Solubility and Phase Transitions of Calcium Sulfate in KCl Solutions

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Ind. Eng. Chem. Res. 2009, 48, 7773–7779

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GENERAL RESEARCH Solubility and Phase Transitions of Calcium Sulfate in KCl Solutions between 85 and 100 °C Liuchun Yang, Baohong Guan,* Zhongbiao Wu, and Xianfa Ma Department of EnVironmental Engineering, Zhejiang UniVersity, Hangzhou 310027, China

The solubility(s) of the three phases of CaSO4, namely, CaSO4 · 2H2O (DH), CaSO4 · 0.5H2O (R-HH), and CaSO4 (AH II), in 0.0-18.0 wt % KCl solutions were systemically investigated at temperatures ranging from 85 to 100 °C. At fixed temperature, the solubility(s) of the three phases all change with KCl concentration and possess a maximum value. The relative magnitudes of the variance of solubility for AH and R-HH are larger than that for DH. This was considered to be correlated to the combined effects of the temperature and concentration of KCl solution on the activity coefficients and water activity. The phase transition behaviors of R-HH and DH are presented with possible intermediate phases, which can be well-explained by the solubility difference of the three forms of CaSO4 and the tendency of forming go¨rgeyite (K2Ca5(SO4)6 · H2O). 1. Introduction Calcium sulfate occurs mainly in three forms: calcium sulfate dihydrate (DH), calcium sulfate hemihydrate (HH), and calcium sulfate anhydrite (AH). β-HH and R-HH are the two forms of HH and have the same X-ray diffraction (XRD) pattern but different differential scanning calorimetry (DSC) characteristics.1 The former type of HH is usually used in the construction industry, while the latter has applications in many areas such as precision instrument molds, ceramics, industrial arts, and architecture due to its superior workability and high strength. There are three forms of AH: soluble AH (γ-AH or AH III), insoluble AH (β-AH or AH II), and high temperature AH (R-AH or AH I).2 In the presence of certain accelerators (Na2SO4, K2SO4, CuSO4, CaO, ash, etc.), insoluble AH can be activated and used as a binder (AH cement) in the building industry.3 Therefore, it also represents very important construction material. Since solubility(s) and phase transitions of calcium sulfates are of great importance for many industry processes, such as building material production and application, desalination, and scaling prevention, a large volume of work has been carried out, among which, the CaSO4-H2O system,4,5 CaSO4-seawater system,6,7 and CaSO4-H3PO4-H2O system8 have been extensively focused on. There have been some efforts to investigate other electrolyte aqueous solutions. Ling9 gained the solubility data of calcium sulfates and the phase equilibrium diagrams in the H2SO4-H2O system. Li10,11 developed a chemical model of the thermodynamics of calcium sulfate on the basis of experimentally determined solubility(s) and calculated the phase transition diagrams in concentrated aqueous HCl-CaCl2 solutions. All of the above-mentioned studies determined the transition condition via the dissolution equilibrium method, and the transition condition has been defined as the one under which the solubility of the two phases involved was equal. In order to find a cost-effective method to utilize large quantities of industrial solid byproduct-synthetic gypsum (i.e., flue gas desulfurization gypsum), we have earlier investigated * To whom correspondence should be addressed. Tel.: +86 571 88273650. Fax: +86 571 88273687. E-mail address: guanbaohong@ zju.edu.cn.

the crystallization routes of R-HH in KCl solutions and indicated that KCl solution could be a potential dehydration medium for R-HH preparation from DH.12 To obtain higher conversion efficiency in R-HH preparation with an aqueous salt solution method, a temperature range around the boiling point is usually favorable. However, little has been addressed on experimentally determined solubility(s) and phase transition behaviors of calcium sulfate in KCl solution at this temperature range. Therefore, in the present work, the solubility(s) of DH, R-HH, and AH in KCl solutions (0.0-18.0 wt %) between 85 and 100 °C have been examined with the dissolution equilibrium method, and the related phase transitions were further discussed for potential production of two useful products, R-HH and AH II. 2. Experimental Section 2.1. Chemicals. Reagent grade KCl (Sinopharm Chemical Reagent Co. Ltd.) and deionized water were used to prepare the solutions in all experiments. Three forms of calcium sulfate with different morphology (see the Supporting Information) were employed as the saturating solid phase. DH is reagent grade with 99.0 wt % purity and has a volume mean particle size of about 10 µm. R-HH is a commercial product with 95 wt % purity and has a volume mean particle size of 30 µm, AH was obtained by calcining the reagent grade DH in an oven at 450 °C for 3 h. The AH prepared by this method was believed to be AH II (insoluble AH). The differential scanning calorimeter (DSC) curves for the three solid substances are given in Figure 1. The endothermic peaks at around 154.5 and 177.2 °C on the DSC curve of DH denote the two-step release of its two water molecules in the crystal when heated, while the exothermic peaks at around 373.7 °C corresponds to the conversion of AH III to AH II.1 It is evident that R-HH possesses an exothermic peak at around 174.0 °C immediately following an endothermic peak at around 165.6 °C on its DSC curve and that no endothermic or exothermic peaks can be found on the DSC curve of AH. 2.2. Procedures. There are usually two ways to determine the solubility: one is the precipitation method and the other the dissolution equilibrium method. The dissolution equilibrium method is thought more reliable because it avoids the complica-

10.1021/ie900372j CCC: $40.75  2009 American Chemical Society Published on Web 07/15/2009

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Ind. Eng. Chem. Res., Vol. 48, No. 16, 2009 Table 1. Equilibrated CaSO4 Phases in KCl Solutiona temperature/°C 85

90

95

100 Figure 1. DSC curves for calcium sulfate dihydrate (DH), R-calcium sulfate hemihydrate (R-HH), and calcium sulfate anhydrite (AH).

tions arising from a reaction process such as precipitation. Thus, the solubility determination becomes possible when equilibrium of unstable and metastable solids in solutions is reached before phase transformation occurs. The experiments were carried out in a 1.5 dm3 doublejacketed glass reactor. The reactor was equipped with a reflux condenser for vapor reflux and a Teflon stirrer for obtaining a homogeneous suspension. Heat was supplied by circulating hot oil through the double walls of the reactor. The temperature of the suspension was monitored by a thermometer inserted into the reactor and controlled at an expected value with a deviation of (0.3 °C. First, an aqueous KCl solution was added into the reactor and preheated to the expected temperature. Then, 20 g solid substance (DH, R-HH, or AH) preheated at 60 °C in an oven was quickly put into the reactor (the solid substance was excessive for saturation). After that, the mixture was kept at a constant temperature until the dissolution process completed. During the batch operation, 15 cm3 sample was withdrawn at different time intervals and immediately filtered with a 0.45 µm cellulose filter. The clear filtrate was taken by a 5 cm3 glass pipet and diluted with deionized water in a 50 cm3 volumetric flask. The solubility(s) of calcium sulfate, expressed as grams of CaSO4/100 cm3, was obtained from the sulfate ion analysis. Sulfate ion analysis was determined by PCMultiDirect COD Vario Moving Laboratories (ET99731, Tintometer GmbH Germany) at a wavelength of 450 nm according to a turbidity method. The solid phases were quickly washed with boiling deionized water three times and separated by vacuum filtration, rinsed with acetone once, and dried at 60 °C in an oven for 4 h. Then, the solid phases were examined by a metallographic microscope to investigate the crystal morphology. X-ray diffraction (XRD, D/Max-2550 pc, Rigaku, Inc., Japan) analysis was conducted using Cu KR radiation at a scanning rate of 8°/min in the 2θ range from 5° to 85°. Thermal analysis was performed on a simultaneous thermogravimetry and differential scanning calorimeter (TG-DSC, NETZSCH STA 409PC Luxx, Selb/Bavaria, Germany) to determine the crystal water contents of the solid phase. 2.3. Equilibration Time and Solubility Determination. The equilibration time between calcium sulfates and KCl solutions should be predetermined to make sure the solubility value can be collected before phase transition occurred. Since each of the three forms of calcium sulfate has a typical DSC curve (see Figure 1) and a theoretical crystal water content (20.93, 6.20, and 0 wt % for DH, HH, and AH, respectively), we could

a

KCl concentration/wt %

equilibration time/min

equilibrated solid phase(s)

1.0-10.0 12.0-14.0 16.0 18.0 1.0-4.0 5.0-14.0 16.0 18.0 1.0-6.0 7.0-10.0 12.0-16.0 18.0 1.0-5.0 6.0-8.0 10.0 12.0 14.0-16.0 18.0

90 60 30 10 90 60 30 10 90 50 15 5 90 70 50 20 10 5

DH DH DH DH DH DH DH+traceAH DH DH DH DH+traceAH DH+traceAH DH DH DH DH DH+traceAH DH+traceAH

The starting material was DH.

determine whether or not phase transition had occurred according to the TG-DSC test of a series of solid samples withdrawn at certain time intervals during a dissolution process. The time interval between the start of the dissolution and the reaction time of the solid sample withdrawn before the onset of phase transition was defined as the equilibration time. However, DH in relatively low concentration KCl solutions (for example, in less than 5 wt % KCl solution at 80-100 °C) showed long time matastability. It was found that sulfate ion concentration achieved equilibration within 90 min. Therefore, according to the specific temperature and KCl concentration, we used 5-90 min as the equilibration time. The metastable R-HH was found to hydrate at 85 °C in low concentration KCl solutions and quickly dehydrate when KCl concentration was above 12 wt %. Therefore, the equilibration time for R-HH solubility measurement in our study varies from 1 to 90 min, depending on temperature and KCl concentration. Although AH is the stable phase at the temperature range investigated, its hydration has been reported to be activated by K+ ions.13 To obtain equilibrium and avoid obvious reaction of K+ ions with AH, we used 30-90 min as the equilibration time for AH solubility measurement. The equilibration time thus determined for the three CaSO4 phases in KCl solutions and the corresponding equilibrated solid phase(s) are listed in Tables 1-3. The solubility of DH and R-HH in pure water was determined to verify the accuracy of the adopted procedure in this paper. An equilibration time of 90 min was found to be appropriate for the solubility of DH determination. However, for HH in pure water, 10 min was adopted as equilibration time at 80 and 85 °C due to a tendency to hydrate, and 40 min was adopted at temperatures higher than 85 °C. Our solubility data and literature values are illustrated in Figure 2. The uncertainty of the measured solubility values is within ( 0.004 g/100 cm3, and the relative deviation is 2.0%. Therefore, the data collected in this study compare reasonably well with published data. 3. Results and Discussion 3.1. CaSO4 Solubility in KCl Solutions: Effects of KCl Concentration and Temperature. The solubility of DH, R-HH, and AH in KCl aqueous solutions was measured from 85 to 100 °C under atmospheric pressure. The investigated concentration of KCl solutions was 0.0-18.0 wt %.

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a

Table 2. Equilibrated CaSO4 Phases in KCl Solution temperature/°C 85

90

95

100

a

KCl concentration/wt %

equilibration time/min

equilibrated solid phase(s)

1.0 3.0-8.0 10.0-12.0 14.0-16.0 18.0 1.0 3.0-5.0 6.0-8.0 10.0-12.0 14.0 16.0-18.0 1.0 3.0-5.0 6.0 7.0-10.0 12.0 14.0 16.0-18.0 1.0 1.0-2.0 3.0-5.0 6.0-10.0 12.0 14.0-18.0

10 15 10 2 1 40 20 10 5 2 1 40 20 15 10 4 2 1 90 60 15 6 3 1

HH HH+traceDH HH HH+traceAH HH+traceAH HH HH HH HH+traceAH HH+traceAH HH+traceAH HH HH HH HH+traceAH HH+traceAH HH+traceAH HH+traceAH HH HH HH HH HH+traceAH HH+traceAH

Figure 3. Solubility of calcium sulfate dihydrate (DH) in KCl solution at 85-100 °C: (O) 85, (0) 90, (g) 100 °C. Lines are fits of corresponding data.

The starting material was R-HH.

Table 3. Equilibrated CaSO4 Phases in KCl Solutiona temperaure/°C 85

90

95 100

a

KCl concentration/wt %

equilibration time/min

equilibrated solid phase(s)

1.0 3.0-12.0 14.0 16.0-18.0 1.0 3.0-5.0 6.0-8.0 10.0-12.0 14.0-18.0 1.0-5.0 6.0-10.0 12.0-18.0 1.0 3.0-10.0 12.0 14.0-18.0

30 90 30 45 40 90 40 60 30 90 60 30 60 90 60 30

AH AH AH AH AH AH AH AH AH AH+traceDH AH AH AH AH AH AH

The starting material was AH.

Figure 2. Solubility of calcium sulfate dihydrate (DH) and R-calcium sulfate hemihydrate (R-HH) in water: (O) DH, data of Partridge and White,5 (∆) HH, data of Hulett and Allen,4 (b) DH, this work, (2) HH, this work.

The solubility of DH in KCl solutions is given in Figure 3. It is evident that the solubility of DH increases with concentration of KCl solution within most part of the range

Figure 4. Solubility of R-calcium sulfate hemihydrate (R-HH) in KCl solution at 85-100 °C: (O) 85, (0) 90, (g) 100 °C. Lines are fits of corresponding data.

investigated. However, maximum values can still be identified from the fit lines, being about 0.80 g/100, 0.82 g/100, and 0.83 g/100 cm3 for 85, 90, and 100 °C, respectively. The concentration of KCl solution corresponding to the maximum solubility decreases with temperature, being 16.3, 14.4, and 13.9 wt %, respectively. Therefore, the solubility of DH seems to be improved a little by increased temperature in a large part of the KCl concentration range, i.e. from 2.6 to about 16.8 wt %. In more concentrated KCl solutions, the solubility of DH tends to decrease, and this trend is enhanced at higher temperature (i.g. 100 °C). The solubility of R-HH as a function of concentration of KCl solution at 85, 90, and 100 °C is given in Figure 4. The dependence of R-HH solubility on concentration of KCl solution is similar to that for DH. The solubility increases with increasing KCl concentration and turns to decrease after passing a maximum value. However, when the effect of temperature is taken into consideration, totally different behavior can be seen from the curves in Figure 4, in which the solubility of R-HH is decreased by elevated temperature. This phenomenon is obvious over most of the KCl concentration range. R-HH obtains a maximum solubility of 0.93 g/100 cm3 in about 12.3 wt % KCl solution at 85 °C, and this maximum value is lowered to 0.88 g/100 cm3 at 100 °C corresponding to a KCl concentration of 12.4 wt %. The variation of solubility of AH with concentration of KCl solution as a function of temperature is illustrated in Figure 5.

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Figure 5. Solubility of calcium sulfate anhydrite (AH) in KCl solution at 85-100 °C: (O) 85, (0) 90, (g) 100 °C. Lines are fits of corresponding data.

AH is the most stable one among the three forms of calcium sulfate and usually has the lowest solubility. AH solubility varies obviously with KCl concentration as DH and HH do, reaching a maximum value of about 0.60 g/100, 0.58 g/100, and, 0.54 g/100 cm3 at 85, 90, and 100 °C, respectively, and then decreases with further increasing KCl concentration. The similar dependence of calcium sulfate solubility on KCl concentration can be explained according to the following thermodynamic analysis. The dissolution equilibrium for CaSO4 solids in electrolyte aqueous solutions is expressed by CaSO4 · nH2O ) Ca2+ + SO24 + nH2O

(1)

In the system of CaSO4 · nH2O(s)-KCl-H2O, the solubility of calcium sulfate is expressed by the molal concentration of Ca2+ ion or sulfate ion: s ) mCa2+ ) mSO42-

(2)

The thermodynamic solubility product for calcium sulfate is given by14 Ksp ) aCa2+aSO42-aH2On ) γCa2+mCa2+γSO42-mSO42-aH2On

(3)

where n is 2, 0.5, or 0 corresponding to DH, HH, and AH. ai and γi are the activity and activity coefficient of the corresponding species. mi is the molal concentration of calcium or sulfate ion. The following relationship can be obtained: s ) mCa2+ ) mSO42- )

(

Ksp γCa2+γSO42-aH2On

)

1/2

(4)

As shown in Figures 3-5, the solubility of calcium sulfate increases with increasing KCl concentration in a wide range. The decrease of the activity of water (n ) 2 or 0.5) and the ion activity coefficients product (γCa2+γSO42-) with increasing KCl concentration could significantly contribute to the solubility increase according to eq 4. In more concentrated KCl solutions, however, the opposite effect on the solubility of calcium sulfate has been observed since apparently the ion activity coefficients product become to have reverse trend with further increasing of electrolyte concentration. Similar effects of the ion activity ) on the solubility of gypsum in coefficients product (γCa2+γSO24 the CaSO4-H3PO4-H2O15 and CaSO4-HCl-H2O10 systems have been reported.

Figure 6. Hydration of calcium sulfate hemihydrate (R-HH) in 5 wt % KCl solution at 85 °C.

3.2. Phase Transitions of Calcium Sulfate in KCl Solution. Recently, we investigated the phase transition routes of R-HH in KCl solutions according to optical observations and solid phase analysis and found that R-HH tended to dehydrate to AH directly or indirectly through an initial hydration and a subsequent dehydration. Therefore, a narrow region of KCl concentration and temperature for R-HH to be metastable has been depicted.12 An example of the hydration processes of R-HH is illustrated in Figure 6. At a temperature of 85 °C, R-HH begins to hydrate within 30 min after mixing with 5 wt % KCl solution and has a typical DSC curve of DH after 1 h, indicating a complete hydration from HH to DH. The short columnar R-HH crystals change to long and slim DH crystals (see the Supporting Information). In fact, the DH thus formed will subsequently dehydrate as the reaction time is prolonged, for example, up to 4-8 h.12 Obviously, AH is the most stable calcium sulfate phase in the aqueous solutions at the temperature range investigated and R-HH should ultimately transform into AH probably via a dissolutionrecrystallization mechanism. However, the exact explanation for this transformation in KCl solutions has not been welldocumented in any of the literature. To further investigate the phase transition behaviors of calcium sulfate in KCl solutions, DH was used as the starting material to be mixed with solutions of different KCl concentration. The phase changes were examined by thermal analysis and demonstrated in Figure 7. As can be seen from the gradually shrinking and finally disappearing endothermic peaks at around 153 and 174 °C on the DSC curves in Figure 7a, DH is dehydrated into AH within 150 min at 95 °C in 10 wt % KCl solution. Furthermore, the crystal water contents obtained from thermal gravimetry (TG) data for solid samples at 30, 110, 130, and 150 min are 20.15%, 13.37%, 6.06%, and 0.08%, respectively. When the KCl concentration is increased up to 16 wt %, the dehydration of DH is significantly enhanced and DH quickly transforms into AH after 75 min of reaction (see Figure 7b). The AH thus obtained in KCl solution is very probably insoluble AH, because soluble AH or AH III can not exist in aqueous medium due to its hydration tendency toward hemihydrate. It should be noted that the reaction product of DH in 10 wt % KCl solution for 130 min has a crystal water content close to that of hemihydrate, but the corresponding DSC curve demonstrates none of the characteristics of R-HH which has a distinctive exothermal peak at about 174 °C. The solid product

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Figure 7. Dehydration of calcium sulfate dihydrate (DH) in (a) 10 and (b) 16 wt % KCl solution at 95 °C.

Figure 8. XRD pattern of the 130 min reaction product of calcium sulfate dihydrate (DH) in 10 wt % KCl solution at 95 °C: (D) DH, (A) AH, (G) go¨rgeyite.

Figure 9. Transformation of calcium sulfate dihydrate (DH) in 18 wt % KCl solution at 95 °C: (D) DH, (G) go¨rgeyite.

at 130 min could be a mixture of AH and DH. Thus, XRD was conducted and the pattern was illustrated in Figure 8. However, the result shows that this solid product is composed of DH, go¨rgeyite, and AH. The absence of HH (R-HH or β-HH) indicates that the transition of DH into AH in the KCl solution may not be a process of gradual release of water molecule from DH crystals, i.e. dehydration. Calcium sulfate can react with KCl and transform to syngenite or go¨rgeyite probably according to the following equations:16,17 2KCl + 2CaSO4 · 2H2O f K2Ca(SO4)2 · H2O(syngenite) + CaCl2 + H2O (5) 5K2Ca(SO4)2 · H2O f K2Ca5(SO4)6 · H2O(görgeyite) + 4K2SO4 + 4H2O

(6)

This transformation is easy to identify through XRD patterns of samples withdrawn from more concentrated KCl solutions. As the XRD patterns in Figure 9 revealed, DH reacts with KCl to produce go¨rgeyite and complete this process within about 80 min in 18 wt % KCl solution at 95 °C. The go¨rgeyite obtained is short, sticklike, and open-ended (see the Supporting Information). Therefore, it seems to be difficult to obtain pure calcium sulfate phase from high concentration KCl solutions. 3.3. Inter-relationship between Phase Transitions and Solubility of Calcium Sulfate in KCl Solution. The interesting thing is that we did not find R-HH as the conversion product of DH in the KCl solution over the concentration and temperature

Figure 10. Solubility(s) of calcium sulfate dihydrate (DH), R-calcium sulfate hemihydrate (R-HH), and calcium sulfate anhydrite (AH) in KCl solution at 95 °C: (0) DH, (O) HH, (∆) AH.

range investigated. However, Marinkovic´18 obtained β-HH in 1.0 M KCl solution and R-HH mixed with K2Ca5(SO4)6 · H2O in 1.5-4.0 M KCl solutions. Although detailed explanation has not been given, the authors suggested that solubility could be responsible for the phenomenon. This prompted us to make a more specific comparison of the solubility(s) of calcium sulfate phases under the conditions investigated. In Figure 10, the solubilities of DH, R-HH, and AH in KCl solution at 95 °C were presented together. It is evident that R-HH has the highest solubility in most of the KCl

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Figure 11. Solubility(s) of calcium sulfate dihydrate (DH), R-calcium sulfate hemihydrate (R-HH), and calcium sulfate anhydrite (AH) in KCl solution at 100 °C: (0) DH, (O) HH, (∆) AH.

concentration range, while AH has the lowest. Therefore, both DH and R-HH are unstable in the KCl solutions and are inclined to transform into the most stable phase, AH. The intersection point of solubility lines of DH and R-HH at 95 °C is at about 17.2 wt % KCl concentration, and this is assumed to be the condition under which DH T R-HH happens. According to this diagram, when the starting material is DH, it should transform into AH in KCl solution with the salt concentration being less than 17.2 wt %. During this transformation, DH is not likely to release water molecules directly from its crystal but more likely to undergo a dissolution and recrystallization process. When the starting material is R-HH, it is also going to transform into AH at the end, but is likely to present an intermediate phase, DH, according to a dissolution-recrystallization mechanism. When KCl concentration is above 17.2 wt %, however, the solubility of DH is higher than that of R-HH. Hence, R-HH tends to transform directly into AH while DH is likely to transform into AH via R-HH (as the intermediate phase) because of sufficiency of aqueous calcium sulfate to precipitate R-HH. At a higher temperature of 100 °C, the solubility curves of DH, R-HH, and AH have a similar trend as that at 95 °C (see Figure 11). However, the intersection point of solubility curves of DH and R-HH shifts to a lower KCl concentration of about 16.8 wt %. We deduced that this intersection point could be shifted to even lower KCl concentration when the temperature was further elevated. Then, Marinkovic´’s report on obtaining R-HH in boiling KCl solutions with salt concentrations varying from 1.5 to 4 M seems to be consistent with our solubility data and phase transfer phenomena, because when transformed to mass percent, 1.5 M is close to 11 wt % which could be near the phase transition point of DH and R-HH at the boiling point. It should be noted that most of our DH dehydration experiments were carried out in the KCl concentration range not surpassing the phase transition point of DH and R-HH. At the same time, R-HH shows an obvious dehydration tendency in solutions with KCl concentration surpassing 8 wt % and temperature ranging from 85 to 100 °C. This fact and the strong tendency of forming double salt could be responsible for the absence of R-HH in our limited DH dehydration experiments in KCl solutions. From the equilibrium diagram of the CaSO4-H2O system,5 it is evident that DH has the lowest solubility below 38 °C, while AH has the lowest solubility above 38 °C and HH has a solubility between that of DH and AH above 98 °C. Thus, according to Ostwald’s rule, DH in water between 38 and 98

°C should transform into AH directly, while below 38 °C HH in contact with water should transform into DH through AH as an intermediate phase and above 98 °C DH should transform into AH through HH as an intermediate phase. In practice, the transformation of HH into DH proceeding via AH had never happened due to the fact that the growth kinetics of AH are far too slow. Therefore, true nucleation and growth kinetics is in fact what makes a crystal phase appear or disappear. The phase transition points defined by the intersections of the solubility curves of the three CaSO4 phases have been reported to shift when some other electrolytes are in solution.11 In a similar way, the differences in the solubility(s) of the three CaSO4 phases in KCl solution as well as true crystallization kinetics should be responsible for the phase transition behaviors between these phases in the CaSO4-KCl-H2O system. DH and HH are unstable in KCl solutions between 85 and 100 °C. The solubility curves of DH and R-HH intersect at high KCl concentration ranges where the thermodynamic driving forces for precipitation of AH and the double salt are significantly enhanced. This suggests that it would be prudent to prepare R-HH or AH II by using KCl solution as the hydrothermal media of DH. 4. Conclusions Solubilities of the three CaSO4 phases, namely, DH, R-HH, and AH II, and the related phase transition behaviors in KCl solution were systematically investigated. The following conclusions have been drawn: With KCl concentrations increasing from 0.0 to 18.0 wt %, the solubilities of the three CaSO4 phases all increase up to a maximum value and, thereafter, decrease. This dependence of solubility on KCl concentration is substantially enhanced when the temperature is elevated from 85 to 100 °C. According to the solubility difference between the three phases of CaSO4, R-HH and DH tend to transform into AH directly or indirectly via certain intermediate phases through a dissolution-recrystallization mechanism in KCl solution. Due to the substantially lowered solubility of AH and the strong tendency of double salt formation, R-HH as the intermediate phase of the conversion of DH could basically not be detected. Acknowledgment We gratefully appreciate financial support from the MOST of the People’s Republic of China through the National HighTech R&D Program (2006AA06Z385) and the Science and Technology Department of Zhejiang Province, China, through the project Science and Technology Plan of Zhejiang Province (2007C23055). Supporting Information Available: Morphology of DH, R-HH, and AH that were used for the solubility determination and morphology of the conversion products of DH and R-HH in KCl solution. This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited (1) Hand, R. J. Calcium sulphate hydrates: a review. Br. Ceram. Trans. 1997, 96, 116. (2) Hanic, F.; Galikova, L.; Havlica, J.; Kapralic, I.; Ambruz, V. Kinetics of the thermal decomposition of CaSO4 in air. Trans. J. Br. Ceram. Soc. 1985, 84, 22. (3) Marinkovic´, S.; Kostic´-Pulek A. Popov, S.; Djinovic´, J.; Trifunovic´, P. The possibility of obtaining beta-anhydrite from waste nitrogypsum. J. Min. Met. 2004, 40B, 89.

Ind. Eng. Chem. Res., Vol. 48, No. 16, 2009 (4) Hulett, G. A.; Allen, L. E. The solubility of gypsum. J. Am. Chem. Soc. 1902, 24, 667. (5) Partridge, E. P.; White, A. H. The solubility of calcium sulfate from 0 to 200°. J. Am. Chem. Soc. 1929, 51, 360. (6) Marshall, W. L.; Slusher, R. Aqueous systems at high temperature. Solubility to 200.degree. of calcium sulfate and its hydrates in sea water and saline water concentrates, and temperature-concentration limits. J. Chem. Eng. Data 1968, 13, 83. (7) Furby, E.; Glueckauf, E.; McDonald, L. A. The solubility of calcium sulphate in sodium chloride and sea salt solutions. Desalination 1968, 4, 264. (8) Sullivan, J. M.; Kohler, J. J.; Grinstead, J. H. Solubility of R-calcium sulfate hemihydrate in 40, 50, and 55% P2O5 phosphoric acid solution at 80, 90, 100, and 110 °C. J. Chem. Eng. Data 1988, 33, 367. (9) Ling, Y.; Demopoulos, G. P. Solubility of calcium sulfate hydrates in (0 to 3.5) mol · kg-1 sulfuric acid solutions at 100 °C. J. Chem. Eng. Data 2004, 49, 1263. (10) Li, Z.; Demopoulos, G. P. Development of an improved chemical model for the estimation of CaSO4 solubilities in the HCl-CaCl2-H2O system up to 100 °C. Ind. Eng. Chem. Res. 2006, 45, 2914. (11) Li, Z.; Demopoulos, G. P. Model-based construction of calcium sulfate phase-transition diagrams in the HCl-CaCl2-H2O system between 0 to 100 °C. Ind. Eng. Chem. Res. 2006, 45, 4517.

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(12) Guan, B.; Ma, X.; Wu, Z.; Yang, L.; Shen, Z. Crystallization routes and metastability of calcium sulfate hemihydrate in potassium chloride solutions under atmospheric pressure. J. Chem. Eng. Data 2009, 54, 719. (13) Singh, N. B. The activation effect of K2SO4 on the hydration of gypsum anhydrite, CaSO4 (II). J. Am. Ceram. Soc. 2005, 88, 196. (14) Li, Z.; Demopoulos, G. P. Solubility of CaSO4 phases in aqueous HCl + CaCl2 solutions from 283 to 353 K. J. Chem. Eng. Data 2005, 50, 1971. (15) Messnaoui, B.; Bounahmidi, T. On the modeling of calcium sulfate solubility in aqueous solutions. Fluid Phase Equilib. 2006, 244, 117. (16) Abu-Eishah, S. I.; Bani-Kananeh, A. A.; Allawzi, M. A. K2SO4 production via the double decomposition reaction of KCl and phosphogypsum. Chem. Eng. J. 2000, 76, 197. (17) Kloprogge, J. T.; Hickey, L.; Duong, L. V.; Martens, W. N.; Frost, R. L. Synthesis and characterization of K2Ca5(SO4)6 · H2O, the equivalent of go¨rgeyite, a rare evaporite mineral. Am. Miner. 2004, 89, 266. (18) Marinkovic´, S.; Kostic-Pulek, A. Products of hydrothermal treatment of selenite in potassium chloride solutions. J. Therm. Anal. Calor. 1999, 57, 559.

ReceiVed for reView March 7, 2009 ReVised manuscript receiVed June 22, 2009 Accepted June 30, 2009 IE900372J