Study of the Solubility, Supersolubility and Metastable Zone Width of

Article Cite This: J. Chem. Eng. Data XXXX, XXX, XXX−XXX

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Study of the Solubility, Supersolubility and Metastable Zone Width of Li2CO3 in the LiCl−NaCl−KCl−Na2SO4 System from 293.15 to 353.15K Huaiyou Wang,†,‡ Baoqiang Du,†,‡ and Min Wang*,†,‡ †

Key Laboratory of Comprehensive and Highly Efficient Utilization of Salt Lake Resources, Qinghai Institute of Salt Lakes, Chinese Academy of Sciences, Xining 810008, China ‡ Key Laboratory of Salt Lake Resources Chemistry of Qinghai Province, Xining 810008, China ABSTRACT: Isothermal dissolution method and turbidity technology were employed to measure the solubility, the supersolubility, and the metastable zone width of Li2CO3 in mixed LiCl−NaCl−KCl−Na2SO4 solutions at temperatures ranging from 293.15 to 353.15 K. It was found that the solubility values of Li2CO3 increased as the Na2SO4 concentration increased and decreased as the NaCl and KCl concentrations increased. In addition, these two opposite effects both affected the solubility of Li2CO3 in mixed electrolyte solutions. NaCl and KCl impurities decreased Li2CO3 supersolubility, which could be overcome by adding a small amount of Na2SO4 salt. The influence of the impurities on the metastable zone width differed from their influence on the solubility. The metastable zone width in electrolyte solutions was mainly affected by Na2SO4 salt, where it increased with the increasing Na2SO4 concentration. Meanwhile, NaCl and KCl salts had little impact on the metastable zone width. Overall, this study provides the basic data for the crystallization process of the lithium carbonate from salt lake brines in the industrial production process.

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solution dynamics. Sunat al.3 investigated the dependence of the Li2CO3 supersolubility on the temperature, impurities, and additives and found that the supersolubility decreases with the addition of NaCl, KCl, NaNO3, and NaBr, while it increases in the presence of Na2SO4,CH4N2O, NH4Cl, (NH4)2SO4, and EDTA disodium. Cheng et al.16 measured the solubility of Li2CO3 in Na−K−Li-Cl brines at temperatures ranging from 293 to 363.15 K and found that the solubility of Li2CO3 in solutions initially increases, reaching a maximum value, and then decreases with the increasing solution concentration. Meanwhile in the mixed NaCl−KCl solution, the solubility gradually decreased with the increasing salt concentrations. He et al.17 and Song et al.18 investigated the Li2CO3 supersaturation and solubility by precipitation from Zhabuye salt lake brine at temperatures ranging from 268.15 to 318.15 K and in aqueous solutions at temperatures ranging from 283.15 to 318.15 K, respectively. In Qinghai-Tibet Plateau, China, there are hundreds of salt lakes which are abundant in lithium resources. Impurities, such as NaCl, KCl, and Na2SO4, coexisting in brine would have different influences on the crystallization behavior of Li2CO3. Therefore, in this present study, the solubility, supersolubility, and MZW of Li2CO3 in mixed LiCl−NaCl−KCl−Na2SO4 solutions were investigated using turbidity technology.

INTRODUCTION Lithium and lithium compounds can be used in a wide range of applications due to their special physical and chemical properties. Li2CO3 is an important commercial lithium compound that can be used as the raw material for the chemical production of a number of lithium products.1−4 For instance, lithium-ion batteries typically used in electronics and lithium metal are mainly produced from Li2CO3.5−7 Lithium resources are found in nature as a solid and a liquid, such as lithium ores and salt lake brine, respectively, where about 70% of the total reserve is found in the liquid form.8−11 Extracting lithium from salt lake brine is an important means of producing Li2CO3. Due to the recent increase in the use of lithium battery materials, the quality of Li2CO3 products has recently drawn notable attention.12−15 Impurities, such as Na+, K+, SO42−, and Cl−, can greatly affect the electrochemical performance of lithium batteries. Reactive crystallization, performed by adding sodium carbonate to a concentrated lithium solution, is the key method of precipitating Li2CO3. Therefore, reactive crystallization of Li2CO3 has attracted significant interest from researchers. Generally, the solubility, supersolubility, and metastable zone width (MZW) are key parameters to be considered during crystallization of Li2CO3. Industrial crystallization is often controlled in the metastable zone to obtain products with a high quality of purity and yield, satisfactory morphology, and size distribution. The supersolubility and MZW are often affected by the temperature, impurities, feeding rate, and © XXXX American Chemical Society

Received: November 18, 2017 Accepted: April 17, 2018

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DOI: 10.1021/acs.jced.7b01012 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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Article

EXPERIMENTAL SECTION Materials and Apparatus. All of the chemical reagents used in this study are listed in Table 1, and were used without

Li2CO3 supersolubility, via a liquid dosing system at a feeding rate of 0.8 mL·min−1. When nucleation had occurred, as indicated by a straight rise in the turbidity curve, the pumping of the solution was stopped, and the time interval between the initial pumping and nucleation was recorded as the pumping time. Supersolubility was determined based on the added amount of Na 2 CO 3 , which was calculated based on concentration, feeding rate, and pumping time. Each measurement was repeated twice to verify the experimental reproducibility. When determining the solubility, about 60 mL of electrolyte solutions of NaCl, KCl, LiCl, and Na2SO4 of known concentrations was introduced into the jacketed glass vessel, which was then sealed with a rubber stopper. The solution was heated using a thermostatic circulator and monitored with a thermometer. Once the desired temperature had been reached, excess Li2CO3 powder was added into the electrolyte solutions, which were stirred at a rate of 300 rpm. On the basis of the published literature,16 once the equilibration time lasted for 48 h, the supernatant solution was immediately filtered through syringe filters with a bore diameter of 0.22 μm. The content of Li2CO3 in the solution was titrated using a dilute HCl solution.3,20 The supersolubility and solubility of Li2CO3 in electrolyte solutions are represented byc super and c sol , respectively. The difference between the supersolubility and solubility is defined as the MZW (Δc = csuper − csol).

Table 1. Chemical Reagents Employed in the Experiment chemical name lithium chloride anhydrous sodium chloride potassium chloride lithium carbonate sodium carbonate anhydrous Sodium sulfate anhydrous

formula LiCl NaCl KCl Li2CO3 Na2CO3 Na2SO4

provider (Kalamar) Shanghai Puzhen Biological Technology Co., Ltd. Tianjin Kemiou Chemical Reagent Co., Ltd. Tianjin Kemiou Chemical Reagent Co., Ltd. Sinopharm Chemical Reagent Co., Ltd. (Kalamar) Shanghai Puzhen Biological Technology Co., Ltd. (Kalamar) Shanghai Puzhen Biological Technology Co., Ltd.

purity (%) ≥99.0 ≥99.95 ≥99.95 ≥99.99 ≥99.8 ≥99.0

further purification. Deionized water (resistivity, 18.25 MΩ·cm) was prepared by a water purification system (UPT-II-20T, Chengdu Ultrapure Technology Co., Ltd.). Figure 1a displays the Crystal SCAN PolyBlock system (E1061, HEL, Ltd.)19 used for measuring the supersolubility. The crystallizer was a 100 mL glass reactor with an internal overhead stirrer and temperature and turbidity sensors controlled by the PolyBlock. The turbidity sensor detected the crystal nucleus with the IR laser reflected by the optical lens. The temperature was controlled both by the PolyBlock and a thermostatic bath (FP-50, JULABO Labortechnik GmbH). Figure 1b shows the apparatus for measuring the solubility using the isothermal dissolution method. The vessel was a 100 mL triple-jacketed glass vessel with an internal diameter of 35 mm. The temperature was controlled using a programmable thermostatic water bath (F12-ED, JULABO Labortechnik GmbH). A magnetic stirrer was used to ensure adequate mixing. The accuracy of the temperature control throughout the experiment was ±0.1 K. Supersolubility and Solubility Determination. To measure supersolubility,3 70.0 g of the LiCl solution or mixed LiCl−NaCl−KCl−Na2SO4 solutions was settled in the crystallizer, and the overhead stirrer, temperature probe, and turbidity probe were immersed in the solution. Then the Crystal SCAN PolyBlock system and thermostatic bath were initiated, where the stirring speed was 500 rpm, and the solution temperature was kept constant. A 9.65% Na2CO3 solution was pumped into the LiCl solution, except for when assessing the impact of the Na2CO3 concentration on the

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RESULTS AND DISCUSSION Solubility of Li2CO3 in LiCl−NaCl−KCl−Na2SO4 Solutions. To verify the measurements, the solubility data of Li2CO3 in pure water taken in mass faction were obtained and are shown in Table 2. As shown in Table 2, the relative deviation (RD) was less than 0.024, suggesting that the solubility values obtained agree well with the reported data.21 Table 2. Solubility of Li2CO3 in Pure Water (Experimental Pressure p = 0.1 MPa)a solubility of Li2CO3 in pure water (100w) temperature (K)

experimental data

literature data21

100 RD

293.15 303.15 313.15 323.15 333.15 343.15 353.15

1.342 1.223 1.135 1.087 1.015 0.959 0.864

1.313 1.244 1.156

2.16 −1.72 −1.85

1.000

1.48

0.843

2.43

Standard uncertainties, u, are u(T) = 0.06 K,u(w) = 2.3 × 10−5, and u(p) = 0.05.

a

Figure 1. (a) Schematic of the supersolubility measurements. (1) Programmable thermostatic bath; (2) turbidity probes; (3) overhead stirrer; (4) temperature probes; (5) crystallizer; (6) PolyBlock; (7) control PC; and (8) liquid dosing system. (b) Schematic of the solubility measurements. (1) Electromagnetic stirrer; (2) jacketed glass vessel; (3) precise thermometer; and (4) programmable thermostatic bath. B



Article

Table 3. Comparison of the Solubility of Li2CO3 in LiCl Solution at 323.15 and 343.15 K solubility of Li2CO3 from ref 16. molarity (mol·kg−1)

mass fraction (%)

concentration of LiCl (mol·L−1)

323.15 K

343.15 K

323.15 K

343.15 K

2.5 3.0 3.3 3.5 4.0

0.00521 0.00383

0.00675 0.00519

0.0346 0.0250

0.0449 0.0337

0.00427 0.00301

0.00263 0.00317

0.0271 0.0188

0.0167 0.0197

solubility of Li2CO3 obtained in this work (%) 323.15 K

343.15 K

0.01647

0.01575

The solubility data in the LiCl solution obtained in this work were also compared with those from ref 16 (see Table 3 and Figure 2) in the mass fraction. It was found that the solubility

Figure 3. Solubility of Li2CO3 in the electrolyte solutions. ■,13.05% LiCl; ▶, 12.81%LiCl−1.39%NaCl−1.80%KCl; ▲, 12.75%LiCl− 1.39%NaCl−1.79%KCl−0.65%Na 2 SO 4 ; ▼ , 12.68%LiCl−1.38% NaCl−1.78%KCl−1.42%Na2SO4; ◀, 12.49%LiCl−1.36%NaCl− 1.75%KCl−3.30%Na2SO4; ●, 12.19%LiCl−5.35%NaCl−1.71%KCl− 3.22%Na2SO4; and ⧫, 11.90%LiCl−9.45%NaCl−1.67%KCl−3.14% Na2SO4.

Figure 2. A comparison of the solubility of Li2CO3 in the LiCl solution at 323.15 and 343.15 K. Ref 16 data: ■, 323.15 K and ●, 343.15 K. Experimental data: ▲, 323.15 K and ▼, 343.15 K.

solubility of Li2CO3 compared with the solubility in the 13.05% LiCl solution (Figure 3), which was in good agreement with the literature.3,16 However, when the Na2SO4 salt was added to the mixed LiCl−NaCl−KCl solution, the salt-in and -out effect influenced the solubility of Li2CO3 differently and changed with salt concentrations. For example, the solubility of Li2CO3 in a 13.05% LiCl solution was higher than the Li2CO3 solubility in a mixed solution of 12.68%LiCl−1.38%NaCl−1.78%KCl−1.42% Na2SO4. The salt-out effect was dominant. While the Na2SO4 concentration was higher than 3.14%, this salt was the primary influence on the solubility, indicating that the salt-in effect was dominant. As the NaCl concentration increased to 5.35%, the solubility of Li2CO3 in a mixed 12.19%LiCl−5.35%NaCl− 1.71%KCl−3.22%Na2SO4 solution was similar to the solubility in a 12.81%LiCl−1.39%NaCl−1.80%KCl solution. The salt-in

value of Li2CO3 deviates from the reported data. As shown in Figure 2, the reported solubility at 343.15 K was higher than that obtained at 323.15 K, which can explain the deviation from this work. Furthermore, the solubility data changed differently when the concentration of LiCl was more than 3.0 mol·L−1. This can also increase the deviation. So we concluded that the solubility data obtained in this work were similar to those in the 3.5 mol·L−1 LiCl solution from ref 16, except for the experimental error. The solubility of Li2CO3 in different electrolyte solutions is summarized in Table 4 and Figure 3. It has been previously shown that NaCl and KCl salts have a salt-out effect on the Li2CO3 solubility, while Na2SO4 has a salt-in effect.3 In this study, the addition of NaCl and KCl salts decreased the

Table 4. Solubility of Li2CO3 in LiCl−NaCl−KCl−Na2SO4 Solutions (Experimental Pressure p = 0.1 MPa)a salt concentration/100w1 LiCl 13.05 12.81 12.75 12.68 12.49 12.19 11.90 a

KCl 1.80 1.79 1.78 1.75 1.71 1.67

NaCl 1.39 1.39 1.38 1.36 5.35 9.45

mass fraction of Li2CO3 in solutions (100w2) Na2SO4

293.15 K

303.15 K

313.15 K

323.15 K

333.15 K

343.15 K

353.15 K

0.65 1.42 3.30 3.22 3.14

0.01886 0.01632 0.01776 0.01834 0.01927 0.01431 0.01106

0.01777 0.01456 0.01641 0.01677 0.01844 0.01347 0.00993

0.01668 0.01367 0.01543 0.01568 0.01734 0.01308 0.00883

0.01647 0.01309 0.01467 0.01515 0.01690 0.01269 0.00855

0.01586 0.01314 0.01480 0.01491 0.01619 0.01231 0.00847

0.01575 0.01321 0.01457 0.01481 0.01596 0.01283 0.00858

0.01579 0.01337 0.01463 0.01499 0.01584 0.01265 0.00869

Standard uncertainties, u, are u(T) = 0.06 K, u(w1) = 0.02, u(w2) = 3.3 × 10−7, and u(p) = 0.05. C



Article

effect of the Na2SO4 in solution was abrogated by the presence of the NaCl salt. When the NaCl salt concentration was increased to 9.45%, the salt-out effect predominated, and the solubility of Li2CO3 decreased rapidly. Therefore, it is necessary to control the amount of sodium sulfate existing in Na−K−Li− SO4−Cl brines during the precipitation crystallization of lithium carbonate. A small amount of sodium sulfate can enlarge the solubility of Li2CO3 and the abrogative effect of a large number of NaCl, which is not beneficial to the reaction crystallization of Li2CO3. Supersolubility and MZW in LiCl−NaCl−KCl−Na2SO4 Solutions. Since the reactive crystallization of Li2CO3 is important during production, it is necessary to investigate the impact of the Na2CO3 concentration on the supersolubility of Li2CO3. The LiCl concentration, feeding rate of Na2CO3, and stirring speed used were 13.05%, 0.8 mL·min−1, and 500 rpm, respectively.3 As shown in Figure 4, the supersolubility initially

Figure 5. The supersolubility and solubility of Li2CO3 in the 13.05% LiCl solution. ●, supersolubility and ■, solubility.

decreased Li2CO3 supersolubility, while solutions containing Na2SO4 had increased supersolubility, which is in accord with previous publications.3 In Figure 6a, it is noticeable that the supersolubility value in a mixed 12.75%LiCl−1.39%NaCl− 1.79%KCl−0.65%Na2SO4 solution was equivalent to that in a LiCl solution at a lower temperature, suggesting that the negative effects of NaCl and KCl on the supersolubility can be overcome by adding a small amount of Na2SO4 salt. When the NaCl concentration was increased to 9.45% (Figure 6b), Na2SO4 had to be added to a concentration of 3.14% to suppress these negative effects. This means that Li2CO3 supersolubility can be controlled by changing the concentration of impurities, such as NaCl, KCl, and Na2SO4, in concentrated lithium solutions during crystallization. The MZW (Δc = csuper − csol) of Li2CO3 in the electrolyte solutions is shown in Figure 7. Notably, the MZW of Li2CO3 decreased as the temperature increased. The impurities had different levels of influence on the MZW and solubility. A wider MZW was observed after Na2SO4 was added compared to when the salt was absent, suggesting that the MZW of Li2CO3 was controlled by the Na2SO4 salt. When the concentration of Na2SO4 was 3.14%, the MZW decreased as the NaCl salt concentrations increased, but it changed a little at the lower temperature.

Figure 4. Impact of the Na2CO3 concentrations on the supersolubility of Li2CO3 at 313.15 K.

increased until it reached its peak with the increasing Na2CO3 concentrations and then decreased gradually once the Na2CO3 concentrations were higher than 9.65%. This is because the concentration is a decisive kinetic factor. At lower level concentrations of Na2CO3, the crystal nucleus of Li2CO3 would dissolve into the mixed solutions before nucleating to a visible size determined by the turbidity probe. However, when the concentration increased, the high concentration led to a much more rapid reaction and nucleation rate; as a result, the nucleus grew rapidly and was determined by the probe, thus obviously reducing the supersolubility. Generally, a wider MZW is preferred during crystallization, so the reactant concentration used in these experiments was 9.65%. Figure 5 presents the supersolubility and solubility curves of Li2CO3 in a 13.05% LiCl solution. The supersolubility was greatly affected by the temperature, where it decreased rapidly as the temperature increased. By contrast, the solubility was minimally affected by these temperature changes. The supersolubility is a decisive kinetic factor in crystallization that is often dependent on the temperature; solute concentration; and stirring, feeding, and cooling rates. High temperatures can accelerate ion movement and increase both ion collision and mass transportation. Therefore, a lower supersolubility was enough to facilitate nucleation. The supersolubility of Li2CO3 in LiCl−NaCl−KCl−Na2SO4 solutions is shown in Figure 6. Compared to pure LiCl solutions, solutions containing NaCl and KCl displayed

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CONCLUSION The solubility, supersolubility, and MZW of Li2CO3 in mixed LiCl−NaCl−KCl−Na2SO4 solutions were measured at temperatures ranging from 293.15 to 353.15 K using turbidity technology. Results showed that the solubility of Li2CO3 was affected by both the salt-out effect of NaCl and KCl salts and the salt-in effect of the Na2SO4 salt in solution. The solubility varied with changes in the salt concentrations. With increasing Na2CO3 concentrations, the supersolubility increased initially until it peaked at its maximum value and then decreased. The temperature more strongly influenced the supersolubility than the solubility because the ion movement and mass transportation in the solution were dependent on the temperature during nucleation. Impurities of NaCl and KCl decreased the supersolubility of Li2CO3, while Na2SO4 increased the supersolubility. Therefore, the supersolubility of Li2CO3 can be controlled by changing the concentrations of impurities in the concentrated lithium solutions during crystallization. The influence of the impurities on MZW was distinct from the influence on the supersolubility. The MZW was primarily D



Article

Figure 6. The supersolubility of Li2CO3 in electrolyte solutions. ■, 13.05%LiCl; ▶, 12.81%LiCl−1.39%NaCl−1.80%KCl; ▲, 12.75%LiCl−1.39% NaCl−1.79%KCl−0.65%Na2SO4; ▼, 12.68%LiCl−1.38%NaCl−1.78%KCl−1.42%Na2SO4; ◀, 12.49%LiCl−1.36%NaCl−1.75%KCl−3.30% Na2SO4; ●, 12.19%LiCl−5.35%NaCl−1.71%KCl−3.22%Na2SO4; and ⧫, 11.90%LiCl−9.45%NaCl−1.67%KCl−3.14%Na2SO4.

Figure 7. The MZW of Li2CO3 in the electrolyte solutions. ■, 13.05%LiCl; ▶, 12.81%LiCl−1.39%NaCl−1.80%KCl; ▲, 12.75%LiCl−1.39%NaCl− 1.79%KCl−0.65%Na2SO4; ▼, 12.68%LiCl−1.38%NaCl−1.78%KCl−1.42%Na2SO4; ◀, 12.49%LiCl−1.36%NaCl−1.75%KCl−3.30%Na2SO4; ●, 12.19%LiCl−5.35%NaCl−1.71%KCl−3.22%Na2SO4; and ⧫, 11.90%LiCl−9.45%NaCl−1.67%KCl−3.14%Na2SO4. (4) Yin, W. T.; Yan, C. Y.; Ma, P. H.; Li, F. Q. Crystallization kinetics of lithium carbonate. Chem.Engin.(China) 2009, 37 (12), 16−19. (5) BU-103: Global Battery Markets, http://batteryuniversity.com/, 2015. (6) Goonan, T.G. Lithium Use in Batteries: U.S. Geological Survey Circular 1371; U.S. Geological Survey, Reston: VA, 2012. (7) Chagnes, A.; Pospiech, B. A brief review on hydrometallurgical technologies for recycling spent lithium-ion batteries. J. Chem. Technol. Biotechnol. 2013, 88, 1191−1199. (8) Song, P. S.; Xiang, R. J. Utilization and exploitation of lithium resources in salt lakes and some suggestions concerning development of Li industries in China. Mineral Deposits 2014, 5, 977−992. (9) Gruber, P. W.; Medina, P. A.; Keoleian, G. A.; Kesler, S. E.; Everson, M. P.; Wallington, T. J. Global lithium availability. J. Ind. Ecol. 2011, 15, 760−775. (10) Vikström, H.; Davidsson, S.; Höök, M. Lithium availability and future production outlooks. Appl. Energy 2013, 110, 252−266. (11) Naumov, A. V.; Naumova, M. A. Modern state of the world lithium market. Russ. J. Nonferr. Met. 2010, 51, 324−330. (12) Guo, X. Y.; Cao, X.; Huang, G. Y.; Tian, Q. H.; Sun, H. Y. Recovery of lithium from the effluent obtained in the process of spent lithium-ion batteries recycling. J. Environ. Manage. 2017, 198, 84−89. (13) Weng, Y. Q.; Xu, S. M.; Huang, G. Y.; Jiang, C. Y. Synthesis and performance of Li[(Ni1/3Co1/3Mn1/3)1−xMgx]O2 prepared from spent lithium ion batteries. J. Hazard. Mater. 2013, 246−247, 163−172. (14) Chinese National Standard. Battery grade lithium carbonate, YS/ T 582-2006, 2017. (15) Xu, Z. H.; Zhang, H. J.; Wang, R. Y.; Gui, W. J.; Liu, G. F.; Yang, Y. Systemic and direct production of battery-grade lithium carbonate from a saline lake. Ind. Eng. Chem. Res. 2014, 53, 16502−16507.

affected by the Na2SO4 salt in the electrolyte solutions, where it increased when the concentration of Na2SO4 increased.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]; Phone: 86-9716313624. ORCID

Min Wang: 0000-0001-8011-083X Funding

This work was financially supported by the National Natural Science Fund projects of China (U1507202 and 21601197) and Natural Science Foundation for the Youth of Qinghai Province, China (2015-ZJ-934Q) Notes

The authors declare no competing financial interest.

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Study of the Solubility, Supersolubility and Metastable Zone Width of

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