CHARACTERIZATION OF ARGENTOJAROSITE SYNTHESIZED WITH
BIOLOGICALLY PRODUCED FERRIC SULFATE SOLUTIONS
THESIS
Presented in Partial Fulfillment of the Requirements for the Degree Master of Science in
the Graduate School of The Ohio State University
By
Chiranjit Mukherjee
Graduate Program in Microbiology
The Ohio State University
2013
Master's Examination Committee:
Dr. Olli H. Tuovinen, Advisor
Dr. Nicholas T. Basta
Dr. Jerry M. Bigham
Dr. Charles J. Daniels
Abstract
Argentojarosite (AgFe3(SO4)2(OH)6) has an important catalytic role in
biohydrometallurgical circuits, but to date very little is known about silver-containing
compounds with jarosite structure. The purpose of this study was to characterize
argentojarosite synthesized via biological oxidation of ferrous sulfate by
Acidithiobacillus ferrooxidans. The contact time and Ag+ concentration (added as
AgNO3) were varied in these experiments. Synthesis of argentojarosite was confirmed
through X-ray diffraction analysis. Additional analysis of solid-phase oxidation products
included elemental composition, color and specific surface area. The sample synthesized
in the presence of 40 mM Ag+ and with 14 days contact time showed the most crystalline
X-ray pattern, and its elemental composition matched well that of ideal argentojarosite.
The color and surface area of the remaining samples seemed to be influenced by the
presence of a poorly crystalline Fe(III) phase believed to be precursor to schwertmannite
(Fe8O8(OH)6SO4), which otherwise remained stable over the time course of 14 days when
no Ag+ was present in the system.
ii
Acknowledgments
I am extremely grateful to my advisor, Dr. Olli H. Tuovinen, for giving me the
opportunity to conduct research under his advice. His insights, expertise and knowledge
have represented an invaluable contribution to my work. The guidance I have received
from him at every instant when I needed it and his attention to detail has helped me grow
not only as a researcher, but also overall as a human being.
I extend my gratitude to my committee members, Dr. Jerry Bigham and Dr. Charles
Daniels, for their guidance and advice related to designing my experiments and analyzing
the results. Special thanks goes to Dr. Nicholas Basta for allowing me to use his lab
facilities without which the analytical work for this study could not have been possible.
I extend a very special note of gratitude to F. Sandy Jones, who has guided me in
every step of my analytical research work and has been a great teacher as well. I also
thank Kevin Jewell, Lab Manager, Service Testing and Research Laboratory, for the ICP
analysis of my mineral samples.
Finally, a big thank you to Dr. Abhijit Mukherjee and Shukla Mukherjee, for being
the most supportive parents ever. Without you, I would have never been where I am
today.
iii
Vita
2010................................................................Bachelor of Technology (Biotechnology),
West Bengal University of Technology,
Kolkata, India.
2011 to present ..............................................Graduate Teaching Associate, Department
of Microbiology, and Center for Life
Science Education, The Ohio State
University, Columbus, Ohio, U.S.A.
Fields of Study
Major Field: Microbiology
Study in: Environmental Microbiology
iv
Table of Contents
Abstract ............................................................................................................................... ii
Acknowledgments ............................................................................................................. iii
Vita..................................................................................................................................... iv
Fields of Study ................................................................................................................... iv
List of Tables .................................................................................................................... vii
List of Figures .................................................................................................................. viii
1. Introduction……………………...……………………………………………….. 1
1.1 Acid mine drainage – microbiology and precipitation of Fe(III) and sulfate……..1
1.2 Schwertmannite……………………………………………………………………8
1.3 Jarosite……………………………………………………...……………………14
1.4 Argentojarosite…………………………………………………………………..19
1.5 Silver catalyzed bioleaching of copper from chalcopyrite………………………22
1.6 Objective of this study…………………………………………………………...27
2. Materials and Methods……………………………………………………………….. 28
2.1 Bacterial cultures and biological synthesis of Fe(III) precipitates……………….28
2.2 Analytical methods………………………………………………………………29
3. Results and Discussion………………………………………………..……………....30
v
3.1 Identification of biologically synthesized argentojarosite…………….………....30
3.2 Effect of concentration of Ag+ and contact time on Fe(III) precipitates………...33
3.3 Separation of the solid and solution phases……………………………………...45
3.4 Synthetic argentojarosite……………………………………………………..…..48
3.5 Preparation of argentojarosite from ferric sulfate……………………………..…50
4. Conclusions and Future Directions………..……………………………..…….....….. 52
References……………...………………………………………..……………………… 54
vi
List of Tables
Table 1. Summary of major Fe(III) minerals………………………………………..…….6
Table 2. Chemical composition of schwertmannite from different sources...............…...11
Table 3. Naturally occurring jarosites and their synthetic equivalents (modified from
Dutrizac and Jambor 2000)………………………………………… ………...................14
Table 4. Color of biologically produced jarosites (source: Sasaki et al. 1995, Gramp et al.
2008)……………………………………………………………………………………..16
Table 5. Elemental composition of various biologically produced jarosites (source: Sasaki
et al. 1995, Gramp et al. 2008)………………………………………………...…………16
Table 6. Summary of physical and chemical properties of various argentojarosites….....21
Table 7. Comparison of full width at half maximum values for 40 mM Ag+ concentration
at 8 h, 5 d and 14 day (a) and for 5, 20 and 40 mM concentration of Ag+ at 14 day (b)...37
Table 8. Comparison of color and specific surface area of the samples collected……....38
Table 9. Elemental analysis of the samples……………………………………………...40
Table 10. Unit cell parameters of the 14-day samples…………………………………...40
vii
List of Figures
Fig. 1.General scheme of Fe2+ oxidation, leading to jarosite and schwertmannite
precipitation……………………………………….………………………………………7
Fig. 2.XRD pattern of biologically produced schwertmannite (source: Wang et al. 2006).
……………………………………………………………………………………………13
Fig. 3.XRD pattern of type specimen of schwertmannite (source: Bigham et al. 1994)...13
Fig. 4. Peak positions for argentojarosite (source: PDF2 41-1398)…………………......22
Fig. 5. Scheme of argentojarosite preparation…………………………………….……..31
Fig. 6. XRD patterns of precipitates collected after 8 h and 5 days, containing 0 and 40
mM Ag+……………………………………………………………….…………...…….32
Fig. 7. XRD patterns of precipitates collected after 8 hours, 5 days and 14 days,
containing 0 mM Ag+…………………………………………………….…………..…..34
Fig. 8. XRD patterns of precipitates collected after 14 days, containing 0, 5, 20 and 40
mM Ag+…………………………………………………………….……………………35
Fig. 9. XRD patterns of precipitates collected after 8 hours, 5 days and 14 days,
containing 40 mM Ag+…………………………………………………………………...36
Fig. 10. Variation in Ag content of precipitates formed in the presence of 0-40 mM
Ag+……………………………………………………………………………….………42
viii
Fig. 11. Variation in Fe content of precipitates formed in the presence of 0-40 mM
Ag+……………………………………………………………………………………….43
Fig. 12. Variation in relative intensity of the 3.07 Å peak in the presence of 0-40 mM
Ag+………..………………………………………………………………….………….44
Fig. 13. Variation in FWHM values for the 2.7 Å peak in the presence of 5-40 mM Ag+.
……………………………………………………………………………………………45
Fig. 14. Scheme for separation of solid and solution phases……………………......…...47
Fig. 15. XRD pattern for precipitates treated with 20 mM Ag+…………………………48
Fig. 16. XRD patterns for synthetic schwertmannite samples (a) and the same samples
treated with 20 mM Ag+ for 14 days (b)…………………………..…………….………50
ix
1. Introduction
1.1 Acid mine drainage –microbiology and precipitation of Fe(III) and
sulfate
One of the environmental hazards associated with mining is acid mine drainage
(AMD), the drainage of acidic water from old, abandoned, or even currently active mine
sites. AMD can be devastating for rivers, streams, and other surface waters. The
acidification of surrounding soil and receiving waterways and sediments is not, however,
the only environmental problem caused by AMD. AMD has the potential to mobilize
toxic metals such as As, Cd, Cu, Ni, Pb and Cr from exposed ores and coal seams, which
can be disseminated into the environment with acid mine waters. The nature of AMD and
its potential for ecotoxicity is a function of the mineralogy of the local rock material,
availability of air and water, climate, and nature of microorganism present, all of which
are highly site-specific, making predictions and modeling extremely challenging.
Oxidation of sulfide minerals is the main source of sulfuric acid generation. Fesulfides such as pyrite (FeS2) and pyrrhotite (Fe1-xS) are most common sulfide minerals,
but other sulfide minerals may also produce AMD. Although the weathering of rocks
generating acid can occur naturally, the mining of coal and sulfide ores exposes sulfide
rich materials to air and water. The primary sources for generation of acids are sulfide
minerals, water or moisture in the air and an oxidant, which can be dissolved O2 or ferric
1
iron (Akcil and Koldas, 2006). Acidophilic iron and sulfur oxidizing microorganisms are
invariably involved in the formation of AMD through the oxidation of Fe(II) and S
entities in sulfide minerals and regeneration of ferric iron in the solution phase.
AMD can generate very high levels of acidity and elevated concentrations of
dissolved iron and sulfate. The pH values at extreme AMD sites can be on the negative
scale as is the case in the Iron Mine, a Superfund site in California where dissolved iron
and sulfate concentrations are as high as 200 g/L and 760 g/L, respectively (Nordstrom et
al. 2000, Coggon et al. 2012). AMD associated with coal mines is typically in the range
of pH 2-4 and dissolved iron and sulfate levels are usually below 1.5 g/l. AMD has the
potential to leach other metals and metalloids from associated minerals and thus AMD is
considered to be one of the most important sources of heavy metal pollution of the
environment (Sheoran and Sheoran 2006). AMD is detrimental to aquatic life, with AMD
precipitating as ‘yellow boy’ in receiving waters. This term refers to a mixture of various
Fe(III)-hydroxysulfates that vary in composition depending on the chemistry of the site.
Poorly crystalline Fe(III) precipitates such as schwertmannite make the bulk of this mass
(Gazea et al. 1996, Coggon et al. 2012).
A diverse range of acidophilic microorganisms are known to be present in AMD
environments. These include bacteria, archaea and also some eukaryotes. Variations in
temperature, pH, redox potential, oxygenation, dissolved metals and sulfate levels, and
ionic strength create specialized biological niches within the AMD system, characterized
by the presence of specific microbial communities. 16S rRNA gene sequencing of
microorganisms has helped identify much of the microbial flora of AMD, although many
2
sequences remain unknown because of unculturability of AMD microorganisms in pure
or defined mixed cultures. Bacteria commonly found in AMD include proteobacteria
such as Acidithiobacillus, Thiomonas, and Acidiphilum spp.; Leptospirillum group
bacteria such as L. ferrooxidans and L. ferriphilum; Firmicutes such as Acidimicrobium
ferrooxidans and Ferromicrobium acidophilus, and gram positives such as Sulfobacillus
spp. Archaea are found in thermally altered AMD (such as the Iron Mine site) and they
belong to the Thermoplasmatales and Sulfolobales; e.g., Thermoplasma, Acidianus,
Metallosphaera, and Sulfolobus spp. (Johnson and Hallberg 2001, Baker and Banfield,
2003). The unifying feature for all these acidophiles is CO2 fixation, and use of ferrous
iron and/or reduced compounds of inorganic sulfur as energy sources.
Pyrite is the most common of Fe-sulfide mineral present in rocks and coal seams. The
oxidation of pyrite by dissolved O2 is the first step in the weathering process of pyrite.
When pyrite is exposed to water containing dissolved O2, the following reaction takes
place:
2 FeS2 + 7 O2 + 2 H2O  2 Fe2+ + 4 SO42- + 4 H+
(1)
Ferric iron is also a chemical oxidant of pyrite (eqn. 2), generated from the initial
oxidation products (eqn. 3):
FeS2 + 14 Fe3+ + 8 H2O  15 Fe2+ + 2 SO42- + 16 H+
(2)
4 Fe2+ + O2 + 4 H+  4 Fe3+ + 2 H2O
(3)
This reaction is negligibly slow at pH values < 3 under abiotic conditions and at ambient
temperatures. Fe-oxidizing acidophiles (e.g., A. ferrooxidans or L. ferriphilum) can
3
increase the reaction rates by up to 106 fold depending on the environmental conditions
(MacDonaldand Clark 1970).
The acidophilic bacteria play both a direct and an indirect role in pyrite oxidation. In
direct contact, the bacteria attach to the surface of the pyrite and oxidize pyrite by direct
electron transfer between the bacterial cells and pyrite surface, thereby releasing Fe2+
ions. The indirect contact mechanism refers to the oxidation of pyrite by the action of
Fe3+ generated by bacterial oxidation of Fe2+ in the solution phase (eqn. 2 and 3).
Ferric ion is more effective in the oxidation of sulfide minerals than dissolved O2, and
so eqn. 3 followed by eqn. 2 is proposed as the dominant pathway for pyrite oxidation
(Baker et al. 2003). A. ferrooxidans is considered as the most important organism for
pyrite oxidation, and its role along with L. ferrooxidans and L. ferriphilum in the
bioleaching of sulfide ores has been well established (Rawlings 2005, Donati and Sand
2007).
The oxidation of sulfide minerals such as pyrite, pyrrhotite and chalcopyrite (CuFeS2)
releases Fe2+ which is then concurrently oxidized to Fe3+ through bacterial action. Ferric
iron undergoes hydrolysis under acidic conditions, leading to the precipitation of
schwertmannite and jarosite type phases in sulfate-rich environments, depending on
specific conditions such as pH and monovalent cations (Fig. 1). Secondary Fe(III)minerals including jarosite, schwertmannite, ferrihydrite and goethite are precipitated
with distance from the source of AMD. Gagliano et al. (2004) showed that there was an
initial accumulation of schwertmannite in a constructed wetland receiving AMD from an
abandoned coal mine. Poorly crystalline schwertmannite eventually partially transformed
4
to goethite with time. Jarosite occurred mostly as interspersed thin lenses in these AMD
impacted sediments because of the lack of K+, Na+ and NH4+ in the system, except for
their leaching from the surrounding soil. Bigham et al. (1996b) developed a model for
mineral speciation in AMD conditions. They showed that precipitates began to form in
the pH range of 2.8 to 4.5. They were mostly schwertmannite, with traces of goethite
and/or jarosite, while those formed at 6.5 or higher were mainly composed of ferrihydrite
or a mixture of ferrihydrite and goethite. The formation of jarosites and schwertmannite
can be summarized by the following equations, where M is a monovalent cation:
3 Fe3+ + M+ + 2 SO42- + 6 H2O MFe3(SO4)2 (OH)6 + 6 H+
(4)
8 Fe3+ + SO42- + 14 H2O Fe8O8(OH)6(SO4) + 22 H+
(5)
Figure 1 provides a general scheme of Fe2+oxidation as it relates to jarosite and
schwertmannite precipitation. The major Fe(III) minerals discussed here are summarized
in Table 1.
It is clear that iron-oxidizing acidophiles such as A. ferrooxidans are intimately
associated with AMD as they can maintain a high Fe3+/Fe2+ ratio of dissolved iron. Thus
they facilitate the precipitation of ferric iron in sulfate-rich solutions, leading to the
formation of Fe(III)-hydroxysulfates. Without the biological regeneration of ferrous iron
to ferric iron, AMD would be mostly depleted of ferric iron and the bulk of dissolved Fe
would be in the reduced form, ferrous iron. Under such circumstances, Fe(III)
precipitation would be minimal. Thus jarosite and schwertmannite precipitation in AMD
environments is based on the preceding biological oxidation and regeneration of ferric
iron. Because these precipitates greatly impact sedimentary processes in receiving
5
environments, it was of interest in this study to synthesize Fe(III)-hydroxysulfates
involving A. ferrooxidans cultures under defined laboratory conditions for physical and
chemical characterization. There already are published protocols for the synthesis of
schwertmannite and K-, NH4, and Na-jarosites in acidic ferric sulfate solutions produced
with A. ferrooxidans (Wang et al. 2006, Gramp et al. 2008). To add to this inventory, the
present study focused on Ag-jarosite which has natural occurrence in sulfide tailings
rather than AMD although the role of acidophiles is very comparable. This type of
jarosite was also of interest because of its formation in silver-catalyzed bioleaching of
chalcopyrite, another example of acid system that involves iron and sulfur oxidizing
microorganisms.
Mineral
schwertmannite
jarosite
Crystallinity
poorly crystalline
Chemical formula
Brief description
Fe8O8(OH)8-2x (SO4)x∙nH2O
Fe(III)-
(1 ≤ x ≤ 1.75)
hydroxysulfate
MFe3(SO4)2(OH)6
Fe(III)-
(M = monovalent ions; e.g.,
hydroxysulfate
crystalline
K+, NH4+, Na+ and H3O+)
ferrihydrite
poorly crystalline
Fe2O3∙0.5H2O
hydrous ferric oxide
goethite
crystalline
FeO(OH)
Fe(III)-oxyhydroxide
akaganéite
crystalline
β-Fe3+O(OH,Cl)
Fe(III)-oxyhydroxide
Table 1. Summary of major Fe(III) minerals.
6
Fig. 1. General scheme of Fe2+ oxidation, leading to jarosite and schwertmannite
precipitation.
7
1.2 Schwertmannite
Oxidation of sulfide minerals such as pyrite and chalcopyrite in rocks, soils,
sediments, and industrial wastes produces a number of Fe(III) oxides, oxyhydroxides and
hydroxysulfates. Precipitates which are primarily composed of Fe(III) compounds are red
to yellow in color. The major significance of this precipitate formation is as follows:
1. Adds to suspended sediment load
2. Reduces effectiveness of wetlands and other mine drainage abatements
3. Acts as a sink for heavy metals (Bigham and Nordstrom 2000).
Owing to its unique characteristics such as poor crystallinity, meta-stability, nanoscale size and association with other closely related, more crystalline phases such as
jarosites, goethite and ferrihydrite, schwertmannite was not identified as a mineral until
the 1990’s. In their attempts to identify and characterize colloidal Fe(III) precipitate
formed in a stream impacted by AMD, Brady et al. (1986) suggested that, based on the
minor peaks present in their X-ray diffraction (XRD) data, poorly crystallized akaganéite
might be one of the constituents. Bigham et al. (1990) showed that the chief compound
formed by bacterial oxidation of Fe in acid sulfate systems within the pH range of 2.5-4.0
was a poorly crystallized Fe(III) “oxyhydroxysulfate.” Subsequent to this report, it was
finally characterized as a new mineral and named schwertmannite (Bigham et al. 1994).
Schwertmannite has a “tunnel” structure similar to that of akaganéite (βFe3+O(OH,Cl)). Recent studies, utilizing pair distribution function (PDF) data, XRD
analyses, and density functional theory calculations have shown that the schwertmannite
octahedral framework is made up of “a highly defective entangled network of structural
8
motifs” (Fernandez-Martinez et al. 2010). This structure also provides a basis for the
transformation of schwertmannite to goethite, which has an orthorhombic unit cell,
through release of sulfate and two octahedral irons.
Schwertmannite has a high specific surface area, usually in the range of 100 to 200
m2/g (Bigham et al. 1994). It is brownish yellow in color, and its color can be represented
using the Munsell color system. The Munsell color system is based on a 3-D model
depicted in the Munsell color tree, in which each color is comprised of three attributes.
These are hue (the color itself, such as red, orange, yellow), value (the lightness/darkness
of the color) and chroma (the saturation or brilliance of the color). Hue (H), value (V) and
chroma (C) are often depicted as H V/C when referring to the color of a solid
(http://munsell.com/about-munsell-color/how-color-notation-works/). In the Munsell
color scheme, the schwertmannite type specimen described by Bigham et al. (1994) had a
color of 8.0 YR 5.3/8.1. It has a calculated density of 3.77 to 3.99 g/cm3 and a unique
“pin-cushion,” “ball and whisker,” or “hedge-hog” morphology due to needle-like
structures radiating from the surface of nanoparticles. Most specimens consist of
aggregates of spherical to ellipsoidal particles that are 200 to 500 nm in diameter
(Bigham et al. 1994).
Schwertmannite has been shown to have a variable chemical formula,
Fe8O8(OH)8-2x(SO4)x∙nH2O (1 ≤ x ≤ 1.75), depending upon the degree to which tunnel
and surface sites are saturated with sulfate (Bigham et al. 1994, 1996b). At the time
when its discovery was reported, the type specimen of schwertmannite was represented
by the empirical unit cell formulae of Fe16O16(OH)9.6(SO4)3.2∙10H2O while a synthetic
9
specimen was represented by Fe16O16(OH)9.4(SO4)3.3∙12H2O (Bigham et al. 1994). The
chemical composition of schwertmannite from different sources, a naturally available
specimen, a synthetic specimen and a biologically produced specimen, are listed in Table
2.
As suggested by these data, schwertmannite can be prepared both via biotic and
abiotic synthesis. Bigham et al. (1990) synthesized abiotic samples of a poorly crystalline
Fe(III)-hydroxysulfate (which was later identified as schwertmannite) by hydrolyzing
FeCl3 solutions containing different concentrations of SO42- for 12 min at 60 °C. The
suspensions were then cooled to room temperature, dialyzed for 30 days against
deionized water using cellulose membranes, and finally freeze-dried.
10
Specimen type
Composition (% wt.)
Fe
Source
S
natural specimen
41.9
5.50
French et al. (2012)
synthetic specimen (abiotic)
47.2
3.50
Bigham et al. (1996a)
biologically synthesized
47.2
5.80
Wang et al. (2006)
57.81
4.15
calculated
specimen
theoretical
Table 2. Chemical composition of schwertmannite from different sources.
Bigham et al. (1990) also achieved the “biotic” synthesis of schwertmannite via
oxidation of FeSO4 solutions at around pH 2.0 in a bioreactor inoculated with
Acidithiobacillus ferrooxidans. Their media contained (per liter) 0.4 g (NH4)2SO4 , 0.02 g
KCl, 0.05 g MgSO4, 0.05 g K2SO4, 0.002 g Ca(NO3)2 and 40 g FeSO4. Wang et al.
(2006) prepared and characterized schwertmannite using a medium which contained 1.6
mM MgSO4∙7H2O, 120 mM FeSO4∙7H2O, and a range of NH4H2PO4 in 4.05 mM H2SO4,
with a varying incubation time, at 36° C. When incubated with 5.4 mM NH4+ for 7 days,
the resultant precipitate had a specific surface area (SSA) of 12.78 m2/g and on the
Munsell color chart its H V/C reading was 6.3YR 4.3/8.6. The Fe and S content of this
sample is presented in Table 2. The relatively low SSA suggests that the schwertmannite
sample poorly matched the reference schwertmannite. Longer incubation time (19 days)
and increased concentration of NH4+(11.4 mM) resulted in a more yellow product (9.2
11
YR 6.5/9.0) with a specific surface area of 28.5 m2/g. The Fe:S molar ratio for this
sample was 4.6:1, which is comparable to the type specimen (4.9) and a synthetic
specimen (4.8) as noted by Bigham et al. (1994). The XRD pattern for this sample is
presented in Fig. 2, which shows the characteristic 8 broad peaks of schwertmannite as
seen for the type specimen (Fig. 3).
Iron oxides may exist in various forms, ranging from well crystallized hematite
(Fe2O3) and goethite to poorly crystalline ferrihydrite, schwertmannite and feroxyhyte (δFe3+O(OH)) and the poorly crystalline iron oxides act as precursor to the well-crystallized
ones. In AMD situations (i.e., sulfate-rich environments), schwertmannite is known to coexist with jarosites. The abundance of one over the other in fresh precipitates depends on
the concentration of monovalent ions and the pH of the source solutions. In the absence
of monovalent cations and in the pH range of 3-4.5, schwertmannite is the dominant
phase, whereas at a lower pH and in the presence of monovalent cations, jarosite
formation takes place, with the soluble cations determining the specific type of jarosite
(Bigham et al. 1996, Wang et al. 2006).
12
Fig. 2. XRD pattern of biologically produced schwertmannite (source: Wang et al. 2006).
Numerical values for XRD peaks are in Ångström units, calculated from º2θ CuKα scale.
Fig. 3. XRD pattern of type specimen of schwertmannite (source: Bigham et al. 1994).
Numerical value for XRD peaks are in Ångström units, calculated from º2θ CoKα scale.
13
1.3 Jarosite
August Breithaupt, a German mineralogist, first discovered in 1852 a yellowishbrown iron-sulfate-hydroxide mineral in the Jaroso Ravine in Sierra Almagrera in Spain
and named it “jarosite” (Swayze et al. 2008). Jarosites belong to a larger family of
minerals represented by the formula AB3(XO4)2(OH)n∙mH2O. For jarosites, B = Fe3+, X =
S, n = 6 and the formula has no crystal water. Different jarosite group minerals are
formed based on substitution of A with cations. Only a few naturally occurring jarosites
are known. Their mineral names, synthetic equivalents, and substitution of A are listed in
Table 3. Traditionally, the term “jarosite” refers to potassium jarosite, KFe3(SO4)2(OH)6,
but all the minerals in this group are referred to as jarosites for the purpose of this review.
Cation
Mineral name
Synthetic equivalent
K+
jarosite
potassium jarosite
ammoniojarosites
ammonium jarosite
H3O
hydronium jarosite
hydronium jarosite
Na+
natrojarosite
sodium jarosite
Ag+
argentojarosite
silver jarosite
Tl+
dorallcharite
thallium jarosite
0.5 Pb2+
plumbojarosite
lead jarosite
0.5 Hg2+
no mineral equivalent
mercury jarosite
NH4+
+
Table 3. Naturally occurring jarosites and their synthetic equivalents (modified from
Dutrizac and Jambor 2000).
14
Parameter ain the unit cell is presented as 7.3 Å but it varies with synthesis
conditions. Parameter c ranges from 16.5 to 17.4 Å for monovalent cation substitution
and 33.7 Å for divalent cation substitution such as plumbojarosites (Das et al. 1996).
Jarosites are found in acid sulfate soils, AMD sediments, mine tailings of sulfide ores,
or as byproducts in refining and hydrometallurgy processes. Pure jarosite samples have
been synthesized both chemically (Brown 1970, Dutrizac and Kaiman 1976, Bigham et
al. 2010) and biologically (Ivarson et al. 1979, Sasaki et al. 1995, Sasaki and Konno
2000, Gramp et al. 2008). Although natural, chemically synthesized and biologically
produced (through hydrolysis of biologically oxidized Fe(III) ) samples show some
variation in their physical properties, chemical analysis helps to identify and categorize
each type. While extensive literature is available on preparation and characterization of
synthetic jarosites, comparatively less is known about jarosites produced from
biologically oxidized ferric iron. Ochreous precipitates are formed near mines where acid
mine drainage comes in contact with fresh water from streams and rivers. The jarosite
group of minerals has a characteristic yellowish color. Table 4 lists some of the
biologically produced jarosites, and their respective colors. The solution phase
concentration of the monovalent cation has a considerable effect on the jarosite color
data.
Elemental analysis for specific cations can be used to establish the specific jarosite
type (Table 5). Jarosites in the environment are usually mixtures (solid solutions) with
various cationic substitutions rather than end member compositions.
15
Cation
Solution phase concentration
Munsell color for jarosite
(mM)
(H V/C)
Na
500
9.6 YR 5.9/10.1
K+
12
9.6 YR 6.2/8.3
31
3.0 Y 7.9/6.3
160
8.2 YR 5.0/10.0
320
9.3 YR 6.2/9.8
53
5.0Y 8.5/11
+
NH4+
+
Ag
Table 4. Color of biologically produced jarosites (source: Sasaki et al. 1995, Gramp et al.
2008).
Cations
Elemental composition (% wt.)
Mol Ratio
Na
K
Ag
Fe
S
N
Fe/S
Na+
4.74
0
0
34.6
13.3
0
1.5
K+
0
7.82
0
33.5
12.9
0
1.5
NH4+
0
0
0
35.0
13.4
2.93
1.5
Ag+
0
0
14.6
29.1
11.2
0
1.5
0
0
0
34.9
13.4
0
1.5
+
H3O
Table 5. Elemental composition of various biologically produced jarosites (source: Sasaki
et al. 1995, Gramp et al. 2008).
16
Ivarson (1973) was among the first to synthesize jarosite by microbial oxidation of
ferrous sulfate. Since then, several studies shown that jarosites can be produced from
ferrous sulfate at ambient temperature using acidophilic iron oxidizing bacteria such as A.
ferrooxidans (Tuovinen and Carlson 1979, Lazaroff et al. 1982, 1985, Sasaki et al. 1995,
Sasaki and Konno 2000, Gramp et al. 2008).
Jarosite formation begins with oxidation of Fe2+ ions to Fe3+, followed by formation
of crystal nuclei and finally the growth of complete crystals (Sasaki and Konno 2000).
The method and rate of Fe2+ oxidation is known to have an effect on the characteristics of
the jarosite formed. Based on their studies with argentojarosite, Sasaki et al. (1995)
concluded that biological synthesis did not have any direct contribution to the
crystallization process of jarosites. However, Sasaki and Konno (2000) showed that
biologically synthesized jarosites had significantly different morphology compared to
chemically synthesized products. Previously, it had been reported that extracellular
polysaccharides may cause an increase in the adhesion of bacterial cells to the jarosite
particles (Sadowski 1999). Sasaki and Konno (2000) concluded that the aggregates they
observed in the biologically synthesized jarosites were due to extracellular substances
secreted by microorganisms.
Usually, microbial cells are removed from the media after partial or complete
oxidation of ferrous sulfate, before monovalent cations are added to the spent media to
obtain specific jarosites (Lazaroff et al. 1982, 1985; Sasaki et al. 1995). Wang et al.
(2006) showed that schwertmannite and jarosite can be precipitated from ferrous sulfatemineral salt media inoculated with A. ferrooxidans, and that elevated temperature and
17
higher concentrations of ammonium ions led to jarosite precipitation while lower
temperature and ionic concentrations precipitated schwertmannite. Wang et al. (2006)
also observed the transformation of schwertmannite into jarosite in the presence of
monovalent ions, but the elemental composition, color and specific surface area of these
“transformed” jarosites were different from those jarosites that were directly precipitated
upon bacterial iron oxidation.
Wang et al. (2006) demonstrated that biological oxidation of iron at 22 and 36 °C
favored the precipitation of schwertmannite when NH4+ concentration was less than 10
mM and no other monovalent cations were present in the media. Ammoniojarosite was
the sole product when NH4+ concentration was higher than 165 mM, with intermediate
concentrations giving rise to a mixed precipitate of schwertmannite and ammoniojarosite.
A consistent difference in structure, elemental composition, color or specific surface area
between the biological products and chemically synthesized samples was not apparent.
The primary role of microorganisms in the synthesis of ammoniojarosite was that of the
oxidation of Fe2+ at pH ~ 1.9 (Wang et al. 2007).
Gramp et al. (2008) synthesized and characterized different types of jarosites in liquid
media inoculated with Acidithiobacillus and having a range Na+, K+ and NH4+
concentrations in order to determine the practical concentrations of monovalent cations
required for jarosite precipitation. The concentration of monovalent cations required
varied depending on the type of jarosite. Potassium jarosite required the lowest level of
cation and natrojarosites the highest level. Schwertmannite was detected based on XRD
18
patterns and SEM micrographs even at the highest concentration of Na+ (500 mM) and
NH4+ (320 mM) (Gramp et al. 2008).
Daoud and Karamanev (2006) determined jarosite precipitation under different pH
conditions in A. ferrooxidans cultures. The maximum precipitation of 0.10-0.12 g/l
occurred at pH range of 2.5-3.0. To date, no kinetic analyses have been reported for
biologically produced jarosites.
Significance and application of jarosites
AMD and the precipitation of schwertmannite and/or jarosite have been largely
considered as negative impacts on the surrounding environments of old mining sites. In
contrast, argentojarosite has been known to be sufficiently abundant in some localities to
be considered as a source of silver (Schempp 1923). However, the most significant
interest in jarosite arises from a hydrometallurgical perspective. Jarosite precipitation has
been used as a means for controlling iron and sulfate concentrations in the
hydrometallurgical circuits for the past 40 years. In zinc industry, the jarosite process is
used to precipitate iron from hot acid solutions in the processing circuit after addition of
alkali ions (NH3, Na2CO3 or Na2SO4) (Dutrizac and Jambor 2000).
1.4 Argentojarosite
The formation of argentojarosite is mostly seen as a problem in hydrometallurgical
processes. Precipitation of silver as silver jarosite reduces its recovery during the
conventional cyanidation process for silver extraction from its ores (May et al. 1973). In
an attempt to learn more about the nature of argentojarosite, May et al. (1973)
19
synthesized argentojarosite chemically. They prepared sufficient amounts to enable
chemical analysis and to study thermal decomposition and optical and physical
properties. Synthesis of argentojarosite was previously reported by Fairchild (1933) but
the yield was too low for any further analysis. Subsequently, Dutrizac and Kaiman (1976)
also synthesized argentojarosites.
May et al. (1973) used a boiling solution of Ag2SO4 with Fe2(SO4)3 and either nitric
or sulfuric acid to prepare agentojarosite. The resultant mixture was refluxed at 97 ºC
for about 200 hours and the mustard-yellow precipitate thus formed was collected after
filtration, washing and drying. Dutrizac and Kaiman (1976) used a similar process where
they heated a solution of ferric sulfate with sulfuric acid and Ag2SO4 in an autoclave at
140 ºC for 2 hours, and the precipitate was collected after washing and drying.
Sasaki et al. (1995) were the first to produce argentojarosites from culture solutions of
Acidithiobacilli. Their media contained 160 mM Fe2+ at pH 2.2. Silver in the form of
AgNO3 (53 mM Ag+) was added and a contact time of 168 hours at 30 ºC was allowed
for aging of the precipitate, after which it was collected using 0.20 µm filters.
Physical and chemical properties of argentojarosite
A summary of the elemental composition and physical properties of argentojarosite
from published data is presented in Table 6.
20
Sample
Elemental composition
SSA
Color
Lattice
(% wt.)
(m2g-1)
H V/C
parameter
(Å)
Ag
Fe
S
theoretical
18.94
29.41
11.26
NAa
standard1
17.3
29.8
11.9
biological1
14.6
29.1
chemical1
15.5
18.16
2
synthetic
a
a
c
NA
7.35
16.58
2.7
5.0 Y 8.5/11
7.35
16.55
11.2
1.7
5.0 Y 8.5/11
7.35
16.56
31.4
11.7
1.2
5.0 Y 8.5/11
7.35
16.55
28.6
NA
NA
NA
NA
NA
Not available. 1Sasaki et al. 2000. 2Dutrizac and Kaiman 1976.
Table 6. Summary of physical and chemical properties of various argentojarosites.
Dutrizac and Kaiman (1976) noted a deficiency of Ag in the elemental composition of
their argentojarosite, and attributed this to hydronium ion substitution for the metal. They
also noted a “slight but consistent” deficiency in iron. Their argentojarosite synthesized at
a lower temperature (75 ºC) had even lower Ag content, 16% wt., indicating that elevated
temperatures inhibited hydronium substitution.
X-ray diffraction pattern
Reference XRD peak positions and d-values for argentojarosite are presented in Fig. 4.
21
2 θ CuKα
Fig. 4. Peak positions for argentojarosite (source: PDF2 41-1398). Numerical value for
XRD peaks are in Å units.
1.5 Silver-catalyzed bioleaching of copper from chalcopyrite
Chalcopyrite is the most abundant copper sulfide mineral accounting for about 70%
of the world’s known copper reserves (Wang 2005). Low grade Cu-ores have been
successfully leached by heap bioleaching processes, but only from copper oxides and
secondary copper sulfides. The primary copper sulfide, chalcopyrite, is extremely
recalcitrant to bioleaching. Low leaching rates have been a recurring problem for low
22
grade copper ores containing chalcopyrite, as less than 50% of the copper is leached even
after many years of heap or dump leaching (Muñoz et al. 2007).
During chalcopyrite bioleaching, both sulfur and iron containing insoluble byproducts
are formed on the mineral surface, and they have a negative effect on the leaching rates.
These boundary layers on the surface of the chalcopyrite particles limit ionic transport
and thus slow down the reaction considerably (Muñoz et al. 2007), leading to passivation
of chalcopyrite bioleaching. Surface analytical techniques such as X-ray photoelectron
spectroscopy have identified several sulfur containing species on leached chalcopyrite
surfaces. These include elemental sulfur, basic ferric sulfate akin to jarosites, a disulfide
phase, Fe-deficient Cu-sulfides and an unreacted chalcopyrite phase (Watling 2006,
Parker et al. 2003, Klauber et al. 2001). While there is a general consensus that the
decrease in copper extraction rate from chalcopyrite during bioleaching is due to
formation of non-reactive layers on the surface of the mineral, there are different views
regarding the true nature of these layers (Debernardi and Carlesi 2013). Some authors
attribute the passivation of chalcopyrite bioleaching to the formation of a non-porous
elemental sulfur layer which blocks diffusion of ions between the leach solution and the
ore (Muñoz et al. 1979, Klauber et al. 2001), while others attribute this to the
precipitation of a jarosite phase on the mineral surface. Both causes are concomitant with
the passivation of chalcopyrite bioleaching (Stott et al. 2000, Sandström et al. 2005).
Several studies have addressed operational conditions involving bacteria or archaea
and catalysts in efforts to remove, alleviate, or modify the passivating boundary layer on
23
chalcopyrite surface (Muñoz et al. 2007). Chalcopyrite dissolution can be ideally
represented as follows (Li et al. 2013):
CuFeS2+ 4 H+ + O2  Cu2+ + Fe2+ + 2 S0 + 2 H2O
(6)
Iron and sulfur oxidizing microorganisms can enhance the bioleaching process
through the oxidation of ferrous iron (eqn. 3) and sulfur (Li et al. 2013):.
2 S0 + 3 O2 + 2 H2O  2 H2SO4
(7)
Iron oxidation regenerates ferric iron which reacts with chalcopyrite and is reduced to
ferrous iron.
CuFeS2+ 4 Fe3+  Cu2+ + 5 Fe2+ + 2 S0
(8)
Thus microorganisms influence the conditions that affect the passivation boundary layer,
i.e., jarosite precipitation and formation of elemental sulfur (Li et al. 2013).
Addition of silver as a catalyst has been shown to have a considerable effect on
improving the rate of chalcopyrite bioleaching (Ahonen and Tuovinen 1990, Sato et al.
2000, Wang 2005, Muñoz et al. 2007, Feng et al. 2013). Gómez et al. (1999) reported a
3-fold increase in copper yields from the bioleaching of chalcopyrite at 45 °C as
compared to the no-silver control, but reported abundant jarosite precipitation at higher
temperatures (55 °C) with silver catalysis.
In the presence of silver, the following reaction takes place:
CuFeS2 + 4 Ag+ Cu2+ + Fe2+ + 2 Ag2S
(9)
Ferric ion acts as an oxidizing agent and helps to dissolve the silver sulfide on the mineral
surface, thus regenarating Ag ions for further attack of chalcopyrite.
Ag2S + 2 Fe3+ 2 Ag+ + 2 Fe2+ + S0
(10)
24
Silver has a catalytic effect because the main product formed on the mineral surface,
a mixture of Ag2S and S, is relatively porous and does not block the chalcopyrite surface,
as compared to the non-porous elemental sulfur layer which forms during the uncatalyzed
reaction (Miller et al. 1981). It was also noted that the products had a higher electrical
conductivity leading to enhanced transport of electrons to the chalcopyrite surface,
thereby facilitating the leaching reactions (Price and Warren 1986).
In general, elemental silver is known to inhibit the growth of microorganisms, and
soluble silver is toxic to Acidithiobacillus cultures even at concentrations as low as 5 ppm
(Hoffman and Hendrix 1976). Tuovinen et al. (1985) showed that while prolonged lag
phases resulted from exposure to 0.01 mM Ag+, once growth of A. ferrooxidans started
the oxidation rate of Fe2+ showed little change. The concentration of dissolved silver in
the bioleaching solution decreases considerably, possibly due to interaction of silver ions
with chalcopyrite, thus alleviating the potential toxic effect during the bioleaching. Sato
et al. (2000) demonstrated that iron oxidation by A. ferrooxidans was not inhibited by the
addition of silver chloride because of its low solubility, but the oxidation was inhibited by
silver sulfate because it does not form a poorly soluble complex in the culture solution,
using 200 mg/L as silver for both of them. Feng et al (2013) reported that pH 1.3, silver
ion concentration of 2.0 mg/l, and 2.5 g/l chloride ion concentration was the most
efficient combination for the bioleaching of chalcopyrite. Their findings are difficult to
explain because silver is sequestered as AgCl and is not available a catalyst.
Ahonen and Tuovinen (1990) showed that the effect of silver catalysis (up to 30 mg/l
as a sulfate or nitrate salt) on chalcopyrite bioleaching was transient, and that the rate of
25
the reaction, after attaining an initial peak, declined to the levels seen before addition of
silver. The catalyst effect of silver is specific for chalcopyrite in bioleaching systems.
Muñoz et al. (2007) optimized the concentration of silver for enhancing the bioleaching
of copper from chalcopyrite in shake flask experiments. The catalytic effect of silver
disappeared above pH ~3.0, possibly due to inhibition of bacterial activity as silver in
solution was detected at that pH value. Muñoz et al. (2007) also noted that an increased
ferric ion concentration ([Fe3+]initial > 1 g/l) negatively impacted the silver catalyzed
bioleaching, while the same condition positively impacted the silver-catalyzed chemical
leaching of their low-grade ore. These findings suggest that the presence of bacterial
cells, under high ferric ion concentration, would enhance jarosite formation, which could
sequester Ag catalyst in the bioleaching of chalcopyrite. This theory is possibly supported
by the findings of Ahonen and Tuovinen (1990), who had noted that only a small fraction
of the silver added to the bioleaching of chalcopyrite containing sulfide ore as a catalyst
was found in the leach solution, while the jarosite fraction contained up to 50 ppm silver,
amounting to a 140-fold enrichment of silver in that fraction.
Thus silver jarosite (argentojarosite) is a sink for much of the silver that is added as a
catalyst for enhancing the bioleaching of chalcopyrite. Commercial applications of silver
catalyst are not prospective because of the irreversible loss of the catalyst from the
system due to argentojarosite precipitation. Argentojarosite is poorly soluble with
solubility product log K of -11.55 (Gaboreau and Veillard 2004), calculated from
AgFe3(SO4)2(OH)6 + 6 H+ → Ag+(aq) + 3 Fe3+(aq) + 2 SO42-(aq) + 6 H2O(l). While the loss of
silver through incorporation into the jarosite fraction is well recognized in
26
biohydrometallurgy, only few studies have been published on its physical and chemical
characteristic and laboratory synthesis involving solutions such as biologically produced
acidic ferric sulfate that simulate bioleaching conditions. Its role in making silver
unavailable for catalysis of chalcopyrite bioleaching prompted the present study of
synthesis and characterization of argentojarosite under well-defined laboratory
conditions.
1.6 Objective of the study
The limited amount of information available on argentojarosite prompted the present
study. The objective was to link the synthesis of argentojarosite to biological oxidation of
iron, as relevant in bioleaching processes. A two-stage protocol was established for the
synthesis, involving ferrous sulfate oxidation and Fe(III) precipitation, which was
directed toward argentojarosite formation with Ag+ addition. Argentojarosite thus formed
was characterized by elemental composition, X-ray diffraction, color and specific surface
area.
27
2. Materials and Methods
2.1 Bacterial cultures and biological synthesis of Fe(III) precipitates
A composite mixed culture of the iron-oxidizing bacterium A. ferrooxidans was
maintained in a medium which contained 3.8 mM (NH4)2SO4, 3.6 mM KH2PO4, 1.9 mM
MgSO4·7H2O, 4.0 mM H2SO4 and 120 mM FeSO4·7H2O, pH 2.0±0.1. For
schwertmannite biosynthesis, the medium was modified by replacing KH2PO4 with 3.6
mM H3PO4 and (NH4)2SO4 with 7.6 mM HNO3 to prevent the formation of potassium
and ammonium jarosites, respectively. The stock cultures were maintained with 5%
inoculum in shake flasks (180 rev/min) at 22±2 °C. 250 ml culture flasks with 100 ml
media were used for growing the cultures for schwertmannite and argentojarosite
formation.
For synthesis of argentojarosite, AgNO3 was added in 0-40 mM concentrations to the
schwertmannite cultures after Fe2+ oxidation was completed, as evident through
appearance of a dark red color (5-7 days). The cultures were then incubated further for 8
hours, 5 days or 14 days. The Fe(III) precipitate thus produced was collected by
centrifugation ( 15,000 g), washed twice with 1 mM H2SO4, and air dried at 22±2 °C for
72 hours before characterization by powder X-ray diffraction (XRD), SSA, color and
elemental analyses.
28
2.2 Analytical methods
XRD analysis was carried out using a Bruker D8 Advance Series II X-ray diffraction
system (Karlsruhe, Germany) with CuKα radiation. The results were analyzed and the
full width at half maximum (FWHM) and unit cell dimensions were calculated using
EVA X-ray data evaluation software. All samples were packed into a quartz sample
holder and step-scanned from 2º to 70º2θ using a step interval of 0.05º2θ and a counting
time of 4 s.
SSA of the samples was determined by using a Micromeritics FlowSorb II 2300
(Micromeritics, Norcross, GA) instrument and the single-point Brunauer-Emmett-Teller
method with N2 gas as the adsorbate. Color measurements of the samples were performed
with a Minolta CR-300 Chroma Meter (Konica Minolta Photo Imaging, Mahwah, NJ)
and color results were recorded in the Munsell color system with Hue, Value, and
Chroma (HVC) notations. Total sulfur content of the samples was determined by
iodometric titration of SO2 evolved using a Model 518 semi-automatic titrator following
combustion of the samples at 800 °C in a Leco (LECO Corporation, St. Joseph, MI)
induction furnace.
Total Ag, Fe, K, and P were determined by inductively coupled plasma (ICP)
emission spectrometry using a Prodigy dual view ICP (Leeman Labs, Hudson, NH)
following complete digestion of precipitate samples in 70% HNO3 over a 24 h period.
29
3. Results and Discussion
3.1 Identification of biologically synthesized argentojarosite
Because silver is highly toxic to A. ferrooxidans, the synthesis of argentojarosite was
based on a two-step protocol. First, ferrous sulfate (120 mM) was oxidized by growing
cultures of A. ferrooxidans to ferric sulfate. In this first step, some of the ferric iron
precipitated as a poorly crystalline phase. In the second step, 0-40 mM silver was the
biologically produced ferric sulfate solution (spent medium), subsequently, precipitating
argentojarosite over time. Samples were collected at 8 h, 5 d and 14 d post addition of
AgNO3 and analyzed by XRD.
The XRD patterns for samples from 0 mM Ag after 8 h and 5 d equilibration showed
broad peaks with relatively high background, suggesting poor crystallinity of the
samples. In contrast, the 40 mM Ag, 8 h sample showed distinct peaks that were better
defined upon further contact time of 5 d (Fig. 6).
30
Fig. 5.Scheme of argentojarosite preparation.
31
Fig. 6. XRD patterns of precipitates collected after 8 h and 5 days, containing 0 and 40
mM Ag+. The peak positions and relative intensities for reference argentojarosite peaks
from the database (source: PDF2 41-1398) are also shown.
32
3.2 Effect of concentration of Ag+ and contact time on Fe(III)
precipitates
The effects of Ag+ concentration and contact time on the Fe(III) precipitates were
examined. Previous work with transformation of schwertmannite to jarosite showed that
increase in concentrations of the monovalent cation and long contact times result in well
crystalline jarosites, which are more yellow and have an elemental composition closer to
the theoretically calculated values (Wang et al. 2006, Gramp et al. 2008). Sasaki et al.
(1995) used 53 mM of Ag+ for argentojarosite synthesis with biologically produced ferric
sulfate solution. A graded concentration range of 0-40 mM Ag+ was selected in the
present study to narrow down the specific concentration under which argentojarosite
would be formed. Contact times of 8 hours, 5 days and 14 days were selected based on a
preliminary study, which suggested that up to 10 days are required for complete
crystallization of argentojarosite. XRD patterns for the control samples (0 mM Ag+)
showed broad peaks characteristic of poorly crystalline precipitates, suggesting the
presence of a precursor to schwertmannite (Fig. 7). The patterns remained largely similar
over the 14-day period, suggesting that there was no crystallization happening in the
absence of Ag+.
The effect of Ag+ concentration can be seen in Fig. 8, where the XRD patterns clearly
show that among the samples collected on the 14th day, the sharpest peaks were observed
for precipitates formed in the presence of 40 mM Ag+. At 5 mM, argentojarosite peaks
had started to form, while at 20 mM, the peaks were at intermediate height. At 0 mM, no
such peaks could be observed. The peaks seen in the XRD patterns for 5, 20 and 40 mM
33
concentrations once again coincided with the characteristic peaks of argentojarosite from
the database.
Fig. 7. XRD patterns of precipitates collected after 8 hours, 5 days and 14 days,
containing 0 mM Ag+.
34
Fig. 8. XRD patterns of precipitates collected after 14 days, containing 0, 5, 20 and 40
mM Ag+.
The effect of contact time can be seen in Fig. 9, where the XRD patterns for samples
collected at 8 h, 5 days and 14 days after addition of 40 mM Ag+ are compared. The
highest intensity peaks were observed for the sample having contact time of 14 days.
Once again, these peak positions coincided with the characteristic peaks of
argentojarosite from the database.
35
Fig. 9. XRD patterns of precipitates collected after 8 hours, 5 days and 14 days, from
solutions containing 40 mM Ag+.
The FWHM values for the Kα1 peak (2.7 Å) was calculated utilizing a more thorough
scan setting (10 s step and 0.01°2θ step interval). The data are presented in Table 7. The
reduction in FWHM with time and increasing concentration indicates that these led to
sharper, less broad peaks with longer contact time and increased Ag+ concentration.
36
Time
FWHM (°2ϴ)
Peak positions
(101)
(110)
(20-1)
(113)
(006)
8h
0.094
0.078
0.116
0.185
0.555
5d
0.059
0.065
0.062
0.067
0.083
14 d
0.059
0.063
0.059
0.064
0.079
a.
Ag+
FWHM (°2ϴ)
(mM)
Peak positions
(101)
(110)
(20-1)
(113)
(006)
5
0.083
0.089
0.090
0.097
0.130
20
0.056
0.055
0.062
0.064
0.098
40
0.059
0.063
0.059
0.064
0.079
b.
Table 7. Comparison of full width at half maximum values for 40 mM Ag+ concentration
at 8 h, 5 d and 14 day (a) and for 5, 20 and 40 mM concentration of Ag+ at 14 day (b).
The d spacings for the Miller index peaks are: 101 = 5.9 Å, 110 = 3.6 Å, 20-1 = 3.1 Å,
113 = 3.0 Å, 006 = 2.7 Å.
Color and specific surface area
Color and specific surface area data from the samples collected are presented in Table
8. The samples without any added silver had a more reddish orange color (5.9 – 6.7 YR),
37
as compared to the more yellow color (6.9 YR – 0.1Y) of the samples with silver added.
The most crystalline argentojarosite at 40 mM 14 day contact time had a color values of
0.1Y 6.6/7.7, which is more red compared to the color reading of 5.0Y 8.5/11 for the
argentojarosite sample synthesized by Sasaki et al. (1995). This is possibly due to trace
amounts of schwertmannite or its precursors remaining in the sample.
Ag+
Contact
(mM)
time
0
8h
5.9 YR
4.0
7.7
0.77
5d
6.5 YR
4.4
7.8
1.50
14 d
6.7 YR
3.4
8.9
4.84
8h
7.8 YR
5.1
6.2
N.D.a
5d
7.7 YR
4.1
11.1
4.80
14 d
6.9 YR
4.8
7.7
1.01
8h
7.0 YR
5.0
7.6
2.16
5d
0.1 Y
6.8
6.8
0.48
14 d
9.6 YR
6.2
8.1
0.58
8h
7.4 YR
5.6
7.0
2.70
5d
9.2 YR
5.7
8.8
0.81
14 d
0.1 Y
6.6
7.7
0.64
5
20
40
Hue
Value
Chroma
SSA
(m2/g)
aN.D., Not determined.
Table 8. Comparison of color and specific surface area of the samples collected.
38
Elemental analysis
Partial elemental composition of the precipitates measured with ICP is listed in Table
9.The most crystalline was the 40 mM Ag+ /14 day sample, with Ag, Fe and S
composition of 18.42, 28.99 and 11.07 % wt., respectively. This is similar to the
calculated composition of argentojarosite based on its ideal formula, 18.94, 29.42 and
11.26 wt. for Ag, Fe and S, respectively. The Ag content is considerably closer to the
ideal composition as compared to 14.6% wt. Ag in argentojarosite samples previously
produced by Sasaki et al. (1995).
The unit cell parameters were determined for the 14 day argentojarosite samples
(Table 10). The parameters were relatively constant regardless of the concentration of Ag
in solution and in the solid phase. Parameters a0 and c0 agree well with previously
published data (Sasaki and Konno 2000).
39
Ag+
Contact time
Elemental composition (% wt.)
(mM)
S
0
Fe
Ag
P
K
0.37
<0.0003
0.01
<0.04
8h
N.D.
5d
N.D.
38.05
0.006
0.73
0.04
14 d
0.61
39.01
0.01
0.72
0.45
8h
N.D.
27.81
0.36
0.76
0.04
5d
N.D.
37.41
0.06
0.96
0.43
14 d
9.91
36.97
6.05
0.73
0.34
8h
N.D.
90.33
2.36
2.41
0.08
5d
12.16
25.66
7.39
0.39
0.16
14 d
11.40
32.36
16.12
0.44
<0.04
8h
N.D.
38.46
7.96
0.98
0.38
5d
11.12
31.72
16.55
0.40
<0.04
14 d
11.07
28.99
18.42
0.31
<0.04
20b
14 d
N.D.
42.37
0.35
0.88
0.15
argentojarosite
AgFe3(SO4)2(OH)6
11.26
29.41
18.94
N.D
N.D.
schwertmannite
Fe8O8(OH)6SO4
4.15
57.81
N.D.
N.D
N.D.
5
20
40
a
a
N.D., Not determined. bFrom precipitate separation experiment.
Table 9. Elemental analysis of the samples.
Ag concentration
Parameter (Å)
Volume (Å3)
Solution phase (mM)
Solid phase (% wt.)
a0
c0
5
6.05
7.35
16.57
774.66
20
16.12
7.35
16.53
773.26
40
18.42
7.35
16.54
772.91
Table 10. Unit cell parameters of the 14-day samples.
40
Comparative data analysis
The ICP data for Fe and Ag in the samples showed marked variation with contact
time and Ag concentration. The results showed an increase in Ag content of the
precipitates with time (Fig. 10), which is in agreement with the XRD patterns of the
argentojarosite. The corresponding trends in Fe content with time (Fig. 11) were,
however, inconclusive, for each of 5, 20 and 40 mM concentrations of Ag+ added. The
trends in % wt. of Fe with time are inconclusive (Fig. 11). The relative intensities showed
a positive trend with Ag in solution (Fig. 12). Conversely, a negative association was
observed with the FWHM and Ag in solution, in accordance with sharp, intense peaks for
well-defined argentojarosite (Fig 13). These graphs remain similar if the Ag in solution is
replaced by the Ag content of precipitates.
41
20
18
16
14
% wt. Ag
12
8h
10
5d
8
14 d
6
4
2
0
0
5
20
40
Ag (mM)
Fig. 10. Variation in Ag content of precipitates formed in the presence of 0-40 mM Ag+.
42
100
90
80
70
% wt. Fe
60
8h
50
5d
40
14 d
30
20
10
0
0
5
20
40
Ag (mM)
Fig. 11. Variation in Fe content of precipitates formed in the presence of 0-40 mM Ag+.
43
800
700
Relative intensity
600
500
400
300
200
100
0
0
5
20
40
Ag (mM)
Fig. 12.Variation in relative intensity of the 3.07 Å peak in the presence of 0-40 mM Ag+.
Contact time 14 days.
44
0.14
0.12
FWHM
0.1
0.08
0.06
0.04
0.02
0
5
20
40
Ag (mM)
Fig. 13. Variation in FWHM values for the 2.7 Å peak in the presence of 5-40 mM Ag+.
Contact time 14 days.
3.3 Separation of the solid and solution phases
The specific role played by the solution and solid phases during transformation of
schwertmannite-like precursor to argentojarosite was studied by separating the two
phases and adding 20 mM Ag+ to each phase (Fig. 14). For the solution phase, even when
the samples were collected for analysis after 14 days of incubation (pH 2.0±0.1) since
45
addition of Ag+, there was not enough precipitate for further analysis. For the solid phase,
the precipitate was collected after 14 days of incubation with 20 mM Ag+ and the XRD
data are presented in Fig. 15. The nature of the XRD pattern suggests the presence of
poorly crystalline compounds, and traces of an unidentified phase with three sharp peaks.
Argentojarosite was not identified in this XRD pattern. This sample contained 42.37% Fe
and 0.35% Ag (Table 9).
These findings suggest that either the solid or the solution phase was not sufficient for
argentojarosite synthesis within the 14-day time course. The solution phase may need a
seeding effect to overcome the slow kinetics and promote jarosite precipitation, and this
has been shown to enhance the rate of jarosite precipitation (Dutrizac 1996). When the
solution phase (i.e., biologically produced ferric sulfate) was held at room temperature
beyond 14 days, precipitation started over the next several weeks. This precipitate was
not analyzed any further in the present study.
46
Fig. 14. Scheme for separation of solid and solution phases.
47
Fig. 15. XRD pattern for precipitates treated with 20 mM Ag+.
3.4 Synthetic argentojarosite
To further investigate the role of schwertmannite in the formation of argentojarosite,
the effect of silver addition on a reference schwertmannite sample was examined. For this
purpose, 20 mM Ag+ was added to 100 mg of chemically synthesized produced
schwertmannite (sample X-11, Bigham et al. 1996a) and incubated for 14 days at pH
2.0±0.1 conditions. The suspended solids precipitate were collected and analyzed by
XRD. Fig. 16 shows a comparison of the XRD pattern of the chemically synthesized
argentojarosite sample with and without the silver treatment. The untreated sample
showed mostly broad peaks, resembling the broad peaks seen in schwertmannite XRD
patterns. The post-treatment pattern shows the formation of sharper peaks. Most of these
48
coincided with the peak positions of reference argentojarosite, as indicated by the
matching d values.
A sharp peak at 4.18 Å was observed before and after silver treatment, and could be
identified as corresponding to goethite. The relatively long storage time (> 18 years) of
the sample and the meta-stable nature of schwertmannite may explain this finding.
49
b
a
Fig. 16. XRD patterns for a synthetic schwertmannite sample (a) and the same sample
treated with 20 mM Ag+ for 14 days (b). The peak positions are marked for the latter,
with AJ indicating matches with the references argentojarosite pattern.
3.5 Preparation of argentojarosite from ferric sulfate
Preparation of argentojarosite was also attempted by mixing 120 mM ferric sulfate
with 40 mM Ag+. The resultant mix was incubated at room temperature in shake flasks
50
for 14 days. However, after 14 days, separation of precipitates from the solution phase
was not successful through centrifugation. Chemical synthesis in previous reports have
involved elevated temperatures in the range of 90-140 ºC (May et al. 1973, Dutrizac and
Kaiman 1976), suggesting that the rate of formation is slow at ambient conditions,
possibly also involving a relatively high activation energy for the reaction.
51
4. Conclusions and Future Directions
Argentojarosite was successfully synthesized using ferric sulfate solutions produced
by Acidithiobacillus ferrooxidans and subsequently amended with Ag-nitrate. Because
dissolved silver is highly toxic to A. ferrooxidans, a two-step protocol for the synthesis of
argentojarosite was developed to separate the precipitation of argentojarosite from the
oxidation step. XRD analysis showed that the sample precipitates matched perfectly with
the argentojarosite standard in the database. The role of the bacteria remains equivocal in
these experiments. While bacteria oxidized ferrous iron to ferric iron, their specific role
in the precipitation of argentojarosite is not clear.
Increased contact time and Ag+ concentration improved the crystallization of
argentojarosite, evident from sharper XRD peaks and decreased the FWHM values. The
sample with the highest Ag concentration (40 mM) and the longest contact time (14 days)
produced the most crystalline specimen of argentojarosite. Its elemental composition was
closest to those of the theoretically calculated values and it was also the most yellow in
color. The color and the surface area were influenced by the presence of poorly
crystalline Fe(III) precipitates even at the highest Ag+ concentration and the longest
contact time.
52
Synthesis of argentojarosite through silver treatment of 120 mM ferric sulfate
solution did not yield desired precipitates over the period of 14 days. Argentojarosite
could not be successfully precipitated within 14 days when the solution and solid phases
were separated after the generation of the amorphous Fe(III) phase. This observation
suggests that both the phases are required for argentojarosite formation. Silver treatment
(20 mM, 14 days) of the authentic schwertmannite sample produced a mixed solid phase
that showed partial transformation of schwertmannite to argentojarosite.
Future experiments could be designed with longer contact times in order to reach
equilibrium in argentojarosite formation. Determination of mass balance under those
conditions would provide useful information on the quantitative aspects of argentojarosite
formation. Experiments leading to a study of the kinetics of the precipitation reaction
would help understand the limiting conditions for the precipitation. Combining the
information from mass balance and kinetic experiments could be used for modeling the
order of the reaction and calculating thermodynamic properties such as solubility
constants and saturation index values.
53
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