Phase Equilibrium of the MgSO4â(NH4)2SO4 ... - ACS Publications

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

Phase Equilibrium of the MgSO4−(NH4)2SO4−H2O Ternary System: Effects of Sulfuric Acid and Iron Sulfate and Its Application in Mineral Carbonation of Serpentine Weizao Liu, Guanrun Chu, Hairong Yue, Bin Liang, Dongmei Luo, and Chun Li* College of Chemical Engineering, Sichuan University, Chengdu 610065, China S Supporting Information *

ABSTRACT: The carbonation Mg-rich natural minerals or industrial wastes is an attractive route to store CO2. Recently, an approach involving the indirect mineral carbonation of serpentine with recyclable (NH4)2SO4 or NH4HSO4 is receiving widespread attention. In this study, the solubilities associated with the mineral process (ternary system of (NH4)2SO4−MgSO4− H2O) were measured and calculated using the isothermal method and Pitzer model, respectively. The effects of adding small amounts of H2SO4 and iron sulfate on the solubilities were evaluated. The results showed that the crystalline region of MgSO4· (NH4)2SO4·6H2O (boussingaultite) is larger than those of the other species, which indicates that boussingaultite is crystallized out easily. The presence of H2SO4 at a concentration of up to 10 wt % and iron sulfate of 2.5 wt % almost had no effect of the solubilities. The minimal ratios of liquid to solid during the leaching unit were calculated as 3.45, 2.19, and 1.36 mL/g at 25, 55, and 80 °C, respectively. Compared with the case of 25 °C, the energy consumption for the evaporation could be reduced by 36.5% at 55 °C and 60.6% at 80 °C. During the crystallization of ammonium sulfate, high purity (NH4)2SO4 is hardly obtained due to the small crystalline field.

1. INTRODUCTION The increasing energy demand and the excessive use of fossil fuels have increased the atmospheric concentration of CO2, resulting in global warming. Carbon dioxide capture and storage (CCS) is considered one of the main strategies to reduce CO2 emissions. Among the CCS strategies, carbon dioxide capture and geological storage (CCGS) is being widely advocated and tested, for example, in the oil industry for enhanced oil recovery (EOR).1,2 However, CCGS is a very location-dependent technology and faces high costs of poststorage monitoring. Compared with CCGS, CO2 capture and storage by mineralization (CCSM) offers a safer, monitoring-free alternative for permanently storing CO2. CCMS can be divided into direct and indirect mineral carbonation. Direct carbonation usually involves only a onestep reaction of gaseous CO2 with solid minerals or alkaline wastes and is thus the most straightforward approach for CCS; however, it suffers from slow reaction kinetics and low efficiency as well as impure (and, thus, low value-added) products.3,4 Indirect carbonation ordinarily is comprised of two © XXXX American Chemical Society

successive reactions, which use various chemicals for the initial extraction of Ca or Mg from minerals and the subsequent carbonation of the extracted Ca and Mg. This method is now receiving widespread attention because of its relatively mild reaction conditions, high carbonation conversions of Ca and Mg, and purer (and, thus, more valuable) byproducts.5,6 The key issues facing the indirect carbonation method are how to reduce both the energy consumption (especially from regeneration of the chemicals) and the cost. To date, chemicals such as NH4Cl,7 CH3COOH,8,9 HNO3,10,11 HCl,10,11 NaOH,12 and NH 3 5,6 have been investigated as extractants or precipitators; however, none of them could be regenerated easily or be used at low energy consumption, implying that these processes cannot easily realize a net reduction of CO2 emissions. Recently, mineral carbonation of abundantly available serpentine on earth with recyclable ammonium sulfate Received: December 25, 2017 Accepted: March 30, 2018

A

DOI: 10.1021/acs.jced.7b01113 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

was proposed by Zevenhoven.13−18 In the process shown in Figure 1, serpentine is roasted together with ammonium sulfate

2. EXPERIMENTAL SECTION 2.1. Materials. All chemical reagents used, including magnesium sulfate anhydrous (MgSO4), ammonium sulfate ((NH4)2SO4), sulfuric acid (H2SO4), ferrous sulfate heptahydrate (FeSO4·7H2O), and sodium hydroxide (NaOH) were of analytical grade (AR). Because of high hygroscopicity, ferric sulfate (Fe2(SO4)3) was first dried at 300 °C for 120 min to remove possible water absorption. Their source, purity and CAS number are presented in Table 1. All solutions were prepared with distilled water. Table 1. Chemicals, Source, CAS Number. and Purity name magnesium sulfate anhydrous ammonium sulfate sulfuric acid

Figure 1. Zevenhoven’s process for CO2 mineralization via serpentine and (NH4)2SO4.

(AS) at 400−500 °C, and the mixture is leached using water to obtain MgSO4 solution. Because a certain amount of iron is always present in serpentine and approximately 60% of NH3 in the AS enters the roasted slag in the form of NH4HSO4, (NH4)3H(SO4)2, and (NH4)2SO4,15 the leaching solution actually could be considered a mixed solution containing MgSO4, AS, and small amounts of H2SO4 and FeSO4 or Fe2(SO4)3. After the stepwise precipitation of FeOOH and Mg(OH)2 (which was used to mineralize CO2) with ammonia from the roasting flue gas, an AS-rich mother liquor was achieved. The liquor must be evaporated for the recovery and recycling of AS. The evaporation is an energy intensive process. Clearly, the amount of water being added during the leaching significantly affects the evaporation energy consumption. Knowledge of the solubility in the (NH4)2SO4−MgSO4−H2O ternary system and the effects of adding a small amount of H2SO4 and FeSO4 or Fe2(SO4)3 on the solubility are essential for determining the optimum leaching parameters. In addition, the mother liquor after precipitation of Mg(OH)2 was a mixture solution of (NH4)2SO4 and MgSO4 because the Mg could not be thoroughly precipitated out due to the formation of the NH3−(NH4)2SO4 buffer solution with a low pH value during precipitation. Therefore, phase diagrams for the (NH4)2SO4−MgSO4−H2O system should be established to provide guidance on recovering (NH4)2SO4 from the mother liquor. Unfortunately, only solubility rules regarding salt lakes systems (Mg2+, NH4+//Cl−, SO42−; Na+, Mg2+, NH4+// SO42−; K+, Mg2+, NH4+// SO42−)19−22 have been reported, and the presented systems, especially at high temperatures, are incomplete. In this work, the phase equilibrium data of ternary system (NH4)2SO4−MgSO4−H2O and the effects of a small amount of H2SO4 and FeSO4 or Fe2(SO4)3 on the equilibrium were tested to minimize the water dosage for dissolving the roasting slag of serpentine and guide the recovery of (NH4)2SO4 from the mother liquor. The Pitzer electrolyte solution theory model was used to determine the activity coefficient of the electrolytes involved in the ternary system ((NH4)2SO4−MgSO4−H2O) and the electrolyte solubilities in the solution were calculated.

purity

source

>99%

Chron Chemicals Chron Chemicals Chron Chemicals Chron Chemicals Chron Chemicals Chron Chemicals

>99% 95−98%

ferrous sulfate heptahydrate ferric sulfate

>99%

sodium hydroxide

>98%

>98.5%

CAS nr 7487-88-9 7783-20-2 7664-93-9 7782-93-0 10028-22-5 1310-73-2

2.2. Apparatus and Experimental Procedure. The phase equilibrium data were measured by isothermal method at atmospheric pressure with temperatures of 25, 55, and 80 °C, respectively. All experiments were conducted in a constant temperature water bath with accuracy of ±0.5 °C. The equilibrium container was a three-necked glass reactor fitted with a magnetic stirrer, a thermometer and a reflux condenser. In a ternary system ((NH4)2SO4−MgSO4−H2O) phase equilibrium experiment, a certain proportion of MgSO4, (NH4)2SO4 and 100 mL of H2O was added to the flask, which was tightly sealed and immersed in the water bath. In a quaternary system ((NH4)2SO4−MgSO4−H2SO4−H2O) phase equilibrium experiment, MgSO 4 and (NH 4 ) 2 SO 4 were dissolved in 100 mL of 2.5 wt % or 10 wt % H2SO4. In order to determine the effect of iron sulfates on the solubilities, MgSO4 and (NH4)2SO4 were dissolved in a 100 mL solution contained 2.5 wt % iron sulfate and 2.5 wt % H2SO4. The experimental results showed that the compositions of the liquid phase were unchanged when the equilibrium time was extended to 4 h. Thus, the equilibration time, that is, the time required to reach the equilibrium concentrations, was at least 4 h. When the phase equilibrium was achieved, the stirrer was stopped, and the equilibrium system was allowed to rest for 30 min to layer the equilibrium solution and solid. An injector was used to collect the equilibrium solutions and determine the concentrations of (NH4)2SO4, MgSO4, and H2SO4. In this manner, one point in the liquidus can be obtained. The determination of the solid phase in equilibrium is performed using the wet dregs method23 with X-ray diffraction (XRD) analysis. A certain amount of wet residues (containing equilibrium solution and solid) were removed and then dissolved with distilled water to determine its constituent. According to the lever principle, a line drawn from the constituent of equilibrium solution through the wet residue should, if extended, pass through the constituent of equilibrium solid. In addition, the equilibrium solid was collected and dried at room temperature for XRD analysis. Thus, the constituent of equilibrium solid can be B



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Table 2. Phase Equilibrium Data of (NH4)2SO4−MgSO4−H2O System at 25± 0.5°C and 0.1 ± 0.001 MPaa compositions of liquid phase/wt %

compositions of wet dregs/wt %

exp

MgSO4

(NH4)2SO4

H2O

MgSO4

(NH4)2SO4

H2O

equilibrium solid phaseb

1 2 4 5 6 7 8 9 11 12 13 14 15 16 17 18

26.38 26.27 23.45 19.59 13.55 6.93 2.58 1.77 1.06 0.69 0.47 0.32 0.23 0.18 0.18 0.00

0 2.24 2.88 4.02 5.92 9.35 16.36 20.61 26.11 30.03 34.75 37.44 40.70 42.76 42.78 43.62

73.62 71.49 73.67 76.39 80.53 83.73 81.07 77.62 72.84 69.28 64.78 62.24 59.07 57.05 57.04 56.38

33.63 27.14 25.01 24.40 22.42 20.25 19.17 17.10 15.19 14.83 12.14 10.25 8.28 6.22

5.47 16.00 17.99 23.24 25.54 28.02 29.38 31.62 33.19 35.56 37.25 39.49 44.67 51.04

60.90 56.86 57.00 52.36 52.04 51.73 51.45 51.28 51.62 49.60 50.61 50.26 47.05 42.74

M7 MN + M7 MN MN MN MN MN MN MN MN MN MN MN MN + N MN + N N

a Standard uncertainties (abbreviated as u) for temperature and pressure are u(T) = 0.5 °C, u(P) = 0.001 MPa; relative standard uncertainties ur(w(X)) = 0.0036, respectively, where X can be MgSO4 or (NH4)2SO4. bM7 = MgSO4·7H2O, MN = MgSO4·(NH4)2SO4·6H2O, N = (NH4)2SO4

Table 3. Phase Equilibrium Data of (NH4)2SO4−MgSO4−H2O System at 55 ± 0.5 °C and 0.1 ± 0.001 MPaa compositions of liquid phase/wt %


exp

MgSO4

(NH4)2SO4

H2O

MgSO4

(NH4)2SO4

H2O


1 2 3 4 5 6 7 8 9 10 11 12 13 14

34.07 33.30 28.75 26.15 21.77 18.22 12.77 7.20 4.02 2.39 1.58 1.10 0.76 0.00

0.00 2.60 4.03 4.83 6.64 8.33 11.65 18.10 24.89 31.99 36.88 41.26 45.89 46.23

65.93 64.11 67.23 69.02 71.59 73.45 75.58 74.70 71.09 65.62 61.54 57.64 53.35 53.77

34.12 30.76 29.82 27.38 26.62 24.84 21.66 18.84 18.08 17.22 16.93 8.78

7.05 16.14 20.85 21.09 24.20 25.88 28.35 30.85 34.11 36.85 39.22 54.68

58.82 53.11 49.33 51.53 49.18 49.27 49.99 50.30 47.81 45.94 43.85 36.55

M6 MN + M6 MN MN MN MN MN MN MN MN MN MN MN + N N

a Standard uncertainties (abbreviated as u) for temperature and pressure are u(T) = 0.5 °C, u(P) = 0.001 MPa; relative standard uncertainties ur(w(X)) = 0.0036, respectively, where X can be MgSO4 or (NH4)2SO4. bM6 = MgSO4·6H2O, MN = MgSO4·(NH4)2SO4·6H2O, N = (NH4)2SO4

3. RESULTS AND DISCUSSION

determined by combining the results of the wet dregs method and XRD analysis. By changing the original proportion of MgSO4 and (NH4)2SO4, other points in the liquids and their equilibrium solid can be determined, thus allowing the phase diagraph to be plotted. 2.3. Analysis and Characterization. The concentrations of MgSO4 in the equilibrium solution and wet residue were determined by the titration of ethylenediaminetetraacetic acid (EDTA) standard. The concentrations of ammonium sulfate were determined by a distillation method, in which the ammonia liberated from the solution sample by NaOH was absorbed in a known volume of H2SO4 solution. XRD analyses were performed using a DX-2007 X-ray diffraction spectrometer (Danton, China) operating with a Cu Kα radiation source that was filtered using a graphite monochromator at a frequency of λ = 1.54 nm. The voltage and anode current were 40 kV and 30 mA, respectively. The continuous scanning mode with a 0.03 s interval and 0.05 s set time was used to collect the XRD patterns.

3.1. Solubilities and Phase Diagrams of the (NH4)2SO4−MgSO4−H2O System at Different Temperatures. The solubilities of the ternary (NH4)2SO4−MgSO4− H2O system at 25, 55, and 80 °C and under atmospheric pressure are given in Tables 2, 3 and 4, respectively. As seen from the results of exp 1 in each table, with increasing equilibrium temperature, the solubility of isolated MgSO4 rose monotonously from 26.38% at 25 °C to 35.40% at 80 °C. The solubility of MgSO4 decreased when (NH4)2SO4 was added to the system. Figure 2 shows the corresponding phase diagrams, which were plotted according to the solubility data in Tables 2−4 and literature data at several temperatures.22,23 The solid phase compositions were determined by combining the wet dregs method with XRD analysis; the results are also shown in Tables 2−4. The straight line drawn from the constituent of equilibrium solution to the wet residue should, if extended, pass through the constituent of equilibrium solid. If the extension C



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Table 4. Phase Equilibrium Data of (NH4)2SO4−MgSO4−H2O System at 80± 0.5°C and 0.1 ± 0.001 MPaa compositions of liquid phase/wt %


exp

MgSO4

(NH4)2SO4

H2O

MgSO4

(NH4)2SO4

H2O


1 2 3 4 5 6 7 8 9 10 11 12 13 14

35.40 33.33 29.37 23.79 20.30 17.32 13.17 8.72 6.21 5.41 4.40 2.91 2.46 0.00

0.00 6.16 7.61 10.97 13.29 16.16 20.15 26.46 30.64 34.00 37.11 43.43 46.43 48.50

64.6 60.51 63.02 65.24 66.41 66.52 66.68 64.82 63.14 60.59 58.49 53.67 51.10 51.50

34.26 31.54 28.10 26.99 26.56 24.32 22.17 21.37 17.65 14.68 11.86 9.16

12.57 20.95 22.21 25.13 27.62 29.19 32.06 33.92 34.88 36.78 39.88 57.29

53.16 47.51 49.70 47.88 45.81 46.49 45.77 44.70 47.47 48.54 48.26 33.55

M1 MN + M1 MN MN MN MN MN MN MN MN MN MN MN + N N

Standard uncertainties (abbreviated as u) for temperature and pressure are u(T) = 0.5 °C, u(P) = 0.001 MPa; relative standard uncertainties ur(w(X)) = 0.0036, respectively, where X can be MgSO4 or (NH4)2SO4. bM1 = MgSO4·H2O, MN = MgSO4·(NH4)2SO4·6H2O, N = (NH4)2SO4 a

Figure 2. Phase diagrams of the (NH4)2SO4−MgSO4−H2O system at (a) 25 °C, (b) 55 °C, (c) 80 °C, (the black spots ● represent the compositions of liquid phase, and the blue circles ○ represent the compositions of wet dregs), and (d) comparison of the solubility data between this research and open literature.22,24

D



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3.2. Effects of H2SO4 and Iron Sulfate on the Solubilities of (NH4)2SO4 and MgSO4. According to Zevenhoven’s process,15,25 approximately 60% of NH4+ remains in the roasting slag in the form of ammonium salts, such as NH4HSO4, (NH4)3H(SO4)2, (NH4)2SO4, and (NH4)2Mg(SO4)2. Therefore, the leaching solution from the roasting slag contains bisulfate and is usually acidic, with pH varying from 1 to 6, depending on the roasting conditions. Given the addition of ammonium bisulfate will introduce extra NH4+, H2SO4 was selected to add into the solution to form bisulfate. At the same time, serpentine from different regions usually contains considerable amounts of iron. Iron can be present in different iron oxides forms, such as wüstite, hematite, and magnetite. After roasting with AS, approximately 20% of the iron is dissolved in the form of FeSO4 or Fe2(SO4)3 or as Fex(NH4)ySO4 intermediate.25 It can be calculated that the concentration of iron sulfate in the leaching solution is approximately 2.5%. Thus, the effects of a small amount of H2SO4 and iron sulfate on the solubilities of (NH4)2SO4 and MgSO4 must be evaluated. In this study, a certain amount of MgSO4 and (NH4)2SO4 was dissolved in 100 mL of 2.5 wt % or 10 wt % H2SO4 at 55 °C. The measured equilibrium solubilities are given in Table 5.

line intersects with a straight line connected to two solid phase points, then the equilibrium solid phase consists of the two solid phases. If several extension lines intersect at the same point, then the point is the compositions of single equilibrium solid. Point D in Figure 2 represents the composition point of MgSO4·(NH4)2SO4·6H2O. It was found that most of the extension lines intersected at point D, independent of the temperature, indicating that MgSO4·(NH4)2SO4·6H2O was crystallized out easily. The patterns of XRD for the equilibrium solid phases are shown in Figure 3. From the phase diagrams

Table 5. Equilibrium Compositions of the (NH4)2SO4− MgSO4−H2SO4−H2O System with Initial H2SO4 Concentrations of 2.5% and 10% at 55± 0.5 °C and 0.1 ± 0.001 MPaa

Figure 3. XRD patterns of the equilibrium solids corresponding to the representative experiment at 25 °C.

equilibrium compositions of liquid phase/wt %

and XRD analysis, two isothermal invariant points C and H can be confirmed. The invariant point C corresponds to the coexistence of solids MgSO4·xH2O and MgSO4·(NH4)2SO4· 6H2O with the saturated solution, and point H corresponds to the coexistence of solids (NH4)2SO4 and MgSO4·(NH4)2SO4· 6H2O with the saturated solution. Points B and G represent the solubility points of the single MgSO4 and (NH4)2SO4, respectively. Point A represents the equilibrium solid compositions of MgSO4−H2O system. As seen, the equilibrium solids change with temperature and are MgSO4·7H2O, MgSO4· 6H2O, and MgSO4·H2O at 25, 55, and 80 °C, respectively. Comparison of the solubility data between this research and open literature22,24 is shown in Figure 2d. As seen, the measured data at 25 °C were very close to the reported data and the data at 55 and 80 °C fell in between 30−60 °C and 60−96 °C, respectively, as predicted. The results confirmed that the experimental method used in this research was reliable. There are five crystallization fields: field △ABC corresponds to the equilibrium of crystal MgSO4·xH2O with the saturated solution; field △ACD corresponds to the equilibrium of crystals MgSO4·xH2O and MgSO4·(NH4)2SO4·6H2O with the saturated solution; field DHMC corresponds to the equilibrium of crystal MgSO4·(NH4)2SO4·6H2O with the saturated solution; field △DFH corresponds to the equilibrium of crystals MgSO4·(NH4)2SO4·6H2O and (NH4)2SO4 with the saturated solution; and field △FGH corresponds to the equilibrium of crystal (NH4)2SO4 with the saturated solution. The curve BCMHG represents the liquidus and is far from the H2O point with an increase in temperature, indicating higher solubilities of MgSO4 and (NH4)2SO4. Surprisingly, it was found that with an increasing equilibrium temperature, field △FGH was enlarged, thus indicating the possibility to crystallize (NH4)2SO4 solely at high temperature.

MgSO4

(NH4)2SO4

H2SO4

initial H2SO4/%

equilibrium solid phasea

32.18 28.19 20.92 12.54 8.91 6.00 2.91 1.60 1.00 0.69 27.81 20.70 12.58 9.28 6.02 2.62 1.35 0.86 0.47

2.93 4.02 6.78 11.79 15.10 19.17 27.43 34.41 39.85 44.50 3.88 6.77 11.52 14.93 18.61 26.57 32.93 38.40 44.97

1.79 1.84 2.07 2.23 2.25 2.13 1.97 1.73 1.55 1.48 7.18 7.93 8.43 8.63 8.50 7.87 6.70 6.34 5.59

2.48 2.48 2.48 2.48 2.48 2.48 2.48 2.48 2.48 2.48 9.84 9.84 9.84 9.84 9.84 9.84 9.84 9.84 9.84

MN MN MN MN MN MN MN MN MN MN + M6 MN MN MN MN MN MN MN MN MN + M6

a

Standard uncertainties (abbreviated as u) for temperature and pressure are u(T) = 0.5 °C, u(P) = 0.001 MPa; relative standard uncertainties ur(w(X)) = 0.0036, respectively, where X can be MgSO4 or (NH4)2SO4.

As seen, the equilibrium H2SO4 concentrations were less than the initial ones, and this phenomenon was more noticeable when more MgSO4 and (NH4)2SO4 were dissolved because of the increase in the salting-out effect. A comparison of the solubilities with the ones without addition of H2SO4 is shown in Figure 4. Clearly, the presence of sulfuric acid decreased the solubilities of MgSO4 and (NH4)2SO4 only slightly. When the E



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model26,27 is often used to calculate the activity of each ion in the aqueous phase of the salt water system, which contains only electrolytes. The model can be applied to more than 280 types of high concentration solutions with up to 6 mol electrolytes per kilogram of water of electrolyte solution. Thus, the Pitzer model can be used to calculate the solubility of the ternary (NH4)2SO4−MgSO4−H2O system. According to the above results, three types of equilibrium solids are always present in the ternary system of (NH4)2SO4− MgSO4−H2O at each equilibrium temperature. Taking 25 °C as an example, the equilibrium solids are (NH4)2SO4, MgSO4· 7H2O, and MgSO4·(NH4)2SO4·6H2O. The solubility equilibrium of MgSO4·(NH4)2SO4·6H2O in electrolyte solutions can be described by the following dissolution reaction MgSO4 ·(NH4)2 SO4 ·6H 2O(s) = 6H 2O + 2NH4 + + Mg 2 + + 2SO4 2 −

(1)

The thermodynamic equilibrium constant of MgSO4· (NH4)2SO4·6H2O (K) is given by following equations

Figure 4. Effects of H2SO4 on the solubilities of MgSO4 and (NH4)2SO4 at 55 °C

2 2 ·a 6 K = a NH + · a Mg 2 + · a SO 2 − H2O 4

mole ratio of Mg/NH4 was higher than 0.5, the solubilities in the quaternary system were almost equal to the ternary. Alternatively, at low Mg/NH4 ratios (usually less than 0.5), the solubilities were suppressed by the H2SO4 addition, and this suppression became more obvious at higher concentrations of H2SO4. However, compared with the effect of temperature, the produced suppression, even by 10% H2SO4, can also be neglected. Figure 5 (the data of solubilities are listed in the Table S1 in the Supporting Information) displays the effect of FeSO4 and

4

2 2 = m NH ·γ 3 + · m Mg 2 + · m SO 2 − (NH 4

4

4)2 SO4

2 ·γMgSO ·a H6 2O

(2)

4

where K is constant and only depends on temperature, and γ(NH4)2SO4 and γMgSO4 are the average activity coefficients of MgSO4 and (NH4)2SO4, respectively, for the (NH4)2SO4− MgSO4−H2O mixed electrolyte solution and can be calculated using the Pitzer equation ln γMX =

⎛ 2v ⎞ 1 |zMzX |f ′(I ) + ⎜ M ⎟∑ ma⎡⎣BMa + ⎝ v ⎠ a 2 +

⎡ ⎛ 2vX ⎞ ⎜ ⎟∑ m ⎢B + ⎝ v ⎠ c c ⎢⎣ cX

(∑ mz)CMa⎤⎦ ⎛v ⎞

⎤

(∑ mz)CcX + ⎜⎝ vM ⎟⎠θMc(I)⎥⎥⎦ X

⎡ ⎤ 1 + ∑ ∑ mc ma⎢|ZMZX |Bca′ + (2vMzMCca + vMψMca)⎥ ⎣ ⎦ v c a +

⎡⎛ v ⎞ ⎤ 1 ∑ ∑ mcmc′⎢⎣⎝⎜ X ⎠⎟ψcc′ X + |zMzX|θcc′ ′(I )⎥⎦ v 2 c c′

(3)

Here, subscripts c and a represent cations and anions, respectively, m represents the molality of the ions, Σmz = Σmczc = Σma|za|. BMX, B′MX, and CMX are functions of the Pitzer (1) (2) φ parameters, β(0) MX, βMX, βMX and CMX. θ and θ′ and Ψ are considered the binary and ternary interaction parameters of ionic−ionic interaction, respectively. f(I) represents the longrange electrostatic interaction term, and f ′(I) is the derivative of f(I) f (I ) = −2Aφ Figure 5. Effects of iron sulfates on the solubilities of MgSO4 and (NH4)2SO4 at 55 °C.

2I ln(1 + bI1/2) b

(4)

For the 2:1 electrolyte (0) BMX = βMX +

Fe2(SO4)3 on the solubilities of MgSO4 and (NH4)2SO4 in 2.5 wt % H2SO4 at 55 °C. Clearly, the addition of a small amount of iron sulfates (2.5%) to the quaternary (NH4)2SO4−MgSO4− H2SO4−H2O solution also decreased the solubilities of MgSO4 and (NH4)2SO4 slightly. The decrease in solubility should be related to the salting-out effect of iron sulfates. XRD analysis indicated that no iron-bearing phase was present in the equilibrium solids. 3.3. Solubility Prediction of (NH4)2SO4−MgSO4−H2O System. Currently, the Pitzer electrolyte solution theory

′ = BMX

(1) 2βMX

α 2I

[1 − (1 + αI1/2)exp(−αI1/2)]

(1) 2βMX ⎤ ⎡ ⎛ 1 ⎞ −1 + ⎜1 + αI1/2 + α 2I ⎟exp( −α1/2)⎥ 2 2 ⎢ ⎝ ⎦ 2 ⎠ αI ⎣

CMX =

φ CMX

2 |zMzX |1/2 (5)

For the 2:2 electrolyte F


Journal of Chemical & Engineering Data (0) BMX = βMX +

+

′ = BMX

(2) 2βMX

α22I

(1) 2βMX

α12I

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A least-squares optimization was first used to obtain the parameters θ, θ′ and Ψ in eq 9 via regression. The following objective function F was established

[1 − (1 + α1I1/2)exp(−α1I1/2)]

16

[1 − (1 + α2I1/2)exp(−α2I1/2)]

F=

⎤ ⎛ ⎞ 1 + ⎜1 + α1I1/2 + α12I ⎟exp( −α1I1/2)⎥ ⎝ ⎠ ⎦ 2

where, a0 is the reference activity of MN in the solution (here, exp 17 in Table 2 was selected), and ai represents the activity of any solutions with solid phase MN crystallized, which corresponds to exps. Two to 16 in Table 2. Obviously, F is a function of θ, θ′, and Ψ. The minimal value can be obtained when all partial differential with respect to θ, θ′, and Ψ are equal to zero. Therefore, a system of ternary linear equations can be obtained, and θ, θ′, and Ψ can be calculated. The results are given in Table 7. Therefore, the activity of MgSO4· (NH4)2SO4·6H2O was calculated as 0.0256 by using eq 10

(2) 2βMX ⎡ ⎤ ⎛ ⎞ 1 + 2 2 ⎢ −1 + ⎜1 + α2I1/2 + α2 2I ⎟exp( −α2I1/2)⎥ ⎝ ⎠ ⎣ ⎦ 2 α2 I

CMX =

φ CMX

2 |zMzX |1/2 (6)

The Pitzer parameters of the electrolytes (NH4)2SO4 and MgSO4 involved in the present study at various temperatures can be found in the literature28 and are listed in Table 6.

Table 7. Particle Interaction Parameters of (NH4)2SO4− MgSO4−H2O System at 25 °C

Table 6. Pitzer Parameters of MgSO4 and (NH4)2SO4

MgSO4

(NH4)2SO4

(10)

i=2

(1) 2βMX ⎡ −1 2 2⎢ α1 I ⎣

Electrolyte

∑ (ln ai − ln a0)2

T/ °C

β(0) MX

β(1) MX

β(2) MX

Cφ MX

25 55 80 25 55 80

0.221 0.2003 0.18305 0.04088 0.08408 0.12008

3.343 3.802 4.1845 0.6585 0.85938 1.02678

−37.23 −44.82 −51.145

0.025 0.0406 0.0536 −0.00116 −0.00116 −0.00116

θ(0) Mg−NH4

θ(1) Mg−NH4

ΨMg−NH4−SO4

0.5742

1.9502

−0.0804

Using the Pitzer parameters in Table 6 and the interaction parameters in Table 7, the solubilities of the ternary system can be calculated iteratively using firstOpt software. The calculated results presented in Figure 6 (the detail data of solubilities and

The activity of water is related to the osmotic coefficient ϕ can be obtained using the following equation ln a H2O =

∑i mi 55.51

ϕ

(7)

The osmotic coefficient ϕ can be obtained using the following equation ϕ−1=

1 ∑i mi ⎡ ⎢If ′(I ) − f (I ) + 2∑ ∑ mc ma⎡⎣Bca + IBca′ + 2 ⎢⎣ c a

⎤

(∑ mz)Cca⎤⎦⎥⎥⎦ (8) Figure 6. Experimental and calculated values of solubility for the MgSO4−(NH4)2SO4−H2O system at 25 °C.

According to eqs 3−8, when the molality of (NH4)2SO4 and MgSO4 in the solution are determined, the average activity of electrolytes is a function of θ, θ′, and Ψ and can be expressed as aMX = P(θ , θ′, ψ )

relative error showed in the Table S2 in the Supporting Information) are very close to the experimental data. This result indicates that using the Pitzer model to calculate the liquid−solid phase equilibrium of this system is feasible. 3.4. Applications of the Phase Diagrams of the (NH4)2SO4−MgSO4−H2O System. 3.4.1. Guidance on the Process Parameter Optimization for Dissolution of AS + Serpentine Roasting Slag. The above results demonstrated that the effect of a small amount of acidity and iron sulfates on the solubilities of MgSO4 and (NH4)2SO4 was negligible. Therefore, the ternary (NH4)2SO4−MgSO4−H2O system can be directly used for guidance on the parameters determination (such as leaching temperature and liquid-to-solid ratio) of the

(9)

On the basis of the phase equilibrium principle, the activity with a same solid phase crystallized in the system is the same as its subsystems. Thus, the activity of MgSO4·7H2O and (NH4)2SO4 is calculated using the solubility data of binary systems MgSO4−H2O and (NH4)2SO4−H2O, respectively. The solubility data employed in this study are from the literature,29 and the activity of MgSO4·7H2O and (NH4)2SO4 was calculated as 0.0813 and 0.6362 at 25 °C, respectively. The activity of MgSO4·(NH4)2SO4·6H2O (MN) cannot be obtained from the literature; as a result, we calculate it as an unknown quantity in this paper. G



Article

dissolution unit of Zevenhoven’s process. As shown in Figure 1, approximately 40% NH3 in raw material AS entered the roasting exhaust, and the rest remained in the roasted slag in the form of ammonium salts, such as NH4HSO4, (NH4)3H(SO4)2, (NH4)2SO4, and (NH4)2Mg(SO4)2.15 The extraction of magnesium from serpentine (containing 21.8 wt % of magnesium) was approximately 65%. Assuming that all the NH4+ and extractable magnesium in the roasting slag (3.42 g) were considered (NH4)2SO4 and MgSO4, respectively, 1.8 g of (NH4)2SO4 and 0.7 g of MgSO4 were produced. The mass ratio of (NH4)2SO4 to MgSO4 is equal to 72:28. Figure 7 shows the

the carbonation of the Mg-rich leaching solution with NH4HCO3/(NH4)2CO3, an AS-rich mother liquor was obtained. It was found that the conversion of Mg was only 95%, even excess ammonium carbonate added. A considerable amount of MgSO4 remained in the mother liquor. The residual MgSO4 may have adverse effect on the recovery and recycling of AS. A phase diagram of ternary (NH4)2SO4−MgSO4−H2O system can be used to guide this process. Assuming that the evaporation was conducted under vacuum at 80 °C, the phase diagram of ternary (NH4)2SO4−MgSO4− H2O system at 80 °C shown in Figure 8 was used to analyze the

Figure 7. Schematic diagram for guidance on the dissolution process.

Figure 8. Schematic diagram for guidance on the recovery of (NH4)2SO4.

phase diagrams of the ternary (NH4)2SO4−MgSO4−H2O system at different temperatures. The point K represented the composition of slag containing 72 wt % (NH4)2SO4 and 28 wt % MgSO4; the system point moved along KI from K to I during the dissolution process. The line KI intersected the liquidus of the ternary system at points a, b, and c at 25, 55, and 80 °C, respectively, and the minimum addition amount of water required to completely dissolve the (NH4)2SO4 and MgSO4 in the slag was 82.5, 75, and 65 wt % at 25, 55, and 80 °C, respectively. Thus, the water amount for dissolution of 2.5 g of (NH4)2SO4 and MgSO4 can be calculated as 11.79, 7.5, and 4.64 g at 25, 55, and 80 °C, respectively, which correspond to a liquid to solid (roasting slag) ratio of 3.45, 2.19, and 1.36 mL/g at 25, 55, and 80 °C, respectively. Compared with the case of 25 °C, the liquid-to-solid ratio can be reduced by 36.5% and 60.6% at 55 and 80 °C, respectively. Thus, a marked energy savings can be realized during the subsequent evaporation of the (NH4)2SO4 mother liquor for the recovery and recycling of AS. 3.4.2. Guidance on the Evaporation of the AS Mother Liquor for Recovery of AS. There are two types of processes concerning the recovery of AS from the mother liquor of mineral carbonation of serpentine. In Zevenhoven’s process,12 after precipitation of magnesium hydroxide using aqueous ammonia from the leaching liquor derived from AS + serpentine roasting slag, an AS-rich mother liquor was obtained. Because a NH3−(NH4)2SO4 buffer solution formed during the precipitation that inhibited the increase in pH value of the solution, magnesium could not be thoroughly precipitated, resulting in appreciable quantity of MgSO4 remaining in the AS-rich mother liquor. In Maroto-Valer’s process,30−32 serpentine was leached by using recyclable NH4HSO4. After

evaporation. As seen, two operation lines for the evaporation, lines Ie and Ie′, might occur, depending on the content of MgSO4 in the solution. When the initial MgSO4 content is high (point a represents the initial solution composition), the system composition will move along the line abcde from point a to e during the evaporation. The evaporation can be divided into four stages. The first stage is the evaporation of H2O without crystallizing out any solid when the system point moves from a to b. During the second stage, from b to c, the system point passes through the crystalline region of MgSO4·(NH4)2SO4· 6H2O (MN). MN is crystallized out, and the corresponding equilibrium liquid composition changes along the line bH from point b to H. During the third stage, from c to d, the system point passes through the cocrystalline region of MN and (NH4)2SO4. (NH4)2SO4 starts to crystallize out, accompanied by MN, and the corresponding equilibrium solid composition changes along the line Dd from point D to d while the equilibrium liquid remains unchanged at point H and finally disappears. The last stage is the decomposition of MN into (NH4)2SO4 and MgSO4 accompanied by the dehydration of the crystal water as the corresponding system point and solid point moves from d to e. Clearly, the AS purity of the product obtained after the complete evaporation of the mother solution, which corresponds to point e, is not high (approximately 85%). To improve the purity of (NH4)2SO4, a small amount of crystal product can be first separated when the system point moves to c. The separated crystal is MN, which can be used as N−Mg compound fertilizers. The equilibrium liquid at point H is further evaporated, and a crystal product with AS purity of 95% can be achieved. H



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Alternatively, when the initial MgSO4 content is low (point a′ represents the initial solution composition), the system composition will move along the line a′b′c′d′e′ from point a′ to e′ during the evaporation. The operation line passes through the crystalline region of (NH4)2SO4. Thus, a pure (NH4)2SO4 product can be obtained if a separation is conducted at point c′. The purity of final evaporative crystallization product is over 95%.

ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.7b01113. Additional tables (PDF)

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REFERENCES

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4. CONCLUSIONS In this paper, the solubilities of ternary system (NH4)2SO4− MgSO4−H2O at 25, 55, and 80 °C were measured using the isothermal method. The effects that small amounts of H2SO4 and iron sulfate have on the solubilities were evaluated. The analysis of the phase diagram shows that the crystalline region of MgSO4·(NH4)2SO4·6H2O is greater than those of the other species, which indicates that it is crystallized out easily. The effects of acidity and iron sulfates on the solubilities of MgSO4 and (NH4)2SO4 were tested; the results indicated that they can be neglected. The presence of H2SO4 at a concentration of up to 10 wt % and iron sulfate of 2.5 wt % almost had no effect of the solubilities of the ternary system. The minimal ratios of liquid to solid during the leaching unit in Zevenhoven’s process were calculated as 3.45, 2.19, and 1.36 mL/g at 25, 55, and 80 °C, respectively. Compared with the case of 25 °C, the energy consumption for the evaporation could be reduced by 36.5% at 55 °C and 60.6% at 80 °C. During the crystallization of ammonium sulfate, high purity (NH4)2SO4 is hardly obtained because of the small crystalline field of (NH4)2SO4. On the basis of the Pitzer model of electrolyte solution theory, the interaction parameters were regressed, and the solubilities of the MgSO 4 −(NH 4 ) 2 SO 4 −H 2 O system at 25 °C were calculated. The results show that the calculated values were consistent with the experimental results.

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Article

AUTHOR INFORMATION

Corresponding Author

*Tel.: +86 028 85408056. Fax: +86 02885461108. E-mail: lic@ scu.edu.cn. ORCID

Weizao Liu: 0000-0003-0976-2864 Hairong Yue: 0000-0002-9558-0516 Bin Liang: 0000-0003-2942-4686 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors are grateful for the financial support of the National Key Projects for Fundamental Research and Development of China (2016YFB0600904) and The Fundamental Research Funds for the Central Universities (2012017yjsy112). The authors also thank Professor Jilin Cao from Hebei University of Technology for providing guidance on the calculations of the solubilities using the Pitzer model. I



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