Gypsum Precipitation Kinetics and Solubility in the NaClâMgCl2

Article pubs.acs.org/IECR

Gypsum Precipitation Kinetics and Solubility in the NaCl−MgCl2− CaSO4−H2O System Samia Ben Ahmed,† Mohamed Mouldi Tlili,†,* Mongi Amami,‡ and Mohamed Ben Amor† †

Laboratoire de Traitement des Eaux Naturelles, Centre de Recherches et Technologies des Eaux, BP 273, Soliman, 8020 Tunisia Institut Préparatoire aux Etudes d’Ingénieur, Campus Universitaire, Merazka, 8000 Nabeul, Tunisia

‡

ABSTRACT: Crystallization of calcium sulfate dihydrate (CaSO4·2H2O, gypsum) in sodium chloride solutions has been studied in the presence of magnesium ion. It was shown that Mg2+ has an inhibiting effect on the two stages of precipitation: nucleation and growth. Kinetic results were explained on the basis of the strong effect of magnesium ions on gypsum solubility. In the studied range of the ionic strength (IS), it was found that in the presence of NaCl the solubility increase stops beyond IS = 0.31 mol·L−1 whereas it continues to rise in the presence of MgCl2. X-ray diffraction analyses showed that magnesium ions act by substituting the calcium ones or by insertion in the defects of the gypsum crystal lattice. minimum at 0.125 mol·kg−1. It rises thereafter to reach a maximum for 1.2 mol MgSO4/kg before decreasing again until saturation with MgSO4. A similar Mg effect was also found by Azimi et al.43,44 who have calculated the DH solubility in CaSO4−MgSO4−H2O43 and CaSO4−Na2SO4−MgCl2−H2O44 systems. From this literature review, it can be concluded that the calcium sulfate solubility is closely related to the magnesium content. However, to our knowledge, no data were reported on the influence of Mg2+ on gypsum kinetic precipitation and microstructure as for the case of calcium carbonate main component of scales deposited in drinking and industrial water plants.45−50 Thus, it was shown that the presence of magnesium makes difficult the nucleation of CaCO3, affects its kinetic crystallization, and promotes the aragonite variety instead of calcite. It has been suggested that Mg2+ acts by adsorption onto CaCO3 clusters formed during the nucleation step and thereafter onto the crystal active growth sites. It was also shown that it can insert or substitute calcium in the CaCO3 lattice to form magnesian calcite or aragonite: Ca1−xMgxCO3.5,48−51 The present research aims to study the effect of magnesium ions added under MgCl2 to an NaCl-CaSO4 solution. This study will be treated according to three axes: kinetic by the measurement of the induction period tind and the maximum growth rate V, thermodynamic by calculating the solubility of gypsum in the presence of Mg2+, and microstructural by analyses of the X-ray diffraction (XRD) patterns of the obtained precipitates.

1. INTRODUCTION Crystallization of calcium sulfate dihydrate is of importance in view of its application in a number of industrial and environmental processes.1 However, it can induce important damage in brackish and saline water desalination plants. Its undesirable precipitation reduces the reverse osmosis membrane life, plugs pipes and reaction vessels, and shortens the thermal transfer of the heat exchangers.2 Therefore, numerous investigations on the crystallization of its dihydrate (DH) form, the most frequent phase, from aqueous systems has been well documented, covering the effect of a large number of organic and inorganic additives.3−11 It was shown that these additives have a significant inhibition effect on the nucleation of DH.12−18 They act by adsorption on a specific crystal face or by chelation between their functional groups and calcium ions in the lattice.15,19,20 Inorganic substances, such as Na, Mg, Al, Cu, Cd, and Cr ions, exhibit a much more complicated effect. Besides their inhibiting effect, also explained by adsorption on active crystal sites,21−28 it was reported that some of them have likewise, depending on the operational conditions, an accelerating effects.29,30 It has been suggested that both thermodynamic and kinetic factors could be responsible for the phenomena. The analysis of the solubility is, therefore, essential to the exploration of the underlying mechanisms.13,29 The solubility of gypsum increases with temperature until 50 °C before gradually decreasing at least to 100 °C.31−33 Moreover, it is strongly dependent to the solution ionic strength (IS)32−39 and to the inhibitory additives (they do not change the IS) kind and content. It was shown by Najibi et al.40 that the solubility of calcium sulfate increases sharply in a 5% NaCl solution before decreasing when [NaCl] reaches 10%. Recently, Tlili et al.41 have shown, by calculating the solubility product of gypsum K′s, that the calcium sulfate solubility increased remarkably for ionic strength up to 0.17 mol·L−1. Because of its abundance in natural waters, the effect of Mg2+ on DH crystallization and microstructure has been the focus of several studies.42−44 Thus, Wollmann and Voigt42 reported that the gypsum solubility is Mg content-dependent. It decreases with increasing magnesium, added under MgSO4, to reach a © 2014 American Chemical Society

2. EXPERIMENTAL SECTION All crystallization tests were carried out in a 0.350 dm3 thermostatic double-walled Pyrex vessel at 50 °C. To avoid the effect of evaporation, the system was fitted with a waterReceived: Revised: Accepted: Published: 9554

February 10, 2014 April 29, 2014 May 1, 2014 May 1, 2014 dx.doi.org/10.1021/ie5004224 | Ind. Eng. Chem. Res. 2014, 53, 9554−9560

Industrial & Engineering Chemistry Research

Article

and the normalized intensity of each peak were determined using “X’Pert HighScore Plus” software.

cooled condenser. The working solution was prepared by mixing, under agitation, equal volume (100 dm3) of equimolar CaCl2 (S1) and Na2SO4 (S2) solutions added by MgCl2 or NaCl previously kept at the work temperature. Then, two sets of solutions, with and without Mg2+, were prepared at fixed ionic strengths. In the first set, MgCl2 was added into S1 at varied concentrations (0−0.024 mol·L−1). In the second set, NaCl (0.086−0.174 mol·L−1) was added instead of MgCl2 to kept the ionic strength constant. During the precipitation test, samples of 1 mL were withdrawn to be analyzed for Ca2+ determination by EDTA complexometry titration. The experiment was stopped when the concentration of free calcium ion had no more significant change. All experiments were replicated three times, from which average values were calculated (3%). Figure 1 shows a typical example of the temporal evolution of Ca2+ concentration. This precipitation curve shows two stages:

3. RESULTS 3.1. Precipitates Identification. For all experiments recovered precipitates were identified as calcium sulfate under its gypsum variety (CaSO4·2H2O). For example, Figure 2 shows the XRD pattern of the precipitated obtained at 50 °C in the absence of magnesium.

Figure 2. X-ray diffraction patterns of deposit scale at 50 °C in the absence of any additives.

In the presence of Mg, a slight difference in the microstructure of gypsum was found; results will be detailed later. 3.2. Mg Effect on the Precipitation Kinetics. From Figure 3, which illustrates the evolution of tind and V vs MgCl2

Figure 1. Ca2+ concentration evolution as a function of time.

The first one, where the dissolved calcium remains constant, corresponds to the nucleation step. Therefore, the nucleation time tind is determined from the slope change. The second step, where free calcium ions gradually decrease, indicates the growth of the already formed nucleus. According to Walton,52 the slope of the linear part in this second precipitation step (after tind) can be considered as the maximum growth rate. Below, this rate will be denoted (V). 2.1. Solubility Determination. At the end of the precipitation test, the precipitation rate becomes infinitely slow and a solide/liquide equilibrium must be established (Figure 1). Therefore, the solubility of calcium sulfate can be determined by dosing the total calcium ions which remain in the solution (s = [Ca2+]total). Seeing that a fraction of calcium can be associated with sulfate as ion pair CaSO4053,54 (s = [Ca2+]free + [CaSO04]), the EDTA complexometry titration method was employed to reach the [Ca2+]total. Indeed, CaEDTA2− complex constant (K = 1012,4 at 25 °C55) is greater than the one of CaSO40 (K = 9,61 × 103 at 25 °C56). 2.2. Precipitate Identification. At the end of each experiment, precipitate was recovered by filtering the precipitating solution through a 0.45 μm membrane. Samples were dried at room temperature before X-ray diffraction analyses. XRD was carried out at room temperature with a Philips X’PERT PRO diffractometer in step scanning mode using Co Kα radiation. The XRD patterns were recorded in the scanning range 2θ = 10−55°. A small angular step of 2θ = 0.017° and fixed counting time of 4s were used. The d spacing

Figure 3. Effect of Mg2+ on calcium sulfate precipitation (induction time and maximum growth rate).

content, it can be concluded that the presence of this salt strongly affects precipitation kinetics by prolonging the nucleation time and decreasing the maximum growth rate V according to the following equations: t ind = 17.4 + 1468 × [MgCl 2]

(1)

ln(V ) = −9.21 − 79.82 × [MgCl 2]

(2)

The induction time increases from 15 to 40 min when the concentration of MgCl2 passes from 0 to 0.014 mol·L−1 and reaches 52 min for 0.024 mol·L−1. The maximum growth rate V 9555

dx.doi.org/10.1021/ie5004224 | Ind. Eng. Chem. Res. 2014, 53, 9554−9560


Article

decreases from 1.5 × 10−4 to 1.76 × 10−5 mol·L−1·min−1 in the absence and presence of 0.024 mol·L−1 of Mg2+. This effect on the precipitation kinetics can be attributed to both the inhibitory effect of Mg2+ and the ionic strength, which play an important role on reducing the ionic activities. To specify the magnesium ion effect, the same experiments were conducted with NaCl instead of MgCl2 to avoid the IS influence. Figures 4 and 5 show the tind and V evolution vs the ionic strength.

Figure 6. Effect of Mg2+ on gypsum solubility.

Analyzing the solubility curve vs Mg content (Figure 7), it can also be shown that the solubility increases with MgCl2

Figure 4. Effect of Mg2+ on nucleation time.

Figure 7. Gypsum solubility as a function of MgCl2 concentration.

concentration according two half straight lines having two different slopes. The first is large and it is equal to 171.67 g· mol−1. If the Mg2+ concentration exceeds 0.005 mol·L−1 (IS = 0.275 mol·L−1), the effect on solubility decreases and the slope is equal to 30.41 g·mol−1. In the absence of magnesium (Figure 8), the solubility increase, with NaCl concentration, is rather exponential until IS = 0.31 mol·L−1: s = 2.175 exp(5.388[NaCl]). From this equation the gypsum solubility in pure water ([NaCl] = 0 mol·L−1) was estimated to be 2.17 g·L−1 close to the one given by Patridge and White57 (2.01 g·L−1 at 50 °C). The observed drop of the solubility values in the

Figure 5. Effect of Mg2+ on maximum growth rate.

The solution behavior regarding the induction time is very different in the presence of Mg ions especially when I exceeds 0.31 mol·L−1. Without magnesium, the induction time tind decreases after reaching a maximum of 32 min. However, it increases continually in the presence of Mg2+. Figure 4 clearly shows that regardless the IS value the presence of Mg2+ prolongs the nucleation time. For example, for IS = 0.31 mol·L−1, tind passes from 32 to 40 min in absence and presence of magnesium. As for tind, the presence of Mg reduces more the gypsum growth rate whatever the solution IS; especially for IS > 0.31 mol·L−1 (Figure 5). 3.3. Mg Effect on Calcium Sulfate Solubility. Figure 6 shows that, for the same IS, calcium sulfate is more soluble in the presence of magnesium. As for the kinetics parameters, the effect of Mg2+ on calcium sulfate solubility is more pronounced when IS exceeds 0.31 mol·L−1.

Figure 8. Gypsum solubility as a function of NaCl concentration. 9556



Article

Figure 9. (020) XRD peak profiles of the gypsum obtained for different concentrations of NaCl.

presence of NaCl beyond IS = 0.31 mol·L−1 was not mentioned in the published works on the effect of ionic strength on gypsum solubility. Here, we must noted that the studied range of the ionic strength (0.25 < IS (mol L−1) < 0.34) is very small with respect to those studied in the literature where IS can reach 6 mol·kg−1.40,58

As the Mg2+ concentration increases, the ionic product (PI) decreases and consequently the supersaturation coefficient is lower. However, it was shown that the solubility increases with the supersaturation of the medium studied. This suggests that besides its effect on the ionic strength, the presence of magnesium may have other roles in the medium reaction. Indeed, according to Mullin,65 the effect of additives is generally interpreted by a solubility change, by adsorption or chemisorption on nuclei or heteronuclei, and by chemical reaction or complex formation in the solution. The adsorption phenomenon, which is linked to the increase of the crystal−solution interfacial tension, has been widely elaborated. Söhnel and Mullin,66 Pokrovsky,46 and Tai and Chien47 found that the increase of the induction time vs [Mg2+]/[Ca2+] molar ratio is due to an increase of the crystal− solution interfacial tension. Recently, Guan et al.67 have shown for α-hemihydrate calcium sulfate that the presence of Mg2+ affects its nucleation kinetics by increasing the interfacial tension between the calcium sulfate crystal and the aqueous solution. This was interpreted by a possible adsorption of Mg2+ on the active sites of the formed nucleus. Other studies on calcium carbonate precipitation showed that the adsorption can be followed by a substitution or insertion of Mg to form a solid solution Ca1−xMgxCO3.5,45−51 Tlili et al.5 showed, by Raman spectrometry, that the substitution rate of Ca by Mg can reach 10% in calcite variety to form magnesian calcite. This same tendency to substitution in calcite was already highlighted in the works of Katz48 and Mucci.49,50 Other workers have mentioned that substitution is also possible in the aragonite lattice.51 If a similar phenomenon occurs in the case of gypsum, it permits us to explain the role of magnesium on the crystallization kinetics and on the solubility of gypsum. XRD analyses of recovered gypsum precipitated in the NaCl−CaSO4−H2O medium show similar patterns. Figure 9 shows the (020) peak profiles, located at 2θ = 11.82°, for each ionic strength. The intensity of this peak was normalized to 100%. Nevertheless, in the presence of magnesium a shift was observed for all peaks. Figure 10 shows the normalized (020)

4. DISCUSSION For the crystallographic point of view, results show that calcium sulfate precipitates under its dihydrate variety whatever the reaction medium NaCl−CaSO4−H2O or NaCl−MgCl2− CaSO4−H2O. Nevertheless, the precipitation kinetics was clearly affected. The increasing of the IS prolongs the nucleation time and decelerates its growth rate until IS = 0.31 mol·L−1. Beyond this IS value, the precipitation rate (V) decreases in the absence of Mg ions and increases in its presence (Figure 5). The effect of the ionic strength on precipitation kinetics can be attributed to the solubility variation. Indeed, Figure 6 shows that, in the absence of Mg2+, the solubility increases until IS = 0.31 mol·L−1 and then it drops whereas it increases continuously with Mg 2+ concentration. The increase can be attributed to the presence of foreign salts to the studied system.59,60 Indeed, Roques60 has shown that the companions of crystallization interfere by increasing the ionic strength, which results in an increase of the solubility. In the absence of magnesium, the decrease of the nucleation time for IS > 0.31 mol·L−1 has been interpreted by Madgin and Swales61 at 25 °C, by Bock62 for gypsum and anhydrite at 25, 30, 40, and 50 °C, and by Power and Fabuss63 and Furby et al.64 as being the result of the solubility decrease. Nevertheless, for the same ionic strength (IS = 0.278 mol·L−1), the solubility value is more important in the presence of magnesium (s = 4.23 g·L−1) than in the presence of NaCl (s = 3.9 g·L−1). This can be explained by the presence of the MgSO04 ion pair in addition of CaSO04. The calcium and sulfate concentration are, in this case, as follows: [Ca 2 +]free = [Ca 2 +]total free − [CaSO04 ] [SO4 2 −]free = [SO4 2 −]total free − [CaSO04 ] − [MgSO04 ] 9557



Article

Figure 10. (020) XRD peak profiles of gypsum obtained for (a) [Mg2+] < 0.005 mol·L−1 and (b) [Mg2+] > 0.005 mol·L−1.

peak profiles. Until 0.005 mol·L−1 of added Mg2+, the peaks shift toward higher 2θ values (Figure 10a, Table 1). Beyond

tensile stresses, the origin of the increase of the unit cell volume, and the peak shift to lower 2θ angles.

Table 1. (020) Peak Shift in Function of MgCl2 Concentration MgCl2 (mol L−1) 2θ (deg) Δ2θ

0 11.828 0

0.0024 11.912 0.084

0.005 11.971 0.143

0.01 11.690 −0.138

5. CONCLUSION The study of the magnesium effect on gypsum crystallization is elaborated in this paper. It was shown that Mg2+ has an inhibiting effect on the two stages of precipitation: nucleation and growth. The effect on gypsum solubility was also evaluated. It was found that the influence of this ion depends on the concentration initially introduced. For Mg2+ concentrations lower or equal to 0.005 mol·L−1, the effect on gypsum solubility is significant. Microstructural analyses showed that XRD peaks shifted to higher 2θ angles and the unit cell volume decreases. In this gap of magnesium concentration, it was therefore suggested that magnesium acts by substituting the Ca2+ ions in the gypsum lattice. Beyond 0.005 mol·L−1 of Mg2+, the solubility continues to increase but with a lower slope. In this case, the magnesium ion acts by insertion in the defects of the gypsum crystal lattice.

0.015 11.679 −0.149

0.005 M, a net shift of the peaks in the reverse direction was found (Figure 10b, Table 1). This displacement can be attributed to the substitution of Ca by Mg and/or an insertion and/or adsorption of Mg in the crystal lattice as expected for the case of calcite5,45−51 and aragonite.51,68 The lattice parameters calculated using the software fulproof69 are collected in Table 2. It was shown that the unit cell volume (= ) decreases for [Mg] < 0.005 mol·L−1 and increases beyond this concentration. This allows us to conclude that • For [Mg2+] < 0.005 mol·L−1, the peak shift is due to the lattice distortions caused by the presence of magnesium. Indeed, magnesium can be a part of the gypsum lattice by substituting the Ca ions (rMg2+ = 1.598 Å is weaker than rCa2+ = 1.974 Å). This involves a slight compression of the crystal lattice, which results in a decrease of the unit cell volume and a shift of the XRD peaks to higher 2θ angles. • For [Mg2+] > 0.005 mol·L−1, the unit cell volume = increases. This can be explained by the fact that, in addition to the substitution of the Ca2+, Mg ions could be adsorbed in the gypsum crystal surfaces. This causes

■

AUTHOR INFORMATION

Corresponding Author

*M. M. Tlili: e-mail, [email protected]; phone, +216 79 32 57 50; fax, +216 79 41 28 02. Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

Table 2. Gypsum Cell Parameters Calculated in the Presence of Magnesium

Mg Mg Mg Mg Mg

-

gypsum gypsum gypsum gypsum gypsum

−1

(0 mol·L ) (0.0024 mol·L−1) (0.005 mol·L−1) (0.01 mol·L−1) (0.015 mol·L−1)

a (Å)

b (Å)

c (Å)

β (deg)

V (Å3)*

10.4929 10.4940 10.5518 10.4274 10.3935

15.2027 15.1921 15.1689 15.2434 15.2709

6.2784 6.2784 6.2161 6.2274 6.2343

99.029 101.616 100.616 94.155 92.905

989.1343 980.4521 977.9322 987.264 988.2398

9558



■

Article

(21) Hamza, S. M.; Rahman, A. A.; Nancollas, G. H. The Influence of Magnesium Ion on the Kinetics of Crystal Growth and Dissolution of Strontium Fluoride. J. Cryst. Growth 1985, 73, 245. (22) Rahman, A. A.; Hamza, S. M.; Nancollas, G. H. AIChE Symp. Ser. 1985, 253 (83), 36. (23) Hamza, S. M. The Influence of Some Metal Ions on the Kinetics of Dissolution of Barium Fluoride. J. Cryst. Growth 1990, 102, 303. (24) Hamza, S. M.; Hamdona, S. K. The Kinetics of Dissolution of Calcium Fluoride in the Presence of Some Metal Ions. J. Cryst. Growth 1992, 125, 591. (25) Hamdona, S. K.; Nessim, R. B.; Hamza, S. M. Spontaneous Precipitation of Calcium Sulphate Dihydrate in the Presence of Some Metal Ions. Desalination 1993, 94, 69. (26) Vreugd, C. H.; Witkamp, G. J.; van Rosmalen, G. M. Influence and Incorporation of Lanthanide and Chromium Ions. J. Cryst. Growth 1994, 144, 70. (27) Dalvi, A. I.; Mohammad, M. N.; Sulami, S. A.; Sahul, K.; Rasheed, R. Gulf Water Conference, Bahrain, 2000. (28) Hamdona, S. K.; Al Hadad, U. A. Crystallization of Calcium Sulphate Dihydrate in the Presence of Some Metal Ions. J. Cryst. Growth 2007, 299, 146. (29) He, S.; Oddo, J. E.; Tomson, M. B. The Nucleation Kinetics of Calcium Sulphate Dihydrate in NaCl Solutions up to 6 M and 90°C. J. Colloid Interface Sci. 1994, 162, 197. (30) Rashad, M. M.; Mahmoud, M. H. H.; Ibrahim, I. A.; Abdel-Aal, E. A. Crystallization of Calcium Sulphate Dihydrate Under Simulated Conditions of Phosphoric Acid Production in the Presence of Aluminum and Magnesium Ions. J. Cryst. Growth 2004, 267, 372. (31) Linke, W. F.; Seidell, A. Van Nostrand-Reinhold, New York, 1965. (32) Silcock, H. L. Pergamon Press: New York, 1979. (33) Dutrizac, J. E. Calcium Sulphate Solubilities in Simulated Zinc Processing Solutions. Hydrometallurgy 2002, 65, 109. (34) Templeton, C. C.; Rodgers, J. C. J. Chem. Eng. Data 1967, 12 (4), 536. (35) Marshall, W. L.; Slusher, R. J. Phys. Chem. 1966, 70 (12), 4015. (36) Marshall, W. L.; Slusher, R.; Jones, E. V. J. Chem. Eng. Data 1964, 9 (2), 187. (37) Ostroff, A. G.; Metler, A. V. J. Chem. Eng. Data 1966, 11 (3), 346. (38) Bock, E. Can. J. Chem. 1961, 39, 1746. (39) Linke, W. F.; Seidell, A. Solubilities of Inorganic and Metal− Organic Compounds; American Chemical Society: Washington, DC, 1965. (40) Najibi, S. H.; Miiller-Steinhagen, H.; Jamialahmadi, M. Calcium Sulphate Scale Formation During Subcooled Flow Boiling. Chem. Eng. Sci. 1997, 52, 1265. (41) Tlili, M. M.; Korchef, A.; Ben Amor, M. Effect of Scalant and Antiscalant Concentrations on Fouling in a Solar Desalination Unit. Chem. Eng. Process 2006, in press. (42) Wollmann, G.; Voigt, W. Solubility of Gypsum in MSO4 Solutions (M = Mg, Mn, Co, Ni, Cu, Zn) at 298.15 and 313.15 K. J. Chem. Eng. Data 2008, 53 (6), 1375. (43) Azimi, G.; Papangelakis, V. G.; Dutrizac, J. E. Development of an MSE-Based Chemical Model for the Solubility of Calcium Sulphate in Mixed Chloride−Sulphate Solutions. Fluid Phase Equilib. 2008, 266, 172. (44) Azimi, G.; Papangelakis, V. G. Thermodynamic Modeling and Experimental Measurement of Calcium Sulphate in Complex Aqueous Solutions. Fluid Phase Equilib. 2009, in press. (45) Sohnel, O.; Mullin, J. W. Precipitation of Calcium Carbonate. J. Cryst. Growth 1982, 60, 239. (46) Pokrovsky, O. S. Precipitation of Calcium and Magnesium Carbonates from Homogeneous Supersaturated Solutions. J. Cryst. Growth. 1998, 186, 233. (47) Tai, C. Y.; Chien, W. C. Interpreting the Effects of Operating Variables on the Induction Period of CaCl2-Na2CO3 System by a Cluster Coagulation Model. Chem. Eng. Sci. 2003, 58, 3233.

ABBREVIATIONS AND SYMBOLS DH, dihydrate; IS, ionic strength; XRD, X-ray diffraction; tind, induction time; V, maximum growth rate; s, solubility; PI, ionic product; = , unit cell volume

■

REFERENCES

(1) Oster, J. D.; Frenkel, H. The Chemistry of the Reclamation of Sodic Soils with Gypsum and Lime. Soil. Sci. Soc. Am. J. 1980, 44, 4. (2) Linnikov, O. D. Kinetics of Calcium Sulphate Crystal Growth at its Crystallization on a Heat-Exchange Surface. Desalination 2000, 128, 35. (3) McCartney, E. R.; Alexander, A. E. The Effect of Additives on the Process of Crystallization of Calcium Sulphate. J. Colloid Interface Sci. 1958, 13, 383. (4) 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. (5) Tlili, M. M.; Ben Amor, M.; Gabrielli, C.; Joiret, S.; Maurin, G.; Rousseau, P. Study of Electrochemical Deposition of CaCO3 by In SituRaman Spectroscopy II. Influence of the Solution Composition. J. Electrochem. Soc. 2003, 150 (7), 485. (6) Le Gouellec, Y. A.; Elimelech, M. Calcium Sulphate (Gypsum) Scaling in Nanofiltration of Agricultural Drainage Water. J. Membr. Sci. 2002, 1, 5310. (7) Bramson, D.; Hasson, D.; Semiat, R. The Role of Gas Bubbling, Wall Crystallization and Particulate Deposition in CaSO4 Scale Formation. Desalination 1995, 105. (8) Chong, T. H.; Sheikholeslami, R. Thermodynamics and Kinetics for Mixed Calcium Carbonate and Calcium Sulphate Precipitation. Chem. Eng. Sci. 2001, 56, 5391. (9) Sudmalis, M.; Sheikholeslami, R. Precipitation and Coprecipitation of CaCO3 and CaSO4. Can. J. Chem. Eng. 2000, 78, 21. (10) Marshall, W. L.; Slusher, R. Thermodynamics of Calcium Sulphate Dihydrate in Aqueous Sodium Chloride Solution, 0−110°C. J. Phys. Chem. 1966, 70 (12), 4015. (11) Prisciandaro, M.; Lancia, A.; Musmarra, D. Calcium Sulphate Dihydrate Nucleation in the Presence of Calcium and Sodium Chloride Salts. Ind. Eng. Chem. Res. 2001, 40, 2335. (12) Prisciandaro, M.; Lancia, A.; Musmarra, D. The Retarding Effect of Citric Acid on Calcium Sulphate Nucleation Kinetics. Ind. Eng. Chem. Res. 2003, 42, 6647. (13) 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. (14) Boisvert, J. P.; Domenech, M.; Foissy, A.; Persello, J.; Mutin, J. C. Hydration of Calcium Sulphate 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. (15) Klepetsanis, P. G.; Koutsoukos, P. G. Kinetics of Calcium Sulphate Formation in Aqueous Media: Effect of OrganoPhosphorus Compounds. J. Cryst. Growth 1998, 193, 156. (16) Oner, M.; Dogan, O.; Oner, G. The Influence of Polyelectrolytes Architecture on Calcium Sulphate Dihydrate Growth Retardation. J. Cryst. Growth 1998, 186, 427. (17) Ben Ahmed, S.; Tlili, M. M.; Ben Amor, M.; Ben Bacha, H.; Elleuch, B. Calcium Sulphate Scale Prevention in a Desalination Unit Using the SMCEC Technique. Desalination 2004, 167, 311. (18) Ben Ahmed, S.; Tlili, M. M.; Ben Amor, M. Optimization of an Antiscale Treatment to Prevent Scaling of a Solar Desalination Unit Using Evaporation-Condensation Principle, Inter. Sci. J. Alter. Ener. Eco. 2008, 5 (61), 32. (19) Weijnen, M. P. C.; Van Rosmalen, G. M. Adsorption of Phosphonate on Gypsum Crystals. J. Cryst. Growth. 1986, 79, 157. (20) Ben Ahmed, S.; Tlili, M. M.; Ben Amor, M. Influence of a Polyacrylate Antiscalant on Gypsum Nucleation and Growth. Cryst. Res. Technol. 2008, 1. 9559



Article

(48) Katz, A. The Interaction of Magnesian with Calcite during Crystal Growth at 25−90°C and one atmosphere. Geochim. Cosmochim. Acta 1973, 37, 1563. (49) Mucci, A.; Morse, J. W. The Incorporation of Mg2+ and Sr2+ into Calcite Overgrowths. Influence of Growth Rate and Solution Composition. Geochim. Cosmochim. Acta 1983, 47, 217. (50) Mucci, A. Growth Kinetic and Composition of Magnesian Calcite Overgrowths Precipated from Sea Water: Quantitative Influence of Orthophosphate Ions. Geochim. Cosmochim. Acta 1986, 50, 2255. (51) Karoui, H.; Korchef, A.; Tlili, M. M.; Mosrati, H.; Gil, O.; Mosrati, R.; Amor, M. B. Effects of Mg2+, Ca2+ and SO42‑ Ions on Precipitation Kinetics and Microstrucutre of Aragonite. Ann. Chim. Sci. Mater. 2008, 33, 123. (52) Walton, A. G. The Formation and Properties of Precipitates; Wiley: New York, 1967. (53) Azimi, G.; Papangelakis, V. G.; Dutrizac, J. E. Development of an MSE-Based Chemical Model for the Solubility of Calcium Sulphate in Mixed Chloride−Sulphate Solutions. Fluid Phase Equilib. 2008, 266, 172. (54) Azimi, G.; Papangelakis, V. G. Thermodynamic Modeling and Experimental Measurement of Calcium Sulphate in Complex Aqueous Solutions. Fluid Phase Equilib. 2009, in press. (55) Westin, K. J.; Rasmuson, A. C. Crystal Growth of Aragonite and Calcite in Presence of Critic Acid, DTPA, EDTA and Pyromellitic Acid. J. Colloid Interface Sci. 2005, 282, 359. (56) Maurel, A.; Audier, Ph.; Barnier, H.; Duffau, C. Etude de la Solubilité du Sulphate de Calcium en Vue de l′Application au Dessalement de l′Eau de Mer et des Eaux Saumâtres. SCA/GETN. 1982, in press. (57) Patridge, E. P.; White, A. H. The Solubility Of Calcium Sulfate From 0 to 200°C. J. Am. Chem. Soc. 1929, 51, 360. (58) Schausberger, P.; Mustafa, G. M.; Leslie, G.; Friedl, A. Scaling Prediction Based on Thermodynamic Equilirum Calculation- Scopes and Limitations. Desalination 2009, 244, 31. (59) Barba, D.; Brandani, V.; Giacomo, G. Solubility of Calcium Sulphate Dihydrate in the System Na2SO4-MgCl2-H2O. J. Chem. Eng. Data 1984, 29, 42. (60) Roques, H. Sur une Nouvelle Méthode Graphique d’Etude des Eaux Naturelles. Ann. Speleol. 1972, 27 (1), 79. (61) Madgin, W. M.; Swales, D. A. Solubilities in the System CaSO4NaCl-H2O at 25°C and 35°C. J. Appl. Chem. 1956, 6, 482. (62) Bock, E. On The Solubility Of Anhydrous Calcium Sulphate And Of Gypsum In Concentrated Solutions Of Sodium Chloride At 25 °C, 30 °C, 40 °C, and 50 °C. Can. J. Chem. 1961, 39 (9), 1746. (63) Power, W. H.; Fabuss, B. Interpretation of the Solubility of Calcium Sulphates in Water for Temperatures between 20°C and 80°C. J. Chem. Eng. Data 1964, 9, 437. (64) Furby, E.; Glueckauf, E.; McDonald, L. A. The Solubility Of Calcium Sulphate In Sodium Chloride And Sea Salt Solutions. Desalination 1968, 4 (2), 264. (65) Mullin, J. W. Crystallisation; Butterworth-Heinemann: Oxford, U.K., 1993. (66) Söhnel, O.; Mullin, J. W. Precipitation of Calcium Carbonate. J. Cryst. Growth 1982, 60, 239. (67) Guan, B.; Vang, L.; Wu, Z. Effect of Mg2+ Ions on the Nucleation Kinetics of Calcium Sulphate in Concentrated Calcium Chloride Solutions. Ind. Eng. Chem. Res. 2010, 49, 5569. (68) Dietzel, M.; Gussone, N.; Eisenhauer, A. Co-precipitation of Sr2+ and Ba2+ with Aragonite by Membrane Diffusion of CO2 Between 10° to 50°C. Chem. Geol. 2004, 203, 139. (69) Rodriguez-Carvajal, J. Collected Abstracts of Powder Diffraction Meeting; Toulouse (France), 1990.

9560


Gypsum Precipitation Kinetics and Solubility in the NaClâMgCl2

Recommend Documents

Gypsum Precipitation Kinetics and Solubility in the NaClâMgCl2