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Influence of Impurities on Crystallization Kinetics of Calcium Sulfate Dihydrate and Hemihydrate in Strong HCl-CaCl2 Solutions Thomas Feldmann* and George P. Demopoulos* McGill University, Montreal, QC, Canada, H3A 0C5 ABSTRACT: The effects of inorganic impurities on the crystallization of calcium sulfates in strong HCl (6.3 mol L−1)-CaCl2 (1.8 mol L−1) solutions were investigated. The impurities considered relate to hydrochloric acid leaching of apatite-type ores for the extraction of rare earth elements. The impurities investigated were K+, Mg2+, Sr2+, Ba2+, Al3+, Fe2+, Fe3+, La3+, Y3+, F− (fluoride), and PO43− (phosphate). The investigation was done in the context of a continuous steady-state crystallization process. Therefore, temperature-controlled, semibatch crystal growth experiments with regulated reagent addition, to ensure nearly constant supersaturation, were performed. The experiments were conducted at 40 and 80 °C corresponding, respectively, to crystallization of calcium sulfate dihydrate (DH) and calcium sulfate hemihydrate (HH). Among all impurities investigated, phosphate and strontium were found to have the most significant effects, with La3+ and Y3+ having some modest effects. Phosphate (added as phosphoric acid) was found to accelerate the growth kinetics of dihydrate up to a certain concentration level (0.3 mol L−1), subsequently causing a retardation effect over the concentration range from 0.3 to 1.0 mol L−1. In contrast, phosphate had no effect on the growth kinetics of hemihydrate. In the meantime, phosphate uptake increased with increasing impurity concentration in the range up to 0.2 mol L−1 and then plateaued at 0.02 molphosphate molsolid−1. X-ray photoelectron spectroscopy (XPS) analysis provided evidence of the presence of a surface calcium phosphate species. On the other hand, the uptake of Sr2+ by dihydrate was much more extensive than that of phosphate (≈5−10×). In this case, substitution rather than adsorption was the mechanism of uptake, reflecting the similar ionic radii between calcium and strontium. At phosphate and strontium concentrations >≈ 0.2 mol L−1, partial transformation of dihydrate to hemihydrate was induced. Finally, La3+ and Y3+ were found to be incorporated at trace level amounts into dihydrate crystals causing crystal morphology changes but not promoting phase transformation.

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solutions6 and associated phase transformation kinetics have been investigated in detail previously,7 no detailed study on crystallization kinetic effects of inorganic impurity species present in the CaCl2-HCl-CaSO4-H2O system has been reported. Previous investigations by other authors focused on the effects of impurities under conditions of phosphoric acid production8−11 or were done in low acid aqueous solutions.12−14 It is indeed the scope of this paper to investigate impurity effects during crystallization of dihydrate and hemihydrate in industrially relevant strong acidic CaCl2(1.8 mol L−1)-HCl(6.3 mol L−1) solutions. In particular in the present study, the crystallization kinetic tests are carried out in the presence of several industrially important impurities under nearly constant supersaturation reflecting continuous steadystate process conditions. It is well-known that impurities present in solution may transfer to the crystallizing solid. The transfer may take place via interstitial uptake between regular lattice positions, coprecipitation, or isomorphous substitution of one of the ions in the host lattice by the impurity ion or by adsorption.8 Furthermore, the adsorption of impurities on the crystal surface may influence the precipitate morphology and/or the crystal growth rate.15,16 As such, they can impact the end use

INTRODUCTION Many industrially relevant crystallization processes are influenced by the presence of impurity elements. The negative effects of impurities range from crystal product contamination and alteration of morphology to the formation of other than the intended phase. An example of such an industrially relevant crystallization process is encountered in the hydrochloric acid processing of raw materials as is the extraction of valuable metals from ores and mineral concentrates. An integral part of such a process is the regeneration of the lixiviant (leaching reagent), namely, high strength (typically 6 mol L−1) HCl. Calcium-containing raw materials, as is phosphate rock (apatite), can be leached effectively with HCl,1 and the use of lime (CaO) as neutralizing agent yields highly concentrated CaCl2 solutions from which HCl has to be regenerated.2 As has been shown previously,3,4 the lixiviant can be conveniently regenerated from CaCl2 solutions by reactive crystallization of calcium sulfates making controlled use of concentrated (2.3−8 mol L−1)3 sulfuric acid (see eq 1). CaCl 2(aq) + H 2SO4(aq) + x H 2O ⇌ 2HCl(aq) + CaSO4 ·x H 2O(sol)

(1)

The calcium sulfate phase formed depends on the amount of crystal water x. Depending on temperature, concentration, and residence time conditions, calcium sulfate dihydrate (DH) (x = 2), α-hemihydrate (HH) (x = 0.5), or anhydrite (AH) (x = 0) is formed.5 While the growth kinetics of HH in impurity free © 2013 American Chemical Society

Received: Revised: Accepted: Published: 6540

October 25, 2012 February 19, 2013 April 17, 2013 April 17, 2013 dx.doi.org/10.1021/ie302933v | Ind. Eng. Chem. Res. 2013, 52, 6540−6549

Industrial & Engineering Chemistry Research

Article

characteristics of the DH and HH sulfate products.17 In the case of cationic impurities, it can be helpful to group them by valency in order to explain their effects. Another influencing factor is the ionic radius. For example, the amount of impurity uptake into HH precipitated in a phosphoric acid environment correlated with the ionic radius of the impurity in comparison to the ionic radius of calcium ions.8 The larger ions showed a higher uptake with the valency also exhibiting a slight effect. Interestingly, the authors found also that for many trivalent ions the uptake in HH was ten times higher than in DH or AH expressed in terms of a partition coefficient, which suggests that the crystal lattice configuration plays an important role too.8 Wang and Yue18 and Yang et al.19 investigated the effect of + K and Na+ on the morphology and growth rate of calcium sulfate hemihydrate in strong CaCl2 salt solutions, finding that K+ ions increase the growth rate only at certain specific concentrations, for example 0.01 mol L−1 and a relative supersaturation level of 0.69. There were contradictory or no effects at all with other conditions. The presence of these ions led however to more elongated crystals upon phase transformation of DH to HH in comparison to impurity free solutions. Furthermore, an effect of valency was reported in the way that di- and trivalent cations led to more elongated crystals compared to univalent ions. The presence of Mg2+ ions in unseeded growth experiments is reported to lead to elongated crystals and reduced growth rate of calcium sulfate dihydrate.11,13,20 The retarding effect of other divalent cations (Cu2+, Fe2+, Cd2+) increased with their increasing concentration in solution.13 Among possible trivalent impurities, the presence of aluminum ions has been well investigated under conditions of phosphoric acid production. It was found that depending on the concentration of Al3+ the growth rate of dihydrate increased and that the morphology of the crystals is affected.11,21 For example, Hasson et al.22 found the presence of Al3+ (0.13 to 0.54 mol L−1) or Fe3+ (0.06 to 0.35 mol L−1) to change the crystal morphology from thin plates to thick dihydrate crystals. Another well investigated group of trivalent impurities are the rare earth elements. A study by Vreugd et al.14 found that rare earth elements in comparison to trivalent chromium ions had a particular growth retarding effect on DH crystallization in solutions with NaNO3 in the range of 0.1−1.0 mol L−1. At an impurity concentration of 3 × 10−4 mol L−1, the growth retardation effectiveness gave the following order La3+ = Ce3+ = Eu3+ > Er3+ > Cr3+ > Cr2O72−. Also, the uptake of lanthanide ions increased with an increase in ionic radius. The uptake was suggested to take place via adsorption. Another important impurity influencing the crystallization of DH in the phosphoric acid process is fluoride (F−); especially, the effect of the Al/F ratio was investigated in the past. The presence of fluoride ions leads to the formation of a number of different, negatively charged Al−F complexes. It was reported that with increasing concentrations of fluoride the growth rate of DH was reduced while the crystal morphology changed from plate-like to irregular, rough, spherical agglomerates.23,24 The overview presented above shows the importance of understanding the effect of impurities in industrial crystallization systems. In this context, the present work investigates the influence of a wide range of impurities that are encountered in a novel process that seeks to recover rare earth elements from phosphate ore via hydrochloric acid leaching. In this process, HCl is regenerated via reactive crystallization of CaSO4

phases in unusually highly concentrated HCl-CaCl2 media. The impurities investigated are K+, Mg2+, Sr2+, Ba2+, Al3+, Fe2+, Fe3+, La3+, Y3+, F− (fluoride), and PO43− (phosphate). Specifically, the crystal growth kinetics, impurity uptake, and phase transformation conditions were studied under a nearly constant supersaturation regime corresponding to steady-state continuous, industrial crystallization processes.

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EXPERIMENTAL SECTION The crystallization experiments were carried out in a closed, baffled, 2 L semibatch, flat-bottom glass reactor with temperature control and a 45° pitched blade impeller with three blades each at two different height levels 3 cm apart from each other. Reagents were supplied by a peristaltic pump feed system. The experiments involved the equimolar addition of 100 mL CaCl2 (3 mol L−1) and 200 mL Na2SO4 (1.5 mol L−1) solutions to 1 L of a background electrolyte containing calcium chloride (1.8 mol L−1) and hydrochloric acid (6.3 mol L−1). Each experiment was conducted within 1 h, and the reagent addition rate was 0.3 mol L−1 facilitating significant crystal growth. Sodium sulfate instead of sulfuric acid was used as a source of sulfate ions in order to ensure nearly constant supersaturation as it has been demonstrated and described in a previous publication.6 Use of H2SO4 instead of Na2SO4 would have resulted in gradually increasing HCl concentration, and since the latter strongly influences calcium sulfate solubility, supersaturation would have changed too.25 All solutions were prepared from deionized water; Fisher Scientific and Sigma Aldrich reagent grade chemicals were used. The experiments were performed at 40 °C in the case of DH growth and 80 °C in the case of HH growth; impurities were tested one at a time. The cationic impurities were added in the form of chloride salts, and the anionic impurities were added in the form of their respective acids. An overview of the tested impurity concentration levels is given in Table 1. These were Table 1. Initial Concentration of Impurity Ions in Experimental Solutions initial concentration [mol L−1]

impurity species 2+

3+

2+

3+

Mg , Al , Fe , Fe , K La3+, Y3+ Sr2+ Ba2+ H3PO4 HF

+

0.001, 0.01, 0.1 0.0001, 0.001, 0.01 0.001, 0.01, 0.1, 0.2 0.0001, 0.0002, 0.0003, 0.0004, 0.0005, 0.001 0.001, 0.01, 0.1, 0.2, 0.3, 0.4, 0.6, 0.8, 1.0 0.001, 0.01, 0.1

chosen to cover a wide range of conditions taking into account solubility limits; e.g., the presence of barium could only be investigated at very low concentrations as it would otherwise precipitate as barium sulfate. Similarly for fluoride, the upper limit of its concentration was set at 0.1 mol L−1 to formation of CaF2. The solubility limit for each impurity was either checked using the OLI Stream Analyzer26 or via quick a priori dissolution tests. The concentrations of the selected rare earth element ions were relatively low, because they represent the target elements of the overall flowsheet of which the HCl regeneration/calcium sulfate crystallization process is a unit operation. Unlike other impurities, they will enter the acid regeneration circuit at low concentration levels only. The presence of phosphate ions was investigated, as these may still be present in the solution that enters the regeneration circuit caused by leaching calcium phosphate-based minerals. 6541

dx.doi.org/10.1021/ie302933v | Ind. Eng. Chem. Res. 2013, 52, 6540−6549


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At the beginning of each experiment, 30 g of lab-made dihydrate or hemihydrate seed was added to the reactor. Slurry samples were taken at the beginning of each experiment and then every 15 min and filtered with 0.2 μm pore size syringe filters. The filter cake was washed with 30 mL of calcium sulfate saturated water followed by a wash with 10 mL of isopropanol. This was done to avoid complications from possible rehydration/transformation or crystal dissolution, at the same time ensuring efficient wash out of remaining impurity laden solution. XRD analysis before and after the washing confirmed that the selected washing procedure did not induce any phase changes as also previosly established in our laboratory.25 The samples were stored at room temperature immersed in isopropanol as a preservation measure until measurements were performed. In order to determine the growth rate, volume-based particle size distributions were measured with a Horiba Laser Scattering Particle Size Distribution Analyzer LA-920, by dispersing the solid in isopropanol, to avoid dissolution of small particles during the measurement. Brunauer-Emmett-Teller (BET) surface area measurements of the seed crystals were undertaken with a Micromeritics Tristar Surface Area Analyzer, whose results were used to calculate a surface area normalized growth rate. A detailed description of this approach is given elsewhere.6 The crystal morphology was analyzed with a Philips XL30 field emission gun (FEG) scanning electron microscope (SEM) after sputter coating the samples with a thin layer of AuPd. X-ray powder diffraction analysis was carried with a Bruker AXS powder diffractometer. Cu Kα radiation with a wavelength of 1.506 Å was employed. Differential scanning calorimetry (DSC) analysis was carried out on a TGA Instruments Q2000 apparatus, with a heating rate of 10 K min−1 and a nitrogen gas flow of 50 mL min−1 in closed aluminum crucibles. Chemical analysis was conducted with a Dionex ICS 5000 ion chromatograph and a Thermo Scientific iCAP 6000 ICP-MS apparatus. Finally, X-ray photoelectron spectroscopy (XPS) measurements were conducted on a Thermo Scientific KAlpha instrument, using an Al Kα X-ray source at 1486.6 eV. Spectra were generated at a perpendicular takeoff angle, using a pass energy of 20 eV and steps of 0.1 eV. During analysis, the pressure was in the order of ≈1.33 × 10−9 Pa. As an internal reference for the absolute binding energies, the Au (4f7/2) peak was used. The experimental spectra were deconvoluted after subtraction of the Shirley background using the VG Avantage program. Thermodynamic calculations were performed using the OLI Stream Analyzer.26 To ensure reproducibility, individual experiments were repeated and reproducibility was found to be good.

Figure 1. Evolution of relative supersaturation in a typical growth experiment at 80 °C, 0.3 mol L−1 reagent inflow, seeded with HH (average and standard deviation based on 3 repeats).

Σ=

aCa 2+aSO4 2−a H2O x (aCa 2+aSO4 2−a H2O x)eq.

−1 (2)

where a refers to the activity of the ionic species/molecule in solution, the “eq.” subscript refers to equilibrium values, and lastly x is determined by the phase that is present; it takes the value 2, 0.5, or 0 for the dihydrate, hemihydrate, or anhydrite form, respectively. The ion activities, ai = γici, were calculated using the activity coefficients estimated with the help of OLI Stream Analyzer version 9.0.26 This solution thermodynamics tool has been shown to provide good estimates of the properties of CaCl2HCl-CaSO4-H2O aqueous electrolyte systems over a wide range of concentrations in good agreement with experimental values.27,28 Activity coefficient estimation by OLI is made on the basis of the excess Gibbs energy (GE); the latter is calculated with the mixed-solvent electrolyte (MSE) model.29 The model accounts for all ion and molecular interactions, namely, long-range electrostatic interactions expressed by the Pitzer-Debye-Hückel equation; the short-range interactions are calculated with the UNIQUAC equation and, finally, the middle-range interactions. Furthermore, OLI relies on the Helgeson-Kirkham-Flowers equation to calculate the standardstate thermodynamic properties of all species involved in reactions.29 The crystal growth rates were determined in a direct way from particle size distribution measurements by a method discussed previously.6 The average surface area of the seed material was 0.555 ± 0.16 m2 g−1 for dihydrate and 0.554 ± 0.07 m2 g−1for hemihydrate. These values were used to normalize the calculated growth rates. Influence of Impurities on Crystal Growth Rates. A summary of the results of the influence of different initial concentrations of impurities on the crystal growth rates can be seen in Figures 2a,b and 3a for the growth of DH and in Figure 3b for HH. They show the measured values in comparison to the growth rate of crystals in impurity-free solution (given as average ± standard deviation by the dashed and fine dashed lines, respectively). Within the wide range of impurity type and concentration investigated, only the presence of phosphate ions showed a significant effect and trend concerning the crystallization of DH. With increasing phosphate concentration, the growth rate increased at first up to 0.2 mol L−1 from 0.29 to 0.39 nm s−1 m−2; upon further increase, the trend reverted and

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RESULTS AND DISCUSSION Crystal Growth under Nearly Constant Supersaturation. Industrial crystallization operations typically employ continuous reactors, which are characterized by a constant supersaturation level as a result of constant composition at steady state. Nearly constant supersaturation, approximating steady-state operation, was maintained in the present work by conducting the experiments in a semibatch reactor with constant reagent (Ca2+ and SO42−) addition rate, as per the procedure described in a previous publication.6 The nearly constant supersaturation obtained throughout the crystallization experiment can be verified with the typical supersaturation profile presented in Figure 1. The relative supersaturation of a calcium sulfate phase is given by eq 2:15 6542



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Figure 2. Influence of cationic impurities on the crystal growth rate of calcium sulfate dihydrate grown at 40 °C, 0.3 mol L−1 reagent inflow.

Figure 3. Influence of impurities on the crystal growth rate of (a) calcium sulfate dihydrate and (b) hemihydrate grown at 40 °C (DH) and 80 °C (HH) at 0.3 mol L−1 reagent inflow.

growth rates declined as can be seen in Figure 3a. However, it is interesting to note that the phosphate presence had no effect on the growth rate of HH. Impurity Uptake. In order to quantify the uptake tendency of impurities by calcium sulfate crystals, an uptake ratio given by eq 3 is used, where N is the number of moles. This represents the amount of impurity that was taken up by the crystals after filtration and washing.

U=

Nimpurity Nprecipitate

(3)

Again, the behavior of phosphate is of interest. As it can be seen in Figure 4, the amount of phosphate uptake in dihydrate crystals increased with increasing concentration up to 0.3 mol L−1. At higher concentrations, the phosphate amount in the solids plateaued. The uptake of other impurities is summarized in Figure 5a,b. Among the investigated cationic impurities, potassium, strontium, yttrium, lanthanum, and barium were taken up by DH crystals but in very minute amounts. The amount of impurity uptake by the solid followed a linear trend for strontium and the REE elements for increasing initial impurity concentration. However, this was not true for samples undergoing phase transformations from DH to HH as can be deduced from the data of Figure 6 for the case of strontium. In the case of the data described in Figure 6, partial phase transformation of DH to HH occurred for the test at 0.2 mol

Figure 4. Influence of initial concentration on phosphate and flouride ion uptake by DH at 40 °C, 0.3 mol L−1 reagent inflow.

L−1. Secondly, in Figure 4 up to the onset of the phase transformation, a linear increase in impurity uptake by DH was detected; at impurity levels higher than 0.3 mol L−1, this trend did not continue. Interestingly, the type of calcium sulfate phase seems to have an effect on the amount of impurity uptake. This was especially true in the case of strontium, where HH was observed to pick up considerably more Sr2+ compared to DH. The experiment with 0.2 mol L−1 Sr2+ was associated with partial phase 6543



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Figure 5. Influence of initial concentration on the uptake of impurities by DH at 40 °C, 0.3 mol L−1 reagent inflow.

morphology. Thus, as it can be appreciated from Figure 7, the presence of Y and La led to more elongated, needle shaped crystals of DH. Even though this morphology resembles that typical for HH, no phase change was evident by XRD analysis (data not shown). Interestingly, although the pickup ratio of strontium was the highest of all impurities investigated, it did not lead to any change in crystal morphology of the DH or HH samples. However, at 0.2 mol L−1 Sr, the appearance of fine needle shaped crystals could be seen alongside the typical DH plates (Figure 8) indicating partial phase transformation of DH to HH, which was also detected by XRD (data not shown). The biggest effect on crystal morphology was caused by increasing levels of phosphate as can be seen from the SEM images in Figure 9. An increasing phosphate concentration led to a deterioration of crystal morphology from compact platelike crystals to slaty crystals with a number of imperfections such as holes. This can be seen in Figure 9a,b. Further increase of phosphate presence resulted in the appearance of fine, needle-shaped crystals alongside with DH crystals as can be seen in Figure 9c. These were confirmed to be HH; in other words, as with the case of strontium, phosphate seems to have induced the partial transformation of DH to HH. Discussion of the Effect of Phosphate. The increased phosphate uptake with increasing initial concentration did not lead to the formation of any distinct calcium phosphate phase. However, it triggered a deterioration of the crystal morphology and induced the partial phase transformation of DH to HH. In order to understand this effect, the speciation of phosphoric

Figure 6. Influence of initial concentration on Sr2+ ion uptake by DH at 40 °C and HH at 80 °C, 0.3 mol L−1 reagent inflow.

transformation of DH to HH (≈10−20%). Even at that limited degree of phase transformation, a strong increase in impurity uptake was observed as is evident by the marked deviation from the linear trend observed over the 0.001 to 0.1 mol L−1 range. Influence of Impurities on Crystal Morphology. Calcium sulfate dihydrate typically crystallizes in a plate-like morphology often associated with twinning in CaCl2-HCl solutions, 3 whereas HH shows an elongated, acicular morphology.6 An evaluation of SEM images showed that among the impurities investigated only rare earth elements (Y, La), phosphate, and strontium had an effect on crystal

Figure 7. Morphology comparison of DH grown in solution containing (a) no impurities, 0.01 mol L−1 (b) Y3+ or (c) La3+ at 40 °C, 0.3 mol L−1 reagent inflow (sample taken at t = 60 min). 6544



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Figure 8. Change in morphology of DH grown in solution containing (a) 0.01 mol L−1, (b) 0.1 mol L−1, and (c) 0.2 mol L−1 Sr2+ at 40 °C, 0.3 mol L−1 reagent inflow (sample taken at t = 60 min).

Figure 9. Change in morphology of DH grown in solution containing increasing levels of phosphoric acid at 40 °C, 0.3 mol L−1 reagent inflow (sample taken at t = 60 min).

acid in the 1.8 mol L−1 CaCl2-6.3 mol L−1 HCl solution (in the presence of 30 g L−1 DH solid) was considered with the help of OLI Stream Analyzer. This analysis revealed that almost all (>99%) of the phosphate exists in the form of the CaH2PO41+ complex as it can be seen from Figure 12. By extrapolation, surface complexes of this type may be thought to form on the DH crystal surface interfering with crystal growth and inducing/activating the DH to HH phase transformation. The presence of a calcium phosphate species on the surface of calcium sulfate crystals was confirmed by XPS analysis as presented later. Furthermore, the CaH2PO41+ formation effectively implies the release of protons into the solution from the phosphoric acid that would result in increased HCl concentration and lower calcium chloride concentration. The increased acid concentration can in turn catalyze the phase transformation of DH to HH as found previously in impurityfree solutions.7 Interestingly, the amount of phosphate present in the calcium sulfate crystals remained nearly constant under the conditions where DH to HH phase transformation was observed (cH3PO4 = 0.3−1.0 mol L−1) to occur. A comparison of the XRD patterns showed that the position of the peaks did not change in comparison to the pure forms (see Figure 10) and the product was a simple mixture of DH and HH crystals. This indicates that most likely the uptake of phosphate was not associated with incorporation into the crystal lattice or with the formation of calcium phosphate phases. Therefore, it seems plausible that the uptake of phosphate took place by surface

Figure 10. XRD pattern of crystals grown in solution with 1 mol L−1 H3PO4 at 40 °C, 0.3 mol L−1 reagent inflow in comparison to impurity-free DH and HH crystals (sample taken at 60 min).

adsorption, similar to phosphonates or phosphate esters that have been shown to adsorb preferentially on growth sites of calcium sulfate DH crystals.30,31 This hypothesis can further explain the observation that at phosphate levels greater than 0.3 mol L−1 the phosphate reporting to the solid remained relatively constant (see Figure 4), this being apparently due to the phase transformation of DH to HH following a dissolution−precipitation type mechanism.7 During phase transformation, the surface of DH changes; it undergoes 6545



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Figure 11. XPS data and fitted Gaussian distributions for DH samples with initial different phosphate levels (a) 0.01 mol L−1 H3PO4, (b) 0.3 mol L−1 H3PO4, (c) 1 mol L−1 H3PO4 at 40 °C, 0.3 mol L−1 reagent inflow (samples taken at 60 min, # = measurement number).

unchanged; e.g., no other peaks appeared meaning no distinct strontium sulfate phase formed. This suggests the substitution of Sr2+ ions for Ca2+ ions in the lattice during the growth process. The peak shift can be appreciated by examining Figure 13a, which shows the pattern of HH crystals grown at 80 °C in a solution containing 0.2 mol L−1 Sr2+ ions at the end of the experiment in comparison to the strontium-free seed material. The newly precipitated material grown on the seed crystals shows peaks that are shifted to lower 2θ values, while the peaks originating from the HH seed remained at their original position. This means that the initial strontium-free seed material is overgrown by a newly formed strontiumincorporating calcium sulfate crystal layer. Further evidence is given in Figure 13b, which shows the evolution of the mole ratio (moles Ca, Sr per mole of DH or HH) in the solid samples over time. During the growth process, the mole ratio of Ca2+ in the crystal decreased, while that of Sr2+ increased correspondingly over time. The value for sulfate on the other hand did not show any trend and remained constant over time. It can be concluded that the likely reason for the substitution seen in the case of strontium, but not in the case of the other cations investigated, originates from the similarity of the ionic radii of the two elements.8 For a coordination number of 8, as is the case for calcium in DH and HH, the ionic radii are Ca2+ = 1.12 Å, Sr2+ = 1.26 Å, Mg2+ = 0.89 Å, and Ba2+ = 1.42 Å.34 It is obvious that the difference in the ionic radius is least between calcium and strontium ions. The radius difference of 0.14 Å also matches well with the peak shift seen in the samples which was approximately 0.11 Å. These results are in agreement with results published by Pouria et al.,35 who found a similar peak shift for gypsum crystals grown from a hemihydrate slurry containing strontium nitrate. Similar results of high foreign cation uptake referring to ionic radii were reported for the uptake of Ce3+, whose radius is also very close to the radius of Ca2+ ions.14 A high uptake of strontium ions, in comparison to other ions, was also observed by Kushnir,36 who investigated the partition of Sr, Mg, Na, and K ions present in seawater brines during the phase transformation between calcium sulfate phases. Another interesting observation, requiring discussion, was the fact that HH took up more strontium in comparison to DH at the same solution impurity concentration. A similar observation was made by Martynowicz et al.8 for DH and HH grown in phosphoric acid solutions. The authors suggest a lower lattice deformation energy in the HH lattice could be responsible for this phenomenon.

dissolution during the process, hence affecting its ability to accommodate the adsorption of ions. The adsorption of phosphate (most likely calcium dihydrogen phosphate complex in correspondence to the speciation data presented in Figure 12) as already mentioned can also be held responsible for the observed crystal morphology deterioration. It is reasonable to assume that the surface calcium phosphate complexes effectively block access of the SO42− growth units to the crystal surface and hence holes and other imperfections developed as a result. XPS measurements of the binding energy for phosphate confirmed the presence of phosphate on the crystal surface (see Figure 11a− c). A comparison of the average P2p level binding energy measured by fitting Gaussian distribution functions to the data give a value of 132.94 ± 0.3 eV. On the basis of the values published by Franke et al.,32 it is possible that a Ca3(PO4)2 (binding energy of 133.1 eV) compound may have locally formed on the surface of the crystal as surface precipitate layer. However, by extrapolation from the calculated speciation data (Figure 12), it is more likely that instead CaHPO4 or Ca(H2PO4)2, which show a binding energy of 133.1 or 132.6 eV for the P2p level,33 respectively, have formed on the crystal surface.

Figure 12. Speciation depending on initial H3PO4 concentration present in 1.8 mol L−1 CaCl2, 6.3 mol L−1 HCl solution with 30 g L−1 DH solid at 40 °C.

Discussion of the Effect of Strontium. Among the investigated cationic impurities, strontium showed the most significant uptake during crystal growth. In order to explain the high strontium uptake, XRD measurements were made. From these, it was seen that the characteristic peaks of DH and HH shifted to lower 2θ values, while the pattern itself remained 6546



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Figure 13. XRD patterns and mole ratio evolution of crystals grown in solution with 0.2 mol L−1 Sr2+ and reagent inflow of 0.3 mol L−1.

Figure 14. DSC patterns of HH crystals grown at 80 °C, reagent inflow of 0.3 mol h−1 at different impurity concentrations: (a) no impurities and 0.1 mol L−1 Sr2+ and (b) 0.2 mol L−1 Sr2+ (two measurements of each sample are shown).

Å). This could explain the observed higher uptake of lanthanum compared to yttrium. However, due to the different valency of these ions, direct substitution for calcium is more difficult without the incorporation of additional charge compensating ions. It was reported that lanthanide sulfates can be incorporated in HH (but less in DH)40 by heterovalent isomorphic inclusion in the presence of alkali metal ions acting as vacancy compensators. The use of Na2SO4 as a source for sulfate ions during the growth process in the present study makes such charge compensating ions available. The measured uptake data summarized in Table 2 conform indeed to the same incorporation model 41 with HH accommodating more lanthanide incorporation than DH, although relatively minor in comparison to Sr2+ incorporation. In addition to the reported uptake, the present work found yttrium and lanthanum to result in DH crystal morphology changes, without this being associated with phase transformation according to XRD measurements. The resulting

Supplementary DSC analysis shown in Figure 14 involved comparison of impurity-free HH crystals with HH crystals grown in solution containing strontium at two different levels. All DSC patterns exhibited the characteristic exothermic peak that identifies the α-form of calcium sulfate hemihydrate.37−39 However, the crystals that were grown in the strontium solution developed an additional shoulder on the exothermic peak in comparison to the Sr-free crystals. This behavior is caused by the fact that the incorporated Sr2+ ions influence the crystal lattice arrangement and hence the observed effect upon heating of the sample. Furthermore, the crystals with the highest strontium content also showed a different behavior in the endothermic part of the thermogram. The endothermic peak associated with the dehydration of the crystals is seen to have almost disappeared and been replaced by two small peaks instead of a single larger one. This might be explained by the fact that the crystal lattice structure of the impure solid is weaker and the energy required to break the crystal water bonds is lower compared to Sr-free crystals. Discussion of the Effect of Rare Earth Elements. The results relating to the uptake of the studied rare earths, Y3+ and La3+, are quite similar to other studies that found these elements to incorporate to some extent into solid calcium sulfate phases. One possible explanation for this behavior may be based on the ionic radii of these ions (1.019 and 1.16 Å for Y3+ and La3+, respectively) in comparison to calcium ions (1.12

Table 2. Comparison of Impurity Uptake U of Y and La in Calcium Sulfate Dihydrate and Hemihydrate for the Initial Impurity Concentration of 0.01 mol L−1 uptake in... DH HH 6547

Y [molY molsolid−1] −4

2.15 × 10 1.3 × 10−3

La [molLa molsolid−1] 3.4 × 10−4 4.2 × 10−3



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elongated crystals can be explained by preferential adsorption of REE ions and their subsequent overgrowth. Apparently, these ions seem to be active only at certain crystal faces causing preferential growth to occur resulting in DH crystals with an elongated shape as opposed to their typically found plate-like morphology. These results also agree with findings of other researchers, who worked on the crystallization of calcium sulfates in the phosphoric acid system. Thus, the presence of La3+ has been reported to retard growth of DH crystals on all faces except the {011} and {11̅ 1} faces.8,14 This most likely caused, as well in the present work, the observed elongated morphology of DH crystals in the presence of higher concentrations (0.001 to 0.01 mol L−1) of lanthanum or yttrium ions. Finally, for such substitution to take place, the impurity cation should have a higher affinity for the anion of the host crystals (SO42−) than Ca2+ itself. This can be evaluated on the basis of the stability constants of Ca2+, Sr2+, La3+, and Y3+ with SO42− anions in solution. Upon review of the relevant data, it was indeed noted that Ca2+ ions have lower stability constants for complex formation with SO42− compared to Sr2+, La3+, and Y3+.42,43 Hence, the postulated incorporation of these cations is justified also on thermodynamic grounds.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (T.F.); george. [email protected] (G.P.D.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS

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REFERENCES

The financial support of Arafura Resources, Ltd., Perth, WA, Australia for this project is gratefully acknowledged.

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CONCLUSIONS The influence of inorganic solute impurities on the crystallization kinetics of calcium sulfate dihydrate and hemihydrate under nearly constant supersaturation relevant to industrial hydrochloric acid regeneration, from CaCl2 solutions obtained during HCl leaching of rare earth-containing apatite ores, has been studied. Impurities investigated were K+, Mg2+, Sr2+, Ba2+, Al3+, Fe2+, Fe3+, La3+, Y3+, F−, and PO43−. The experiments were conducted at 40 and 80 °C in concentrated CaCl2 (1.8 mol L−1)-HCl (6.3 mol L−1) solutions relevant to HCl regeneration by reactive crystallization of calcium sulfate DH and HH. The most significant findings are the following: (1) Calcium sulfate dihydrate took up Sr2+ and phosphate and to much lesser extent La3+ and Y3+ during crystal growth. At the same molar concentrations, the uptake of strontium was the highest. (2) Except in the case of phosphate, no significant influence on the crystal growth rate of DH was seen. For phosphate concentrations from 0.001 to 0.2 mol L−1, the crystal growth rate increased from 0.29 to 0.39 nm s−1 m−2. Concentrations of 0.3 up to 1.0 mol L−1 led to partial phase transformation of DH to HH. Similar to phosphate, Sr2+ was also found to induce a similar phase transformation from DH to HH. The DH to HH transformation process was found to agree with a dissolution-precipitation-based phase transformation mechanism. (3) The morphology of DH was influenced by the presence of lanthanum and yttrium ions toward more elongated crystal shapes. At lower phosphate ion concentrations, the DH morphology was found to deteriorate by developing holes and a slaty structure with many imperfections. (4) The uptake of Sr2+ ions is suggested to happen via a substitution mechanism during crystal growth. Similar ionic radii between Ca2+ and Sr2+ ease the substitution process without noticeably influencing the crystal morphology. (5) The uptake of phosphate by the calcium sulfate crystals is suggested to take place via an adsorption process on active sites on the crystal surface. This was confirmed by XPS measurements that provided evidence of a CaHPO4/Ca(H2PO4)2 surface species. 6548



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