Cl

Article pubs.acs.org/IECR

Solubility Phase Diagram of the Quaternary System Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K: Experimental Redetermination and Model Simulation Hongxia Li,† Dewen Zeng,*,†,‡ Yan Yao,† Xia Yin,§ Dongdong Li,† Haijun Han,† and Hongyan Zhou† †

Key Laboratory of Salt Lake Resources and Chemistry, Qinghai Institute of Salt Lakes, Chinese Academy of Sciences, Xining 810008, P. R. China ‡ College of Chemistry and Chemical Engineering, Central South University, Changsha 410083, P. R. China § College of Chemistry and Chemical Engineering, Hunan University, Changsha 410082, P. R. China S Supporting Information *

ABSTRACT: Solubility isotherms for the ternary system MgCl2−MgSO4−H2O and the quaternary reciprocal system Li+, Mg2+//Cl−, SO42−−H2O were determined at 298.15 K by an isothermal dissolution method. In the ternary phase diagram, there are six solubility branches corresponding to the solid phases MgSO4·nH2O(s) (n = 7, 6, 5, 4, 1) and MgCl2·6H2O(s). In the quaternary equilibrium phase diagram, there are 16 solubility co-saturated lines corresponding to the solid phases MgSO4· nH2O(s) (n = 7, 6, 5, 4, 1), MgCl2·6H2O(s), Li2SO4·H2O(s), LiCl·MgCl2·7H2O(s), and LiCl·H2O(s). This report describes for the first time that the equilibrium solid phases MgSO4·H2O(s) and MgSO4·4H2O(s) have been found to exist in this quaternary system. However, the phase field of MgSO4·H2O(s) overlaps with the phase fields of MgSO4·4H2O(s) and MgSO4·5H2O(s), which indicates that MgSO4·4H2O(s) and MgSO4·5H2O(s) are metastable phases; MgSO4·H2O(s) is a relatively more stable phase in both the ternary and quaternary systems. A Pitzer−Simonson−Clegg thermodynamic model was used to simulate the properties of the sub-binary and subternary systems and to predict the solubility phase diagram of the quaternary system. The results of the modeling are in reasonable agreement with the experimental data.

1. INTRODUCTION

In fact, this quaternary system and its ternary subsystems have been widely investigated by several research groups. However, many of the solubility phenomena relevant to the system are still not well understood. While most authors1−6 take it for granted that MgCl2·6H2O(s) and MgSO4·nH2O(s) (n = 7, 6, 5, 4) are the only stable solid phases in the ternary system MgCl2−MgSO4−H2O at 298 K, other researchers7−10 report that another stable phase, MgSO4·H2O(s), exists in the system at 298 K. The experimental data supporting either conclusion remain very limited. Bursa and Stanisz-Lewicka9 provide only one solubility datum point for MgSO4·H2O(s). Shojhet7 reported solubility data only for MgSO4·H2O(s) (cosaturated with MgCl2·6H2O(s) and co-saturated with MgSO4· 6H2O(s), respectively). Van’t Hoff7 reported that the MgSO4· H2O(s) phase definitely coexists with MgCl2·6H2O(s) in the ternary system. However, they reported no composition of the invariant point, needless to say, those of the MgSO4·H2O(s) solubility isotherm. Later studies11−13 simulated the formation field of MgSO4·H2O(s) based on the limited experimental data available. As for the solubility phase diagram of the quaternary system Li+, Mg2+//Cl−, SO42−−H2O at 298 K, Kydynov et al.14 experimentally found no phase zones for the solid phases MgSO4·nH2O(s) (n = 6, 5, 4, 1). Ren and Song15 similarly found

Isothermal evaporation crystallization is a simple and costeffective technique that can be used to extract valuable resources from salt lakes containing the ions Li+, Na+, K+, Mg2+, Cl−, SO42−, and B4O72−. This method takes advantage of the natural climate in the area of the salt lakes, i.e., low humidity and abundant solar energy. The theoretical basis of isothermal evaporation crystallization lies in the thermodynamic equilibrium and nonequilibrium solubility phase diagrams of related multi-component systems. After several steps of solar pond evaporation, most Na+ ions crystallize as NaCl or Na2SO4. As water activity decreases, the complicated salt lake system becomes a relatively simple system that can be approximately represented by Li+, Mg2+//Cl−, SO42−−H2O. In the last stage of the natural evaporation process, magnesium ions salt out as the hydrated sulfate MgSO4·nH2O, and the mass ratio of Mg:Li in the solution decreases to approximately 20. If lithium salts out as Li2SO4·H2O(s) along with MgSO4·nH2O(s) during the evaporation process, the loss of lithium salt will increase abruptly. Chemical engineers who work with salt lakes informed us that the loss ratio of lithium could be as high as 50% in a realistic solar pond process (personal communication). To understand the loss mechanism and to develop a new evaporation process to avoid this substantial loss of lithium, we need a more thorough understanding of the phase diagram of the quaternary reciprocal system Li+, Mg2+//Cl−, SO42−−H2O, especially in the zone describing the formation of solid Li2SO4· H2O(s). © 2014 American Chemical Society

Received: Revised: Accepted: Published: 7579

December 31, 2013 March 30, 2014 April 1, 2014 April 1, 2014 dx.doi.org/10.1021/ie404445t | Ind. Eng. Chem. Res. 2014, 53, 7579−7590

Industrial & Engineering Chemistry Research

Article

Table 1. Impurity Contents in the Reagents Prepared in This Work Impurity contents (w × 106) Regents

Ca

Fe

K

Na

Mn

Li

Mg

LiCl·H2O Li2SO4·H2O MgCl2·6H2O MgSO4·7H2O

17.034 1.871 1.313 19.24

 1.696 ··· ···

26.575 51.90 7.627 ···

54.69 2.229 2.203 8.40

   0.237

  0.459 

1.225 2.539  

···, below detectable limit; , has not been detected.

no formation fields for the solid phases MgSO4·nH2O(s) (n = 4, 1). Kwok et al.16 calculated the phase diagram of the quaternary system and took the solid phase MgSO4·H2O(s) as the stable phase but did not provide experimental support for this choice. In this study, we ask whether the phase MgSO4·H2O(s) is present and stable in the ternary system MgCl2−MgSO4−H2O and the quaternary system Li+, Mg2+//Cl−, SO42−−H2O. The answer will help us to define the formation zone of Li2SO4· H 2 O (s) , which could aid the development of a new crystallization course that avoids lithium loss as Li2SO4· H2O(s). We present phase diagrams of the ternary system MgCl2−MgSO4−H2O and the quaternary system Li+,Mg2+// Cl−,SO42−−H2O that have been comprehensively experimentally determined and simulated by a thermodynamic model.

2. EXPERIMENTAL SECTION 2.1. Materials. Lithium chloride monohydrate was prepared by neutralizing lithium carbonate (purity by mass fraction >0.999, Shanghai China-Lithium Industry Co., Ltd.) with hydrochloric acid (G. R.). Lithium sulfate monohydrate was prepared by neutralizing lithium carbonate with sulfuric acid (G. R.). Magnesium chloride hexahydrate (purity by mass fraction >0.98, Sinopharm Chemical Reagent Co., Ltd.) and magnesium sulfate heptahydrate (purity by mass fraction >0.99, Tianjin Kermel Chemical Reagent Co., Ltd.) were analytical grade reagents. These salt products were further purified two or three times by crystallization, with half salt recovery in each case. After recrystallization, impurities were detected by ICPAES (ICAP 6500 DUO, Thermo Scientific), and the total impurity concentrations were ≤100 ppm, as shown in Table 1. Kieserite (MgSO4·H2O(s)) was prepared by baking the purified MgSO4·7H2O(s) in an electrothermal constant-temperature drybox (202−0, Beijing Kewei Yongxing Co., Ltd.) for 24 h at 160 °C, and the resulting solid was identified by X-ray diffraction (XRD) as shown in Figure 1a. It seems that this is a mixture of MgSO4·H2O(s) and MgSO4·1.25H2O(s). Doubly distilled water, with conductivity less than 1.2 × 10−4 S m−1, was used to prepare samples for phase equilibrium experiments and for chemical analysis. Ethylenediamine tetraacetic acid disodium salt (A. R. reagent, Tianjin Kermel Chemical Reagent Co., Ltd.), mercury nitrate (A. R. reagent, Shanghai Zhongqin Chemical Reagent Co., Ltd.), and barium chloride dihydrate (A. R. reagent, Tianjin Binhai Kedi Chemical Reagent Co., Ltd.) were used for chemical analysis. 2.2. Apparatus and Improvement. Because the investigated systems contain strongly hydrophilic salts and a variety of hydrated salts, the ambient temperature and humidity greatly influence the phase equilibrium results. To obtain accurate results, we improved the phase equilibrium experimental apparatus, as shown in Figure 2. (1) In the constant temperature system, solubility measurements were carried out in a thermostat bath (LAUDA E200, Germany) with a

Figure 1. X-ray diffraction pattern of the kieserite obtained by drying MgSO4·7H2O(s) at 160 °C (a) and equilibrium with MgCl2 aqueous solution using the previous kieserite as initial crystal (b).

temperature stability of ±0.01 K. The bath temperature was monitored and recorded by means of a precision digital display thermal resistance thermometer (JW-1, Hebei Xinghua Electronic Instruments Plant) with an accuracy of ±0.001 K. The digital thermometer was tested by the Chinese National Institute of Metrology with a platinum resistance thermometer standard device, which provided the temperature correction (RGcf 2009−10001, Test Certificate). (2) In the equilibrium system, to prevent effects due to the water vapor from the water bath, we equipped the thermostat bath with a transparent organic glass plate with six holes (four for equilibrium bottles, the other two for sampling bottles). Every equilibrium bottle 7580

dx.doi.org/10.1021/ie404445t | Ind. Eng. Chem. Res. 2014, 53, 7579−7590


Article

Figure 2. Experimental apparatus diagram of the phase equilibrium determination: (1) Vacuum pump. (2) Drying tower. (3) Three-way valve. (4) Organic glass plate. (5) Water bath. (6) Sampling bottle. (7) Weighing bottle. (8) Glass sand core filter tube. (9) Equilibrium bottle. (10) Magneton. (11) Thermostat control head. (12) Organic glass pedestal. (13) Magnetic stirrer. (14) Bulb. (15) Mercury thermometer. (16) Point contact thermometer. (17) Electric relay. (18) Air constant temperature box. (19) Condenser.

mL cylindrical glass bottle, which was immersed in the water bath. The solution and solid in the glass bottle were stirred with a magnetic stirrer placed under the thermostat bath. Each sample was stirred at 298.15 K for 7−9 days and then kept static for approximately 8 h. A certain amount of liquid sample was pumped into a preweighed 15 mL bottle through a #4 glass sand core filter tube, which was connected to a sealed vacuum pump. The sample was then cooled, weighed, diluted, and analyzed. The solid and remaining liquid solution were stirred continuously for two more days, and another liquid sample was taken out for analysis. When the difference between the two samplings was less than 1%, the latter results were taken as the solubility. The wet solid phase was quickly transferred by glass colander into a preweighed 20 mL glass-stoppered bottle, and it was analyzed according to the analytical method described above. Another part of the wet solid phase was pumped into a double-jacketed sampling bottle filled with water from the bath to keep the sample temperature constant. A glass sand core filter tube was put into the sample to pump away most of the mother liquor, and the dried solid was identified quickly by XRD or by using a polarizing microscope. If the type of solid phase was determined to be the same by Schreinemakers’ method19 and by XRD, the results were considered reliable. The next group of equilibrium experiments was started by adding a small amount of salt or water to the remaining sample of solid and solution. To determine the time needed for the equilibrium between the solid phase kieserite and the electrolyte in this ternary system at 298.15 K, we prepared two initial solutions with different MgSO4 concentration, whose compositions are presented as points in Figure 3 marked “0 day”. A certain amount of kieserite, obtained by dehydration from MgSO4· 7H2O(s), was added to the two solutions as the initial crystal seed. The two samples were then placed in the water bath and stirred. The liquid phase compositions of the samples were

was supported by an organic glass pedestal with some holes to keep the water cycle unimpeded and to keep the temperature uniform between the equilibrium bottle and the bath. (3) In the sampling system, to prevent the solid phase from precipitating out of a saturated solution in the sampling process, we adjusted the ambient temperature (around the sampling system) and the temperature of the sampling devices before every sampling event, so that the temperatures were all near that of the baths. 2.3. Analysis Method and Improvement. The compositions of the liquid phases and their corresponding wet solid phases were analyzed by mass titration17 instead of volumetric methods. The relative error could be controlled below 0.2%. The Mg2+ ion concentration was determined by EDTA complexometric titration at pH 9.5−10 (ammonia buffer). Eriochrome black T was used as the indicator and the interference of the Li+ ion was eliminated by adding different amounts of mixed alcohol. The volume ratio of n-butyl alcohol and anhydrous ethanol was 1:9. This detailed analysis method and its results were described in our previous work.18 The Cl− ion concentration was determined by Hg(NO3)2 complexometric titration at pH 3−3.5, using a mixture of phenylazoformic acid 2-phenylhydrazide and bromophenol blue as the indicator. The SO42− ion concentration was determined by a gravimetric method, BaSO4 precipitation with an uncertainty of ±0.0005 in the mass fraction, which was burned to constant weight at 800 °C. The Li+ ion concentration was determined by the subtraction method. The type of solid phase was determined by Schreinemakers’ method,19 X-ray diffraction (Panalycal X’pert Pro, Holland) and a polarizing microscope (Leitz Dialux-Pol, Germany). A Sartorius balance (A200S) was used for weighing, with an error of ±0.1 mg. 2.4. Experimental Method and Procedures. 2.4.1. Ternary System MgCl2−MgSO4−H2O at 298.15 K. The solid− liquid equilibrium experiments were carried out by adding a second salt to the invariant points of the binary system in a 270 7581



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determined to be the equilibrium time for the kieserite phase in the following experiments. 2.4.2. Quaternary Reciprocal System Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K. The solid−liquid equilibrium experiments were carried out by adding a third salt to the invariant points of the four subternary systems, which were contained in equilibrium bottles immersed in the water bath in the thermostat with an equilibrium time of 7−10 days. Other details of the process were the same as described in section 2.4.1. The equilibrium solid phases were identified by XRD combined with Schreinemakers’ method. For the quaternary system, we used a spatial coordinate system in Schreinemakers’ method.

3. EXPERIMENTAL RESULTS AND DISCUSSION 3.1. Ternary System MgCl2−MgSO4−H2O at 298.15 K. The measured solubility results are presented in Table 2 and plotted in Figure 4. The equilibrium solid phases present in the ternary system are MgSO4·nH2O(s) (n = 7, 6, 5, 4, 1) and MgCl2·6H2O(s). The phase field of MgSO4·H2O(s) was definitively found to exist in this system at 298.15 K, and its solubility isotherm covers those of MgSO4·4H2O(s) and MgSO4·5H2O(s). Thus, the solid phase MgSO4·H2O(s) is the stable phase, and MgSO4·4H2O(s) and MgSO4·5H2O(s) are metastable phases.

Figure 3. MgSO4 concentration as a function of time in the ternary system MgCl2−MgSO4−H2O at 298.15 K. Symbols: experimental data in this work. Solid line: captured isotherm solubility for MgSO4·H2O(s) from Figure 7 in the literature.12

measured for different time intervals, and the analyzed results are shown in Figure 3. We found that the equilibrium could be approximately reached after 3−10 days. The crystal type of the equilibrium solid phase was identified again by XRD and is presented in Figure 1b. It seems that the MgSO4·1.25H2O(s) disappears in the equilibrium kieserite phase. Thus, 10 days was

Table 2. Solubility Data of the Ternary System MgCl2−MgSO4−H2O at 298.15 K composition of solution (100 wa)

a

composition of wet solid phase (100 wa)

no.

MgSO4

MgCl2

H2O

MgSO4

MgCl2

H2O

solid phaseb

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29

26.76 22.81 15.86 9.39 8.41 6.67 5.32 4.82 4.94 4.57 4.109 3.91 3.86 3.77 3.40 3.42 3.22 2.64 2.55 2.36 2.31 2.25 2.22 3.705 3.111 2.666 1.531 0.818 0

0 2.79 8.46 15.64 17.19 20.32 24.35 25.99 26.20 26.84 28.775 29.72 30.26 29.78 30.65 31.04 31.42 32.77 32.84 34.27 34.51 34.52 34.47 29.496 31.349 32.535 34.298 35.062 35.64

73.24 74.40 75.68 74.97 74.40 73.01 70.33 69.19 68.86 68.59 67.116 66.37 65.88 66.45 65.95 65.54 65.36 64.59 64.61 63.37 63.18 63.23 63.31 66.799 65.54 64.8 64.171 64.12 64.36

 31.87 31.27 29.68  30.53 30.01 29.53 29.94 29.79  30.44 29.11 32.27 28.59 28.90 31.38 32.82 35.47 32.35 20.24 3.49 28.42   16.599  23.58 

 1.88 4.53 7.67  8.85 10.48 11.44 11.93 12.78  13.73 14.99 13.94 16.91 16.52 16.53 16.18 14.84 17.11 25.98 39.16 20.11   26.664  30.34 

 66.25 64.20 62.65  60.62 59.51 59.03 58.13 57.43  55.83 55.90 53.79 54.50 54.58 52.09 51.00 49.69 50.54 53.78 57.35 51.47   56.737  46.07 

Ep Ep Ep Ep Ep Ep Ep Ep Ep + Hx Hx Hx Hx Hx + Pt Pt Pt Pt + Te Te Te Te Te Te + Bis Te + Bis Te + Bis Hx + Kie Kie Kie Kie Kie + Bis Bis

w: mass fraction. bEp: MgSO4·7H2O(s). Hx: MgSO4·6H2O(s). Pt: MgSO4·5H2O(s). Te: MgSO4·4H2O(s). Bis: MgCl2·6H2O(s). Kie: MgSO4·H2O(s). 7582



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Figure 4. Solubility isotherms determined in this work in the ternary system MgCl2−MgSO4−H2O at 298.15 K: ○, Ep; ◓, Ep + Hx; □, Hx; ⬒, Hx + Pt; △, Pt; tilted square solid top half, Pt + Te; ◇, Te; tilted square solid bottom half, Te + Bis; ▽, Kie; tilted square solid left half, Hx + Kie; tilted square solid right half, Kie + Bis; ☆, Bis; full filled symbols are for wet solid phases.

We compared our solubility data for the phases MgSO4· nH2O(s) (n = 7, 6, 5, 4) and MgCl2·6H2O(s) to solubilities reported in the literature4,6 as shown in Figure 5. It was found

recognized MgSO4·4H2O(s) and MgSO4·5H2O(s) as metastable phases in the ternary system at 298.15 K. We compared the solubility isotherm for the phase MgSO4·H2O(s) measured in this work with those reported previously. All of the data are presented in Figure 6. Shojhet8 reported two solubility

Figure 6. Comparison of solubility isotherm for the solid phase MgSO4·H2O(s) in the ternary system MgCl2−MgSO4−H2O at 298.15 K: tilted square solid left half, Hx + Kie in this work; ▽, Kie in this work; tilted square solid right half, Kie + Bis in this work; ⬒, Hx + Kie;8 ⬓, Kie + Bis;8 ■, Kie;9 ---, captured model values;10 ···, captured model values;11 , captured model values.12

measurements for MgSO4·H2O(s), co-saturated with MgSO4· 6H2O(s) and MgCl2·6H2O(s). That invariable point for MgSO4· H2O(s) with MgSO4·6H2O(s) agrees with our result very well, but the solubility for MgSO4·H2O(s) with MgCl2·6H2O(s) differs greatly from our measured solubility. Bursa and StaniszLewicka9 reported only one experimental solubility value for MgSO4·H2O(s), which many later researchers10−12 used for simulations of the solubility isotherms (Figure 6). Our experimental solubility data for MgSO4·H2O(s) are generally higher than the solubility reported in the literature.10−12 One reason for the difference could be that the prepared kieserite in this work contained some MgSO4·1.25H2O(s). However, to what extent the impurity MgSO4·1.25H2O(s) increases the solubility of MgSO4·H2O(s) is unknown yet. On the other hand, except for the single experimental datum reported by Bursa and Stanisz-Lewicka,9 no other experimental solubility data of MgSO4·H2O(s) in the ternary system can support the simulated results in the literature,10−12 as we know. 3.2. Quaternary Reciprocal System Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K. The experimental solubility data are

Figure 5. Comparison of the solubility data of the solid phases MgSO4·7H2O(s), MgSO4·6H2O(s), MgSO4·5H2O(s), MgSO4·4H2O(s), and MgCl2·6H2O(s) in the ternary system MgCl2−MgSO4−H2O at 298.15 K determined in this work with those from various literature: ○, this work; Δ, Du et al.;4 □, Butchkareva.6

that our data are almost identical to previously reported results.6 This might indicate that the data obtained in this work are reliable. However, no solubility data were reported for MgSO4·H2O(s) in the previous studies.4,6 We noticed that several research groups8−13 did recognize the phases MgSO4·nH2O(s) (n = 7, 6, 1) as stable phases and 7583



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Table 3. Solubility Data of the Quaternary Reciprocal System Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K composition of solution (100 wa)

composition of wet solid phase (100 wa)

Jänecke coordinateb

no.

MgSO4

MgCl2

LiCl

Li2SO4

MgSO4

MgCl2

LiCl

Li2SO4

Li22+

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50

18.499 18.294 18.232 18.278 18.419 16.057 12.011 8.746 7.238 5.07 5.044 4.94 3.908 3.86 5.044 6.582 6.563 5.69 5.355 3.4 3.39 3.505 3.479 4.661 4.269 3.919 2.31 2.301 2.488 2.98 0.902 0.446 0.077 0.019 0.008 0.01 0 0 0.014 0.028 0 4.109 4.522 4.71 2.813 1.892 0.746 0.802 0.832 0.818

0 0 0 0 0 2.79 7.752 13.204 17.343 24.716 25.116 26.2 28.633 30.26 24.531 19.384 19.461 22.2 23.346 30.65 30.57 29.214 29.569 25.355 26.534 27.112 34.51 33.131 31.777 29.81 31.421 30.165 24.807 14.705 14.662 14.346 14.665 5.69 5.705 1.657 0 28.775 25.361 25.25 26.667 28.299 31.223 32.273 34.467 35.062

0 1.074 4.027 7.089 8.071 9.33 8.473 7.207 5.984 2.294 1.062 0 1.388 0 3.989 5.348 5.326 4.49 4.31 0 0.74 1.824 2.782 3.779 3.575 3.62 0 1.298 2.728 3.03 4.143 5.905 12.732 27.303 27.326 27.71 27.312 40.068 40.108 44.14 45.757 0 2.732 4.31 4.304 4.523 4.783 3.72 1.667 0

14.220 12.686 8.458 4.056 2.572 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0.028 0 0 0 0 0 0 0 0 0

27.708 31.550 28.665 17.814 24.723 37.986 36.106 31.713 32.219 9.213 22.345 29.94 30.166 29.11 34.683 11.382 31.93 33.26 33.206 28.59 34.65 32.149 9.844 31.108 36.359  20.24 2.577 2.44 21.83 5.039 4.187 3.735 3.536 6.618 7.01 0 0 5.141 1.525 0      13.409   

0 0 0 0 0 0.853 0 1.975 0 22.694 16.095 11.93 14.698 14.99 3.579 16.071 1.314 7.18 6.104 16.91 12.98 14.877 23.901 10.824 5.952  25.98 38.282 38.17 18.61 34.32 29.332 26.635 27.309 22.474 16.962 25.801 13.65 8.117 0 0      26.956   

0 0.656 1.641 4.664 4.215 3.344 6.472 5.561 8.786 2.015 0.842 0 0.154 0 9.113 6.433 11.448 4.42 6.213 0 0.33 1.084 5.292 4.195 7.216  0 0.688 1.563 4.08 5.098 7.167 11.629 15.19 18.706 25.803 19.81 39.531 41.062 43.829 43.972      5.563   

20.20 10.92 17.52 22.43 16.78 0 0.517 0 7.301 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0  0 0 0 0 0 0 0 0 0 0 0 0 0 10.868 13.24      0   

45.701 45.728 45.098 44.244 43.659 40.346 35.547 28.685 22.559 8.231 3.937 0 4.682 0 13.573 19.629 19.523 15.88 14.928 0 2.44 6.018 8.814 12.75 11.835 11.864 0 4.002 8.323 9.57 12.647 17.85 36.504 67.563 67.659 68.432 67.65 88.77 88.739 96.723 100 0 9.584 14.31 14.33 14.56 14.44 11.265 5.06 0

SO22+

H2O

solid phasec

100 95.475 82.787 69.301 64.955 48.911 35.493 24.519 19.221 12.811 13.169 12.98 9.287 9.17 12.09 17.016 16.946 14.18 13.064 8.07 7.87 8.146 7.764 11.076 9.953 9.04 5.03 4.999 5.347 6.63 1.939 0.95 0.155 0.033 0.013 0.018 0 0 0.022 0.043 0.05 10.15 11.175 11.02 6.599 4.29 1.586 1.71 1.779 1.812

1320.58 1347.94 1395.1 1439.64 1451.01 1463.01 1418.16 1328.07 1232.98 1147.74 1200.7 1209.78 1050.09 1046.05 1064.91 1187.45 1185.46 1126.8 1092.91 1046.34 1013.44 1017.31 956.814 1052.09 1023.14 1008.4 919.69 919.22 905.42 954.41 913.55 903.94 842.54 675.76 676.495 673.95 676.983 566.098 564.598 559.2 557.84 1107.61 1112.6 1096.12 1038.52 990.208 899.725 901.503 901.168 948.81

Ls + Ep Ls + Ep Ls + Ep Ls + Ep Ls + Ep Ls + Ep Ls + Ep Ls + Ep Ls + Ep + Hx Ep + Hx Ep + Hx Ep + Hx Hx + Pt Hx + Pt Ls + Hx + Pt Ls + Hx Ls + Hx Ls + Hx Ls + Hx Pt + Te Pt + Te Pt + Te Ls + Pt + Te Ls + Pt Ls + Pt Ls + Pt Te + Bis Te + Bis Ls + Te + Bis Ls + Te Ls + Bis Ls + Bis Ls + Bis Ls + Bis Ls + Bis Ls + Bis +LiC Bis +LiC LiC + Lc Ls + LiC + Lc Ls + Lc Ls + Lc Hx + Kie Hx + Kie Hx + Kie + Ls Kie + Ls Kie + Ls Kie + Ls + Bis Kie + Bis Kie + Bis Kie + Bis

w: mass fraction. bxs,SO42− = 100zSO42−nSO42−/D, xs,Li22+ = 100zLi22+nLi22+/D, xs,H2O = 100nH2O/D, D = ∑zini = ∑zjnj, i, j, z, and n stand for anions, cations, charge number, and amount of substance, respectively. cLs: Li2SO4·H2O(s). Ep: MgSO4·7H2O(s). Hx: MgSO4·6H2O(s). Pt: MgSO4·5H2O(s). Te: MgSO4·4H2O(s). Bis: MgCl2·6H2O(s). LiC: LiCl·MgCl2·7H2O(s). Lc: LiCl·H2O(s). Kie: MgSO4·H2O(s). a

in the reciprocal system for the first time. As in the ternary system MgCl2−MgSO4−H2O, the formation field of MgSO4· H2O(s) covers those for the two phase fields of MgSO4·4H2O(s) and MgSO4·5H2O(s) in the quaternary reciprocal system. Both of the results in the ternary and quaternary systems indicate

given in Table 3. In the quaternary reciprocal system, the solid phases in equilibrium with saturated solutions were MgSO4· nH2O(s) (n = 7, 6, 5, 4, 1), MgCl2·6H2O(s), LiCl·MgCl2· 7H2O(s), LiCl·H2O(s), and Li2SO4·H2O(s). The phase fields of MgSO4·H2O(s) and MgSO4·4H2O(s) were found experimentally 7584



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Table 4. Binary Pitzer−Simonson−Clegg Model Parameters solute

Bmx

Bmx1

Wmx

Umx

Vmx

a

a1

source of water activity for the parametrization

LiCl MgCl2 Li2SO4 MgSO4

233.956 −53.558 43.02 145.295

0 0 0 −43862.73

0.156 99 −47.823 −10.482 −42.682

29.398 −75.619 −13.332 −46.490

−34.09 41.335 4.543 0

13 13 13 7

0 0 0 95

Gibbard and Scatchard25 Rard and Miller26 Rard et al.27 Archer and Rard28

that MgSO4·H2O(s) is the stable phase, while MgSO4·4H2O(s) and MgSO4·5H2O(s) are the metastable phases. To justify the equilibrium solid phases, we present some representative XRD patterns of the equilibrium solids containing the kieserite in the supporting files. Our experimental results are quite different from those reported in the literature14,15 in the solid phase types and their sizes of phase formation field, which will be discussed in detail in this work later. One point worth mentioning here, now that the kieserite is found stable in the reciprocal quaternary system at 298.15 K, is that it should be expected to exist in the system at 323 K, too. However, Vaisfel’d and Shevchuk20 found no kieserite phase in the system at 323 K. Their data could be wrong.

4.2. Ternary Parameter Determination. To calculate the ternary phase diagrams, the solubility product for each solid phase should be determined in advance. The binary solubilities for LiCl(aq), MgCl2(aq), and MgSO4(aq) were taken from reference data.29 The data for Li2SO4(aq) were taken from another source.30 The solubility product based on the molar fraction of each solid phase was obtained by calculating the activities of ions and water at the solubility points in each binary system according to eq 1 using the parameters in Table 4. The resulting ln K values for LiCl·H2O(s), Li2SO4·H2O(s), MgSO4·7H2O(s), and MgCl2·6H2O(s) are shown in Table 5. M v+X v _ × nH 2O(s) = v+ M+(aq) + v− X−(aq) + nH 2O(aq)

(

+

−

)

v v n + a − a ln K = ln a M (aq) X (aq) H 2O(aq)

4. MODELING As we discussed previously,21−23 the Pitzer−Simonson−Clegg model,13,24 a mole fraction-based ion interaction model, is especially good for describing the properties of our aqueous system containing highly soluble salts. The activity of the water and activities of the ions as functions of salt concentration in the framework of the model were described in our previous work.21−23 4.1. Binary Parameter Determination. Binary model parameters for LiCl(aq), MgCl2(aq), Li2SO4(aq), and MgSO4(aq) were obtained by fitting to the water activity data25−28 at 298.15 K. All parameters are summarized in Table 4. The binary model parameters α and α1 for LiCl(aq), MgCl2(aq), and Li2SO4(aq) were taken to be 13.0 and 0, respectively. As recommended by Clegg et al.,13 we set the parameters α and α1 for the 2−2 electrolyte MgSO4(aq) to 7 and 95, respectively. The calculated water activities agree well with the experimental values. The results are shown in Figure 7.

(1)

Table 5. Solubility Products (ln K) of the Solid Phases in the System Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K solid phase

ln K

references of solubility data for the parametrization of ln K

LiCl·H2O Li2SO4·H2O MgSO4·7H2O MgCl2·6H2O MgSO4·6H2O MgSO4·5H2O MgSO4·4H2O MgSO4·H2O LiCl·MgCl2·7H2O

4.85 −11.02 −14.43 −2.14 −13.766 −12.908 −11.964 −9.598 1.52

Linke and Seidell29 Li et al.30 Linke and Seidell29 Linke and Seidell29 this work this work this work this work Voskresenskaya and Yanat’eva32

Using the binary parameters in Tables 4 and 5, we predicted the solubility isotherms of the four ternary systems and found that the predicted results (dashed lines in Figures 8, 9, 10, and 11) deviate from the experimental values reported in the literature30−32 and our own experimental values. Thus, the ternary mixing parameters Wmnx, Qmnx, and Umnx must be introduced into the model. We fitted the mixing parameters to the solubility isotherms for the four ternary systems described in the literature30−32 and in this work, and we obtained the mixing model parameters listed in Table 6. By using the binary and ternary parameters in Tables 4, 5, and 6, we calculated the solubility isotherms of the four ternary systems at 298.15 K (solid lines in Figures 8, 9, 10, and 11), which show satisfactory agreement with the experimental data. In addition, we predicted the water activities in the four ternary systems using our model and compared them with the experimental values.33−36 Figure 12 shows that the predictions become more consistent with the experimental values when the mixing parameters are used in the calculation. 4.3. Solubility Prediction for the Quaternary System Li+,Mg2+//Cl−,SO42−−H2O. 4.3.1. Comparison with Experimental Data. By using the binary and ternary parameters listed in Tables 4−6, we predicted the solubility isotherms of the quaternary system at 298.15 K and compared them to the

Figure 7. Comparison of calculated water activities for the system LiCl−H2O, MgCl2−H2O, Li2SO4−H2O, and MgSO4−H2O with literature data at 298.15 K: ○, LiCl(aq);25 □, MgCl2(aq);26 △, Li2SO4(aq);27 ◇, MgSO4(aq);28 , model values. 7585



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Figure 8. Comparison of calculated (lines) and experimental (symbols) solubility data30 for the ternary system Li2SO4−MgSO4− H2O at 298.15 K: ○, Ls; ●, Ls + Ep; △, Ep; , model values with binary and ternary parameters; ---, model values with binary parameters only.

Figure 10. Comparison of calculated (lines) and experimental (symbols) solubility data32 for the ternary system LiCl−MgCl2− H2O at 298.15 K: ○, Lc; ●, Lc + LiC; □, LiC; ■, LiC + Bis; Δ, Bis; , model values with binary and ternary parameters; ---, model values with binary parameters only.

Figure 9. Comparison of calculated (lines) and experimental (symbols) solubility data31 for the ternary system LiCl−Li2SO4− H2O at 298.15 K:, △, Lc; ●, Lc + Ls; ○, Ls; , model values with binary and ternary parameters; ---, model values with binary parameters only.

Figure 11. Comparison of calculated (lines) and experimental (symbols) solubility data for the ternary system MgCl2−MgSO4− H2O at 298.15 K in this work: ○, Ep; ●, Ep + Hx; □, Hx; ■, Hx + Pt; △, Pt; ▲, Pt + Te; ◇, Te; ⧫, Te + Bis; ▼, Kie + Hx; ▽, Kie; ★, Kie + Bis; ☆, Bis; , model values with binary and ternary parameters; ---, model values with binary parameters only.

experimental data measured in this work, as shown in Figure 13. Both the experimental and predicted results indicate that the formation fields of LiCl·H2O(s), LiCl·MgCl2·7H2O(s), and MgCl2·6H2O(s) are quite narrow, and the phases of MgSO4· 4H2O(s) and MgSO4·5H2O(s) are metastable in the quaternary reciprocal system, as they are in the ternary system MgCl2− MgSO4−H2O. The predicted solubility isotherms of MgSO4· H2O(s) agree with our experimental results, which indicate that both the experiments and modeling predictions have some degree of reliability. 4.3.2. Comparison with Literature Data. As mentioned above, the phase diagram of the reciprocal system has been described by many research groups.14−16 We collected results and model values from the literature for comparison (Figure 14). There are remarkable differences in the reported phase field sizes, solid phase types, and stabilities for magnesium sulfate hydrates. Kydynov et al.14 reported that there are only five phase fields present in the system at 298 K: LiCl·H2O(s), LiCl·MgCl2·7H2O(s), MgCl2·6H2O(s), Li2SO4·H2O(s), and

Table 6. Ternary Pitzer−Simonson−Clegg Model Parameters mixture parameters

systems

Wmnx

Qmnx

Umnx

Li2SO4−MgSO4−H2O LiCl−Li2SO4−H2O LiCl−MgCl2−H2O

−9.539 −61.862 −56.072

−3.164 40.219 19.828

0 −9.771 −2.744

MgCl2−MgSO4−H2O

−16.0

17.252

−18.0

sources of solubility data for parametrization Li et al.30 Plyushchev et al.31 Voskresenskaya and Yanat’eva32 this work

MgSO4·7H2O(s) (Figure 14a). They found no formation fields for the solid phases MgSO4·6H2O(s), MgSO4·5H2O(s), MgSO4· 4H2O(s), and MgSO4·H2O(s) as either stable or metastable phases. Additionally, these authors reported a formation field for the solid phase MgCl2·6H2O(s) that is inconceivably large and very different from other reported formation fields15,16 and 7586



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Figure 12. Deviations of model and experimental values in water activity in the four ternary systems. Δaw = aw(exp) − aw(model). Hollow symbols: model values calculated by binary parameters. Full-filled symbols: model values calculated by binary parameters and mixing parameters. Experimental values: ○ ●, Zhang et al.;33 □ ■, Song et al.;34 Δ ▲, Yao et al.;35◇ ⧫, Rard et al.36

Figure 13. Dry salt phase diagram of the quaternary system Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K: ○, Ep + Ls; ●, Ep + Ls + Hx; hexagon solid top half, Hx + Ls; hexagon solid, Hx + Ls + Pt ; ⊗, Pt + Ls; ▲, Pt + Ls + Te; ▽, Te + Ls; ■, Te + Ls + Bis; ⊕, Ep + Hx; △, Hx + Pt; upright triangle solid top half, Pt + Te; □, Te + Bis; star solid top half, Hx + Kie; ⬒, Hx + Kie + Ls; upside down triangle solid top half, Kie + Ls; ☆, Kie + Bis; ◑, Kie + Ls + Bis; ◇, Ls + Bis; ⊞, LiC + Bis; ⧫, Ls + Bis + LiC; open box with x, Lc + LiC; pentagon with vertical line, Lc + Ls; ★, Ls + LiC + Lc; , predicted stable invariable lines; ---, predicted metastable invariable lines.

and MgSO4·5H2O(s)) not reported elsewhere in the literature14 (Figure 14b). We identified the solid phases MgSO4·4H2O(s) and MgSO4·H2O(s) as metastable and stable, respectively, in the ternary system MgCl2−MgSO4−H2O, so they should exist in

the one that we determined. We can safely conclude that their results14 are incorrect. Ren and Song15 reported seven phase fields in this reciprocal system at 298 K, including two phase fields (MgSO4·6H2O(s) 7587



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Figure 14. Comparisons of the phase diagrams for the quaternary reciprocal system Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K: (a) Kydynov et al.,14 (b) Ren and Song,15 (c) Kwok et al.,16 (d) this work.

H2O(s). We note that Kwok et al.16 did not measure the solubility isotherms for the solid phase MgSO4·H2O(s). Figure 14c shows only the predicted formation field of MgSO4·H2O(s), which was not verified by any of the authors’ experimental work. Furthermore, Kwok et al.16 did not explain how they determined the solubility product parameter for the solid phase MgSO4·H2O(s), and this parameter determines the size of the calculated formation field for the solid phase. As we mentioned previously, the effect of the possible impurity MgSO4· 1.25H2O(s) on the measured solubility value of MgSO4· H2O(s) in this work is yet to be determined. 4.3.3. Application of the Predicted Phase Diagram. To understand the crystallization route for the title system, we have calculated the equal-scale lines of the water content and presented them in Figure 15. As water content decreases, the system composition moves from right to left on the dry salt scale, all the way to the lower left corner in most cases. Generally, it is preferable for the ratio wMg:wLi of naturally evaporated salt lake brine to be as low as possible. On the basis of the configuration in the lower left corner of this phase diagram, one can determine the ratio wMg:wLi for the salt lake brine in the last stage of solar pond evaporation. When the phase MgSO4·H2O(s) does not crystallize as usual because of kinetic reasons, salt lake brine with a composition in the dry salt scale below the line “ac” in Figure 15 can reach point “a” without losing Li as Li2SO4·H2O(s). At point “a”, the solution is co-saturated with the solid phases Li2SO4·H2O(s), MgSO4·

the quaternary reciprocal system; however, they are not reported by Ren and Song.15 We note that the phase field of MgCl2·6H2O(s) reported by Ren and Song15 (Figure 14b) is much larger than the one we determined in this work. Ren and Song15 analyzed the Mg2+ ion concentration, eliminating the interference of the Li+ ion by adding a mixture of n-butyl alcohol and ethanol. However, they15 did not report the optimal volume ratio of n-butyl alcohol and ethanol or the optimal amount of the mixture to add to the sample, which we found in our previous study18 can greatly affect the accuracy of Mg2+ analysis. We used those results to analyze Mg2+ in this work. Ren and Song15 also analyzed the SO42− ion by using a turbidimetric method when the concentration of sulfate salts was low, which could have caused non-negligible errors. Kwok et al.16 predicted the phase diagram of the quaternary reciprocal system based on their limited experimental data, and they recognized the stable formation field of the solid phase MgSO4·H2O(s) in the system, as shown in Figure 14c. Their predicted phase field for MgSO4·H2O(s) extended to the region of xs,Li22+ > 20, which is remarkably larger than what we determined in this work: xs,Li22+ < 15. As described in the Introduction, determining the boundary between the solid phases MgSO4·H2O(s) and Li2SO4·H2O(s) is the most important part of describing the phase diagram of the system because that boundary is important for designing a MgSO 4 ·nH 2 O (s) crystallization process to avoid the loss of Li as Li2SO4· 7588



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crystallization route can be designed to separate the Mg and Li salts. A brine with a wMg:wLi ratio as low as 10 can be obtained in the solar pond process by adding MgSO4·H2O(s) as a crystal initiator.

■

ASSOCIATED CONTENT

S Supporting Information *

Information as mentioned in the text. This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel: +86 13618496806. Fax: +86 731 88879616. Notes

The authors declare no competing financial interest.

■ ■

Figure 15. Equal-scale lines of water content in the quaternary reciprocal system Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K. All lines are predicted: , stable invariable lines; ---, metastable invariable lines; ···, equal-scale lines of water content, numbers indicate mole ratio of water and salts, rH2O, salts.

ACKNOWLEDGMENTS This work was financially supported by the 100 Top Talents Project of the Chinese Academy of Sciences.

4H2O(s), and MgCl2·6H2O(s), and the water activity, mole ratio of water and salts, rH2O, salts, and wMg:wLi are approximately 0.29, 890, and 23, respectively. When the water activity and wMg:wLi are kept above 0.29 and 23, respectively, the solid phase Li2SO4·H2O(s) cannot crystallize along with MgSO4·4H2O(s) and MgCl2·6H2O(s), and Li loss is avoided. If some MgSO4·H2O(s) is added as an initiator in the last stage of evaporation, the metastable state of equilibrium can be broken. Any brine with a composition below line “bc” in Figure 15 can reach point “b” without forming Li2SO4·H2O(s), and at that point, the water activity, J(H2O), and wMg:wLi are approximately 0.31, 925, and 10, respectively. The higher water activity means that the evaporation condition is easier to achieve, and the lower wMg:wLi value of the evaporated brine possesses higher economic value.

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NOMENCLATURE aw = water activity w = mass fraction of substance Bmx, Bmx1, Wmx, Umx, Vmx = binary parameters of the model PSC for electrolyte MX a, a1 = adjustable model parameters ln K = solubility product of electrolyte MX Wmnx, Qmnx, Umnx = ternary parameters of the model PSC for system MX−NX−H2O T = temperature, K m = molality concentration, mol kg−1 n = amount of substance for ions i or j, mol z = charge number for ions i or j REFERENCES

(1) Takegami, S. Reciprocal salt pairs Na2Cl2 + MgSO4 = Na2SO4 + MgCl2 at 25 °C. Mem. Coll. Sci., Kyoto Imp. Univ. 1921, 4, 317. (2) Rode, T. Vapor pressure and solubility of ternary aqueous systems formed by sodium and magnesium chlorides and sulfates at 25 °C. Izv. Sect. Fiz. Khim. Anal. Acad. Nauk SSSR 1941, 14, 402. (3) Kurnakov, N. S.; Kuznetzov, V. G. Meta-stable hydrates of magnesium sulfate in the system: Magnesium chloride−magnesium sulfate−water. Izv. Sect. Fiz. Khim. Anal. Acad. Nauk SSSR 1935, 7, 186. (4) Du, X. H.; Song, P. S.; Zhang, J. T. Phase equilibrium in the quaternary system MgB4O7−MgSO4−MgCl2−H2O at 25 °C. Wuhan Huagong Xueyuan Xuebao (China) 2000, 13−17. (5) Balarew, C.; Tepavitcharova, S.; Rabadjieva, D.; Voigt, W. Solubility and crystallization in the system MgCl2−MgSO4−H2O at 50 and 75 °C. J. Solution Chem. 2001, 30, 815−823. (6) Butchkareva, I. N. In Experimental Solubility Data on Multicomponent Aqueous Salt Type Systems, Vol 1; Izd. Khimija: Leningrad, 1973; pp 935−936. (7) Van’t Hoff, J. H. et al. Untersuchungen über die Bildungsverhältnisse der ozeanischen Salzablagerungen; Akademische Verlagsgesellschaft: Leipzig, 1912. (8) Shojhet, D. N. Equilibria at 25 °C in the solutions of the ternary system: Magnesium chloride−magnesium sulfate−water. Izv. Sect. Fiz. Khim. Anal. Acad. Nauk. SSSR 1938, 10, 331. (9) Bursa, S.; Stanisz-Lewicka, M. Liquid−solid phase equilibrium in the MgSO4−MgCl2−H2O system at 25 °C. Chem. Stosow. 1981, 25, 89−93. (10) Filippov, V. K.; Cheremnykh, L. M. Thermodynamic study of the magnesium chloride, sulfate−water system at 25 °C. Ukr. Khim. Zh. 1984, 50, 1027−1032.

5. CONCLUSIONS By improving the experimental setup, analytical method, initial material, and time to equilibrium, we comprehensively determined the solubilities for the ternary system MgCl2− MgSO4−H2O and the quaternary reciprocal system Li+, Mg2+// Cl−, SO42−−H2O at 298.15 K. The formation fields of the solid phases MgSO4·nH2O(s) (n = 4, 1) were determined to exist in the quaternary system for the first time. The phase MgSO4· H2O(s) was determined to be stable; however, MgSO4·nH2O(s) (n = 5, 4) was found to be a metastable phase in the ternary system and quaternary system. A Pitzer−Simonson−Clegg model was used to simulate the properties of the sub-binary and subternary systems and to predict the phase diagram of the system Li+, Mg2+//Cl−, SO42−−H2O at 298.15 K. Reasonable agreement between the experimental and predicted data indicates that the phase diagram constructed in this work can be safely used to design crystallization processes for salt lake brine. On the basis of the phase diagram of the title system determined in this work, one can reasonably explain why an evaporated brine with wMg:wLi = 20−25 is usually obtained in the last stage of a solar pond process. At lower ratios, the valuable Li element will be lost as Li2SO4·H2O(s) along with MgCl2·6H2O(s) and MgSO4·4H2O(s). Importantly, an effective 7589



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(32) Voskresenskaya, N. K.; Yanat’eva, O. K. Heterogeneous equilibria in the ternary system LiCl−MgCl2−H2O. Izv Akad Nauk SSSR Ser. Khim. 1937, 1, 97−120. (33) Zhang, Z.; Yao, Y.; Song, P. S.; Chen, J. Q. Isopiestic determination of the osmotic and activity coefficients of the lithium sulfate−magnesium sulfate−water system. Wuli Huaxue Xuebao (China) 1993, 9, 366−373. (34) Song, P. S.; Yao, Y.; Sun, B.; Li, W. Pitzer model of thermodynamics for the Li+, Na+, K+, Mg2+/Cl−, SO42−−H2O system. Sci. Sinica Chim. (China) 2010, 40, 1286−1296. (35) Yao, Y.; Sun, B.; Song, P. S.; Zhang, Z.; Wang, R. L.; Chen, J. Q. Thermodynamics of lithium-containing aqueous electrolyte solutionisopiestic determination in osmotic and activity coefficients in lithium chloride −magnesium chloride−water system at 25 °C. Huaxue Xuebao (China) 1992, 50, 839−848. (36) Miladinović, J.; Ninković, R.; Todorović, M.; Rard, J. A. Isopiestic investigation of the osmotic and activity coefficients of {yMgCl2 + (1−y)MgSO4}(aq) and the osmotic coefficients of Na2SO4·MgSO4(aq) at 298.15 K. J. Solution Chem. 2008, 37, 307−329.

(11) Harvie, C. E.; Weare, J. H. The prediction of mineral solubilities in natural waters: The Na−K−Mg−Ca−Cl−SO4−H2O system from zero to high concentration at 25 °C. Geochim. Cosmochim. Acta 1980, 44, 981−997. (12) Steiger, M.; Linnow, K.; Ehrhardt, D.; Rohde, M. Decomposition reactions of magnesium sulfate hydrates and phase equilibria in the MgSO4−H2O and Na+−Mg2+−Cl−−SO42−−H2O systems with implications for Mars. Geochim. Cosmochim. Acta 2011, 75, 3600− 3626. (13) Clegg, S. L.; Pitzer, K. S.; Brimblecombe, P. Thermodynamics of multicomponent, miscible, ionic solutions: Mixtures including unsymmetrical electrolytes. J. Phys. Chem. 1992, 96, 9470−9479. (14) Kydynov, M.; Musuraliev, K.; Imanakunov, B. Quaternary system of lithium and magnesium sulfates and chlorides in water at 25 °C. Zh. Prikl. Khim. 1966, 39, 2114−2117. (15) Ren, K. W.; Song, P. S. Solubilities and properties of solution in the Li+, Mg2+//Cl−, SO42−−H2O system at 25 °C. Huaxue Xuebao (China) 1994, 10, 69−74. (16) Kwok, K. S.; Ng, K. M.; Taboada, M. E.; Cisternas, L. A. Thermodynamics of salt lake system: Representation, experiments, and visualization. AIChE J. 2008, 54, 706−727. (17) Li, H. X.; Dong, O. Y.; Yao, Y.; Wang, H. Y.; Zeng, D. W. The mass titration analytical method and its application. Yanhu Yanjiu (China) 2011, 19, 31−36. (18) Li, H. X.; Yao, Y.; Zeng, D. W. Measurement of concentrations of magnesium ion by masking method in presence of the interference lithium ion in salt water system. Yanhu Yanjiu (China) 2012, 20, 24− 30. (19) Schreinemakers, F. A. H. Graphical deductions from the solution isotherms of a double salt and its components. Z. Phys. Chem. 1893, 11, 75−109. (20) Vaisfel’d, M. I.; Shevchuk, V. G. The Li2Cl2 + MgSO4 ⇔ Li2SO4 + MgCl2−H2O system at 50 °C. Russ. J. Inorg. Chem. 1967, 12, 887− 889. (21) Guo, L. J.; Sun, B.; Zeng, D. W.; Yao, Y.; Han, H. J. Isopiestic measurement and solubility evaluation of the ternary system LiCl− SrCl2−H2O at 298.15 K. J. Chem. Eng. Data 2012, 57, 817−827. (22) Ouyang, H. T.; Zeng, D. W.; Zhou, H. Y.; Han, H. J.; Yao, Y. Solubility of the ternary system LiCl + NH4Cl + H2O. J. Chem. Eng. Data 2011, 56, 1096−1104. (23) Yin, X.; Chen, Q. Y.; Zeng, D. W.; Wang, W. L. Phase diagram of the system KNO3 + LiNO3+ Mg(NO3)2 + H2O. CALPHAD: Comput. Coupling Phase Diagrams Thermochem. 2011, 35, 463−472. (24) Clegg, S. L.; Pitzer, K. S. Thermodynamics of multicomponent, miscible, ionic solutions: generalized equations for symmetrical electrolytes. J. Phys. Chem. A 1992, 96, 3513−3520. (25) Gibbard, H. F.; Scatchard, G. Liquid-vapor equilibrium of aqueous lithium chloride, from 25 to 100 °C and from 1.0 to 18.5 molal, and related properties. J. Chem. Eng. Data 1973, 18, 293−298. (26) Rard, J. A.; Miller, D. G. Isopiestic determination of the osmotic and activity coefficients of aqueous MgCl2 solutions at 25 °C. J. Chem. Eng. Data 1981, 26, 38−43. (27) Rard, J. A.; Clegg, S. L.; Palmer, D. A. Isopiestic determination of the osmotic and activity coefficients of Li2SO4 (aq) at T = 298.15 and 323.15 K, and representation with an extended ion-interaction (Pitzer) model. J. Solution Chem. 2007, 36, 1347−1371. (28) Archer, D. G.; Rard, J. A. Isopiestic investigation of the osmotic and activity coefficients of aqueous MgSO4 and the solubility of MgSO4·7H2O (cr) at 298.15 K: Thermodynamic properties of the MgSO4 + H2O system to 440 K. J. Chem. Eng. Data 1998, 43, 791− 806. (29) Linke, W. F.; Seidell, A. Solubilities of Inorganic and Metal Organic Compounds, fourth ed.; American Chemical Society: Washington DC, 1965. (30) Li, B.; Wang, Q. Z.; Li, J.; Fang, C. H.; Song, P. S. Study of the ternary systems Li+, K+/SO42−−H2O and Li+, Mg2+/SO42−−H2O at 25 °C. Wuli Huaxue Xuebao (China) 1994, 10, 536−542. (31) Plyushchev, V. E.; Tulinova, V. B.; Lovetskaya, G. A. The LiCl− Li2SO4−H2O system. Zh. Neorg. Khim. 1959, 4, 1184−1189. 7590


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