Effect of Mg2+ Ions on the Nucleation Kinetics of Calcium Sulfate in

May 19, 2010 - With the fitted curves of turbidity as a function of time, the induction period was determined according to a similar method reported b...
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Ind. Eng. Chem. Res. 2010, 49, 5569–5574

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Effect of Mg2+ Ions on the Nucleation Kinetics of Calcium Sulfate in Concentrated Calcium Chloride Solutions Baohong Guan,† Liuchun Yang,†,‡ and Zhongbiao Wu*,† Department of EnVironmental Engineering, Zhejiang UniVersity, Hangzhou 310027, China, School of CiVil Engineering and Architecture, East China Jiao Tong UniVersity, Nanchang 330013, China

The precipitation kinetics of the metastable calcium sulfate hemihydrate (HH) from aqueous solutions has seldom been addressed. For the first time, the nucleation kinetics of calcium sulfate (probably in R-HH form) in concentrated CaCl2 solutions at 90 °C and the effect of Mg2+ ions were investigated by spontaneous precipitation experiments and a turbidity monitoring method. The results indicate that, at a moderate supersaturation with respect to calcium sulfate, Mg2+ ions retard the nucleation with 0.01 M Mg2+ ions showing the strongest effect; at a higher supersaturation, the inhibition effect decreases when the concentration of Mg2+ ions increases from 0.001 to 0.10 M, and 0.20 M Mg2+ ions show a promoting effect on the precipitation. The interfacial tension values increase with the concentration of Mg2+ ions, resulting in a decrease in nucleation rate in Mg2+ ion-containing solutions compared to Mg2+ ion-free solutions. This may indicate an involvement of Mg2+ ions in the surface reaction during calcium sulfate precipitation. 1. Introduction The crystallization of calcium sulfate, especially in its dihydrate (DH) form, from aqueous systems has been well documented, covering the effect of a large number of organic and inorganic additives on the crystallization kinetics and the product morphology.1,2 Carboxylic acids and phosphonic acid derivatives, such as citric acid and nitrilotrimethylenephosphonic acid, have been extensively investigated and show a significant inhibition effect on the nucleation of DH by extending the induction period.3-6 This has been believed to be relevant with the adsorption behaviors of the additives on particular faces of the crystal or chelation between the organic groups and the lattice ions.5,7 Inorganic substances, such as Na+, Mg2+, Al3+, Cu2+, and Cr3+ ions, exhibit a much more complicated effect: Besides the inhibition effect on the crystallization of DH, which was also explained by adsorption of active species,8 in some cases, accelerating effects were also reported.9,10 It has been suggested that both thermodynamic and kinetic factors could be responsible for the phenomena. The analysis of the important parameters such as solubility and interfacial tension between calcium sulfate crystals and electrolyte solutions is essential to the exploration of the underlying mechanisms.4,9 The precipitation of the metastable form, calcium sulfate hemihydrate, from water of certain degrees of salinities is often observed in nature or wet phosphoric acid processes and has found applications in several attempts to prepare a superior bonding material, R-calcium sulfate hemihydrate (R-HH).11-14 CaSO4-alkali/alkaline earth metal chlorides-H2O systems are the most frequently encountered aqueous systems in nature and industrial processes (e.g., in wet flue gas desulfurization processes which use calcium-based absorbents). As a part of systematic research which is aimed at seeking a cost-effective method to utilize large quantities of synthetic gypsum (i.e., flue gas desulfurization gypsum), we have previously investigated the metastability of R-HH and dehydration behavior of DH in alkali/alkaline earth metal chloride solutions, which could be * To whom correspondence should be addressed. Phone: +86 571 87952459. Fax: +86 571 87953088. E-mail: [email protected]. † Zhejiang University. ‡ East China Jiao Tong University.

used as a potential medium for the preparation of high quality R-HH from synthetic DH.11,15 Nonetheless, little work has been done for homogeneous nucleation kinetics of R-HH. Also, the complex role of inorganic ions such as Na+, K+, and Mg2+ commonly present in aqueous solutions in the crystallization of R-HH has not been well understood.16 The object of the present study is to quantify the effect of Mg2+ ions on the nucleation kinetics of calcium sulfate (probably in R-HH form) in CaCl2 solutions by measuring the induction period and thus to estimate the values of the interfacial tension between the crystals and the supersaturated solution. The information revealed may help to shed light on the mechanisms of R-HH crystallization and provide an instructive control strategy for R-HH preparation. 2. Chemicals and Methods 2.1. Materials. Reagent grade CaCl2 (Sinopharm Chemical Reagent Co. Ltd., China), MgCl2 · 6H2O (Taicang Meida Reagent Co., Ltd., China), and Na2SO4 (Gaojing Chemical Co., Ltd., China) were used as received. 2.2. Procedures. The experiments were carried out in a 1.5 dm3 double-walled glass vessel equipped with a condenser, a thermometer, and a Teflon stirrer. The temperature of the reactor was kept constant by the oil circulation between the double walls and a thermostatted oil bath. In each run, 600 cm3 of CaCl2 stock solution containing certain amounts of MgCl2 was added into the reactor and preheated to a fixed temperature (i.e., 90 °C) at a constant stirring rate of 200 rounds/min. The nucleation experiments were initiated by quickly adding 6-15 cm3 of 0.35 M Na2SO4 solution to the preheated solution. All solutions had been prepared with deionized water and filtered through a 0.45 µm membrane filter before being used. During the experiments, 10 cm3 of the suspension was withdrawn at certain time intervals with a pipet and transferred into a transparent glass pipe, which was subsequently moved into a photoelectric turbidity meter. The turbidity data could be read directly. All of these procedures could be performed within 15 s. With the fitted curves of turbidity as a function of time, the induction period was determined according to a similar method reported by Lancia et al.,17 which is based on the

10.1021/ie902022b  2010 American Chemical Society Published on Web 05/19/2010

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Ind. Eng. Chem. Res., Vol. 49, No. 12, 2010 Table 1. The Solubility (C*) of r-HH in 3.74 M CaCl2 Solutions with Different Mg2+ Ion Concentrations at 90 °C

Figure 1. DSC characteristics of the calcium sulfate crystals precipitated from a 3.74 M CaCl2 solution at a supersaturation ratio of 3.1. (Solid line) Mg2+ ion-free, (dashed line) in the presence of 0.20 M Mg2+ ions.

intersection point of two straight lines. One was drawn on the turbidity curve corresponding to a nearly linear increase of turbidity with time due to precipitation, and the other is the baseline or x axis. At the end of each run, the residual suspensions were filtered immediately, and the solid phases were washed with boiling water, rinsed with acetone, and dried for 1 h in an oven at 60 °C. Thermogravimetry and differential scanning calorimetry (TGDSC, STA 409PC NETZSCH Germany) were employed for the determination of the crystal water content and calcium sulfate phase identification. The sample (15 mg) was sealed in an Al2O3 crucible and scanned at a rate of 10 °C/min under a N2 gas atmosphere. The crystal morphologies were examined by field emission scanning electronic microscopy (FESEM, SIRION 100, FEI, Holand). 2.3. Supersaturation Ratio. The main composition of the background electrolyte was 3.74 M CaCl2, which could guarantee the metastability of the R-HH phase in all of the experiments. The thermal analysis confirmed that the precipitation products are calcium sulfate hemihydrate with crystal water contents close to a theoretical value of 6.21. As can been noted from the DSC curves for the crystals from Mg2+ion-free solution and solution containing 0.20 M Mg2+ ions, the distinctive exothermal peaks at about 185 °C indicate that the crystals are R-HH (see Figure 1). It has been proposed that the growth affinity, β, is thermodynamically more appropriate than the relative supersaturation, σ, in the definition of the driving force.18 However, when Ca2+ is in very large excess and SO42- concentration is low, the growth affinity is close to the relative supersaturation in value because β ) ln

0.5 [Ca2+][SO24 ][H2O] 0.5 [Ca2+]eq[SO24 ]eq[H2O]eq

≈ ln

[SO24 ] [SO24 ]eq

≈ ln

C ≈ C*

C - 1 ) σ(1) C* where [Ca2+], [SO42-], and [H2O] are the activities of corresponding species. The subscript eq refers to equilibrium conditions. Therefore, the discussion will be based on the relative supersaturation. The supersaturation ratio S was defined as C C ) S) C* C*

(2)

Mg2+ ion concentration (M)

C* (mmol/dm3)

0 0.001 0.01 0.10 0.15 0.20

2.11 2.11 2.07 2.04 1.98 1.95

where C and C* are the total sulfate concentration in solution at the initial time and at equilibrium, respectively. C* has previously been determined by a seeded growth equilibration method in the same aqueous system.16 Changes of solution composition could result in variation of the equilibrating concentration of sulfate ions. Thus, we found that the solubility of R-HH determined from the sulfate ion concentration at equilibrium decreased with the introduction of less than 0.20 M Mg2+ ions. The solubilities of R-HH in the 3.74 M CaCl2 solution for different concentrations of Mg2+ ions are given in Table 1. 2.4. Estimation of Interfacial Tension and Nucleation Rate. According to classic nucleation theory, the induction period of nucleation, defined as the time interval between the formation of a supersaturated solution and the onset of the changes of the physical properties due to the initiation of the precipitation or nucleation, can be correlated with the supersaturation ratio by the following equation:9,19 log tind ) A +

B T3(log S)2

(3)

where A is a constant, T is the absolute temperature (K), and B can be determined by B)

βγ3V2mNAf(θ)

(4)

(2.3R)3

where β is a shape factor and has a value of 16π/3 for a spherical particle, γ is the interfacial tension (J/m2), Vm is the molar volume (cm3 mol-1), NA is Avogadro’s number (mol-1), f(θ) is a correction factor which takes into account the nucleation modes, and R is the gas constant (J/mol K). In particular, when purely homogeneous nucleation takes place, f(θ) ) 1, and when heterogeneous nucleation occurs, f(θ) ) 0.01. Plotting of log tind against 1/(log S)2 over a range of supersaturation ratios for a fixed temperature gives a straight line with slope B. It has been proposed that a change in the slope of B may indicate a transition of nucleation mechanism from homogeneous to heterogeneous nucleation.4,19 The nucleation rate, i.e., the number of nuclei formed per unit time per unit volume, can be determined from the following equation: J ) F exp

[

-βγ3V2mNAf(θ) (RT)3 ln2 S

]

(5)

where J is the nucleation rate and F is a frequency constant and is known as the pre-exponential factor having a theoretical value of 1030 nuclei/cm3 · s. 3. Results and Discussions 3.1. Calcium Sulfate Nucleation in CaCl2 Solution in the Absence of Mg2+ Ions. The changes of the turbidity of the CaCl2 solution saturated with CaSO4 were monitored and plotted

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Figure 2. Determination of the induction period for calcium sulfate nucleation in CaCl2 solution at 90 °C: (O) S ) 3.6, (∆) S ) 3.3, (0) S ) 2.8, (]) S ) 2.2.

Figure 3. Induction period for calcium sulfate nucleation as a function of the supersaturation ratio in the absence of Mg2+ ions at 90 °C.

against time to find the induction period for the precipitation of calcium sulfate. Figure 2 shows the data recorded at supersaturation ratios from 2.2 to 3.6 in the absence of Mg2+ ions. The induction period was obtained from the time when the turbidity underwent a sudden increase from the baseline. The supersaturation ratio shows a significant effect on the induction period. At a supersaturation ratio of 2.8, the induction period was 2.0 min, while at a higher supersaturation ratio of 3.6, a shorter induction period of 1.3 min was obtained. The curves also indicate that the precipitation approaches an equilibrium or end within 9.0 to 12.0 min depending on the supersaturation ratio under the investigated conditions. Figure 3 reports the linearization of experimentally determined induction period data at a supersaturation ratio ranging from 1.9 to 3.6 for calcium sulfate nucleation in 3.74 M CaCl2 solution at 90 °C. From the slope of the straight line, the value of interfacial tension has been calculated to be 9.9 mJ/m2, assuming that homogeneous nucleation is dominant and spherical nuclei occur. 3.2. Effect of Mg2+ Ions on the Induction Period and the Interfacial Tension. The effect of Mg2+ ions from 0.001 to 0.20 M on calcium sulfate precipitation was investigated at different supersaturation ratios. Figure 4 shows the turbidity as a function of time at a supersaturation ratio of 2.0 for different concentrations of Mg2+ ions. It can be observed that the introduction of Mg2+ ions leads to the inhibition of calcium sulfate precipitation as compared to the case in the absence of Mg2+ ions. This effect is intensified by the increasing concen-

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Figure 4. Variations of turbidities with time in the absence and presence of Mg2+ ions during calcium sulfate precipitation at a supersaturation ratio of 2.0: (0) 0 M Mg2+ ions, (O) 0.001 M Mg2+ ions, (∆) 0.01 M Mg2+ ions, (]) 0.10 M Mg2+ ions, (b) 0.20 M Mg2+ ions.

tration of Mg2+ ions, with 0.01 M Mg2+ ions showing the strongest inhibition. The corresponding induction period increases from 4.0 min for the Mg2+ ion-free system to 6.0 min for 0.01 M Mg2+ ions. However, a further increase in Mg2+ ion concentration from 0.10 to 0.20 M demonstrates a lesser retarding effect. The linearization of experimental data for calcium sulfate nucleation in the presence of Mg2+ ions at 90 °C was used to estimate the interfacial tension values of corresponding aqueous systems. The results for the precipitation in the presence of 0.01 and 0.10 M Mg2+ ions are given in Figure 5, which shows the fitted lines with changes in the slopes, indicating a transition from a homogeneous to a heterogeneous nucleation mechanism. Therefore, only the slopes of the steeper lines were withdrawn to calculate the γ values according to eq 4. All the γ values obtained are listed in Table 2. It is evident that the interfacial tension increases with the concentration of Mg2+ ions present in the solution. At a concentration of 0.001 M, Mg2+ ions improve the interfacial tension from 9.9 mJ/m2 (for the baseline) to 11.3 mJ/m2, and a further increase of the Mg2+ ion concentration to 0.15 M also improves the γ value but shows a relatively slighter effect. When the Mg2+ ion concentration was added up to 0.20 M, a γ value of 14.0 mJ/m2 was achieved, corresponding to a significant increase of about 41% with respect to the value for the Mg2+ ion-free system. The overall rising trend of the γ values is consistent with the decreasing solubility of R-HH in the presence of Mg2+ ions (see Tables 1 and 2). We have not found any values of γ for R-HH in aqueous solutions reported in the literature. This is perhaps due to the fact that R-HH is generally regarded as a metastable phase. However, much work has been carried out on the precipitation of DH, the usually stable phase of calcium sulfate under ambient conditions. The γ value for DH precipitation from pure water saturated with calcium sulfate at around 90 °C has been reported to be 20.9 mJ/m2.20 This value was increased up to 63 mJ/m2 when 3.0 m NaCl was introduced in the aqueous system as the background electrolyte.9 Since the larger the interfacial tension values are, the higher the free energy that needs to be overcome to form a critical nucleus in a nucleation process is, the significantly larger γ values for DH precipitation with respect to that for R-HH in strong electrolyte solutions justifies the fact that the precipitation of R-HH is more favored than DH in the CaCl2 solution under the investigated conditions (see also Figure 1).

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Figure 5. Induction period as a function of the supersaturation ratio in the presence of Mg2+ ions at 90 °C: (a) 0.01 M Mg2+ ions, (b) 0.10 M Mg2+ ions. Table 2. Calculated Interfacial Tension Values (γ) for Calcium Sulfate in 3.74 M CaCl2 Solutions with Different Mg2+ Ion Concentrations at 90 °C Mg2+ ion concentration (M)

γ (mJ/m2)

0 0.001 0.01 0.10 0.15 0.20

9.9 11.3 11.5 11.9 12.6 14.0

3.3. Effect of Mg2+ Ions on Nucleation Rate of Calcium Sulfate. It can be noted that both interfacial tension and the supersaturation ratio can influence the nucleation kinetics according to eq 5, which predicts the nucleation rate at a high supersaturation ratio. The solubility of R-HH in CaCl2 solution can be altered by the addition of Mg2+ ions, leading to a variation of supersaturation with respect to the Mg2+ion-free system (see Table 1). Therefore, both interfacial tension and the supersaturation ratio should be taken into consideration to quantify the effect of Mg2+ ions on the nucleation of calcium sulfate in the aqueous system. Basically, the variation of interfacial tension shows a stronger effect on the nucleation rate than the changes in supersaturation do, as can be concluded from their orders in eq 5. For a clear demonstration, the calculated nucleation rates at supersaturation ratios ranging from 1.6 to 5.8 for different Mg2+ ion concentrations and hence γ values are illustrated in Figure 6. Generally, higher supersaturation and lower Mg2+ ion concentrations lead to a promoted nucleation rate. Suppose the initial sulfate ion concentration was fixed at a level to provide a supersaturation ratio of 3.0 for Mg2+ ion-free solution, the introduction of 0.20 M Mg2+ ions would lead to an increase of the supersaturation ratio to 3.25 considering the solubility change. The increase of supersaturation would theoretically allow for a higher nucleation and growth rate. Nonetheless, this effect is obviously covered by the significantly increased interfacial tension in the presence of Mg2+ ions, which results in a dramatic decrease in the nucleation rate from 4.4 × 1029 down to 1.3 × 1029 nuclei/cm3 · s, as indicated by the circles in Figure 6. It has been proved that Mg2+ ions may be adsorbed onto the steps’ edges and kinks and incorporated into the crystal lattice, which leads to enhanced CaCO3 solubility and a lower crystal growth rate.21,22 Similarly, adsorption and incorporation of metal ions, such as Mg2+, Cr3+, Fe3+, Cu2+, and Cd2+, by CaSO4 crystals, has also been suggested by some investigators.8,23 Moreover, it has been reported that the coprecipitation of Sr, Mg, Na, and K with R-HH is more favored than with DH and

Figure 6. Nucleation rates of calcium sulfate in CaCl2 solution with the supersaturation ratio as a function of Mg2+ ion concentrations: (9) 0 M Mg2+ ions, (b) 0.001 M Mg2+ ions, (2) 0.10 M Mg2+ ions, (1) 0.15 M Mg2+ ions, (() 0.20 M Mg2+ ions.

calcium sulfate anhydrite (AH) in seawater.24 So, it is very likely that the adsorption or incorporation of Mg2+ ions and subsequently poisoning of the active sites of the calcium sulfate nucleus may give a reasonable explaination for the inhibited nucleation kinetics in the present study. However, it was noted that Mg2+ ions do not always show an inhibiting effect on the precipitation of calcium sulfate. On the contrary, we found that a high concentration of Mg2+ ions exhibited a promoting effect, as reflected by the reduced induction period according to the turbidity data for the precipitation in the presence of 0.20 M Mg2+ ions (see Figure 7). Since it is difficult to discriminate the nucleation and growth process in spontaneous precipitation experiments, the variation of turbidity with time as a function of the concentration of Mg2+ ions in fact reflects the overall kinetics of the nucleation and the growth. Therefore, it seems likely that Mg2+ ions show a strong retarding effect on the nucleation and growth of calcium sulfate at low concentrations but a promoting effect on the crystallization at higher concentrations. This is consistent with what we have found in a recent study on R-HH seeded growth kinetics in the aqueous system with the same compositions.16 Since the increased interfacial tension in the presence of Mg2+ ions exerts a strong negative effect on the nucleation and growth of calcium sulfate in the solution, there should be some other factors causing the elevated calcium sulfate precipitation rate. It is known that Ca2+ ion desolvation could be the rate limiting step for its adsorption and incorporation into sulfates during

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Figure 7. Variations of turbidity with time in the absence and presence of Mg2+ ions during calcium sulfate precipitation at a supersaturation ratio of 3.1: (0) 0 M Mg2+ ions, (O) 0.001 M Mg2+ ions, (∆) 0.01 M Mg2+ ions, (]) 0.10 M Mg2+ ions, (b) 0.20 M Mg2+ ions.

the crystallization.25 However, due to the greater hydration enthalpy of Mg2+ ions (-1921 kJ mol-1) than that of Ca2+ ions (-1592 kJ mol-1), Mg2+ ions could compete for hydration water between cations and the water adsorbed onto the calciumbased mineral surface.26,27 By molecular dynamics (MD) calculation, Kerisit and Parker have shown that the residence time of a water molecule in the Mg2+ ions primary hydration shell (≈2.8 × 10-6 s) is several orders of magnitude higher than that of Ca2+ ions (≈38 × 10-12 s).28 Such a strong Mg2+ ion-water interaction indicates that Mg2+ ions can disrupt the surface hydration layer and can lead to surface destabilization and, consequently, favor the attachment of lattice ions in

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crystallization. Moreover, the introduction of a relatively high concentration of Mg2+ ions results in a lower water activity and decreased number of water molecules surrounding the lattice ions and hence reduced dimensions of the hydrated ions. The adsorption and incorporation of the hydrated Ca2+ ions on the surface of calcium sulfate crystallites, as well as their dehydration, may be facilitated in strong electrolyte solutions. On the basis of the above discussions, we propose that the increased local supersaturation and enhanced surface reaction process could be responsible for the increase of the precipitation rate of calcium sulfate in the presence of a high concentration of Mg2+ ions. 3.4. Morphologies and Thermal Analysis of the Precipitation Products. The morphologies of the solid phase from 3.74 M CaCl2 solution were investigated by SEM, and the results are shown in Figure 8. Columnar crystals of about 4-7 µm in diameter and 30-70 µm in length were obtained for Mg2+ ionfree and 0.01 M Mg2+ ion solutions. However, crystals precipitated in the presence of 0.20 M Mg2+ ions show much diversity in the diameter and longitude dimensions. Acicular crystals with a high aspect ratio of about 100 occur together with short columnar crystals (7-10 in aspect ratio). This may indicate a substantial increase but uneven distribution in the local supersaturation or an enhancement in the surface reaction rate under high ionic strength conditions. Moreover, the differences in crystal size and morphologies suggest that much more nucleus could form from the CaCl2 solution in the presence of 0.20 M Mg2+ ions. These analyses may support the experimental results where the turbidity increases faster at high supersaturation ratios for 0.20 M Mg2+ ions. It should be pointed out here that these crystals are obviously not the nucleus but

Figure 8. Morphologies of R-HH crystals precipitated from 3.74 M CaCl2 solution at a supersaturation ratio of 3.1: (a) Mg2+ ion-free, (b) in the presence of 0.01 M Mg2+ ions, (c) in the presence of 0.20 M Mg2+ ions.

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the precipitation product that can be detected by the turbidity meter after a certain degree of outgrowth. Therefore, what has been revealed from these morphologies should be viewed as the overall results of the effect of impurities on the nucleation and the growth process. However, the solid samples were withdrawn at the end of the reaction in each run so as to obtain a sufficient solid for phase determination by TG-DSC. Since the reaction time in the precipitation system ranges from 15 to 35 min (depending on the supersaturation), it is not clear whether or not R-HH is the only phase that occurrs, though we had determined the solid phase by TG-DSC as R-HH. In fact, one could easily propose a short-lived DH nucleation followed by R-HH transformation under given conditions. Due to the metastability of R-HH, which could trigger phase transitions during nucleation and the reaction time before sample withdrawal, further investigation using timeresolved online techniques should be carried out to confirm the nucleation of R-HH. 4. Conclusion From the experiments on spontaneous precipitation of calcium sulfate (probably in R-HH form) in concentrated CaCl2 solutions at 90 °C, the following conclusions could be drawn: The induction period for the spontaneous precipitation of calcium sulfate in 3.74 M CaCl2 solution has great dependence on the supersaturation ratios, and the interfacial tension between calcium sulfate and the CaCl2 solution at 90 °C is 9.9 mJ/m2. The presence of Mg2+ ions in the CaCl2 solutions increases the interfacial tension between the calcium sulfate crystal and the aqueous solution. This effect and possible adsorption of Mg2+ ions on the active sites of the nucleus may be responsible for the retarded nucleation kinetics at relatively lower supersaturation ratios. A high concentration of Mg2+ ions (i.e., 0.20 M) shows a promoting effect on the precipitation at supersaturation ratios higher than 3.0. This could be attributed to the promoted desolvation of Ca2+ ions and surface reaction during calcium sulfate formation in aqueous solution with stronger ionic strength. Mg2+ ions seem to facilitate acicular calcium sulfate crystals to precipitate from the CaCl2 solution, and this effect as well as the effect on the overall kinetics can be used to manipulate the preparation of R-HH with a desired morphology. 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 Science and Technology Department of Zhejiang Province, China, through the project of Science and Technology Plan of Zhejiang Province (2007C23055). The authors wish to thank the two anonymous reviewers for their insightful discussion and instructive suggestions. Literature Cited (1) McCartney, E. R.; Alexander, A. E. The effect of additives on the process of crystallization of calcium sulfate. I. Crystallization of calcium sulfate. J. Colloid Interface Sci. 1958, 13, 383. (2) Edinger, S. E. An investigation of the factors which affect the size and growth rates of the habit faces of gypsum. J. Cryst. Growth 1973, 18, 217. (3) Prisciandaro, M.; Lancia, A.; Musmarra, D. The retarding effect of citric acid on calcium sulfate nucleation kinetics. Ind. Eng. Chem. Res. 2003, 42, 6647.

(4) Prisciandaro, M.; Olivieri, E.; Lancia, A.; Musmarra, D. Gypsum precipitation from an aqueous solution in the presence of Nitrilotrimethylenephosphonic acid. Ind. Eng. Chem. Res. 2006, 45, 2070. (5) Boisvert, J. P.; Domenech, M.; Foissy, A.; Persello, J.; Mutin, J. C. Hydration of calcium sulfate hemihydrate (CaSO4 · 1/2H2O) into gypsum (CaSO4 · 2H2O). The influence of the sodium poly(acrylate)/surface interaction and molecular weight. J. Cryst. Growth 2000, 220, 579. (6) Klepetsanis, P. G.; Koutsoukos, P. G. Kinetics of calcium sulphate formation in aqueous media: effect of organophosphorus compounds. J. Cryst. Growth 1998, 193, 156. (7) Weijnen, M. P. C.; Van Rosmalen, G. M. Adsorption of phosphonate on gypsum crystals. J. Cryst. Growth 1986, 79, 157. (8) Hamdona, S. K.; Al Hadad, U. A. Crystallization of calcium sulfate dihydrate in the presence of some metal ions. J. Cryst. Growth 2007, 299, 146. (9) He, S.; Oddo, J. E.; Tomson, M. B. The nucleation kinetics of calcium sulfate dihydrate in NaCl solutions up to 6 m and 90 °C. J. Colloid Interface Sci. 1994, 162, 297. (10) Rashad, M. M.; Mahmoud, M. H. H.; Ibrahim, I. A.; Abdel-Aal, E. A. Crystallization of calcium sulfate dihydrate under simulated condition of phosphoric acid production in the presence of aluminum and magnesium ions. J. Cryst. Growth 2004, 267, 372. (11) 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. (12) Marshall, W. L.; Slusher, R. Aqueous systems at high temperature. Solubility to 200 °C of calcium sulfate and its hydrates in sea water and saline water concentrates, and temperature-concentration limits. J. Chem. Eng. Data 1968, 13, 83. (13) Sullivan, J. M.; Kohler, J. J.; Grinstead, J. H. Solubility of R-calcium sulfate hemihydrate in 40, 50 and 55% P2O5 phosphoric acid solutions at 80,90,100, and 110 °C. J. Chem. Eng. Data 1988, 33, 367. (14) Zu¨rz, A.; Older, I.; Thiemann, F.; Bergho¨fer, K. Autoclave-free formation of R-hemihydrate gypsum. J. Am. Ceram. Soc. 1991, 74, 1117. (15) Guan, B.; Yang, L.; Wu, Z.; Shen, Z.; Ma, X.; Ye, Q. Preparation of R-calcium sulfate hemihydrate from FGD gypsum in K, Mg-containing concentrated CaCl2 solution under mild conditions. Fuel 2009, 88, 1286. (16) Yang, L.; Wu, Z.; Guan, B.; Fu, H.; Ye, Q. Growth rate of R-calcium sulfate hemihydrate in K-Ca-Mg-Cl-H2O systems at elevated temperature. J. Cryst. Growth 2009, 311, 4518. (17) Lancia, A.; Musmarra, D.; Prisciandaro, M. Measuring induction period for calcium sulfate dihydrate precipitation. AIChE J. 1999, 45, 390. (18) Witkamp, G. J.; Van der Eerden, J. P.; Van Rosmalen, G. M. Growth of gypsum 1. Kinetics. J. Cryst. Growth 1990, 102, 281. (19) So¨hnel, O.; Mullin, J. W. Interpretation of crystallization induction periods. J. Colloid Interface Sci. 1988, 123, 43. (20) Ahmed, S. B.; Tlili, M. M.; Amor, M. B. Influence of a polyacrylate antiscalant on gypsum nucleation and growth. Cryst. Res. Technol. 2008, 43, 935. (21) Davis, K. J.; Dove, P. M.; De Yoreo, J. J. The role of Mg2+ as an impurity in calcite growth. Science 2000, 290, 1134. (22) Chen, T.; Neville, A.; Yuan, M. Assessing the effect of Mg2+ on CaCO3 scale formation-bulk precipitation and surface deposition. J. Cryst. Growth 2005, 275, 1341. (23) Sluis, S.; Witkamp, G. J.; Van Rosmalen, G. M. Crystallization of calcium sulfate in concentrated phosphoric acid. J. Cryst. Growth 1986, 79, 620. (24) Kushnir, J. The coprecipitation of strontium, magnesium, sodium, potassium and chloride ions with gypsum. An experimental study. Geochim. Cosmochim. Acta 1980, 44, 1471. (25) Nielsen, A. E. J. Cryst. Growth 1984, 67, 289. (26) Phillips, C. S. G.; Williams, R. J. P. Inorganic Chemistry 1; Oxford University Press: New York, 1965. (27) Arvidson, R. S.; Collier, M.; Davis, K. J.; Vinson, M. D.; Amonette, J. E.; Luttge, A. Magnesium inhibition of calcite dissolution kinetics. Geochim. Cosmochim. Acta 2006, 70, 583. (28) Kerisit, S.; Parker, S. C. Free energy of adsorption of water and metal ions on the {101j4} calcite surface. J. Am. Chem. Soc. 2004, 126, 10152.

ReceiVed for reView December 21, 2009 ReVised manuscript receiVed March 27, 2010 Accepted April 26, 2010 IE902022B