Ind. Eng. Chem. Res. 2010, 49, 4981–4988
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Solubilities of NaCl, KCl, LiCl, and LiBr in Methanol, Ethanol, Acetone, and Mixed Solvents and Correlation Using the LIQUAC Model Miyi Li,†,‡ Dana Constantinescu,‡ Lisheng Wang,† Andre´ Mohs,‡ and Ju¨rgen Gmehling‡,* School of Chemical Engineering & EnVironment, Beijing Institute of Technology, 100081 Beijing, China, Technische Chemie, Institut fu¨r Reine und Angewandte Chemie, Carl Von Ossietzky UniVersita¨t Oldenburg, D-26111, Oldenburg, Germany
The solubilities of NaCl, KCl, LiCl, and LiBr in pure methanol, ethanol, and acetone were measured over a temperature range from 293.15 to 333.15 K. Furthermore salt solubilities in the mixed solvents (water + methanol, water + ethanol, water + acetone, methanol + ethanol, methanol + acetone, ethanol + acetone) were determined at 313.15 K. For a few systems solubility data are reported for the first time. In a few cases a comparison with published data stored in the Dortmund Data Bank (DDB)1 showed disagreement. The LIQUAC model was used to correlate the experimental data. The calculated salt solubilities are in good agreement with the experimental results for the systems NaCl + water + methanol and KCl + water + methanol. 2. Experimental Section
1. Introduction The knowledge of salt solubilities in pure organic and mixed solvent electrolyte systems is of great importance for the design and simulation of unit operations such as crystallization, liquid-liquid extraction, and other industrial processes.2 The data are also required in connection with theoretical studies concerning the liquid phase structure and its thermodynamic properties. Accurate solubility data are also of great interest for the development of electrolyte models. For the semiempirical LIQUAC3,4 model, a large database was used for optimizing the required parameters. LIQUAC can be used to correlate and predict salt solubilities (SLE), liquid-liquid equilibria (LLE), mean ion activity coefficients, vapor-liquid equilibria (VLE), and osmotic coefficients for electrolyte solutions. Unfortunately, most of the published data are only available for aqueous systems. For pure organic or mixed solvent electrolyte systems the number of available data is much smaller, and often the published data show large scattering. More reliable data are required in order to enlarge the database for fitting the required parameters of electrolyte models.
2.1. Chemicals. Sodium chloride and potassium chloride with purities higher than 99.7% were obtained from VWR international bvba/spr. Lithium chloride and lithium bromide with minimum purities of 99% were supplied by Sigma-Aldrich Inc. Prior to the measurements, the salts were dried in an oven at 433 K for 2 days. Acetone with a purity of 99.98% was supplied by Carl Roth GmbH & Co. Ethanol and methanol with purities greater than 99.8% were supplied by VWR. The organic solvents were not further purified. Doubly distilled water was used for the measurements. 2.2. Apparatus and Procedure. The apparatus used in this work is shown in Figure 1. The experiments were carried out in a jacketed glass cell with a volume of 140 cm3. The temperature of the cell is controlled by circulating water from a temperature-controlled bath. The cell was first loaded with a small excess of salt in the chosen solvent. Then the pure organic
In this work, salt solubilities in pure organic solvents were measured as a function of temperature. Furthermore salt solubilities in mixtures were investigated for different solvent compositions. Four salts (NaCl, KCl, LiCl, and LiBr) in three organic solvents (methanol, ethanol, and acetone) and their binary mixtures were measured over the whole solvent composition range. A well-designed procedure was implemented to measure 7 binary systems and 12 ternary systems. Some of the systems investigated were measured for the first time. The data measured were used to extend the database for modeling work. Finally, the experimental results were compared with the results predicted by the LIQUAC model using the already available parameters. * To whom correspondence should be addressed. Tel.: +49-441798-3831. Fax: +49-441-798-3330. URL: http//www.uni-oldenburg.de/ tchemie. E-mail:
[email protected]. † Beijing Institute of Technology. ‡ Carl von Ossietzky Universita¨t Oldenburg.
Figure 1. The apparatus applied for solubility measurements: (1) thermostatted syringe; (2) digital temperature display; (3) Pt-100 thermometer; (4) magnetic stirring rod; (5) jacketed glass cell; (6) magnetic stirrer; (7) temperature-controlled bath; (8) pump.
10.1021/ie100027c 2010 American Chemical Society Published on Web 04/13/2010
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Table 1. Salt Solubilities (mol · kg-1) in Organic Solvents at Different Temperatures mNaCl
mKCl
mLiCl
mLiBr
T (K)
methanol
methanol
ethanol
methanol
ethanol
acetone
methanol
293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15
0.241 0.238 0.235 0.231 0.225 0.219 0.217 0.212 0.209
0.0708 0.0736 0.0754 0.0780 0.0811 0.0833 0.0858 0.0888 0.0907
0.0069 0.0064 0.0061 0.0059 0.0054 0.0051 0.0050 0.0046 0.0044
10.26 10.28 10.29 10.30 10.30 10.31 10.32 10.33 10.33
5.901 5.840 5.800 5.737 5.707 5.643 5.628 5.587 5.556
0.2776 0.2556 0.2236 0.1931 0.1730 0.1509 0.1240 0.1037
16.39 16.44 16.47 16.53 16.60 16.67 16.76 16.83 16.90
solvents (methanol, ethanol, and acetone) or binary mixture with the desired composition were added. The cell was tightly closed during the measurement and care was taken to ensure that the composition of the mixed solvents was not changed because of evaporation by leaving only 3-5 mL of gas-phase in the cell. Cell and bath temperatures were measured by precision Pt 100 thermometers with an accuracy of (0.01 K. The binary solvent mixtures (water + methanol, water + ethanol, water + acetone, methanol + ethanol, methanol + acetone, and ethanol + acetone) were prepared using a balance with an uncertainty of (0.0001 g (Sartorius A200S). The experimental points for mixed solvents were arranged in 10% steps by varying the salt-free mass fraction. To avoid the formation of microcrystals and supersaturation during the measurements, the solutions were stirred at a speed of around 600 rpm for approximately 12 h in the case of organic solvents or organic solvent mixtures and 6 h for water-organic electrolyte systems. This ensured an intensive contact between the solid and the liquid phase. After sedimentation for 24 h in the case of the organic or mixed organic electrolyte systems and 12 h for water-organic electrolyte systems, three liquid samples of about 3 mL were taken by using a syringe equipped with a 0.45 µm filter and transferred to capped vials with a volume of 15 mL. Prior to sampling, syringe and filter were thermostatted to a temperature 5 K above the temperature of the solution. The mass of the empty vial (W3) and the mass of the sample together with the vial (W1) were determined by using an electric balance (Sartorius CP225D) with an uncertainty of (0.00001 g. The liquid samples were first dried in an oven at 353 K for 2 days, and then at 433 K for at least 24 h. The mass of solid together with the vial (W2) was weighed by the same balance ((0.00001 g). The drying of the samples was continued until a constant mass was reached. The solubility of the salt can then be calculated by the following relation: solubility [mol · kg-1] )
(
)
W2 - W 3 1 · W1 - W2 Msalt
salt from the saturated solution. The solubility of each sample was calculated by eq 1. As solubility the mean value of the three samples was chosen. When the relative standard deviation of one of the samples was greater than 0.5%, the measurement was repeated, whereby the relative standard deviation (RSD) within a set of different experimental results was defined as
RSD % )
[
1 n-1
i
i)1
jx
]
0.5
n
∑ (x
- jx)2
100
(2)
where xi is the experimental solubility of sample i and jx is the mean solubility of n measurements. In the case of solubilities less than 0.1 mol · kg-1, the criterion was extended to 3%. 3. Solubility Data The salt solubilities measured in organic solvents are listed in Table 1. To avoid the evaporation of acetone in the system LiCl + acetone, a maximum temperature of 328.15 K was used. The systems with NaCl or KCl in ethanol or acetone were not investigated, since the accuracy of the balance was not sufficient to determine the small solubilities with the required accuracy. A comparison of the KCl solubilities in methanol measured in this work and those reported by Pinho and Macedo2 show good agreement regarding the absolute solubilities and the temperature dependence, as can be seen from Figure 2. But the solubilities measured in this work are systematically 1.5% higher than the
(1)
where Msalt [mol · kg-1] is the molar mass of the salt. The salt solubility measurements in pure organic solvents were carried out over a temperature range from 293 to 333 K in 5 K steps, whereby the measurements always were started at the highest temperature. For each experimental point, the stirring temperature was set slightly above the equilibrium temperature in order to avoid the formation of microcrystals. Then the solubilities were measured at the desired temperature after adequate sedimentation. For the measurements of the mixed solvent electrolyte systems a temperature of 313.15 K was chosen. For each experimental point, three samples were taken by using the syringe equipped with a filter, to avoid the dragging of small particles of the
Figure 2. Comparison of the KCl solubility in methanol at different temperatures: (9) this work; (O) Pinho and Macedo;2 (4,3) further published solubility data stored in DDB.10-12 Line represents the average solubilities.
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a relative error of (1% can be assumed at high solubilities. But the error increases with decreasing solubility. That is the reason why the solubilities of NaCl and KCl in pure organic solvents could not be measured. 4. Solid-Liquid Equilibria Modeling
Figure 3. Comparison of the LiCl solubility in ethanol at different temperatures: (9) this work. The other symbols represent further published solubility data stored in DDB.13-16
values reported by Pinho and Macedo, although the procedure was similar to those used by Pinho and Macedo.5 On the basis of the reproducibility of the experimental results and a comparison with the already published data, it can be concluded that the procedure used in this work provides reliable solubilities. Solubilities for LiCl in ethanol are shown in Figure 3. It can be seen that a reliable description of the temperature dependence of the solubility of LiCl in ethanol was achieved when compared with the solubility data reported by other authors. From Figure 3 it can be seen that the available data show large scattering and in some cases even a different temperature dependence. Because most of the available salt solubility data in organic solvents are quite old or questionable, an adequate evaluation of the data is not possible. The solubilities of LiCl and LiBr in methanol, ethanol, or acetone investigated in this work are listed in Table 1. It can be seen, that the solubilities of LiCl and LiBr are considerably higher than for the other alkali-metal halogenides. Even in ethanol and acetone the solubility is measurable, while for NaCl and KCl the solubility in these solvents is too small for our measurement procedure. In general the salt solubility in organic solvents is a lot lower than in water because of the lower polarity and dielectric constant of the solvents. The salt solubilities determined for mixed solvents at 313.15 K are given in Table 2. The solubilities for all the systems are expressed on molality scale, while the solvent composition is expressed in mass % (w %) on the salt-free basis. Since the solubilities of NaCl and KCl in pure ethanol or acetone are very low (
