VaporâLiquid Equilibria for the ZnSO4âH2SO4âH2O and MgSO4

Article pubs.acs.org/jced

Vapor−Liquid Equilibria for the ZnSO4−H2SO4−H2O and MgSO4− H2SO4−H2O Systems at (30, 60, 90, and 101.3) kPa Geng Li,† Yan Zhang,† Edouard Asselin,‡ and Zhibao Li*,† †

Key Laboratory of Green Process and Engineering, Institute of Process Engineering, Chinese Academy of Sciences, Beijing 100190, China ‡ Department of Materials Engineering, The University of British Columbia, Vancouver, British Columbia V6T 1Z4, Canada ABSTRACT: Vapor−liquid equilibria (boiling points) data for the ZnSO4−H2SO4−H2O and the MgSO4−H2SO4−H2O systems were determined by the quasi-static ebulliometric method at (30, 60, 90, and 101.3) kPa. The boiling points of the ZnSO4−H2O and the ZnSO4− H2SO4−H2O systems were used to regress new chemical model parameters with the average absolute deviation of only 0.08 K. The model was verified by comparing its predictions of the solubility of zinc sulfate in water from (273.15 to 310.15) K and the activity of water in zinc sulfate solution at 298.15 K with literature data. Furthermore, the distributions of the ZnSO4(aq) (ZnSO4 neutral species), Zn2+, MgSO4(aq) (MgSO4 neutral species) and Mg2+ species were predicted. The ZnSO4−H2SO4−H2O system speciation provided a thermodynamic basis for the empirically optimized zinc sulfate concentrations currently employed in the Zn electrowinning industry.

■

formation of the ZnSO4 ion pair,18 Wang and Dreisinger19 analyzed the effect of temperature on the pH of the ZnSO4− H2SO4−H2O system and built a model capable of closely predicting measurements. There are also several research studies of the MgSO4−H2SO4−H2O system. These studies showed that zinc ferrite containing relatively small amounts of magnesium exhibit significantly lower dissolution rates than pure zinc ferrite.20 Barton and Scott21 developed an empirical expression for calculating the specific conductivity of acidic zinc sulfate electrolytes, which includes the effect of Mg2+. In the present work, we aim to establish speciation-based chemical models for the ZnSO4−H2SO4−H2O and the MgSO4−H2SO4−H2O systems. Specifically, the boiling points of the ZnSO4−H2O and MgSO4−H2O binary systems, as well as the ZnSO4−H 2SO4−H2O and MgSO4−H 2SO4−H2O ternary systems, are determined at (30, 60, 90, and 101.3) kPa. The dissociation constant of ZnSO4(aq) is estimated by Helgeson’s method, and then a comprehensive thermodynamic model for the ZnSO4−H2SO4−H2O system is developed on the basis of the Bromley−Zemaitis model embedded in OLI (OLI Simulation Software, version 9.0). The parameters for the model are obtained by regression of the experimental data for the ZnSO4−H2O and ZnSO4−H2SO4−H2O systems. Default OLI parameters for the MgSO4−H2O and MgSO4−H2SO4− H2O systems were used for calculating boiling points, and the results are compared to the newly measured experimental data.

INTRODUCTION Zinc is produced from a sulfate solution generated by the leaching of oxidized zinc sulfide concentrates through the conventional Roast−Leach−Electrowin (RLE) route or via pressure leaching, which is also followed by elecrowinning.1−4 There is also an increasing demand for metallic zinc originating from secondary resources, for example, zinc ash,5 brass smelting,2 flue dusts from electric arc furnaces,6,7 zinc plant residues,8 and batteries9,10 etc. In all cases, metallic zinc is almost always recovered via electrowinning.11 Thus, the ability to predict the equilibrium data for the ZnSO4−H2SO4−H2O system is important for the production of zinc. The influence of MgSO4 in an H2SO4 medium on the RLE route should also be investigated because the Mg2+ ion is a frequent impurity. The ZnSO4−H2SO4−H2O and MgSO4−H2SO4−H2O systems have been investigated extensively over the years. Copeland and Short12 systematically studied the ternary system from 268.15 K to 373.15 K and drew the diagram of ZnSO4 solubility with temperature at various H2SO4 concentrations. Tartar et al.13 carried out a thermodynamic study of the ZnSO4−H2SO4−H2O system at 298.15 K and first determined the electrolysis energy efficiencies of electrolytic zinc plants by means of the free energy. Liu and Papangelakis14 developed a model for the O2−ZnSO4−H2SO4−H2O system that was able to give an accurate prediction of the system’s speciation. Guerra and Bestetti15 determined a new set of Pitzer model parameters for the ZnSO4−H2SO4−H2O system by correlation with existing isobaric osmotic coefficient data. However, the dissociation of the neutral ZnSO4(aq) species16,17 was not considered in their work. Taking into consideration the © 2014 American Chemical Society

Received: May 13, 2014 Accepted: August 14, 2014 Published: August 22, 2014 3449

dx.doi.org/10.1021/je5004252 | J. Chem. Eng. Data 2014, 59, 3449−3460

Journal of Chemical & Engineering Data

Article

Kwong (SRK) equation of state.28 xi is the liquid phase mole fraction and yi is the liquid phase activity coefficient for component i, calculated by the aqueous activity coefficient model. PSi is the vapor pressure of pure component i at the system temperature T, calculated by the Antoine equation using the parameters listed in Table 1. φSi is the vapor fugacity coefficient of pure component i at T and PSi . θSi is the Poynting pressure correction from PSi to P. The Poynting pressure correction θSi , φi, and φSi could be reasonably assumed to be unity at low equilibrium pressures, and then eq 1 can be simplified to

The model equipped with these new parameters is then used to predict the speciation of the binary and ternary ZnSO4− H2SO4−H2O systems and to explain the optimum conditions for zinc electrowinning.

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EXPERIMENTAL SECTION Experimental Materials. Analytical grade ZnSO4·7H2O and MgSO4·7H2O was used in the experiments and supplied by Xilong Chemical Co., Ltd. with minimum purities of 99.5 % and 99.0 %, respectively. Sulfuric acid was supplied by Beijing Chemical Plant with a mass fraction purity of 95%−98%. All reagents were used without further purification. The water used in the study was deionized water produced in a local laboratory. Experimental Apparatus. The boiling points of the systems were determined by the quasi-static ebulliometric method.22−24 The experimental apparatus22,23 is mainly composed of two ebulliometers, one working ebulliometer filled with the experimental sample of known composition and another reference ebulliometer filled with deionized water. The boiling points were measured by two calibrated microthermometers. The temperature uncertainty was considered to be within ± 0.15 K, including a reading error of ± 0.05 K and a nonimmersed stem correction error of ± 0.10 K. The temperature variation may cause a pressure uncertainty of ± 0.6 kPa. The vapor phases were condensed with glycol solution at 280 K to minimize the loss of vapor. The total fluctuation of the feed liquid composition was estimated to be 0.002 based on the holdup on the inner face of the ebulliometer.25 Experimental Procedure. The experimental procedure was as follows:24 the experimental solution of known composition and the deionized water were charged into separate ebulliometers. The composition of the experimental solution was obtained gravimetrically with a digital analytical balance (model JA5003) with an uncertainty of ± 0.001 g. The system pressure was carefully adjusted to ensure that the expected boiling point of water was reached when the equilibrium was achieved. The boiling point of the experimental solution was then recorded, and the measurement was carried out again at a higher water temperature. The system pressure of (30, 60, 90, and 101.3) kPa was determined by controlling the water temperature at (342.35, 359.15, 369.85, and 373.15) K, respectively, according to the Antoine equation.26 The Antoine equation constants for water are listed in Table 1. The experimental apparatus and procedure were verified by comparing experimental data for the butyl acetate + cyclohexane binary mixture with literature data.22

Pyi = xiγiPiS

The vapor−liquid equilibria for the systems in our work can be described by the following relationship: water (vapor) = water (aq)

a

A

B

C

water

16.3872

3885.70

230.170

(3)

When the system achieves thermodynamic equilibrium, the chemical potential of the vapor and aqueous species should be equal: μv = μvo + RT ln fv = μaq = μaqo + RT ln aaq

(4)

where μV and μaq are the chemical potentials of water in the vapor and liquid phase; μ0vi and μ0aqi are the standard state chemical potentials; R is the gas constant (8.314 J·mol−1·K−1); T is the temperature (K); aaq is the activity of water. Speciation Chemistry for the ZnSO4−H2SO4−H2O and the MgSO4−H2SO4−H2O Systems. The thermodynamic modeling of electrolyte solutions requires suitable representation of all chemical reactions and resulting species.29 The Bromley−Zemaitis model embedded in the OLI software was chosen for the ZnSO4−H2SO4−H2O and the MgSO4−H2SO4− H2O systems. The dominant species, including ZnSO4(aq), Zn2+, SO42−, and HSO4− for the ZnSO4−H2SO4−H2O system, and MgSO4(aq), Mg2+, SO42−, and HSO4− for the MgSO4− H2SO4−H2O system, and the equilibrium reactions among them are listed in Table 2. The other zinc containing species Table 2. Dominant Chemical Species and Their Dissociation Reactions in the ZnSO4−H2SO4−H2O and MgSO4−H2SO4− H2O Systems

Table 1. Antoine Equationa Constants for Water component

(2)

species

dissociation reactions

H2O ZnSO4(aq) MgSO4(aq) H2SO4(aq) H2SO4(aq) HSO4−

H2O = H+ +OH− ZnSO4(aq) = Zn2+ +SO2‑ 4 MgSO4(aq) = Mg2+ +SO2‑ 4 H2SO4(aq) = H+ +HSO‑4 H2SO4 = SO3 +H2O HSO‑4 = H+ +SO2‑ 4

Antoine equation: ln(psi /kPa) = A − B/[(T − 273.15)/K + C].

(Zn(OH)2(aq), Zn(OH)3−, Zn(OH)42−, and ZnOH+) and magnesium species (MgOH+) are present in negligible quantities in the strong acidic solution used here, thus they are neglected. The dissociation reaction of ZnSO4(aq) is presented below as an example:

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MODELING METHODOLOGY Vapor−Liquid Equilibria. In general vapor−liquid equilibria can be expressed as follows:27 Pyi φi = xiγiPiSφiSθiS

ZnSO4 (aq) = Zn 2 + + SO4 2 −

(1)

where P is the system pressure, yi is the vapor phase mole fraction of component i, and φi is the vapor phase fugacity coefficient of component i calculated from the Soave−Redlich−

(5)

The thermodynamic equilibrium constant of this equation is molality-based and is defined as follows: 3450


Journal of Chemical & Engineering Data K=

a Zn 2+ ·aSO24− a ZnSO4 (aq)

=

Article

where β0(m−m) and β0(m−s) are the adjustable parameters for molecule−molecule and molecule−ion interactions, respectively. mm is the concentration of neutral species, and ms is the concentration of ion species. As for the water activity in the system, the formula proposed by Meissner and Kusik34 is adopted in the thermodynamic framework of OLI:

(mZn 2+ ·γZn 2+)(mSO24− ·γSO2−) 4

(mZnSO4 (aq)·γZnSO (aq))

(6)

4

where ai, mi, and γi represent the activity, concentration in molality, and the activity coefficient of species i for the dissociation reaction, respectively. Aqueous Activity Coefficient. To determine the vapor− liquid equilibrium and the equilibrium constants of a given reaction, the relevant activity coefficients are required. The Bromley−Zemaitis equation incorporated in the aqueous model of the OLI platform was selected to calculate the ion activity coefficients, while the Pitzer model was used to describe the ion and molecule activity coefficients in the ZnSO4− H2SO4−H2O and MgSO4−H2SO4−H2O systems. The Bromley−Zemaitis equation is used to represent the ion−ion interaction in this work. The equation was first developed by Bromley30 and empirically modified by Zemaitis31 by adding two new terms. It has been successfully applied to electrolytes with concentrations up to 30 M and temperatures up to 473.15 K. For the cations the activity coefficient is expressed by eq 7:

log γi =

log(a w )mix =

i

1+

⎤ ⎥⎛ | Z | + | Z | ⎞ 2 i j + CijI + DijI 2 ⎥⎜ ⎟ mj ⎥⎝ 2 ⎠ ⎥ ⎦

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RESULTS AND DISCUSSION Boiling Point Measurement for the ZnSO4−H2SO4− H2O and the MgSO4−H2SO4−H2O Systems. The boiling points for the ZnSO4−H2O and the MgSO4−H2O binary systems as well as the ZnSO4−H2SO4−H2O and the MgSO4− H2SO4−H2O ternary systems were determined at (30, 60, 90, and 101.3) kPa. These data for ZnSO4−H2O and ZnSO4− H2SO4−H2O systems are shown in Tables 3 and 4 and Table 3. Boiling Points for the ZnSO4−H2O Binary System at Different Pressuresa mZnSO4 no. 1 2 3 4 5 6 7 8 9 10

(7)

where j indicates all anions in solution, A is the Debye−Huckel parameter, and I is the ionic strength of the solution. B, C, and D are temperature-dependent empirical coefficients, where Bij = B1ij + B2ijt + B3ijt2 (t is the temperature in centigrade) and the other coefficients C and D have the same form. Zi and Zj are the cation and anion charges, respectively. The Pitzer model28 is a general model that describes the activity coefficients as a function of the concentration of ions in aqueous electrolyte solutions, and it can be applied to dilute and concentrated aqueous electrolyte solutions. The Pitzer model is used to describe the molecule−molecule and molecule−ion interactions in this work. The expression of the ionic activity coefficient is32,33 j

+

j

k

k

∑ (∑ mk|Zi|)Cijmj + 1/2 ∑ ∑ Ψijkmjmk j

k

j

k

(8)

where B, C, θ, Ψ are interaction parameters. f is an electrostatic term as a function of ionic strength, and mj is the molality of ion j. The second-order parameters Bij are extended to include molecule−molecule and molecule−ion interaction parameters. In the OLI platform, the Pitzer model for neutral species is written as ln γaq = 2β0(m−m)mm + 2β0(m−s)ms

0.0000 0.2500 0.5000 0.7500 1.0000 1.5000 2.0000 2.5000 3.0000 3.5000

30 kPa

60 kPa

90 kPa

101.3 kPa

342.22 342.32 342.42 342.52 342.62 342.83 343.08 343.38 343.94 344.40

359.13 359.18 359.28 359.33 359.43 359.64 359.89 360.25 360.76 361.32

369.89 369.94 369.99 370.09 370.24 370.35 370.55 370.96 371.47 371.93

373.16 373.21 373.26 373.36 373.51 373.72 373.87 374.36 374.87 375.28

compared to the newly developed model in Figures 1 and 2, respectively. For the ZnSO4−H2O system, the boiling point increases only slightly with increasing molality of ZnSO4 throughout the concentration range investigated. It can be seen that the maximum boiling point temperature increase is only 2.18 K when the molality of ZnSO4 increases from 0 mol· kg−1 to 3.5 mol·kg−1. As to the ZnSO4−H2SO4−H2O system, the measurement was carried out at fixed H2SO4 concentrations of 1 mol·kg−1 to 4 mol·kg−1. The boiling point also slightly increases with increasing concentrations of both ZnSO4 and H2SO4. The maximum temperature increase is 3.66 K at the H2SO4 concentration of 4 mol·kg−1 when the molality of ZnSO4 increases from 0 mol·kg−1 to 2 mol·kg−1. The boiling points for the MgSO4−H2O and the MgSO4−H2SO4−H2O systems are shown in Table 5 and compared to the calculated data in Figure 3. The calculated values are in good agreement with the experimental data, indicating that the default interaction parameters in OLI, listed in Table 6, are reliable in the calculation of boiling points for the MgSO4−H2O and the MgSO4−H2SO4−H2O systems.

+ 2 ∑ θijmj + 1/2|zi| ∑ ∑ Cjkmjmk j

mol·kg

boiling point/K

−1

a Standard uncertainties u are u(T) = 0.15 K, u(p) = 0.6 kPa, and u(m) = 0.0001 mol·kg−1.

ln γm , j = 1/2Zi2f ′ + 2 ∑ mjBij + 1/2Zi2 ∑ ∑ B′jk mjmk j

(10)

j

where (aw0) is the hypothetical water activity of pure electrolyte ij (where i is an odd number for all cations and j is an even number for all anions), Xi represents the cationic fraction (Xi = Ii /Ic), and Yj represents the anionic fraction (Yj = Ij/Ia).

⎡ ⎢ (0.06 + 0.6B )|Z Z | I ij i j + ∑⎢ + Bij 2 ⎢ I ⎛ ⎞ 1.5I j ⎜1 + ⎟ ⎢ |ZiZj| ⎠ ⎝ ⎣

−AZi2

∑ ∑ XiYj log(a w 0)ij

(9) 3451



Article

Table 4. Boiling Points for the ZnSO4−H2SO4−H2O Ternary System at Different Pressuresa ZnSO4 no.

(mol·kg−1)

1 2 3 4 5 6 7 8 9 10

0.0000 0.2500 0.5000 0.7500 1.0000 1.2500 1.5000 1.7500 2.0000 2.5000

1 2 3 4 5 6 7 8 9

0.0000 0.2500 0.5000 0.7500 1.0000 1.2500 1.5000 1.7500 2.0000

1 2 3 4 5 6 7 8 9

0.0000 0.2500 0.5000 0.7500 1.0000 1.2500 1.5000 1.7500 2.0000

1 2 3 4 5 6 7 8 9

0.0000 0.2500 0.5000 0.7500 1.0000 1.2500 1.5000 1.7500 2.0000

boiling point (K) 30 kPa

60 kPa

H2SO4 = 343.18 343.28 343.43 343.64 343.89 344.04 344.24 344.45 344.65 345.31 H2SO4 = 344.40 344.60 344.85 345.07 345.41 345.71 346.02 346.32 346.63 H2SO4 = 345.97 346.32 346.63 346.98 347.34 347.69 348.10 348.50 349.01 H2SO4 = 347.89 348.35 348.71 349.16 349.62 350.08 350.53 350.99 351.40

1.0000 mol·kg−1 360.04 360.20 360.30 360.50 360.71 360.91 361.11 361.32 361.52 362.18 2.0000 mol·kg−1 361.32 361.52 361.72 362.00 362.34 362.64 362.95 363.20 363.56 3.0000 mol·kg−1 363.00 363.30 363.66 363.97 364.32 364.68 365.09 365.55 366.01 4.0000 mol·kg−1 364.88 365.39 365.75 366.21 366.67 367.13 367.59 368.10 368.56

90 kPa

101.3 kPa

370.91 371.01 371.16 371.37 371.57 371.73 371.93 372.13 372.39 373.00

374.21 374.36 374.46 374.72 374.92 375.07 375.28 375.53 375.78 376.39

372.24 372.44 372.70 372.92 373.31 373.62 373.91 374.21 374.51

375.63 375.83 376.04 376.29 376.64 376.95 377.25 377.56 377.86

374.01 374.31 374.67 374.97 375.33 375.68 376.14 376.54 377.05

377.36 377.66 378.01 378.32 378.67 379.03 379.49 379.89 380.40

375.83 376.34 376.80 377.30 377.71 378.17 378.67 379.03 379.64

379.33 379.79 380.20 380.66 381.11 381.52 382.08 382.48 382.99

Figure 1. Isobaric boiling points for ZnSO4 aqueous solutions. Points present experimental data; the solid line represents the values regressed by OLI. 0 0 Δr GT0 = Δr H298.15K − Δr S298.15K {298.15

− θ1[1 − exp[exp(θ2 + θ3T ) + θ4 + (T − 298.15)/ θ5]]}

where the θs are constants with the following values: θ1 = 218.312, θ2 = −12.741, θ3 = 0.018751 K−1, θ4 = −7.84·10−4, and θ5 = 219 K. 0 The standard molar enthalpy ΔrH298.15K and entropy 0 ΔrS298.15K of the reaction are determined by the following equations: 0 0 0 Δr H298.15K = ΔHf,298.15K (Zn 2 +) + ΔHf,298.15K (SO24 −) 0 − ΔHf,298.15K (ZnSO4 )

0 − Sf,298.15K (ZnSO4 )

(14)

The required standard-state thermodynamic data of the equilibrium species used in the calculation were taken from the literature,37,38 and they are listed in Table 7. By combining eqs 11 to 14, the dissociation constants of ZnSO4(aq) at various temperatures were computed. The values of the equilibrium constants obtained were then fit to the following mathematical expression, which is the same as that used in the OLI software: log K = A + B /T + CT + DT 2

Calculation of the Equilibrium Constant. The equilibrium constants (K) of the dissociation reactions, which are sensitive to the variation of chemical species at equilibrium, need to be known prior to system modeling. The equilibrium constant for the dissociation of ZnSO4(aq) is not in the OLI default databank and needs to be determined. The equilibrium constant K can be calculated by eq 11: Δr GT0 RT

(13)

0 0 0 Δr S298.15K = Sf,298.15K (Zn 2 +) + Sf,298.15K (SO24 −)

a Standard uncertainties u are u(T) = 0.15 K, u(p) = 0.6 kPa, and u(m) = 0.0001 mol·kg−1.

ln K = −

(12)

(15)

where A, B, C, and D are the empirical parameters listed in Table 8, and T is the temperature in Kelvin. A comparison between the calculated and the literature39 equilibrium constant of eq 3 is shown in Figure 4. The calculated values are in agreement with the literature data, which indicates that the Helgeson method is reliable for the calculation of the equilibrium constant of the ZnSO4(aq) dissociation reaction. The calculated equilibrium constant can then be used in the model parametrization. Model Parameterization. For the purpose of improving the predictive ability of OLI, new model parameters including Bromley−Zemaitis equation parameters and Pitzer model

(11) 17,35,36

The Helgeson method was adopted to estimate the Gibbs free energy of the reaction: 3452



Article

Figure 2. Isobaric boiling points for ZnSO4−H2SO4 aqueous solutions. Points present experimental data; the solid line represents the values regressed by OLI.

Table 5. Boiling Points for the MgSO4−H2SO4−H2O System at Different Pressuresa MgSO4 no.

−1

(mol·kg )

1 2 3 4 5 6 7

0.5000 1.0000 1.5000 2.0000 2.5000 3.0000 3.5000

1 2 3 4 5 6 7 8 9 10

0.2500 0.5000 0.7500 1.0000 1.2500 1.5000 1.7500 2.0000 2.2500 2.5000

boiling point (K) 30 kPa

60 kPa

H2SO4 = 0.0000 mol·kg−1 342.47 359.22 342.52 359.38 342.59 359.48 342.67 359.53 343.13 359.99 343.38 360.20 343.84 360.76 H2SO4= 1.0000 mol·kg−1 342.52 359.44 343.03 359.60 343.13 359.85 343.14 359.83 343.39 359.79 343.44 360.21 343.84 360.50 344.03 360.74 344.24 360.96 344.45 361.22

90 kPa

101.3 kPa

369.98 370.19 370.29 370.40 370.86 371.01 371.57

373.26 373.36 373.57 373.62 373.96 374.11 374.72

370.21 370.41 370.59 370.54 370.70 371.12 371.42 371.61 371.83 372.09

373.58 373.79 374.01 373.93 373.91 374.22 374.47 374.72 374.92 375.13

Figure 3. Isobaric boiling points for MgSO4−H2SO4 aqueous solutions. Points present experimental data; the solid line represents the values calculated by OLI.

Table 6. Bromley−Zemaitis Parameters for the MgSO4− H2O and the MgSO4−H2SO4−H2O Systems Embedded in OLI

a

Standard uncertainties u are u(T) = 0.15 K, u(p) = 0.6 kPa, and u(m) = 0.0001 mol·kg−1.

parameters B1 B2 B3 C1 C2 C3 D1 D2 D3

parameters were determined via regression of the experimental boiling point data, and these are listed in Tables 9 and 10, respectively. The Bromley−Zemaitis equation parameters for SO42−−Zn2+ interaction as well as the Pitzer model parameters for ZnSO4(aq)−ZnSO4(aq), ZnSO4(aq)−Zn2+, and ZnSO4(aq)− SO42− interactions were regressed with the boiling point data for the ZnSO4−H2O system. The results of the regression are shown in Figure 1, indicating that good agreement was achieved with an overall average absolute deviation for temperature (AAD (T)) of 0.07 K and maximum absolute deviation for temperature (MAD (T)) of 0.19 K. Incorporated with the newly obtained parameters, the Bromley−Zemaitis equation parameters for HSO4−−Zn2+ interaction and Pitzer model parameters for ZnSO4(aq)−HSO4− interaction were

HSO−4 −Mg2+ −1

1.255·10 −1.518·10−5 1.4594·10−6 −1.212·10−3 −6.347·10−6 −7.12·10−8 1.6236·10−6 1.1688·10−7 1.037·10−9

2+ SO2− 4 −Mg

−5.601·10−4 2.3539·10−4 2.4126·10−6 1.538·10−3 −1.171·10−5 −1.12·10−7 −2.852·10−5 1.7038·10−7 1.6039·10−9

obtained by regression of the boiling point data for the ZnSO4−H2SO4−H2O system. The regressed results are in accordance with the experimental data as shown in Figure 2. The deviations between the regressed values and the experimental data were only 0.08 K for AAD (T) and 0.58 K for MAD (T), and these are also listed in Table 11. 3453



Article

Table 7. Relevant Thermodynamic Data

Table 9. Newly Obtained Parameters for the Bromley− Zemaitis Equation

ΔG0f,298.15K

ΔH0f,298.15K

S0f,298.15K

species

kJ·mol−1

kJ·mol−1

J·mol−1·K−1

refs

parameters

Zn2+ SO42− ZnSO4(aq)

−147.277 −744.459 −904.900

−153.385 −909.602 −1047.30

−109.621 18.828 5.020

35 35 36

B1 B2 B3 C1 C2 C3 D1 D2 D3

Table 8. Equilibrium Constant Parameters for ZnSO4(aq) Dissociation species

A

B

C

D

ZnSO4(aq)

−4.80374

457.019

0.01073

−2.49721·10−5

Verification of the Model. The model equipped with new parameters was verified by comparison of predicted results for the solubility of ZnSO4 in water and the activity of water in the ZnSO4−H2O system with literature data. Solubility of ZnSO4 in Water. The solubility of ZnSO4 in water from (273.15 to 310.15) K is predicted and compared with the literature40−42 in Figure 5, and it is demonstrated that the predicted solubility is consistent with the literature data. Activity of Water. The water activity for the ZnSO4−H2O system was also predicted as a function of ZnSO4 concentration at 298.15 K, 101.3 KPa, and it is shown in Figure 6. The predicted water activity by the new model is consistent with the literature,13,43 especially when the concentration of ZnSO4 is less than 2 mol·kg−1. The comparison between the calculation and literature data indicates that the chemical model established is reliable and that it can be used to predict the distribution of species. Species Distribution for the ZnSO4−H2SO4−H2O and MgSO4−H2SO4−H2O Systems. Equipped with the new parameters obtained in the present work, OLI’s Analyzer (9.0) software was adopted to analyze the concentration and temperature effects on the distribution of species in ZnSO4− H2SO4 solutions. The default parameters in OLI were used for the analysis of species distribution for the MgSO4−H2SO4− H2O system. Temperature Effect. The proportions of the dominant species (ZnSO4(aq) and Zn2+) for the ZnSO4−H2O system at 101.3 kPa as a function of temperature are illustrated in Figure 7. It is clear that the fraction of ZnSO4(aq) increases as the

HSO−4 −Zn2+ −1

−3.17485·10 7.56923·10−5 6.07274·10−5 1.113687·10−1 −1.62810·10−4 5.94289·10−7 −2.10959·10−3 −1.79036·10−5 −1.55749·10−6

2+ SO2− 4 −Zn

−2.81710·10−2 −4.42201·10−4 −1.40951·10−5 4.46218·10−3 −1.31407·10−5 6.40664·10−6 4.83579·10−4 −2.21527·10−5 9.48599·10−8

temperature increases, indicating a decrease of the equilibrium constant for the ZnSO4(aq) dissociation reaction. Simultaneously, the proportion of the Zn2+ species decreases with temperature. The relative concentration of ZnSO4(aq) and Zn2+ as a function of total dissolved Zn are 20.72% and 79.28%, respectively, at 273.15 K, but 36.51% and 63.49%, respectively, at 373.15 K. It can be seen in Figure 8 that the fraction of MgSO4(aq) increases with temperature below 353.15 K, and it slightly decreases with temperature above 353.15 K. The opposite tendency can be seen for Mg2+. The other zinc containing species (Zn(OH)2(aq), Zn(OH)3−, Zn(OH)42−, and ZnOH+) and magnesium containing species (MgOH+) in the system represent less than 0.02%, hence, they are neglected. Concentration Effect. Figures 9 and 10 show the distribution of ZnSO4(aq) and Zn2+ species for the ZnSO4− H2O system and MgSO4(aq) and Mg2+ species for the MgSO4−H2O system at 101.3 kPa as functions of ZnSO4 and MgSO4 concentrations, respectively. The proportion of ZnSO4(aq) first decreases to a minimum, then consistently increases, as the concentration of ZnSO4 increases, at both temperatures ((298.15 and 373.15) K). The fraction of Zn2+ species consistently decreases after passing through a maximum as the concentration of ZnSO4 increases. For the MgSO4−H2O system, the proportion of MgSO4(aq) first increases to a maximum, then decreases slightly and consistently, as the concentration of MgSO4 increases, at both temperatures ((298.15 and 373.15) K). An opposite tendency can be observed for Mg2+.

Figure 4. Comparison between calculated and literature equilibrium constant for ZnSO4(aq) dissociation. 3454



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Table 10. Newly Obtained Pitzer Model Parameters parameters

b01

b02(·10−3)

b03(·10−5)

b11

b12(·10−2)

b13(·10−4)

β0(ZnSO4 (aq)−ZnSO4 (aq))

0.171412

−9.97663

7.46288

−0.0996528

−1.10513

3.48378

β0(ZnSO4 (aq)−Zn2+)

0.0243919

−2.11134

−1.27362

−0.186894

−5.07225

−2.30637

β0(ZnSO

0.0257728

−1.98209

−0.353051

−0.128088

−4.81423

−1.67677

β0(ZnSO

1.18856

−6.62322

−0.444864

−3.97087

−4.93223

2.23027

2− 4 (aq) − SO4 )

− 4 (aq) − HSO4 )

Table 11. Deviations (ΔT) between the Experimental Data and Calculated Boiling Points for the ZnSO4−H2SO4−H2O and MgSO4−H2SO4−H2O Systemsa p/kPa 30 60 90 101.3 overall 30 60 90 101.3 overall

|ΔT|av/K ZnSO4−H2O 0.09 0.08 0.06 0.11 0.08 MgSO4−H2O 0.14 0.27 0.34 0.49 0.31

|ΔT|max/K 0.16 0.20 0.17 0.22 0.22 0.32 0.41 0.55 0.62 0.48

p/kPa

|ΔT|av/K

|ΔT|max/K

ZnSO4−H2SO4−H2O 30 0.10 0.43 60 0.08 0.18 90 0.04 0.13 101.3 0.05 0.15 overall 0.07 0.43 MgSO4−H2SO4−H2O 30 0.28 0.66 60 0.48 0.85 90 0.57 0.92 101.3 0.70 0.82 overall 0.51 0.81

a |ΔT| = |Tcal − Texp|; |ΔT|av = (∑ni − 1 |ΔT|k)/n, where n is the number of data points.

Figure 6. Comparison of predicted water activity of zinc sulfate solution at 298.15 K with literature.

Figure 5. Comparison of predicted solubility of zinc sulfate in water at 273.15 K to 310.15 K with literature.

Figure 7. Zinc sulfate speciation as a function of temperature at the zinc sulfate concentration of 1 mol·kg−1.

The distribution of ZnSO4(aq) and Zn2+ in the ZnSO4− H2SO4−H2O system and MgSO4(aq) and Mg2+ species for the MgSO4−H2SO4−H2O system at (298.15 and 373.15) K (101.3 kPa), as a function of H2SO4 concentration is shown in Figures 11 and 12. The fraction of ZnSO4(aq) first decreases to a minimum, then increases significantly as the H2SO4 concentration increases at 298.15 K. When the H2SO4 concentration is larger than 4 mol·kg−1, the proportion of ZnSO4(aq) is 82.35%. At the temperature of 373.15 K, the distribution of ZnSO4(aq) consistently increases as the H2SO4 concentration increases. When the H2SO4 concentration is larger than 3.5 mol·kg−1, the proportion of ZnSO4(aq) is close to 100 %. The reason for this behavior is that at low temperature (298.15 K), the proportion

of SO42− increases when the H2SO4 concentration is increased. The large amount of SO42− in solution strongly affects the equilibrium for ZnSO4(aq) and prevents it from dissociating. However, at high temperature (373.15 K), the equilibrium constant for ZnSO4(aq) dissociation is much smaller than that at low temperature (298.15 K). It can be seen in Figure 12 that the fraction of MgSO4(aq) increases with increasing H2SO4 concentration at 298.15 K. However, at 373.15 K, the distribution of MgSO4(aq) decreases slightly and consistently after passing a maximum at the concentration of 2 mol·kg−1 H2SO4. Predicting Optimum Zinc Electrowinning Conditions. As the standard electrode potential of Zn2+/Zn0 is −0.763 V, 3455



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Figure 8. Magnesium sulfate speciation as a function of temperature at the magnesium sulfate concentration of 1 mol·kg−1.

deposition is possible because of the high hydrogen overpotential at the zinc electrode, which renders the kinetics of Zn0 deposition more favorable.44 Zinc metal is typically produced from solution containing zinc sulfate ((55 to 75) g/L Zn2+) and sulfuric acid ((125 to 175) g/L H2SO4) at (308.15 to 318.15) K with current densities ranging from (400 to 650) A·m−2. These ranges of operating conditions have been determined through years of trial and error.15 Maximizing the potential of the Zn2+/ Zn0 electrode not only reduces the thermodynamic driving force for hydrogen evolution relative to zinc deposition (thus saving power through increased current efficiency) but it also directly reduces the total power consumption of the electrolytic cell by reducing its potential.45 The optimum concentration Zn2+ for electrowinning may be explained by the distribution of Zn species as calculated by the newly obtained chemical model. Figure 13 shows the concentration of Zn2+ as a function of ZnSO4 molality at 313.15 K with the H2SO4 concentration of 1.5 mol·kg−1 (151 g/L). It can be observed that the Zn2+ concentration first increases rapidly to 0.8 mol/kg, then slowly increases, while the ZnSO4 concentration consistently increases from 0 mol·kg−1 to 2 mol·kg−1. With the Zn2+ activity calculation (Figure 14)

Figure 9. Zinc sulfate speciation as a function of zinc sulfate concentration at (298.15 and 373.15) K.

hydrogen should thermodynamically be evolved prior to zinc deposition when a solution of zinc sulfate is electrolyzed. Zinc

Figure 10. Magnesium sulfate speciation as a function of magnesium sulfate concentration at (298.15 and 373.15) K. 3456



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Figure 11. Zinc sulfate speciation as a function of sulfuric acid concentration at (298.15 and 373.15) K with a zinc sulfate concentration of 1 mol· kg−1.

obtained by the new model, the electrode potential for the Zn2+/Zn0 electrode at various ZnSO4 concentrations was predicted by the Nernst equation as depicted in Figure 15. It is clear that the maximum Zn2+/Zn0 electrode potential (−0.7917 V) is obtained at 1.2 mol·kg−1 ZnSO4 (75 g/L Zn2+). A 3-dimensional plot for the potential of the Zn2+/Zn0 electrode as a function of zinc sulfate and sulfuric acid concentrations at 313.15 K is shown in Figure 16. When the sulfuric acid concentration is (125 to 175) g/L H2SO4, the maximum range for the Zn2+/Zn0 electrode potential is in the ZnSO4 concentration range of (0.8 to 1.2) mol·kg−1 ((55 to 75) g/L Zn2+), that is, the optimum Zn electrowinning conditions used in industry.

■

CONCLUSIONS

The ebulliometric method and the apparatus used herein are suitable for measuring the boiling point of the ZnSO4−H2SO4− H2O and the MgSO4−H2SO4−H2O systems. Isobaric boiling point data were determined for the ZnSO4−H2O, MgSO4− H2O binary and the ZnSO4−H2SO4−H2O, MgSO4−H2SO4− H2O ternary systems at (30, 60, 90, and 101.3) kPa. The equilibrium constant of the ZnSO4(aq) dissociation reaction

Figure 12. Magnesium sulfate speciation as a function of sulfuric acid concentration at (298.15 and 373.15) K with a magnesium sulfate concentration of 1 mol·kg−1.

Figure 13. Concentration of Zn2+ as a function of zinc sulfate concentration at 313.15 K with a sulfuric acid concentration of 1.5 mol·kg−1. 3457



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Figure 14. Activity of Zn2+ as a function of zinc sulfate concentration at 313.15 K with a sulfuric acid concentration of 1.5 mol·kg−1.

Figure 15. EZn2+/Zn0 as a function of zinc sulfate concentration at 313.15 K with a sulfuric acid concentration of 1.5 mol·kg−1.

made to it in this work, is a good tool to describe the ZnSO4− H2SO4−H2O and the MgSO4−H2SO4−H2O systems. The newly obtained model successfully predicts the solubility of zinc sulfate in water from (273.15 to 310.15) K and the activity of water in aqueous zinc sulfate solution at 298.15 K. The species distribution of ZnSO4 (MgSO4) varies with temperature, ZnSO4 (MgSO4) concentration, and H2SO4 concentration. The chemical speciation model developed herein enabled an explanation of the rationale for the industrial use of Zn2+ concentrations in the range of (55−75) g/L during Zn electrowinning.

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AUTHOR INFORMATION

Corresponding Author

*Tel/fax: +86-10-62551557. E-mail: [email protected]. Funding

The support of the National Basic Research Development Program of China (973 Program with Grant No. 2013CB632605), Key Program in Science & Technology of Qinghai Province (2012-G-213A), and National Natural Science Foundation of China (21206165) are gratefully acknowledged.

Figure 16. EZn2+/Zn0 as a function of zinc sulfate and sulfuric acid concentrations at 313.15 K.

was calculated by Helgeson’s method. A speciation-based chemical model was established, and the model parameters were obtained via regression of the experimental data. The modeling results indicate that the model, after the improvement

Notes

The authors declare no competing financial interest. 3458



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(20) Filippou, D.; Demopoulos, G. P. A reaction kinetic model for the leaching of industrial zinc ferrite particulates in sulphuric acid media. Can. Metall. Q. 1992, 31, 44−54. (21) Barton, G. W.; Scott, A. C. A validated mathematical model for a zinc electrowinning cell. J. Appl. Electrochem. 1992, 22, 104−115. (22) Li, G.; Li, Z. Determination and prediction of vapor−liquid equilibria for a system containing water + butyl acetate + cyclohexane + ethanol. J. Chem. Eng. Data 2012, 57, 2543−2548. (23) Li, G.; Li, Z.; Asselin, E. Determination and modeling of vapor− liquid equilibria for the sulfuric acid + water + butyl acetate + ethanol system. Ind. Eng. Chem. Res. 2013, 52, 3481−3489. (24) Li, G.; Li, Z. Measurement and modeling of vapor−liquid equilibria for the octane + sulfuric acid + water + ethanol system. J. Chem. Eng. Data 2013, 58, 2044−2050. (25) Zhao, J.; Jiang, X. C.; Li, C. X.; Wang, Z. H. Vapor pressure measurement for binary and ternary systems containing a phosphoric ionic liquid. Fluid Phase Equilib. 2006, 247, 190−198. (26) Poling, B. E.; Prausnitz, J. M.; O’Connell, J. P. The Properties of Gases and Liquids, 5th ed.; McGraw-Hill: New York, 2001. (27) Smith, J. M.; Van Ness, H. C.; Abbott, M. M. Introduction to Chemical Engineering Thermodynamics, 7th ed.; McGraw-Hill: New York, 2004. (28) Pitzer, K. S. Activity Coefficients in Electrolyte Solutions, 2nd ed.; CRC Press: Boca Raton, FL, 1991. (29) Chen, C. C.; Mathias, P. M.; Orbey, H. Use of hydration and dissociation chemistries with the electrolyte-NRTL model. AIChE J. 1999, 45, 1576−1586. (30) Bromley, L. A. Thermodynamic properties of strong electrolytes in aqueous solutions. AIChE J. 1973, 19, 313−320. (31) Zemaitis, J. F. Predicting vapor−liquid equilibria in multicomponent aqueous solutions of electrolytes. In: Newman, S. A., Barner, H. E., Klein, M. Eds. Thermodynamics of Aqueous Systems with Industrial Applications.; American Chemical Society: Washington D.C., 1980. (32) Pitzer, K. S. Thermodynamics of electrolytes. I. Theoretical basis and general equations. J. Phys. Chem. 1973, 77, 268−277. (33) Fürst, W.; Renon, H. Effect of the various parameters in the application of Pitzer’s model to solid−liquid equilibrium. Preliminary study for strong 1−1 electrolytes. Ind. Eng. Chem. Process Des. Dev. 1982, 21, 396−400. (34) Meissner, H. P.; Kusik, C. L. Aqueous solutions of two or more strong electrolytes: Vapor pressure and solubilities. Ind. Eng. Chem. Proc. Des. Dev. 1973, 12, 205−208. (35) Helgeson, H. C. Thermodynamics of complex dissociation in aqueous solution at elevated temperatures. J. Phys. Chem. 1967, 71, 3121−3136. (36) Jackson, K. J.; Helgeson, H. C. Chemical and thermodynamic constraints on the hydrothermal transport and deposition of tin: I. Calculation of the stability of cassiterite at high pressures and temperatures. Geochim. Casmochim. Ac. 1985, 49, 1−22. (37) OLI Systems, Inc. OLI Esp 9.0; OLI Systems, Inc.: NJ, 1999. (38) Bale, C. W.; Pelton, A. D.; Thompson, W. T. Facility for the Analysis of Chemical Thermodynamics (F*A*C*T); Thermofact Ltd./ Ltée & McGill Univ.: Montreal Canada, 1991. (39) Nair, V. S. K.; Nancollas, G. H. Thermodynamics of ion association. Part IV. Magnesium and zinc sulphates. J. Chem. Soc. 1958, 3706−3710. (40) Rudolph, W. W.; Brooker, M. H.; Tremaine, P. R. Raman spectroscopy of aqueous ZnSO4 solutions under hydrothermal conditions: Solubility, hydrolysis, and sulfate ion pairing. J. Solution Chem. 1999, 28, 621−630. (41) Agde, G.; Schimmel, F. Concerning a proof of existence of iron chloride hexahydrates. Z. Anorg. Allg. Chem. 1928, 173, 111−114. (42) Schroder, W. Regarding the relationships of berylium with the group of the vitriol creators and the earth alkali metals. I. Z. Anorg. Allg. Chem. 1936, 228, 129−159. (43) Albright, J. G.; Rard, J. A.; Serna, S.; Summers, E. E.; Yang, M. C. Isopiestic determination of the osmotic and activity coefficients of ZnSO4(aq) at T = 298.15 K, and the standard potential of the

REFERENCES

(1) Corrious, J.-P.; Gely, R.; Viers, P. Thermodynamic and kinetic study of the pressure leaching of zinc sulfide in aqueous sulfuric acid. Hydrometallurgy 1988, 21, 85−102. (2) Jha, M. K.; Kumar, V.; Singh, R. J. Review of hydrometallurgical recovery of zinc from industrial wastes. Resour. Conserv. Recycl. 2001, 33, 1−22. (3) Raghavan, R.; Mohana, P. K.; Verma, S. K. Modified zinc sulphate solution purification technique to obtain low levels of cobalt for the zinc electrowinning process. Hydrometallurgy 1999, 51, 187−206. (4) Gnoinski, J.; Sole, K. C. The influence and benefits of an upstream solvent-extraction circuit on the electrowinning of zinc in sulfate media: the scorpion zinc process. In: Young, C. A., Taylor, P. R., Anderson, C. G., Choi, Y. Eds. Hydrometallurgy 2008: Proceedings of the Sixth International Symposium; Society for Mining, Metallurgy, and Exploration, Inc.: Phoenix, AZ, 2008; pp 600−607. (5) Barakat, M. A.; Mahmoud, M. H. H.; Shehata, M. Hydrometallurgical recovery of zinc from fine blend of galvanization processes. Sep. Sci. Technol. 2006, 41, 1757−1772. (6) Cruells, M.; Roca, A.; Núñez, C. Electric arc furnace flue dusts: characterization and leaching with sulphuric acid. Hydrometallurgy 1992, 31, 213−231. (7) Havlik, T.; Turzakova, M.; Stopic, S.; Friedrich, B. Atmospheric leaching of EAF dust with diluted sulphuric acid. Hydrometallurgy 2005, 77, 41−50. (8) Safarzadeh, M. S.; Moradkhani, D.; Ashtari, P. Recovery of zinc from Cd−Ni zinc plant residues. Hydrometallurgy 2009, 97, 67−72. (9) Sayilgan, E.; Kukrer, T.; Civelekoglu, G.; Ferella, F.; Akcil, A.; Veglio, F.; Kitis, M. A review of technologies for the recovery of metals from spent alkaline and zinc−carbon batteries. Hydrometallurgy 2009, 97, 158−166. (10) Senanayake, G.; Shin, S.-M.; Senaputra, A.; Winn, A.; Pudaev, D.; Avraamides, J.; Sohn, J.-S.; Kim, D.-J. Comparative leaching of spent zinc−manganese−carbon batteries using sulfur dioxide in ammoniacal and sulfuric acid solutions. Hydrometallurgy 2010, 105, 36−41. (11) Herrero, D.; Arias, P. L.; Guemez, B.; Barrio, V. L.; Cambra, J. F.; Requies, J. Hydrometallurgical process development for the production of a zinc sulphate liquor suitable for electrowinning. Miner. Eng. 2010, 23, 511−517. (12) Copeland, L. C.; Short, O. A. Studies of the ternary system ZnSO4−H2SO4−H2O from −5 to 70° and ZnO−SO3−H2O at 25°. J. Am. Chem. Soc. 1940, 62, 3285−3291. (13) Tartar, H. V.; Newschwander, W. W.; Ness, A. T. A thermodynamic study of the system zinc sulfate−sulfuric acid−water at 25°. J. Am. Chem. Soc. 1941, 63, 28−36. (14) Liu, H.; Papangelakis, V. G. Thermodynamic equilibrium of the O2−ZnSO4−H2SO4−H2O system from 25 to 250 °C. Fluid Phase Equilib. 2005, 234, 122−130. (15) Guerra, E.; Bestetti, M. Physicochemical properties of ZnSO4− H2SO4−H2O electrolytes of relevance to zinc electrowinning. J. Chem. Eng. Data 2006, 51, 1491−1497. (16) Filippou, D.; Demopoulos, G. P.; Papangelakis, V. G. A new method for the estimation of acid activities through speciation, and its application to process solutions of ZnSO4−FeSO4−Fe2(SO4)3− H2SO4. In Hydrometallurgy: Fundamentals, Technology and Innovations; Hiskey, H. B., Warren, G. W., Eds.; Soc. for Mining, Metallurgy and Exploration: Littleton, CO, 1993; pp 223−240. (17) Filippou, D.; Demopoulos, G. P.; Papangelakis, V. G. Hydrogen ion activities and species distribution in mixed metal sulfate aqueous systems. AIChE J. 1995, 41, 171−184. (18) Hayes, A. C.; Kruus, P.; Adams, W. A. Raman spectroscopic study of aqueous (NH4)2SO4 and ZnSO4 solutions. J. Solution Chem. 1984, 13, 61−75. (19) Wang, W.; Dreisinger, D. B. The acid−base behavior of zinc sulfate electrolytes: The temperature effect. Metall. Mater. Trans. B 1998, 29, 1157−1166. 3459



Article

electrochemical cell ZnHg x (two phase)|ZnSO 4 (aq)|PbSO 4 (s)| PbHgx(two phase). J. Chem. Thermodyn. 2000, 32, 1447−1487. (44) Graf, G. G. Zinc. Ullmann’s Encyclopedia of Industrial Chemistry (electronic release); Wiley-VCH: Weinheim, Germany, 2007. (45) Parada, T. F.; Asselin, E. Reducing power consumption in zinc electrowinning. J. Miner., Metals, Mater. Soc. 2009, 61, 54−58.

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VaporâLiquid Equilibria for the ZnSO4âH2SO4âH2O and MgSO4

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