Uranium acid mine drainage refers to acidic water released from a uranium mining site using processes like underground mining and in-situ leaching.[1] Underground, the ores are not as reactive due to isolation from atmospheric oxygen and water. When uranium ores are mined, the ores are crushed into a powdery substance, thus increasing surface area to easily extract uranium. The ores, along with nearby rocks, may also contain sulfides. Once exposed to the atmosphere, the powdered tailings react with atmospheric oxygen and water. After uranium extraction, sulfide minerals in uranium tailings facilitates the release of uranium radionuclides into the environment, which can undergo further radioactive decay while lowering the pH of a solution.
Uranium may exist naturally as U+6 in ores but also forms the water-soluble uranyl ion UO2+2 when uranium tailings are oxidized by atmospheric oxygen in the following reaction.[2]
U+6 + O2 → UO2+2
The solubility of uranium increases under similar oxidizing conditions when it forms uranyl carbonate complexes in the following reaction.
U+6 + O2 + 2CO32−→ [UO<sub>2</sub>(CO<sub>3</sub>)<sub>2</sub>]2+
Extraction of uranium from the ore may occur under acid or alkaline leaching processes using sulfuric acid and sodium carbonate respectively. If leached with sulfuric acid, uranyl forms a soluble uranyl sulfate complex in the following reaction. Hydrogen ions in solution react with water to produce hydronium ions which lowers a solution's pH making it more acidic.
UO2 + 3H2SO4 + 1/2 O2 → [UO<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>]4− + H2O + 4H+
H+(aq) + H2O(l) → H3O+(aq)
During in-situ leaching uranyl reacts with iron, a common natural oxidant, to produce uranyl trioxide which is further oxidized then leached using alkaline sodium carbonate in the following reactions.
UO2 + 2Fe3+ → UO2+2 + 2Fe2+
UO2 + 1/2 O2 → UO3
UO3 + 3Na2CO3 + H2O → [UO<sub>2</sub>(CO<sub>3</sub>)<sub>3</sub>]4+ + 4Na+ + 2NaOH
When considering the formation secondary uranium minerals, as discussed in the case study section below, the pH of the solution that contains uranophane is one of determining factors of how much of the uranophane is in mineral form or in the form of its ions. Shown in figure 2, from a study performed by Tatiana Shvareva et al. in 2011, is the dissolution of uranophane in pH of 3 (Figure 3b) and pH of 4 (Figure 3a). The graphs demonstrate that in a more acidic environment, the concentrations of Ca, U, and Si are more likely to be more abundant in more basic environments where it is more likely that they will form minerals.[3] This is more likely to happen when the acidic mine drainage is released into rivers or large water deposits and they become diluted to a pH closer to that of water.[4]
The enthalpies of formation (from elements and from oxide species) and Gibbs free energies of formation (from elements) of the uranium minerals boltwoodite, Na-boltwoodite, and uranophane are shown in Table 1. Solubility constants (dissociation of minerals to ions) of the same minerals, determined using a bomb calorimeter in a study by Shvareva, Tatiana et al. in 2011, are shown in Table 2. The Gibbs free energies of formation show that the process, when the reactions from the individual elements to the oxides are taken into account, is spontaneous. The enthalpies of formation, when only considering the reaction from the oxides to the mineral, suggest a relatively high probability for their Gibbs free energy of formation values to also be spontaneous.