electron and proton transferring properties of vitamin k1 across a self-assembled phospholipid...

Journal of Electroanalytical Chemistry 445 (1998) 71–80

Electron and proton transferring properties of vitamin K1 across aself-assembled phospholipid monolayer

Roberto Herrero 1, Francesco Tadini Buoninsegni, Lucia Becucci, Maria Rosa Moncelli *

Chemistry Department, Uni6ersity of Florence, Via G. Capponi, 9, 50121-Florence, Italy

Received 3 June 1997

Abstract

The electrochemical behaviour of vitamin K1 (VK1) incorporated in a self-assembled monolayer of dioleoylphosphatidylcholine(DOPC) deposited on a hanging mercury drop electrode was investigated by a computerized chronocoulometric technique. Thekinetics of VK1 electroreduction to the corresponding quinol, VK1H2, and that of VK1H2 reoxidation to VK1 were examined forreactant concentrations ranging from 0.5 to 2 mol% by varying the pH from 5.5 to 9 with phosphate and borate buffers. On thebasis of a general kinetic approach it was concluded that the reduction of VK1 to VK1H2 in a DOPC monolayer takes place viathe reversible uptake of one electron, yielding the semiquinone radical anion VK1�−, followed by the rate determining protonationof the latter. On the other hand, the oxidation of VK1H2 to VK1 takes place via the reversible release of one electron, yieldingthe semiquinone radical cation VK1H2

+ �, followed by the rate determining deprotonation of the latter. The only effective protondonors in VK1 reduction are the protons, whereas the main proton acceptors in VK1H2 oxidation are the water molecules.© 1998 Elsevier Science S.A. All rights reserved.

Keywords: Vitamin K; Membrane model; Phospholipid monolayer; Electron transfer

1. Introduction

K vitamins are present in vegetables, where theyseem to play an active role in the photosyntheticmechanism [1]. They also play a role in cellular res-piration as electron carriers and in oxidative phos-phorylation ([2] and references therein). These proper-ties are ascribable to the reversible character of thisquinone�hydroquinone redox system. These vitaminshave long been recognized as essential constituents forblood coagulation [3,4]. They are effective in preventinghemorrhage in humans afflicted with hypothrombine-mia. It was suggested that the activity of vitamin K isassociated with the maintenance of sufficient concentra-tions of the blood clotting accelerator prothrombin, aprotein playing a main role in clot formation. As far as

is known, prothrombin does not contain vitamin K orother quinones, but its formation in the liver is gov-erned by vitamin K. Thus, vitamin K deficiency isalways associated with diminished plasma prothrombinlevels [4]. There is experimental evidence that not onlyprothrombin, but also the VII, IX and X factors in-volved in clot formation, need vitamin K for theirbiosynthesis. Vitamin K seems to play a major role inthe carboxylation of glutamic acid residues in theseproteins: when vitamin K is lacking, the carboxylationstep does not take place, and the proteins so synthe-sized do not have clotting activity. A second modifica-tion of some clotting proteins has also been observed,which depends on the presence of vitamin K and con-sists in the hydroxylation of an aspartic acid residue[5]. In the human body, vitamin K seems to be syn-thesized in the required physiological quantities bymicroorganisms in the intestine [6,7]. K vitamins existin three series, namely the phylloquinone (K1), the

* Corresponding author.1 On leave from University of La Coruna (Spain).

0022-0728/98/$19.00 © 1998 Elsevier Science S.A. All rights reserved.PII S 0 022 -0728 (97 )00537 -8

R. Herrero et al. / Journal of Electroanalytical Chemistry 445 (1998) 71–8072

menaquinone (K2) and the menadione (K3) series. Wehave limited our analysis to the highly lipophilic vita-min K1 (VK1).

In spite of the biological relevance of the electronand proton transferring properties of the vitamin Kgroup, electrochemical investigations of these propertiesare scarce. Polarographic [8–12] and cyclic voltammet-ric [13,14] studies of this class of compounds in aqueousand non aqueous media have been focused mainly onthe reduction of their quinone structure [2,8,15,16].More recently, the electrochemical properties of vita-min K have been investigated on gold and platinumelectrodes modified with cysteamine and cystamine [17]and with n-alkanethiols [18]. The behaviour of vitaminK embedded in soy bean lecithin liposomes [19] hasalso been studied.

This work describes the electrochemical behaviour ofthe liposoluble vitamin K1 (VK1) incorporated in amembrane model consisting of a phospholipid mono-layer deposited on a hanging mercury drop electrode[20–23]. The phospholipid coating is obtained byspreading a solution of the lipid in a suitable solvent(e.g. pentane or hexane) on the surface of an aqueouselectrolyte, allowing the solvent to evaporate, and im-mersing a hanging mercury drop electrode in the elec-trolyte. Self-assembled phospholipid monolayers soobtained, albeit poorer membrane models than themore widely employed BLMs, offer a number of advan-tages over the latter. Thus, the electric potential and theflux of electroreducible species across these monolayerscan be controlled more accurately and more directlythan across bilayer lipid membranes (BLMs). As dis-tinct from usual planar BLMs, these lipid monolayersare completely solvent-free, and hence have a very lowcompressibility coefficient [24], which prevents de-tectable changes in their thickness (and hence in theirdifferential capacity) with a change in the potentialdifference across the film. Moreover, these mercury-supported monolayers provide an inherent mechanicalstability and a resistance to high electric fields that arenot shared by BLMs. One major advantage of lipidmonolayers deposited on mercury over those depositedon solid electrodes is represented by the perfectlysmooth and defect-free support provided by liquid mer-cury to the lipid film. Solid electrodes exhibit a highdensity of surface defects that produce ‘pinholes’ in thelipid film: these may provide direct access of electroac-tive species to the electrode surface thus affecting exper-imental results to a notable extent.

This work is part of a project aimed at investigatingthe electron and proton transferring properties of phys-iological quinones incorporated in biomimetic mem-branes supported by electrodes. A first investigation hasconcerned the reduction mechanism of ubiquinone-10incorporated in a self-assembled monolayer of di-oleoylphosphatidylcholine (DOPC) deposited on mer-

cury [25]. In the present work the behaviour of VK1and of the corresponding quinol, VK1H2, was investi-gated as a function of their concentration in the lipidmonolayer, the solution pH and the buffer concentra-tion. This has permitted conclusions to be drawn aboutthe mechanism of VK1 reduction and VK1H2 oxidationin terms of the rate determining steps and of the natureof the main proton donor and acceptor.

2. Experimental

The self-assembled monolayers of dioleoylphos-phatidylcholine (DOPC) on mercury were prepared asdescribed elsewhere [20–22]. The water used was ob-tained from light mineral water by distilling it once,and then by distilling the resulting water from alkalinepermanganate. Merck reagent grade KCl was baked at500°C before use to remove any organic impurities.DOPC was obtained from Lipid Products (SouthNutfield, Surrey, England). Vitamin K1 (VK1) fromFluka was used without further purification.The oxi-dised form of VK1 is highly lipophilic, as indicated bylog P=3.27, where P is its partition coefficient in thecyclohexane�water system [26]; the reduced form,VK1H2, is likely to be more lipophilic than VK1, assuggested by the P value of hydroquinone being higherthan that of quinone in the octanol�water system [26].VK1 was dissolved in pentane and stored at −20°C.Working solutions of DOPC and VK1 for spreading onthe surface of the aqueous electrolyte in the cell wereprepared every 3 days and stored at −20°C. All mea-surements were carried out in aqueous 0.1 M KCl at25°C. The solution pH was controlled with aNaH2PO4+Na2HPO4 buffer from 5.5 to 7.5 and with aHBO2+NaBO2 buffer from 7.5 to 9. The overall con-centration of the acidic and basic components of thebuffer is referred to as the buffer concentration.

The home-made hanging mercury drop electrode(HMDE) employed in the measurements, the cell andthe detailed procedure to produce self-assembled lipidmonolayers deposited on mercury were described else-where [21,22]. Differential capacity measurements werecarried out using a Metrohm Polarecord E506. The acsignal had a 10 mV amplitude and a 75 Hz frequency.All potentials were measured versus a saturated calomelelectrode (SCE). A chronocoulometric procedure de-scribed elsewhere [27] and making use of a whollycomputerized apparatus [28] was employed to studyelectron transfer to the VK1 molecule across the lipidmonolayer. The microprocessor used to control all theoperations was a Model NOVA 4X from Data General,whereas an Amel Model 551 fast rise potentiostat witha rise time of 0.1 ms was employed for the potentiostaticcontrol of the three-electrode system. The detailedscheme of the home-made electronic current integrator

R. Herrero et al. / Journal of Electroanalytical Chemistry 445 (1998) 71–80 73

Fig. 1. Q(E,t) vs. t curves for 1 mol% VK1 reduction in a 0.01 M borate buffer of pH 7.8, as obtained by stepping the potential from a fixedinitial value Ei= −0.150 V to different final values E. From bottom to top, the E values are −0.300, −0.350, −0.400, −0.450, −0.500,−0.520, −0.540, −0.560, −0.580, −0.600, −0.620, −0.640, −0.660, −0.680, −0.700, −0.720 and −0.740 V.

working under microprocessor control is described else-where [29].

Each chronocoulomogram consisted of a series ofconsecutive potential steps of progressively increasingheight from a fixed initial value Ei, which was set equalto −150 mV/(SCE) for VK1 electroreduction and to−750 mV for VK1H2 electrooxidation, to differentfinal values E ; the latter values ranged from −250 to−800 mV for VK1 electroreduction and from −600 to−100 mV for VK1H2 electrooxidation. Each chrono-coulomogram was recorded on a single lipid coatedmercury drop. The charge Q(t) following each potentialstep Ei�E was recorded against the time t elapsedfrom the instant of the step for 100 ms, after which thepotential was stepped back to Ei where it remained for3 s. During this period, the reaction product of VK1(i.e. VK1H2) or else of VK1H2 (i.e. VK1) was com-pletely reconverted into the corresponding reactant.Thus, an increase in the rest time at Ei beyond 3 s leftthe charge Q(t) practically unaltered. It should benoted that the initial potential Ei cannot be made morepositive than −100 mV nor more negative than −800mV, because outside this potential range the lipidmonolayer becomes permeable to inorganic ions andcan no longer be regarded as a biomimetic membrane.This limits the accessible pH range to pH values ]5.5,because at lower pH values the reoxidation of VK1H2

to VK1 would require an initial potential Ei\−100mV. On the other hand, pH values \9 cannot be

employed, because they cause a gradual deteriorationof VK1.

The differential capacity C of the DOPC monolayerincorporating VK1 was constantly measured againstthe applied potential both before and after the record-ing of each chronocoulomogram, in order to assess thegood quality of the film. Each series of chronocoulo-metric measurements at different buffer concentrationsand constant pH or at different pH values and constantconcentration of the acidic component of the bufferwas carried out by using a newly formed drop for eachchronocoulomogram.

The charge density sM on the lipid coated mercurydrop at constant applied potential was estimated bycontracting the drop surface by an accurately measuredamount while keeping its neck in contact with the(DOPC+VK1) layer on the surface of the aqueouselectrolyte, and by recording the charge flowing alongthe external circuit as a consequence of this contrac-tion. By so doing, the lipid monolayer that coats thedrop surface remains in equilibrium with the ‘lipidreservoir’ on the solution surface, and hence retains itsproperties upon contraction of the drop surface. Conse-quently, the charge flowing along the external circuit asa consequence of this contraction provides a measure ofthe charge density sM on the mercury surface, with anaccuracy of 0.02 mC cm−2, once it is divided by thechange in drop area [30].


Fig. 2. Q(E,t) vs. t curves for 1 mol% VK1H2 oxidation in a 0.01 M borate buffer of pH 7.8, as obtained by stepping the potential from a fixedinitial value Ei= −0.750 V to different final values E. From bottom to top, the E values are −0.600, −0.550, −0.500, −0.450, −0.400,−0.350, −0.300, −0.275, −0.250, −0.225, −0.200, −0.175, −0.150, −0.125 and −0.100 V.

3. Results

Figs. 1 and 2 show two series of Q(t) vs. t curves fora DOPC monolayer containing 1 mol% VK1, as ob-tained by performing negative potential steps from−150 mV and positive potential steps from −750 mVto different final potentials E. As concerns the electrore-duction of VK1 in Fig. 1, at the less negative E valuesat which VK1 is still electroinactive the charge Q(t)increases abruptly in less than 1 ms, due to the flow ofthe capacitive current that is required to charge theinterphase, and then attains a time independent value.At more negative E values Q(t) increases in time firstabruptly, due to the capacitive contribution, and thenmore slowly, due to the gradual electroreduction ofVK1 in time. With a further gradual shift of E towardsnegative values the electroreduction rate of VK1 in-creases progressively, until ultimately the VK1 incorpo-rated in the lipid monolayer is completely reduced inless than 1 ms one, after which Q(t) attains once againa time independent value. The Q(t) vs. t curves in Fig.2, which refer to the electrooxidation of VK1H2, aresimilar in shape to those in Fig. 1.

Fig. 3 shows a series of curves of Q(t=100 ms) vs. Eat pH 8 for different VK1 concentrations ranging from0.5 to 2 mol%. These curves are sigmoidal in shape,with a rising portion preceded by a sloping foot andfollowed by a sloping plateau. The common slope ofthe foot and of the plateau is a measure of the differen-tial capacity of the lipid monolayer. The faradaic con-

tribution Qf(t) to Q(t) due to the reduction of VK1 isestimated by measuring the charge from the straightline obtained by extrapolation of the foot of the Q(t)vs. E curve (dashed line in Fig. 3). The maximumlimiting value, Qf,max, attained by Qf(t) along theplateau is clearly independent of the electrolysis time tand measures the maximum charge involved in thereduction of VK1 incorporated in the film. The inset ofFig. 3 shows a plot of Qf,max vs. the VK1 concentration.The experimental values are in good agreement with thestraight line calculated for a complete two-electronreduction of VK1 to VK1H2 upon assuming that boththe lipid and the VK1 molecules occupy a surface areaof 65 A.

The curve of the differential capacity C vs. E asprovided by a DOPC film incorporating 1 mol% VK1practically coincides with that of a pure DOPC film; inparticular, the differential capacity along the flat mini-mum in the C vs. E curve assumes the same value ofabout 1.75 mF cm−2 in both cases.

Figs. 4 and 5 are relative to the electroreductionprocess of VK1: the first one shows a series of Q(t) vs.E curves at different electrolysis times t and at constantpH and buffer concentration, whereas the second oneshows curves at different pH values while keeping botht and the concentration of the acidic component of theborate buffer constant. It is apparent that an increase inpH causes a negative shift in the reduction potential ofVK1. Over the whole pH range investigated a change inthe buffer concentration from 5×10−4 to 5×10−2 M


Fig. 3. Q(t=100 ms) vs. E plots for VK1 reduction in a 0.01 M borate buffer of pH 8.0. Numbers on each curve denote the mol% VK1 in theDOPC monolayer. The dashed line was obtained by extrapolating the foot of the Q(t=100 ms) vs. E curves. The inset shows a plot of Qf,max

vs. the VK1 concentration: the straight line was calculated as described in the text.

at constant pH has no appreciable effect on VK1reduction. The corresponding plots relative to VK1H2

electrooxidation are shown in Figs. 6 and 7.The charge density sM on the DOPC coated mercury

drop was measured both in the absence and in thepresence of VK1 at potentials along the foot, the risingportion and the plateau of the Q vs. E curves. At allpotentials investigated the presence of 1 mol% VK1 wasfound to have no effect on the sM value within theaccuracy of our measurements (:0.02 mC cm−2). Asalready pointed out in connection with ubiquinone-10reduction [25], this implies that the electroactive moi-eties of the products of partial and total reduction ofVK1 are either uncharged or else localized somewherein the polar head region of the lipid monolayer.

4. Discussion

The experimental behaviour will be interpreted onthe basis of a general approach outlined in Ref. [25](see also Refs. [31,32]), which holds strictly for a se-quence of elementary electron transfer and heteroge-neous chemical steps involving the variousintermediates as well as any proton donors or accep-tors, with a single rate determining step (rds). Withreference to an electroreduction process, the algebraicsum of all the elementary steps preceding the rds isexpressed by the general equation:

rR+hH+ +n� e−UI (1)

Here R is the reactant (VK1 in the present case), r andh are stoichiometric coefficients, n� is the number ofelectrons exchanged before the rds and I is the interme-diate involved in the rds, which is given the generalform:

iI+h %H+ +de−�products (2)

Here i and h % are the molecularities of the rds withrespect to the intermediate I and to the hydrogen ions,respectively. Eq. (2) includes as particular cases anelectron transfer step, a protonation step or a dimeriza-tion step. In the case of a rate determining protonationstep we must consider that, in view of the principles ofgeneral acid-base catalysis, all proton donors present inthe solution may contribute to the rds of Eq. (2) inaddition to the hydrogen ion.

The experimental strategy consists in carrying out theanalysis of the experimental data while keeping con-stant the concentration GR(t=0) of the reactant incor-porated in the lipid film prior to the electrolysis as wellas the faradaic charge Qf(t) that follows the potentialstep Ei−E and consumes the reactant R. If these tworequirements are satisfied, by straightforward consider-ations it can be readily shown [25] that the relationshipbetween the electrolysis time t and the applied potentialE at constant pH takes this simple form:


Fig. 4. Q(t) vs. E curves for 1 mol% VK1 reduction at different electrolysis times t in a 0.01 M borate buffer of pH 8.0. From right to left, t takesthe values 10, 20, 30, 40, 60, 80 and 100 ms. The inset shows a plot of log t vs. FE/(2.3RT) as obtained at Qf= −0.30 mC cm−2 from the Q(t)vs. E curves in the figure. The straight line in the inset has a slope equal to 0.80.

RTF�( ln t(E

�pH=const

= in� +db (3)

The parameter, b, which comes into play only if the rdsis an electron transfer step (namely if d=1), expressesthe ‘symmetry factor’ for this step and is expected totake values close to 0.5. Moreover, once GR(0) andQf(t) are kept constant, the relationship between theapplied potential E and the pH at constant electrolysistime t is given by

FRT

� (E(pH

�t=const

= −2.303ih

in� +db(4)

provided that the rds is not a protonation step. In viewof the principles of general acid-base catalysis, thepresence of a rate determining protonation step is re-vealed by a change in E when the buffer concentrationis varied while keeping pH, GR(0), Qf and t constant,unless the hydrogen ion itself is the only effectiveproton donor. In the latter case the rate of change of Ewith varying pH is given by:

FRT

� (E(pH

�t=const

= −2.303ih+h %in� +db

(5)

The inset of Fig. 4 shows a plot of log t vs. FE/(2.3RT)over the t range from 10 to 100 ms as obtained forQf= −0.30 mC cm−2 from the Q(t) vs. E curves in thesame figure. This roughly linear plot exhibits a practi-cally unit slope. The same behaviour was observed atall pH values investigated. The plot of −FE/(2.3RT)

vs. pH as obtained for Qf= −0.30 mC cm−2 andt=100 ms from a series of Q vs. E curves in differentHBO2+NaBO2 buffered solutions of constant HBO2

concentration is also approximately linear, with a slopeabout equal to unity (see the inset of Fig. 5).

In view of Eq. (3), the unit slope of the experimentallog t vs. FE/(2.3RT) plots indicates that the quantityin� +db is also equal to unity. Since b is expected to beclose to 0.5, we must conclude that n� = i=1 and d=0,and hence that the rds is a chemical step following thereversible uptake of the first transferring electron. Arate determining chemical step consisting of a dimeriza-tion of an intermediate I resulting from the uptake ofone electron must be excluded, since otherwise wewould have in� +db=2. The slope of the −FE/(2.3RT) vs. pH plot being close to unity can be justifiedeither on the basis of Eq. (4) with h= i=1, or else onthe basis of Eq. (5) with h=0 and h %=1. However, inthe former case the rds would be preceded by both anelectron transfer and a protonation step (see Eq. (1)),and we would be left with a rate determining chemicalstep that can be neither a dimerization step nor aprotonation step, since in the latter case the −FE/(2.3RT) vs. pH plot would have a slope equal to 2. Inview of the difficulty of envisaging another plausiblechemical rds, the most logical conclusion consists inregarding h=0 and h %=1 in Eq. (5). This implies thatthe chemical rds is a protonation step involving theproton as the only effective proton donor. The mecha-


Fig. 5. Q(t=100 ms) vs. E curves for 1 mol% VK1 reduction in borate buffers of constant HBO2 concentration (0.01 M) and different pH values.From left to right, the pH increases from 7.5 to 9.0 by 0.5 increments. The inset shows a plot of −FE/(2.3RT) vs. pH as obtained at Qf= −0.30mC cm−2 from the Q(t= l00 ms) vs. E curves in the figure. The straight line in the inset has unit slope.

nism resulting from the above considerations is there-fore:

VK1+elVK1�−;

VK1�− +H+rds�

VK1H�;

VK1H�+H+ +elVK1H2 (6)

The reduction of VK1 incorporated in a DOPC mono-layer does not depend on the buffer concentration atconstant pH, and hence does not satisfy the principlesof general acid-base catalysis. This behaviour can bejustified by assuming that the protonation takes placewell inside the polar head region of the DOPC mono-layer, which is practically impermeable to the HBO2

molecules on the time scale of our measurements whileit may be permeated by the proton itself. In this casethe role of the buffer is exclusively that of maintainingthe pH just outside the lipid layer constant during VK1reduction, via a dissociation reaction in quasi equi-librium. Hence, the rate determining protonation step isaffected by a change in pH, but not by a change in thebuffer concentration at constant pH.

The general approach leading to Eqs. (3)–(5) relieson the assumption that the elementary steps precedingthe rds are in quasi equilibrium at the location wherethe rds takes place. If we assume that the rate determin-ing protonation step takes place well inside the polarhead region, and hence close to the boundary between

this region and the adjacent hydrocarbon tail region,then we must also assume that any translocation of theelectroactive moiety of the VK1 molecule, namely itsquinone ring, across the hydrocarbon tail region is fastin the time scale of our chronocoulometric measure-ments so as to be regarded as in quasi equilibrium. Thisis consistent with the observation that the translocationof several lipophilic ions across lipid monolayers andbilayers takes place in less than 1 ms [33].

The electroreduction behaviour of VK1 is similar tothat shown by ubiquinone molecules incorporated in aDOPC monolayer supported by mercury [25]. It isreasonable to suppose that the observed sequence ofchemical and electrochemical steps is shared by most ofthe biological quinones involved in the electron trans-port chain in biomembranes. At any rate, some cautionmust be used in transferring these conclusions to a‘native’ environment, since the thermodynamic proper-ties of biological quinones in biomembranes can benotably affected by noncovalent binding interactions toproteins.

An important difference in the electrochemical be-haviour of VK1 and ubiquinone is represented by themuch higher reoxidation rate of VK1H2 with respect tothat of ubiquinol; this allowed us to investigate thekinetics of VK1H2 oxidation on the time scale of ourchronocoulometric measurements, since in the timewindow from 0 to 100 ms the oxidation of VK1H2 toVK1 is practically complete.


Fig. 6. Q(t) vs. E curves for 1 mol% VK1H2 oxidation at different electrolysis times t in a 0.005 M phosphate buffer of pH 7.0. From right toleft, t takes the values 10, 20, 30, 40, 60, 80 and 100 ms. The inset shows a plot of log t vs. FE/(2.3RT) as obtained at Qf=0.11 mC cm−2 fromthe Q(t) vs. E curves in the figure. The straight line in the inset has unit slope.

For the kinetic analysis of VK1H2 electrooxidation,the same approach described for VK1 electroreductionwas employed. In the present case Eq. (1) becomes:

roR+hoA−n� oe−lIo (1b)

where the reactant R is now represented by VK1H2 andA denotes some proton acceptor; here as well as in thefollowing the symbols with a subscript ‘o’ denote quan-tities analogous to those already introduced for VK1reduction, but referring to VK1H2 oxidation.Analogously, Eq. (2) becomes:

ioIo+h %oA−doe−�products (2b)

By considerations entirely analogous to those leading toEqs. (3) and (4), we now obtain:

RTF�( ln t(E

�pH=const

= − [ion� o+do(1−bo)] (3b)

and:

FRT

� (E(pH

�t=const

= −2.303ioho+h %o

ion� o+do(1−bo)(4b)

Here bo is once again the symmetry factor for the rds ifthis is an electron transfer step, namely if do=1. Theinset of Fig. 6 shows a plot of log t vs. FE/(2.3RT) atconstant pH and for a faradaic charge Qf=0.11 mCcm−2. The slope of this plot is about equal to unity.The inset of Fig. 7 shows a plot of −FE/(2.3RT) vs.pH as obtained for Qf=0.30 mC cm−2, while keepingthe concentration of the basic component of the borate

buffer constant. As distinct from the pronounced pHdependence of VK1 reduction, the pH dependence ofVK1H2 oxidation is very weak, although the positiveslope of 0.20 in the plot of the inset in Fig. 7 isdefinitely greater than the uncertainty in the experimen-tal measurements.

In view of Eq. (3b), the unit slope of the experimentalplot of − log t vs. FE/(2.3RT) indicates that the quan-tity n� oio+ (1−bo)do is also equal to unity. This impliesthat n� o= io=1 and do=0, and hence that the rds is achemical step following the reversible release of the firsttransferring electron. In spite of the very weak pHdependence of the kinetics of VK1H2 oxidation, themost probable rds is represented by the deprotonationof the radical cation VK1H2

· + resulting from the re-lease of one electron by VK1H2. Such a very weak pHdependence can be explained by assuming that by farthe major proton acceptor is represented by water,whose concentration inside the polar head region ispractically independent of pH.

The most reasonable mechanism for VK1H2 elec-trooxidation is therefore:

VK1H2lVK1H�+2 +e−;

VK1H�+2 +H2O�VK1H�+H3O+;

VK1H�+H2OlVK1+H3O+ +e− (7)

The oxidation behaviour of VK1H2 differs from that ofubiquinol in that the main proton acceptor in ubiquinoloxidation is the hydroxyl ion [34]. This difference in


Fig. 7. Q(t=100 ms) vs. Ef curves for 1 mol% VK1 oxidation in borate buffers of constant concentration, 0.01 M, of HBO2 and different pHvalues. From left to right the pH varies from 7.5 to 9.0 by 0.5 increments. The inset shows a plot of −FE/(2.3RT) vs. pH as obtained at Qf=0.30mC cm−2 from the chronocoulomograms in the figure. The straight line in the inset has a slope equal to 0.20.

behaviour can be tentatively explained by postulat-ing that the deprotonation of the semiquinone radicalcation VK1H2

�+ of VK1 takes place deeper insidethe polar head region with respect to that of thesemiubiquinone radical cation: the permeability of theneutral water molecules in this region is expected tobe higher than that of the charged hydroxyl ions.The practical insensitivity of VK1H2 oxidation to pHis expected to make it a poor proton carrier acrossbiomembranes. On the other hand, the high rate ofVK1H2 oxidation with respect to that of ubiquinolmay bear some relation to the role played by thisreduced species as an oxidant in the carboxylation ofthe glutamate residues of prothrombin in blood clotting[4].

Acknowledgements

The authors wish to thank Professor RolandoGuidelli for his very useful comments and Mr LucianoRigheschi and Mr Riccardo Collini for valuable techni-cal assistance. Thanks are due to the Ministerio deEducacion y Cultura, Spain, for a fellowship to R.H.and ENEA, Italy, for a PhD fellowship to F.T.-B.during the tenure of which the present results wereobtained. The financial support of the Ministero dell’U-niversita’ e della Ricerca Scientifica e Tecnologica andof the Consiglio Nazionale delle Ricerche is gratefullyacknowledged.

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