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\DOI{10.5802/crchim.321}
\datereceived{2023-12-14}
\daterevised{2024-04-26}
\datererevised{2024-05-25}
\dateaccepted{2024-05-27}
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\begin{table*}[t!]
\caption{(continued)}
\tabcolsep=8.5pt
\begin{tabular}{ccccccc}
\hline
$a_{\mathrm{H}_2\mathrm{O}}$ &amp; 
Initial pH &amp; 
$T$~(K) &amp; 
Solution &amp; 
$\log r^{\mathrm{a}}$ &amp;
Authors &amp; 
Element \\ 
\hline}
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\caption{(continued)}
\fontsize{8}{10}\selectfont\tabcolsep1.2pt
\begin{tabular}{@{}cccccccc@{}}  
\hline
Authors using a basic medium &amp; 
Jarosite type &amp; 
Leaching medium &amp; 
$T$~(K) &amp; 
$E_{\mathrm{a},\mathrm{pro}}$~(kJ/mol) &amp; 
$n_{\mathrm{pro}}$ &amp; 
$E_{\mathrm{a},\mathrm{ind}}$~(kJ/mol) &amp; 
$n_{\mathrm{ind}}$\\
\hline
}
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%\makeatletter
%\def\TITREspecial{\relax}
%\def\cdr@specialtitle@english{GDR Prom\'eth\'ee - French Research Network on \textit{Hydrometallurgical Processes for Primary and Secondary Resources}}
%\def\cdr@specialtitle@french{GDR Prom\'eth\'ee - Groupement de Recherche sur les Proc\'ed\'es Hydrom\'etallurgiques pour la Gestion Int\'egr\'ee des Ressources Primaires ...}
%\makeatother

\CDRsetmeta{articletype}{review}

\title{Dissolution and precipitation of jarosite-type compounds: state
of the art and future perspectives}

\alttitle{Dissolution et pr\'{e}cipitation des compos\'{e}s de type
jarosite : \'{E}tat de l'art et perspectives futures}

\author{\firstname{Mateus} \lastname{De Souza Buriti}}
\address{Universit\'{e} de Pau et des Pays de l'Adour, 
E2S UPPA, LaTEP, Rue Jules Ferry BP 7511 64 075 PAU Cedex, Pau, France}
\email[M. De Souza Buriti]{mdsburiti@univ-pau.fr}

\author{\firstname{Pierre} \lastname{C\'{e}zac}\CDRorcid{0000-0003-1363-7740}}
\addressSameAs{1}{Universit\'{e} de Pau et des Pays de l'Adour, 
E2S UPPA, LaTEP, Rue Jules Ferry BP 7511 64 075 PAU Cedex, Pau, France}
\email[P. C\'{e}zac]{pierre.cezac@univ-pau.fr}

\author{\firstname{Lidia} \lastname{Cas\'{a}s}\CDRorcid{0000-0002-9638-1244}\IsCorresp}
\addressSameAs{1}{Universit\'{e} de Pau et des Pays de l'Adour, 
E2S UPPA, LaTEP, Rue Jules Ferry BP 7511 64 075 PAU Cedex, Pau, France}
\email[L. Cas\'{a}s]{lidia.casas@univ-pau.fr}

\shortrunauthors

\keywords{\kwd{Jarosite}
\kwd{Kinetics}
\kwd{Thermodynamics}
\kwd{Environment}
\kwd{Industrial processes}}

\altkeywords{\kwd{Jarosite}
\kwd{Cin\'{e}tique}
\kwd{Thermodynamique}
\kwd{Environnement}
\kwd{Proc\'{e}d\'{e}s industriels}}

\begin{abstract}
Jarosite-type compounds are complex-structured minerals that are widely
distributed in natural environments, utilized in industrial processes,
or formed as byproducts of these processes. Consequently, a
comprehensive understanding of the kinetics and thermodynamics of
jarosite precipitation and dissolution is important for the effective
control and management of this mineral. However, literature data is
lacking or inconsistent regarding these phenomena. The availability of
accurate and reliable data is of utmost importance for the development
and improvement of robust hydrometallurgical processes and for
regulating the environmental impact of this mineral. In light of this
importance, this work aims to review the most important studies
published up to 2023 related to obtaining data on jarosite dissolution
and precipitation. Therefore, this article contributes to a detailed
analysis of the current state of the art and outlines future prospects
for research on this secondary mineral.
\end{abstract}

\begin{altabstract}
Les compos\'{e}s de type jarosite sont des min\'{e}raux \`{a} structure
complexe largement pr\'{e}sents dans les environnements naturels,
utilis\'{e}s dans des proc\'{e}d\'{e}s industriels ou form\'{e}s en
tant que sous-produits de ces proc\'{e}d\'{e}s. Par cons\'{e}quent, une
compr\'{e}hension approfondie de la cin\'{e}tique et de la
thermodynamique de la pr\'{e}cipitation et de la dissolution de la
jarosite est importante pour le contr\^{o}le et la gestion efficaces de
ce min\'{e}ral. Cependant, il existe un manque de donn\'{e}es
bibliographiques ou des incoh\'{e}rences concernant ces
ph\'{e}nom\`{e}nes. La disponibilit\'{e} de donn\'{e}es pr\'{e}cises et
fiables est d'une importance capitale pour le d\'{e}veloppement et
l'am\'{e}lioration de proc\'{e}d\'{e}s hydrom\'{e}tallurgiques robustes
et pour r\'{e}guler l'impact environnemental de ce min\'{e}ral. Dans
cette optique, ce travail vise \`{a} examiner les \'{e}tudes les plus
importantes publi\'{e}es jusqu'en 2023 relatives \`{a} l'obtention de
donn\'{e}es sur la dissolution et la pr\'{e}cipitation de la jarosite.
Par cons\'{e}quent, cet article contribue \`{a} une analyse
d\'{e}taill\'{e}e de l'\'{e}tat actuel de l'art et esquisse les
perspectives futures de la recherche sur ce min\'{e}ral secondaire. 
\end{altabstract}

\maketitle

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\let\rmmu\upmu

\section{Introduction}\label{sec1}
Jarosite-type compounds are minerals belonging to the alunite--jarosite
group presenting an ideally general formula
[\textbf{A}$_{\upalpha}$\textbf{B}$_{\upbeta}$(\textbf{CO}$_{4}$)$_{\upgamma}$(OH)$_{\updelta}$]. 
Several types of jarosites can form depending on the ions occupying the
\textbf{A}-,  \textbf{B}-, and \textbf{C}-sites in the mineral
structure. \textbf{A}-sites are commonly occupied by 
monovalent\unskip\break
(K$^{+}$,  Na$^{+}$, Ag$^{+}$,  Tl$^{+}$, $\mathrm{NH}_{4}^{+}$,  Rb$^{+}$,
H$_{3}$O$^{+}$), divalent  (Cu$^{2+}$, Hg$^{2+}$,  Ca$^{2+}$,
Sr$^{2+}$, Ba$^{2+}$, Pb$^{2+}$), trivalent or quadrivalent (Ga$^{3+}$,
Cr$^{3+}$,  Th$^{+4}$,  U$^{+4}$) cations. The \textbf{B}-site can be
occupied by smaller cations with octahedral coordination (Fe$^{3+}$ 
and Al$^{3+}$), or minor ions such as V$^{3+}$, Fe$^{2+}$, Cu$^{2+}$, 
Zn$^{2+}$, Mg$^{2+}$. Finally, the \textbf{C}-site can be occupied by 
S$^{6+}$,  Cr$^{6+}$, As$^{5+}$ or P$^{5+}$~\cite{1,2,3,4,5,6}. Ion
stoichiometry (${\upalpha}$, ${\upbeta}$, ${\upgamma}$, ${\updelta}$)
will be directly related to the valency of the ions present in the
\textbf{A}-,  \textbf{B}-, and \textbf{C}-sites.

The most important and common jarosite-type mineral found in
terrestrial environments is K-jarosite
[KFe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$]. The structure of jarosite is
based on linear tetrahedral--octahedral--tetrahedral (T--O--T) sheets,
made up from slightly distorted FeO$_{6}$ octahedra and SO$_{4}$
tetrahedra~\cite{5,7}. Authors affirm that K-jarosite structure
generally contains ferric ion vacancies and ``additional water'' in the
form of hydronium. Therefore, the chemical formula is more correctly
written as $\mathrm{H}_{3}\mathrm{O}_{1-x}\mathrm{K}_{x}
\mathrm{Fe}_{3-y}[(\mathrm{OH})_{6-3y}({\mathrm{H}_{2}}
\mathrm{O})_{3y}(\mathrm{SO}_{4})_{2}]$~\cite{3,4,7,8,9}.\looseness=-1

According to Alpers et~al.~\cite{4}, any attempt to define the
thermodynamic properties of minerals in the alunite--jarosite group
must consider issues related to the non-stoichiometry of their
structures. Indeed, the values of the (Al $+$ Fe)/SO$_{4}$ ratio can
vary considerably from the ideal stoichiometry of 3/2 (3 moles of
Fe$^{3+}$ to 2 moles of sulfate), as observed in the work of 
H\"{a}rtig et~al.~\cite{8}, with ratios ranging from 2.20/2 to 2.52/2,
as well as in the research of Ripmeester 
\mbox{et~al.}~\cite{9}, with a ratio
of 2.33/2.

\subsection{Jarosite in the environmental context}\label{sec1.1} 
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:
\begin{itemize}
\item 
Acid sulfate soils formed from pyrite  sediments~\cite{10,11,12,13}; 
\item
Acid sulfate waters~\cite{14}; 
\item
Acid rock drainage (ARD) or acid mine drainage 
(AMD)~\cite{4,15,16,17,18,19,20,21,22,23}; 
\item
Weathering residues from sulfide ore deposits~\cite{24}; 
\item
Alteration of coal from pyritic coal beds~\cite{25}; 
\item
Clay layers and beds~\cite{26,27}; 
\item
Proximity to volcanic fissures~\cite{28,29,30}; 
\item
Oceanic fumaroles~\cite{31,32} and hot springs~\cite{33}; 
\item
Hydrothermal environments~\cite{34,35}, sometimes associated with 
basalt~\cite{36}; 
\item
Hypersaline lakes, such as those in Australia and 
Antarctica~\cite{26,29,37,38,39,40,41,42}. 
\end{itemize}

The formation of jarosite and the release of H$^{+}$ are often
attributed to the oxidation of sulfur-bearing minerals such as pyrite
(FeS$_{2}$), which alters groundwater and effluent 
environments~\cite{43,44}. Ideally, the formation~(\ref{eq1}) and 
dissolving~(\ref{eq2}) equations for K-jarosite are expressed below:
{\begin{eqnarray}
&amp;&amp;
\mathrm{K}^{+}+3\mathrm{Fe}^{3+}+{2\mathrm{SO}}_{4}^{2-}+
6\mathrm{H}_{2}\mathrm{O}\rightarrow \nonumber\\&amp;&amp;\quad
\mathrm{K}\mathrm{Fe}_{3}(\mathrm{SO}_{4})_{2}(\mathrm{OH})_{6}+
6\mathrm{H}^{+}
\label{eq1}\Seqnsplit
&amp;&amp;
\mathrm{K}\mathrm{Fe}_{3}(\mathrm{SO}_{4})_{2}
(\mathrm{OH})_{6}+6\mathrm{H}^{+}\rightarrow \nonumber\\&amp;&amp;\quad
\mathrm{K}^{+}+3\mathrm{Fe}^{3+}+2\mathrm{SO}_{4}^{2-}+
6\mathrm{H}_{2}\mathrm{O}\label{eq2}
\end{eqnarray}}\unskip
On the other hand, the dissolution of jarosite in AMD occurs when the
acid production exceeds the buffering capacity of the surrounding rock
or soil. The pH values in pyritic mine waste and acid lakes typically
range from pH 1.5 to  3~\cite{45}, leading to the release of heavy
metals present in the jarosite structure. Therefore, it is essential to
understand the formation of jarosite and its dissolution, as well as
their impact on soil and water chemistry, in order to effectively
manage soils containing sulfur-bearing
minerals~\cite{46}.

Furthermore, the early 2000s brought the discovery, in situ and from
orbiting spacecraft, of jarosite on Martian 
soil~\cite{47,48,49,50,51,52,53,54,55}. This observation indicates that
oxidizing and acidic fluids were present when jarosite was 
formed~\cite{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~\cite{36,58,59}.
\looseness=-1

\subsection{Jarosite in the hydrometallurgical industrial
context}\label{sec1.2} 
Since the 1960s, the hydrometallurgical industry has been widely using
the precipitation of jarosite compounds in industrial processes for
separation of zinc from  iron~\cite{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~\cite{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\unskip\break 
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 
\mbox{effectively} manage
hazardous waste to create 
\mbox{potential} remediation strategies, diminish
environmental 
\mbox{impact} and improve hydrometallurgical processes to create
economic  profit~\cite{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~\cite{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.

\subsection{Jarosite stability and transformations}\label{sec1.3} 
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~\ref{fig1}.

\begin{figure*}
{\vspace*{2pt}}
\includegraphics{fig01}
{\vspace*{2pt}}
\caption{\label{fig1}Transformations of jarosite compounds into
different ferric phases.}
\vspace*{1pt}
\end{figure*}

It is established that jarosite can easily decompose when removed from
its stability zone, transforming into phases of iron oxide or iron
(III)  oxyhydroxide~\cite{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 Al$^{3+}$ or Fe$^{3+}$  minerals~\cite{46,65,66,69}.
According to  Stoffregen~\cite{70}, under surface conditions, jarosite
tends to transform into goethite through the following reaction when
exposed to diluted water with a higher pH:
{\begin{equation}\label{eq3}
\mathrm{KFe}_{3}(\mathrm{SO}_{4})_{2}(\mathrm{OH})_{6}\rightarrow
3\mathrm{FeOOH}+\mathrm{K}^{+}+{2\mathrm{SO}}_{4}^{2-}+
3\mathrm{H}^{+}
\end{equation}}\unskip

However, this reaction can be complicated by the formation of
metastable phases such as schwertmannite and  ferrihydrite~\cite{71}.
Ferrihydrite is a well-known hydrated iron oxide with chemical formula
$5\mathrm{Fe}_{2}\mathrm{O}_{3}\cdot9\mathrm{H}_{2}\mathrm{O}$~\cite{72}, 
where the hydration level can vary depending on how much water is
present within the molecular  arrangement~\cite{73,74}. Schwertmannite
is an oxyhydroxysulfate with chemical formula
$\mathrm{Fe}_{8}\mathrm{O}_{8}
(\mathrm{OH})_{8-2x}{(\mathrm{SO}_{4}})_{x}\cdot 
n\mathrm{H}_{2}\mathrm{O}$ where $1 \leq x\leq1.75$.

Above the transition temperatures from goethite to hematite, i.e., at
373.15~K, the\vadjust{\pagebreak} 
transformation 
\mbox{reaction} of jarosite becomes: 
{\begin{eqnarray}
&amp;&amp;
\mathrm{K}\mathrm{Fe}_{3}(\mathrm{SO}_{4})_{2}(\mathrm{OH})_{6}\rightarrow
\nonumber\\&amp;&amp;\quad
3/2\mathrm{Fe}_{2}\mathrm{O}_{3}+\mathrm{K}^{+}+{2\mathrm{SO}}_{4}^{2-}+
3\mathrm{H}^{+}+1.5\mathrm{H}_{2}\mathrm{O}
\label{eq4}
\end{eqnarray}}\unskip
The only compound that has appeared in addition to the previous
equation is the formation of Fe(SO$_{4}$)(OH), which does not exist
naturally but has already been reported by Posnjak and 
Merwin~\cite{75} in the study of the
Fe$_{2}$O$_{3}$--H$_{2}$O--SO$_{3}$ system at temperatures higher than
373.15~K:
{\begin{eqnarray}
&amp;&amp;
\mathrm{K}\mathrm{Fe}_{3}(\mathrm{SO}_{4})_{2}(\mathrm{OH})_{6}+
3\mathrm{H}^{+}+{\mathrm{SO}}_{4}^{2-}\rightarrow 
\nonumber\\&amp;&amp;\quad
3\mathrm{Fe}(\mathrm{SO}_{4})(\mathrm{OH})+\mathrm{K}^{+}+
3\mathrm{H}_{2}\mathrm{O}
\label{eq5}
\end{eqnarray}}\unskip

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
(${\upalpha}$-Fe$_{2}$O$_{3}$) and maghemite
(${\upgamma}$-Fe$_{2}$O$_{3}$). Hematite can be formed from magnetite
through the following equation:
{\begin{equation}\label{eq6}
2\mathrm{Fe}_{3}\mathrm{O}_{4}+1/2\mathrm{O}_{2}\rightarrow 
3(\upalpha -\mathrm{Fe}_{2}\mathrm{O}_{3})
\end{equation}}\unskip

Like iron oxides, anhydrous iron oxyhydroxides (FeOOH) can exhibit
several polymorphs. Goethite (${\upalpha}$-FeOOH) is the most
well-known compound, but there is also akaganeite  (${\upbeta}$-FeOOH)
and lepidocrocite (${\upgamma}$-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~\cite{76}.

Bigham et~al.~\cite{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 ${&gt;}$
6.5), 118 to 1065~mV (2.8 ${&lt;}$ pH ${&lt;}$ 4.5), 0 to 1000~mV (4.5 ${&lt;}$
pH ${&lt;}$ 6.5) and 592 to  1184~mV (pH ${&lt;}$ 2.6). 

The continuation of this paper is organized as follows: 
Section~\ref{sec2} presents thermodynamic studies of jarosite,
including those using dissolution, precipitation, and calorimetry. 
Section~\ref{sec3} discusses the kinetic dissolution studies within the
environmental and industrial contexts.  Section~\ref{sec4} 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~\ref{sec5} highlights the most relevant methods used
in the literature to characterize jarosite phases. 

\section{Thermodynamic studies of jarosite}\label{sec2}
Several studies have focused extensively on the study of the
crystallography and structure of jarosite  compounds~\cite{78,79}. The
investigations into the structure of jarosite date back to the mid-20th
century. Initially,  Parker~\cite{80} examined alunite minerals
containing potassium or sodium. Concurrently, Brophy  et~al.~\cite{81}
studied K-jarosite and K-alunite.  Kubisz~\cite{82} dedicated his
research to the study of potassium and hydronium substitution in
jarosite, demonstrating the existence of H$_{3}$O-jarosite.

However, Brophy and Sheridan~\cite{2} were the first to investigate the
substitution of K--Na--$\mathrm{H}_{3}\mathrm{O}^{+}$ mixture in the
jarosite structure under terrestrial surface temperature and pressure
conditions. Additionally, in their study, they asserted that hydronium
substitution by potassium in the jarosite structure is facilitated by
an increase in temperature, in line with 
\mbox{several}
studies conducted in the 1950s on the
composition of natural and synthetic jarosites. 
\mbox{Furthermore,} Ross 
et~al.~\cite{83} showed that H$_{3}$O-jarosite can form in the absence
of alkalis.

Since the studies on the structure and composition of jarosites,
researchers have frequently noted the variability of the solubility
product data. Some authors have delved into dissolution and/or
precipitation phenomena to obtain thermodynamic 
data~\cite{4,7,68,70,84,85,86}, while others have reinterpreted these
results by introducing additional thermodynamic data for equilibrium
calculations, proposing new  values~\cite{22,24,87,88,89,90}.

Nevertheless, the most cited paper and reliable data used nowadays
still is Baron and  Palmer~\cite{68} who studied the thermochemistry
and solubility of synthetic K-Jarosite dissolution at different pH and
near room temperatures using a stirred batch reactor and were the first
to carry out a calorimetric study of K-Jarosite. This study formed a
basis for other similar investigations, like the thermodynamic study by
Smith  et~al.~\cite{7} who calculated the solubility product constant
$(K_{\mathrm{sp}})$, the ion activity product (IAP) and saturation
indices (SI). However, more recently, Lemire  et~al.~\cite{90} from the
OECD Nuclear Energy Agency published a book on the chemical
thermodynamics of iron re-evaluating the results obtained by Baron and 
Palmer~\cite{68}, which will be discussed later in 
\mbox{Section~\ref{sec2.1}.}

Other authors, like Drouet and  Navrotsky~\cite{76}, worked with the
calorimetric study of Na-jarosite, while Majzlan  et~al.~\cite{91}
studied thermodynamic properties of H$_{3}$O-jarosite. Other authors
studied the thermodynamics of jarosite such as Jamieson 
et~al.~\cite{17}, Navrotsky  et~al.~\cite{19}, Casas  et~al.~\cite{92},
and Forray  et~al.~\cite{93}. 

Hence,  Section~\ref{sec2} is divided into two main parts. 
Section~\ref{sec2.1} presents studies using dissolution and
precipitation processes to obtain thermodynamic data, while 
Section~\ref{sec2.2} focuses on thermodynamic data obtained through
calorimetry studies.

\subsection{Thermodynamic studies through dissolution and
precipitation}\label{sec2.1} 
Table~\ref{tab1} 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.

\begin{table*}
\caption{\label{tab1}Solubility constants of jarosites found in the
literature at 298.15~K}
\begin{tabular}{ccccc}
\thead
Authors &amp; Jarosite-type &amp; Method &amp; $\log K_{\mathrm{sp}}$ &amp; 
$\Delta_{f}G{\text{\textdegree}}_{298.15}$~(kJ/mol) \\
\endthead
Brown~\cite{84} &amp; K &amp; Prec. and Diss. &amp; ${-}$9.65 &amp; ${-}$3276 and ${-}$3192\\
Zotov et~al.~\cite{87} &amp; K &amp; Precipitation &amp; ${-}$9.47 &amp;${-}$3305.70\\ 
Van Breemen~\cite{88} &amp; K &amp; Dissolution recalculated &amp; ${-}$9.46 &amp;${-}$3299.00 \\ 
Vlek et~al.~\cite{85} &amp; K &amp; Dissolution &amp; ${-}$9.86 &amp;${-}$3331.40 \\ 
Kashkay et~al.~\cite{86} &amp; K &amp; Precipitation &amp; ${-}$9.32 &amp;${-}$3299.67 \\ 
Lindsay~\cite{96} &amp; K &amp;  &amp; ${-}12.51$ &amp; \\ 
Hladky and Slansky~\cite{89} &amp; K &amp; Calculated &amp; ${-}$9.08 &amp; \\ 
Bladh~\cite{24} &amp; K &amp; Dissolution &amp; ${-}$7.12 &amp;${-}$3287.60\\ 
Chapman et~al.~\cite{22} &amp; K &amp; Dissolution &amp; ${-}$9.21 &amp;\\ 
Alpers et~al.~\cite{4} &amp; K &amp; Precipitation &amp; ${-}$9.35 &amp;${-}$3300.25 \\ 
Stoffregen~\cite{70} &amp; K &amp; Precipitation &amp; ${-}$8.8 &amp; ${-}$3416.3$^{\mathrm{a}}$ \\ 
\"{O}born and Berggren~\cite{97} &amp; K &amp; Dissolution &amp; ${-}$9.42 to ${-}$9.41 &amp;  \\ 
Baron and Palmer~\cite{68} &amp; K &amp; Dissolution &amp; ${-}$11.0 &amp;${-}$3309.8 \\ 
Smith et~al.~\cite{7} &amp; K &amp; Dissolution &amp; ${-}$11.3 &amp; \\ 
Lemire et~al.~\cite{90} &amp; K &amp; Recalculated &amp; ${-}$9.8 &amp; \\ 
Kashkay et~al.~\cite{86} &amp; Na &amp; Precipitation &amp; ${-}$8.93 &amp;${-}$3256.70 \\ 
Hladky and Slansky~\cite{89} &amp; Na &amp; Calculated &amp; ${-}$5.10 &amp; \\ 
Bladh~\cite{24} &amp; Na &amp; Dissolution &amp; ${-}$3.60 &amp; \\ 
Chapman et~al.~\cite{22} &amp; Na &amp; Dissolution &amp; ${-}$5.28 &amp; \\ 
Alpers et~al.~\cite{4} &amp; Na &amp; Precipitation &amp; ${-}$7.18 &amp; \\ 
\"{O}born and Berggren~\cite{97} &amp; Na &amp; Dissolution &amp; ${-}$9.39 to ${-}$9.31
&amp; \\ 
Hladky and Slansky~\cite{89} &amp; Pb &amp; Calculated &amp; ${-}$7.89 &amp; \\ 
Chapman et~al.~\cite{22} &amp; Pb &amp; Dissolution &amp; ${-}$16.28 &amp; \\ 
Bladh~\cite{24} &amp; H$_{3}$O &amp; Dissolution &amp; ${-}$3.57 &amp; \\ 
Alpers et~al.~\cite{4} &amp; H$_{3}$O &amp; Precipitation &amp; ${-}$6.04 &amp; 
\botline 
\end{tabular}
\tabnote{$^{\mathrm{a}}$~Values of $\Delta _{f}G{\text{\textdegree}}$ at 473.15~K
and 10~MPa.}
\vspace*{6pt}
\end{table*}

Brown~\cite{84} investigated the precipitation of K-jarosite at
298.15~K and 0.1~MPa over a six-month period. He used K$_{2}$SO$_{4}$,
hydrated FeSO$_{4}$, and dissolved Fe$_{2}$(SO$_{4}$)$_{3}$ in known
concentrations of H$_{2}$SO$_{4}$. Subsequently, he employed a reverse
approach (dissolution) using 0.5~g of jarosite in 0.005 M
H$_{2}$SO$_{4}$ 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 (${-}3276\pm84$ and 
${-}3192\pm25$~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~\cite{84} were initially
recalculated  (${-}3302\pm84$ and  ${-}3300\pm25$~kJ/mol for precipitation
and dissolution, respectively) by Zotov  et~al.~\cite{87}, who
identified an arithmetic error in the calculations. Then, they used
different thermodynamic data from 
\mbox{Naumov} et~al.~\cite{94}. Furthermore,
Zotov  et~al.~\cite{87} conducted precipitation experiments using a
natural water sample and found the Gibbs free energy to be 
${-}$3305.7~kJ/mol. Subsequently, Van  Breemen~\cite{88} also revised
Brown's~\cite{84} values using other thermodynamic data from Robie and 
Walbaum~\cite{95} and found the Gibbs free energy to be equal 
to ${-}$3299~kJ/mol.
\looseness=-1

Vlek and colleagues~\cite{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
$(\log K_{\mathrm{sp}})$ 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~\cite{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.~\cite{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  ${-}3299.7 \pm 4$~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, Nordstr\"{o}m~\cite{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~\cite{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~\cite{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~\cite{84}, Zotov
et~al.~\cite{87}, Vlek et~al.~\cite{85}, and Van  Breemen~\cite{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 
$(\log K_{\mathrm{sp}})$ to be ${-}$7.12, while that of Na-jarosite was 
${-}$3.60, and H$_{3}$O-jarosite was  ${-}$3.57.

The Gibbs free energy determined by Bladh~\cite{24} for K-jarosite 
(${-}$3287.60~kJ/mol, recalculated by Baron and  Palmer~\cite{68}) from
Posnjak and Merwin~\cite{75} and  Kashkay et~al.~\cite{86} data is
inconsistent with Alpers et~al.~\cite{4} data. However, values reported
by Zotov  et~al.~\cite{87} and Vlek et~al.~\cite{85} are closer. Despite
this, Kashkay  et~al.~\cite{86} were the only ones to provide a set of
data for all three substitution ions (K--Na--H$_{3}$O) in the
\textbf{A}-site of jarosite at that time.

Chapman et~al.~\cite{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 
\mbox{thermodynamic}
parameters $(K_{\mathrm{sp}})$ to predict the saturation state of a
mixture. They claimed that this variation is probably due to the use of
different $\Delta_{f}G{\text{\textdegree}}$ 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 Fe$^{3+}$ have already
been reported. For example, Chapman and  colleagues~\cite{22} chose to
use the data used by Naumov  et~al.~\cite{94}, Kashkay et~al.~\cite{86},
and Zotov~\cite{87} because they considered them coherent with each
other. The results obtained by Nordstr\"{o}m et~al.~\cite{21} regarding
the solubility product varied three to five orders of magnitude below
than those found by Chapman  et~al.~\cite{22}.

Alpers et~al.~\cite{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 
$295\pm3$~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~$\rmmu$m 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~\ref{sec4.3}.

Thus, Alpers et~al.~\cite{4} recommended using a $\Delta_{f}G{\text{\textdegree}}$
value of  $3293.5\pm2.1$~kJ/mol for a solid with composition
K$_{0.77}$Na$_{0.03}$(H$_{3}$O)$_{0.20}$Fe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$.
The average  $\log K_{\mathrm{sp}}$ of  ${-}$9.35 showed good agreement
with previous results from Chapman  et~al.~\cite{22}. Alpers and
colleagues highlighted the need to determine the solubility of jarosite
for more varied compositions to quantify the influence on solubility of
substitution in the \textbf{A}-site (K, Na, H$_{3}$O). Despite the lack
of data, they observed that the K--H$_{3}$O substitution approaches the
ideal substitution compared to Na--K and Na--H$_{3}$O substitutions.
Alpers  et~al.~\cite{4} estimated that additional experimental work was
needed to accurately determine the solubility and free energy of
formation of jarosite.

Stoffregen~\cite{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~\cite{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~\cite{4}.

The Gibbs free energies were calculated by  Stoffregen~\cite{70} at
473.15~K and 10~MPa, as presented in Table~\ref{tab1}. Furthermore,
they suggested that jarosite and Na-jarosite form nearly ideal minerals
(\textbf{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.

\"{O}born and Berggren~\cite{97} 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.~\cite{77} reported solubility values for
K-jarosite in a range from ${-}$9.21
(Kashkay et~al.~\cite{86}; Chapman et~al.~\cite{22}) 
to ${-}$12.51 (Lindsay~\cite{96}).

Today, the results obtained by Baron and Palmer~\cite{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 
\mbox{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~\cite{68} were closer to those of Alpers et~al.~\cite{4}, but one
to two orders of magnitude higher compared to the results of Kashkay 
et~al.~\cite{86} and Chapman et~al.~\cite{22}.
$K_{\mathrm{sp}}$ decreases as the temperature increases,
ranging from  ${-}$11.06 at 277.15~K to ${-}$11.68 at 308.15~K, indicating
a negative  enthalpy~\cite{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~\cite{68}, several reasons explain the
divergences from previous studies:
\begin{enumerate}[(a)]
\item
Inconsistency of thermodynamic data and variations in the values used
for the calculation of free energies and ion activities; 
\item
Substitution of ions in the jarosite structure; 
\item
Variations introduced by different experimental methodologies; 
\item
Analytical uncertainties.
\end{enumerate}

Arslan and Arslan~\cite{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~\cite{68} was considered the starting
point for more complex structures containing mixtures of Fe$^{3+}$ and
Al$^{3+}$ in \textbf{B}-sites, as well as the substitution of S$^{6+}$
and Cr$^{6+}$ in \textbf{C}-sites. Baron and  Palmer~\cite{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.~\cite{7}, aiming to
monitor the ion release during 
\mbox{dissolution} and to characterize the
newly formed solid phases. They calculated the equilibrium constant
$(K_{\mathrm{sp}})$ and solid saturation ratios  $(\Omega)$ at pH 2 and
8 and room temperature (293.15~K). The results are close to those of
Baron and  Palmer~\cite{68}, with $\log K_{\mathrm{sp}}$ 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~\cite{7} was significant as
it was the first to examine the dissolution of jarosite and the
formation of new phases at pH 8, where Fe$^{3+}$ 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~\cite{68}. This incongruence
(non-stoichiometry of the solid) is due to the strong stability of the
FeO$_{6}$ octahedron compared to the weak stability of KFe(OH)$_{4}$.

Smith et~al.~\cite{99} also investigated the dissolution of  Pb-jarosite
[(H$_{3}$O)$_{0.74}$Pb$_{0.13}$Fe$_{2.92}$(SO$_{4}$)$_{2}$(OH)$_{5.76}$] 
and arsenic-substituted  Pb-jarosite
[(H$_{3}$O)$_{0.68}$Pb$_{0.32}${\ubreak}Fe2.86(SO$_{4}$)$_{1.69}$(AsO$_{4}$)$_{0.31}$(OH)$_{5.59}$]
at pH 2 and 8. The  $\log K_{\mathrm{sp}}$ 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~\cite{70} since the solid
stability limits assume that the dissolution reaction occurs
congruently (stoichiometrically).

Casas et~al.~\cite{92} studied solubility through the precipitation of
Na-jarosite and the speciation of an
Fe(III)--Na--H$_{2}$SO$_{4}$--H$_{2}$O 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(SO$_{4}$)$_{2}$. 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(SO$_{4}$)$_{2}$,
$\mathrm{FeHSO}_{4}^{2+}$, 
$\mathrm{FeSO}_{4}^{+}$, $\mathrm{HSO}_{4}^{-}$, 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  
\mbox{$16.5\pm0.3$.}\looseness=-1

As mentioned earlier in this section, Lemire et~al.~\cite{90}
re-evaluated the  $\log K_{\mathrm{sp}}$ 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.~\cite{7})
render the estimations provided by Baron and  Palmer~\cite{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.~\cite{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~\cite{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.~\cite{7} and
were further corroborated by calorimetric studies described in 
Section~\ref{sec2.2} which affirmed that Baron and Palmer~\cite{68} had
achieved equilibrium.\looseness=-1

\subsection{Thermodynamic studies through\newline calorimetry}\label{sec2.2} 
Drouet and Navrostky~\cite{76} conducted a calorimetric study to obtain
formation enthalpies through dissolution calorimetry in molten mixtures
at high temperatures. In contrast to  Stoffregen~\cite{70}, who used
estimations, the data were obtained directly. However, the results of 
Stoffregen~\cite{70} are 
not consistent with the equation
{\begin{equation}\label{eq7}
\Delta _{\mathrm{f}}G{\text{\textdegree}}=
\Delta _{\mathrm{f}}H{\text{\textdegree}}-
T\Delta _{\mathrm{f}}S{\text{\textdegree}}
\end{equation}}\unskip
On the other hand, Drouet and Navrotsky~\cite{76} also determined the
heat capacity using Differential Scanning Calorimetry (DSC), and it
appears to agree with the results obtained by  Stoffregen~\cite{70}.
The recommended values for Gibbs free energy  (${-}$3318~kJ/mol) from
calorimetric data are similar to the findings in previous studies,
indicating that the redox/pH potential diagrams no longer need
adjustments and that equilibrium was achieved in Baron and Palmer's 
study~\cite{68}.

Subsequently,  Drouet et~al.~\cite{100} conducted calorimetric studies
to determine the enthalpy of jarosite containing chromate (CrO$_{4}$).
Furthermore,  Drouet et~al.~\cite{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.~\cite{91} focused on
the thermodynamic properties of H$_{3}$O-jarosite and obtained data on
enthalpy, heat capacity, and entropy. Then, Majzlan  et~al.~\cite{102}
continued their research on different jarosite-type compounds (K, Na,
Rb, and NH$_{4}$), reporting their heat capacities and entropies. They
calculated entropies from experimentally measured \textit{Cp} data using an
adiabatic calorimeter, unlike  Stoffregen~\cite{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.~\cite{103} studied the formation of yavapaiite
KFe(SO$_{4}$)$_{2}$ 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~\cite{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--SO$_{4}$--H$_{2}$O diagrams as well as
the estimation of the equilibrium constant  ($\log K_{\mathrm{sp}}$
ranging from  ${-}$24.3 to  ${-}$26.2). Furthermore, Forray 
et~al.~\cite{104} continued their work on the synthesis and
thermodynamics of Pb--As, Pb--Cu, and Pb--Zn jarosite mixtures,
resulting in the following  $\log K_{\mathrm{sp}}$ values: 
${-}13.94\pm1.89$;  ${-}4.38\pm1.81$;  ${-}3.75\pm1.80$, respectively. They
concluded that the transformations and dissolution reactions of
arsenic-containing jarosites are much more complex.

\section{Kinetic studies of jarosite dissolution\newline 
processes}\label{sec3}
The most used tool to model dissolution and precipitation processes in
the geochemistry field is thermodynamics. Nonetheless, not every
mineral can attain equilibrium in a rapid way, especially when using
conditions at low temperatures and neutral pH. These minerals
presenting slow reaction rates often keep geochemical processes from
reaching equilibrium and this makes geochemical kinetics another
essential tool for modeling these  situations~\cite{105}. Therefore,
elucidating the behavior of jarosite dissolution and precipitation can
aid in the development of robust models which describe the geochemical
cycling of these elements in the environment and in hydrometallurgical
processes.

The importance of studying the dissolution kinetics of jarosite
compounds gained academic significance after the 1990s. Kinetic data
have been found and modeled by several authors. The results generally
converge in the same direction; however, differences appear mainly due
to the various parameters used in the studies. Three main models have
been used to mathematically represent the process and the effects of
variables: the derivative kinetic model (DVKM), the Noyes--Whitney
kinetic model (NWKM), and the shrinking core kinetic model (SCKM).

The DVKM entails obtaining initial rate constants derived through the
fitting of initial kinetic data to a second-order polynomial equation
and calculating the slope, considering the beginning of the curve as
time zero.

 While the NWKM describes dissolution as a diffusion process,
proportional to the difference between the solute's saturation
concentration in the solvent (solubility) and the solute's
concentration in the solution. This concentration gradient-based
approach is inspired by Fick's law and has influenced several
diffusion-based dissolution  models~\cite{106}. This equation is widely
used in the study of dissolution in the pharmaceutical field.

The SCKM, initially developed by Yagi and  Kunii~\cite{107} is
described by Levenspiel~\cite{108}. It considers mass transfer
mechanisms in the studies of the 
\mbox{dissolution} reaction of a solid in a
fluid. In this model, particle size decreases over time, and particles
eventually disappear as the reaction progresses. 
\mbox{During} dissolution,
two main steps are observed: (i)~the induction period, where the
reaction front is forming and the concentration increases very slowly,
and (ii)~the progressive period, where the reaction front is already
established and the concentrations increase faster until reaching
equilibrium. 

In this section, the studies will be divided into two categories: those
intended for the study in environmental settings 
(Section~\ref{sec3.1}) and those in the hydrometallurgical industry
conditions (Section~\ref{sec3.2}). Additionally,  Section~\ref{sec3.3}
discusses the nature of the jarosite-type compound samples used in the
thermodynamic studies of  Section~\ref{sec2} and in the kinetic studies
of  Sections~\ref{sec3.1} and~\ref{sec3.2}, whether they are natural or
synthetic and how this distinction may influence the outcomes.

\subsection{Kinetics of jarosite environmental\newline
dissolution}\label{sec3.1} 
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~\cite{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 
$7.9\pm0.5\times10^{-7}~\mathrm{s}^{-1}$. As discussed earlier, in the
thermodynamic part studied by these authors, their study served as a
basis for subsequent research.

Gasharova et~al.~\cite{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 
$\mathrm{H}_{3}\mathrm{O}^{+}$. The kinetic rates found for each jarosite
compound were two orders of magnitude higher than those of Baron and 
Palmer~\cite{68} using the DVKM model, ${-}$6.84 and ${-}$6.35 for K$^{+}$
and $\mathrm{H}_{3}\mathrm{O}^{+}$, respectively.\looseness=-1

The dissolution rates of K-jarosite were also 
\mbox{studied} by Smith 
et~al.~\cite{7} to simulate conditions similar to ARD/AMD sites and
remediated sites, using pH 
\mbox{values} of 2 and 8 at 293.15~K. They noticed
that Baron and  Palmer~\cite{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~\ref{fig2}, equilibrium was reached by both Baron and Palmer 
et~al.~\cite{68} and Smith et~al.~\cite{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.~\cite{7} and 40 days for Baron and  Palmer~\cite{68}. The
concentrations at the end of dissolution were similar for iron (Fe),
potassium (K), and sulfate 
$(\mathrm{SO}_{4}^{2-})$, 0.390, 0.204, and
0.330~mmol/L, respectively, for Smith  et~al.~\cite{7} compared to
0.434, 0.178, and 0.332~mmol/L for Baron and  Palmer~\cite{68}.
However, the dissolution curves indicate a significantly faster
attainment of equilibrium for the earlier study. The kinetic data 
$(\log r)$ turned out to be ${-}$8.51 for Baron and  Palmer~\cite{68} and
${-}$8.80 for Smith  et~al.~\cite{7}. This difference can be partially
attributed to the 5~K temperature difference between the experiments. 

\begin{figure*}
\vspace*{2pt}
\includegraphics{fig02}
\caption{\label{fig2}Dissolution curves of K-jarosite for (a) Baron and
Palmer~\cite{68} and (b) Smith et~al.~\cite{7}  (Figures reprinted with
permission from Elsevier).}
\vspace*{2pt}
\end{figure*}

Welch et~al.~\cite{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~\ref{fig3} shows dissolution kinetic curves at pH 3. 
Figure~\ref{fig3}(a) shows kinetic data utilizing HCl, which promotes
the release of elements in larger quantities, while H$_{2}$SO$_{4}$ 
(Figure~\ref{fig3}b) produces a less pronounced release.

\begin{figure*}
\vspace*{2pt}
\includegraphics{fig03}
\caption{\label{fig3}Dissolution curves of K-jarosite according to
Welch et~al.~\cite{43} at pH 3 using (a)~HCl and  (b)~H$_{2}$SO$_{4}$
(Figures reprinted with permission from Elsevier).}
\vspace*{2pt}
\end{figure*}

Elwood Madden et~al.~\cite{109}, Pritchett et~al.~\cite{111}, and Legett 
et~al.~\cite{112} simulated conditions closely resembling those found on
Mars with different pH, temperatures, and salinities for the
dissolution of K-jarosite, and Zahai  et~al.~\cite{113} for Na-jarosite.

Elwood Madden et~al.~\cite{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  $E_{\mathrm{a}} =79$~kJ/mol,
applying the Arrhenius equation. 
\mbox{Moreover,} a Type V behavior was found
where the lowest rate is at pH 3.5, but the rate increases either by
increasing $a_{\mathrm{H}+}$ (pH ${&lt;}$ 3.5) or increasing
$a_{\mathrm{OH}}$- (pH ${&gt;}$ 3.5). A general trend of the  Elwood-Madden
et~al.\ data~\cite{109} can be observed in  Figure~\ref{fig4}. 

\begin{figure}
\includegraphics{fig04}
\caption{\label{fig4}Variation of the dissolution rate (mol/m$^{2}$/s)
of K-jarosite with pH at room temperature (293.15--298.15~K).}
\end{figure}

Kendall et~al.~\cite{44} used the same conditions as Smith 
et~al.~\cite{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 Fe$^{3+}$, its release
proved faster after the first day, exceeding the potassium
concentration. Despite the effect observed by Welch  et~al.~\cite{43},
where dissolution kinetics are faster for chloride environments than
for sulfuric environments, Kendall  et~al.~\cite{44} reached equilibrium
in 14 days of experimentation, much faster than Baron and 
Palmer~\cite{68} or Smith et~al.~\cite{7}, challenging the accuracy or
reliability of the results obtained by Kendall and\break collaborators. 

Similar to Baron and Palmer~\cite{68}, the dissolution rate studies
mentioned above used batch reactors that do not take reactive transport
into account, a critical aspect in environmental 
geochemistry~\cite{46,114}. 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.~\cite{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.~\cite{115} 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.~\cite{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 
\mbox{because} of
differing experimental conditions or mineral 
\mbox{samples.}

As an example,  Figure~\ref{fig4} illustrates the variation in
dissolution rate values of K-jarosite with pH for 
\mbox{reactions} in batch
reactors at room temperature (293.15--298.15~K).

Figure~\ref{fig4} shows the variation in rate even for similar
conditions of pH and temperature, the most 
\mbox{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.~\cite{109} or Kendall et~al.~\cite{44}, have used only the first
two hours, which increases the rate for potassium ions). In 
Table~\ref{tab2}, the different studies in batch reactors (or AFM
study, used by Gasharova et~al.~\cite{45}) are gathered along with their
main parameters and results.

\begin{table*}
\caption{\label{tab2}Kinetic data $(\log r)$ of jarosite dissolution}
\begin{tabular}{ccccccc}
\thead
$a_{\mathrm{H}_2\mathrm{O}}$ &amp; Initial pH &amp; $T$~(K) &amp; Solution &amp; $\log r^{\mathrm{a}}$ &amp;
Authors &amp; Element \\ 
\endthead
\multicolumn{7}{c}{K-jarosite}\vspace*{6pt} \\
1 &amp;  1.97 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$11.18 &amp; Dixon et~al.~\cite{42} &amp;  K$^{+}$ \\  
1 &amp;  1.97 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$11.19 &amp; Dixon et~al.~\cite{42} &amp; Fe$^{3+}$ \\ 
1 &amp;  8 &amp;  296.15 &amp;  Tris &amp;  ${-}$7.95 (0.13) &amp; Kendall et~al.~\cite{44} &amp; K$^{+}$ \\ 
1 &amp;  6 &amp;  296.15 &amp;  Water &amp;  ${-}$9.20 (0.05) &amp;Kendall et~al.~\cite{44} &amp; K$^{+}$ \\ 
1 &amp;  2 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$9.10 (0.00) &amp; Kendall et~al.~\cite{44} &amp; Fe$^{3+}$ \\ 
1 &amp;  2 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$8.42 (0.03) &amp; Kendall et~al.~\cite{44} &amp; K$^{+}$ \\ 
0.75 &amp;  N.D. &amp;  273.15 &amp;  NaCl &amp;  ${-}$9.94(0.09) &amp; Pritchett et~al.~\cite{111} &amp; K$^{+}$ \\ 
0.35 &amp;  N.D. &amp;  273.15 &amp;  CaCl$_{2}$ &amp;  ${-}$11.51(0.05) &amp; Pritchett et~al.~\cite{111} &amp; K$^{+}$ \\ 
0.75 &amp;  N.D. &amp;  263.15 &amp;  NaCl &amp;  ${-}$11.55(0.00) &amp; Pritchett et~al.~\cite{111} &amp; K$^{+}$ \\ 
1 &amp;  6.4 &amp;  280.15 &amp;  Water &amp;  ${-}$9.71 (0.16) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  6.8 &amp;  296.15 &amp;  Water &amp;  ${-}$9.2 (0.05) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  5.9 &amp;  313.15 &amp;  Water &amp;  ${-}$8.59 (0.03) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  6.4 &amp;  323.15 &amp;  Water &amp;  ${-}$7.15 (0.08) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  0.9 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$7.18 (0.06) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  2 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$8.54 (0.06) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  3 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$9.08 (0.05) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  4 &amp;  296.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$9.28 (0.02) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  5.9 &amp;  296.15 &amp;  NaOH &amp;  ${-}$9.22 (0.18) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  6.4 &amp;  296.15 &amp;  NaOH &amp;  ${-}$9.20 (0.10) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  7.1 &amp;  296.15 &amp;  NaOH &amp;  ${-}$8.92 (0.05) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  9.6 &amp;  296.15 &amp;  NaOH &amp;  ${-}$8.71 (0.09) &amp; Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  10.9 &amp;  296.15 &amp;  NaOH &amp;  ${-}$6.60 (0.07) &amp;Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  6.3 &amp;  296.15 &amp;  MES &amp;  ${-}$8.49 (0.07) &amp;  Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  7.9 &amp;  296.15 &amp;  MES &amp;  ${-}$7.95 (0.13) &amp;  Elwood Madden et~al.~\cite{109} &amp; K$^{+}$ \\ 
1 &amp;  4.3 &amp;  296.15 &amp;  Water &amp;  ${-}$9.1 (0.4) &amp;  Elwood Madden et~al.~\cite{110} &amp;  K$^{+}$ \\ 
1 &amp;  4.9 &amp;  296.15 &amp;  Water &amp;  ${-}$8.2 (0.5) &amp;  Elwood Madden et~al.~\cite{110} &amp; K$^{+}$ \\ 
1 &amp;  2 &amp;  293.15 &amp;  HClO$_{4}$ &amp; ${-}$8.8 &amp; Smith et~al.~\cite{7} &amp; K$^{+}$ \\ 
1 &amp;  8 &amp;  293.15 &amp;  Ca(OH)$_{2}$ &amp; ${-}$9.1 &amp;Smith et~al.~\cite{7} &amp; K$^{+}$ \\ 
1 &amp;  3.76 &amp;  298.15 &amp;  Water &amp;  ${-}$11.2 (0.1) &amp; Welch et~al.~\cite{43} &amp;  K$^{+}$ \\ 
1 &amp;  4 &amp;  298.15 &amp;  HCl &amp;  ${-}$11.3 (0.9) &amp; Welch et~al.~\cite{43} &amp; K$^{+}$ \\ 
1 &amp;  3 &amp;  298.15 &amp;  HCl &amp;  ${-}$11.4 (0.4) &amp; Welch et~al.~\cite{43} &amp; K$^{+}$ \\ 
1 &amp;  4 &amp;  298.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$11.3 (0.5) &amp; Welch et~al.~\cite{43} &amp; K$^{+}$ \\ 
1 &amp;  3 &amp;  298.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$11.8 (0.3) &amp; Welch et~al.~\cite{43} &amp; K$^{+}$ \\ 
1 &amp;  5.5 &amp;  298.15 &amp;  Water &amp;  ${-}$6.8 (0.4) &amp; Gasharova et~al.~\cite{45} &amp; K$^{+}$ \\ 
1 &amp;  2 &amp;  298.15 &amp;  HClO$_{4}$ &amp;  ${-}$8.5 (0.03) &amp; Baron and Palmer~\cite{68} &amp; K$^{+}$
\vspace*{6pt}\\ 
 &amp; \multicolumn{6}{c}{Na-jarosite}\vspace*{6pt} \\ 
N.D. &amp;  5.7 &amp;  277.15 &amp;  Water &amp;  ${-}$9.29(0.20) &amp; Zahrai et~al.~\cite{113} &amp; Na$^{+}$ \\ 
N.D. &amp;  8.1 &amp;  295.15 &amp;  Water &amp;  ${-}$8.85(0.06) &amp; Zahrai et~al.~\cite{113} &amp; Na$^{+}$ \\ 
\tabtwosplittabular
N.D. &amp;  1 &amp;  295.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$6.98(0.06) &amp; Zahrai et~al.~\cite{113} &amp; Na$^{+}$ \\ 
N.D. &amp;  2 &amp;  295.15 &amp;  H$_{2}$SO$_{4}$ &amp;  ${-}$6.85(0.35) &amp; Zahrai et~al.~\cite{113} &amp; Na$^{+}$ \\ 
N.D. &amp;  8 &amp;  295.15 &amp;  Tris$+$HCl &amp;  ${-}$7.22(0.16) &amp; Zahrai et~al.~\cite{113} &amp; Na$^{+}$ \\ 
N.D. &amp;  10 &amp;  295.15 &amp;  Tris &amp;  ${-}$6.93(0.06) &amp; Zahrai et~al.~\cite{113} &amp; Na$^{+}$ \\ 
N.D. &amp;  5.7 &amp;  323.15 &amp;  Water &amp;  ${-}$7.89(0.35) &amp; Zahrai et~al.~\cite{113} &amp; Na$^{+}$ \\ 
0.75 &amp;  N.D. &amp;  273.15 &amp;  CaCl$_{2}$ &amp;  ${-}$10.58(0.38) &amp; Pritchett et~al.~\cite{111} &amp; Na$^{+}$ \\ 
0.35 &amp;  N.D. &amp;  273.15 &amp;  CaCl$_{2}$ &amp;  ${-}$11.19(0.17) &amp; Pritchett et~al.~\cite{111} &amp; Na$^{+}$
\botline 
\end{tabular}
\xxtabnote{$^{\mathrm{a}}$The values in parentheses represent the standard
deviation among repeated experiments.}
\xxtabnote{Buffers: Tris  (Tris(Hydroxymethyl)aminomethane) and MES
(2-($N$-morpholino)ethanesulfonic acid).}
\vspace*{-4pt}
\end{table*}

\subsection{Kinetics of jarosite industrial dissolution}\label{sec3.2} 
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~\cite{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 
\mbox{recover} high-value metals such as
silver. Primarily, researchers have used an alkaline medium as
lixiviant in numerous studies regarding jarosites containing 
silver~\cite{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~\cite{125},
arsenic~\cite{126,127,128},  chromium~\cite{129,130}
mercury~\cite{131}, and  thallium~\cite{67}. However, in recent years,
attention has also been given to the use of various acids as a
decomposition medium, as evidenced in  Table~\ref{tab3} where
conditions are displayed, as well as the activation energies and
reaction order for both the induction and progressive periods.

\begin{sidewaystable*}
\vspthree
\caption{\label{tab3}Literature using the shrinking core kinetic model
to describe the dissolution of jarosite-type compounds}
\fontsize{8}{10}\selectfont\tabcolsep1.2pt
\begin{tabular}{cccccccc}  
\thead
Authors using a basic medium &amp; 
Jarosite type &amp; 
Leaching medium &amp; 
$T$~(K) &amp; 
$E_{\mathrm{a},\mathrm{pro}}$~(kJ/mol) &amp; 
$n_{\mathrm{pro}}$ &amp; 
$E_{\mathrm{a},\mathrm{ind}}$~(kJ/mol) &amp; 
$n_{\mathrm{ind}}$\\ 
\endthead
\morerows{1}{Roca et~al.~\cite{117}} &amp; 
\morerows{1}{[Ag$_{0.78}$(H$_{3}$O)$_{0.22}$]Fe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$}  &amp; 
NaOH &amp; 323.15 &amp; 42.00 &amp; 0.50 &amp; 71.00 &amp; 0.50 \\
 &amp; &amp; Ca(OH)$_{2}$ &amp; 323.15 &amp; 42.00 &amp; 0.00 &amp; 71.00 &amp; 0.00\vspace*{4pt}\\ 
\morerows{1}{Pati\~{n}o et~al.~\cite{119}}  &amp; 
\morerows{1}{Pb$_{0.32}$Ag$_{0.011}$(H$_{3}$O)$_{0.35}$]Fe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$} &amp; 
NaOH &amp; 323.15 &amp; 53.00 &amp; 1.00 &amp; N.D. &amp; N.D. \\
 &amp; &amp;  Ca(OH)$_{2}$ &amp; 323.15 &amp; 47.00 &amp; 0.00 &amp; N.D. &amp; N.D.\vspace*{4pt}\\ 
\morerows{1}{Pati\~{n}o et~al.~\cite{120}}  &amp; 
\morerows{1}{[Na$_{0.675}$Ag$_{0.005}$(H$_{3}$O)$_{0.32}$]Fe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$} &amp; 
NaOH &amp; 308.15 &amp; 96.00 &amp; 0.40 &amp; 84.00 &amp; 0.90 \\
 &amp; &amp; Ca(OH)$_{2}$ &amp; 308.15 &amp; 47.00 &amp; 0.50 &amp; 84.00 &amp; 0.70\vspace*{4pt}\\ 
\morerows{1}{Cruells et~al.~\cite{122}}  &amp; 
\morerows{1}{[K$_{0.91}$Ag$_{0.007}$(H$_{3}$O)$_{0.83}$]Fe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$} &amp; 
NaOH &amp; 303.15 &amp; 43.00 &amp; 0.60 &amp; 86.00 &amp; 0.70\\
 &amp; &amp; Ca(OH)$_2$ &amp; 303.15 &amp; 80.00 &amp; 0.50 &amp; 36.00 &amp; 0.20\vspace*{4pt}\\ 
\morerows{1}{Salinas et~al.~\cite{123}}  &amp;
\morerows{1}{[Ag$_{0.001}$Na$_{0.07}$K$_{0.02}$Pb$_{0.007}$(NH$_{4}$)$_{0.59}$(H$_{3}$O)$_{0.31}$]Fe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$} &amp; 
NaOH &amp; 308.15 &amp; 77.00 &amp; 1.10 &amp; 88.00 &amp; 1.40\\
 &amp; &amp; NaOH/NaCN &amp; 308.15 &amp; 46.00 &amp; 0.80 &amp; 154.00 &amp; 3.20\vspace*{4pt}\\ 
\morerows{1}{Pati\~{n}o et~al.~\cite{121}} &amp; 
\morerows{1}{[(NH$_{4}$)$_{0.71}$Ag$_{0.040}$(H$_{3}$O)$_{0.24}$]Fe$_{2.85}$(SO$_{4}$)$_{2}$(OH)$_{5.50}$}  &amp; 
Ca(OH)$_{2}$ &amp; 308.15 &amp; 70.00 &amp; 0.40 &amp; N.D. &amp; N.D.\\
 &amp; &amp; Ca(OH)$_{2}$/NaCN &amp; 308.15 &amp; 29.00 &amp; 0.65 &amp; N.D. &amp; N.D.\vspace*{4pt}\\ 
Roca et~al.~\cite{118} &amp; 
[(NH$_{4}$)$_{0.71}$Ag$_{0.040}$(H$_{3}$O)$_{0.25}$]Fe$_{2.85}$(SO$_{4}$)$_{2}$(OH)$_{5.50}$ &amp; 
NaOH &amp; \morerows{1}{303.15/323.15} &amp; 60.00 &amp; 0.60 &amp; 80.00 &amp; 0.75\\
 &amp; &amp; NaOH/NaCN &amp; &amp; 52.00 &amp; 0.00 &amp; 54.00 &amp; 0.50\vspace*{4pt}\\ 
\morerows{1}{Reyes et~al.~\cite{132}} &amp; 
\morerows{1}{[Na$_{0.87}$(H$_{3}$O)$_{0.13}$]Fe$_{2.5}$(SO$_{4}$)$_{1.95}$(AsO$_{4}$)$_{0.05}$[(OH)$_{4.45}$(H$_{2}$O)$_{1.55}$]} &amp; 
NaOH &amp; 303.15 &amp; 57.10 &amp; 0.70 &amp; 120.40 &amp; 1.00 \\ 
 &amp; &amp; Ca(OH)$_{2}$ &amp; 303.15 &amp; 48.60 &amp; 1.51 &amp; 79.40 &amp; 0/1\vspace*{4pt}\\ 
Perez-Labra et~al.~\cite{125} &amp; 
Rb$_{0.9432}$Fe$_{3}$(SO$_{4}$)$_{2.1245}$(OH)$_{6}$ &amp; 
Ca(OH)$_{2}$ &amp; 298.15 &amp; 98.70 &amp; 0.43 &amp; N.D. &amp; N.D. \\ 
\morerows{1}{Flores et~al.~\cite{126}} &amp; 
\morerows{1}{[K$_{0.75}$(H$_{3}$O)$_{0.25}$]Fe$_{1.84}$(SO$_{4}$)$_{1.82}$(AsO$_{4}$)$_{0.18}$[(OH)$_{2.34}$(H$_{2}$O)$_{3.66}$]} &amp; 
NaOH &amp; &amp; 60.30 &amp; 1.86 &amp; 84.70 &amp; 2.65 \\
 &amp; &amp; Ca(OH)$_{2}$ &amp; &amp; 74.40 &amp; 1.14 &amp; 88.30 &amp; 0.24 \\ 
\morerows{1}{Pati\~{n}o et~al.~\cite{127}} &amp; 
\morerows{1}{[Na$_{0.87}$(H$_{3}$O)$_{0.13}$]Fe$_{2.50}$(SO$_{4}$)$_{1.95}$(AsO$_{4}$)$_{0.05}$[(OH)$_{4.45}$(H$_{2}$O)$_{1.55}$]} &amp; 
NaOH &amp; 303.15 &amp; 57.11 &amp; 0.746/0 &amp; 118.11 &amp; 1.06\\
 &amp; &amp; Ca(OH)$_{2}$ &amp; 303.15 &amp; 48.22 &amp; 1.56 &amp; 73.50 &amp; 1.03\vspace*{4pt}\\ 
\morerows{1}{Pati\~{n}o et~al.~\cite{133}} &amp;
\morerows{1}{KFe$_{3}$(SO$_{4}$)$_{1.82}$(AsO$_{4}$)$_{0.18}$(OH)$_{5.82}$} &amp; 
NaOH &amp; 303.15 &amp; 60.30 &amp; 1.86 &amp; 84.60 &amp; 0.23/2.65\\
 &amp; &amp; Ca(OH)$_{2}$ &amp; 303.15 &amp; 74.40 &amp; 1.15 &amp; 88.20 &amp; 0.24/0\vspace*{4pt}\\
M\'{e}ndez et~al.~\cite{128} &amp; 
[(NH$_{4}$)$_{0.23}$K$_{0.77}$]Fe$_{2.47}$(SO$_{4}$)$_{1.69}$(AsO$_{4}$)$_{0.31}$[(OH)$_{4.01}$(H$_{2}$O)$_{1.59}$] &amp; 
NaOH &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. \\ 
Islas et~al.~\cite{141} &amp; 
Tl$_{0.86}$(H$_{3}$O)$_{0.25}$Fe$_{3.11}$(SO$_{4}$)$_{2.11}$[(OH)$_{6.11}$(H$_{2}$O)$_{0.44}$] &amp; 
NaOH &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. \\
Mireles et~al.~\cite{129} &amp; 
[K$_{0.86}$(H$_{3}$O)$_{0.14}$]Fe$_{2.67}$[(SO$_{4}$)$_{1.23}$(CrO$_{4}$)$_{0.77}$][(OH)$_{5.01}$(H$_{2}$O)$_{0.99}$] &amp; 
Ca(OH)$_{2}$ &amp; &amp; 51.56 &amp; 1.99 &amp; 63.75 &amp; 0.67 \\\splittabular   
\morerows{1}{Gonz\'{a}lez-Ibarra et~al.~\cite{134}}  &amp; 
\morerows{1}{AgFe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$}  &amp; 
NaOH &amp; 303.15 &amp; 40.42 &amp; N.D. &amp; N.D. &amp; N.D.\\
 &amp; &amp; Ca(OH)$_2$ &amp; 303.15 &amp; 21.72 &amp; N.D. &amp; N.D. &amp; N.D.\vspace*{4pt}\\ 
Reyes et~al.~\cite{130} &amp; 
[K. (H$_{3}$O)]Fe$_{3}$[(SO$_{4}$)$_{2-x}$(CrO$_{4}$)$_{x}$][(OH)$_{6}$(H$_{2}$O)$_{n}$] &amp; 
NaOH &amp; &amp; 75.70 &amp; 1.10 &amp; 82.70 &amp; 1.10 \\ 
\morerows{1}{Ordonez et~al.~\cite{131}} &amp; 
Hg$_{0.39}$(H$_{3}$O)$_{0.22}$Fe$_{2.71}$(SO$_{4}$)$_{2.17}$[(OH)$_{4.79}$(H$_{2}$O)$_{2.09}$] &amp; 
NaOH &amp; 303.15 &amp; 56.90 &amp; 0.99/0 &amp; 81.70 &amp; 0.52/1.48 \\
 &amp; 
Tl$_{0.86}$(H$_{3}$O)$_{0.14}$Fe$_{3.11}$(SO$_{4}$)$_{2.11}$(OH)$_{6.11}$(H$_{2}$O)$_{0.44}$ &amp; 
NaOH &amp; 303.15 &amp; 91.90 &amp; 1.10 &amp; 101.60 &amp; 1.10 \\
\morerows{2}{Islas et~al.~\cite{67}} &amp; 
Hg$_{0.39}$(H$_{3}$O)$_{0.22}$Fe$_{2.71}$(SO$_{4}$)$_{2.17}$(OH)$_{4.79}$(H$_{2}$O)$_{2.09}$ &amp; 
NaOH &amp; 303.15 &amp; 56.90 &amp; 0.99 &amp; 81.70 &amp; 1.48\vspace*{4pt}\\
 &amp; Pb$_{0.48}$(H$_{3}$O)$_{0.04}$Fe$_{2.97}$(SO$_{4}$)$_{1.93}$(OH)$_{6.05}$(H$_{2}$O)$_{0.6}$ &amp; 
NaOH &amp; 303.15 &amp; 85.20 &amp; 0.74 &amp; 66.40 &amp; 1.19 \\
 &amp; Pb$_{0.39}$(H$_{3}$O)$_{0.22}$Fe$_{2.89}$(S$_{0.92}$As$_{0.08}$O$_{4}$)$_{2}$(OH)$_{5.51}$(H$_{2}$O)$_{0.49}$ &amp; 
NaOH &amp; 303.15 &amp; 66.10 &amp; 0.99 &amp; 96.90 &amp; 2.40\vspace*{4pt}\\
\morerows{1}{Islas et~al.~\cite{135}} &amp; 
\morerows{1}{(Na$_{0.96}$Ag$_{0.04}$)Fe$_{2.82}$(SO$_{4}$)$_{2.01}$(OH)$_{5.36}$(H$_{2}$O)$_{0.86}$} &amp; 
NaOH/CN$^{-}$ &amp; 303.15 &amp; 46.30 &amp; 0.61 &amp; 69.80 &amp; 0.66 \\
 &amp; &amp; $\mathrm{NaOH}/\mathrm{S}_{2}\mathrm{O}_{3}^{2-}$ &amp; 
303.15 &amp; 53.30 &amp; 0.51 &amp; 70.10 &amp; 0.51\\
 &amp; &amp; NaOH/SCN$^-$ &amp; 303.15 &amp; 45.10 &amp; 0.60 &amp; 71.50 &amp; 0.59 \\ 
Flores et~al.~\cite{136} &amp; 
\parbox[t]{12.5pc}{\centering
[(NH$_{4}$)$_{0.72}$Na$_{0.06}$(H$_{3}$O)$_{0.21}$]Fe$_{3(2.52)}$(SO$_{4}$)$_{1.85}${\ubreak}(AsO$_{4}$)0.15[(OH)$_{4.41}$(H$_{2}$O)$_{1.59}$]} &amp; 
NaOH/CN &amp; 303.15 &amp; 51.60 &amp; 0.94 &amp; 65.00 &amp; 1.22\vspace*{6pt} \\\xbotline 
Authors using an acidic medium &amp; 
Jarosite type &amp; 
Leaching medium &amp; 
$T$~(K) &amp; 
$E_{\mathrm{a},\mathrm{pro}}$~(kJ/mol) &amp; 
$n_{\mathrm{pro}}$ &amp; 
$E_{\mathrm{a},\mathrm{ind}}$~(kJ/mol) &amp; 
$n_{\mathrm{ind}}$ \\\xbotline
Calla-Choque~\cite{137} &amp; 
Ag/Cu &amp; Thiourea &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. \\ 
Reyes et~al.~\cite{130} &amp; 
[K. (H$_{3}$O)]Fe$_{3}$[(SO$_{4}$)$_{2-x}$(CrO$_{4}$)$_{x}$][(OH)$_{6}$(H$_{2}$O)$_{n}$] &amp; 
HCl &amp; 303.15/323.15 &amp; 68.03 &amp; 1.2 &amp; 109.4 &amp; 1.05/0.47 \\ 
\morerows{1}{Reyes et~al.~\cite{61}} &amp; 
[(M. H$_{3}$O)]Fe$_{3}$[(S$_{1-x}$As$_{x}$O$_{4}$)$_{2}$][(OH)$_{6}$(H$_{2}$O)$_{n}$] &amp; 
H$_{2}$SO$_{4}$ &amp; 303.15/323.15 &amp; 77.9 &amp; 1.2 &amp; 66.8 &amp; 1 \\ 
&amp; 
Tl$_{0.86}$(H$_{3}$O)$_{0.14}$Fe$_{3.11}$(SO$_{4}$)$_{2.11}$(OH)$_{6.11}$(H$_{2}$O)$_{0.44}$ &amp; 
H$_{2}$SO$_{4}$ &amp; 303.15 &amp; 58.4 &amp; 1.07 &amp; 55.9 &amp; 0.9 \\
\morerows{2}{Islas et~al.~\cite{67}} 
 &amp; Hg$_{0.39}$(H$_{3}$O)$_{0.22}$Fe$_{2.71}$(SO$_{4}$)$_{2.17}$(OH)$_{4.79}$(H$_{2}$O)$_{2.09}$ &amp; 
H$_{2}$SO$_{4}$ &amp; 303.15 &amp; 83.5 &amp; 0.98 &amp; 61.2 &amp; 0.98 \\
 &amp; Pb$_{0.48}$(H$_{3}$O)$_{0.04}$Fe$_{2.97}$(SO$_{4}$)$_{1.93}$(OH)$_{6.05}$(H$_{2}$O)$_{0.6}$ &amp; 
H$_{2}$SO$_{4}$ &amp; 303.15 &amp; 63.3 &amp; 1.01 &amp; 55.9 &amp; 1.1 \\
 &amp; Pb$_{0.39}$(H$_{3}$O)$_{0.22}$Fe$_{2.89}$(S$_{0.92}$As$_{0.08}$O$_{4}$)$_{2}$(OH)$_{5.51}$(H$_{2}$O)$_{0.49}$ &amp; 
H$_{2}$SO$_{4}$ &amp; 303.15 &amp; 77.3 &amp; 1.05 &amp; 35.1 &amp; 0.99 \vspace*{6pt}\\ 
Calla-Choque and Lapidus~\cite{124} &amp; Ag &amp; Thiourea et Oxalate &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. &amp; N.D. \\ 
Calla-Choque and Lapidus~\cite{138} &amp; 
NaFe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$, Ag &amp; 
Thiourea et Oxalate &amp; 303.15 &amp; 44 &amp; N.D. &amp; N.D. &amp; N.D. \\ 
Nolasco et~al.~\cite{140} &amp; 
[Na$_{0.70}$(H$_{3}$O)$_{0.30}$][Fe$_{2.95}$Cu$_{0.042}$](SO$_{4}$)$_{2.25}$[(OH)$_{5.46}$(H$_{2}$O)$_{0.54}$] &amp; 
H$_{2}$SO$_{4}$ &amp; 323.15 &amp; 57,9 &amp; 1,1 &amp; N.D. &amp; N.D.
\botline 
\end{tabular}
\end{sidewaystable*}

\subsubsection{Dissolution through alkaline medium}\label{sec3.2.1} 
One of the first team to study this dissolution was Roca 
et~al.~\cite{117}, who utilized NaOH and Ca(OH)$_{2}$ as basic medium,
temperatures ranging from 298.15 to 333.15~K, and NaCN concentrations
from 5 to 30~mM. Subsequently, Pati\~{n}o et~al.~\cite{119} investigated
the cyanidation--decomposition of Pb-jarosite containing silver, also
using both alkaline media because this type of jarosite is commonly
found in Nature and industrial processes.

Later, Pati\~{n}o et~al.~\cite{120} examined the decomposition of
silver-containing Na-jarosite with cyanidation. They obtained
activation energies of  96~kJ/mol for the NaOH medium and 39.6~kJ/mol
for the Ca(OH)$_{2}$ medium. Additionally, they found that the addition
of salts such as NaCl or Na$_{2}$SO$_{4}$ increased the reaction rate
for the experimental conditions used.

Cruells et~al.~\cite{122} studied the decomposition of silver-containing
K-jarosite using the same media. They obtained activation energies of 
86~kJ/mol for the NaOH medium and 36~kJ/mol for the Ca(OH)$_{2}$
medium. When coupling with cyanidation (CN$^-$), the activation energy
decreased to  43~kJ/mol for NaOH and increased to 80~kJ/mol with
Ca(OH)$_{2}$.

Salinas et~al.~\cite{123} investigated alkaline (NaOH) decomposition
coupled with cyanidation of NH$_{4}$-jarosites containing silver. When
only NaOH was utilized, the activation energy rose to  77~kJ/mol, while
coupling with CN$^-$ reduced the energy to\break 46~kJ/mol.

Pati\~{n}o et~al.~\cite{121} studied silver-containing NH$_{4}$-jarosite
using Ca(OH)$_{2}$. They found an activation 
\mbox{energy} of  70~kJ/mol using
only the base and  29~kJ/mol adding CN$^-$. The decomposition rates
were higher than those of jarosites produced during the zinc
hydrometallurgical process. Roca  et~al.~\cite{118} used the same type
of jarosite as Pati\~{n}o  et~al.~\cite{121} and studied its
decomposition in a cyanide and alkaline environment using NaOH instead
of Ca(OH)$_{2}$. The activation energy for the induction period was
higher than for the progressive conversion period. The activation
energies for the progressive period were  60~kJ/mol without cyanide and
52~kJ/mol with cyanide. They confirmed a trend of reduced activation
energy using cyanidation, as already observed by 
\mbox{Salinas 
et~al.~\cite{123}.}

Gonz\'{a}lez-Ibarra et~al.~\cite{134} investigated the kinetic
decomposition of an industrial jarosite using both NaOH and
Ca(OH)$_{2}$ as well as cyanidation. They succeeded in increasing
silver recovery by coupling alkalinity with cyanidation. When using
NaOH, kinetics were controlled by the chemical reaction stage, while
using Ca(OH)$_{2}$, diffusion through a passivation layer of CaCO$_{3}$
played a crucial role, and the activation energies were 40.42 and
21.72~kJ/mol,
\mbox{respectively.}

Islas et~al.~\cite{135} studied the leaching of silver in jarosites
using non-cyanide complexing agents such as thiosulfate
$(\mathrm{S}_{2}\mathrm{O}_{3}^{2-})$ and thiocyanate (SCN$^-$), as
they are less harmful to human health. The proposed overall kinetic
equations allow optimizing jarosite leaching using these complexing
agents as alternatives to cyanide.

Reyes et~al.~\cite{132} studied the reactivity of 
Na-jarosites
containing arsenic. They found 
\mbox{activation} energies of  57.11~kJ/mol for
NaOH and 48.61~kJ/mol for Ca(OH)$_{2}$.  Flores et~al.~\cite{126}
investigated the decomposition of K-jarosites containing arsenic using
NaOH and Ca(OH)$_{2}$ media, with activation energies of 60.3~kJ/mol
and  74.4~kJ/mol, respectively. This difference in activation energy
between K- and Na-jarosite may be related to the fact that K-jarosite
is more stable during dissolution than Na-jarosite.

Pati\~{n}o et~al.~\cite{127,133} examined the decomposition of
Na--As-jarosite in NaOH and Ca(OH)$_{2}$ media, the former through a
design of experiments approach and the latter using only the SCKM.

The first study reported the presence of unreacted As-jarosite and the
formation of halos of Fe(OH)$_{3}$ particles with adsorbed AsO$_{4}$.
Temperature and OH$^-$ concentration, as well as their combination,
were the most influential parameters on the kinetics. The second study
found activation energies of 60.3~kJ/mol and 74.4~kJ/mol for NaOH and
Ca(OH)$_{2}$ media, respectively. This corroborates the claim that
activation energies are lower with NaOH than with Ca(OH)$_{2}$,
indicating easier dissolution when NaOH is used.\looseness=-1

M\'{e}ndez et~al.~\cite{128} studied NH$_{4}$--K-jarosite containing
arsenic, and Flores  et~al.~\cite{136} investigated the decomposition of
arsenic-containing NH$_{4}$--Na-jarosite, both using NaOH medium. The
\mbox{controlling} mechanism was the chemical reaction, and there was the
formation of Fe(OH)$_{3}$ with adsorbed arsenate. Flores and
collaborators emphasized that using this type of jarosite is an
alternative to arsenic removal from water because, even after
decomposition, arsenic is retained in the residual solid.

P\'{e}rez-Labra et~al.~\cite{125} synthesized rubidium jarosite and
studied its decomposition in an alkaline solution containing five
different concentrations of Ca(OH)$_{2}$. The obtained solid residues
incorporated calcium and formed goethite. The kinetic analysis could be
modeled by the SCKM with the chemical reaction control stage, with an
activation energy of  98.70~kJ/mol.

Islas et~al.~\cite{141} studied the dissolution of thallium jarosite
using NaOH, leading to the formation of an amorphous solid around the
particle, likely containing hydroxide and thallium. The decomposition
of this type of jarosite seems to be described by a constant particle
size model, controlled by the chemical reaction.

Mireles et~al.~\cite{129} analyzed the decomposition kinetics of
chromium-containing jarosite in a Ca(OH)$_{2}$ medium. Experimental
data show that the reaction follows the SCKM with chemical reaction
control. The activation energy was 63.75~kJ/mol for the induction
period and 51.56~kJ/mol for the progressive period, and the reaction
results in the formation of Fe(OH)$_{3}$ residues. 

Ordo\~{n}ez et~al.~\cite{131} dedicated their study to the decomposition
of mercury jarosite in a NaOH medium. The activation energy was 
81.70~kJ/mol for the induction period and 56.90~kJ/mol for the
progressive period.

This section illustrates that the dissolution of various types of
jarosite has been extensively studied, mainly utilizing NaOH and
Ca(OH)$_{2}$ as basic mediums. A trend was observed where reactions
occurring in NaOH medium were less energy-dependent compared to those
in Ca(OH)$_{2}$ medium. Cyanidation was also employed to enhance the
dissolution rate in basic medium. Moreover, other complexing agents
(such as $\mathrm{SCN}^-/\mathrm{S}_{2}\mathrm{O}_{3}^{2-}$) were
studied as alternatives to CN$^-$. Passivation layers and the formation
of other mineral phases were identified depending on the basic medium
and conditions employed. The SCKM revealed that the induction period
exhibits a higher dependency on temperature (high $E_{\mathrm{a}}$)
than the progressive period.

\subsubsection{Dissolution through acidic medium}\label{sec3.2.2} 
More recently, acid dissolution has also attracted attention, given
that jarosite is an acidic compound and requires a large amount of base
for its dissolution process in a basic environment. 
Calla Choque 
\mbox{et~al.~\cite{137}} were the first researchers to apply the SCKM to the
acid decomposition of jarosite, coupling 
\mbox{leaching} with thiourea as a
complexing agent to study the effect of different cations (Cu$^{2+}$,
Fe$^{3+}$, Zn$^{2+}$, Fe$^{2+}$) in this system for silver recovery.
They \mbox{reported} that acid decomposition would be an interesting
alternative as jarosites are naturally acidic, meaning that process
costs would be reduced.

Subsequently, Reyes et~al.~\cite{130} studied the dissolution rates of
K--Cr(VI)-jarosite in acidic (HCl) and basic (NaOH) environments to
investigate the mobility of chromium under more extreme conditions
(mostly at pH ${&lt;}$ 1 and $T&gt;303.15$~K). They suggested studying
reactions at an intermediate pH (2 ${\leq}$ pH ${\leq}$ 10) because
unreacted solids remain, and calculations must consider equilibrium.

While Reyes et~al.~\cite{61} studied the dissolution kinetics of Na--As-,
K--As-, and NH$_{4}$--As-jarosites in sulfuric acid to determine the
factors influencing the recovery of high-value metals by varying pH,
temperature, particle size, and substitution at \textbf{A}- and
\textbf{C}-sites of jarosite-type compounds. Sulfuric acid was chosen
due to its lower cost and better compatibility with other reducing
agents. Similar to their previous research, Reyes and collaborators
used more extreme conditions to calculate condition parameters (mostly
pH ${&lt;}$ 1 and  $T&gt;303.15$~K) to obtain the final model. For pH ${&lt;}$
1, they reported that the decomposition process follows the SCKM
controlled by the chemical reaction stage. The authors demonstrated
that, for the acidity levels used, ion substitution does not influence
the dissolution rate.

Islas et~al.~\cite{67} studied the dissolution of several jarosites
containing potentially hazardous elements (Pb, As, Tl, Cr) in different
conditions using the SCKM in basic (NaOH) and acidic (H$_{2}$SO$_{4}$)
media. They demonstrated that the kinetics-controlling step is the
chemical reaction at extreme pH levels (pH ${&lt;}$ 1), and mass transfer
in the liquid phase when pH deviates towards less extreme values (1
${&lt;}$ pH ${&lt;}$ 1.68).

Calla-Choque and Lapidus~\cite{138} undertook a study to analyze the
effects of pH, temperature, and solid/liquid ratio on silver recovery
from jarosite, employing complexing agents such as thiourea and
oxalate. Their research highlights the use of oxalate in stabilizing
thiourea within an acidic environment. The decomposition of jarosite
releases powerful oxidants, mainly iron, inducing chemical instability
and thiourea oxidation. Results showed improved recovery with a
temperature increase from 323.15 to 333.15~K, pH reduction, and
adjusting the solid/liquid ratio. In addition, when oxalate was used,
thiourea degradation was minimized, \mbox{consequently} decreasing the
temperature required for jarosite dissolution.

Calla-Choque and Lapidus~\cite{124} studied the dissolution rate with a
non-homogeneous particle size system, unlike the approach in Reyes 
et~al.~\cite{61}, which used specific particle sizes, in the presence of
oxalate and thiourea at different pH and temperatures. They modified
the SCKM to account for particle size variation. According to Gbor and 
Jia~\cite{139}, particle size distribution is of fundamental importance
when evaluating kinetics with the SCKM. However, it is often
disregarded, leading to erroneous conclusions about the control regime
of the process.

Nolasco et~al.~\cite{140} conducted jarosite dissolution in acidic
conditions and investigated how the incorporation of divalent cations,
such as Cu$^{2+}$ and Fe$^{2+}$, into jarosite structure affects the
reaction rate. The SKCM was able to satisfactorily describe the
leaching behavior of Na-jarosite containing copper under acidic
conditions. According to authors, the nature of the central atom
bonding (Fe--OH) within the jarosite structure plays an important role
in acid dissolution. They suggested that the complete replacement of
\textbf{B}-site by other trivalent cations could likely alter the
dissolution rate. 

Studies on dissolution rates in acidic medium are relatively recent in
the literature but have reached a higher level of scientific maturity,
partly due to the groundwork laid by studies conducted in basic
mediums. Temperature, pH, particle size, ion substitution into jarosite
structure, utilization of complexing agents, and use of different acids
are the main parameters investigated. These studies reveal certain
trends, such as the chemical kinetic control of the progressive period
when extreme pH levels (${&lt;}$1.5) and temperatures  (${&gt;}$323~K) are
employed. However, further experiments are required to elucidate the
behavior of the SCKM under conditions of less acidic pH (${&gt;}$1.5) and
ambient temperature (${\sim}$295~K) during the experiments.

\subsection{Natural and synthetic jarosites}\label{sec3.3} 
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. 

\subsubsection{Synthetic jarosites}\label{sec3.3.1} 
The article by Driscoll and Leinze~\cite{142} is one of the most
frequently used as a guideline for the synthesis of different types of
jarosites within the scope of their work at the United States
Geological Survey (USGS). Their method is based on the synthesis of
K-jarosite as proposed by Baron and  Palmer~\cite{68}. They mix 17.2~g
of hydrated ferric sulfate 
$(\mathrm{Fe}_{2}(\mathrm{SO}_{4})_{3}\cdot n\mathrm{H}_{2}\mathrm{O})$
and 5.6~g of potassium hydroxide (KOH) dissolved in 100~ml of
$18~\mathrm{M}\Omega{\cdot}\mathrm{cm}$ ultrapure water. The solution
is prepared on a hot plate at 368.15~K with stirring. Moderate boiling
is maintained for  4~h; then, the solution is removed from the hot
plate, and after decantation, the liquid phase is rinsed with deionized
water. The washings are discarded, and the remaining solid is placed in
a drying oven at 383.15~K for  24~h. After drying, the substance is
recovered, in the form of a yellow-brown precipitate. Smith 
et~al.~\cite{7}, Elwood Madden et~al.~\cite{109}, Zahrai
et~al.~\cite{113}, Pritchett  et~al.~\cite{111}, Kendall et~al.~\cite{44},
and Dixon  et~al.~\cite{42} have used the method validated by Driscoll
and  Leinze~\cite{142}.

As explained by Baron and Palmer~\cite{68}, synthetic jarosite exhibits
a potassium deficiency due to the substitution by hydronium ions
$(\mathrm{H}_{3}\mathrm{O}^{+})$ in the alkali ion position. Various factors control
this hydronium substitution, including pH and the activity of K$^{+}$.
To minimize this substitution and synthesize a product with a
composition close to the ideal, it is necessary to add an excess amount
of KOH, thereby increasing the activity of K$^{+}$ and simultaneously
reducing the concentration of H$^{+}$.

Other authors have employed different methods for the production of
K-jarosite. Gasharova  et~al.~\cite{45} based their synthesis on the one
suggested by Fairchild  et~al.~\cite{143}, in which crystallization at
423.15~K and 0.5~MPa occurs from an aqueous solution of K$_{2}$SO$_{4}$
and Fe$_{2}$(SO$_{4}$)$_{3}$ in 1 N sulfuric acid. Jones 
et~al.~\cite{144} synthesized jarosite according to the method of
Dutrizac  et~al.~\cite{145}, Dutrizac and Jambor~\cite{27}, and Crabbe 
et~al.~\cite{146},
dissolving 6.4~g 
$\mathrm{Fe}_{2}(\mathrm{SO}_{4})_{3}\cdot n\mathrm{H}_{2}\mathrm{O}$
and 24~g HNO$_{3}$ in 800~g deionized water, then adding 0.8~ml
H$_{2}$SO$_{4}$ and heating at 363.15~K for  24~h. Brand 
et~al.~\cite{147} used 1.74~g of
$\mathrm{Fe}_{2}(\mathrm{SO}_{4})_{3}\cdot9\mathrm{H}_{2}\mathrm{O}$ and
0.70~g of K$_{2}$SO$_{4}$ diluted in deionized water, reacting at
368.15~K for  5~h, then filtering through a syringe filter, obtaining a
calculated composition of
K$_{0.9}$(H$_{3}$O)$_{0.1}$Fe$_{2.39}$(SO$_{4}$)$_{2}$OH$_{4.17}$(H$_{2}$O)$_{1.83}$. Reyes 
et~al.~\cite{61} synthesized K-jarosite using the method of
Pati\~{n}o~\cite{127}, which was based on that of Dutrizac 
et~al.~\cite{148} for Na-jarosite. In this method, 1 L solution was used
containing 0.3~mol/L of 
$\mathrm{Fe}_{2}(\mathrm{SO}_{4})_{3}\cdot n\mathrm{H}_{2}\mathrm{O}$,
0.01~mol/L of H$_{2}$SO$_{4}$, 0.027~mol/L of 
$\mathrm{Na}_{2}\mathrm{HAsO}_{4}\cdot7\mathrm{H}_{2}\mathrm{O}$, and
0.2~mol/L of Z$_{2}$SO$_{4}$, where Z represents monovalent alkaline
species Na$^{+}$, K$^{+}$,  $\mathrm{NH}_{4}^{+}$. Jarosite was formed
by dissolving the reagents, adjusting the pH with H$_{2}$SO$_{4}$ or
MgCO$_{3}$, then heating a 2 L reactor for  24~h at  370~K. At the end
of the reaction, slurries were filtered, washed in hot water and dried
at 383~K.

\subsubsection{Natural jarosites}\label{sec3.3.2} 
Welch  et~al.~\cite{43} used natural jarosite samples [(K,
H$_{3}$O)Fe$_{3}$(SO$_{4}$)$_{2}$(OH)$_{6}$] from a degraded acidic
sulfate soils site on the north coast of New South Wales, Australia, in
order to determine experimentally the dissolution kinetics. Similarly,
Qian  et~al.~\cite{115} and Trueman et~al.~\cite{46} also used natural
jarosite. Calla-Choque and Lapidus~\cite{124} used jarosite from a zinc
leaching plant in Mexico.

It was evident that Welch and colleagues' jarosite dissolved very
slowly compared to synthetically prepared material in the study by
Smith  et~al.~\cite{7}. Indeed, the solids presented different
morphologies:  euhedral~\cite{45} for natural jarosite or 
anhedral~\cite{7} for synthesized jarosite. The reactivity differences
between synthetic mineral and naturally weathered samples have been
documented for  sulfates~\cite{149,150} and silicate 
minerals~\cite{151,152,153}, as indicated by
\mbox{Welch et~al.~\cite{43}.}

According to White and Brantley~\cite{153}, when studying silicate
dissolution,
the time dependency of silicate weathering can be
discussed in intrinsic terms, as the surface area increases with the
duration of weathering, accumulation of leached layers, and secondary
precipitates. Alternatively, it can be discussed in extrinsic terms, as
low permeability, high mineral/fluid ratios and increased solute
concentrations, produce weathering reactions close to thermodynamic
equilibrium under field conditions compared to highly unsaturated
conditions during experimental reaction. These factors could explain
the differences between dissolution rates of natural and synthesized
samples.

\section{Precipitation kinetics of jarosite-type\newline 
compounds}\label{sec4}

\subsection{Studies in abiotic environments at high\newline
temperatures}\label{sec4.1}
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~\cite{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
\mbox{environments.}

Dutrizac and Kaiman~\cite{154} initially worked on the synthesis of
rubidium and thallium jarosite. Subsequently, Dutrizac 
et~al.~\cite{145} investigated the conditions affecting the formation of
Pb-jarosite, while  Dutrizac~\cite{155} studied 
their formation in 
\mbox{hydrochloric} acid. Dutrizac and  Dinardo~\cite{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~\cite{157}, mercury~\cite{158}, cesium and 
lithium~\cite{159}, silver~\cite{160}, arsenic at 370.15~K from sulfate
and hydrochloric acid  medium~\cite{161} as well as at
423.15~K~\cite{148}, cadmium~\cite{162}, thiocyanate and
cyanate~\cite{163}, gallium~\cite{164}, lanthanides~\cite{165}, 
vanadium~\cite{166}, rare earth elements~\cite{167}, thallium
(III)~\cite{168}, chromium (III)~\cite{169,170},  
germanium and
tin~\cite{171}, \mbox{alkaline-earth} elements (beryllium and 
radium)~\cite{172}, halides (fluoride, chloride, bromide, and 
iodide)~\cite{173}, scandium, yttrium, and uranium~\cite{174}, selenium
(VI) and selenium  (IV)~\cite{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: $\mathrm{K} &gt; \mathrm{NH}_{4}^{+} &gt; \mathrm{Na}$. 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~\cite{62}.

Limpo et~al.~\cite{176} studied the influence of acidity, Fe$^{3+}$
concentration, ammonia concentration, and seeding on the precipitation
rate of Fe$^{3+}$ as NH$_{4}$-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 
$\mathrm{Fe}_{2}(\mathrm{OH})_{4}^{2+}$ 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~\cite{177} conducted batch experiments to study the effect of
seeding on the initial precipitation rate of NH$_{4}$-jarosite,
noticing a linear increase with the growing quantity of seeds. This
trend proved valid for both high and low concentrations of Fe$^{3+}$,
$\mathrm{NH}_{4}^{+}$, and 
$\mathrm{Mn}_{2}^{+}$.

Dutrizac~\cite{178} investigated the effect of seeding on the
precipitation of NH$_{4}$-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(SO$_{4}$)$_{1.5}$, 0.30~M Na$_{2}$SO$_{4}$, 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\break 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~\ref{sec3}.

Dutrizac~\cite{178} also compiled apparent activation energy data
published by different  authors~\cite{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\unskip\break
ranges.

Dutrizac~\cite{182} examined the factors influencing the precipitation
of K-jarosite in a hydrochloric and sulfate medium. Subsequently, 
Dutrizac~\cite{183} compared the precipitation rates of
NH$_{4}$-jarosite and Na-jarosite in a ferric sulfate medium (0.30~M
Fe(SO$_{4}$)$_{1.5}$/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 NH$_{4}$-jarosite compared 
to\unskip\break
Na-jarosite.

In the hydrometallurgical industry, significant efforts have been made
to study the precipitation of jarosite at high temperatures in abiotic
environments (${&gt;}$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. 

\subsection{Studies in biotic environments at low\newline
temperatures}\label{sec4.2} 
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~\cite{184}, suggest the formation
of jarosite at a 303.15~K in environments rich in
\textit{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\unskip\break 
growth.

Wang et~al.~\cite{185} synthesized schwertmannite and jarosite from
solutions containing ferrous iron sulfate, inoculated with
\textit{Acidithiobacillus ferrooxidans}. Schwertmannite transformed
into H$_{3}$O--NH$_{4}$-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.~\cite{186} synthesized NH$_{4}$-jarosite in the presence of
microorganisms (\textit{A.~ferrooxidans} between 295.15 and 309.15~K, 
Sulfolobus metallicus at 318.15~K, 
and \textit{Acidithiobacillus
caldus} and 
\textit{\mbox{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.~\cite{187} investigated arsenic trapping during Fe$^{2+}$
oxidation in a medium containing \textit{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~\cite{188}
as well as  Kaksonen~\cite{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.

\subsection{Studies in abiotic environments at low\newline
temperatures}\label{sec4.3} 
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.~\cite{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$^{+}$, Mg$^{2+}$, Ca$^{2+}$, NO$^{3-}$, and
$\mathrm{PO}_{4}^{3-}$. The duration of the experiments was 10 days. They
observed a decrease in Fe$^{3+}$ concentration in the solution as the
experiment temperature increased. The phases formed were primarily
composed of H$_{3}$O-jarosite. However, a phase similar to Maus salt
(K$_{5}$Fe$_{3}$(SO$_{4}$)$_{6}$(OH)$_{2.8}$H$_{2}$O) was identified at
temperatures below 298.15~K. When potassium was replaced by hydronium
in Maus salt, metavoltine was  produced~\cite{191}. The proportion of
poorly crystallized or non-crystallized crystals decreased with
increasing temperature.

According to Sasaki and Konno~\cite{192} and  Wang~\cite{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 \mbox{resemble} previous descriptions of  schwertmannite
[Fe$_{8}$O$_{8}$(OH)$_{6}$(SO$_{4}$)] \cite{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.~\cite{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~\ref{sec2.1}, 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 
\mbox{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
\mbox{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.

\section{Characterization of the solid phases of jarosite}\label{sec5} 
Solid analyses are crucial in evaluating jarosite-type compounds, given
the diversity of jarosites that can exist both in Nature and in
industrial processes. This variety leads to complex solid structures,
different morphologies, and surfaces. The inherent complexity of these
variations makes the identification of these compounds challenging.

Among the commonly employed methods for the identification and
assessment of jarosites, scanning electron microscopy (SEM) and X-ray
diffraction (XRD) appear indispensable. 

The study of the dissolution and precipitation of jarosite-type
compounds has resulted in different X-ray diffractograms and varied SEM
images. As an example, in  Figure~\ref{fig5}, two completely different
jarosite morphologies are presented.

\begin{figure*}
\includegraphics{fig05}
\caption{\label{fig5}SEM of synthetic anhedral jarosite (a)~by Smith 
et~al.~\cite{7} and natural euhedral jarosite (b)~by Welch 
et~al.~\cite{43} (Figures reprinted with permission from Elsevier).}
\end{figure*}

In Figure~\ref{fig5}(a), jarosite exhibits an irregular and globular
shape with various grain sizes, most measuring between 1 and 
5~$\rmmu$m. Some grains show cleaved sides, which may be signs of
mechanical abrasion resulting from agitation during 
synthesis~\cite{7}. Similar particle morphologies for synthetic
jarosite have been reported  elsewhere~\cite{194}. On the other hand,
in  Figure~\ref{fig5}(b), using a natural sample, Welch  et~al.~\cite{43}
observed a relatively euhedral shape for jarosite, with a grain size
ranging from 0.5 to 5~$\rmmu$m. Qian  et~al.~\cite{115} also used
euhedral Na-jarosite, indicating a highly crystalline phase, similar to
Gasharova  et~al.~\cite{45}.

According to Sasaki and Konno~\cite{192}, particle morphology is also
influenced by cations in the \textbf{A}-site. K-jarosite exhibits a
round shape, while NH$_{4}$-jarosite manifests as an agglomeration of
pseudo cubic particles. The anhedral and euhedral natures of different
samples can lead to varying dissolution rates. Another example
illustrating these differences concerns the evolution of the morphology
of jarosites synthesized by Reyes  et~al.~\cite{130} as the amount of
chromium in the structure increases. The morphology evolves from 
Figure~\ref{fig6}a/b to Figure~\ref{fig6}c/d. 

\begin{figure*}
\includegraphics{fig06}
\caption{\label{fig6}Cr-jarosites synthesized with increasing amounts
of chromium by Reyes  et~al.~\cite{130} (Figures reprinted from open
access article).}
\end{figure*}

Due to these differences, the use of solid analyses is essential to
understand the morphologies and structures of this type of mineral, as
well as to assess how they may influence the thermodynamic and kinetic
data of precipitation/dissolution of these\unskip\break 
solids.

\section{Conclusion and perspectives}\label{sec6}
The bibliographic review presented in this article traces the evolution
of the existing research on jarosite dissolution and precipitation,
initially aimed at understanding the structure of jarosite and its
presence in the environment. Subsequently, these studies focused on
acquiring thermodynamic and kinetic data related to the formation and
dissolution of jarosite compounds.

The entirety of this literature review highlights a marked diversity in
existing results. Among these, significant divergences and gaps have
been 
\mbox{identified:}
\begin{itemize}
\item 
A major issue arising from these published data is the difficulty in
obtaining thermodynamic and kinetic data due to the time required to
reach equilibrium under the conditions used. To circumvent
long-duration experiments, a trend has emerged favoring studies under
more extreme conditions (high temperature and low pH) that promote
faster dissolution or precipitation.
\item 
When less favorable reaction conditions were employed, shortcuts such
as using the first few hours of the reaction were employed to describe
the kinetics. This time lapse, varying according to the application and
conditions, proves insufficient to predict the \mbox{behavior} of jarosite
compound kinetics. This is due to both reactions (dissolution and
precipitation) requiring extended times, leading to solution evolution,
and the transformation of solids in terms of structure and morphology.
\item 
The transformations undergone can result sometimes in poorly
crystalline solids with different morphologies, which can influence the
data obtained. Therefore, the use of solid analysis methods is crucial
for a correct understanding of the results and for comparison with
previous findings.
\item 
The disparity in kinetic data highlights that using different
precipitation methodologies in order to produce jarosite solids for
dissolution studies could introduce errors, as the solids may vary, and
complicate the comparison between studies.
\item 
The diversity of jarosite structures observed in the environment and in
industry also places constraints on comparisons between experiments, as
ions incorporated into their structure can alter kinetics and
thermodynamic properties.
\item 
Comparing these results is further complicated by the choice of
thermodynamic and kinetic models, along with their inherent limitations
in accurately characterizing the intricate reactive mechanisms of
jarosite compounds.
\end{itemize}

However, where difficulties persist, opportunities to discover
mechanisms, characterize behaviors, improve existing data, and resolve
contradictions emerge. 

\begin{itemize}
\item 
The challenge for further kinetic studies of jarosite lies primarily in
understanding the dissolution and precipitation reaction mechanisms at
temperatures closer to room temperature, conditions under which most
jarosites form in the environment and where some industrial processes
take place.
\item 
The studies must take into account the evolution of precipitated and
dissolved solids over time.
\item 
It seems fundamental to understand how the precipitation methodology
can influence the synthesized solid and how morphology can play a
significant role in dissolution.
\item 
The use of different models to describe both phenomena
(dissolution/precipitation) is important. The use of models that
connect the thermodynamic and the kinetic aspects, such as the
Transition State Theory, widely used by geochemists, would be of great
value to facilitate the control of the kinetics using the
thermodynamics of solutions.
\end{itemize}

\section*{Declaration of interests}
The authors do not work for, advise, own
shares in, or receive funds from any organization that could benefit
from this article, and have declared no affiliations other than their
research organizations.

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