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Electrochimica Acta 49 (2004) 1699–1709

Surface electrochemistry of UO2 in dilute alkalinehydrogen peroxide solutions

J.S. Goldika, H.W. Nesbittb, J.J. Noëla, D.W. Shoesmitha,∗a Department of Chemistry, The University of Western Ontario, London, Ont., Canada N6A 5B7

b Department of Earth Sciences, The University of Western Ontario, London, Ont., Canada N6A 5B7

Received 25 August 2003; received in revised form 24 November 2003; accepted 28 November 2003

Abstract

The reaction of H2O2 on SIMFUEL electrodes has been studied electrochemically and under open circuit conditions in 0.1 mol l−1 NaCl(pH 9.8). The composition of the oxidized UO2 surface was determined by X-ray photoelectron spectroscopy (XPS). Peroxide reductionwas found to be catalyzed by the formation of a mixed UIV /UV (UO2+x) surface layer, but to be blocked by the formation of UVI (UO2

2+)species on the electrode surface. The formation of this UVI layer blocks both H2O2 reduction and oxidation, thereby inhibiting the potentiallyrapid H2O2 decomposition process to H2O and O2. Decomposition is found to proceed at a rate controlled by desorption or reduction of theadsorbed O2 species. Reduction of O2 is coupled to the slow oxidative dissolution of UO2 and formation of a corrosion product deposit ofUO3·yH2O.© 2004 Elsevier Ltd. All rights reserved.

Keywords: Uranium dioxide; Hydrogen peroxide; Cyclic voltammetry; Corrosion potential; Nuclear waste disposal

1. Introduction

A primary requirement in the assessment of long-term nu-clear waste disposal scenarios is the development of mod-els to predict the corrosion/dissolution of the spent fuelwaste form within a failed nuclear waste container, sincethis process will control the source term for the release ofradionuclides into groundwater systems. The prospects forlong-term containment using the proposed Canadian wastecontainer material (Cu) and design are very good, the greatmajority of containers being expected to survive well be-yond the period when radiation fields can produce oxidizingconditions[1]. However, it is judicious to assume that con-tainment will not be perfect and that some containers willfail before radiation fields are totally innocuous. If failureof the container occurs while�/� radiation fields are signif-icant (∼300–1000 years), then substantial corrosion of thefuel is possible[2]. A more reasonable assumption is thatcontainer failure leading to wetting of the fuel and, hence,the onset of its corrosion, would not occur until�/� radia-tion fields had decayed to insignificant levels. Under these

∗ Corresponding author.E-mail address: [email protected] (D.W. Shoesmith).

conditions, only the effects of�-radiolysis of water on fuelcorrosion should be important.

Despite extensive study[3–14], a full mechanistic under-standing of UO2 corrosion in the presence of�-radiolysishas proven elusive. A review of previous studies is given in[5]. The primary reason for this is that a clear understandingof the influence of hydrogen peroxide, the primary oxidizingproduct of�-radiolysis, on UO2 corrosion is still unavail-able[15]. The interaction of H2O2 with UO2 is complicatedby its decomposition to produce O2, which can also act as acathodic reagent for fuel dissolution. An attempt to illustratethese possibilities is given inFig. 1.

Recent studies[15,16] show that the corrosion behaviorof UO2 changes substantially with peroxide concentration.For concentrations in the range from 10−5 to 10−2 mol l−1,the steady-state corrosion potential (ECORR)SS, measured inslightly alkaline solution (pH∼ 9.5), eventually achievesa value independent of [H2O2] suggesting a condition ofredox buffering in which the potential is controlled by thekinetics of the peroxide decomposition process,

H2O2 + 2e− → 2OH− (1)

H2O2 → O2 + 2H+ + 2e− (2)

0013-4686/$ – see front matter © 2004 Elsevier Ltd. All rights reserved.doi:10.1016/j.electacta.2003.11.029

1700 J.S. Goldik et al. / Electrochimica Acta 49 (2004) 1699–1709

Fig. 1. Schematic illustrating the local radiochemistry and corrosion pro-cesses at the UO2–water interface.

For [H2O2] > 10−2 mol l−1, (ECORR)SS increases markedlywith [H2O2], and XPS analysis indicates the accumulationof a UVI corrosion product deposit on the UO2 electrodesurface. It was speculated that the presence of these depositsblocks H2O2 decomposition, but allows peroxide-driven dis-solution to occur. However, no convincing evidence was of-fered to support this claim.

A clear understanding of this mechanism is essential if arecently published mixed potential model[17] for fuel dis-solution is to be improved and verified. The primary goalof this paper is to determine the mechanism of interactionof H2O2 with UO2 in more detail. A primary emphasis isplaced on determining the chemical state of the UO2 sur-face, since it is expected that this will control the kineticbalance between peroxide-driven UO2 dissolution and per-oxide decomposition[18]. A second, longer-term goal, is themeasurement of kinetic parameters (Tafel slopes, standardrate constants, etc.) that can be used in the mixed potentialmodel[17].

2. Experimental

2.1. Electrode materials and preparation

Electrodes were cut from SIMFUEL pellets fabricated byAtomic Energy of Canada Limited (Chalk River, Ontario,Canada). SIMFUEL is an unirradiated analogue of used nu-clear fuel, produced by doping the UO2 matrix with a seriesof stable elements (Ba, Ce, La, Mo, Sr, Y, Rh, Pd, Ru, Nd, Zr)in proportions appropriate to simulate the chemical effectscaused by in-reactor irradiation[19,20]. The microstructureof SIMFUEL faithfully reflects that of typical CANDU fuelwith polygonal, equiaxed UO2 grains, 8–15�m in size and adensity 97% of theoretical. As a consequence of this dopingprocedure, holes are injected into the U5f band, due to thesubstitution of trivalent rare-earth species for UIV in the UO2fluorite lattice. This leads to an increase in oxide conductiv-ity. The noble metal elements (Mo, Ru, Rh, Pd), insoluble

in the oxide lattice, congregate in metallicε-particles. Thisphase consists of small, spherical precipitates (0.5–1.5�mdiameter) uniformly distributed in the UO2 matrix [19].These materials have proven useful in replicating the chem-ical effects caused by in-reactor irradiation. The SIMFUELused in these studies mimics UO2 fuel irradiated to 1.5 at.%burn-up.

Electrodes were fabricated from slices cut from these pel-lets using the procedure previously described[18]. Eachelectrode was approximately 3 mm thick and 1.2 cm in di-ameter. The resistivity of the electrodes was measured us-ing electrochemical impedance spectroscopy, and values of50–60� cm (determined from the high frequency real inter-cept of a Nyquist plot) were typical.

2.2. Electrochemical cell and equipment

The cell was a standard three-electrode, three-compartmentdesign. A saturated calomel reference electrode (SCE) wasused in all experiments, and all potentials are quoted againstthis scale. The working electrode was connected to a PineInstruments model AFASR analytical rotator to provideforced solution convection. The counter electrode was a5 cm2 Pt sheet spot-welded to a Pt wire. Electrode compart-ments were separated by sintered glass frits to minimizecontamination of the working electrode compartment. ALuggin capillary was employed to minimize the ohmic po-tential drop due to solution resistance between the referenceand working electrodes.

A Solartron model 1287 potentiostat was used to con-trol applied potentials and to record current responses.CorrwareTM (supplied by Scribner Associates) softwarewas used to control the instruments and to analyze thedata. The cell was housed in a grounded Faraday cage tominimize external sources of noise. The current interruptmethod was used to compensate for the potential drop dueto the electrode resistance.

2.3. Electrode preparation and solutions

Electrodes were polished on wet 1200 SiC paper andrinsed with methanol and distilled water prior to use. Onimmersion in the cell, a cathodic potential of−1.6 V wasapplied for 10 min to remove air-formed oxides. On com-pletion of an electrochemical measurement, the electrodewas repolished prior to further use. Electrodes prepared forXPS analysis were removed from the cell, washed in deaer-ated Millipore water, dried immediately, and then examinedby XPS. All solutions were prepared with Millipore water(ρ = 18.2 M� cm), and deaerated with UHP grade argonprior to, and during, all experiments. The electrolyte usedwas 0.1 mol l−1 NaCl adjusted to pH 9.8 with NaOH. Hydro-gen peroxide (3%, w/v, supplied by Fisher Scientific) wasadded to the cell just prior to electrochemical experiments.The peroxide concentration in the cell was determined bytitration against standardized, acidified KMnO4.

J.S. Goldik et al. / Electrochimica Acta 49 (2004) 1699–1709 1701

Table 1U4f7/2 XPS fitting parameters: binding energiesEb and full widths athalf height�E1/2

Species Eb (eV) �E1/2 (eV)

U(IV) 379.8 ± 0.2 1.70± 0.05U(V) 380.4 ± 0.2 1.70± 0.05U(VI) 381.0 ± 0.2 1.70± 0.05

2.4. XPS measurements

XPS spectra were recorded on a Surface Science SSX-100spectrometer using Al K� radiation. The U4f, including theassociated satellite peaks on the high binding energy side ofthe U4f5/2 signal, and Cls spectral regions were recorded.The U4f7/2 peak was deconvoluted into contributions fromthe uranium IV, V, and VI oxidation states using publishedbinding energies[21–27], and a peak shape with 70% Gaus-sian and 30% Lorentzian contributions. The fitting param-eters (binding energies,Eb, and full width at half max-ima,�E1/2) are summarized inTable 1. The satellite struc-ture was used to confirm the presence of the various ox-idation states of uranium. The Cls band at 284.8 eV wasused to determine, and correct for, the extent of samplecharging.

3. Results and discussion

Fig. 2 shows a cyclic voltammogram recorded at a scanrate of 10 mV s−1, and illustrates the stages of oxidation andreduction observed on the SIMFUEL electrode as a func-tion of potential. Only a very minor surface oxidation is ob-served for the potential range≥ −100 mV, consistent with

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H2O H

2

UO2.33

UO2+x

UO3·yH

2O UO

2+x

UO2 UO

2+x

UO2+x

UO3·yH

2O

A

I / m

A·c

m-2

E / V

Fig. 2. Cyclic voltammogram of 1.5 at.% SIMFUEL in an Ar-purged 0.1 mol l−1 NaCl solution at pH 9.8. Scan rate = 10 mV s−1. Electrode rotationrate = 16.6 Hz.

previous observations on SIMFUEL [28,29]. This shallowoxidation current has been shown to be attributable to theoxidation of the UO2 surface to a mixed UIV/UV oxide [18],a process which occurs via the injection of O2− ions into in-terstitial sites in the UO2 fluorite lattice. The large increasein current for E > 400 mV is due to the further oxidation ofthis UO2+x layer to produce soluble uranyl species, UO2

2+,which have limited solubility at this pH [30] leading to theformation of a UVI deposit (possibly UO3·yH2O) on theelectrode surface [18]. On the reverse cathodic scan boththese layers, UO2+x/UO3·yH2O are reduced in the potentialrange from −600 to −900 mV, region A in Fig. 2. The verylarge current at more negative potentials has been shown tobe due to the reduction of H2O to H2, and appears to occurpredominantly on the ε-particles. Fig. 3 shows a series ofvoltammograms recorded in a 2 × 10−3 mol l−1 H2O2 solu-tion from −1200 mV (i.e. the transport-limited current re-gion for H2O2 reduction) to various anodic limits. The plotsare offset on the vertical axis (by 1 mA cm−2) for clarity.The bar marked A (from Fig. 2) shows the potential rangewithin which the oxides formed on the anodic scan (withinthe range indicated by the arrow B) are reduced on the re-verse anodic scan. For an anodic scan limit of −100 mV, theforward and reverse scans are reversible, indicating no sig-nificant influence of surface state on H2O2 reduction. Thisis consistent with the voltammogram in Fig. 2, which showsno significant surface oxidation up to this potential. Whenthe anodic limit is extended to +100 mV, an enhancement ofthe current is observed on the reverse scan, suggesting thatthe cathodic reduction of H2O2 is catalyzed by the anodicformation of an oxidized surface layer on the forward scan,Fig. 2.

Our recent XPS examination of anodically formed oxidefilms shows that this surface layer is a mixed UIV/UV oxide


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2 mA·cm-2

B

A

Cur

rent

E / V

Fig. 3. Cyclic voltammograms of 1.5 at.% SIMFUEL to different anodiclimits in a 2 × 10−3 mol l−1 H2O2 solution. The curves are offset by1 mA cm−2. The scan rate used was 10 mV s−1 and the arrows indicatethe direction of the scan.

[18]. This evidence, coupled with the large separation be-tween the potential for formation of UV and the potential forits reduction (Fig. 2) is consistent with a solid-state oxida-tion process. Oxidation to create a UV species can be con-sidered as the creation of a hole in the narrow occupied U5fband of UO2. The migration of this hole via a polaron hop-ping process with a low activation energy (∼20 kJ mol−1)leads to a substantial increase in conductivity of the elec-trode surface ([31] and references therein). Chemically, thisinvolves the injection of an O2− ion into an interstitial site inthe UO2 fluorite lattice accompanied by formation of adja-cent UV species. Such an oxidation process has been shownto catalyze O2 reduction on UO2 [32,33] via the ability ofadjacent UIV and UV species to act as donor–acceptor sitesaccording to the scheme proposed by Presnov and Trunov[34,35]. Potentially, such a non-stoichiometric surface layerwould also be expected to catalyze H2O2 decomposition, asillustrated schematically in Fig. 4.

A further extension of the anodic limit to +300 mV pro-duces a slight inhibition of the H2O2 reduction current onthe reverse scan for E > −800 mV. This inhibition is verysignificantly enhanced when the anodic limit is extended to+500 mV (Fig. 3). This inhibition persists until the surface

Fig. 4. Schematic illustrating the H2O2 decomposition reaction, catalyzedby a non-stoichiometric UO2+x surface layer.

Fig. 5. Cyclic voltammograms of 1.5 at.% SIMFUEL in a 2×10−3 mol l−1

H2O2 solution. The black trace is the first scan to an anodic limit of500 mV and back, while the gray trace is the subsequent scan to the sameanodic limit and back. The scan rate used was 10 mV s−1 and the arrowsindicate the scan direction.

films formed on the anodic scan are reduced in potentialrange A (Figs. 2 and 3). Fig. 5 shows two consecutive scansto an anodic potential limit of 500 mV without any reprepa-ration of the electrode surface between scans. Cathodic re-duction does not totally revive the electrode reactivity ob-served on the first scan at anodic potentials. The differencein reactivity at anodic potentials on the first and second scansis also observed when peroxide is not present and cannot beattributed to the anodic oxidation of H2O2.

Fig. 6 shows the voltammetric scan to an anodic limit of100 mV plotted as log I versus E. No clear Tafel region forH2O2 reduction is observed. Also, the current falls precip-itously as the anodic limit of +100 mV is approached. Asimilar suppression of H2O2 reduction has been observedon a UO2 (as opposed to SIMFUEL) electrode [29]. Theabsence of a clear Tafel region may reflect the absence of asteady-state under voltammetric conditions and a more ex-tensive investigation of the kinetics of H2O2 reduction isunderway [36]. Previous XPS studies, however, have shownthat around +100 mV, the composition of the SIMFUEL sur-face becomes dominated by UVI species [18], due to anodicdissolution as UO2

2+, and the subsequent deposition of ahydrated UVI oxide (probably UO3·yH2O) on the electrode.Since a predominantly UVI surface would have insulatingproperties, this layer blocks electron transfer from the un-derlying conductive UO2+x, which most likely accounts forthe suppression of H2O2 reduction on the reverse scan. Thiscurrent is only revived once these surface layers are reducedin the potential region A, Fig. 3.

While we have no experimental evidence to confirm theformation of the phase UO3·yH2O, its formation would beconsistent with the low solubility of UVI at a pH of 9.5 [30].Our XPS evidence [18] shows a clear correlation betweenthe coverage of the electrode surface by UVI and a changein the O1s peak, indicating the incorporation of H2O and/or


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log

(-I

/ A·c

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)

E / V

Fig. 6. Cyclic voltammogram of 1.5 at.% SIMFUEL in a 2 × 10−3 mol l−1 H2O2 solution to an anodic limit of 100 mV and back, plotted as log I vs. E.

OH− species into the surface. In the short-term experimentsperformed in this study there is no visual evidence for theformation of the yellow UO3·yH2O. However, the accumu-lation of such a yellow deposit is observed in longer anodicoxidations at more positive potentials [18]. One interestingobservation (Fig. 3) is that the oxidation–reduction cycle onthe electrode leads to an enhancement of the H2O2 cathodicreduction in the transport-limiting region (E < −800 mV).A possible explanation for this behavior is that reductionof the anodically formed UIV/UV oxide and UO3·yH2O de-posits are incomplete (i.e. only proceed to UO2+x) and thesurface retains some of its ability to catalyze H2O2 reduction.

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OR

R / V

Fig. 7. Corrosion potentials as a function of time for a 1.5 at.% SIMFUEL electrode exposed to solutions of varying H2O2 concentration: (a)4.8 × 10−4 mol l−1; (b) 2.4 × 10−3 mol l−1; (c) 8.0 × 10−3 mol l−1; (d) 1.7 × 10−2 mol l−1.

Fig. 7 shows a series of corrosion potential (ECORR)measurements as a function of time in solutions contain-ing various concentrations of H2O2. For all concentrationsemployed, ECORR rises rapidly and achieves a steady-statevalue. Fig. 8 shows the reciprocal of the time requiredfor ECORR to achieve a value of −100 mV, (t−100)−1, andFig. 9 shows the steady-state values of corrosion potential,(ECORR)SS, as a function of [H2O2].

Previously, we have shown that (t−100)−1 is proportionalto the concentration of dissolved O2 [37], the gamma irradi-ation dose rate to water [38,39], and the strength of an alpharadiation source exposed to water [40]. The use of (t−100)−1


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log{

(t -100

/s)-1

}

log{[H2O

2] / mol·L-1}

Fig. 8. Reciprocal of the time taken for the corrosion potential to achieve a value of −100 mV in open circuit experiments plotted logarithmically againstperoxide concentration. The slope of the line of best fit is ∼ 1.1.

as a kinetic parameter is based on the assumption that sur-face oxidation by reaction with H2O2 will have produced asimilar extent of oxidation by −100 mV irrespective of theconcentration of H2O2 used. This extent of oxidation canbe expressed as a charge Q (in coulombs) and, hence, theintegrated oxidation rate, R = Q(t−100)

−1, will be directlyproportional to (t−100)−1. Our previous results [18] showthat, for SIMFUEL, based on electrochemical and XPS mea-surements, the surface composition at −100 mV will be amixed UIV/UV oxide (UO2+x). Despite the scatter, the datain Fig. 8 clearly show that the rate of oxidation of the sur-face to UO2+x is rapid and proportional to [H2O2].

0.000 0.005 0.010 0.01560

70

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130

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150

(EC

OR

R) SS /

mV

[H2O

2] / mol·L-1

Fig. 9. Steady-state corrosion potentials ((ECORR)SS) as a function of hydrogen peroxide concentration in 0.1 mol l−1 NaCl solution at pH 9.8. Theelectrode was rotated at 16.6 Hz.

The values of (t−100)−1 recorded here on SIMFUEL aresimilar to those recorded previously on UO2 (CANDU fuel)specimens exposed to H2O2-containing solutions [15,16].This was not the case for values recorded on UO2 [37] andSIMFUEL [28] specimens in solutions containing equivalentconcentrations of O2. In the case of O2, the rate of surfaceoxidation is >102 times slower than for H2O2, but substan-tially faster on SIMFUEL than on UO2. This was attributedto the increased number of donor–acceptor, UIV/UV, sitespresent in SIMFUEL (due to the RE3+ cation doping) whichcatalyzed the cathodic reduction of O2 thereby accelerat-ing the coupled UO2 oxidation process. However, a second


feature not previously recognized is that the extent of sur-face oxidation of UO2 is significantly greater than that whichoccurs on SIMFUEL. This feature is thought to be relatedto enhanced oxidation of non-stoichiometric grain bound-aries in UO2 formed during the fuel sintering process [41].The absence of such boundaries would account for the muchshallower oxidation process observed on SIMFUEL. This isreflected in the very small oxidation current observed in thevoltammogram in Fig. 2.

The small reduction peak in region A of Fig. 2 is a measureof the extent of surface oxidation. However, electrochemi-cally we cannot differentiate between a change in stoichiom-etry (x in UO2+x) and an increase in oxide film thickness.Despite this difficulty, the charge involved in irreversiblyoxidizing the surface is very small, suggesting minimal pen-etration of UO2+x into the surface of the electrode. Thus, toa first approximation, the slope of Fig. 8 can be consideredas a reaction order, indicating that the UO2 surface reacts toUO2+x with first-order kinetics.

The values of (ECORR)SS, Fig. 9, are more positivethan those previously measured on UO2 for [H2O2] ≤5×10−3 mol l−1, the region classified as the redox-bufferedregion [13,15]. The data recorded throughout this regionare generally more scattered than previously observed onUO2, suggesting a significant variability in the steady-statecondition of the UO2+x surface and/or the rates of elec-trochemical reactions occurring on this surface. On UO2,(ECORR)SS varied from specimen to specimen due to slightvariations in the properties of the commercially producedfuel pellets from which they were cut, but were reproducibleto within 20 mV on individual electrodes.

An independence of (ECORR)SS on [H2O2] in theredox-buffered region has previously been attributed to thecontrol of the potential by the coupling together of thehalf-reactions (1) and (2). For this to be the case, the ki-netics of this combination would have to be faster than thecoupling of half-reaction (1) with half-reaction (3),

UO2+x + 2xH+ → (UO2)2+ + xH2O + (2 − 2x)e− (3)

The rates of half-reactions (1) and (2) would be expected todepend on [H2O2]. Therefore, for ECORR to remain constantas [H2O2] changed would require that both the reactionsrespond identically to changes in [H2O2]. It would also meanthat, despite the constancy of (ECORR)SS, the decompositionrate would have to increase with [H2O2].

The equilibrium potentials (in volts) for reactions (1) and(2) are given by [42]

(Ee)1 = 1.45 − 0.059 pH + 0.0295 log[H2O2] (4)

and

(Ee)2 = 0.46 − 0.059 pH − 0.0295 log[H2O2]

+0.0295 log pO2 (5)

Where pO2 is the partial pressure of O2. Assuming pO2 isequal to one atmosphere, we can calculate the Ee values

(versus SCE) for our pH 9.8 and a [H2O2] concentration of10−3 mol l−1 to be

(Ee)1 = 783 mV

(Ee)2 = −29 mV

When compared to the (ECORR)SS value for this concen-tration (97–130 mV, Fig. 9), it is apparent that half-reaction(1) possesses a significantly greater overpotential thanhalf-reaction (2). Thus, if H2O2 decomposition is occurringon the electrode surface, then the H2O2 oxidation reactionshould be kinetically more facile.

It would be anticipated that, if H2O2 decomposition ki-netics are to proceed on the electrode surface, then a mixedUIV/UV oxide surface composition would be required, sincesuch a surface would contain the donor–acceptor sites ex-pected to catalyze decomposition as illustrated in Fig. 4.Despite this, no evidence for H2O2 oxidation is seen for po-tentials anodic to (ECORR)SS in voltammetric experiments(Fig. 3), indicating that this reaction does not occur at anysignificant rate.

We suggested previously that this was due to the blockingof the electrode surface by insulating UVI species, eitherabsorbed or in the form of a UO3·yH2O corrosion productdeposit. To determine whether this was indeed the case, weperformed XPS experiments on surfaces allowed to achievea steady state on open circuit; i.e., on the electrodes forwhich ECORR–time curves are plotted in Fig. 7. As examples,Fig. 10 shows U4f7/2 photoelectron spectra recorded on,(a) a cathodically-reduced SIMFUEL surface and, (b) ona surface oxidized on open circuit in a solution containing4×10−3 mol l−1 of H2O2. Also shown is the deconvolutionof the spectra into contributions from UIV, UV, and UVI. It isapparent from this figure that oxidation in a H2O2 solutionleads to a surface dominated by UVI species.

Fig. 11 shows the fractions of the three oxidation statesplotted as a function of [H2O2] (A) and (ECORR)SS (B).While the correlation with [H2O2] is poor, there is a clearcorrelation of composition with (ECORR)SS; as (ECORR)SSincreases the surface becomes increasingly covered with UVI

species. The lack of any firm correlation with [H2O2] ex-plains the variability observed in Fig. 9, while the clear cor-relation with (ECORR)SS confirms the conclusion from thevoltammetric results in Fig. 3, that the presence of UVI sur-face species blocks H2O2 anodic oxidation at positive po-tentials and H2O2 reduction at negative potentials. The elec-trode is reactivated for H2O2 reduction only when the UVI

surface layer is cathodically reduced in region A, Fig. 5.For these relatively short-term experiments the exact na-

ture of the UVI species blocking the electrode surface isuncertain. However, oxidation to the UVI state is incom-patible with the maintenance of the UO2+x fluorite lattice.Consequently, such a species would be expected to form anadsorbed surface species (possibly UO2(OH)2). Since themixed UIV/UV surface film is very thin, the formation of anadsorbed UVI species would impede H2O2 decomposition in


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1.2x105

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nsity

/ co

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Eb / eV

(a)

(b)

Fig. 10. X-ray photoelectron spectrum of the U4f7/2 state for a 1.5 at.% SIMFUEL electrode: (a) cathodically reduced at −1.6 V for 10 min; (b) allowedto reach a steady-state corrosion potential of ∼110 mV in a 4×10−3 mol l−1 H2O2 solution. Also shown is the deconvolution of the peak into componentsdue to the uranium IV, V, and VI oxidation states (see Table 1 for peak positions).

two ways: (i) by preventing the formation of UV in the cat-alytic UO2+x layer, and (ii) by blocking the donor–acceptor(UIV/UV) sites required to catalyze decomposition. As(ECORR)SS increases, the coverage of the surface with insu-lating UVI species becomes effectively complete, Fig. 11a,and any oxidation current for H2O2 oxidation becomesblocked. The filled data points in Fig. 11a show the fractionsof the three oxidation states in the electrode surface after1 h of electrochemical oxidation at 100 mV in the absenceof H2O2 (from [18]). The much lower UVI content of thesurface after electrochemical oxidation indicates that oxi-dation in H2O2 solutions involves a more rapid extractionfrom the fluorite lattice of UVI to produce adsorbed uranylspecies than is achieved electrochemically.

An intriguing observation is that the value of (ECORR)SSover the H2O2 concentration range from 10−4 to 10−2

mol l−1 is in the same range as that achieved if the electrodewere exposed to an air-saturated solution in the absence ofH2O2. It was suggested previously [15,16] that this couldindicate redox buffering involving rapid H2O2 decomposi-tion with the UO2 anodic dissolution reaction subsequentlydriven by the cathodic reduction of the decomposition prod-uct, O2. This mechanism could still apply but the presentresults show that H2O2 decomposition is blocked.

The decomposition of H2O2 and the oxidation of the SIM-FUEL by H2O2 produce two products which could blockthe subsequent reactivity of the surface; O2 by half-reaction(2), and, as demonstrated in this paper, surface UVI species.


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[H2O

2] / mol·L-1

(a)

(b)

Fig. 11. Oxidation state fraction on the surface of a 1.5 at.% SIMFUEL electrode, obtained from deconvolution of the U4f7/2 XPS peak: (a) as a functionof the steady-state corrosion potentials achieved in H2O2 solutions of different concentrations, (�) UIV, (�)UV, (�) UVI; (b) as a function of the H2O2

concentration in solution, (�) UIV, (�) UV, (�) UVI. Also shown, in (a), are the surface oxidation state fractions achieved after potentiostatic oxidationat 100 mV: (�) UIV; (�) UV; (�) UVI.

At the positive (ECORR)SS values attained, the reduction ki-netics of O2 will be very slow, and at pH 9.8 in the absenceof UO2

2+ complexing agents, the release of UO22+ to so-

lution would also be expected to be kinetically very slow.Thus, the rate of the dissolution process would be controlledby either the rate of release of UVI species from the surfaceor the rate of reduction of adsorbed O2 molecules coupledto the production of UVI species. Both these slow reactionsteps would liberate the blocked surface sites required forthe continuation of the potentially rapid H2O2 decompo-

sition reaction. Thus, H2O2 decomposition would be con-trolled at the rate of the subsequent O2-driven dissolutionprocess. Since UVI, as UO2

2+, has a low solubility at thepH employed in these studies, continuation of this sequenceof reaction steps would, albeit very slowly, lead to the ac-cumulation on the SIMFUEL surface of a corrosion prod-uct deposit, UO3·yH2O, as demonstrated electrochemically[18] and shown to occur under natural corrosion conditionson UO2 [15,16]. An attempt to illustrate this mechanismschematically is shown in Fig. 12.


Fig. 12. Proposed mechanism for the reaction of H2O2 with UO2 surfacesin alkaline solutions.

This mechanism also offers a plausible explanation forthe behavior observed for [H2O2] > 10−2 mol l−1, when(ECORR)SS increases to more positive values, Fig. 9. Thisincrease is more obvious in previous studies in which(ECORR)SS values have been recorded at higher [H2O2]than employed in the present study. At these higher poten-tials the rate of the overall corrosion process (UO2+x →UO2.33 → UO3·yH2O) is unsustainable by O2 reduction.A possibility is that the overall corrosion reaction becomesdriven by H2O2 reduction since, at these higher [H2O2],H2O2 can displace (O2)ads from the fuel surface, thereby ac-celerating the oxidative formation of the corrosion productdeposit.

4. Summary and conclusions

The cathodic reduction of H2O2 has been studied on SIM-FUEL (doped UO2) surfaces and shown to be catalyzedby the formation of a thin surface layer of UO2+x (mixedUIV/UV oxide), but kinetically blocked once surface ad-sorbed UVI species begin to form. The presence of UVI onthe oxidized surface has been demonstrated by XPS.

This leads to the situation where the potentially rapidH2O2 decomposition process to H2O and O2 is kineticallyblocked by coverage of the catalytic UO2+x surface by theseUVI species. This would account for the absence of a currentfor H2O2 oxidation at anodic potentials.

Under these circumstances, H2O2 decomposition can onlyproceed at the rate at which catalytic surface sites are lib-erated by the desorption of adsorbed O2 or UVI species,or by the oxidative dissolution of the surface by the ca-thodic reduction of adsorbed O2. This process will be ki-netically very slow, and oxidative conversion of UO2 toa UO3·yH2O corrosion product deposit will proceed onlyslowly.

Acknowledgements

This research was funded under the Industrial ResearchChair agreement between the Canadian Natural Sciences andEngineering Research Council (NSERC) and Ontario PowerGeneration, Toronto, Canada.

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