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

Abstract

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.

1. Introduction

Sandstone uranium deposits are of great economic value and strategic significance, and they have been the major object of uranium mining in China nowadays [1], due to their characteristics of shallow burial, large reserves, suitable for in situ leaching and low cost [2,3]. The composition of a sandstone-type uranium ore is complex, including pitchblende (UO₂), coffinite (U(SiO₄)_1−x(OH)_4x, (x < 0.5)), ningyoite (Ca_2−xU_x(PO₄)₂·nH₂O, (x ≤ 1, n = 1, 2)) and autunite (Ca(UO₂)(PO₄)₂·(10–12)H₂O). UO₂ often occurs in symbiosis with pyrite (FeS₂), chalcopyrite (CuFeS₂) and marcasite (FeS₂), among which pyrite (FeS₂) is the most common symbiosis mineral in UO₂ deposits [4].

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 (H₂SO₄) 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 H₂SO₄ and dissolve, while only a small part of U(IV) can dissolve in H₂SO₄ under natural conditions. However, most of the uranium in sandstone uranium deposits is UO₂ [8]. The leaching of UO₂ first requires oxidizing it into the U(VI) redox state, where it can more readily dissolve, as shown in Equation (1) [9]. Hydrogen peroxide (H₂O₂) is commonly used as the oxidant [10].

UO₂ + H₂SO₄ + O ⇌ UO₂SO₄ + H₂O

(1)

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 (H₂O₂) and iron (Fe²⁺/Fe³⁺) on UO₂ dissolution in an oxygen-free batch reactor. The interaction in the absence of UO₂ gave a stoichiometric redox reaction of Fe²⁺ and H₂O₂ when the H₂O₂ and Fe²⁺ were present in equal concentrations; however, the predomination of H₂O₂ resulted in the delayed catalytic decomposition of H₂O₂. With UO₂ present, it either dissolved slowly or precipitated as uranium peroxide (UO₄·nH₂O), depending strongly on the ratio of H₂O₂ and Fe²⁺. Zhou et al. [12] studied the kinetics of uranium dissolution and migration under the action of an acidic solution containing Fe³⁺ and its relationship with Fe³⁺. They found that the uranium oxidized by Fe³⁺ 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 Fe²⁺ and Fe³⁺. When the transformation rate of Fe³⁺ to Fe²⁺ 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 Fe³⁺, v(Fe³⁺), in the solution was shown to follow v(U) = 0.0206 + 0.0429 exp [−v(Fe³⁺)/5.07]. Filippov [13] studied the manganese dioxide oxidation of UO₂ 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 Fe³⁺, the dissolution percentage of UO₂ and the redox potential rise sharply, which proves that iron ion plays a catalytic role in the process of oxidizing UO₂. 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 FeS₂ coexistence in sandstone uranium ore on the leaching of UO₂ under the conditions of a strong acid and oxidant in acid in situ leaching mining. Therefore, this study explores the influence of FeS₂ on the leaching of U from FeS₂ to obtain the basic kinetics of H₂SO₄.

2. Experiment

2.1. Materials

Both UO₂ and FeS₂ were powders below 74 μm (passed through a 74-μm sieve) from 272 Uranium Industry Co. Ltd., China National Nuclear Corporation. H₂SO₄ and H₂O₂ (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.

2.2. Methods

A 250-mL mixed solution of 5 g/L H₂SO₄ and 0.06 mol/L H₂O₂ 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 ℃). Different amounts of FeS₂ (0.1, 0.4, 0.8 and 1.2 g) and 0.2 g UO₂ 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 Fe²⁺ concentrations were analyzed by a UV spectrophotometer (Beifen-Ruili, Beijing, China); the difference between them was the concentration of Fe³⁺. 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,

$$\mathsf{\eta }=\frac{cV}{m_0}\times 100\%$$

(2)

where c is the concentration of U in the solution (mg/L), V is the volume of the solution (mL) and m₀ is the initial mass of U in UO₂ (mg).

3. 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 influence of a pyrite addition amount on different systems of H₂SO₄-H₂O₂, H₂SO₄-H₂O₂-FeS₂ was investigated. The concentrations of U and Fe under different FeS₂ additions are shown in Figure 1. Without the addition of FeS₂, the leaching rate of UO₂ 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 FeS₂, 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 FeS₂, and the final leaching rate was 48.40% with a U concentration of 341.31 mg/L. With further increases of the FeS₂ mass, the Fe³⁺ 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 UO₂ 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 UO₂ reached the maximum. The slope of the U concentration versus time gradually became smaller, indicating that the reaction rate of UO₂ 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 FeS₂ amounts 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 of the solution at 360 min versus the quantity of FeS₂. 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 H₂O₂, may be the reason why the slope of the uranium leaching rate curve slowed down at the end of the experiment.

4. Reaction Mechanism

The following reactions may occur in the reaction system [16,17,18,19]:

H₂O₂ ⇌ •OH + •OH

(3)

UO₂ + 2•OH+ H₂SO₄ ⇌ UO₂²⁺ + 2H₂O + SO₄²⁻

(4)

FeS₂ + 14•OH ⇌ Fe²⁺ + 2SO₄²⁻ + 6H₂O + 2H⁺

(5)

Fe²⁺ + •OH + H⁺ ⇌ Fe³⁺ + H₂O

(6)

UO₂ + 2Fe³⁺ ⇌ UO₂²⁺ + 2Fe²⁺

(7)

Hydrogen peroxide first dissociates into •OH (Equation (3)) and participates in the reactions. In the absence of FeS₂, part of UO₂ can be oxidized directly by •OH (Equation (4)). With the addition of FeS₂, Fe plays an intermediary role in the leaching of UO₂, as shown in Equations (5)–(7). Here, •OH first oxidizes FeS₂ into 2SO₄²⁻ and Fe²⁺ (Equation (5)) and then Fe²⁺ to Fe³⁺ (Equation (6)), which, in turn, oxidizes UO₂ into the more soluble uranyl (UO₂²⁺) ion (Equation (7)), resulting in U leaching into the surrounding solution. The oxidation of Fe²⁺ by •OH (Equation (6)) regenerates the Fe³⁺ ion concentration, enabling the further leaching of U through Equation (7).

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 Fe³⁺ 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 Fe²⁺ ions (Equation (5)), which oxidize to Fe³⁺ (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.

5. 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:

$$1-2\mathsf{\alpha }/3-{\left(1-\mathsf{\alpha }\right)}^{2/3}=kt$$

(9)

However, if the leaching process is mainly determined by the fluid–solid chemical reaction, then the rate expression becomes:

$$1-{\left(1-\mathsf{\alpha }\right)}^{1/3}=kt$$

(10)

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 g of FeS₂ and 0.06 mol/L H₂O₂ 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 R² 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 R² 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.

The reaction rate constant k obtained from the time-dependent gradients in Figure 3b was substituted into the following Arrhenius equation [10,14]:

$$\ln k=\ln A-\frac{Ea}{\mathrm{RT}}$$

(11)

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:

$$k_m=A\exp \left(-\frac{Ea}{\mathrm{RT}}\right)C_0{m}^p$$

(12)

where k_m 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 FeS₂.

Set $A\exp \left(-\frac{Ea}{\mathrm{RT}}\right)C_0$ as k’, and Equation (12) can then be simplified as

$$k_m=k^{\prime }{m}^p$$

(13)

Figure 5a displays the 1 − (1 − α)^1/3 versus time with different FeS₂ additions. The slopes of these data are the reaction rate constant of UO₂, which increases with the increasing amount of the coexistent Fe. This further demonstrates that the addition can accelerate the dissolution of UO₂. We then plotted the reaction rate constant at different FeS₂ additions in Figure 5b, showing that the influence of a FeS₂ 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.

6. Conclusions

The experiment of UO₂ dissolution in H₂SO₄-H₂O₂, H₂SO₄-H₂O₂-FeS₂ and H₂SO₄-FeS₂ 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 Fe³⁺ is 5 mg/L, the promoting effect of Fe³⁺ can be observed. With the increase of the concentration of Fe³⁺ in the solution, the promoting effect on the dissolution of UO₂ will be more obvious. When Fe³⁺ was 20 mg/L, the reaction rate of UO₂ reached the maximum, and any further increase of the Fe³⁺ concentration could not increase the reaction rate of UO₂. When the mass of FeS₂ increased from 0 g to 1.5 g, the uranium leaching rate increased by 45.7%. However, FeS₂ cannot promote the dissolution of UO₂ in the absence of an oxidant.

The dissolution of UO₂ 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 FeS₂ was 0.36 at 25 °C.