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Journal of Electroanalytical Chemistry 499 (2001) 129–135

Reduction of peroxodisulfate on gold(111) covered by surfaceoxides: inhibition and coupling between two oxide reduction

processes

Z. Samec 1, K. Krischer, K. Doblhofer *Fritz-Haber-Institut der Max-Planck-Gesellschaft, Abteilung Physikalische Chemie, Faradayweg 4–6, D-14195, Berlin, Germany

Received 13 July 2000; received in revised form 6 November 2000; accepted 6 November 2000

Abstract

Cyclic voltammetry and potentiostatic current transient measurements with the rotating disc electrode are used to study thereduction of peroxodisulfate (PDS) on Au(111) covered by surface oxides. A strong inhibition is observed, but the PDS reductionrate does not correlate with the oxide surface coverage through a simple geometric blocking factor. Evidence is provided that theinhibition is related to the coverage of Au(111) by the more stable oxide, which blocks the PDS reaction sites located probablyon the surface step edges. The difference between the kinetics of the reduction of the two main oxide species on Au(111) isascribed to a coupling between the two processes. A model of the coupling based on the diffusion controlled proton catalysis ofthe oxide reduction is proposed and verified by numerical simulation. © 2001 Elsevier Science B.V. All rights reserved.

Keywords: Reaction mechanism; Peroxodisulfate reduction; Au(111) electrode; Au oxides

1. Introduction

In previous work we investigated the mechanism ofperoxodisulfate (PDS) reduction on polycrystalline [1]and single crystal Au electrodes [2,3] in acidic solutions.The reaction was shown to proceed via two parallelpathways, i.e. the outer-sphere electron transfer to theS2O8

2− anion in the solution

S2O82− +2e−�2SO4

2− (1)

and the electrocatalytic reduction comprising electrontransfer to the adsorbed anion or to products of itsdissociation in the adsorbed state

S2O82−� (S2O8

2−)ads (2a)

(S2O82−)ads+2e−�2SO4

2− (2b)

The former pathway prevails at potentials negative tothe potential of zero charge (pzc) and, due to a repul-sive (Frumkin) double layer effect [4], the reduction

current decreases with decreasing potential, which givesrise to a negative charge transfer impedance and tooscillatory phenomena [5]. On the contrary, the electro-catalytic reduction dominates at potentials more posi-tive than the pzc, and its kinetics are related closely tothe charge-dependent competitive adsorption of S2O8

2−

and SO42− anions. Since both reduction pathways have

a maximum close to the pzc, their overlap results in apeak-shaped voltammogram that is currently observedon the oxide free gold electrodes [1–3].

The main aim of this work has been to study thepossible involvement of surface oxides. A relevantmechanism comprising the chemical oxidation of themetal surface by PDS followed by electrochemical re-duction of surface oxides has been proposed previouslyfor the polycrystalline Au electrode in alkaline solutions[6]. Since the surface oxides formed on Au(111) andAu(110) electrodes in acidic solutions (pHB2) are re-duced at a potential that is much more positive than thepzc [7], we carried out the electrochemical measure-ments in less acidic solutions (pH 3.8) where a morepronounced effect of the surface oxides can be foreseen,e.g. the oxides formed on Au(111) in neutral solutionsare reduced close to the pzc [8,9]. We show that the

* Corresponding author.E-mail address: doblhofer@fhi-berlin.mpg.de (K. Doblhofer).1 Present address: J. Heyrovsky Institute of Physical Chemistry,

Dolejskova 3, 18223 Prague 8, Czech Republic.

0022-0728/01/$ - see front matter © 2001 Elsevier Science B.V. All rights reserved.PII: S 0022 -0728 (00 )00499 -X

Z. Samec et al. / Journal of Electroanalytical Chemistry 499 (2001) 129–135130

presence of the surface oxides on Au(111) in fact leadsto a strong inhibition of the reaction which, however,does not correlate with the surface oxide coverage in anexpected way. Kinetic analyses of the PDS and oxidereduction processes at a constant potential allow us topropose a novel model of the electrochemical reductionof surface oxides on Au(111) and to identify the reac-tions sites for PDS.

2. Experimental

Electrochemical measurements were carried out on asingle-crystal Au(111) electrode, which was purchasedfrom MaTeck (Julich, Germany). The electrode wasprepared from 5N gold, oriented (B1°), polished (B0.03 mm) and cut in the form of a disc 1 mm thick and9 mm in diameter (geometric area 0.74 cm2). The discwas cleaned in a mixture of sulfuric and nitric acid(1:1), rinsed with and stored in triply distilled water.Prior to each experiment, the surface of the electrodewas flame annealed in a butane+air flame until itshowed a dark red glow for 2 s, cooled for 30 s in air,and then quenched in triply distilled water. Ex situSTM topography of the Au(111) surface after anneal-ing and cycling in a perchloric acid solution have beenreported [3]. Supporting electrolyte solutions were pre-pared from triply distilled water and HClO4, Na2SO4

(Suprapure, Merck) or NaClO4 (p.a., Merck). Small

amounts of HClO4 or NaOH (Suprapure, Merck) wereadded to the solutions of Na2SO4 or NaClO4 to adjusttheir pH. All other experimental conditions, instrumen-tation and procedures have been described previously[2]. Ohmic potential drop compensation was usedthroughout. Electrode potentials are referred to a mer-curosulfate electrode (+0.65 V vs. SHE).

3. Results and discussion

3.1. Inhibition of peroxodisulfate reduction by surfaceoxides

Fig. 1 shows the series of voltammograms of therotating Au(111) electrode in the absence (dashed line)and presence (solid line) of PDS in solutions of variouscomposition and pH. Using the accepted designation[10,11], the peaks of the surface oxidation are labelledOA1/2, OA3 and OA4, and the associated reductionpeaks are labelled OC2 and OC3. On sweeping thepotential from −0.8 V towards positive values, thePDS reduction gives rise to the current peak at apotential close to the pzc (−0.08 V). This peak isobviously associated with the reaction that proceeds onthe oxide-free Au(111) surface. The reduction current ispractically independent of the sweep rate (5–100 mVs−1), i.e. the measured current corresponds to a chargetransfer controlled process. An excess of sulfate, i.e thePDS reduction product, causes only a small decrease inthe reduction rate at potentials more positive than thepzc (panel C), which has been linked to the increasedadsorption of sulfate anion competing with that of PDS[3]. Strong adsorption of both anions also leads to theshift and/or suppression of the oxidation peaks OA1–OA3, an effect that has been observed and discussedpreviously [2,3].

A comparison of voltammograms on panels A and Bof Fig. 1 reveals that the effect of pH on the rate of thePDS reduction on the oxide-free Au(111) is rathernegligible, while the rate of oxide reduction changesconsiderably, cf. the position of OC2 or OC3 peak onthe dashed curve. Fig. 2 shows the dependence of thepeak potential Ep for the reduction of OC3 species onpH. In acidic solutions, the slope dEp/dpH attains thevalue of about −70 mV. At around pH 3, Ep shiftsnegatively by ca. 400 mV to a potential which isnegative to the pzc and becomes practically indepen-dent of pH. Therefore, at pH\3.5 there exists a rangeof potentials at which the reduction of the surfaceoxides is relatively slow, and their effect on the PDSreduction can be easily evaluated.

It was expected that the rate of PDS reduction on theoxidised surface is proportional to the geometric block-ing factor (1−uT) which is determined by the net oxidecoverage uT. However, when the potential sweep is

Fig. 1. Cyclic voltammograms of the rotating Au(111) electrode (1200rpm) in the absence (dashed curve) and presence (solid curve) of 0.4mM Na2S2O8 in various supporting electrolytes: (A) 1 mM HClO4,pH 2.6, (B) 5 mM NaClO4, pH 4.6, and (C) 2.5 mM Na2SO4, pH 4.1.Sweep rate: 100 mV s−1.

Z. Samec et al. / Journal of Electroanalytical Chemistry 499 (2001) 129–135 131

Fig. 2. Effect of the solution pH on the potential Ep of the oxidereduction peak OC3 measured on the rotating Au(111) electrode(1200 rpm) in solutions containing x mM HClO4 or pH adjusted with5 mM NaClO4. Sweep rate: 100 mV s−1.

gated under potentiostatic conditions. In these experi-ments, the cyclic polarisation of the electrode (100 mVs−1) was stopped at E=0.9 V for 30 s to cover theAu(111) surface by oxides, and then the potential wasstepped to E=0.075 V. The current transients follow-ing the potential step have a sigmoidal shape (Fig. 3),i.e. after the initial delay when the PDS reductioncurrent is close to zero, a relatively fast rise to thestationary current value is observed. The transitiontime ta, which can be defined as the time correspondingto one half of the stationary current value, increaseswith the solution pH, but it is independent of the PDSconcentration and it is not affected by the presence ofsulfate (Fig. 4).

These observations suggest that the extended inhibi-tion of the PDS reduction as it emerges from voltam-metric (Fig. 1) or transient (Fig. 3) measurements isrelated exclusively to the kinetics of the oxide reduc-tion. Indeed, the effect of pH on the transition time ta

can be understood as being a consequence of the de-creasing rate of oxide reduction with increasing pH [7].On the other hand, it appears that not all the oxidespecies are involved in the inhibitory effect. Therefore,the current transients displayed in Fig. 3 should becompared with the time decay of various oxides onAu(111) at a constant potential.

3.2. Relation to kinetics of the surface oxide reduction

Kinetics of the oxide reduction were studied by po-tentiostatic transient measurements of the oxide reduc-tion charge in the absence of PDS. Cyclic polarisation(100 mV s−1) of the rotating Au(111) electrode wasstopped at E=0.9 V for 30 s. The potential was then

Fig. 3. Current transients of the peroxodisulfate reduction on therotating Au(111) electrode (1200 rpm) at the potential E=0.075 Vreached stepwise after potentiostatic oxidation of the electrode sur-face at E=0.9 V for 30 s. Solution composition: 0.4 mM Na2S2O8+5 mM NaClO4 with pH adjusted to 4.3 (1), 4.6 (2), 7 (3) and 7.5 (4)

Fig. 4. Effect of the solution pH on the transition times ta for theOC2 oxide reduction (full points) and peroxodisulfate reduction(open points) at 1200 rpm. Solution composition: 0.4 mMNa2S2O8+5 mM NaClO4 (�), 1.6 mM Na2S2O8+5 mM NaClO4

( ), 0.4 mM Na2S2O8+2.5 mM Na2SO4 (�) and 5 mM NaClO4

().

reversed (or the sweep towards negative potentials isinitiated) at a potential corresponding to the oxidisedAu(111) surface, e.g. at ca. 0.85 V in 1 mM HClO4, thePDS reduction proceeds at a negligible rate until themajor part of the surface oxides is removed (panelsA–C). In order to establish clearly the relation betweenthe reduction rate and time, the inhibition was investi-

Z. Samec et al. / Journal of Electroanalytical Chemistry 499 (2001) 129–135132

Fig. 5. Cyclic voltammogram of the rotating Au(111) electrode (1200rpm) in the solution containing 5 mM NaClO4, pH 5.0 and the oxidereduction current during the sequence of potential sweeps fromE=0.075 V to E= −0.78 V following the potential step fromE=0.9 V (30 s) to E=0.075 V with various delay times prior to thesweep initiation. Inset: resolution of the reduction peaks OC2 andOC3 for the delay time of 30 s.

sured current to two overlapping peak functions of thetype

I=4Ipexp(− fDE)

(1+exp(− fDE))2 (3)

where DE=E−Ep, the fitting parameters Ep, Ip and frepresent the peak potential, peak current and thecoefficient F/RT, respectively. Results of the fitting areillustrated by the inset in Fig. 5. Fig. 6 shows thedependences of the reduction charges Qred for the peaksOC3 and OC2, as obtained by integration of the tworesolved peaks, on the delay time at E=0.075 V. Thesum of the maximum OC2 and OC3 reduction charge inacidic solution, i.e. ca. 570 mC cm−2, agrees with thecharge under the peaks OA1–OA4 in panel A of Fig. 1,i.e. 540 mC cm−2, which corresponds to the roughnessfactor of 1.2, on assuming that the charge required foran ideal Au(111) surface is 444 mC cm−2 [11].

Apparently, the coverage of the Au(111) surface bythe OC3 oxide species follows an exponential decay,with the initial reduction charge of ca. 500–650 mCcm−2. On the other hand, the dependence of the surfacecoverage by the OC2 oxide species on the delay time hasa sigmoidal shape, i.e. the initial reduction charge of ca.70 mC cm−2 remains constant until the major portionthe OC3 oxide species is removed. Time evolution of thesurface released by the OC2 oxide species clearly corre-lates with the PDS current transients displayed in Fig. 3.Indeed, the transition time ta for reduction of the OC2oxide species, i.e. the time corresponding to u2=Qred/Q red

max=0.5, correlates with the transition time ta forPDS reduction, cf. the full points in Fig. 4. We concludethat the rate of PDS reduction is proportional to thefraction (1−u2) of the Au(111) surface that has beenoccupied by the OC2 oxide species.

changed stepwise to E=0.075 V and, after a defineddelay, swept from 0.075 V to −0.78 V at 100 mV s−1.Fig. 5 shows the cyclic voltammogram and the currentversus potential profile of the surface oxide reductioncorresponding to various delay times at E=0.075. Thetime evolution of this profile indicates that the contribu-tions of the reduction peaks OC2 and OC3 to the netreduction current vary differently with the delay time.These contributions can be resolved by fitting the mea-

Fig. 6. Charge Qred under the oxide reduction peaks OC3 (A) and OC2 (B) as a function of the delay time t at E=0.075 V at pH 4.3 (�), 5.0( ) and 8.3 (�). For other conditions see the legend to Fig. 5.

Z. Samec et al. / Journal of Electroanalytical Chemistry 499 (2001) 129–135 133

Fig. 7. Charge Qred under the oxide reduction peaks OC3 (A) and OC2 (B) as a function of the delay time t at E=0.075 V at pH 5.0 and therotation speed 3600 rpm (�), 1200 rpm ( ) and 300 rpm (�). For other conditions see the legend to Fig. 5.

3.3. Coupling between the oxide reduction processes

A remarkable difference between the kinetics of thereduction of the OC2 and OC3 oxide species points toa coupling between the two processes, which is respon-sible for the delay in the reduction of the OC2 oxidespecies, during which the OC3 species decay with time.We have developed a model of such a coupling that isbased on proton assisted oxide reduction.

The basic assumption of the model is that the ratedetermining step in the reduction of both OC2 andOC3 oxide species in acidic solution is electrochemicaldesorption [10]

AuOH+e−+H+�Au +H2O (4)

The rate of the two reduction processes can then bedescribed by the kinetic equations

−du2

dt=k2(E)u2cH+(t) (5)

−du3

dt=k3(E)u3cH+(t) (6)

where cH+ is the time dependent proton concentrationat the electrode surface, k2 and k3 are the potentialdependent rate constants, ui represents the relative sur-face oxide coverage, which can be expressed as the ratioof the reduction charge Qred to maximum reductioncharge Q red

max. When the bulk proton concentration cH+b

is high, the rate of proton transport from the solutionto the electrode surface can be sufficient to maintain theproton concentration at the electrode close to the bulkvalue, and the time decay of both the OC2 and OC3species at a constant potential E can follow an expo-nential law, i.e. ui=exp(−kicH+

b t). On the other hand,when the bulk proton concentration is low, then de-

pending on the ratio of the rate constants k2/k3, theprotons at the surface are consumed mainly by thefaster process. Since this ratio is less than unity, thereshould be a lack of protons for the reduction of theOC2 species, which at high pH becomes controlled bythe much slower reaction

AuOH+e−�Au+OH− (7)

as can also be seen in the position of the voltammetricpeaks (Fig. 2). Experimental evidence for the diffusioncontrol of H+ transport is provided by kinetic measure-ments at various rotation speeds. As can be seen fromFig. 7, the characteristic time constants for both decayprocesses increase with decreasing rotation speed. Thisbehaviour is rationalised for the reduction of the OC2species in Fig. 8, which shows that the plot of thetransition time ta versus the inverse square root of therotation speed f is linear, cf. full points in Fig. 8. It is tobe noted that an analogous dependence can be inferredfrom the PDS current transients measured at variousrotation speeds, cf. the open points in Fig. 8.

This model was verified by solving numerically Eqs.(5) and (6) together with the equation describing thechange of the proton concentration at the electrodesurface in the limit of the linear diffusion layer approx-imation [12]

dcH+

dt=

2Sd

duT

dt+

2DH+

d2 (cH+b −cH+) (8)

where uT=u2+u3, S is the scaling factor (2×10−9

mol cm−2), DH+ is the proton diffusion coefficient andd is the thickness of the diffusion layer. The calculatedtemporal evolutions of u2, u3 and cH+ for three differ-ent rotation rates at pH 5 are reproduced in Fig. 9 andcan be compared directly with the experimental resultsof Fig. 7. Obviously, the model captures correctly the

Z. Samec et al. / Journal of Electroanalytical Chemistry 499 (2001) 129–135134

Fig. 8. Effect of the rotation speed f on the transition time ta for theOC2 oxide reduction on Au(111) in 5 mM NaClO4, pH 5.0 () andperoxodisulfate reduction on Au(111) in 0.4 mM Na2S2O8+5 mMNaClO4, pH 6.9 (�).

experiments and which manifests itself in the lineardecrease of u3. Furthermore, from Fig. 9 it becomesapparent that the sigmoidal shape of u2(t) is a conse-quence of the sigmoidal increase of cH+(t) which sets inwhen OC3 is nearly completely reduced.

Quantitatively, the calculated transition times ta areonly about a third of the measured ones. The reason forthis is at the moment unclear as ta depends only on theknown parameters determining the mass transfer, i.e.DH+, d and the bulk pH, but it does not depend on thetwo rate constants, which are difficult to deduce fromthe experiments. Furthermore, the model predicts alinear relation between transition time and proton con-centration. In the experiments a much weaker thanlinear dependence of ta on pH is observed (Fig. 6). Thereason for this is the dominance of Eq. (7) which is nottaken into account in the simulations. Despite thesequantitative deviations of ta between experiments andsimulations, the calculations support strongly the inter-pretation that the reduction kinetics of the two oxidespecies OC3 and OC2 are coupled through the protonconcentration at the electrode.

3.4. Nature of the reaction sites for PDS

A serious implication of the present analysis is thatPDS is reduced only on the surface sites that can beoccupied by the OC2 oxide species. Unfortunately, thesignificance of the OC2 reduction peak is rather unclear[13]. The charge under the OC2 reduction peak reachesthe saturation level of ca. 50 mC cm−2 and corresponds

delayed onset and the sigmoidal decrease of the reduc-tion current of OC2 with time. Generally, this be-haviour is found in the mass transfer limited region ofthe reduction of OC3, which is in accordance with the

Fig. 9. Calculated temporal evolution of the variables according to Eqs. (5), (6) and (8); parameter values: k3=3×105 s−1, k2=0.1×k3, pH 5,D=10−4 cm2 s−1, d=18 mm (a), 30 mm (b) and 60 mm (c) (corresponding to 3600, 1200 and 300 rpm). The coverages (solid lines) are scaledto the initial (maximum) coverages (u3(t=0)=7.14u2(t=0)) and cH+ (dashed lines) to the bulk proton concentration.

Z. Samec et al. / Journal of Electroanalytical Chemistry 499 (2001) 129–135 135

to 25% of the surface sites on the basis of 1 e−/sitestoichiometry [14]. The resolution of the OC2 reductionpeaks in Fig. 5 yields a comparable value of ca. 70 mCcm−2. The more stable OC2 oxide species are formedprior to OC3 oxide species, and are supposed to berepresented by a sublattice of adsorbed OH species onthe anion-free surface [10]. The OA1/2 and OC2 peaksare only resolved in the perchlorate solutions, not insulfate or hydrogensulfate and, besides, hydrogensul-fate added after film formation in dilute perchloric acidalso leads to the absence of the resolved OC2 state [13].As has been noted earlier [11], the effect of sulfatesdepends on their concentration, and at low sulfateconcentrations the OC2 peak can still be seen, cf. panelC in Fig. 1. The absence of the OC2 peak was ascribedto strong anion re-adsorption as free metal sites aregenerated upon oxide film reduction [13]. However, asomewhat different interpretation of the initial stages ofsurface oxidation has been provided in the STM/AFMstudy of Au(111) surfaces in sulfuric acid solutions [15].The pronounced dependence of the OA1–OA3 peakson the defects in the surface topography has led theseauthors to conclude that the OA1–OA3 peaks are notassociated with the electrosorption of OH species in thesulfate adlayer [10,11], but rather with AuOH turnoverprocesses at step edges. Evidence indicating the signifi-cance of the initial adsorption of OH species on thesurface steps has been put forward earlier [16,17]. Con-sequently, the OC2 peak can be associated with thereduction of the turned-over oxide species on the stepedges, which thus should represent also the preferablereaction sites for PDS reduction.

4. Conclusions

The presence of surface oxides on Au(111) leads to astrong inhibition of PDS reduction. However, the re-duction rate is not proportional to the geometric block-ing factor (1−uT) which is determined by the totaloxide coverage uT. Experimental data obtained from

potentiostatic transient measurements of the oxide re-duction charge suggest that the inhibition is relatedclosely to the coverage of Au(111) by the more stableOC2 oxides, which block the PDS reaction sites locatedprobably on the surface step edges. The differencebetween the kinetics of the reduction of the OC2 andOC3 oxide species points to a coupling between the twoprocesses, which is responsible for the delay in thereduction of the OC2 oxide species, during which theOC3 species decay exponentially with time. A model ofthe coupling based on diffusion controlled proton catal-ysis of oxide reduction has been proposed and verifiedby numerical simulation.

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