Electrochimica Acta 51 (2006) 3278–3286

Surface electrochemistry of UO2 in dilute alkalinehydrogen peroxide solutions

Part II. Effects of carbonate ions

J.S. Goldik, J.J. Noel, D.W. Shoesmith ∗,1

Department of Chemistry, The University of Western Ontario, 1151 Richmond Street, London, Ont., Canada N6A 5B7

Received 29 June 2005; received in revised form 9 September 2005; accepted 16 September 2005Available online 28 October 2005

Abstract

The electrochemical reduction of hydrogen peroxide has been studied on uranium dioxide electrodes. The reduction kinetics are found to beinfluenced by dissolved carbonate/bicarbonate ions. The formation of hydrated UVI species on the electrode surface is avoided in carbonate solutions,allowing H2O2 reduction to proceed at less cathodic potentials than in carbonate-free solutions. At more cathodic potentials, the adsorption ofccsO©

K

1

clowrtatoce

iw

d

0d

arbonate ions on the active reduction sites inhibits the H2O2 reduction reaction. Over a narrow potential region, the reduction of peroxide isatalyzed by coadsoption of H2O2 and HCO3

−/CO32−. The pH dependence of the H2O2 reduction reaction appears to be stronger in carbonate

olutions than in solutions that do not contain carbonate. This can be attributed to the displacement of inhibiting CO32−/HCO3

− adsorbed ions byH−.2005 Elsevier Ltd. All rights reserved.

eywords: Uranium dioxide; Hydrogen peroxide reduction; Mechanism; Corrosion; Nuclear waste disposal

. Introduction

The development of models to predict nuclear fuel (UO2)orrosion rates is a primary requirement in the assessment ofong-term nuclear waste disposal scenarios [1]. In the eventf container failure, the fuel surface could come into contactith groundwater, and the corrosion rate would be controlled by

edox conditions in the disposal vault. It is reasonable to assumehat failure should not occur during the time period when betand gamma radiation fields are high (<103 years). Beyond thisime, only the alpha radiation fields persist, and thus the effectsf the �-radiolysis of water on fuel corrosion are the primaryoncern. As the major oxidizing product of �-radiolysis, H2O2xerts the greatest influence on UO2 corrosion [2].

Our previous study [3] examined the surface electrochem-stry of UO2 electrodes in dilute alkaline H2O2 solutions. Itas found that the activity of the electrode surface for perox-

∗ Corresponding author. Tel.: +1 519 661 2111x86366; fax: +1 519 661 3022.E-mail addresses: jgoldik@uwo.ca (J.S. Goldik), jjnoel@uwo.ca (J.J. Noel),

wshoesm@uwo.ca (D.W. Shoesmith).

ide reduction was related to its surface chemical state. MixedUIV/UV (i.e., UO2+x) oxide surfaces were able to support thereduction process, while an accumulation of UVI species, suchas (UO2

2+)ads, blocked the reaction. In this study, we examine theeffects of carbonate and bicarbonate ions, which are both impor-tant groundwater species known to strongly complex UO2

2+ [4].The influence of carbonate/bicarbonate ions on the oxidative

dissolution of UO2 has been studied by several authors. In O2containing solutions, Grandstaff [5] found that the amount ofdissolved uranium increased with [HCO3

−] and PCO2 . At lowtotal carbonate concentrations, the dissolution rate was approx-imately first order in [HCO3

−]. Nicol and Needes [6] studiedthe anodic reaction using electrochemical techniques, and con-firmed that UO2 dissolution proceeds at more cathodic potentialsin the presence of carbonate, as compared to a carbonate-freeelectrolyte. Shoesmith et al. [7,8] used a combination of elec-trochemical and XPS techniques to confirm this mechanism forUO2 oxidation and dissolution in carbonate solutions. Hiskey[9] determined a rate equation r = k[H2O2]1/2[CO3]1/2 for UO2leaching by H2O2 in ammonium carbonate solutions over the pHrange 9–11. These results were interpreted as being consistent

1 ISE member. with the electrochemical mechanism of Nicol and Needes.

013-4686/$ – see front matter © 2005 Elsevier Ltd. All rights reserved.oi:10.1016/j.electacta.2005.09.019

J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286 3279

More recently, de Pablo et al. [10], have shown that, incarbonate-free solution at pH 6, the H2O2 consumption ratewas greater than the UO2 dissolution rate. This was attributedto the formation of surface oxides on UO2 (possibly schoepite,[(UVIO2)8O2(OH)12]·12H2O). The authors suggested that theseoxides could catalyze the H2O2 decomposition process, whichmay also account for some of the difference in the rates of H2O2consumption and UO2 dissolution. The same authors concludedthat the surface oxide phase was not formed in carbonate solu-tion, consistent with the XPS evidence of Shoesmith and Sunder[7,8]. Ekeroth and Jonsson [11] found that HCO3

− increasedthe H2O2 consumption rate as well as the UO2

2+ release rateto solution. A Fenton-like mechanism, involving a UV interme-diate, was proposed for the reaction between H2O2 and UO2:

UO2 + H2O2 → UO2(surf)+ + OH

• + OH−

followed by either

2UO2(surf)+ → UO2

2+ + UO2

or

UO2(surf)+ + OH

• → UO22+ + OH−

All of these previous studies dealt with either the rate of theanodic dissolution reaction, or with the kinetics of the open-circuit corrosion process, and while these revealed how theUthcdrf

2

iaT5rfLi

i(spld1aadi

current interrupt method was used to compensate for the ohmicpotential drop, which occurs mainly in the (SIMFUEL) workingelectrode, due to its small resistivity. Previously, the electroderesistivity has been measured to be ∼60 � cm [3].

SIMFUEL electrodes were cut from pellets fabricated byAtomic Energy of Canada Limited (Chalk River, Ont., Canada).The composition and microstructure of this material has beendescribed in previous publications [13,14]. The SIMFUEL usedin this study replicates UO2 fuel taken to 1.5 at.% burn-up, avalue which would be a little low for real used CANDU fuel. Theelectrodes were approximately 3 mm thick and 1.18 cm in diam-eter. The SIMFUEL was polished on wet 1200 grit SiC paperand rinsed with methanol and deionized water prior to use. Uponimmersion in the cell, a potential of −1.6 V was applied to theelectrode for 10 min in order to remove any air-formed oxides.Previous experiments and XPS evidence confirm that this stepproduces a clean surface [15].

The base electrolyte used was 0.1 mol L−1 NaCl (ACP chem-icals). The carbonate/bicarbonate concentrations were adjustedwith Na2CO3 and NaHCO3 (Merck). The pH values of the result-ing solutions were all 9.7 ± 0.2, unless stated otherwise, and thetotal carbonate concentration ([CO3]tot = [CO3

2−] + [HCO3−])

ranged from 10−3 to 0.2 mol L−1. The solution pH was con-firmed using an Orion 250A+ pH meter with an Orion 91-07Triode pH/ATC probe. The solutions were prepared from Milli-pore water (ρ = 18.2 M� cm), and were purged with UHP argongtdpbdt

3

3

S(t0oSlasstfroidfd

O2 dissolution process depends on H2O2, a full mechanis-ic interpretation has not yet been attempted. In this work, weave studied the cathodic reduction of H2O2 by electrochemi-al means. A primary rationale for studying this reaction is toevelop a current–potential relationship that may be incorpo-ated into a mixed potential model to predict the rate of nuclearuel dissolution under disposal conditions [12].

. Experimental

A standard three-electrode, three-compartment cell was usedn all experiments. A saturated calomel electrode (sce) was useds the reference, and all potentials are reported against this scale.he counter electrode was a platinum sheet (approximatelycm2) spot-welded to a 15 cm platinum wire (Alfa Aesar). The

eference and counter electrode compartments were separatedrom the main compartment of the cell by sintered glass frits. Auggin capillary was used to minimize the ohmic potential drop

n the electrolyte solution.A Solartron model 1287 potentiostat was used in all exper-

ments. The data were analyzed using CorrwareTM softwaresupplied by Scribner Associates). Electrochemical impedancepectroscopy (eis) was performed using the above setup cou-led with a Solatron model 1255B frequency response ana-yzer. A 10 mV sinusoidal potential waveform was applied, andata accumulated as a function of frequency from 100 kHz to0−2 Hz. ZplotTM software was used to collect the eis resultsnd to perform equivalent circuit fitting. That the system wast steady state was checked by recording a small number ofata points on a reverse frequency scan. The cell was housedn a grounded Faraday cage to minimize external noise. The

as (BOC gases) for at least 20 min prior to experiments in ordero expel dissolved O2. The Ar purging was continued for theuration of all experiments. Hydrogen peroxide (3% w/v, sup-lied by Fisher Scientific) was introduced to the cell immediatelyefore each experiment. The H2O2 concentration in the cell wasetermined upon completion of each experiment by redox titra-ion against standardized KMnO4 (Aldrich).

. Results

.1. SIMFUEL electrochemistry

Fig. 1 compares cyclic voltammograms obtained on a 1.5 at.%IMFUEL electrode in (a) a 0.1 mol L−1 NaCl solution, andb) a 0.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 solu-ion. The very small anodic currents for potentials −200 tomV are attributed to surface oxidation to a mixed UIV/UV

xide (UO2+x), consistent with our previous observations onIMFUEL [3]. Large anodic currents for the oxidative disso-

ution of this UO2+x layer to soluble uranyl (UO22+) species

re observed at higher potentials (>300 mV). This oxidative dis-olution reaction is clearly accelerated in carbonate containingolution; this is indicative of the ability of HCO3

−/CO32− ions

o strongly complex UO22+. In the absence of carbonate, sur-

ace oxidation to UO2+x on the anodic scan produces a smalleduction peak (A) in the potential range −600 to −900 mVn the cathodic scan. This reduction peak is noticeably dimin-shed in the carbonate solution. The formation of hydrated UVI

eposits (UO3·yH2O), which has been confirmed in carbonate-ree solutions [3], is inhibited when the electrolyte containsissolved carbonate. The large increase in cathodic current for

3280 J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286

Fig. 1. Cyclic voltammograms of 1.5 at.% SIMFUEL in 0.1 mol L−1 NaCl(dashed line), and 0.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 (solidline), both solutions at pH 9.7. ω = 16.7 Hz; ν = 10 mV s−1.

potentials negative to −1.0 V is due to H2O reduction, a processbelieved to be catalyzed by the metallic �-phase in the SIMFUEL[14].

3.2. H2O2 reduction on SIMFUEL

Fig. 2 shows voltammograms on a SIMFUEL electrode in an8 × 10−3 mol L−1 H2O2 solution, with (a) 0.1 mol L−1 NaCl,and (b) 0.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 asthe supporting electrolyte. In our previous study [3], we exam-ined the cyclic voltammetric behaviour of SIMFUEL in slightlyalkaline NaCl solutions containing H2O2. The currents observedon the cathodic scan are suppressed relative to those recordedon the anodic scan, Fig. 2. This was attributed to the formationon the forward scan of a UO3·yH2O layer by precipitation ofUO2

2+ from solution. This layer has insulating properties and

FH0al

Fig. 3. Cyclic voltammograms of 1.5 at.% SIMFUEL taken to different anodiclimits in a 0.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 solution (pH 9.7)with [H2O2] = 8 × 10−3 mol L−1. The curves are offset by 2 mA cm−2 for clarity.ω = 16.7 Hz; ν = 10 mV s−1. The markers (+) indicate the zero line for each plot.

the current for electron transfer to H2O2 is blocked, and is not‘revived’ until this layer is reduced in the potential range −600to −900 mV (region A in Fig. 1). The presence of carbonate ionsclearly influences the shape of the current–potential profile. Cur-rents on the anodic scan are suppressed for potentials cathodicto −200 mV, while those on the cathodic scan are suppressedfor potentials negative to −700 mV. The dotted line shows thetheoretical diffusion-limited current value (jL), calculated usingthe expression [16]:

jL = ζnFcbD2/3v−1/6ω1/2 (1)

where n is the number of electrons transferred in the reductionprocess (n = 2), F the Faraday constant, cb the bulk concentra-tion of H2O2 in mol cm−3, D the diffusion coefficient of H2O2(D = 1.71 × 10−5 cm2 s−1 [17]1), v the kinematic viscosity ofthe solution (assumed to be the same for both solutions; a valueof 1.0 × 10−2 cm2 s−1 was used), ω the rotation frequency of theelectrode (ω = 16.7 Hz) and ζ is a numerical coefficient given byNewman’s expression [19]:

ζ = 1.5553

1 + 0.2980(Sc)−1/3 + 0.14514(Sc)−2/3 (2)

where Sc is the Schmidt number (Sc = v/D). While the peroxidereduction current approaches the diffusion-limited value in NaClsolution, the presence of carbonate ions causes a suppression oft

icta

bcc

ig. 2. Cyclic voltammograms of 1.5 at.% SIMFUEL in an 8 × 10−3 mol L−1

2O2 solution. The electrolytes were 0.1 mol L−1 NaCl (dashed line) and.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 (solid line), both solutionst pH 9.7. The dotted line shows the theoretical diffusion limited current calcu-ated from Eq. (1). ω = 16.7 Hz; ν = 10 mV s−1.

he current below the theoretical diffusion-limited value.In the potential region −200 to +100 mV (labeled region B

n Fig. 2) there is a modest enhancement of the current on theathodic scan in the carbonate-containing solution compared tohe carbonate-free solution. Fig. 3 shows the effect of varying thenodic limit (Ean) on the voltammetric scan. The catalytic effect

1 The equilibrium H2O2 + HCO3− = HCO4

− + H2O has been determined toe significant at 25 ◦C [18]. How this affects the value of the effective diffusionoefficient is not currently known. Furthermore, the higher ionic strength of thearbonate solution should also lower the diffusion coefficient of H2O2.

J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286 3281

Fig. 4. Levich-type plots for H2O2 reduction currents recorded during acyclic voltammetric experiment: (�) −800 mV, forward scan; (�) −800 mV,reverse scan; (�) −200 mV, forward scan; (©) −200 mV, reverse scan.[H2O2] = 2 × 10−3 mol L−1.

in region B is only observed when Ean is >100 mV. The currentsrecorded in region B on the reverse scan are only weakly depen-dent on both potential and the rotation frequency of the electrode.Fig. 4 shows a Levich plot (−j versus ω1/2) for currents recordedat −800 and −200 mV (within region B) on both the forwardand reverse potential scans. For the more negative potential, thecurrents are clearly rotation rate dependent, consistent with atransport contribution to the kinetics of H2O2 reduction in thisregion. The very weak current dependence on ω at −200 mVsuggests that the reduction process occurring in region B is notsignificantly influenced by diffusion.

In order to gain a better understanding of the nature of thecatalytic effect in region B, the current on the anodic scan (jf)was subtracted from the current on the cathodic scan (jr). Thisenables us to separate the catalytic process, observed only onthe cathodic scan, from the non-catalyzed process, observed onboth the anodic and cathodic scans. Fig. 5 shows plots of thiscurrent difference (jr − jf) for a series of potential scan rates, ν

(mV s−1). Both the peak potential (EP) and the peak current aftersubtraction (jP) depend on �. Fig. 6 shows that the plot of jP ver-sus ν1/2 is linear, while EP is linear when plotted against log(ν).Both these dependencies are diagnostic for an irreversible elec-trochemical process [16]. The integrated charge (Q) in regionB, obtained by integrating (jr − jf) over the potential range forwhich jr is greater than (as an absolute value) jf, is dependenton the bulk H O concentration. This is shown in Fig. 7 fortcepaeciUl

Fig. 5. Plots of the residual current for H2O2 reduction (jr − jf) in region Bagainst the applied potential on a 1.5 at.% SIMFUEL electrode in a 0.1 mol L−1

NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 solution. The potential scan rates (ν) are:(�) 30 mV s−1; (�) 25 mV s−1; (�) 20 mV s−1; (©) 15 mV s−1; (+) 10 mV s−1;(×) 5 mV s−1. [H2O2] = 8 × 10−3 mol L−1 and ω = 16.7 Hz.

Fig. 6. Plots of the peak current (jP) and peak potential (EP) from Fig. 5 asfunctions of the potential scan rate (ν).

2 2otal carbonate concentrations of 0.1, 0.13 and 0.2 mol L−1, andonfirms that the peak involves H2O2 reduction. For the low-st carbonate concentration the total reduction charge is directlyroportional to [H2O2]. Evidently, Q is smaller at higher carbon-te concentrations. Experiments performed on a split SIMFUELlectrode (identical to 1.5 at.% SIMFUEL except that it does notontain the noble metal �-phase) also showed the catalytic effectn region B, confirming that this catalytic effect is related to theO2 lattice and is not associated with the �-particles [unpub-

ished results]. A more extensive study of the H2O2 reduction

3282 J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286

Fig. 7. Plots of the integrated charge (Q) in region B against the hydrogen perox-ide concentration. Values of [CO3]tot are: (�) 0.1 mol L−1; (©) 0.13 mol L−1;(×) 0.2 mol L−1.

reaction on SIMFUELs with different dopant concentrations isunderway.

Fig. 8 shows a cyclic voltammogram for the reduction ofH2O2 in which the initial cathodic scan (immediately follow-ing the anodic scan to +300 mV) is reversed in direction at−300 mV (i.e., just after scanning cathodically through potentialregion B). The following anodic scan back to the anodic limitdoes not follow the ‘catalytic’ path, indicating that the anodi-cally formed surface state responsible for the catalysis of H2O2reduction is destroyed once the potential is made more negativethan −300 mV (or more negative than the potentials in regionB). As shown by the current on the final cathodic scan (+300 to−1200 mV, solid line), this catalytic state is regenerated by thesecond scan to +300 mV.

F2Ni(−

Fig. 9. Tafel-type plots for the reduction of H2O2 on 1.5 at.% SIMFUEL insolutions of different total carbonate concentration. [H2O2] = 8 × 10−3 mol L−1.The values of [CO3]tot are: (�) 1 × 10−3 mol L−1; (�) 2 × 10−3 mol L−1;(�) 2 × 10−2 mol L−1; (©) 5 × 10−2 mol L−1; (+) 1 × 10−1 mol L−1; (×)2 × 10−1 mol L−1.

Fig. 9 shows a series of Tafel-type plots for H2O2 reduction asa function of [CO3]tot. The kinetic reduction currents (jk) weredetermined from the anodic scans of voltammograms (identicalto Fig. 2) by varying ω and extrapolating to ω−1/2 = 0 accordingto the Koutecky–Levich equation [16]:

1

−j= 1

−jk+ 1

−jL(3)

where jL is given by Eq. (1). That the currents obtained inthis manner are steady-state kinetic values was demonstratedby a limited number of potentiostatic experiments in which thecurrent at individual potentials was recorded. In these moretedious experiments, a steady-state current was establishedalmost immediately. The Tafel slopes are greater than −0.4 V perdecade over the entire potential region for all [CO3]tot, consistentwith our prior observations in alkaline NaCl solutions [3,20]. Atpotentials <−0.5 V, the reduction current at a constant [H2O2] issuppressed as [CO3]tot is increased. For potentials >−0.3 V thecurrent is enhanced by carbonate, consistent with the catalyticeffect described above.

3.3. Effects of pH

Voltammograms for the reduction of H2O2 on SIMFUEL incTwmitiiBlo

ig. 8. Cyclic voltammograms at 10 mV s−1 for the reduction of× 10−3 mol L−1 H2O2 on a 1.5 at.% SIMFUEL electrode in 0.1 mol L−1

aCl + 0.1 mol L−1 Na2CO3/NaHCO3. The potential profile is shown in thenset. The dotted line shows the current recorded on the first section of the scanto −300 mV) and the solid line shows the second section of the scan (from300 mV).

arbonate solutions of pH 8.6 and 13.0 are compared in Fig. 10.he H2O2 reduction wave is shifted in the cathodic directionith an increase in pH. The magnitude of this shift is approxi-ately 61 mV/pH for j = 0 on the forward scan (points C and D

n Fig. 10). Fig. 11 shows a series of Tafel plots for H2O2 reduc-ion at different pH values. The kinetic currents were obtainedn the same manner as in Fig. 9. Between pH 8.6 and 10.6, theres a steady increase in the cathodic current with increasing pH.etween pH 10.6 and 11.4, the H2O2 reduction rates show very

ittle pH-dependence. At pH 13.0, there is a significant inhibitionf the H2O2 reduction currents.

J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286 3283

Fig. 10. CVs of 1.5 at.% SIMFUEL in a 0.1 mol L−1 NaCl + 0.1 molL−1 Na2CO3/NaHCO3 solution with [H2O2] = 8 × 10−3 mol L−1 at differentpH (indicated in the figure). The scan rate was 10 mV s−1 and the electroderotation frequency was 16.7 Hz.

3.4. EIS results

Fig. 12 shows eis results (presented as a Nyquist plot)obtained on a SIMFUEL electrode polarized at +300 mVin a 0.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3solution (pH 9.7) with (a) no added H2O2, and (b)[H2O2] = 8 × 10−3 mol L−1. The two time constant equiv-alent circuit shown in Fig. 13 was used to fit the data. Inthis circuit Rct and Cdl represent the resistance to chargetransfer at the electrode solution interface and the double layercapacitance, respectively: Rads and Cads represent the resistanceand capacitance of adsorbed intermediates involved in theanodic dissolution process, and Rs is the solution resistance.The values obtained by fitting this circuit to the data are listedin Table 1. The capacitive circuit elements were modeled asconstant phase elements to compensate for non-idealities in theimpedance responses.

FSda

Fig. 12. Nyquist plot of EIS data obtained on a SIMFUEL electrode polarizedat +300 mV in a 0.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 solution atpH 9.7. (�) [H2O2] = 0; (©) [H2O2] = 8 × 10−3 mol L−1.

Fig. 13. Equivalent circuit model used in the fitting of the EIS data.

Table 1Values of the circuit elements shown in Fig. 13 obtained by fitting to the data inFig. 12

Circuit element [H2O2] = 0 [H2O2] = 8 × 10−3 mol L−1

Rs (� cm2) 51.6 ± 0.1 52.2 ± 0.2Cdl (F cm−2) 0.000397 ± 0.000004 0.00054 ± 0.00001Cdl (exponent) 0.84 ± 0.02 0.808 ± 0.005Rct (� cm2) 248 ± 1 180 ± 2Cads (F cm−2) 0.0185 ± 0.0006 0.0204 ± 0.0005Cads (exponent) 0.95 ± 0.02 0.93 ± 0.01Rads (� cm2) 129 ± 3 421 ± 9

4. Discussion

Our previous study [3] of the surface electrochemistry ofSIMFUEL in H2O2 solutions showed that the catalytic activityof UO2 for H2O2 reduction depends strongly on the chemicalstate of the electrode surface. It was found that mixed UIV/UV

surfaces are able to support the reduction process, whereas UVI

surfaces (i.e., UO3·yH2O-covered) are insulating, and thereforeblock electron transfer. H2O2 reduction is believed to occur atdonor–acceptor redox (DAR) sites on the electrode surface [20].According to the theory of Presnov and Trunov [21,22], electrontransfer to the adsorbed molecule occurs at active sites formedby adjacent cations in different valence states. In the case ofUO2+x, the DAR sites are most likely adjacent UIV/UV cations.XPS evidence [15] confirms that a mixed UIV/UV oxide surface

ig. 11. Tafel plots for the reduction of 8 × 10−3 mol L−1 H2O2 on a 1.5 at.%IMFUEL electrode in 0.1 mol L−1 NaCl + 0.1 mol L−1 Na2CO3/NaHCO3 atifferent pH values. Kinetic currents were obtained using Eq. (3). The pH valuesre: (�) 8.6; (�) 9.6; (�) 10.6; (�) 11.2; (×) 11.4; (+) 13.0.

3284 J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286

is present in the potential region where H2O2 reduction currentsare observed.

Fig. 1 shows that CO32−/HCO3

− ions in the electrolyte solu-tion have a significant effect on the extent of anodic processesduring the cyclic voltammetric measurement. The initial stagesof UO2 oxidation (UO2 → UO2+x) in the potential region −300to +100 mV on the anodic scan appear to be unaffected by car-bonate. In the absence of carbonate, the UO2

2+ species producedat more positive potentials can exceed the solubility limit andprecipitate on the electrode surface as UO3·yH2O [4]. Both theUO2+x layer and the UO3·yH2O precipitate are reduced in thepotential region A (−600 to −900 mV) on the cathodic scan.The considerably smaller reduction peak in the carbonate solu-tion indicates that the formation of UO3·yH2O is prevented dueto complexation of UO2

2+. XPS data [10] indicate that the UO2+x

layer is still retained in carbonate solution, although it is muchthinner than in a carbonate-free electrolyte. The sharp increasein anodic current for E > +200 mV can thus be attributed to anenhanced oxidative dissolution process:

UO2+x + aHCO3− + xH2O

→ [UO2(HCO3)a]2−a + 2xOH− + (2− 2x)e− (4)

where the anodic charge is not recovered on the subsequentcathodic scan. This is consistent with the ability of dissolvedC[p

oFsp−daoikTsHdaHSd

Botptgoo

via the reactions:

UO2 + HCO3− → (UO2HCO3)ads + e− (5)

UO2(HCO3)ads + OH− → (UO2CO3)ads + e− + H2O (6)

i.e., HCO3− stabilizes a UV species on the electrode surface

as an intermediate in the overall anodic dissolution processto yield [UO2(CO3)2]2−. Based on dissolution experimentsin H2O2 solutions, the dissolution product has been shownby UV–vis spectrophotometry [26] to be a species such as[UO2(O2)b(CO3)c]2−2b−2c, formed via the reaction:

UO2 + cCO32− + bH2O2 + 2bOH−

→ [UO2(O2)b(CO3)c]2−2b−2c + 2bH2O + 2e− (7)

This claim is consistent with solution phase evidence [27,28]suggesting the formation of stable complexes between H2O2and uranyl carbonate species.

Our eis results, Fig. 12, show that H2O2 not only is ableto form stable mixed peroxide–carbonate complexes with dis-solved UO2

2+ species, but is also involved as a surface intermedi-ate in the anodic dissolution reaction. In the absence of peroxide,the two time constant response exhibited in the impedance spec-trum, Fig. 12, confirms that an adsorbed surface species involv-ing CO3

2−/HCO3− is an intermediate in the overall dissolution

process. The spectra in Fig. 12 and the derived parameters inTsRtept

atbtat[pmH

o(

FU

O32−/HCO3

− ions to form strong complexes with UVI species4], facilitating UO2 dissolution, and preventing any subsequentrecipitation.

The presence of carbonate ions also has a significant impactn the kinetics of H2O2 reduction on SIMFUEL. As shown inig. 2, the currents for H2O2 reduction on the forward anodiccan (i.e., starting from an unoxidized UO2 surface) are sup-ressed in the carbonate solution for potentials cathodic to200 mV. Studies of the electrochemical reduction of O2 on

ifferent types of UO2 electrodes have revealed similar carbon-te effects [23–25]. These have been attributed to adsorptionf CO3

2− (or HCO3−) on the active DAR sites, thereby block-

ng O2 adsorption and reduction. This is consistent with thenown ability of CO3

2−/HCO3− to coordinate UV/UVI species.

he fact that we observe a similar current suppression hereuggests that similar DAR sites are involved in both O2 and2O2 reduction. The Tafel plots in Fig. 9 confirm that theegree of inhibition increases with increased carbonate content,nd for [CO3]tot ≤ 10−3 mol L−1 any influence of carbonate on2O2 reduction becomes negligible at very negative potentials.imultaneously the catalytic effect in potential region B alsoisappears, Fig. 9.

The catalytic reduction of H2O2 observed in potential regionis clearly associated with the UO2 surface and not the surface

f the noble metal �-particles present in the SIMFUEL, sincehe effect is observed irrespective of whether such particles areresent in the material. The need to scan to very positive poten-ials in order to observe this effect, Fig. 3, indicates that it isenerated due to an anodically produced surface species. Basedn eis data [4], it has been proposed that adsorbed UO2HCO3r UO2CO3 species are formed on UO2+x at anodic potentials

able 1 confirm that the nature of this adsorbed intermediate istrongly influenced by the presence of H2O2. The decrease inct and increase in Cads indicate that the charge transfer process

o produce the adsorbed intermediate is accelerated in the pres-nce of CO3

2−/HCO3−. However the increase in Rads shows the

resence of H2O2 stabilizes the adsorbed surface and retards theransfer of UVI to solution.

It is the formation of this peroxide-containing surfacedsorbed species that accounts for the catalytic H2O2 reduc-ion process observed in cathodic scans, region B in Fig. 2,y providing an active DAR site for H2O2 reduction, as illus-rated in Fig. 14. In this figure, we indicate the involvement ofUV species since this has been proposed to be the identity of

he adsorbed U species in anodic dissolution in CO32−/HCO3

−4,29]. The significant charges involved in the H2O2 reductionrocess in region B, Fig. 7, may indicate that this surface inter-ediate can sustain charge transfer to a sequence of adsorbed2O2 molecules before it is cathodically destroyed.Once this site is destroyed by the electrochemical reduction

f UV, the catalyzed reduction of H2O2 is no longer observedFig. 8). A quantitative measure of the variation in the charge

ig. 14. Schematic showing the possible involvement of an adsorbedVO2(HCO3) species in the electron transfer process to H2O2.

J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286 3285

(Q) associated with H2O2 reduction in region B with [CO3]totis difficult since the peroxide reduction currents on both theforward (anodic) and reverse (cathodic) scans vary with [H2O2].An increase in current on the forward scan (jf), attributable tothe maintenance of a more catalytic surface in the presence ofcarbonate, would lead to a decrease in (jr − jf), and hence, tothe observed decrease in Q with [CO3]tot, Fig. 7. An alternative,and probably more plausible, explanation for the suppressionof Q with [CO3]tot could be the increased tendency to forman insulating UVIO2CO3 layer at very positive potentials andhigher [CO3]tot. In this case, the rate of UO2 dissolution becomeschemically controlled [30]:

UO2CO3 + CO32− → [UO2(CO3)2]2− (8)

and the surface coverage by catalytic UVO2HCO3 sitesdecreases.

The Tafel plots in Fig. 9 all have slopes in excess of −400 mVper decade, consistent with our previous observations in theabsence of carbonate [20]. These very large values were inter-preted in terms of a chemical–electrochemical (CE) mechanism,in which the small potential-dependence of the current is a con-sequence of the potential-dependent surface coverage of activeDAR sites. The proposed cycle involves chemical oxidation ofthe DAR site (represented as {UIV − UV}) by H2O2:

2{UIV − UV} + H O → 2{UV − UV} + 2OH− (9)

f

2

Oispafap

tTtipbiboitpieitai

Fig. 15. Plots of the Tafel slope (taken from the potential region −600 to−300 mV in Fig. 9) against the total carbonate concentration. The error barsrepresent ±2σ, where σ is the standard deviation of the slope.

current to become increasingly less dependent on potential asthe applied potential approaches −1000 mV, Fig. 9. The mostlikely function of CO3

2− (or HCO3−) is that it is chemisorbed

on the UO2 surface, thereby blocking the DAR sites utilized byH2O2 in oxidizing the surface.

In our previous work [20], we observed a weak pH depen-dence of the H2O2 reduction rate between pH 10 and 12,with a significant inhibition for pH ≥ 13. The results presentedhere (Fig. 11) show a more conspicuous dependence on pH.It was previously suggested that the decline in the reduc-tion current at very alkaline pH was related to the acid–baseproperties of the H2O2 molecule (pKa(H2O2) = 11.65 [33]),or the hydrated DAR structure, or both. The bicarbonate ion(HCO3

−) has a pKa value of 10.377 at 20 ◦C [34], and thechange in the ratio [HCO3

−]/[CO32−] with pH may explain

the more dramatic variation in H2O2 reduction rate observedhere.

The changes in Tafel behaviour apparent in Fig. 11 allowfor more extensive analysis. In the potential region labeled 1(E ≥ −600 mV), the Tafel slope is effectively independent ofpH, indicating that the mechanism of H2O2 reduction, and theinfluence of carbonate on it, are not influenced by [OH−], exceptfor a slight increase in the overall current as pH increases.However, in region 2 (E ≤ −600 mV) the relationship betweencurrent and potential changes. Once the pH exceeds 10, thesteeper dependence of current on potential indicates that thektpsi

dte(

2 2

ollowed by its electrochemical regeneration:

{UV − UV} + 2e− → 2{UIV − UV} (10)

ur results demonstrate that for potentials cathodic to −300 mV,.e., for potentials at which the catalytic UVO2HCO3 surfacetate no longer exists, the H2O2 reduction currents are sup-ressed as the [CO3]tot is increased. This inhibition can bettributed to a competition between H2O2 and CO3

2−/HCO3−

or DAR sites on the electrode surface. The affinity of carbon-te (or bicarbonate) for oxidized UO2 surfaces has been verifiedreviously, even at negative potentials [7,8].

Inspection of Fig. 9 reveals that, for the potential region −600o −300 mV, the Tafel slopes are also dependent on [CO3]tot.hese slopes are plotted against [CO3]tot in Fig. 15, along with

he estimated error (it should be noted that these calculationsnvolve less than an order of magnitude change in current). Inrevious work [20], we adapted the kinetic analysis developedy Calvo and Schiffrin [31,32] for H2O2 reduction on Cu2O tonterpret the mechanism on SIMFUEL. By this analysis it cane shown that the Tafel slopes are related to the relative ratesf reactions (9) and (10). Assuming both reaction steps to berreversible, then an increase in the Tafel slope, i.e., a tendencyowards chemical as opposed to electrochemical control, can beroduced by either a decrease in the rate of reaction (9) or anncrease in the rate of reaction (10). Since carbonate would bexpected to stabilize the oxidized UV state, and hence to kinet-cally hinder its reduction, a carbonate-induced acceleration ofhe electron transfer reaction (10) seems unlikely. That carbon-te most likely suppresses the rate of the chemical reaction (9)s further indicated by the tendency of the peroxide reduction

inetics are becoming more electrochemically influenced. Ifhe presence of adsorbed carbonate is responsible for the sup-ression in the rate of the first chemical step, then this shiftuggests that this suppression becomes less important as the pHncreases.

A possibility is that as the pH increases, OH− ions are able toisplace CO3

2− from surface sites, while in less alkaline solu-ions (i.e., lower OH− concentrations) HCO3

− ions are not asasily displaced. Consequently, the H2O2 reaction to create UV

reaction (9)) is not as severely inhibited. This is supported by the

3286 J.S. Goldik et al. / Electrochimica Acta 51 (2006) 3278–3286

observation that once pH > pKa(HCO3−), the current becomes

effectively independent of pH. Interestingly, the change in Tafelslope between regions 1 and 2 suggests that displacement ofadsorbed carbonate only occurs for E ≤ −0.6 V. Reference toFig. 1 shows that it is around this potential that the oxidized sur-face states likely to strongly coordinate carbonate are removed.Clearly for pH 13, there is a very significant change in mecha-nism consistent with previous arguments [20,35] that HO2

− isnot as readily reduced as H2O2.

5. Conclusions

(i) Carbonate ions have a significant catalytic effect on theanodic dissolution of UO2.

(ii) The influence of carbonate on the cathodic reduction ofH2O2 on UO2 surfaces depends on both potential and car-bonate concentration.

(iii) H2O2 reduction can be catalyzed at less negative potentialsby coadsorption with an anodically formed UO2HCO3 sur-face species. This species appears to act as a donor–acceptorsite, allowing electron transfer to H2O2.

(iv) The ability of carbonate to suppress H2O2 reductiondecreases as the pH decreases. The most likely reason isthe lower strength of surface coordination of CO3

2− (com-pared to HCO3

−) and its displacement from surface sites.

A

a

R

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[

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cknowledgment

This work was funded under the industrial research chairgreement between NSERC and Ontario Power Generation.

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