The oxidative dissolution mechanism of uranium dioxide. I. The effect of temperature in hydrogen carbonate medium (original) (raw)

Abstract

The oxidative dissolution of uranium (IV) dioxide has been experimentally investigated as a function of hydrogen carbonate concentration at 4 different temperatures (10, 25, 45, and 60°C) by using a continuous thin-layer flow-through reactor. The experimental results have been interpreted as evidence for a bicarbonate-promoted oxidative dissolution mechanism which can be differentiated in to 3 steps: 1) initial oxidation of the uranium dioxide solid surface; 2) binding of HCO 3

FAQs

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What is the proposed mechanism for uranium dioxide oxidative dissolution in bicarbonate?add

The study proposes a three-step mechanism: oxidation of U(IV), bicarbonate surface coordination, and product detachment.

How does temperature affect the oxidative dissolution rate of UO2?add

The dissolution rate increases with temperature, displaying an activation energy range of 30-80 kJ mol-1.

What concentrations of hydrogen carbonate were tested in the dissolution experiments?add

Hydrogen carbonate concentrations ranged from 10^-4 to 0.05 mol dm^-3 during the experiments.

How was the experimental dissolution rate of UO2 calculated?add

The rate was calculated using flow rate, uranium concentration, and surface area, expressed in mol m^-2 s^-1.

What discrepancies were noted between predicted and measured dissolution rates?add

Discrepancies suggest that oxidant production by radiolysis affects oxidative dissolution at low concentrations.

Figures (9)

The experimental results in terms of rates of dissolution, tem- perature and bicarbonate concentration are collected in Table 1 and shown in Fig. 1. Preliminary results presented in de Pablo and colleagues (1997) under the same experimental conditions have been also included in Table 1. From these results, an experimental rate equation for the dissolution of UO.(s) in bicarbonate solution can be derived as: Fig. 1. Dissolution rates as a function of bicarbonate concentration at various temperatures compared with the dissolution rates predicted by Eqn. (15) (solid lines) using the constants reported in Table 4. @ 10°C, © 25°C, ™ 45°, 0 60°.

The experimental results in terms of rates of dissolution, tem- perature and bicarbonate concentration are collected in Table 1 and shown in Fig. 1. Preliminary results presented in de Pablo and colleagues (1997) under the same experimental conditions have been also included in Table 1. From these results, an experimental rate equation for the dissolution of UO.(s) in bicarbonate solution can be derived as: Fig. 1. Dissolution rates as a function of bicarbonate concentration at various temperatures compared with the dissolution rates predicted by Eqn. (15) (solid lines) using the constants reported in Table 4. @ 10°C, © 25°C, ™ 45°, 0 60°.

Table 1. Experimental dissolution rates as a function of temperature and bicarbonate concentration. * Data from de Pablo et al. (1997).

Table 1. Experimental dissolution rates as a function of temperature and bicarbonate concentration. * Data from de Pablo et al. (1997).

Table 3. Variation of log k and n (see Eqn. 1) with temperature.

Table 3. Variation of log k and n (see Eqn. 1) with temperature.

Fig. 3. Peak fit for the Uyg,7/2) peak obtained by X-ray Photoelectron Spectroscopy (XPS). Only a contribution of U(IV) can be observed. The assumption of a fast detachment of the carbonate surface complex (step 3) is supported by results from X-Ray Photo- electron Spectroscopy analysis of surface. In Fig. 3, the fit for the Une7/2) peak of the final solid surface from experiments at 25°C shows that U is observable only in the IV oxidation state, suggesting that the U(VI) remaining at the surface is present only at concentrations below detection, or that it has been rapidly dissolved by the hydrogen carbonate, as previously shown in Bruno and colleagues (1995). The same observation has been reported at acidic pH when the dissolution is enhanced (Sunder et al., 1991; Torrero et al., 1997), while at alkaline pH final solid surface is close to UO. 55, indicating a slower

Fig. 3. Peak fit for the Uyg,7/2) peak obtained by X-ray Photoelectron Spectroscopy (XPS). Only a contribution of U(IV) can be observed. The assumption of a fast detachment of the carbonate surface complex (step 3) is supported by results from X-Ray Photo- electron Spectroscopy analysis of surface. In Fig. 3, the fit for the Une7/2) peak of the final solid surface from experiments at 25°C shows that U is observable only in the IV oxidation state, suggesting that the U(VI) remaining at the surface is present only at concentrations below detection, or that it has been rapidly dissolved by the hydrogen carbonate, as previously shown in Bruno and colleagues (1995). The same observation has been reported at acidic pH when the dissolution is enhanced (Sunder et al., 1991; Torrero et al., 1997), while at alkaline pH final solid surface is close to UO. 55, indicating a slower

Fig. 2. Dissolution rate as a function of 1/T at various bicarbonate concentrations. From the slopes of these curves, apparent activation energies (E,,,) were obtained by applying the Arrhenius equation. @5-10-? mol dm~* 0 10~? mol dm~*, gi 10~? mol dm~°, 0 10-4 mol dm~3

Fig. 2. Dissolution rate as a function of 1/T at various bicarbonate concentrations. From the slopes of these curves, apparent activation energies (E,,,) were obtained by applying the Arrhenius equation. @5-10-? mol dm~* 0 10~? mol dm~*, gi 10~? mol dm~°, 0 10-4 mol dm~3

Table 2. pH values at various experimental conditions. where the values of n and k are a function of temperature. The dependence of the rate on pH was not considered because in the pH range of this study (7.5-8.5; see Table 2), it has been proven that dissolution rate is not dependent on pH (Torrero et al., 1997). The values for k and n empirically obtained from the determination of the rate of dissolution are shown in Table 3.

Table 2. pH values at various experimental conditions. where the values of n and k are a function of temperature. The dependence of the rate on pH was not considered because in the pH range of this study (7.5-8.5; see Table 2), it has been proven that dissolution rate is not dependent on pH (Torrero et al., 1997). The values for k and n empirically obtained from the determination of the rate of dissolution are shown in Table 3.

Fig. 4. Kinetic constants as a function of 1/T. From the slopes of these curves, activation energies (E,) were obtained by applying the Arrhenius equation.

Fig. 4. Kinetic constants as a function of 1/T. From the slopes of these curves, activation energies (E,) were obtained by applying the Arrhenius equation.

Fig. 5. Comparison of the dissolution rates calculated by usinf Eqn. (15) (solid line) and those obtained experimentally by several authors using unirradiated UO, (@, O, MU, A) and natural uraninites (A). @ Bruno and colleagues (1995), O Gray and colleagues (1994), mi Stew- ard and Gray (1994), 1 Grambow (1989), & Gray and Wilson (1995), A Grandstaff (1976).

Fig. 5. Comparison of the dissolution rates calculated by usinf Eqn. (15) (solid line) and those obtained experimentally by several authors using unirradiated UO, (@, O, MU, A) and natural uraninites (A). @ Bruno and colleagues (1995), O Gray and colleagues (1994), mi Stew- ard and Gray (1994), 1 Grambow (1989), & Gray and Wilson (1995), A Grandstaff (1976).

Table 5. Comparison between the dissolution rate calculated with Eqn. (15) and the experimental data obtained by Gray and Wilson (1995). For the rate predictions at 25°C the constants k,, k, and k_, reported in Table 4 have been used. However in all other cases, these constants have been calculated at different temperatures using the E,, values reported in the text.

Table 5. Comparison between the dissolution rate calculated with Eqn. (15) and the experimental data obtained by Gray and Wilson (1995). For the rate predictions at 25°C the constants k,, k, and k_, reported in Table 4 have been used. However in all other cases, these constants have been calculated at different temperatures using the E,, values reported in the text.

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