Lead–acid batteries are important to modern society because of their wide usage and low cost. The primary source for production of new lead–acid batteries is from recycling spent lead–acid batteries. In spent lead–acid batteries, lead is primarily present as lead pastes. In lead pastes, the dominant component is lead sulfate (PbSO4, mineral name anglesite) and lead oxide sulfate (PbO•PbSO4, mineral name lanarkite), which accounts for more than 60% of lead pastes. In the recycling process for lead–acid batteries, the desulphurization of lead sulfate is the key part to the overall process. In this work, the thermodynamic constraints for desulphurization via the hydrometallurgical route for recycling lead pastes are presented. The thermodynamic constraints are established according to the thermodynamic model that is applicable and important to recycling of lead pastes via hydrometallurgical routes in high ionic strength solutions that are expected to be in industrial processes. The thermodynamic database is based on the Pitzer equations for calculations of activity coefficients of aqueous species. The desulphurization of lead sulfates represented by PbSO4 can be achieved through the following routes. (1) conversion to lead oxalate in oxalate-bearing solutions; (2) conversion to lead monoxide in alkaline solutions; and (3) conversion to lead carbonate in carbonate solutions. Among the above three routes, the conversion to lead oxalate is environmentally friendly and has a strong thermodynamic driving force. Oxalate-bearing solutions such as oxalic acid and potassium oxalate solutions will provide high activities of oxalate that are many orders of magnitude higher than those required for conversion of anglesite or lanarkite to lead oxalate, in accordance with the thermodynamic model established for the oxalate system. An additional advantage of the oxalate conversion route is that no additional reductant is needed to reduce lead dioxide to lead oxide or lead sulfate, as there is a strong thermodynamic force to convert lead dioxide directly to lead oxalate. As lanarkite is an important sulfate-bearing phase in lead pastes, this study evaluates the solubility constant for lanarkite regarding the following reaction, based on the solubility data, PbO•PbSO4 + 2H+ ⇌ 2Pb2+ + SO42– + H2O(l).
This report summarizes the 2021 fiscal year (FY21) status of ongoing borehole heater tests in salt funded by the disposal research and development (R&D) program of the Office of Spent Fuel & Waste Science and Technology (SFWST) of the US Department of Energy’s Office of Nuclear Energy’s (DOE-NE) Office of Spent Fuel and Waste Disposition (SFWD). This report satisfies SFWST milestone M2SF- 21SN010303052 by summarizing test activities and data collected during FY21. The Brine Availability Test in Salt (BATS) is fielded in a pair of similar arrays of horizontal boreholes in an experimental area at the Waste Isolation Pilot Plant (WIPP). One array is heated, the other unheated. Each array consists of 14 boreholes, including a central borehole with gas circulation to measure water production, a cement seal exposure test, thermocouples to measure temperature, electrodes to infer resistivity, a packer-isolated borehole to add tracers, fiber optics to measure temperature and strain, and piezoelectric transducers to measure acoustic emissions. The key new data collected during FY21 include a series of gas tracer tests (BATS phase 1b), a pair of liquid tracer tests (BATS phase 1c), and data collected under ambient conditions (including a period with limited access due to the ongoing pandemic) since BATS phase 1a in 2020. A comparison of heated and unheated gas tracer test results clearly shows a decrease in permeability of the salt upon heating (i.e., thermal expansion closes fractures, which reduces permeability).
Actinide oxalates are chemical compounds important to nuclear industry, ranging from actinide separation in waste reprocessing, to production of specialty actinides, and to disposal of high level nuclear waste (HLW) and spent nuclear fuel (SNF). In this study, the solubility constants for Pr2(C2O4)3·10H2O and Nd2(C2O4)3·10H2O by performing solubility experiments in HNO3 and mixtures of HNO3 and H2C2O4 at 23.0 ± 0.2 °C have been determined. The targeted starting materials, Pr2(C2O4)3·10H2O and Nd2(C2O4)3·10H2O, were successfully synthesized at room temperature using PrCl3, NdCl3 and oxalic acid as the source metrials. Then, we utilized the targeted solubility-controlling phases to conduct solubility measurements. There was no phase change over the entire periods of experiments, demonstrating that Pr2(C2O4)3·10H2O and Nd2(C2O4)3·10H2O were the solubility-controlling phases in our respective experiments. Based on our experimental data, we have developed a thermodynamic model for Pr2(C2O4)3·10H2O and Nd2(C2O4)3·10H2O in the mixtures of HNO3 and H2C2O4 to high ionic strengths. The model for Pr2(C2O4)3·10H2O reproduces well the reported experimental data for Pu2(C2O4)3·10H2O, which are not utilized for the model development, demonstrating that Pr(III) is an excellent analog for Pu(III). Similarly, the model for Nd2(C2O4)3·10H2O reproduces the solubility of Am2(C2O4)3·10H2O and Cm2(C2O4)3·10H2O. The Pitzer model was used for the calculation of activity coefficients. Based on the published, well established model for dissociation constants for oxalic acid and stability constants for actinide-oxalate complexes [i.e., AmC2O4+, and Am(C2O4)2−] to high ionic strengths, we have obtained the solubility constants (log10K0) for the following reactions at 25 °C,Pr2(C2O4)3·10H2O ⇌ 2Pr3+ + 3C2O42− + 10H2O(l)Nd2(C2O4)3·10H2O ⇌ 2Nd3+ + 3C2O42− + 10H2O(l) to be −30.82 ± 0.30 (2σ), and −31.14 ± 0.35 (2σ), respectively. These values can be directly applied to Pu2(C2O4)3·10H2O, Am2(C2O4)3·10H2O and Cm2(C2O4)3·10H2O. The model established for actinide oxalates by this study provides the needed knowledge with regard to solubilities of actinide/REE oxalates at various ionic strengths, and is expected to find applications in many fields, including the geological disposal of nuclear waste and the mobility of REE under the surface conditions, as Pr2(C2O4)3·10H2O and Nd2(C2O4)3·10H2O can be regarded as the pure Pr and Nd end-members of deveroite, a recently discovered natural REE oxalate with the following stoichiometry, (Ce1.01Nd0.33La0.32Pr0.11Y0.11Sm0.01Pb0.04U0.03Th0.01Ca0.04)2.01(C2O4)2.99·9.99H2O. Regarding its importance in the geological disposal of nuclear waste, Am2(C2O4)3·10H2O/Pu2(C2O4)3·10H2O/Cm2(C2O4)3·10H2O can be the source-term phase for actinides, as demonstrated by the instance in the disposal in clay/shale formations. This is exemplified by the stability of Am2(C2O4)3·10H2O in comparison with Am(OH)3(am), Am(OH)3(s) and AmCO3(OH)(s) under the relevant geological repository conditions.
In this study, I present experimental results on the equilibrium between boracite [Mg3B7O13Cl(cr)] and kurnakovite [chemical formula, Mg2B6O11.15H2O(cr); structural formula, MgB3O3(OH)5.5H2O(cr)] at 22.5 ± 0.5 °C from a long-term experiment up to 1629 days, approaching equilibrium from the direction of supersaturation, Mg3B7O13Cl(cr) + H+ + 2B(OH)4 + 18H2O(1) . 3MgB3O3(OH)5.5H2O(cr) + Cl . Based on solubility measurements, the 10-based logarithm of the equilibrium constant for the above reaction at 25 °C is determined to be 12.83 ± 0.08 (2s). Based on the equilibrium constant for dissolution of boracite, Mg3B7O13Cl(cr) + 15H2O(l) = 3Mg2+ + 7B(OH)4 + Cl + 2H+ at 25 °C measured previously (Xiong et al. 2018) and that for the reaction between boracite and kurnakovite determined here, the equilibrium constant for dissolution of kurnakovite, MgB3O3(OH)5.5H2O(cr) = Mg2+ + 3B(OH)4 + H+ + H2O(1) is derived as 14.11 ± 0.40 (2s). Using the equilibrium constant for dissolution of kurnakovite obtained in this study and the experimental enthalpy of formation for kurnakovite from the literature, a set of thermodynamic properties for kurnakovite at 25 °C and 1 bar is recommended as follows: ΔH0f = 4813.24 ± 4.92 kJ/mol, .G0f = 4232.0 ± 2.3 kJ/mol, and S0 = 414.3 ± 0.9 J/(mol.K). Among them, the Gibbs energy of formation is based on the equilibrium constant for kurnakovite determined in this study; the enthalpy of formation is from the literature (Li et al. 1997), and the standard entropy is calculated internally with the Gibbs-Helmholtz equation in this work. The thermodynamic properties of kurnakovite estimated using the group contribution method for borate minerals based on the sums of contributions from the cations, borate polyanions, and structural water to the thermodynamic properties from the literature (Li et al. 2000) are consistent, within their uncertainties, with the values listed above. Since kurnakovite usually forms in salt lakes rich in sulfate, studying the interactions of borate with sulfate is important to modeling kurnakovite in salt lakes. For this purpose, I have re-calibrated our previous model (Xiong et al. 2013) describing the interactions of borate with sulfate based on the new solubility data for borax in Na2SO4 solutions presented here.
The US Department of Energy Office of Nuclear Energy is conducting a brine availability heater test to characterize the thermal, mechanical, hydrological and chemical response of salt at elevated temperatures. In the heater test, brines will be collected and analyzed for chemical compositions. In order to support the geochemical modeling of chemical evolutions of the brines during the heater test, we are recalibrating and validating the solubility models for the mineral constituents in salt formations up to 100°C, based on the solubility data in multiple component systems as well as simple systems from literature. In this work, we systematically compare the model-predicted values based on the various solubility models related to the constituents of salt formations, with the experimental data. As halite is the dominant constituent in salt formations, we first test the halite solubility model in the Na-Mg-Cl dominated brines. We find the existing halite solubility model systematically over-predict the solubility of halite. We recalibrate the halite model, which can reproduce halite solubilities in Na-Mg-Cl dominated brines well. As gypsum/anhydrite in salt formations controls the sulfate concentrations in associated brines, we test the gypsum solubility model in NaCl solutions up to 5.87 mol•kg-1 from 25°C to 50°C. The testing shows that the current gypsum solubility model reproduces the experimental data well when NaCl concentrations are less than 1 mol•kg-1. However, at NaCl concentrations higher than 1, the model systematically overpredicts the solubility of gypsum. In the Na - Cl - SO4 - CO3 system, the validation tests up to 100°C demonstrate that the model excellently reproduces the experimental data for the solution compositions equilibrated with one single phase such as halite (NaCl) or thenardite (Na2SO4), with deviations equal to, or less than, 1.5 %. The model is much less ideal in reproducing the compositions in equilibrium with the assemblages of halite + thenardite, and of halite + thermonatrite (Na2CO3•H2O), with deviations up to 31 %. The high deviations from the experimental data for the multiple assemblages in this system at elevated temperatures may be attributed to the facts that the database has the Pitzer interaction parameters for Cl - CO3 and SO4 - CO3 only at 25°C. In the Na - Ca - SO4 - HCO3 system, the validation tests also demonstrate that the model reproduces the equilibrium compositions for one single phase such as gypsum better than the assemblages of more than one phase.
Montmorillonite with an empirical formula of Na0.2Ca0.1Al2Si4O10(OH)2(H2O)10 is a di-octahedral smectite. Montmorillonite-rich bentonite is a primary buffer candidate for high level nuclear waste (HLW) and used nuclear fuel to be disposed in mild environments. In such environments, temperatures are expected to be ≤ 90oC, the solutions are of low ionic strengths, and pH is close to neutral. Under the conditions outside the above parameters, the performance of montmorillonite-rich bentonite is deteriorated because of collapse of swelling particles as a result of illitization, and dissolution of the swelling clay minerals followed by precipitation of non-swelling minerals. It has been well known that tri-octahedral smectites such as saponite, with an ideal formula of Mg3(Si, Al)4O10(OH)2•4H2O for an Mg-end member (saponite-15A), are less susceptible to alteration under harsh conditions. Recently, Mg-bearing saponite has been favorably considered as a preferable engineered buffer material for the Swedish very deep holes (VDH) disposal concept in crystalline rock formations. In the VDH, HLW is disposed in deep holes at depth between 2,000 m and 4,000 m. At such deployment depths, the temperatures are expected to be between 100oC and 150oC, and the groundwater is of high ionic strength. The harsh chemical conditions of high pH are also introduced by the repository designs in which concretes and cements are used as plugs and buffers. In addition, harsh chemical conditions introduced by high ionic strength solutions are also present in repository designs in salt formations and sedimentary basins. For instance, the two brines associated with the salt formations for the Waste Isolation Pilot Plant (WIPP) in USA have ionic strengths of 5.82 mol•kg-1 (ERDA-6) and 8.26 mol•kg-1 (GWB). In the Asse site proposed for a geological repository in salt formations in Germany, the Q-brine has an ionic strength of ~13 mol•kg-1. In this work, we present our investigations regarding the stability of saponite under hydrothermal conditions in harsh environments.
Methane (CH4) and carbon dioxide (CO2), the two major components generated from kerogen maturation, are stored dominantly in nanometer-sized pores in shale matrix as (1) a compressed gas, (2) an adsorbed surface species and/or (3) a species dissolved in pore water (H2O). In addition, supercritical CO2 has been proposed as a fracturing fluid for simultaneous enhanced oil/gas recovery (EOR) and carbon sequestration. A mechanistic understanding of CH4-CO2-H2O interactions in shale nanopores is critical for designing effective operational processes. Using molecular simulations, we show that kerogen preferentially retains CO2 over CH4 and that the majority of CO2 either generated during kerogen maturation or injected in EOR will remain trapped in the kerogen matrix. The trapped CO2 may be released only if the reservoir pressure drops below the supercritical CO2 pressure. When water is present in the kerogen matrix, it may block CH4 release. However, the addition of CO2 may enhance CH4 release because CO2 can diffuse through water and exchange for adsorbed methane in the kerogen nanopores.
In this paper, a solubility study on brucite [Mg(OH)2(cr)] in Na2SO4 solutions ranging from 0.01 to 1.8 mol•kg–1, with 0.001 mol•kg–1 borate, has been conducted at 22.5°C. Based on the solubility data, the Pitzer interaction parameters for MgB(OH)4+—SO42– and MgB(OH)4+—Na+ along with the formation constant for MgSO4(aq) are evaluated using the Pitzer model. The formation constant (log10β10= 2.38 ± 0.08) for MgSO4(aq) at 25°C and infinite dilution obtained in this study is in excellent agreement with the literature values. The experimental data on the solubility of gypsum (CaSO4•2H2O), at 25°C, in aqueous solutions of MgSO4 with ionic strengths up to ~11 mol•kg–1 were analyzed using models with and without considering the MgSO4(aq) species. The model incorporating MgSO4(aq) fits better to the experimental data than the model without MgSO4(aq), especially in the ionic strength range beyond ~4 mol•kg–1, demonstrating the need for incorporation of MgSO4(aq) into the model to improve the accuracy.
In this work, solubility measurements regarding boracite [Mg3B7O13Cl(cr)] and aksaite [MgB6O7(OH)6·2H2O(cr)] from the direction of supersaturation were conducted at 22.5 ± 0.5 °C. The equilibrium constant (log10K0) for boracite in terms of the following reaction, Mg3B7O13Cl(cr) + 15H2O(l) ⇌ 3Mg2+ + 7B(OH)4– + Cl– + 2H+ is determined as -29.49 ± 0.39 (2σ) in this study. The equilibrium constant for aksaite according to the following reaction, MgB6O7(OH)6•2H2O(cr) + 9H2O(l) ⇌ Mg2+ + 6B(OH)4– + 4H+ is determined as -44.41 ± 0.41 (2σ) in this work. This work recommends a set of thermodynamic properties for aksaite at 25 °C and 1 bar as follows: ΔH$0\atop{f}$ =-6063.70 ± 4.85 kJ·mol-1, ΔG =-5492.55 ± 2.32 kJ·mol-1, and S0 = 344.62 ± 1.85 J·mol-1·K-1. Among them, ΔG$0\atop{f}$ is derived from the equilibrium constant for aksaite determined by this study; ΔH$0\atop{f}$ is from the literature, determined by calorimetry; and S0 is computed in the present work from ΔG$0\atop{f}$ and ΔH$0\atop{f}$. This investigation also recommends a set of thermodynamic properties for boracite at 25 °C and 1 bar as follows: ΔH$0\atop{f}$ =-6575.02 ± 2.25 kJ·mol-1, ΔG$0\atop{f}$ =-6178.35 ± 2.25 kJ·mol-1, and S0 = 253.6 ± 0.5 J·mol-1·K-1. Among them, ΔG$0\atop{f}$ is derived from the equilibrium constant for boracite determined by this study; S0 is from the literature, determined by calorimetry; and ΔH$0\atop{f}$ is computed in this work from ΔG$0\atop{f}$ and S0. The thermodynamic properties determined in this study can find applications in many fields. For instance, in the field of material science, boracite has many useful properties including ferroelectric and ferroelastic properties. The equilibrium constant of boracite determined in this work will provide guidance for economic synthesis of boracite in an aqueous medium. Similarly, in the field of nuclear waste management, iodide boracite [Mg3B7O13I(cr)] is proposed as a waste form for radioactive 129I. Therefore, the solubility constant for chloride boracite [Mg3B7O13Cl(cr)] will provide the guidance for the performance of iodide boracite in geological repositories. Boracite/aksaite themselves in geological repositories in salt formations may be solubility-controlling phase(s) for borate. Finally, solubility constants of boracite and aksaite will enable researchers to predict borate concentrations in equilibrium with boracite/aksaite in salt formations.