CHEMISTRY OF METALS IN BIOLOGICAL

SYSTEMS

Direct electrochemistry of redox proteins

Christophe Leger, Fraser Armstrong,?

Inorganic Chemistry Laboratory, Oxford OX1 3QR, U.K.

Louvain, May 2002

Contents

1 Background: redox thermodynamics 21.1 The Nernst equation . . .. . . . . . . . . . . . . . . . . . . . . . . . 21.2 Reference potential and reference electrodes . . . . .. . . . . . . . . 41.3 Influence of ligand binding on reduction potentials .. . . . . . . . . 4

2 Mediated potentiometry 7

3 Introduction to direct methods 83.1 Electrodes for electron transfer to/from proteins . . .. . . . . . . . . 83.2 Electrochemical equipment . . .. . . . . . . . . . . . . . . . . . . . 93.3 Vocab and conventions .. . . . . . . . . . . . . . . . . . . . . . . . 10

4 Bulk Voltammetry 104.1 : : : at stationary electrodes . . .. . . . . . . . . . . . . . . . . . . . 104.2 : : : at rotating electrodes. . . . . . . . . . . . . . . . . . . . . . . . 124.3 : : : at stationary micro-electrodes . . .. . . . . . . . . . . . . . . . 13

5 Protein Film Voltammetry 145.1 Non-catalytic voltammetry at slow scan rates . . . .. . . . . . . . . 145.2 Fast scan voltammetry .. . . . . . . . . . . . . . . . . . . . . . . . 175.3 Catalytic voltammetry with adsorbed redox enzymes. . . . . . . . . 20

6 PFV Quiz 25

References 26? fraser.armstrong@chem.ox.ac.uk

1 Background: redox thermodynamics

1.1 The Nernst equation

Consider a reaction mixture containing the oxidised and reduced forms of two differentspecies (1 and 2):

Ox1 +Red2 Red1 +Ox2 (1)

The free energy of the reaction (�rG) is given by

�rG = �rG0 +RT ln

[Red1][Ox2]

[Ox1][Red2]

R is the gas constant,T the absolute temperature and�rG0 is the (tabulated) standard1

free energy of the reaction.If equilibrium between the different species is reached,�rG = 0, and the ratio of

concentrations (the reaction quotient) is linked to�rG0 by the relation:

Keq =[Red1]eq[Ox2]eq

[Ox1]eq[Red2]eq= exp

���rG

0

RT

�(2)

Initially upon mixing Ox1 and Red2, the concentrations of various species in thesolution do not satisfy eq. (2), and�rG is non zero: the system is not at equilibrium.Thermodynamics predicts2 that reaction (1) will proceed in the forward direction untilthe reaction quotient equalsKeq.

During the reaction, Red2 is oxidised and gives electrons to Ox1.

Ox1 + ne� ! Red1

Red2 ! Ox2 + ne�

It is possible to measure the flux of electrons from Red2 to Ox1, and therefore therate of the overall reaction, by placing the species in a two-compartment cell (Ox1 inone compartment and Red2 in the other). Place in each side an electrode at which thespecies can interact. When the two electrodes are connected together, a current flows,while the system evolves towards equilibrium.

The current flow between the two electrodes results from a potential differencebetween them, the value of which can be predicted using theNernst3 equation. The

1“Standard conditions” means that the activity of all constituants is unity, and the pressure equals oneatmosphere.

2Thermodynamics predicts only the direction of the reaction. In reality, a “favorable” reaction might nothappen because its rate is very small, in which case the equilibrium cannot be reached.

3Walther Hermann Nernst (1864-1941) was born in West Prussia. He obtained his Ph.D. at Wrzburg(Friedrich Kohlrausch), beginning his career as a physicist.

He graduated in 1887 with a thesis on electromotive forces produced by magnetism in heated metal plates.

2

potentialsE1 andE2 of the two electrodes are given by:

E1 = E01 +

RT

nF

ln[Ox1]

[Red1](3a)

E2 = E02 +

RT

nF

ln[Ox2]

[Red2](3b)

E01 is the standard reduction potential of the redox couple Ox1/Red1. F = 96500C

is the Faraday4 constant. The potential difference between the electrodes in compart-ments 1 and 2 is therefore:

V = E2 �E1 = E02 �E

01 +

RT

nF

ln[Ox2][Red1]

[Red2][Ox1](4)

The electrons are going to flow from the cell whose electrode potential is the lowest tothe other until the potential differenceV is zero and the concentrations satisfy

Keq =[Red1]eq[Ox2]eq

[Ox1]eq[Red2]eq= exp

��nF

RT

(E02 � E

01)

�

This is equivalent to eq. (2), since reduction potentials and free energies are linkedby

�rG0 = �nFE0

r

Reaction (1) will proceed forward significantly (Keq will be large) only ifE01 >

E02 . (If E0

1 is “high”, Ox1 is called a strong oxidant and Red1 is a weak reductant.)The Nernst equation can therefore be used to determine the direction in which a redoxreaction will proceed spontaneously.5

Nernst’s first outstanding work was his theory of the electromotive force of the voltaic cell (1888). Nernstapplied the principles of thermodynamics to the chemical reactions proceeding in a battery. In 1889, heshowed how the characteristics of the current produced could be used to calculate the free energy change inthe chemical reaction producing the current. The Nernst Equation, which related the voltage of a cell to itsproperties, also has very important biophysical implications. In 1905, he was appointed Professor of physicalchemistry in the University of Berlin. His heat theorem, known as the Third Law of Thermodynamics, wasdeveloped in 1906.

In 1914 Walther Hermann Nernst and Max Planck succeeded in bringing Albert Einstein to Berlin.4Michael Faraday (1791-1867). His scientific work laid the foundations of all subsequent electro-

technology. From his experiments came devices which led directly to the modern electric motor, generatorand transformer. Faraday was also the greatest scientific lecturer of his day, who did much to publicize thegreat advances of nineteenth-century science and technology. During the 1830s, Faraday worked on devel-oping his ideas on electricity. He enunciated a new theory of electro-chemical action between 1832 and 1834one of the results of which was that he coined many of the words now so familiar - electrode, electrolyte,anode, cathode and ion to name but five. His work led him to reject the traditional theory that electricity wasan imponderable fluid or fluids. Instead, he proposed that electricity was a form of force that passed fromparticle to particle of matter.

5It may not be enough to compare standard reduction potentials however, since the sign ofln[Ox2][Red1][Ox1][Red2]

in eq. (4) can change the sign ofV (and�rG) and thus the direction of the current flow.

3

Figure 1: Thebiological redoxspectrum at pH 7.

1.2 Reference potential and reference electrodes

If you are interested in studying only one half-reaction, it is convenient to make surethat the potential difference between the two electrodes reflects the potential of theelectrode you’re interested in. This can be done by using in the other compartment anelectrode designed to have a constant potential: this is called a “reference electrode”.

TheStandard Hydrogen Electrodeis one of these. It consists of a platinum elec-trode immersed in a pH=0 electrolyte below one atmosphere of H2.

By convention, the potential of the Standard Hydrogen Electrode is zero. All re-duction potentials tabulated in the literature are quoted versus the SHE.

A real SHE is rarely practical. Instead one uses convenient reference electrodessuch as:

�Hg/Hg2Cl2/KCl (Standard Calomel Electrode, SCE, E(SCE)=241mVvs. SHE.) Thisis the most commonly used reference electrode

� Hg/Hg2SO4/K2SO4, used when chloride ions must be avoided. E(SCE)=615mVvs.SHE.

� Ag/AgCl (E=800mVvs. SHE)

Because standard conditions1 (pH 0!) are not suitable for biological reactions,reduction potentials are usually stated for pH 7, and termedE

00

orEm;7.

Figure 1 gives an idea of the range of reduction potentials spanned by biologicallyimportant redox couples.

1.3 Influence of ligand binding on reduction potentials

Redox reactions can be coupled to other chemical equilibria such as ligand binding(e.g. protons, substrate, inhibitor: : : ) or conformational changes.

4

�rG03 = �nFE

03

Red:H

�rG02 =

-2.3RTlog10KRed

Red�rG

01 = �nFE

01

+e�

+e�

+H++H+�rG04 =

-2.3RTlog10KOx

Ox

Ox:H

Figure 2: Squarescheme for aprotonation reactioncoupled to a redoxprocess.

A very common (and physiologically important) coupled reaction is proton transfer,as represented in the square scheme fig. 2.KOx andKRed are the acidity constantsfor Ox and Red. Utilising the principle of thermodynamic cycles (the sum of�rG

0

values round the square is zero), these acidity constants can be linked to the reductionpotentials of the protonated and un-protonated redox couples. Note that potentialsalone cannot be summed; they must be scaled byn.

Usually Red is a much better base than Ox, so it has a higherpK, i.e. pKOx <

pKRed.� ForpH < pKOx, both Ox and Red are protonated.

Ox : H + ne� ! Red : H

The reduction potential ispH independent, and equalsEacid (E03 in fig. 2).

� ForpH betweenpKOx andpKRed, Ox is not protonated but Red is.

Ox + ne� +H

+ ! Red : H

The reduction potential is pH dependent: it decreases by2:3RT=nF Volts per pH unit( 59n

mV/pH at 25�C).In general, for a redox process involvingn electrons andm protons, the maximal pH-dependence is

2:3RT

F

m

n

V=pH (5)

� ForpH > pKRed, neither Ox nor Red are protonated, the redox process is

Ox + ne� ! Red

The reduction potential ispH independent, and equalsEalk (E01 in fig. 2).

For ann-electron, one-proton process, the whole pH-dependence of the reductionpotential is given by:

E00

([H+]) = Ealk +2:3RT

nF

log10

0@1 + [H+]

KRed

1 + [H+]

KOx

1A (6)

5

Figure 3: (A) Simplified Pourbaix diagram fora 1e�:1H+ reaction. (B) Reduction potentialvs.pH for amicyanin from Paracoccus versatusat 2�C. The values ofpKRed andpKOx caneasily be measured from the pH dependence.Line is best fit to eq.(6), with n = 1.Amicyanin is a type I blue copper proteinwhich exhibits a CuII /CuI redox process.From ref. [1].

Figure 4: Pourbaix-like diagram for thechange in the reduction potential of the flavincofactor of E. colifumarate reductase as afunction of the succinate concentration(increasing from right to left) at 20�C, pH 7.Line is best fit to eq.(6), with n = 2, but with[succinate] instead of [H+]. From ref. [2].

Redox-linked protonations are conveniently represented by a Pourbaix diagram,6 aplot ofE00

as a function of pH, as schematised in fig. 3A and illustrated with real datain fig. 3B.

These thermodynamic considerations apply for any ligand which binds a redox cen-ter: provided that the dissociation constants from the reduced and oxidised forms aredifferent, the reduction potential depends on the concentration of ligand, and the disso-ciation constants can be measured from the concentration dependence of the reductionpotential.

6Marcel Pourbaix (1904-1998). Marcel Pourbaix was born in Myshega (Russia). He studied in Brusselsand graduated from the Faculty of Applied Sciences of the Universit´e Libre de Bruxelles in 1927, with whichhe was associated from 1937 for the rest of his career.

By 1938, he had devised the potential-pH diagrams for which he became famous. In 1939, just beforethe outbreak of World War II, he presented to the Faculty his doctoral dissertation, accompanied by a thesisentitled “Thermodynamics of Dilute Aqueous Solutions. Graphical Representation of the Role of pH andPotential”. The war and some confusion among the jury on the sign of electrode potential impeded thecompletion of the graduation process and, on the suggestion of Prof F.E.C. Scheffer (Delft), the thesis waspresented to the Technical University Delft. This Delft thesis had a major influence on corrosion science.During the fifties and early sixties, Marcel Pourbaix and his collaborators produced potential-pH diagramsfor all the elements and published the “Atlas of Electrochemical Equilibria” [3] in French in 1963 and inEnglish in 1966.

6

Mediator E0 (mV vsSHE)

Benzyl viologen -360Lapachol -172Methylene blue 11Ferricyanide (K3Fe(CN)6) 360Hexachloroiridate (Na2IrCl6) 870

Table 1: Standardreduction potentialsat pH 7, 25�C, ofcommon mediators.

As an example, fig. 4 illustrates the succinate-concentration-dependence of the re-duction potential of the flavin inE. coli fumarate reductase. In this case, the reductionpotential decreases as the concentration of succinate is raised (from right to left in thisfigure) because binding of succinate to the oxidised form of the enzyme is tighter thanto the reduced form (pKOx > pKRed).

Seee.g. refs. [4–7] for the effect of exogenous thiolate or metal binding on thereduction potential of Fe-S clusters.

2 Mediated potentiometry

The reduction potential of a redox couple is determined by recording the ratio[Ox]/[Red] observed after allowing the system to equilibrate with the electrode at dif-ferent potentials.

The potential is usually varied by adding titrants of an oxidant or a reductant. Un-like small molecules, protein redox centers do not generally react rapidly with the mea-suring electrode and equilibrium is not established quickly. To overcome this problem,small redox agents called mediators are added to the solution to transport electronsbetween the active site and the electrode. For best results, these should have reduc-tion potentials close to that of the active site being studied; mixtures of mediators areoften employed to cover a wide range. (A short list of mediators and their reductionpotentials is given in table 1.)

The ratios [Ox]/[Red] are typically determined by examination of the optical orEPR spectra. According to the Nernst equation (3a), a plot ofE versus log[Ox]/[Red]should be a linear, with a slope2:3RT=nF . The electrode potential equalsE00

when[Ox]/[Red]= 1.

Very often, the concentration of Ox or Red is measured as a function of the elec-trode potential, and the plot is fitted with:

[Red] /1

1 + exp�nF

RT(E �E

00)� (7a)

[Ox] /exp

�nF

RT(E �E

00

)�

1 + exp�nF

RT(E �E

00

)� (7b)

7

Figure 5: Potentiometric titration ofPseudomonas aeruginosaHis117Gly azurin, atype-I blue copper protein (0.1mM protein,20mM MES, pH 6, 33mM Na2SO4, 1M NaCl,20�C.) The different symbols correspond tostep-wise reduction and step-wise reoxidation.Plain line is best fit with eq. 7b withn = 1.The adsorption at 638nm is characteristic ofthe oxidised (blue) form of the copper center.Mediators: K3Fe(CN)6, Na2IrCl6,1,2-ferrocene dicarboxyl acid. From ref. [8].

This is illustrated in fig. 5.

3 Introduction to direct methods

3.1 Electrodes for electron transfer to/from proteins

In an increasing number of cases, direct electrochemistry of redox proteins (withoutthe need for mediators) is possible, and this opens the way for detailed studies not onlyof the thermodynamics but also the kinetics of reactions. To be successful, electrodesmust exchange electrons quickly with the proteins, and preserve their native proper-ties. These electrodes may resemble natural environments or reaction partners for theprotein.

Interactions between protein and electrode may be tailored to be weak or strong.Weak interactions might ideally give rise to diffusion-controlled voltammetry (sec-tion 4), whereas with strong interactions, the experiment may address just a smallsample (“film”) of protein molecules on the electrode (section 5).

The various types of electrode surfaces for which protein voltammetry is commonlyobserved are:� Metal (Au, Pt, Ag) surfaces on which a monolayer of adsorbate is self assembled(“Self Assembled Monolayers” or SAMs). The adsorbate is a bifunctional molecule ofthe type X-Y, where X is a substituent that anchors the molecule on the metal electrodesurface (e.g. a thiol) and Y is a functional group that interacts with the protein (typi-cally carboxyl for cytochromesc, or amino for acidic proteins such as plastocyanin orferredoxins.)� Various electrode materials onto which a surfactant film or layers of polyions in-corporate large membrane-bound proteins. This has been successful with bacterialphotosynthetic reaction centers [9,10].� Pyrolytic graphite edge or basal plane electrodes provide hydrophylic or hydrophobicinteractions, respectively. The former is often used with co-adsorbates (aminocyclitols,

8

Figure 6: Electrochemical set up.

polymixyn, polylysine) which probably form cross-linkages between the protein andthe electrode surface.

3.2 Electrochemical equipment

The experiment is carried out using an electrochemical analyser in conjunction with thecell. The cell consists of three electrodes. The reference electrode is often containedin a side arm linked to the main compartment by a capillary tip called a Luggin (afterthe glassblower who invented it). The tip is positioned close to the working electrode.To avoid passing current through the reference electrode,7 a third electrode, calledauxiliary or counter electrode, is used. The working electrode can also be rotated(section 4.2) to control mass transport of solution species. See fig. 6.

The analyser measures the current registered in response to the potential that is ap-plied. In general, the potential of the working electrode (versus the reference electrode)is modulated (e.g. in a linear sweep) and the current flowing between the workingelectrode and the counter electrode is recorded. Since the electrode potential is sweptforward and back, the technique is called “cyclic voltammetry”.8 A voltammogram isa plot of the current as a function of the electrode potential.

7This would change the potential and also damage the electrode.8There are many alternative ways of changing the potential with time to generate and record a cur-

rent [11].

9

3.3 Vocab and conventions

A cathodic process is a reduction, the cathode is the electrode onto which the reductionoccurs. A anodic process is an oxidation,: : : occuring at the anode.

Usually, a cathodic current is counted as negative and an oxidation results in apositive current. American people and software often use the opposite convention forthe sign of the current (See e.g. fig. 8).

For the following descriptions, we will assume that electron transfer is fast, inwhich case the redox reaction is said to be reversible, and the Nernst equation canbe used despite the fact that the system is not at equilibrium.

4 Bulk Voltammetry

4.1 Bulk voltammetry at stationary electrodes

Imagine the case of a solution containing only the reduced form of an electroactivespecies. The potential is swept linearly in time, as shown in fig. 7A, starting from alow potential. The current response as a function of time is plotted in fig. 7B, and thecyclic voltammogram (current against potential) in fig. 7C.

While the potential is lower thanE00

, no oxidation occurs and no current is mea-sured [see (a) in figs. 7]. When the electrode potential approachesE

00

, Red starts beingoxidised into Ox, giving electrons to the electrode. This is measured as a (positive) cur-rent which increases as more and more species are oxidised (b). However, the electrodeoxidises only species close to it, and the interface is soon depleted. The current thengoes through a maximum before decreasing (c). It doesn’t go to zero since diffusionstill brings (slowly) Red from the bulk of the solution.

While a positive current is being measured, Ox produced by the reaction accumu-lates near the electrode and diffuses slowly towards the bulk.

After the scan is reversed, (d), the current is still positive and decreasing: Redspecies are still being oxidised since the electrode potential is above the reduction po-tential. Near the reduction potential, the Ox species which have accumulated are nowreduced and a negative current is observed (e) until the concentration of Red near theinterface drops down (f), and so does the current.

This results in a peak-like response in both directions.9

9The modeling of voltammograms might be complex because the current response depends on the diffu-sion of the species in the solution, and the equations coupling reaction of the species and their diffusion intothe solution are in general difficult to solve.

10

Figure 7: Cyclic voltammetry for aredoxspecies free to diffuse in solution, at astationary electrode. In a typical voltametricexperiment, the electrode potential is sweptlinearly in time (A), and the current recordedas a function of time (B). A convenient andusual way of displaying the results is to plotthe current as a function of the potential (C).The labels (a) to (f) are referred to in the text.“arb. u.” means “arbitrary units”.

Figure 8: Data for diffusion limitedvoltammetry of a redox protein, thecytochromec2 from Rps. palustris.C = 0:2mM, cell volume 0.5mL: 100nM ofprotein. From ref. [12].Mind the axis! Thevoltammogram is plotted with the americanconvention: the high electrode potentials are onthe left, and the current is positive for areduction.

For ann-electron reaction, the separation between cathodic (reduction) and anodic(oxidation) peaks is given by

�Ep = Ep;a �Ep;c = 2:218RT

nF

=57

n

mV at 298K (8)

� The peak current is proportional to the bulk concentration of speciesC, and to thesquare root of the scan rate�:

ip = 2:69 105n3=2AD1=2 � C

p� (9)

This is the Randles-Sevcik equation.A is the electrode surface,D is a diffusion coef-

11

Figure 9: Cyclic voltammetry at arotatingdisc electrode of anelectroactive species insolution. The electrode potential is sweptlinearly in time (A), and the current recordedas a function of time (B). A convenient way ofdisplaying the results is to plot the current as afunction of the potential (C).

ficient. A linear plot ofip againstp� is the criterion used to identify when the redox

species are diffusing from the bulk to the electrode.� The reduction potential is given by the average of the cathodic and anodic peakpotentials.

E00

�Ep;a +Ep;c

2(10)

A system that conforms to these criteria is said to bereversibleanddiffusion con-trolled.

4.2 Bulk voltammetry at rotating electrodes

The peak shape of the diffusion-limited voltammogram at a macro electrode is due tothe depletion of electroactive species near the electrode surface as they are consumedby the redox reaction.

There are many electrochemical techniques in which the electrode moves with re-spect to the solution. In the most popular, the electrode (called a “rotating disc elec-trode”) is rotating along its axis in the solution. This introduces aconvectivemovementof the solution which increases the efficiency of the transport of species from the bulk

12

Figure 10: Cyclic voltammograms ofamicyanin from P. denitrificansat a 3�m goldmicroelectrode.� = 10mV/s. The amicyaninconcentration in 50mM phosphate buffer (pH7)were (a) 130�M, (b) 90�M, (c) 50�M and (d)0�M. Note the sigmoidal shape of thevoltammogram despite the fact that theelectrode is stationary, and the increase inlimiting current when the concentration ofprotein is raised. From ref. [14].

towards the electrode. Because the depletion layer10 cannot spread anymore in thesolution, the current reaches a limiting valueilim at high driving force, fig. 9.� The Levich equation predicts that the limiting current is proportional to the concen-tration of electroactive speciesC and to the square root of the rotation rate!:

ilim = 0:620nFA��1=6s

D2=3 � C

p! (11)

In this equation,�s is the kinematic viscosity of the solution.11

A plot of i�1lim against!�1=2 is called a Koutecky-Levich plot.� The reduction potentialE00

is simply given by the half-wave potentialE1=2, thepotential at which the current reaches half its limiting value.12

E00

= E1=2 (12)

The scan rate does not enter the measurement if it is small because the current dependson the electrode potential but is independent of time: the voltammogram is said to beatsteady state.

4.3 Bulk voltammetry at stationary micro-electrodes

The considerations above apply only for “macro”-electrodes (i.e. when the diameterof the electrode is larger than the typical size of the diffusion layer). With a “micro”-

10The zone of the solution adjacent to the electrode where the concentration of species differ from that inthe bulk.

11The kinematic viscosity is the ratio of the viscosity over the density.E.g. the viscosity of pure water at20oC is 103�Pa�s, and its density� 1g/cm3 ; this gives a kinematic viscosity�s � 10�2cm2/s.

12A Heyrovsky-Ilkovich plot,log10[(ilim� i)=i] againstE, is usually used to analyze sigmoidal polaro-graphicwavesrecorded under steady-statediffusion-limitedconditions [13]. For ann-electron electrochem-ical reaction that is reversible (i.e. the Nernst equation is obeyed at all times), the plot is linear with a slope�nF=2:3RT , and the semi-logarithmic plot intercepts the line for(ilim � i)=i = 1 atE = E00 .

13

Figure 11: Cyclic voltammetry for aredoxspecies adsorbed on an electrode surface.Rotation of the electrode makes no difference.In a typical voltametric experiment, theelectrode potential is swept linearly in time(A), and the current recorded as a function oftime (B). Cyclic voltammogram: plot thecurrent as a function of the potential (C).

electrode, whose typical size is of the order of a few micrometers, the voltammogramhas a sigmoidal shape despite the fact that the transport is diffusion-limited, fig. 10.

Microelectrodes have been used in the context of cell biology; their size make themsuitable to detect electroactive species released by a single cell. See e.g. refs [15,16].

5 Voltammetry of adsorbed proteins: Protein FilmVoltammetry

When the protein sample is immobilised on the electrode surface, diffusion can beeliminated and much greater thermodynamic and kinetic resolution can be obtainedwith extremely small sample quantities.

5.1 “Non-catalytic” (or “stoichiometric”) voltammetry at slowscan rates

Figure 11 shows the shape of a cyclic voltammogram when the redox species is ad-sorbed onto an electrode.

14

Figure 12: Cyclic voltammogram forPseudomonas aeruginosaazurin adsorbed at apyrolytic graphite electrode. The dashed line isthe baseline which has to be subtracted toremove the contribution of the non-faradaiccurrent. Inset is the baseline subtracted current(the faradaic current).0�C, pH 8.5.� =20mV/s.�A �5.5pmol.

Starting again at an electrode potential lower thanE0, the redox centers are fully

reduced [see (a) in fig. 11]. Sweeping the potential towards high values, the proteinstarts being oxidised when the electrode potential approaches the reduction potential(b); a positive current is then measured, which drops down to zero at high potentialwhen all the adsorbed proteins are oxidised (c). On the reverse scan, a reductive (neg-ative) current is observed when the electrode potential matches the reduction potentialof the protein (d) until all the sample is reduced and the current vanishes.� For an ideal, reversible system, the signal consists of symmetrical oxidation andreduction peaks centered at the reduction potentialE

00

.

E00

= Ep;c = Ep;a (13)

� The area under the peak gives the charged passed for that redox couple (nF electronsper mole of adsorbed centers). It should be the same for the oxidative and for thereductive peaks.� The peak current is proportional to the scan rate�, to the surface concentration ofelectroactive species13 � and to the square ofn.

ip =n2F2A��

4RT(14)

A linear plot of ip against� proves that the redox species are adsorbed onto the elec-trode.� The peak width at half height,�, is:

� � 3:53RT

nF

(15)

( 91n

mV at 25�C, or 83n

mV at 0�C.) Therefore cooperative two-electron transfers givesignals with up to four times the height and half the width of one electron transfers;they are therefore more easily distinguished.

In experiments, there is an additional contribution to the current due to the “charg-ing of the electrode”. Figure 12 shows a non-catalytic voltammogram allowing the

13Therefore, the electroactive coverage must be high enough for a current to be observed. Typically, acoverage higher than 5pmol/cm2 will suffice.

15

Figure 13: Cyclic voltammogram forSulfolobus acidocaldarius7Fe ferredoxinadsorbed at a pyrolytic graphite electrode. Thedashed line is the baseline which has to besubtracted to remove the contribution of thenon-faradaic current.0�C, pH 8.5,� = 20mV/s.�A �3.5pmol. [J. McEvoy andF. Armstrong, unpublished results].

Figure 14: Cyclic voltammograms (raw dataout of scale, base-line subtracted anddeconvoluted signals) for E. colifumaratereductase (FrdAB) adsorbed at a pyrolyticgraphite edge electrode. This enzyme contains3 Fe-S clusters and a flavin cofactor.20�C, pH7 & 9, � = 10mV/s. @pH 7, FADox/FADred�50mV vsSHE,[2Fe2S]2+=+ �41mV ,[4Fe4S]2+=+ �307mV , [3Fe4S]+=0 �67mV .�A �0.4pmol. Note the strong pH dependenceof the FAD signal.

determination of the redox potential of the copper site of azurin. The dashed line isthe baseline which has to be subtracted to remove the contribution of the non-faradaiccurrent.

The 7Fe ferredoxin ofSulfolobus acidocaldariuscontains one [3Fe-4S] cluster,with redox transitions[3Fe4S]+=0 and [3Fe4S]0=2� and one [4Fe-4S] cluster with a2+/+ redox transition, fig. 13. When different centers are present in a protein, thedifferent redox transitions appear as multiple peaks, the areas of which reveal the sto-ichiometry of the redox processes.E.g. the area under the low potential[3Fe4S]0=2�

peak is twice as much as that of the high potential[3Fe4S]+=0 peak. This figure illus-trates the fact that 2-electron redox processes give prominent signals.

The fumarate reductase fromE. coli contains 3 Fe-S clusters and one flavin cofac-tor. The signal associated with the flavin is easily distinguished from the three one-electron peaks due to the Fe-S clusters (fig. 14). The reduction potential of the FAD isstrongly pH dependent, as expected for a reduction process coupled to protonation.

Advantages of PFV versus equilibrium titration & bulk voltammetry. Voltam-metry is now becoming a routine technique to measure reduction potentials, and offersseveral advantages over bulk titrations.

� No need for a spectroscopic “handle”.There is no requirement for a distinctive

16

and unambiguous change in some spectroscopic parameter. The counterpart of thisadvantage is obvious: voltammetry provides no structural information.

� Sample economy.A few pmol of protein/enzyme are adsorbed onto the electrode(although making a film requires a larger amount than that, typically ten to one hundredtimes as much, this is still much less than required for bulk titration)

�Quick measurements.Recording a cyclic voltammogram takes usually a few secondsto a few minutes.

� Instantaneous dialysis.The electrode can be transfered in a solution of different com-position/pH and the measurement repeated with the same sample. The protein mighteven survive long enough in a hostile environment for measurements to be performedbefore the electrode is taken back into a more gentle solution.

� Easier analysis & modeling of data.The baseline (the charging current) is easily re-moved, and interpretation of non-catalytic data might not require solving any transport(diffusion/convection) equation.

� In situ measurements.Reduction potentials can be measured as a function of temper-ature [17–20], or even pressure [21,22].

5.2 Fast scan voltammetry: equilibrium and beyond.

The measurement of a reduction potential (an equilibrium property) theoretically re-quires that the system is at equilibrium. This appears to be in contradiction to mea-suring a current: indeed, the flow of electrons when the protein is oxidised or reducedresults from the fact that the system is driven out of equilibrium when the electrodepotential is changed around the (equilibrium) reduction potential.

In practice, this does not matter too much if the scan rate (and therefore the current)is slow enough (typically lower than�10mV/s) that the system isnearlyat equilibrium.

When the scan rate is raised, however, departure from equilibrium can be observed,and a great deal of information about thekineticsof redox processes can be gained bylooking at the scan rate dependence of the voltammograms.

Effect of electron transfer rate. Theoretical voltammograms for an uncoupledone-electron redox process are plotted in fig. 15A. At slow scan rate, the oxidationpeak occurs atE00

. If the scan rate is high, because it takes time for the electrontransfer between the electrode and the redox center to occur, the maximal current isobservedafter the redox potential has been reached, i.e. at higher electrode potential.The same reasoning applied to the reductive process predicts thatEp;c is lower thanE00

.

When measurements are performed over a large range of scan rates, the resultscan be displayed by plottingEp;a andEp;c as a function of the scan rate, on a logscale (fig. 15B). This is called a trumpet plot.14 The more efficient the electron transfer

14Guess why. Do you want a clue? Nobody in our research group is named Trumpet.

17

Figure 15: Effect of scan rate on thevoltammetry of a redox species undergoing aone-electron, no-proton redox process.A: (theoretical) cyclic voltammograms atdifferent scan rates (from 1000V/s to 10mV/s).The difference betweenEp;a andEp;c

increases as the scan rate is raised.B: “Trumpet plot”:Ep;a (filled squares) andEp;c (emty squares) as a function of the log ofthe scan rate. The lower an electron transferrate is, the trumpet plot will shift to low scanrates.

kon = ko�[H+]

Ka

[3Fe-4S]0:H

E00

ko�

[3Fe-4S]0

[3Fe-4S]+

Figure 16: L-shape scheme used to interpretthe fast scan voltammetry of a [3Fe-4S] cluster.

between the electrode and the active site, the higher the scan rate at which the oxidativeand reductive peaks start to separate.15

Effect of coupled reactions. The rate of electron transfer between the electrodeand the active site might not have physiological relevance; however, fast scan voltam-metry can be used to resolve and quantify the chemical processes coupled to electrontransfer. From a biological point of view, proton transfer is certainly the most importantreaction coupled to electron transfer, and determining the mechanism by which light-or redox-driven pumps translocate protons is an exciting challenge.

Several examples of voltammetric studies of coupled reactions, involving cy-tochromes and Fe-S clusters, have recently been reported (see [29] and referencestherein).

As an example, fig. 17 illustrates the voltammetric study of the [3Fe-4S]+=0 one-electron one-proton reaction (fig. 16), for a mutant ofA. vinelandiiferredoxin I. High-resolution crystal structures reveal that the [3Fe-4S] is buried with no access to watermolecules, and that a carboxylate group from an aspartate (D15) is located close to thecluster on the protein surface. The mechanism of proton transfer between the clusterand solvent has been determined using PFV in conjunction with site-directed mutage-nesis [25–28].

The experiments depicted in fig. 17 were performed at pH 5.4, much higher than

15For the effect of distance between the electrode and the redox site on the rate of electron transfer, seee.g.refs. [23,24].

18

Figure 17: Effect of scan rate on the voltammetry of a redox species undergoing aone-electron, one-proton redox process. The data are for the [3Fe-4S]+=0 redoxcouple of a slow proton-transfer mutant of Azodobacter vinelandiiferredoxin I(D15E) [25–27]. pH 5.4, 0�C. From [28].

19

pKOx and 1.3 pH unit lower than thepKRed = 6:7 of the [3Fe-4S] cluster, the reducedform of which is therefore protonated at equilibrium.� At slow scan rates, [label (1) in fig. 17], oxidation and reduction peaks for the [3Fe-4S]+=0 appear at the same electrode potential. Under these “close-to-equilibrium”conditions, the reduction is followed by protonation, and oxidation proceeds along thereverse route.�When the scan rate is increased, oxidation and reduction peaks start to separate.� (2) At scan rates between 1 and 10V/s (the scans were started from the high potentiallimit, and only the first scan is considered), the reductive peak is still clearly visible, butthe oxidation peak vanishes, because the cluster is trapped in the protonated form: it isquickly protonated upon reduction, but the rate of de-protonation,ko� in fig. 16, is toosmall for the cluster to be de-protonated during the oxidative scan. The deprotonationof the cluster gates its reoxidation.16

� (3) At very high scan rates, 20V/s and above, both peaks are observed, but the averagereduction potential is lower than at low scan rate, and matches the alkaline limit: thisis because the scan is reversed before the reduced cluster is protonated.

These experiments allow the determination of the rates of (de)protonation, becausethe time scale of potential modulation can be changed over orders of magnitude (1 minat 10mV/s to 1 ms at 1000V/s) to match that of the chemical events.

5.3 Catalytic voltammetry with adsorbed redox enzymes

In the absence of substrate and at sufficiently high coverage, a redox enzyme immo-bilized onto an electrode may give peak-like signals corresponding to the reversibletransformation of its redox centers (fig. 14). Upon adding substrate, the non-turnoverpeaks are transformed to sizeable waves: reaction with substrate transforms the activesite, which is regenerated by electron exchange with the electrode in a succession ofcatalytic cycles. Current equates directly to rate of turnover, and so the relationship be-tween driving force (potential) and catalytic activity is traced in a single voltammetricexperiment. Note that catalysis may be studied even if coverage is too low to observenon-catalytic signals (this is the case for succinate dehydrogenase, the mitochondialcounterpart of fumarate reductase [30], for DMSO reductase fromE. coli [31], and fornitrate reductase [32]).

As an example, the non-catalytic voltammetry ofE. coli fumarate reductase hasbeen shown in fig. 14.In vivo, this enzyme reduces fumarate to succinate,

C CTT

H

���OOC

TTH

��COO�

+2H++ 2e

� �����! �����

""�OOC

bb COO�bb

16We will try and get an estimate ofko� during the practicals.

20

2e to electrode

succinate

catalysis

fumarate

FADRedFADOx

Figure 18: A schematic catalytic cycleillustrating the flux of electrons from succinateto the electrode via the FAD cofactor offumarate reductase.

Figure 19: (A) Non-catalytic voltammogramobtained for E. colifumarate reductase(FrdAB) adsorbed at a pyrolytic graphite edgeelectrode in the absence of substrate. The rawvoltammogram (outer dash-dot line) is not toscale. Inset: background corrected current(small dots) and deconvoluted data (dashedlines).� = 10mV/s,20�C, pH 7. (B) Catalyticwave showing reversible succinate oxidationand fumarate reduction by adsorbed FrdAB ina solution containing succinate and fumarate.� = 1mV/s,20�C, pH 7,! = 3000rpm. Fromref [2].

but it can also oxidize succinate in the presence of an articial electron partner.If this enzyme is adsorbed onto an electrode, in a solution containing succinate,

and if the electrode potential is high enough, the flavin cofactor is oxidised, giving 2electrons to the electrode, fig. 18.

FADRed ! FADOx + 2e�

The oxidised enzyme can bind succinate, and the oxidation of succinate results in thereduction of the flavin:

FADOx + succinate! FADRed + fumarate

The reduced FAD can be re-oxidised giving electrons to the electrode and so on. Thisresults in a steady-state flux of electrons from succinate in solution to the electrode,via the adsorbed enzyme, which is measured as a steady-state current. This catalyticcurrent is directly proportional to the rate of turnover, and is a direct measure of theactivity of the enzyme.

At low electrode potential, in a solution containing fumarate, a reductive (negative)current is observed that is proportional to the rate of fumarate reduction by the enzyme.

Figure 19B illustrates this bi-directional catalysis byE. coli fumarate reductasewhen the solution contains both fumarate and succinate.

21

Figure 20: Substrate-concentrationdependence of catalytic voltammetry forsuccinate oxidation by E. colifumaratereductase (FrdAB). Raw catalyticvoltammograms (A), and base-line subtracteddata (B). pH 7.5, scan rate� = 1mV/s,electrode rotation rate! = 3000rpm(increasing it further makes no difference),T = 20�C, succinate concentration[S] = 20�M–16mM (no fumarate). Note theincrease in limiting current (at high potential)as the succinate concentration is raised. Fromref. [2].

In an assay of the enzyme activity in solution, the enzyme is oxidised by a high po-tential dye (e.g.ferricyanide). The rate of turnover can be measured as a function of theconcentration of substrate, and the results allow the determination of Michaelis-Mentenparameters (kcat andKm). The same parameters can be determined from voltammet-ric experiments looking at the substrate-concentration dependence of the limiting cur-rent [2,33], see fig. 20,e.g.for succinate oxidation,

ilim = nFA�kcat

1 +Ksuccm

[succinate]

(16)

In practice, this kind of measurement might be far from easy.17

Mass-transport controlled catalytic voltammetry. Figure 21A illustrates thevoltammetry for hydrogen oxidation byAllochromatium vinosum[Ni-Fe] hydrogenase.In this case, the limiting current increases dramatically as the rotation rate is raised be-cause during turnover the concentration of hydrogen near the electrode decreases, theenzyme is able to consume H2 faster than it is brought to the electrode by the convec-tive motion of the solution. The higher the rotation rate, the more efficient the transportof hydrogen from the bulk solution towards the enzyme, and the higher the current. Atinfinite rotation rate, the catalytic current is finite: mass transport is no longer rate lim-iting, and the extrapolated current reveals the intrinsic efficiency of the enzyme. The

17(i) The limiting current is proportional toA�, the total amount of enzyme adsorbed. This can be de-termined (with a very relative accuracy) only when the electrode coverage is high enough for non-catalyticsignals to be measured in the absence of substrate. (ii) The measurement ofKm could be performed withoutknowing the exact electroactive coverage. This requires however that the adsorbed film is stable enough asa function of time for the coverage to be constant when limiting currents are measured with the same film insolutions of different substrate concentrations. (iii) Last, eq. (16) does not take into account mass transportof substrate in solution; this is a sensible approach if there is no depletion of substrate near the electrode.This should be checked for by looking at the rotation rate dependence of the limiting current, or using therotation rate dependence to determineilim(! =1):

22

Figure 21: (A) Influence of electrode rotationrate on the catalytic current measured forhydrogen oxidation byAllochromatium vinosum[Ni-Fe] hydrogenaseat different electrode rotation rates (!, given inunits Revolution Per Minute). (B)Koutecky-Levich plot (eq. 17) showing smallbut non-zero intercept at infinite rotation rate,allowingkcat to be estimated. 0.1 bar H2,� =100mV/s,T = 45�C, pH=6.5. Fromref. [34,35].

Figure 22: Reversible oxidative inactivation ofhydrogen oxidation by A. vinosum[Ni-Fe]hydrogenase. 1 bar H2, � =0.3mV/s,T = 45�C, pH=9. From ref. [38].

Koutecky-Levich plot in fig 21B therefore follows:

1

ilim(!)=

1

nFA�kcat+

constantp!

(17)

This equation handles departure from mass-transport control at high! [compare toeq. (11), page 13].

Interestingly, this study showed that the turnover number of the enzyme is signif-icantly higher than that observed in solution assays, using oxidizing dyes [35, 36]; inthe latter case, it becomes evident that turnover in solution is limited by electron trans-fer from the soluble electron partner, and that the electrode is a much faster electronacceptor than the soluble dye [37].

Potential dependent catalytic activity. The [Ni-Fe] hydrogenase fromA. vinosumis known to deactivate in oxidizing conditions, (e.g. after exposure to oxygen) due tothe over-oxidation of the active site. The enzyme has to be reduced (e.g. incubatedwith hydrogen) for the activity to be recovered. This makes tricky the assay of theenzyme in solution, since it involves keeping the enzyme in very oxidising conditions;the enzyme might therefore inactivate over the time course of the solution assay!

This inactivation was not observed in the electrochemical experiments depicted infig. 21 because the scan rate was fast enough: the enzyme is taken to high potential

23

Figure 23: Catalytic voltammetry of E. coliDMSO reductase (DmsABC): reduction ofDMSO @pH 6.5 and 9.20mM DMSO(100�Km), � = 5mVs�1, 298K, 0rpm. Thismembrane-bound oxidoreductase is involvedduring anaerobic respiration. The active sitefor DMSO reduction contains aMo-bismolybdopterin guanine dinucleotide.The potential window of activity coincideswith the appearance of the Mo(V) EPR signalobserved in potentiometric titrations,suggesting that crucial stages of catalysis arefacilitated while the active site is in theintermediate Mo(V) oxidation state. From [31].

(where the limiting current is measured) and back to low potential before the the inac-tivation reaction can proceed.

Figure 22 is a catalytic voltammogram recorded essentially in the same conditionsas in fig. 21 but for a slow scan rate (recording a whole catalytic voltammogram at thisscan rate takes more than an hour and a half.) The activity of the enzyme, measuredas a positive current, decreases above -200mV. This kind of study therefore allows theenergetics of the oxidative reaction causing the “switch off” to be probed (the potentialof the switch corresponds to that of the redox transition towards an oxidised, inactiveactive site). Moreover, one can design experiments in which a high electrode potentialis held, and the slow decrease in current resulting from the inactivation measured: thisallows the kinetics of the inactivation to be studied, too.

Potential dependent catalytic activity. Ad lib. Surprisingly, among the differentredox enzymes which have been studied by PFV, quite a few of them exhibit complexactivity profiles as a function of the driving force (the electrode potential) (See ref. [39]and references therein).

Succinate dehydrogenase, the membrane-extrinsic domain of Complex II, effi-ciently catalyzes fumarate reduction, but only over a narrow potential range becauseactivity decreases abruptly once a critical driving force is reached,i.e., at sufficientlylow potential [30, 40–43]. This reversible switch (of possible physiological relevance)may arise from a conformational change linked to the oxidation state of FAD. By con-trast, the fumarate reduction activities of several fumarate reductases [33, 44–46] aremarked by complex sigmoidal increases upon successive reductions of the prostheticgroups.

As a last example, fig. 23 shows catalytic voltammograms corresponding to thereduction of dimethyl sulfoxide byE. coli DMSO reductase [31]. At high pH, theactivity of the enzyme drops down below -300mV, but this effect is independent of

24

Figure 24: Effect of electrode rotation on thecatalytic voltammetry of E. colifumaratereductase (FrdAB) in a solution initially free offumarate.[S] =50mM�200�K succinate

m .

scan rate: the switch on and off is fast on the time scale of the experiment. Althoughthe concept of pH optima for enzyme activity is well established, the possibility thatactivity of redox enzymes might be optimised within a finite, and perhaps narrow, rangeof potential has not been explored. While pH optima give mechanistic informationon protonation equilibria during catalysis, potential optima provide information on theroles played by different oxidation states of redox-active sites. Certain redox transitionsmay regulate electron flow to or from the active site. Or substrate binding or atomtransfer may occur only when the active site is in a certain oxidation state. Theserelationships are difficult to observe by conventional techniques but can be revealedby protein film voltammetry, due to its ability to measure subtle changes in catalyticactivity as the electrode potential is varied.

6 PFV Quiz

In fig. 24, I have plotted two voltammograms recorded with a film ofE. coli fu-marate reductase in contact with a solution containing only succinate at a concentration[S] =50mM (approximately 200 times the Michaelis constant). The plain line corre-sponds to a stationary (non-rotating) electrode, whereas the dashed votammogram wasrecorded with the electrode rotating at 3000rpm. (i) Why is the current at high potentialrotation-rate independent in this case? (ii) Why are the shapes of the voltammogramsdifferent?

25

An extensive description of most of the electrochemical techniques will be foundin refs [11,13].

Refs [4, 29, 47] are reviews of the use of voltammetry to probe the energetics anddynamics of redox proteins and enzymes. In refs. [4,48], emphasis is made on ProteinFilm Voltammetry. The mediated protein electron transfer voltammetry of enzymes,and its application to biosensors are discussed in [29]. Ref. [47] goes beyond the scopeof faradaic bioelectrochemistry.

References

[1] Lars J. C. Jeuken, R. Camba, F. A. Armstrong, and G. W. Canters. The pH-dependentredox inactivation of amicyanin fromParacoccus versatusas studied by rapid protein-filmvoltammetry.J. Biol. Inorg. Chem., 2001.

[2] Christophe Leger, Kerensa Heffron, Harsh R. Pershad, Elena Maklashina, C´esar Luna-Chavez, Gary Cecchini, Brian A. C. Ackrell, and Fraser A. Armstrong. Enzyme electroki-netics: energetics of succinate oxidation by fumarate reductase and succinate dehydroge-nase.Biochemistry, 40:11234–11245, 2001.

[3] M. Pourbaix.Atlas d’equilibreselectrochimiques. Gauthier-Villars, 1963.

[4] Fraser A. Armstrong, Hendrik A. Heering, and Judy Hirst. Reactions of complex metallo-proteins studied by protein film voltammetry.Chem. Soc. Rev., 26:169–179, 1997.

[5] J. N. Butt, A. Sucheta, F. A. Armstrong, J. Breton, A. J. Thomson, and E. C. Hatchikian.Voltammetric characterization of rapid and reversible binding of an exogenous thiolateligand at a [4Fe-4S] cluster in ferredoxin-iii from desulfovibrio africanus.J. Am. Chem.Soc., 115(4):1413–1421, 1993.

[6] J. N. Butt, S. E. J. Fawcett, J. Breton, A. J. Thomson, and F. A. Armstrong. Electrochemicalpotential and ph dependences of [3Fe-4S][M3Fe-4S] cluster transformations (M = Fe,Zn, Co, and Cd) in ferredoxin III fromDesulfovibrio africanusand detection of a clusterwith M = Pb. J. Am. Chem. Soc., 119(41):9729–9737, 1997.

[7] S. E. J. Fawcett, D. Davis, J. L. Breton, A. J. Thomson, and F. A. Armstrong. Voltam-metric studies of the reactions of iron-sulphur clusters ([3Fe-4S] or [M3Fe-4S]) formed inPyrococcus furiosusferredoxin.Biochem. J., 335(Pt2):357–368, 1998.

[8] L. J. C. Jeuken, Pieter van Vliet, M. P. Verbeet, R. Camba, J.P. McEvoy, F.A. Arm-strong, and G.W. Canters. Role of the surface-exposed and copper-coordinating histidinein blue copper proteins: the electron-transfer and redox-coupled ligand binding propertiesof His117Gly azurin.J. Am. Chem. Soc., 122(49):12186–12194, 2000.

[9] James F. Rusling. Enzyme bioelectrochemistry in cast biomembrane-like films.Acc. Chem.Res., 31:363–369, 1998.

[10] J. K. Kong, W. Sun, X. Wu, J Deng, Z. Lu, Y. Lvov, R. Z. B. Desamero, H. A. Frank, andJ. F. Rusling. Fast reversible electron transfer for photosynthetic reaction center from WTRhodobacter sphaeroidesre-consituted in polycation sandwiched monolayer film.Bioelec-trochem. Bioenerg. (now Bioelectrochemistry), 48:101–107, 1999.

26

[11] Allen J. Bard and Larry R. Faulkner.Electrochemical methods. Fundamental and applica-tions. John Wiley & Sons, Inc., 2001.

[12] Gianantonio Battistuzzi, Marco Borsari, Marco Sola, and Francesco Francia. Redoxthermodynamics of the native and alkaline forms of eukaryotic and bacterial class I cy-tochromesc. Biochemistry, 36:16247, 1997.

[13] Bond, A. M. (1980)Modern polarographic methods in analytical electrochemistry,MarcelDekker, Inc.

[14] M. Kudera, A. Aitken, L. Jiang, S. Kaneko, H. A. O. Hill, P. J. Dobson, P. A. Leigh,and W. S. McIntire. Electron transfer processes of redox proteins at inherently modifiedmicroelectrode array devices.J. Electroanal. Chem., 495:36–41, 2000.

[15] S. Arbault, P. Pantano, J. A. Jankowski, M. Vuillaume, and C. Amatore. Monitoring anoxidative stress mechanism at a single human fibroblast.Anal. Chem., 67(19):3382–3390,1995.

[16] T. J. Schroeder, R. Borges, J. M. Finnegan, K. Pihel, C. Amatore, and R. M. Wightman.Temporally resolved, independent stages of individual exocyotic secretion events.Biophys.J., 70(2):1061–1068, 1996.

[17] P. L. Hagedoorn, M. C. P. F. Driessen, M. vandenBosch, I. Landa, and W. R. Hagen. Hy-perthermophilic redox chemistry: a re-evaluation.FEBS Lett., 440(3):311–314, 1998.

[18] J. B. Park, C. L. Fan, B. M. Hoffman, and M. W. W. Adams. Potentiometric and electronnuclear double-resonance properties of the 2 spin forms of the [4fe-4s]+ cluster in thenovel ferredoxin from the hyperthermophilic archaebacteriumPyrococcus furiosus. J. Biol.Chem., 266(29):19351–19356, 1991.

[19] P. S. Brereton, M. F. J. M. Verhagen, Z. H. Zhou, and M. W. W. Adams. Effect of iron-sulfurcluster environment in modulating the thermodynamic properties and biological functionof ferredoxin fromPyrococcus furiosus. Biochemistry, 37(20):7351–7362, 1998.

[20] J. P. McEvoy and F. A. Armstrong. Protein film cryovoltammetry - demonstration with a7Fe ferredoxin.Chem. Comm., pages 1635–1636, 1999.

[21] M. T. Cruanes, K. K. Rodgers, and S. G. Sligar.J. Am. Chem. Soc., 114:9660, 1992.

[22] L. D. Gilles de Pelichy and E. T. Smith. Redox properties of mesophilic and hyperther-mophilic rubredoxins as a function of pressure and temperature.Biochemistry, 38:7874,1999.

[23] Z. Q. Feng, S. Imabayashi, T. Kakiuchi, and K. Niki. Electroreflectance spectroscopicstudy of the electron transfer rate of cytochrome c electrostatically immobilised on the!-carboxyl alkanethiol monolayer modified gold electrode.J. Electroanal. Chem., 394:149–154, 1995.

[24] Q. Chi, J. Zhang, J. E. T. Anderson, and J. Ulstrup. Ordered assembly and controlledelectron transfer of the blue copper protein azurin at gold (111) single-crystal substrates.J.Phys. Chem. B, 105(20):4669–4679, 2001.

[25] J. Hirst, J. L. C. Duff, G. N. L. Jameson, M. A. Kemper, B. K. Burgess, and F. A. Arm-strong. Kinetic and mechanism of redox-coupled, long-range proton transfer in an fe-sprotein. investigation by fast-scan voltammetry.J. Am. Chem. Soc., 120(28):7085–7094,1998.

27

[26] K. Chen, J. Hirst, R. Camba, C. A. Bonagura, C. D. Stout, B. K. Burgess, and F. A. Arm-strong. Atomically defined mechanism for electron-coupled proton transfer to a buriedredox centre in a protein.Nature, 405:814–817, 2000.

[27] R. CambaAcosta, J. Hirst, A. Norris, K. S. Cheng, B. K. Burgess, and F. A. Armstrong.Detailed studies of the mechanism of redox coupled proton transfer in proteins.J. Inorg.Bioch., 74(0162-0134):87, 1999.

[28] Raul Camba.Reaction mechanism of iron-sulfur proteins studied by protein film voltam-metry. PhD thesis, Oxford University, 2000.

[29] Fraser A. Armstrong and George S. Wilson. Recent developments in faradaic bioelectro-chemistry.Electrochim. Acta, 45:2623–2645, 2000.

[30] Judy Hirst, Arthur Sucheta, Brian A. C. Ackrell, and Fraser A. Armstrong. Electrocatalyticvoltammetry of succinate dehydrogenase: direct quantification of the catalytic propertiesof a complex electron-transport enzyme.J. Am. Chem. Soc., 118(21):5031–5038, 1996.

[31] K. Heffron, C. Leger, R. A. Rothery, J. H. Weiner, and F. A. Armstrong. Determination ofan optimal potential window for catalysis byEscherichia colidimethyl sulfoxide reductase,and hypothesis on the role of MoV in the reaction pathway.Biochemistry, 40(10):3117–3126, 2001.

[32] Lee J. Anderson, David J. Richardson, and Julea N. Butt. Catalytic protein film voltam-metry from a respiratory nitrate reductase provides evidence for complex electrochemicalmodulation of enzyme activity.Biochemistry, 40(38):11294–11307, 2001.

[33] Arthur Sucheta, Richard Cammack, Joel Weiner, and Fraser A. Armstrong. Reversible elec-trochemistry of fumarate reductase immobilized on an electrode surface. direct voltammet-ric observations of redox centers and their participation in rapid catalytic electron transport.Biochemistry, 32(20):5455–5465, 1993.

[34] H. R. Pershad, J. L. C. Duff, H. A. Heering, E. C. Duin, S. P. J. Albracht, and F. A. Arm-strong. A voltammetric appraisal of rapid catalytic electron transport inChromatium vi-nosum[NiFe]-hydrogenase.J. Inorg. Biochem., 74(4):264–264, 1999.

[35] Harsh R. Pershad, Jillian L. C. Duff, Hendrik A. Heering, Evert C. Duin, Simon P. J.Albracht, and Fraser A. Armstrong. Catalytic electron transport inChromatium vinosum[Ni-Fe]-hydrogenase: application of voltammetry in detecting redox-active centers andestablishing that hydrogen oxidation is very fast even at potentials close to the reversibleH+/H2 value.Biochemistry, 38(28):8992–8999, 1999.

[36] Anne K. Jones, Emma Sillery, Simon P. J. Albracht, and Fraser A. Armstrong. Electrocat-alytic oxidation of hydrogen by an enzyme at rates matching a platinum catalytst.Chem.Com., in press, 2001.

[37] Patrick Bertrand, Francois Dole, Marcel Asso, and Bruno Guigliarelli. Is there a rate-limiting step in the catalytic cycle of [Ni-Fe] hydrogenases?J. Biol. Inorg. Chem.,5(6):682–691, 2000.

[38] A. K. Jones, H. R. Pershad, S. P. J. Albracht, and F. A. Armstrong. Switching between ac-tive and inactive states of [NiFe]-hydrogenases: Investigation of the kinetics and potentialsby protein film voltammetry. In preparation.

28

[39] S. J. Elliott, C. Leger, H. R. Pershad, J. Hirst, K. Heffron, F. Blasco, R. Rothery, J. Weiner,and F. A. Armstrong. Detection and interpretation of redox potential optima in the catalyticactivity of enzymes.Biochim. Biophys. Acta, special issue EBEC 2002, 2002.

[40] Arthur Sucheta, Brian A. C. Ackrell, Bruce Cochran, and Fraser A. Armstrong. Diode-likebehaviour of a mitochondrial electron-transport enzyme.Nature, 356:361–362, 1992.

[41] Judy Hirst, Brian A. C. Ackrell, and Fraser A. Armstrong. Global observation of hydro-gen/deuterium isotope effect on bidirectional catalytic electron transport in an enzyme:direct measurement by protein-film voltammetry.J. Am. Chem. Soc., 119(32):7434–7439,1997.

[42] Harsh R. Pershad, Judy Hirst, Bruce Cochran, Brian A. C. Ackrell, and Fraser A. Arm-strong. Voltammetric studies of bidirectional catalytic electron transport inEscherichiacoli dehydrogenase: comparison with the enzyme from beef heart mitochondria.Biochim.Biophys. Acta, 1412:262–272, 1999.

[43] Brian A. C. Ackrell, Fraser A. Armstrong, Bruce Cochran, Arthur Sucheta, and Tao Yu.Classification of fumarate reductase and succinate dehydrogenases based upon their con-trasting behaviour in the reduced benzylviologen/fumarate assay.FEBS Lett., 326(1–3):92–94, 1993.

[44] Hendrik A. Heering, Joel H. Weiner, and Fraser A. Armstrong. Direct detection and mea-surement of electron relays in a multicentered enzyme: voltammetry of electrode-surfacefilms of Escherichia colifumarate reductase, an iron-sulfur flavoprotein.J. Am. Chem.Soc., 119(48):11628–11638, 1997.

[45] Julea N. Butt, Jeremy Thornton, David J. Richardson, and Paul S. Dobbin. Voltammetryof a flavocytochromec3: the lowest potential heme modulates fumarate reduction rates.Biophys. J., 78(2):1001–1009, 2000.

[46] Anne K. Jones, Ra´ul Camba, Graeme A. Reid, Stephen K. Chapman, and Fraser A. Arm-strong. Interruption and time resolution of catalysis by a flavoenzyme using fast scanprotein film voltammetry.J. Am. Chem. Soc., 122(27):6494–6495, 2000.

[47] R. Guidelli, G. Aloisi, L. Becucci, A. Dolfi, M. R. Moncelli, and F. T. Buoninsegni. Bio-electrochemistry at the metal/water interface.J. Electroanal. Chem., 504:1–28, 2001.

[48] F. A. Armstrong. Insights from protein film voltammetry into mechanisms of complexbiological electron-transfer reactions.J. Chem. Soc., Dalton trans., pages 661–671, 2002.

29