution)
+2[dmg]red
(4.81)
If this mechanism is correct, the Co2+ ions released after the reduction
rocess should then react with bulk dmgH2 and the coordination-reductioncoordination process proceed in cycles until all of the dmgH2 becomes
reduced.
The reduction of cobalt as well as nickel dimethylglyoxime complexes was
also suggested in a number of studies to involve catalytic hydrogen evolution,
304
Electvodepvocesses
according to the following kind of scheme
(metal) -I- [dmgH,] (adsorbed) ---+
c o (0)
[C o (0)
(dmgH)]- (adsorbed)
+ H+(solution)
[CO(0)
(dmgH)]- (adsorbed)
(4.83)
+ BH+ (solution)
{ [CO(0)
(dmgH)]- (adsorbed)HC
+ e -+
H, (gas)
where BH' is the source of the donor proton, for example, NH;.
As recommended in Chapter 2, establishment of the details of the cobalt (and
nickel) dimethylglyoxime reduction process at mercury electrodes requires the
use of a wide range of techniques and strategies with spectroscopic characterization of reaction pathways being mandatory. The inherent difficulty associated
with establishing mechanisms associated with surface-based processes is illustrated by the discrepancies and controversies surrounding the nature of an
electrode process that is widely employed in trace analysis [69,70].
5.4.1 Voltammetric reduction ofthe cobalt dimethylglyoxime system at
mercury electrodes in aqueous media
Studies on solutions containing only non-coordinated dimethylglyoxime (added
as dmgH,) provide useful reference data for the studies on the reduction of the
cobalt dimethylglyoxime complex. D C polarograms of a solution containing
1x
M free ligand in an aqueous 0.1 M ammonia buffer medium,,' show
one reduction wave with an El/, value of - 1.55 V versus Ag/AgCl when the
drop time is 0.6 s. Figure 4.3 1(a) shows a polarogram in 0.1 M ammonia for
reduction of dmgH, compared to that of the well-documented two-electron
cadmium reduction cd2+(solution) 2e- + Cd(ama1garn) at the same concentration of 1 x
M and with all other experimentally controlled parameters
being kept constant. From the data it is evident from the ratio of the diffusioncontrolled limiting currents, id(dmgH,)/ id(cd2+),that the number of electrons
(n) transferred during the course of reduction of dmgH, is significantly greater
than the known value of n = 2 for the cadmium ion, assuming
- that the diffusion
coefficients (D) of the cadmium ion and dmgH, are similar.21
Cyclic voltammetric experiments at a hanging mercury drop electrode show
that the reduction wave of the free ligand is situated on the shoulder of the
+
200.1M ammonium chloride/O. 1M ammonia solution.
2 1 ~ hdiffusion-controlled
e
limiting current, Id, in a DC polarogram is given by the Ilkovic
equation, which means Id oc n ~ ' / See
~ . references [46-481 and Section 7.1 in Chapter 2
for details.
Stripping voltarnrnetvy
305
M c d 2 + with
. 4.31 (a) D C polarogram for the reduction of 1 x lop3 M dmgH2 and 1 x
a drop time of 0.6 s. (b) Cyclic voltammograms of 1 x loF3M dmgH2 at a hanging mercury drop
electrode with a scan rate of 200 mV s-l, (i) first scan; (ii) second scan. The temperature was 20°C
and 0.1 M ammonia buffer was the electrolyte. Adapted from: Anal. Chem. 70 (1998) 1312.
rocess giving rise to the aqueous solvent limit (Fig. 4.31 (b)) at the mercury electrode. However, on the reverse scan, a new oxidation wave at about -0.50 V
versus Ag/AgCl is observed. This wave is only present after the free ligand
has been reduced and therefore represents an oxidation process associated with
the generation of a reduced form of the ligand. In the second and subsequent
cycles, a new reduction wave is evident at about -0.60 V versus Ag/AgCl as
shown in Fig. 4.31 (b(ii)). This reduction process is present only subsequent to
the oxidation process at -0.50 V versus Ag/AgCl, and this chemically reversible
306
Electrode processes
redox coude is therefore the result of an initial reduction of the free ligand. The
reduction'of d m g ~ has
, been proposed to lead to the formation (eqn 4.77) of
2,3-di(hydroxylamino)butane (DHAB). If this is correct, the reversible redox
couple observed on the second and the subsequent cycles is likely to be the result
of the formation of a mercury complex with DHAB at the electrode surface.
The following scheme for the voltammetry of dmgH2 at a mercury electrode
therefore is proposed:
2DHAB (solution)
+ Hg(e1ectrode) +Hg (DHAB)2(solution) + 2e- (-0.55 V)
(4.87)
Importantly, the reversible redox couple observed on second and subsequent
cycles can be used as an indicator to determine whether the reduction of the
ligand is involved in the reduction process observed in the presence of cobalt
(and nickel).
Figure 4.32 shows D C polarograms obtained for c o 2 + ions (Fig. 4.32(a))
and then with increasing additions of dmgH2 to the 0.1 M ammonia buffer
solution (Figs 4.32(b-d)). The polarographic reduction of co2+in the absence of
dmgH2 has a limiting current value which is the result of a diffusion-controlled
Fig. 4.32 DC polarograms in 0.1 M ammonia buffer at 2OoC of. (a) 1 x lop3M co2+; (b) with
addition of 2 x lo-' M dmgH2; (c) with addition of 5 x lo-' M dmgH2; and (d) with addition of
1 x lop3 M dmgH2. Adapted from: Anal. Chew. 70 (1998) 1312. Copyright, American Chemical
Society.
Stripping voltammetry
307
3 3 Linear sweep voltammogram at a scan rate of 1000 mV s-' for the reduction of adsorbed
Co(dmgH)2 and Ni(dmgH)2 complexes at a 0.24 cm2 hangin mercury drop electrode. Experimental
conditions: 4 M NH3/NHrC1 buffer (pH 8.8); 10 pg L-'CO
and ~ i ~ 20°C;
' ;
2x
MdqH2
ligand; accumulation potential -300 mV versus Ag/AgCl; accumulation time 60 s. Provided by
courtesy: R.W. Knight, Deakin University, Victoria, Australia.
5'
two-electron reduction to C o ( 0 )(metal) under the conditions of Fig. 4.32.
The initial addition of dmgH2 to the solution results in an increase in the
limiting current as well as the appearance of a maximum.22With a considerable
excess of dmgH2, the limiting current attains a constant value which is about
four times greater [70] than the limiting current for the reduction of noncomplexed co2+. This result again indicates that a multi-electron reduction
process occurs at negative potentials where the number of electrons partaking in
the reduction reaction in the diffusion-controlled region is about eight electrons,
assuming equal diffusion coefficients for complexed and non-complexed forms
of cobalt(I1). This increase in the polarographic limiting current is similar to
at reported by Stromberg almost sixty years ago 1671.
Linear sweep voltammograms of solutions containing c o 2 + and ~ i and
~
excess dmgH2 at a hanging mercury drop electrode reveal large peaks attributable to reduction of adsorbed cobalt or nickel complex (Fig. 4.33). The
orptive stripping peak currents [72] are linearly proportional to scan rate
-1000 mV s-l), which is characteristic ofthe reduction of an adsorbed species
,48,49,72] in the 'thin film' configuration (Section 18 in Chapter 2). The
symmetrical shape of the process is also consistent with reduction of a surfaceconfined species (Section 18 in Chapter 2). At more negative potentials, the
reduction of the free ligand is encountered due to an excess being present in
the solution. Reduction of this excess ligand under conditions of cyclic voltammetry (nickel case only shown in Fig. 4.34) gives rise to the observation of a
chemically reversible process on the second cycle as expected when the ligand
is reduced (compare Figs 4.31(b) and 4.34). However, even when the switching potential was - 1-20V versus Ag/AgCl, thus eliminating the possibility of
reducing the free ligand, the same chemically reversible redox couple is still
evident (Fig. 4.34). This redox couple present on second and subsequent cycles
results from the formation of reduced dmgH,, and therefore the reduction of
22~olarographic
maxima are attributed to differences in interfacial tension around the
mercury drop causing solution streaming (see reference [711 for details).
+
308
Electrode processes
(4
Fig. 4.34 Cyclic voltammograms at a scan rate of 100 mV s-I and 20°C for the Ni(dmgH)2 cornplex generated in situ with excess dimethylglyoxime at a hanging mercury drop electrode in 0.1 M
ammonia buffer: (a) first scan and (b) second scan, (c) second scan, but after switching the potential at -1.2 V versus Ag/AgCl. Adapted from: Anal. Chem. 70 (1998) 1312. Copyright, American
Chemical Society.
the Co (dmgH), and Ni (dmgH), complexes also must involve ligand reduction.
Analogous cyclic voltammetric results were observed for both the nickel and
cobalt dimethylglyoxime systems. The voltammetric data therefore imply that
the electrochemistry of cobalt and nickel complexes in the presence of excess
ligand involves strong adsorption at a hanging mercury drop electrode and that
reduction of the adsorbed complex in a thin film format and also from bulk
solution represents a multi-electron system at both the dropping and hanging
mercury drop electrodes.
5.4.2 Bulk electrolysis and coulometn'c experiments i n aqueous media at
dropping mercury and mercury pool electrodes
For microelectrolysis experiments, a cell was designed which enabled a small
volume of the C ~ ( d m g Hsolution
)~
(1-2 mL) to be exhaustively reduced over
a 10-h period of time with a dropping mercury electrode (DME) having a drop
life of 0.5 s [70]. Since the limiting current in D C polarograms was reached over
the range from - 1.2 to -1.4 V versus Ag/AgCl (Fig. 4.32), the electrolysis was
carried out at constant potential in this region, with stirring provided by a
stream of nitrogen or argon gas which also removed oxygen. During the course
of electrolysis, D C polarograms were periodically recorded. From the plot ofthe
polarographic limiting current versus microelectrolysis time, the charge needed
for complete reduction was calculated by extrapolating to the time predicted
Stripping voltammetry
309
for completion of the electrolysis. The number of electrons taking part in the
reduction process at the DME was subsequently calculated from the derived
value of the total charge for exhaustive electrolysis.
The determination of the number of the electrons transferred as described
gave an n-value of 10.2 k 0.9 for 11 experiments [70]. The value was
endent of the ligand concentration when the dmgH, to Co(I1) concentration ratio was in excess of 2 : 1. This independence on the determined number
of electrons on the dmgH, excess, provides strong evidence that no cyclic
catalytic electrode reaction occurs at the dropping mercury electrode under
diffusion-controlled conditions.
To confirm the microcoulometric results at a dropping mercury electrode
are valid at stationary mercury electrodes, larger scale bulk-CPE (controlled
potential electrolysis) experiments were performed at a mercury pool electrode.
he results for the exhaustive electrolysis (coulometry) with different ratios of
gH2 and Co(I1) are summarized in Table 4.10. The Co(I1) concentration
m these experiments was constant at either 3.4 x lop5 or 5 x lop5M while
dmgH, concentration varied and it was found that the number of electrons
sferred (n) increases with increasing concentration of dmgH,, from the
value n % 2 for Co(I1) reduction in the absence of dmgH,, reaching n
10 for
mgH, concentrations higher than the stoichiometric 2 : 1 concentration ratio.
Table 4.10 The number of electrons (n) determined via
coulometry (bulk-CPE at a mercury pool electrode at 20°C)
for the Co(dmgH)2 reduction process in 0.1 M ammonia buffer
for different concentration ratios of dmgH2 to C0(11)~
-
Coniposition of
electrolysed solution
-
Concentration ratio
[dmgH,] / [Co(11)]
3.4 x
M Co(11)
0
3.4 x ~ O - ~ M C O ( I I ) 1.18
4x
M dmgH,
5.0 x
MCo(11)
2
1.0 x lop4M dmgH,
2.94
3.4 x loM5M Co(11)
M dmgH,
1.0 x 1o - ~
5.0 x low5M Co(11)
10
~ d m ~ ~ ,
5.0 x
M Co(11)
29.41
3.4 x
1.0 x 1o - ~M dmgH,
M CO(II)
100
5.0 x
5.0 x lop3 MdmgH,
-
Calculated number of
electrons, n
2.3
5.18
11.6 (1.5)
9.43
10.2 (1.7)
10.39
9.0 (2.3)
'Values in parenthesis are standard deviations from triplicate experiments; data
taken from reference [70].
3 10
Electrode processes
These results imply that the reduction of the Co(dmgH)2 complex, whether
present in bulk solution or adsorbed onto a mercury surface, still represents a
10-electron reduction process per molecule of Co(dmgH)2 complex.
During the reductive electrolysis at a mercury pool electrode, the solution
changed from a deep brown colour to clear, indicating complete loss of complex.
Very little Co(dmgH)2 could be voltammetrically detected in the colourless
solution. Voltammetric monitoring at very negative potentials also revealed that
the excess dmgH, is not consumed during the reduction. This result indicates
that free Co(I1) ions are not a product of electrolysis as they would immediately
react with the excess of dmgH2 in the solution and the brown colour would
persist. That is, the overall process does not involve pure ligand reduction of
the kind Co (dmgH)2--+ Co (11) reduced ligand giving rise to a cyclic catalytic
reduction process as represented in eqns (4.79)-(4.81).
Evidence to support reduction to cobalt metal was obtained via analysis of the
mercury pool electrode. After exhaustive electrolysis experiments, the mercury
pool electrode was dissolved in nitric acid and the resulting solution examined
by the cobalt dimethylglyoxime adsorptive stripping method. The expected
amount of cobalt was detected after cobalt metal present in or on the mercury
surface has oxidized to Co(I1) by nitric acid.
A CPE experiment also was performed at the mercury pool electrode on a
solution made from a sample of synthesized Co(dmgH)2. Coulometry showed
that the number of electrons transferred in the exhaustive reduction process
was 9.5 k 0.5 [70]. All coulometric evidence therefore suggests that both the
Co(I1) metal centre and the ligand are reduced when C ~ ( d m g His) ~
reduced at
a mercury electrode.
+
5.4.3 Constant current reductive coulometrlc stripping at
mercury thinrfilm electrodes
Ma et al., [69] have described a coulometric method for the determination
of the number of electrons involved in the reduction of Co(dmgH)2 and
Ni(dmgH)2 which is based on the analysis of chronopotentiometric (E - t )
data.23In this method, cobalt and nickel as their dimethylglyoxime complexes
are first quantitatively accumulated onto a mercury thin-film electrode as in
adsorptive stripping voltammetry, but they are then exhaustively stripped by
means of a controlled current rather than by sweeping the potential. The
plot of E versus t (Fig. 4.35) or, more usefully in the analytical sense, E versus dt/E (Fig. 4.36 (a)), produces the adsorptive stripping chronopotentiogram
in the technique of chronopotentiometric adsorptive stripping analysis. Figure 4.36(b) illustrates the addition of cobalt and nickel standard solutions to
dimethylglyoxime when the derivative form of readout is used.
Reductive coulometric stripping chronopotentiometry, using adsorptive
accumulation of metal complexes, M,L,, (Co(dmgH)2 and Ni(dmgH)2 are
3 detailed
~
description of the techniques of chronopotentiometry and constant current
chronopotentiometric stripping analysis is contained in reference [46].
2
Stripping voltarnrnetr,
311
Time
ig. 4.35 Chronopotentiometric (potential-time) curve. t represents the transition time [46] or
ts,,p in the coulometric stripping method [69].
Fig. 4.36 (a) Chronopotentiometric stripping analysis showing transformation of E versus t into
dt/dE versus E profile; (b) dt/dE versus E form of chronopotentiometric stripping curves obtained
after 180 s of accumulation followed by reduction with a constant current of 50pA in 24.8 pL
volumes of samples containing 0.10mM dimethylglyoxime (pH 9.2) and to which 0, 1, 2, 4, and
~ (ii)+co2+ have been added. Adapted from Anal. Chem. 69 (1997) 1782.
8 pg L-' (i) ~ i and
3 12
Electrode processes
the cases of interest in this book) is based on the quantitative adsorption of the
complexes onto a mercury drop or film electrode from very small solution volumes, typically [69] 10-30 pL, samples. Thus the initial stage of the electrode
process is
MmLp(solution)--+ MmLP(adsorbed)
(4.88)
After adsorption, the metal complex is reduced by means of an applied constant
current, Istrip,
according to the reaction scheme
in which the metal ion(s), the ligand(s), or both are reduced. From known
values for the sample volume, V, the number of electrons involved in the
reduction, n, and the time needed for quantitative reduction of the adsorbed
complex, tstfiP,
the sample concentration of complex CcompIex,
can be calculated
from Faraday's law as
(4.90)
Ccomplex = tstrip Istrip / n F v
However, any other reducible species present in the sample will also diffuse
to the electrode during reduction in quiescent conditions. This, as well as the
double-layer current requirements will lead to a fraction of the applied current
being consumed by alternative pathways [69]. Denoting this 'chemical current7,
ichem,
allows eqn (4.90) to be modified to
The chemical current, Ichem,which is normally at least one order of magnican be determined by repeating the adsorption/stripping
tude smaller than Istrip,
process under identical conditions, with the exception that different constant
are applied during reductive stripping. By plotting l/tstn, versus
values of Istrip
I,,,, a straight line is obtained, the intercept on the current axis24being equal to
Ichem.
Using this value of Ichem,the concentration of the complex in the sample
can be determined from eqn (4.91) and, obviously, vice versa, if the sample
concentration is known, the number of electrons, n, can be determined.
In the experimental studies described in reference [69] exhaustive adsorption
of Ni(I1) or Co(I1) complexes, in the 0-4 pg L-I concentration range, was
achieved by vibrationally promoted adsorptive accumulation of complex for
3 min at -0.30 V for Ni(dmgH)2 and -0.75 V versus Ag/AgCl for Co (dmgH)z
using E25 pL volume samples, hanging in drop form under a 3-mm diameter
GC mercury thin-plated working electrode [69] in a nitrogen atmosphere. The
adsorbed complexes were reduced by means of a constant current. Additional
details of the use of these very small volumes of solution and the enhanced mass
24~amples
containing either 1pg L-' Ni(I1) or 1 pg L-' Co(I1)were analysed using constant
reductive currents equal to 10, 30, and 50 pA. By plotting the inverse of the stripping times
versus the stripping currents, chemical currents equal to 3.1 pA for nickel and 3.4 pA for
cobalt were obtained by linear extrapolation.
Stripping voltammetry
313
ig. 4.37 Values of tSmp obtained by chronopotentiometric stripping voltammetry after 180s
of accumulation followed by reduction with a constant current of 50pA on samples containing 0-400 pg L-' nickel(I1). Reproduced by courtesy: Anal. Chem. 69 (1997) 1782. Copyright,
American Chemical Society.
transport achieved by vibration of the small volume are available in references
[69,73,74].
The coulometric form of analysis requires quantitative adsorption of the
complex and a linear relationship between metal complex concentration and
stripping signals, or stripping signals and electrolysis time. For a surface-based
process of the kind utilized in adsorptive stripping voltammetry, these linear
relationship conditions can be expected only at relatively low electrode surface loadings, (typically less than 5 per cent). The required linear relationship is
obtained in the 0-4 pg L-' concentration range for both nickel and cobalt (see
Fig. 4.37 for the nickel case).
By repetitive analysis of samples containing 1, 2, or 4 pg L-' Ni(I1) or Co(I1)
the number of electrons, n , involved in the reduction of Ni(dmgH)2 and
Co(dmgH)2 could be determined using eqn (4.91). The results are summarized in Table 4.11. The mean n values, 10.1 for Ni(I1) and 9.9 for Co(II),
with an estimated standard deviation of 0.2, clearly show that there are 10 electrons involved in the reduction at thin-film electrodes also as deduced above by
coulometry and CPE at dropping mercury and mercury pool electrodes. This
combination of coulometric studies almost certainly established that an overall
10-electron process involving both metal and ligand reduction occurs in the
reduction of Co(dmgH)2 and Ni(dmgH)z complexes at mercury electrodes.
Examination of electrolysis products generated by
exhaustive electrolysis at a mercury pool electrode
Infired spectroscopic study
5.4.4
Examination of products formed after large-scale reductive electrolysis at a
mercury pool electrode is difficult with conventional infrared spectroscopic
techniques due to the presence of strong absorption from the aqueous solvent.
3 14
Electrode processes
Table 4.11 Number of electrons involved in the
reduction of Ni(dmgH)2 and C ~ ( d m g H calculated
)~
according to eqn (4.91) and use of the technique of
coulometric stripping chronopo tentiometrya
'Potentiostatic adsorption from a 24.8 pL sample drop in a nitrogen
atmosphere for 180s prior to reductive stripping with a constant
current of 50 pA; data taken from reference [69].
b~oncentration.
However, it is possible to examine these solutions by Fourier Transform Attenuated Total Reflectance Infra-red (ATR) Spectroscopy using a water insoluble
zinc selenide crystal cell for the measurements. In order to obtain any information relevant to the electrolysis products, background subtraction was required.
Thus, negative bands are associated with consumption of the starting material and positive bands correspond to formation of new products. The major
new bands are located at 1630 cm-' and a group of peaks in the region around
3300 cm-l. These data are now considered with respect to each of the products
considered likely to be formed after an eight-electron ligand reduction process.
1. Reduction ofthe carbon-nitrogen double bonds ofthe ligand according to eqn (4.77)
A total of four electrons per ligand and a total of eight electrons for the two
ligands in the complex are expected according to the interpretation given by
eqn (4.77). I R evidence for 2,3-di(hydroxy1amino)butane being a product of
electrolysis is provided by the occurrence of new I R bands attributable to N H
groups in the 3300 cmwlregion and at about 1630 cm-l. The shape ofthe bands
in the 3200-3500 cm-' range suggests that hydrogen intramolecular bands, such
as N-H . . - 0or N-H . N may be present in the spectrum. A negative I R peak
would be expected for the consumption of the double bond at approximately
1600 cm-l. Unfortunately, this is difficult to detect in the 1R spectrum because
it is masked by the large positive peak at about 1630 cm-l, which is assigned
to the formation of an NH group as expected according to eqn (4.77). The
Stn'pping voltammetry
3 15
negative band located in the 1400 cm-' region may be assigned to the loss of
the symmetric C = N stretching band [75].
2. Reduction of the oxime group on the ligand Equation (4.92) involving
of both ligands
would give rise to an overall eight-electron ligand-based reduction process.
owever, this seems unlikely to be the major pathway because there is no
evidence of hydroxyl reduction in the I R spectrum.
3. Complete reduction of only one of the ligands in the complex to form the amine
according to eqn (4.78) In the scheme given in eqn (4.78), one ligand (eight
electrons) and the metal centre would be completely reduced leaving one unreduced oxime group. However, again no IR evidence to support this proposal
was found, since if the oxime groups were to be reduced to amine or imine
groups, the decay of their bands appearing above 3000 cm-' would have been
expected.
The great similarity between the I R spectra of the products resulting from the
reduction of Co (dmgH)2(electrolysispotential = - 1.25 V versus Ag/AgCl)
and reduction of free dmgH2 ligand (electrolysis potential = - 1.8 V versus
Ag/AgCl) provides evidence to support the proposal that the ligand undergoes reduction in the complex, and that in both cases the same product(s) is
(are) obtained. After eliminating 2,3-diiminobutane and 2,3-diaminobutane as
major products, it can be proposed that 2,3-di(hydroxy1amino)butane is a major
roduct of reductive electrolysis of the C ~ ( d m g Hcomplex
)~
in ammonia buffer.
ss spectromehy study
In the mass spectra of both dmgH2 and C ~ ( d m g H )the
~ , parent peak corresonding to dmgH2 ( m / z = 116, % Int = loo), as well as peaks corresponding
to fragments of the decomposition of dmgH2 ( m l z = 99, % Int = 69,
C 4 H 7 N 2 0m
; / z = 84, % Int = 9, C4H6NO;m / r = 68, % Int = 21, C4H6N;
m / z = 58, % Int = 24, C2H4NO; m / z = 42, % Int = 28,CNO) were
identified [70].
In the mass spectra of the electrolysis product(s), the five most intensive
peaks have the following characteristics: m / z = 87, % Int = 27, C3H7N20;
m / z = 85, % Int = 6, C3H5N20; m / z = 73, % Int = 81, C 3 H 7 N 0
or C2H50N2;m / z = 60, %C2H6NO; m / z = 45, % Int = 13, C H 3 N 0 .
Thus, they contain in their composition the hydroxylamine (-CH-NHOH)
groups. The peak ( m l z = 60, % Int = 100, C2H6NO),which is consistent
with half of the 2,3-di(hydroxy1amino)butane molecule, provided the highest
intensity in the mass spectrum. The fact that the peak is not observed in the
dmgH2 nor in the Co(dmgH)2 mass spectra implies that a major product of
3 16
Electrode processes
bulk electrolysis is 2,3-di(hydroxylamino)butane.The absence of a parent peak
for 2,3-di(hydroxy1amino)butane ( m / z = 120, C4HI2N202)in the spectrum
of the electrolysis product is attributed to the ready and symmetrical splitting
of the parent molecule under the conditions of mass spectrometry employed in
these studies [70]. It can also be noted that the mass spectrum of the electrolysis
product does not fit that expected for other hypothesized products of electrolysis such as 2,3-diaminobutane or oxygen free 2,3-diiminobutane, which might
have been formed as a result of eqns (4.78) and (4.92). The results of the mass
spectrometric analysis of solid dmgH2, Co(dmgH)2, and the product of electrolysis therefore also supports the conclusion that a major reduction product is
2,3-di(hydroxy1amino)butane.
5-4.5 Detection of one-electron reduced intermediates via
electrochemical studies i n aprotic organic solvents
The electrode process assigned to reduction of C ~ ( d r n g Hand
) ~ the nickel analogue has been concluded to involve an overall 10-electron reduction process.
The role of water and pH is always likely to be significant in complex reaction
schemes of this kind, so that studies in aprotic organic solvents may reveal intermediates in the reaction scheme. To overcome solubility problems associated
with Co(dmgH)2 and Ni(dmgH)2 in aprotic organic solvents, a dimethylglyoxime analogue was examined in which a six-carbon backbone chain replaced
the two methyl groups. The basic structure of the cobalt and nickel complexes
of 1,2-cyclooctanedione dioxime is shown in Fig. 4.30(b). These derivatives
may be referred to as carbocyclic bis-dioxime complexes25and are abbreviated
as C ~ ( c g d o Hand
) ~ Ni(C8doH)2.
Co (dmgH)2
The sparing solubility of C ~ ( d m g Hin
) ~dichloromethane allows a comparative
electrochemical investigation to be undertaken in this solvent in order to veri@
that it is valid to extrapolate data obtained from the C O ( C ~ ~ Oanalogue.
H)~
The initial reduction of the Co(dmgH)2 complex at a hanging mercury drop
electrode in dichloromethane (0.1M Bu4NPF6)which occurs at about -2.3 V
versus FC/FC+is barely resolved from the solvent limit at 20°C. However, when
the temperature is lowered to -50°C, when the mercury electrode is solid rather
than liquid, two redox processes having some degree of chemical reversibility are
observed (Fig. 4.38(a)), although the second reduction process is not resolved
from the solvent limit. At -70°C, both reduction processes are extremely
well defined (Fig. 4.38(b)). The reversible EIl2 values are -2.25 V for the first
process and -2.40V versus ~ c / F c +for the second process. Similar potentials
were obtained when acetone (0.1 MBu4NPF6)was used as the solvent.
2 5 ~ h ecomplexes are actually isolated as their dihydrates C O ( C ~ ~ O H ) ~ and
-~H~O
N i ( c , d ~ H -)2~H 2 0 , so on dissolution in an organic solvent neither completely anhydrous
or aprotic conditons are achieved. The presence of the waters of solvation is neglected in
subsequent discussion.
Stripping voltammetry
. 4.38 Cyclic voltammograms for reduction of 1 x loW3M C ~ ( d m g H in
) ~ dichlorornethane
(0.1M Bu4NPF6)at a hanging mercury drop electrode and using a scan rate of 1V s-' at (a) -50°C
and (b) -70°C. Adaptedfrom: Anal. Chem. 70 (1998) 1312. Copyright, American Chemical Society.
Voltammetric studies of the reduction of Co(dmgH)2 at mercury electrodes
in aprotic solvents therefore indicate that two primary one-electron transfers,
which could be metal or ligand based, are available:
Co (dmgH)2(solution) f e[Co(dm@) 21 - (solution)
+e
[Co(dmgH)21 - (solution)
(4.93)
[CO( d m g ~21)2- (solution)
vidence for a one-electron reduced nickel dimethylglyoxime complex was
also obtained in dichloromethane by cyclic voltammetry and also simultaneous electrochemical-ESR experiments [70]. Consequently, one-electron and
two-electron reduced intermediates probably do exist in the overall-electron
reduction process.
CO(C~~OH)~
The reduction of C O ( C ~ ~ Oin
H )dichloromethane,
~
as is the case with
C ~ ( d m g H )also
~ , occurs close to the solvent limit and again is only well defined
at low temperatures at a solid frozen mercury electrode. In dichloromethane,
at a hanging mercury drop electrode at -40°C, two reduction waves are again
observed, with the second not being fully resolved from the solvent limit. The
reversible Elj2 value for the first reduction process is -2.35 V versus ~ c / F c +
which is 100 mV more negative than for reduction of Co(dmgH)2. However,
clearly the voltammetry of the two complexes is closely related.
An investigation of the effect of water addition on the reduction process was
undertaken in acetone 20°C. Unlike the case with dichloromethane, water is
very soluble in acetone so the water addition experiments are readily undertaken in this solvent. Voltammograms obtained at a scan rate of 500mVs-'
3 18
Electrode processes
M C ~ ( c ~ d o at
H a) hanging
~
mercury drop electrode
for reduction of 1 x
in acetone, in the absence of deliberately added water, and with 2 per cent
water added, showed that the peak potential shifts reduction by about 300 m ~ ,
from -2.45 to -2.15 V versus FC/FC+, and the peak height increases by a
factor of three. Furthermore, in the presence of water, a new oxidation peak is
generated on the reverse scan (-0.8 V versus FC/FC+ in presence of 1 per cent
water) which has an oxidation potential which is dependent on the amount
of water added and has similar characteristics to the oxidation peak observed
after reduction of dimethylglyoxime in water. Similar results were obtained
when a pH 9 ammonia buffer solution was added instead of pure aqueous
solvent. The voltammogram for reduction of C O ( C ~ ~ OinHa) 100
~ per cent
water solution containing 0.1 M ammonia buffer at pH 9 exhibits only one
large reduction peak and has most of the features observed for reduction of
C ~ ( d m g H in
) ~aqueous media. The peak potential for the reduction process
is shifted by about 600mV to less negative potentials in changing from pure
organic to pure aqueous solvent.
Analogous experiments based on addition of 0.2, 0.6, and 1 per cent water
to acetone solutions were undertaken under polarographic conditions with a
drop time of 0.5 s. With increasing water content, not only was there a shift
in the half-wave potential towards less negative potentials (-2.40 to -2.20 V
versus FC/FC+ on addition of 1 per cent water) and an increase in the limiting current (50 per cent increase on addition of 1 per cent water), but the
formation of a polarographic maximum (see Fig. 4.32) was also observed.
Polarograms in acetone with addition of small amounts of added water therefore have most of the characteristics observed for reduction of C ~ ( d m g Hin
)~
aqueous media. Clearly water and acid-base reactions play an important role
in the reduction of C ~ ( d m g Hand
) ~ presumably they are also critical in reactions with the previously identified one- and two-electron reduced forms of
the complex.
Conclusions related to the reduction ofcobalt (and nickel) i n
the presence of dimethylglyoxime
5.4.6
The results described above, taken from references [69,70] suggest that the
electrochemical reduction of cobalt and nickel dimethylglyoxime complexes in
aqueous media involves the overall reduction ofboth the central metal atom and
the surrounding ligands in an overall 10-electron reduction process. In aprotic
solvents, both the cobalt and nickel complexes undergo an initial one-electron
probably followed by a second reduction one-electron step. However, in water,
electron transfer is coupled with proton transfer giving rise to a stable electrolysis
product, 2,3-di(hydroxy1amino)butane. In the case of the cobalt complexes
this is evidenced by the number of electrons obtained during coulometric and
microcoulometric experiments, identification of electrolysis products by I
spectroscopy and mass spectrometry, as well as the voltammetric observation
of a reduced ligand-based process after reduction. Apparently, reduction and
subsequent disruption of the ligand liberates c o 2 + , which must then lead to
Glucose biosensovs
3 19
formation of C o ( 0 ) at the prevailing potential. Hence, during the determination
of cobalt by adsorptive stripping voltammetry, the predominant overall process
that occurs at the mercury electrode surface is believed to be as follows:
c o 2 +(solution)
+ 2dmgH2(solution)-+ [Co(11)(dmgH)J (solution) + 2 ~ (solution)
'
(4.95)
o (11)(dm@)
(solution)
+ Hg(1iquid electrode) --+[CO(11)(dmgH)2]Hg(a&orbed)
(4.96)
+ 1 0 e + 1OH' (solution)
+CO(Hg)(amalgam or metal) + 2DHAB (solution)
[Co(II)(dmgH),]Hg(adsorbed)
(4.97)
In summary, there is no detailed theory-experiment comparison available
with the adsorptive stripping method, unlike the case with anodic stripping
voltammetry. It has in fact taken about fifty years of study on the cobalt
dimethylglyoxime system to achieve consensus on the nature of the overall
rocess. Indeed, even after employment of many modern voltammetric and
spectroelectrochemical studies, significant gaps in our knowledge of the details
of all steps associated with the mechanism still exist. Again the conclusion is
reached that the difficulty involved in establishing the details ofprocesses involving a complex mixture of solution phase and surface-based processes should
never be underestimated.
Voltammetric techniques frequently do not have adequate resolution for direct
application in biologically important fluids such as blood or urine. The problem
is that literally hundreds of electroactive species encompassing a wide concentration range may be present, which almost guarantees that overlap of processes
will occur within the small 2-3 V potential range available under ideal conditions in aqueous media. However, electrochemical detection of the reactions
between an enzyme and its biological partner, whose concentration needs to be
known, provides a powerful approach to selective and sensitive determination
of some biologically important compounds. The sensors prepared in this manner constitute the electrochemical class of the biosensor field of analysis, and
the series of reactions utilized illustrate the need to understand the details of a
complex series of reactions that enable these devices to operate. The voltammetric principles are in fact no different from those described above for carbonyl
complexes, photovoltaic dye-sensitizers, or stripping voltammetry, with a combination of homogeneous and heterogeneous processes being involved in all of
the mechanisms.
In 1962, the working principle of an electrochemically based biosensor was
demonstrated by Clark and Lyons [76]. In these pioneering studies [77], glucose oxidase (GOD) was fixed to a dialysis membrane which was then placed
320
Electrode processes
,
over the measuring section of an oxygen electrode. The device constructed ;
this manner may be termed a glucose biosensor and its use for the determination of glucose relied on the reaction in eqn (4.98) and measurement of the
consumption of oxygen via the oxygen e l e ~ t r o d e . ~ ~
glucose
+ oxygen +gluconolactone + hydrogen peroxide
GOD
The diagnosis and effective management of diabetes requires the convenient,
rapid, and precise determination of blood glucose. Consequently, the search to
fabricate commercially viable glucose biosensors has been the focus of intensive research during the last twenty years [78-811. While the initial protocols for
development of electrochemically based glucose biosensors employed eqn (4.98)
and measurement of oxygen consumption [75,82] or hydrogen peroxide formation [83], in commonly available hand-held devices, the chemical redox reaction
of glucose with oxygen is replaced by the reaction of GOD with so-called electrochemically detectable mediators such as ferrocene and its derivatives [79],
ferricyanide [ F ~ ( c N ) ~ ] [84]
~ - or other suitable redox reagents [85,86]. The
mediated enzymatic reaction sequence utilized with an oxido-reductase such
as GOD which contains a flavin adenine dinucleotide (FAD) coenzyme centre
can be represented [78] by the reaction sequence
I
+ Substrate -%Enzyme - FADH2 + Product (4.99)
Enzyme - FADHz + 2Med,, ----+ Enzyme - FAD + 2MedXd+ 2 ~ '
Enzyme - FAD
I
k2
I
(4.100)
I
with the reaction detected at an electrode by a mediator M being associated
with the process
In the reaction process given by the equation sequence (4.99-4.101), Medred
and Med,, represent the reduced and oxidized form of the mediator respectively and the current associated with catalysed process (4.101) is related to the
concentration of Medredwhich is in turn proportional to the substrate concentration which is glucose in the case of the glucose biosensor (Fig. 4.39). In terms
of the notation given in Section 7.3 in Chapter 2, the reaction sequence, when
all chemical species are water soluble, is represented by two chemical ( C )steps
(eqns 4.99 and 4.100) having homogeneous rate constants kl and k2 respectively,
and an electron-transfer (E) step having a formal potential, E!, a heterogeneous
I
II
I
11
I
8
I
2 6 ~ hoxygen
e
electrode, often known as the Clark electrode, is another example of an
important electrochemically based sensor. The principles of this device are available in
most textbooks on analytical chemistry. See for example, D.A. Skoog, E.J. Holler, and
T.A. Nieman, Princ@les of Instrumental Analysis, 5th edn, Saunders College Publishing,
Philadelphia, 1998.
I
1
1I
i
Glucose biosensors
32 1
Electrode
surface
. 4.39 Reaction scheme applicable to an electrochemical biosensor which depicts the role of
the mediator in the enzyme-catalysed reaction, where E is the enzyme and Med the mediator and
the subscripts ox and red represent their oxidized and reduced forms respectively.
charge-transfer rate constant k0 and charge-transfer coefficient a. The voltammetric theory associated with this sequence of reactions will be considered
ater.
Oxygen consumption can be monitored using a Clark oxygen electrode
which consists of a platinum working electrode, covered with an oxygenermeable membrane (e.g. Teflon), held at a potential of about -0.8 V versus
Ag/AgCl [76]. The current that flows under these conditions is associated with
the reaction
0 2
+ 4 ~ +' 4e-
Pt
P
2H20
-0.8 V versus Ag/AgCl
(4.102)
and the measurement of the reduction current as a function of time monitors
the depletion of oxygen.
Monitoring of H 2 0 2 generated by eqn (4.98) can be made by poising a
platinum electrode at +0.7 V versus Ag/AgC1 [78] and measurement of the
oxidation current derived from the reaction
H202
Pt
f0.7 V versus Ag/AgCl
+ + 2e-
2 ~ ' O2
(4.103)
as a function of time. Figure 4.40 summarizes the principles of the two main
classes of electrochemical glucose biosensors which utilize either oxygen as the
electron acceptor or an artificial electron acceptor.
Numerous biosensors based on O2 consumption or H 2 0 2 generation have
been reported. However, a potential difficulty in using oxido-reductase-oxygen
reactions is the stoichiometric limitation which arises under circumstances
where the oxygen concentration is lower than that of the analyte (e.g. glucose) concentration [78]. The now highly popular mediator-based biosensors
eliminate the oxygen dependency and also minimize interferences from ascorbic and uric acid [79-81,84471. Clearly the mediator should react rapidly with
the enzyme, exhibit reversible heterogeneous charge-transfer electrode kinetics
(large value of k0 in eqn (4.99)) and be stable with respect to pH, temperature,
redox state, and 0 2 .
322
Electrode processes
+
Glucose + Electron acceptor
Glucose oxidase
'---r-J
I
Electron acceptor = Oxygen
Artificial electron acceptor
1
Electrochemical transducer
(i.e. electrode)
Fig. 4.40 Schematic representation on the two methods available for development of an electrochemical glucose biosensor utilizing enzymatic oxidation of glucose by GOD.
According to the biennial reviews published in Analytical Chemistry, the
number ofpapers and patents based on electrochemical methods represent about
three-quarters of the total literature on chemical biosensors [88]. Thus, electrochemical biosensors hold a leading position in this area of research activity. In
the commercial sense, there are several forms of biosensors used in medicinal
chemistry [78,89-971. A remarkable feature of the commercial market is that
85 per cent of biosensor sales are for a single analyte-glucose biosensor [98]. The
success ofthis market predominantly can be attributed to the introduction of the
ExacTech system27(Fig. 4.41) which is a disposable glucose biosensor produced
by a combination of microelectric and screen printing techniques [81,99,100]
and which uses a ferrocene-mediated system. In practice, a single drop of blood
obtained by a pin prick from a finger is placed onto a sensor (Fig. 4.42) and afier
a few seconds the diabetic sufferer has access to their required glucose concentration level. Presently, more than a billion of these disposable single use sensors
are mass produced annually so this is a major success story of electroanalytical
chemist*.
The principles and applications of biosensors have been described in numerous books and reviews (see references [88,101-1111 for example). However,
the dominance of the ferrocene-mediated glucose sensor in the market place
and the elegance of the chemistry provide a compelling reason for choosing to
assess this device in terms of the principles established in Chapter 2.
Ferrocene (Fc) voltammetry presented in other parts of this book has been
shown to present an almost ideal example of a chemically and electrochemically
27~anufactured
by MediSense Inc., Abingdon, England and Cambridge, Mass., USA.
28~dapted
in part with permission from H.A.O. Hill and G.S. Sanghera, Biosensovs-A
Practical Approach (ed. A.E.G. Cass), Oxford University Press, Oxford, 1990, pp. 19-46.
Glucose biosenson
-41 The MediSense range of blood glucose electrochemical biosensors. Provided by courtesy
of MediSense, Abingdon, England and Cambridge, Mass., USA.
4.42 Measurement of glucose concentration by placement of a drop of blood onto a strip
sensing electrode which is coupled to a battery-operated hand-held instrument [81,98]. Provided by
courtesy of MediSense, Abingdon, England and Cambridge, Mass., USA.
Electrode processes
Fig. 4.43 Structure of ferrocene derivatives (FcR) and representation of the ferrocene electrode
process employed in ferrocene-mediated glucose electrochemical biosensors.
le 4.12 Reversible potential and second-order rate
constants for the oxidation of glucose oxidase by the
ferricinium derivative of the various ferrocenes as measured
at pH 7.0 and 25"Ca
"
Ferrocene desivative
E: (mv)
loP5k , , , ( ~ - ' sec-')
1,1'-dimethyl
Ferrocene
Vinyl
Carboxy
1,lr-dicarboxy
(Dimethylamino)methyl
100
165
250
289
403
3 86
0.77
0.26
0.30
2.01
0.26
5.26
"Data taken from references [79, 811.
b ~ e l a t i vto
e a SCE.
+
+
reversible Fc
FC+ e- process (e.g. Section 3.2 in Chapter 2). In this
sense, it is not surprising that electrochemical investigations using cyclic voltammetry also have shown [79,81,112-1 151 ferrocene, (bis(~5-cyclopentadienyl)
iron) to be an excellent mediator for use in electrochemical biosensors. Ferrocene itself [79,81] exhibits a reversible
value of 165 mV versus a SCE
and with the added advantage that the many derivatives (FcR) are available if
fine tuning of properties is required (Fig. 4.43, Table 4.12). In this form of
glucose biosensor, the ferrocene mediator replaces dioxygen as the cofactor for
GOD (Fig. 4.40). Thus, at a suitable electrode, held at a potential where FcR
is oxidized to FCR', the following reaction scheme occurs [79,81]
+
+ 2FcR + 2H'
k
Glucose + GOD,, ---+ GODred+ gluconolactone
GODred 2FcRf k'"' GOD,,
(4.105)
(4.106)
in which catalytic oxidation of the ferrocene derivative occurs to give very large
and easily measured currents that are proportional to the glucose concentration.
Glucose biosensors
325
The design and response optimization of a ferrocene-based biosensor for glucose
escribed in the following section.
Optimization of the pe f o formance $ a solution-phase
electvochemical glucose biosensor
Voltammetric studies on water-soluble forms of ferrocene can be employed
to demonstrate the interaction of ferrocene and the GOD-glucose system29.
errocene and its derivatives exhibit a reversible one-electron charge-transfer
rocess under conditions of cyclic voltammetry. Figure 4.44 shows a typical
cyclic voltammogram at a pyrolytic graphite electrode for 0.5 rnM water-soluble
ferrocene monocarboxylic acid (FMCA) in the presence of glucose. As required
for a diffusion-controlled reversible process (Section 8.1 in Chapter 2), the peak~ / ~(Ipis
to-peak separation, AEp is close to 60 mV at 25°C and ~ ~ / is vconstant
peak current and v the scan rate) over the scan rate of 1-100 mV s-'. From
ot of the peak current values against the square root of scan rate and use of
0.0
0.1
0.2
0.3
0.4
Potential, (V) vs SCE
0.5
Fig. 4.44 (a) Cyclic voltammograms at a scan rate of 1 mV sec-' for oxidation of 0.5 mM FMCA
in the presence of 50 rnM glucose; (b) as for (a) with the addition of 10.9 pM GOD. Reproduced
by courtesy: Anal. Chern. 56 (1984) 667. Copyright, American Chemical Society.
"' he
commercially available systems shown in Fig. 4.41 utilize principles related to those
described in Section 6.2, but are based on electrode immobilized (rather than solutionsoluble ferrocene) and GOD reagents in contact with both the worhng electrode and the
drop of blood to be determined for its glucose concentration (Fig. 4.42).
326
Electrode processes
the Randles-Sevzik equation (Section 8.1 in Chapter 2),
where [FMCA] is the concentration ofthe ferrocene monocarboxylic acid in the
bulk solution and the other symbols are as defined in Section 8.1 in Chapter 2,
the diffusion coefficient of FMCA can be determined to be 3 x
cm2 sec-I.
The change in shape of the voltammogran~upon the addition of 10.9 pM GOD
can be clearly seen by examination of Fig. 4.44. Thus, the peaks evident with the
reversible voltammetry of FMCA completely disappear on addition of GOD
and instead a large catalytic current flows at positive potentials, due to the
regeneration of ferrocene according to the equation sequence (4.108)-(4.110)
MEd(solution) ,
e
M,, (solution)
Ered(solution)
+ M,, (solution)
kcat
i
E,, (solution)
+ e-
+ Mred(solution)
(4.108)
(4.109)
A reversible electron-transfer process, which is followed by a catalytic reaction, as occurs in redox mediation by a mediator M with an enzyme E,
such as GOD, can be modelled by a reaction scheme of the kind shown in
eqns (4.108) to (4.110). The theoretical treatment of this reaction mechanism
has been reviewed extensively by Bartlett et al. [I 151. Under conditions where
eqn (4.109) represents the rate-determining step (k fast relative to kc,,), the
reaction schemes can be treated as a simple catalytic process represented as in
Chapter 2 by the notation
where Z is a component which regenerates A. The theoretical analysis for such
a system has been provided by Nicholson and Shain [I161 for the case where the
concentration of Z > A. From their treatment, different values for k:,,/a, where
a = nFv/RT and kLat is the pseudo-first-order rate constant, may be obtained. If
kLat/a, termed the kinetic parameter is small, then the cyclic voltammogram will
approximate that obtained for a simple reversible electron transfer as depicted
in Fig. 4.44(a). Conversely, if kL,/a is large and the reduced A (FcR) species is
continually replenished at the electrode, then a limiting or plateau current (him)
is observed as in Fig. 4.44(b). Consequently, no reduction peak is observed
since the concentration of species B (FCR') in the region close to the electrode
surface will be negligible. An analytical solution to the equation describing the
process under these pseudo-first-order conditions and assuming equal diffusion
Glucose biosensors
327
coefficients, D, for all species is given by
where [Ale is the concentration of A in the bulk solution (Section 8.2 in
Chapter 2) and Ell2 is the half-wave potential. At very negative potentials where
I = him, eqn (4.113) reduces to
which means that the limiting current is independent of scan rate for large values
of kLat/a. This is the reason why 4i, is referred to as a steady-state limiting
current.
Alternatively, quantitative kinetic information may be obtained by making
use of the working curve provided in reference [I 161 which relates Ic/Idto
the kinetic parameter k:,,/a, where Idis the diffusion current calculated as the
average of IP/v1l2(constant for a reversible system) and I,, the catalytic current,
is calculated from the increase in current upon the addition of enzyme. From a
set of Ic/Iddata points, values of k:, /a may be determined for several scan rates.
he calculated values of kLa,/a are then replotted against the inverse of scan
rate (v-I), for a series of GOD concentrations as depicted in Fig. 4.45(a). From
the slope of each curve, a good estimate of k:, for each GOD concentration is
obtained. A plot of k:, as a function of GOD concentration will have a slope
equal to the second-order homogeneous rate constant (kc,,) for the reaction
etween FMCA and GOD (in this example kc,, = 2.01 x lo5M-I sec-l) . This
inetic analysis provides an excellent method for choosing a good mediator for
a given enzyme system. A range of ferrocene derivatives and the rate of GOD
oxidation determined by this method are given in Table 4.12. Obviously, a
full simulation of the reaction scheme could be generated and kc,, calculated
by comparison of simulated and experimental voltammograms as described in
ection 12 in Chapter 2.
In glucose electrochemical biosensors, the limiting current value is measured
at a constant potential, e.g. 0.5 V versus SCE, based on the FMCA case given
in Fig. 4.44. Thus, the sensing is really of the amperometric (current at constant
otential) rather than voltammetric kind. In a practical enzyme electrode, where
rrocene is immobilized on the electrode surface rather than being present in the
solution phase as in the above discussion, other criteria must also be considered.
ost importantly, the solubility of the reduced form of the ferrocene derivative
must be low in aqueous solutions to aid entrapment within the electrode. In this
context 1, 1'-dimethylferrocene is an excellent mediator in terms of rate of oxidation of the enzyme and desirable physical characteristics for immobilization.
1, 1'-dimethylferrocene has an ~ F ~ v a l uofe 100mV versus SCE (Table 4.12)
and consequently for amperometric experiments the electrode may be poised
at 160mV versus SCE. The enzyme electrode is calibrated over the range
1-30 mM glucose (stirred solutions). Figure 4.45(b) depicts a typical calibration curve with the background current (-- 1.5 PA) subtracted and demonstrates
328
Electvode pvocesres
l/(Scan rate mVs-l)
0
5
10
15
20
25
30
35
Glucose (mM)
Fig. 4.45 (a) The kinetic parameter, kL,,/a, as a function of the inverse scan rate (mV s-I) for GOD
concentrations of ( ID ) 10.9, ( ) 20.6, ( ) 29.3, ( o ) 37.2 pM; (b) calibration curve for the glucose
electrode in (*) argon-, (0)air- and (+) dioqgen-saturated buffer. Reproduced by courtesy: Anal.
Chem. 56 (1984) 667. Copyright, American Chemical Society.
+
*
that a linear current versus concentration response is obtained in the range of
0-30 mM glucose. Above this concentration range the response is non-linear,
becoming insensitive to additional amounts of glucose above 70 rnM.
An important application of a glucose biosensor is in the clinical assay of
whole blood. In order to confirm that no problem arises from the presence of
oxygen dissolved in blood, the performance of the sensor has been tested in
air, argon, and oxygen-saturated buffer (Fig. 4.4507)). While there is very little
difference between the response in air and under argon, under pure oxygen,
there is a significant difference. However, since whole blood contains less than
200 pM dioxygen, this is not a significant problem. Analysis ofbuffered solutions
containing 7 rnM glucose and a range of metabolites commonly found in blood
shows that only ascorbic acid at 0.13 mM gives any increase in current [811.
Glucose biosensors
329
.3 Fabvication of a glucose bioelectvochernical sensor employing glucose
~ x i d a s eimmobilized onto a n electvode suvface
hile the details of commercially available glucose biosensors are propriety
nowledge, a laboratory bench-based experiment suitable for a student practical
exercise or evaluation experiments based on the following enzyme immobilization method has been described in reference [$I]. A brief outline ofthis method
rovides the concepts needed to produce a practical sensor.
(1) Graphite electrodes onto which GOD can be immobilized are constructed
by cutting a 4-mm diameter disc from a rod. The disc is then sealed into a
glass tube with epoxy resin, with a connection to the external circuit being
made by a wire, bonded with silver araldite, which is attached to the back
of the electrode. After electrode pre-conditioning at 100°C for 40 h, and
cooling in air, the enzyme, and ferrocene mediator may be immobilized by
the following steps.
) Deposit 15 pL of a 0.1 M ferrocene solution (1,1'-dimethylferrocene dissolved in toluene) on to the electrode surface and allow it to air dry.
(3) Place the electrode in 1 mL of a 0.15 M solution of l-cyclohexyl-3(2-morpholinoethyl) carbodiimide-p-methyltoluenesulphonate in 0.1 M
acetate buffer (pH 4.5) and incubate for $0 min at 20°C.
) Thoroughly rinse the electrode in water and then place it in a stirred 0.1 M
acetate pH 4.5 buffer solution containing 12.5 mg m ~ - 'GOD. Incubate at
20°C for 90 min.
) Rinse the electrode with the same acetate buffer, cover it with 0.03 pm
polycarbonate membrane (Nucleopore) and store in buffer containing 1 rnM
glucose.
(6) Prior to use, condition the electrode to give a stable current. This is achieved
by maintaining the electrode potential at 160 mV versus SCE for 10 h in a
7 mM glucose solution.
4 Glucose analysis ofwhole blood with a cornrnercially available
ucose bioelectvochernical sensor
The transformation of the ferrocene-based glucose enzyme electrode from a
laboratory bench version to a hand-held commercial device has been achieved
in the instrument marketed worldwide as the ExacTech glucose meter (Figs 4.41
and 4.42). The meter, in one form, comprises a pen-sized potentiostat (length
136 mm), weighing less than 30 g, with an LCD display for the glucose reading
ig. 4.42). The electrode, disposable test strip, incorporates an immobied layer of GOD and l ,lr-dimethyl-3-(1-hydroxy-2-aminoethyl)ferrocene,
ated with a hydrophilic membrane to attract the blood sample (one drop).
ch strip also contains its own reference electrode. The ExacTech blood glucose meter kit includes an automatic lancing device and glucose control solutions
for testing the meter, in an easy to carry travel case. The test strips are calibrated
on the ExacTech system with fresh whole capillary blood and the meter is able
330
Electrodeprocesses
to operate between 18°C and 30°C and 20-80 per cent humidity, producing- a
glucose reading in mg L-' in 30 s.
Assays of blood plasma samples with the enzyme electrode have been compared with results obtained with a Yellow Springs Instrument glucose analyser
routinely used in hospitals [81]. The latter device also incorporates the use of
GOD but is based on the detection of hydrogen peroxide in pre-diluted plasma.
Results for a sample size (n) of 23 gives a correlation coefficient between the
two assays of 0.98. An assay of whole blood, after the addition of heparin as an
anti-coagulating agent, when compared with plasma glucose levels measured
using the enzyme electrode yielded a correlation coefficient for the two assays
of 0.99 (for n = 10).
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In Chapter 2, the basic theory describing the voltammetry associated with oxidation or reduction of an ideal thin layer was considered. Unfortunately, such
ideally behaved thin films are rarely encountered in practice and cannot even be
prepared for many redox active systems. Thus, concepts based on these principles have limited applicability. As an alternative to preparation of film modified
electrodes, a general form of attachment of solids to electrode surfaces has arisen
from the work of Scholz et al. [I-41. In the common form of this attachment
method, the face of a disc electrode is placed in contact with crushed or powdered microparticles and on pressing the electrode against the solid, mechanical
transfer of small quantities of material occurs. Other methods of transferring
microparticles to electrode surfaces, and which achieve the same end result,
have been employed in more recent studies [4]. In this chapter devoted to electrochemical studies of solids, the electrode onto which has been placed an array
of microparticles is transferred to a solution (electrolyte) in which the solid is
insoluble (sparingly soluble). An attractive feature of this method is that the
voltammetry of almost any redox active solid can be studied, irrespective of
whether the oxidized or reduced forms of the solid are conducting or nonconducting, or whether 'thin film' or 'thick film', or other forms of adhered
solid conditions are achieved in the theoretical sense. Furthermore, and just as
importantly in the educational sense, almost all the basic features relevant to
all forms of solid-state electrochemistry, including those relevant to the thin
film case can be illustrated via a detailed description of a few carefully selected
examples of voltammetric studies of redox active solids adhered to electrode
surfaces as arrays of microparticles.
Under the conditions relevant to studies on microparticles considered in this
chapter, the interfaces formed by an individual solid particle present on an
electrode surface (as a hemisphere) can be described by the diagram in Fig. 5.1.
As revealed by casual inspection of this schematic diagram, an electrochemical
solid-state redox reaction, in even its simplest form, represents an inherently
complex multiphase problem. However, the processes that must occur are not
too difficult to envisage in a conceptual manner. For example, assuming both
Introduction
335
Phase 4
Electrode
. 5.1 Schematic representation of an example of (a) a three-phase system initially present prior to
commencement of electrolysis and (b) a four-phase interface system generated when a (hemispherical)
solid particle is oxidized or reduced after being adhered to an electrode surface which is then placed
in contact with solution (electrolyte). Phases 1 and 2 are the oxidized and reduced forms of the solid.
Phase 3 is the solution (electrolyte). Phase 4 is the electrode surface.
the oxidized and reduced forms of a species adhered to an electrode surface
are insoluble in the solution (electrolyte) phase, the oxidation process, when
neutral species A(so1id) is attached to the electrode surface and oxidized in a
one-electron process to [A+][X-] (solid), can often be described by eqn. (5.1)
+
+
[A+][X-] (solid) eA(so1id) f X- (solution)
(5.1)
X- is the anion from the electrolyte that must be transported across the
n-solid interface and then within the solid, as a means of achieving charge
neutralization as the electron is transported from the solid to the electrode.
n order to understand the processes that occur when A(so1id) is converted
to [A'] [x-](solid) at an electrode surface, as for example in eqn. (5.1), many
of the spectroscopic techniques used to elucidate features of the solution-phase
voltammetry (Section 16 in Chapter 2) will need to be applied in their solid-state
relevant formats. Additionally, surface analysis techniques that enable changes
in chemical composition of a solid to be detected, X-ray diffraction techniques
t allow the phases of A(so1id) and [A+][X-](solid) to be characterized, and
situ and ex situ forms of microscopy (electron scanning, atomic force, etc. as
scribed in Section 19.1 in Chapter 2) will emerge as important techniques
t need to be used to understand the important morphological changes that
take place when A(so1id)is converted to [A'] [X-] (solid).Additionally, since mass
changes almost invariably accompany a solid-state electrochemical process, the
use ofthe EQCM technique (Section 18.2 in Chapter 2) becomes very valuable.
A long list of variables have already been identified as being present in the
solid-state voltammetry of a process as simple as that given in eqn (5.1). However, to understand the full level of complexity, it needs to be recognized that
solids may be attached to electrode surfaces in numerous ways and that it is not
always obvious whether the attached solid exists as a thin film, a thick film, an
array of microcrystalline or amorphous solid particles, or combinations of these
morphologies. Essentially, it must be recognized that the nature of the surfaceattached solid does matter and that even the voltammetry itself may induce a
otential dependent change in the nature of the adhered solid.
Instead of having a film or an array of microparticles on an electrode surface,
a solid-electrode interface also may be established by mixing solid sample with,
say, carbon paste and making a composite electrode or by incorporating the
redox solid into a conducting polymer electrode system. However, in essence,
in all forms of solid-state electrochemistry, a complex interfacial, multi-phase
system is inherently always present, irrespective of the format in which the
solid is adhered to an electrode surface. Consequently, since the charge and
ion-transport processes and morphological changes also are likely to be cornplicated, it is not surprising that the time (scan rate) dependence of solid-state
voltammograms is enormously variable, and far more difficult to quantify than
when the oxidized and reduced forms of an electrode process are both soluble
in the solution phase.
It follows from the above discussion that matters to be addressed in electrochemical studies of solid microparticles adhered to an electrode surface,
and illustrated in the examples provided in this chapter, must include the
following:
(1) What is the mechanism of the electron-transfer process?
(2) How is the charge neutralization achieved to counterbalance the charge
associated with the transport of the electrons?
(3) IS the size (thickness, length, width, volume) and morphology of the
microparticles important?
(4) What is the reversible potential of the process described in eqn (5.1) and how
should the activities of species involved in the reaction and the nature of the
different phases formed huring the course of a voltammetric experiment be
defined?
(5) What is the relationship of the voltammetry of an array of microparticles to
that of an ideal thin film considered in Section 17 in Chapter 2?
(6) What is the relationship to the voltammetry of microp&-ticles to that of
redox active centres incorporated into conducting polymers grown onto
electrode surfaces and to solids mixed with carbon paste to make modified
carbon paste electrodes?
(7) Are dissolution processes inherently likely to be associated with the voltammetry of solids attached to electrode surfaces? This point needs to be
considered since neither A(so1id) nor [A+][X-](solid) are truly insoluble in
the solvent (electrolyte) of interest and it also needs to be noted that during
the initial stages, [A+][X-] (solid), generated by oxidation of A(solid), must
be present in quantities below the solubility level.
In view of their inherent complexity, a fully comprehensive account of the
subject of solid-state voltammetry (electrochemistry)is beyond the scope of this
Mechanistic aspects of transport processes
337
book. For a more detailed account on the subject of solid-state electrochemistry, the reader is referred to references [4-81. Results obtained for the redox
active solid systems presented in this chapter illustrate only the common features that are associated with voltammetric studies of solids. Importantly, the
coupling of spectroscopic, microscopic, surface science, and microgravimetric
techniques with electrochemical measurements is highlighted to emphasize that
voltammetric data alone are even less likely to be adequate to characterize a
redox response that occurs at an electrode surface, than is the case when redox
active species are solution soluble. A major problem in solid-state studies is that
the measured current may only reflect the 'average' of many non-time-resolved
events rather than the microscopic details that are required to provide a detailed
en arrays of nonorder to introduce the concept of voltammetry of non-conducting
croparticles attached to electrode surfaces in contact with solution (electrolyte), the oxidation of decamethylferrocene, ~ e ( q ~ - ~ ~ ~ e ~ ) ~to( s o l i d ) ,
its cationic, [ ~ e ( q ~ - ~ ~ ~ salt
e ~ and
) ~ ]the
+ , oxidation of tranr- and
~i+Cr(CQ),(dpe)~(solid)
and reduction of the tranr-[Cr(CQ)2(dpe)2]X salt
e = Ph3PCH2CH2PPh2,X- = anion) are considered as typical examples
of this form of solid-state voltammetry.
T h e oxidation of decamethylfewocenel
As noted in Section 3.2 in Chapter 2, the organic solvent solution-phase
~ Hfor
~ reference
) ~
voltammetry for the oxidation of ferrocene, F ~ ( ~ ~ - isc used
ectrode calibration because it represents an example of an almost ideal model
of an electrochemically and chemically reversible process. In Section 3.1, the
voltammetry of the related decamethylferrocene compound, Fe(q5-C5Me5)2,
is shown to represent an ideal process to illustrate many of the principles of the
voltammetry of adhered microparticles.
3.1.1
Voltammetry of decamethyljirrocene dissolved in dichloromethane
The voltammetry of F e ( q ' - ~ ~ M ewhen
~ ) ~ dissolved in organic solvents
corresponds to an essentially reversible one-electron charge-transfer process
19,101.
+ [ ~ e ( ~ ~ - ~ ~ M e ~ ) ~ ] + ( s+o le-u t i o n )(5.2)
' ~ d a ~ t with
e d permission from r. Electroanal. Chem. 372 (1994) 125. Copyright, Elsevier.
5
Fe(q -C5Me5)2(solution)
0.5
E (V) vs Ag/AgCI
Fig. 5.2 Cyclic voltammograms obtained at 22°C with a basal-plane pyrolytic y p h i t e electrode
at scan rates of 20 (smallest current), 50, 100, and 200 (largest current) mVs- for oxidation of
1 x 1oP3 M decamethylferrocene dissolved in dichloromethane (0.1 M Bu4NPF6).Adapted from:
J. Electroanal. Chem. 372 (1994) 125.
M decamethylferrocene disA cyclic voltammogram for oxidation of 1x
solved in dichloromethane (0.1 M Bu4NPF6)using a basal-plane graphite electrode is shown in Fig. 5.2. The voltammetric response is chemically reversible
and corresponds to a reversible one-electron diffusion-controlled process. Ideally, the cyclic voltammetric peak potentials for both oxidation and reduction for
a reversible process should be independent of scan rate (Chapter 2), and the slight
dependence of peak potentials on scan rate observed in Fig. 5.2 is attributable
to a small amount of uncompensated resistance (Section 6 in Chapter 2).
General considerations related to the voltammetry of microcrystals of
decamethyljifewocene attached to a n electrode surface
3.1-2
Since decamethylferrocene is insoluble and does not react with water, the
voltammetry of this compound when attached to an electrode surface in
contact with aqueous (electrolyte) media, could be expected to provide a simple system to explore some of the factors that govern the charge-transport
process and the incorporation of electrolyte anion (X-) when water insoluble
[ F ~ ( $ - c ~ M ~[XI
~ )(solid)
~ ] is formed via the process described in eqn (5.3).
A major difficulty that normally arises in voltammetric studies of surfaceconfined non-conducting solids like decamethylferrocene is associated with the
large IRu (ohmic) drop and the associated large value of the R,C time constant
(R, = uncompensated resistance, C = capacitance), which makes it difficult
to couple electron transfer and ion transport in a sufficiently fast manner to
achieve measurable currents within the available potential range. However,
when the solid is attached to the electrode as an array of microparticles, a relatively large surface area of sample is exposed to the electrolyte solution at the
Mechanisticaspectsoftransportprocesses
339
electrode-solution interface. This solid-electrode-solvent (electrolyte)interface
generated during a voltarnrnetric experiment (Fig. 5.1) plays an important role in
achieving well-defined voltammograms with non-conducting solid compounds
by minimizing the IR, drop and time constant, allowing rapid coupling of electron and ion-transport processes within the solid and allowing rapid transport
of ions across the solid-solvent (electrolyte) boundary.
3.1.3 Method of attachment of decamethyyeferrocene to the electrode suface
working electrode used to describe the solid-state voltammetry of
methylferrocene at aqueous (electrolyte)interfaces in this chapter was made
from basal-plane pyrolytic graphite (5-mm diameter). In order to transfer solid to
the surface of the carbon electrode, 1-3 mg amounts of crystalline decamethylferrocene complex were placed on a filter paper. After grinding the sample
to microcrystalline size using the flat side of a spatula, the carbon electrode
was pressed onto and rubbed over the crushed decamethylferrocene, thereby
causing some of the compound to adhere to the electrode surface as an array of
microcrystalline particles. For electrochemical measurements, the electrode was
transferred into a conventional electrochemical cell which contained the aqueous (electrolyte) solution of interest. The electrode surface was renewed after
each measurement by dipping into dichloromethane and allowing the comund to dissolve in the organic solvent or alternatively, a clean surface was
tained by cutting the soft graphite electrode surface with a razor blade.
3.1.4 Generalfeatures of the voltammetry of solid decamethyyefenocene attached to a
graphite electrode which is placed i n contact with aqueous electrolyte
A voltammogram for the oxidation of solid microcrystalline decamethylferrocene attached to a basal-plane graphite electrode and then subsequently
immersed in aqueous 0.1 M NaC104 electrolyte solution is shown in Fig. 5.3 as
a function of scan rate.2 The symmetrical responses are clearly different in shape
to those obtained in solution-phase voltammetry (Fig. 5.2). The areas under
reduction and oxidation signals are similar and for the voltammograms shown
in Fig. 5.3 the charge corresponds to 96 pC at a scan rate of 10 mV s-'. Scanning electron microscopy (SEM) data reveals the presence of microcrystalline
material on the surface with a particle size range for the majority of particles
eing 0.3-2 pm (Fig. 5.4). Using a very crude estimate of the electrode surface
coverage of approximately 1 per cent for 1 pm particles (200 ng decamethylferrocene over the 0.5 cm electrode diameter), the magnitude of the current for
exhaustive electrolysis should produce about 60 pC of charge. This calculation
suggests that the compound is significantly electrolysed during the course of
the slow scan rate voltammetric experiment. However, the appearance of the
microcrystalline compound after oxidation, as established by SEM, reveals no
'under some circumstances, particularly at very high scan rates or with very high surface
coverage, peak splitting, and additional processes may be detected. This more complex
situation is only briefly considered in this book. The interested reader should consult
references [I 1,121 for a more detailed discussion of these phenomena.
E (v)'vs Ag/AgCl
E (V) vs Ag/AgC1
Fig. 5.3 Cyclic voltammograms obtained at 22'C with scan rates of 10, 20, 50, and 100 mV s-I
for oxidation of microcrystalline decamethylferrocene attached to a basal-plane pyrolytic graphite
electrode placed in contact with aqueous 0.1 M NaCIOi electrolyte. Adapted from: J. Electroanal.
Chem. 372 (1994) 125.
Fig. 5.4 Scanning electron micrograph of microcrystalline decamethylferrocene attached to a
basal-plane pyrolytic graphite electrode. Reproduced by courtesy: J. Electroanal. Chem. 372
(1994) 125. Copyright, Elsevier.
dramatic changes in the number of particles and their sizes. In contrast, it will
emerge in other systems considered later in this chapter, that significant changes
in particle shapes and sizes may occur, and in other cases only a small fractions of the solid are oxidized or reduced during the course of a voltammetric
experiment.
Varying the scan rate in NaC104 electrolyte leads to a marked change in
both peak width and peak separation. At higher scan rates, the signals broaden
Mechanistic aspects of transport processes
34 1
5.1 Cyclic voltammetric data obtained as a function of scan rate at 22°C
fir oxidation of microcrystalline decarnethylferrocene attached to a basal-plane
py~olyticgraphite electrode which is in contact with aqueous (0.1 M NaC104)
electr~l~te"~~
7
Scan rate (mV s-l)
EpO" (mVc)
I3rd(mVc)
AEp (mV)
WP;2 (mV)
W$ (rnv)
--
a l l a h taken from J. Electroanal. Chern. 372 (1994) 125.
h ~ , " " represents the oxidation peak potential; E F ~the
, reduction peak potential; AEP = E r - Ered;
P
W G , the peak width at half-height for the oxidation process; w$, the peak width at half-height
for the reduction process.
?nmV versus Ag/AgCl.
considerably (Fig. 5.3, Table 5.1). At a scan rate of 10 mV s-l, the peak width
at half-height of about 30mV is clearly much lower than the value of about
90 mV predicted for the ideal thin film case (Section 18 in Chapter 2). The
agnitude of peak currents decreases slightly after the first cycle. However, the
rocess remains visible over prolonged periods of cycling of the potential.
Voltammograms obtained with a rotating disc electrode in the presence of
0.1 M NaC104 electrolyte (Fig. 5.5) show the same general characteristics as
with a stationary electrode, although the small change in wave shape of the
reduction process and decrease of the current with repetitive cycling suggest
that minor loss of material may be occurring into the solution phase by product
dissol~tion.~
Comparison of results at stationary and rotated electrodes leads to
the conclusion that the influence of mass transport as a rate-determining step
is confined predominantly to the solid phase or the solid-solvent (electrolyte)
interface.
3.1.5 Analysis of composition o f solid before and after oxidative electrolysis
lectron probe analyses of samples of adhered decamethylferrocene electrolysed
in the presence of 0.1 M KC104 electrolyte at different potentials enables the
3~oltamrnograms
obtained with 0.1 M KPF6 as the electrolyte show no wave shape change
and no loss of oxidation or reduction current within experimental error irrespective of
whether the electrode is rotated or stationary. The almost indefinitely stable response leads
to the conclusion that no loss of oxidized or reduced decamethylferrocene compound
occurs by dissolution into aqueous 0.1 M KPF6 electrolyte.
E (V) vs Ag/AgCl
Fig. 5.5 Voltammetry (three consecutive potential cycles) at 22O C of microcrystalline decamethylferrocene adhered to a rotated basal-plane pyrolytic graphite electrode placed in contact with aqueous
0.1 M NaC104 electrolyte. Scan rate = 100 mV s-' . Rotation rate = 1000 rpm. Adapted from: J .
Electroanal. Chem. 372 (1994) 125.
change in elemental composition of the solid to be m ~ n i t o r e d Prior
. ~ to electrolysis, the carbon background is clean with respect to both iron, chlorine, and
potassium (Fig. 5.6(a)). In the case of decamethylferrocene attached to a basalplane graphite disc electrode, oxidation for 2 min at 500 mV versus Ag/AgCl
in the presence of 0.1 M KC104 as the electrolyte leads to the smaller particles
on the electrode surface having close to a 1 : 1 iron to chlorine composition
(Fig. 5.6(b)), as required for formation of [ ~ e ( q ~ - ~ ~[C104]
M e ~(solid).
) ~ ]This
suggests that surface oxidation occurs with uptake of one perchlorate anion for
each electron transferred. Larger particles in the order of 10 prn show a reduced
chlorine signal implying that oxidation is favoured at the surface or edge of the
crystals and that unoxidized material remains at the centre of the larger crystals on the time-scale of this experiment. A sample oxidized first at +500 mV
and then reduced at -500 rnV versus Ag/AgCl shows only iron to be present
(Fig. 5.6(c))as expected when [ ~ e ( q ~ - ~ ~(solid)
M eis~reformed
) ~ ] after expulsion
of C l o y incorporated during the initial oxidation step.
The use of a wide range of electrolytes in the aqueous solution phase was found
to produce drastic changes in the voltammetric behaviour which, as might be
expected if eqn (5.3) is correct, are almost exclusively anion related. Figure 5.7
shows the voltammetry of decamethylferrocene when the electrode to which it
is attached is placed in contact with HC104, NaC104, and KC104 electrolyte
solution. No major change is found when C l o y remains as the anion and only
"he NaC104 electrolyte was replaced by KC104 in these experiments because of the
easier detection of potassium, instead of sodium, by the electron probe analysis technique.
In principle, K+ cations could become incorporated into the solid during the course of
potential cycling experiments so the failure to detect the element potassium eliminates this
possible mechanism.
Mechanistic aspects of transport processes
343
. 5.6 Electron probe X-ray microanalysis of decamethylferrocene attached to a basal-plane
pyrolytic graphite electrode placed in contact with aqueous 0.1 M KC104 electrolyte. (a) Electrode
background; (b) after oxidative electrolysis at +500 mV; (c) after oxidative electrolysis at +500 mV
and then reductive electrolysis at -500 mV versus Ag/AgCl. Reproduced by courtesy: J. Electroanal.
Chern. 372 (1994) 125. Copyright, Elsevier.
the cation is varied. Thus, changes in the cation are concluded to be of minor
importance.
In Fig. 5.8, voltammograms obtained in different electrolyte media are comared when the electrolyte anions are varied. A very wide range of potentials
observed as the identity of the anion is changed. Other differences are
so present. The distinctively different voltammetric data obtained when the
drophobic tetraphenylborate anion (BPhJ (Fig. 5.8 (a)) and the carboborane anion ([CB1I Hlz]-) (Fig. 5.8(b)) relative to hydrophilic PF, (Fig. 5.8(c))
and C l o y (Fig. 5.8(d)) support the concept that hydrophobicity is important,
although the size ofthe anion also may be relevant. With BPh, and [CBIIHl2]the current magnitude decreases relative to the C104 case and there is a negative shift in potential of more than 400mV. In perchlorate electrolyte, both
oxidation and reduction peak potentials have their most positive values. The
relationship of the voltammetry in C10, media to that observed in the presence
of the slightly more hydrophobic anion PF, is interesting. The oxidation process
shows a similar symmetrical response at low scan rates in both cases, whereas
the reduction process in PF, electrolyte is broader and less symmetric (compare Fig. 5.8(c) and (d)). When fluoride (Fig. 5.8(e)) or chloride (Fig. 5.8(f))
are present in the electrolyte, the response rapidly decreases in current magnitude and almost disappears after 10-20 cycles indicating dissolution of oxidized
product. Peak splitting is sometimes observed under higher coverage conditions
than those used to obtain the voltammogram shown in Fig. 5.7. Hence, in
344
Solid-electvode-solvent intefaces
NaC104
(a)
-0
E (V) vs Ag/AgCl
Fig. 5.7 Cyclic voltammograms obtained at 22°C and with a scan rate of 100 mV s-' for microcrystalline decamethylferrocene attached to a basal-plane pyrolytic graphite electrode and then placed in
contact with aqueous media containing (a) 0.1 M NaC104, (b) 0.1 M KC1O4, and (c) 0.1 M HC104
electrolyte. Adapted from: J. Electroanal. Chem. 372 (1994) 125.
NaClO
E (V) vs Ag/AgCl
E (V) vs Ag/AgCl
Fig. 5.8 Cyclic voltammograms obtained at 22°C and with a scan rate of 100 mV s-' (except for (f)
at 200 mV s-') for microcrystalline decamethylferrocene attached to a basal-plane pyrolytic graphite
electrode and placed in contact with aqueous (0.1M) electrolyte (a) NaBPh4, (b) CsCBl1Hl2, (c)
KPF6, (d) NaC104, (e) NaF, and (f) NaC1. Adapted from: J. Electroanal. Chem. 372 (1994) 125.
Mechanisticaspectsoftransportprocesses
345
E (V) vs Ag/AgCl
. 5.9 Cyclic voltammograms obtained at a scan rate of 10 mV s-' for oxidation ofmicrocrystallir~e
decamethylferrocene attached mechanically to a basal-plane pyrolytic graphite electrode and placed
in contact with aqueous 0.1 M KPF6 electrolyte at temperatures of (a) 22" C, (b) 52" C, and (c)
82" C. Adapted from: J. Electroanal. Chem. 372 (1994) 125.
summary, the identity of the anions as well as particle size, electrode coverage,
and scan rate all clearly influence the voltammetry.
3.1.7 Efect of temperature
th a kinetically controlled process, the rate usually increases with temature. The voltammetric response of decamethylferrocene attached to a
basal-plane graphite electrode immersed in 0.1 M aqueous KPF6 electrolyte
Is (Fig. 5.9) that a step of 30°C to higher temperature results in an
ximately three-fold increase of peak current. The current increase with
temperature is much more pronounced than in solution-phase voltammetry and
comparable to data reported in electrochemical studies with polymer films [13].
Electrochemistry of microparticles of
r a n ~ - C r ( C O ) ~ ( d p ter) a~ n, ~ - [ C r ( C O ) ~ ( d p[XI
e ) ~salts,
]
and
C ~ ( C o ) ~ ( d p (dpe
e ) ~=
, bidentate P h 2 P C H 2 CH2PPh2, X - = anion)
attached to a n electvode surface5
th trans- and ~ i s - C r ( C O ) ~ ( d p esolids
) ~ are non-conducting and insolue in water, so that studies analogous to those on decamethylferrocene
possible. However, significant additional information is available with
t r a n s - [ ~ r ( ~ ~ ) ~ ( d ~reaction
e ) ~ ] ~ /via
+ use of the EQCM and other
methods [I 2,141, while structural changes accompanying oxidation of solid
~ i s - C r ( C O ) ~ ( d pto
e ) ~the trans-[Cr(C0)2(dpe)2][X] salt can be probed by a
combination of voltammetry and specular reflectance IR spectroscopy [15].
' ~ d a ~ t ewith
d permission from J. Amer. Chem. Soc. 115 (1993) 9556; Organometallics,13
(1994) 5 122; J.Electvoanal. Chem. 404 (1996) 227. Copyright,American Chemical Society
and Elsevier.
346
Solid-electrode-solvent intefaces
Fig. 5 -10 Structural representation of (a) cis- and (b) trans-Cr(CO)z ( d ~ e ) ~ .
3.2.1 Structural aspects and solution-phase voltammetry
Cr(CO)2(dpe)2,unlike decamethylferrocene, has cis and trans isomeric structural forms (Fig. 5.10). After chemical or electrochemical oxidation, a wide
range of stable trans-[C~-(Co)~(dpe)~]
[XI (X- = anion) salts may be isolated, but the cis analogue of the oxidized form is unknown as a solid. Each
and trans-[Cr(C0)2(dpe)2][X] are soluble in
of cis- and tran~-Cr(CO)~(dpe)~
non-polar organic solvents and the organic solution-phase voltammetry of this
chromium carbonyl system has been studied in considerable detail [16-181.
If t r a n s - [ ~ r ( ~ ~ ) ~ ( d pise present
) ~ ] + in bulk solution, then the one-electron
reversible reduction process
+ e- +trans-Cr(C0)2(d~e)~(solution)
($91
trans-[~r( ~ (dpe)2]'
0 ) ~(solution)
(5.4)
is observed.
As expected, if tran~-Cr(CO)~(dpe)~
is in bulk solution, then the reverse of
eqn (5.4) occurs:
1
tran~-Cr(C0)~(dpe)~(solutiotl)
,
&t r a n s - [ C r ( ~(dpe)2]+
~ ) ~ (solution)
+ e(5.5)
However, if ~ i s - - C r ( C O ) ~ ( d is
p ein) ~bulk solution, the more complex reaction
scheme given in eqn (5.6) is operative on the voltammetric time-scale.
+c i s - [ ~ r ( ~ ~ ) ~ ( d ~ e ) ~ ] ~ ( s+o le-u t i o n )
( ~ 3 2
cis-Cr(CO)2(dpe)2(solution)
J, fast
(q') 1
+
trans-Cr(C0j2(dpe)2(solution) Ff trans-[~r( ~ (dpe)2]+(sol~ti~n)
0 ) ~
e(5.6)
3.2.2 Voltammetry of solid trans- Cr(CO)z(dpe)z attached to a graphite electrode
Since many aspects of the solid-state t r a n s - [ ~ r ( ~ ~ ) ~ ( d p e )process
~]~"[12,14,15] are closely related to those of the [ F ~ ( ~ I ~ - C ~ M ~couple
~)]O/'
[I 11, only facets of the voltammetry of solid trans-Cr(C0)2 (dpe)z and
Mechanistic aspects of transport processes
347
t~ans-[Cr(CO), (dpe)2][XI adhered to a graphite electrode that require amplification or have not been featured in Section 3.1 need to be presented in any detail.
Electrolyte concentration dependence
~oltatnmo~ram
obtained
s
with rnicroparticles of trans-Cr-(C0)2(dpe)2attached
to a basal-plane graphite electrode and placed in aqueous NaCl solution at
30°C are shown as a function of electrolyte concentration over the range of
0.1-3.5 M in Fig. 5.11. This series of experiments leads to the conclusion that the
anion concentration, while not strongly influencing the current magnitude, does
have a small influence on the peak potentials, which approximates a Nernstian
shift of 2.303RTIF volt per decade change in concentration. If the value of
(E,"" ~ ; ~ ) /is2 assumed to be the reversible potential, then this result is as
ected6 for an equation of the kind given in eqn (5.7, assuming the activity
+
E (V) vs Ag/AgCl
E (V) vs Ag/AgCl
5.11 Cyclic voltammograms obtained at 30° C and with a scan rate of 200 mV s-' for oxidation
of trans-Cr(CO)z(dpe)z adhered to a basal-plane pyrolytic graphite electrode and then placed in
contact with (a) 0.1 M, (b) 0.3 M, (c) 1.0 M, and (d) 3.5 M aqueous NaCl electrolyte. Reproduced
by courtesy: Organometallics 13 (1994) 5 122. Copyright, American Chemical Society.
Nernstian form of dependence of potential on electrolyte concentration has been
detected with a range of solid-state systems, but exceptions to this result are also commonly observed, so that even the thermodynamics of the processes are not always well
established 141.
6~
Table 5.2 Cyclic voltammetric dataa obtained at 50°C for oxidation of microcrystalline trans-Cr(CO)z(dpe) adhered to a basal
pyrolytic graphite electrode placed in contact with aqueous
(electrolyte) media as a function of electrolyte anion.
OHF-
sotc1BrNO,
IClO,
~ 1 0 , ~
~r~
"Data taken from Organornetalliis 13 (1994) 5122. Symbols E r and
with
units of mV versus Ag/AgCl, and AEp are as defined in Table 5.1. Ep values are
obtained by extrapolation to zero current of data obtained at low scan rates.
b ~ h aqueous
e
electrolytes used were 0.1 M NaOH, NaF, NaC1, KBr, KI,
Na2S04,NaN03, and NaC104.
' A G p = Gibbs energy of partition of anions for a water-1,2-dichloroethane
interface [191.
20 cycles by oxidation of cis-Cr(CO)z(dpe)a.
d ~ e n e r a t e after
d
of the solids is unity, which of course need not be the case (see later).
trans-Cr(C0)2( d ~ e(solid)
)~
-
A
+ C1- (solution)
trans-[Cr (C0)z( d ~ e )[Cl]
~ ](solid)
+ e-
(5-7)
Electrolyte anion and cation dependence
The oxidation of solid trans-Cr(CO)z(dpe)2as a function of different electrolyte
anions has been studied in detail at 50°C by cyclic voltammetry at a range of
scan rates [12]. Under these conditions, changing the anion from perchlorate
to hydroxide produces a potential shift7 of almost 500 mV (Table 5.2). A large
7 ~20-30°C,
t
as used for most studies, currents are difficult to detect above the background
value with some electrolytes. At 50°C used for these comparative studies, voltammetric
responses are always very well defined. An additional process also may be detected at
elevated temperature, at high surface coverages and fast scan rates. When peak splitting
occurs, or more than one process is detected, the peak potentials cited correspond to the
dominant signal pair. The influence of different electrode coverage and particle sizes in these
anion and cation dependence studies was addressed by extrapolation of peak potential data
obtained as a function of scan rate to zero current and using these zero current potentials
as the basis of comparison [15].
Mechanistic aspects
oftransport processes
349
electrolyte anion effect was also observed for oxidation of decamethylferrocene.
The comparison ofpeak potentials with the Gibbs energies for partition of anions
[I 91 at a 1,2-dichloromethane-waterinterface (Table 5.2), suggests that the free
energy of transfer of the particular anion from solution into the solid phase is an
ortant term in determining the peak potential. The influence of cations on
~otentialof voltammetric response for oxidation of microparticles of solid
t~ans-Cr(CO)~(dpe)~
attached to graphite electrodes, as was the case with the
decamethylferrocene system, was not significant [12].
Electvochemical quartz cryrtal microbalance studies
The EQCM technique has been highly useful in the investigation of surlacebased electrode processes associated, for example, with monolayers [20],
multilayer deposition and dissolution [21], mass transport in polymer films
[22], corrosion processes [23], and electroless deposition and mass changes
caused by adsorption [24] (see Section 19 in Chapter 2 for details). Not sur~]+'~
prisingly, important mechanistic aspects of the t v a n ~ - [ C r ( C O ) ~ ( d ~ e )redox
reaction related to the charge neutralization process also can be probed conveniently via studies of the oxidation of tran~-Cr(CO)~(dpe)~
and reduction of
t~ans-[Cr(CO)~(dpe)~]
[XI microparticles attached to a gold Q C electrode and
use of the EQCM method [12].
icrocrystalline tran~-Cr(CO)~(dpe)~
was attached to the Q C electrode by
ng a cotton bud containing a small amount of the solid onto the gold
electrode. The trans-[Cr(C0)2(dpe)2][X] salts were not mechanically
ached to the gold Q C in this manner because of a tendency of the gold
er to be scratched from the Q C during the rubbing process. Thus, the
tvan~-[Cr(CO)~(dpe)~]
[XI salts were attached to the electrode by dissolving the
solid in acetone and then placing a drop of this solution onto the electrode.
Arrays of microcrystalline deposits of trans- [Cr (C0)2(dpe)2][XI were formed
gold electrode after the solvent was evaporated.
n order to obtain well-defined voltammograms, an adequate coverage
rocrystalline particles must be achieved. A scanning electron microobtained for solid t~ans-Cr(CO)~(dpe)~
attached to a gold Q C electrode
ns crystalline particles of approximately 20 pm in length (Fig. 5.12).
ough trans-[Cr(CO)2(dpe)2][Cl] attached to a gold Q C produced microtals which, on average, had a smaller surface area relative to those for attached
tral compound (compare Fig. 5.12(a) and (b)),particles were detected with
lengths ranging from 0.1 to 20 pm.
CM studies ofsolid tran~-Cr(CO)~(dpe)~
attached to a gold electrode
igure 5.13(a) shows a cyclic voltarnrnogram obtained at a scan rate of 50 mV s-'
for the oxidation of microcrystalline particles of t~ans-Cr(CO)~(dpe)~
mechanically attached to a gold Q C electrode over a potential range of -800-600 mV
versus Ag/AgCl in 0.1 M KC1 electrolyte. A well-defined process was observed
which will be shown to correspond to the general process given in eqn (5.8),
Fig.
5.12 Scanning
electron
micrographs
of
(a)
tran~-Cr(CO)~(dpe)~and
(b) trans-[Cr(CO)z(dpe)2]Cl attached to a gold QC. Reproduced by courtesy: J. Electroanal.
Chem. 404 (1996) 227. Copyright, Elsevier.
but where X - , the electrolyte anion, is Cltrans-Cr (CO), (dpe), (solid)+X- (solution)
Figure 5.13(b) shows the relative mass change that accompanies the oxidation and reduction processes at the gold-coated Q C electrode as a function
of potential. The mass change data indicate that the oxidation of solid
tran~-Cr(CO)~(dpe)~
is accompanied by an increase in mass which is consistent
with incorporation of chloride ions from the 0.1 M KC1 electrolyte. However,
the mass continues to increase even on the reverse scan of the cyclic voltammogram, until the potential is reached where solid trans-[Cr ( C 0 ) 2(dpe)z]Cl
(formed via oxidation in the first half-cycle) is reduced. Thus, oxidation of
microcrystalline trans-Cr(CO), (dpe)? in 0.1 M KC1 solution, which requires
incorporation of C1- into the lattice, is found to be a very slow process. In
Mechanistic aspects of tvanspovt processes
-800
-600
-400 -200
0
200
Potential (mV) vs Ag/AgCl
400
35 1
600
Potential (mV) vs Ag/AgCl
-13 (a) Cyclic voltammogram (scan rate 50 rnV s-') and (b) the corresponding mass change
diagram for oxidation of microcrystalline particles of t r a n ~ - C r ( C O ) ~ ( d p eattached
)~
to a gold
Q C which has been placed in contact with aqueous 0.1 MKC1 electrolyte. Reproduced by
courtesy: J. Electroanal. Chem. 404 (1996) 227. Copyright, Elsevier.
contrast, the expulsion of chloride ions, which accompanies the reduction
rocess on the reverse scan of a cyclic voltammogram, is associated with both a
sharp peak and a rapid decrease in mass. The EQCM diagrams therefore imply
that the rate-determining step in the oxidation of solid t~ans-Cr(CO)~(dpe)~
in
contact with 0.1 M KC1 electrolyte is the very slow reaction involving incorporation of anion into the solid. This explains why the oxidation peak current is
relatively small and the oxidation current fails to decay to zero on the time-scale
of the cyclic voltammetric measurements.
Figure 5.13(b) also shows that an overall increase in mass is evident after completion of each potential cycle involving the oxidation of trans-Cr(CO)2(dpe)2
and then its reformation by reduction. This result implies that some chloride
ions inserted via oxidation remain inside the solid on this time-scale even after
the reductive electrochemical treatment. Presumably, the reverse reductive scan
leads to rapid expulsion only of those chloride ions that are incorporated at or
352
Solid-electrode-solvent intefaces
Table 5.3 Theoretical and experimentally determined molar
masses (M) obtained from EQCM experiments when solid
tran~-Cr(CO)~(dpe)~
is adhered to a gold electrode which is placed in
contact with different electrolytesa
Electrolyte (0.1 M)
Experimental (M/g mol-')
heo ore tical ( M / mol-')c
~
aData taken from J. Electroanal. Chew. 404 (1996) 227; scan rate = 50 mVsC1.
b ~ r r oisr standard deviation (10 experiments using three different gold QC electrodes).
Calculations are based on the reduction process observed during the reverse scan of cyclic
voltammograms.
Talues are calculated assuming that only the non-solvated electrolyte anion is associated
with the charge neutralization process.
near the surface of the solid on the time-scale of cyclic voltammetric experiments. With continuous cycling of the potential, the mass changes associated
with oxidation and reduction processes become constant.
Quantitative analysis of the microbalance data in terms of charge and mass
change relationships enables the molar mass of the anion expelled from the solid
to be calculated. Since it is very difficult to calculate accurately the mass change
from the oxidation reaction, because the process occurs over a wide potential
range, all molar mass data have been calculated from the better defined and
sharper reduction process. Table 5.3 contains the experimental and theoretical
molar masses of the species being expelled from trans- [Cr (CO)z( d ~ e )[XI
~ ](solid)
formed by oxidation of trans-Cr(CO)2(dpe)2(~olid)
when the electrode is in contact with a variety of electrolytes. Clearly, the species being expelled in all
experiments corresponds to the electrolyte anion. Electron microprobe X-ray
analysis experiments analogous to those shown in Fig. 5.6 for decamethylferrocene also support the hypothesis that the oxidation of tran~-Cr(CO)~(dpe)~
involves the uptake of the anion from electrolyte solution, and the expulsion
of the anion during reduction.
Figure 5.14 shows the result of a double-potential step experiment in which
(dpe)2 attached to a gold electrode
microcrystalline particles of tran~-Cr(CO)~
were alternatively oxidized and then reduced in the presence of 0.1 M KC1 solution. The potential was initially held at -600 mV versus Ag/AgCl. After 1000 s,
the potential was stepped from -600 to 0 mV (Fig. 5.14(a)),and this resulted in
to trans-[Cr(C0)2 (dpe)d [Cl](solid).
the oxidation of trans-Cr(C0)2(dpe)2(~~lid)
Figure 5.14(b) shows the very rapid development of an oxidation current
which subsequently decayed rapidly and then slowly as a function of time.
Figure 5.14(c) shows that there was also an increase in mass for about 6000s,
after which time there was no further mass change (or apparent further oxidation
of the neutral compound).
Time (s)
0
5000
10000
Time (s)
15000
20000
Time (s)
ig. 5.14 A double-potential step experiment for microcrystalline particles of trans-Cr(CO)a(dpe)a
attached to a gold Q C which was then placed in contact with aqueous 0.1 MKC1 electrolyte.
(a) potential-time sequence; (b) current-time data; (c) relative mass change as a function of time.
Reproduced by courtesy: J. Electroanal. Chem. 404 (1996) 227. Copyright, Elsevier.
After 10 000 s, the potential was stepped back to -600 mV versus Ag/Agcl,
in order to reform neutral tvans-Cr(C0)2(dpe)2(solid).The mass decrease was
indicative of the expulsion of chloride ions from the crystal lattice. However,
even on this time-scale, not all of the chloride ions incorporated into the microcrystalline particles of tvan~-Cr(CO)~(dpe)~
during oxidation appear to have
been expelled from the solid during the reverse process. The double-potential
step data again suggest that, while the slow oxidation process involves the conversion of trans-Cr(C0)2 (dpe)2(solid) to trans- [Cr (CO)2(dpe)2][XI (solid), not all
the neutral tvans complex is recovered during the relatively rapid component of
the reverse reduction process.
[XI salts attached to a gold electvode
EQCM studies of solid tran~-[Cv(CO)~(dpe)~]
Figure 5.15(a) shows a cyclic voltammogram over the potential range 400 to
-900 mV versus Ag/AgCl for the reduction of solid tvans-[Cr(CO)2(dpe)2][Cl]
Potential (mV) vs /AgC1
-1000
-800
-600
-400
-200
0
200
400
Potential (mV) vs Ag /AgCl
Fig. 5.15 (a) Cyclic voltammogram (scan rate 50 mV s-l) and (b) the corresponding mass change
diagram for microcrystalline particles of trans-[Cr(C0)2(dpe)z][Cl] attached to a gold QC which was
then placed in contact with 0.1 M KC1 aqueous electrolyte. Reproduced by courtesy: J Electroanal.
Chew. 404 (1996) 227. Copyright, Elsevier.
Mechanistic aspects of transport processes
355
in 0.1 M KC1 solution at a scan rate of 50 mV s-' . In this particular experiment,
the compound and electrolyte anions are the same, although this of course need
not be the case.
The well-defined reduction response at -400 mV versus A g / ~ g l
is due to the reduction of tran~-[Cr(CO)~
(dpe)2][Cl](solid) to neutral
tmns-Cr(C0)2 (dpe)2(solid). The plot of relative mass as a function of potential (Fig. 5.15(b))indicates that no mass change occurs until reduction current is
observed. The mass decrease at potentials more negative than -400 mV versus
A ~ / A is
~ due
C ~to the reduction of trans-[Cr (C0)2(dpe)2][Cl](solid) to neutral
trans-Cr(C0)2(dpe)2(solid)with C1- ions being expelled from the crystal lattice.
n the reverse scan, the oxidation peak for trans-Cr(C0)2(dpe), at -100 m~
versus Ag/AgCl also is smaller than the reduction peak (as previously obsemed
13) when using solid trans-Cr(C0)2(dpe)2.
a potential of -300 mV is reached on the reverse scan, there is a conincrease in mass due to the reformation of the chloride salt. The mass
ontinues to increase until a potential of 400 mV versus Ag/AgCl is reached, and
resumably would continue to increase at more positive potentials. However,
in order to prevent gold electrode oxidation, the potential range was limited
the
to 400 mV. As was the case when starting with trans-Cr(CO)2(dpe)2(~ahd),
oxidation process was found to be very slow compared with the reduction
of trans-[Cr(CO)2( d ~ e )[Cl]
~ ] (solid). Other trans-[Cr(CO)2(dpe),] [XI salts were
also investigated in a variety of electrolytes and under circumstances where the
solid anion may or may not be the same as the electrolyte anion. The molar
masses of the ions expelled from the salts during the reduction process were
calculated and the values are listed in Table 5.4. The results suggest that, in cases
where the salt anion is different from the electrolyte anion, that an ion-exchange
.4 Theoretical and experimentally determined molar
masses (M) obtained by EQCM experiments when solid
trans-[Cr(CO)2(dpe)2][X] is attached to a gold electrode
which is placed in contact with different electrolytesa
X in trans-[Cr(CO)2
(dpehl [XI
Electrolyte
(0.1 M)
Experimental
(M/g rnol-)
Theoretical
( M / mol-I)'
~
'Data taken from J. Electroanal. Chem. 404 (1996) 227; scan rate 50 nlV s-' .
' ~ r r o ris standard deviation (six experiments using three different gold Q C
electrodes). Calculations are based on the reduction process observed in cyclic
voltammetric experiments.
'Values are calculated assuming that only the non-solvated electrolyte anion
is associated with the charge neutralization process.
356
Solid-electuode-solvent intefaces
reaction takes place prior to the reduction of tuan~-[Cr(CO)~(dpe)~]
[XI. That
is, in these cases, the molar mass for the reduction of trans-[Cr(CO)2(dpe)2][XI
to tuan~-Cr(CO)~(dpe)~
is not equivalent to the molar mass of the anion in the
salt. Rather, the experimentally determined molar mass approximates that of the
anion in the electrolyte. For example, reduction of trans-[Cr(C0)2(dpe),1[Br]
in 0.1 M KBr has an experimental molar mass of 73.1 g mol-' , which correlates with the expulsion of bromide ions from the crystal lattice. However,
reduction of initially tuans-[Cr(C0)2(dpe)2][Br](solid) in contact with 0.1 M KC1
electrolyte produces an experimental molar mass of 27.8 g mol-I, instead of
79.9 gmol-l which implies that tuans-[Cr(C0)2(dpe)z][Cl](solid)was formed
rapidly by an ion-exchange process and that the chloride ion and not the
bromide ion was expelled from the solid during the reduction step.
Results obtained from electron microprobe X-ray analysis experiments also indicate that an ion exchange reaction occurs prior to
electrolysis. Figure 5.16(a) represents the electron microprobe analysis of
trans-[Cr(C0)2(dpe)~[Cl](solid)and the expected P, Cr, and C1 peaks are
Fig. 5.16 Electron microprobe X-ray analysis of (a) solid trans-[C~(CO)~(dpe)~][cl]
attached to
a graphite electrode, (b) solid trans-[Cr(CO)2(dpe)2][C1] attached to a graphite electrode which
was held at 200 mV versus Ag/AgCl for 2 min in 0.1 M KBr aqueous electrolyte. Reproduced by
courtesy: J. Electroanal. Chem. 404 (1996) 227.
evident. trans-[Cr (CO)2(dpe)2][Cl](solid) was then attached to a graphite electrode which was then dipped into a 0.1 M KBr electrolyte for 2 min with the
electrode being held at 200 mV versus Ag/AgCl to prevent any redox processes
horn occurring. The electron microprobe analysis shows the disappearance of
C 1 ions from the solid and the incorporation ofBr- ions (Fig. 5.16(b)).These
results are consistent with the occurrence of an ion-exchange process. However, it was also observed that larger particles of trans-[Cr(CO)2(dpe)2][Cl]
(above 5 pm) are not completely converted to trans-[Cr(C0)2(dpe)2][Br].
quation (5.9) summarizes the ion-exchange process that is believed to occur
at the electrode surface at open-circuit potential:
tm~s[Cr(CO)2 (dpe)21 [Cl](solid)
+ Br (solution)
-
A double-potential step experiment in 0.1 MKC1 solution is shown for
reduction of trans-[Cr (CO), (dpe)2][Cl](solid) in Fig. 5.17. The potential
was initially held at -20mV versus Ag/AgCl for 300 s. At this potential,
trans- [Cr(C0)2(dpe)2][Cl] is stable in the redox sense. The absence of any mass
change confirms that there was no Faradaic process during the course of this part
of the experiment (see Fig. 5.17(c)).After 300 s, the potential was stepped from
-20 to -600 mV versus Ag/AgCl and held at the latter potential for 2000 s. At
-600 mV versus Ag/AgCl, the salt was reduced to the neutral compound and
the charge-transfer process was accompanied by a very rapid mass decrease due
to the expulsion of chloride ions from the salt, as shown in Fig. 5.17(c). A constant mass was achieved after 60 s and no further detectable mass change occurred
during the remainder of the experiment. The current response shows analogous
behaviour to the mass change (Fig. 5.17(b)). After a total time of 2300 s had
elapsed, the potential was then stepped back to -20 mV versus Ag/AgCl for
2000 s which produced an oxidation current that decayed rapidly as a function
of time, and an increase in mass due to the reincorporation of chloride ions into
the solid. The oxidation of trans-Cr(C0)2(dpe)2(so~id)to achieve the starting
mass condition for trans-[Cr(CO)2(dpe),] [Cl](solid) took approximately 1300 s.
This experiment again confirms that the incorporation of chloride ions is slower
than the expulsion process. The major difference with the potential step experiments for trans-[Cr (CO), (dpe)2][Cl](solid) and trans-Cr(C0)2(dpe)2(solid) was
that C1- ion expulsion and uptake proceed to the same extent in the former case.
The fact that reduction of tran~-[Cr(CO)~(dpe)~]
[Cl], which is attached to the
electrode by evaporation of a sample dissolved in a drop of organic solvent, seems
to lead to a more extensive level of electrolysis than when this compound is
may
formed in situ by oxidation of mechanically attached t~ans-Cr(CO)~(dpe)~,
be attributed to the fact that the particle size of solid attached to the electrode
using a solvent evaporation technique is smaller than is achieved with direct
attachment of solid trans-Cr(C0)2(dpe),.
Time (s)
Fig. 5.17 A double-potential step experiment for solid trans-[Cr(C0)2(dpe)2][C1]attached to a
gold electrode which was then placed in contact with 0.1 M KC1 aqueous electrolyte. (a) Potential,
(b) current, and (c) relative mass change as a function oftime. Reproduced by courtesy: J. Electvoanal.
Chem. 404 (1996) 227. Copyright, Elsevier.
Mechanistic aspects of transport processes
359
3.2.3 Voltammetric and FTIR studies on solid cis-Cr(C0)2(dpe)2
attached to a graphite electrode
The eighteen electron c i ~ - C r ( C O ) ~ ( d p is
e ) ~known to be the thermodynamically favoured isomer. In contrast, observation of isomerization of
the transiently formed ci~-[Cr(CO)~(dpe)~]+
reveals that the seventeen electron tran~-[Cr(CO)~(dpe)~]+
cation is strongly favoured in both thermoynamic and kinetic senses [16]. Accordingly, the solution phase voltammetry of c i ~ - C r ( C O ) ~ ( d pis
e ) governed
~
by eqn (5.6) with the short lived
c i s - [ ~ r ( C O ) ~ ( d p e )cation
~ ] + being difficult to detect spectroscopically (Table
.5). For convenience, the solid-state voltammetry may be written in shorthand
form as
cisO(solid)4 cis+(solid)
I
trans+(solid)
where ciso = solid c i ~ - C r ( C O ) ~ ( d pcis'
e ) ~ ,= solid c i s - [ C r ( C ~ ) ~ ( d p e )and
~]+
and the electrolyte anion required for
trans' = solid tran~-[Cr(CO)~(dpe)~]+
charge neutralization has been omitted.
In Fig. 5.18, and in accordance with eqn (5.lo), the voltammetric response
obtained at 50°C when cis-Cr(C0)2(dpe)2(solid)attached to a graphite electrode
?
Cvcle 3-1 0
J.
E (V) vs Ag/AgCl
Fig. 5.18 Cyclic voltammograms obtained at 50°C with a scan rate of 200 mV s-' over 10 potential
cycles for oxidation of solid cis-Cr(CO)2(dpe)2, mechanically attached to a basal-plane pyrolytic
graphite electrode and then placed in contact with aqueous 0.1 M NaCl electrolyte. Reproduced by
courtesy: Organornetallics 13 (1994) 5122. Copyright, American Chemical Society.
360
Solid-electrode-solvent intefaces
is in contact with aqueous 0.1 M NaCl can be seen to change from an initial
response associated with the oxidation of the cis isomer, to that for the solidstate t r a n s - [ ~ r ( ~ ~ ) ~ ( d ~redox
e ) ~ ]couple
~ / + described in Section 3.2.2. The
voltammogram of the solid-state trans+I0 process is in fact clearly evident after
2 cycles (Fig. 5.18). With 0.1 M NaC104 electrolyte, the data obtained after
10 cycles when commencing with solid ~ i s - C r ( C O ) ~ ( d pise )indistinguishable
~
from that obtained with solid trans-Cr(C0)2 (dpe)2 (Table 5.2).
A useful technique to monitor the course of processes that take place on
an electrode surface is specular reflectance FTIR spectroscopy. Good FTIR
signal-to-noise ratios are possible after about 5-40 scans of the I R spectrum
when solids are attached as arrays of microcrystals to basal-plane graphite electrodes because the modified graphite surface still reflects sufficient light intensity.
Effective background subtraction is possible by measuring the initial response
prior to the start of electrolysis and then plotting only the relative change in
intensity of the I R absorbance during the course of electrolysis. An I R spectrum obtained after electrochemical oxidation of solid trans-Cr(CO)2(dpe),
to solid trans-[~r(CO)~(dpe)~]+
cation at a graphite electrode in contact
with aqueous 0.1 M NaCl is shown in Fig. 5.19(a). In this experiment, the
tran~-Cr(CO)~(dpe)~
isomer with an I R band at 1796 cm-* (compare diffuse
reflectance KBr powder I R spectrum 1783 cm-') is converted into the trans+
cation form which has an I R band at 1846 cm-' (compare diffuse reflectance
KBr powder I R spectrum 1845 cm-') when oxidation is undertaken at 500 mV
versus Ag/AgCl.
The spectroscopic identification of the electrochemical generation of solid
cis-[Cr(C0)2 (dpe)21f via oxidation of solid cis-Cr(C0)2 (dpe)2 and proof of
the subsequent isomerization step which generates trans-[Cr (CO)2(dpe)2]+also
is achieved by use of in situ specular reflectance FTIR spectroelectrochemical
experiments (Fig. 5.19(b) and (c)). After oxidation of solid cis-Cr(CO)2(dpe)2
attached to an electrode at 500 mV versus Ag/AgCl, the detection of new I
bands at 1846 and 1933 cm-' shows that some ofthe ciso starting material (diffuse
reflectance I R bands for cis-Cr(C0)2(dpe)2(solid)occur at 1780 and 1844 cm-')
has been oxidized. The I R band at 1844 cm-' corresponds predominantly to
formation ofthe transf isomer. The I R band at 1933 cm-I (Table 5.5) is attributable to the formation of the oxidized cis+ isomer, which is too unstable to be
detected in solution-phase spectroelectrochemical studies. The second I R band
expected for cis' is unresolved from the trans+ band (Table 5.5). The shift of
89 cm-' for comparable I R bands of the cisf species relative to the cis species
(Table 5.5) is of the order expected for the formation of cis+ and is associated with the increased positive charge on the metal centre. The I R band at
1933 cm-' shown in Fig. 5.19(b) decays slowly (50 per cent after 60 rnin) when
monitored at room temperature. Evidently, constraints posed by the solid matrix
allow the achievement of kinetic stabilization of a thermodynamically unstable
isomer. The I R detection of the trans+ isomer after a potential of 500 mV versus
Ag/AgCl is applied for 2 min is achieved with considerably decreased spectral
overlap with the cis+ isomer (Fig. 5.19(c)).Thus, while the processes leading to
generation of cisf and subsequent isomerization to transf may be complex, it is
-
Mechanistic aspects of transport processes
36 1
trans
. 5.19 Specular reflectance FTIR spectra obtained for cis- and trans-Cr(C0)a ( d ~ emechanically
)~
attached to a basal-plane pyrolytic graphite electrode, after being placed in contact with aqueous
0.1 M NaCl electrolyte, electrochemically oxidized at ambient temperature, and then dried in air after
removal of the electrolyte solution: (a) trans-Cr(CO)a(dpe)a afier the potential was held at 500 mV
versus Ag/AgCl for 2rnin; (b) c i ~ - C r ( C O ) ~ ( d p eafier
) ~ the potential was held at 500mV versus
Ag/AgCl for 5 s; (c) the same system as in (b) after 2 min. The spectra were background-corrected
by subtraction of the response prior to electrolysis. Reproduced by courtesy: Oqanornetallics 13
(1994) 5 122. Copyright, American Chemical Society.
clear that the basic mechanism has a solid-state parallel to that identified in the
solution-phase voltammetry. A summary of surface attached and other forms of
spectra is contained in Table 5.5.
3.3 Overview offactors that influence the voltammetry of
decamethylferrocene and trans-Cr ( C O ) 2
attached to an
electrode suface
The data obtained for the voltammetry of solid decamethylferrocene and
trans-Cr(CO)z(dpe)2, which both involve an overall one-electron oxidation
process to generate a close to isostructural cation, show that: (a) a wide range of
responses can be observed from microcrystalline samples of the non-conducting
solid which depend on the crystal size and surface coverage, and also the
medium in contact with the solid; (b) voltammetric wave shapes may have
Table 5.5 I R data obtained in the carbonyl region (V (CO))for cisand tvans-[cr ( ~ (dpe)2]o/+
0 ) ~ species in solution (CH2C12),in the
solid state by the diffuse reflectance method (KBr powder) and by
the specular reflectance method before and after electrolysis when
solid is adhered to a basal-plane pyrolytic graphite electrode
Technique (medium)
v (CO) (cm-')
trans
trans+
cis
Solution-phase (CH2C12)
1792
1850
Diffuse reflectance (KBr powder)
1783
1845
Specular reflectance (graphite electrode)
1796
I846
1846
1772
1830
1764
1844
1780
cisS
1933
"Data taken from Organometallics 13 (1994) 5122.
b~somerizationto transsi to rapid for detection.
'Second band not resolved from trans' band.
different characteristics, depending on the electrolyte, solvent, and scan rate; (c)
a large separation between oxidation and reduction peak potentials may occur;
(d) oxidation and reduction components of the processes associated with cyclic
voltammograms may exhibit different characteristics; (e) anions may be incorporated from the electrolyte into the solid during oxidation and expelled during
reduction.
The primary questions to resolve are how the electron-transfer step occurs
in the oxidation of these non-conducting solids and how the mass-transport
mechanism of anions occurs to achieve charge neutrality. The microcrystalline
material has small areas where simultaneous contact with both the electrode
and solution occurs. Starting from the three-phase boundary contact position
(Fig. 5. I), it is proposed that a current flow occurs via a coupled electron-transfer
self-exchange and ion-exchange charge neutralization mechanism within the
solid or along the solid-solution boundary. Figure 5.20 gives a schematic form
of the mechanism that may apply for the t v a n s - [ ~ r ( ~ ~ ) ~ ( d ~couple.
e ) ~ ] ~The
/+
mechanism of surface conduction, which is accompanied by transport of ions
across a solid-solution interface would, as observed, be expected to depend
very much on electrolyte, temperature, and solvent. In the case of oxidation
of decamethylferrocene, a flow of charges up to the order of millicoulombs
is observed, which requires that oxidation of more than just the surface of
8 0 n some occasions a counter intuitive charge neutralization process may occur. For
example, in the oxidation of the [BU~N]+salt of [ C r ( C 0 ) 5 1 ] to give Cr(CO)51(solid),
when the modified electrode is in contact with perchlorate containing electrolytes, charge
neutralization may occur by C10& uptake into the solid phase rather than the expected
expulsion of Bu4N+ into the solution phase [25].
Mechanistic aspects of transport processes
363
Solution
C1-
K+
Electrode
Neutral
trans-Cr (CO)2(dpe),
Oxidation of
trans-Cr (CO)2(dpe),
[Cr (CO), (dpe),l C1
'
\
I
I
Reduction of
ig. 5.20 A schematic representation of the redox chemistry of solid particles of
trans-Cr(C0)2(d~e)~
and trans-[Cr(CO)2(dpe)2][C1] attached to a gold electrode which has been
placed in contact with 0.1 MKCI aqueous electrolyte. Reproduced by courtesy: J. Electroanal.
Chem. 404 (1996) 227. Copyright, Elsevier.
the rnicrocrystals occurs. This postulate is supported by electron probe X-ray
analysis experiments, which demonstrate that a substantial part of the surfaceattached decamethylferrocene is oxidized at positive potentials and reduced at
negative potentials.
Of particular significance is the observation that the cation in the electrolyte is not crucial to the voltammetry of solid decamethylferrocene or
t~ans-Cr(CO)~(dpe)~
whereas a wide range of voltammetric behaviour can be
found by varying the electrolyte anion. In these cases, this sensitivity to the
anion would appear to be associated with the fact that oxidation takes place with
anion uptake and reduction occurs with anion expulsion. The transport of ions
between two immiscible (solid-solution) phases is clearly a complex process,
but could be expected to depend on the free energy of anion exchange, that is
on the anion distribution between the two phases.
Solid-solution
interaction
t
........
Solid dissolves
....................... in electrolyte
Surface
conduction
'..'...-:...
...
... . ..'.. .
.....
,. . .
". .
.a. . . .
2 dim.
1 d m . '"......
3 dim?. .....
":.
.....Intercalation'.. ....
b
Temperature
Fig. 5.21 Schematic representation of some of the different phenomena related to ion transport that
may take place in the solid-state voltammetry of microparticles attached to an electrode and placed
in contact with a solvent (electrolyte) medium. Reproduced by courtesy: J. Electroanal. Chem. 372
(1994) 125. Copyright, Elsevier.
The peak splittings most commonly observed under conditions of high surface coverage and fast scan rate suggest that at least two processes may exist which
have different scan rate and temperature dependencies. Figure 5.21 schematically illustrates some oithe phenomena that are believed to occur with the ion
transport part of the mechanism associated with the redox chemistry of microcrystalline solids. However, clearly there are many features not predicted on the
basis of a thin film model considered in Section 18 in Chapter 2.
In the case of oxidation of microcrystalline trans-Cr(CO)z(dpe)2 to
trans-[Cr(CO)2(dpe)2]+, a variation of potential of more than 400mV is
observed when 0.1 M NaF is used instead of 0.1 M NaC104 as the electrolyte (Table 5.2). A similarly large anion dependence is found for oxidation of decarnethylferrocene (Fig. 5.8). The correlation of these reversible
potential data with free energies of partition for anion transport across a
water-dichloromethane interface (Table 5.2), coupled with voltammetric and
spectroscopic data indicate that the overall oxidation reaction for the oxidation
of trans-Cr(CO)z(dpe)z(solid) may be summarized by eqns (5.1I)-@. 13)
trans0(solid)
trans+ (solid) -k e-
X- (solution)
+X- (solid)
trans' (solid)
+ X- (solid) ---\ (trans' - X- (solid))
(5.11)
(5.13)
where X- denotes the electrolyte anion, and the trans0/+ nomenclature represents the trans isomer of the [ ~ r ( ~ ~ ) ~ ( d couple
~ e ) ] in
~ the
/ + appropriate
oxidation state, and where an analogous set of equations also probably applies
for oxidation of decamethylferrocene.
T o achieve the oxidation of the solid, both electron and ion transport
are necessary and the ions have to be transported to sites within the solid.
This may occur via diffusion of the ions within the solid, which is coupled
Mechanistic aspects of transport processes
365
with an electron-transport process involving electron hopping. Additionally, a
nucleation-growth mechanism may be needed to convert one solid phase to
another and this step or the coupled electron-transport-ion-transport process
may be rate determining.
Simultaneous electrochemical and Q C microbalance studies at the solidelectrode-aqueous solution (electrolyte) interface confirm that oxidation of
microcrystalline t~ans-Cr(CO)~(dpe)~
mechanically attached to a gold electrode
involves the formation of a t r a n s - [ C r ( ~ ~ ) ~ ( d ~salt
e ) via
~ ] +incorporation of
anions from the supporting electrolyte in order to maintain charge neutrality. O n the reverse scans of cyclic voltammograms and in the several stages of
double-potential step experiments, reduction occurs and anions are expelled.
owever, not all of the anions incorporated during the oxidation step are necessarily expelled from the crystal lattice on the time-scale of fast sweep cyclic
voltammetric experiments. This result suggests that the reduction reaction at
short time domains predominantly occurs in a spatial region relatively close to
the surface of the solid. Evidence also has been provided for a slower process,
which probably occurs at greater depths within the microcrystal.
.4 Problems with a theovetical descviption o f the
voltammetry of non-conducting microcr$als
-
The vast majority of solution-phase voltammetric experiments utilize conditions
with a well-defined electrode-solution interface consisting of just two phases.
n the solid-state electrochemical reactions considered so far, the process always
involves a multiphase system. In the most likely case, there will be four phases.
ne phase is the electrode, serving both as the electrical conductor and also
as a surface on which the solid sample particles are immobilized. The second
and third solid phases are the oxidized and reduced forms of the solid, although
this can be considered as a single phase if the oxidized and reduced forms
of the solid can exist as a continuous phase. The fourth phase is the solution
(electrolyte). The electronic conductivity of the sample phase will determine the
potential difference between the sample phase and the electrolyte solution. If this
conductivity is very high, the potential difference will be equal to that existing
between the electrode and the solution. In this case, the electrode process can,
if not otherwise inhibited, proceed on the entire surface of the solid and within
the bulk solid. Metal electrodes and adhered alloy rnicroparticles will most likely
react in this way [4].
However, when the conductivity of the sample is low, one may wonder
how an electrochemical reaction can proceed at all. Experimental evidence
and
available for oxidation of decamethylferrocene and tr~ns-Cr(CO)~(dpe)~
other systems (41 would indicate that the only possible initial reaction place is
the three-phase boundary between the electrode, the sample, and the solution (electrolyte) (Fig. 5: 1(b)). At this three-phase bound&y, electrons can
be exchanged between the electrode and the sample, and ions between the
sample and the solution (Fig. 5.22). In recent theoretical studies [26-281, the
Electrolyte solution
Insulator
Ion exchange
Electron exchange
h
Electrode
Fig. 5.22 Schematic representation of a simple three-phase electrochemical system supporting both
ion transfer between the solid compound and the solvent (electrolyte), and electron transfer between
the solid and the electrode. Adapted from: Electroanal. Chem. 20 (1998) 1.
propagation of an electrochemical reaction through a solid (the single phase
simplification) was considered for a case involving both electronic and ionic
conductivity of the solid. In this theoretical description, coupled diffusion of
electrons and ions within the crystal lattice was assumed and the redox reaction
was initiated at the three-phase boundary, where the solid was in contact with
both the electrode and the solution. From this contact point, the redox reaction
advances either along the surface of the solid or into the body of the solid by
diffusion of ions and the transport of electrons. This theoretical treatment of a
very simple model when a microparticle is attached to a solid electrode surface,
produced the following conclusions:
(1) The three-phase boundary is always the starting point for the reaction front,
independent of the geometry of the particle and its conductivity.
(2) The reaction will be surface confined if the diffusion of ions into the bulk
of the crystal is impossible or very slow.
(3) Generally, the net current is the sum of the current from the surface and
bulk processes.
(4) If both the surface and the bulk reactions proceed at comparable rates, the
reaction front expands in three dimensions from the three-phase boundary.
(5) The surface current will be negligible in cases where the bulk reaction is
dominating.
The voltammetric behaviour of azobenzene microcrystals is considered to
be an example where the electrode reaction is confined to the surface of
the solid materials [29], whereas the oxidation of decamethylferrocene and
tv~ns-Cr(CO)~(dpe)~
is clearly more complex and may occur via a range of
mechanisms.
A major problem to be encountered in a realistic theoretical description of
a solid-state electrochemical process is that the morphology of the sample distributed on the electrode surface contributes to the voltammetry and is often
complex. For example, microcrystals of the sample are often scattered on a solid
Voltammetuy of TCNQ
367
electrode and may have a certain size distribution as well as a preferential orientation due to the immobilization procedures. The microcrystals will expose
their different faces to the solution and also to the solid support so that in principle the polyciystallinity can decisively influence the electrochemical properties.
urthermore, the morphology can be potential dependent.
A major difficulty in providing a theoretical description also arises if the generation of an additional solid phase occurs with the direct conversion of one
solid phase into another, as will commonly be the case (see Section 4 below).
n this case, the number of phases varies during the course of an electrochenileal reaction, and the interfacial areas are no longer independent of time. The
transformation of one solid phase into another cannot always proceed along a
continuous series of mixed crystal phases as assumed in the simple theory considered above. A new phase is generated when the starting and the product
ses are unable to form mixed crystals at all potentials in the voltammetric
eriment. This situation can lead to a considerable splitting of the oxidation
reduction peaks in cyclic voltammetry (see reference [30] and the T C N Q
case considered below). However, it is also possible that a continuous transformation from one phase to the other occurs through mixed-phase formation
I], so that numerous possible scenarios are available in attempting to develop
eoretical models. The examples of voltammetry presented in the remainder
f this chapter represent cases where the rate-determining steps are nucleationrowth processes (Section 4) or coupled electron and ion transport within the
microparticle to give thin film or diffusion-like (Section 5) behaviour.
7,7,8,8 tetracyanoquinodimethane, (Fig. 5.23) usually designated as T C N
-, is well known as an
more particularly its one-electron reduced salt, T C
whereas pure T
example of an organic semiconductor [32,33]. Th
ssesses a rather low conductivity of 3 x 10-l2 S cm-I [34], the potassium
S cm-' [35], which is typ] has a conductivity of 2 x
nductor. Consequently, unlike the case with non-conducting
) ~ ] ~ /transport
+
of ions required
[ ~ e ( q ~ - ~ ~ ~ e ~ ) ~ ~] )~~"( -d ~ e systems,
(solid) process should be able to occur rapidly
to accompany the [
in three dimensions within the solid adhered to an electrode surface, even
with thick microparticles or films. Therefore, surface diffusion should not play
a major role in the voltammetry, and exhaustive electrolysis of all adhered
solid T C N Q should be rapidly achieved on the voltammetric time-scale. Not
surprisingly, because of the semiconducting properties of TCNQ- salts, the
368
Solid-electrode-solvent
intefaces
Fig. 5.23 Structure of 7,7,8,8 tetracyanoquinodimethane (TCNQ).
Process I1
I
I
-1.0
-0.5
I
I
I
0.0
0.5
1.O
E (V) vs Ag (s)1 AgCl (s)lKC1 (aq), 3 M
Fig. 5.24 Voltarnmograms for reduction of a 1 x lop3 M solution of T C N Q in acetonitrile
(0.1 M Bu4NPF6).(a) Steady-state voltamniograrn at a 10-pm diameter Pt microdisc electrode (scan
rate 10mVs-l). (b) Transient cyclic voltammogram at a 1-mrn diameter Pt macrodisc electrode
(scan rate 200 r n s-l).
~ Reproduced by courtesy: J. Chem. Soc., Faraday. Trans. 92 (1996) 3925.
Copyright, Royal Society of Chemistry.
[ T C N Q ] ~ /reaction is one of the most widely studied electrochemical systems
in both solution and solid-state phases. The wide variety of data available can
therefore be scrutinized to illustrate how numerous surface science and electrochemical techniques may be coupled with the electrochemistry in order to
identify the nuances that accompany the voltammetry of T C N Q in particular,
and probably the voltammetry of many other solid-state systems.
4.1
Solution-phase voltammetry of TCNQ
Typical solution-phase voltammograms obtained when T C N Q is dissolved in
an organic solvent such as acetonitrile are shown in Fig. 5.24. That is, two chemically and electrochemically reversible one-electron reductions are observed
under both transient (cyclic voltammetry) and steady-state (microdisc electrode) conditions [36] and they correspond to the reduction processes given
in eqns (5.14) and (5.15).
+ e- 4 TCNQ-(solution)
Process 11 TCNQ- (solution) + e- + TCNQ~-(solution)
Process I
TCNQ(so1ution)
(5.14)
(5.15)
Further reduction processes are also possible, but are not of interest in this
discussion.
r (dpe)2]0/+solidIn studies on the [ F ~ ( T , I ~ - c ~ M ~and
~ ) ~trans] O / +[ ~(co),
state redox reactions, the electrochemical responses have been attributed to
Voltammetry of TCNQ
369
reactions that occur rapidly on the surface of the non-conducting microcrystals
and more slowly within the bulk solid. However, an obviously distinguishing
feature of T C N Q is the conducting nature of the one-electron reduced salts
which should drastically aid the transport of ions required to achieve charge
neutralization within the bulk solid.
Voltamwzetvic studies o n micvocvystals of TCNQ adheved to electrode
su faces i n contact with N a + , K+, RD', and CS+containing electrolytes9
Q has been adhered to electrodes in a variety of ways. Initially, studng arrays of microcrystals are described [37,38] which are analogous to
eported above with ~ e ( q ~ - C ~ M
ande C
~ r)( ~
c O ) , ( d ~ e )Subsequently,
~.
voltammetric studies on other forms of adhered T C N Q will be considered.
n the studies described in 4.2, microcrystals of T C N Q were adhered to
the electrode surface as follows. First a sample of 1-3 mg of crystalline TCNQ
owder was placed on a coarse-grade filter paper. The electrode was then
gently rubbed into the powder until it became thoroughly coated with TCNQ.
uccessful coating was-evidenced by observing the shiny blue colour present
when the electrode was observed at a low angle. A cyclic voltammogram of
surface-immobilized microcrystals of T C N Q on a basal-plane pyrolytic graphite
electrode immersed in aqueous 0.1 M KC1 solution is shown in Fig. 5.25. As
for the voltammetry of ~ e ( q ~ - ~ and
~ Mtran~-Cr(CO)~(dpe)~,
e ~ ) ~
the nature
of the solid-state voltammograms differs considerably from those observed in
the solution phase (Fig. 5.24). In particular, an 'inert zone' appears between
the principal peaks. Figure 5.26 is a schematic form of representation of the
first reduction process in which the principal peaks (I,, and Ired) are equal
I
-1 .O
I
I
I
-0.5
0.0
0.5
E (V) vs Ag(s)IAgCl(s)II<Cl(aq),3 M
I
1.O
ig. 5.25 Cyclic voltammograms obtained at a scan rate of 100 mV s-* for microcrystals of TCNQ
immobilized on a basal-plane pyrolytic graphite electrode placed in contact with aqueous 0.1 M KC1
electrolyte (dashed: potential scanned only over region of process I). Reproduced by courtesy:
J. Chem. Soc., Faraday. Trans. 92 (1996) 3925. Copyright, Royal Society of Chemistry.
' ~ d a ~ t ewith
d permission from A.M. Bond, S. Fletcher, F. Marken, SJ. Shaw, and P.G.
Symons, J. Chem. Soc. Faraday Trans. 92 (1996) 3925 and A.M. Bond, S. Fletcher, and
P.G. Symons, Analyst 123 (1998). Copyright, Royal Society of Chemistry.
Fig. 5.26 Schematic representation ofprocess I (eqn 5.15) for reduction of microcrystals of TCNQ
adhered to an electrode.
as required for a chemically reversible reaction. The reduction process will
be shown below to be accompanied by insertion of cations from the electrolyte solution, [M+][X-] (solution), into the solid material to maintain charge
neutrality, implying an overall general reaction for process I of the form
xTCNQ(so1id)
+ ye- + y ~ + ( s o l u t i o n )
The second signal pair in Fig. 5.25 (process 11) corresponds to the further reaction
Significant chemical irreversibility is evident in process 11, which is partially
dianion dissolution. Consequently, because of this
accounted for by T C N Q ~
and other complications only process I is considered in the discussion that
follows.
The oxidation level of the T C N Q solid before and after the first reduction
process at a basal-plane pyrolytic graphite electrode in contact with 0.1 M KC1
can be quantitatively determined using the microelectrode method illustrated in
Fig. 5.27. In this experiment, the potential of the electrode containing adhered
solid T C N Q was held at either +0.50 or -0.10 V versus Ag/AgCl for 5 min.
The resulting material was then dissolved in a minimum amount of acetonitrile containing 0.1 M Bu4NPF6as supporting electrolyte, and the steady-state
voltammogram was recorded at a 10-pm diameter platinum microdisc electrode.
From analysis of the results presented in Fig. 5.27, with respect to the position of
zero current, it immediately becomes clear that the solid material present when
the potential is maintained at f0.50 V versus Ag/AgCl contains predominantly
neutral TCNQ, whereas the solid material present when the electrode potential is changed to -0.10 V versus Ag/AgCl contains predominantly an anionic
TCNQ- salt.
Voltammetry of TCNQ
371
F l v e d TCNQ-
I
-0.7
I
I
-0.1
I
I
0.5
E (V) vs Ag (s) I AgC1 (s)I KC1 (aq), 3 M
ig. 5.27 Steady-state voltammograms (obtained at a scan rate of 10 mV s-' using a 10-pm diameter
Pt microdisc electrode) of solid dissolved into minimal volumes of acetonitrile (0.1 M Bu4NPF6)after
holding the potential of a T C N Q coated basal-plane pyrolytic graphite electrode in contact with
0.1 MKCl for 5 min at different potentials. Top curve: TCNQ- generated in the solid state at
-0.10V versus Ag/AgCl. Bottom curve: T C N Q regenerated in the solid state at f O.5OV versus
Ag/AgCl. Reproduced by courtesy: J. Chern. Soc., Faraday. Tmns. 92 (1996) 3925. Copyright,
Royal Society of Chemistry.
-100
100
300
500
E (mV) vs Ag(s)l AgCl(s)l KC1 (aq), 3 M
Fig. 5.28 Test of the effect of the electrode material. Cyclic voltammograms obtained at a scan
rate of 100 mV s-' for T C N Q microcrystals immobilized on a 1-mm carbon electrode, a 0.5-mm
platinum electrode, a 0.4-mm gold electrode, and a RAMTMelectrode in contact with aqueous
0.1 M KC1 electrolyte. The currents are normalized to compensate for the different reacting masses.
It is evident that, within experimental error, the normalized responses are essentially independent of
electrode material. Reproduced by courtesy: Analyst 123 (1998) 1891. Copyright, Royal Society of
Chemistry.
T o test if the underlying electrode material had any effect on the voltammetry
of the T C N Q microcrystals, 1 mm GC, 0.5 m platinum, 0.4 mm radius gold
macrodisc electrodes and a RAMTMmicroarray electrodelo were used as substrates. In all cases the normalized responses (Fig. 5.28) were virtually identical,
' O ~ h eRAM^^ array electrode had a working surface of 28 mm2 and was inlaid with about
3200 carbon microdiscs (7 ym diameter) embedded in epoxy resin. About one-third of
the microdiscs were electrically connected, so the nearest-neighbour distance of active
microdiscs was approximately 70 ym. See references [39,40] for further details.
372
Solid-electvode-solvent intefaces
Table 5.6 Comparison of charge densities, critical potentials, inert zone widths
(qi) and peak shape parameters obtained at a scan rate of 100 mV s-' froln
cyclic voltammetric experiments with T C N Q microcrystals adhered to different
electrode materials in contact with 0.1 M KCla
Electrodes
GC
Parameter measuredb
500
Platinum
250
Gold
200
1000 microdiscY. 5
2600
10000
5500
6 900
224.5
215.5
228
228
24
24.5
27
29.5
200.5
191
210
198.5
33
35
31
23
27
42.5
27
24
RAM'"
"Data taken from reference [37].
' r is the electrode radius, q,d, charge density associated with reduction process, E$: and E:$, the
critical potential associated with the onset of oxidation and reduction processes, respectively, qi, the
width of inert zone between oxidation and reduction processes, AEiT4 and A E G ~are the widths at
three-quarters of wave height for oxidation and reduction processes respectively.
5 e e footnote 10 on previous page.
with voltammetric steady states achieved after 10 cycles being used as the basis
of comparison. It is therefore concluded that the underlying electrode materials
exert no significant effect on the voltammetry of the microcrystals. Table 5.6
summarizes numerical data derived from Fig. 5.28. For all four electrode materials, the charges passed in redox cycling experiments are considerably larger
than those expected from purely surface reactions, confirming that a significant
component of the redox reactions occurs inside bulk TCNQ(so1id) material.
Moreover, within experimental error, the critical potentials for the onset of the
redox reactions are also independent of the underlying electrode materials, so
this parameter also must be associated with phenomena occurring inside the
bulk of the solid TCNQ.
In order to undertake comprehensive studies on the influence of the electrolyte, studies at a RAM^^ electrode can be undertaken in a flowing solution
using the cell illustrated in Fig. 5.29. The major advantage of using a flowing
solution is that different electrolyte solutions can be rapidly exchanged and the
effects on the voltammetry at a given electrode noted under conditions where
the nature of the electrode-TCNQ(so1id) interface remains constant.
It can be noted that some T C N Q salts have a small degree of solubility
in water and that this may lead to a time-dependent voltammetric response.
The problems of the small solubility of T C N Q salts and mechanical stability of adhered crystals in flowing solution may be overcome by applying a
thin overcoat of sparsely cross-linked ~ a f i o n @
over the array of microcrystals
adhered to the electrode surface. The ~ a f i o n ' coating is thin, ionomeric, and
not highly cross-linked and as a result, it swells. An important consequence of
Voltammetry of TCNQ
Solut~onoutlet
373
-
RAM^^
working
electrode
Solution
inlet
/
/
Stainless steel
counter electrode
Teflon
spacer
\
EPOXY
resin
block
2 9 Schematic representation of a thin-layer flow cell used to study the electrolyte dependence
ofthe voltammetry of microcrystals of T C N Q adhered to a RAMTMelectrode. The stainless steel cell
body forms the counter electrode, while the RAMTMelectrode forms the working electrode. The
reference electrode is located downstream. This apparatus permits the rapid exchange of electrolyte
solutions without disturbing the electrode surface. Reproduced by courtesy: Analyst 123 (1998) 1891.
Copyright, Royal Society of Chemistry.
ese characteristics is that the NafionB coating does not act as a microporous
(electrolyte-excluding) ion-selective membrane, but instead acts as a macroporous (electrolyte-including) gel, preventing convection in the immediate
vicinity of the TCNQ(solid), but not preventing the solid from experiencing
e bulk concentration of cations. This means that the T C N Q microcrystals
neath the ~ a f i o n @
coating are bathed in the same concentration of KC1 as
t found in the bulk of solution, as required if the electrolyte dependence is
to be assessed.
Figure 5.30 shows voltammograms of T C N Q microcrystals immobilized on
afion-coated RAM^^ electrode in contact with 0.1 M aqueous solutions of
KCl, K N 0 3 , KBr, KF, and KI. There is very little effect of changing the identity
of the anion. Within experimental error, the critical potentials for oxidation and
reduction are independent of the identity of the anion. These results powerfully
suggest that the anions play no role in the electrochemical reaction, and justify
their omission from eqn 5.16.
By contrast, the identity of the cation and its concentration is crucial. If
the concentration of potassium chloride in the bulk of solution is increased in
decade increments from 0.001 M to 1 M, then the entire voltammetric response
shifts by about 60 mV decade-' (Fig. 5.31). Figure 5.32 shows the voltammetric
'
hi his result should be contrasted with the conclusion reached in studies on the oxidation
of decamethylferrocene and trans-Cr(C0)2( d ~ ewhere
) ~ the voltammetry is independent
of the identity of the cation, but depends on the nature and concentration of the anion.
374
Solid-electrode-solvent
intefaces
I
I
I
I
I
0
200
400
600
E (mV) vs Ag (s)1 AgCl (s)I KC1 (aq), 3 M
Fig. 5.30 Cyclic voltammograms of T C N Q microcrystals adhered to a ~afion@-coated
RAMTM
electrode (scan rate 1 0 0 m ~ s - ' )in contact with aqueous 0.1 MKCI, K N 0 3 , KBr, KF, and KI
electrolytes. The responses are similar because the anions do not take part in the charge neutralization
process. Reproduced by courtesy: Analyst 123 (1998) 1891. Copyright, Royal Society of Chemistry.
I
-0.3
I
I
-0.1
I
I
0.1
I
I
I
0.3
I
0.5
E (V) vs Ag (s)/AgCl (s)I KC1 (aq), 4 M
Fig. 5.31 The dependence of cyclic voltammograms (scan rate 100 mV s-l) for T C N Q microcrystals immobilized on a ~afion@-coated
RAMTMelectrode in contact with aqueous KC1 electrolyte of
different concentrations confirms the dependence of the cation on the charge neutralization process.
Reproduced by courtesy: Analyst 123 (1998) 1891. Copyright, Royal Society of Chemistry.
responses of T C N Q microcrystals in 0.1 M NaC1, 0.1 M KC1, 0.1 M RbNQ3,
and 0.1 M CsC1, all recorded on the same electrode. It is clear that the critical
potentials, the widths of the inert zones and the peak widths all depend upon
the identity of the cation. Numerical data are collected in Table 5.7. These
data confirm that the cation plays an important role in the charge neutralization
process and shows why the cation is included in eqn 5.16.
Voltammetry of TCNQ
375
0
200
400
E (mV) vs Ag (s) l AgCl (s) I KC1 (aq), 3 M
. 5.32 Cyclic voltammograms (scan rate 50mVs-l) obtained for T C N Q immobilized on a
RAMTMelectrode in contact with aqueous 0.1 M NaCl, 0.1 M KC1,O. 1M R b N 0 3 ,
~afion@-coated
and 0.1 M CsCl electrolytes demonstrate the different widths of inert zones for different cations and
confirm that the cation is associated with the charge neutralization process. Reproduced by courtesy:
Analyst 123 (1998) 1891. Copyright, Royal Society of Chemistry.
.7 Effect of changing the cation on the critical oxidation
and reduction potentials, inert zone widths (vi), and peak shape
parameters of cyclic voltammograms of T C N Q microcrystals on a
NafionB-coated RAMTM
electrodea
0.1 M NaCl
0.1 M KC1
0.1MRbN03
0.1 M CsCl
244
212
154
240
24
32
14
4
220
180
140
236
20
20
21
28
20
23
27
17
"Data from Fig. 5.32 (reference [38]). Clearly the cations have a significant effect on the
kinetics of the phase transformations.
4.2.1
Voltammetn'c evidencefor a nucleation-growth mechanism
e effect of scan rate on the reduction T C N Q microcrystals on a lXllMTM
trode placed in contact with 0.1 M NaCl, 0.1 M KCl, 0.1 M R b N 0 3 , and
0.1 M CsCl is shown in Fig. 5.33, where the peaks at lower scan rates can be
seen to emerge from the rising portions of the voltammograms at higher scan
rates. This is a tell-tale sign of a rate-determining nucleation and growth mechanism [41]. The effect occurs because of the special mathematical relationship
which couples nucleation to growth, which causes current-voltage curves to
have steeper gradients at low scan rates than at high scan rates. Further evidence
of nucleation is provided by Fig. 5.34, in which the scan directions are reversed
in the foot of each peak. Current maxima occur on the reverse scans, which
Solid-electvode-solvent intefaces
TCNQ -+(N~+)(TCNQ-)
Fig. 5.33 Effect of scan rate ( m ~ s - I )on the reduction of T C N Q microcrystals adhered
to a ~afion@-coatedRAMTM electrode in contact with aqueous 0.1 M NaC1, 0.1 M KC1,
0.1 M R b N 0 3 , and 0.1 M CsCl electrolytes. The emergence of the voltammetric peaks at low
scan rates from the voltammetric peaks at high scan rates is characteristic of a nucleation-growth
mechanism. Reproduced by courtesy: Analyst 123 (1998) 1891. Copyright, Royal Society of
Chemistry.
are diagnostic of rate-determining nucleation and growth kinetics. The 'triggering' of nucleation at high overpotentials results in an enhanced current at
low overpotentials when the scan direction is reversed as in Fig. 5.34. This is
in stark contrast to diffusion-controlled processes, which inevitably show lower
currents on reverse scans due to the depletion of reactant. Diffusion control can
also be excluded because of the absence of a 'tail' in the voltammograrns caused
by the necessary waiting for the reactant to arrive (or depart) from the reaction
site. Thus, nucleation is implicated in both the reduction (Fig. 5.34(a)) and
the oxidation processes (Fig. 5.34(b)),indicating a reversible solidsolid phase
transformation.
Definitive evidence for nucleation-growth kinetics also is provided by
chronoamperometric (double-potential step) experiments [41]. In Fig. 5.35,
the results of two successive potential steps experiments are shown for 0.1 M
aqueous solutions of NaC1, KC1, R b N O z , and CsCl: the first one reductive
and the second one oxidative. The step durations are 10 s each, except in the
case of CsC1, where 20 s is used. Note that the current-time transients exhibit
well-defined peaks (not to be confused with the vertical spikes, which are caused
by the rapid onset and decay of the capacitance charging currents at short times).
The existence ofwell-defined peaks in response to potential steps is the standard
'text-book' test for the existence of nucleation-growth kinetics. The presence
Voltammetry of T C N Q
-0.1
0.15
377
0.40
E (V) vs Ag (s)IAgCl (s)I KC1 (aq), 3 M
Fig. 5.34 Cyclic voltammograms (scan rate 100mVs-')of T C N Q microcrystals adhered to a
~afion'-coated RAMTMelectrode in contact with aqueous 0.1 MNaC1 electrolyte. Two cycles
are shown in which the scan directions are reversed at potentials corresponding to the foot of (a) the
reduction and (b) the oxidation process. Current maxima observed on scan reversal are diagnostic
of nucleation-growth and are observed in both directions. Reproduced by courtesy: Analyst 123
(1998) 1891. Copyright, Royal Society of Chemistry.
of nucleation and growth now explains the phenomenon of the 'inert zone'
shown in the cyclic voltammogram in Fig. 5.26. If every reduction and oxidation event is driven to 100 per cent completion, then a nucleation overpotential
must be supplied on each subsequent scan to restart the phase transformation
process.
In summary, electrochemical experiments reveal that T C N Q microcrystals
immobilized on the surface of an electrode and then placed in an aqueous solution containing a Group I cation exhibit responses which are consistent with
chemically reversible solid-solid phase transformations occurring via nucleationgrowth mechanisms. The electrochemical responses are essentially independent
of the identity of the anions but strongly dependent on the identity and concentration of the cations. This is because the cations rather than the anions are
intercalated and de-intercalated within solid TCNQ.
4.2.2 Theoretical thewnodynamic considerations velevant to a
nucleation-growth model
A model of the form shown in Fig. 5.1(b) can be used to explain the principal experimental features of the solidsolid transformations in T C N Q [37].
In this model, a T C N Q nanocrystal (phase 1) is assumed to be a hemisphere
whose volume does not change during the phase transformation to its cation
400
(N~+)(TCNQ-)-----> TCNQ
(K+)(TCNQ-) ---) TCNQ
I TCNQ d (R~+)(TCNQ)
Fig. 5.35 Double-potential-step chronoamperograms for T C N Q microcrystals adhered to a
~afion@-coated
RAMTMelectrode in contact with aqueous 0.1 M NaC1, KC1, R b N 0 3 , and CsCl
electrolytes. Potentials are cited in mV versus Ag/AgCl. The existence of peaks in the current-time
transients is further confirmation of a nucleation-growth mechanism. Reproduced by courtesy:
Analyst 123 (1998) 1891. Copyright, Royal Society of Chemistry.
salt (phase 2). The aqueous solution and the electrode constitute phases 3 and 4,
respectively. Though geometrically idealized, the assumption of hemispherical
geometry actually does not greatly affect the conclusions, but it does fix some
geometric constants at certain values.
The reversible work required for formation of unit area of interface between
two phases p and q is called the specific interfacial free energy y,, . This parameter
will be shown to play a controlling role in the solidsolid transformation. From
the theory of the thermodynamics of solids [42], it is expected that the specific
interfacial free energy y12between the solid phases 1and 2 will be high (hundreds
of mJ mp2)owing to lattice mismatch. However, no such mismatch should exist
at the interfaces (1,3) and (2,3) due to the liquid nature of phase 3. In the case of
T C N Q , there is probably significant lattice mismatch at the interface (1,2) due
to the different arrangements of the T C N Q stacks in the anionic and neutral
forms.
In principle, the first nucleus of the new phase 2 arising from reduction of
phase 1 could appear at four different locations: inside phase 1, at the two-phase
boundaries (1,3) or (1,4), or at the three-phase boundary (1,3,4). In an obvious
notation, the total work done to form a small crystal of the new phase 2 in each
Voltanlnzetry of TCNQ
1 symmetry):
locations
C1 = -
In these equations, AG, is the change in volume free energy per unit volume
of reactant and is given by AG, = nFp,,,q, where p,, is the molar density of
and q is the overpotential. However, since the terms in parentheses
tend to be self-cancelling, while yl2 is large, it is expected that
~t therefore follows that nucleation is most likely to occur at the three-phase
ary (1,3,4)since this is the location that minimizes the emergent area of the
nergy (1,2) interface in the critical nucleus. Furthermore, consideration
of conductivity and transport pathways also suggests the three-phase boundary
e region in which an electrochemical reaction is most likely to occur.
after, attention is restricted to this case.
y a standard argument of nucleation theory [43] it can be noted that A G1z,4
passes through a maximum A GT3,4at a value of r denoted r*, known as the critical radius. Evidently A G[3,4is the energy barrier to nucleation. Differentiating
eqn (5.20) and setting the result equal to zero yields
ince the nucleation rate is of the form [43]
a = a0 exp
-4"
Yf2
3kT(nF~mr)~
This formula clearly reveals the relationship between the nucleation rate a ,
the overpotential q, and the specific interfacial fiee energy y12 of the twophase boundary between TCNQ and its cation salt during the transformation.
However, what is less clear is how a critical overpotential qcn, arises, below
380
Solid-electrode-solvent intefaces
which the system is inert and above which the phase transformation occurrapidly (Fig. 5.26). To understand this feature, it may be first of all supposed
that the minimum detectable rate of nucleation in a statistically large field of
independent nanocrystals is x nuclei cm-2 s-'. (Note: although this rate is
obviously dependent on system size, it will conveniently emerge that the critical
overpotential is only a very weak function of x .) Next, it is necessary to explore
conditions for which
a=x
(5.27)
because this equality is first satisfied at q = qcn, Finally, substituting for a from
eqn (5.26) and inverting the expression yields
which is the required result. Equation 5.28 demonstrates that the greater the specific interfacial free energy ylz of the two-phase boundary in the critical nucleus,
the greater the critical overpotential for the nucleation process. Moreover, for
x << ao, which corresponds to the case in which the minimum observable rate
is much less than the maximum possible rate (the usual case), the system size
dependence of x has only a small influence on qcrit.For example, for a0 = 102'
and x = 10, 100, and 1000, the term in brackets in eqn 5.28 has the values 0.151 [ 4 n / 3 k ~ ] ' / ~0.155[4n/3kt]'/~,
,
and 0.160[4n/3k~]
'I2, respectively,
which is a constant within 6 per cent. Note, also, that for large values of y12
there exists a large range of q values (up to qcrit)where the rate of reaction is
vanishingly small, that is, the system is inert. The origin of this phenomenon
clearly lies in the non-linear dependence of a on q given by eqn 5.26, and it is
this which causes the 'inert zones7in the voltammograms.
The above model also explains why, in the absence of mass-transport or
ohmic control, the peak widths in voltammograms are narrow: the growth
rate of phase 2 at the expense of phase 1 is expected to increase exponentially
(or greater) with overpotential. Hence, the interface (1,2) is predicted to sweep
through phase 1 at a very high rate after it has been nucleated, which is precisely
what is observed.
It has been shown above that the voltammetry of nanocrystals of T C N
immobilized on an electrode surface exhibits the response expected for a compound undergoing a solid-solid phase transformation under rate control by
nucleation and growth. Characteristic features of this transformation are (a)
the existence of an inert zone between large reduction and oxidation peaks,
(b) narrow peak widths, (c) peaked current-time transients in response to large
amplitude potential steps, and (d) loops in cyclic voltammograms.
Although the theoretical model is successful in explaining the principal experimental features of the voltammograms, full quantitative application of the
model must be approached with caution. The reasons are two-fold. First,
the model makes no provision for the development of strain at sites of lattice mismatch. Second, the model assumes an insignificant change in volume
Voltammetry of TCNQ
381
during uptake and expulsion of cations, and this is strictly true only12 in the case
n general, the occurrence of strain at sites of lattice mismatch depends
on whether the interface is coherent or incoherent, and in the experiments
described above, this has not been ascertained. If an interface were coherent,
large strains might develop and then it would be necessary to modify the theory
(or, at least, the interpretation of yI2)because of the additional work needed
to form the interface. Alternatively, if an interface were incoherent, which is
ht to be the case in most reconstructive first-order phase transformations
ere would be no crystallographic continuity and hence no strain.
he effect of a small change in volume during uptake and expulsion of cations
ess worrisome, because the resulting stress can be relieved by plastic flow or
motion of lattice defects. Moreover, the use of widely spaced nanocrystals
in the present work also prevents the pile-up of large-scale stresses such as one
see in densely packed metallurgical specimens. O n the other hand, large
es in volume (say, >10 per cent) may be problematic, and in such cases
some modification of the theory might be needed.
2 . 3 Exploration of the nucleation-growthprocess and
her mechanistic details by ancillary techniques
voltammetry of surface-immobilized T C N Q provides evidence of solidphase transformations controlled by nucleation and growth mechanism.
ta obtained by ancillary techniques such as in situ optical microscopy, ex situ
scanning electron microscopy (SEM) and X-ray diffractometry help to further
elucidate the details of [ T C N Q ] ~ ' process.
Optical microscopy
n order to monitor colour changes that occur during potential cycling experiments on solid TCNQ, an electrochemical cell was fitted with a quartz-window
ase, and the working electrode was attached to a screw thread which allowed
t to be positioned in the focal plane of an inverted metallurgical microscope
371. Optical data were then recorded as video images or stored as single frames
in a computer [37]. Using this approach, a colour change from yellow to bluegreen was observed upon electrochemical reduction of T C N Q microcrystals
adhered to an electrode in contact with 0.1 M KC1. Interestingly, the colour
change occurred on the time-scale of seconds for small crystals (< 1ym) but on
the time-scale of minutes for larger crystals. This observation suggests that the
voltammetric responses recorded on short time-scales are primarily associated
with crystals in the nanometer size range.
1 2 ~ h eunit cell volume per molecule of T C N Q is 256 A3 in both T C N Q and
[ ~ a ' ] [TCNQ-] . For [K'] [TCNQ-] , [ ~ b ' ] [TCNQ] and [Cs+I2[TCNQ-12 [TCNQ],
values are 280, 287, and 277 A3 respectively (see Table 5.9 and J. Chem. Soc., Faraday
Trans. 92 (1996) 3925).
Scanning electyon microscopy and electron probe mic~oanalysis
The SEM and electron probe microanalysis (EPM) experiments were undertaken with an in-lens field emission scanning electron microscope (FESEM).
To facilitate the transfer of samples from the electrochemical cell to the FESEM
chamber, dual-purpose gold substrates were employed that acted both as the
electrode and the FESEM sample holder. As with regular electrodes, these substrates were mechanically coated with T C N Q microcrystals before immersion
in the electrochemical cell. The T C N Q salts were then generated in situ by
electrolysis in 0.1 M Na+, K+, ~ b ' , and Cs+ chloride solutions. After removal
from the electrochemical cell, the now T C N Q salt-coated substrates were rinsed
with de-ionized water and dried in air at room temperature prior to obtaining
FESEM images [37,38].
On a gold
An FESEM image of some freshly deposited T C Q microc~stals
substrate is shown in Fig. 5.36. The majority of the microcrystals are between
100 and 500 nm in diameter. There are also a few crystals (not shown) of slightly
larger dimensions, but as mentioned earlier, these react slowly on the time-scale
of voltammetric experiments, and are probably of little statistical significance.
Four FESEM images, each recorded after nine and a half continuous voltammetric cycles of the potential in a fferent electrolyte solution9 are shown in
Fig. 5.37. The electrolytes are 0.1 N a a 0.1
KC17 0CsC1. It is clear that significan earrangement of solid
formation of needle-shaped microcrystals occurs upon cy
FESEM micrograph of T C N Q adhered to a gold substrate. The majority of particles are
between 100 and 500 nm in diameter. Reproduced by courtesy: Analyst 123 (1998) 1891. Copyright,
Royal Society of Chemistry.
Voltar-lznzetry of TCNQ
383
-37 FESEM micrographs of a gold electrode surface containing adhered T C N Q in conith aqueous 0.1 M NaC1, 0.1 M KC1, 0.1 M R b N 0 3 , and 0.1 M CsC1 electrolytes after
nine-and-a-half voltammetric cycles of the potential. In all cases needle-shaped highly crystalline
phologies have developed from that initially seen in Fig. 5.36. Reproduced by courtesy: Analyst
(1998) 1891. Copyright, Royal Society of Chemistry.
alyses indicate that the microcrystals contain the corresponding Group
from the electrolyte solutions. Anions are absent, demonstrating once
ke is negligible.
e average lengths, widths, volumes, and aspect ratios of the
crystals in the four cases considered in Fig. 5.37. Since the
microcrystals are generated by rapid solid-solid phase transe surprising if they were equilibrium morphologies. Most
the morphologies are simply those that permit the most
reactlon between the oxidized and the reduced forms of the TC
needles having average aspect ratios great
to-volume ratios higher than the equilibrium ones.
en the microcrystals are returned to the electrochemical cell and re-
persist even when experiments are repeated many times. It is concluded from
these data that after the first few voltammetric cycles, the dominant process is
384
Solid-electvode-solvent
intefaces
Table 5.8 Average lengths, widths, volumes, and aspect ratios ofneedle-shaped
microcrystals of T C N Q salts formed by potential cycling experiments. All the
microcrystals are needle-shaped with aspect ratios exceeding 4. See reference
1381 for further details.
T C N Q salt
Average length
(nm)
Average width
(nm)
Average volume
(lo6nm3)
Average
aspect ratio
the reversible transformation of the T C N Q microcrystals to the TCNQ-salt
rnicrocrystals with minimal change of bulk morpholo&.
Even though the majority of the voltammetric response is associated with
smaller microcrystals, larger rnicrocrystals (> 1 pm) are not completely inert.
particular, outgrowths occur, as shown in Fig. 5.38. Presumably these originate
from twin plane defects which commonly arise in high-speed crystal growth.
Lacunae can also be observed. These are hollow depressions in the centres of
crystal faces caused by the inability of surface diffus;on to transport molecules
fist enough from the rapidly growing crystal edges. The fact that many of the
outgrowths are hollow is clearly shown in Fig. 5.39. It is evident that electrolyte
solution can be trapped inside these structures.
The main conclusion from the FESEM studies is that large changes in microcrystal morphology occur during the early stages of voltammetric cycling,
leading to the formation of stable, needle-shaped crystals that are retained on
the surface of the electrode. Under the same conditions, larger crystals are less
morphologically stable, and tend to form hollow outgrowths.
X-ray dlffractometry
Dual-purpose substrates for electrochemistry and e x situ X-ray diffraction experiments were fabricated from microfoils of 0.05 pm gold on 13 pm mylar.
To increase the amount of material available for X-ray analysis, T C N Q was
deposited on the gold by evaporation of an ethanolic solution rather than by
mechanical attachment. After electrolysis to form the T C N Q salts, each surface was rinsed with de-ionized water and allowed to dry in air. Typically, the
diffraction angle was varied from 3" to 32" with a step interval of 0.025".
The chemical formulae and crystal structures of the electrogenerated T C N
salts were readily established by comparing experimental X-ray diffractograms
with a library of diffractograms obtained from the JCPDSACDD Powder
Diffraction File [45] or back-computed from published structures [38].
Figure 5.40 is an experimental X-ray diffractogram of T C N Q on a goldmylar substrate prior to electrolysis, and it can be seen that several peaks
Voltammetry of TCNQ
385
Fig. 5.38 FESEM images of outgrowths formed during the course of reduction of large T C N Q
crystals (> 1 pm) adhered to a gold electrode in contact with (a) 0.1 M NaCl and (b) 0.1 M KC1
aqueous electrolytes. Clearly, the bulk mass of the larger crystals react more slowly than the smaller
crystals. Reproduced by courtesy: Analyst 123 (1998) 1891. Copyright, Royal Society of Chemistry.
due to TCNQ are prominent above the background. Indeed, most of the
mylar response is concentrated between 22" and 30" and therefore it is easy to
exclude this from analysis. Peaks from gold occur above 40" and do not interfere. Figures 5.41(a) and 5.42(a) show the X-ray diffiactograms of TCNQ on
gold-mylar substrates before and after electrolysis in solutions of 0.1 M NaCl
386
Solid-electuode-solvent intefaces
.39 FESEM images of lacunae formed at the end of needle-shaped outgrowths during the
course of reduction of large TCNQ clystals adhered to a gold electrode in contact with 0.1 M NaCI.
(b) A close-up image of part of (a) which shows the details of a single lacunae which could trap
electrolyte solution. Reproduced by courtesy: Analyst 123 (1998) 1891. Copyright, Royal Society
of Chemistry.
and 0.1 M R b N 0 3 , respectively. Equivalent data in 0.1 M KC1 and 0.1 M CsCl
have also been obtained and are reported in reference [38].Some peaks due to
unreacted T C N Q are evident in all cases, but there are a sufficient number of
product peaks to distinguish the freshly formed T C N Q salts from each other
and from the inorganic salts in solution. For example, in the case of sodium
a
e
t
Before Electrolysis
T C N Q from JCPDS
2-0 angle (deg)
ig. 5.40 X-ray diffractograms of T C N Q microcrystals on a gold-mylar substrate before electrolysis.
Although the mylar contributes a very broad peak between 22" and 30°, the T C N Q peaks are still
prominent. Note that the T C N Q peaks are sharp, indicating good internal crystallinity. Reproduced
by courtesy: Analyst 123 (1998) 1891. Copyright, Royal Society of Chemistry.
Before Electrolysis
After Electrolysis
(Na+)(TCNQ-) from JCPDS
T C N Q from JCPDS
10
15
2-0 angle (deg)
20
10
15
2-0 angle (deg)
20
Fig. 5.41 X-ray diffractograms observed from T C N Q rnicrocrystals adhered to a gold-mylar
substrate in contact with aqueous 0.1 MNaC1 electrolyte before and after reductive electrolysis.
(a) Individual intensities. (b) Difference in intensities. Reproduced by courtesy: Analyst 123 (1998)
1891. Copyright, Royal Society of Chemistry.
(Fig. 5.41 (a)),the product peaks match those of [ ~ a +[TCNQ-1.
]
This is clearly
seen in the difference curves between the diffractograms shown in Fig. 5.41 (b)
after background subtraction. A decrease in the amount of T C N Q salts is indicated by the positive peaks. Small discrepancies in the relative peak heights
Before Electrolysis
After Electrolysis
Computed ( ~ b ' )(TCNQ-)
T C N Q from JCPDS
10
15
20
2-0 angle (deg)
10
15
20
2-0 angle (deg)
Fig. 5.42 X-ray diffractograms observed from TCNQ microcrystals adhered to a gold-mylar substrate in contact with aqueous 0.1 M RbCl electrolyte before and after reductive electrolysis.
(a) Individual intensities. (b) Difference in intensities. Reproduced by courtesy: Analyst 123 (1998)
1891. Copyright, Royal Society of Chemistry.
between the experimental data and the JCPDS data are not significant and are
due to the preferential orientation of crystals. As already shown in the FESEM
images, the microcrystals are lying flat on the electrode surface. Consequently,
some peaks are stronger while others are weaker than in the case of random
orientation. Analogous results are obtained for all the other salts.
The most surprising X-ray diffractometryresult is that of rubidium (Fig. 5.42).
Of the four known structures for rubidium T C N Q salts, it is the 1 : 1 salt
normally found at low temperature (113 K) that forms electrochemically. This
result strikingly illustrates the dominance of kinetics over thermodynamics in
rapid phase transformation processes.
Table 5.9 summarizes the structural parameters of the electrochemically generated T C N Q salts, and Fig. 5.43 show computer-generated perspective views
of these structures. Close inspection reveals that three different kinds of lattice rearrangement accompany electrochemical cation ingress: the separation of
T C N Q anion stacks (Cs'); the 90" rotation of alternate T C N Q anion stacks
accompanied by some lateral compression ( ~ b ' ) , and the 90" rotation of alternative T C N Q anion stacks with some flattening of the layers (Na+ and Kf).
Note also that in the cases of ~ a +K+,
, and ~ b ' , the reaction stoichiometries
are x = 1 and y = 1 (defined in eqn 5.16), whereas in the case of CS+they are
x = 3 and y = 2. The reaction for Cs' is therefore
+ 2~s+(solution)+ 2 e + [Cs+12[TCNQ-12 [TCNQ] (solid)
3TCNQ(solid)
(5.29)
rather than
+
TC~Q(so1id) CS+(solution)
+ e + [cs'] [TCNQ-](solid)
(5.30)
Voltammetry of TCNQ
389
5.9 Structural parameters derived from X-ray diffraction measurements
N Q and its salts of Group I cations, as reported in the literaturea
pace group
TCNQ
monoclinic
[~a+]
[TCNQ-]
triclinic
[K+]
[R~+I
[TCNQ-] [TCNQ-I
monoclinic monoclinic
[Cs+]2
[TCNQPl2
[TCNQ]
monoclinic
-
a (A)
16.415
b (A)
7.06
c
8.95
a (degree)
90.00
B (degree)
98.65
Y (degree)
90.00
Z
4
Unit cell volume (A3) 1024
Unit cell volume per
256
molecule of
TCNQ (A3)
Cation radius (A)
-
(9
'These structures are the ones observed electrochemically at 295 K. Note the complex stoichiometry
of the caesiurn salt. The unit cells contain 2, formula units and have dimensions a, b, c with opposite
angles of a , p , y , respectively. See reference 1381 for further details.
as might have been anticipated. The reason for this is unclear, but is probably
related to the fact that CS+has a larger radius than the other cations in Table 5.9,
which makes this cation more difficult to accommodate.
Structural representations contained in Fig. 5.43 indicate that cation ingress
into T C N Q is always accompanied by a significant (though reversible) structural
rearrangement of the [TCNQ-] anion stacks. Moreover, there is no evidence of
non-integer values of x and y,which would be expected if classical intercalation
were occurring. The question therefore arises of whether the electrochemical
reaction between Group I cations and T C N Q can be classified as intercalation at
all. In the older literature [46], intercalation seems to have meant the reversible
insertion of a guest species into a layered host with no orientational rearrangement of the structural features of the host, whereas in the case of T C N Q the
reversible insertion of a guest species into a channel host occurs with significant
orientational rearrangement of the structural features of the host. This would
appear to rule out the use of the word 'intercalation'. However, the definition
of intercalation has expanded in the recent literature to include all reversible
insertion reactions, even those in which the host and guest experience some
degree of perturbation (from subtle to extreme) in their geometric, chemical,
electronic, and optical properties [46,47]. If this revised definition is accepted,
then eqn. 5.16 would constitute an example of an intercalation reaction that is
also a nucleation and growth reaction.
Face View
Edge View
TCNQ
Fig. 5.43 Perspective views of the structures of T C N Q and its salts formed electrochemically at
2 2 ° C Atoms not drawn to scale. Reproduced by courtesy: J. Chem. Soc., Faraday Tram. 92 (1996)
3925. Copyright, Royal Society of Chemistry.
Voltammetry of TCNQ
39 1
Visualization of the cation channels in [TCNQ-] salts can be achieved by
molecular simulation based on the atomic coordinates from X-ray analysis [38].
he results are shown in Fig. 5.44, where it can be seen that even cations hav(see Table 5.9) as large as Cs' (1.67 A) can be accommodated in the
-1 channels. It is interesting to note that the cations are unsolvated in
the [TCNQ-] channels which implies that they have lost their inner solvation
spheres of water molecules at the mouths of the channels. Inside the solid phase,
the cations are, of course, stabilized by a 'coordination sphere' of [TCNQ-]
anions.
Fig. 5.44 Visualization of the cation channels in T C N Q salts made by molecular simulation using
atomic coordinates from X-ray analogues. The size of the cations are contained in Table 5.9.
Reproduced by courtesy: Analyst 223 (1998) 1891. Copyright, Royal Society of Chemistry.
4.3 Electrochemically driven tran$ormation of microcrystalline TCNQ to
tetraalkylamrnonium [TCNQ-] salts13
The metal cations inserted into the [TCNQIOI- system during the course of
redox cycling experiments described in Section 5 are relatively small. Questions
therefore arise as to whether voltammetry would still be possible when larger
)
(Fig. 5.45) have to be transported across
tetraalkylammonium ( N R ~ +cations
the solid-solution (electrolyte)interface and as to whether systematic thermodynamic and/or kinetic effects attributable to cation size effects can be interrogated
by studying the voltammetry of T C N Q adhered to an electrode in contact with
aqueous N R 4 + containing electrolytes.
4.3.1
Voltammetry
Systematic trends with parameters such as the peak potentials, the peak
height, and the gap between reduction and oxidation response are observed
in the voltammetry of T C N Q as the chain length of the R group in
the NR4+ cation is increased. Voltammograms obtained at a scan rate
of 20 mVs-' in the flowing solution configuration with a RAM^^ electrode (Fig. 5.46) clearly confirm that the peak current magnitudes for
this electrode-cell configuration are in the order N H ~ ' > NMe4+ >
N E ~ ~>+ NPr4+. Indeed, for the NPr4+ case, the current magnitude is
very small and more akin in magnitude to that observed with the solid
Fig. 5.45 comparison of sizes of some ions that have been used to form T C N Q - salts. Provided
by courtesy of P.G. Symons, Monash University, Victoria, Australia.
1 3 ~ .Bond,
~ .
S. Fletcher, F. Marken, and P.G. Symons, unpublished studies, 1996-98.
Voltammetry of TCNQ
393
NH,+
5.46 Voltammograms of microcrystals of T C N Q immobilized on a Nafion-coated R A M T M
electrode in contact with aqueous 0.1 M NR4C1 electrolyte solutions. Current scales are different for each case. $ed: NH4+, 1.1 PA; NMQ+, 360nA; NEtr+, 210nA; NPr4+, 25nA. (Scan
rate 20 mV s-l, third cycle of potential recorded.) Provided by courtesy of P.G. Symons, Monarh
University, Victoria, Australia.
trarzs-[~r(C0)~
(dpe)$/+ process. For the aqueous 0.1 M NH4Cl electrode
voltammograms have similar characteristics to those described
case, (Fig. 5.46),
reviously in the presence of potassium containing electrolytes (Fig. 5.30). The
almost symmetrical reduction and oxidation responses with a large peak-to-peak
separation and large peak current values and charges are indicative of extensive
solidsolid phase conversion. The data are most simply interpreted in terms
of a one-electron reduction accompanied by incorporation of N H ~ +into the
solid lattice (eqn 5.31), although XRD data are not available to confirm this
stoichiometry.
TCNQ(so1id)
+ e- + N H ~ + ( s o ~ u ~6
~ o ~ ) [ N H ~ +[TCNQ-]
]
(solid) (5.31)
or the other electrolytes considered in Fig. 5.46, the extent of electrolysis on
the voltammetric time-scale of 20 mV s-' must be far from exhaustive.
The critical potentials for the onset of the reduction and re-oxidation response
both shift to more positive values upon increasing the size of the electrolyte
cation, and the peak separation between the reduction and oxidation components increases slightly with change in cation size. The shifi in ( E F ~ EF)/2
can again be attributed to a change in the formal thermodynamic (reversible)
redox potential, for a reaction of the kind given in eqn 5.32.
+
It therefore may be proposed that a hydrophobic cation, such as NPr4+, will
be removed from the aqueous to the solid phase at a more positive potential
394
Solid-electrode-solvent
intevfaces
than the less hydrophobic NH4+ in accord with data available for the transfer
of cations from aqueous into organic phases [19]. Thus, neglecting any solidstate effects, the trend in peak potentials can be explained in terms of increasing
to Npr4+ cations. The equivalent effect
hydrophobicity in changing from N H ~ +
was noted for the t r a n s - [ ~ r ( ~ ~ ) ~ ( d ~process
e ) ~ ] ~which
'+
involved insertion
of anions into the structure.
The gap between the reduction and re-oxidation processes, or the so-called
'inert zone', when alkaline metals are the cations, has been attributed to the need
to provide a potential to overcome interiacial free energy between differences in
the two solid phases by a nucleation-controlled mechanism. Assuming that the
formation of solid [NR4+][TCNQ-] requires an increased volume change as
the size of the NR4+ cation increases, and therefore a more pronounced lattice
mismatch, allows the present results to be explained by a nucleation-growth
process. However, details concerning the crystallographic changes accompanying reduction and incorporation of the cation at an electrode surface are
series of salts.
unknown for the N R ~ +
A new feature introduced in the solid-state voltammetry of T C N Q in aqueous 0.1 M N B u ~ +containing electrolyte solution (Fig. 5.47) is the splitting of
the reduction response which may be explained in terms of the formation of
different phases with distinctly different energies being postulated. The origins
of the responses for the first reduction process (IaRdand Ia,,) and the second
reduction process (Ibredand Ib,,) (see Fig. 5.47) are summarized in eqns (5.33)
and (5.34).
A
[NBu4+][TCNQ] [ T C N Q ](solid)
[NBU~+]
[TCNQ] [TCNQ]- (solid)
(533)
+ e + NBu4+(solution)
(5.34)
- 2[NBu4+][TCNQ-] (solid)
A
I
I
I
0.0
0.5
1.O
E (V) vs Ag /AgCI (3 M KCI)
Fig. 5.47 Cyclic voltammograms obtained at a scan rate of 1 rnV s-' and at 22OC for reduction
of T C N Q adhered to a basal-plane pyrolytic graphite electrode placed in contact with aqueous
0.1 M NBu4C1 electrolyte. Provided by courtesy of F. Marken, La Trobe University, Victoria,
Australia.
Voltawznzetry of TCNQ
4.3.2
395
Scanning electron microscopy
canning electron micrographs were obtained after electrochemical reduction
of solid T C N Q adhered to a gold electrode in contact with 0.1 M NH4+,
~ e ~ NEt4+,
' ,
and NBu~' aqueous electrolytes. In the presence of NH4+,
ell-defined, elongated, needle-shaped microcrystals of up to 5-pm length were
observed after reduction. In NMe4+ containing electrolyte, smaller crystals of
approximately 1-pm length, but with a similar shape to the NH4' case were
ined. In the presence of N ~ t 4 +
even smaller crystals were produced and with
4' the solid formed after reduction retains its non-crystalline appearance.
These observations are consistent with a far greater extent of electrolysis in
the presence of the smaller N H ~ ' cation and that a rate-determining step is
controlled by the cation insertion process in conjunction with crystal growth.
.4 Dissolution of solid TCN
from electvode sufaces
The change from weakly diffracting solid to needle-shaped crystals during the
course of potential cycling and/or T C N Q reduction experiments implies that
dissolution and probably precipitation reactions are highly significant in early
stages of the cyclic voltammetry of microcrystals of T C N Q adhered to electrode
surfaces. Indeed, the initial formation of any new solid at an electrode-solidsolution (electrolyte) interface produces very small amounts of material which
can be expected to dissolve into the solution phase because the solubility will
not be exceeded, unlike the case at later times in the experiment when large
amounts of solid are generated. Thus, contribution of dissolution processes need
to be considered in any reaction where a solid adhered to an electrode is placed
in contact with an electrolyte.
Use ofrotating ring-disc electrode and in situ electron spin
resonance techniques to detect dissolution processes that accompany the
voltammetry ofsolid TCNQ adhered to electrode s ~ f a c e s ' ~
n principle, the use of i n situ rotating ring-disc electrode (RRDE) and
4.4.1
simultaneous electrochemical-electron spin resonance (SEESR) techniques
(Section 16.1 in Chapter 2) should enable the detection of solution-soluble
products created during electrochemical experiments on T C N Q attached to an
electrode surface.
Direct detection of products of electrode processes by scan reversal techniques, as used in cyclic voltammetry at a stationary macrodisc electrode, are
obviously not available with the rotating disc electrode (RDE), since the product of the electrode reaction is continuously swept away from the surface of the
disc (see Section 9.1 in Chapter 2). Thus, at the RDE, reversal of the direction
of the potential sweep, under conditions where the scan rate is sufficiently slow
1 4 ~ d a p t e with
d
permission from J . Electrochem. Soc. 144 (1997) 1566. Copyright, The
Electrochemical Society.
396
Solid-electrode-solvent intefaces
compared to the rotation rate, will just retrace the curve obtained in the forward scan. Information equivalent to that obtained by cyclic voltammetry at a
stationary electrode is obtained in the RDE method by addition of an independent ring electrode surrounding the disc [48-501 to give the so-called RRDE
methodMeasurement of the current at the ring electrode with the potential maintained at a given value enables knowledge to be obtained about what is occurring
at the disc electrode surface. For example, if the potential of the ring is held at
a value at the foot of a reversible reduction wave, any soluble product formed
at the disc will be swept over to the ring by the radial flow streams where it
will be oxidized back to the starting material or 'collected'. The mass transfer
to a ring electrode is larger than that to a disc at a given rotation rate, because
flow of fresh solution to it occurs radially from the area adjacent to the ring, as
well as normally from the bulk solution. However, the theoretical treatment of
ring electrodes is more complicated than that of the RDE, since the radial mass
transfer term must be included in the convective-diffusion equation. While
the mathematics may be difficult, the results are simple to understand. In the
R R D E method, as applied to studies of T C N Q , the solid may be attached
to the disc, or so-called generator electrode, and the ring electrode, which is
separated from the disc electrode by the solution phase, can then be used as
the detector electrode to establish the identity of soluble species that have been
transported across the solution phase.
In SEESR experiments, compounds to be studied may be attached to working electrode contained in a quartz flat cell which is then used within an electron
spin resonance (ESR) cavity [51]. While R R D E experiments allow studies to
be undertaken on the voltammetry of components formed either transiently
or permanently in solution, ESR measurements should enable a distinction
to be made between solution and solid-state surface-confined paramagnetic
species. Thus, the combination of these in situ methods enables features of the
complex electrode-solid-solvent (electrolyte) interface to be probed. In particular, these techniques may reveal whether extensive or transient dissolution of
solids occurs during transformations between oxidized (reduced) and reduced
(oxidized) states.
RRDE studies on TCNQ
Figure 5.48 shows the platinum disc and ring (Pt/Pt) electrode responses (versus
disc potential) for reduction of solid microcrystalline T C N Q attached to the disc
electrode under conditions of hydrodynamic voltammetry when the electrode
is placed in Li+-buffer.15 In this medium, a reduction reaction yielding waterM ~ is) observed
soluble [L~+][TCNQ-] (solubility product [52]: 2.2 x
with a peak potential Ep = -0.18 V versus Ag/AgCl (Fig. 5.48(a)). The conclusion that dissolved [TCNQ-](solution) is formed is reached on the basis that
1 5 ~ hbuffers
e
used in the R R D E studies were Britton-Robinson p H 10.0 buffers containing 0.02 M LiOH or K O H which are referred to in the text as ~i+-bufferand K+-buffer,
respectively.
Voltammetry of TCNQ
397
. 5.48 Pt/Pt R R D E cyclic voltammetry of T C N Q attached to the disk electrode in contact
with Li+-buffer using a scan rate of 20 rnV s-I and frequency of rotation of 1000 min-I. (a) Disc
current versus disc potential; (b) Ring current versus disc potential, with a ring potential of 0.2 V
versus Ag/AgCl; (c) Ring current (five times) versus disc potential, with a ring potential of -0.2 V
versus Ag/AgCl. Reproduced by courtesy: J. Electrochem. Soc. 144 (1997) 1566. Copyright, The
Electrochemical Society.
no oxidation reaction is detected at the disc electrode during the reverse sweep
(Fig. 5.48 (a))and by noting that dissolved solution-soluble [TCNQ-] (solution) is
in fact swept past the ring electrode where it is detected via its oxidation response
when the ring electrode is held at a potential of f0.2 V versus Ag/AgCl. The
observed behaviour at the RRDE is consistent with the reaction sequence
+ ~i+(solution)+ e- --+ [~i'] [TCNQ-] (solid)
[~i'] [TCNQ-] (solid) -+ ~i'(so1urion) + [TCNQ-] (solution)
TCNQ(so1id)
(5.35)
(5.36)
No significant response on the ring electrode was detected during a complete cyclic voltammetric sweep when its potential was held at -0.2 V versus
Ag/AgCl (Fig. 5.48(c)), where detection of T C N Q could be expected. This
is consistent with the absence of transfer of solid T C N Q from the disc to the
ring electrode, and that complete dissolution of solid [~i'] [TCNQ-] occurs
during one sweep of the potential under conditions of RRDE cyclic voltammetry. Control experiments confirmed that T C N Q remained firmly attached
to the disc electrode surface when no reduction potential was applied, even at
frequencies of rotation of up to lo4 min-I .
Pt/Pt KRDE cyclic voltammetry of T C N Q attached to the disc electrode in contact wit],
K+-buffer. Three consecutive scans shown with a scan rate of 20 rnV s-I and frequency of rotation
of 1000 min-l. (a) Disc current versus disc ~otential;(b) Ring current ( ~ 2 0versus
)
disc potential,
with a ring potential of 0.2 V versus Ag/AgC1; (c) Ring current (x20) versus disc potential, with a
ring potential of -0.3 V versus Ag/AgC1. Reproduced by courtesy: J. Electrochem. Soc. 144 (1997)
1566. Copyright, The Electrochemical Society.
is significantly different when
E voltammetric response of T C
, rather than ~ i ' , is present in the electrolyte solution. Figure 5.49 shows
disc and ring detector electrod
ses during the first three voltammetric sweeps when surlace-attached
is reduced when the electrode is in
'-buffer. In this c
] salt is sparingly soluble (solueven though the
bility product 5 x 1 0 - l ~
ely small level of dissolution of
reduced solid occurs. W
is present in the electrolyte, the disc electrode
response under hydrody
nditions (Fig. 5.49(a)) is in fact closely related
to that found with static electrodes with a well-defined apparently inert region
being detected between the reduction and oxidation peaks, and with a slow
decrease in peak height as cycling of the potential occurs.
equency of electrode rotation is 3 1000 min-' , a significant amount of
] is detected during the initial stages of the reduction of solid T
wever, when repetitive cyclic voltammetric experiments are und
(Fig. 5.49(a))only very minor ring electrode responses are prese
Apparently, in the initial cycle of the potential, reduction of T
-1 which is predominantly 0x1
formation of solid [K+][TC
structurally modified or differently adhered form of solid T C N Q , although some
dissolution of the adhered salt does occur. In second and subsequent cycles ofthe
Voltammetry of TCNQ
399
ptential and after modification of adhered TCNQ, the nucleation and growth
reaction interconverting the T C N Q and [K'] [TCNQ-] solids are dominant
relative to the dissolution process (eqn 5.37).
[K'] [TCNQ-] (solid)
+ K' (solution) + [TCNQ-] (solution)
(5.37)
Interestingly, when the ring electrode potential corresponding to the reduction of T C N Q is set at -0.30V versus Ag/AgCl (Fig. 5.49(c)), a small yet
significant response indicates that during the first reductive sweep, TCNQ is
expelled from the surface of the disc electrode at the same time as reduction
of solid T C N Q occurs. T C N Q is also expelled from the sudace during oxidation of solid [K'] [TCNQ-] to TCNQ. Significant expulsion of T C N Q seems
to accompany only the initial reductive cycle of the potential (Fig. 5.49(c)).
contrast, some T C N Q always appears to be expelled during oxidation of
+][TCNQ-] (solid) back to TCNQ(so1id)on the disc electrode during the posive potential oxidative voltammetric scan, (Fig. 5.49(c)). This loss of T C N Q
accounts for the decrease in peak height observed in second and subsequent
otential cycles of the disc electrode (Fig. 5.49(a)) even though no soluble
[TCNQ-] (solution) is detected at the ring electrode (Fig. 5.49(b)). Clearly,
the RRDE experiments reveal that significant rearrangement in the nature of
adhered T C N Q occurs during potential cycling experiments, as is also demonstrated in the SEM studies described in Section 4 under the headings of scanning
electron microscopy and electron probe microanalysis. However, many of the
details of the loss of solid accompanying the electrochemistry of T C N Q adhered
to electrode surfaces are unknown.
SEESR studies on TCNQ
The ESR detection of electrochemically generated radical anions (SEESR
technique) should be sensitive to the phase of a particular paramagnetic comound. In the case of electrochemically generated [TCNQ-I, a multiple-line
SR spectrum would be expected from dissolved [TCNQ-](solution) [54],
and a single-line spectrum for a solid phase [TCNQ-] salt [55]. Figure 5.50
shows the solution phase ESR spectrum obtained from chemically synthesized [54] [Li'] [TCNQ-] salt dissolved in Li+-buffer. The simulated ESR
spectrum with g = 2.0023, a ( 1 4 ~=
) 0.0985 mT, and a('H) = 0.1435 mT,
also included in Fig. 5.50, is fully consistent with published data [54]. The
in situ SEESR spectrum shown in Fig. 5.50 was recorded during the course
of one-electron controlled-potential electrolysis of surface-attached microcrystalline T C N Q adhered to an electrode in contact with ~i+-buffer.Clearly,
the ESR spectra shown in Fig. 5.50(a) and (b) are indistinguishable and there
is no ESR evidence to indicate the presence of solid [Li'] [TCNQ-1. Under
the thin-layer conditions of the SEESR flat cell used in these experiments, the
(solution) and oxidareduction of TCNQ(so1id)to yield dissolved [ ~ i +[TCNQ-]
]
tion of dissolved [Li'] [TCNQ-] (solution) back to suriace-attached TCNQ(so1id)
is chemically reversible. Figure 5.51 shows a series of ESR spectra recorded
Field (mT)
Fig. 5.50 Solution-phase ESR spectra of TCNQ-; g = 2.0023, a ( " ~ ) = 0.0985mT,
a('~=
) 0.1435 mT. (a) Solid line is the experimental spectrum obtained from 0.1 M [Li+][TCNQ-]
in Li+-buffer and the dotted line, which almost completely overlaps experimental spectrum is the
simulated spectrum; (b) in situ ESR spectrum recorded during the course of electrochemical reduction of T C N Q adhered to a Pt electrode in contact with Li+-buffer. Reproduced by courtesy: J.
Electrochem. Soc. 144 (1997) 1566. Copyright, The Electrochemical Society.
of (a) the initial solution, (b) fully reduced solid T C N Q , (c) fully re-oxidized
[TCNQ-](solution), and (d) again fully reduced solid T C N Q in a thin-layer
SEESR cell. Optical monitoring of the electrode surface confirms that deposition of a thin yellow layer ofTCNQ(so1id)on the electrode surface occurs upon
re-oxidation of the blue solution-soluble [TCNQ-] (solution).
When K+-buffer is used as the electrolyte solution, distinctly different SEES
behaviour is observed. Figure 5.52 shows the ESR spectra obtained via constant potential electrolysis at four different potential regions. The solution-phase
ESR spectrum of [TCNQ-](solution) is observed when the potential is held at
the foot of the reduction waves (i.e. prior to the reduction peak potential).
The double integral of the spectrum is consistent with the concentration of
[TCNQ-](solution) being below that derived from the solubility product of
[K'] [TCNQ-] . No clear indication of a solid-state spectrum is evident when
Voltammetry of TCNQ
401
Field (mT)
Fig. 5.51 In situ generation of solution-phase ESR spectrum of T C N Q - in ~i+-bufferat a Pt
electrode. ESR spectrum (a) of initially surface-attached T C N Q ; (b) after complete reduction of
TCNQ(so1id) to [TCNQ-](solution); (c) after complete regeneration of TCNQ(so1id); (d) after
complete reduction to TCNQ-(solution) after regeneration of T C N Q as in (c). Reproduced by
courtesy: J. Electrochem. Soc. 144 (1997) 1566. Copyright, The Electrochemical Society.
electrolysis is undertaken at this potential, even after leaving the cell switched
on for several minutes. However, after applying a potential corresponding to
the reduction peak potential (Fig. 5.52(b)),a distortion is observed in the central
region of the solution-phase ESR signal. The dominant feature is a single-line
ESR spectrum when the potential is adjusted to values more negative than
the peak potential (Fig. 5.52(c) and (d)). Furthermore, when the potential was
applied at these negative values for increased periods of time, the intensity of
the solution-phase ESR spectrum remained constant, whereas the intensity of
the solid-state spectrum increased. A comparison with the single-line solidstate ESR spectrum of [K+][TCNQ-] (Fig. 5.53) enables the conclusion to be
reached that the in situ ESR spectrum of reduced [TCNQ-] results from the
superimposition of solution and solid-state spectra of [ T C N Q - ] ( s o l u t i o n ) and
[K'] [TCNQ-] (solid), respectively.
In summary, for the [TCNQ]'' system, the combination of R R D E and
SEESR experiments shows that microcrystalline T C N Q , when attached to
Field (mT)
Fig. 5.52 ESR spectra of in situ electrochemically generated TCNQ-(solution) and/or
[K+][TCNQ-](solid) from TCNQ(so1id) adhered to a Pt electrode in contact with Kf -buffer.
Dependence of ESR response on applied potential is shown (a) prior to peak; (b) at peak; (c) on
decreasing part of peak; (d) after peak. The electrolysis potentials were determined from a voltammogram obtained in the SEESR cell at a scan rate of 100mVs-l. Reproduced by courtesy: J.
Electrochem. Soc. 144 (1997) 1566. Copyright, The Electrochemical Society.
an electrode surface, is reduced in the solid-state phase. Dissolution may
take place up to the point where the solubility product is not exceeded,
and when the applied potential is not sufficiently negative to enable nucleation and crystal growth to occur. In the case of ~i+-containingelectrolyte
solutions, the solubility product is never exceeded. In contrast, the limited solubility of [K'] [TCNQ-] allows the simultaneous observation of both solid-state
reduction and dissolution pathways to be observed.
4.4.2 RRDE and SEESR studies on adhered solid trans- C T ( C O () ~d ~ e ) ~
As shown in Section 3.2.2, although, neither tr~ns-Cr(CO)~(dpe)~
nor
the trans-[Cr(CO)2(dpe)2]+ salt show any significant conductivity, the
t r a n s - [ ~ r ( ~ ~ ) , ( d ~ eprocess
) ~ ] ~ / in
+ the surface-attached state is well defined. It
is therefore interesting to compare the results from RRDE and SEESR studies
with the [TCNQ]O/- process, where semiconducting properties and dissolution
processes have been detected, with those for t r a n s - [ ~ r ( ~ 0 ) ~ ( d ~ e ) z 1 0 / + .
Figure 5.54 shows GC/GC RRDE voltammograms in 0.1 M aqueous KC1
when microcrystalline tran~-Cr(CO)~(dpe)~
is attached to a GC disc electrode
which is rotated at a frequency of 1000 min-'. Clearly, no ring electrode
response is detected at a ring potential of either 0.1 or -0.5 V versus Ag/AgC1
Voltammetvy of TCNQ
403
1 gz2.0023
Field (mT)
ig. 5.53 Solid-state ESR spectra of [K+][TcNQ-](solid). (a) [K+][TCNQ-] prepared from
reaction of KI and T C N Q in acetonitrile; (b) [K' ] [TCNQ-] (solid) prepared by electrochemical reduction of surface-attached T C N Q in contact with K' -buffer. Reproduced by courtesy: J .
Electrochem. Soc. 144 (1997) 1566. Copyright, The Electrochemical Society.
during the course of either the oxidation or subsequent reduction component
of the experiment. This result indicates that neither solid trnn~-Cr(CO)~(dpe)~
nor the trans-[Cr(CO)2(dpe)2]+salt dissolve to any measurable extent so that
a dissolution/precipitation mechanism is unlikely to accompany the solid-state
voltammetry of the trans-[c~(co), (dpe)2]0"- process. However, the solubility of these chromium carbonyl solids may be enhanced by the addition of
organic solvent. Figure 5.55 shows the GC/GC RRDE response for oxiwhen the electrode is placed in a
dation of solid trans-Cr(CO)z(dpe)2(~~hd)
1 : 1 aqueous: acetonitrile mixture containing 0.1 M KCI. Upon oxidation
of solid tran~-Cr(CO)~](dpe)~
in this mixed solvent medium, quantitative
removal of solid t r a n r - [ ~ r ( ~ ~ ) , ( d ~ via
e ) ~dissolution
]+
is now observed at
a ring electrode potential of -0.5 V versus Ag/AgCl [51]. Partial dissolution
of tvan~-[[Cr(CO),(d~e)~]+
can be observed at lower acetonitrile concentrations. At acetonitrile content of up to 50 per cent, no dissolution of
neutral t~ans-Cr(CO)~(dpe)~
could be detected on the ring electrode when
the ring potential was adjusted to 0.1 V versus Ag/AgCl and the frequency of
rotation varied between 100 and 1o4 min-' . Therefore, the electrochemical
oxidation of t~ans-Cr(CO)~(dpe)~
in aqueous acetonitrile mixture generates a
tvan~-[Cr(CO)~(dpe)~]+
salt which is followed by dissolution of this species.
404
Solid-electuode-solvent intefaces
Fig. 5.54 GC/GC R R D E voltammograrns obtained at a scan rate of 2 0 m ~ s - *and frequency
of rotation of 1000 min-' for t r ~ n + C r ( C O ) ~ ( d p attached
e)~
to the disc electrode in contact with
aqueous 0.1 MKCI electrolyte. (a) Disc current versus disc potential; (b) ring current versus disc
potential, with a ring potential of 0.1 V versus Ag/AgCl; (c) ring current versus disc potential, with
a ring potential of -0.5 V versus Ag/AgCl. Reproduced by courtesy: J. Electrochem. Soc 144 (1997)
1566. Copyright, The Electrochemical Society.
Fig. 5.55 GC/GC R R D E voltamn~ogramsobtained at a scan rate of 20 m~ s-* and frequency of
rotation of 1000 min-' for tr~ns-Cr(CO)~(dpe)~
attached to the disc electrode in contact with a 1 : 1
mixture ofwater : acetonitrile containing 0.1 M KC1 electrolyte. (a) Disc current versus disc potential;
(b) ring current versus disc potential, with a ring potential of 0.1 V versus Ag/AgCl; (c) ring current
versus disc potential, with a ring potential of -0.5 V versus Ag/AgCl. Reproduced by courtesy: J.
Electrochem. Soc. 144 (1997) 1566. Copyright, The Electrochemical Society.
Voltammetvy of TCNQ
405
Field (mT)
Fig. 5.56 In situ electrochemical generation of the solution-phase ESR spectrum of
tran~-[Cr(CO)~(dpe)~]'
by oxidation of tran~-Cr(CO)~(dpe)~
at a Pt electrode in a 1 : 1 mixture
of acetonitrile and aqueous K'-buffer. (a) Surface-attached tran~-Cr(CO)~(dpe)~
prior to oxidation; (b) oxidation of trans-Cr(CO)a (dpe)z to trans-[Cr(CO)z (dpe)2]+; (c) regeneration of solid
tran~-Cr(CO)~(dpe)~
by electrochemical reduction of trans-[Cr(C0)2(dpe)2]; (d) re-oxidation of
tran~-Cr(CO)~
( d ~ e to
) ~trans-[Cr(C0)2 (dpe)2If after experiment (c). Reproduced by courtesy: J.
Electrochem. Soc. 144 (1997) 1566. Copyright, The Electrochemical Society.
The solution-phase ESR spectra of trans-[Cr(CO),(dpe)2]+ in acetonitrile
is essentially independent of the water content and the origin of the cation
(synthesized or generated electrochemically). The g value is 2.012 and the
hyperfine splitting parameters a ( 3 1 ~ )= 2.85 mT and a ( 5 3 ~ r )= 0.4 rnT.
Under the thin-layer conditions of the SEESR flat cell, the solid-state
oxidation of trans-Cr(CO), (dpe), to yield dissolved trans- [Cr (CO), (dpe)2]
is chemically reversible in a 1 : 1 mixture of acetonitrile and Kf -buffer.
Figure 5.56 shows a series of ESR spectra recorded from (a) the initial solid
which prowhen surface attached, (b) fully oxidized tr~ns-Cr(CO)~(dpe)~
duces solution-soluble t r a n s - [ ~ r ( ~ ~ ) ~ d p(c)
e ]fully
+ , reduced solution-soluble
t r a n s - [ C r ( ~ O ) ~ ( d ~ ewhich
) ~ ] + produces insoluble trans-Cr(C0)2(d~e)~,
and
(d) again fully oxidized trans-cr(CO),(d~e)~
produced in (c). As in the case of
the [TCNQ]'' process, the dissolved product of the electrochemical reaction
in this case, trans-[Cr(CO),(dpe)2]+, can be almost fully restored to insoluble
tran~-Cr(CO)~(dpe)~
under thin-layer electrochemical cell conditions to produce a thin orange-red coating of solid on the Pt electrode, resulting in almost
complete loss of ESR activity. After oxidation, the full intensity of the ESR
spectrum can be recovered.
+
Solid-electrode-solvent interfaces
Field (mT)
Fig. 5.57 ESR spectra of t r a n ~ - [ C r ( C O ) ~ ( d ~ e ) ~and/or
]+
trans-[Cr(C0)2(dpe)2][A]
(A- = electrolyte anion) generated electrochemically in K+-buffer at a Pt electrode. Dependence
of acetonitrile content. (a) 0% acetonitrile, sensitivity x10; (b) 10% acetonitrile x2; (c) 50%
acetonitrile, x 1. Reproduced by courtesy: J. Electrochem. Soc. 144 (1997) 1566. Copyright, The
Electrochemical Society.
The SEESR behaviour of surface-attached tran~-Cr(CO)~(dpe)~
in different acetonitrile-water mixtures is shown in Fig. 5.57. The potential of the Pt
working electrode was held at a value required for the one-electron oxidation
of tr~ns-Cr(CO)~(dpe)~.
When no acetonitrile is present, the observed broad
single-line ESR spectrum (Fig. 5.57(a)) is consistent with an oxidation process
proceeding solely in the solid state. No evidence of the characteristic solutionis obvious. However,
phase ESR spectrum of tran~-[Cr(CO)~(dpe)~]'(~ol~tion)
the situation changes upon increasing the acetonitrile content. At 10 per
cent acetonitrile content, a spectrum containing characteristics of both the
solid-state and solution-phase ESR spectra is observed (Fig. 5.57 (b)). Upon
further increasing the acetonitrile content to 50 per cent, no solid-state ESR
signal can be detected (Fig. 5.57(c)) since, as shown above, the oxidized
compound dissolves completely. Figure 5.58 shows the single-line solidstate ESR spectra of (a) chemically synthesized trans-Cr(C0)2(dpe)z][I]and
(b) trans- [Cr (CO)z( d ~ e )[A]
~ ](A- = electrolyte anion) electrochemically generated by oxidation of solid tr~ns-Cr(CO)~(dpe)~
when attached to a Pt flag
electrode in a quartz flat cell. Both ESR spectra are similar, having g values of
2.012. Effects of different molecular packing densities in the differently synthesized solid materials are not as pronounced as in the case of the T C N Q system.
Voltammetry of TCNQ
I
I
I
I
I
330
340
350
360
370
407
Field (mT)
ig. 5.58 Solid-state ESR spectra of trans-[Cr(CO)2(dpe)a][A], A- = I- or anion from the electrolyte. (a) Synthesized sample of trans-[Cr(CO)a(dpe)2][I](solid); (b) tran~-[Cr(CO)~
(dpe)2][A](solid)
prepared by electrochemical oxidation of surface-attached trans-Cr(CO)2 (dpe)2 at a Pt electrode in
contact with K+-buffer, sensitivity x 10. Reproduced by courtesy: J. Electrochem. Soc. 144 (1997)
1566. Copyright, The Electrochemical Society.
In summary, the detection of a single-line ESR spectrum observed upon oxidation of surface-attached tr~ns-Cr(CO)~(dpe)~
without any acetonitrile being
~ ] is formed
present simply leads to the conclusion that trans-[Cr ( C 0 ) 2( d ~ e )[Cl]
as product through a solid-state electrochemical process.
In situ R R D E and SEESR measurements demonstrate that microcrystals of
organic T C N Q and organometallic tr~ns-Cr(CO)~(dpe)~
mechanically attached
to electrodes, in contact with solution (electrolyte) media, follow analogous
overall pathways during their solid-state voltammetric reduction and oxidation
processes, despite the substantially different physical properties of the solids.
Some or all of either compound may be dissolved upon electrochemical reaction, depending on the solubility of the product in the solvent and supporting
electrolyte combination. A nucleation/growth-type mechanism may control
the electrochemical behaviour of both surface-confined processes, although the
rate of electrolysis is probably controlled by the conductivity ofthe solids and the
process is probably more confined to the surface in the case of non-conducting
crystals.
408
Solid-electrode-solvent
intefaces
4.4.3 Use o f i n situ atomic force microscopy (2lFh.I) to detect dissolution and
solid-state redistribution veactions that occur during the initial stage of redox cycling
experiments at a solid TCNQ-glassy carbon electrode-aqueous (electrolyte) intefafaceJh
In order to further characterize the processes that occur during the initial stages
of the solid-state T C N Q voltammetry, i n situ rather than e x situ forms of
microscopy are needed. The in situ technique of AFM described in Section 19.1
in Chapter 2 is ideally suited to examine the morphological changes that occur
on the electrode surface in the initial stage of potential cycling experiments.
In the studies described in detail in reference [56], an AFM, operating in
the contact mode, was employed to image T C N Q mechanically attached
to a 3-mm diameter GC electrode. In all cases, simultaneous recordings of
the cyclic voltammograms and conventional AFM topographical images were
undertaken using microcrystals of T C N Q attached to a GC electrode. The
TCNQ-modified electrode was transferred to the AFM electrochemical cell
and the electrochemical reaction was carried out in the presence of aqueous
0.1 M KC1, CsC1, or Et4NC1 electrolyte. During a typical experiment, an
electrode area of between 1 x 1pm2 and 50 x 50 pm2 was imaged.
Figure 5.59 shows an AFM image (in air) of the particles of T C N Q present
after adherence to a GC electrode. The range of sizes of the attached solid
0Pm
5 pm
10 pm
Fig. 5.59 AFM image obtained in air of almost amorphous TCNQ particles attached to a
GC electrode. Reproduced by courtesy: J. Solid State Electrochem. 4 (1999) 24. Copyright,
Springer-Verlag.
'%Adapted with permission from J.Solid State Electrochem. 4 ( 1999) 24. Copyright, SpringerVerlag.
Voltammetry of TCNQ
(d)
q
0.1
0
-0.1
-0.2
-0.3
-0.4
-200
0
200
409
400
mv
(f)
0.02
0.01
2 -0.01O
-0.02
-0.03
. 5.60 Voltammograms obtained at a scan rate of 100 mV s-"after almost amorphous T C N Q
adhered to a GC electrode is placed in contact with 0.1 M KC1 (a-c) [(a) first, (b) the third, and (c)
the tenth scan respectively] or 0.1 M CsCl aqueous electrolyte (d-!(I [(d) first, (e) the third, and (!(Ithe
sixth scan respectively]. The reference electrode was Ag/AgCl. Reproduced by courtesy: J. Solid
State Electrochem. 4 (1999) 24. Copyright, Springer-Verlag.
particles over the area of electrode imaged is predominantly between 0.1 and
1 pm, but the coverage is lower than used in voltammetric studies described
previously (see SEM image shown in Fig. 5.36).
Figure 5.60(a)-(c) contains a series of voltammograms obtained in the AFM
electrochemical cell for the [ T C N Q ] ~ ' process when the GC electrode is in
contact with 0.1 M KC1 as the electrolyte. Clearly, the initial reduction stages
with this low coverage produces a complex (and non-reproducible) voltammetric response with current crossover being evident afier reversing the scan
direction. However, a stable voltammetric response begins to emerge afier about
five cycles. Figure 5.60(d)-(f) show a series of voltammograms obtained during
the course of redox cycling experiments with solid T C N Q and with 0.1 M
CsCl as the electrolyte. Again the initial voltammograms are very complex, but
after six cycles of the potential the voltammetric response expected on the basis
Fig. 5.61 Sequences of in situ AFM images obtained when the solid [TCNQ]'' system is attached
to a GC electrode in contact with (a-c) 0.1 M CsCl aqueous electrolyte after 0, 2, and 4 potential
cycles between 0.4 V and -0.2 V versus Ag/AgCl at a scan rate of 0.1 V s-'. Reproduced by courtesy:
J. Solid State Electrochem. 4 (1999) 24. Copyright, Springer-Verlag.
of the nucleation-growth mechanism is observed. Additional potential cycling
leads to a small decrease in the magnitude of the current, but the principal
features of the voltammogram are retained.
The AFM images obtained during redox cycling of T C N Q with 0.1 M KC1
as the electrolyte reveal that after two potential cycles between 0.4 and -0.2 V
versus A ~ / A & ~ that
,
the larger sizea particles have either dissolved and/or
been transformed to nanometer-sized oarticles. Figure 5.61 shows a seauence
of AFM images of T C N Q particles ahhered to ;he electrode surface &ring
redox cycling in 0.1 M CsCl. These results show that the three particles in
the centre of the figure progressively decrease in size while at the same time,
new material becomes evident at the top area of this figure. It is therefore
clear that both dissolution and solid-phase redistribution processes, which could
involve reprecipitation or surface diffusion are involved at early stages of electrochemical experiments. After 20 cycles of the potential with 0.1 M CsCl as
the electrolyte, needle-like crystals similar to those seen previously by the ex situ
electron scanning method were evident on the electrode surface as shown in
Fig. 5.62(a) and (b).
Voltammetry of TCNQ
41 1
. 5.62 (a) Topographic and (b) shaded in situ AFM images of the needle-like crystals formed
when the [TCNQ]'/- system is attached to a G C electrode followed by 20 cycles of the potential
between 0.4 V and -0.2 V versus Ag/AgCl with aqueous 0.1 M CsCl as the electrolyte. Reproduced
by courtesy: J. Solid State Electrochem. 4 (1999) 24. Copyright, Springer-Verlag.
(a)
$
0.015
0.010
0.005
0.000
-0.005
-0.010
-0.015
(b)
2
(c)
3
0.015
0.010
0.005
0.000
-0.005
-0.010
-0.015
0.015
0.010
0.005
0.000
-0.005
-0.010
-0.015
ig. 5.63 Voltammograms obtained at a scan rate of 0.1 V s-' after almost amorphous T C N Q
is mechanically attached to a GC electrode which is then placed in contact with 0.1 M Et4NC1
aqueous electrolyte (a) first, (b) third, and (c) tenth cycle of the potential respectively). The reference
electrode was Ag/AgCl. Reproduced by courtesy: J. Solid State Electrochem. 4 (1999) 24. Copyright,
Springer-Verlag.
When 0.1 M Et4NCl is the electrolyte, significantly different voltammetric
data (Fig. 5.63) were obtained, relative to the case when a metal ion containing electrolyte was used. In this case, the very first scan is exceptionally well
defined and only a very small decrease in the peak current was observed on
subsequent cycling of the potential. AFM data also revealed that for this system
predominantly only the smaller particles were voltammetrically electroactive
(Fig. 5.64). The reactivity of the particles is random, but data suggest that the
probability of reaction of the smaller sized particles is higher than for the larger
sized particles. When the potential was maintained at +0.5 V versus Ag/AgCl,
0
17.12
urn
34.25
Fig. 5.64 Irr situ AFM images ofthe solid [TCNQ]'' system attached to a GC electrode in contact
with 0.1 M Et4NC1 aqueous electrolyte after (a) 0, (b) 1, and (c) 2 cycles of the potential at a scan
rate of 0.1 V s-* between 0.5 and -0.2 V versus Ag/AgCl. Reproduced by courtesy: J. Solid State
Electrochem. 4 (1999) 24. Copyright, Springer-Verlag.
very little change occurred in the AFM images over very long time periods
(about half an hour). In contrast, when the potential was maintained at -0.2 V
versus Ag/AgCl, the larger particles are removed from the surface, but only
on the minutes rather than seconds time-scale of voltammetry (Fig. 5.65). The
is the
C electrolyte
~
different electrochemical behaviour of T C N Q when E ~ ~ N
is attributed to very low solubility of the [Et4N+][TCNQ-] salt relative to the
metal ion salts.
A detailed examination of the electrode surface, with high resolution AF
conditions, revealed that in all cases T C N Q particles of submicron size are
transformed into a thin layer of material during the initial stages of the redox
cycling. The structure of this thin layer of material depends on the electrolyte
used. When KC1 is used as an electrolyte, a very regular surface coverage develops over the electrode ('bricks in a wall' model). When CsCl is the electrolyte a
very different morphology was observed (dendritic shape). Apparently, during
the course of redox cycling experiments on a GC electrode, the initially unstable solid phase of T C N Q is transformed into at least two more energetically
stable solid phases of TCNQ; initially a solid thin layer is formed from which
the crystalline form with a needle-like shape is evolved.
7
Voltammetry of TCNQ
41 3
ig. 5.65 Sequence of shaded in situ AFM images obtained after (a) 0 s, (b) 162 s, (c) 324 s, and
(d) 486 s when initially almost amorphous T C N Q is attached to a GC electrode in contact with
0.1 MEt4NCl is the electrolyte and the potential is stepped from 0.5 to -0.2V versus Ag/AgCl.
Reproduced by courtesy: J. Solid State Electrochem. 4 (1999) 24. Copyright, Springer-Verlag.
The simultaneous in situ AFM experiments demonstrate that a dissolution
reaction makes a significant contribution to the early stage of voltammetric
redox cycling experiments involving the chemically reversible conversion of
T C N Q to the reduced salt. Thus, features of the voltammetry of T C N Q
mechanically attached to an electrode surface may be rationalized as follows:
(1) The non-reproducibility found with the first few cycles of the potential is
associated with the initial formation and then partial dissolution and reprecipitation of reduced [TCNQ-I . This dissolution-reprecipitation process
enables the transformation to occur from a relatively higher energy surface
state of immobilized and almost amorphous solid to a lower energy surface
state of attached rnicrocrystals. The latter form of attached microcrystals
undergo significantly less dissolution during the course of redox cycling at a
scan rate of 0.1 V s-' and provide the ideal voltammetric response associated
with the nucleation-growth solid-state transformation shown in Fig. 5.26.
(2) After conversion to the microcrystalline state, voltammograms are stable in
the sense that peak positions and wave shape become almost independent of
the number of redox cycles. However, a decrease in the current magnitude
of the response may still arise from loss of material to the bulk solution.
provides a mechanism for the growth of highly
(3) The dissolution
crystalline material on an electrode surface. Initially, immobilized material
is prepared and attached to an electrode via a non-electrochemical process.
The stronger form of adhesion of the electrochemicallyprepared microcrystals minimizes the extent of the dissolution process to a level that at some
electrode suriaces it may become lower than that which is voltammetrically
detectable.
(4) After a significant level of dissolution occurred in the initial stages of redox
cycling, diffusion of reduced and dissolved [TCNQ-] (solution) into the bulk
solution occurred in a competitive process with precipitation of energetically modified suriace-attached microcrystals. Thus, the relative rates of the
physical process of crystallization and diffusion, the strength of adhesion of
solid to the electrode surface and the voltammetric scan rate are important
factors contributing to the nature of the observed voltammetric response.
(5) The presence of an initial dissolution process is consistent with both solution and solid-state phase TCNQ- being detected in the interfacial region
during RRDE and in situ electrochemical-ESR spectroelectrochemical
experiments.
r
Diffusion layer
Amorphous solid T C N Q
Electrochemical dissolution
Electrode
surface
T C N Q + e-
TCNQ-
Diffusion of
TCNQ-
Bulk
solution
Fig. 5.66 Schematic form of representation of the mechanism proposed to account for the electrochemical conversion of almost amorphous particles of solid T C N Q to needle-shaped microcrystals
during the initial stages of redox cycling experiments at a GC electrode. Reproduced by courtesy:
J. Solid State Electrochem. 4 (1999) 24. Copyright, Springer-Verlag.
Voltammetvy of T C N Q
41 5
(6) Figure 5.66 provides a model that can be used to explain the unstable voltammetric responses frequently observed at the early stages of potential cycling
experiments.
(7) After a period of redox cycling, the response shown in Fig. 5.26 can be
achieved. This response can be explained adequately by a simple nucleationgrowth mechanism, where the reduced and oxidized phases can be interconverted without the need to involve the presence of a dissolution process.
(8) The time (cycle number) dependence of the voltammetric response on the
electrolyte composition is also explained by the presence of an initial dissolution process and the subsequent formation of a stable solid phase on
the electrode surface, which is more structurally suited to electrochemical
interconversion between the oxidized and reduced forms of the solid.
4.5 Compavison of electvochemical data with micvocvystals and othevforms of
suface-conjned T C N Q
n addition to studies described above with microcrystals adhered to electrode
surfaces, voltammetric studies have been undertaken on single crystals of T C N Q
and electrochemically prepared T C N Q 'thin films'. Data obtained with these
different forms of adhered solid are presented below and are shown to have
a great deal of similarity, implying that a nucleation-growth mechanism is
operative in all cases.
4.5.1 Voltammetry ofsingle crystals of T T F . T C N Q
One-dimensional organic conductors, such as TTF . T C N Q (where TTF .
T C N Q is tetrathiafulvalenium tetracyanoquinodimethanide), can be used as
electrode materials [57]. In the study of Bartlett [58] single crystals were grown
by interdiffusion of solutions of TTF (tetrathiafulvalene)and T C N Q in acetonitrile and used as electrodes. Figure 5.67 shows a typical cyclic voltammogram in
background electrolyte for the single crystal electrode in contact with the aqueous electrolyte. The electrode response is stable when cycled between -200
and +425 mV versus SCE. Potential excursions outside this range lead to oxiand
dation (generates T C N Q and TTF~') or reduction (generates T C N Q ~
TTF) of the electrode itself and the deposition of insoluble products ( T C N Q or
TTF as appropriate) on the electrode surface. The peaks observed in the cyclic
voltammogram shown in Fig. 5.66(b) have been assigned to the oxidation of
insoluble T C N Q to insoluble [ ~ a ' ] [TCNQ-] at +267 mV and the reduction
of [Na'] [TCNQ-J to T C N Q at -20 mV. Studies revealed that many multilayers of material are involved. Thus, after generation of TCNQ(so1id)by oxidation
of the TTF . T C N Q semiconductor electrode, most features observed in the
voltammetry correspond to those found in voltammetric studies when microcrystals of T C N Q are directly attached to an electrode as described above.
Indeed, cyclic voltammetric and chronoamperometric studies by Scaboo and
Chambers [59] on electrodes prepared from needles of T C N Q . TTF single
r indicate the presence of nucleation
crystals in contact with aqueous 1M
416
Solid-electrode-solvent intefaces
Fig. 5.67 Cyclic voltammograms obtained at a scan rate of 20 rnV s-' for a single crystal
TTF . T C N Q electrode in a pH 7.2 phosphate buffer containing 0.15 M NaC1. (a) Clean electrode;
(b) with insoluble [ ~ a +[TCNQ-]/TCNQ
]
present on the surface formed by a potential excursion
outside the stable range. Potential axis is V vs Sm.Reproduced by courtesy: J. Electroanal. Chem.
300 (1991) 175. Copyright, Elsevier.
and growth mechanisms in the electrochemistry of the [TCNQ]'/- process. In
another detailed study with single crystal electrodes, Bartlett and Tong examined the voltammetry in LiCl media and other conditions where dissolution is
expected and employed in situ electrochemical STM to explore the details ofthe
dissolution processes [60]. In general terms, the voltammograms obtained when
TCNQ(so1id)is generated from T C N Q .TTF single crystals have the same characteristics as those obtained when microcrystals of T C N Q are adhered directly
to the electrode surface.
4.5.2 Voltammetry of TCNQ prepared by oxidation of
films of 9-aminoacridine
Chambers et al. [611 have also extensively examined the voltammetry of 'films'
of solid T C N Q adhered to electrodes in contact with aqueous potassium electrolyte solutions. In these studies, T C N Q was generated by electrochemical
oxidation of polycrystalline films of a 9-aminoacridine hydrochloride salt in
contact with a carbon electrode substrate. Oxidation of these films at 0.6V
versus SCE leads to expulsion of hydrogen ions to form TCNQ(so1id) and to
the now familiar voltammograms of the kind shown in Fig. 5.26. The process
in the 'thin film7-typeterminology considered in Section 18 in Chapter 2 may
be written as
TCNQ(surf)
+ solution) + e ,& [K'] [TCNQ-] (surf)
(5.38)
Voltammetvy of TCNQ
4 17
This T C N Q form of 'film' electrochemistry has been studied by cyclic voltammetry [61], UV-visible spectroelectrochemisty [62] and Q C microgravimetr[63,64]. At slow scan rates, the charges under the cyclic voltammetric waves
correspond to electrolysis of up to 100 monolayers of T C N Q sites, as is the case
with studies emanating from adhered microcrystals of T C N Q . Thus, irrespective of whether T C N Q is adhered to an electrode surface as a microcrystal or
generated electrochemically from 9-aminoacridine, or from a single crystal of
TTF . TCNQ, closely related voltammetry is observed.
Chambers et al. [61] have modelled the 'film7 behaviour in several ways.
Initially, they considered the possibility that a square scheme is operative.
E: (surf)
A+e-
$
B
E; (surf)
C-t-e-
D
EY(~~~Q
The assumption of a square scheme [61,62] with
< ~ i ( s u r t ) provides
,
a
possible means to rationalize the large peak potential separations, while incorporation of phenomenological 'interaction parameters7into the calculation may
be used to account for the narrow peak widths, relative to values expected
with conventional thin film theory (Section 18 in Chapter 2). Effectively, this
approach extends the treatment of Brown and Anson [65], who considered
the extension of a reversible uncomplicated surface-confined redox couple to a
square scheme mechanism.
The second approach used by Chambers et al. assumed that the wave shapes
are a result ofnucleation phenomena and utilized the theory of Camacho and coworkers [66,67] who incorporated nucleation-growth-collision theory into a
treatment of surface cyclic voltammetry. Both the square scheme and nucleation
approaches predict, under suitable conditions, narrow peak-shaped voltarnmograms that display significant hysteresis between the reduction and oxidation
components. Finally, comparison of these interpretations with those emerging
from the unusual quasi reversibility (UQR) concept of Feldberg and Rubinstein
[68] was considered.
Theoretical principles associated with the d f l e m t models considered to
explain the voltammetry of T C N Q (thinfilms'
The voltammograms modelled via the surface version of the square scheme
(eqn 5.39) were calculated using a procedure based on that described in
Section 17 in Chapter 2 when the Nernst and kinetic conditions were adjusted to
their new values after each measurement of the electrode potential. That is, after
each potential step, A E , the surface form of the Nernst condition was satisfied
and then the surface concentrations were adjusted for first-order kinetic decay
of the intermediate species B and C. The rate constants kl and k2 in eqn (5.39)
were chosen to be large enough to produce irreversible voltammograms at the
fastest sweep rate employed. In this case, the peak width of a surface-confined
voltammetric process is 66/n mV at 25OC. Quasi-reversible voltammograms
were calculated in a similar fashion with the assumption of Butler-Volmer
kinetics (Section 18 in Chapter 2). In order to match the experimental and
calculated voltammograms, it was necessary to assume different values of the
~
double-layer capacity in the oxidized, C:;, and the reduced, C T states.
Following Brown and Anson [65], non-ideality effects were Introduced via
interaction parameters ei = R T r T r i / F with units of volts in the expression
In this equation for the A/B couple, the summation of the surface coverage is
taken over all four species, A, B, C, and D in the square scheme.
is the total
surface coverage, Qi is the fractional coverage for species i, and ri is an adjustable
parameter that identifies the magnitude and nature (attractiveor repulsive) of the
interaction. A similar expression was used for the C/D couple with a different
E'(SUI+)
value. The critical condition that leads to narrow voltammetric waves in
this model occurs when species in the same oxidation state have greater attractive
interactions than species in different oxidation states. This requires that the r
values are greater than zero. In the calculations presented by Chambers et al. [61],
a single interaction parameter, ci,was used to define the surface voltammograms.
As expected, when the interaction parameters are set to zero, the solution to the
equations matches the analytical solution of Laviron [69] which was described in
Section 17 in Chapter 2. A slightly different treatment of the interaction model
has been provided by Chidsey and Murray [70] who employed a statistical
mechanics approach to introduce the interaction parameter.
In the second model considered by Chambers, the voltammograms were
calculated using eqn (5.43) described in reference [GI.
1=
*=P
RTv
exp
* [$]
In this equation, which is derived [67], assuming a two-dimensional surfacenucleation process, & is the charge required to electrolyse a surface monolayer,
b is a kinetic parameter, v is the sweep rate, and X = ( F / R T ) (E - EO (surf)). The
kinetic parameter is derived assuming that the surface redox reaction is kinetically controlled by Butler-Volmer kinetics and that the reaction proceeds at the
perimeter of expanding nucleation centres. The phases that grow correspond to
the two redox states and hence when the growing centres intersect, the current
falls back to the baseline. A unity n value is assumed, other terms have their
usual significance and the positive sign corresponds to oxidation and the negative sign to reduction. In applying this equation to the T C N Q voltammetry,
Qm was set equal to the charge under the peaks in voltammograms. However,
it needs to be noted that the two-dimensional nucleation assumption employed
in this theory is only approached at very low surface coverages.
Voltammetry of TCNQ
419
Theory-experiment comparisons based on voltammetric models
igure 5.68 shows an example of a voltammogram calculated via the 'square
scheme' model described in eqn (5.39) using positive interaction parameters.
The open circles are a 'best fit' calculated voltammogram obtained by trial-and) , (surf), Ci;, c g d ,tox,
and tred
for a
error manipulation of the charge, ~ P ( s ~ r fE:
articular set of experimental conditions. Further details of the calculations and
L
values of the parameters are available in reference [61].
The agreement between the experimental and calculated ('square scheme'
model) voltammograms clearly is not completely satisfactory because the theoretically derived voltammograms exhibited greater currents on the rising part of
the waves. Thus, to match the peak height and the peak width at half-height, the
values of the total charge used in the calculations were arbitrarily adjusted to be
significantly larger than the experimental values [61]. Furthermore, the t values,
which are due to attractive forces between redox sites of like oxidation state,
and which should not be sweep-rate dependent, varied in a systematic fashion
with sweep rate. This variation of the calculated E values with sweep rate must
be regarded as a failure of the model to account for the voltammetric behaviour.
However, it should also be noted that the fit for a classical quasi-reversible 'thin
film' voltammogram with n = 1, assuming Butler-Volmer kinetics (Section 18
in Chapter 2) iseven poorer than that se& in Fig. 5.68 [61].
I
E (mV versus SCE)
Fig. 5.68 Comparison of cyclic voltammograms obtained for the solid [TCNQIOI- process
(rT= 7.3 f 0.7 x lops mol ~ m - at~ a) scan rate of 10 mV s-' (filled circles) with a 'film' adhered
to a GC electrode of area 0.0707 cm2 in contact with 1.0 M potassium acetate with theory based on
the square scheme. Open circles are calculated current values using the 'square' scheme model with
p d - 39 mV and &Ox = 36mV. Reproduced by courtesy: J. Electrochem. Soc. 143 (1996) 3039.
Copyright, The Electrochemical Society.
The second explanation for the narrow peak widths and large peak separations considered by Chambers et al. [61] is the two-dimensional nucleation
model. For two-dimensional nucleation, and for overpotentials of less than
k 5 0 mV, the theoretically predicted dependence of the peak current, the peak
width, and peak separation on the scan rate are u0.67L J O . ~ , and vO.', respectively.
Comparison with the experimental data [61] reveals that peak current scan rate
dependence (approximately
approaches that of theory. However, the peak
half-height widths and especially the peak potential separations are significantly
less dependent on sweep rate than predicted by this theory [61].
Figure 5.69(a) and (b) shows examples of the reduction and oxidation components of typical voltammograms calculated using two-dimensional nucleation
theory. The open circles are best-fit obtained by manipulation of the following
parameters: charge, EY(surt), E;(;urt), C:;', C E ~ ,box7 and bred. The agreement
between experiment and theory is considerably better than seen for the voltammograms c&ulated assuming the 'square7scheme model. The agreement was
especially good for the oxidation process (Figure 5.69(b)),and for both the oxidation and reduction processes the agreement on the rising part of the processes
is excellent. Detailed listings of parameters extracted from the experimental
voltammograms are available in reference [61] and reveal that the nucleation
model does account for the potential dependence of the individual oxidation
and reduction process in that the calculated ~ ' ( s u r t ) values are constant. However, unfortunately, the b values are not independent of either the sweep rate or
(b)
300
200
n
-3
'r
100
I
I
300
200
E (mV versus SCE)
100
0
E (mV versus SCE)
Fig. 5.69 Comparison of experimental and theoretical (nucleation model) voltammograms obtained
for the solid [TCNQ]'/- grocess. (a) Filled circles: reduction component, scan rate = 5 mV s-',
r~ = 11.5 f 0.6 x 10- mol crnp2; open circles: calculated using two-dimensional nucleation
theory, ele surf) = 55 mV, b = 7.0 x loP5v2sP2. (b) Filled circles: oxidation component,
scan rate = 5 mV s-', rT = 11.5 f 0.6 x 1 0 - h o l crnp2; open circles: calculated using
two-dimensional nucleation theory, E; (surf) = 210.5 mV, b = 8.5 x 10-6 v2sP2. Reproduced
by courtesy: J. Electrochem. Soc. 143 (1996) 3039. Copyright, The Electrochemical Society.
-
Vo'oltammetvy of TCNQ
I
42 1
I
0
1
Fraction oxidized
Fig. 5.70 Plot ofpotential versus fraction of [TCNQ-] (solid)oxidized to TCNQ(so1id)via the square
scheme with surf surf) = 60 rnV and ~&(surt) = 220 mV. The dotted curves were calculated using
these E' (surf) values and two-dimensional nucleation theory with b =
v2s - ~ .The curves
labelled with N are Nernstian and the dashed curve corresponds to the hypothetical N-shaped
free-energy curve of Feldberg and Rubinstein [68]. Reproduced by courtesy: J. Electrochem. Soc.
143 (1996) 3039. Copyright, The Electrochemical Society.
the surface coverage, so in this context and others, as noted above, this model
does not appear to adequately explain all features of the voltammetry.
Feldberg and Rubinstein [68] introduced the idea of U Q R , and this is the
third model considered by Chambers et al. [61]. This concept attempts to
rationalize cyclic voltammograms with significant hysteresis between coupled
oxidation and reduction peaks by invoking an N-shaped free-energy curve.
This pathway is depicted in Fig. 5.70 by the dashed line. This figure also shows
the Nernstian pathway (labelled N) for a square scheme along with a nucleation
pathway calculated using typical b values.
Figure 5.70 clearly shows that the nucleation pathway is distinctly different
from a hypothetical N-shaped free-energy curve. For a constant value of the
kinetic parameter b, the nucleation pathway predicts a v0.4 dependence of the
peak potential separation on the scan rate that arises from the assumption of
Butler-Volmer kinetics [66]. In contrast, the UQR model predicts that the
value of E,O" is independent of scan rate. Inspection of the results [61]
shows that' the experimental result is intermediate between these expectations
so that again, agreement between theory and experiment with the UQR model
is not adequate.
~r~
Compavisons ofpotential step (chronoamperometric) expen'ments and theory
based on two- and three-dimensional nucleation models
Use of the chronoamperometric current-time response resulting from a potential step avoids the need to consider terms associated with electron-transfer
kinetics and this simplification allows different nucleation mechanisms to be
more readily distinguished than under voltammetric conditions [711. Figure 5.71
shows a typical current transient experimental response for the reduction of a
'thin' T C N Q film prepared as described above by Chambers et al. [61], after
subtraction of the charging current. The signature nucleation response, where
the Faradaic current increases from zero at t = 0 to a maximum, then falls back
towards the base line, is clearly evident in the 'thin' film case as is the case when
microcrystals of T C N Q are adhered to an electrode surface.
The experimental current-time transients have been evaluated by Chambers
et al. [61] using two- and three-dimensional nucleation models (Fig. 5.71) and
it can be noted that a significantly improved fit with the data was obtained using
the three-dimensional instantaneous nucleation theory of Isaev and Baraboshkin
[72] (Fig. 5.71).
Fig. 5.71 Chronoamperometric responses when the potential is stepped from 350 to -90 mV versus
SCE for the solid-state [TCNQ]'' reduction process (after subtraction of charging current) when a
'film' of T C N Q is adhered to a GC electrode and which is then placed in contact with 1 M potassium
acetate. Filled circles are the experimental data, open circles, which closely match the experimental
data, are calculated data using a three-dimensional nucleation model, and the triangles are calculated
data using a two-dimensional nucleation model. Reproduced by courtesy: J. Electrochem. Soc. 143
(1996) 3039. Copyright, The Electrochemical Society.
Voltammetry of TCNQ
423
Conclusions related to the electrochemistry of TCNQ
d to electrode sufaces
A]] electrochemical data obtained on the [TCNQ]'/-
couple are similar irreve of whether adhered rnicrocrystals, a TTF T C N Q electrode or
ochernical oxidation of 9-aminoacridine is used as the source of solid
Q . The electrochemical data coupled with spectroelectrochemical and
microscopy data all imply that a nucleation-growth process is associated with
the rate-determining step, although many complexities are associated with
conversion of solid T C N Q to a [TCNQ-] salt and vice versa.
In the case of T C N Q voltammetry, it is doubtful that a 'thin film' description
of the solid adhered to an electrode surface is adequate and clearly dissolution processes and redistribution reactions occur during the course of potential
ling experiments. The modelling of results obtained by cyclic voltammeare therefore exceptionally demanding. Thus, applications of the theory of
anchez-Maestre et al. used in Section 4.5 involves assumptions that do not
correspond exactly to experiments with TCNQ. Thus, this theory is only valid
for relatively low overpotentials of f5OmV and not the large overpotentials
encountered experimentally. Furthermore, it is unlikely that the nucleation
henomenon that is operative is two-dimensional as assumed in this theory.
ather, it is reasonable to assume that the nucleation sites originate at the
de-microcrystal interface and grow into the bulk of the solid in a threedimensional manner that allows access for the ion motion which accompanies
the electron transfer processes. O n this basis it is expected that incorporation
of three-dimensional instantaneous nucleation assumptions into a theory for
cyclic voltammetry is required. Even then, it is doubtful that such a model
would provide 'perfect' agreement with experiment. Thus, significantly, more
development of the theory is required to fully model the plethora of events that
occur at the solid-electrode-solvent (electrolyte) interface under conditions of
cyclic voltammetry and potential cycling experiments.
The discovery of nucleation-growth mechanisms in TCNQ(so1id) electrochemical processes and details of morphology changes illustrate the complexity
of the structural factors that may be involved in the electrochemistry of micro or
even nanocrystals adhered to a surface. The rich variety of physical phenomena
that may be difficult to unravel from sole reliance on solid-state voltammetric
experiments is also clearly revealed in the TCNQ(so11d)studies. It is anticipated
that in future, the combination of immobilized microcrystal voltammetry, flowcell reaction order measurements, double-potential-step experiments, XRD,
and microscopy will be successfully applied to other systems, and thereby
uncover the existence of many more nucleative solidsolid phase transformations, as well as shedding further light on the relationship between voltammetric
inert zones, specific interfacial free energies, and interfacial structure. For example, nucleation and growth kinetics are probably present in many systems in
which they have not been previously suspected. Cases in point include the electrochemical intercalation of metal ions into solid fullerenes [73] where again a
large voltammetric inert zone is present and in the oxidation of tetrathiafulvalene [74]. The same class of mechanism would also now appear to be associated
with the decamethylferrocene and Cr(C0)2(dpe)2electrochemistries and may
also apply to insertion of ~ i that
+ is important in a range of batteries [75] as well
as many other systems where inert zones have been reported.
ion transport wit
determining
where coupled electro
microparticle are rate
If nucleation growth is not rate determining, then a mechanism that is likely to
limit the overall rate of charge transport within a solid is the electron hopping
or counterion diffusion/migration processes that must occur to achieve charge
neutralization. Studies on adhered solids that probably have layered structures
have revealed behaviour converging towards 'thick film' three-dimensional
diffusion where coupled electron and ion movement within the solid is rate
determining [76,77].
(Fig. 5.72(a)) complex, where
In the case of the [O~(bpy)~-4-tet-Cl][ClO~]
microparbpy is 2, 2'-bipyridyl and 4-tet is 3,6-bis(4-pyridy1)-1,2,4,5-tetrazine,
ticles that have a layered structure have been formed on graphite macrodisc
and platinum microdisc electrodes [76]. Images of the solid obtained by SEM
reveal that repeated voltammetric cycling of the potential, when the modified
electrode is in contact with aqueous sodium perchlorate electrolyte, induces
1
n+
II
I
N /N
6
Fig. 5.72 Structures of (a) [Os(bpy)a-4-tet-C1](C104) where bpy is 2, 2'-bipyridyl and 4-tet is
3,6-bis(4-pyridy1)-1,2,4,5-tetrazine
and (b) [ { ~ - ( b ~ ~ ) ~ } {( p~- '~ () ]bn +~for~ the
) ~ particular
}
case
when M = R u and M' = R u or more generally M or M' = R u or 0 s ; bpy = 2, 2'-bipyridyl; L =
1$-dihydroxy-2,5-bis(pyrazo1-1'-yl)benzene dianion) .
Voltammetric studies on systems
425
some crystallization on the electrode surface. However, the voltammetric
response does not indicate the presence of a nucleation-crystal growth process
when rnicroparticles of solid [Os(bp~)~-4-tet-C1]
[C104]( 0 s 2 + ) is oxidized to
solid [O~(bpy)~-4-tet-C1]
( ~ 1 0(0s3+).
~ ) ~ In contrast, with perchloric acid as
the electrolyte, the rnicroparticles remain apparently 'amorphous' even after
several thousand voltammetric cycles of the potential have been undertaken.
[C104]
The voltammetric responses associated with the [O~(bpy)~-4-tet-C1]
complex are close to ideal apparently diffusion-controlled processes when
undertaken with the solid adhered to a platinum microdisc electrode. Thus,
with a 5-pm radius electrode and at fast scan rates, 'linear diffusion' is dominant, as evidenced by the observation of peak-shaped processes with a square
root dependence on scan rate (Fig. 5.73), and the shapes and general characteristics of cvclic voltarnmo~ramsare similar to those observed for com~letelv
solution-phase reactions described in Section 8 in Chapter 2 with stationary
rnacrodisc electrodes. However, in this case, the current response is controlled
by coupled electron transport and diffusion of ClO; within the solid17 rather
than diffusion of the anion within the solution phase. As expected with a 2-pm
radius electrode and at low scan rates, the current is controlled by a 'radial diffusion' as evidenced by the sigmoidal shape near steady-state respbnse (Fig. 5.74)
which is also characteristic of solution-~hasevoltammetrv at microdisc electrodes. The hysteresis in these slow scan rate data shown in Fig. 5.74 may be
indicative of small changes in the film structure.
Using diffusion-controlled theory, the data obtained from the responses
shown in Figs 5.73 and 5.74 (linear and radial diffusion) may be combined
to determine the effective concentrations, Ceff, of redox centres within the
'film' (1.8 and 1.6 M for NaC104 and HC104 electrolytes, respectively) and
the 'apparent' diffusion coefficients,18 D,,,, for homogeneous charge (coupled
electron and ion) transport (2.7 x lo-" and 5.0 x lo-" cm2 s-' for NaC104
and HC104 electrolytes, respectively). This requires adaption of eqns (2.34)
(Randle-Sevcik theory) and (2.48) (microelectrode theory) that have previously
been presented in Chapter 2 to g v e eqns (5.42) and (5.43)
J
I
U
I
J
I
where Dappis the apparent diffusion coefficient and Ceffis the effective concentration of 0 s 2 + centres adhered to the electrode surface, I,, is the steady-state
limiting current and all other terms have their usual meaning given in Chapter 2.
Thus, in principle, it is possible, when microparticles are adhered to an electrode
surface, to have a model of charge transport that resembles behaviour also found
17within the solid structure, electron transport by electron hopping and ion movement
must be coupled to achieve charge neutrality. Thus, the apparent diffusion coefficients
represent the coupling of the electron and ion charge-transport processes.
''see n. 17.
426
Solid-electrode-solvent
(a)
intefaces
0.31
Fig. 5.73 Scan rate dependence of the voltammetric response when [O~(bpy)~-4-tet-Cl]ClO~
mechanically attached to a 5-pm radius Pt microdisc electrode is placed in contact with aqueous (a)
1.0M HC104, (b) 1.0 M NaC104 electrolyte. From top to bottom, the scan rates are 1000, 500,
200, and 100 mV s-l. Potential axes are versus SCE. Plots of peak current versus scan rate are linear
and pass through the origin as required for a diffusion-controlled process. Reproduced by courtesy:
J. Phys. Chem. B 104 (2000) 6389. Copyright, American Chemical Society.
in conducting polymers containing redox active centres [78]. Surface coverages
studies correspond to several
calculated in the [O~(bpy)~-4-tet-Cl]ClO~(solid)
hundred monolayer equivalents and almost 100 per cent of the osmium centres
are active on the voltammetric time-scale.
Raman spectroscopy can be employed to identify the nature of the changes
accompanying the surface-confined
oxidation process. Figure 5.75(a)
and (b) shows the Raman spectra obtained when solid [O~(bpy)~-4-tet-Cl]ClQ~
Voltammetvic studies on systems
427
ig. 5.74 Steady-state voltammograms recorded at a scan rate of 1 rnVsfl when
[Os(bpy)z-4-tet-C1]C104 adhered to a 2-pm radius Pt microdisc electrode is placed in contact with 1.0 M HC104 and 1.0 M NaClO4. Potential axis is versus SCE. Reproduced by courtesy:
J. Phys. Chem. B 104 (2000) 6389. Copyright, American Chemical Society.
Fig. 5.75 Raman spectra obtained when [O~(bpy)~-4-tet-Cl]ClO~
is mechanically attached to a
gold disc electrode and the potential is held at (a) -0.2 V in 0.1 M NaC104, (b) 0.4 V in 0.1 M
NaC104, (c) -0.2 V in 0.1 M HC104, and (d) 0.4V in 0.1 M HC104. All potentials are versus SCE.
The laser excitation wavelength is 632.8 nm. Reproduced by courtesy: J. Phys. Chem. B 104 (2000)
6389. Copyright, American Chemical Society.
428
Solid-electrode-solvent intefaces
is immobilized on a gold electrode and the potential is held at -0.2 and 0.4 V
versus SCE, respectively while the modified electrode is in contact with aqueous 0.1 M NaC104 electrolyte. The exciting HeNe laser (632.8 nm) used in
these Raman experiments is expected to be preresonant with the 0s2+-bpyn*
metal to ligand charge transfer (MLCT) transition and this is confirmed with
the enhancement of features at 1604, 1550, 1480, 1320, 1268, 1167, and
1015 cm-' , all being associated with the bipyridyl moieties. A weaker feature
at 373 cm-' is associated with the 0 s 2 + - N stretch. Two very intense features at
811 and 934 cm-' are attributed to surface-enhanced transitions, possibly due to
the tetrazine or its associated free pyridine. These bands are not observed when
the bare solid is examined on glass or when the bare electrode is examined,
but are observed when the material is adhered to a gold surface and are largely
unaffected by switching the oxidation state of the layer. Figure 5.75(c) and (d),
illustrates the Raman spectra of the solid material when the electrode potential
is held at -0.2 and 0.4 V versus SCE respectively, when aqueous 0.1 A4 HClO,
is the electrolyte. In the acid solution, the pyridine group is expected to be
fully protonated. The only significant difference between the acidic and neutral medium is the 0s2+-N mode which has shifted to 383 cm-' under acidic
conditions. In each case, oxidation of the solid material results in loss of the
low-frequency 0s2'-N vibration, consistent with metal oxidation. If oxidation
of the metal results in the loss of the MLCT it might be expected that the resonance condition would be lost. However, after oxidation, the laser becomes
preresonant with a bpy (n)-0s ligand to metal charge transfer (LMCT) transition and the bpy-based bands continue to be observed even for the oxidized
layers, albeit with different relative intensities.
While apparently 'diffusion'-controlled processes are observed for the oxidation of the [O~(bpy)~-4-tet-Cl]ClO~
complex, adhered to an electrode, the electrochemistry of a series of 'electrochemically open'
[{M(bpy)2){M'(bpy)21
(p-L)](PF,)n complexes (M, M' = R u , 0 s ; bpy =
2, 2'-bipyridyl; L = 1,4-dihydroxy-2,5- bis (pyrazol-1-yl)benzene dianion;
n = 2 or 3) at a solid-electrode-aqueous electrolyte interface produces what apparently approaches 'thin film' behaviour [79]. Synthesis of [{M(bpy)2){M'(bpy)z}
(p-L)](PF& compounds (structure shown in
Fig. 5.72(b)) does not yield well-defined crystalline compounds, and inclusion
of solvent and even salts (KN03)has been detected [80]. SEM images (Fig. 5.76)
reveal that these solids readily disintegrate into small submicrometre-sized
~articles.to ~ i v ea 'thin film' or laver-like structure.
The solid-state voltammetry bf the [ { R ~ ( b p y (p-L)]
) ~ ) ~ (PF6)2 (RuRu),
[{os(b~~)2}2(~-L)1(PF6)3
(OSOS), and [{Ru(bpy)2){os(bpy)2)(~-L)](PF6)3
(OsRu) complexes is associated with reversible hydroquinone, semi-quinone,
J
U
Voltammetric studies on adheved micvopavticles
429
. 5.76 SEM images of surface-attached solid [{R~(bpy)~)~(p-L)l(PF~)~
shown in (a) high and
(b) lower magnification. Reproduced by courtesy: J. Phys. Chew. B 104 (2000) 1977. Copyright,
American Chemical Society.
and quinone interconversions of the bridging dioxolene ligand and not a metalbased process as is the case with the [O~(bpy)~-4-tet-Cl]ClO~
compound
[77]. Cyclic voltammograms obtained when the complexes are mechanically
attached to a basal-plane pyrolytic graphite electrode, which is placed in contact with aqueous 0.1 M KPF6 electrolyte, exhibit two extremely well-defined
rocesses which may be modelled by two overlapping Gaussian-shaped or 'thin
layer'-type curves (Fig. 5.77). In general, these processes may be written as in
eqns (5.44) and (5.45).
The individual peak potentials change with scan rate, but as expected, the
reversible half-wave potentials, EiI2, measured as the average of the oxidation and reduction peak potentials remain essentially constant [79]. The
peak currents increase almost linearly with the scan rate as expected with
a 'thin film' and not on the square root of scan rate as in the 'diffusion'controlled [O~(bpy)~-4-tet-C1]
[C104] oxidation process. In conjunction with
results from Raman spectroscopic [79] analysis of the voltammetric response
implies that extensive electrolysis of a thin particulate deposit of compound
occurs at the electrode surface on the voltammetric time-scale. Support for
this interpretation is also gained by noting that the calculated integrated
charge for both R u R u processes (Fig. 5.77) is 0.8 mC (corresponding to 6 pg
irrespective of the scan rate over the range of
of [{R~(bpy)~}~(p-L)l(PF,),
1mV s-' to 0.1 V s-l.
Data obtained with adhered solid and from dimethylformamide [80] solutionphase voltammetric experiments are compared in Table 5.10. Significant
differences are detected and the dependence of E;,2 for the solid-state process
on the identity and concentration of the electrolyte is characteristic of solid
I
I
I
I
-0.5
0.0
0.5
1.0
E (V vs SCE)
I
I
I
-0.6 -0.4-0.2
I
I
I
I
I
0.0 0.2 0.4 0.6 0.8
E (V vs SCE)
Fig. 5.77 Cyclic voltammograms (scan rates 100, 50, 20, 10mVs-l) for the oxidation of solid (a) [{Ru(bpy)2)2(p-L)](PF6)2, (b) [{Os(bpy)2){Ru(bpy)2)2(P-L)](PF6)3, and
(c) [{0~(bpy)~}~(p-L)](PF~)~
attached to basal-plane pyrolytic graphite electrodes and then
placed in contact with aqueous 0.1 M KPF6 electrolyte. Cyclic voltammograms (scan rates
10 mV s-*) with Gaussian curves fitted for the oxidation of solid (d) [{RU(~~~)~}~(~-L)](PF~)~
(e)
]
Reproduced by
[{Os(bpy)21{Ru(bpy)21(P-L)](PF6)3 and (g [{Os(bpy)212{Ru(bpy)212 ( P - ~ )(PF6)3.
courtesy: J. Phys Chem. B 104 (2000) 1977. Copyright, American Chemical Society.
state rather than solution-phase [79] processes. Furthermore, considerable differences in the level of stabilization of the semi-quinone intermediate are seen
in the solution-phase voltammetry and the solid-state case as evidenced by the
significant variation in AEiIz values obtained on comparison of these data in
the different phases.
(p-L)] (PF,), solids
In summary, the conclusion that [{M(bpy)2J{M'(bpy)2)
are 'electrochemically open' has been reached on the basis of the high currents
observed from essentially bulk conversion of the solid on the voltammetric timescale and the absence of rate-determining nucleation phenomena associated
with the formation of new phases. The form of structural features responsible for the fast electron and ion transport and the behaviour approaching that
expected for 'thin layer'-type voltammetry appear to be a molecular structure that readily accommodates rapid electron and ion transfer, large molecular
ions of low symmetry which aid Coulombic cation-anion-electron interaction
and ion exchange across the solid-aqueous solution boundaries, and structural
packing which favours salt or solvent inclusion [79].
e 5.10 Reversible potentialsa obtained from cyclic voltammograms of [{M(bpy)2]{M'(bpy)2}(p-L)](PF6),
dissolved in dimethylformamide (0.1 M Bu4NBF4)at a platinum disc electrode and as a solid adhered to a basal-plane
pyrolytic graphite electrode in contact with aqueous 0.1 M KPF6 electrolyte.
Compound
Solution phaseb
2
Solid phasec
h
Ru Ru
0 s Ru
0 s0 s
Process I
E;/2(V)
Process I1
Ei12(V)
AEl12
(V)
0.13
0.19
0.23
0.54
0.44
0.42
0.41
0.25
0.19
Process I
Process 11
AE1/2
$
AEJJ(V)
E;/, (V)
w / 2 (V)
AEp (V)
E;/2 (V)
w / 2 (V)
(V)
3
2.
0.17
0.13
0.12
0.09
-0.01
-0.01
0.19
0.23
0.15
0.13
0.12
0.13
0.38
0.31
0.12
0.19
0.23
0.15
0.29
0.32
0.13
2
X
"Reversible half-wave potentials (Eii2) are V vs Ag/AgCl and calculated as the average of the oxidation and reduction peak potentials, AE1/2 =
separation between processes 1 and 2, WIi2 is the half-peak-height width, AEp is the separation of the oxidation and reduction peak
potentials.
' ~ a t ataken from reference [79]. Scan rate = 0.1 V s-'.
'Data taken from reference [80]. Scan rate = 0.01 V s-l.
3.
n
%
2
0
X
Q
3
rc
%
Y
3
F;-
24
2
.
n
E+
6.1 Corrzparisorz with the voltammetry ofazurin adhered to
an electrode as an ideal thin j l v n
Clearly the voltammetry of [{M(bpy)2) {M'(bpy)2) (p-L)](PF6),, while having
many of the characteristics of 'thin' film behaviour, is adhered to the electrode
in the form of arrays of microparticles (Fig. 5.76) and not as a continuous film.
It is therefore interesting to see how the behaviour of this system mimics one
that is classically regarded as corresponding to genuinely thin film voltammetry.
As noted in Section 18 of Chapter 2, metalloproteins attached to electrode
surfaces provide close to ideal 'thin film' voltammetric responses. However, as
is the case with voltammetry at microparticle-electrode-electrolyte interfaces,
nuances are present in these so-called ideal 'thin film' systems that also have their
counterparts in the voltammetry of arrays of [{M(bpy)z}{M'(bpy)2}(p-L)] (PF6),
microparticles. T o explore both the excellent conformance to theory, but yet
subtle deviation from ideality, similar to that found with the microparticle 'thin
film' case, the reduction of azurin, considered in Section 18 in Chapter 2 as a
model 'thin film' voltammetric system, is now considered in more detail [81].
A range of methods of forming films of metalloproteins have been described
in the literature [81]. Commonly, a few microlitres of dilute protein solution
are drawn up into a fine capillary and applied directly onto a freshly polished
graphite electrode surface. Scratching the surface with the glass tip induces
adsorption, probably by exposing a fresh 'virgin' carbon surface. Co-adsorbates
may also be added to both the protein and the buffered electrolyte solution
in order to stabilize the adsorbed film, presumably by using positively charged
amino groups to form ternary salt bridges between the protein and the negatively
charged electrode suriace. For studies on the voltammetry of azurin, polymyxin
B sulphate (a complex polypeptide, carrying several amine groups) has been used
in some experiments as a co-adsorbate [81]. If polymyxin was not included, the
voltammetric peaks are broader, indicating that the co-adsorbate plays a role
in directing the formation of a uniform monolayer. An alternative method of
preparing a thin film is to simply place the graphite electrode in contact with
a dilute solution of azurin and allow spontaneous adsorption to occur. These
methods of forming adhered solid are of course significantly different from
~ } (PF6),.
mechanical attachment of microparticles of [{M(bpy),}{ M ' ( b ~ y )(p-L)]
Figure 5.78 shows cyclic voltammograms obtained as a function of scan rate
for azurin surface confined on a graphite electrode in contact with 0.1 M NaCl
electrolyte. At slow scan rates of 10 mV s-' , the voltammetric response appears
to be close to that predicted for an ideal reversible 'thin film' process, since the
reduction and oxidation peak potentials are very similar in value (Section 18.2.1
in Chapter 2). In contrast, at fast scan rates, the shapes and characteristics associated with an irreversible 'thin film' process are evident (Fig. 5.78) and the
resemblance to data for the voltammetry of [{M(bpy)2}{M'(bpy),}(p-L)] (PF6),
(Fig. 5.77) is obvious.
As the scan rate increases, both Butler-Volmer and Marcus theories predict
(as described in Section 18.2.3 in Chapter 2) that the reduction and oxidation
peaks will separate in potential, and also broaden. Voltammograms in Fig. 5.78
Voltammetricstudies on adhered microparticles
433
Fig. 5.78 Voltammograms obtained at O°C for a thin film of azurin adhered to an edge-plane
pyrolytic graphite electrode in contact with 0.1 M NaCl (pH 8.5) at scan rates of 10 mV s-l, 1V s-l,
and 100 V s-' (top to bottom). The y axis has the current scale for the background-corrected peaks
removed for clarity. The x axis is V versus SHE. Provided by courtesy: F.A. Armstrong and J. Hirst,
University of Oxford, England.
confirm these predictions as clearly the peaks separate and broaden when the scan
analogous
rate is increased. In the case of [{M(bpy)2}{M'(bpy)2}(p-L)](PF6),
behaviour is observed as the scan rate increases [79].
Figure 5.79 shows the change in the cyclic voltammograms, at a scan rate of
100 V s-' , for reduction of azurin that is surface confined on a graphite electrode
and when the electrolyte concentration is changed from 0.1 to 2 M NaCl, while
Fig. 5.80 shows the electrolyte concentration dependence for the oxidation of
microparticles of [ { R ~ ( b p y )(p-L)]
~ } (PF6)2.In both cases, analogous behaviour
is seen to be observed. For azurin, the reversible potential is 40 mV more negative in 0.1 M KC1 than 0.1 M NaC1, and also more negative in 0.1 M NaCIOl
than 0.1 M NaCl by 50 mV. A similar variation with electrolyte composition is
observed for oxidation of [{Ru(bpy)2}(p-L)] (PF6)2[79].
Apart from the consequences of the Butler-Volmer and Marcus theories
and uncompensated resistance, there are at least two other possible reasons
why the peaks may display broadening at higher scan rates [81]. The first is
thermodynamic dispersion, since different redox centres may have different
environments and therefore slightly different free energies and hence different
reversible potentials for individual molecules or ions. Analogously, there may be
434
Solid-electvode-solvent
intefacer
Fig. 5.79 Cyclic voltarnrnograms obtained at O°C for a thin film of azurin adhered to an edge-plane
pyrolytic graphite electrode in contact with 0.1 M NaCl and 2 M NaCl (buffered at pH 8.5) at a scan
rate of 100 V s-' . The peak separation in 2 M NaCl (34 mV) is less than that in 0.1 M NaCl (86 mV).
K is the normalized current relative to that expected for a reversible process (Section 18 in Chapter 2)
and the ootential axis is V versus SHE. Provided by courtesy: F.A. Armstrong and J. Hirst, University
of 0xfo;d, England.
I
-0.2
I
0.3
E ( V vs SCE)
I
0.8
Fig. 5.80 (a) Cyclic voltammograms (scan rate 1OmVs-l) for the oxidation of solid
[ { R ~ ( b p y )(p-L)]
~ } ~ (PF6)z attached to basal-plane pyrolytic graphite electrodes and then placed in
contact with aqueous 0.1, 0.02, and 0.005 M KPFh electrolyte. (b) Plot of the peak potentials E r
(process 1 (m) and process 2 (A)) and E r d (process 1 (0)and process 2 (A)) versus the logarithm
of the concentration of KPF6. The dashed lines show the shift in the reversible half-wave potential
(E:~ E,OX)/2. Reproduced by courtesy: J. P h p . Chem. B 104 (2000) 1977. Copyright, American
~ h e m i c Society.
d
+
Voltammetricstudies on adhered microparticles
435
kinetic dispersion in which some individual components can transfer electrons
at faster rates than others. Kinetic dispersion cannot be readily distinguished
(if at all) from Marcus broadening. For azurin, the peak-to-peak separation at
the low scan rate limit is about 10mV rather than OmV and the peak width
is 18 mV greater than predicted by the Nernst or equilibrium calculation [81].
Thus, perfect agreement is not achieved with theoretical expectations even for
this ostensibly well-behaved 'thin film' voltammetry.
In order to model the azurin 'thin film' voltammetry with Butler-Volmer
or Marcus models, corrections have been introduced to account for the nonidealities not accommodated by an electron-transfer-only model [81]. After
introduction of these corrections, the values of the first-order rate constant
for the electron-transfer process (KO) for 'thin film' voltammetry of azurin is
approximately 1000 s-' in 0.1 M NaCl and 5000 s-' in 2 M NaCl (Fig. 5.8 1).
owever, again after correction for non-ideality, agreement between theory
and experiment is still not perfect [81]. Rigorous modelling of the 'thin film'
version of voltammetry when microparticles are adhered to an electrode surface, has yet to be undertaken, but probably would produce an analogous set of
difficulties if a solely electron-transfer model is used, to those encountered with
azurin that have been elegantly addressed by Hirst and Armstrong [811. The suggestion that emerges from the study with [{M(bpy)2}{M'(bpy)2}(p-L)] (PF6)ais
that 'thin film' behaviour can be found in cases where arrays of microparticles are present on a surface. The observation of apparently parallel forms of
utler -Volmer
0.4
0.6
Potential V
Fig. 5.81 Comparison of an experimental voltammogram obtained for azurin which is surface
confined on an edge-plane pyrolytic graphite electrode in contact with aqueous 2 M NaCl electrolyte
and those predicted using the Butler-Volmer theory (ko = 5000 s-I, a = 0.5) and Marcus theory
(A = 0.2 eV and k, = 1.8 x lo5 s-I). The modelled results have been corrected to the same peak
height and the same position as the experimental data to allow an exact comparison of wave shape.
It is apparent that extra broadening is present in the experimental data and that better agreement
is obtained using Marcus theory (with a low reorganization energy or h value). Reproduced by
courtesy: Anal. Chem. 70 (1998) 5062. Copyright, American Chemical Society. This reference may
be consulted for further details of the significance of the kinetic parameters and for further details of
the modelling procedures.
- 4 0 1 ,
-0.6
I
-0.4
-0.2
0.0
0.2
Potential (V)
0.4
1
0.6
Fig. 5.82 Calculated reversible (Ej = OV) cyclic voltammograms of differently shaped cuboid
particles of constant volume Vnl with height (H) and a square base of length (L) and breadth (B),
but of different size. The diffusion coefficients of the electron and cation from the electrolyte used
for charge neutralization are both lo-' cm2 s-l, V, = 153.8 cm3 mol-', (a) L = B = 28 prn,
H = 10 pm, (b) L = B = H = 20 pm, (c) L = B = 16 pm, H = 31 pm. Reproduced by courtesy:
J. Solid State Electrochem. 4 (2000) 314. Copyright, Springer-Verlag.
voltammetry, under what appears at first sight to be inherently different forms
of surface adherence, raises questions as to whether many so-called thin films,
in a microscopic sense, are actually equivalent to arrays of microcrystals, and
as to whether examples of genuine thin film formation on electrode surfaces
are actually quite rare. Finally, it may be noted that in their description of the
voltammetry of a cube-shaped crystal adhered to an electrode surface, Schroder
et al. [28] note that the dimensions of the crystal are predicted to be critical in
determining the shape of a voltammogram when uptake or expulsion of electrons and electrolyte ions are coupled and their mass transport within the solid
is governed by Fick's law of diffusion (see Fig. 5.82).
The electrochemical studies described in this chapter that are related to
rnicroparticles adhered to electrode surfaces reveal that an order of magnitude of
additional complexity is present relative to the case when all species involved in
an electrochemical process are solution soluble. Thus, while qualitative understanding of the processes is relatively readily achieved, quantitative comparisons
with theory frequently tend to be somewhat disappointing if a wide range of
time domains (e.g. scan rates in cyclic voltammetry) are used.
The key processes in the dynamic electrochemistry of solid microparticles,
adhered to electrode surfaces placed in contact with electrolyte, are the coupling of electron and ion transport required to achieve charge neutralization, the
A n overview of techniques
437
movement of ions within and between the solid and solution phases, and any
accompanying morphology or other changes in the two solid phases associated
with the solid-state redox-based conversion process that takes place. Ratedetermining steps identified in this chapter may involve nucleation-growth
processes in the interconversion of oxidized and reduced forms of solids (e.g.
the [TCNQ]'' process), and the diffusion of ions, coupled to electron transport (e.g. the 0s2+I3+process). Even 'thin film' behaviour may exist. Numerous
possibilities of dissolution, redistribution, and morphological changes were also
identified as accompanying or potentially accompanying the solid-solid interconversion of adhered microparticles. The technological importance of studies
of adhered microparticles is obvious in electrochromic devices, photovoltaic
cells, and other solid-state devices such as batteries and fuel cells (Chapter I), so
the rapid increase in studies of solids adhered to surfaces is likely to continue.
The voltammetric and related electrochemical techniques employed in studies
described in Chapter 5 are generally similar to those used in the solution-phase
studies reviewed in Chapters 3 and 4. The techniques and examples of their
uses include:
(1) transient cyclic voltammetry-useful
(2) microelectrode voltammetry-useful
for general mechanistic diagnosis;
for mechanistic diagnosis, particularly
when data are interpreted in conjunction with those obtained from transient
techniques at macrodisc electrodes;
(3) potential step methods-particularly
useful for identifying nucleationgrowth processes;
(4) CPE-used to interconvert oxidized and reduced forms of microparticles
on long time domains;
(5) coulometry-useful for qualifying 'amounts7 of solid being interconverted
on an electrode sudace during the course of voltammetric or potential step
experiments;
(6) RRDE-useful for detection of dissolution processes that may accompany
redox-based solid-solid interconversion processes;
(7) EQCM-provides useful information on mass changes that accompany
surface-based redox processes when ions are involved in the charge
neutralization process and when dissolution of solid occurs.
Characterization of surface-confined processes by in situ and ex situ techniques
used in conjunction with electrochemical methodologies is essentially mandatory if any detailed understanding of inherently complex solid-state processes
is to be obtained. Examples of dual electrochemical and non-electrochemical
techniques presented in Chapter 5 and their uses include:
(1) EPM-qualitative and quantitative elemental surface analysis;
(2) X-ray diffraction-structural characterization of oxidized and reduced forms
of microparticles;
(3) SEM-imaging of crystal morphology changes accompanying solid-solid
interconversion processes;
438
Solid-electrode-solvent interfaces
(4) video camera imaging-monitoring
of colour changes and crystal size and
shape changes of large particles;
(5) infrared, Raman and ESR spectroelectrochemistry-structural information
associated with redox-based solid-solid transformations;
(6) AFM-real-time information on crystal growth and/or dissolution processes
associated with the electrochemistry of microparticles.
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Introduction
As noted in Chapters 3 and 4, voltammetric studies on small redox active
inorganic, organometallic, and organic molecules frequently provide an efficient method of obtaining important thermodynamic and kinetic information.
Logically, therefore, it would be expected that extensive solution-phase
voltammetric studies on important electron-transfer metalloproteins such as
the cytochromes, ferredoxins, rubredoxins, plastocyanins, and azurins would
have been undertaken almost as soon as these water-soluble biologically important molecules had been isolated and characterized. However, it was in fact
only in 1977 that the first reports of well-defined solution-phase voltammograms, which could be used to calculate a reversible potential consistent with
that expected on the basis of potentiometric measurements, were published for
cytochrome c [I ,2]. Prior to that time, metalloproteins were either reported to
be electroinactive under voltammetric conditions or else their voltammetry was
described as so complex (highly irreversible, complicated by adsorption) that no
useful information could be obtained 13-51.
The conundrum of lack of, or ill-defined, solution-phase voltammetry for
metalloproteins that are designed for fast and specific electron-transfer reactions
in biological systems, was explained initially in terms of protein adsorption or
even denaturation at electrode surfaces, very small diffusion coefficients and
hence very low currents that are difficult to measure, or an inability of the
metal centre to approach close enough to an electrode surface to transfer electrons [3-51. However, after about twenty-five years of intensive research, it
has become apparent that voltammetric studies on all electron-transfer metalloproteins, and indeed other highly important redox active metalloproteins
such as myoglobin, in both the solution phase or as thin films (Chapters 2
and 5) can in fact produce equally well-defined chemically and electrochemically reversible responses as expected for small organic and inorganic molecules
where both halves ofthe redox couple are chemically stable. The 'trick' to obtain
well-defined metalloprotein voltammetry is to understand and then model
442
Metalloprotein voltammetry
or minimize the influence of undesirable sudace-based processes commonly
encountered with these highly surface-active biologically important molecules,
as indeed is the case with surface-active ruthenium dyes used in photovoltaic
cells (see Section 4 in Chapter 4).
In this chapter, some of the key results obtained after twenty-five years of
intensive investigation into the electrochemistry of metalloproteins are summarized to highlight both 'heroic failures' and outstanding achievements in this
field. Via this approach, the numerous subtleties and nuances of the technique
of voltammetry (beneficial and deleterious) not considered in Chapters 1-5 can
be introduced. What is noteworthy, is the significant learning that takes place
while understanding the progression from initial literature reports of 'bad' or
no electrochemistry, to a second batch of publications containing phraseology
such as 'transient', 'unstable', 'quasi-reversible', or 'fair' electrochemistry until
ultimately, the latest publications are said to contain 'ideal' or almost 'reversible'
voltammetric behaviour under carefully specified conditions. The evolutionary pathway towards the attainment of 'ideal' voltammetry is presented in the
belief that it provides an ideal teaching tool to elucidate the concept of blocked
electrodes, the use of chemically modified electrodes, the importance of functional groups on carbon electrodes and microscopic aspects of voltammetry
which always need to be considered when a voltammetric response is found to
be unexpectedly complex.
Metalloproteins are characterized frequently by high charges (e.g. +7 to +8
for cytochrome c, -8 to -9 for plastocyanin, and -18 to -19 for ferredoxin)
with an asymmetric distribution so that they often have large dipole moments
[6]. Typically, the metal containing redox active centre is iron or copper which
is located within the protein at some distance from the surface. Thus, electron
transfer to and from an electrode surface may have to occur over distances of
several Angstrom (A).
The most widely studied electron-transfer metalloprotein is undoubtedly
cytochrome c [7,8]. As noted by Bowden [9], the quintessential cytochrome c
protein, whose crystallographic structure was first reported three decades ago
[lo], has served as a stalwart model, playground, and testbench for chemists
interested in understanding fundamental aspects of biological electron transfer [6-81. Not surprisingly, it has also become a benchmark system in protein
electrochemistry [9].
Cytochrome c, an essentially globular protein with a diameter of approximately 30 A. Horse-heart cytochrome c has a molecular weight of 12.4 kDa
and contains 104 amino acid residues. It is a very basic protein with an overall
charge in the range of +7 to +8 at neutral pH [6]. The protein possesses a
Studies on protein-suface attachment
443
heterogeneous surface charge with a predominance of positively charged lysine
residues in the vicinity of the partially exposed haem-edge providing 0.6 per
cent of the total surface area of the protein. The exposed haem edge is believed
to be the site of electron transfer, and electrostatic interactions between the
positively charged lysine side chains and, complementary, negatively charged
groups on its physiological redox partners (or an electrode surface), induce the
electron-transfer process. Some other metalloproteins such as plastocyanins and
ferredoxins are highly negatively rather than positively charged in their native
forms at neutral pH [6], so that as will be shown later their voltammetry may
be different from cytochrome c in situations where the sign of the charges on
the protein and on chemically modified electrode surfaces are significant. However, with respect to other structural features, generally analogous characteristics
apply to all metalloproteins.
Proteins, in general, adsorb strongly to many surfaces, including materials used as
electrodes. Solution-phase voltammetric studies on these biologically important
molecules are therefore likely to be accompanied by surface interaction of material with the electrode surface. Clearly, the nature and significance of surface
interactions will need to be understood in order to elucidate the voltammetry
of metalloproteins.
In the case of azurin, Hill et al. [ I l l have provided direct evidence of the
adsorption of individual molecules of this metalloprotein onto a gold electrode
surface by in situ STM (see Section 19.1 in Chapter 2 for details of this scanning probe microscopy (SPM) method). Since gold has been widely used as
an electrode material in voltammetric studies of metalloproteins [3,4], significant information relevant to metalloprotein electrochemistry can be expected
to emerge in the future from consideration of data obtained from these in situ
scanning probe studies.
In order to rigidly attach azurin to an atomically flat gold substrate, which
constitutes an ideal surface for high resolution SPM, Hill et al. [I 11 replaced the
serine in position 118 by cysteine to give the azurin (Pseudomonas aeruginosa)
mutant S118C. This modification, which generated a well-defined gold-sulfur
interaction on the surface, represents a conservative mutation in that no significant protein structure change is likely to have occurred. Thus, electrochemical
activity should be, and indeed is, retained, as will be shown later.
The gold [I 111 surface,' onto which azurin S118C could be rigidly attached
via anchoring of the thiol group, was prepared by evaporation of gold onto
freshly cleaned mica. For STM imaging in air, freshly annealed gold substrate
was immersed in a 150 pM azurin solution for 3-12 h, gently rinsed, dried with
h he observation [I 11 of atomically flat, hexagonal facets confirmed the presence of a gold
[I 111 surface.
444
Metallopvotein voltammetry
argon, and immediately imaged. Solution studies were carried out in the same
way, with the exception that the gold substrate was removed from the protein
solution, gently rinsed, then immediately immersed under pure water or buffer,
and then imaged [I 11.
Imaging in air of gold substrate that had been placed in contact with azurin
S118C demonstrated that the surface was covered with a significant amount
of material. At high magnification, some of this material could be resolved
into regularly sized, quasi-spherical molecules [I 11. Since such features were
not observed if the substrate was immersed in protein-free deionized water and
imaged under the same conditions, these are in all probability, azurin molecules.
This conclusion is further supported by the fact that the smallest, clearly resolved
features were approximately 4 nm in diameter (the long axis of the azurin molecule is about 4 nm in length). Images of similarly treated substrate taken under
solution were even sharper than those obtained in air and under these conditions individual molecules, and/or small aggregates could be resolved [ l I]. The
detection of aggregates is, of course, noteworthy as it raises the question as to
whether these could give rise to different voltammetry relative to that observed
with individual protein molecules, as assumed to be present in 'thin film7studies
on adsorbed azurin (Section 18 in Chapter 2 and Section 4.8.1 in Chapter 5).
Figure 6.l(a-g) shows images taken under solution at specified times after
the gold substrate was in contact with a solution containing approximately
70 yM protein. The sequence of images shows a clear evolution over time,
with the amount of adsorbed azurin steadily increasing. Images of individual
isolated molecules are presented in Fig. 6.2.2 A comparison with a space-filling
computer graphical model of a single molecule is presented in Fig. 6.2(c). If
the models are orientated so as to direct the exposed surface cysteine residue
towards an underlying substrate, the correlation, in terms ofboth size and shape,
with the STM images is excellent.
The STM results on azurin demonstrate that the S118C mutant has a high
affinity for a gold electrode surface. Friis et al. [12] have published a combined
AFM/STM (Section 19.1 in Chapter 2) study of wild type or native azurin
on a gold surface and concluded that the protein chemisorbs on the substrate
surface via its disulfide group (Cys3-Cys26). However, since this interpretation may not represent the complete story [Ill, the full relevance of these
open-circuit potential ex situ images to voltammetric data has yet to emerge.
The difficulty in assigning the full details of the form of surface attachment is
that the adsorption and immobilization of protein at a solid/liquid interface
remains a poorly understood process [13,14] in which van der Waals, electrostatic and hydrophobic forces, amongst others, play a role [15-171. At the
present time, few analytical methods exist that enable adsorption process to
be studied at a molecular level [18]. However, SPM now offers the ability
to carry out such studies in an in situ mode and at a truly molecular level, so
that in future, the molecular level details required to fully understand what
'see http://www.rsc.org/suppdata/njc/l998/ 1119 for full colour venions of Fig. 6.2(a-c).
Fig. 6.1 Constant-current STM images of the in situ adsorption of azurin S118C on an Au[l 111
surface under pH 7 phosphate buffer. (a) The annealed gold substrate surface prior to the introduction
of protein, x-scale 0-2.5 nm; (b) 3 min after the injection of 100 p1 of 100 pM S118C solution;
(c) 5 min after protein injection; (d) 10 min after protein injection; (e) 15 min after protein injection;
(018 min after protein injection; (g) 32 min after protein injection. Reproduced by courtesy: New
J. Chem. (1998) 1119. Copyright, Royal Society of Chemistry.
446
Metalloprotein voltammetry
Fig. 6.2 STM images of immobilized azurin S118C on an A u [ l l l ] surface. (a) In situ constant
current image; (b) a three-dimensional, constant-current image; (c) a comparison of constant
height images of two azurin Sl l 8 C molecules and orientated space-filling models of the protein.
Reproduced by courtesy: New J. Chem. (1998) 1119. Copyright, Royal Society of Chemistry.
occurs at gold and other electrode surfaces as a function of potential should
be revealed. While the above report focuses on adsorption of azurin on gold,
molecular details on azurin, plastocyanin, ferredoxin, cytochrome c, and other
metalloproteins attached to gold and other surfaces such as carbon ultimately
should become available by scanning probe techniques. At present, all that can
be noted from observations made with other proteins, such as cytochrome c
[I2,191 is that the affinity for a gold substrate varies significantly with the nature
of the metalloprotein. Nevertheless, a fact to emerge from all the studies is that
metalloproteins in contact with a gold or other surface will almost invariably
become attached as both single molecules and as aggregates and that surface
interactions related to these observations almost certainly will influence the
voltammetry.
Influence ofru$ace attachment
447
f a metalloprotein in solution is strongly adsorbed or chemically attached
Q electrode surfaces, as frequently appears to be the case, then clearly the
voltammetric response should be modified from that expected for a purely
iffusion-controlled solution-phase response. Three likely scenarios arising
under these circumstances are:
If essentially all of the metalloprotein becomes surface attached and retains
the structure of the native form, then ideal 'thin film' voltammetry of the
kind exhibited by azurin at carbon electrodes (Section 18 in Chapter 2 and
Section 4.8.1 in Chapter 5) may be anticipated.
If the surface-attached metalloprotein exhibits a modified structure3 from
that of the native form present in the solution phase and if the modified form
(e.g. an aggregate) is electroinactive in the reversible potential region of the
native form, then electrode blockage could occur, so that no voltammetry
is observed. However, in this case, it is also possible that the voltammetric
response for the modified surface-attached metalloprotein may be detected
at a different potential to that of the native form of the metalloprotein.
If both solution phase and surface-confined metalloprotein are present
in their native form, then mixtures of thin film and diffusion-controlled
voltammetry could be expected, with the relative contributions being a
complex function of voltammetric technique (time domain) and protein
concentration (Section 18.4 in Chapter 2).
General features of voltammetry of metalloproteins at
bare (unmodijied) gold electrodes
At an edge-plane graphite electrode, the voltammetry ofboth native and S118C
forms of azurin exhibit essentially classical diffusion-controlled solution-phase
responses (Fig. 6.3(a) and (b)). That is, the peak height is proportional to
the square root of scan rate (Fig. 6.3(c)) and the reversible potential is about
90 mV versus SCE. Equation 6.1 describes the azurin voltammetry at graphite
electrodes.
Az [Cu (11)](solution)
+ e- + Az [Cu(I)](solution)
(6.1)
In contrast, at bare gold electrodes (Fig. 6.4(a) and (b)),but not at a modified gold
electrode (Fig.6.4(c))the processes are significantly more complex than implied
by eqn (6.1). This raises the question of why the use of carbon or chemically
modified gold surfaces simplify the voltammetry relative to that observed at a
'bare' gold metal electrode.
3~nteraction
with the electrode surface may, for example, lead to unfolding of the structure
and subsequently to denaturation of the rnetalloprotein.
448
Metallopotein voltammetvy
-200
0
200
mV vs SCE
0
0.05
0.1
0.15
0.2
0.25
Scan ratex (V/s)%
mV vs SCE
Fig. 6.3 Close to ideal reversible voltammetry obtained for azurin at a polished edge-plane graphite
electrode (a) cyclic voltammogram of 100pM wild-type azurin (pH 8, 0.1 M KC1, scan rate
10 mV s-l); (b) cyclic voltammetry of 60 yM S118C azurin (pH 7, 0.1 M KC1, scan rate 20 mV s-l);
(c) plot of peak height versus square root of scan rate for 76 yM wild-type azurin (pH 7, 0.1 M KCl).
Adapted from: Coord. Chem. Rev.200-202 (2000) 41 1. Copyright, Elsevier.
The SPM clearly reveals that azurin and the S118C mutant become attached
to a gold substrate and presumably this factor is related to the observation of their
complex voltammetry at gold electrodes. The voltammetry of native or wildtype azurin at a 'bare' gold electrode, as shown in Fig. 6.4(a), is very transitory
in nature and also occurs at more positive potentials than the reversible process
detected at other surfaces (Figs 6.3(a) and 6.4(c)). In the case of the S1l 8 C
mutant, a stable, but not diffusion-controlled (shape and scan rate dependence
consistent with surface confinement) cyclic voltammetric response shown in
Fig. 6.4(b) is observed at a potential that is again significantly more positive
than expected for a reversible diffusion-controlled process (compare Figs 6.3(b)
with 6.4(b)).The role of the surface cysteine present in the azurin S118C in this
behaviour may be demonstrated through the addition of B-alenoyl maleimide
to the cell. Removal of the free surface thiol, through its reaction with the
maleimide results in a dramatic change in electrochemical response to give a
cyclic voltammogram that resembles the transient behaviour observed with the
wild-type protein [I I].
Since the voltammetric response of the mutant protein is stable to continual scanning of the potential, the surface attachment occurs with retention of
electrochemical activity. In contrast, wild-type azurin, which presumably is
attached to the gold surface in a different manner, becomes electroinactive in
the Faradaic sense on cycling of the potential. However, note in Fig. 6.4(c),
449
Influence ofsuface attachment
(b)
I
I
0
I
I
I
400
-100
I
200
(mV) vs SCE
I
-100
I
I
0
100
(mV) vs SCE
I
I
100
300
(mV) vs SCE
I
500
I
200
Fig. 6.4 Voltammograms of azurin at a gold electrode (a) 76 pM wild-type azurin (pH 7, 0.1 M
KCl, scan rate 20 mV s-') showing the first (I), second (2), and fifth (5) cycles ofpotential; (b) 60 pM
S118C azurin (pH 7, 0.1 M KC1, scan rate 20 mV s-l) showing the initial 20 cycles of potential;
(c)wild-type azurin at an SS-bpy-modified gold electrode (pH 6, 50 mM phosphate buffer, scan rate
20 mV s-I). Adapted from: Coord. Chem. Rev. 200-202 (2000) 41 1. Copyright, Elsevier.
that at a SS-bpy (4,4'-bipyridyl disulfide or bis(Cpyridy1 disulfide) modified
gold electrode, a diffusion-controlled ideal process is observed as is the case
at an edge-plane graphite electrode (Fig. 6.3(a)). The voltammetry of ferredoxins, rubredoxins, plastocyanins, and myoglobins on 'bare' gold electrodes is
essentially the same as that for wild-type azurin in the sense that no stable welldefined solution-phase diffusion-controlled processes have been reported unless
chemically modified gold or other electrode surfaces are used. The electrode
dependence clearly needs to be explained.
While images of azurin on gold surfaces are impressive, inadequate data are
available to elucidate all the nuances of azurin voltammetry. In contrast, imaging
data on cytochrome c is not yet of the same quality as for azurin but a great deal
more is known about the voltammetry of cytochrome c. Importantly, many
features analogous to those described above for azurin, are associated with the
cytochrome c process given in eqn (6.2)
cyt c[Fe(III)](solution)
+ e + cyt c[Fe(II)](solution)
(6.2)
450
Metalloprotein voltammetry
Consequently, a detailed overview of the voltammetry of this metalloprotein is
presented, with the presumption being that analogous considerations apply to
the surface interactions and voltarnrnetry of azurin.
4.2 The transient nature of the voltammetry ofcytochrome c at
'bare' gold electrodes: an explanation based on a
se& blocking mechanism4
Recently, Hill et al. [20] have introduced the concept of a self-blocking mechanism to explain the difficulty associated with obtaining well-defined and stable
electrochemical responses for cytochrome c at bare gold electrodes.
4.2.1 Important experimental details
In order to obtain the results reported below, extensive cytochrome c purification, specific gold electrode treatment, and specific cytochrome c transferral
procedures to the electrochemical cell were used. Brief details of these procedures relevant to the voltammetry described at bare gold electrodes are as
follows:
(1) Tuna-heart cytochrome c and horse-heart cytochrome c, unless otherwise
specified, were purified by fast protein liquid chromatography.
(2) Before each set of experiments, the 4-mm diameter gold disc electrode
was electrochemically cleaned by cycling the potential between -0.45 and
+1.45 V versus SCE in 1 M sulfuric acid solution for 10 min at a scan rate
of 0.1 v s - l .
(3) Before polishing, the electrode was initially dipped into concentrated nitric
acid for 30 s and then in ethanol for 3 min.
(4) A mirror-like surface was obtained by polishing on a Buehler ~ i c r o c l o t h ~ ~
with a water-alumina (0.03pm) slurry, followed by 3 min of sonication.
(5) A drop of water was left on the surface of an inverted electrode during
transfer but was shaken off before the electrode was placed in the cell.
(6) Experiments were undertaken at 21°C.
4.2.2 Studies with potassium chloride and sodiumjuoride electrolytes
When voltammetric experiments are undertaken using conventional conditions (gold electrode placed in solution and experiments are commenced after
the electrode has been in contact with the solution for a period in excess of a
minute), no significant Faradaic current is observed at the reversible potential
of about OV versus SCE for the process given in eqn (6.2) when potassium
chloride is the electrolyte [20]. That is, the response is essentially indistinguishable from that obtained with the electrolyte alone (compare Fig. 6.5(a)
and (b)), except for a slight suppression of the background current which may
be attributed to adsorption of native cytochrome c or other material. This is,
4 ~ d a p t e with
d
permission from J. Electroanal. Chem. 436 (1997) 17. Copyright, Elsevier.
InJuence ofsurface attachment
I
-60
I
I
I
I
I
I
I
0
60
mV vs SCE
I
I
1
120
-60
I
1
1
1
0
1
1
1
1
60
mV vs SCE
1
1
1
120
I
-60
I
I
I
I
I
I
I
0
60
mV vs SCE
45 1
I
I
I
120
Fig. 6.5 Cyclic voltammograms obtained at a scan rate of 20 mV s-I in 0.08 M potassium chloride
solution (a) electrolyte alone, (b)-(h) electrolyte containing 110 pM horse-heart cytochrome c and a
time delay after the gold electrode came into contact with the solution before scanning commenced
of (b) 120 s, (c) 2 s, (d) 10 s, (e) 20 s, (f) 30 s, (g) 40 s, (h) 60 s. The numerals 1, 2, 3, etc. indicate
the number of cycles. Reproduced by courtesy: J. Elect~oanal.Chem. 436 (1997) 17. Copyright,
Elsevier.
data are obtained which are equivalent to the (lack of) voltammetric responses
reported in studies prior to 1977. However, if the same 'bare' gold electrode
is treated and repolished according to the procedure described above and the
potential scan is commenced within 2 s of the electrode coming into contact
with the cytochrome c solution, then the voltammogram seen in Fig. 6.5(c) is
obtained. Despite the fact that the response is sigmoidal rather than the predicted
peak shape, and is only observed for about 3 cycles of the potential, the halfwave potential does occur at potentials that are close to the thermodynamically
expected reversible value of about 0 V versus SCE.
Clearly, the almost sigmoidal shape of the voltammetric response differs considerably from the peak-shaped response predicted for a diffusion-controlled
process at a fully electroactive 4-mm diameter gold macrodisc electrode. Electrode blockage is implicated. Figure 6.6 shows how radial diffusion and hence
a sigmoidal-shaped voltammogram may occur with high blockages when each
bare electrode site remaining is small and no overlap of diffusion layers occurs.
However, if only a small extent of blockage occurs or long time domain experiments are undertaken, then overlap of diffusion layers may occur to give a
peak-shaped curve associated with linear diffusion.
The maximum current observed for the transient response is actually only
about 35 per cent of the value calculated on the basis of the reported diffusion
I
Electrode
Redox
active
site
Redox
inactive
part of
electrode
Fig. 6.6 Illustrations of the overlap of diffusion layers which takes place (a) as time increases or
(b) the spacing between electroactive sites decreases in a situation where electrode blockage leads to
formation of an array of electroactive microelectrodes. 6 is the diffusion-layer thickness.
coefficients for cytochrome c and assuming that the entire geometric area is
electroactive and that linear diffusion is the dominant mode of mass transport
(Table 6.1). Thus, substantial electrode blockage appears to have occurred. A
time dependence study (Fig. 6.5(d-h)) reveals no detectable response if the gold
electrode is left in contact with the solution for about 30 s prior to commencing
the scan. Results obtained with horse-heart and tuna-heart cytochrome c, used
as supplied by the manufacturer, differ to those obtained for purified samples,
in that the duration for which the transient response is observed is longer in
the latter case [20]. Additionally, if the extensive cleaning, maintenance, and
electrode transferral procedures detailed in Section 4.2.1 are not used, then the
duration of the measurable response also decreases.
Adsorption of solution contaminants could represent one form of electrode
blockage [21]. Deaminated and oligomeric forms of cytochrome c adsorbed
on the electrode surface also may lead to electrode blockage [22] as has been
elegantly demonstrated in voltarnmetric studies on cytochrome c at indium
oxide electrodes. However, it may be noted that at least one adsorbed form
of cytochrome c, which will have a different structural form to the soluble
native protein [23-261, and which may be related to the reduced high p H form
[27], can undergo irreversible electron transfer at gold electrodes at a potential of about -400mV versus SCE which is quite negative with respect to
Influence ofsuface attachment
Table 6.1 Half-wave potentials, peak-to-peak separations, and current ratios obtained from voltammetric experiments on a gold electrode using a scan rate of 20 mV s-' in different electrolytes containing
0.1 mM cytochrome ca
Electrolyte
0.08 M KC1 (SS-bpy-modified electrode)
0.08 M K2HP04-KH2P04'
0.10 M tris-cacodylatee
0.08 M KCle
0.10 M NaFe
30
35
35
40
35
60
65
70
70
75
1.OO
0.83
0.64
0.34
0.24
'Data taken from reference [20].
b ~ l (+5
p mV) calculated from the average of the oxidation and reduction peak potentials when peak-shaped curves are observed, or from the potential at half the limiting
current value when sigmoidal-shaped curves are observed.
'versus SCE.
d ~ u r r e nobserved
t
relative to that expected for an entirely electroactive electrode surface area and linear diffusion.
eFirst full cycle.
the reversible potential of the native protein [27,28]. Figure 6.7 reveals that if
the potential is scanned to this negative potential region so that reduction of
adsorbed, presumably non-native form of cytochrome c occurs, then apparently
some displacement of material is achieved thereby partially restoring the voltammetry of native cytochrome c at about 0 V versus SCE. A large positive charge
(about +8 for the native horse-heart ferricytochrome c) would be expected for
both native and adsorbed forms of cytochrome c, although the distribution of
charge may differ. Blocking of the electrode surface may occur in a more uniform manner with highly charged species than with non-charged species [29],
since coulombic forces will influence the orientation and spacing. These electrostatic constraints between adsorbed and native forms may also limit the number
of surface-active sites, which can become occupied by cytochrome molecules,
and determine their spacing, although eventually sufficient blocking may be
achieved so that current flow at the formal potential of cytochrome c, cannot
be detected. The lowering of the background current on the gold electrode in
the presence of cytochrome c (Fig. 6.5(a) and (b)) shows that the capacitance
and the surface area is reduced [30-321 probably due to adsorption of material
on the electrode surface.
A self-blocking model, also explains how the sigmoidal-shaped response associated with a partially blocked electrode, may be transformed to an ideal linear
diffusion model peak-shaped response (Fig. 6.4(c)) by addition of surface-active
SS-bpy. SS-bpy which was used to minimize surface interactions of ruthenium
photovoltaic dye-sensitizers in Section 2.5 in Chapter 4 is covalently bound to
gold surfaces but, in contrast to adsorbed cytochrome c material, does not block
electron transfer [32].
454
Metallopvotein voltammetvy
I
I
I
I
I
I
I
-400
-300
-200
-100
0
100
200
mV vs SCE
Fig. 6.7 Cyclic voltammograms recorded at a gold electrode using a scan rate of 20 mV s-' in
(a) 0.08 M potassium phosphate buffer solution, (b) with 120 pM horse-heart cytochrome c added
and cycling over the potential range of -100-220 mV versus SCE, with scanning of the potential
commencing immediately after the electrode comes into contact with the solution, (c) as for (b),
but potential range extended to encompass the range -500 to 250mV versus SCE, (d) response
obtained over the extended potential range after five cycles. Reproduced by courtesy: J. Electroanal.
Chem. 436 ( 1997) 17. Copyright, Elsevier.
Figure 6.8 shows the shape-time dependence of the voltammetry on addition
of SS-bpy to a blocked gold electrode, which initially exhibits no detectable
Faradaic response. O n addition of SS-bpy, a sigmoidal response becomes apparent after a few minutes and, in time, this becomes peak-shaped in character. O n
repolishing the electrode, an even more pronounced peak-shaped response is
observed immediately (Fig. 6.8). Apparently, polishing instantaneously removes
the strongly adsorbed molecules and provides electroactive sites to which the
SS-bpy may bind and hence protect the gold surface from further adsorption of
cytochrome c.
If electrolyte anion adsorption is an important parameter in the observed
voltammetric response, then replacing the chloride containing electrolyte with
Influence of sugace attachment
-100
-50
0
50
100
455
150
mV vs SCE
ig. 6.8 Cyclic voltammograms for 110 pM horse-heart cytochrome c at a gold electrode obtained
in 0.08 M potassium chloride electrolyte at a scan rate of 20 mV s-' after addition of SS-bpy. The
transient response obtained without SS-bpy was first allowed to decay (not shown), then 20 cycles
(solid line) of the potential were recorded after addition of 10 mM SS-bpy. Finally, the electrode was
repolished to give the response shown with the dashed line. Reproduced by courtesy: J. Electroanal.
Chem. 436 (1997) 17. Copyright, Elsevier.
one containing fluoride would be expected to lead to a decrease in the stability
of the transient cytochrome c electrochemical response since the fluoride ion is
known to be even more weakly adsorbed than chloride at a gold surface [33].
The cyclic voltammetry of cytochrome c in a sodium fluoride electrolyte is
shown in Fig. 6.9. The response is indeed more transient than when chloride
electrolyte is employed and, at a scan rate of 20 mV s-' , not identifiable within
two cycles. The peak current obtained from the first scan in the fluoride
electrolyte is only about a quarter that expected if the electrode were fully
electroactive and mass transport were to occur solely by linear diffusion (see
Table 6.1).
4.2.3
Studies with phosphate and tris-cacodylate buffered electrolytes
The self-blocking array-type model suggests that a more stable response would
be achieved by addition of an electrolyte containing a highly charged anion
which may be adsorbed onto the electrode or present in the double-layer region
of the electrode-solution interface. The voltammetry of cytochrome c when
456
Metallopvotein voltammetry
I
I
-50
0
I
I
50
100
mV vs SCE
Fig. 6.9 Cyclic voltammograms (first 5 cycles) of 95 pM horse-heart cytochrome c in 0.1 M sodium
fluoride, recorded with a scan rate of 20 mV s-I at a gold electrode. Scanning commenced immediately after the electrode came into contact with the solution. Reproduced by courtesy: J. Electroanal.
Chem. 436 (1997) 17. Copyright, Elsevier.
the potassium chloride electrolyte is replaced with a phosphate buffer5 is shown
in Fig. 6.10. Instead of the first scan being small and sigmoidal shaped, it now
possesses a peak current about 80 per cent that expected for a fully electroactive
electrode under reversible conditions with linear diffusion (Table 6.1). While
subsequent scans still eventually become sigmoidal shaped, the response for an
approximately 100 pM cytochrome c solution is now identifiable above the
background for many cycles (also see references [33-361). Employing a triscacodylate electrolyte (the anion present in solution being [Me2As02]-), as in
the work of Bowden et al. [24], produces (Fig. 6.1 1) a similar kind of response
for cytochrome c as with the use of a phosphate buffer. The cacodylate anion
would be expected to show enhanced adsorption properties relative to phosphate at a gold surface, but is only singly charged. Clearly the nature of the
electrolyte is significant in voltammetric studies of cytochrome c as is adsorption
and the presence of electrode modifiers such as SS-bpy.
4.2.4 Conclusions devivedfrom voltammetn'c studies at 'bare' gold electrodes
The transient nature of the diffusion-controlled component of the electrochemistry of cytochrome c observed at a 'bare' macrodisc gold electrode may be
explained in terms of a microscopic model of electron transfer occurring at
sites remaining unblocked by time-dependent adsorption of positively-charged
non-native forms of cytochrome molecules.
he phosphate ion is present as a mixture of [ H P O ~ ] ~and
- [H2P04]- at pH 7 and
therefore possesses a slightly greater charge per unit concentration than the chloride or
fluoride ions.
Influence of suIface attachment
457
inset
I
I
I
-150
-100
50
0
I
I
I
50
100
150
mV vs SCE
ig. 6.10 Cyclic voltammograms (first 20 cycles) of 100pM horse-heart cytochrome c solution
recorded with a scan rate of 20 mV s-I at a gold electrode in a 0.08 M potassiunl phosphate buffer.
Scanning commenced immediately after the electrode came into contact with the solution. Inset:
first 4 scans over a limited potential range detailing the peak-to-sigmoidal shape progression of the
response and concomitant shift in the peak potential. Reproduced by courtesy: J. Electroanal. C h e m .
436 ( 1997) 17. Copyright, Elsevier.
The proposed self-blocking model also explains why not even a transient
response due to cytochrome c could be detected at the reversible potential at
a gold disc microelectrode (5 pm radius) [37] since almost complete blocking
of the small electrode surface by adsorbed protein molecules would be very
rapidly achieved and the current at the residual part of the bare gold surface of
the microelectrode would be too small to measure.
The work of Shibata et al. [38] reports the enhancement of the electrochemical response of cytochrome c at a gold electrode by the addition of ad-atoms to
the electrode surface which suppress the gold-protein interactions and simultaneously decrease the electrostatic repulsion between native protein in solution
and a partially covered electrode surface. This current enhancement is accompanied by a change from sigmoidal to peak-shaped curves. The concept of
ad-atoms constituting unblocked or arrays of electroactive sites is consistent
with the blocked electrode model. Similarly, the microscopic model may also
be applied to explain the observation of transient electrochemical responses
reported with cytochrome c at other unmodified metal [21,24,39] and some
metal oxide [40] electrodes. Effectively it has been proposed [20] that fast electron transfer takes place at suitable sites and no electron transfer at blocked sites.
458
Metalloprotein voltammetry
I
I
I
I
I
-50
0
50
100
150
mV vs SCE
Fig. 6.11 Cyclic voltammograms (first 5 cycles) of a 107 pM horse-heart cytochrome c solution,
recorded with a scan rate of 20 mV s-' at a gold electrode in a 0.1 M tris-cacodylate buffer. Scanning
commenced immediately after the electrode came into contact with the solution. Reproduced by
courtesy: J. Electroanal. Chem. 436 (1997) 17. Copyright, Elsevier.
Of course, localized variation in electron-transfer kinetics has of course been
considered in other systems where it has been pointed out that measurement of
total current only provides an average picture of the surface heterogeneity [41].
Finally, it may be noted that voltammetric responses at variable phosphate buffer
concentrations [35,36] also reveal the importance of electrolyte concentration
and oligomers, with again the shape of many of the curves being explicable
in terms of a partially blocked electrode surface. In the future, it is probable
that these surface-based reactions will be directly observable using SPM, so that
more detailed explanations of the surface-based processes will become available.
4.2.5 Electrode blockage at indium oxide electrodes
Taniguchi [42] has recently stressed the role of the electrode and protein interactions that may give rise to a partially blocked electrode. In the case of
cytochrome c, indium oxide is another electrode surface at which ideal essentially diffusion-controlled voltammetric responses can be observed provided
purified samples are used [22]. In contrast, on In203, as with a bare gold
electrode, commercially available samples of cytochrome c only yield unstable,
ill-defined voltammograms [22]. Voltammograms obtained with commercially
available and purified samples are shown in Fig. 6.12. Deliberate addition of
deamidated and oligomeric components, which are contained in commercially
available samples in small amounts, leads to a change in wave shape (Fig. 6.13),
and ultimately, if sufficient of these compounds are added, the voltammetric
electrode is eliminated [22]. In contrast, at an SS-bpy
response on an h203
modified electrode, no change in the voltammetry occurs on addition of these
structurally different forms of cytochrome c [22]. These non-native components adsorb so strongly onto the electrode surface that the electrode reaction
of native cytochrome c is blocked. In contrast, at the chemically modified
Metalloproteins at gold electrodes
459
(4
0.2 yA
E (V) vs
Fig. 6.12 Cyclic voltammograms (-)
obtained at a scan rate of 20 mV s-' of (a) commercially
available cytochrome c, (b) approximately 0.1 rnM purified cytochrome c, in a pH 7 phosphate buffer
solution containing 0.1 M NaCl at 25°C. (- - - - -) represents the background response. Adapted
from: Denki Kagaku 60 (1992) 1043.
Fig. 6.13 Changes in shape of the cyclic voltammograms obtained at a scan rate of 20 mV s-' for
(-)
approximately 0.1 mM purified cytochrome c after adding (- - - - -) (a) approximately 20 pM
deamidated and (b) approximately 10 pM oligomeric cytochrome c in a phosphate buffer solution
containing 0.1 M NaCl (pH 7) at 25OC. Adapted from: Denki Kagaku 60 (1992) 1043.
electrode, SS-bpy molecules adsorb more strongly onto the gold electrode surface than any component of cytochrome c and, thus, no strong adsorption of
deamidated and oligomeric cytochrome c occurs. This observation further supports the concept that exclusion of strong adsorption of non-native forms of
cytochrome c onto the electrode is one of the important requirements to obtain
a completely electroactive surface. Similar behaviour indicating the importance
of purification of sample was also observed for myoglobin electrochemistry at
electrode [42].
an h203
As noted above, the difficulty in undertaking voltammetric studies of metalloproteins on a 'bare' gold electrode has been attributed to adsorption of the
protein itselc a denatured form of the protein or impurity present in the electrolyte solution or in the protein, which produces either a completely blocked
and electroinactive gold electrode surface, or a partially blocked electrode with
6~ections5 and 6 have been adapted with permission from Anal. Proc. 29 (1992) 132,
30 (1993) 218 and Coord. Chem. Rev. 200-202 (2000) 411. Copyright, Royal Society of
Chemistry and Elsevier.
460
Metallopvotein voltammetry
an array of electroactive sites. Commonly, in order to overcome problems with
electrode blockage, a chemically modified electrode surface is used to achieve
stable well-defined voltammetry at or near to the reversible potential. SS-bpy has
already been shown above to enable well-defined voltammetry to be observed,
even when voltammetric studies are undertaken on unpurified cytochrome c
solutions.
Central to the action of these surface-modifying molecules appears to be
the ability to prevent direct adsorption of the metalloprotein on to the bare
electrode suriace 1431 as well as the maintenance of the protein structure in
a configuration which allows rapid electron transfer between the chemically
modified surface and the protein redox site. In achieving these two functions,
the chemically modified surface appears to 'promote' or facilitate the observation of the direct electrochemistry of metalloproteins. A basic mechanism for
the interaction between the 'promoter' (facilitator) and cytochrome c [44-471
is schematically reproduced in Fig. 6.14.~It encompasses diffusion of metalloprotein to the surface, adsorption to the facilitating adlayer in a favourable
orientation, electron transfer, desorption, and then diffusion of the protein
away from the electrode surface.
Figure 6.15 shows STM images of a single-crystal gold SS-bpy modified
electrode surfaces which clearly change on addition of cytochrome. O n some
occasions in these experiments, which are described in more detail in reference [47], it was possible to resolve small structures, which could be assigned
to individual cytochrome molecules. Though these observations are consistent with the electrode-facilitator-protein model shown in Fig. 6.14, it is not
clear whether the adsorbed metalloprotein sits on top of or slowly displaces the
facilitating adlayer of SS-bpy.
As anticipated, polycrystalline gold electrodes, after polishing with an alumina
slurry (as typically used in electrochemical studies), are seen to be far from flat on
a micrometre scale. In addition to the presence of 'rounded hill' surface features,
100-200 nm in diameter (r.m.s. roughness 15-30 nm), defects and polishing
scratches are also evident (Fig. 6.16). In view of the relative complexity and
roughness of these surfaces, the adsorption of protein at polycrystalline gold
electrodes cannot be studied directly. However, one would expect [47] that
protein adsorption/aggregation on these rough surfaces would be even greater
than that observed on atomically flat single crystalline electrodes of the same
nominal geometric area.
Figure 6.17 illustrates the cyclic voltammogram obtained at a gold electrode in
the presence of the modifier 4,4'-bip~ridine.This form of surface modification
was used in one of the first publications where well-defined voltammetry was
reported [2] for cytochrome c. The contrast with voltammetric data contained at
' ~ tshould be noted that figures used to depict these adlayer-protein interactions are
obviously idealized models. The electrode surface, in reality, is a complex mixture of
topographies. Protein-surface interactions will similarly be highly variant. The experimentally determined electrode potentials and heterogeneous electron-transfer rate constants
will therefore reflect a distribution of all those Val& actually present [47].
Metallopvoteins at gold electrodes
461
0
Horse-heart cytochrome c
Adsorption
_______)
Electron transfer
'QL
cytochrome c
Fig. 6.14 (a) Schematic representations of the interaction between a 4-pyridyl thiolate modified
gold electrode surface and a cytochrome c nlolecule. Shown on the right (b) is an STM image
of a facilitating 2-mercaptopyrimidine adlayer on gold (with the protein structure superimposed).
(c) A schematic depiction of the processes associated with redox protein electron transfer at an
electrode/electrolyte interface. Reproduced by courtesy: Coord. Chem. Rev. 200-202 (2000) 41 1.
Copyright, Elsevier.
'bare' gold electrodes in Section 4 is most pronounced. Subsequently, numerous
organic compounds have been examined [48] specifically in terms of their ability
to achieve ideal electrochemistry of horse-heart cytochrome c at a modified gold
electrode. Consequently, the structural requirements for successful promotion
of the voltammetric response are now well known [49]. Figure 6.18 provides a
schematic illustration of the nature of the modified electrode surface as deduced
from the various studies [49]. Effectively, an array of adsorbed molecules at
the electrode interface provides the electroactive sites rather than the bare gold
electrode. The electrochemical implication of this arrangement is illustrated in
Fig. 6.19. Thus, linear diffusion of cytochrome c to and from what is effectively
a fully electroactive macrodisc electrode may only occur if a monolayer of
greater coverage of surface modifier is present or if the sizes of the newly created
electroactive sites are large so that the spacing between electroactive sites is small.
It therefore follows that the model of electron transfer of proteins at chemically
modified gold electrodes may be based on radial diffusion to microscopic active
462
Metalloprotein voltammetry
rnV vs SCE
Fig. 6. I§ (a) STM image of an SS-bpy adlayer on Au[l 11] in 0.05 M HC1O4 solution. (b) Constant
current electrochemical STM image of a SS-bpy modified Au[l11] electrode about 3 min after the
addition of approximately 100p1 of 2OOpM cytochrome c solution to the STM cell. Scale bar
4 nm. (c) The simultaneously recorded voltammetric response obtained at a scan rate of 21 mV s-l.
Reproduced by courtesy: Coord. Chern. Rev. 200-202 (2000) 41 1. Copyright, Elsevier.
sites in instances where surface coverage by the modifier is incomplete, and
that electron transfer does not occur to any significant extent at the unmodified
part of the gold electrode, which is presumably blocked by adsorbed protein
material.
Figure 6.20 contains examples of the time dependence of the voltammetry
of cytochrome c after a gold electrode modified ex situ with SS-bpy (an excellent promoter of cytochrome c voltammetry) is placed in a solution of another
modifier (relatively poor promoter). The transition from a peak to sigmoidalshaped voltammogram followed by a decrease in the limiting current is readily
understood in terms of gradual displacement of adsorbed SS-bpy from the electrode surface coupled with the implications that are inherent in the microscopic
model. An analogous time dependence is observed when an ex situ modified
Metallopvoteins at gold electvodes
463
Fig. 6.16 Ambient contact AFM image of a polished polycrystalline gold electrode surface. x-range
0-400 nm, scan size 3.6 x 3.6 p.m. Reproduced by courtesy: Coord. Chem. Rev.200-202 (2000) 41 1.
Copyright, Elsevier.
gold electrode is placed in solution and a change in wave shape and decrease in
peak height occur as the modifier dissolves from the sudace of the electrode.
In the case of azurin, chemical modification of a gold electrode with SS-bpy
also produces a stable almost ideal chemically and electrochemically reversible
voltammetric process (Fig. 6.4(c)).The reduction of spinach plastocyanin at gold
electrodes modified with 2-mercaptoethylamine, 2,2'-dithiobis(ethy1amine)
and other compounds has been described in considerable detail [50,51]. Application of the linear diffusion model, which is valid at high coverage (see above),
leads to the conclusion that fast electron transfer is associated with the voltammetry of plastocyanin at a suitably modified electrode. However, under some
conditions, reduction of plastocyanin exhibits characteristics of a contribution
from radial diffusion, as shown by a full analysis of the wave shape and scan-rate
dependence. For example, time-dependent [50] studies of plastocyanin voltammograms at a PATS-2-modified electrode (PATS-2 being 2-pyridylaldehyde
thiosemicarbazone), prepared in the ex situ mode, reveals a change from a
peak-shaped to a sigmoidal-shaped curve (Fig. 6.21(a)), which is attributable
to gradual loss of the modifier. In contrast, if PATS-2 is present in the bulk
solution, a persistent response is observed (Fig. 6.21 (b)), corresponding closely
to the linear diffusion mass-transport model at a highly covered surface [50].
Finally, it may be noted that the electrochemistry of the very negatively (- 19)
charged ferredoxin at PATS-modified gold electrodes has also been interpreted
as poor [50], since the Faradaic response is hardly, if at all, detectable above
464
Metallovrotein voltammetry
Fig. 6.17 Cyclic voltammograms obtained at a gold electrode for horse-heart cytochrome c
(5 r n g m ~ ~ in
' ) 0.1 M NaC104-0.02M phosphate buffer at pH 7.0, in the presence of 10 rnM
4,4'-bipyridyl over the potential range from +0.20 to -0.20 V versus SCE, at scan rates of (1) 20,
(2) 50, and (3) 100 mV s-' . Provided by courtesty: H.A.O. Hill, University of Oxford, England.
the background current (Fig. 6.22). However, after correcting for the background current, the response again can be regarded as reversible according to a
model in which electron transfer only occurs at microscopically-sized chemically
modified electroactive sites.
6 Voltammetry o f
erately functionalize
s at naturally an
on electrodes
In Section 5, well-defined electrochemistry at gold macrodisc electrodes has
been shown to be routinely achieved, provided, modified surfaces are prepared
with suitable materials. At carbon electrodes, metalloprotein voltammetry without deliberate addition of modifier is often possible, but a wide range of shapes
and scan-rate dependence are observed depending on the exact nature of the
electrode and its treatment, and the metalloprotein. Figure 6.23 shows, for
example that the details of the voltammetric shape depends on the identity of
the protein, whether edge or basal-plane pyrolytic graphite is used, the pH and
the nature of the modification.
Under some conditions, gven in detail in the captions to Figs 6.23 and 6.24 it
can be seen that sigmoidal rather than peak-shaped voltammograms are observed
and that the dependence ofthe peak current is not always a function of the square
root of scan rate, as required by the macroscopic model where mass transport is
assumed to occur by linear diffusion and the entire electrode surface is assumed
to be electroactive. Thus, a close relationship is noted with voltammetry at
modified gold electrodes.
Metalloproteins at carbon electvodes
465
Fig. 6.18 Schematic illustration of an array of adsorbed molecules on a chemically modified gold
electrode surfice. The proven or anticipated mode of adsorption is given for each of the surface modifiers: (a) 1,2-bis(4-pyridy1)ethylene; (b) SS-bpy (4,4-bipyridyl disulfide); (c) PATS-4
(4-pyridylaldehydethiosemicarbazone); (d) CATS-4 (thiodemicarbazone; 4-carboxybenzaldehyde);
(e) thiophenol; (0 2-mercaptoethylamine. Reproduced by courtesy: J. Electroanal. Chem. 217
(1987) 141. Copyright, Elsevier.
The need to again introduce the concept of a microscopic electrode array
model can be appreciated readily by noting the remarkably close relationship of
metalloprotein electrochemistry at carbon as well as gold electrode surfaces with
voltammograms obtained for the oxidation of ferrocene at arrays of platinum
microelectrodes having both spacing between electrodes and sizes of the electrodes themselves in the micrometre size range (Fig. 6.25). It may be deduced
from this comparison, that if protein electrochemistry occurs only at electroactive sites of micrometre size with non-electroactive sites constituting the
remainder of the surface, then a better theoretical description of the observed
voltammetry will be obtained than is the case when the entire electrode surface
is assumed to be electroactive.
The use of an electrode microarray-type model of course poses the question
of why only parts of the electrode surface are electroactive and other parts are
466
Metallopvotein voltammetry
Electrode sudace
Surface potential
\ E734
\
otential drop across the adlayer
rotein molecule
Facilitating fun'ctionalized adlayer
(b)
Electrode surface
\
Partial coverage of facilitating adlayer
Fig. 6.19 Metalloprotein electron transfer processor at a chemically modified electrode. Model (a)
illustrates the situation where the electrode surface area is both flat and uniformly electroactive over
its entirety. In the case if the surface is modified with a monolayer coverage of facilitating adlayer,
the drop in potential across this must be taken into account when interpreting subsequently obtained
protein voltammetry. Model (b) illustrates the situation where the electrode surface is functionalized,
but not uniformly. Electron transfer accordingly takes place preferentially at localized sites on the surface. A schematic illustration of an array of adsorbed facilitating molecules on a gold electrode surface
for each of the surface modifiers shown in Fig. 6.18: clockwise from top centre SS-bpy, PATS-4,
CATS-4, 2-mercaptopyrimidine, 2-mercaptoethylamine, 1,2-bis(4-pyridy1)ethylene.Reproduced
by courtesy: Coord. Chem. Rev.200-202 (2000) 41 1. Copyright, Elsevier.
blocked at a carbon electrode. In Section 4.1 when describing the voltammetry
of metalloproteins at bare gold electrodes it was implied that the electrode was
being blocked by adsorbed or denatured proteins or impurities. However, details
of the nature of the postulated microscopically small active sites that represent
Metalloproteins at carbon electrodes
467
0.0 rnV vs SCE
Thiophenol
v
2-mercaptoethylarnine
Fig. 6.20 Cyclic voltammetry with a scan rate of 20 mV sf1 for 0.4 mM cytochrome c in a pH 7,
20 mM phosphate buffer containing 100 mM NaC104, using a gold electrode modified ex situ with
SS-bpy (a) and the effect of increasing exposure time to thiophenol (b,c) and 2-mercaptoethylamine
(d,e). Reproduced by courtesy: Eur. J. Biochem. 191 (1990) 737. Copyright, Blackwell Science.
the unblocked part of a carbon electrode surface now need to be considered.
Because of its highly ordered structure, pyrolytic graphite is anisotropic and
has two distinctive types of surfaces, depending on the plane along which it
is cleaved [52]. The basal plane formed by cutting along the aromatic ring
is relatively inert. In contrast, cutting the graphite across the aromatic rings
results in an edge face that is relatively reactive. In the presence of oxygen, this
leads to the formation of surface sites associated with various carbon-oxygen
(C-0) functional groups. The surface-oxidized functional groups, which are
ideally localized at the edge plane, impart considerable hydrophilicity and ionic
character to the surface and may form the sites where protein electrochemistry
takes place, either because protein adsorption is minimal at such sites or because
specific interactions occur at these sites which enable electron transfer to take
place. From the point of view of the microscopic theory of electron transfer, the
new feature being introduced is the assumption that electron transfer occurs only
at the functionalized (unblocked) sites and that these sites are microscopically
468
Metalloprotein voltammetry
I
I
-100 rnV
+300 mV
E (vs SCE)
Fig. 6.21 Voltammetry at a scan rate of 20 mV S-I for 0.25 mM spinach plastocyanin in a pH 8.0
buffer. (a) Gold electrode modified ex situ with PATS-2; (b) as for (a) but in presence of 0.1 mM
PATS-2 in the bulk solution. Reproduced by courtesy: Eur. J. Biochern. 191 (1990) 737. Copyright,
Blackwell Science.
small. The functionality may be naturally imparted by spontaneous formation
of functional groups on carbon or by deliberate modification of gold surfaces.
It may be argued that with the localized active site model, reduction of a protein such as cytochrome c should not occur at the basal-plane electrode because
there should be (ideally) no electroactive sites at which the electron transfer can
take place. That is, because this electrode should be purely carbon in nature,
adsorption of protein should occur on all points on this interface, and no reduction of native protein should take place at this fully blocked surface. In contrast,
the observed electrochemistry at the edge plane would depend on the number
of redox active sites per unit area that were not blocked. In reality, a basal-plane
graphite electrode will have a more dilute number of active sites than the edgeplane electrode because of non-idealities introduced either in its manufacture
or during the cleavage process. Thus, the inherent difference between a basal
plane and an edge-plane electrode in a real experimental situation should be the
surface density of electroactive sites.
Data obtained from X-ray photoelectron spectroscopy and other studies
have established the presence of C - 0 functional groups on the graphite
Metalloproteins at carbon electrodes
Bare Au
0.4
469
p~l
-350 mV vs SCE
Fig. 6.22 The voltammetry of Clostridiumpasteurianum ferredoxin at (a) a bare gold electrode and
(b,c,d) the same electrode modified ex situ with different isomers of PATS. Reproduced by courtesy:
Eur. J. Biochem. 191 (1990) 737. Copyright, Blackwell Science.
surface: carbonyl, phenolic, carboxylic, and ether-like groups are all thought
to be present. At the basal-plane electrode, the oxygen-to-carbon ratio at the
surface is very low (O/C = 0.02) [53]. O n polishing, this increases to 0.1 1
[53] and, with an edge plane, the ratio may be as high as 0.33 [53]. The surface
density of active sites and the shape of the cyclic voltammograms is therefore
predicted to be a function of the method of fabrication and history of the
electrode according to the microscopic model.
Highly oriented pyrolytic graphite (HOPG) is a substrate commonly used in
SPM because it is easily cleavable to produce a surface which is atomically flat
470
Metalloprotein voltammetry
PGE, pH 5.0
PGE
+
PGE
PGE
PGE, pH 4.0
10.25 1 1 ~
Fig. 6.23 Modifications of the edge-plane pyrolytic graphite electrode (PGE)-solution interface
and their effect on the voltammetry oE (a) 150 pM cytochrome c in 5 mM Tricine-100 mM NaCl
buffer (pH 8.0) at Ph3Si-modified PGE electrode; scan rate 20 mV s-'; temperature 20°C; (b) as in
(a) at bare PGE electrode; (c) 50 pM cytochrome c in 5 mM Tricine-100 mM NaCl buffer (pH 8.0)
at a [ ~ r ( N ~ ~ ) ~ ] ~ + - r n o PGE
d i f i eelectrode;
d
scan rate 20 mV s-'; temperature 25°C); (d) as in (c) at
bare PGE electrode; (e) 25 pM plastocyanin in 5 rnM acetate-1 mM KC1 buffer (pH 5.0) at bare PGE
electrode; scan rate 20 mV s-I; temperature 3°C); (0as in (e) at pH 4.0; (g) 30 yM plastocyanin in
5 mM HEPES-100 mM KC1 buffer (pH 7.0) at bare PGE electrode; scan rate 20 mV s-' ; temperature
20°C; (h) as in (g) at [ ~ r ( N ~ ~ ) ~ ] ~ + - r n o PGE
d i f i eelectrode.
d
The marked region (f) corresponds
to a potential of 0 V versus SCE and zero current. Reproduced by courtesy: J.Am. Chem. Soc. 111
(1989) 9185. Copyright, American Chemical Society.
over several square micrometres (Fig. 6.26(a)). HOPG has been used successfully in the AFM/STM imaging of proteins and enzymes and is equivalent to the
basal-plane graphite surfaces commonly used in electrochemical investigations
[3,4]. High-resolution STM images of horse-heart cytochrome c, deposited on
a HOPG surface, are presented in Fig. 6.27. Individual molecules (with dimensions very close to those determined by X-ray diffraction) of the cytochrome
are clearly evident. Interestingly, and of likely electrochemical significance,
the protein molecules preferentially aggregate at defect sites and step edges in
the surface (Fig. 6.27(a)) which are likely to possess hydrophilic, 'anchoring'
oxygen-based functionalities. An equivalent situation is thereby observed, on
at least part of the electrode surface, to that at a deliberately modified gold
electrode surface.
Metallovvoteins at carbon electvodes
d
I
0
471
0.2
0.4
Scan rate 'V (V s-')'
The radial limit
I
Fig. 6.24 Cyclic voltammograms (scan rate = 20 mV s-I) for reduction of 150 pM cytochrome c
(pH = 8.0) at (a) a polished edge-plane pyrolytic graphite electrode, and (b) a freshly cleaved
basal-plane pyrolytic graphite electrode; scan rate 20 mV s-I; temperature 2 0 ° C Also shown are a
plot of peak current for the reduction process at the edge-plane electrode versus the square root of
the scan rate, d l 2 , verieing the predominance oflinear diffusion for (a) and a plot of log[(^,,, - I)/I]
versus the potential, E, at the basal-plane electrode, verieing the predominance of radial diffusion for
(b). Reproduced by courtesy: J. Am. Chem. Soc. 111 (1989) 9185. Copyright, American Chemical
Society.
Edge-plane graphite electrodes used in voltammetric studies are, in stark
contrast to HOPG, very rough (Fig. 6.26(b)).The act of polishing adds further
to this in the creation of surface scratches of variable diameter and depth. Further
to this, edge-plane graphite surfaces, polished in air, possess a high density
of various carbon oxide functionalities and approach the situation depicted in
Fig. 6.14 at a modified gold electrode.
6.1 Cytochrorne c
The observation of an almost sigmoidal-shaped curve for reduction of
cytochrome c at a basal-plane electrode (Fig. 6.24(b)) with an ElI2 value equal
to the standard redox potential E ~ O , and a linear plot of E versus
- I)/I]
(E denotes potential, I, the current, and him, the limiting current) with a slope
0.002 V) over the temperature range 15--25"C, is consistent
of 2.303 R T I F (f
with a reversible process at an array of microscopically small and well-spaced
472
Metalloprotein voltammetry
Macroelectrode
r=0.5
~ ~ 0 . 1
pm UMA pm UMA
Macroelectrode
rz0.5
r=O.l
pm U M A
pm UMA
l
Potential (V) vs SCE
I
Potential (V) vs SCE
Potential (V) vs SCE
log (v)/log(mV s-I)
Fig. 6.25 Cyclic voltammograms for oxidation of low3M ferrocene at arrays of Pt microelectrodes and a conventionally sized Pt disc electrode of the same geometric area in acetonitrile
(0.1 M Bu4NC104).(a) Scan rate 5 mV s-' for typical array elements of radius 0.1 (right) and 0.5 pm
(centre) with a total area of 0.5 cm2 and for a macro-sized Pt electrode (left) of area 0.5 cm2. (b) Influence of scan rate at the 0.5 pm array electrode. (c) as for (a) but with a scan rate, V,of 20 V s-I where
Ip is the peak oxidation current. (d) Plot of log Ip versus log v for data obtained in parts (a), (b), and
(c). Macroelectrode; A radius of array element = 0.1 pm; 0radius of array element = 0.5 pm.
UMA represents ultramicroelectrode array. Reproduced by courtesy: Anal. Chem. 59 (1987) 2625.
Copyright, American Chemical Society.
electroactive sites. In contrast, the peak-shaped response at the edge-plane
electrode (Fig. 6.24(a)), with an average of the reduction and oxidation peak
potentials giving an E ~ O value equal to the value obtained for Eli2 at a basalplane electrode (within the experimental error of A5 mV) also is consistent with
a reversible process, but under conditions where diffusion layers overlap to give
the equivalent of linear diffusion. The polished basal-plane electrode represents an intermediate case of behaviour, but the electron-transfer process always
appears to be reversible on the time-scale of the voltammetric experiments.
T o support the applicability of the microscopic model at carbon electrodes,
the obvious experiments to consider are those which deliberately alter the surface density of electroactive sites at the edge-plane electrode. For cytochrome c,
it has been proposed that positively charged groups around the active haem edge
site interact with deprotonated C - 0 functional groups at the electrode surface
[53]. If this idea is correct, then blocking agents which can bind preferentially
to the oxygen groups on the surface decrease the surface density of suitable
sites at which electron transfer to the protein can take place. Comparison of
Fig. 6.23(a) and (b) shows the influence on the electrochemistry of cytochrome c
when trimethylsilane or triphenylsilane blocking groups are attached [54]to a
Metalloproteins at carbon electrodes
473
Fig. 6.26 (a) A constant-current STM image of the single crystal HOPG electrode surface. x-range
0-4 nm, scan size 2 x 2 pm. (b) Low-resolution contact AFM image of a polished edge-plane graphite
electrode surface. x-range 0-250 nm, scan size 6 x 6 pm. Reproduced by courtesy: Coord. Chern.
Rev.200-202 (2000) 41 1. Copyright, Elsevier.
carbon electrode. According to the microscopic model, the almost sigmoidal,
rather than peak-shaped, curves that are observed in the presence of silane blocking groups are not the result of slow electron transfer as previously proposed
[54]. Rather, a decrease in the number of active sites leads to a predominance
of radial diffusion over linear diffusion but the rate of electron transfer remains
extremely fast. Thus, a plot of E versus log[&, - I)/I] is linear, with a slope of
2.303 RT/F 0.002V) at 20" C, and Ellz = E ~ O with an experiment error of
f5 mV at the silane-derivatized electrodes. This latter result is closely related to
the reduction of cytochrome c at a basal-plane electrode, where the surface density of active sites, present as electrode imperfections, is also postulated to be low.
+
0
10.0
20.0
nM
Scale bar 3.5 nm
Fig. 6.27 Constant-current ambient STM images of molecules of horse-heart cytochrome c
deposited from aqueous solution onto a freshly cleaved highly ordered pyrolytic graphite electrode surface. (a) Illustrates the obvious preferential accumulation of the protein at step edges on
the, otherwise pristine, surface. x-range 0-5 nm, scan size 80 x 80 nm. (b) Shows three individual
cytochrome molecules. x-range 0-5 nm, scan size 25 x 25 nm. Scale bar 3.5 nm. Reproduced by
courtesy: Coord. Chem. Rev. 200-202 (2000) 41 1. Copyright, Elsevier.
Armstrong et al. also showed [55] that it is possible to attach positively
charged chromium(1II) complexes to the C - 0 functional groups at an edgeplane graphite electrode. Figure 6.23(c) and (d) shows the cyclic voltammograms
for reduction of cytochrome c in the presence and absence of a positively charged
[ c ~ ( N H ~ ) ~modified
] ~ + surface. As was the case after the introduction of a silane
blocking agent, the peak-shaped voltammogram observed in the absence of
modification is transformed into a fully reversible steady-state sigmoidal-shaped
response corresponding to reversible charge transfer at a low density of electroactive sites of microscopic dimension. Cytochrome c is a positively charged
protein and addition of groups of the same charge to the electrode surface has
the effect of decreasing the surface density of electroactive sites. The addition
or P~(NH~):+
also causes a change in shape from the reversible linear
of &lg2+
diffusion model to the reversible radial diffusion model.
Metalloproteins at carbon electrodes
475
6.2 Plastocyanin
The copper protein plastocyanin (PC),which like azurin exhibits a CU"-CU'
redox process, has an overall highly negative charge of about -9 at pH 7 in
contrast with cytochrome c, which has a very positive charge of approximately
+8 at the same pH value. Effects related to charge therefore should occur in an
opposite sense if the microscopic model is valid. The functional groups, where it
is possible that the electron transfer takes place at the pyrolytic graphite electrode,
are pH dependent. At pH values less than about 5, it has been suggested that
considerable protonation of electrode surface functional groups takes place [53].
Thus, if electrostatic terms are important in the negatively charged plastocyanin,
the electrochemistry should appear to be pH dependent.
Figure 6.23(e) and (f) shows the cyclic voltammetric response for reduction
ofplastocyanin as the pH is varied. The sigmoidal shape of the response at pH 5
(Fig. 6.23(e)) can be attributed to a low concentration ofwell-separated suitable
sites for electron transfer which are of microscopic dimension in size, whereas
the peak shape at pH 4 (Fig. 6.23(f)) may be associated with an increase in the
density of suitable surface-active sites, giving rise to a predominance of linear
diffusion rather than radial diffusion.
Figure 6.23(g) shows the response of the negatively charged plastocyanin in
the presence of positively charged chromium ammine complexes. The situation is now reversed from that observed for positively charged cytochrome c
(Fig. 6.23(c) and (d)), where the addition of [ c ~ ( N H ~ ) ~leads
] ~ + to the
change from a peak to a sigmoidal-shaped curve. The positively charged
domains formed from [ c ~ ( N H ~ ) ~treatment
]~'
of the electrode surface promote reversible binding with negatively charged plastocyanin, so that a high
density of active sites is achieved and a peak-shaped curve is found instead
of the sigmoidal response shape for plastocyanin at an unmodified electrode
(Fig. 6.23(g)).
The iron-containing ferredoxing proteins are extremely negatively charged
and the shapes of cyclic voltammograms can also be extensively modified by
addition of positively charged species such as [ c ~ ( N H ~ ) ~to
] ~the
+ surface
(Fig. 6.28), in an analogous way to that previously described for negatively
charged plastocyanin.
The electrochemistry of ferredoxin at a freshly cleaved basal-plane graphite
electrode in the presence of [ c ~ ( N H ~ ) ~ions
] ~ is
+ reversible, as is shown by the
R T I n F slope of the plot oflog[(Ili, -I)/I] versus the potential, E , (Fig. 6.29(a)),
although the dependence of the current on scan rate, v , is clearly sensitive to
the density of surface-active sites (compare Fig. 6.29(a) and (b) at polished edge
and basal-plane electrodes).
476
Metallopvotein voltammetry
1rnM
-350 mV vs NHE
I
NaCl
100 mV
n
Fig. 6.28 Cyclic voltammograms for reduction of 67 pM Clostridium yasteurianum ferredoxin in a
pH 6.0 buffer containing 1 rnM NaCl at an edge-plane pyrolytic graphite electrode in the presence
of increasing concentrations of [ c ~ ( N H ~ ) ~ ] Scan
~ ' . rate 20 rnV s-' ; temperature 25°C. Reproduced
by courtesy: J. Electroanal. Chem. 217 (1987) 331. Copyright, Elsevier.
6.4 General conclusions concerning the voltammetry of
metalloproteins at carbon electrodes
Numerous studies are available that describe the influence of added charged
species on the voltammetry of redox proteins at graphite and other forms of
carbon electrode. The general observation can be made, using the microscopic
model, that the observed response lies within the limit expected for 100 per cent
radial (microscopicmodel) or 100 per cent linear (macroscopicmodel) diffusion
with reversible electron transfer [56,57].
uantitative use of a microscopic model to
in the unusual features of metalloprotein
voltammetry at car
An array of microelectrodes represents a demanding electrode configuration to
model, but significant advances in the area have been made which enable many
aspects of metalloprotein solution-phase voltammetry to be quantified [58,59].
' ~ d a ~ t ewith
d permision from J. Electroanal. Chem. 314 (1991) 191. Copyright, Elsevier.
Also see Inorg. Chem. 31 (1992) 5007.
Quantitative use of micvoscopic model
477
Scan rate v112/(vs-')ll2
I
-400 mV
vs SCE
2 0.6
\
?
' j 0.5
4w
5 0.4
lz
0.3
d
.2 0.2
1
a
2 0.1
I
0.0
0.1 0.2 0.3 0.4
Scan rate v112/(vs-')lI2
Fig. 6.29 Cyclic voltammograms for 50 pM Clostridiurn pasteurianurn ferredoxin in a pH 6.4 20 mM
HEPES/100 rnM NaCl buffer (6 mM [ c ~ ( N H ~ ) ~ ]scan
~ + rate
;
20 mV s-l; temperature 20°C) at
(a) freshly cleaved and (b) polished basal-plane graphite electrodes. Plots of reduction current, Ired,
versus square root of scan rate, v1I2, (shown as e) and, in the case of (a), a plot of log[(hi, - I)/I]
versus the potential E (shown as o) is also included. Reproduced by courtesy: J. Am. Chern. Soc.
111 (1989) 9185. Copyright, American Chemical Society.
7.1 Cytochrome c voltammetry at carbon macrodisc electrodes
In cyclic voltammetry, and if the process were in fact quasi-reversible, as has
been assumed in most reports on the voltammetry of cytochrome c [3,4], the separation of the reduction and oxidation peak potentials A Ep and the wave shape at
suitably modified macrodisc electrodes would be expected to be independent of
the solution concentration of the electroactive species (Section 8 in Chapter 2).
However, for cytochrome c, a significant concentration dependence is observed
at graphite macrodisc electrodes which is not explained by the macroscopic linear diffusion model. The peak separation value, AEp = (E;~ - EpO"),which has
commonly been used to determine the standard heterogeneous charge-transfer
rate constant, ko, at E: when mass transport occurs solely by linear diffusion
(Section 8 in Chapter 2), actually increases as the bulk cytochrome c concentration increases. For example, with a 2.4-mm radius basal-plane graphite disc
electrode a low concentration of cytochrome c (cc,,,) ( t 2 0 pM), the peak separation (AE,) = 58 f 2 mV (see Figs 6.30 and 6.31), which is, in fact, the
value expected for a reversible electron transfer when mass transport occurs
478
Metallopvo tein voltammetvy
I
-0.3
0.0
I
I
0.3
-0.3
I
I
I
0.0
I
I
I
0.3
Potential (V) (vs Ag/AgCl)
Fig. 6.30 Cyclic voltammograms of cytochrome c at a polished basal-plane pyrolytic graphite electrode in a 100 mM NaC104 5 mM phosphate buffer (pH 7) as a function of scan rate. (a) c,,,., 860 pM;
(b) c,,,,, 21 pM. Reproduced by courtesy: J. Electroanal. Chem. 314 (1991) 191. Copyright, Elsevier.
log v (mV s-I)
Fig. 6.31 Reduction and oxidation peak separations (AEp values) obtained from cyclic voltammograms of cytochrome c at a polished basal-plane pyrolytic graphite electrode in a 100 mM
NaC104 5 rnM phosphate buffer (pH 7), as a function of scan rate. Symbols (experimental data):
A, c,,,,, 860 pM; B, c,,,.,, 380 pm; C, icy,,,, 122 pm; D, cCYt., 41 pm; E, cCyt,,, 17 pm. Solid lines
(simulated data): A, Q = 0.992; B, 6' = 0.975; C, Q = 0.85; D, 6' = 0.35; E, Q = 0.2. Other simulation parameters are in Table 6.2. Reproduced by courtesy: J. Electroanal. Chem. 314 (1991) 191.
Copyright, Elsevier.
by linear diffusion (Section 8 in Chapter 2). With increasing concentrations
(20 pM < cCytc < 900 pM), enhancement of the peak separation for a given scan
rate is observed and AE, values of up to 200 mV (see Fig. 6.31) can be obtained
under some conditions. Typically, a change from peak to sigmoidal-shaped
voltammograms is observed with scan rates v 3 0.5 V s-' and for concentrations
of cCy,., > 500 pM.
The concentration dependence of AE, and voltammetric shape cannot be
explained with a conventional model for a charge-transfer-controlled process
where mass transport to the electrode surface occurs via linear diffusion.
Such an interpretation would imply different charge-transfer rates (ko values,
see Section 7.2 in Chapter 2) for different concentrations (slower rates for
Quantitative use of micvoscopic model
479
higher concentrations), a proposal which is inconsistent with a first-order
heterogeneous charge-transfer process.9
An alternative interpretation of the experimental data, which does not require
the value of ko to be a function of concentration, can be obtained if the surface of the graphite electrode is assumed to have the property ofheterogeneous
electroactivity, which is a function of the bulk cytochrome c concentration.
That is, if only part of the surface is electroactive, non-linear mass-transport
patterns (e.g. microscopic radial diffusion) may be induced. Neglect of these
non-linear diffusion terms can strongly influence the apparent kinetics of the
electrode interface if the purely linear diffusion model is applied erroneously. In
summary, if a model is developed in which the cytochrome c bulk concentration
is linked to the fraction of blocked electrode surface, with the higher concentration leading to a higher fraction of surface being blocked, then it follows that
the larger peak separation values at higher concentration can be attributed to
a higher fraction of blocked electrode sudace rather than to a change in the
heterogeneous rate of electron transfer (ko value).
The model for partially blocked electrodes, proposed by Matsuda et al. [60,61]
can be adapted relatively easily to quantify the relation between surface coverage
and cytochrome c concentration. This model is based on a regular hexagonal
blocking pattern with circular active sites,'' and the geometry of the pattern is
defined by the radius of the active site r;, and the half-distance between active
sites ro (see Fig. 6.32). The fractional coverage 8 is given by the expression
- <)/$. ASshown in reference [60], a set of values for the parameters, Ti, Q,
and 8 , can be found which give excellent agreement1' between theory and
experiment (Fig. 6.33 and Table 6.2).
The results in Table 6.2 are consistent with the active sites having average
sizes in the micrometre and sub-micrometre range. It must be emphasized that
these calculated sizes are actually average values, for the time-scale of slow scan
rate cyclic voltammetric experiments, and are expected to be dependent on
the nature of the model. Realistically, random sizes, forms, and spacing rather
than a uniform surface are present. The fractional coverage, 8 , exhibits values
close to 1 for concentrations of c,,,., > 300 pM. This situation corresponds to
an almost completely blocked electrode surface. With very dilute solutions of
cytochrome c, 8 has a value close to zero, which corresponds to a situation
where most of the surface is electroactive.
Simulated cyclic voltammograms for a fractional coverage of 8 = 0.922 and
a rate constant ko = 1.0 cm s-' are shown in Fig. 6.33 as a function of a scan
(vi
9k0has the units of cm s-' (Section 7.2 in Chapter 2) and is therefore a first-order process.
'ONthough the regularity of the blocking pattern of this model may be a simplification
of the randomly blocked surface situation applying to cytochrome c, differences in the
(microscopic) non-linear diffusion patterns between model and real surfaces have only a
minor influence on calculated parameters [58,62].
he diffusion coefficient used in these calculations was D = 8 x low7cm2 s-' because
this value provided excellent agreement between thory and experiment. See reference [63]
for a discussion of the values reported in the literature.
480
Metallopvotein voltammetvy
(4
Fig. 6.32 Modelling of partially blocked electrode surfaces: (a) infinitely extending pattern of electroactive disc sites with radius ri and half average distance ro, as used in the model proposed by Matsuda
et al. (see text for details) for conventionally sized electrodes; (b) a finite number of non-interacting
electroactive microdisc sites of radius rSwithin a microdisc electrode of r a d m r: shaded part of the
microelectrode is electroinactive. Reproduced by courtesy: J. Electroanal. Chem. 3 14 (1991) 191.
Copyright, Elsevier.
rate.12 The change in shape from peaked to sigmoidal form with increasing scan
rate closely matches the experimental data for a cytochrome c concentration of
860 pM (Fig. 6.30(b)). Simulated and experimental voltammograms obtained
at other concentrations are also in excellent agreement over the scan rate range
of 10-1000 mV sP1. The rate of electron transfer used in these calculations
essentially corresponds to a reversible process, implying that the rate of electron
transfer is very fast at electroactive sites and zero at blocked sites which are
electrochemically insulating.
7.2
Cytochrome c voltammetry at carbon microdisc electrodes
The electrochemical properties of cytochrome c also have been investigated at
carbon disc microelectrodes which had an electrochemically calibrated radius
of 6.3 pm. In cyclic voltammetric experiments at these rnicroelectrodes, no
Faradaic response was detected when solutions were prepared directly from
commercially supplied sample [58]. Due to its much smaller surface area, a
microelectrode is expected to be far more sensitive to the presence of surfaceaffecting impurities present in the bulk solution than are normal-sized electrodes.
Therefore, the cytochrome c solutions were purified by charcoal filtration to
remove surface-active trace impurities estimated to be present in all lyophilized
samples [21,39,40].
At the microdisc electrode, under conditions of cyclic voltammetry with
purified samples, bulk cytochrome c concentration-dependent voltammetric
behaviour was observed, as was the case for normal-sized electrodes. For con< 250 pM and with scan rates 0.05 ( v 5 1.O V s-', the
centrations of c,,,
system showed the peak-shaped curves (see Fig. 6.34(a)) theoretically expected
for the given disc size and diffusion coefficient where all the electrode is
1 2 ~ h double
e
layer capacitance of 8 pF ~ m added
- ~ to the simulated data was estimated
from experimental curves.
Quantitative use of microscopic model
48 1
Fig. 6.33 Simulations of cyclic voltammograms at a partially covered electrode as a function of scan
rate, using the model of Matsuda et al. (see text for details). Parameters used in calculations: 19,0.992;
Q, 4.0~"; D, 8.0 x lov7 cm2 spl; kO, 1.0cms-l, a , 0.5; electrode area, 1.0crn2; concentration
of redox active species, 1.0 rnM; double-layer capacity, 8 pF cmP2. Reproduced by courtesy: J.
Electroanal. Chem. 314 (1991) 191. Copyright, Elsevier.
Table 6.2 Calculated geometric parameters 8 , ro, and
ri for a partially blocked graphite electrode surfacea
obtained as a function of cytochrome c concentration.
Values of parameters used in the calculation are D =
8.0 x loy7cm2s-'; k0 = 1.0 cm s-'; a = 0.5
-
aData obtained from reference [58]. Symbols are defined in the text; a
is the charge-transfer coefficient.
b ~ deviation
o
from an ideal reversible system (linear diffusion) within
experimental error.
482
Metalloprotein voltammetry
Potential (V vs Ag/AgCl)
Fig. 6.34 Cyclic voltammograms for reduction of cytochrome c at a carbon microdisc electrode
with r = 6.3 pm in a 100 mM NaC104 5 pm phosphate buffer (pH 7), as a function of scan rate.
For clarity only the (reduction) negative scan direction is shown, (a) cCyt.,, 150 yM; (b) cCy,.,,660 pM.
Scan rates: (al)v = 50mVs-l; (a2,bl)v = 1 0 0 m V ~ - ~(as,
; ba)v = 500rnVs-I. Reproduced by
courtesy: J. Electroanal. Chem. 3 14 (199 1) 191. Copyright, Elsevier.
surface active. However, for concentrations ccyt., > 500 pM, and for the
same scan rates, steady-state behaviour (sigmoidal-shaped curves) was observed
(Fig. 6.34(b)).Analysis of these steady-state curves (ccyt., = 660 pM) gave values
for the shift ofthe half-wave potential, E l p , from the reversible value (E: -E1p)
of (5 & 4) mV and for the Tomes potential criteria (EIl4- E3/4) of (63 & 4) mV.
As with the conventionally sized graphite disc electrode, this concentrationdependent voltarnmetric behaviour can be explained by partial self-blocking of
the electrode surface. Because there is no rigorous theoretical treatment available to describe the diffusion properties of partially blocked microelectrodes, a
simple approximation may be made to describe the observed response. It can be
assumed that by blocking parts of the microelectrode surface an array of noninteracting smaller electroactive sites of circular shape remains (Fig. 6.32 (b))
with the mass transport to each site being by radial diffusion. With this model,
the number and size of these microactive sites, and therefore the limiting current, depends on the cytochrome c bulk concentration because the blocked
fraction of the surface is a function of this concentration.
Theoretical calculations using the above approximate model were made by
digital simulation [58] (also see Section 12.2 in Chapter 2). Simulated curves
for different k0 values and electroactive sizes are shown in Fig. 6.35. For a
microelectrode of given radius, the number of active sites within the partially
covered rnicrodisc can be determined from the limiting current and must be
equal, according to the proposed model, to the sum of the calculated limiting
currents ofthe individual active sites. The limiting current to a single disc-shaped
Quantitative use of microscopic model
483
Fig. 6.35 Simulations of cyclic voltammograms at microdisc electrodes with radii (a) r, = 0.1 pm,
(b) r, = 5 pm, as a function of rate constant. For clarity, on1 the (reduction) negative potential scan
direction is shown: rate constants (al,bl) kO = oo; (aa,b2) k = 1.0 cm s-l; (as,bS) ko = 0.1 cm s-l;
cm2 s-l ; a , 0.5; scan rate,
(a4,b4) kO = 0.01 cm s-l . Other simulation parameters: D, 8.0 x
0.5 v s-' . Reproduced by courtesy: J. Electroanal. Chern. 314 (1991) 191. Copyright, Elsevier.
7i
active site is of course independent ofthe electrode kinetics and can be calculated
according to the equation (Section 10 in Chapter 2)
where I& is the limiting current, n the number of electrons (1for cytochrome c),
D the diffusion coefficient, F the Faraday constant and rs is the radius of the
electroactive site. The number of active sites j, is therefore given by
where I;: is the experimentally determined limited current. All other symbols have their normal meaning. The fractional coverage, 6 , of the microdisc
electrode can be obtained from
where r is the radius of the microdisc.
The fractional coverage of the normal-sized graphite electrodes has been
determined as a function of the cytochrome c bulk concentration (see Table 6.2).
If similar fractional coverages are also assumed to apply for the partially blocked
microdisc electrode, then by use of eqns (6.3) and (6.4), the size of the active
sites in the microdisc can be calculated for experimentally determined values
484
Metallopvotein voltammetvy
1
t
-0.3
t
0
Potential (V vs Ag/AgCI)
I
0.3
Fig. 6.36 Cyclic voltammograms at microdisc electrodes; experimental (+) and simulated (-)
data. For clarity only the (reduction) negative scan direction is shown. Experimental conditions:
660 pm cytochrome c in 100 rnM NaC104, 5 rnM phosphate buffer (pH 7) at a carbon microdisc
electrode with r , 6.3 pm; scan rate v , 0.5 V s-l. Simulation parameters: 5, 0.14 pm; js, 25; D,
8.0 x
cm2s-'; a , 0.5; v , 0 . 5 ~ s - l ;rate constants A, kO = m; B , k0 = l.0cms-l; C,
kO = 0.1 cms-I; D, kO = 0.01 cm s-l. Reproduced by courtesy: J. Electroanal. Chem. 314
(1991) 191. Copyright, Elsevier.
4:;.
of
Values lie within the range calculated at the large electrodes. The size of
the active sites is important because the kinetic sensitivity of the electrode to the
rate of electron transfer is directly related to it. The smaller the active sites, the
higher the kinetic sensitivity of the electrode (Section 14 in Chapter 2). This
fact is illustrated in Fig. 6.36, where it is shown that for a site size v, = 5 pm,
the voltammetric curves are indistinguishable for rate constants oo 2 ko 2
0.1 cm s-' , whereas when r, = 0.1 prn, even the curve for 1.0 cm s-' is well
resolved from the reversible case.
The value of ko as a function of vs [64] is calculated by
where K O is a dimensionless parameter and has been defined [64] as a function
of the half-wave potential (E: - Ell2)and the Tomes potential criterion
E3/4). For the experimental values cyt.c = 660 pM, (E: = (5 f.4) mV
and (E1/4-E3/4) = (63k4) mV, the parameter K O is given by log K O = 0.9A10.3.
The resulting calculated heterogeneous charge-transfer rate constants, ko, for a
range of fractional coverages similar to those obtained at the conventionally
sized graphite electrodes are presented in Table 6.3.
Figure 6.36 shows that an excellent fit of the proposed model with experimental data is obtained for a very fast rate of electron transfer in the sense
Quantitative use of micvoscopic model
48 5
Table 6.3 Heterogeneous charge-transfer rate constants
calculated for horse-heart cytochrome c for a range of
surface parameters of a partially blocked carbon disc
microelectrode"
"Data obtained from reference [58].
b ~ a d i uof
s the active site.
"umber of active sites calculated by use of eqn (6.4) and an experimentafly
determined limiting current of 72pA (660pM cytochrome c in 0.1 M
NaC104, 5 mM phosphate buffer, p H 7).
d~alculated
by use of eqn (6.5).
eCalculated by use of equation (6.6) for D = 8.0 x lop7 cm2 s-' and
l o g ~ O= 0.9.
f 19value obtained by interpolating values in Table 6.2 for cCY,,, = 660 pM.
that the experimental curve lies very close to the simulated curves for ko = oo
(fully reversible) and ko = 1.0 cm s-', but is considerably removed for the cases
ko = 0.1 and ko = 0.01 cm s-l. The conclusion is therefore reached that the
process is close to reversible and certainly much faster than 0.1 cm s-l.
The value for ko calculated assuming a similar coverage as at a large graphite
electrode is >0.4 cms-', which, as shown from data contained in Table 6.3
is about two orders of magnitude higher than the values previously calculated by using models in which ideal linear mass transport to a fully active
electrode surface is assumed [58]. Further support for the presence of partially blocked microelectrodes also comes from data provided by Hill et al.
[65] who studied the voltammetry of cytochrome at a very large assembly of
more than 8000 carbon-fibre microdisc electrodes. At this very large array
of microdisc electrodes, high quality near steady-state voltammograms were
observed (Fig. 6.37) down to cytochrome c concentrations of 1 pM. However,
the response deteriorated as cycling of the potential progressed as expected when
time-dependent adsorption of denatured or other forms of cytochrome c occurs.
Furthermore, curvature is evident in the limiting current versus concentration
plot (Fig. 6.37) again as expected when cytochrome c concentration-dependent
electrode blockage occurs.
7.3 Conclusions derived fvom modelling the voltammetry of
cytochrome c at carbon electrodes
The electrochemistry of cytochrome c at carbon electrodes can be explained in
terms of a model where the electrode surface is partially blocked, presumably
Metalloprotein voltanzmetry
-150 0
-300-200-100 0 100 200 300 400
Potential (mV)
60
Concentration (pM)
Fig. 6.37 The non-linear diffusion-limited current versus concentration plot obtained from voltammograms of horse-heart cytochrome c (5 mM sodium phosphate buffer (pH 7.0)/100 mM NaC1)
using a large assembly of carbon-fibre microdisc electrodes (scan rate 20 mV s-l). Inset: A cyclic
voltammogram obtained at the same electrode assembly with a scan rate of 10 mV s-' for the reduction of 49.6 yM horse-heart cytochrome c and a cyclic voltammogram of the buffer solution alone.
Adapted from: Lab on a Chip. 1 (2001) 127.
by the presence of some form of adsorbed protein. Under conditions where
electrode blockage is substantial, this self-inhibition process may dominate the
voltammetric behaviour of the interiace because the Faradaic process can take
place only at the few remaining active sites of the electrode surface. The nonlinear mass transport to these extremely small active sites then contributes
significantly to the voltammetry observed for cytochrome c. An equivalent model to that used to explain cytochrome c voltammetry also has been
successfully applied to the voltammetry of plastocyanin at carbon electrodes [59].
ence that chernica
e surface can a
ification of the
In the preceding part of this chapter, the voltammetric response of metalloproteins has been described at chemically modified or functionalized electrode
surfaces. Until now, the question as to whether the compounds used to modify the electrode surface are innocent in the thermodynamic sense has not
been raised. However, it will now be demonstrated that positively charged
13~dapted
from Electrochem. Commun. 1 (1999) 309
Chemical modification of electrode su$ace
487
compounds used to modify graphite electrode surfaces, may alter the measured
redox potential of both negatively changed recombinant and mutant forms of
rubredoxin and ferredoxin from Clostridium pasteurianum (Cp). This view contrasts to the frequently held assumption that modifiers are thermodynamically
innocent.
8.1 The thermodynamic effects of chemical modijcation of
graphite electrodes on rubredoxin electrochemistry
No Faradaic voltammetric response is observed when a freshly polished pyrolytic
graphite electrode is placed in contact with a recombinant rubredoxin (Rd) solution. However, after the electrode is modified ex situ by dipping into a 10 rnM
solution of highly charged cationic poly(L-lysine) (degree ofpolymerization, 20)
for 2min, followed by rinsing carefully with distilled water, a very welldefined voltammogram is observed in the reversible potential region. When the
electrode is modified ex situ with 160 mM tris(l,2-diaminoethane)chromium
chloride, [Cr (en)3]C13l 4 solution, again a well-defined voltammogram is
observed. In contrast, after using ex situ modifications with very high concentrations of Mg2+, only a very weak response is observed. Figure 6.38 contains
a comparison of square-wave voltammograms15obtained when poly(L-lysine),
[ ~ r ( e n ) ~ ]and
~ +Mg2+
,
ex situ modified graphite electrodes are placed in contact
with a rubredoxin solution.
The fact that these charged electrode surface modifiers are not thermodynamically innocent is revealed by noting that the reversible potentials obtained
by square-wave voltammetry using in situ modified electrodes depend on the
molar ratio of cationic modifier to rubredoxin. A positive potential shift is
always observed with increasing concentrations of modifier but the magnitude
of the shifts varies with the identity of the cation. The shifts in reversible potential are significant at low poly(L-lysine) to rubredoxin ratios, but approach a
limiting value as the ratio is increased (Fig. 6.39(a)). The peak potentials are
also positively shifted as the concentration of [Cr(en)3I3+ is increased. However, the magnitude of the shifts are much smaller than when poly(L-lysine) is
used under the same conditions (Fig. 6.39(b)). The measured peak potentials
for [modifier]/ [Rd] extrapolated to zero modifier concentration are essentially
identical for both poly(L-lysine) and [cr(en)313+ (Fig. 6.39(b)). The extrapolated value is believed to represent the 'correct' E,O value of the rubredoxin
redox process.
14[cr(en)313+in water undergoes hydrolysis to form an equilibrium containing mononuclear and binuclear ions [66]. However, for conciseness, this form of electrode modification
will be referred to as [cr(en)313+.
1 5 ~ hsquare-wave
e
technique is described in Section 4.1 in Chapter 2 and for a reversible
process the peak position is located at E ~ O [67]. Furthermore, the same shape applies irrespective of whether a macroscopic or microscopic electrode model applies, provided the
process is reversible as appears to be the case with square-wave voltammograms for both
rubredoxin and ferredoxin [68].
488
Metallopvotein voltammetvy
-500
-480
-360
-240
-120
0
Potential (mV) vs Ag/AgCl
Fig. 6.38 Square-wave voltammograms of 80 pM recombinant Clostridium Pasteurianum rubredoxin
in a 30 mM tris-HC1/0.1 M NaCl (pH 7.4) buffer at an edge-plane pyrolytic graphite electrode
modified by dipping into a solution of (a) 2mM poly(L-lysine); (b) 0.16M [Cr(en)3]C13; and
(c) 0.20 M MgC12. Square-wave amplitude 50 mV, frequency 43 Hz. Provided by courtesy: Z. Xiao
and A.G. Wedd, University of Melbourne, Australia.
8.2 Thevmodynamic effects of chemical mod$cation of
graphite electrodes onfervedoxin electvochemistry
Observations on the electrochemistry of recombinant ferredoxin Fd at an in situ
modified pyrolytic graphite electrode in the presence of modifiers are similar
to the rubredoxin case (compare Figs 6.39(a) and 6.40). The shift in reversible
potential of recombinant ferredoxin with increased modifiers concentration was
found to be greater in magnitude with poly(L-lysine) as the modifiers than with
[cr(en)3I3+.Total shifts were found to be f 3 0 and +15 mV respectively for
[promoter]/[Fd] = 12.0 (Fig. 6.41).
Reversible electrochemistry can also be obtained by ex situ modification of
the electrode by dipping into a concentrated solution of 160 mM [cr(en)313'
or 10 mM poly(L-lysine). Under these conditions, the peak current is constant
for up to 20 scans when poly(L-lysine) is used but decreases rapidly with repeat
scans in the case of [ ~ r ( e n ) ~ ]That
~ + .is, [cr(en)313+incorporated onto/into the
graphite electrode slowly dissolves from the surface into the interfacial region
and the bulk solution and it is the surface-confined state of the complex that
leads to the promotion of the voltammetry of ferredoxin.
Chemical modification of electrode
Potential (mV) vs SHE
(b)
-80
1
I
I
I
0
5
10
15
Ipromoter]:[r-Rd]
Fig. 6.39 Dependence of the reversible square-wave peak potentials (Epvalue) of 80 pM recombinant rubredoxin (r-Rd) obtained at an in situ poly(L-1ysine)-modified pyrolytic graphite electrode
on the solution-phase concentrations of modifier and protein: (a) square-wave voltammograms at
different solution-phase poly(L-lysine) concentrations; (b) plots of Ep versus [modifier]/[protein]
concentration ratio. Voltammetric conditions are as in Fig. 6.38. Reproduced by courtesy:
Electrochem. Comm. 1 (1999) 309. Copyright, Elsevier.
The electrochemistry of mutant D33,35,39N ferredoxin (charge -8 relative to -11 on the recombinant form) was also measured [68] to determine
whether a change in total protein charge would influence the efficiency of
promotion or the magnitude of potential shifts. [cr(en)313+was added in the
same [protein]/[modifier] ratios as for the native protein (Fig. 6-41(c)). The
Metallopvotein voltammetvy
Potential (mV) vs Ag/AgCl
Fig. 6.40 Square-wave voltammograms of 80 yM recombinant Clostridiumpasteurianurn ferredoxin
in a 30 rnM tris HCV0.1 M NaC1, (pH 7.4) buffer at a pyrolytic graphite electrode, showing
the dependence of the reversible peak potentials (EP)on poly(L-lysine): R d molar ratios; where
[poly(L-lysine)]/[Rd] = 0.25; 0.50; 1.O; 2.0; 4.0; 8.0; 12. Other voltammetric conditions are as in
Fig. 6.38. Provided by courtesy: Z. Xiao and A.G. Wedd, University of Melbourne, Australia.
maximum change in measured potential was approximately +6 mV in this case
which is significantly smaller than obtained with recombinant rubredoxin. Also,
in contrast to the work on the rubredoxin protein, the measured potential was
4.0
d ]and above.
constant for [ ~ r ( e n ) ~ ] ~ + /=[ ~
8.3 Conclusions concerning the dependence of the reversible potential
on the presence of a su?face modijier
The surface of a freshly polished pyrolytic graphite electrode contains negative patches due to the formation of carboxylate functionalities on the surface
that result from oxidation of carbon in the presence of air. Consequently, the
electrode can be modified by the incorporation of cationic species such as
[cr(en)3I3+and poly(L-lysine) onto the surface to give favourable electroactive
sites that do not readily become blocked by adsorption of protein. Modification
of the electrode can be achieved by dipping it into a concentrated solution of
cations. This contact with positive-charged species apparently results in association of the cation onto the electrode surface causing the electrode surface
Chemical mod$cation of electrode surface
49 1
10
[Modifier]/ [Fd]
Fig. 6.41 Dependence of reversible square-wave voltammetric peak potentials (Ep)on modifier:
ferredoxin molar concentration ratios for (a) 80 pM recombinant ferredoxin in the presence of
poly(L-lysine), (b) 80 pM recombinant ferredoxin in the presence of [ ~ r ( e n ) ~ and
] ~ +(c) 80 pM
D33/35/39N ferredoxin in the presence of [cr(en)313+.Voltammetric conditions as in Fig. 6.38.
Provided by courtesy: Z. Xiao and A.G. Wedd, University of Melbourne, Australia.
to have favourable electroactive sites with which the negatively charged proteins can interact, without encountering electrostatic repulsion or electrode
blockage. Reversible electrochemistry results at these favourable sites. However, it is important to note that the use of these modifiers may cause shifts
in the measured redox potentials of proteins, with the magnitude of the shifts
being dependent on the molar ratio of modifier to protein when an in situ form
of electrode modification is used.
It is possible that these shifts in potential are due to ion-pair formation between
the protein and modifier molecules such that the reduced form of the protein is
stabilized (causingthe protein to be more difficult to oxidize and thus making the
reversible potential more positive). As may be expected ifbinding is electrostatic
in origin, the strength of binding of modifier to protein and hence the shift in
potential as a function of modifier concentration seems to depend on the protein
charge. This is shown by comparison of data obtained in the [ ~ r ( e n ) ~titration
]~+
experiments with the D33/35/39N Fd mutant and recombinant ferredoxin
[68]. The mutant has a total charge of -8 at pH 7.5 as compared to the charge on
the native protein of - 11. The addition of [cr(en)3I3+caused a smaller shift and
]
ratios. The
approaches a limiting value at lower [ ~(en):'
r ] / [ ~ dconcentration
dependence of modifiers on the reversible potential of other metalloproteins also
492
Metalloprotein voltammetry
has been noted [69-721 and it always must be remembered that modification
of the electrode surface cannot be assumed to be an innocent activity in the
thermodynamic sense.
9 Long-range electron-transfer effects encountered in
cytochrome c voltammetry at long-chain alka
modified electrodes
In Section 18.2.3 in Chapter 2 it was noted that electron-transfer reactions
that occur over long distances lead to slow rates of electron transfer that may be
analysed in terms ofMarcus theory. In principle, voltammetry of metalloproteins
at suitable chemically modified electrode surfaces could lead to this situation
being achieved. For example, the schematic diagrams used in Figs 6.14 and 6.19
to explain the voltammetry of cytochrome c at partially blocked electrodes
could lead to electron transfer over long distances if the electrode modifier had
had significantly larger dimensions than SS-bpy and other electrode modifiers
considered so far. In Section 18 in Chapter 2, the voltammetry of cytochrome c
at SAMs was considered when long-chain alkane thiolates were attached to
gold electrodes. It can now be noted that use of these SAMs, discussed in
Chapter 2, in the case of metalloprotein electrochemistry, represents in reality
a special case of chemically modified electrodes that both inhibits problems
with electrode blockage, and which concomitantly leads to electron transfer
over long distances. This aspect of cytochrome c electrochemistry considered
in Chapter 2, can now be revisited to further illustrate that metalloprotein
voltammetry, after allowing for the nuances, is in fact not significantly different
from that obtained with smaller molecules. In particular, additional discussion
of Figs 2.92-2.95 of Chapter 2 will be considered in this context.
If cytochrome c becomes attached to the SAM, as in the studies of Bowden
et al. [9], cytochrome c electrochemistry, instead of being undertaken in diffusional mode as described above and with all the electrode-blockage problems,
can be performed in a diffusionless mode, in which a monolayer or submonolayer of the protein is immobilized on top of the alkane thiolate modified
electrode surface (Fig. 2.93). As discussed in Chapter 2 and elsewhere [73],
reaction schemes become considerably simplified upon exclusion of rate limitations arising from mass transfer and adsorption/desorption processes, leaving
a particularly simple overall reaction to characterize
in which the measured rate constants (ko), with units of s-l, is that of the
fundamental electron-transfer step itself (assuming that conformational changes
or other complications are not a factor), and surf denotes a surface-confined
species. Alternatively, of course, if cytochrome c is not immobilized onto the
modified surface, then diffusion-controlled theory can apply, but the heterogeneous electron-transfer rate will again be slow with ko values having the units of
Long-range electron-transfer effects
493
centimetre per second being measurable. In both the diffusional and diffusional
modes of metalloprotein voltammetry at these chemically modified electrodes,
Marcus theory can now be expected to be used when highly irreversible voltammograms are observed as a result of very slow electron transfer being introduced
by the long distance required for electron transfer.
Figure 2.94 shows some typical cyclic voltammograms (CV) obtained from
diffusionless cytochrome c voltammetry experiments using SAMs of two different thicknesses constructed using COOH-terminated alkanethiols. As noted in
Chapter 2, the area under the voltammetric peaks provides a direct measure of
the number of electroactive cytochrome molecules on the surface r(mol/cm2);
the sudace formal potential ([Efo](suq)can be obtained from reversible or nearreversible cyclic voltammograms as the midpoint between the peak potentials;
the peak separation can be used to determine the standard electron-transfer rate
constant (ko) for reaction (6.7). As noted in Chapter 2, there are always complications associated with these analyses, since the redox proteins are confined
to heterogeneous solid surfaces, thus giving rise to thermodynamic and kinetic
dispersions and other phenomena [74,75]. Nevertheless, for voltammograms of
the quality shown in Fig. 2.94, the complications are not overly serious, and
f , ko can readily be obtained using the fairly
accurate values for I?, [ ~ f o ] ~and
simple physical models presented in Chapter 2.
The two cyclic voltammograms shown in Fig. 2.94 show that as the alkyl
chain of the thiolate molecule is lengthened, the peak separation increases due to
a decrease in electron-transfer rate. The distance dependence of the electrontransfer rate for the cytochrome c modified gold electrode system has been
studied [76-781 by systematically varying the number of methylenes (n) in the
COOH-terminated alkanethiol spacer molecule. For n > 8, the exponential
decay of electron-transfer rate with distance shown in Fig. 2.95, which signifies
a non-adiabatic electron-transfer process involving long range electron transfer.
The tunnelling decay parameter obtained from these experiments [9] are consistent with through-bond tunnelling of electrons along the alkane chains forming
the modified gold electrode and are consistent with comparable studies of small
molecule electroactive SAM [79]. However, for cytochrome c, adsorbed on
thinner alkane thiolates (n 5 8), there is disagreement at present as to whether
the electron-transfer rate continues to decay exponentially or levels off as the
distance continues to decrease [9].
As noted in Chapter 2, peak separations for cyclic voltammograms such as
shown in Fig. 2.94 can be used with the simple Butler-Volmer model, described
in the seminal work by Laviron [go], to readily give ko values used to construct
Fig. 2.98. These values of ko (s-l) are uniquely those for zero overpotential,
[Efolsurf, where the rate is controlled by the intrinsic free-energy barrier [81].
Also as noted in Chapter 2, the other activation issue to consider is the driving
force or potential dependence of the rate of electron transfer. In order to extract
the reorganization energy (A)and the intrinsic activation barrier for the redox
protein, theoretical linear sweep voltammetric treatments of the kind described
in Chapter 2 and references [82,83] have been applied to cytochrome c adsorbed
494
Metallopvotein voltammetvy
on self-assembled alkane thiolate monolayers [84]. A value for h of 0.3 eV
was determined based on fitting the theoretical analytical expression to the
voltammetric data, but the presence of non-ideal peak broadening dictates that
this value be considered as a lower limit to the true h [9].
A second approach to achieve long-distance electron transfer, developed
by Miller's group [85,86], does not rely on cytochrome c immobilization
to circumvent the mass-transfer problem, but using the hydroxy-terminated
long-chain alkane thiol surface modification to slow the normally rapid heterogeneous electron-transfer kinetics (rate constants having cm s-' units) to
sufficiently low values. This strategy has the effect of shifting the voltammetry
wave to irreversible potentials (Fig. 2.92), well into the Marcus inverted region,
whereupon current-voltage data can be acquired at the foot of the voltammetric wave, where mass-transport limitations are minimal (see Fig. 2.92). This
approach has been used to measure the reorganization energy for several species
of cytochromes and values of around 0.6 eV have been found [86]. Determination of reorganization energies for redox proteins has been a challenge for
decades, as evidenced by a lack of consensus regarding such a fundamental
kinetic parameter as the reorganization energy for cytochrome c [9]. With
these electrochemical approaches using chemically modified electrodes, the
prospects for this important measurement becoming routine have brightened
considerably [9].
10 Voltammetry of metalloprotei
surfactant environments
In the preceding part of this chapter, almost ideal voltammograms of metalloproteins have been obtained using highly purified protein solutions or chemically
modified or functionalized electrode surfaces. An alternative method of achieving well-defined voltammetry for metalloproteins is to undertake studies with
material trapped in surfactants that are cast as thin films on electrode surfaces
[87]. Again, as with the case of a chemically modified electrode, in this environment, no direct interaction of a metalloprotein with an electrode surface occurs
to block electron transfer and reversible well-defined electrode processes are
observed. This surfactant actually provides a medium akin to the lipid bilayer
biomembrane-like environment [87] in which many metalloproteins exist in
nature [88].
In living organisms, biomembranes are about half-phospholipids and halfprotein and generally exist in a partly fluid, semi-permeable state [87-891. In
the early 1980s, Kunitake and co-workers showed that ordered films of waterinsoluble surfactants can be cast onto solid surfaces from organic solvents or
aqueous vesicle dispersions [87,90]. Evaporation of solvent leaves thin, selfassembled multi-bilayer films, similar to stacks of biomembranes. Similar films
can be made with surfactants and polymers [91-931. Several research groups have
described the reversible electrochemistry of small molecules in the liquid crystal
Metallopvoteins in sufactant envi~onments
495
state present in these cast multi-Maya surfactantfilms [%-I021 and this concept
has been extended to achieve well-defined metalloprotein voltammetry. Waterinsoluble surfactants are used in these studies, usually with two alkyl chains, with
an example being the surfactant DDAB (Fig. 6.42).
The work of Hawkridge and co-workers [87] shows that myoglobin (Mb)
in solution-phase voltammetry, like cytochrome c and other metalloproteins,
gives ill-defined voltammetry on bare metal electrodes, although it is detected
with reasonable definition on indium tin oxide electrodes [103-1051. In contrast, direct reversible electron transfer between pyrolytic graphite, Au, or Pt
electrodes and myoglobin is readily obtained in water-insoluble films of DDAB
(Fig. 6.43) [I O4,106,lO7] immobilized onto the electrode surface, which is in
turn in contrast with an aqueous electrolyte media. In order to achieve charge
balance the electrode reaction requires a proton from the solution phase, and
the process can be represented as
[Fe(II)](surfactant)
(6.8)
didodecyldimethylammonium bromide (DDAB)
Fig. 6.42 Structure of the surfactant didodecyldimethylammonium bromide (DDAB).
I
I
I
I
I
-0.40
-0.20
0.00
E (V) vs SCE
0.20
0.40
Fig. 6.43 Cyclic voltammograms at 100rnVs-l in pH 5.5 buffer: (a) pH 5.5 buffer on a bare
pyrolytic graphite electrode; (b) 0.5 mM m~oglobinin buffer on bare pyrolytic graphite; (c) myoglobin-DDAB film on pyrolytic graphite in buffer, no myoglobin in solution. Adapted from: Anal.
Chem. 67 (1995) 2386; Electrochem. Inteface 6(4) (1997) 26.
496
Metallopvotein voltammetry
In a qpical experiment [87] a solution of DDAB in chloroform is spread on
an electrode surface and the solvent is allowed to evaporate. The surfactantcoated electrode is then placed in an electrochemical cell containing a solution
of myoglobin, and the protein is taken up rapidly into the film. Diffusion of
myoglobin within the film accounts for the fast uptake. Diffusion of protons
also accounts for proton charge transport during cyclic voltammetry [I061 so
that a diffusion-controlled form of voltammetry is observed as also can occur
with microcrystals adhered to electrode surfaces and in contact with aqueous
electrolyte (Chapter 5).
A wide range of other water-insoluble surfactants can be used to make
liquid crystal Glms in which myoglobin gives reversible electron transfer
[87,106-112]. Use of insoluble surfactant films appears to be a relatively general
way of obtaining well-defined voltammetry for metalloproteins as evidenced
by the measurement of reversible processes for cytochrome P450,,, (Fig. 6.44),
haemoglobin, cytochrome c, and Chlorella ferredoxin in films of various surfactants [112-1161 or composites of DDAB and Nafion [117]. Bianco et al.
reported reversible electrochemistry for cytochrome c, cytochrome c3, and
cytochrome ~ 5 5 3in lipid films doped with lauric acid [I 18-1 201, and for spinach
ferredoxin in phosphatidylcholine-cholesterol films doped with dodecylamine
or didodecyldimethylammonium bromide [I 2 11.
A question that arises is: why is voltammetry so well defined and reversible
when proteins are in surfactant films compared to proteins in solution with a bare
electrode? Again to answer probably is associated with avoidance of electrode
blockage. Reversible electron transfer is found on bare electrodes only when
-10 1
I
-0.80
I
I
I
I
-0.60
-0.40
-0.20
0.00
E (V) vs SCE
Fig. 6.44 Cyclic voltammograms on basal-plane pyrolytic graphite electrodes obtained at a scan rate
of 100 mV s-' in pH 7 buffer containing 0.1 M KC1: (a) substrate-free cyt P450Cam-DMPC film
in oxygen-free buffer containing no enzyme; (b) bare electrode in oxygen-free buf-fer containing
40 pM cyt P450Cam;
and (c) DMPC film in oxygen-free buffer containing no enzyme. Adapted from:
J. Chem. Soc., Faraday Trans. 93 (1997) 1769; Electrochem. Inteface 6(4) (1997) 26.
References
497
myoglobin is purified by chromatography immediately before the experiment,
and only on hydrophilic tin-doped indium oxide [l 03,1041 or pyrolytic graphite
electrodes [122]. As is the case with cytochrome c (Section 4.2.5) exposure of
these electrodes to constituents present in partly purified protein solutions blocks
electron transfer to myoglobin (Fig. 6.43(b)) [l 221. Reflectance-absorbance
FTIR and X-ray photoelectron spectroscopy reveals the presence of protein
containing adsorbed material on pyrolytic graphite and Pt electrodes exposed
to these partly purified solutions of myoglobin [122]. Adsorption of this material
blocks electron transfer to myoglobin and even ferricyanide [122]. Exposure of the same electrodes with solutions of cationic surfactants removes the
macromolecular adsorbates, allowing well-defined voltammetric peaks for Mb
to be observed. Thus, electron transfer in myoglobin containing surfactant
films is facilitated partly by strong adsorption of surfactants at the electrodefilm interface, which in turn inhibits adsorption of protein macromolecules
from myoglobin solutions, thereby preventing the blockage of electron transfer
between the metalloprotein and the electrode [l 031.
11 Conclusions related to t
etalloproteins
The major factor that inhibits electron transfer between electrodes and metalloproteins is adsorption onto electrodes of macromolecular impurities or
denatured or non-native forms of the metalloprotein itself. This adsorption can
create a passive layer that blocks electron transfer. Clearly, the most straightforward way to overcome this problem, at least in principle, is to use highly purified
protein solutions and well-cleaned electrodes. Alternatively, the electrode can
be coated with a suitable monolayer of a redox-inactive surface modifier, or
naturally functionalized electrodes containing suitable non-blocked electrontransfer sites can be employed. A novel extension of the surface modification
method employs long-chain alkylthiol monolayers on gold electrodes with end
functional groups that may even be capable of binding proteins with strong
electrostatic interactions or even covalent bonds. Finally, use of surfactants may
also lead to prevention of electrode blockage and well-defined voltammetry.
In summary, after accounting for all the nuances, the voltammetry of
metalloproteins can be conducted under conditions that are characterized by
well-defined current-voltage curves that are characterized by rapid electrontransfer reactions, as would be expected from electron transfer and structurally
related metalloproteins.
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ndex
AC, see alternating current voltammetry
accumulation potential 282, 307, 312
activation parameters 124
active sites 483
activity 11, 14, 17, 336, 347
activity coeffkient 14, 15
ad-atoms 457
adhered solid 334, 339, 369
adlayers 460, 461, 462
adsorption 130, 149, 151, 168, 249, 265, 277,
312, 314, 432, 450
array of molecules 461, 465
azurin 444
cytochrome c 153, 454
facilitating molecules 466
influence on voltammetry 150
isotherm 149
metal complexes 283, 312
prevention 460
protein 441, 486
solution contaminants 452
adsorptive stripping chronopotentiometry 310
adsorptive stripping voltammetry 283, 301, 307,
310, 319
AFM
polished edge-plane graphite electrode image
473
polished polycrystalline gold electrode image
463
topographical images of T C N Q 408, 410, 41 1,
412,413
see also atomic force microscopy
A ~ / A ~reference
+
electrode 39
Ag/AgCl reference electrode 38
alkane thiols 139, 140, 150, 494
alternating current voltammetry 33, 35, 41
amalgams 282, 284, 286, 290, 295, 297, 298
analysis of whole blood 328, 329
angular velocity 68, 71
anode 6,24, 26,27
anodic stripping voltammetry 282-301
Anson plot 266
apparatus considerations 101
array of adsorbed molecules 461, 465
array of electrodes
carbon 286, 287, 293, 294
disc 24, 284, 288
platinum 465, 472
see also microarray electrodes
array of microcrystals or microparticles 267, 336,
338, 369
Arrhenius equation 136
aspect ratios of crystals 383, 384
Atkins, P.W. 3, 4
atomic force microscopy 156, 159, 160, 408, 438
see also AFM
atomically flat gold 443
attachment of solids to electrodes 334
auxiliary electrodes, see counter electrodes
Azurin 131, 146, 443-450, 463
background current 35, 36, 46, 48, 103, 286,
450,463
see also non-Faradaic current
background subtraction 281, 286
band electrodes 79
Barker, G.C. 2
basal plane graphite electrodes 338, 369, 464,
468, 471, 473, 477
batteries 1, 5, 9, 10, 23, 25, 26, 437
Becquerel, E. 2, 28
Berzelius, J.J. 2, 11, 13, 14
biologically important processes or molecules
131, 441, 442,443
biosensors 249, 319, 321, 322
see also glucose biosensor
4,4'-bipyridyl disulfide, see SS-bpy
blocking of electrodes 442, 453, 479, 482, 485,
497
see also self-blocking processes
Boltzmann constant 141
boundary conditions 60, 62, 90
bulk electrolysis 41, 100, 102-107, 115, 212, 253,
256,257,308-310
at dropping mercury electrodes 308
at gauze or mesh electrodes 107
at mercury pool electrodes 308
cells 104-106, 208, 212
see also exhaustive electrolysis
n B ~ 3 P234,241-243
Butler-Volmer theory 54, 63, 64, 88, 136, 138,
140, 141, 142, 143, 144, 146, 418,
419,432,433
cadmium amalgam 299
cadmium ion reduction 305
calibration curves in analysis 285
Index
calomel reference electrode 38, 39
capacitance 4, 5, 6, 38, 82, 192, 220, 338, 453
see also double layer
capacitance current 47, 48, 58, 65, 68, 77, 101,
152
carbon electrodes 4, 33, 36
fibre 24, 25
microdisc 24
see also basal plane, diamond, edge plane, glassy
carbon, graphite, HOPG and pyrolytic
electrodes
carboxylate functionalities on electrode surfaces
490
Carlisle, A. 2, 27
catalytic processes 26, 130, 146, 148, 149, 178,
255, 256,303, 324, 326,327
catalytic electrodes 26
cathode 6, 23, 24, 26, 27
cell time constant 82
cells, see electrochemical cells
centre line velocity parameter 72
channel electrodes 52, 71-75, 90, 94, 98, 108,
109, 110, 112, 113, 117,230-235,
240, 244
channel flow-spectroelectrochemical cells 108
charge 84, 102, 164, 336, 352
see also coulometry
charge density 194, 372
charge neutralization process 129, 336, 436
charge on the electron 141
charge-transfer process 11, 14, 83, 478
charge-transfer coefficient 54, 141, 182
charging current 47, 78, 80
see also background current and double layer
current
chemical current 312
chemical energy 26
chemical potential 21
chemical reactions coupled to electron-transfer
processes 65, 146
chemical reversibility 55, 57, 83
chemical synthesis based on knowledge of
electrochemical data 213
chemically irreversible processes 57,
chemically modified electrodes 33, 130, 442, 443,
460,465,486, 487,492,494
see also functionalised electrodes and modified
electrodes
chronoamperometry 376
see also double-potential step experiments
chronocoulometry 266, 267
chronopotentiometry 310
chronopotentiometric adsorptive stripping analysis
310
Clark oxygen electrode, see oxygen electrodes
closed system 21
cobalt dimethylglyoxime 301-319
cobalt ion reduction 306
collection of products at ring electrodes 396
combinations of techniques used to elucidate
mechanisms 243-245
comparison of experimental and theoretical
voltammograms 181, 186, 187, 189,
191, 222,261, 293,299
comparison of voltammetric techniques 95-100
composite electrodes 4, 336
concentration distributions 88, 90, 194
concentration gradients 52
conditional potential 15
see also formal potential
conducting polymers 130, 168, 336, 425
constant current electrolysis 103, 310, 312, 314
constant potential electrolysis, see controlled
potential electrolysis
controlled potential electrolysis 102, 103, 208,
227, 309,310, 400, 437
convection 10, 49, 51, 52, 68, 69, 75, 77, 89, 93,
109,284,396
see also natural convection
convergent diffusion 79
see also radial diffusion
convolution voltammetry 83, 88
see also Semi-integration
copper deposition 23-25, 92, 93
copper nail experiment 7, 19
Cottrell equation 95
Cottrell, F.G. 2
coulometric stripping chronopotentiometry
310-314
coulometric titration 210, 21 1
coulometry 102, 104, 116,227,308-313,437
microcoulometry 308, 309
counter electrodes 7, 34, 35, 106, 407
counterion diffusion/migration 424
coupled electron and ion transport 362, 365, 367,
424, 425,436
coupled electron and proton-transfer 179, 195
CPE, see controlled potential electrolysis
Cr(C0)3(C6Me6) 114, 115
cis and trans-Cr(C0)2 (Ph2PCH2CH2PPh2)2
345-365,402-407
[cr(en)313+-modified electrode 488-491
[ ~ r ( ~ H ~ ) ~ ] ~ + - m o electrode
d i f i e d 470
critical potentials 372, 373, 374, 375, 379, 380
critical radius 379
cross redox reactions 188, 224, 250
Cruikshanks, W. 2
crystals and their role in electrochemistry 361,
382-395,408-416, 437,438
crystal sensitivity factor 166
Index
current density 79, 80, 287
current-time relationships 103, 104, 422
cyclic voltammetry 41, 43, 45, 46, 47, 56, 57-68,
98, 101, 197-203, 244
chemical reactions coupled to electron transfer
65-68
depletion 75
derivative 88, 89
fast scan rate 80, 82, 83, 113, 280, 281
irreversible processes 62-64
quasi-reversible processes 64, 65, 85
reversible processes 60-62, 84, 85, 86
theory 59-68, 128-152
1,2-cyclooctanedione dioxime 3 16-3 18
cylindrical electrodes 78
cytochrome c 150, 152, 153, 154, 442, 449,
450-459,460,461, 464, 467,468,
470, 471, 474,477-486, 492-494,496
cytochrome c3 496
cytochrome P450 496
Dalton 14
data analysis service 91
Davy, H. 2, 11, 13, 14
decameth~lferrocene213-21 5, 337-345, 361
decay constant 143
defect sites 460, 470
density of electronic states 143
depletion 75, 288
deposition step 282
derivative cyclic voltammetry, see cyclic
voltammetry
2,3-diaminobutane 303, 315
diamond electrodes 4
didodecyldimethylammonium bromide 495, 496
differential pulse voltammetry 41, 42
diffusion 10, 49, 50, 52, 62, 68, 75, 79, 89, 90,
93, 161
coefficients 49, 184, 193, 205, 270, 276, 284,
425
coupled with electron transport 425
layer 81, 161, 452
of ions within a solid 364, 425
semi-infinite 50, 52, 84
surface 367
three-dimensional 52, 424
two-dimensional 52, 101
diffusionless voltammetry 129, 133, 144, 154, 493
digital simulations, see simulations
2,3-di(hydroxy1amino)butane 303, 306, 315, 319
dimensionless variables 96, 98, 130, 290
dimethylglyoxime 30 1-306
disc electrodes 37
505
see also microdisc, rotated disc and inlaid disc
electrodes
disproportionation 18, 19, 111, 188, 192, 207,
238,243,250, 264
dissolution 168, 336, 343, 367, 395, 396, 397,
398, 413, 414, 423, 437, 438
and precipitation processes 403, 405, 413
and redistribution reactions 408, 410
distance dependence of the rate of electron
transfer 140, 154
see also Marcus theory
distance scale 286, 287
DME, see dropping mercury electrodes
double layer 35, 82, 152
capacitance 46, 65, 182, 260
current 46, 312
see also charging current
effects 204
double-potential step experiments 266, 352, 353,
354, 357, 358, 376,378
drop time of mercury electrode 305, 308, 318
dropping mercury electrodes 293, 308
dye sensitized cells 28, 29
E~O,see formal potential
EiI2, see reversible half-wave potential
Ell2 87, 91, 96, 102, 202, 218, 274
see also half-wave potential
E1/4 76, 87
E3/4 76, 87
E3/4-E1/4 or E1/4-E3/4 76, 204, 218, 271, 276,
482, 484
edge-plane graphite electrode 433, 448, 464, 468,
470, 471
electrochemical cells 7, 9, 11, 16, 21, 22, 23, 26
high pressure 129
voltammetric 34-36
see also bulk electrolysis and non-isothermal
electrochemical cleaning 450
electrochemical-ESR experiments 108-1 13, 317
see also ESR, EPR and SEESR
electrochemical quartz crystal microbalance 155,
162, 164-171, 264-266, 269,
276-278, 335, 349-358, 437
electrochemical rate constants dependence on
potential, see Butler-Volmer theory
and Marcus theory
electrochemical reversibility 55, 56
electrochemical series 11, 12, 13, 14
electrochemical stripping techniques 283
see also stripping voltammetry
electrochemical synthesis 6, 100, 104, 196, 21 1,
212,226-230
see also bulk electrolysis
Index
electrochemical technology 29
electrochromic devices 437
electrocrystallization 152
electrodes 4, 6
see also band, bulk electrolysis, carbon, channel,
composite, diamond, dropping
mercury, gauze, glassy carbon, gold,
graphite, hanging mercury drop,
hemi-cylinder, hemispherical,
hydrodynamic, indium oxide,
ion-selective, lead, macrodisc,
mercury, mercury pool, mercury
thin-film, mesh, metal, micro jet,
microarray, microband, microdisc,
microhemispherical, microring,
modified, optically transparent,
oxygen, planar, platinum,
polycrystalline gold, polymer modified,
pyrolytic graphite, quasi-reference,
RAM, reference, ring, rotating disc,
rotating ring-disc, sacrificial,
semiconducting, silver, smart,
spherical, titanium, wall-jet working
electrode area 47, 54, 76, 80, 95, 103, 182, 189,
220,286,288
see also many equations that contain this term
electrode blockage 204, 447, 451, 452, 458, 460,
486
electrode edge effects 287
electrode fouling 4, 6, 203, 204
see also blocking
electrode functionality 442
electrode kinetic effect 204
electrode mechanisms 43, 68
CE 73
CECE 234, 235
CECEC 233
C2E 240
EC 55, 57, 66, 82, 96, 111, 126, 130, 146,
217, 259
EC2 83
EC2E 96, 97, 100
ECE 55, 56, 66, 67, 68, 69, 78, 80, 81, 82, 96,
99, 102, 104, 111, 149,217
Eccatalpic 130
EE 130,217
ErevCirrevErev 55
see also square reaction schemes
electrode passivation 232, 234
see also blocking of electrode
electrode potentials 3, 19
see also standard potentials
electrode radius 79, 287
electrode rotation 69-71
electrode size 3, 4, 80
electrode surface loadings 313
see also surface coverage
electrodes, practical considerations 101
electrodeposition 23, 24, 25, 287
electrode-solution interface 91, 92
electrolysis 3, 11, 14,24,33,79, 104, 113, 115,
116, 122, 313
see also bulk electrolysis
electrolyte 27, 28, 113, 205, 342-345, 347-349
electrolytic cell 7, 10, 31
electron hopping 365, 424
electron microprobe X-ray analysis, see electron
probe analysis
electron paramagnetic resonance 108-1 13, 207,
208, 210, 211, 395, 396, 399-403,
405-407, 438
electron probe analysis 341, 342, 352, 356, 382
electron scanning microscopy 145, 367, 382-386
see also scanning electron microscopy
electron spin resonance, see electron paramagnetic
resonance
electron transfer 52-54
long-range 492-494
electronic conductor 92
electronic coupling 143
electronic spectra 226
see also UV-visible spectroelectrochemistry
electronic states 144
electron-transfer processes 14, 23, 48, 62, 90,
131, 140,436
electrorefining 1, 23
electrospray mass spectrometry 119, 122, 123,
258
electrostatic force microscope 156
electrosynthesis, see el