The kinetics and mechanism of H₂O₂ decomposition at the U₃O₈ surface in bicarbonate solution

Abstract

In the event of nuclear waste canister failure in a deep geological repository, groundwater interaction with spent fuel will lead to dissolution of uranium (U) into the environment. The rate of U dissolution is affected by bicarbonate (HCO₃⁻) concentrations in the groundwater, as well as H₂O₂ produced by water radiolysis. To understand the dissolution of U₃O₈ by H₂O₂ in bicarbonate solution (0.1–50 mM), dissolved U concentrations were measured upon H₂O₂ addition (300 μM) to U₃O₈/bicarbonate mixtures. As the H₂O₂ decomposition mechanism is integral to the dissolution of U₃O₈, the kinetics and mechanism of H₂O₂ decomposition at the U₃O₈ surface was investigated. The dissolution of U₃O₈ increased with bicarbonate concentration which was attributed to a change in the H₂O₂ decomposition mechanism from catalytic at low bicarbonate (≤5 mM HCO₃⁻) to oxidative at high bicarbonate (≥10 mM HCO₃⁻). Catalytic decomposition of H₂O₂ at low bicarbonate was attributed to the formation of an oxidised surface layer. Second-order rate constants for the catalytic and oxidative decomposition of H₂O₂ at the U₃O₈ surface were 4.24 × 10⁻⁸ m s⁻¹ and 7.66 × 10⁻⁹ m s⁻¹ respectively. A pathway to explain both the observed U₃O₈ dissolution behaviour and H₂O₂ decomposition as a function of bicarbonate concentration was proposed.

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Introduction

The current strategy for the disposal of spent nuclear fuel is in a deep geological repository according to the majority of the international community. The repositories provide a long-term storage solution, yet the release of radioactive species from spent nuclear fuel into the environment from the repository is projected to occur in the future upon failure of the repository barriers. Therefore, it is necessary to develop safety models for the repositories to predict their performance when failure occurs and nuclear material is exposed to the local environment. The main pathway for radionuclide release is predicted to be caused by the ingress of groundwater into the repository and interaction of the groundwater with the surface of the spent fuel. Understanding the reaction mechanisms between groundwater and spent fuel is integral to the development of safety models. Such interactions between the groundwater and spent fuel will lead to dissolution of the UO₂ matrix which constitutes the majority of the spent fuel.¹ The solubility of U in groundwater is governed by the form of U (U^(IV), U^(V) and U^(VI)), with the hexavalent U^(VI) form being more soluble than U^(IV) and U^(V).^2–4 Therefore, the presence of U^(VI) facilitates U dissolution into the groundwater upon canister failure.

Under the reducing, anoxic conditions typically found in groundwater at repository depths, the solubility of U^(IV) is very low,^5–7 and so significant dissolution of the UO₂ spent fuel may not be expected. However, radiation from the spent fuel will cause radiolysis of fuel adjacent water leading to the formation of a complex water chemistry involving radical, ionic and molecular species in the form of both reductants (e_aq−, H^·, H₂) and oxidants (OH^·, H₂O₂).⁸ This will significantly affect the local redox chemistry of the water and the oxidation state of U.

Of the oxidants generated by radiolysis, it has been shown that H₂O₂ is the dominant species in regards to U dissolution under deep geological repository conditions.^9,10 The interaction of H₂O₂ with the UO₂ surface has been thoroughly studied due to its importance for U dissolution, and it has been proposed that there are two competing pathways, both of which involve the decomposition of H₂O₂ at the UO₂ surface.^11,12 The first involves catalytic H₂O₂ decomposition forming H₂O₂ and O₂ where the UO₂ surface acts as a catalyst (ads = adsorbed):

(H₂O₂)_ads → 2(OH˙)_ads (1)

In this case, the H₂O₂ decomposition does not directly cause U dissolution. The second is an oxidative decomposition reaction where H₂O₂ oxidises U^(IV) to U^(V) (eqn (4)) and U^(V) to U^(VI) (eqn (5)) while itself being reduced to OH⁻ (eqn (6)) in a redox couple.

U^(IV)O₂ → U^(V)O₂+ + e⁻ (4)

U^(V)O₂ → U^(VI)O₂²⁺ + e⁻ (5)

Typically, groundwater also contains bicarbonate (HCO₃⁻) which has been shown to enhance the dissolution of U due to favourable complexation with U^(VI) and stabilisation of the dissolution products:^13–16

U^IV + HCO₃⁻ → U^V(HCO₃)_ads + e⁻ (7)

U^V(HCO₃)_ads + OH⁻ → U^VI(CO₃)_ads + e⁻ + H₂O (8)

U^VI(CO₃)_ads + HCO⁻₃ → (U^VIO₂(CO₃)₂)²⁻ + H⁺ (9)

Therefore, the concentration of bicarbonate is believed to have a significant effect on U dissolution and the rate of H₂O₂ decomposition at the UO₂ surface.

Due to the importance of developing models to predict U dissolution into groundwater, various studies have been undertaken with UO₂ in simulated groundwater. A recent study by Kumagai et al.¹⁷ has shown that increasing the oxygen content from UO₂ to UO_2.3 increased U dissolution and reduced the rate of H₂O₂ decomposition at the oxide surface. As the dissolution of U is governed by the redox behaviour of the U atoms, it follows that the ratio of U^(IV), U^(V) and U^(VI) will have a significant impact on both U dissolution as well as the H₂O₂ decomposition pathway. Therefore, the form of uranium oxide that exists on the spent fuel oxide will have a large effect on the dissolution of U into the environment. Due to the radiolysis of spent fuel surface adjacent groundwater and the elevated temperatures from spent fuel decay, the formation of highly oxidised forms of U is expected i.e., where x > 0.3 for UO_2+x. However, there is still a lack of knowledge regarding the impact of higher oxidised forms of U on the mechanism of U dissolution and H₂O₂ decomposition.

To investigate this, we adopted U₃O₈ as an extreme case, which corresponds to UO_2.66 containing two U^(V) atoms and one U^(VI) atom.^18–20 U₃O₈ has been observed on used nuclear fuel both in wet²¹ and air^22,23 environments, and can be used as a highly oxidised form of uranium oxide for an examination of the effects of U valence on U dissolution. As the complexation of bicarbonate with U^(VI) is thought to drive U dissolution by favourable complexation, the effect of U oxidation state on the dissolution of U in bicarbonate solution can be investigated by using U₃O₈. The H₂O₂ decomposition mechanism is dependent on U oxidation and so can also be investigated using U₃O₈ for comparison with UO₂. The concentration of bicarbonate in groundwater is dependent on the location of the deep geological repository, and can range from ∼10⁻⁴ M (Tono, Japan),²⁴ to ∼10⁻³ M (Daejeon, South Korea),^25,26 to ∼10⁻² M (Forsmark, Sweden)²⁷ and so it is necessary to understand U dissolution and H₂O₂ decomposition at uranium oxide surfaces over a range of bicarbonate concentrations.

Therefore, in this work, U dissolution from U₃O₈ suspensions with H₂O₂ as a function of sodium bicarbonate (NaHCO₃) has been investigated, and the mechanism of H₂O₂ decomposition at the U₃O₈ surface has been elucidated.

Experimental

Materials

Two samples of U₃O₈ powder were used in this study to investigate the reproducibility of the U dissolution tests. The first U₃O₈ powder (sample 1) was prepared by heating UO₂ powder to 750 °C for 3 hours under a continuous flow of air. The second (sample 2) was prepared by dissolving U metal in 13 M HNO₃ (Fujifilm Wako Pure Chemical, 60%) to form UO₂(NO₃)₂·(H₂O)_n, which was then heated under identical conditions to give a 96% yield of U₃O₈. The formation of U₃O₈ was confirmed by XRD and the data was refined using the Rietveld method.²⁸ The average crystallite size was measured using the Scherrer equation:²⁹where d is the mean crystallite size, λ is the X-ray wavelength (1.5406 Å), b is the full width at half maximum value, and θ is the diffraction peak position. The crystallite sizes were calculated as 47 and 46 nm for sample 1 and 2 respectively. The orthorhombic lattice constants were also calculated from the diffractograms using Bragg's law³⁰ for orthorhombic structures (1/d_hkl² = h²/a² + k²/b² + l²/c²) giving values of a = 6.72, b = 11.96 and c = 4.15 Å for sample 1 and a = 6.71, b = 11.95 and c = 4.14 Å for sample 2. The lattice constants were consistent with those for U₃O₈ (a = 6.72, b = 11.96 and c = 4.15 Å).³¹ This indicated that the structure of each U₃O₈ sample prepared via different methods was almost identical. Sample 1 was used for determination of the pseudo-first order rate constants for H₂O₂ decomposition to investigate the mechanism of decomposition. Sample 2 was used for determination of the second order rate constants for H₂O₂ decomposition at the U₃O₈ surface for comparison with UO₂. This assignment was solely due to the amount of sample required to conduct each set of experiments – there was an insufficient amount of sample 1 for the second order experiments. The reproducibility of the dissolution tests with different U₃O₈ samples could then be analysed by comparison of H₂O₂ decomposition rates on each sample. The specific surface area of the powders was measured for calculation of the second order rate constants. Specific surface areas were measured by the Brunauer–Emmett–Teller method³² of adsorption/desorption using Kr gas with a Micromeritics Tristar II instrument. This method involves the adsorption of a monolayer of gas onto the surface of the powder at cryogenic temperatures, and the volume of adsorbed gas provides surface area information. Values of 1.20 m² g⁻¹ and 2.52 ± 0.2 m² g⁻¹ for sample 1 and sample 2 were obtained respectively. After the immersion tests, the U₃O₈ powder was dried under vacuum and analysed by Raman spectroscopy to investigate alterations to the U₃O₈ surface.

Dissolution experiments

The effect of NaHCO₃ (Alfa Aesar) on the dissolution of U₃O₈ powder by reaction with H₂O₂ (Fujifilm Wako Pure Chemical, 30%) was investigated by monitoring the U and H₂O₂ concentration as a function of reaction time. A suspension of U₃O₈ was prepared at concentrations of NaHCO₃ between 0.1–50 mM (pH 8.2–9.7), and the suspensions were purged with Ar for approximately 18 hours to ensure removal of O₂ to imitate the anoxic conditions of groundwater. Into the suspension, H₂O₂ was added to initiate the reaction. The concentration of H₂O₂ added was 300 μM which has been shown to be optimal to study oxidative dissolution on UO₂.¹⁷ Ar purging was continued throughout the experiment. Experiments were conducted under atmospheric pressure at a temperature of 25 °C which was maintained with a coolant system. Samples of the suspension were taken at intervals over the course of the reaction. The samples were immediately filtered through a 0.45 μm filter to stop the reaction, and then analysed for H₂O₂ and U. For determination of the pseudo-first order rate constants, 50 mg of U₃O₈ (sample 1) was added to 50 ml bicarbonate solution, whilst for the second-order rate constant measurements, 50, 100, 150 and 200 mg of U₃O₈ (sample 2) were added to 70 ml bicarbonate solution. Error in the experimental methodology was estimated as <5% by conducting a set of dissolution experiments in triplicate and taking the standard deviation of the H₂O₂ pseudo-first order decay constants.

Analytical techniques

The concentration of U was measured by ICP-OES using a PerkinElmer Avio-200 spectrometer. Calibration was conducted using U standards and measurements were done in triplicate. The standard deviations of the measurements were typically <1% of the measured values. The concentration of H₂O₂ was measured by the Ghormley triiodide method where the iodide ion (I⁻) reacts with H₂O₂ and is converted to triiodide (I₃⁻) using ammonium heptamolybdate ((NH₄)₆Mo₇O₂₄) and an acidic buffer (KHC₈H₄O₄).^33,34 The concentration of H₂O₂ was then determined from the absorbance spectra of I₃⁻ at 350 nm using a Shimadzu UV-3600 Plus UV-Vis-NIR spectrophotometer. Raman analysis of the oxide surface was conducted with a JASCO NRS-4500 Raman spectrometer. A 532 nm laser was introduced through a 20× objective lens, and 3 spectra of 10 seconds each were recorded and averaged for each sample.

Results and discussion

U dissolution

The dissolution of U upon addition of H₂O₂ to U₃O₈ suspensions as a function of NaHCO₃ was investigated by measuring dissolved U concentrations over the reaction time. Fig. 1 shows the dissolved uranium (U) minus the dissolved U concentration prior to H₂O₂ addition (U₀). The dissolution of U changed over the experimental time and showed a clear effect of bicarbonate on the dissolution of U₃O₈. At 0.1 < [NaHCO₃] < 5 mM, the dissolution was low. At > 5 mM the extent of dissolution significantly increased with bicarbonate. This is due to a change in the H₂O₂ decomposition mechanism as discussed later. The magnitude of U dissolution decreased with increasing bicarbonate concentration (i.e., from 5 to 10 mM bicarbonate the increase in U dissolution was ∼0.3 mM, and from 20 to 50 mM was ∼0.1 mM) indicating a complex relationship.

Fig. 1 The dissolution of U as a function of time in a 50 mg suspension of U₃O₈ (sample 1) in (a) 0.1, 1 and 5 mM bicarbonate and (b) 10, 20 and 50 mM bicarbonate solution after addition of 300 μM H₂O₂.

At t = 1 min, the value of [U–U₀] became negative at certain bicarbonate concentrations, and so the measured values of [U–U₀]_t=1min were plotted as a function of bicarbonate concentration (Fig. 2). A decrease in the concentration of dissolved U can be seen in 5, 10 and 20 mM NaHCO₃ solution, suggesting deposition from solution of U onto the U₃O₈ surface after the initial addition of H₂O₂ as highlighted by the second y-axis. As the extent of deposition increased with bicarbonate, it can be predicted that the deposits are uranium carbonates. Under the experimental conditions, the stable form of U in solution is UO₂(CO₃)₃⁴⁻ and so the deposits may be UO₂(CO₃)₃. Another possibility is the formation of uranyl peroxide (UO₂(O₂)), where the increase in deposition with bicarbonate is due to an increase in dissolved U with bicarbonate and, therefore, an increase in uranyl peroxide.

Fig. 2 The initial change in U concentration one minute after H₂O₂ addition to the U₃O₈ suspension as a function of bicarbonate concentration.

Kinetics of H₂O₂ decomposition

The dissolution of U from U₃O₈ was induced by the addition of H₂O₂ to the bicarbonate solution. Therefore, to understand the observed U dissolution behaviour from U₃O₈, the kinetics and mechanism of H₂O₂ decomposition at the U₃O₈ surface was studied. The kinetics of the reaction between H₂O₂ and U₃O₈ as a function of NaHCO₃ were investigated by measuring the concentration of H₂O₂ over the reaction time. Fig. 3 shows the concentration of H₂O₂ after adding 300 μM to a 50 mg suspension of U₃O₈ at different bicarbonate concentrations as a function of time. The H₂O₂ concentration decreased quickly at low bicarbonate concentration (0.1 mM), but the decomposition slowed down as the bicarbonate concentration increased to 5 mM. Further increases in the bicarbonate concentration up to 50 mM caused the rate of H₂O₂ decomposition to gradually increase again, until the H₂O₂ concentration profile in 50 mM bicarbonate was similar to that in 0.1 mM bicarbonate. The results in Fig. 3 clearly show an effect of bicarbonate on H₂O₂ decomposition on U₃O₈.

Fig. 3 The concentration of H₂O₂ as a function of time in a 50 mg suspension of U₃O₈ (sample 1) in (a) 0.1, 1 and 5 mM bicarbonate and (b) 10, 20 and 50 mM bicarbonate solution after addition of 300 μM H₂O₂.

To investigate the mechanism of H₂O₂ decomposition at the U₃O₈ surface, the kinetics of decomposition was investigated. Previous studies on H₂O₂ decomposition at the surface of uranium oxides have shown that the reaction follows first order kinetics with respect to H₂O₂. As the U₃O₈ surface is in excess relative to H₂O₂, the reaction can be modelled as a pseudo-first order reaction. Therefore, the rate of H₂O₂ decomposition can be explained by,

and the kinetics of H₂O₂ decomposition can be investigated by plotting the ln[H₂O₂] vs. time, where the gradient of the resulting straight-line plot gives the pseudo-first order rate constant, k, for the reaction (Fig. 4). The plots exhibited non-linear behaviour after the initial addition of H₂O₂ to the U₃O₈/bicarbonate mixtures indicating an initial reaction of H₂O₂ with the U₃O₈ surface. This initial fast decomposition of H₂O₂ is attributed to the formation of a surface layer and is further discussed later.

Fig. 4 (a) A plot of ln[H₂O₂] vs. time as a function of bicarbonate concentration with 50 mg U₃O₈ (sample 1) in (a) 0.1, 1 and 5 mM bicarbonate and (b) 10, 20 and 50 mM bicarbonate solution showing pseudo-first order behaviour.

The calculated values of k from the linear region (from t = 2 hours to the experiment end) are plotted against bicarbonate concentration in Fig. 5 and the value of k was found to be in the range between 0.4 to 1.6 × 10⁻⁵ s⁻¹. The decrease in the pseudo-first order rate constant coincided with U deposition from solution indicating that the secondary phases that deposit on the surface of the U₃O₈ may block the approach of H₂O₂ to the surface. As the plots in Fig. 4 show linear behaviour, this suggests that these deposits are stable over the experimental timescale.

Fig. 5 The pseudo-first order rate constants for H₂O₂ decomposition.

U dissolution with H₂O₂ decomposition

The mechanism of U dissolution via H₂O₂ decomposition can be investigated by analysing the extent of U dissolution as a function of H₂O₂ decomposition. To illustrate this, Fig. 6 shows a plot of the amount of dissolved U against the amount of consumed H₂O₂ for each bicarbonate concentration. The U dissolution per H₂O₂ decomposition shows linear behaviour. If we consider the oxidative pathway for H₂O₂ decomposition at the U₃O₈ surface, U^(V) is oxidised to U^(VI) leading to decomposition of the U₃O₈ unit since the net charge in the lattice is no longer neutral. Therefore, each H₂O₂ decomposition event via oxidative decomposition will lead to a U₃O₈ dissolution event (U₃O₈ + H₂O₂ → 3UO₂²⁺_(aq)), and a gradient of 3 may be expected from the plots in Fig. 6. If we consider only catalytic decomposition of H₂O₂, no U dissolution would occur giving an ideal gradient of 0. Therefore, the measured gradients provide a ratio of oxidative to catalytic H₂O₂ decomposition at each bicarbonate concentration, assuming these pathways are the only pathways for H₂O₂ decomposition (i.e. for 50 mM bicarbonate the ratio of oxidative dissolution is and catalytic decomposition is ). The dissolution of U from U₃O₈ in 10 mM bicarbonate gave a gradient of 2.68. By comparison with the gradients measured from the dissolution of U from UO₂ (∼0.4) and UO_2.3 (∼1) with H₂O₂ addition in 10 mM bicarbonate,¹⁷ this shows that the oxidative dissolution of U increases with increased oxidation state of the uranium oxide. As the complexation of U^(VI) with bicarbonate drives the dissolution of U from the surface, it follows that the two-step oxidation of U^(IV) → U^(V) → U^(VI) for UO₂ would result in less dissolution of U than the one-step oxidation for U^(V) → U^(VI) in U₃O_8, and for intermediate UO_2+x stoichiometries the dissolution rate would increase with increasing values of x. Another point of consideration regarding U dissolution is the crystal structure of UO₂ (cubic fluorite) and U₃O₈ (orthorhombic). As the crystal structures are different, the number of surface sites for H₂O₂ decomposition will impact U dissolution. The surface site densities for UO₂ and U₃O₈ have been reported as between 126 (ref. 35) to 165 (ref. 36) sites per nm² for UO₂ and 48 (ref. 36) sites per nm² for U₃O₈. As the surface sites are ∼3 times lower for U₃O₈, the observed increase in U dissolution from U₃O₈ relative to UO₂ is more pronounced than the measured dissolved U concentrations suggest.

Fig. 6 U dissolution from a 50 mg U₃O₈ (sample 1) suspension as a function of consumed H₂O₂ in (a) 0.1, 1 and 5 mM bicarbonate and (b) 10, 20 and 50 mM bicarbonate solution.

Using the ratios taken from the gradients, the contributions of catalytic (k_cat) and oxidative (k_ox) decomposition to the measured pseudo-first order rate constant can be found and are plotted in Fig. 7 as a function of bicarbonate.

Fig. 7 The catalytic (k_cat) and oxidative (k_ox) pseudo-first order rate constants for H₂O₂ decomposition on U₃O₈ as a function of bicarbonate concentration.

At low bicarbonate concentrations, the main pathway for H₂O₂ decomposition is the catalytic decomposition mechanism as there is little U dissolution associated with H₂O₂ decomposition. As k_ox is low, this indicates that the U₃O₈ is protected from H₂O₂ by a surface layer. It is postulated that upon addition of H₂O₂ to the bicarbonate solution, oxidative dissolution proceeds on the bare U₃O₈ surface (Fig. 8). As this involves the oxidation of U^(V) to U^(VI), it is likely that U^(VI) forms a surface layer on the U₃O₈ which protects against further oxidative dissolution as U^(VI) is already fully oxidised, and due to the low concentration of bicarbonate the surface layer is stable.

Fig. 8 The formation of a protective surface layer on the surface of U₃O₈ due to oxidative decomposition of H₂O₂.

The composition of the surface layer is thought to be in the hydroxide form (UO₂(OH)₂·xH₂O) due to the formation of hydroxide from the oxidative decomposition of H₂O₂. Raman analysis of the U₃O₈ surface after removal from solution and vacuum drying showed spectra representative of U₃O₈ only (Fig. 9). Peaks relating to U₃O₈ were observed including the U–O A1_g stretching modes at 335 and 410 cm⁻¹, and the U–O E_g stretching mode at 475 cm⁻¹.³⁷ As U₃O₈ was the only phase observed, any surface layer that formed had been removed prior to Raman analysis. If the surface layer is in the hydroxide form, it is expected to decompose upon drying which would explain the observed results. Further studies are required to elucidate the composition of the surface.

Fig. 9 Raman spectra of U₃O₈ (sample 1) after the dissolution tests for different concentrations of NaHCO₃.

As the bicarbonate concentration increases from 0.1 to 5 mM, the rate of catalytic H₂O₂ decomposition decreases. This is caused by an increase in deposition from solution as seen in Fig. 2. As the pseudo-first order rate constant decreases up to 5 mM, it can be said that the deposits do not catalyse H₂O₂ decomposition to the extent that U₃O₈ does.

Increasing the bicarbonate concentration > 5 mM changes the main H₂O₂ decomposition mechanism pathway from catalytic to oxidative. At NaHCO₃ concentrations of 10 mM and above, at least 90% of the H₂O₂ decomposed via oxidation of U^(V) to U^(VI). This is due to increased dissolution of the U^(VI) surface layer and exposure of the U₃O₈ surface beneath leading an increase in k_ox. As the value of k increases with bicarbonate, this suggests that the dissolution step is rate determining rather than the redox reaction. Therefore, dissolution experiments in solutions of higher bicarbonate concentrations are required to elucidate the true value of k for oxidative dissolution in this system. Interestingly, a study by Nilsson et al. on UO₂ dissolution in 10 mM NaHCO₃ with H₂O₂ addition using pellets showed that ∼14% of H₂O₂ decomposition events occurred via oxidative dissolution while the value was even lower (∼2%) on SIMFUEL.³⁸ This suggests that k_cat is high in the case of the pellets indicating that the surface oxide that forms on the pellets is more protective than on the powders. The large discrepancy between the H₂O₂ decomposition behaviour between UO₂ pellets and U₃O₈ powder (and UO₂ powder) is a point that requires investigation.

To clarify the dependence of the catalytic and oxidative mechanisms on U₃O₈, the second order rate constants for H₂O₂ decomposition were obtained for 0.1 mM and 50 mM solutions with U₃O₈ (sample 2). The second order rate equation,

can be used to obtain the second order rate constant by plotting the pseudo-first order rate constant against the U₃O₈ surface area to total solution volume ratio (Fig. 10). The second order rate constant in 0.1 mM bicarbonate was 4.24 × 10⁻⁸ m s⁻¹. At this concentration, the decomposition was shown to be almost completely catalytic, and so this can be attributed to the catalytic decomposition reaction pathway shown in eqn (1)–(3). At 50 mM, the value of the measured second order rate constant was 8.44 × 10⁻⁹ m s⁻¹, and as the ratio of oxidative decomposition was ∼90%, we can estimate the oxidative decomposition rate constant to be 7.60 × 10⁻⁹ m s⁻¹ for the pathway shown in eqn (5) and (6). These values are within the range described in the literature for catalytic decomposition (3.6 × 10⁻⁸ to 5 × 10⁻¹¹ m s⁻¹) and oxidative decomposition (1.4 × 10⁻⁷ to 2.0 × 10⁻¹⁰ m s⁻¹) of H₂O₂ at the UO₂ surface.³⁹ The pseudo-first order rate constant measurement for 0.1 mM and 50 mM bicarbonate solutions using U₃O₈ sample 1 (shown in Fig. 5) are included in Fig. 10 showing the reproducibility of the data using different U₃O₈ powders.

Fig. 10 The pseudo-first order rate constant, k, plotted against the U₃O₈ (sample 2) surface area to solution volume ratio (m⁻¹) for 0.1 mM and 50 mM HCO₃⁻ solutions (70 ml solution). The data for 0.1 mM (□) and 50 mM (x) calculated using U₃O₈ sample 1 shown in Fig. 5 is included to show reproducibility (50 mg U₃O₈ in 50 ml solution).

Proposed pathway for U₃O₈ dissolution by H₂O₂ in NaHCO₃ solution

From the experimental results, a proposed pathway to explain the observed behaviour of U₃O₈ in bicarbonate solution with H₂O₂ is summarized, and a schematic is provided in Fig. 11. At low bicarbonate concentrations upon H₂O₂ addition, oxidative decomposition of H₂O₂ occurs at the exposed U₃O₈ surface forming a surface layer comprised of U^(VI) that provides protection against further oxidative dissolution. The decomposition of H₂O₂ proceeds via catalytic decomposition, and so the rate of U dissolution is low. The surface layer protects the U₃O₈ in bicarbonate concentrations up to 5 mM, and the H₂O₂ decomposition mechanism remains catalytic and U dissolution remains low. At 10 mM bicarbonate, the concentration of bicarbonate is sufficient to induce dissolution of the surface layer, and the surface layer does not fully protect the U₃O₈ which is exposed leading to oxidative decomposition of H₂O₂ and an increase in U dissolution. At higher bicarbonate concentrations, the surface layer is further dissolved, and oxidative decomposition of H₂O₂ and dissolution of U proceeds at higher rates.

Fig. 11 The proposed pathway for U₃O₈ dissolution upon H₂O₂ addition as a function of bicarbonate concentration.

Conclusions

Based on the presented results, the effect of bicarbonate on U dissolution from U₃O₈ with H₂O₂ addition can be split into 3 sections:

(1) [NaHCO₃] < 5 mM: H₂O₂ decomposition occurs via catalytic decomposition at the U₃O₈ surface, and the dissolution of U into solution is low.

(2) 5 mM < [NaHCO₃] < 20 mM: secondary phases deposit onto the surface of the U₃O₈ upon H₂O₂ addition. The mechanism of decomposition changes from catalytic to oxidative, causing dissolution of U.

(3) [NaHCO₃] > 20 mM: the decomposition mechanism of H₂O₂ is >90% oxidative, leading to significant dissolution of U.

The concentration of bicarbonate and form of uranium oxide has a large influence on U dissolution and H₂O₂ decomposition. Significant dissolution of U from U₃O₈ was observed at bicarbonate concentrations > 5 mM, and the extent of U dissolution was found to be larger on U₃O₈ than for UO₂ which was attributed to the one-step oxidation of U^(V) to U^(VI) for U₃O₈ compared to the two-step oxidation for UO₂ from U^(IV) to U^(V) to U^(VI). The rate of H₂O₂ decomposition on U₃O₈ was comparable to literature data for UO₂. However, the mechanism of H₂O₂ decomposition on U₃O₈ showed a strong dependence on the concentration of bicarbonate in solution with catalytic preferred at low bicarbonate and oxidative at high bicarbonate. The increase in catalytic activity at low bicarbonate was attributed to oxidation of the U₃O₈ surface and formation of a surface oxide.

Predicting the dissolution behaviour of spent fuel in the far future upon deep geological repository failure is a challenging task that requires significant experimental data for the development of accurate predictive models. In this work, elucidation of the mechanism of H₂O₂ decomposition on U₃O₈ and its effect on U dissolution was achieved, along with H₂O₂ decay constants as a function of simulated groundwater bicarbonate concentration. In groundwater containing high bicarbonate concentrations, significant dissolution of U from U₃O₈ is expected. This provides contributions to the development of such models for safety assessment of deep geological repositories. By demonstrating that the form of U will play a major role in the rate of U dissolution into the environment, the need for further studies regarding the effect of spent fuel composition on radionuclide dissolution into groundwater has been highlighted.

Conflicts of interest

There are no conflicts to declare.

Acknowledgements

This study was performed as part of “Project on Research and Development of Spent Fuel Direct Disposal as an Alternative Disposal Option (2020FY)” funded by the Ministry of Economy, Trade and Industry of Japan.

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