1.1. Jarosite in the environmental context

Jarosite-type compounds are largely found in the terrestrial ecosystem as ferric sulfate salt that often forms in acidic and oxidizing environments enriched in sulfate. Different studies have shown its presence in:

• Acid sulfate soils formed from pyrite sediments [10, 11, 12, 13];
• Acid sulfate waters [14];
• Acid rock drainage (ARD) or acid mine drainage (AMD) [4, 15, 16, 17, 18, 19, 20, 21, 22, 23];
• Weathering residues from sulfide ore deposits [24];
• Alteration of coal from pyritic coal beds [25];
• Clay layers and beds [26, 27];
• Proximity to volcanic fissures [28, 29, 30];
• Oceanic fumaroles [31, 32] and hot springs [33];
• Hydrothermal environments [34, 35], sometimes associated with basalt [36];
• Hypersaline lakes, such as those in Australia and Antarctica [26, 29, 37, 38, 39, 40, 41, 42].

The formation of jarosite and the release of H⁺ are often attributed to the oxidation of sulfur-bearing minerals such as pyrite (FeS₂), which alters groundwater and effluent environments [43, 44]. Ideally, the formation (1) and dissolving (2) equations for K-jarosite are expressed below:

Furthermore, the early 2000s brought the discovery, in situ and from orbiting spacecraft, of jarosite on Martian soil [47, 48, 49, 50, 51, 52, 53, 54, 55]. This observation indicates that oxidizing and acidic fluids were present when jarosite was formed [48, 56, 57], and consequently, suggests the possibility of water existence on Martian soil. Other evidence has shown that hematite found on Mars has also been formed from jarosite [36, 58, 59].

1.2. Jarosite in the hydrometallurgical industrial context

Since the 1960s, the hydrometallurgical industry has been widely using the precipitation of jarosite compounds in industrial processes for separation of zinc from iron [60]. The production of zinc generates a large amount of jarosite-type compounds, which represents a potential environmental problem, since the jarosite scavenger property allows the presence of toxic elements in the iron residues. On the other hand, these residues may present economic value with the recovery of critical and valuable metals from zinc production waste. The study of jarosite leaching conditions constitutes, therefore, an attractive research area to reduce environmental issues and produce new resources [61].

Due to the presence of jarosite in industrial processes, it is of great importance to manage its residues, as its formation and dissolution mechanisms can alter the chemical environmental of rivers and groundwater by releasing the retained acidity, sulfate, hazardous elements, and secondary mineral waste depending on the conditions. Therefore, understanding the immobilization property of jarosite-type compounds is useful in forming stable minerals to reduce their spread in soils and aquifers and to help manage waste materials from mining tailings.

Therefore, having in mind the wide range of applications, rigorously understanding and evaluating jarosite dissolution and stability is important to predict the effect of long-term jarosite behavior on soil, water, and industrial processes, as well as to effectively manage hazardous waste to create potential remediation strategies, diminish environmental impact and improve hydrometallurgical processes to create economic profit [43, 44, 46, 62, 63, 64, 65, 66, 67].

With that aim, the collection and critical assessment of thermodynamic and kinetic data for jarosite dissolution and precipitation processes are necessary. Therefore, numerous scientists and research groups have made efforts to elucidate the behavior of jarosite-type compounds under different conditions over the past decades. Regardless, the results described in the literature concerning the solubility and kinetics of this compound are highly variable [46, 68]. Consequently, this review dedicates the following sections to shedding light on the works conducted up to this date and explaining the potential sources of disagreement among them.

1.3. Jarosite stability and transformations

Ferric phases undergo transformations into alternative phases upon removal from their stability zones. Jarosite-type compounds, for instance, are documented to go through a transformation resulting in the different phases, as illustrated in Figure 1.

It is established that jarosite can easily decompose when removed from its stability zone, transforming into phases of iron oxide or iron (III) oxyhydroxide [5, 35]. These transformations can also result in the release of acidity. Jarosite is considered a retained acidity (RA) resource in these environments, a form of acidity potentially released through the hydrolysis of poorly soluble and poorly crystalline Al³⁺ or Fe³⁺ minerals [46, 65, 66, 69]. According to Stoffregen [70], under surface conditions, jarosite tends to transform into goethite through the following reaction when exposed to diluted water with a higher pH:

However, this reaction can be complicated by the formation of metastable phases such as schwertmannite and ferrihydrite [71]. Ferrihydrite is a well-known hydrated iron oxide with chemical formula 5Fe₂O₃ ⋅ 9H₂O [72], where the hydration level can vary depending on how much water is present within the molecular arrangement [73, 74]. Schwertmannite is an oxyhydroxysulfate with chemical formula Fe₈O₈(OH)_8−2x(SO₄)_x ⋅ nH₂O where 1 ⩽ x ⩽ 1.75.

Above the transition temperatures from goethite to hematite, i.e., at 373.15 K, the transformation reaction of jarosite becomes:

The anhydrous forms of iron oxide can also exhibit different polymorphs, meaning the same chemical structure but different crystalline systems. The most well-known forms are hematite (α-Fe₂O₃) and maghemite (γ-Fe₂O₃). Hematite can be formed from magnetite through the following equation:

Like iron oxides, anhydrous iron oxyhydroxides (FeOOH) can exhibit several polymorphs. Goethite (α-FeOOH) is the most well-known compound, but there is also akaganeite (β-FeOOH) and lepidocrocite (γ-FeOOH). Redox-pH potential diagrams have been constructed showing the jarosite–hematite boundary between pH 2 and 4, and jarosite–goethite between pH 2 and 3 [76].

Bigham et al. [77] analyzed ochreous sediments as well as their associated solutions from twenty-eight AMD sites ranging from 279 to 300 K. The results revealed a change in solid composition depending on the pH of the mixture. Precipitates formed above pH 6.5 were mainly composed of ferrihydrite or a mixture of ferrihydrite and goethite. In contrast, precipitates with pH between 2.8 and 4.5 were mainly composed of schwertmannite, with traces of goethite and/or jarosite. Solutions with intermediate pH (between 4.5 and 6.5) produced a mixture of ferrihydrite and schwertmannite. Only the solution at pH 2.6 contained a considerable amount of jarosite. Bigham and his colleagues also developed a revised model for mineral speciation at 298 K and redox potential ranging, according to the pH, from −237 to 947 mV (pH > 6.5), 118 to 1065 mV (2.8 < pH < 4.5), 0 to 1000 mV (4.5 < pH < 6.5) and 592 to 1184 mV (pH < 2.6).

The continuation of this paper is organized as follows: Section 2 presents thermodynamic studies of jarosite, including those using dissolution, precipitation, and calorimetry. Section 3 discusses the kinetic dissolution studies within the environmental and industrial contexts. Section 4 focuses on the kinetic studies of jarosite-type compound precipitation at high and low temperatures, both with and without the presence of microorganisms. Finally, Section 5 highlights the most relevant methods used in the literature to characterize jarosite phases.

2.1. Thermodynamic studies through dissolution and precipitation

Table 1 presents the thermodynamic solubility constants, mostly at 298.15 K, reported in the literature, which vary considerably from one study to another. For example, for K-jarosite, the logarithm of the equilibrium constant ranges from −7.12 to −12.51, a difference of five orders of magnitude. These discrepancies are primarily explained by variations in the methodologies used, which will be described in this section.

Brown [84] investigated the precipitation of K-jarosite at 298.15 K and 0.1 MPa over a six-month period. He used K₂SO₄, hydrated FeSO₄, and dissolved Fe₂(SO₄)₃ in known concentrations of H₂SO₄. Subsequently, he employed a reverse approach (dissolution) using 0.5 g of jarosite in 0.005 M H₂SO₄ solution for six months. Brown considered the potential presence of additional compounds after precipitation and dissolution experiments. Consequently, X-ray diffraction (XRD) analysis was employed to address this question. However, all the X-ray diffraction patterns could only be identified as jarosite lines. Through the calculations of the free energy of formation, he demonstrated that equilibrium was practically reached by both methods with a slight difference in free energy of formation (−3276 ± 84 and −3192 ± 25 kJ/mol for precipitation and dissolution, respectively). Additionally, he calculated redox potential (E_h), pH, and plotted the activity logarithm diagram for the jarosite–goethite–pyrite system, highlighting the equilibrium of these minerals under ambient conditions.

The Gibbs free energies found by Brown [84] were initially recalculated (−3302 ± 84 and −3300 ± 25 kJ/mol for precipitation and dissolution, respectively) by Zotov et al. [87], who identified an arithmetic error in the calculations. Then, they used different thermodynamic data from Naumov et al. [96]. Furthermore, Zotov et al. [87] conducted precipitation experiments using a natural water sample and found the Gibbs free energy to be −3305.7 kJ/mol. Subsequently, Van Breemen [88] also revised Brown’s [84] values using other thermodynamic data from Robie and Walbaum [97] and found the Gibbs free energy to be equal to −3299 kJ/mol.

Vlek and colleagues [85] determined the solubility product of jarosite using the chelation method. Given that the solubility of Fe (III) compounds was low under their conditions, solubility measurements to detect the presence of iron were limited to highly acidic solutions. They developed a new method involving the use of EDTA as a chelating agent to increase the concentrations of Fe (III) from a natural jarosite sample, taken from a sulfate-acidic soil, in suspension. In this way, dissolution was studied, yielding a Gibbs free energy of −3310.4 kJ/mol and a logarithm of the solubility product constant (logK_sp) of −9.86. They acknowledged that the results obtained in their study could be subject to various errors resulting from pH measurements, iron concentrations, as well as incorrect estimations of constants and ionic strength. Moreover, they did not mention the use of XRD to identify the remaining solids after the experiments. Baron and Palmer [68] subsequently recalculated their Gibbs free energy with different free energies for the ions and found a value of −3331.4 kJ/mol.

Kashkay et al. [86] conducted three precipitation experiments from supersaturated solutions of K-jarosite, Na-jarosite, and Pb-jarosite over a 15-month period at room temperature. Additionally, they performed precipitations at 373.15 K. The Gibbs free energy for K-jarosite was −3299.7 ± 4 kJ/mol at 298.15 K and −3184 kJ/mol at 373.15 K. The solubility found at 298.15 K was greater for Na-jarosite than for K-jarosite.

In his thesis, Nordström [21] studied hydrogeochemical and microbiological parameters influencing the chemical speciation of heavy metals in AMD effluents in Northern California. His calculations revealed that the samples were in a state of supersaturation with respect to jarosite.

Hladky and Slansky [89] investigated the stability of minerals in the alunite–jarosite group in aqueous solutions at room temperature and pressure. Stability domains concerning pH were investigated for K-jarosite, Na-jarosite, and Pb-jarosite using Gibbs free energy values. It was revealed that pH constitutes the main control factor for the stability of minerals in the alunite group under the studied conditions.

Bladh [24] simulated the chemical erosion of felsic rocks to study the formation of goethite, jarosite and alunite. He recalculated the thermodynamic data previously obtained by Brown [84], Zotov et al. [87], Vlek et al. [85], and Van Breemen [88]. According to Bladh, the ambiguities present in those studies arose from the varied data regarding the free energies of ions that were utilized in the calculation. He calculated the solubility product (logK_sp) to be −7.12, while that of Na-jarosite was −3.60, and H₃O-jarosite was −3.57.

The Gibbs free energy determined by Bladh [24] for K-jarosite (−3287.60 kJ/mol, recalculated by Baron and Palmer [68]) from Posnjak and Merwin [75] and Kashkay et al. [86] data is inconsistent with Alpers et al. [4] data. However, values reported by Zotov et al. [87] and Vlek et al. [85] are closer. Despite this, Kashkay et al. [86] were the only ones to provide a set of data for all three substitution ions (K–Na–H₃O) in the A-site of jarosite at that time.

Chapman et al. [22] studied two AMD samples to understand the processes controlling the transfer and attenuation of heavy metals. They highlighted the variation in jarosite solubility thermodynamic data, making it difficult to select the necessary thermodynamic parameters (K_sp) to predict the saturation state of a mixture. They claimed that this variation is probably due to the use of different Δ_fG° values. They utilized X-ray powder diffraction to identify the solids. However, the results were limited to identifying the mineral phases since most of the precipitates were amorphous.

Several values for the of Fe³⁺ have already been reported. For example, Chapman and colleagues [22] chose to use the data used by Naumov et al. [96], Kashkay et al. [86], and Zotov [87] because they considered them coherent with each other. The results obtained by Nordström et al. [21] regarding the solubility product varied three to five orders of magnitude below than those found by Chapman et al. [22].

Alpers et al. [4] studied precipitates from AMD solutions in California, USA. In their experiments, seven samples with initial pH between 0.8 and 1.4 were kept for 11 to 13 years at a temperature of 295 ± 3 K. At the time of sampling, they were undersaturated with respect to jarosite. The oxidation of ferrous iron during storage led to saturation and precipitation of the ore. These minerals were characterized by solid analysis techniques such as XRD, scanning electron microscopy (SEM)-energy dispersive X-ray spectroscopy (EDS), and thermogravimetric analysis. The crystals had a diameter between 5 and 10 μm and a rhombohedral shape. The authors state that the well-crystallized nature indicated that the AMD had reached a thermodynamic equilibrium with jarosite. The relationship between crystallinity and equilibrium will be further addressed in Section 4.3.

Thus, Alpers et al. [4] recommended using a Δ_f G° value of 3293.5 ± 2.1 kJ/mol for a solid with composition K_0.77Na_0.03(H₃O)_0.20Fe₃(SO₄)₂(OH)₆. The average logK_sp of −9.35 showed good agreement with previous results from Chapman et al. [22]. Alpers and colleagues highlighted the need to determine the solubility of jarosite for more varied compositions to quantify the influence on solubility of substitution in the A-site (K, Na, H₃O). Despite the lack of data, they observed that the K–H₃O substitution approaches the ideal substitution compared to Na–K and Na–H₃O substitutions. Alpers et al. [4] estimated that additional experimental work was needed to accurately determine the solubility and free energy of formation of jarosite.

Stoffregen [70] studied the stability relationships between jarosite and Na-jarosite at temperatures ranging from 423.15 to 523.15 K through hydrothermal experiments lasting 1 to 11 weeks. X-ray powder diffraction was used to identify phases present in the products after the experiments. Experimental study of jarosite at high temperatures has advantages: (i) there are only three stable phases at temperatures above 423.15 K, compared to seven phases reported at 298.15 K by Posnjak and Merwin [75], (ii) reaction rates are higher, and (iii) hydronium substitution, which is common in natural jarosites and complicates the determination of terminal member properties, is negligible [4].

The Gibbs free energies were calculated by Stoffregen [70] at 473.15 K and 10 MPa, as presented in Table 1. Furthermore, they suggested that jarosite and Na-jarosite form nearly ideal minerals (A-site substituted only for K or Na) at 473.15 K, and there is complete miscibility between the two types of jarosite at 298.15 K.

Öborn and Berggren [95] studied two types of jarosite (K- and Na-jarosite) from sulfide sedimentary deposits in the Baltic Sea. The jarosites were characterized by XRD after dissolution, and the solubility product was determined from the dissolution of samples that reached equilibrium at different time intervals (from 6 h to 8 days). The solubility constants corresponded to −9.41 to −9.42. Additionally, Bigham et al. [77] reported solubility values for K-jarosite in a range from −9.21 (Kashkay et al. [86]; Chapman et al. [22]) to −12.51 (Lindsay [94]).

Today, the results obtained by Baron and Palmer [68] are most commonly used to predict the dissolution/precipitation of jarosite. Indeed, their work enabled the calculation of Gibbs free energies, enthalpy, entropy, and heat capacity. They were the first researchers to conduct a calorimetric study of the free energy of formation. Dissolution experiments of synthetic K-jarosite were conducted at different pH values (1.5–3.0) and temperatures close to room temperature (277.15–308.15 K) using a batch reactor. Two methodological approaches were considered for dissolution: first, the temperature at 298.15 K was kept constant while the pH varied from 1.5 to 3; second, the temperature varied from 277.15 to 308.15 K with a constant pH of 2. The composition (chemical formula) of jarosite was quasi-ideal, and the solids remained in the reactor for up to six months. XRD analyses were not conducted to characterize the reaction product solids in detail. Their study served as the basis for subsequent similar studies.

They observed that equilibrium was reached after 40 to 125 days of dissolution at pH 2 and 298.15 K, with an initial concentration of 0.2 g of jarosite/L of solution. The results obtained by Baron and Palmer [68] were closer to those of Alpers et al. [4], but one to two orders of magnitude higher compared to the results of Kashkay et al. [86] and Chapman et al. [22]. K_sp decreases as the temperature increases, ranging from −11.06 at 277.15 K to −11.68 at 308.15 K, indicating a negative enthalpy [68]. In other words, the addition of heat does not favor the reaction. However, this dependence in this temperature range is not linear, implying a variation in the enthalpy of reaction as the temperature changes.

According to Baron and Palmer [68], several reasons explain the divergences from previous studies:

1. Inconsistency of thermodynamic data and variations in the values used for the calculation of free energies and ion activities;
2. Substitution of ions in the jarosite structure;
3. Variations introduced by different experimental methodologies;
4. Analytical uncertainties.

Arslan and Arslan [60] conducted a review of the stability regions between jarosite and goethite at 298.15 and 368.15 K. They observed that the stability region of jarosite not only shifts toward a more acidic and oxidizing zone as the temperature increases but also becomes more restricted. However, they emphasized that, in addition to these thermodynamic considerations, the reaction kinetics play a significant role in the precipitation of iron compounds.

The work of Baron and Palmer [68] was considered the starting point for more complex structures containing mixtures of Fe³⁺ and Al³⁺ in B-sites, as well as the substitution of S⁶⁺ and Cr⁶⁺ in C-sites. Baron and Palmer [98] also studied the dissolution of analog jarosite with Cr substitution in the sulfur site. These studies were further pursued based on calorimetric investigations.

Thermodynamics was also studied by Smith et al. [7], aiming to monitor the ion release during dissolution and to characterize the newly formed solid phases. They calculated the equilibrium constant (K_sp) and solid saturation ratios (Ω) at pH 2 and 8 and room temperature (293.15 K). The results are close to those of Baron and Palmer [68], with logK_sp equal to −11.30. According to these authors, positive saturation ratios for hematite and goethite indicate the likely appearance of these minerals after a prolonged dissolution period. They utilized XRD analysis to identify the solids remaining after dissolution.

The study conducted by Smith and colleagues [7] was significant as it was the first to examine the dissolution of jarosite and the formation of new phases at pH 8, where Fe³⁺ reprecipitates. This is likely responsible for preventing the continuation of jarosite dissolution and reinforces the incongruence of jarosite dissolution already observed by Baron and Palmer [68]. This incongruence (non-stoichiometry of the solid) is due to the strong stability of the FeO₆ octahedron compared to the weak stability of KFe(OH)₄.

Smith et al. [99] also investigated the dissolution of Pb-jarosite [(H₃O)_0.74Pb_0.13Fe_2.92(SO₄)₂(OH)_5.76] and arsenic-substituted Pb-jarosite [(H₃O)_0.68Pb_0.32Fe2.86(SO₄)_1.69(AsO₄)_0.31(OH)_5.59] at pH 2 and 8. The logK_sp results showed a decrease (−9.35 to −10.65) compared to the standard jarosite structure. Their study demonstrated that, since jarosite dissolution is non-stoichiometric under the conditions studied, it is important to step back from diagrams such as Stoffregen [70] since the solid stability limits assume that the dissolution reaction occurs congruently (stoichiometrically).

Casas et al. [92] studied solubility through the precipitation of Na-jarosite and the speciation of an Fe(III)–Na–H₂SO₄–H₂O system at 343.15 K. They reported a lack of information on Na-jarosite solubility. Solubility data available in the literature up to that point had been obtained at room temperature or above 363.15 K. There was no information in the range 333.15–353.15 K, corresponding to the temperature range in which thermophilic microorganisms are active. The synthesis was carried out in reactors using a mixture of NaOH and FeH(SO₄)₂. Different mixtures of sulfuric acid and ferric iron were prepared and left to reach equilibrium for 45 days in a rotary oven at 343.15 K. The formation of Na-jarosite was observed at pH above 1.72 with the presence of goethite in the mixture after XRD analysis of the remaining solids. In contrast, for pH below 1.42, Na-jarosite was pure. Research by Casas and colleagues highlighted the dependence of ionic balance on bisulfates and sulfate complexes such as FeH(SO₄)₂, ${\mathrm{FeHSO}}_4^{2+}$, ${\mathrm{FeSO}}_4^{+}$, ${\mathrm{HSO}}_4^{-}$, and H⁺. The concentrations of these species are strongly influenced by acidity and temperature. The equilibrium constant found for Na-jarosite at 343.15 K was 16.5 ± 0.3.

As mentioned earlier in this section, Lemire et al. [90] re-evaluated the logK_sp value at 298.15 K, revising from −11.0 to −9.8. According to them, uncertainties in the stoichiometry, particle size, crystallinity variation, and the incongruence of dissolution (as reinforced by Smith et al. [7]) render the estimations provided by Baron and Palmer [68] merely speculative.

This detailed section illustrates that estimating the solubility of jarosite-type compounds is a complex task that has been attempted by numerous scientists over the past 50 years. The biggest challenges in estimating this constant are related to the variety of substitutions within the mineral, the variability of experimental methodologies employed, and discrepancies in the thermodynamic data regarding the free energy of ions. Despite the revaluation of Lemire et al. [90], who considered the calculated solubility product constant found by Baron and Palmer as speculative, their study has remained the most widely utilized since due to its detailed methodology and well-explained results. The techniques utilized by Baron and Palmer [68], including the precipitation methodology to obtain the solids used in the dissolution experiments, continue to be employed. The principal trends were also confirmed by Smith et al. [7] and were further corroborated by calorimetric studies described in Section 2.2 which affirmed that Baron and Palmer [68] had achieved equilibrium.

2.2. Thermodynamic studies through calorimetry

Drouet and Navrostky [76] conducted a calorimetric study to obtain formation enthalpies through dissolution calorimetry in molten mixtures at high temperatures. In contrast to Stoffregen [70], who used estimations, the data were obtained directly. However, the results of Stoffregen [70] are not consistent with the equation

Subsequently, Drouet et al. [100] conducted calorimetric studies to determine the enthalpy of jarosite containing chromate (CrO₄). Furthermore, Drouet et al. [101] also investigated solid mixtures between Na-jarosite and Na-alunite. From these studies, they determined the enthalpy of formation and were able to recommend values for Gibbs free energy and entropy.

Similar to Drouet and colleagues, Majzlan et al. [91] focused on the thermodynamic properties of H₃O-jarosite and obtained data on enthalpy, heat capacity, and entropy. Then, Majzlan et al. [102] continued their research on different jarosite-type compounds (K, Na, Rb, and NH₄), reporting their heat capacities and entropies. They calculated entropies from experimentally measured Cp data using an adiabatic calorimeter, unlike Stoffregen [70] who extrapolated data from experiments at high temperatures. Majzlan and collaborators considered their data more reliable. According to them, the significant variations found in the values of thermodynamic properties of jarosite indicated that the study of these properties would need to be addressed more reliably in future work. Additionally, with the data on enthalpy and entropy, it would be possible to calculate phase diagrams for the phase equilibria of this complex and challenging mineralogical group.

Forray et al. [103] studied the formation of yavapaiite KFe(SO₄)₂ resulting from the heating of K-jarosite at temperatures above 673.15 K to better understand the decomposition of this mineral. Then, in 2010, Forray and colleagues [93] synthesized, characterized, and studied the thermodynamics of Pb-jarosite, providing the enthalpy of formation of the compound and enabling the calculation of Pb–Fe–SO₄–H₂O diagrams as well as the estimation of the equilibrium constant (logK_sp ranging from −24.3 to −26.2). Furthermore, Forray et al. [104] continued their work on the synthesis and thermodynamics of Pb–As, Pb–Cu, and Pb–Zn jarosite mixtures, resulting in the following logK_sp values: −13.94 ± 1.89; −4.38 ± 1.81; −3.75 ± 1.80, respectively. They concluded that the transformations and dissolution reactions of arsenic-containing jarosites are much more complex.

3.1. Kinetics of jarosite environmental dissolution

Many researchers started studying jarosite dissolution rates with different conditions and methodologies. The first robust study of the dissolution kinetics of jarosite was conducted by Baron and Palmer [68], focusing on the release of ions in an acidic environmental medium. They used the Noyes–Whitney model to fit the data from dissolution at pH 2 and 298.15 K, resulting in a first-order kinetics and a kinetic constant K equal to 7.9 ± 0.5 × 10⁻⁷  s⁻¹. As discussed earlier, in the thermodynamic part studied by these authors, their study served as a basis for subsequent research.

Gasharova et al. [45] investigated the surface behavior of jarosite using Atomic Force Microscopy (AFM) at pH 5.5. The goal was to directly observe morphological changes on the surfaces of jarosite depending on interactions with different aqueous solutions, to study the mechanisms of dissolution and to compare dissolution rates of jarosites containing K⁺ and H₃O⁺. The kinetic rates found for each jarosite compound were two orders of magnitude higher than those of Baron and Palmer [68] using the DVKM model, −6.84 and −6.35 for K⁺ and H₃O⁺, respectively.

The dissolution rates of K-jarosite were also studied by Smith et al. [7] to simulate conditions similar to ARD/AMD sites and remediated sites, using pH values of 2 and 8 at 293.15 K. They noticed that Baron and Palmer [68] had studied dissolutions at low temperatures but without characterizing the solids produced at the end of the reaction. Even without investigating temperature and agitation, they suggested that the dissolution was controlled by mass transfer.

In Figure 2, equilibrium was reached by both Baron and Palmer et al. [68] and Smith et al. [7] with the same acid medium (HCl), same pH (2) and initial concentration (0.2 g/L), and similar temperatures (293.15 and 298.15 K, respectively). However, the time required to reach this quasi-equilibrium was approximately 80 days for Smith et al. [7] and 40 days for Baron and Palmer [68]. The concentrations at the end of dissolution were similar for iron (Fe), potassium (K), and sulfate $\left({\mathrm{SO}}_4^{2-}\right)$, 0.390, 0.204, and 0.330 mmol/L, respectively, for Smith et al. [7] compared to 0.434, 0.178, and 0.332 mmol/L for Baron and Palmer [68]. However, the dissolution curves indicate a significantly faster attainment of equilibrium for the earlier study. The kinetic data (logr) turned out to be −8.51 for Baron and Palmer [68] and −8.80 for Smith et al. [7]. This difference can be partially attributed to the 5 K temperature difference between the experiments.

Welch et al. [43] used higher pH values (3 and 4) to simulate ARD conditions in Australia using a natural sample. According to their observations, equilibrium was not reached after 12 days. Jarosite continued to dissolve for the rest of the year, though at a slower rate. Figure 3 shows dissolution kinetic curves at pH 3. Figure 3(a) shows kinetic data utilizing HCl, which promotes the release of elements in larger quantities, while H₂SO₄ (Figure 3b) produces a less pronounced release.

Elwood Madden et al. [109], Pritchett et al. [110], and Legett et al. [111] simulated conditions closely resembling those found on Mars with different pH, temperatures, and salinities for the dissolution of K-jarosite, and Zahai et al. [112] for Na-jarosite.

Elwood Madden et al. [109] studied the effect of pH (1 to 10) and temperature (277.15 to 323.15 K) on the initial kinetic rates of K-jarosite dissolution. Their objective was also to investigate morphologies and reaction products. Linear regression on rates found from experiments resulted in a value of E_a = 79 kJ/mol, applying the Arrhenius equation. Moreover, a Type V behavior was found where the lowest rate is at pH 3.5, but the rate increases either by increasing a_H+ (pH < 3.5) or increasing a_OH- (pH > 3.5). A general trend of the Elwood-Madden et al. data [109] can be observed in Figure 4.

Kendall et al. [44] used the same conditions as Smith et al. [7] but added a third scenario in which a sudden input of fresh water (ultrapure water) was considered, using As-jarosite to compare with K-jarosite dissolution in a sulfuric acid medium. They observed a faster release of K⁺ during the first two hours, followed by a decrease in speed. Regarding Fe³⁺, its release proved faster after the first day, exceeding the potassium concentration. Despite the effect observed by Welch et al. [43], where dissolution kinetics are faster for chloride environments than for sulfuric environments, Kendall et al. [44] reached equilibrium in 14 days of experimentation, much faster than Baron and Palmer [68] or Smith et al. [7], challenging the accuracy or reliability of the results obtained by Kendall and collaborators.

Similar to Baron and Palmer [68], the dissolution rate studies mentioned above used batch reactors that do not take reactive transport into account, a critical aspect in environmental geochemistry [46, 113]. Therefore, more recent research in this field has studied the dissolution process in continuous flow reactors to consider the effect of liquid flowing.

Dixon et al. [42] studied the dissolution rates of jarosite in circulation reactors, showing that dissolution rates are not affected by the flow rate, but they did not explore the relationship with pH. Meanwhile, Qian et al. [114] studied the dissolution of natural jarosite under quiescent solution conditions using non-stirred batch reactors and column leaching experiments. They reported slower dissolution rates than those observed using dissolution methods in stirred reactors.

However, the use of gravity-driven column leaching experiments, while more realistic, did not allow control over mass transport conditions within the column. Noticing this lack of information, Trueman et al. [46] investigated the dissolution of jarosite in a circulation reactor with a constant flow rate and at various pH levels.

All these studies provided useful insights under a range of different conditions such as pH, temperature, and salinity. However, varying orders of magnitude were found for the dissolution rate because of differing experimental conditions or mineral samples.

As an example, Figure 4 illustrates the variation in dissolution rate values of K-jarosite with pH for reactions in batch reactors at room temperature (293.15–298.15 K).

Figure 4 shows the variation in rate even for similar conditions of pH and temperature, the most important parameters for jarosite dissolution. These variations are likely due to the nature of jarosite (natural or synthetic), its structure and ion substitution, differing solid/liquid ratios, stirring speeds, and the duration of dissolution monitoring (several authors, such as Elwood Madden et al. [109] or Kendall et al. [44], have used only the first two hours, which increases the rate for potassium ions). In Table 2, the different studies in batch reactors (or AFM study, used by Gasharova et al. [45]) are gathered along with their main parameters and results.

3.2. Kinetics of jarosite industrial dissolution

Numerous efforts have been made to examine the impacts of chemical composition, temperature, pH, and particle size on dissolution rates in various reactive industrial environments. Due to the growing interest in the recovery of strategic metals from industrial residues containing Zn, as well as other high-value metals such as Pb, Cu, Ag, Cd, Co [61, 116], several leaching agents have been used as promising alternatives to recovery processes. Meanwhile, others authors have focused on avoiding the release of toxic metals in industrially precipitated jarosite.

The shrinking core kinetic model (SCKM) has been widely utilized over the past thirty years to describe the decomposition kinetics of jarosite-type compounds by evaluating decomposition conditions in order to prevent environmental issues and recover high-value metals such as silver. Primarily, researchers have used an alkaline medium as lixiviant in numerous studies regarding jarosites containing silver [117, 118, 119, 120, 121, 122, 123, 124], prompting many studies in alkaline environments (NaOH/Ca(OH)₂). In these dissolutions, jarosite decomposition typically occurs between 323.15 and 363.15 K, and silver is leached with or without cyanide. Other studies have also examined the release of minerals such as rubidium [125], arsenic [126, 127, 128], chromium [129, 130] mercury [131], and thallium [67]. However, in recent years, attention has also been given to the use of various acids as a decomposition medium, as evidenced in Table 3 where conditions are displayed, as well as the activation energies and reaction order for both the induction and progressive periods.

3.3. Natural and synthetic jarosites

Based on the kinetic studies previously presented in this section, it is noteworthy that the nature of the jarosite-type compound can impact the results. For this reason, this additional subsection on the dissolution kinetics presents some kinetic studies that synthesized jarosite-type compounds and the methodology used to achieve this synthesis. The use of synthesized methodologies allows other scientists to replicate the experiments and compare them with new dissolution conditions, for example. On the other hand, some authors have preferred the use of natural jarosite samples to replicate processes as they occur in the environment. However, natural jarosite often contains different amounts of other species that will influence the concentration of elements in aqueous solution. Additionally, the effect of the temporal alteration of rocks and minerals can also reduce the dissolution rate of this mineral.

4.1. Studies in abiotic environments at high temperatures

The precipitation of jarosite is of significant industrial interest due to the adoption of the industrial process known as the “Jarosite Process”. This process is employed for controlling hydrometallurgical impurities, primarily in zinc production. During its course, ferric iron precipitates in the form of jarosite, and zinc can be recovered. These processes are known to occur at elevated temperatures exceeding 363.15 K [27].

Dutrizac is one of the authors who has extensively researched the effects of conditions on the precipitation of various jarosite-compounds types in hydrometallurgical processes and their behavior in the presence of other elements in different reaction environments.

Dutrizac and Kaiman [154] initially worked on the synthesis of rubidium and thallium jarosite. Subsequently, Dutrizac et al. [145] investigated the conditions affecting the formation of Pb-jarosite, while Dutrizac [155] studied their formation in hydrochloric acid. Dutrizac and Dinardo [156] examined the coprecipitation of copper and zinc with Pb-jarosite.

Dutrizac and collaborators have also studied the behavior of several elements during the precipitation of jarosite compounds, including: selenate [157], mercury [158], cesium and lithium [159], silver [160], arsenic at 370.15 K from sulfate and hydrochloric acid medium [161] as well as at 423.15 K [148], cadmium [162], thiocyanate and cyanate [163], gallium [164], lanthanides [165], vanadium [166], rare earth elements [167], thallium (III) [168], chromium (III) [169, 170], germanium and tin [171], alkaline-earth elements (beryllium and radium) [172], halides (fluoride, chloride, bromide, and iodide) [173], scandium, yttrium, and uranium [174], selenium (VI) and selenium (IV) [175].

Among Dutrizac’s most important findings, in his study of the effects of alkaline precipitation of Na-jarosite, he demonstrates that alkali substitution in the jarosite structure and solid formation yields increase with an increase in alkali concentration in the precipitation solution. According to the author, seeding and ionic strength had little influence on the yield since for a 24 h reaction the iron-precipitated % remained almost constant. The author established that the stability of different precipitated jarosite types follows the order: $\mathrm{K}>{\mathrm{NH}}_4^{+}>\mathrm{Na}$. Agitation proved to be partially important in this process, meaning that it has a positive influence as it is necessary to keep the solids in suspension. However, additional agitation does not have a significant effect, which in general leads to conclude that the process is probably limited by the surface reaction rather than diffusion [62].

Limpo et al. [176] studied the influence of acidity, Fe³⁺ concentration, ammonia concentration, and seeding on the precipitation rate of Fe³⁺ as NH₄-jarosite from solutions containing sulfuric acid at temperatures between 363.15 and 373.15 K. These authors observed that the control step was crystal growth, and the nucleation rate was proportional to the surface area of the solids introduced into the system, the ammonia concentration, and the concentration of the ${\mathrm{Fe}}_2{\left(\mathrm{OH}\right)}_4^{2+}$ complex produced in the solution. This study demonstrated that the acidity of the solution had a significant, inversely proportional influence on the precipitation rate due to competition between hydrogen ions and ammonium ions to be adsorbed on the jarosite surface. Furthermore, pH appeared to regulate the concentration of the complex, making it a crucial element to ensure the success of the iron precipitation.

The study of seeding at different temperatures was then undertaken and showed positive effects on precipitation rate. Teixeira and Tavares [177] conducted batch experiments to study the effect of seeding on the initial precipitation rate of NH₄-jarosite, noticing a linear increase with the growing quantity of seeds. This trend proved valid for both high and low concentrations of Fe³⁺, ${\mathrm{NH}}_4^{+}$, and ${\mathrm{Mn}}_2^{+}$.

Dutrizac [178] investigated the effect of seeding on the precipitation of NH₄-jarosite and Na-jarosite, connecting it to other factors such as pH, temperature, and concentrations of ferric sulfate. The experiments were conducted with a solution containing 0.30 M Fe(SO₄)_1.5, 0.30 M Na₂SO₄, a pH of 1.5, a temperature of 371.15 K, and agitation at 500 rpm, lasting around 8 h. His conclusions indicated that seeding promotes precipitation, where the initial rate increases linearly in the presence of seeds (from 0 to 100 g/L). The presence of seeds also broadens the jarosite precipitation range to lower pH and temperatures, thereby eliminating the induction time.

Among the discoveries he made, he determined that the activation energy of Na-jarosite in the presence of seeds was equal to 106 kJ/mol. The increase in temperature (from 353.15 to 373.15 K) and pH (from 1.10 to 2.00) favors precipitation. In order to calculate the initial kinetic rate, he fitted the data to a second-order polynomial and calculated the slope using the derivative kinetic method (DVKM) as explained in Section 3.

Dutrizac [178] also compiled apparent activation energy data published by different authors [179, 180, 181] up to that year, demonstrating the variety of results for different types of jarosite but always at high temperatures and using restricted temperature ranges.

Dutrizac [182] examined the factors influencing the precipitation of K-jarosite in a hydrochloric and sulfate medium. Subsequently, Dutrizac [183] compared the precipitation rates of NH₄-jarosite and Na-jarosite in a ferric sulfate medium (0.30 M Fe(SO₄)_1.5/sulfuric acid) at a temperature between 343.15 and 373.15 K. The results showed an activation energy of 87 kJ/mol and indicated a faster precipitation of NH₄-jarosite compared to Na-jarosite.

In the hydrometallurgical industry, significant efforts have been made to study the precipitation of jarosite at high temperatures in abiotic environments (>363.15 K), given that the majority of microorganisms are not capable of thriving at such high temperatures. However, the development of hydrometallurgy for the recovery of high-value metals has led to the creation of new processes operating at more moderate temperatures due to the use of a biotic medium. In this way, additional research has emerged to consider lower temperatures as can be seen in the next section.

4.2. Studies in biotic environments at low temperatures

Investigations aiming to understand the formation of jarosite under biotic environments and low temperatures, as well as the rate at which this reaction takes place, are necessary. Its precipitation can compromise the performance of recovering high-value metals. Additionally, understanding precipitation is very important since it can lead to the formation of diffusion barriers on mineral surfaces, impacting dissolution rate, reducing the concentration of ferric iron, hindering ore attack, and potentially trapping high-value metals, such as silver.

In this regard, researchers have dedicated their work to study the control of mineral formation, including jarosite, in the presence of microorganisms. The experimental results obtained in sulfate-rich bioleaching systems, studied by Toro [184], suggest the formation of jarosite at a 303.15 K in environments rich in Acidithiobacillus bacteria. In these systems, jarosite precipitation exhibits very favorable and non-selective kinetics regarding pH. This is attributable to the high concentration of iron (III), which promotes the formation of precipitation intermediates of this phase, as well as the presence of bacterial cells serving as surface for crystalline growth.

Wang et al. [185] synthesized schwertmannite and jarosite from solutions containing ferrous iron sulfate, inoculated with Acidithiobacillus ferrooxidans. Schwertmannite transformed into H₃O–NH₄-jarosite when the reaction time increased from 19 to 40 days, and the reaction temperature from 309.15 to 318.15 K, in acidic medium at pH 2.

Wang et al. [186] synthesized NH₄-jarosite in the presence of microorganisms (A. ferrooxidans between 295.15 and 309.15 K, Sulfolobus metallicus at 318.15 K, and Acidithiobacillus caldus and Sulfobacillus spp. at 338.15 K) at temperatures ranging from 295.15 to 338.15 K and ferrous sulfate oxidation in pH ranging from 2 to 3, with varying concentrations of ${\mathrm{NH}}_4^{+}$. Schwertmannite served as a solid precursor for jarosite formation. An increase in ammonium concentration favored the formation of jarosite rather than schwertmannite.

Egal et al. [187] investigated arsenic trapping during Fe²⁺ oxidation in a medium containing Acidithiobacillus ferrooxidans. They found that the type of bacteria used favored the formation of tooeleite ore rather than schwertmannite, while jarosite formed without the need for bacteria at the end of the experiments (after 15 days).

Recently, process studies involving the oxidation of iron by archaea at 343.15 K in continuous reactor were undertaken by Kaksonen [188] as well as Kaksonen [189] in a two-stage airlift bioreactor, aiming to assess the performance of these reactors and their ability to separate the jarosite produced during the reaction.

4.3. Studies in abiotic environments at low temperatures

Despite the advancement of research in jarosite precipitation at low temperatures under biotic conditions, little data was found to understand jarosite precipitation at low temperatures without the influence of microorganisms. Consequently, new studies under these conditions have been initiated in the last few years.

Bigham et al. [190] characterized ferric iron precipitates in batch experiments conducted at temperatures ranging from 295.15 to 313.15 K, using ferric sulfate and mineral salts in sulfuric acid. The goal of these experiments was to simulate bioleaching solutions in an oxidizing environment containing K⁺, Mg²⁺, Ca²⁺, NO³⁻, and ${\mathrm{PO}}_4^{3-}$. The duration of the experiments was 10 days. They observed a decrease in Fe³⁺ concentration in the solution as the experiment temperature increased. The phases formed were primarily composed of H₃O-jarosite. However, a phase similar to Maus salt (K₅Fe₃(SO₄)₆(OH)_2.8H₂O) was identified at temperatures below 298.15 K. When potassium was replaced by hydronium in Maus salt, metavoltine was produced [191]. The proportion of poorly crystallized or non-crystallized crystals decreased with increasing temperature.

According to Sasaki and Konno [192] and Wang [186], synthesis conditions play a crucial role in the structure and morphology of jarosite samples. They observed a tendency for the formation of more individual particles at higher temperatures. Some particles exhibited needle-shaped surface characteristics, which could be attributed to metavoltine formation.

However, based on the observation of surface morphology, these characteristics also resemble previous descriptions of schwertmannite [Fe₈O₈(OH)₆(SO₄)] [185, 186, 193]. The presence of schwertmannite is not observed in the diffractograms because this mineral is poorly crystallized, and its characteristic peaks are broad and masked by the sharper jarosite peaks. Similarly, the background noise in the diffractograms is directly proportional to the amount of poorly crystallized materials, including potentially schwertmannite.

As part of the study on jarosite formation under ambient-like conditions, Kim et al. [194] investigated the effects of the presence of oxyanions (${\mathrm{AsO}}_4^{-}$, ${\mathrm{SeO}}_3^{-}$, ${\mathrm{SeO}}_4^{-}$, ${\mathrm{MoO}}_4^{-}$, ${\mathrm{CrO}}_4^{-}$) in a solution of ferric sulfate (III) and potassium hydroxide, varying concentrations and aging times (1 h to 40 days). They utilized techniques such as XRD, SEM, and chemical analysis, to follow the precipitation chemical reaction. Initially, an amorphous phase formed for all samples. As the aging time increased, K-jarosite and oxyanions-containing jarosite began to precipitate at room temperature, with varying precipitation rates and levels of crystallinity. As stated before in Section 2.1, crystallinity increases over time as equilibrium is achieved.

In jarosite samples with low precipitation rates and low crystallinity, the amorphous phase contained high concentrations of oxyanions, likely due to the rapid precipitation of the amorphous phase of iron oxyanions. The results demonstrated that jarosite coprecipitation could play a more significant role in controlling the behavior of ${\mathrm{CrO}}_4^{-}$ than ${\mathrm{AsO}}_4^{-}$ in the context of acid drainage from mining operations. Oxyanions-containing jarosites exhibited oxyanion concentrations similar to those in the starting solution, suggesting that amorphous phases dissolve and recrystallize in equilibrium with the liquid phase over the aging period.

These studies conducted at low temperatures and in the absence of microorganisms show that the parameters employed during precipitation experiments have an impact not only on the crystallinity and surface characteristics of particles, but also on the formation of additional mineral phases. It indicates the importance of solid analysis to understand the phenomenon and highlights the challenge in characterizing ferric iron phases. Further research is warranted to evaluate the kinetics of jarosite formation and improve the comprehension of its formation reaction.