Impact of Stoichiometry on the Mechanism and Kinetics of Oxidative

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Impact of Stoichiometry on the Mechanism and Kinetics of Oxidative Dissolution of UO2 Induced by H2O2 and #-Irradiation Yuta Kumagai, Alexandre Barreiro Fidalgo, and Mats Jonsson J. Phys. Chem. C, Just Accepted Manuscript • Publication Date (Web): 26 Mar 2019 Downloaded from http://pubs.acs.org on March 27, 2019

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The Journal of Physical Chemistry

Impact of Stoichiometry on the Mechanism and Kinetics of Oxidative Dissolution of UO2 Induced by H2O2 and γ-Irradiation

Yuta Kumagaia, Alexandre Barreiro Fidalgob, Mats Jonssonb* a

Nuclear Science and Engineering Center, Japan Atomic Energy Agency, 2-4 Shirane Shirakata,

Tokai-mura, Nakagun, Ibaraki 319-1195, Japan b

Department of Chemistry, Applied Physical Chemistry, KTH Royal Institute of Technology,

SE-100 44 Stockholm, Sweden

Corresponding author Tel: +46 8 790 9123 E-mail: [email protected]

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Abstract Radiation-induced oxidative dissolution of uranium dioxide (UO2) is one of the most important chemical processes of U driven by redox reactions. We have examined the effect of UO2 stoichiometry on the oxidative dissolution of UO2 in aqueous sodium bicarbonate solution induced by hydrogen peroxide (H2O2) and γ-ray irradiation. By comparing the reaction kinetics of H2O2 between stoichiometric UO2.0 and hyper-stoichiometric UO2.3, we observed a significant difference in reaction speed and U dissolution kinetics. The stoichiometric UO2.0 reacted with H2O2 much faster than the hyper-stoichiometric UO2.3. The U dissolution from UO2.0 was initially much lower than that from UO2.3, but gradually increased as the oxidation by H2O2 proceeded. Increase in the initial H2O2 concentration caused decrease in the U dissolution yield with respect to the H2O2 consumption both for UO2.0 and UO2.3. This decrease in the U dissolution yield is attributed to the catalytic decomposition of H2O2 on the surface of UO2. The γ-ray irradiation induced the U dissolution that is analogous to the kinetics by the exposure to a low concentration (2 × 10−4 mol dm−3) of H2O2. The exposure to higher H2O2 concentrations caused lower U dissolution and resulted in deviation from the U dissolution behavior by γ-ray irradiation.


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Introduction The redox chemistry of uranium (U) between tetravalent and hexavalent largely governs the behavior of U in the environment.1-6 This can mainly be attributed to the difference in solubility in water between U(IV) and U(VI). Therefore, understanding of the redox chemistry is the key to evaluate and control the migration of U in the environment. The oxidation of U(IV) to U(VI) increases the solubility of U and facilitate the migration in the environment.1,6 Conversely, the reduction of U(VI) to U(IV) decreases the solubility and immobilize U by causing precipitation2. A redox reaction of particular importance is the oxidative dissolution of U(IV) oxide (UO2) induced by ionizing radiation, because it is an essential process for estimating the risk of spent UO2 nuclear fuel exposure to the environment.7,8 The exposure of the fuel to the environment is anticipated in the deep geological repository for spent nuclear fuel as a result of canister failure,9-11 and also anticipated in severe nuclear reactor accidents where all the barriers between the fuel and the environment have collapsed.12,13 An intrinsic property of spent nuclear fuel is its radioactivity. The ionizing radiation emitted from the fuel causes radiolysis of water, and eventually results in the dissolution of the UO2 fuel matrix.9-11 The water radiolysis produces oxidants (OH, H2O2) and reductants (eaq-, H, H2) that can significantly alter the redox conditions. For kinetic reasons, the oxidants will, at least initially, dominate the surface reactions resulting in net oxidation of the UO2 surface. This process has been studied extensively over many decades and it has been found that H2O2 is the radiolytic oxidant that dominates the oxidation process under conditions relevant for the 3 ACS Paragon Plus Environment


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safety assessment of a deep geological repository for spent nuclear fuel. Oxidation of UO2 by H2O2 is rather complicated mainly for two reasons:14-21 1) the oxidation process contains multiple elementary reaction steps, and 2) the redox property of the UO2 matrix is strongly dependent on the composition. The reaction between H2O2 and UO2 has for long been proposed to proceed via two competing pathways where one leads to oxidation of the surface and the other one is catalytic decomposition of H2O2 leaving the UO2 surface unaltered.9-11 In a very recent work we concluded that the competition between the two reaction pathways is governed by the concentration of H2O2.22 We proposed the surface-bound hydroxyl radical as a key-intermediate in the reaction process. The mechanism is depicted below.

1 2H2O2 +UO2→UO2⋯•OH H2O2 +UO2⋯•OH→HO2• + H2O + UO2 2 HO2•→H2O2 + O2 UO2⋯•OH→UO2+ +OH ―

(1) (2) (3) (4)

When the process initiated by reaction (1) goes through the reactions (2) and (3), the H2O2 decomposes to H2O and O2 where the UO2 surface acts as a catalyst. Reaction (4) leads to oxidation and eventually dissolution of UO2. In addition, numerous previous studies have demonstrated that 4 ACS Paragon Plus Environment

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the process of oxidative dissolution of UO2 is significantly affected by the presence of fission products.9-11,17-21 In general, trivalent fission products present as oxides in solid solution appear to reduce the redox reactivity of the material while metallic particles catalyze both oxidation and reduction. Given the impact of the presence of the fission products, the oxidative dissolution of spent fuel must be considered different from that of UO2. However, our understanding of the dissolution of UO2 still needs to be deepened as the basis on which the effects of the fission products should be added. Especially, the kinetics of the UO2 dissolution needs to be further investigated. A recent review by T. E. Eriksen et al. reported a considerable variation in the reported rate constants for the reactions of UO2 with the radiolytic oxidants.11 For example, the reported rate constants for the reaction of UO2 with H2O2 differ by more than 3 orders of magnitude. This large variation cannot be justified merely by experimental error. One possible parameter that most likely affects the rate constants is the stoichiometry of UO2. The hyper-stoichiometric nature of most UO2 powders could play a major role.9 Electrochemical studies have pointed out formation of hyper-stoichiometric UO2+x on the surface during the course of the oxidative UO2 dissolution.23,24 The formation of a UO2+x layer is reported to have an impact on the kinetics of UO2 dissolution.25 However, our understanding of the chemical kinetics for oxidant consumption as well as uranium dissolution does not account for stoichiometry.10 In this work, we have investigated the effect of the hyper-stoichiometry of UO2 on the oxidative dissolution kinetics 5 ACS Paragon Plus Environment


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by comparing UO2.0 and UO2.3. We examined the oxidative dissolution of UO2 in aqueous solution upon exposure to H2O2 and to γ-radiation.

Experimental Materials UO2 powder provided by Westinghouse AB was used in this study. Stoichiometric UO2 samples were prepared from the UO2 powder by reducing it under continuous flow of 5 % H2 in Ar at 450 °C for 9 hours.26 The O/U values of the UO2 powders were determined from the weight gain by oxidizing aliquots of the powders to U3O8 in air at 400 °C for 16 hours.27 The O/U values of the as-received and the reduced UO2 powders were 2.3 and 2.0, respectively. The specific surface areas of the powders were measured by the BET method of isothermal adsorption and desorption of a gaseous mixture consisting of 30% N2 and 70% He in a Micrometrics Flowsorb II 2300 instrument. The resulting values for UO2.0 and UO2.3 were 5.4 and 5.8 ± 0.2 m2 g−1, respectively.

H2O2 experiments The dissolution of the UO2 powders by the reaction of H2O2 was investigated by monitoring the concentrations of U and H2O2 as a function of reaction time. The reaction was performed in 1.0 × 10−2 mol dm−3 aqueous sodium bicarbonate (NaHCO3) solution (pH 9.2 - 9.5) in order to accelerate the dissolution of oxidized U(VI) on the surface of UO2 powder and to prevent precipitation of 6 ACS Paragon Plus Environment

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uranyl peroxides (UO2O2‧4H2O / UO2O2‧2H2O).

28,29

The suspension of the UO2 powder was

prepared and kept under N2 atmosphere to avoid oxidization of the surface of the UO2 powder by atmospheric oxygen (O2). Aqueous H2O2 solution was then added to the suspension under magnetic stirring and N2 purging. After the addition of H2O2, the suspension was sampled at time intervals. The sampled suspension was filtered and then analyzed by spectrophotometry (JASCO, V630) to determine the concentrations of H2O2 and dissolved U. The analytical methods of the spectrophotometry are described later.

Irradiation experiments The dissolution of UO2 powders induced by water radiolysis was examined by exposure to γ-ray irradiation. The suspension of the UO2 powder in aqueous 1.0 × 10−2 mol dm−3 of NaHCO3 solution was prepared in test tubes sealed with septa in an Ar-purged glove box. The suspension was then irradiated for up to 24 hours. During the irradiation, the suspension was agitated by continuous inflow of N2 gas. The N2 gas was washed with an aqueous solution of the same composition as the reaction solution before flowing into the suspension to prevent loss of CO2 and water. This gas washing adequately decreased the CO2 loss and the pH value of the samples after 24 hours was kept down to 9.9-10.1. After the irradiation, the suspension was filtered with 0.2 μm membrane and the filtered solution was analyzed for dissolved U and H2O2. The dose rate was measured by Fricke



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dosimetry.30 The measured dose rate was typically 0.16 Gy s−1 and the difference in dose rate throughout the series of irradiations was less than 10 %.

Analytical methods for spectrophotometry H2O2 was measured by using the Ghormley triiodide method.31,32 In the method, H2O2 is converted to triiodide ion (I3−) by reaction with the iodide ion (I−) using ammonium heptamolybdate ((NH4)6Mo7O24) as a catalyst at slightly acidic pH. Absorption spectrum of I3− was measured and the concentration of H2O2 was determined from absorbance at 350 nm. Dissolved U was measured using the

arsenazo

III

method.33,34

The

method

uses

1,8-dihydroxynaphtalene-3,6-disulfonic

acid-2,7-bis[(azo-2)-phenylarsonic acid] (arsenazo III) as a chromogenic reagent for U(VI). The absorbance at 653 nm was used to determine the concentration.

Results and discussion Kinetics We examined the reaction of H2O2 with UO2.0 by measuring the concentrations of H2O2 and dissolved U as a function of reaction time. The result obtained by the addition of 2 × 10−4 mol dm−3 of H2O2 is shown in Figure 1 along with the result of the corresponding experiment using UO2.3. The H2O2 concentration decreased much faster by the reaction with UO2.0 than by the reaction with UO2.3. Moreover, the U dissolution from UO2.0 was much lower than from UO2.3. 8 ACS Paragon Plus Environment

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Figure 1. Concentrations of H2O2 and U as a function of time for (a) UO2.0 and (b) UO2.3 exposed to H2O2 in aqueous suspensions containing 10 mM HCO3−. The first data point at time = 0 s is the initial H2O2 concentration calculated from the added amount of H2O2 and the volume of aqueous suspension. The second data point was the H2O2 concentration measured at 30 s after the H2O2 addition. The solid/liquid ratio was 1.0 mg cm−3

The results in Figure 1 clearly show that the stoichiometry of UO2 has an impact on the dissolution kinetics. In our previous study we examined the dissolution kinetics of UO2.3 and found that the reaction rate at a given H2O2-concentration and the dissolution of U are dependent on the 9 ACS Paragon Plus Environment


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initial concentration of added H2O2.22 Therefore, we also examined the effect of initial H2O2-concentration on the reaction kinetics with UO2.0.

Figure 2. H2O2 consumption by UO2.0 and U dissolution as a function of time for four different initial concentrations of H2O2: (a) 5 ×10−4 mol dm−3, (b) 1.0 ×10−3 mol dm−3, (c) 2.0 ×10−3 mol dm−3, (d) 5.0 ×10−3 mol dm−3. The solid/liquid ratio was 1.0 mg cm−3. The first data point at time = 0 s is the initial H2O2 concentration calculated from the added amount of H2O2 and the volume of aqueous suspension. The second data point was the H2O2 concentration measured at 30 s after the H2O2 addition.


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Similarly to the reaction of UO2.3, the reaction kinetics of UO2.0 is dependent on the initial H2O2-concentration. The results are shown in Figure 2. It is obvious that the reaction took longer time when the initial H2O2-concentration increased. As we will discuss in the following paragraphs, this corresponds to decrease in the reaction rate at a given H2O2-concentration with increasing initial H2O2 concentration. Moreover, the amount of U dissolution was not proportional to the added amount of H2O2. In addition, we observed a very rapid initial H2O2-consumption within a minute after the addition of H2O2 that was particularly pronounced at initial H2O2-concentrations higher than 1 × 10−3 mol dm−3. We analyzed the reaction kinetics by curve fitting. The measured data did not follow the pseudo-first-order reaction kinetics, i.e. single exponential function failed to fit the data. Therefore we evaluated the rate of H2O2 reaction, ―d[H2O2] d𝑡, as a function of the H2O2 concentration, instead of the reaction rate constant. In order to calculate numerically the H2O2 reaction rate, we fit the data with multi-exponential function and obtained smooth curves, 𝑛

𝑐 = 𝐴0 + ∑𝑖 𝐴𝑖exp ( ― 𝑘𝑖𝑡)

(1)

where c denotes the H2O2 concentration, Ai and ki were used as variable parameters and t is the reaction time. The fitted curves are shown in Figure 1 and 2. The calculated H2O2 reaction rates are shown in Figure 3.The reaction rates at the beginning were very high and difficult to analyze quantitatively.



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Figure 3. Reaction rates of H2O2 with UO2.0 obtained as the first derivatives of fit curves shown in Figure 1 and 2. The reaction rates are plotted as a function of the H2O2 concentration during the course of the reaction. The initial H2O2 concentrations are noted in the figure.

As can be seen in Figure 3, when the reaction rates are compared at the same H2O2 concentration, i.e. at the same X value, a higher initial H2O2-concentration resulted in a lower reaction rate at that H2O2 concentration. The decrease in the reaction rate with increasing initial H2O2-concentration indicates that the reactivity of UO2.0 decreased during the course of H2O2 reaction. The same trend was previously observed for UO2.3.22 It should be noted that further studies are needed to unravel the reaction process down to elementary reaction steps and to elucidate the scheme for the reactions of H2O2 on the surface of stoichiometric UO2. As we proposed in the previous paper,22 the reaction scheme is expected to contain a reaction intermediate, which is presumably surface-bound hydroxyl


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radical. Also the scheme must include the formation of oxidized, hyper-stoichiometric layer on the surface, as will be discussed later.

U dissolution yields The U dissolution is also dependent on the initial H2O2 concentration. To compare the U dissolution at different initial H2O2 concentrations, we evaluated the yield of U dissolution with respect to H2O2 consumption. The final U dissolution yield was calculated as [U]f / ([H2O2]i – [H2O2]f), where [H2O2]i is the initial H2O2 concentration and [U]f and [H2O2]f are the final concentrations of U and H2O2 after the reaction was completed, respectively. For the initial H2O2 concentration higher than 1 × 10−3 mol dm−3, the final U and H2O2 concentrations were measured after overnight reaction (data not shown in Fig. 2), to ensure that the reaction was complete. The final H2O2 concentration was less than 2% of the initial concentration in each run. The final U dissolution yields are shown in Figure 4 as a function of initial H2O2 concentration. For comparison, the yields for UO2.3 are also plotted in Figure 4.22



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Figure 4. Final U dissolution yields for UO2.0 and UO2.3. The data for UO2.3 are from Ref 22.

The results in Figure 4 show that the UO2 stoichiometry affects the U dissolution yield at low H2O2 concentrations. The general trend in final U dissolution yield with respect to the initial H2O2-concentration clearly differs between UO2.0 and UO2.3. The U dissolution yield of UO2.3 monotonically decreases with increasing initial H2O2-concentration. This trend is well in line with the mechanism of surface catalyzed decomposition of H2O2 that is expressed by Reactions (1) – (4).22 The dissolution yield for UO2.0 approaches that of UO2.3 at high initial H2O2-concentrations. This implies that the surface of UO2.0 transforms into UO2.3 by H2O2 oxidation. The reaction mechanism involving the surface oxidation seems consistent with the reaction mechanism proposed by D. W. Shoesmith and coworkers, pointing out that the U dissolution becomes significant after the UO2 surface is oxidized to UO2.3.9,23,24 Consequently the low U dissolution yields for UO2.0 at low initial H2O2-concentrations can most probably be rationalized by the fact that a significant part of 14 ACS Paragon Plus Environment

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H2O2 must be consumed for forming hyper-stoichiometric UO2+x at the surface prior to the oxidation of U to the readily soluble hexavalent form. From this follows that the U dissolution yield must change during the course of the reaction, depending on the total amount of H2O2 that UO2.0 has reacted with.

In order to further examine the effect of the total H2O2 exposure, we monitored the reaction of UO2.0 consecutively exposed to low concentrations of H2O2. Figure 5 shows the results obtained by adding 2.0 ×10−4 mol dm−3 H2O2 to UO2.0 consecutively 3 times.



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Figure 5. Evolution of reaction kinetics upon consecutive H2O2 additions to UO2.0; (a) H2O2 consumption, (b) U concentration. The values of U concentration were corrected to [U] = 0 and t = 0 for each a new addition. The initial H2O2 concentration was 2 ×10−4 mol dm−3 and solid/liquid ratio was 1.0 mg cm−3.

We observed that the consecutive H2O2 additions decrease the reaction rate and increase the U dissolution yield. These results show that the total H2O2 exposure as well as the H2O2 concentration affects the kinetics of oxidative UO2 dissolution. To illustrate this, the total U dissolution is plotted as a function of total H2O2 consumption for different sets of the experiments in Figure 6.


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Figure 6. Total U dissolution from UO2.0 as a function of the total H2O2 consumption. The plots are made from the kinetic data shown in Figures 1, 2, and 5. The vertical axis: the amount of U dissolution in total by the consecutive additions of H2O2 (1 – 3 times), the horizontal axis: the total H2O2 consumption in that consecutive additions. The solid/liquid ratio of the samples was 1.0 mg cm−3. The lines are linear fits to the data during the reaction of last 3 × 10−4 mol dm−3 of H2O2 assuming these lines have the same slope.

The momentary U dissolution yield, which corresponds the slope of the plot, is initially low and increases gradually with increasing total H2O2 consumption. Although this trend is expected from the surface oxidation prior to significant U dissolution, the relation between the U dissolution and the H2O2 consumption is clearly affected by the initial H2O2 concentration. The UO2 surface oxidation process alone cannot explain why the amount of H2O2 required for this oxidation process is dependent on the H2O2 concentration. Therefore, in order to understand the relation between U 17 ACS Paragon Plus Environment


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dissolution and H2O2 consumption (Figure 6), both of the surface oxidation process and the catalytic H2O2 decomposition should be taken into consideration.

As can be seen, all the plots in Figure 6 have similar slopes during the consumption of the last 3 to 4 × 10−4 mol dm−3 of H2O2. The difference in the initial H2O2 concentration becomes insignificant when the H2O2 concentration decreases below 3 × 10−4 mol dm−3. The average momentary U dissolution yield was evaluated based on the last 3 × 10−4 mol dm−3 consumption of H2O2 by a linear fit. The evaluated momentary yield is 77 %. The result for the addition of 2.0 × 10−3 mol dm−3 H2O2 was not included in the analysis since the number of data points was not enough for analyzing the reaction of the last 3 × 10−4 mol dm−3 of H2O2. The low U dissolution yield at high H2O2 concentrations (>3 × 10−4 mol dm−3) is in agreement with the mechanism of surface catalyzed decomposition of H2O2 as this mechanism contains the reaction of H2O2 with the surface intermediate species, Reaction (2), that is facilitated by a high H2O2 concentration.

The mechanism of surface catalyzed decomposition of H2O2 is expected to have a similar effect on the dissolution kinetics of UO2.3. In order to make this comparison between UO2.0 and UO2.3, the results of UO2.3 were plotted in Figure 7 in the same way as in Figure 6. The decrease in the U dissolution yield at high H2O2 concentrations is observed also for UO2.3. Also similarly to the results of UO2.0, the consumption of the final 3 to 4 × 10−4 mol dm−3 of H2O2 resulted in similar momentary 18 ACS Paragon Plus Environment

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U dissolution yields for the different initial H2O2 concentrations. The momentary yield of U dissolution during the consumption of the final 3 × 10−4 mol dm−3 for UO2.3 was higher than that for UO2.0. The same linear fit analysis was applied to the data of UO2.3 and the momentary U dissolution yield was evaluated to be 120 %. A yield over 100 % seems erroneous, but it simply reflects the hyper-stoichiometry of UO2.3. The average oxidation number of U in UO2.3 is 4.6, and theoretically 1 atom of U in UO2.3 needs only 0.7 molecule of H2O2 to be oxidized to U(VI). Therefore, the upper limit of the U dissolution yield could be 140 %.

Figure 7. Total U dissolution from UO2.3 as a function of the total H2O2 consumption. The solid/liquid ratio was 1.0 mg cm−3

When the results obtained by the addition of 2 × 10−4 mol dm−3 H2O2 are compared between UO2.0 and UO2.3, the U dissolution from UO2.3 increases more promptly than that of UO2.0. Again, this demonstrates that the UO2.0 surface must be oxidized to hyper-stoichiometric UO2+x before the U 19 ACS Paragon Plus Environment


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dissolution becomes significant. Therefore, the reaction kinetics of oxidative UO2 dissolution must involve chemical evolution of UO2 surface coupled with the reaction in aqueous solution.

Radiation-induced dissolution of U The radiation-induced oxidative dissolution of UO2-based nuclear fuel under the conditions of a geological spent fuel disposal is expected to be characterized by continuous exposure to low steady-state concentrations of oxidants produced by water radiolysis and H2O2 is assumed to have the highest impact among the oxidants on the oxidative dissolution of the fuel.35,36 According to this understanding, the radiation-induced U dissolution can basically be simulated by the reaction with H2O2 at low concentration. However, this should be also evaluated separately for UO2.0 because all the results of the H2O2 experiments indicate that the mechanism of the oxidative U dissolution is far more complicated than our previous simplistic model.10 Therefore, γ-ray irradiation experiments were performed and compared to the results of the H2O2 experiments. The results of γ-ray irradiations are shown in Figure 8. It is obvious that the dissolution of U by γ-ray irradiation of UO2.3 is more significant than for UO2.0. This is in line with the results of the H2O2 experiments, and demonstrates that the U dissolution is limited during the oxidation of UO2.0. The concentration of H2O2 was in the order of 10−5 mol dm−3 in the irradiated solutions.


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Figure 8. Concentrations of (a) U and (b) H2O2 as a function of absorbed dose.

To compare the U dissolution induced by γ-ray irradiation and that induced by H2O2 exposure, the results are plotted in Figure 9 on the same scale for U dissolution. In Figure 9a, the U dissolution per unit surface area is shown as a function of the H2O2 consumption per unit surface area. The data were normalized by the solid surface area to solution volume ratio (S/V) to compensate for the difference in S/V between the γ-ray irradiation experiments and the H2O2 exposure experiments. Figure 9b shows the results of γ-ray irradiation experiments normalized in the same way. It is difficult to directly compare these data on the same horizontal axis based on a quantitative correlation of the absorbed dose to the total H2O2 consumption. This is mainly because the oxidation of UO2 by carbonate radical (CO3•−) is not negligible in the U dissolution by γ-ray irradiation.36 In 21 ACS Paragon Plus Environment


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aqueous NaHCO3 solution CO3•− is generated from the reaction of hydroxyl radical (•OH) with HCO3−/CO32−. CO3•− is involved in the interfacial reaction. Nevertheless, the data show analogous dependence of the U dissolution on the progress of oxidation. The correlation between the total H2O2 consumption and the absorbed dose becomes c.a. 4 × 10−8 mol J−1 when it is estimated from this analogy. This value is reasonable because it is comparable with the reported yield of H2O2 (3 × 10−8 mol J−1) in γ-radiolysis of 5× 10−2 mol dm−3 aqueous NaHCO3 solution.37 This analogy indicates that the UO2 oxidation proceeds through the same reaction mechanism between the exposure to H2O2 and γ-radiation even though these two processes experimentally had different kinetics, and therefore indicates that the radiolytic oxidative dissolution process can be simulated by H2O2 reaction when the concentration of H2O2 is 2 × 10−4 mol dm−3 or lower.

Figure 9. Comparison of U dissolution between H2O2 reaction and γ-ray irradiation. The values are normalized by the surface/volume ratios (S/V).


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Conclusions We have demonstrated that the reaction kinetics of the oxidative UO2 dissolution is significantly affected by the stoichiometry of UO2. When stoichiometric UO2.0 and hyper-stoichiometric UO2.3 are compared, the reaction of UO2.0 with H2O2 was much faster than that of UO2.3. At the same time, more U is dissolved from UO2.3 than from U2.0. The results of consecutive exposure of UO2.0 to H2O2 and the results of γ-ray irradiation of UO2.0 consistently shows that the U dissolution yield significantly increases when the surface of UO2.0 has been oxidized. In addition, the radiolytic dissolution of UO2 can be simulated by the H2O2 exposure only when the concentration of H2O2 is low enough. The results of this study suggest that 3 × 10−4 mol dm−3 is the upper limit of the working concentration of H2O2 for such studies.

Acknowledgement The Swedish Nuclear Fuel and Waste Company (SKB) is gratefully acknowledged for financial support. The authors gratefully acknowledge Dr. Inna Soroka and Ms. Annika C. Maier for their help in characterization of UO2 powders.

References



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TOC Graphic


Impact of Stoichiometry on the Mechanism and Kinetics of Oxidative

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