# Dissolution, Solubility, and Stability of the Basic Ferric Sulfate-Arsenates [Fe(SO4)x(AsO4)y(OH)z·nH2O] at 25–45°C and pH 2–10

Authors: Zongqiang Zhu; Hongqu Yang; Jie Liu; Yinian Zhu; Shen Tang; Lihao Zhang; Xingxing Wang
Journal: Journal of Chemistry
 (2021)

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## Abstract

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. The 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 AsO43− > SO42− > Fe3+ in 1–3 h, in the order of SO42− > AsO43− > Fe3+ from 1–3 h to 12–24 h, and finally in the order of SO42− > Fe3+ > AsO43−. The released iron, sulfate, and arsenate existed dominantly as Fe3+/Fe(OH)2+/FeSO4+, HSO4−/SO42−/FeSO4+, and H3AsO40/H2AsO4−, respectively. The higher initial pHs (6 and 10) could obviously inhibit the release of Fe3+ from solid into solution, and the solid components were released in the order of SO42− > AsO43− > Fe3+. The crystal tops were first dissolved, and the crystal surfaces were gradually smoothed/rounded until all edges and corners disappeared. The 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 0.045 mg/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 (ΔGfo) of −914.06 ± 0.03 kJ/mol.


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## Body

## 1. Introduction

Arsenic is an extremely toxic byproduct of the mining and smelting of precious and nonferrous metals [1–5] 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]. The 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].The 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)-AsO4 compounds coprecipitated during usual neutralization of hydrometallurgical effluents [12], the controlled or autoclave processing resulted in the formation of well-crystallized precipitates [3]. The basic ferric sulfate-arsenate [FeSAsOH, Fe(SO4)x(AsO4)1−x(OH)x·wH2O] [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)-SO4-AsO4 solution (150–230°C), on which recent characterizations indicated that arsenic was immobilized in the FeSOH [Basic ferric sulfate]–FeSAsOH solid solution [4].The crystalline phase “Type 2” [Fe4(SO4)y(AsO4)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 [FeAsO4·2H2O] (<0.8 mg/L As) and could meet the TCLP leachability criterion (<5 ppm) [13].The well-crystallized ferric arsenate could be successfully prepared by the hydrothermal precipitation method from the Fe-SO4-AsO4-H2O solution at pH<2.5 and >210°C [10]. The crystalline “Phase 3” [Fex(SO4, AsO4) (OH)y·nH2O] was a monoclinic polytype of the basic ferric sulfate and it formed at 175–225°C through the isomorphic replacement of AsO4 for SO4 with the OH decrease to keep the charge balances, for example, the Fe(SO4)0.6(AsO4)0.4(OH)0.6·0.4H2O 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 FeAsO4·0.75H2O and only the phase FeAsO4·0.75H2O formed, respectively [11]. The 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(SO4)0.7−0.2(AsO4)0.2−0.7(OH)0.7−0.2] was proposed to be constructed of the FeO6 octahedral chains that were cross-linked through the SO4/AsO4 tetrahedra, which formed alternative layers of SO4/AsO4 tetrahedra and FeO6 octahedra [14]. The “Type 2” [Fe4(SO4)y(AsO4)3(OH)x] was the same as the “Phase 3” [Fex(SO4, AsO4) (OH)y·nH2O], although their crystal structure and stability remained unknown [3, 4].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)-SO4-AsO4 solutions are prepared by a simple hydrothermal method. The dissolution mechanism, long-term solubility, and stability of the FeSAsOH solids at different solution pHs and temperatures are examined. The 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.


## 2. Materials and Methods

### 2.1. Synthesis

To prepare the starting solutions, the analytic-reagent grade Fe2(SO4)3·9H2O (Shanghai Aladdin Biochemical Technology Co., Ltd.) and As2O5 (Hengyang Industrial Corporation, Shuikoushan Mining Administration, Hunan, China) were dissolved in ultrapure water in the chosen mole proportion to give various starting Fe(III)/AsO4 and SO4/AsO4 mole ratios (Table 1). The 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.Table 1
Preparation and component of the synthetical basic ferric sulfate-arsenates.
SamplePrecursor solution (100 mL)Solid composition [Fe(SO4)x(AsO4)y(OH)z·nH2O]Fe(III) + SO4 solutionH3AsO4 solutionMole ratioFe2(SO4)3·9H2O (g)H2O (mL)1 M H3AsO4 (mL)H2O (mL)Fe/AsO4SO4/AsO4FeSAsOH-15.62015010402.003.00Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2OFeSAsOH-28.43025010403.004.50Fe(SO4)0.37(AsO4)0.59(OH)0.49·0.28H2OFeSAsOH-38.43025015352.003.00Fe(SO4)0.27(AsO4)0.72(OH)0.30·0.21H2OFeSAsOH-48.43025020301.502.25Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O + FeAsO4·0.75H2O mixtureFeSAsOH-511.24025020302.003.00Fe(SO4)0.24(AsO4)0.70(OH)0.42·0.16H2O


### 2.2. 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 HNO3 solution. The 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. The H2O 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 1.540598 Å (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. The 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.


### 2.3. Dissolution Tests

Five grams of each dried basic ferric sulfate-arsenate was added to 0.1 L of HNO3 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). The 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% HNO3 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 HNO3 or NaOH solutions were supplemented after each sampling. The 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.


## 2.1. Synthesis

To prepare the starting solutions, the analytic-reagent grade Fe2(SO4)3·9H2O (Shanghai Aladdin Biochemical Technology Co., Ltd.) and As2O5 (Hengyang Industrial Corporation, Shuikoushan Mining Administration, Hunan, China) were dissolved in ultrapure water in the chosen mole proportion to give various starting Fe(III)/AsO4 and SO4/AsO4 mole ratios (Table 1). The 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.Table 1
Preparation and component of the synthetical basic ferric sulfate-arsenates.
SamplePrecursor solution (100 mL)Solid composition [Fe(SO4)x(AsO4)y(OH)z·nH2O]Fe(III) + SO4 solutionH3AsO4 solutionMole ratioFe2(SO4)3·9H2O (g)H2O (mL)1 M H3AsO4 (mL)H2O (mL)Fe/AsO4SO4/AsO4FeSAsOH-15.62015010402.003.00Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2OFeSAsOH-28.43025010403.004.50Fe(SO4)0.37(AsO4)0.59(OH)0.49·0.28H2OFeSAsOH-38.43025015352.003.00Fe(SO4)0.27(AsO4)0.72(OH)0.30·0.21H2OFeSAsOH-48.43025020301.502.25Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O + FeAsO4·0.75H2O mixtureFeSAsOH-511.24025020302.003.00Fe(SO4)0.24(AsO4)0.70(OH)0.42·0.16H2O


## 2.2. 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 HNO3 solution. The 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. The H2O 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 1.540598 Å (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. The 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.


## 2.3. Dissolution Tests

Five grams of each dried basic ferric sulfate-arsenate was added to 0.1 L of HNO3 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). The 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% HNO3 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 HNO3 or NaOH solutions were supplemented after each sampling. The 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.


## 3. Results and Discussion

### 3.1. Solid Characterization

#### 3.1.1. Solid Composition

The ICP-OES and TGA analysis of the synthetic basic ferric sulfate-arsenates (FeSAsOH) enabled the proportion determination of the main constituents to be described as follows: 41.49 wt% Fe2O3, 11.23 wt% SO3, 43.59 wt% As2O5, and 3.68 wt% H2O for the sample FeSAsOH-1, which corresponded approximately to the values for the chemical compound with the formula of Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O when it was normalized to Fe = 1.00, and the increase in AsO4 in the synthetic solids decreased the requirement for OH− to keep the charge balances; i.e., OH− ions were applied to balance the charges (Table 1). These data were also in good agreement with the empirical formulas of the arsenic-containing compound “Type 2” [Fe4(SO4)y(AsO4)3(OH)x] [13], the precipitate “Phase 3” [Fex(SO4, AsO4) (OH)y·nH2O] [11], the arsenical compound “Ba-5” [Fe(SO4)x(AsO4)1-x(OH)x·(1-x) H2O] [3, 4], and the FeSAsOH [Fe(SO4)0.7−0.2(AsO4)0.2−0.7(OH)0.7−0.2] [9]. Accordingly, the formulas for the samples FeSAsOH-2, FeSAsOH-3, and FeSAsOH-5 were also determined (Table 1). The sample FeSAsOH-4 was a mixture of the FeSAsOH with the formula Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O and the ferric orthoarsenate subhydrate [FeAsO4·0.75H2O] (Table 1).


#### 3.1.2. XRD

The XRD spectra of the basic ferric sulfate-arsenates are shown in Figure1 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.Figure 1
XRD spectra of the synthetic basic ferric sulfate-arsenates before and after dissolution for 330 days.The 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 (dobs, Iobs) of 3.38(100), 3.26(44), 2.64(15), 1.63(11), 2.33(11), and 2.06(10).The 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). The major XRD peaks showed that the sample FeSAsOH-4 was identified as a mixture of the basic ferric arsenate sulfate [13] and ferric orthoarsenate subhydrate [FeAsO4·0.75H2O] (Figure 1, 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 trace-to-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. The acicular crystals of the ferric orthoarsenate subhydrate (FeAsO4·0.75H2O) 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).


#### 3.1.3. FT-IR

The FT-IR spectra for the basic ferric sulfate-arsenates before and after dissolution for 330 d were recorded (Figure2, Figure S2 in Supplementary Materials) and assigned with reference to the literature [2–4, 17, 18]. No significant variations could be viewed in the spectra after dissolution. It could serve as a simple method to identify the functional groups on the precipitate surface. All spectra showed the vibrations of SO42−, AsO43−, OH−, and M-O [19, 20]. No split was observed for the stretching vibrations of the FeO6 octahedra [17].Figure 2
FT-IR spectra of the basic ferric sulfate-arsenate [FeSAsOH-1] before and after dissolution for 330 days.The OH stretching regions from 2800 cm−1 to 3800 cm−1 were characterized by the bands at 3537–3552, 3448–3471, and 3408–3415 cm−1 for the sample FeSAsOH-1. The OH stretching bands at 3408–3415 cm−1 indicated the existence of the strong H-bonds. The higher wavenumber vibration is regularly assigned to the weak H-bonding. The H2O bending (1626–1639 cm−1) vibration is unrecognized in the vibrational structure when the hydration degree is insignificant [2–4]. The 1626–1639 cm−1 bands were readily ascribed to be bending of adsorbed water, i.e., the framework deformation vibrations of H-bonded H2O molecules [15].The free SO42− ions have the same tetrahedral symmetry as the free AsO43− ions, and both show four fundamental vibrations, i.e., v1, v2, v3, and v4 at 983, 450, 1105, and 611 cm−1 for SO42− and at 818, 350, 786, and 405 cm−1 for AsO43−, respectively [2–4]. In the present work, the structural SO42− exhibited the v1, v2, v3, and v4 vibrations at 1032–1034, 434–438, 1130–1134, and 640 cm−1, respectively [19]. The 849–852 and 785–791 cm−1 bands were the v1 and v3 vibrations of AsO43−, respectively [2–4]. The 582–584 cm−1 bands were possibly the OH bending vibrations or the stretching vibrations of Fe–OH2 [2]. The 501–505 cm−1 vibrations were ascribed to Fe-O-As.The FT-IR spectra of FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 were very similar. The FT-IR spectra of FeSAsOH-2 showed a strongerv1(SO4) vibration at 1032–1034 cm−1, a weaker v4(SO4) vibration at 640 cm−1, and a weaker Fe–OH2 stretching vibration at 582–584. The FT-IR spectra of FeSAsOH-4 showed a weaker v3(SO4) vibration at 1128–1136 cm−1 and stronger v2(SO4) vibration at 442–457 cm−1 in comparison to the sample FeSAsOH-1. The weak v1(AsO4) vibration at 804–808 cm−1 and strong v1(AsO4) vibration at 858 cm−1 in the FT-IR spectra of FeSAsOH-4 were considered as the contributions of the ferric orthoarsenate subhydrate and FeSAsOH, respectively (Figure 2, Figure S2 in Supplementary Materials).


#### 3.1.4. 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]. The 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).Figure 3
FE-SEM-EDS results of the basic ferric sulfate-arsenate [FeSAsOH-1] showing the (a) well-shaped monoclinic crystals and the crystals (b) before and (c) after dissolution.
(a)(b)(c)Figure 4
FE-SEM images of the basic ferric sulfate-arsenate [FeSAsOH-1] before and after dissolution for 330 days.Figure 5
FE-SEM-EDS results of (a) basic ferric sulfate-arsenate [FeSAsOH-4, Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O] and (b) ferric orthoarsenate subhydrate [FeAsO4·0.75H2O].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 <5μm (Figure 3). The 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 well-shaped monoclinic prismatic crystals and the rounded type particles (Figure5). The EDS results confirmed that the two shapes were identical, with one grown to its full monoclinic 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 [FeAsO4·0.75H2O] 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 (FigureS3, Supplementary material).


### 3.2. Solution Evolution during Dissolution

#### 3.2.1. FeSAsOH-1

For the dissolution of the sample FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] at 25°C and initial pH 2 (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.Figure 6
Solution evolution during the dissolution of the basic ferric sulfate-arsenate [FeSAsOH-1, Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] at 25–45°C and initial pH 2–10 for 330 days.The released Fe3+ concentrations rose gradually to 0.122–0.123 mmol/L after 5760 h. The final Fe3+ 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).The dissolved SO42− 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. The final SO42− 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.The dissolved AsO43− 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). The final dissolved AsO43− 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.


#### 3.2.2. FeSAsOH-2–FeSAsOH-5

For the dissolution of the FeSAsOH-3 [Fe1.00(SO4)0.27(AsO4)0.72(OH)0.30·0.21H2O] and FeSAsOH-5 [Fe1.00(SO4)0.24(AsO4)0.70(OH)0.42·0.16H2O] samples at different initial pHs and temperatures (Figure S4 in Supplementary Materials), the aqueous pHs and the elemental concentrations of iron, arsenate, and sulfate showed almost the same evolution trends as in the sample FeSAsOH-1 [Fe1.00(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O], because the three samples had almost the same components.For the dissolution of the FeSAsOH-2 [Fe1.00(SO4)0.37(AsO4)0.59(OH)0.49·0.28H2O] sample at different initial pHs and temperatures (Figure S4 in Supplementary Materials), the aqueous pHs and the elemental concentrations of iron, arsenate, and sulfate showed almost the same evolution trends as in the FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 samples, but with higher concentrations of the final dissolved Fe3+ and SO42− and lower concentration of the final dissolved AsO43−. As assumed, it could be related to the fact that the sample FeSAsOH-2 had a lower AsO4/SO4 mole ratio in comparison to the FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 samples.For the dissolution of the sample FeSAsOH-4 [Fe1.00(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O + FeAsO4·0.75H2O mixture] at 25°C and initial pH 2 (Figure S4 in Supplementary Materials), the aqueous pHs and elemental concentrations of iron, sulfate, and arsenate showed almost the similar evolution trends as in the samples FeSAsOH-1, FeSAsOH-2, FeSAsOH-3, and FeSAsOH-5, although the sample FeSAsOH-4 contained trace amount of FeAsO4·0.75H2O.The concentrations of the released arsenate showed the lowest value of 0.006837–0.006851 mmol/L (∼0.51 mg/L As), 0.000604–0.000608 mmol/L (∼0.045 mg/L As), 0.009817–0.010044 mmol/L (∼0.75 mg/L As), 0.005991–0.006026 mmol/L (∼0.45 mg/L As), and 0.001844–0.001853 mmol/L (∼0.139 mg/L As) for the dissolution of FeSAsOH-1–FeSAsOH-5 at different temperatures and pHs, respectively (Figure6, 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]. The “Type 2” with an XAs mole fraction of 0.75 produced a similar As concentration (i.e., 0.8 mg/L) in the TCLP leaching [13].


### 3.3. Dissolution Mechanism

#### 3.3.1. FeSAsOH-1

The dissolution of the FeSAsOH-1 sample [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] is expressed by (1)FeSO40.27AsO40.73OH0.27⋅0.26H2O=1.00Fe3++0.27SO42−+0.73AsO43−+0.27OH−+0.26H2OAccordingly, one mole of the FeSAsOH-1 sample would release 0.27 mol OH−. Thus, this dissolution in stronger acidic solution could deplete H+ ions causing an increase in solution pH, while the dissolution in alkaline medium could deplete OH− ions causing a decrease of solution pH. The complexations would control the speciations of all released constituents (equation (2)–equation (4)).(2)mFe3++nOH−=FemOHn3m−n+,m=3,n=4;m=2,n=2,m=1,n=1−4,(3)H++SO42−=HSO4−(4)nH++AsO43−=HnAsO43−n−,n=1−3,(5)H++OH−=H2OFor the dissolution of the sample FeSAsOH-1 at initial pH 2 (Figure6), the aqueous AsO4/Fe molar ratios attained 1.53–22.64 within 1 h and then declined to the solid stoichiometric AsO4/Fe molar ratio of 0.73 in 12–24 h and finally to 0.49–0.52 (25°C), 0.24–0.28 (35°C), and 0.13–0.14 (45°C) after 5760 h of dissolution. This means that, at the beginning, AsO43− was favorably released from solid into solution with respect to Fe3+.The solution SO4/Fe molar ratios varied between 0.50 and 3.46 and approached continuously to 0.59–0.60 (25°C), 0.89–0.90 (35°C), and 0.91 (45°C) after 5760 h (240 d) of dissolution. The ratio was always higher than 0.27, the stoichiometric SO4/Fe molar ratio of the solid; i.e., SO42− was also preferentially released into solution with respect to Fe3+.The solution AsO4/SO4 molar ratios attained 3.92–8.15 within 1 h and afterward decreased to the solid stoichiometric ratio of 2.70 in 1–3 h and finally to 0.86–0.87 (25°C), 0.30–0.31 (35°C), and 0.14–0.15 (45°C) after 5760 h. AsO43− ions were preferentially dissolved into solution in comparison with SO42− at the beginning of dissolution.All constituents were preferentially dissolved in the order of AsO43− > SO42− > Fe3+ within 1–3 h, in the order of SO42− > AsO43− > Fe3+ from 1–3 h to 12–24 h, and finally in the order of SO42− > Fe3+ > AsO43−, indicating a nonstoichiometric dissolution of the basic ferric sulfate-arsenate FeSAsOH-1 and/or formation of AsO43−/Fe-rich residuals at initial pH 2. Additionally, the higher initial pHs (6 and 10) could obviously inhibit the release of Fe3+ from solid into solution and prefer to form iron-rich residuals, and the solid components were released in the order of SO42− > AsO43− > Fe3+.The 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 Fe3+/Fe(OH)2+/FeSO4+, H3AsO40/H2AsO4−, and HSO4−/SO42−/FeSO4+ for iron, arsenate, and sulfate, respectively. For the dissolution of FeSAsOH-1 at different temperatures and pHs (Figure 6), all solutions were undersaturated with respect to Fe2(SO4)3, FeAsO4·2H2O, 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).The FeSAsOH structure is constructed from layered Fe-O6 octahedra that are cross-linked by AsO4 and SO4 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. The reaction was expressed by the favorable release of AsO43−, followed by SO42− and Fe3+ from solid into solution, whereas Fe3+ was preferentially left behind to form a residual octahedral outlier (Figure 7). The 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 Fe3+ cations were also preferentially removed from the aqueous solution, while sulfate ions were preferentially left in the aqueous solution. The dissolution was finally restricted by the Fe-O(H) breakdown in the FeO6 octahedra and obstructed by the OH− and AsO4 tetrahedra outliers.Figure 7
Dissolution mechanism of the basic ferric sulfate-arsenates in aqueous solution.This 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 (AsO4+SO4)/Fe, AsO4/Fe, and SO4/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. The AsO4/(AsO4+SO4) mole ratios on the solid surface showed no obvious variation after dissolution. In consideration of the crystal structure of FeSO4(OH), the nonstoichiometric dissolution behavior can be explained by the precipitation of a ferric oxyhydroxide phase or by the preferential release of SO4 leaving behind chains of FeO6 octahedra [9].


#### 3.3.2. 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 AsO43− > SO42− > Fe3+ within 1–3 h, in the order of SO42− > AsO43− > Fe3+ from 1 to 3 h to ∼24 h, and finally, in the order of SO42− > Fe3+ > AsO43−, suggesting a nonstoichiometric dissolution and/or formation of AsO43−/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 Fe3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO42− > AsO43− > Fe3+.For the dissolution of FeSAsOH-4, all components were preferentially released in the order of AsO43− > SO42− > Fe3+ in ∼1 h, in the order of SO42− > AsO43− > Fe3+ from ∼1 h to ∼45 d, and finally, in the order of SO42− > Fe3+ > AsO43− after ∼45 d, suggesting a nonstoichiometric dissolution and/or formation of AsO43−/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 Fe3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO42− > AsO43− > Fe3+. However, it was also observed that the aqueous solutions for the FeSAsOH-4 sample showed a slower evolution than those for FeSAsOH-1 and FeSAsOH-3, which was related to the higher H2O mole ratio in FeSAsOH-4.


### 3.4. Solubility Calculation

Based on the batch dissolution results of the synthetic basic ferric sulfate-arsenates, the computing of the aqueous Fe3+, SO42−, AsO43−, 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]. The ion activity products (logˍIAP) for the basic ferric sulfate-arsenates were calculated, which were equal to their solubility products (logˍKsp) at the dissolution equilibrium. The key speciations are summarized in Table S1 in Supplementary Materials.The dissolution of the basic ferric sulfate-arsenates [Fe(SO4)x(AsO4)y(OH)z·nH2O] is described by(6)FeSO4xAsO4yOHz⋅nH2O=1.00Fe3++xSO42−+yAsO43−+zOH−+nH2OThe ion activity product (IAP) is expressed according to(7)IAP=Fe3+1.00SO42−xAsO43−yOH−z,where {} is the aqueous activity of Fe3+, SO42−, AsO43−, or OH−.By using the Gibbs free energies of formation (ΔGfo), ΔGfo[Fe3+] = −4.6 kJ/mol, ΔGfo[SO42−] = −744.6 kJ/mol, ΔGfo[AsO43−] = −647.618 kJ/mol, ΔGfo[OH−] = −157.3 kJ/mol, and ΔGfo[H2O] = −237.18 kJ/mol [24, 25], the ΔGfo values for the synthetic basic ferric sulfate-arsenates [Fe(SO4)x(AsO4)y(OH)z·nH2O] were calculated, suggesting that the IAPs to the end of dissolution (5760–7920 h) were equal to the Ksp values for the basic ferric sulfate-arsenates. At the thermodynamic standard state, the standard free energy of reaction (ΔGro) is estimated according to (8)ΔGro=−5.708log_IAP.For equation (6),(9)ΔGro=1.00ΔGfoFe3++xΔGfoSO42−+yΔGfoAsO43−+zΔGfoOH−+nΔGfoH2O−ΔGfoFeSO4xAsO4yOHz⋅nH2O.Rearranging,(10)ΔGfoFeSO4xAsO4yOHz⋅nH2O=1.00ΔGfoFe3++xΔGfoSO42−+yΔGfoAsO43−+zΔGfoOH−+nΔGfoH2O−ΔGro.The experimental data of the dissolution for 240 d (5760 h), 270 d (6480 h), 7200 h (300 d), and 7920 h (330 d) and the calculated thermodynamic properties for the basic ferric sulfate-arsenates are presented in Tables2–6. The logˍIAP and ΔGfo values were −23.04 ± 0.01 and −914.06 ± 0.03 kJ/mol for the basic ferric sulfate-arsenate FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O], respectively (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. The solubilities decreased slightly with the increasing initial pH and were not obviously affected by the system temperature. The experimental results on dissolution were highly reproducible (Table 2–Table 6).Table 2
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257601.980.1223880.0991700.060330−23.04−23.04−914.05−914.0664801.980.1220300.0968320.059796−23.05±0.01−914.09±0.0372001.980.1233730.0994820.059996−23.04−914.0479201.980.1227460.0991700.059863−23.04−914.0625657603.680.0002520.0690760.007664−25.19−25.22−926.33−926.4764803.680.0002530.0687640.006838−25.23±0.03−926.53±0.1472003.680.0002570.0692320.006851−25.22−926.4979203.680.0002540.0684530.006837−25.22−926.52251057604.030.0002170.0721950.010698−25.22−25.24−926.50−926.5964804.040.0002030.0721950.010925−25.25±0.02−926.64±0.0972004.040.0002070.0726630.010887−25.24−926.5979204.040.0002030.0723510.010961−25.24−926.6325b257601.980.1309830.0790560.068739−22.99−22.97−913.76−913.6464801.990.1322360.0785880.069339−22.96±0.02−913.60±0.1272001.990.1325940.0799910.068538−22.96−913.6179201.990.1336690.0795230.068605−22.96−913.5935257601.940.1522020.1358140.041910−23.06−23.0464801.950.1536340.1381530.042044−23.03±0.0272001.950.1538130.1365930.041510−23.0479201.950.1535450.1367490.041910−23.0445257601.910.2674270.2444960.036772−22.86−22.8664801.920.2723510.2466790.035837−22.84±0.0272001.910.2708290.2473020.035103−22.8779201.910.2711870.2462110.035771−22.86a,bDissolution tests in duplicate.Table 3
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-2 [Fe(SO4)0.37(AsO4)0.59(OH)0.49·0.28H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257602.081.4665071.2337050.077881−22.03−22.01−931.40−931.3164802.091.4691931.2220110.078749−22.00±0.02−931.28±0.0972002.091.4684761.2220110.078682−22.00−931.2879202.091.4695511.2204520.078349−22.00−931.2925657603.000.0659840.6344730.000728−23.46−23.45−939.59−939.5664803.000.0663420.6421130.000745−23.45±0.01−939.54±0.0372003.000.0664320.6410220.000730−23.46−939.5779203.000.0663420.6464790.000735−23.45−939.55251057603.050.0528230.6011040.000611−23.59−23.59−940.35−940.3364803.040.0539870.6045340.000604−23.59±0.00−940.33±0.0272003.040.0540760.6021950.000606−23.59−940.3279203.040.0540410.6015720.000608−23.59−940.3225b257602.061.4638211.2449320.073477−22.08−22.06−931.70−931.6264802.071.4620301.2351090.074745−22.06±0.02−931.58±0.0872002.071.4616721.2263770.074011−22.06−931.5979202.071.4629251.2288720.073677−22.06−931.6035257602.021.4208461.4039790.060930−22.11−22.1064802.021.4235321.4050710.060797−22.11±0.0172002.031.4262181.3964950.060130−22.0979202.021.4251441.4011730.060730−22.1145257601.971.2928181.5982660.057994−22.17−22.1964801.961.2838651.6263330.058461−22.19±0.0272001.961.2775981.6158860.058261−22.1979201.961.2829701.6116760.058394−22.19a,b Dissolution tests in duplicate.Table 4
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-3 [Fe(SO4)0.27(AsO4)0.72(OH)0.30·0.21H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257601.970.1443230.1033810.059262−23.16−23.14−901.14−901.0264801.980.1454870.1040040.059462−23.14±0.02−900.99±0.1272001.980.1457550.1041600.059329−23.14−900.9979201.980.1456660.1044720.059596−23.14−900.9825657603.910.0004200.0687640.010056−25.07−25.08−912.05−912.0864803.900.0004170.0675170.009817−25.08±0.01−912.11±0.0372003.900.0004140.0670490.010044−25.08−912.0979203.900.0004160.0678290.010053−25.08−912.07251057604.220.0003750.0831100.012148−25.12−25.13−912.28−912.3564804.210.0003670.0840450.011985−25.13±0.01−912.34±0.0772004.210.0003630.0848250.011878−25.13−912.3879204.210.0003580.0852930.011962−25.14−912.4025b257601.970.1432490.1118010.057060−23.17−23.18−901.19−901.2164801.970.1457550.1114890.055992−23.17±0.01−901.19±0.1172001.960.1441440.1101630.056659−23.19−901.3279201.970.1453080.1113330.056793−23.17−901.1635257601.950.2130820.2013040.041310−23.04−23.0564801.950.2163950.2036420.039374−23.05±0.0172001.950.2172900.2014590.039641−23.0479201.950.2176480.2017710.039308−23.0545257601.930.3144300.2850370.035971−22.90−22.8964801.930.3229360.2870640.036238−22.89±0.0172001.930.3228460.2865960.036572−22.8979201.930.3229360.2881560.036638−22.89a,bDissolution tests in duplicate.Table 5
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-4 [Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O and FeAsO4·0.75H2O mixture].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2OFeAsO4·0.75H2OpHFeSO4AsO4log_IAPMean log_IAPlog_IAPMean log_IAP25a257601.960.1318780.2823860.054991−22.66−22.64−23.89−23.8764801.970.1334900.2716270.054323−22.64±0.02−23.86±0.0272001.970.1339370.2766170.054857−22.63−23.8679201.970.1338480.2753700.054257−22.63−23.8625657603.200.0005660.3407040.005979−24.39−24.40−25.53−25.5464803.200.0005480.3378970.005991−24.40±0.01−25.54±0.0172003.200.0005360.3388320.006026−24.41−25.5579203.200.0005390.3375850.006021−24.41−25.55251057603.200.0004680.3173140.006248−24.47−24.48−25.59−25.6164803.200.0004570.3138840.006032−24.49±0.01−25.62±0.0272003.200.0004610.3134160.006066−24.48−25.6179203.200.0004580.3123250.006059−24.49−25.6125b257601.950.1329530.2819190.055792−22.67−22.66−23.91−23.9064801.950.1394880.2837900.055591−22.65±0.01−23.89±0.0172001.940.1411000.2800470.055925−22.66−23.9079201.940.1400250.2833220.055391−22.67−23.9135257601.940.1952660.3895090.031099−22.66−22.66−24.03−24.0364801.950.2012640.3838960.030899−22.63±0.03−24.00±0.0372001.930.1997420.3821810.030565−22.68−24.0579201.930.2003690.3824920.030565−22.68−24.0545257601.890.2751270.4849370.024692−22.67−22.64−24.12−24.0964801.900.2889140.4872760.024359−22.63±0.03−24.09±0.0372001.900.2915110.4846250.024359−22.63−24.0879201.900.2899890.4886800.024626−22.62−24.08a,b Dissolution tests in duplicate.Table 6
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-5 [Fe(SO4)0.24(AsO4)0.70(OH)0.42·0.16H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257601.960.1566780.1203770.028163−24.29−24.27−879.30−879.2064801.970.1577520.1217800.028096−24.27±0.02−879.16±0.1072001.970.1578420.1220920.027896−24.27−879.1779201.970.1574840.1236510.028163−24.27−879.1625657603.900.0009130.0796790.001860−26.01−26.01−889.09−889.1264803.890.0009020.0810830.001853−26.01±0.00−889.11±0.0372003.890.0008940.0810830.001844−26.01−889.1479203.890.0008970.0812390.001849−26.01−889.12251057604.160.0005840.0935570.002324−26.16−26.17−889.95−890.0564804.150.0005730.0907500.002227−26.18±0.01−890.08±0.1072004.150.0005700.0905940.002232−26.18−890.0979204.150.0005720.0913740.002228−26.18−890.0825b257601.960.1600800.1252110.027562−24.29−24.26−879.28−879.1564801.970.1633930.1272380.027095−24.26±0.03−879.13±0.1372001.970.1637510.1258340.027762−24.25−879.0879201.970.1634820.1270820.027562−24.25−879.1035257601.940.2402990.2087880.006020−24.49−24.4864801.940.2479090.2129980.005877−24.48±0.0172001.940.2469250.2083200.005908−24.4879201.940.2470140.2128420.005897−24.4845257601.880.3731620.3428870.004635−24.42−24.4064801.890.3737000.3450700.004664−24.39±0.0272001.890.3735210.3442900.004661−24.3979201.890.3741470.3458490.004678−24.39a,bDissolution tests in duplicate.Using the result of the solubility experiments at initial pH 2 for 240 d (5760 h), 270 d (6480 h), 7200 h (300 d), and 7920 h (330 d), the logˍIAP for FeAsO4·0.75H2O is determined after the following reaction:(11)FeAsO4⋅0.75H2O=Fe3++AsO43−+0.75H2OThe calculated logˍIAP of −23.87 ± 0.02 (Table5) is in good agreement with the logˍKsp of −26.54 ± 0.35 [26].


## 3.1. Solid Characterization

### 3.1.1. Solid Composition

The ICP-OES and TGA analysis of the synthetic basic ferric sulfate-arsenates (FeSAsOH) enabled the proportion determination of the main constituents to be described as follows: 41.49 wt% Fe2O3, 11.23 wt% SO3, 43.59 wt% As2O5, and 3.68 wt% H2O for the sample FeSAsOH-1, which corresponded approximately to the values for the chemical compound with the formula of Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O when it was normalized to Fe = 1.00, and the increase in AsO4 in the synthetic solids decreased the requirement for OH− to keep the charge balances; i.e., OH− ions were applied to balance the charges (Table 1). These data were also in good agreement with the empirical formulas of the arsenic-containing compound “Type 2” [Fe4(SO4)y(AsO4)3(OH)x] [13], the precipitate “Phase 3” [Fex(SO4, AsO4) (OH)y·nH2O] [11], the arsenical compound “Ba-5” [Fe(SO4)x(AsO4)1-x(OH)x·(1-x) H2O] [3, 4], and the FeSAsOH [Fe(SO4)0.7−0.2(AsO4)0.2−0.7(OH)0.7−0.2] [9]. Accordingly, the formulas for the samples FeSAsOH-2, FeSAsOH-3, and FeSAsOH-5 were also determined (Table 1). The sample FeSAsOH-4 was a mixture of the FeSAsOH with the formula Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O and the ferric orthoarsenate subhydrate [FeAsO4·0.75H2O] (Table 1).


### 3.1.2. XRD

The XRD spectra of the basic ferric sulfate-arsenates are shown in Figure1 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.Figure 1
XRD spectra of the synthetic basic ferric sulfate-arsenates before and after dissolution for 330 days.The 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 (dobs, Iobs) of 3.38(100), 3.26(44), 2.64(15), 1.63(11), 2.33(11), and 2.06(10).The 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). The major XRD peaks showed that the sample FeSAsOH-4 was identified as a mixture of the basic ferric arsenate sulfate [13] and ferric orthoarsenate subhydrate [FeAsO4·0.75H2O] (Figure 1, 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 trace-to-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. The acicular crystals of the ferric orthoarsenate subhydrate (FeAsO4·0.75H2O) 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).


### 3.1.3. FT-IR

The FT-IR spectra for the basic ferric sulfate-arsenates before and after dissolution for 330 d were recorded (Figure2, Figure S2 in Supplementary Materials) and assigned with reference to the literature [2–4, 17, 18]. No significant variations could be viewed in the spectra after dissolution. It could serve as a simple method to identify the functional groups on the precipitate surface. All spectra showed the vibrations of SO42−, AsO43−, OH−, and M-O [19, 20]. No split was observed for the stretching vibrations of the FeO6 octahedra [17].Figure 2
FT-IR spectra of the basic ferric sulfate-arsenate [FeSAsOH-1] before and after dissolution for 330 days.The OH stretching regions from 2800 cm−1 to 3800 cm−1 were characterized by the bands at 3537–3552, 3448–3471, and 3408–3415 cm−1 for the sample FeSAsOH-1. The OH stretching bands at 3408–3415 cm−1 indicated the existence of the strong H-bonds. The higher wavenumber vibration is regularly assigned to the weak H-bonding. The H2O bending (1626–1639 cm−1) vibration is unrecognized in the vibrational structure when the hydration degree is insignificant [2–4]. The 1626–1639 cm−1 bands were readily ascribed to be bending of adsorbed water, i.e., the framework deformation vibrations of H-bonded H2O molecules [15].The free SO42− ions have the same tetrahedral symmetry as the free AsO43− ions, and both show four fundamental vibrations, i.e., v1, v2, v3, and v4 at 983, 450, 1105, and 611 cm−1 for SO42− and at 818, 350, 786, and 405 cm−1 for AsO43−, respectively [2–4]. In the present work, the structural SO42− exhibited the v1, v2, v3, and v4 vibrations at 1032–1034, 434–438, 1130–1134, and 640 cm−1, respectively [19]. The 849–852 and 785–791 cm−1 bands were the v1 and v3 vibrations of AsO43−, respectively [2–4]. The 582–584 cm−1 bands were possibly the OH bending vibrations or the stretching vibrations of Fe–OH2 [2]. The 501–505 cm−1 vibrations were ascribed to Fe-O-As.The FT-IR spectra of FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 were very similar. The FT-IR spectra of FeSAsOH-2 showed a strongerv1(SO4) vibration at 1032–1034 cm−1, a weaker v4(SO4) vibration at 640 cm−1, and a weaker Fe–OH2 stretching vibration at 582–584. The FT-IR spectra of FeSAsOH-4 showed a weaker v3(SO4) vibration at 1128–1136 cm−1 and stronger v2(SO4) vibration at 442–457 cm−1 in comparison to the sample FeSAsOH-1. The weak v1(AsO4) vibration at 804–808 cm−1 and strong v1(AsO4) vibration at 858 cm−1 in the FT-IR spectra of FeSAsOH-4 were considered as the contributions of the ferric orthoarsenate subhydrate and FeSAsOH, respectively (Figure 2, Figure S2 in Supplementary Materials).


### 3.1.4. 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]. The 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).Figure 3
FE-SEM-EDS results of the basic ferric sulfate-arsenate [FeSAsOH-1] showing the (a) well-shaped monoclinic crystals and the crystals (b) before and (c) after dissolution.
(a)(b)(c)Figure 4
FE-SEM images of the basic ferric sulfate-arsenate [FeSAsOH-1] before and after dissolution for 330 days.Figure 5
FE-SEM-EDS results of (a) basic ferric sulfate-arsenate [FeSAsOH-4, Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O] and (b) ferric orthoarsenate subhydrate [FeAsO4·0.75H2O].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 <5μm (Figure 3). The 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 well-shaped monoclinic prismatic crystals and the rounded type particles (Figure5). The EDS results confirmed that the two shapes were identical, with one grown to its full monoclinic 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 [FeAsO4·0.75H2O] 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 (FigureS3, Supplementary material).


## 3.1.1. Solid Composition

The ICP-OES and TGA analysis of the synthetic basic ferric sulfate-arsenates (FeSAsOH) enabled the proportion determination of the main constituents to be described as follows: 41.49 wt% Fe2O3, 11.23 wt% SO3, 43.59 wt% As2O5, and 3.68 wt% H2O for the sample FeSAsOH-1, which corresponded approximately to the values for the chemical compound with the formula of Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O when it was normalized to Fe = 1.00, and the increase in AsO4 in the synthetic solids decreased the requirement for OH− to keep the charge balances; i.e., OH− ions were applied to balance the charges (Table 1). These data were also in good agreement with the empirical formulas of the arsenic-containing compound “Type 2” [Fe4(SO4)y(AsO4)3(OH)x] [13], the precipitate “Phase 3” [Fex(SO4, AsO4) (OH)y·nH2O] [11], the arsenical compound “Ba-5” [Fe(SO4)x(AsO4)1-x(OH)x·(1-x) H2O] [3, 4], and the FeSAsOH [Fe(SO4)0.7−0.2(AsO4)0.2−0.7(OH)0.7−0.2] [9]. Accordingly, the formulas for the samples FeSAsOH-2, FeSAsOH-3, and FeSAsOH-5 were also determined (Table 1). The sample FeSAsOH-4 was a mixture of the FeSAsOH with the formula Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O and the ferric orthoarsenate subhydrate [FeAsO4·0.75H2O] (Table 1).


## 3.1.2. XRD

The XRD spectra of the basic ferric sulfate-arsenates are shown in Figure1 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.Figure 1
XRD spectra of the synthetic basic ferric sulfate-arsenates before and after dissolution for 330 days.The 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 (dobs, Iobs) of 3.38(100), 3.26(44), 2.64(15), 1.63(11), 2.33(11), and 2.06(10).The 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). The major XRD peaks showed that the sample FeSAsOH-4 was identified as a mixture of the basic ferric arsenate sulfate [13] and ferric orthoarsenate subhydrate [FeAsO4·0.75H2O] (Figure 1, 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 trace-to-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. The acicular crystals of the ferric orthoarsenate subhydrate (FeAsO4·0.75H2O) 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).


## 3.1.3. FT-IR

The FT-IR spectra for the basic ferric sulfate-arsenates before and after dissolution for 330 d were recorded (Figure2, Figure S2 in Supplementary Materials) and assigned with reference to the literature [2–4, 17, 18]. No significant variations could be viewed in the spectra after dissolution. It could serve as a simple method to identify the functional groups on the precipitate surface. All spectra showed the vibrations of SO42−, AsO43−, OH−, and M-O [19, 20]. No split was observed for the stretching vibrations of the FeO6 octahedra [17].Figure 2
FT-IR spectra of the basic ferric sulfate-arsenate [FeSAsOH-1] before and after dissolution for 330 days.The OH stretching regions from 2800 cm−1 to 3800 cm−1 were characterized by the bands at 3537–3552, 3448–3471, and 3408–3415 cm−1 for the sample FeSAsOH-1. The OH stretching bands at 3408–3415 cm−1 indicated the existence of the strong H-bonds. The higher wavenumber vibration is regularly assigned to the weak H-bonding. The H2O bending (1626–1639 cm−1) vibration is unrecognized in the vibrational structure when the hydration degree is insignificant [2–4]. The 1626–1639 cm−1 bands were readily ascribed to be bending of adsorbed water, i.e., the framework deformation vibrations of H-bonded H2O molecules [15].The free SO42− ions have the same tetrahedral symmetry as the free AsO43− ions, and both show four fundamental vibrations, i.e., v1, v2, v3, and v4 at 983, 450, 1105, and 611 cm−1 for SO42− and at 818, 350, 786, and 405 cm−1 for AsO43−, respectively [2–4]. In the present work, the structural SO42− exhibited the v1, v2, v3, and v4 vibrations at 1032–1034, 434–438, 1130–1134, and 640 cm−1, respectively [19]. The 849–852 and 785–791 cm−1 bands were the v1 and v3 vibrations of AsO43−, respectively [2–4]. The 582–584 cm−1 bands were possibly the OH bending vibrations or the stretching vibrations of Fe–OH2 [2]. The 501–505 cm−1 vibrations were ascribed to Fe-O-As.The FT-IR spectra of FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 were very similar. The FT-IR spectra of FeSAsOH-2 showed a strongerv1(SO4) vibration at 1032–1034 cm−1, a weaker v4(SO4) vibration at 640 cm−1, and a weaker Fe–OH2 stretching vibration at 582–584. The FT-IR spectra of FeSAsOH-4 showed a weaker v3(SO4) vibration at 1128–1136 cm−1 and stronger v2(SO4) vibration at 442–457 cm−1 in comparison to the sample FeSAsOH-1. The weak v1(AsO4) vibration at 804–808 cm−1 and strong v1(AsO4) vibration at 858 cm−1 in the FT-IR spectra of FeSAsOH-4 were considered as the contributions of the ferric orthoarsenate subhydrate and FeSAsOH, respectively (Figure 2, Figure S2 in Supplementary Materials).


## 3.1.4. 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]. The 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).Figure 3
FE-SEM-EDS results of the basic ferric sulfate-arsenate [FeSAsOH-1] showing the (a) well-shaped monoclinic crystals and the crystals (b) before and (c) after dissolution.
(a)(b)(c)Figure 4
FE-SEM images of the basic ferric sulfate-arsenate [FeSAsOH-1] before and after dissolution for 330 days.Figure 5
FE-SEM-EDS results of (a) basic ferric sulfate-arsenate [FeSAsOH-4, Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O] and (b) ferric orthoarsenate subhydrate [FeAsO4·0.75H2O].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 <5μm (Figure 3). The 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 well-shaped monoclinic prismatic crystals and the rounded type particles (Figure5). The EDS results confirmed that the two shapes were identical, with one grown to its full monoclinic 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 [FeAsO4·0.75H2O] 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 (FigureS3, Supplementary material).


## 3.2. Solution Evolution during Dissolution

### 3.2.1. FeSAsOH-1

For the dissolution of the sample FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] at 25°C and initial pH 2 (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.Figure 6
Solution evolution during the dissolution of the basic ferric sulfate-arsenate [FeSAsOH-1, Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] at 25–45°C and initial pH 2–10 for 330 days.The released Fe3+ concentrations rose gradually to 0.122–0.123 mmol/L after 5760 h. The final Fe3+ 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).The dissolved SO42− 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. The final SO42− 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.The dissolved AsO43− 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). The final dissolved AsO43− 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.


### 3.2.2. FeSAsOH-2–FeSAsOH-5

For the dissolution of the FeSAsOH-3 [Fe1.00(SO4)0.27(AsO4)0.72(OH)0.30·0.21H2O] and FeSAsOH-5 [Fe1.00(SO4)0.24(AsO4)0.70(OH)0.42·0.16H2O] samples at different initial pHs and temperatures (Figure S4 in Supplementary Materials), the aqueous pHs and the elemental concentrations of iron, arsenate, and sulfate showed almost the same evolution trends as in the sample FeSAsOH-1 [Fe1.00(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O], because the three samples had almost the same components.For the dissolution of the FeSAsOH-2 [Fe1.00(SO4)0.37(AsO4)0.59(OH)0.49·0.28H2O] sample at different initial pHs and temperatures (Figure S4 in Supplementary Materials), the aqueous pHs and the elemental concentrations of iron, arsenate, and sulfate showed almost the same evolution trends as in the FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 samples, but with higher concentrations of the final dissolved Fe3+ and SO42− and lower concentration of the final dissolved AsO43−. As assumed, it could be related to the fact that the sample FeSAsOH-2 had a lower AsO4/SO4 mole ratio in comparison to the FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 samples.For the dissolution of the sample FeSAsOH-4 [Fe1.00(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O + FeAsO4·0.75H2O mixture] at 25°C and initial pH 2 (Figure S4 in Supplementary Materials), the aqueous pHs and elemental concentrations of iron, sulfate, and arsenate showed almost the similar evolution trends as in the samples FeSAsOH-1, FeSAsOH-2, FeSAsOH-3, and FeSAsOH-5, although the sample FeSAsOH-4 contained trace amount of FeAsO4·0.75H2O.The concentrations of the released arsenate showed the lowest value of 0.006837–0.006851 mmol/L (∼0.51 mg/L As), 0.000604–0.000608 mmol/L (∼0.045 mg/L As), 0.009817–0.010044 mmol/L (∼0.75 mg/L As), 0.005991–0.006026 mmol/L (∼0.45 mg/L As), and 0.001844–0.001853 mmol/L (∼0.139 mg/L As) for the dissolution of FeSAsOH-1–FeSAsOH-5 at different temperatures and pHs, respectively (Figure6, 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]. The “Type 2” with an XAs mole fraction of 0.75 produced a similar As concentration (i.e., 0.8 mg/L) in the TCLP leaching [13].


## 3.2.1. FeSAsOH-1

For the dissolution of the sample FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] at 25°C and initial pH 2 (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.Figure 6
Solution evolution during the dissolution of the basic ferric sulfate-arsenate [FeSAsOH-1, Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] at 25–45°C and initial pH 2–10 for 330 days.The released Fe3+ concentrations rose gradually to 0.122–0.123 mmol/L after 5760 h. The final Fe3+ 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).The dissolved SO42− 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. The final SO42− 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.The dissolved AsO43− 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). The final dissolved AsO43− 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.


## 3.2.2. FeSAsOH-2–FeSAsOH-5

For the dissolution of the FeSAsOH-3 [Fe1.00(SO4)0.27(AsO4)0.72(OH)0.30·0.21H2O] and FeSAsOH-5 [Fe1.00(SO4)0.24(AsO4)0.70(OH)0.42·0.16H2O] samples at different initial pHs and temperatures (Figure S4 in Supplementary Materials), the aqueous pHs and the elemental concentrations of iron, arsenate, and sulfate showed almost the same evolution trends as in the sample FeSAsOH-1 [Fe1.00(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O], because the three samples had almost the same components.For the dissolution of the FeSAsOH-2 [Fe1.00(SO4)0.37(AsO4)0.59(OH)0.49·0.28H2O] sample at different initial pHs and temperatures (Figure S4 in Supplementary Materials), the aqueous pHs and the elemental concentrations of iron, arsenate, and sulfate showed almost the same evolution trends as in the FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 samples, but with higher concentrations of the final dissolved Fe3+ and SO42− and lower concentration of the final dissolved AsO43−. As assumed, it could be related to the fact that the sample FeSAsOH-2 had a lower AsO4/SO4 mole ratio in comparison to the FeSAsOH-1, FeSAsOH-3, and FeSAsOH-5 samples.For the dissolution of the sample FeSAsOH-4 [Fe1.00(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O + FeAsO4·0.75H2O mixture] at 25°C and initial pH 2 (Figure S4 in Supplementary Materials), the aqueous pHs and elemental concentrations of iron, sulfate, and arsenate showed almost the similar evolution trends as in the samples FeSAsOH-1, FeSAsOH-2, FeSAsOH-3, and FeSAsOH-5, although the sample FeSAsOH-4 contained trace amount of FeAsO4·0.75H2O.The concentrations of the released arsenate showed the lowest value of 0.006837–0.006851 mmol/L (∼0.51 mg/L As), 0.000604–0.000608 mmol/L (∼0.045 mg/L As), 0.009817–0.010044 mmol/L (∼0.75 mg/L As), 0.005991–0.006026 mmol/L (∼0.45 mg/L As), and 0.001844–0.001853 mmol/L (∼0.139 mg/L As) for the dissolution of FeSAsOH-1–FeSAsOH-5 at different temperatures and pHs, respectively (Figure6, 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]. The “Type 2” with an XAs mole fraction of 0.75 produced a similar As concentration (i.e., 0.8 mg/L) in the TCLP leaching [13].


## 3.3. Dissolution Mechanism

### 3.3.1. FeSAsOH-1

The dissolution of the FeSAsOH-1 sample [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] is expressed by (1)FeSO40.27AsO40.73OH0.27⋅0.26H2O=1.00Fe3++0.27SO42−+0.73AsO43−+0.27OH−+0.26H2OAccordingly, one mole of the FeSAsOH-1 sample would release 0.27 mol OH−. Thus, this dissolution in stronger acidic solution could deplete H+ ions causing an increase in solution pH, while the dissolution in alkaline medium could deplete OH− ions causing a decrease of solution pH. The complexations would control the speciations of all released constituents (equation (2)–equation (4)).(2)mFe3++nOH−=FemOHn3m−n+,m=3,n=4;m=2,n=2,m=1,n=1−4,(3)H++SO42−=HSO4−(4)nH++AsO43−=HnAsO43−n−,n=1−3,(5)H++OH−=H2OFor the dissolution of the sample FeSAsOH-1 at initial pH 2 (Figure6), the aqueous AsO4/Fe molar ratios attained 1.53–22.64 within 1 h and then declined to the solid stoichiometric AsO4/Fe molar ratio of 0.73 in 12–24 h and finally to 0.49–0.52 (25°C), 0.24–0.28 (35°C), and 0.13–0.14 (45°C) after 5760 h of dissolution. This means that, at the beginning, AsO43− was favorably released from solid into solution with respect to Fe3+.The solution SO4/Fe molar ratios varied between 0.50 and 3.46 and approached continuously to 0.59–0.60 (25°C), 0.89–0.90 (35°C), and 0.91 (45°C) after 5760 h (240 d) of dissolution. The ratio was always higher than 0.27, the stoichiometric SO4/Fe molar ratio of the solid; i.e., SO42− was also preferentially released into solution with respect to Fe3+.The solution AsO4/SO4 molar ratios attained 3.92–8.15 within 1 h and afterward decreased to the solid stoichiometric ratio of 2.70 in 1–3 h and finally to 0.86–0.87 (25°C), 0.30–0.31 (35°C), and 0.14–0.15 (45°C) after 5760 h. AsO43− ions were preferentially dissolved into solution in comparison with SO42− at the beginning of dissolution.All constituents were preferentially dissolved in the order of AsO43− > SO42− > Fe3+ within 1–3 h, in the order of SO42− > AsO43− > Fe3+ from 1–3 h to 12–24 h, and finally in the order of SO42− > Fe3+ > AsO43−, indicating a nonstoichiometric dissolution of the basic ferric sulfate-arsenate FeSAsOH-1 and/or formation of AsO43−/Fe-rich residuals at initial pH 2. Additionally, the higher initial pHs (6 and 10) could obviously inhibit the release of Fe3+ from solid into solution and prefer to form iron-rich residuals, and the solid components were released in the order of SO42− > AsO43− > Fe3+.The 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 Fe3+/Fe(OH)2+/FeSO4+, H3AsO40/H2AsO4−, and HSO4−/SO42−/FeSO4+ for iron, arsenate, and sulfate, respectively. For the dissolution of FeSAsOH-1 at different temperatures and pHs (Figure 6), all solutions were undersaturated with respect to Fe2(SO4)3, FeAsO4·2H2O, 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).The FeSAsOH structure is constructed from layered Fe-O6 octahedra that are cross-linked by AsO4 and SO4 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. The reaction was expressed by the favorable release of AsO43−, followed by SO42− and Fe3+ from solid into solution, whereas Fe3+ was preferentially left behind to form a residual octahedral outlier (Figure 7). The 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 Fe3+ cations were also preferentially removed from the aqueous solution, while sulfate ions were preferentially left in the aqueous solution. The dissolution was finally restricted by the Fe-O(H) breakdown in the FeO6 octahedra and obstructed by the OH− and AsO4 tetrahedra outliers.Figure 7
Dissolution mechanism of the basic ferric sulfate-arsenates in aqueous solution.This 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 (AsO4+SO4)/Fe, AsO4/Fe, and SO4/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. The AsO4/(AsO4+SO4) mole ratios on the solid surface showed no obvious variation after dissolution. In consideration of the crystal structure of FeSO4(OH), the nonstoichiometric dissolution behavior can be explained by the precipitation of a ferric oxyhydroxide phase or by the preferential release of SO4 leaving behind chains of FeO6 octahedra [9].


### 3.3.2. 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 AsO43− > SO42− > Fe3+ within 1–3 h, in the order of SO42− > AsO43− > Fe3+ from 1 to 3 h to ∼24 h, and finally, in the order of SO42− > Fe3+ > AsO43−, suggesting a nonstoichiometric dissolution and/or formation of AsO43−/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 Fe3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO42− > AsO43− > Fe3+.For the dissolution of FeSAsOH-4, all components were preferentially released in the order of AsO43− > SO42− > Fe3+ in ∼1 h, in the order of SO42− > AsO43− > Fe3+ from ∼1 h to ∼45 d, and finally, in the order of SO42− > Fe3+ > AsO43− after ∼45 d, suggesting a nonstoichiometric dissolution and/or formation of AsO43−/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 Fe3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO42− > AsO43− > Fe3+. However, it was also observed that the aqueous solutions for the FeSAsOH-4 sample showed a slower evolution than those for FeSAsOH-1 and FeSAsOH-3, which was related to the higher H2O mole ratio in FeSAsOH-4.


## 3.3.1. FeSAsOH-1

The dissolution of the FeSAsOH-1 sample [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] is expressed by (1)FeSO40.27AsO40.73OH0.27⋅0.26H2O=1.00Fe3++0.27SO42−+0.73AsO43−+0.27OH−+0.26H2OAccordingly, one mole of the FeSAsOH-1 sample would release 0.27 mol OH−. Thus, this dissolution in stronger acidic solution could deplete H+ ions causing an increase in solution pH, while the dissolution in alkaline medium could deplete OH− ions causing a decrease of solution pH. The complexations would control the speciations of all released constituents (equation (2)–equation (4)).(2)mFe3++nOH−=FemOHn3m−n+,m=3,n=4;m=2,n=2,m=1,n=1−4,(3)H++SO42−=HSO4−(4)nH++AsO43−=HnAsO43−n−,n=1−3,(5)H++OH−=H2OFor the dissolution of the sample FeSAsOH-1 at initial pH 2 (Figure6), the aqueous AsO4/Fe molar ratios attained 1.53–22.64 within 1 h and then declined to the solid stoichiometric AsO4/Fe molar ratio of 0.73 in 12–24 h and finally to 0.49–0.52 (25°C), 0.24–0.28 (35°C), and 0.13–0.14 (45°C) after 5760 h of dissolution. This means that, at the beginning, AsO43− was favorably released from solid into solution with respect to Fe3+.The solution SO4/Fe molar ratios varied between 0.50 and 3.46 and approached continuously to 0.59–0.60 (25°C), 0.89–0.90 (35°C), and 0.91 (45°C) after 5760 h (240 d) of dissolution. The ratio was always higher than 0.27, the stoichiometric SO4/Fe molar ratio of the solid; i.e., SO42− was also preferentially released into solution with respect to Fe3+.The solution AsO4/SO4 molar ratios attained 3.92–8.15 within 1 h and afterward decreased to the solid stoichiometric ratio of 2.70 in 1–3 h and finally to 0.86–0.87 (25°C), 0.30–0.31 (35°C), and 0.14–0.15 (45°C) after 5760 h. AsO43− ions were preferentially dissolved into solution in comparison with SO42− at the beginning of dissolution.All constituents were preferentially dissolved in the order of AsO43− > SO42− > Fe3+ within 1–3 h, in the order of SO42− > AsO43− > Fe3+ from 1–3 h to 12–24 h, and finally in the order of SO42− > Fe3+ > AsO43−, indicating a nonstoichiometric dissolution of the basic ferric sulfate-arsenate FeSAsOH-1 and/or formation of AsO43−/Fe-rich residuals at initial pH 2. Additionally, the higher initial pHs (6 and 10) could obviously inhibit the release of Fe3+ from solid into solution and prefer to form iron-rich residuals, and the solid components were released in the order of SO42− > AsO43− > Fe3+.The 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 Fe3+/Fe(OH)2+/FeSO4+, H3AsO40/H2AsO4−, and HSO4−/SO42−/FeSO4+ for iron, arsenate, and sulfate, respectively. For the dissolution of FeSAsOH-1 at different temperatures and pHs (Figure 6), all solutions were undersaturated with respect to Fe2(SO4)3, FeAsO4·2H2O, 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).The FeSAsOH structure is constructed from layered Fe-O6 octahedra that are cross-linked by AsO4 and SO4 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. The reaction was expressed by the favorable release of AsO43−, followed by SO42− and Fe3+ from solid into solution, whereas Fe3+ was preferentially left behind to form a residual octahedral outlier (Figure 7). The 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 Fe3+ cations were also preferentially removed from the aqueous solution, while sulfate ions were preferentially left in the aqueous solution. The dissolution was finally restricted by the Fe-O(H) breakdown in the FeO6 octahedra and obstructed by the OH− and AsO4 tetrahedra outliers.Figure 7
Dissolution mechanism of the basic ferric sulfate-arsenates in aqueous solution.This 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 (AsO4+SO4)/Fe, AsO4/Fe, and SO4/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. The AsO4/(AsO4+SO4) mole ratios on the solid surface showed no obvious variation after dissolution. In consideration of the crystal structure of FeSO4(OH), the nonstoichiometric dissolution behavior can be explained by the precipitation of a ferric oxyhydroxide phase or by the preferential release of SO4 leaving behind chains of FeO6 octahedra [9].


## 3.3.2. 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 AsO43− > SO42− > Fe3+ within 1–3 h, in the order of SO42− > AsO43− > Fe3+ from 1 to 3 h to ∼24 h, and finally, in the order of SO42− > Fe3+ > AsO43−, suggesting a nonstoichiometric dissolution and/or formation of AsO43−/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 Fe3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO42− > AsO43− > Fe3+.For the dissolution of FeSAsOH-4, all components were preferentially released in the order of AsO43− > SO42− > Fe3+ in ∼1 h, in the order of SO42− > AsO43− > Fe3+ from ∼1 h to ∼45 d, and finally, in the order of SO42− > Fe3+ > AsO43− after ∼45 d, suggesting a nonstoichiometric dissolution and/or formation of AsO43−/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 Fe3+ from solid into solution and prefer to form iron-rich residual, and the solid components were released in the order of SO42− > AsO43− > Fe3+. However, it was also observed that the aqueous solutions for the FeSAsOH-4 sample showed a slower evolution than those for FeSAsOH-1 and FeSAsOH-3, which was related to the higher H2O mole ratio in FeSAsOH-4.


## 3.4. Solubility Calculation

Based on the batch dissolution results of the synthetic basic ferric sulfate-arsenates, the computing of the aqueous Fe3+, SO42−, AsO43−, 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]. The ion activity products (logˍIAP) for the basic ferric sulfate-arsenates were calculated, which were equal to their solubility products (logˍKsp) at the dissolution equilibrium. The key speciations are summarized in Table S1 in Supplementary Materials.The dissolution of the basic ferric sulfate-arsenates [Fe(SO4)x(AsO4)y(OH)z·nH2O] is described by(6)FeSO4xAsO4yOHz⋅nH2O=1.00Fe3++xSO42−+yAsO43−+zOH−+nH2OThe ion activity product (IAP) is expressed according to(7)IAP=Fe3+1.00SO42−xAsO43−yOH−z,where {} is the aqueous activity of Fe3+, SO42−, AsO43−, or OH−.By using the Gibbs free energies of formation (ΔGfo), ΔGfo[Fe3+] = −4.6 kJ/mol, ΔGfo[SO42−] = −744.6 kJ/mol, ΔGfo[AsO43−] = −647.618 kJ/mol, ΔGfo[OH−] = −157.3 kJ/mol, and ΔGfo[H2O] = −237.18 kJ/mol [24, 25], the ΔGfo values for the synthetic basic ferric sulfate-arsenates [Fe(SO4)x(AsO4)y(OH)z·nH2O] were calculated, suggesting that the IAPs to the end of dissolution (5760–7920 h) were equal to the Ksp values for the basic ferric sulfate-arsenates. At the thermodynamic standard state, the standard free energy of reaction (ΔGro) is estimated according to (8)ΔGro=−5.708log_IAP.For equation (6),(9)ΔGro=1.00ΔGfoFe3++xΔGfoSO42−+yΔGfoAsO43−+zΔGfoOH−+nΔGfoH2O−ΔGfoFeSO4xAsO4yOHz⋅nH2O.Rearranging,(10)ΔGfoFeSO4xAsO4yOHz⋅nH2O=1.00ΔGfoFe3++xΔGfoSO42−+yΔGfoAsO43−+zΔGfoOH−+nΔGfoH2O−ΔGro.The experimental data of the dissolution for 240 d (5760 h), 270 d (6480 h), 7200 h (300 d), and 7920 h (330 d) and the calculated thermodynamic properties for the basic ferric sulfate-arsenates are presented in Tables2–6. The logˍIAP and ΔGfo values were −23.04 ± 0.01 and −914.06 ± 0.03 kJ/mol for the basic ferric sulfate-arsenate FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O], respectively (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. The solubilities decreased slightly with the increasing initial pH and were not obviously affected by the system temperature. The experimental results on dissolution were highly reproducible (Table 2–Table 6).Table 2
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257601.980.1223880.0991700.060330−23.04−23.04−914.05−914.0664801.980.1220300.0968320.059796−23.05±0.01−914.09±0.0372001.980.1233730.0994820.059996−23.04−914.0479201.980.1227460.0991700.059863−23.04−914.0625657603.680.0002520.0690760.007664−25.19−25.22−926.33−926.4764803.680.0002530.0687640.006838−25.23±0.03−926.53±0.1472003.680.0002570.0692320.006851−25.22−926.4979203.680.0002540.0684530.006837−25.22−926.52251057604.030.0002170.0721950.010698−25.22−25.24−926.50−926.5964804.040.0002030.0721950.010925−25.25±0.02−926.64±0.0972004.040.0002070.0726630.010887−25.24−926.5979204.040.0002030.0723510.010961−25.24−926.6325b257601.980.1309830.0790560.068739−22.99−22.97−913.76−913.6464801.990.1322360.0785880.069339−22.96±0.02−913.60±0.1272001.990.1325940.0799910.068538−22.96−913.6179201.990.1336690.0795230.068605−22.96−913.5935257601.940.1522020.1358140.041910−23.06−23.0464801.950.1536340.1381530.042044−23.03±0.0272001.950.1538130.1365930.041510−23.0479201.950.1535450.1367490.041910−23.0445257601.910.2674270.2444960.036772−22.86−22.8664801.920.2723510.2466790.035837−22.84±0.0272001.910.2708290.2473020.035103−22.8779201.910.2711870.2462110.035771−22.86a,bDissolution tests in duplicate.Table 3
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-2 [Fe(SO4)0.37(AsO4)0.59(OH)0.49·0.28H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257602.081.4665071.2337050.077881−22.03−22.01−931.40−931.3164802.091.4691931.2220110.078749−22.00±0.02−931.28±0.0972002.091.4684761.2220110.078682−22.00−931.2879202.091.4695511.2204520.078349−22.00−931.2925657603.000.0659840.6344730.000728−23.46−23.45−939.59−939.5664803.000.0663420.6421130.000745−23.45±0.01−939.54±0.0372003.000.0664320.6410220.000730−23.46−939.5779203.000.0663420.6464790.000735−23.45−939.55251057603.050.0528230.6011040.000611−23.59−23.59−940.35−940.3364803.040.0539870.6045340.000604−23.59±0.00−940.33±0.0272003.040.0540760.6021950.000606−23.59−940.3279203.040.0540410.6015720.000608−23.59−940.3225b257602.061.4638211.2449320.073477−22.08−22.06−931.70−931.6264802.071.4620301.2351090.074745−22.06±0.02−931.58±0.0872002.071.4616721.2263770.074011−22.06−931.5979202.071.4629251.2288720.073677−22.06−931.6035257602.021.4208461.4039790.060930−22.11−22.1064802.021.4235321.4050710.060797−22.11±0.0172002.031.4262181.3964950.060130−22.0979202.021.4251441.4011730.060730−22.1145257601.971.2928181.5982660.057994−22.17−22.1964801.961.2838651.6263330.058461−22.19±0.0272001.961.2775981.6158860.058261−22.1979201.961.2829701.6116760.058394−22.19a,b Dissolution tests in duplicate.Table 4
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-3 [Fe(SO4)0.27(AsO4)0.72(OH)0.30·0.21H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257601.970.1443230.1033810.059262−23.16−23.14−901.14−901.0264801.980.1454870.1040040.059462−23.14±0.02−900.99±0.1272001.980.1457550.1041600.059329−23.14−900.9979201.980.1456660.1044720.059596−23.14−900.9825657603.910.0004200.0687640.010056−25.07−25.08−912.05−912.0864803.900.0004170.0675170.009817−25.08±0.01−912.11±0.0372003.900.0004140.0670490.010044−25.08−912.0979203.900.0004160.0678290.010053−25.08−912.07251057604.220.0003750.0831100.012148−25.12−25.13−912.28−912.3564804.210.0003670.0840450.011985−25.13±0.01−912.34±0.0772004.210.0003630.0848250.011878−25.13−912.3879204.210.0003580.0852930.011962−25.14−912.4025b257601.970.1432490.1118010.057060−23.17−23.18−901.19−901.2164801.970.1457550.1114890.055992−23.17±0.01−901.19±0.1172001.960.1441440.1101630.056659−23.19−901.3279201.970.1453080.1113330.056793−23.17−901.1635257601.950.2130820.2013040.041310−23.04−23.0564801.950.2163950.2036420.039374−23.05±0.0172001.950.2172900.2014590.039641−23.0479201.950.2176480.2017710.039308−23.0545257601.930.3144300.2850370.035971−22.90−22.8964801.930.3229360.2870640.036238−22.89±0.0172001.930.3228460.2865960.036572−22.8979201.930.3229360.2881560.036638−22.89a,bDissolution tests in duplicate.Table 5
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-4 [Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2O and FeAsO4·0.75H2O mixture].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)Fe(SO4)0.28(AsO4)0.74(OH)0.22·0.28H2OFeAsO4·0.75H2OpHFeSO4AsO4log_IAPMean log_IAPlog_IAPMean log_IAP25a257601.960.1318780.2823860.054991−22.66−22.64−23.89−23.8764801.970.1334900.2716270.054323−22.64±0.02−23.86±0.0272001.970.1339370.2766170.054857−22.63−23.8679201.970.1338480.2753700.054257−22.63−23.8625657603.200.0005660.3407040.005979−24.39−24.40−25.53−25.5464803.200.0005480.3378970.005991−24.40±0.01−25.54±0.0172003.200.0005360.3388320.006026−24.41−25.5579203.200.0005390.3375850.006021−24.41−25.55251057603.200.0004680.3173140.006248−24.47−24.48−25.59−25.6164803.200.0004570.3138840.006032−24.49±0.01−25.62±0.0272003.200.0004610.3134160.006066−24.48−25.6179203.200.0004580.3123250.006059−24.49−25.6125b257601.950.1329530.2819190.055792−22.67−22.66−23.91−23.9064801.950.1394880.2837900.055591−22.65±0.01−23.89±0.0172001.940.1411000.2800470.055925−22.66−23.9079201.940.1400250.2833220.055391−22.67−23.9135257601.940.1952660.3895090.031099−22.66−22.66−24.03−24.0364801.950.2012640.3838960.030899−22.63±0.03−24.00±0.0372001.930.1997420.3821810.030565−22.68−24.0579201.930.2003690.3824920.030565−22.68−24.0545257601.890.2751270.4849370.024692−22.67−22.64−24.12−24.0964801.900.2889140.4872760.024359−22.63±0.03−24.09±0.0372001.900.2915110.4846250.024359−22.63−24.0879201.900.2899890.4886800.024626−22.62−24.08a,b Dissolution tests in duplicate.Table 6
Experimental data and solubility determination for the basic ferric sulfate-arsenate FeSAsOH-5 [Fe(SO4)0.24(AsO4)0.70(OH)0.42·0.16H2O].
Temp (°C)Initial pHTime (h)Dissolution data (mmol/L)log_IAPMean log_IAPΔGfo (kJ/mol)Mean ΔGfo (kJ/mol)pHFeSO4AsO425a257601.960.1566780.1203770.028163−24.29−24.27−879.30−879.2064801.970.1577520.1217800.028096−24.27±0.02−879.16±0.1072001.970.1578420.1220920.027896−24.27−879.1779201.970.1574840.1236510.028163−24.27−879.1625657603.900.0009130.0796790.001860−26.01−26.01−889.09−889.1264803.890.0009020.0810830.001853−26.01±0.00−889.11±0.0372003.890.0008940.0810830.001844−26.01−889.1479203.890.0008970.0812390.001849−26.01−889.12251057604.160.0005840.0935570.002324−26.16−26.17−889.95−890.0564804.150.0005730.0907500.002227−26.18±0.01−890.08±0.1072004.150.0005700.0905940.002232−26.18−890.0979204.150.0005720.0913740.002228−26.18−890.0825b257601.960.1600800.1252110.027562−24.29−24.26−879.28−879.1564801.970.1633930.1272380.027095−24.26±0.03−879.13±0.1372001.970.1637510.1258340.027762−24.25−879.0879201.970.1634820.1270820.027562−24.25−879.1035257601.940.2402990.2087880.006020−24.49−24.4864801.940.2479090.2129980.005877−24.48±0.0172001.940.2469250.2083200.005908−24.4879201.940.2470140.2128420.005897−24.4845257601.880.3731620.3428870.004635−24.42−24.4064801.890.3737000.3450700.004664−24.39±0.0272001.890.3735210.3442900.004661−24.3979201.890.3741470.3458490.004678−24.39a,bDissolution tests in duplicate.Using the result of the solubility experiments at initial pH 2 for 240 d (5760 h), 270 d (6480 h), 7200 h (300 d), and 7920 h (330 d), the logˍIAP for FeAsO4·0.75H2O is determined after the following reaction:(11)FeAsO4⋅0.75H2O=Fe3++AsO43−+0.75H2OThe calculated logˍIAP of −23.87 ± 0.02 (Table5) is in good agreement with the logˍKsp of −26.54 ± 0.35 [26].


## 4. Conclusions

The basic ferric sulfate-arsenates [FeSAsOH, Fe(SO4)x(AsO4)y(OH)z·nH2O] were prepared and characterized to investigate their potential fixation of arsenic in the oxidizing and acidic environment through a dissolution for days. The synthetic solids were generally the well-shaped prismatic crystals of the monoclinic polytypes. In the dissolution of the FeSAsOH-1 [Fe(SO4)0.27(AsO4)0.73(OH)0.27·0.26H2O] sample at 25–45°C and pH 2–10, the constituents were released in the order of AsO4 > SO4 > Fe within 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. The released iron, sulfate, and arsenate existed dominantly as Fe3+/Fe(OH)2+/FeSO4+, HSO4−/SO42−/FeSO4+, and H3AsO40/H2AsO4−, respectively. The 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 SO4 > AsO4 > Fe. The detaching was 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 0.045 mg/L. The 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 (Ksp) for 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 (ΔGfo) of −914.06 ± 0.03 kJ/mol. For all the five synthetic basic ferric sulfate-arsenates, the logˍIAPs varied between −26.17 ± 0.01 and −22.01 ± 0.02. The logˍIAP for FeAsO4·0.75H2O was estimated to be −23.87 ± 0.02 from the result of the solubility experiments at initial pH 2 for 240–330 d.



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*Source: 5556295-2021-06-07.xml*