Effect of Pyrite on the Leaching Kinetics of Pitchblende in the Process of Acid In Situ Leaching of Uranium

: In the process of acid in situ leaching of sandstone uranium ore, pyrite, which is a common associated mineral of pitchblende, would inevitably participate in the reaction. Therefore, it is important to study the inﬂuence of pyrite on the leaching kinetics of pitchblende. In this study, we compared the difference leaching rates of pitchblende in the systems of sulfuric acid–hydrogen peroxide, sulfuric acid–hydrogen peroxide–pyrite and sulfuric acid–pyrite and studied the inﬂuence of temperature and pyrite quantity on the leaching rate of pitchblende. The results show that the leaching process of pitchblende follows the shrinking particle model controlled by a chemical reaction, and the apparent activation energy Ea of the leaching reaction is (3.74 ± 0.40) × 10 kJ/mol. Pyrite itself cannot promote the dissolution of pitchblende; however, it can promote the leaching of pitchblende in the presence of an oxidizer. Increasing the quantity of pyrite in a certain range can increase the leaching rate of pitchblende, and the reaction order of pyrite is 0.36.

Acid in situ leaching uranium mining is an important uranium mining method, in which the leaching solution is directly injected into the underground ore-bearing strata through drilling, and uranium is obtained by the chemical reaction of the mineral and aqueous solution. Sulfuric acid (H 2 SO 4 ) is widely used in acid in situ leaching, because it is cheaply priced and has a quick reaction [5][6][7]. In spite of the advantages, such as low production cost and reduced damage to the surface of the ecological environment, a large amount of sulfuric acid injected will reach the underground ore aquifer and cause serious pollution to the groundwater environment. Furthermore, sulfuric acid leaching is a nonselective process resulting in other minerals being dissolved into the groundwater and affects the in situ leaching process. Uranium minerals containing uranyl (e.g., autunite) in the ore-bearing strata can react directly with H 2 SO 4 and dissolve, while only a small part of U(IV) can dissolve in H 2 SO 4 under natural conditions. However, most of the uranium in sandstone uranium deposits is UO 2 [8]. The leaching of UO 2 first requires oxidizing it 2 of 7 into the U(VI) redox state, where it can more readily dissolve, as shown in Equation (1) [9]. Hydrogen peroxide (H 2 O 2 ) is commonly used as the oxidant [10].
Some research on the effect of iron on the in situ leaching of uranium has been studied. Amme [11] investigated the impact of the reactions between hydrogen peroxide (H 2 O 2 ) and iron (Fe 2+ /Fe 3+ ) on UO 2 dissolution in an oxygen-free batch reactor. The interaction in the absence of UO 2 gave a stoichiometric redox reaction of Fe 2+ [12] studied the kinetics of uranium dissolution and migration under the action of an acidic solution containing Fe 3+ and its relationship with Fe 3+ . They found that the uranium oxidized by Fe 3+ migrates from the ore to the solution within 10 h; in addition, the reaction rate of uranium was positively correlated with the transformation rate of Fe 2+ and Fe 3+ . When the transformation rate of Fe 3+ to Fe 2+ reached zero, the oxidation and dissolution of uranium nearly ceased, and the uranium concentration in the solution achieved an equilibrium. The reaction rate of uranium, v(U), with respect to Fe 3+ , v(Fe 3+ ), in the solution was shown to follow v(U) = 0.0206 + 0.0429 exp [−v(Fe 3+ )/5.07]. Filippov [13] studied the manganese dioxide oxidation of UO 2 in the absence of iron ions, showing that the redox potential cannot be used as the only standard to judge the oxidation rate, and the real reaction rate depends on the reaction mechanism. After the addition of Fe 3+ , the dissolution percentage of UO 2 and the redox potential rise sharply, which proves that iron ion plays a catalytic role in the process of oxidizing UO 2 . Kinetics can be interpreted as a tool for investigating the rates of chemical reactions and understanding the ways different processes are affected. The most commonly employed method for analyzing the kinetics of uranium leaching is shrinking particles with the shrinking core model [14,15]. There has been no report on the influence of the FeS 2 coexistence in sandstone uranium ore on the leaching of UO 2 under the conditions of a strong acid and oxidant in acid in situ leaching mining. Therefore, this study explores the influence of FeS 2 on the leaching of U from FeS 2 to obtain the basic kinetics of H 2 SO 4 .

Materials
Both UO 2 and FeS 2 were powders below 74 µm (passed through a 74-µm sieve) from 272 Uranium Industry Co. Ltd., China National Nuclear Corporation. H 2 SO 4 and H 2 O 2 (30%) were purchased from Hengyang Kaixin Chemical Reagent Co., Ltd (Hengyang, China). The reagents used in the experiment were all analytically pure, and the water used was deionized water.

Methods
A 250-mL mixed solution of 5 g/L H 2 SO 4 and 0.06 mol/L H 2 O 2 was added into a 500-mL three-neck flask equipped with a condenser and was heated in an electric thermostatic water bath (Shanghai Kuntian, Shanghai, China) to the desired reaction temperature (15-45°C). Different amounts of FeS 2 (0.1, 0.4, 0.8 and 1.2 g) and 0.2 g UO 2 were added to the flask. The supernatant was extracted over a range of leaching times and filtered to obtain a 1-mL solution. The uranium concentration was analyzed by an atomic absorption spectrophotometer (Thermo Fisher, Waltham, America). The total iron and Fe 2+ concentrations were analyzed by a UV spectrophotometer (Beifen-Ruili, Beijing, China); the difference between them was the concentration of Fe 3+ . The Eh value was measured by a redox potentiometer from Mettler (Zurich, Switzerland). Each experiment was repeated twice, and the average value was used. Equation (2) was used to calculate the leaching rate of UO 2, where c is the concentration of U in the solution (mg/L), V is the volume of the solution (mL) and m 0 is the initial mass of U in UO 2 (mg).

Results and Discussion
In the cases of 5 g/L H 2 SO 4 , 0.06 mol/L H 2 O 2 and 0.2 g UO 2 at 25 • C, the influence of a pyrite addition amount on different systems of H 2 SO 4 -H 2 O 2 , H 2 SO 4 -H 2 O 2 -FeS 2 was investigated. The concentrations of U and Fe under different FeS 2 additions are shown in Figure 1. Without the addition of FeS 2 , the leaching rate of UO 2 was slow. The maximum leaching rate was only 36.88%, with a U concentration of 260.06 mg/L at 360 min. After the addition of 0.1 g FeS 2 , there was no obvious rate change in the initial stage of the reaction; however, the relative rate of the reaction slowly increased after 120 min compared to that without FeS 2 , and the final leaching rate was 48.40% with a U concentration of 341.31 mg/L. With further increases of the FeS 2 mass, the Fe 3+ ion concentration quickly increased, resulting in an increase of the U concentration. When the Fe concentration was less than 5 mg/L, it had no obvious effect on the UO 2 leaching. When the Fe concentration reached 5 mg/L, the reaction rate of U was obviously higher than that without Fe. When the Fe concentration was about 20 mg/L, the reaction rate of UO 2 reached the maximum. The slope of the U concentration versus time gradually became smaller, indicating that the reaction rate of UO 2 gradually slowed down at the end of the experiment, and the final leaching rates of U were 64.79%, 76.34% and 79.58%, corresponding to the FeS 2 amounts of 0.4 g, 0.8 g and 1.2 g, respectively.
Minerals 2022, 12, x FOR PEER REVIEW Equation (2) was used to calculate the leaching rate of UO2, where c is the concentration of U in the solution (mg/L), V is the volume of the solu (mL) and m0 is the initial mass of U in UO2 (mg).

Results and Discussion
In the cases of 5 g/L H2SO4, 0.06 mol/L H2O2 and 0.2 g UO2 at 25 °C, the influen a pyrite addition amount on different systems of H2SO4-H2O2, H2SO4-H2O2-FeS2 wa vestigated. The concentrations of U and Fe under different FeS2 additions are show Figure 1. Without the addition of FeS2, the leaching rate of UO2 was slow. The maxim leaching rate was only 36.88%, with a U concentration of 260.06 mg/L at 360 min. A the addition of 0.1 g FeS2, there was no obvious rate change in the initial stage of the tion; however, the relative rate of the reaction slowly increased after 120 min compare that without FeS2, and the final leaching rate was 48.40% with a U concentration of 34 mg/L. With further increases of the FeS2 mass, the Fe 3+ ion concentration quickly increa resulting in an increase of the U concentration. When the Fe concentration was less 5 mg/L, it had no obvious effect on the UO2 leaching. When the Fe concentration rea 5 mg/L, the reaction rate of U was obviously higher than that without Fe. When th concentration was about 20 mg/L, the reaction rate of UO2 reached the maximum. slope of the U concentration versus time gradually became smaller, indicating tha reaction rate of UO2 gradually slowed down at the end of the experiment, and the leaching rates of U were 64.79%, 76.34% and 79.58%, corresponding to the FeS2 amo of 0.4 g, 0.8 g and 1.2 g, respectively.  Figure 2 shows the fitted curves of the uranium leaching rate and the Eh value o solution at 360 min versus the quantity of FeS2. Generally, the uranium leaching rat creased with the pyrite quantity added; however, the acceleration of the pyrite qua on the leaching rate became small at 360 min. The Eh value decreased with the increa pyrite, which may have been caused by the consumption of hydrogen peroxide in dissolution of pyrite. The decrease of Eh, which was attributed to the decrease in the centration of H2O2, may be the reason why the slope of the uranium leaching rate c slowed down at the end of the experiment.  Figure 2 shows the fitted curves of the uranium leaching rate and the Eh value of the solution at 360 min versus the quantity of FeS 2 . Generally, the uranium leaching rate increased with the pyrite quantity added; however, the acceleration of the pyrite quantity on the leaching rate became small at 360 min. The Eh value decreased with the increase of pyrite, which may have been caused by the consumption of hydrogen peroxide in the dissolution of pyrite. The decrease of Eh, which was attributed to the decrease in the concentration of H 2 O 2 , may be the reason why the slope of the uranium leaching rate curve slowed down at the end of the experiment.
To prove the above speculation, we analyzed the FeS2 particles after leaching under the conditions of 5 g/L H2SO4, 0.2 g UO2 and 0.8 g of FeS2 without H2O2 at 25 °C. As we suspected, neither U (VI) nor Fe 3+ were observed in the solution without the presence of an oxidizer. It shows that FeS2 itself cannot oxidize UO2. Since FeS2 does not contain any oxygen, it cannot be oxidized directly into another species. The only role of FeS2 in this study was to provide a source of Fe 2+ ions (Equation (5)), which oxidize to Fe 3+ (Equation (6)) through the reaction with •OH from the decomposition of H2O2 (Equation (3)). FeS2 promotes the leaching of UO2 only in the presence of an oxidizing agent.

Apparent Activation Energy and Kinetics Model
In the leaching reaction, the particles shrink, and the surface is not covered with other solids, which conforms to the shrinking particle model (SPM) [14]. According to the leaching kinetics model, uranium leaching is controlled by reactant diffusion and/or a surface chemical reaction. SPM was used to fit the leaching data at different temperatures with the kinetic reaction model [20][21][22]. For the following reactions:
To prove the above speculation, we analyzed the FeS 2 particles after leaching under the conditions of 5 g/L H 2 SO 4 , 0.2 g UO 2 and 0.8 g of FeS 2 without H 2 O 2 at 25 • C. As we suspected, neither U (VI) nor Fe 3+ were observed in the solution without the presence of an oxidizer. It shows that FeS 2 itself cannot oxidize UO 2 . Since FeS 2 does not contain any oxygen, it cannot be oxidized directly into another species. The only role of FeS 2 in this study was to provide a source of Fe 2+ ions (Equation (5)), which oxidize to Fe 3+ (Equation (6)) through the reaction with •OH from the decomposition of H 2 O 2 (Equation (3)). FeS 2 promotes the leaching of UO 2 only in the presence of an oxidizing agent.

Apparent Activation Energy and Kinetics Model
In the leaching reaction, the particles shrink, and the surface is not covered with other solids, which conforms to the shrinking particle model (SPM) [14]. According to the leaching kinetics model, uranium leaching is controlled by reactant diffusion and/or a surface chemical reaction. SPM was used to fit the leaching data at different temperatures with the kinetic reaction model [20][21][22]. For the following reactions: aA(fluid) + bB(solid) → Products (8) If the leaching process is mainly determined by diffusion of the reactant inside the solid, the rate expression is: However, if the leaching process is mainly determined by the fluid-solid chemical reaction, then the rate expression becomes: where k is the apparent reaction rate constant, min −1 , and α is the reaction fraction. The reaction fractions of UO 2 under the conditions of 5 g/L H 2 SO 4 , 0.2 g UO 2 , 0.8 g of FeS 2 and 0.06 mol/L H 2 O 2 at different temperatures versus time are fitted in Figure 3. It can be seen that Equation (10) can better fit the experiment data, as the maximum R 2 was 0.97 in the fitting results controlled by diffusion of the reactant inside the solid (i.e., Equation (9) and Figure 3a), but all the values of R 2 were greater than 0.99 in the fitting results controlled by the chemical reaction (i.e., Equation (10) and Figure 3b). Therefore, the leaching process was controlled by the chemical reactions.
If the leaching process is mainly determined by diffusion of the reactant inside solid, the rate expression is: However, if the leaching process is mainly determined by the fluid-solid chem reaction, then the rate expression becomes: where k is the apparent reaction rate constant, min −1 , and α is the reaction fraction.
The reaction fractions of UO2 under the conditions of 5 g/L H2SO4, 0.2 g UO2, 0.8 FeS2 and 0.06 mol/L H2O2 at different temperatures versus time are fitted in Figure  can be seen that Equation (10) can better fit the experiment data, as the maximum R 2 0.97 in the fitting results controlled by diffusion of the reactant inside the solid (i.e., E tion (9) and Figure 3a), but all the values of R 2 were greater than 0.99 in the fitting res controlled by the chemical reaction (i.e., Equation (10) and Figure 3b). Therefore, the le ing process was controlled by the chemical reactions.  The reaction rate constant k obtained from the time-dependent gradients in Figur was substituted into the following Arrhenius equation [10,14]: where A is the pre-index factor; Ea is the apparent activation energy, kJ/mol; T is the t modynamic temperature of the reaction, K and R is the molar gas constant (in J mol −1 · According to the fitting results in Figure 4, we calculated the pre-exponential fa A to be e (9.15 ± 1.56) min −1 and the apparent activation energy Ea to be (3.74 ± 0.40) × 10 kJ/ The evaluated activation energies were lower than 4.86 × 10 kJ/mol calculated by Par al. [23]. The reaction rate constant k obtained from the time-dependent gradients in Figure 3b was substituted into the following Arrhenius equation [10,14]: where A is the pre-index factor; E a is the apparent activation energy, kJ/mol; T is the thermodynamic temperature of the reaction, K and R is the molar gas constant (in J mol −1 ·K −1 ). According to the fitting results in Figure 4, we calculated the pre-exponential factor A to be e (9.15 ± 1.56) min −1 and the apparent activation energy Ea to be (3.74 ± 0.40) × 10 kJ/mol. The evaluated activation energies were lower than 4.86 × 10 kJ/mol calculated by Park et al. [23]. The relationship between the leaching rate constant of UO2 and the quantity of FeS2 in the control stage of the chemical reactions can be expressed as: where km is the rate constant based on the amount of pyrite added. C0 is the constant of the other experimental parameters, m is the mass of FeS2 and p is the reaction order of The relationship between the leaching rate constant of UO 2 and the quantity of FeS 2 in the control stage of the chemical reactions can be expressed as: where k m is the rate constant based on the amount of pyrite added. C 0 is the constant of the other experimental parameters, m is the mass of FeS 2 and p is the reaction order of FeS 2 .
Set A exp − Ea RT C 0 as k , and Equation (12) can then be simplified as Figure 5a displays the 1 − (1 − α) 1/3 versus time with different FeS 2 additions. The slopes of these data are the reaction rate constant of UO 2 , which increases with the increasing amount of the coexistent Fe. This further demonstrates that the addition can accelerate the dissolution of UO 2 . We then plotted the reaction rate constant at different FeS 2 additions in Figure 5b, showing that the influence of a FeS 2 addition on the reaction rate becomes less obvious as the Fe amount increases; this is probably due to the leaching reaction tending towards completion, with an order of the reaction of 0.36. The relationship between the leaching rate constant of UO2 and the quantity o in the control stage of the chemical reactions can be expressed as: where km is the rate constant based on the amount of pyrite added. C0 is the consta the other experimental parameters, m is the mass of FeS2 and p is the reaction ord FeS2.
Set exp − as k', and Equation (12) can then be simplified as = ' Figure 5a displays the 1 − (1 − α) 1/3 versus time with different FeS2 additions slopes of these data are the reaction rate constant of UO2, which increases with the inc ing amount of the coexistent Fe. This further demonstrates that the addition can acce the dissolution of UO2. We then plotted the reaction rate constant at different FeS2 tions in Figure 5b, showing that the influence of a FeS2 addition on the reaction ra comes less obvious as the Fe amount increases; this is probably due to the leaching tion tending towards completion, with an order of the reaction of 0.36.

Conclusions
The experiment of UO2 dissolution in H2SO4-H2O2, H2SO4-H2O2-FeS2 and H2SO was conducted, and the effect of FeS2 on the dissolution of UO2 was investigated.
FeS2 can promote the dissolution of UO2 well in the presence of H2O2. When Fe mg/L, the promoting effect of Fe 3+ can be observed. With the increase of the concentr

Conclusions
The experiment of UO 2 dissolution in H 2 SO 4 -H 2 O 2 , H 2 SO 4 -H 2 O 2 -FeS 2 and H 2 SO 4 -FeS 2 was conducted, and the effect of FeS 2 on the dissolution of UO 2 was investigated.
FeS 2 can promote the dissolution of UO 2 well in the presence of H 2 O 2 . When Fe 3+ is 5 mg/L, the promoting effect of Fe 3+ can be observed. With the increase of the concentration of Fe 3+ in the solution, the promoting effect on the dissolution of UO 2 will be more obvious. When Fe 3+ was 20 mg/L, the reaction rate of UO 2 reached the maximum, and any further increase of the Fe 3+ concentration could not increase the reaction rate of UO 2 . When the mass of FeS 2 increased from 0 g to 1.5 g, the uranium leaching rate increased by 45.7%. However, FeS 2 cannot promote the dissolution of UO 2 in the absence of an oxidant.
The dissolution of UO 2 was controlled by a chemical reaction, and the apparent activation energy Ea was (3.74 ± 0.40) × 10 kJ/mol. The leaching process followed the shrinking particle model controlled by the chemical reaction. The reaction order of FeS 2 was 0.36 at 25 • C.

Data Availability Statement:
The data presented in this study are available on request from the corresponding author.

Conflicts of Interest:
The authors declare no conflict of interest.