Solid–Liquid Stable Equilibrium of the Quaternary System NaCl–KCl

Article pubs.acs.org/jced

Solid−Liquid Stable Equilibrium of the Quaternary System NaCl− KCl−CaCl2−H2O at 303.15 K Dongchan Li,†,‡ Junsheng Yuan,*,†,‡ Fei Li,§ and Shenyu Li§ †

School of Marine Science and Engineering, ‡Engineering Research Center of Seawater Utilization Technology, Ministry of Education, §School of Chemical Engineering, Hebei University of Technology, 300130 Tianjin, China ABSTRACT: The solubility and the physicochemical properties corresponding to densities and pH values in the solid−liquid system (NaCl−KCl−CaCl2−H2O) at 303.15 K have been studied with the method of isothermal dissolution equilibrium. The diagrams of dry-salt phase, water, and physicochemical properties versus composition in the system were plotted. The dry-salt phase diagram of the system includes one threesalt cosaturated point, three isotherm dissolution curves, and three crystallization regions corresponding to sodium chloride, potassium chloride, and calcium chloride tetrahydrate. Neither solid solution nor double salts were found. On the basis of the extended HW model and its temperature-dependent equation, the values of the Pitzer parameters of β(0), β(1), and Cϕ of NaCl, KCl, and CaCl2, the mixed ion-interaction parameters of θNa,K, θNa,Ca, θK,Ca, ΨNa,K,Cl, ΨNa,Ca,Cl, ΨK,Ca,Cl, Debye−Hückel parameter Aϕ, and the solubility equilibrium constants of the solid phase in the quaternary system at 303.15 K were obtained, and the solubilities of the mentioned quaternary system at 303.15 K were calculated. The results show that the predicted solubilities agree well with the experimental data.

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INTRODUCTION The phase equilibrium of the hexary oceanic system (Na+, K+, Ca2+, Mg2+//Cl−, SO42−−H2O) is involved in desalination and dehydration processes, production of potash fertilizer, formation of natural evaporate deposits,1,2 disposal of nuclear and toxic wastes in rock salt formations,3,4 interpretation of recent salt discoveries on Mars,5,6 corrosion of building materials,7 and so on. It provides the foundations for many fields of research and technology such as oceanology, atmospheric chemistry, environmentology, geochemistry, and especially the chemical engineering process. Recently, in the Nanyishan Section of the Qaidam Basin, China, a huge store of oilfield brine was discovered, in which the hydrochemistry is the calcium chloride type with high concentrations of sodium, potassium, and calcium.8 The brine mostly belongs to the complex system of (Li−Na−K−Ca−Sr−Cl−B4O7−H2O). For comprehensive utilization of this valuable oilfield brine, the solubility data and phase diagram of the relative systems are urgently needed. They can reveal the interaction between brine and minerals and show the crystallization path of the various salts. These data will play a significant role in describing geochemical evolution and guiding the industrial process. The quaternary system (NaCl−KCl−CaCl2−H2O) is a subsystem of the complex system (Li−Na−K−Ca−Sr−Cl− B4O7−H2O). Regarding to thermodynamic equilibrium studies, Van’t Hoff was the first to report a stable phase diagram at 293.15 K using the isothermal dissolution method.1 Although some subsystem research on phase equilibria in this hexary system has been reported,9−13 the knowledge is still not sufficient enough in certain subsystems, particularly in CaCl2containing systems, and the parametrization of thermodynamic © XXXX American Chemical Society

models describing the solid−liquid equilibria in these systems requires a large number of reliable experimental data as well. In this paper, the solubility and physicochemical property (density and pH) data of the quaternary system (NaCl−KCl−CaCl2− H2O) at 303.15 K were determined with the isothermal dissolution method. Furthermore, the predictive solubilities were calculated on the basis of Pitzer’s model and its extended HW model.

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EXPERIMENTAL SECTION Apparatus and Reagents. The instruments in this equilibrium experiment mainly include an equilibrium tube, an electric stirrer, and a thermostat. The electric stirrer is placed in close to the bottom of the equilibrium tube. When the electric stirrer began to work, the mixed solution of salt and water in the sealed equilibrium tube was fully stirred for enough time until the solid−liquid equilibrium was reached. The entire experiment process was carried out in the thermostat which can automatically control the temperature at (303.15 ± 0.1) K. The solid phase minerals were identified with a combination of a XP-300 digital polarizing microscope (Shanghai Caikon Optical Instrument Co., Ltd., China) and an X-ray diffractometer (X’pert PRO, Spectris. Pte. Ltd., The Netherlands). The pH values of the equilibrium solutions were measured by the PHS-3C digital acidometer (Shanghai Precision & Scientific Instrument Co., Ltd.) with a precision of ± 0.01. The Received: February 3, 2015 Accepted: May 15, 2015

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DOI: 10.1021/acs.jced.5b00109 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

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Table 1. Solubility and Physicochemical Property Data of the Quaternary System (NaCl + KCl + CaCl2 + H2O) at 303.15 K and Pressure p = 0.1 MPaa composition of liquid phase, w(B)·102

composition of liquid phase, Jb (g/100 g S)

density ρ

no.

NaCl

KCl

CaCl2

NaCl

CaCl2

H2O

g·cm−3

pH

equilibrium solid phase

1,E1 2 3 4 5 6 7 8 9 10 11 12 13 14,E2 15 16 17 18,E 19,E3 20 21

19.59 18.03 16.79 15.51 14.03 11.86 10.22 8.15 6.21 4.38 2.60 1.26 0.77 0.55 0.54 0.52 0.50 0.48 0.00 0.23 0.36

11.72 10.38 9.76 9.19 8.35 7.38 6.47 5.60 4.84 3.99 3.47 3.33 3.30 0.00 1.09 2.02 3.09 4.42 4.34 4.24 4.32

0.00 2.77 4.46 6.39 8.30 11.07 14.16 17.46 21.30 25.25 28.84 32.57 34.60 48.42 48.64 48.84 49.04 49.17 49.25 49.21 49.24

62.57 57.83 54.14 49.89 45.73 39.13 33.13 26.11 19.20 13.03 7.45 3.39 1.99 1.12 1.07 1.01 0.95 0.89 0.00 0.43 0.67

0.00 8.88 14.38 20.55 27.05 36.52 45.90 55.94 65.84 75.10 82.61 87.65 89.48 98.87 96.76 95.06 93.18 90.94 91.90 91.67 91.32

219.39 220.72 222.48 221.65 225.95 229.92 224.15 220.41 209.12 197.44 186.45 169.11 158.60 104.21 98.93 94.63 90.01 84.69 86.60 86.29 85.46

1.2201 1.2247 1.2321 1.2389 1.2426 1.2514 1.2593 1.2725 1.2896 1.3132 1.3312 1.3612 1.3801 1.5216 1.5223 1.5237 1.5273 1.5294 1.5201 1.5227 1.5266

7.03 6.89 6.73 6.6 6.46 6.32 6.16 6.01 5.9 5.7 5.61 5.5 5.45 4.24 4.28 4.35 4.64 4.96 4.43 4.57 4.81

NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + KCl NaCl + C4 NaCl + C4 NaCl + C4 NaCl + C4 NaCl + KCl + C4 KCl + C4 KCl + C4 KCl + C4

Note: Standard uncertainties u are u(T) = 0.10 K, ur(p) = 0.05; ur(Na+) = 0.01, ur(K+) = 0.0010, ur(Ca2+) = 0.0030, ur(Cl−) = 0.0030; C4, CaCl2· 4H2O.

a

densities (ρ) of the equilibrium solution were determined by a density bottle method with a precision of ± 0.0002 g cm−3. The physicochemical parameters of density and pH were all placed in a thermostat that electronically controlled the temperature, which ensured that the physicochemical parameters were strictly determined at (303.15 ± 0.1) K. Chemicals. The chemicals used in dissolution experiment were of analytical-purity grade and recrystallized before use. The amounts of impurities in the compounds were less than 1.0 % (KCl, 99.5 % (w/w), NaCl, 99.5 % (w/w), and CaCl2·6H2O, 99.0 % (w/w)). The distilled water with electrical conductivity less than 1.0 × 10−4 S·m−1 and pH 6.60 was used in the experiment of artificial brine preparation and chemical analysis. Experimental Methods. The solid−liquid stable equilibrium was studied with the isothermal dissolution method. First, according to phase equilibrium compositions, the appropriate salts and distilled water were placed and sealed in the equilibrium tube. Then the electric stirrer in the bottom of the equilibrium tube began to work until the solid−liquid equilibrium was reached. The entire experiment was carried out in the thermostat, which can automatically control the temperature at (303.15 ± 0.1) K. The liquid phases were sampled every 2 h and analyzed. If two samples gave the identical analysis results, then the solid−liquid equilibrium was reached. A preliminary experiment confirmed that 48 h is the optimum time to ensure equilibrium for the system NaCl− KCl−CaCl2−H2O at 303.15 K. After the mixture was stirred for 48 h, it was allowed to settle and separate for at least 4 h. Then the samples of the equilibrium saturated solutions and wet residues were analyzed, respectively. The liquid samples were taken from the solution with the syringes and equipped with a filter, which was previously heated at 303.15 K slightly to prevent precipitation. Then the samples were weighed and diluted the liquid to 100 mL in volumetric flasks with distilled

water for the chemical analysis. The corresponding wet residues of the samples were separated with vacuum filtration and leached with an inert solvent, which can replace the residual saturated liquid and guarantee the residue be dried quickly in the vacuum filtrating process. Then part of the residue samples were weighed and diluted for the chemical analysis. The other part of the residue samples would be kept in sealed sample bottles and observed as soon as possible using a XP-300D digital polarizing microscope with an oil immersion and further identified using X-ray diffraction. During the XRD identification, the temperature and humidity were correspondingly controlled by the air condition and dehumidifier. Analytical Methods. The composition of K+ in the liquid phases and wet residues was determined by gravimetric methods of sodium tetraphenyl borate with precision ± 0.10 %. The Ca2+ concentration was measured by titration with a standard solution of EDTA in the presence of alkali and Caindicator (precision: ± 0.30 %). The Cl− concentration was titrated with the standard HgNO3 solution in the presence of diphenylcarbazone and bromophenol blue as the mixed indicator (precision: ± 0.30 %),14 and the Na+ concentration was determined with atomic absorption spectrophotometry (AAS), or calculated by subtraction via ion equilibrium. The relative error of the Na+ concentration with AAS method was estimated to be less than 1.0 %.

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RESULTS AND DISCUSSION Quaternary System NaCl−KCl−CaCl2−H2O at 303.15 K. The solubility data and the relevant physical chemical properties (density and pH value) of the (NaCl−KCl−CaCl2− H2O) system at 303.15 K were measured and tabulated in Table 1. The solubilities of salts in the equilibrium solutions were expressed as weight percentage wb and Jänecke index Jb (g/100 g dry salt) respectively. On the basis of the Jänecke B

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index Jb in Table 1, the experimental stable phase diagram and the enlarged partial phase diagram of the system at 303.15 K were plotted, as shown in Figure 1.

Figure 1. Equilibrium phase diagram of the quaternary system (NaCl + KCl + CaCl2 + H2O) at 303.15 K: ●, experimental point; , experimental isotherm curve.

Figure 2. Identification for the minerals of the invariant point E in the system (NaCl + KCl + CaCl2·4H2O) at 303.15 K with polarizing microscopy using an oil immersion method: (a) invariant point E; (b) conoscopic interference pattern of CaCl2·4H2O in invariant point E.

The phase diagram in Figure 1 includes one invariant point (point E), three univariant curves, and three crystallized regions of single salts. Point E is the invariant point of three cosaturated salts, and the compositions of NaCl, KCl, CaCl2 in the liquid phase with Jänecke index Jb are 0.89, 8.17, and 90.94, respectively. The three-univariant curves are corresponding to E1E, E2E, and E3E, indicating cosaturation of two salts. The three crystallization regions are corresponding to sodium chloride (NaCl), potassium chloride (KCl), and calcium chloride tetrahydrate (CaCl2·4H2O). The crystallization area of calcium chloride tetrahydrate is the smallest because of its high solubility, whereas the crystallization areas of sodium chloride and potassium chloride are larger, which indicates that sodium chloride and potassium chloride are more easily saturated and crystallize from the solution in this system as a result of the strong salting-out effect of calcium chloride tetrahydrate. For the solid mineral identification, first, the mineral CaCl2· 4H2O is a biaxial crystal and optically negative, that is, 2ν (−). Second, the minerals NaCl and KCl are of cubic system and optically isotropic, but they have a different value of refractive index. The refractive index of KCl is lower than that of NaCl; therefore, these two minerals can be distinguished through the property of refractive index. With an XP-300D digital polarizing microscope using the oil immersion method, the crystal photos of single and orthogonal polarized light for the minerals in invariant point E (NaCl + KCl + CaCl2·4H2O) and moreover the interference pattern of convergent light for the CaCl2·4H2O in invariant point E were both observed in Figure 2. The solid phases are further confirmed with the X-ray diffraction method, and the XRD pattern in Figure 3 demonstrated that three salts NaCl, KCl, and CaCl2·4H2O coexisted in the wet residue of the invariant point E. There are no solid solutions or double salts found in the quaternary system at 303.15 K. The water diagram of the system at 303.15 K is shown in Figure 4. This diagram (diagram of Jänecke) represents the projection of the univariant curves of the system (NaCl + KCl + CaCl2 + H2O) at 303.15K. It shows that the Jänecke index of water gradually increase with the increasing NaCl index. At the

invariant point E, the Jänecke index of J (H2O) is the smallest, which indicates that the water activity is quite small. A comparison of the equilibria phase diagrams for the quaternary system at (288.1513 and 303.15) K is shown in Figure 5. It was found that (1) the mineral of CaCl2·6H2O transformed into CaCl2·4H2O with the increase of temperature from (288.15 to 303.15) K, and (2) the crystallized regions of NaCl and CaCl2·4H2O increased obviously and the area of KCl crystallization region was decreased significantly, which implies that it is suitable to use the temperature variation effects to separate the potassium-containing mixed salts. On the basis of the experimental data in Table 1, the relationship diagrams of the solution physicochemical properties (densities and pH value) vs the composition of calcium chloride in weight percentage (wCaCl2) are shown in Figure 6. It can be found that the physicochemical properties of the equilibrium solution change regularly with the changing calcium chloride concentration in the system. The densities of the aqueous solutions gradually increase with the increasing calcium chloride concentration, and reach maximum values at invariant point E in Figure 6a, whereas the pH values of the equilibrium solution decreased gradually with increasing calcium chloride concentration on the whole, and reach a singularity point at the cosaturation point E with pH 4.96 in Figure 6b. Solubility Prediction. Ion-Interaction Model. Pitzer and co-workers have developed an ion interaction theory and published a series of papers,15,16 which gave a set of expressions for osmotic coefficients of the solution and mean activity coefficient of electrolytes in the solution. The Pitzer ion− interaction model and its extended Harvie−Weare (HW) model have been successfully used in calculating thermodynamic properties and the solubilities of electrolytes.17,18 Using the activity coefficients and the solubility products of the equilibrium solid phases allowed us to identify the coexisting solid phases and their compositions at equilibrium. Additional work has focused on developing variable-temperature models, C

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Figure 3. X-ray diffraction pattern of the invariant point E (NaCl + KCl + CaCl2·4H2O).

Figure 4. Water phase diagram of the system (NaCl + KCl + CaCl2 + H2O) at 303.15 K.

Figure 6. Relationship between the physicochemical properties and composition of calcium chlorite in the quaternary system (NaCl + KCl + CaCl2 + H2O) at 303.15 K: ▲, experimental value; , experimental relationship diagram.

Figure 5. Comparison of phase diagram of the quaternary system (NaCl + KCl + CaCl2 + H2O) at (288.1513 and 303.15) K: −●−, at 303.15 K; ···○···, at 288.15K.

D

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Table 2. Single Salt and Mixing Ion-Interaction Parameters in the Solution of the Quaternary System at 303.15 K and Pressure p = 0.1 MPa single salts

β(0)

β(1)

NaCl KCl CaCl2

0.07873197 0.05074728 0.30618476 θNa,Ca

0.28124611 0.22481123 1.72604892 ΨNa,K,Cl

Aϕ

θNa,K −0.00395625

0.394819

0.0500

θK,Ca 0.115600

Table 3. Solubility Equilibrium Constant of Minerals in the Quaternary System at 303.15 K and Pressure p = 0.1 MPa NaCl

KCl

CaCl2·4H2O

Ksp

39.152352

9.022468

1690481.052114

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ΨNa,Ca,Cl −0.00300

0.0089770 −0.00107540 0.00151465 ΨK,Ca,Cl −0.00417353

CONCLUSION The solubilities and the relevant physicochemical properties (densities and pH value) of the quaternary system (NaCl + KCl + CaCl2 + H2O) at 303.15 K were determined using the isothermal dissolution method. The solid minerals were identified using a XP-300D digital polarizing microscope with an oil immersion, and further using X-ray diffraction. According to the experimental data and the identification result, the isothermal phase diagram and the physicochemical property (densities and pH value) versus composition diagrams were plotted. The phase diagram includes one invariant point, three univariant curves and three crystallized regions of single salts. There are no solid solutions or double salts formed in the system. With the increase of temperature from (288.15 to 303.15) K, the mineral of CaCl2·6H2O transformed into CaCl2· 4H2O, and the crystallization region of KCl was increased obviously, which implies that it is suitable to use the temperature variation effects to separate the potassiumcontaining mixture salts. On the basis of the extended HW model and temperature-dependent equation, the values of the Pitzer parameters and the solubility equilibrium constants Ksp of the solid minerals in the quaternary system at 303.15 K were calculated. The predictive solubilities agree well with the experimental data, which confirms that the obtained Pitzer parameters and the solubility equilibrium constants are reliable for predicting the solubilities of the quaternary system (NaCl + KCl + CaCl2 + H2O) at 303.15 K.

which will increase the applicability to a number of diverse geochemical systems. The primary focus has been to broaden the models in order to generate parameters at higher or lower temperatures.19−21 Model Parametrization. Pitzer’s single salt parameters β(0), β(1), Cϕ, of NaCl, KCl, and CaCl2, the mixing ion-interaction parameters of θNa,K, θNa,Ca, θK,Ca, ΨNa,K,Cl, ΨNa,Ca,Cl, ΨK,Ca,Cl, the Debye−Hückel parameter Aϕ were obtained using a temperature-dependent equation from the literature and are shown in Table 2. The solubility equilibrium constants Ksp of solid phases NaCl, KCl, and CaCl2·4H2O in the quaternary system at 303.15 K were calculated and are listed in Table 3.

species

−0.00341758

Cϕ

Calculated Solubility. On the basis of the Pitzer ioninteraction model and its extended HW models for aqueous electrolyte, the solubilities of the quaternary system (NaCl− KCl−CaCl2−H2O) at 303.15 K have been calculated. On the basis of calculated solubilities, a comparison between the experimental and predicted phase diagrams of the quaternary system at 303.15 K was drawn in Figure 7. The predicted

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AUTHOR INFORMATION

Corresponding Author

*Tel: +86(22)60202241. Fax: +86(22)60204598. E-mail: [email protected]. Funding

Financial support from the Program of the National Nature Science Foundation of China (21406048 and 21106103), applied basic research plan of Hebei Province (13963103D), the Natural Science Foundation of Tianjin (Grant No. 15JCQNJC06100), and the Specialized Research Fund for the Doctoral Program of Chinese Higher Education (Grant No. 20111208120003) are greatly acknowledged. Notes

Figure 7. Comparison of experimental and calculated phase diagram of the quaternary system (NaCl−KCl−CaCl2−H2O) at 303.15 K. −●−, experimental point; ···○···, calculated point; C4, CaCl2·4H2O.

The authors declare no competing financial interest.

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REFERENCES

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