Dissolution , Solubility , and Stability of the Basic Ferric Sulfate-Arsenates [ Fe ( SO 4 ) x ( AsO 4 ) y ( OH ) z · nH 2 O ] at 25 – 45 ° C and pH 2 – 10

Basic ferric sulfate-arsenates [FeSAsOH, Fe(SO4)x(AsO4)y(OH)z·nH2O] were prepared and characterized to study their potential fixation of arsenic in the oxidizing and acidic environment through a dissolution for 330d. ,e synthetic solids were well-shaped monoclinic prismatic crystals. For the dissolution of the sample FeSAsOH–1 [Fe(SO4)0.27(AsO4)0.73 (OH)0.27·0.26H2O] at 25–45°C and initial pH 2, all constituents preferred to be dissolved in the order of AsO4> SO4> Fe in 1–3 h, in the order of SO4>AsO4> Fe from 1–3 h to 12–24 h, and finally in the order of SO4> Fe>AsO4. ,e released iron, sulfate, and arsenate existed dominantly as Fe/Fe(OH)/FeSO4, HSO4/SO4/FeSO4, and H3AsO4/H2AsO4, respectively. ,e higher initial pHs (6 and 10) could obviously inhibit the release of Fe from solid into solution, and the solid components were released in the order of SO4>AsO4> Fe. ,e crystal tops were first dissolved, and the crystal surfaces were gradually smoothed/ rounded until all edges and corners disappeared. ,e dissociations were restricted by the Fe-O(H) breakdown in the FeO6 octahedra and obstructed by the OH and AsO4 tetrahedra outliers; the lowest concentration of the dissolved arsenic was 0.045mg/L. Based on the dissolution experiment at 25°C and pH 2, the solubility products (Ksp) for the basic ferric sulfate-arsenate [Fe(SO4)0.27(AsO4)0.73 (OH)0.27·0.26H2O], which are equal to the ion activity products (logˍIAP) at equilibrium, were calculated to be -23.04± 0.01 with the resulting Gibbs free energies of formation (ΔGf) of −914.06± 0.03 kJ/mol.

Basic ferric sulfate-arsenates [FeSAsOH, Fe(SO 4 ) x (AsO 4 ) y (OH) z ·nH 2 O] were prepared and characterized to study their potential fixation of arsenic in the oxidizing and acidic environment through a dissolution for 330d. e synthetic solids were well-shaped monoclinic prismatic crystals. For the dissolution of the sample FeSAsOH-1 [Fe(SO 4 ) 0.27 (AsO 4 ) 0.73 (OH) 0.27 ·0.26H 2 O] at 25-45°C and initial pH 2, all constituents preferred to be dissolved in the order of AsO 4 3− > SO 4 2− > Fe 3+ in 1-3 h, in the order of SO 4 2− > AsO 4 3− > Fe 3+ from 1-3 h to 12-24 h, and finally in the order of SO 4  /FeSO 4 + , and H 3 AsO 4 0 /H 2 AsO 4 − , respectively. e higher initial pHs (6 and 10) could obviously inhibit the release of Fe 3+ from solid into solution, and the solid components were released in the order of SO 4 2− > AsO 4 3− > Fe 3+ . e crystal tops were first dissolved, and the crystal surfaces were gradually smoothed/ rounded until all edges and corners disappeared. e dissociations were restricted by the Fe-O(H) breakdown in the FeO 6 octahedra and obstructed by the OH − and AsO 4 tetrahedra outliers; the lowest concentration of the dissolved arsenic was 0.045 mg/L. Based on the dissolution experiment at 25°C and pH 2, the solubility products (K sp ) for the basic ferric sulfate-arsenate [Fe(SO 4 ) 0.27 (AsO 4 ) 0.73 (OH) 0.27 ·0.26H 2 O], which are equal to the ion activity products (logˍIAP) at equilibrium, were calculated to be -23.04 ± 0.01 with the resulting Gibbs free energies of formation (ΔG f o ) of −914.06 ± 0.03 kJ/mol.

Introduction
Arsenic is an extremely toxic byproduct of the mining and smelting of precious and nonferrous metals [1][2][3][4][5] and a common metalloid element in mineral feedstocks, which could be mobilized/discharged during the metallurgical operation, and it results in a serious environmental problem [6,7]. With the pyrite and arsenopyrite destruction in the Au extraction from refractory Au-sulfide ores, large amounts of arsenic would enter into the residues together with ferric ion and sulfate [8,9]. e characterization of the different residual solids formed in the autoclave system and their solubilities are related to the environmental and metallurgical problems that have not been finally solved, and their physicochemical properties are still not well understood [10,11]. e solubility and stability of arsenical solid wastes depend on the type of arsenic-containing phases and their crystallinity [2]. Contrary to the poorly crystalline Fe(III)-AsO 4 compounds coprecipitated during usual neutralization of hydrometallurgical effluents [12], the controlled or autoclave processing resulted in the formation of wellcrystallized precipitates [3]. e basic ferric sulfate-arsenate [FeSAsOH, Fe(SO 4 ) x (AsO 4 ) 1−x (OH) x ·wH 2 O] [2] was one of the three crystalline ferric sulfate-arsenates, which were detected to crystallize when the arsenic-containing minerals as the raw materials were treated in the autoclave under the hydrothermal condition of the Fe(III)-SO 4 -AsO 4 solution (150-230°C), on which recent characterizations indicated that arsenic was immobilized in the FeSOH [Basic ferric sulfate]-FeSAsOH solid solution [4]. e crystalline phase "Type 2" [Fe 4 (SO 4 ) y (AsO 4 ) 3 (OH) x ] was found to precipitate under the autoclave processing condition of Fe/As ratio <1.5 and 200-225°C, which was tetragonal or monoclinic and exhibited a similar leaching behavior of relatively lower solubility (<0.34 mg/L As) with the amorphous mineral scorodite [FeAsO 4 ·2H 2 O] (<0.8 mg/ L As) and could meet the TCLP leachability criterion (<5 ppm) [13]. e well-crystallized ferric arsenate could be successfully prepared by the hydrothermal precipitation method from the Fe-SO 4 -AsO 4 -H 2 O solution at pH<2.5 and >210°C [10]. e crystalline "Phase 3" [Fe x (SO 4 , AsO 4 ) (OH) y ·nH 2 O] was a monoclinic polytype of the basic ferric sulfate and it formed at 175-225°C through the isomorphic replacement of AsO 4 for SO 4 with the OH decrease to keep the charge balances, for example, the Fe(SO 4 ) 0.6 (AsO 4 ) 0.4 (OH) 0.6 ·0.4H 2 O solid solution [4], although its structure was suggested to be triclinic (pseudoorthorhombic) in the further research. At the Fe/As mole ratios of 3.65-1.77, 1.55-1.20, and 1.00, only the "Phase 3", the mixture of "Phase 3" and minor FeAsO 4 ·0.75H 2 O and only the phase FeA-sO 4 ·0.75H 2 O formed, respectively [11]. e arsenic concentrations leached from "Phase 3" at room temperature for 40 h into the water were steadily <0.1 mg/L, which suggested that "Phase 3" might be a suitable solid phase for arsenic removal [11]. Lately, the crystal structure of the basic ferric sulfate-arsenate [FeSAsOH, Fe(SO 4 ) 0.7−0.2 (AsO 4 ) 0.2−0.7 (OH) 0.7−0.2 ] was proposed to be constructed of the FeO 6 octahedral chains that were cross-linked through the SO 4 / AsO 4 tetrahedra, which formed alternative layers of SO 4 / AsO 4 tetrahedra and FeO 6 octahedra [14].
Precise examination of the solubility and stability of the basic ferric sulfate-arsenates (FeSAsOH) is important to promote the risk evaluation on arsenic-polluted sites and to prevent the arsenic to be released back into the environment [15]. But, until now, most researches were carried out mainly on the forming conditions, structural characterization, and leachability of FeSAsOH [2,4,13], and little information about the dissolution mechanism, solubility, and stability of the basic ferric sulfate-arsenates is nowadays accessible. Although previous works indicated enormously low reactivity and solubility of these compounds in both acid and alkaline solutions [9,11,14], further studies on the arsenic release mechanism and their long-term solubility at various pHs are essential.
In this work, the crystalline FeSAsOH solids from Fe(III)-SO 4 -AsO 4 solutions are prepared by a simple hydrothermal method. e dissolution mechanism, long-term solubility, and stability of the FeSAsOH solids at different solution pHs and temperatures are examined. e structural and morphological variations of the synthetical FeSAsOH phases before and after dissolution are examined using various instruments; besides, the potential for arsenic fixation is discussed.

Synthesis.
To prepare the starting solutions, the analyticreagent grade Fe 2 (SO 4 ) 3 ·9H 2 O (Shanghai Aladdin Biochemical Technology Co., Ltd.) and As 2 O 5 (Hengyang Industrial Corporation, Shuikoushan Mining Administration, Hunan, China) were dissolved in ultrapure water in the chosen mole proportion to give various starting Fe(III)/ AsO 4 and SO 4 /AsO 4 mole ratios (Table 1). e starting solutions were used at their natural pHs (0.49-0.90). Finally, each resulting mixture was vigorously agitated at 600 r/min for 0.5 h and then moved into a 0.2 L stainless-steel autoclave. After heating at 200°C for 1 day, the resulting slurry was cooled and separated using vacuum filtration. Finally, the precipitate was cleaned 3 times by using ultrapure water and dried at 110°C for 1 day.

Characterization.
To determine the bulk compositions, 50 mg of each basic ferric sulfate-arsenate was digested in 20 mL 6 M hydrochloric acid that was diluted to 50 mL by using HNO 3 solution. e iron, sulfur, and arsenic concentrations were analyzed by a Perkin-Elmer Optima 7000DV inductively coupled plasma-optical emission spectrometer (ICP-OES) with the proper reference standards. e H 2 O contents were then estimated by the mass balance based on the thermal analytic results, which were obtained from 30°C to 1135°C in nitrogen gas using a Netzsch STA 409 thermogravimetric analyzer (TGA). All of the prepared solids before and after dissolution were studied by an X'Pert PRO X-ray diffractometer (XRD) with Cu-Kα radiation of 1.540598Å (40 mA and 40 kV) in the 2θ range from 5°to 90°at the scan step of 0.0263°and the scan rate of 5.3333°/min and recognized by comparing the recorded XRD spectra to literature references. e functional groups and the morphologies of the basic ferric sulfate-arsenates were analyzed by a Nicolet Nexus 470 Fourier transform infrared spectrophotometer (FT-IR) over the spectral range from 400 to 4000 cm −1 and a Jeol JSM-7900F field emission scanning electron microscope (FE-SEM) with an energy dispersive spectrometer (EDS), respectively.

Dissolution Tests.
Five grams of each dried basic ferric sulfate-arsenate was added to 0.1 L of HNO 3 solutions (pH 2 and 6) or NaOH solution (pH 10) in a polypropylene bottle, which was capped and put in the temperature-control water bathes (25°C, 35°C, or 45°C). e pHs of the mixing slurries with no adjustment were recorded periodically. 5 mL of the solution from each bottle was collected at the fixed intervals from 1 h to 330 days, filtered and immediately stabilized using 0.2% HNO 3 solution, followed by the measurement for iron, sulfur, and arsenic using ICP-OES or an atomic absorption spectrometer (AAS, Perkin-Elmer AAnalyst 700). To hold the starting volume constant, the equivolume HNO 3 or NaOH solutions were supplemented after each sampling. e effect of this volumetric variation on the elemental concentrations was considered in the following thermodynamic simulation. After 330 days (7920 h), the remaining solids were taken out from the bottles and characterized using various instruments, as described formerly. At initial pH 2, the tests were made twice to check the repeatability.

XRD.
e XRD spectra of the basic ferric sulfatearsenates are shown in Figure 1 and Figure S1 in Supplementary Materials. In the patterns, the major peak positions of the basic ferric sulfate-arsenates "Ba-5" [3], "Phase 3" [11], and "Type 2" [13] were plotted in the higher chart. e samples FeSAsOH-1 and FeSAsOH-3 agreed well with the arsenical compound "Type 2" [13] in the major peak positions in the XRD spectra ( Figure 1) with the strongest lines (d obs , I obs ) of 3.38(100), 3.26(44), 2.64(15), 1.63(11), 2.33(11), and 2.06 (10). e samples FeSAsOH-2 and FeSAsOH-5 were also in agreement with the arsenical compounds "Type 2" [13] in the major peak positions ( Figure 1, Figure S1 in Supplementary Materials). e major XRD peaks showed that the sample FeSAsOH-4 was identified as a mixture of the basic ferric arsenate sulfate [13] and ferric orthoarsenate sub- Figure S1 in Supplementary Materials), which was related to its forming condition [3,16]. It was also reported previously that some of the ferric sulfate-arsenate (FAS) samples contained traceto-minor amount of the basic ferric sulfate (BFS), as indicated by the presence of overlapping peaks at 2θ values of −26.5°and 27.5°, which occurred as the asymmetric and broadened peaks. e acicular crystals of the ferric orthoarsenate subhydrate (FeAsO 4 ·0.75H 2 O) occurred in trace amounts in the samples [9]. No obvious variations were recognized in the XRD spectra after dissolution ( Figure 1, Figure S1 in Supplementary Materials).

FE-SEM.
It was suggested that the crystal structure of the basic ferric sulfate-arsenates could be related to monoclinic polytypes [4,11] and orthorhombic-monoclinic [2] and triclinic (pseudo-orthorhombic) crystal systems [9]. e morphologies of the synthetic basic ferric sulfate-arsenates before and after dissolution for 330 days, which were evaluated by X-ray diffraction, were also examined by FE-SEM with EDS (Figures 3-5, Figure S3 in Supplementary Materials).
For the synthetic solid FeSAsOH-1, the particles consisted of aggregates and the individual crystallites were generally the well-shaped monoclinic prismatic crystals with the size of <5 μm (Figure 3). e EDS analysis showed that the surface had a relatively higher As/(As + S) molar ratio (0.77-0.85) than that of the bulk solid (0.73). After the dissolution at different temperatures and pHs for 330 d, almost all tops of the crystals and some edges and corners were corroded, and finally, the crystals became smoothed/ rounded (Figures 3 and 4).
For the synthetic solid FeSAsOH-4, the basic ferric sulfate-arsenate particles appeared in two shapes: the wellshaped monoclinic prismatic crystals and the rounded type particles ( Figure 5). e EDS results confirmed that the two shapes were identical, with one grown to its full monoclinic  Transmittance (a.u.)    Journal of Chemistry crystal and the other redissolved and/or recrystallized to form its rounded shape having the As/(As + S) molar ratios of 0.85-0.89 and 0.85-0.87, respectively ( Figure 5). Furthermore, the ferric orthoarsenate subhydrate [FeAsO 4 ·0.75H 2 O] aggregates appeared in the well-shaped triclinic prismatic crystals without terminations having the As/(As + S) mole ratios of 0.98-1.00 and with terminations having the As/ (As + S) mole ratios of 0.94-0.97 ( Figure 5), which was well in accordance with the XRD analysis ( Figure 1). Similar to FeSAsOH-1, almost all tops of the crystals and some edges and corners were dissolved, and the crystals became smoothed/rounded after the dissolution at different temperatures and pHs for 330 d ( Figure S3, Supplementary material).  (Figure 6), the aqueous pHs fluctuated between 1.88 and 2.09 and attained a steady state of 1.98 after 5760 h. At initial pH 6, the pHs decreased gradually to 3.51 within 48 h and then increased steadily to 3.68 after 5040 h. At initial pH 10, the pHs decreased gradually to 3.77 within 2880 h and then increased steadily to 4.04 after 5040 h. e released Fe 3+ concentrations rose gradually to 0.122-0.123 mmol/L after 5760 h. e final Fe 3+ concentrations declined from 0.122 to 0.123 mmol/L to 0.000203-0.000217 mmol/L as the initial pH was increased from 2 to 10 for the dissolution at 25°C (Figure 6). e dissolved SO 4 2− concentrations rose quickly to 0.073130 mmol/L within 12 h and then declined/increased with a slight fluctuation to 0.096832-0.099170 mmol/L after 5760 h for the dissolution at 25°C and pH 2. e final SO 4 2− concentrations declined from 0.096832 to 0.099170 mmol/L to 0.072195-0.072663 mmol/L as the initial pH was increased from 2 to 10 for the dissolution at 25°C and rose obviously to 0.244496-0.247302 mmol/L as the temperature was increased from 25°C to 45°C. e dissolved AsO 4 3− concentrations rose quickly to 0.091229 mmol/L after 12 h and then slightly declined to 0.019954-0.026161 mmol/L over 12-120 h, and after that, it increased/decreased once again steadily to 0.059796-0.060330 mmol/L after 5760 h for the dissolution at 25°C and pH 2 ( Figure 6). e final dissolved AsO 4 3− concentrations at 25°C and pH 6 were found to be the lowest of 0.006838-0.006851 mmol/L, and it declined from 0.019954 to 0.026161 to 0.035103-0.036772 mmol/L as the temperature was increased from 25 to 45°C.  Figure S4 in Supplementary Materials). In the previous studies, it was shown that, during short-term environmental stability tests, pure basic ferric sulfate-arsenates release less than 5 mg/L of arsenic into solution [3]. e "Type 2" with an X As mole fraction of 0.75 produced a similar As concentration (i.e., 0.8 mg/L) in the TCLP leaching [13].

Dissolution
For the dissolution of the sample FeSAsOH-1 at initial pH 2 (Figure 6 /Fe-rich residuals at initial pH 2. Additionally, the higher initial pHs (6 and 10) could obviously inhibit the release of Fe 3+ from solid into solution and prefer to form iron-rich residuals, and the solid components were released in the order of SO 4 2− > AsO 4 3− > Fe 3+ . e speciation of toxic metals and metalloids is an important factor for their mobility in the environment [1]. During the dissolution of the basic ferric sulfate-arsenate FeSAsOH-1 at initial pH 2, the released components existed mainly as Fe 3+ /Fe(OH) 2+ (Figure 6), all solutions were undersaturated with respect to Fe 2 (SO 4 ) 3 , FeAsO 4 ·2H 2 O, maghemite, ferrihydrite, and H-jarosite, suggesting that the formation of all these ferric sulfates/arsenates was thermodynamically unfavorable. By contrast, all solutions were very closely saturated or oversaturated in respect to lepidocrocite (SI � −1.90-2.46), goethite (SI � −0.36-3.34), and hematite (SI � 1.77-9.07), suggesting that the formation of these iron-rich precipitates was thermodynamically favorable. Even though the XRD measurement indicated that no other phases than the basic ferric sulfate-arsenates existed, their presence in a small amount under the detection limit could not be excluded (Figure 1). e FeSAsOH structure is constructed from layered Fe-O 6 octahedra that are cross-linked by AsO 4 and SO 4 tetrahedra ( Figure 7) [4,9,18]. For the dissolution of the basic ferric sulfate-arsenate [FeSAsOH-1] at pH ≥2, the gradual decrease in solution pH indicated an OH − depleting. e reaction was expressed by the favorable release of AsO 4 3− , followed by SO 4 2− and Fe 3+ from solid into solution, whereas Fe 3+ was preferentially left behind to form a residual octahedral outlier (Figure 7). e components at the tops of the crystals were first dissolved and the crystal surfaces were gradually smoothed/rounded until all edges and corners disappeared (Figure 7). In the subsequent recrystallization, the arsenate ions and Fe 3+ cations were also preferentially removed from the aqueous solution, while sulfate ions were preferentially left in the aqueous solution. e dissolution was finally restricted by the Fe-O(H) breakdown in the FeO 6 octahedra and obstructed by the OH − and AsO 4 tetrahedra outliers.
is could be also verified by comparing the FE-SEM-EDS measurements on the FeSAsOH-1 surface before and after dissolution at 25°C and initial pH 2 for 7920 h, which showed that the (AsO 4 +SO 4 )/Fe, AsO 4 /Fe, and SO 4 /Fe molar ratios increased from 1.06, 0.85, and 0.20 to 0.99, 0.80, and 0.19 after 330 days of dissolution, respectively. e AsO 4 /(AsO 4 +SO 4 ) mole ratios on the solid surface showed no obvious variation after dissolution. In consideration of the crystal structure of FeSO 4 (OH), the nonstoichiometric dissolution behavior can be explained by the precipitation of a ferric oxyhydroxide phase or by the preferential release of SO 4 leaving behind chains of FeO 6 octahedra [9].

FeSAsOH-2-FeSAsOH-5.
For the dissolution of the basic ferric sulfate-arsenates FeSAsOH-2, FeSAsOH-3, and FeSAsOH-5, all components preferred to be released in the order of AsO 4 3− > SO 4 2− > Fe 3+ within 1-3 h, in the order of SO 4 2− > AsO 4 3− > Fe 3+ from 1 to 3 h to ∼24 h, and finally, in the order of SO 4 2− > Fe 3+ > AsO 4 3− , suggesting a nonstoichiometric dissolution and/or formation of AsO 4 3− /Ferich residuals at initial pH 2 ( Figure S4 in Supplementary Materials). Additionally, the higher initial pHs (6.00 and 10.00) could obviously inhibit the release of Fe 3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO 4 2− > AsO 4 3− > Fe 3+ . For the dissolution of FeSAsOH-4, all components were preferentially released in the order of AsO 4 3− > SO 4 2− > Fe 3+ in ∼1 h, in the order of SO 4 2− > AsO 4 3− > Fe 3+ from ∼1 h to ∼45 d, and finally, in the order of SO 4 2− > Fe 3+ > AsO 4 3− after ∼45 d, suggesting a nonstoichiometric dissolution and/or formation of AsO 4 3− /Fe-rich residuals at initial pH 2 ( Figure S4 in Supplementary Materials). Additionally, the higher initial pHs (6.00 and 10.00) could obviously inhibit the release of Fe 3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO 4 2− > AsO 4 3− > Fe 3+ . However, it was also observed that the aqueous solutions for the 8 Journal of Chemistry FeSAsOH-4 sample showed a slower evolution than those for FeSAsOH-1 and FeSAsOH-3, which was related to the higher H 2 O mole ratio in FeSAsOH-4.

Solubility Calculation.
Based on the batch dissolution results of the synthetic basic ferric sulfate-arsenates, the computing of the aqueous Fe 3+ , SO 4 2− , AsO 4 3− , and OH − activities in the final steady state (5760 h, 6480 h, 7200 h, and 7920 h) was achieved by the PHREEQC program [21] with the minteq.v4.dat database [22] and the thermodynamic properties for aqueous metal-arsenate species [23,24]. e ion activity products (logˍIAP) for the basic ferric sulfatearsenates were calculated, which were equal to their solubility products (logˍK sp ) at the dissolution equilibrium. e key speciations are summarized in Table S1 in Supplementary Materials. e dissolution of the basic ferric sulfate-arsenates [Fe(SO 4 ) x (AsO 4 ) y (OH) z ·nH 2 O] is described by e ion activity product (IAP) is expressed according to For equation (6), Rearranging, (2)

10
Journal of Chemistry   (Table 2). For all the five synthetic basic ferric sulfate-arsenates, the ion activity products (logˍIAPs) varied between −26.17 ± 0.01 and −22.01 ± 0.02. e solubilities decreased slightly with the increasing initial pH and were not obviously affected by the system temperature. e experimental results on dissolution were highly reproducible ( Table 2-Table 6).
Using /H 2 AsO 4 − , respectively. e higher initial pHs (6 and 10) could obviously inhibit the release of Fe(III) from solid into solution, and the solid components were released in the order of SO 4 > AsO 4 > Fe. e detaching was restricted by the Fe-O(H) breakdown in the FeO 6 octahedra and obstructed by the OH − and AsO 4 tetrahedra outliers; the lowest concentration of the dissolved arsenic was 0.045 mg/L. e components at the tops of the crystals were first dissolved, and the crystal surfaces were gradually smoothed/rounded until all edges and corners disappeared.
Based on the dissolution experiment at 25°C and pH 2 for 330 d, the solubility products (K sp ) for Fe(SO 4 ) 0.27 (AsO 4 ) 0.73 (OH) 0.27 ·0.26H 2 O, which are equal to the ion activity products (logˍIAP) at equilibrium, were calculated to be −23.04 ± 0.01 with the resulting Gibbs free energies of formation (ΔG f o ) of −914.06 ± 0.03 kJ/mol. For all Data Availability e powder XRD data in XML format, the FT-IR data in XLSX format, and all solution analytical data in XLSX format used to support the findings of this study are available from the corresponding author upon request.

Conflicts of Interest
e authors declare that there are no conflicts of interest regarding the publication of this paper.