Article pubs.acs.org/IECR
Solubility of Sodium Chloride in Ionic Liquids Fatemeh Saadat Ghareh Bagh,† Farouq S. Mjalli,*,‡ Mohd Ali Hashim,† Mohamed Kamel Omar Hadj-Kali,§ and Inas M. AlNashef§ †
Chemical Engineering Department, University of Malaya, Kuala Lumpur, Malaysia Petroleum and Chemical Engineering Department, Sultan Qaboos University, Muscat, Oman § Chemical Engineering Department, King Saud University, Riyadh, Saudi Arabia ‡
S Supporting Information *
ABSTRACT: In this study, the solubility of sodium chloride in 16 imidazolium-, pyrrolidinium-, pyridinium-, and ammoniumbased ionic liquids at different temperatures was measured. The solubility of sodium chloride was found to increase with temperature in all of the ionic liquids studied. Both cations and anions in the ionic liquid affected solubility, but anions had more of an effect than cations. The highest measured solubility of sodium chloride was 8.96 wt % at 125 °C in 1-ethyl-3methylimidazolium dimethyl phosphate. The nonrandom two-liquid model was used to calculate the solubility of sodium chloride in some ionic liquids. The experimental and calculated values were in good agreement in many cases.
1. INTRODUCTION Alkali metals are highly reactive elements and are not found in nature in their elemental forms. Reducing agents, such as hydrogen, do not reduce alkali metals from their compounds to a metallic state. This is because the alkali metal bonds are too strong to be broken by the reducing agents. In industry, electrical reduction is the most common method for producing sodium and lithium from their salts. Recently, fundamental physicochemical studies have been conducted on electrolytes using nonaqueous organic solvents for alkali metal chlorides to be used in battery applications. Therefore, a safer and more economical electrolytic process must be developed for these applications. The controllability and operability of the process can also be improved.1−4 Sodium metal is an essential alkali metal with broad technical and commercial applications. It is used as a heat-transfer medium in atomic reactors; in the manufacturing of zirconium (Zr), titanium (Ti), sodium cyanide (NaCN), sodium peroxide (Na2O2), and sodium hydride (NaH); and in the manufacturing of batteries. Demand for a power source for electric cars and consumer electronics has stimulated the search for secondary rechargeable batteries. The main criteria for such a secondary battery are high-energy density, low cost, material availability, cyclic life, acceptable temperature range, and safety.5−7 A Downs cell is used to extract sodium metal electrolytically from molten sodium chloride (NaCl). NaCl has a high melting temperature of approximately 800 °C. Thus, a mixture of calcium chloride (CaCl2) and barium chloride (BaCl2) is added to molten NaCl to decrease the operating temperature to approximately 600 °C. A circular iron cathode in the Downs cell surrounds the carbon anode. The electrolysis products, namely, sodium metal and chlorine gas, are separated by a steel mesh or a gauze to prevent them from contacting and reforming NaCl.2,8 The high operating temperature necessitates the concentric cylindrical cell design of the Downs process, producing a cell © 2013 American Chemical Society
with poor space efficiency. This results in high operating costs per production unit. The electrolyte mixture exhibits a melting temperature of approximately 580 °C, making smooth operation of the cell difficult. This causes frequent cell freezeups and other setbacks, increasing the operating and labor costs of the industrial electrolytic process. Therefore, the Downs process is not suitable for automation.2,8 Several concepts have been proposed to reduce emissions and energy consumption, but none have been successfully applied on an industrial scale. Developing an electrolytic process that can be used to produce an alkali metal more economically is increasingly important. The operability of the process must also be improved, making automation possible. Suitable processing techniques are limited because sodium has a strong affinity to oxygen and water and cannot be electrolyzed in aqueous solutions because of its negative reduction potential. All processes must also comply with existing environmental regulations. Therefore, a novel process that electrolyzes sodium at or near ambient temperatures has many industrial applications.2,8,9 The need for low-cost, dependable, high-energy-density, and rechargeable batteries continues to escalate with increasing demand from the electronics and automobile industries. A sodium metal anode has been considered as an alternative to a lithium-based battery because it provides potential advantages. Using sodium in batteries requires an electrolyte that has a low melting temperature, is ionically conductive, and exhibits a broad electrochemical window to permit sodium deposition and a corresponding cathode reaction.10 Ionic liquids (ILs) are salts that exist in liquid form at ambient temperatures (below 100 °C). They are widely used in industry because of their exceptional properties, such as high Received: Revised: Accepted: Published: 11488
April 23, 2013 July 20, 2013 July 29, 2013 July 29, 2013 dx.doi.org/10.1021/ie401282y | Ind. Eng. Chem. Res. 2013, 52, 11488−11493
Industrial & Engineering Chemistry Research
Article
thermal stability to 350 °C, negligible flammability, high conductivity, and negligible vapor pressure. These specific properties make ILs attractive for industrial applications. Replacing conventional solvents with ILs would prevent emissions of volatile organic compounds, a major source of environmental pollution. ILs are not intrinsically environmentally friendly because some are extremely toxic, but they can be designed to be environmentally friendly with numerous potential benefits for sustainable chemistry.11,12 They could be used to replace volatile organic compounds, especially halogenated organic solvents, in certain systems. They have been examined as potential solvents or carriers of metal ions, such as Co2+, Pd2+, Zn2+, Cu2+, and Fe3+, and for use in liquid− liquid extraction. Although they might cause negative environmental effects, such as possible hydrolysis and toxic hydrogen fluoride formation or partial loss to the aqueous phase, they are still interesting and important for metal processing applications. The extraction mechanism with ILs is usually different from and more complex than that with conventional solvents. ILs favor extraction of charged species; therefore, most metal ions are transported to the IL phase based on a cation- or anionexchange mechanism. This is because of the unique ionic solvation environment and the ion-exchange capabilities of ILs influencing their specific extraction behavior. ILs are widely used as solvents in organic reactions, but they are more frequently used to separate metal ions in extraction, membrane, and adsorption systems.13 ILs have been used successfully as green solvents in metal extraction and separation because of their ionic structure. Studies have attempted to use ILs as efficient electrolytes in the Downs process to decrease the operating temperature from 600 to 120 °C. Imidazolium-based ILs are the most commonly used ILs in various metal ion systems.13,14 Ammonium and phosphonium ILs have been used in extraction processes and are considered prospective solvents and carriers in separation techniques.15−17 A study on Li+, Na+, Cs+, Ca2+, Sr2+, and La3+ extraction confirmed that imidazolium ILs with nonafluorobutanesulfonate ions ([NfO−]) exhibit good extraction abilities.13 The study showed that metal ions with more charge are more easily transferred to the [Cnmim][NfO] phase. The IL anion is the most important factor affecting the solubility of metal ions in ILs. For example, imide ILs exhibit the lowest NaCl solubility: less than 0.1 g per 100 g. ILs with methylsulfate anions exhibit the highest solubility. Dimethylimidazolium dimethylphosphate exhibits a remarkable NaCl solubility of 12 g per 100 g, corresponding to a NaCl mole fraction of 0.33. The substituents on the imidazolium cation can affect the solubility of NaCl in ILs. For example, replacing the ethyl group with a butyl group for the 1-ethyl-3-methylimidazolium octylsulfate IL was found to change its solubility from 0.6 to 2.12 g per 100 g. However, in other cases, this effect is not significant.18 Imidazolium-based ILs are better solvents than pyridinium-, ammonium-, sulfonium-, imide-, or phosphonium-based ILs.18 The literature shows that ILs exhibit high electrical conductivities because of their ionic structure. This is a desirable characteristic for electrolytes used in electrochemical processes. This study investigated the ability of ILs to serve as electrolytes for producing sodium metal from an inexpensive source, namely, NaCl. The solubility of NaCl was measured in 16 imidazolium-, pyrrolidinium-, pyridinium-, and ammoniumbased ILs at different temperatures. Table 1S (Supporting Information) provides a summary of the ILs used in this study.
For simplicity, each IL is abbreviated in the table. These ILs were selected because the literature indicated that they exhibit a good negative electrochemical limit, making them potentially suitable for sodium reduction.13 The ability of an IL to serve as an electrolyte is reflected by a high NaCl solubility in that IL. ILs with higher solubilities are more suited to being electrolytes. This study evaluated the solubility of NaCl as a common and inexpensive source of sodium metal in various ILs at different temperature ranges. Another study that measures the electrical conductivities of the solutions and examines the stability of sodium metal in the ILs will be presented in the future.
2. EXPERIMENTS Materials. All of the ILs and the sodium chloride used in this study were supplied by Merck KgaA (Darmstadt, Germany). The manufacturer’s guide stated that the purity of the chemicals was >99%; thus, they were used without further purification. Solubility Measurements. Measuring the solubility of salts in solvents is an important laboratory task. A reliable solubility test requires a pure solvent and solute, reliable and satisfactory separation of the sample from the saturated solution with the undissolved solute, and adequate temperature control. The shake-flask method is usually used to determine equilibrium solubility at a given temperature. According to this method, a compound is added to a solvent and shaken for 24 h or longer. Saturation is confirmed by observing the presence of undissolved material, and saturation is reached if the solvent and excess solute are heated and allowed to cool to a certain temperature. Solutions that hold excess solute in the solvent are called supersaturated solutions. Supersaturation of salts can be avoided by slow cooling and continuous shaking of the sample during cooling. After the slurry has been separated, the sample can be analyzed. The filtration must be performed at the same temperature as the solubility determination and under conditions that minimize the loss of volatile compounds. The sample is usually diluted to prevent crystallization. The shake-flask method used in this study is the most accurate method of determining solubility, but it is timeconsuming. Before measuring the solubility of the salt, the solvent (IL) was dried in a vacuum at 65 °C for a minimum of 24 h to remove any water contaminant present in the IL. A vacuum oven (MTI, DZF-6050) was used for this task. After the vacuum drying, some of the ILs were analyzed for their water content using a C30 Compact Karl Fischer Coulometer. The analysis results showed that the water contents were no more than 0.01% by weight. The solubility of NaCl was determined by placing 0.1 g of the salt (dried) in 5 mL of the IL and stirring vigorously for a minimum of 48 h at a specific temperature. A magnetic hot-plate stirrer (IKA C-MAG HS7 S2 hot-plate stirrer) was used, and the plate temperature was controlled. Different mixing temperatures were used for different measurements. The solubility test was conducted in a glovebox environment (Innovative Technology, Amesbury, MA, model 2GB) with humidity maintained at less than 4 ppm. The IL must reach a saturation state, meaning that it cannot dissolve more NaCl. This was determined visually by adding more NaCl to the solution and allowing time for the solvent to dissolve it. If the NaCl was left suspended in the solution and seen with the naked eye, this indicated that the solvent was saturated. 11489
dx.doi.org/10.1021/ie401282y | Ind. Eng. Chem. Res. 2013, 52, 11488−11493
Industrial & Engineering Chemistry Research
Article
When saturation was reached, the solution was filtered at the same temperature as the solution; that is, if the solution temperature was 110 °C, the filtration temperature was also 110 °C. The filtrate was then diluted with deionized water and analyzed by inductively coupled plasma-atomic emission spectrometry (ICP-AES). Each analysis was repeated three times, and the average was recorded. The estimated relative uncertainties of ICP measurements are 5%, 10%, and >10% for the ranges 1−10, 0.1−1, and
