Electrochimica Acta 49 (2004) 16991709
Surface electrochemistry of UO2 in dilute alkalinehydrogen peroxide solutions
J.S. Goldik a, H.W. Nesbitt b, J.J. Nol a, D.W. Shoesmith a,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 5B7Received 25 August 2003; received in revised form 24 November 2003; accepted 28 November 2003
The reaction of H2O2 on SIMFUEL electrodes has been studied electrochemically and under open circuit conditions in 0.1 mol l1 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 (UO22+)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 ofUO3yH2O. 2004 Elsevier Ltd. All rights reserved.
Keywords: Uranium dioxide; Hydrogen peroxide; Cyclic voltammetry; Corrosion potential; Nuclear waste disposal
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 . 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 (3001000 years), then substantial corrosion of thefuel is possible . 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: firstname.lastname@example.org (D.W. Shoesmith).
conditions, only the effects of -radiolysis of water on fuelcorrosion should be important.
Despite extensive study , a full mechanistic under-standing of UO2 corrosion in the presence of -radiolysishas proven elusive. A review of previous studies is given in. 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 . 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 in Fig. 1.
Recent studies [15,16] show that the corrosion behaviorof UO2 changes substantially with peroxide concentration.For concentrations in the range from 105 to 102 mol l1,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) 16991709
Fig. 1. Schematic illustrating the local radiochemistry and corrosion pro-cesses at the UO2water interface.
For [H2O2] > 102 mol l1, (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  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 . A second, longer-term goal, is themeasurement of kinetic parameters (Tafel slopes, standardrate constants, etc.) that can be used in the mixed potentialmodel .
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, 815m 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 metallic -particles. Thisphase consists of small, spherical precipitates (0.51.5mdiameter) uniformly distributed in the UO2 matrix .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 . 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 of5060 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( = 18.2 M cm), and deaerated with UHP grade argonprior to, and during, all experiments. The electrolyte usedwas 0.1 mol l1 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) 16991709 1701
Table 1U4f7/2 XPS fitting parameters: binding energies Eb 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 , 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 in Table 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 s1, 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
-1.4 -1.2 -1.0 -0.8 -0.6 -0.4 -0.2 0.0 0.2 0.4 0.6-2.5
UO2.33 UO2+xUO3yH2O UO2+x
I / m
E / V
Fig. 2. Cyclic voltammogram of 1.5 at.% SIMFUEL in an Ar-purged 0.1 mol l1 NaCl solution at pH 9.8. Scan rate = 10 mV s1. 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 ,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, UO22+,which have limited solubility at this pH  leading to theformation of a UVI deposit (possibly UO3yH2O) on theelectrode surface . On the reverse cathodic scan boththese layers, UO2+x/UO3yH2O 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 -particles. Fig. 3 shows a series ofvoltammograms recorded in a 2 103 mol l1 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 cm2) 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
1702 J.S. Goldik et al. / Electrochimica Acta 49 (2004) 16991709
-1.2 -0.8 -0.4 0.0 0.4
E / V
Fig. 3. Cyclic voltammograms of 1.5 at.% SIMFUEL to different anodiclimits in a 2 103 mol l1 H2O2 solution. The curves are offset by1 mA cm2. The scan rate used was 10 mV s1 and the arrows indicatethe direction of the scan.
. 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 mol1)leads to a substantial increase in conductivity of the elec-trode surface ( 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 donoracceptor 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 2103 mol l1H2O2 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 s1 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 of10...