Hydrometallurgy 146 (2014) 96–110
Contents lists available at ScienceDirect
Hydrometallurgy
journal homepage: www.elsevier.com/locate/hydromet
Review
Chalcopyrite hydrometallurgy at atmospheric pressure: 2. Review of
acidic chloride process options
H.R. Watling ⁎
CSIRO Minerals Down Under, CSIRO Process Science and Engineering, P.O. Box 7229, Karawara, W.A. 6152, Australia
a r t i c l e
i n f o
Article history:
Received 30 August 2013
Received in revised form 24 February 2014
Accepted 5 March 2014
Available online 29 March 2014
Keywords:
Chalcopyrite dissolution
Chloride lixiviant
Seawater
Heap leaching
Bioleaching
a b s t r a c t
Hydrometallurgical process developments for the extraction of copper from chalcopyrite tend to target complex
concentrates, dirty concentrates that would incur penalties if smelted or low-grade ores that are thus far an uneconomic source of copper. Perceived advantages of chloride systems are the higher solubilities of copper and
iron, the ease of ferrous ion oxidation and faster leaching kinetics of chalcopyrite compared with ferric sulfate
systems, and the generation of sulfur rather than sulfate as the product of sulfide oxidation. Process developments for concentrates employ acidic, oxidising leach media containing sodium or other chloride salts and temperatures up to the boiling points of the high-concentration solutions. In those processes, chloride ion is thought
to be an active agent in the dissolution mechanism. Leaching conditions fall into two groups, those targeting
Cu(II) and those targeting Cu(I) in pregnant leach solutions. For low grade ores, usually processed in heaps,
the use of seawater or other naturally saline water in leaching operations may be an ‘economic’ choice to overcome the scarcity and/or cost of freshwater. Few studies have been published describing the advantages and disadvantages of seawater substitution for freshwater in leaching processes but, from the sparse information
available, seawater appears to be as efficient a solvent and carrier of acid and oxidant as freshwater. The recent
descriptions of some iron(II)- and sulfur-oxidising, salt-tolerant acidophilic microorganisms indicate that a
diverse group of microorganisms that could function in sulfide heaps irrigated with seawater await discovery.
With regard to processing using seawater instead of freshwater, the salt content in seawater would impact
directly on solution transport costs to and round a mine (through increased solution viscosity and specific
gravity) and could adversely affect product and by-product purity.
© 2014 Elsevier B.V. All rights reserved.
Contents
1.
2.
3.
4.
5.
Introduction . . . . . . . . . . . . . . . . . . . . . . .
1.1.
Sulfate systems reprised . . . . . . . . . . . . . .
1.2.
Scope of this review . . . . . . . . . . . . . . . .
Benefits and disadvantages of chloride leaching . . . . . . .
Acidic chloride systems for chalcopyrite concentrate . . . . .
3.1.
Comparison of studies . . . . . . . . . . . . . . .
3.2.
Ferric chloride as oxidant . . . . . . . . . . . . . .
3.3.
Cupric chloride as oxidant . . . . . . . . . . . . .
3.4.
Mixed ferric chloride–cupric chloride oxidants . . . .
Hybrid chloride–sulfate systems applied to concentrates . . .
4.1.
Oxygen as oxidant in hybrid chloride–sulfate system .
4.2.
Ferric ion as oxidant in hybrid chloride–sulfate system
4.3.
Chlorate as oxidant . . . . . . . . . . . . . . . .
Other oxidants . . . . . . . . . . . . . . . . . . . . . .
5.1.
Cl2–Br2 . . . . . . . . . . . . . . . . . . . . . .
5.2.
High MgCl2–low HCl–FeCl3 . . . . . . . . . . . . .
5.3.
Gaseous Cl2–aqueous Cl2–HClO . . . . . . . . . . .
⁎ Corresponding author.
E-mail address: Helen.Watling@csiro.au.
http://dx.doi.org/10.1016/j.hydromet.2014.03.013
0304-386X/© 2014 Elsevier B.V. All rights reserved.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
97
97
97
97
98
98
99
100
101
102
102
102
103
103
103
103
104
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
5.4.
HClO4 . . . . . . . . . . . . . . . . . . . . . . .
Acidic chloride systems for chalcopyrite ores in heaps or dumps
6.1.
Solution chemistry and secondary reaction products . .
6.2.
Bioleaching in chloride heaps . . . . . . . . . . . . .
6.3.
Use of saline water or seawater in heaps or dumps . . .
7.
Summary . . . . . . . . . . . . . . . . . . . . . . . . .
Acknowledgements . . . . . . . . . . . . . . . . . . . . . . .
References . . . . . . . . . . . . . . . . . . . . . . . . . . .
6.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
1. Introduction
The need to process low-grade and/or complex chalcopyritecontaining ores (see Table 1 for minerals referred to in the text
and their ideal formulae) that cannot be concentrated has been the
main driver for the development of hydrometallurgical processes.
Other drivers are the current imbalance between copper supply and
demand, the overall decline in ore grades and the extensive exploitation of low-grade oxide and secondary sulfide ores that may eventually leave large quantities of low-grade chalcopyrite ores as a major
but, thus far, uneconomic source of copper. Typically, for large
near-surface deposits, average copper cut-off grade for conventional
processing is approximately 0.4% Cu (British Geological Survey,
2007), from which it may be deduced that the term “low grade” refers to ores with b 0.4% Cu.
This review comprises the second part of an update on the status
of copper extraction from chalcopyrite under atmospheric conditions,
either in concentrated form or in low-grade ores. In the first part,
sulfate-based systems operated at atmospheric pressure were described
and compared (Watling, 2013 and references therein), with the aim of
informing researchers, metallurgists and plant operators of the wide variety of chemical systems that might be applied in the future. This second part of the review is focused on the use of chloride systems for
the extraction of copper from chalcopyrite. Developments using chloride fall into two groups, (i) those employing acidic, oxidising leach
media containing sodium or other chloride salts up to concentrations
encountered in brines, at temperatures up to the boiling points of the
selected solution compositions (e.g., Hyvärinen and Hämäläinen,
2005) and (ii) those in which naturally saline water is substituted for
freshwater water in leaching operations where freshwater is scarce
Table 1
Minerals and their ideal formulae.
Copper minerals
Ideal formula
Bornite
Chalcocite
Chalcopyrite
Covellite
Other minerals
Albite
Alunitea
Ferrihydrite
Goethite
Gypsum
Hematite
Jarositea
Magnetite
Marcasite
Muscovite
Phlogopite
Pyrite
Pyrrhotite
Schwertmannite
Silicab
Vermiculite
Zeolite (natrolite)
Cu5FeS4
Cu2S
CuFeS2
CuS
a
b
NaAlSi3O8
KAl3(SO4)2(OH)6
5Fe2O3.9H2O
FeOOH
CaSO4.2H2O
Fe2O3
KFe3(SO4)2(OH)6
Fe3O4
FeS2
KAl2(Si3Al)O10(OH, F)2
KMg3(Si3Al)O10(OH, F)2
FeS2
Fe1 − xS (x = 0–2)
Fe8O8(OH)6(SO4)·nH2O
SiO2
(Mg,Fe,Al)3(Al,Si)4O10(OH)2·4H2O
Na2Al2Si3O10.2H2O
Potassium or other monovalent cations.
Silica varieties include amorphous, colloidal and gel.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
97
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
104
105
105
106
106
107
108
108
(e.g., Dreisinger, 2009). Fundamental studies on the mechanisms of
chloride leaching have been a common feature of process development
among the first group of processes but are largely absent from the few
processes involving the substitution of saline water for freshwater.
1.1. Sulfate systems reprised
The first part of the review covered the extraction of copper from
chalcopyrite using sulfate, sulfate–chloride or sulfate–nitrate leach
media (Watling, 2013). In summary, many studies generated
data consistent with leaching rates being largely independent of acid
concentration beyond that required to solubilise a sufficient concentration of ferric ions to react with available chalcopyrite surfaces, but dependent on sulfate concentration and solution oxidation reduction
potential (ORP), the optimum ORP being dependent on ferrous
ion and cupric ion concentrations. Topics concerning the influence of
chalcopyrite crystallographic structure and the formation of secondary
overlayers on chalcopyrite surfaces were also discussed. Copper extraction rates were enhanced by increased temperature, the presence of
some microorganisms or by the addition of a chloride salt. However, sulfate processes with the addition of nitric acid or a nitrate salt were less
well developed and the potential benefits remain poorly defined at this
time. The efficiencies of sulfate leaching systems with superior-strength
oxidants (compared with ferric ions) were discussed. For the most part,
these were studied at laboratory scale but are yet to be exploited at
commercial scale. The selected alternative oxidants were more costly
than ferric ions, but some offered advantages in terms of extraction
efficiency and kinetics, and further studies are warranted.
1.2. Scope of this review
In this second part of the review on chalcopyrite hydrometallurgy at
atmospheric pressure, chloride-based leaching systems are described
and compared. The majority of those processes targeted the processing
of copper concentrates and were operated at atmospheric pressure and
at temperatures approaching the boiling point of ferric chloride solutions (Table 2). A few higher-temperature processes conducted in pressure vessels have also been described by McDonald and Muir (2007a, b)
and the effects of adding sodium chloride to them investigated. These
pressure oxidation processes are not considered further in this review.
No account is taken of the possible economics of processing, but rather
the aim is to inform researchers, metallurgists and plant operators about
the wide variety of chemical systems that might be applied in the future
when copper demand is higher, ore grades are lower and new technologies, particularly for reagent recovery and recycle, have been developed. While the advantages or disadvantages of current technologies
may be referred to in the context of reported results or applications of
specific systems, detailed accounts of the engineering of such technologies, their management and/or control are outside the scope of the
review.
2. Benefits and disadvantages of chloride leaching
The strong interest in chloride systems resides in: (i) the increased
solubilities of iron and other metals; (ii) enhanced redox properties
98
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
Table 2
Some acidic chloride-based processes operated at atmospheric pressure for the extraction of copper from CuFeS2 or other copper sulfide concentrates indicating the variety of lixiviants/
oxidants employed.
1
2
3
4
Process
Leach reagents and initial concentrations
T
(°C)
Size
(μm)
Scale
U.S. Bureau of Mines (Haver and Wong, 1971)
UBC-Cominco (Milner et al., 1974; Muir and Dixon, 2002)
CYMET (Atwood and Curtis, 1974, 1975)
Minemet Recherche (Demarthe et al., 1976; Guy and
Broadbent, 1983; Guy et al., 1983)
Duval CLEAR process (Dutrizac, 1992; Hoffmann, 1991;
Schweitzer and Livingston, 1982)
Elkem process (Andersen and Boe, 1985)
4 M FeCl3; 0.24 M HCl; possible CaCl2 addition to control sulfate.
1–3.6 M FeCl3; possible CaO or CaCl2 addition to control sulfate.
1.8 M FeCl3–1.1 M CuCl2–3.94 M NaCl.
0.4 M CuCl2–0.1 M HCl–4.3 M NaCl.
106
95–100
95–98
107
−44
−44
n.s.
n.s.
P
L
P
P
104
n.s.
C
105–115 n.s.
P
95
n.s.
105–116 −150
P
P
85
n.s.
D
85–95
−100
D
95
−41
L
110
−44«
C
Two stage leach, 0.15 M CuCl2–0.02 MFeCl3 as oxidants in mixed 1.4 M NaCl + 0.6 M
KCl brine; residual CuFeS2 pressure leached with O2 (dissolution with iron rejection).
6
FeCl3 in brine under controlled redox (430–460 mV vs Pt/calomel electrode) such that
all copper remains in solution as Cu(I).
7 Cuprex (Dalton et al., 1987, 1991; Muir and Dixon, 2002)
Excess FeCl3 in two-stage counter-current leach; soluble copper as Cu(II).
8 Neo-ferric technologies high concentration chloride leach
MgCl2 (2.9–4.3 M)–HCl brine; secondary leach contains FeCl3 derived from primary
(Harris and White, 2008; Harris et al., 2007)
leach dissolution of CuFeS2.
9 Intec process (Intec Ltd., 2008; Moyes and Houllis, 2002; P. Solution of 0.47 M Cu(II), 4.9 M NaCl and 0.27 M NaBr has ORP 950–1000 mV vs
TM
Everett, 1996; P.K. Everett, 1996; Taylor and Jansen, 1999) Ag/AgCl due to formation of BrCl−
)
2 (Halex
10 Outotec hydrocopper process (Hyvärinen, and Hämäläinen, 0.3 M CuCl2 in 4.8 M NaCl at pH 2 (HCl);
1999, 2005; Lundström et al., 2009)
11 Falconbridge process(es) (Liddicoat and Dreisinger, 2007;
FeCl3 (~0.5 M Fe for ‘goethite’ process and ~1.5 M Fe for ‘hematite’ process)–CuCl2
Lu and Dreisinger, 2013a, 2013b)
(~0.8 M Cu) in 0.08 M free HCl (pH 0 or lower) with CaCl2 (2.75 M Ca); (total Cl N5 M).
12 Sumitomo (Imamura et al., 2006; Makino et al., 1996)
A two stage process: (i) Cu(II) and Fe(III) in solution reduced by reaction with CuFeS2;
N5.7 M total chloride; (ii) partially leached CuFeS2 reacted with Cl2 gas.
5
C, commercial; D, demonstration; P, pilot; L, laboratory; n.s., not specified; «, ORP control by fine grinding (increased surface area).
Note: Flowsheets for processes are depicted in cited references as well as some reviews, e.g., Muir and Dixon (2002).
because cupric and cuprous ions were stabilised as chloride complexes
and the Cu(I)/Cu(II) redox couple could contribute to sulfide oxidation
reactions; (iii) the faster leaching kinetics of chalcopyrite compared
with sulfate systems; (iv) the generation of elemental sulfur rather
than sulfate; and (v) low pyrite reactivity in chloride systems
(Demarthe and Georgeaux, 1978; Dutrizac, 1990; Lu and Dreisinger,
2013a; Muir and Dixon, 2002; Senanayake and Muir, 2003).
The production of elemental sulfur rather than sulfate during the
processing of metal sulfide minerals is advantageous because sulfur is
inert and can be stored until market conditions are favourable for its further processing and use. In a hydrometallurgical process, less energy is
used if the sulfide oxidation stops at sulfur rather than progressing to
sulfate production. The adoption of hydrometallurgy rather than pyrometallurgy eliminates the environmental problem of SO2 production
(Spink, 1977). In addition, the production of less sulfuric acid is beneficial to downstream processing of pregnant leach solution (PLS) that
requires less neutralising agents and therefore the generation of lower
volumes of waste materials (e.g. gypsum).
Perceived disadvantages are: (i) the corrosive action of chloride,
thus necessitating the use of more expensive materials of construction
for reactors; (ii) the need for fine grinding for processes operated at atmospheric pressure; (iii) the co-leaching of multiple elements that require additional treatments and (iv) the difficulty of electrowinning
high-grade copper from chloride solutions. Nevertheless, in respect of
the last point, there are four processes in which copper is recovered
from chloride solutions (Lu and Dreisinger, 2013b). Hein and Joly
(2011) surveyed 25 solvent extraction plants and reported that 18 of
them were operated with PLS containing high sulfate, chloride and cation concentrations, four of them with chloride concentrations N35 g L−1
(1 M).
In the Copper Technology Roadmap of 2004 (AMIRA International,
2004), more efficient use of water in the unit processes associated
with extraction, comminution and separation was given “top priority”.
Among proposed strategies for improvements in water efficiencies included in the Roadmap were the use of saline water and recycled
water in processing. The use of seawater, brackish and hypersaline
water and recycled process water have been variously trialled and/or
implemented at mining operations over many years, mainly in
respect of substituting seawater for freshwater in flotation circuits
(e.g., Moreno et al., 2011). There are few reports of leaching with seawater in the public domain (Table 3).
It was noted in the Copper Technology Roadmap that adaptation of
existing processes to different water sources requires better understanding of the impact of water quality on processing operations and
on materials of construction. Some of these have been reviewed by
Philippe et al. (2010). In respect of processing chemistry, they include
viscosity and specific gravity (increased water transport costs), chemical buffering effects (influencing leach chemistry), product and byproduct contamination (clean water required for washing), evaporation
(and therefore water make-up) and capillary forces, and scaling
(secondary precipitation). In addition, Philippe et al. (2010) stated
that the use of sea water in the bioleaching of low-grade sulfide ores
would inhibit beneficial bacterial activity.
3. Acidic chloride systems for chalcopyrite concentrate
The leaching of chalcopyrite in acidic chloride media was studied
extensively in the 1980s and 1990s resulting in a number of proposed
processes accompanied by suggested flowsheets. Winand (1991) and
Senanayake and Muir (2003) reviewed some important fundamentals
of chloride leaching systems and noted that, in multi-component systems such as those discussed in this review, solubilities, complexation
reactions and electrochemical reactions are important and possibly limiting factors in the development of processes. In addition Winand
(1991), and also Hoffmann (1991), Dutrizac (1992) and Muir and
Dixon (2002) reviewed individual process developments with brief
descriptions, critical assessments and/or reasons why processes have
not seen long-term commercial success.
In each section that follows, the chemistry of the system is discussed,
based on fundamental laboratory studies. Then example processes that
utilise the chemistry are described briefly.
3.1. Comparison of studies
While the main focus of this section is on acidic chloride processes
for concentrates, the rare examples of the application of such processes
to ores are included. Two main difficulties arise when comparing the results of independent studies, temperature and chalcopyrite surface area.
As is the case for sulfate leaching systems, copper extraction rates
increase as the temperature is increased (Fig. 1), a characteristic of
chalcopyrite dissolution that has been exploited in many processes at
atmospheric (Table 2) and higher pressures. Thus, where possible,
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
99
Table 3
Saline water use at mining operations.
Project
Zaldivar, Chile
Michilla, Chile
Esperanza, Chile
Boleo, Baja California
Cu
Cu
Cu, Au, Mo
Zn, Co, Mn
Mining company
Water source
Unit processes
References
Placer Dome
Antofagasta Minerals
Antofagasta Minerals
Baja Mining Corporation
Brackish
Seawater
Seawater
Seawater
Leaching
Agglomeration, heap leaching
Concentrate production
Tank leaching
Weston et al. (1995)
Aroca (1999); Wiertz (2009)
Chadwick (2009); Parraguez et al. (2009)
Dreisinger et al. (2005); Dreisinger (2009).
inter-study comparisons should be made between studies conducted at
similar temperatures.
The second variable is chalcopyrite particle or grain size used in different studies. Dutrizac (1981) leached various chalcopyrite sized fractions in 0.2 M FeCl3 and 0.3 M HCl solutions, calculated initial rate
constants from the initial slopes and, by plotting the rate constants
against 1/r, where r was the mean particle size in microns, showed
that rates were directly proportional to the surface area of chalcopyrite
being leached. That result was consistent with data obtained for chalcopyrite leached in ferric sulfate–sulfuric acid medium, though rates in
chloride medium were higher than in sulfate medium. The result was
also consistent with the demonstrated advantage of fine grinding in
some process developments conducted at above atmospheric pressure
(e.g., Activox process, chalcopyrite particles 5–10 μm; Dreisinger,
2006). While a brief account of chalcopyrite preparation and measures
of particle size was provided in many of the studies reviewed in this
work, the relevant surface areas were not reported, making it impossible to normalise the data and so make convincing comparisons. Nevertheless, some data obtained at 85–95 °C for the leaching of chalcopyrite
in acidified chloride or hybrid sulfate–chloride media are compared in
Fig. 2 and data obtained at 23–55 °C for the leaching of chalcopyrite in
chloride–chlorine media are compared in Fig. 3.
3.2. Ferric chloride as oxidant
The oxidation of chalcopyrite in ferric chloride media, while faster
than oxidation in ferric sulfate media, was still generally slow for a metallurgical process, thus driving the assessment of many multi-component
systems. The examples summarised below are for laboratory-scale tests
conducted at atmospheric pressure and temperatures up to approximately 100 °C. Most of these processes were developed for the extraction of
copper from chalcopyrite concentrates.
Dutrizac (1981) examined chalcopyrite dissolution in solutions
containing 0.3 M HCl and 0.2 M FeCl3 at b100 °C (reaction (1)) and
summarised the results as follows: (i) leaching rates were higher in
chloride media than in sulfate media at temperatures N50 °C, with a
three-fold increase in leaching rate for equivalent-sized chalcopyrite
particles; (ii) the rate of chalcopyrite dissolution in ferric chloride
medium was independent of temperature in the range of 45–100 °C, essentially independent of acid concentration, but directly proportional to
chalcopyrite surface area.
CuFeS2 þ 4FeCl3 ¼ CuCl2 þ 5FeCl2 þ 2S
0
ð1Þ
The study by O'Malley and Liddell (1987) was focused on the formation of copper(I) species during the ferric chloride leaching of a ground,
natural chalcopyrite sample. Under test conditions in which ferric ion
concentrations were limiting, they showed that the extent of copper
extraction depended on the initial ferric chloride concentration and
that the total chloride concentration controlled the extent to which
copper(II) was reduced to copper(I) at the chalcopyrite surface (reaction (2)). Wang (2005) also proposed that Cu(I) species had a role in
ferric chloride leaching of chalcopyrite and suggested that reactions (1)
and (3) were both principal reactions in the ferric chloride system.
0
ð2Þ
0
ð3Þ
CuFeS2 þ 3CuCl2 ¼ 4CuCl þ FeCl2 þ 2S
CuFeS2 þ 3FeCl3 ¼ CuCl þ 4FeCl2 þ 2S
The rate of dissolution of monosized chalcopyrite in ferric chloride–
sodium chloride media acidified with hydrochloric acid was studied to
establish the mechanism of leaching (Palmer et al., 1981). The linear
kinetics were consistent with the rate being controlled by surface
phenomena and being dependent on both chloride and ferric ion concentrations. Yoo et al. (2010) also focused on leaching mechanisms
when they investigated the leaching of chalcopyrite in FeCl3–HCl and
FeCl3–H2SO4 media, and mixtures of the two. In rank order, copper
extraction was Cl N Cl–SO4 ≫ SO4 medium. The authors cited the data
of Lin et al. (1991), who noted that copper existed as cupric ions in
solutions with low chloride concentration but that, at high chloride con2−
3−
centrations, cuprous ion species such as CuCl, CuCl−
2 , CuCl3 and CuCl4
were formed and that the standard potential between cuprous and cupric ions increased. The results were consistent with those of O'Malley
and Liddell (1987). After conducting a thermodynamic study of copper
species in chloride solutions, Yoo et al. (2010) concluded that the
100
(5)
20
Cu extraction [%]
15
70
10
60
50
40
20
5
Cu extraction [%]
80
T °C
80
(3)
60
(4)
40
(2)
20
(1)
0
0
1
2
3
Time [hours]
0
0
1
2
3
4
Time [hours]
Fig. 1. Effect of temperature on copper extraction from ground chalcopyrite in 1 M FeCl3–
0.2 M HCl medium (redrawn from Havlik and Kammel, 1995).
Fig. 2. Copper extraction from chalcopyrite leached in chloride media at 85–90 °C. Data
from: (1) Al-Harahsheh et al. (2008), 0.5 M FeCl3, 90 °C; (2) Al-Harahsheh et al. (2008),
0.5 M FeCl3–0.025 M CuCl2, 90 °C; (3) Bonan et al. (1981), 0.5 MCuCl2–0.1 M HCl–4 M
NaCl, 95 °C; (4) Ruiz et al.(2011), 0.2 M H2SO4–0.6 M NaCl–O2 (gas sparge), 90 °C; and
(5) Xian et al. (2012), 0.5 M NaClO3–1 M HCl, 85 °C.
100
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
(1)
(2)
(3)
100
Cu extraction [%]
80
60
40
20
0
0
2
4
6
8
(4)
(5)
10
Time [hours]
Fig. 3. Copper extraction from chalcopyrite leached in chloride/chlorine media at
23–55 °C. Data from: (1) Brocchi and Jena (1992), slurry chlorination, 30 °C; (2) Cho
(1987), 0.228 M HClO, pH 3.6–4.4, 23 °C; (3) Çolak et al. (1987), water saturated with
Cl2, 33 °C; (4) Dutrizac (1982), 0.1 M FeCl3–0.3 M HCl, 55 °C; (5) Velásquez-Yévenes
et al. (2010b), 0.008 M Cu2+–0.2 M HCl–0.34 M NaCl, ORP 550–620, 35 °C.
increase in the critical potential caused by cuprous ion species in a chloride solution was a key parameter in the faster chalcopyrite leaching
rate.
Elemental sulfur is the major sulfur-reaction product when chalcopyrite is oxidised in ferric chloride media (reactions (1)–(3)). Dutrizac
(1990) reported that more than 95% of the sulfide moiety of chalcopyrite was oxidised to elemental sulfur in medium containing 0–2 M
FeCl3 and 0–3 M HCl at 95 °C, with only 5% of the sulfide being oxidised
to sulfate, independent of leaching time (0–90 h). Dutrizac (1990)
noted that small chalcopyrite grains rapidly became enveloped in elemental sulfur and that the sulfur morphology was independent of either
ferric chloride or hydrochloric acid concentrations. In a similar investigation, Rath et al. (1988) conducted dissolution tests at temperatures
up to 100 °C and concluded that the chalcopyrite oxidation rate was directly proportional to the square root of the ferric chloride concentration and inversely proportional to particle diameter. In other studies,
different sulfur morphologies were described (Hirato et al., 1986; Lu
et al., 2000; Majima et al., 1985) but relationships between them and
leaching conditions, particle sizes or retention times are yet to be
elucidated.
Saxena and Mandre (1992) studied the ferric chloride leaching (0.2
M FeCl3) of chalcopyrite in an ore containing 0.75% Cu. Agitated batch
leaching tests of up to 6 h duration were conducted in the temperature
range of 30–90 °C using the −104 + 74 μm size fraction of the ore. In
those tests, the solution and ore were both pre-heated before being
contacted. The results indicated that the process was controlled initially
by chemical reaction between ferric chloride and chalcopyrite but that,
subsequently, the rate was controlled by diffusion through the sulfur
product layer.
Example processes utilising ferric chloride as oxidant include the
USBM process (Haver and Wong, 1971; Haver et al., 1975) and the
UBC-Cominco Process (Milner et al., 1974; Muir and Dixon, 2002).
The purpose of the USBM investigation (Haver and Wong, 1971) was
to develop an alternative method of obtaining copper from chalcopyrite
without the evolution of SO2. Ferric chloride was chosen as the oxidant
because it had the correct oxidation potential to convert sulfide to elemental sulfur (reaction (3)). Conditions were: 4 M FeCl3, with a ratio
of 2.7:1 FeCl3:CuFeS2; 106 °C and 2 hour leach duration. The small
amount of sulfur that was oxidised to sulfate could be precipitated as
gypsum with the addition of CaCl2 to the solution. The leachate
contained 1 M Cu (60:40 Cu(I):Cu(II)), 4 M Fe and only 0.03 M SO2−
4
and the residue contained 16 wt.% Fe, 72 wt.% S (approximately 50% of
which was elemental sulfur) and 0.024 wt.% CaO. The original
flowsheet, utilising solvent extraction to remove the sulfur, cementation to recover copper from solution and chlorination to regenerate
the ferric chloride, was estimated to be too expensive for commercial
acceptance (Haver et al., 1975). Subsequent cost-reducing modifications to the process were the extraction of sulfur from residues using
aqueous ammonium sulfide, direct electrowinning of copper in a diaphragm cell and regeneration of ferric ions by aeration.
Milner et al. (1974) described a closed-cycle hydrometallurgical process in which the high-quality products metallic copper, metallic iron
and elemental sulfur were produced. A solution containing 1–3.6 M Fe
as FeCl3 was used to leach chalcopyrite (9–12 h at 95–100 °C) (reaction (3)) and CaCl2 or CaO addition was used to control sulfate concentration. Thus the leach chemistry of the UBC-Cominco process was
almost identical to that of the USBM process, yielding solutions with a
1:1 ratio Cu(I):Cu(II) (Muir and Dixon, 2002). In the UBC-Cominco process, the hot filtered solution was treated with copper metal to reduce
Cu(II) to Cu(I) and insoluble CuCl, crystallised from the cooled solution,
was converted to high purity copper powder by hydrogen reduction.
3.3. Cupric chloride as oxidant
In the previous section, the results discussed were obtained from
studies in which the starting acidic solutions contained ferric chloride.
The distinction is made in this section, that the studies were initiated
in acidic solutions containing cupric chloride. Clearly, once chalcopyrite
dissolution had commenced (e.g., reactions (2), (4) and (5)), ferrous
ions would be released to solution during chalcopyrite oxidation, potentially become oxidised to ferric ions and increase the contribution of reactions such as (1) and (3) to chalcopyrite dissolution.
CuFeS2 þ H2 SO4 ¼ CuS þ FeSO4 þ H2 S
3
1
ð4Þ
0
CuFeS2 þ CuCl2 þ =4 O2 ¼ 2CuCl þ =2 Fe2 O3 þ 2S
ð5Þ
Nicol and co workers (Miki and Nicol, 2011; Nicol et al., 2010;
Velásquez-Yévenes et al., 2010a, 2010b) conducted a wide ranging
study of the dissolution of chalcopyrite in chloride solutions containing
cupric ions and dissolved oxygen. Initially, it was shown that the rate of
chalcopyrite dissolution was enhanced when leaching was conducted in
an ORP range of 550–620 versus SHE at 35 °C in a solution containing:
0.2 M HCl, 0.008 M Cu2 + (simulating raffinate) and 5–15 mg L− 1
dissolved O2 (Velásquez-Yévenes et al., 2010a). Leaching at ORP
b540 mV caused reduced rates of chalcopyrite dissolution and covellite
or chalcocite formed on some chalcopyrite surfaces.
In the extension of their work, Velásquez-Yévenes et al. (2010b) focused on the kinetics of chalcopyrite dissolution under conditions that
might be expected in a heap leach operation. Experiments were conducted at 35 °C with controlled ORP and initial test solutions were 0.2
M HCl with 0.008 M Cu2 +. However, the addition of ferrous ions or
the conduct of tests at low pH caused difficulties in controlling the
ORP. The presence of a small concentration of cupric ions was essential
but increased copper concentrations did not result in enhanced chalcopyrite dissolution rates. Dissolution rates were not greatly affected by
changes in the total chloride concentration, which was varied by adding
sodium chloride. The strong dependences of dissolution rates on both
temperature and particle size (surface area) were consistent with
many other literature data (e.g., Ikiz et al., 2006; Naderi et al., 2011;
Skrobian et al., 2005; Xian et al., 2012).
The goal of the third part of the study (Nicol et al., 2010) was to describe a reaction mechanism consistent with the newly acquired data
and their previous results. In ancillary experiments it was shown that:
(i) the chalcopyrite surface was converted to a covellite-like phase during leaching at potentials below the potential window; (ii) at ‘normal’
leach solution pH, elemental sulfur formed mainly as isolated globules
and seldom on the chalcopyrite surface, suggesting that an aqueous
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
intermediate sulfur species existed; and (iii) finely ground pyrite
catalysed the reaction and became covered in layers of sulfur. Nicol
et al. (2010) proposed a reaction model in which, within the potential
window, chalcopyrite dissolved partially to form H2S and a covellitelike surface species (reaction (4)) and that the H2S was oxidised by oxygen in a reaction catalysed by cupric ions in two stages (reactions (6)
and (7)). The authors noted that in the proposed reactions, the species
represented as Cu2 + and Cu+ included both aquo-ions and chlorocomplexes that varied according to solution composition. Nicol et al.
(2010) supported the hypothetical reaction model with a detailed
study of the kinetics of copper-catalysed oxidation of H2S. Subsequently,
in the fourth part of their study, Miki and Nicol (2011) modified their
proposed mechanism to account for a newly-detected, intermediate
peroxide species (reaction (7) replaced by reactions (8), (9) and (10)).
2þ
H2 S þ 2Cu
þ
þ
0
¼ 2Cu þ S þ 2H ðrapidÞ
þ
þ
2þ
4Cu þ O2 þ 4H ¼ 4Cu
þ
þ 2H2 OðslowÞ
þ
Cu þ O2 ¼ Cu ·O2
þ
þ
ð7Þ
ð8Þ
þ
2þ
þ HO2
−
ð9Þ
þ
2þ
þ 2H2 O
ð10Þ
Cu ·O2 þ Cu þ H →2Cu
þ
ð6Þ
extraction if it were to form on chalcopyrite surfaces was not observed.
Turkmen et al. (2012) analysed leaching kinetics using the shrinking
core model, but failed to report whether the elemental sulfur had formed
on chalcopyrite surfaces, which is a basic assumption of that model.
Example processes utilising cupric chloride as the initial oxidant include the Minemet Recherche Process (Demarthe et al., 1976, 1977; Guy
and Broadbent, 1983; Guy et al., 1983) and the HydroCopperTM Process
(Hyvärinen and Hämäläinen, 1999, 2005; Lundström et al., 2005, 2009).
The Minemet Recherche process was developed for the extraction of
Cu, Pb and Zn from complex sulfide concentrates (Demarthe et al.,
1977). The sulfide concentrate was leached in 0.4 M CuCl2–0.1 M HCl–
4.3 M NaCl solution (pH ≤1) at temperature 107 °C for 3 h. This reaction
mixture was filtered and separated into two parts. One part of the solution was subjected to air oxidation (pH 1–3, temperature N90 °C);
precipitated goethite was removed before the solution was recycled to
leach fresh concentrate. Copper was partially extracted from the other
portion of the solution using an organic extractant. The partially leached
residue was treated in a second dissolution stage under the same conditions as the first. Advantages of the two stage dissolution process were
iron removal and elemental sulfur as the main sulfur-containing reaction
product, in accord with reactions (2) and (11) (Demarthe et al., 1977).
2þ
2Cu þ H2 O2 þ 2H →2Cu
In a recent mechanistic study Cai et al. (2012) showed that, when
treated in sealed vessels for a month with hydrochloric acid (with and
without cupric ions) at temperatures up to 100 °C, chalcopyrite was
transformed into chloride-rich covellite-like phases of different stoichiometries (e.g., CuS1–0.5xClx; CuS0.5Cl0.5; reaction (4)) and that some of
the iron leached from the chalcopyrite was precipitated as an iron
oxide thought to be hematite (Fe2O3; e.g., reaction (5)).
Key conditions used for the study by Bonan et al. (1981) were as
follows: ground chalcopyrite particles (sized fractions in the range of
25–90 μm, 5 g) were suspended in 500 mL of 0.1–0.5 M Cu(I)–Cu(II)
chloride solution with 0.1 M HCl and 2.5–4 M NaCl in a reactor from
which oxygen was largely eliminated by sparging with nitrogen gas
(see also a similar study by Tchoumou and Roynette, 2007). Removal
of oxygen was necessary to prevent reaction (7) from proceeding, as
this reaction would cause an increase in ORP during the experiment.
Strong chloride solutions were chosen to promote reaction (2) and sta2−
bilise the cuprous chloride complexes, mainly CuCl−
(Lin
2 and CuCl3
et al., 1991); the Cu(II)/Cu(I) ratio was varied in the range of 0.5–7.
Initial ORP values for each solution composition were not reported but
for selected tests, were in the range of 555–585 mV (versus SHE). The
results showed that chalcopyrite leached faster in solutions of higher
Cu(II)/Cu(I) ratio and as the chloride concentration and/or temperature
was increased. Leaching rates were primarily dependent on ORP regardless of the Cu(II)/Cu(I) ratio and chloride concentration that created the
condition. In all tests, sulfide was converted to elemental sulfur with
almost no sulfate formation.
Padilla et al. (1997) gave a good account of the chemistry of the
CuCl2–NaCl–O2 system for chalcopyrite leaching. They studied oxygenated brine solutions containing up to 5 M chloride ions in tests conducted
at atmospheric pressure and temperatures up to 105 °C and reported that
high copper extraction could only be achieved at temperatures near the
boiling points of the solutions. In separate but similar studies Skrobian
et al. (2005) and Turkmen et al. (2012) examined the leaching of chalcopyrite concentrate in CuCl2–NaCl–HCl. The results of both studies indicated that the addition of cupric ions strongly enhanced copper extraction
but that the addition of sodium chloride enhanced copper extraction to
a lesser extent. The nature of the elemental sulfur formed during
CuCl2–HCl leaching of chalcopyrite was not investigated in either study.
Skrobian et al. (2005) noted that it is commonly accepted that sulfur precipitated on chalcopyrite surfaces can slow the rate of extraction, reported that sulfur was partially oxidised to sulfuric acid by cupric chloride,
and ‘concluded’ that the formation of a compact layer that would slow
101
2Fe
þ
2þ
þ 4Cu þ 1:5O2 þ H2 O⇆2FeOðOHÞ þ 4Cu
ð11Þ
Some of the above cited laboratory studies may have contributed to
the development of the HydroCopperTM process, a CuCl2–NaCl system
operated at pH 1.5–2.5 and close to 100 °C (Hämäläinen, 2005;
Hyvärinen and Hämäläinen, 2005; Lundström et al., 2005). Chalcopyrite
is leached by the strong CuCl2 (0.3 M)–NaCl (4.8 M) solution sparged
with oxygen or air to oxidise the product ferrous ions and therefore
facilitate iron(III) rejection as goethite (reaction (11)) or hematite (reaction (5)). The copper is precipitated as cuprous oxide and then converted to metal by hydrogen reduction. A standard chloralkali cell is
used to regenerate the reagents chlorine, caustic and hydrogen.
Hyvärinen and Hämäläinen (2005) make a number of claims in respect
of the process: (i) capital costs mid way between heap leaching and
pressure leaching, depending on ore mineralogy and some external
factors; (ii) suited to smaller operations, e.g., 20–150 thousand tonnes
per annum capacity; (iii) applicable to low quality concentrates; and
(iv) operating costs approximately US 20–30 cents per kilogram
depending on feed mineralogy and energy costs. These attributes, together with efficient extraction of copper and precious metals, iron(III) (and
arsenic) rejection, various process routes for solution purification and
reagent generation, make this a promising technology but the absence
of independent assessments in the public domain and/or commercial development thus far may be indicative of unresolved operational issues.
3.4. Mixed ferric chloride–cupric chloride oxidants
According to Parker et al. (1981), a mixed oxidant of ferric and cupric
ions in chloride medium was an effective oxidant for chalcopyrite
because the reduction of copper(II) to copper(I) occurred faster than
the reduction of iron(III) to iron(II) at the chalcopyrite surface (reactions (2) and (3)). Iron(III) or oxygen subsequently oxidised the
cuprous ions formed in the surface reaction to regenerate the cupric
ion oxidant. In addition to faster reaction kinetics with added chloride,
sulfur crystallinity was increased and there was minimal oxidation of
sulfide or sulfur to sulfate (Senanayake and Muir, 2003). Al-Harahsheh
et al. (2008) examined the catalytic effect of cupric ions on the oxidation
of chalcopyrite in ferric chloride medium, especially noting the effect
of agitation. They explained the observed reduction in chalcopyrite
oxidation rates in agitated systems, compared with stagnant systems,
as being due to the removal of cupric chloride complexes formed at
the interface between chalcopyrite surfaces and ferric ions and proposed that cupric ion was acting as a secondary oxidant in the dissolution reaction. These authors also noted that chalcopyrite oxidation
102
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
was faster in a closed vessel (without agitation) than in an open vessel
but could not provide an explanation for this phenomenon.
The chemical system cupric chloride–ferric chloride–halide (single
or mixed salts) was a key component of the Cymet process (Allen
et al, 1973; Atwood and Curtis, 1975; Kruesi, 1972; Paynter, 1973),
the Duval CLEAR process (Schweitzer and Livingston, 1982) and the
Falconbridge process (Liddicoat and Dreisinger, 2007).
In the CYMET process, described as a “hydrometallurgical process for
pollution-free recovery of metallic copper from chalcopyrite” (Atwood
and Curtis, 1974), copper sulfides were almost completely dissolved
(107 °C) in two stages (oxidation and reduction) with the production
of elemental sulfur (reaction (3)) and concomitant cupric chloride reduction to cuprous chloride. The process required a molar ratio of at least 4:1
FeCl3:CuFeS2. The cuprous chloride was prevented from precipitating in
the strong chloride solution or, in a later process modification (Atwood
and Curtis, 1975), a mixed NaCl–KCl solution, the KCl assisting the rejection of iron and sulfate as K-jarosite (reaction (12)). In a subsequent iron
and sulfate control stage, iron(II) was re-oxidised to iron(III) and excess
amounts precipitated with sulfate ions. The CYMET process was operated
for a period of about 10 years but shut down in 1982, ostensibly because
of a slump in the copper industry (Hoffmann, 1991).
3þ
3Fe
þ
2−
þ 2SO4 þ 6H2 O þ M →MFe3 ðSO4 Þ2 ðOHÞ6 þ 6H
þ
ð12Þ
+
where M = K+, Na+, NH+
4 or H3O
The Duval Clear (Copper Leach, Electrolysis and Regeneration)
process (Schweitzer and Livingston, 1982) applied similar chemistry
for the initial partial leaching of chalcopyrite using Fe(III) and Cu(II) as
oxidants in a mixed NaCl–KCl brine at 104 °C (e.g., reactions (1)–(3)).
In the second stage, the final chalcopyrite leaching and oxy-hydrolysis
of iron were combined in a pressure leach at 150 °C. Part of the iron
was precipitated as K-jarosite (reaction (12)). A commercial plant was
operated in Arizona in the early 1980s for a period of 6 years, producing
approximately 80 tonnes copper per day, but the purity of the copper
was not high enough to justify the cost (Ayres et al., 2002). Muir
and Dixon (2002) briefly described some of the technical problems encountered with the CLEAR process. The recently-described, two-stage
counter-current chloride leach developed for Falconbridge (Liddicoat
and Dreisinger, 2007) also applied similar chemistry at temperatures
near the boiling point of the lixiviant. The feed solution contained high
concentrations of ferric chloride (~0.5 M Fe for the proposed ‘goethite’
process and ~1.5 M Fe for the ‘hematite’ process), ~0.8 M Cu as cupric
chloride, 2.75 M Ca as calcium chloride with 0.08 M free HCl
(pH ≤ 0); the total chloride concentration was N5 M.
4. Hybrid chloride–sulfate systems applied to concentrates
4.1. Oxygen as oxidant in hybrid chloride–sulfate system
The advantage of using sulfuric acid and sodium chloride to create a
chloride lixiviant is that those reagents are cheaper than ferric chloride
or cupric chloride. Nevertheless, an oxidant is still required, in this case
oxygen (reaction (13)). There should be no need to add ferric ions to the
system initially; ferrous ions produced during chalcopyrite leaching
would be oxidised to ferric ions and subsequently contribute to overall
copper extraction (reactions (1) and (3)). With careful selection of
leaching conditions, most of the iron and some sulfate can be precipitated as sodium jarosite (reaction (12)).
þ
2þ
CuFeS2 þ O2 þ 4H →Cu
þ Fe
2þ
0
þ 2S þ 2H2 O
ð13Þ
Li et al. (2010) examined the extraction of copper from a chalcopyrite concentrate (size fraction 38–75 μm) for a suite of leach systems.
In tests conducted at 75 °C, they measured faster kinetics for the NaCl
(0.25 M)–H2SO4 system at pH 2 (97% copper extraction in 170 h) than
at pH 1 (58% copper extraction in 170 h). This difference was attributed
to the increased solubility of iron-containing secondary minerals due to
the formation of FeCl2+ solution species and the consequent lower ferric ion activities in the presence of chloride. Ruiz et al. (2011) reported
that leaching of chalcopyrite concentrate with average particle size of
12 μm in sulfate–chloride solutions was rapid, 90% of the copper being
extracted in 180 min at 100 °C. The presence of 0.5 M chloride ions
(29 g L−1 NaCl) enhanced the leaching rate significantly but the addition of 3 g L−1 Fe3+ caused the ORP to increase and the leaching rate
to slow.
Lu et al. (2000) dissolved finely-ground chalcopyrite concentrate
in solutions of pH b 0.8 (0.8 M H2SO4) containing 1 M NaCl at temperatures in the range of 60–95 °C. They achieved up to 97% copper
extraction in 9-hour tests. Based on their results showing that
chloride concentrations N0.5 M did not enhance the leaching rate
of chalcopyrite (consistent with the results of Palmer et al., 1981),
Lu et al. (2000) concluded that it was important only that there
were sufficient chloride ions present rather than an excess. This finding has commercial implications for geographical areas lacking in
freshwater but able to use saline bore water or seawater; seawater
contains approximately 0.5 M chloride ions. Examination of the
residues showed that the sulfur reaction product obtained in the
presence of 1 M NaCl was crystalline and porous, allowing reactants
to diffuse through the surface product layer to the unreacted mineral
surface.
An application of this leach chemistry is described in a patented
process (Sawyer and Shaw, 1983) for the recovery of copper from a
copper–lead matte using oxygenated acidic sulfate–chloride solutions.
According to Lu et al. (2000), that was possibly the only low-pressure
sulfide leach process to have been commercialised. O'Brien et al.
(1999) piloted a similar process that involved oxygen sparging of an
agitated leach at 80–95 °C and atmospheric pressure. A size fraction
− 100 + 75 μm of the ground chalcopyrite–pyrite ore (4% Cu) was
used for the tests of duration less than 24 h. The pyrite in the ore
promoted chalcopyrite oxidation via galvanic interaction and up to
95% of the chalcopyrite was oxidised compared with only 17% of the
marcasites/pyrite content.
4.2. Ferric ion as oxidant in hybrid chloride–sulfate system
Dutrizac (1981) undertook a critical survey on the ferric ion
leaching of chalcopyrite and augmented the reported data to resolve
some inconsistencies. Using lithium chloride as the additive, Dutrizac
(1981) reported that increased chloride ion concentration in a ferric
sulfate–sulfuric acid leach system resulted in progressively accelerated
copper extractions from chalcopyrite concentrate (− 20 + 14 μm
particle size) at temperatures higher than 50 °C, specifically 2.5 times
at 3–4 M chloride ion, 0.3 M H2SO4 (~ pH 0.2), 0.1 M Fe2(SO4)3 and
90 °C. The addition of lithium chloride was a means of increasing the
chloride concentration without promoting the formation of jarositelike compounds, as would happen if sodium or potassium chlorides
were added. More recently, the influence of sodium chloride on ferric
sulfate oxidation of chalcopyrite was studied at 95 °C using a finelyground chalcopyrite concentrate (5% solids loading; d50 = 5.5 μm)
in 0.9 M Fe2(SO4)3 solution acidified with H2SO4 to pH 0.15 and sparged
with O2 (Carneiro and Leão, 2007). Up to 90% of the copper was extracted with 1 M sodium chloride (58 g L− 1), compared with 45% of the
copper in the absence of sodium chloride. Carneiro and Leão (2007)
attributed the enhanced copper extraction in the presence of sodium
chloride to the reduction of ferric iron concentration (jarosite precipitation; reaction (12)), the formation of cuprous chloride complex ions
and the participation of the Cu(I)/Cu(II) redox couple in oxidation,
and a chloride-induced increase in the surface area and porosity of
the sulfur reaction product. Not specifically mentioned by Carneiro
and Leão (2007) but probably also contributing to enhanced copper extraction would be the acid generated during jarosite precipitation
(reaction (12)).
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
In a 20-day study at 87 °C (Kinnunen and Puhakka, 2004), a much
lower addition of sodium chloride (5 g L− 1; 0.09 M) resulted in enhanced copper extraction in ferric sulfate media initially 0.38 M Fe(III)
and pH 1. Copper extractions of approximately 100% and 80% were
achieved with and without added sodium chloride, respectively. In
that study, ferric ions were regenerated in a secondary bioreactor containing a chloride-tolerant culture of iron(II)-oxidising microorganisms.
In the tests with added sodium chloride, solutions were pH b1, compared with pH 1.2–1.3 in the absence of sodium chloride; the increased
acidity in the chloride-amended test would have contributed to enhanced copper extraction.
4.3. Chlorate as oxidant
Kariuki et al. (2009) studied chalcopyrite leaching in sealed vessels.
They mixed 2 g chalcopyrite concentrate with up to 3 g sodium chlorate
and then added 30 mL of 10 g L−1 H2SO4 (solution composition 0.9 M
NaClO3, 0.1 M H2SO4; ~ pH 0.7). The results showed that chalcopyrite
oxidation (reaction (14)) increased with increased temperature in the
range of 45–100 °C. Assuming a stoichiometric reaction, the end result
would be a solution containing 0.9 M Cl−, sufficient to contribute to enhanced copper extraction through the stabilisation of iron and copper
complex ions but at higher pH than the conditions of Carneiro and
Leão (2007).
6CuFeS2 þ 17NaClO3 þ 3H2 SO4 →3Fe2 ðSO4 Þ3 þ 6CuSO4 þ 17NaCl þ 3H2 O
103
the AuBr−
4 complex once all the copper has been leached (Muir and
Dixon, 2002).
0
4CuFeS2 þ 5O2 þ 20HCl→4CuCl2 þ 4FeCl3 þ 8S þ 10H2 O
0
ð16Þ
2CuFeS2 þ 5NaBrCl2 →2CuCl2 þ 2FeCl3 þ 4S þ 5NaBr
ð17Þ
4CuCl þ O2 þ 4HCl→4CuCl2 þ 2H2 O
ð18Þ
0
CuFeS2 þ 4CuCl2 →5CuCl þ FeCl3 þ 2S
ð19Þ
The technical advantages of the Intec process were marketed as
(i) not requiring an autoclave and the use of inexpensive materials of
construction such as fibre glass and polypropylene (temperature
below solution boiling point), (ii) high intensity electrowinning from
Cu(I) solution (not requiring solvent extraction) to yield high purity
granular copper (meeting LME A grade specification), (iii) ability to
treat low-grade and contaminated concentrates and (iv) lixiviant regeneration (Moyes and Houllis, 2002; P. Everett, 1996; P.K. Everett,
1996). However, according to Muir and Dixon (2002), while the Intec
process showed promise, several issues remained. Gold was encapsulated in sulfur or pyrite, requiring separate treatment for gold recovery,
sulfur and pyrite reported to the leached residue, presenting an environmental issue for residue-disposal, and the amalgamation of silver required strict control with respect to copper product contamination
and the environment. Operating costs were relatively low, but copper
extraction was also low (94%) (Lu and Dreisinger, 2013a).
ð14Þ
Sodium chlorate also enhanced the dissolution of a chalcopyrite
concentrate in hydrochloric acid solutions (Xian et al., 2012). At 45 °C,
copper extractions after 5 h in solutions 0.5 M and 1 M NaClO3 in 1 M
HCl were 45% and 65% respectively (Xian et al., 2012). The leaching
rate accelerated with increased hydrochloric acid up to 1.5 M. In addition to chlorate ion, the reaction products ferric chloride and gaseous
chlorine (detected by odour during leaching; reaction (15)) could
also have oxidised chalcopyrite but the stoichiometry of the reaction
products indicated that the gaseous chlorine was more likely to
have been volatilised. Surface area (grain size) was an important parameter in the study by Xian et al. (2012), consistent with published
data (e.g., Ikiz et al., 2006; Lottering et al., 2008; Naderi et al., 2011;
Watling et al., 2009).
NaClO3 þ 6HCl→3Cl2 þ NaCl þ 3H2 O
ð15Þ
5. Other oxidants
5.1. Cl2–Br2
The halide complex BrCl −
2 is described as having a number of
advantages when applied to mineral leaching. It is known to store
anodic energy in a soluble form, which should therefore not contaminate the metal product of electrolysis, and the resulting high oxidation potential of the anolyte promotes the leaching of specific metals
such as precious metals (P. Everett, 1996). In the INTEC process, the
lixiviant (0.47 M Cu(II), 4.9 M NaCl and 0.27 M NaBr) is generated
at the anode of the electrowinning cell. The complex anion BrCl −
2
(HalexTM) with Cu(II) and O2 are the oxidants during the leaching
of copper concentrate at atmospheric pressure, 80–85 °C and ORP
950–1000 mV vs Ag/AgCl (reactions (16)–(19)) (Intec Ltd., 2008).
The solution is maintained at pH 2–3 during the leach to avoid HCl
or Cl2 volatilisation and to promote goethite precipitation in the
first stage of leaching. In the third stage, the Eh is enhanced by the
BrCl−
2 complex ion, allowing free gold to dissolve and stabilise as
5.2. High MgCl2–low HCl–FeCl3
Harris et al. (2006a, 2006b, 2007, 2008); Harris and White, 2006,
2008) described a method for the leaching of sulfide minerals at atmospheric pressure in a temperature range of 75–115 °C using high chloride salt/low hydrochloric acid concentrations with an oxidant. The
solution, initially 2.9–4.3 M MgCl2 was acidified with HCl to pH ~1. Subsequently, the solution also contained ferric chloride after reaction with
the ore, and was controlled at solution pH ~1 and ORP 200–600 mV versus SHE. The leaching conditions were chosen to promote the formation
of hydrogen sulfide from the base metal sulfide (example reactions (20)
and (21) for chalcopyrite). The hydrogen sulfide was stripped from the
solution, reducing the amount of sulfate generated in the leach to very
low levels.
2CuFeS2 þ 6HCl þ Cl2 →4H2 S þ 2CuCl2 þ 2FeCl3
ð20Þ
4CuFeS2 þ 20HCl þ O2 →8H2 S þ 4CuCl2 þ 4FeCl3 þ 2H2 O
ð21Þ
This process was developed for the Ferguson Lake Project in Canada
(Harris et al., 2006a, 2006b, 2008), where the massive sulfide ore body
containing base and precious metals was not amenable to concentration
by flotation or other physical separation. In this process, the ore was
treated directly in the highly-concentrated chloride brine with its high
proton activity. Reagent recycle was one of the keys to the economics
of the process. Königsberger et al. (2008a) developed a Pitzer model
with which the thermodynamic properties of MgCl2–HCl aqueous solutions up to 350 g L–1 (10 M) chloride in the temperature range of
25–120 °C could be predicted. The authors noted that the model accurately predicted solid–liquid and liquid–vapour equilibria as well as linear trends indicating nearly ideal mixing behaviour in MgCl2–HCl
solutions. In a further study, Königsberger et al. (2008b) showed that
the solubilities of metal(II) chlorides, the solubility constants of iron(III)
oxides and hydroxides, and the redox potentials of Fe(III)/Fe(II)
depended on the MgCl2 concentration in a near-linear relationship. The
authors concluded that these results could be used to predict process parameters for improving yields or controlling iron and other impurities.
104
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
5.3. Gaseous Cl2–aqueous Cl2–HClO
Brocchi and Jena (1992) and Jena and Brocchi (1992) examined slurry chlorination of a chalcopyrite concentrate (~35 wt.% Cu, −75 μm particle size) at 30 °C. An aqueous slurry of the concentrate was prepared
and chlorine gas bubbled through it for the selected time period. The
pH of the slurry before and during chlorination was not reported. It
was found that the rapid reaction was dependent on solids loading.
With a 5% w/v solids loading, 100% copper extraction was possible in
5 min; with a 10% solids loading, the time required for 100% extraction
was 30 min; but with a 15% solids the maximum extraction was 85% Cu.
The possible process was described as “attractive” because of the high
yield, ease of operation, small capital investment, low energy consumption and relatively easy scale-up for commercialisation.
In chlorination tests at 23 °C and solutions of pH 3.6–4.4, conditions
under which hypochlorous acid was the dominant chlorine solution
species (Fig. 4) and cupric ions would remain in solution, Cho (1987)
found that up to 92% Cu was extracted in 1 h from a chalcopyrite
concentrate in a 0.228 M HClO solution. The reaction was pHindependent within the chosen pH range (reaction (22)). Çolak et al.
(1987) studied the dissolution kinetics of chalcopyrite in water saturated with chlorine and represented the overall reaction (23) as the net result of a series of intermediate reactions between Cu2S, CuS, FeS, S, FeCl2
or S2Cl2 with chlorine. Dissolution rates decreased with increased temperature in the range of 10–33 °C, the decrease being attributed to the
lower concentration of dissolved chlorine at higher temperature. Up to
95% Cu was extracted at 10 °C in 80 min. Solution pH values before and
after leaching were not reported but were presumably mildly acidic.
While the proposed overall reaction (23) does not contain any insoluble
compounds, Çolak et al. (1987) noted that the dissolution rate was controlled by diffusion through a product layer on the unreacted chalcopyrite
surface containing sulfur and an undefined ‘ash’ layer of insoluble gangue
minerals. The work by Jackson and Strickland (1958) was carried out
using dilute aqueous chlorine solutions at pH 1. The title of their report
is misleading in that the study was made using individual mineral specimens, not ore as suggested. The authors reported that a layer of elemental
sulfur or sulfur monochloride formed on chalcopyrite surfaces.
2CuFeS2 þ 11HClO→2Cu
2þ
2−
−
0
þ Fe2 O3 þ 2SO4 þ 2S þ 11Cl þ 11H
þ
2CuFeS2 þ 17Cl2 þ 16H2 O→2CuCl2 þ 2FeCl3 þ 4H2 SO4 þ 24HCl
ð22Þ
ð23Þ
In the study described by Ikiz et al. (2006), the active agent
was hypochlorous acid prepared by acidifying commercial grade hypochlorite solution with hydrochloric acid. The optimum pH for copper
extraction from an ore was pH 5 but the pH decreased during the dissolution, consistent with reaction (22). Using a 0.2 M HClO concentration
1
chlorine aq
0.6
hypochlorous
acid
mol
0.8
chloride
0.4
hypochlorite
chlorine vap
0.2
0
0
2
4
6
8
10
pH
Fig. 4. Chlorine speciation as a function of solution pH for a 0.2 M solution of sodium
hypochlorite acidified with hydrochloric acid (OLI Analyzer Studio Version 3.2).
and a 4 g L−1 solids loading, Ikiz et al. (2006) measured copper extractions of 40–80% in 1–15 min. They compared copper extraction from
−153 μm to −212 + 150 μm size fractions and showed that extraction
rates decreased with increased particle size. The results highlight the
importance of chalcopyrite liberation and available sulfide surface area
for larger particles.
Processes employing chlorine or hypochlorous acid to oxidise
chalcopyrite are not new. For example, Slater (1922) described a twostage process in which a copper ore was (i) subjected to a sulfuric acid
leach to remove copper oxides and, after separating the leached ore
from solution and (ii) the residue was reacted with a ‘relatively concentrated’ oxidizing agent, either sodium hypochlorite or calcium hypochlorite in slightly acidic condition. Slater (1922) presented case
studies on several ores which yielded high extractions. The inventor
failed to note the ore crush size, but as liberation was not reported
to be an issue, presumably the ores were milled or pulverised for
the test work. Kissock (1940) employed hypochlorous acid (“chlorine
monoxide”) for the dissolution of mineral sulfides and extracted nickel
and cobalt. The mechanism was described as the oxidation of the metal
sulfides through the release of nascent oxygen which attacked non-iron
containing compounds preferentially. More recently, Welham et al.
(2012) proposed the use of an aqueous solution of chlorine-based
oxidising species to oxidise ore or concentrate containing sulfide mineral phases. The leach process steps were (i) exposure of ore or concentrate to 10 mol% HClO (acidic) and (ii) allowing or facilitating the
oxidation of sulfide phases, in the process decreasing the pH so that
the predominant oxidising species was Cl2.
In the Sumitomo Process, based on the earlier MCLE Process for
leaching nickel matte, copper is leached from chalcopyrite concentrate
using chlorine gas at atmospheric pressure (Imamura et al., 2006). A
two-stage leach is employed. In the first stage, the introduction of chalcopyrite concentrate to a solution containing cupric and ferric chlorides
causes their reduction to cuprous and ferrous ions, respectively,
with chalcopyrite dissolution (reactions (2) and (3)). The almost 6 M
background concentration of chloride prevents the precipitation of
cuprous chloride (CuCl). In the second stage, conducted at 500–520 mV
vs Ag/AgCl, the partially leached chalcopyrite is reacted with chlorine
gas (reactions (24)–(26)) to effect the complete extraction of copper.
CuFeS2 þ 1:5Cl2ðgÞ →CuCl þ FeCl2 þ S2
2þ
2Fe
þ
ð25Þ
−
ð26Þ
þ 2Cl
2þ
þ 2Cl
þ Cl2ðgÞ →2Fe
2Cu þ Cl2ðgÞ →2Cu
ð24Þ
−
3þ
5.4. HClO4
Two studies were found in which the reactions of chalcopyrite with
perchloric acid were examined. The first was a mechanistic study
(Harmer et al., 2006). In that study, perchloric acid was chosen because,
at the concentrations used, it would not act as an oxidising agent, the
perchlorate anion would not complex other solution species, and its
use would facilitate the monitoring of soluble sulfur species arising
from chalcopyrite dissolution. An extensive study of the reaction products was undertaken and a reaction pathway, generally consistent
with that of Hiroyoshi et al. (2000) for ferrous-promoted chalcopyrite
dissolution, was proposed.
In the second study (Li et al., 2010) on chalcopyrite dissolution
rates at a temperature of 75 °C and pH 1 and 2 in different acids, including perchloric acid, the conditions employed were closer to those
applied in extraction studies. Nevertheless, this study also focused on
mechanisms, rather than copper extraction and, with specific exceptions, was most like an acid leaching (ferric ion-deficient) system,
that progressed according to reaction (4) rather than reaction (1). An
interesting result was that perchloric acid was the most effective copper
extractant at pH 1 but the least effective at pH 2. In contrast, the
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
hydrochloric acid-sodium chloride system was the most effective extractant at pH 2 but very poor at pH 1.
6. Acidic chloride systems for chalcopyrite ores in heaps or dumps
Heap leaching is a technology developed for the extraction of metals
from whole ores crushed to particle sizes appropriate for the required
extraction, acid consumption and heap permeability. It is a refinement
of the less-closely controlled in situ, in stope and run-of-mine dump
leaching options, and has been practised for many years as a means of
recovering metal values from ores that are not economic to concentrate
(Domic, 2007; Watling, 2006; and references therein).
There are a number of challenges associated with whole ore leaching
in heaps or other types of reactor, which must be addressed during process development. Those discussed in this review are restricted to differences that might be encountered when saline water is substituted
for freshwater in the leaching process, not including engineering issues
such as those described by Philippe et al. (2010). It should be noted that
Philippe et al. (2010) also concluded that the construction of a desalination plant could be an economic alternative to the delivery and use of
seawater at a mine site, and this conclusion was supported by a recent
report (Nueva Minería y Energía, 2013) in which it was indicated that
the use of desalinated water by the Chilean mining industry was likely
to increase eight-fold by 2022, with 16 ‘official projects’ in the pipeline.
6.1. Solution chemistry and secondary reaction products
The solution chemistry of seawater, its composition, complexes
(inorganic and organic) and ion associations, has been investigated for
many years and comprehensive data are available concerning the properties of seawater (e.g., Millero et al., 2008; Safarov et al., 2012; Turner
and Whitfield, 1987; Turner et al., 1981, and references therein). Not
surprisingly, acidification [from approximately pH 8.2 to pH 7.6] affects
copper, iron and other element speciation in seawater (Breitbarth et al.,
2009; Byrne, 2002; Byrne et al., 1988; Millero et al., 2009; Richards et al.,
2011). Acidification of seawater in mineral leaching is extreme in
comparison and has a major impact on solution speciation. However,
the predicted copper speciation in seawater at pH 6 is essentially the
same as that for pH 1–2, the range commonly employed in heap
leaching (Table 4, models of the inorganic components of a seawater
composition based on Turner et al. (1981), pH adjusted with H2SO4).
Simulated iron speciation evolves over a wider pH range, where the
predicted number of supersaturated solution species diminishes from
30 at pH 8 to eight at pH 6, two at pH 4 and zero species at pH 1–2. In
Table 4
Predicted copper and iron speciation in simulated seawater in the range pH 8 (natural) to
pH 1 (metal extraction).
Distribution (%) of total for element
Copper(II) monobromide ion (+1)
Copper(II) chloride
Copper(II) monochloride ion (+1)
Copper(II) dicarbonate ion (−2)
Copper(II) carbonate
Copper(II) ion (+2)
Copper(II) monohydroxide ion
(+1)
Iron(III) dichloride ion (+1)
Iron(III) monochloride (+2)
Iron(III) ion (+3)
Iron(III) dihydroxide ion (+1)
Iron(III) hydroxide
Iron(III) tetrahydroxide ion (−1)
Iron(III) monohydroxide ion (+2)
Iron(III) sulfate ion (+1)
CuBr+
CuCl2
CuCl+
Cu(CO3)2−
2
CuCO3
Cu2+
Cu(OH)+
a
FeCl+
2
FeCl2+
3+
Fe
Fe(OH)+
2
Fe(OH)3
Fe(OH)−
4
FeOH2+
+
FeSO4
pH 8
pH 6
b0.1
1
8
3
70
17
1
0.1
3
31
0.1
3
31
1
64
65
93
7
18
79
3
pH 4
pH 2
0.4
0.4
18
5
0.4
0.4
94
76
0.1
4
1
respect of acid consumption, the simulation predicted that 10% and 4%
less acid (H2SO4) would be required to achieve pH 1.5 or pH 2.0 set
points, respectively, in seawater in the absence of ore compared with
a freshwater system.
A number of coincident reactions take place when an ore is
contacted with a lixiviant, in this case acid solution, depending on ore
mineralogy. These include acid consumption as a result of the adsorption of protons on gangue mineral surfaces and congruent dissolution
(e.g., muscovite, reaction (27)) or incongruent dissolution of gangue
minerals (e.g., phlogopite alteration to vermiculite, reaction (28))
(Dreier, 1999; Jansen and Taylor, 2003). In combination, the reactions
result in increased element concentrations in solution and/or the
precipitation from solution of a potentially large number of phases
that may become supersaturated (Dreier, 1999; Jansen and Taylor,
2003).
þ
3þ
þ 3SiðOHÞ4
ð27Þ
þ
2KMg3 AlSi3 O10 ðOHÞ2 þ 10H →Mg2 Al2 Si3 O10 ðOHÞ2 þ 3SiðOHÞ4
2þ
þ
þ 4Mg þ 2K
ð28Þ
Two published studies were found in which copper extraction
in freshwater was compared with seawater and more concentrated
‘process solutions’. Torres et al. (2013) identified the main secondary
reaction product in the leached residues as gypsum with, in addition,
up to 8% of poorly-crystalline material typically comprising a mixture
of iron(III)- and Si-rich compounds (unpublished data). Watling et al.
(2014) showed that the amounts of poorly-crystalline iron(III)- and
silica-rich phases formed during leaching increased with increased salinity. In general, the dissolution of silicate minerals with the formation
of amorphous silica (SiO2) or silica gel (e.g., Halinen et al., 2009) or disordered aluminium silicates comprises an important but possibly detrimental suite of reactions. Secondary reaction products like these could
reduce the permeability of the heap bed, and increase the viscosity of
the solutions being recycled around the heap operation. A simulation
of muscovite dissolution in acidified freshwater, seawater and doublestrength seawater indicated that the number of supersaturated
compounds increased with increased salinity and with increased pH
(Fig. 5). At pH 1.5, it was predicted that silica and, in double strength
seawater, gypsum were supersaturated but that iron(III) compounds
were undersaturated; at pH 2, iron (III) oxides, iron(III) hydroxyoxide
and jarosite were supersaturated, together with alunite, albite, gypsum
and silica, depending on the salinity.
A reaction that would impact directly on copper extraction is the
precipitation of ferric ions as insoluble iron(III) compounds such as
ferrihydrite, schwertmannite, goethite, iron(III) hydroxide or one of
several jarosite compounds (reaction (12)), possibly leaving a limiting
8
pH 1
0.4
0.4
98
þ
KAl2 ðAlSi3 O10 ÞðOHÞ2 þ 10H →K þ 3Al
number of supersaturated
compounds
Speciesa
105
pH 1.5
7
pH 2
6
5
4
3
2
1
0
Fresh
1
Calculated using a seawater composition based on Turner et al. (1981) but omitting
organic components, acidified with H2SO4 (OLI Analyzer Studio 3.2 software).
Sea
2 X Sea
Fig. 5. The predicted number of supersaturated species in a simulation of muscovite dissolution (representing low-grade ore) in acidified freshwater, seawater and double-strength
seawater at pH 1.5 (representing heap feed solution) and pH 2 (representing heap
discharge solution) (OLI Analyzer Studio version 3.2).
106
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
concentration of the ferric ion oxidant for sulfide dissolution. A greater
extent of precipitation of iron(III) compounds from seawater leachates,
together with their known ability to sequester trace metals (Burgos
et al., 2012; Scott, 1987) may cause copper hold-up in the heap, in addition to loss of permeability in the heap bed. An increased proportion
of jarosite compounds precipitated, relative to iron (III) oxides and
hydroxides, as a consequence of the high concentration of monovalent
cations in seawater (this laboratory, unpublished data), is consistent
with the predicted supersaturated compounds and with the results of
Carneiro and Leão (2007).
6.2. Bioleaching in chloride heaps
Historically, it has been believed that the presence of chloride at the
concentrations found in seawater (approximately 0.5 M NaCl) would
preclude the bioleaching of sulfide ores, based on published laboratory
studies about the limited salt tolerance of iron- and sulfur-oxidising
bacterial cultures of species known to colonise sulfide heaps (Zammit
et al., 2012 and references therein). Known halophiles, such as
Halothiobacillus spp. (Kelly and Wood, 2000), were not acid-tolerant
and did not oxidise iron(II), a key reaction in the dissolution of sulfide
minerals. A strategy commonly used to overcome the inhibitory chloride effect was to dilute saline water with freshwater. For example,
for the column bioleaching tests that formed part of the development
of heap leaching at Ivan Mine, Chile, underground saline water was
diluted to approximately 0.14 M chloride, so as to permit bacterial
iron- and sulfur-oxidation (Leong et al., 1993).
A second strategy was to adapt microbial cultures to the presence
of chloride. Weston et al. (1995) adapted bacterial cultures to water
containing ~0.08 M chloride, because that was the anticipated chloride
content in the water supply for Zaldivar Mine, Chile. At Pacific Ore Technologies, sulfur-oxidising microbial cultures were developed and
adapted to different levels of salinity up to eight-times that of seawater
for application in a process that was not reliant on iron(II) oxidation
(Williams, 2006); one of those cultures was applied to the bioleaching
of ore from Sherlock Bay (Western Australia) as a process option for
the nickel–copper sulfide deposit near which the underground water
supply had a salinity similar to that of seawater.
In the mean time, the application of molecular microbiological
techniques to mining and acid mine drainage systems is revealing an
increasing number of acidophiles that are halotolerant and possess
the ability to oxidise iron(II) and/or reduced inorganic sulfur compounds (Table 5). For example, a mixed culture, Ni–S-J069B, has been
developed, capable of functioning at pH 3.5 and 50–60 °C in saline
solution of up to 1.4 M chloride, and applied to the leaching of a nickel
sulfide ore, arsenic-rich sulfide tailings and a chalcopyrite concentrate
(McCreddan and Seet, 2013). In respect of chalcopyrite concentrate,
the results reported were unimpressive (91% copper extraction
in 118 days), but the low ORP values recorded during leaching indicated
that this mixed culture did not oxidise iron(II) efficiently. The main interest lies in the putative identification of the species in the culture (Table 5)
because that shows that some well-known genera of bioleaching microorganisms can function in high-chloride environments.
It must be concluded that, far from being inhospitable environments
to microorganisms, chloride heaps host diverse and active iron(II)- and
sulfur-oxidising microbial populations that will be isolated and described in the future. Their expected importance to the mining industry
is reflected by their inclusion in recent patent applications related to
mineral leaching (e.g., Davis-Belmar et al., 2010; Dew and du Plessis,
2002; Ohtsuka and Mitarai, 2007; Rautenbach et al., 2011).
6.3. Use of saline water or seawater in heaps or dumps
Dutrizac and MacDonald (1971) investigated copper extraction
from chalcopyrite ore (3% Cu) under simulated dump leaching conditions. For column tests, the ore was crushed to − 0.63 mm and 3 kg
portions loaded into 5 mm internal diameter columns to give a 1 m
bed depth. When the ore was leached with solution containing 0.1 M
Fe(III) and 0.1 M H2SO4 with/without approximately 0.1 M NaCl,
copper extractions ranged from 1 to 8% in tests conducted in the
range of 25–40 °C but, more importantly for that evaluation, the
presence of sodium chloride suppressed copper extraction. However
Table 5
Halotolerant acidophiles capable of oxidising iron(II) and/or reduced inorganic sulfur compounds.
Microorganism/culture
NaCl (g L−1)
range [optimum]
pH
range [optimum]
Fe(II) OX
S OX
Reference
‘Leptospirillum ferriphilum’-like strain YSK
0–7
[0]
5–30
0–30
[10]
0–60
[0]
25–50
[30]
10–30
[20]
~8–47
[0]
60
5–50
[5]
0–30
[10]
10–45
[20]
–80
1.6
✓
×
Wang et al. (2012)
1–3
[2]
✓a
✓
×
✓
Rautenbach et al. (2011)
Qi et al. (2009); Wang et al. (2012)
1–4.5
[2]
1.7
✓
✓
Huber and Stetter (1989)
✓
✓c
Norris et al. (2010)
2–3
[2]
1.8
✓
✓
Kamimura et al. (2001)
✓
✓
Nicolle et al. (2009); Kelly et al. (2005)
2
1.6–4.5
✓
✓
✓
✓
Norris and Simmons (2004); Kelly et al. (2005)
Holden et al. (2001); Rodgers et al. (2002)
1–4
×
✓
Ohtsuka and Mitarai (2007)
1–5
[4]
3.5
×
✓
Kamimura et al. (2003, 2005)
×
✓
McCreddan and Seet (2013)
Leptospirillum sp.-Cl
Sulfobacillus sp. TPY
Thiobacillus prosperusb strains
Thiobacillus prosperus ‘Milos culture’
‘Thiobacillus’-like strain KU2-11
‘Acidihalobacter aerolicus’ strain V6
Acidihalobacter ferrooxidans (V8)
‘Alicyclobacillus’-like strains
‘Acidithiobacillus’-like strain TTH19A
‘Acidithiobacillus thiooxidans’-like strain SH
Ni–S-J069B mixed cultured
a
Growth with halophilic sulfur oxidising species.
b
Closest relatives now known to be Acidihalobacter spp. with which it is likely to be reclassified as Acidihalobacter prosperus.
c
Tetrathionate 1 mM and CO2 enriched air stimulated iron(II) oxidation.
d
Putatively identified genera were Acidimicrobium, Acidiphilium, Acidithiobacillus, Acidobacterium, Acidocella, Acidisphaera, Alicyclobacillus, Sulfobacillus and a dominant but unidentified
microorganism.
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
at 60 °C, copper extractions were approximately 17% and 10% after
55 days leaching with/without added chloride, respectively. Having
confirmed the greater than 50 °C temperature requirement for enhanced extraction using a second ore, together with ancillary laboratory
tests, Dutrizac and MacDonald (1971) concluded that large chloride additions would not benefit copper extraction in a dump-leach, as temperatures were generally below 50 °C. In contrast, Muñoz-Ribadeneira and
Gomberg (1970, 1971) reported that the addition of as little as 0.1 M
HCl to a sulfuric acid leach of chalcopyrite resulted in increased copper
extraction.
The earliest reference to seawater use in heap leaching found so far
is one by Domic (2007) in which it was noted that in 1991, the first
commercial leaching–solvent extraction–electrowinning plant using
Pudahuel's thin layer leaching technology and license using seawater
for leaching started at Lince (now Mina Michilla). With more than
20 years experience, it can be assumed that seawater is as good a
solvent and carrier as freshwater, of acid, ferric ions and possibly salttolerant microorganisms to the ore and the soluble metals to the next
stage of a heap leach process.
The chloride heap leaching process invented by Muller et al. (2011),
underpinned by the body of work described by Nicol and co-researchers
(Miki and Nicol, 2011; Nicol et al., 2010; Velásquez-Yévenes et al.,
2010a, 2010b), was not per se a seawater heap-leach process, although
the range of chloride (Cl−) concentrations proposed for the method's
greatest efficiency did encompass that of seawater. The strategy
was to control ORP in an acidic mixed sulfate–chloride solution to maintain the chalcopyrite surface potential in the range of 550–600 mV versus SHE. The chloride concentration in the proposed process was
between 5 and 100 g L− 1 (0.14–2.8 M Cl−) and the presence of at
least 1 mg L−1 (ppm) oxygen was required. Muller et al. (2011) proposed that, in the presence of chloride in the suggested concentration
range, (i) the type, morphology and distribution of elemental sulfur
was such that it formed away from a chalcopyrite surface, (ii) Cu(I) species were stabilised allowing the Cu(II)/Cu(I) couple to control ORP,
(iii) the thermodynamics and possibly the rate of the non-oxidative reaction was enhanced (reaction (29) or (4)) and (iv) the formal potential
of the Cu(II)/Cu(I) couple was increased. They also proposed that chloride affected (v) the rate of oxidation of Cu(I) to Cu(II) and the dissolved
O2 concentration, and (vi) resulted in a reduction in the required acid to
achieve the desired pH. The inventors noted that iron had no direct role
in the leaching mechanism, provision of acid drove the non-oxidative
reaction and that ORP determined the mixed potential at the mineral
surface and controlled the mechanism of chalcopyrite dissolution. It
was further suggested that a number of insoluble materials such as
pyrite, magnetite, hematite, activated carbon of coal, zeolites and several
elements (Ag, Bi, Cd or Hg) might enhance the kinetics of H2S oxidation
(see reaction (30)), thus increasing the rate of chalcopyrite dissolution.
þ
2−
CuFeS2 þ 4H þ 2SO4 →CuSO4 þ FeSO4 þ 2H2 SðaqÞ
0
H2 SðaqÞ þ O2 →2S þ 2H2 OðcopperðIIÞ catalysed reactionÞ
ð29Þ
ð30Þ
The CuproChlor process (Aroca et al., 2012; Espejo et al., 2001; Rauld
Faine et al., 2005) was developed to extract copper from copper oxide
and mixed copper oxide–secondary sulfide ores containing chalcocite,
bornite and covellite (Herreros et al., 2006). It is a chloride heap leach
with two main differences; first, fine copper-rich ore or concentrate
particles were bound by salt bridges (gypsum) to host rocks during ‘agglomeration’ in a process akin to the Geocoat technology (Kohr et al.,
2004) and second, the chloride concentration in solution was more
than three times that of seawater. It was proposed that the chloride
(i) stabilised Cu(I) formed during leaching with the result that the
Cu(II)/Cu(I) couple could contribute to sulfide oxidation, (ii) enhanced
sulfide oxidation by ferric ions and (iii) influenced the morphology of
elemental sulfur, facilitating the movement of ferric and ferrous ions
to and from chalcopyrite surfaces. At the Michilla mine, Chile, 65%
107
of the contained copper was extracted in 50 days of leaching and
90% in 110 days; the CuproChlor process was considered to be ‘competitive with bioleaching’ (Herreros et al., 2006). A modification of the
CuproChlor process incorporated the calcium chloride agglomeration
strategy to facilitate the leaching of copper sulfide concentrates (including chalcopyrite) in heaps (Rauld Faine et al., 2005), by which it was
differentiated from other processes for the leaching of concentrates in
heaps, such as the GEOCOAT process described in numerous patents
(e.g., Kohr et al., 2004).
7. Summary
Hydrometallurgical process developments for the extraction of
copper from chalcopyrite, an abundant but refractory mineral with
respect to its dissolution, tend to be targeted at complex or dirty concentrates that would incur penalties were they to be smelted or lowgrade ores that are a major but, thus far, uneconomic source of copper.
The copper values in concentrates permit a degree of flexibility in process development but the lower values in low-grade ores demand that
new technologies be low cost. Perceived advantages of chloride systems, such as the high solubility of copper and ferric iron, the ease of ferrous ion oxidation, faster leaching kinetics of chalcopyrite compared
with ferric sulfate systems, and generation of sulfur rather than sulfate
as the primary product of sulfide oxidation, have driven research and
development over many years. Perceived disadvantages were the corrosive action of chloride, thus necessitating the use of more expensive
materials of construction for reactors, the need for fine grinding for processes operated at atmospheric pressure, and the difficulty of electrowinning high-grade copper from chloride solutions.
Process developments for concentrates employ acidic, oxidising
leach media containing sodium or other chloride salts up to concentrations encountered in brines, at temperatures up to the boiling points of
the selected solution compositions. In these processes, chloride ion is
proposed to be an active agent in the dissolution, which may occur by
different mechanisms depending upon the solution composition. Innovative reactors and flow sheets, the need for process controls and additional reagent costs, are offset by the copper values in concentrates
which, for a number of reasons including undesirable impurities, may
make them unsuitable for smelting. Reagent recovery and recycle is incorporated into some flowsheets, for example the High Concentration
Chloride Leach, the Intec Process and Hydrocopper Process, contributing
to their potential economic and environmental sustainability.
For low grade ores, usually processed in heaps, the use of seawater
or other naturally saline water in leaching operations is an ‘economic’
choice to overcome the scarcity and/or cost of freshwater. Few studies
have been published describing the advantages and disadvantages of
seawater substitution for freshwater in leaching processes but, from
the sparse information available, seawater appears to be as efficient a
solvent and carrier of acid and oxidant as freshwater. In seawater solutions the chloride concentration is too low to stabilise cuprous ions and
cupric ions are the predominant copper species in solution. There is a
view that seawater use for sulfide ore heaps would prohibit the beneficial catalytic action of acidophilic microorganisms. However, the recent
description of some iron(II)- and sulfur-oxidising, salt-tolerant acidophiles suggests that there is a diverse group of ‘still-to-be-described’
microorganisms that could function in sulfide heaps irrigated with
seawater. With regard to processing, the salt content in seawater
would impact directly on solution transport costs to and around a
mine (through increased solution viscosity and specific gravity) and
could impact on product and by-product contamination (requiring
clean water and additional unit process).
Common to high- and low-chloride processes, accessibility of
chalcopyrite grains to the lixiviant and chalcopyrite grain size are key
parameters in the success of a technology and increased temperature
and stronger oxidants than cupric and ferric ions can be employed to increase reaction kinetics.
108
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
Acknowledgements
Dr. R. McDonald is thanked for a careful and critical review of the
typescript. The financial support of the Australian Government through
the CSIRO Minerals Down Under Flagship is gratefully acknowledged.
References
Al-Harahsheh, M., Kingman, S.W., Al-Harahsheh, A., 2008. Ferric chloride leaching of
chalcopyrite: synergetic effect of CuCl2. Hydrometallurgy 91, 89–97.
Allen, E.S., Kruesi, P.R., Lake, J.L., 1973. CYMET process—hydrometallurgical conversion of
base metal sulphides to pure metals. CIM Bull. 66 (734), 81–87.
AMIRA International, 2004. Copper Technology Roadmap. AMIRA International,
Melbourne.
Andersen, E., Boe, G.H., 1985. Hydrometallurgical method of extraction of copper from
sulphide-containing material. U.S. Patent 4552632 (12 November 1985).
Aroca, F., 1999. Plant description and operation of heap leaching, solvent extraction and
electrowinning of copper at Minera Michilla. In: Jergensen, G.V. (Ed.), Copper
Leaching, Solvent Extraction and Electrowinning Technology. SME, Englewood, CO,
pp. 279–294.
Aroca, F., Backit, A., Jacob, J., 2012. CuproChlor®, a hydrometallurgical technology for
sulphides leaching. In: Casas, J.C., Ciminelli, V.S.T., Montes-Atenas, G., Stubina, N.
(Eds.), HydroProcess 2012. Gecamin, Santiago, pp. 97–108.
Atwood, G.E., Curtis, C.H., 1974. Hydrometallurgical process for the production of copper.
U.S. Patent 3,785,944 (15 January 1974).
Atwood, G.E., Curtis, C.H., 1975. Hydrometallurgical process for the production of copper.
U.S. Patent 3,879,272 (22 April 1975).
Ayres, R.U., Ayres, L.W., Rade, I., 2002. The life cycle of copper, its co-products and byproducts. Report commissioned by the Mining, Minerals and Sustainable Development project of the International Institute for Environment and Development Source
(http://pubs.iied.org/pdfs/G00740.pdf).
Bonan, M., Demarthe, J.M., Renon, H., Baratin, F., 1981. Chalcopyrite leaching by CuCl2 in
strong NaCl solutions. Metall. Trans. B 12, 269–274.
Breitbarth, E., Bellerby, R.J., Neill, C.C., Ardelan, M., Meyerhöfer, M., Zöllner, E., Croot, P.L.,
Riebesell, U., 2009. Ocean acidification affects iron speciation in seawater. Biogeosci.
Discuss. 6, 6781–6802 (Source: http://biogeosciences-discuss.net/6/6781/2009/bgd6-6781-2009.pdf).
British Geological Survey, 2007. Copper. BGS Minerals UK Centre for Sustainable Mineral
Development, Nottingham Source: http://www.bgs.ac.uk/mineralsuk/statistics/
mineralProfiles.html.
Brocchi, E.A., Jena, P.K., 1992. Studies on alternate routes for extracting copper from chalcopyrite concentrates. In: Misra, V.N., Halbe, D., Spottiswood, D.J. (Eds.), Proceedings
of the International Conference on the Extractive Metallurgy of Gold and Base Metals.
AusIMM, Melbourne, pp. 281–285.
Burgos, W.D., Borch, T., Troyer, L.D., Luan, F., Larson, L.N., Brown, J.F., Lambson, J., Shimizu,
M., 2012. Schwertmannite and Fe oxides formed by biological low-pH Fe(II) oxidation versus abiotic neutralization: impact on trace metal sequestration. Geochim.
Cosmochim. Acta 76, 29–44.
Byrne, R.H., 2002. Inorganic speciation of dissolved elements in seawater: the influence of
pH on concentration ratios. Geochem. Trans. 3, 11–16.
Byrne, R.H., Kump, L.H., Cantrell, K.J., 1988. The influence of temperature and pH on trace
metal speciation in seawater. Mar. Chem. 25, 163–181.
Cai, Y., Chen, X., Ding, J., Zhou, D., 2012. Leaching mechanism for chalcopyrite in hydrochloric acid. Hydrometallurgy 113–114, 109–118.
Carneiro, M.F.C., Leão, V.A., 2007. The role of sodium chloride on surface properties of
chalcopyrite leached with ferric sulphate. Hydrometallurgy 87, 73–82.
Chadwick, J., 2009. Copper Recovery. International Mining pp. 23–30 (August).
Cho, E.H., 1987. Leaching studies of chalcopyrite and sphalerite with hypochlorous acid.
Metall. Trans. B 18, 315–323.
Çolak, S., Alkan, M., Kocakerim, M.M., 1987. Dissolution kinetics of chalcopyrite containing
pyrite in water saturated with chlorine. Hydrometallurgy 18, 183–193.
Dalton, R.F., Price, R., Hermana, E., Hoffmann, B., 1987. The CUPREX process—a new
chloride-based hydrometallurgical process for the recovery of copper from sulphidic
ores. In: Davies, G.A. (Ed.), Separation Processes in Hydrometallurgy. Ellis Horwood,
Chichester, pp. 466–476.
Dalton, R.F., Diaz, G., Hermana, E., Price, R., Zunkel, A.D., 1991. The Cuprex metal extraction process: recovering copper from sulfide ores. J. Met. 43 (8), 51–56.
Davis-Belmar, C.S., Demergasso, C.S., Rautenbach, G.F., 2010. Treatment of sulfide mineral
or mixed sulfide and oxide mineral involves bioleaching mineral in chloride ion solution with culture consortium Leptospirillum ferriphilum strain and sulfur oxidizing
halophilic or halotolerant microorganism. World Patent Application WO 2010/
009481-A2 (21 January 2010).
Demarthe, J.M., Georgeaux, A., 1978. Hydrometallurgical treatment of complex sulfides.
In: Jones, M.J. (Ed.), Complex Metallurgy '78. The Institution of Mining and Metallurgy,
London, pp. 113–120.
Demarthe, J.M., Gandon, L., Georgeaux, A., 1976. A new hydrometallurgical process
for copper. In: Yannopoulos, J.C., Agarwal, J.C. (Eds.), Extractive Metallurgy of
Copper. Hydrometallurgy and electrowinning, 2. TMS of AIME, New York, pp.
825–847.
Demarthe, J.M., Sonntag, A., Georgeaux, A., 1977. Method for obtaining copper from
copper-bearing ores. US Patent 4,023,964 (17 May 1977).
Dew, D.W., du Plessis, C.A., 2002. Bioleaching of a sulphide concentrate in saline solution.
World Patent Application WO2002/081761-A2 (17 October 2002).
Domic, E.M., 2007. A review of the development and current status of copper bioleaching
operations in Chile: 25 years of successful commercial implementation. In: Rawlings,
D.E., Johnson, D.B. (Eds.), Biomining. Springer, Berlin, pp. 81–95.
Dreier, J., 1999. The Chemistry of Copper Heap Leaching. Source: http://jedreiergeo.com/
copper/article1/Chemistry_of_Copper_Leaching.html.
Dreisinger, D., 2006. Copper leaching from primary sulfides: options for biological and
chemical extraction of copper. Hydrometallurgy 83, 10–20.
Dreisinger, D., 2009. Hydrometallurgical process development for complex ores and concentrates. J. South. Afr. Inst. Min. Metall. 109, 253–271.
Dreisinger, D.B., Murray, W., Baxter, K., Holmes, M., Jacobs, H., Molnar, R., 2005. The
metallurgical development of the El Boleo copper–cobalt–zinc project. ALTA Copper
2005ALTA Metallurgical Services, Melbourne.
Dutrizac, J.E., 1981. The dissolution of chalcopyrite in ferric sulfate and ferric chloride
media. Metall. Trans. B 12, 371–378.
Dutrizac, J.E., 1982. Ferric ion leaching of chalcopyrites from different localities. Metall.
Trans. B 13, 303–309.
Dutrizac, J.E., 1990. Elemental sulfur formation during the ferric chloride leaching of
chalcopyrite. Hydrometallurgy 23, 153–176.
Dutrizac, J.E., 1992. The leaching of sulphide minerals in chloride media. Hydrometallurgy
29, 1–45.
Dutrizac, J.E., MacDonald, R.J.C., 1971. The effect of sodium chloride on the dissolution of chalcopyrite under simulated dump leaching conditions. Metall. Trans. 2,
2310–2312.
Espejo, R. Y Arriagada F., D'amico, J., Jo, M., Rojas, J., Neuberg, H., Ruiz, M., Yanez, H., Reyes,
R., Bustos, S., Montealegre, R., Rauld, J., 2001. Procedimiento para aglomerar
minerales de cobre chancados fino, mediante la adicion de cloruro de calcio y acido
sulfurico y procedimiento de lixiviacion previa aglomeracion del mineral. Chilean
Patent Application No. 40891 (19 May 1997); granted 2001.
Everett, P.K., 1996a. Production of metals from minerals. US Patent 5,487,819 (30 January
1996).
Everett, P., 1996b. The Intec copper process. Low cost treatment of chalcopyrite. ALTA
Copper Hydrometallurgy ForumALTA Metallurgical Services, Melbourne.
Guy, S., Broadbent, C.P., 1983. Formation of copper(I) sulphate during cupric chloride
leaching of a complex Cu/Zn/Pb ore. Hydrometallurgy 11, 277–288.
Guy, S., Broadbent, C.P., Lawson, G.J., 1983. Cupric chloride leaching of a complex copper/
zinc/lead ore. Hydrometallurgy 10, 243–255.
Halinen, A.-K., Rahunen, N., Kaksonen, H., Puhakka, J.A., 2009. Heap bioleaching of a
complex sulfide ore Part 1: effect of pH on metal extraction and microbial composition in pH controlled columns. Hydrometallurgy 98, 92–100.
Hämäläinen, M., 2005. Method for leaching copper concentrate. U.S. Patent 6,929,677 B2
(16 August 2005).
Harmer, S.L., Thomas, J.E., Fornasiero, D., Gerson, A.R., 2006. The evolution of surface
layers formed during chalcopyrite leaching. Geochim. Cosmochim. Acta 70,
4392–4402.
Harris, B., White, C., 2006. Process for recovering iron as hematite from a base metal
containing ore material. PCT International Patent Application from U.S. Provisional
Application 60/752877 (13 December 2006).
Harris, B., White, C., 2008. Recent developments in the high-strength chloride
leaching of base metal sulphide ores. ALTA Nickel/CobaltALTA Metallurgical
Services, Melbourne.
Harris, B., White, C., Ballantyne, B., 2006a. Chloride leaching of polymetallic sulphide ore.
ALTA CopperALTA Metallurgical Services, Melbourne.
Harris, B., White, C., Davis, B., Roy, A., 2006b. Process for recovering a base metal from
a sulfide ore material. PCT International Patent Application from U.S. Provisional
Application 60/752900 (13 December 2006).
Harris, G.B., Lakshmanan, V.I., Sridhar, R., Puvvada, G., 2007. A process for the
recovery of value metals from base metal sulfide ores. Canadian Patent 2478516
(12 November 2007).
Harris, G.B., White, C.W., Demopoulos, G.P., Ballantyne, B., 2008. Recovery of copper from
a massive polymetallic sulphide by high concentration chloride leaching. Can. Metall.
Q. 47, 347–356.
Haver, F.P., Wong, M.M., 1971. Recovering elemental sulfur from non-ferrous metals.
Ferric chloride leaching of chalcopyrite concentrate. Report of Investigations 7474.
US Bureau of Mines.
Haver, F.P., Baker, R.D., Wong, M.M., 1975. Improvements in ferric chloride leaching of
chalcopyrite. Report of Investigations 8007. US Bureau of Mines.
Havlik, T., Kammel, R., 1995. Leaching of chalcopyrite with acidified ferric chloride and
carbon tetrachloride addition. Miner. Eng. 8, 1125–1134.
Hein, H., Joly, P., 2011. Metallurgical performance and characteristics of small and medium size copper SX plants in Chile. In: Velenzuela, L.F., Moyer, B.A. (Eds.), Proceedings
19th International Solvent Extraction Conference. Gecamin, Santiago (8 pp.).
Herreros, O., Quiroz, R., Longueira, H., Fuentes, G., Viñals, J., 2006. Leaching of djurleite in
Cu2+/Cl− media. Hydrometallurgy 82, 32–39.
Hirato, T., Kinoshita, M., Awakura, Y., Majima, H., 1986. The leaching of chalcopyrite with
ferric chloride. Metall. Trans. B 17, 19–28.
Hiroyoshi, N., Miki, H., Hirajima, T., Tsunekawa, M., 2000. A model for ferrous-promoted
chalcopyrite leaching. Hydrometallurgy 57, 31–38.
Hoffmann, J.E., 1991. Winning copper via chloride chemistry—an elusive technology. J.
Met. 43 (8), 48–49.
Holden, P.J., Foster, L.J., Neilan, B.A., Berra, G., Vu, Q.M., 2001. Characterisation of novel salt tolerant iron-oxidising bacteria. In: Ciminelli, V.S.T., Garcia, O. (Eds.), Biohydrometallurgy:
Fundamentals, Technology and Sustainable Development, Part A. Elsevier, Amsterdam,
pp. 283–290.
Huber, H., Stetter, K.O., 1989. Thiobacillus prosperus sp. nov., represents a new group of
halotolerant metal-mobiizing bacteria isolated from a marine geothermal field.
Arch. Microbiol. 151, 479–485.
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
Hyvärinen, O., Hämäläinen, M., 1999. Method for producing copper in hydrometallurgical
process. U.S. Patent 6,007,600 (28 December 1999).
Hyvärinen, O., Hämäläinen, M., 2005. HydrocopperTM—a new technology producing
copper directly from concentrate. Hydrometallurgy 77, 61–65.
Ikiz, D., Gulfen, M., Aydin, A.O., 2006. Dissolution of primary chalcopyrite ore in hypochlorite
solution. Miner. Eng. 19, 972–974.
Imamura, M., Takeda, K., Ando, K., Nagase, N., Ojima, Y., 2006. Iron recovery in the
Sumitomo chlorine leach process for copper concentrates. In: Dutrizac, J.E., Riveros,
P.A. (Eds.), Iron Control Technologies, Proceedings of the Third International Symposium on Iron Control in Hydrometallurgy (Montreal). CIM, Montreal, pp. 191–204.
Intec Ltd., 2008. The Intec Copper Process. Source: http://intec.com.au/technology/copper.
Jackson, K.J., Strickland, J.D.H., 1958. The dissolution of sulfide ores in acid chlorine
solutions; a study of the more common sulfide minerals. Trans. Am. Inst. Min. Metall.
Pet. Eng. 212, 373–379.
Jansen, M., Taylor, A., 2003. Overview of gangue mineralogy issues in oxide copper heap
leaching. ALTA Copper conference 2003. ALTA Metallurgical Services, Melbourne.
Jena, P.K., Brocchi, E.A., 1992. Extraction of copper from dry Salobo (Brazil) chalcopyrite
concentrate and aqueous slurries thereof by chlorination with Cl2. Trans. Inst. Min.
Metall. C 101, 48–51.
Kamimura, K., Kunomura, K., Wakai, S., Murakami, K., Sugio, T., 2001. Some properties of
a novel obligately autotrophic iron-oxidizing bacterium isolated from seawater.
Hydrometallurgy 59, 373–381.
Kamimura, K., Higashino, E., Moriya, S., Sugio, T., 2003. Marine acidophilic sulfur-oxidizing
bacterium requiring salts for the oxidation of reduced inorganic sulfur compounds.
Extremophiles 7, 95–99.
Kamimura, K., Higashino, E., Kanao, T., Sugio, T., 2005. Effects of inhibitors and NaCl on the
oxidation of reduced inorganic sulfur compounds by a marine acidophilic, sulfuroxidizing bacterium, Acidithiobacillus thiooxidans strain SH. Extremophiles 9, 45–51.
Kariuki, S., Moore, C., McDonald, A.M., 2009. Chlorate-based oxidative hydrometallurgical
extraction of copper and zinc from copper concentrate sulfide ores using mild acidic
conditions. Hydrometallurgy 96, 72–76.
Kelly, D.P., Wood, A.P., 2000. Reclassification of some species of Thiobacillus to the newly
designated genera Acidithiobacillus gen. nov., Halothiobacillus gen. nov. and
Thermithiobacillus gen. nov. Int. J. Syst. Evol. Microbiol. 50, 511–516.
Kelly, D.P., Wood, A.P., Stackebrandt, E., 2005. Genus II. Thiobacillus Beijerinck 1904b,
597AL, Bergey′s Manual® of Systematic Bacteriology2nd ed. Springer, N.Y. pp. 764–769.
Kinnunen, P.H.-M., Puhakka, J.A., 2004. Chloride-promoted leaching of chalcopyrite
concentrate by biologically-produced ferric sulfate. J. Chem. Technol. Biotechnol. 79,
830–834.
Kissock, A., 1940. Extraction of metals. U.S. Patent 2205565 (25 June 1940).
Kohr, W.J., Shrader, V., Johansson, C., 2004. High temperature heap bioleaching process.
U.S. Patent 6,802,888 B2 (12 October 2004).
Königsberger, E., May, P., Harris, B., 2008a. Properties of electrolyte solutions relevant to
high concentration chloride leaching. I. Mixed aqueous solutions of hydrochloric
acid and magnesium chloride. Hydrometallurgy 90, 177–191.
Königsberger, E., Königsberger, L.-C., May, P., Harris, B., 2008b. Properties of electrolyte solutions relevant to high concentration chloride leaching. III. Solubility of pertinent
solids and iron(III)/iron(II) redox potential measured in concentrated magnesium
chloride solutions. Hydrometallurgy 90, 192–200.
Kruesi, P.R., 1972. Process for the recovery of metals from sulfide ores through electrolytic
dissociation of the sulfides. U.S. Patent 3673061 (27 June 1972).
Leong, B.J.Y., Dreisinger, D.B., Lo, M., Branion, R., Hackl, R.P., Gormely, L.S., Crombie, D.R.,
1993. The microbiological leaching of a sulphidic copper ore in a strongly saline
medium (I) shake flask and column studies. In: Torma, A.E., Wey, J.E., Lakshmanan,
V.I. (Eds.), Biohydrometallurgical Technologies. TMS, Warrendale, pp. 117–126.
Li, J., Kawashima, N., Kaplun, K., Absalon, V.J., Gerson, A.R., 2010. Chalcopyrite leaching:
the rate controlling factors. Geochim. Cosmochim. Acta 74, 2881–2893.
Liddicoat, J., Dreisinger, D., 2007. Chloride leaching of chalcopyrite. Hydrometallurgy 89,
323–331.
Lin, H.K., Wu, X.J., Rao, P.D., 1991. The electrowinning of copper from a cupric chloride
solution. J. Met. 43 (8), 60–65.
Lottering, M.J., Lorenzen, L., Phala, N.S., Smit, J.T., Schalkwyk, G.A.C., 2008. Mineralogy and
uranium leaching response of low-grade South African ores. Miner. Eng. 21, 16–22.
Lu, J., Dreisinger, D., 2013a. Copper leaching from chalcopyrite concentrate in Cu(II)/
Fe(III) chloride system. Miner. Eng. 45, 185–190.
Lu, J., Dreisinger, D., 2013b. Solvent extraction of copper from chloride solution I: Extraction isotherms. Hydrometallurgy 137, 13–17.
Lu, Z.Y., Jeffrey, M.I., Lawson, F., 2000. The effect of chloride on the dissolution of chalcopyrite in acidic solutions. Hydrometallurgy 56, 189–202.
Lundström, M., Aromaa, J., Forsén, O., Hyvärinen, O., Barker, M.H., 2005. Leaching of
chalcopyrite in cupric chloride solution. Hydrometallurgy 77, 89–95.
Lundström, M., Liipo, J., Karonen, J., Aromaa, J., 2009. Dissolution of six sulfide concentrates in the Hydrocopper environment. The Southern African Institute of Mining
and Metallurgy Base Metals Conference 2009. SAIMM, Marshalltown, South Africa,
pp. 127–138.
Majima, H., Awakura, Y., Hirato, T., Tanaka, T., 1985. The leaching of chalcopyrite in ferric
chloride and ferric sulphate solutions. Can. Metall. Q. 24, 283–291.
Makino, S., Sugimoto, M., Yano, F., Matsumoto, N., 1996. Operation of the MCLE (matte
chlorine leach electrowinning) plant for nickel refining at Sumitomo Metal Mining
Co., Ltd. In: Warren, G.W. (Ed.), EPD Congress 1996. TMS, Warrendale, pp. 297–311.
McCreddan, T., Seet, S., 2013. Update on microbial leaching at elevated pH using
BioHeapTM technology. Proceedings of ALTA 2013. ALTA Metallurgical Services,
Melbourne.
McDonald, R.G., Muir, D.M., 2007a. Pressure leaching of chalcopyrite. Part I. Comparison of
high and low temperature reaction kinetics and products. Hydrometallurgy 86,
191–205.
109
McDonald, R.G., Muir, D.M., 2007b. Pressure leaching of chalcopyrite. Part II. Comparison
of medium temperature kinetics and products and effect of chloride ion. Hydrometallurgy 86, 206–220.
Miki, H., Nicol, M., 2011. The dissolution of chalcopyrite in chloride solutions. Part IV. The
kinetics of the auto-oxidation of copper(I). Hydrometallurgy 105, 246–250.
Millero, F.J., Feistel, R., Wright, D.G.T., McDougall, T.J., 2008. The composition of standard
seawater and the definition of the reference-composition salinity scale. Deep-Sea Res.
I 55, 50–72.
Millero, F.J., Woosley, R., DiTrolio, B., Waters, J., 2009. Effect of ocean acidification on the
speciation of metals in seawater. Oceanography 22 (4), 72–85.
Milner, E.F.G., Peters, E., Swinkels, G.M., Vizsolyi, A.I., 1974. Copper hydrometallurgy. U.S.
Patent 3798026 (19 March 1974).
Moreno, P.A., Aral, H., Cuevas, J., Monardes, A., Adaro, M., Norgate, T., Bruckard, W., 2011.
The use of seawater as process water at Las Luces copper–molybdenum beneficiation
plant in Taltal (Chile). Miner. Eng. 24, 852–858.
Moyes, J., Houllis, F., 2002. Intec base metal processes—releasing the potential of chloride
metallurgy technical update and commercialization status. In: Peek, E., Van Weert, G.
(Eds.), Chloride Metallurgy 2002, vol. 2. The Metallurgical Society of CIM, Montreal,
pp. 577–593.
Muir, D., Dixon, D., 2002. Hydrometallurgical chloride processes: a select overview
and comparison, chapter 3, Chloride and chloride-assisted processes for base metal
sulphides. Chloride Metallurgy 2002 Short CourseCIM, Montreal.
Muller, E.L., Basson, P., Nicol, M.J., 2011. Chloride heap leaching. U.S. Patent 8,070,851 B2
(6 December 2012). (also EP 2024523 B1 granted 21 December 2011).
Muñoz-Ribadeneira, F.J., Gomberg, H.J., 1970. Possibilities of recovering copper from
chalcopyrite by nuclear underground leaching in situ. Trans. Am. Nucl. Soc. 13,
636–637.
Muñoz-Ribadeneira, F.J., Gomberg, H.J., 1971. Leaching of chalcopyrite (CuFeS2) with
sodium chloride sulfuric acid solutions. Nucl. Technol. 11 (3), 367–371.
Naderi, H., Abdollahy, M., Mostoufi, N., Koleini, M.J., Shojaosadati, S.A., Manafi, Z., 2011.
Kinetics of chemical leaching of chalcopyrite from low-grade ore: behaviour of different size fractions. Int. J. Miner. Metall. Mater. 18, 638–645.
Nicol, M., Miki, H., Velásquez-Yévenes, L., 2010. The dissolution of chalcopyrite in chloride
solutions. Part 3. Mechanisms. Hydrometallurgy 103, 86–95.
Nicolle, J. Le C., Simmons, S., Bathe, S., Norris, P.R., 2009. Ferrous iron oxidation
and rusticycanin in halotolerant, acidophilic ‘Thiobacillus prosperus’. Microbiology
155, 1302–1309.
Norris, P.R., Simmons, S., 2004. Pyrite oxidation by halotolerant acidophilic bacteria. In:
Tsezos, M., Hatsikioseyian, A., Remoundaki, E. (Eds.), Proceedings of the 15th International Biohydrometallurgy Symposium (Athens). National Technical University of
Athens, Athens, pp. 1347–1351.
Norris, P.R., Davis-Belmar, C.S., Nicolle, J. Le C., Calvo-Bado, L.A., Angelatou, V., 2010. Pyrite
oxidation and copper sulfide ore leaching by halotolerant, thermotolerant bacteria.
Hydrometallurgy 104, 432–436.
Nueva Minería y Energía, 2013. Mining companies will invest over U.S. $10 billion for seawater plants up to 2022. Source: http://www.nuevamineria.com/revista/2013/06/27/
mining-companies-will-invest-over-u-s-10-billion-for-seawater-plants-up-to-2022/.
O'Brien, R.T., MacDonald, C.A., Meadows, N.E., 1999. Chloride–sulphate leaching
from chalcopyrite ores from Sabah. ALTA Copper Sulphides Symposium and Copper
Hydrometallurgy Forum. ALTA Metallurgical Services, Melbourne.
O'Malley, M.L., Liddell, K.C., 1987. Leaching of CuFeS2 by aqueous FeCl3, HCl, and NaCl:
effects of solution composition and limited oxidant. Metall. Trans. B 18, 505–510.
Ohtsuka, N., Mitarai, T., 2007. Chloride ion-resistant sulfur-oxidizing bacteria. U.S. Patent
Application 2007/0202584 (30 August 2007).
Padilla, R., Lovera, D., Ruiz, M.C., 1997. Leaching of chalcopyrite in CuCl–NaCl–O2 system.
In: Mishra, B. (Ed.), Proceedings of the EPD Congress 1997. TMS, Warrendale,
pp. 167–177.
Palmer, B.R., Nebo, C.R., Rau, M.F., Fuerstenau, M.C., 1981. Rate phenomena involved in the
dissolution of chalcopyrite in chloride-bearing lixiviants. Metall. Trans. B 12, 595–601.
Parker, A.J., Paul, R.L., Power, G.P., 1981. Electrochemical aspects of leaching copper from
chalcopyrite in ferric and cupric salt solutions. Aust. J. Chem. 34, 13–34.
Parraguez, L., Bernal, L., Cartagena, G., 2009. Chemical study for selectivity and recovery
of metals sulphides by flotation using seawater. In: Amelunxen, P., Kracht, W.,
Kuyvenhoven, R. (Eds.), Proceedings of PROCEMIN 2009. Gecamin, Santiago,
pp. 323–333.
Paynter, J.C., 1973. A review of copper hydrometallurgy. J. South. Afr. Inst. Min. Metall. 74,
158–170.
Philippe, R., Dixon, R., Dal Pozzo, S., 2010. Sal or desal: seawater supply options for the
mining industry. In: Wiertz, J. (Ed.), WIM 2010: Proceedings of the 2nd International
Congress on Water Management in the Mining Industry. Gecamin, Santiago, pp. 13–20.
Qi, H., Chen, H., Ao, J., Zhou, H., Chen, X., 2009. Isolation and identification of a strain of
moderate thermophilic and acidophilic bacterium from deep sea. Acta Oceanol. Sin.
31, 152–158 (in Chinese).
Rath, P.C., Paramguru, R.K., Jena, P.K., 1988. Kinetics of dissolution of sulphide minerals in
ferric chloride solution, 1: dissolution of galena, sphalerite and chalcopyrite. Trans.
Inst. Min. Metall. C 97, C150–C158.
Rauld Faine, J., Aroca Alfaro, F., Montealegre Jullian, R., Backit Gutierrez, A., 2005. Nonbiochemical method to heap leach copper concentrates. U.S. Patent 6926753
(9 August 2005).
Rautenbach, G.F., Davis-Belmar, C.S., Demergasso, C.S., 2011. Method of treating a
sulphide mineral. U.S. Patent Application 2011/0201095-A1 (18 August 2011).
Richards, R., Chaloupka, M., Sano, M., Tomlinson, R., 2011. Modelling the effects of ‘coastal’
acidification on copper speciation. Ecol. Model. 222, 3559–3567.
Rodgers, L.E., Holden, P.J., Foster, L.J.R., 2002. Culture of Acidiphilium cryptum BV1 with
halotolerant Alicyclobacillus-like app.: effects on cell growth and iron oxidation.
Biotechnol. Lett. 24, 1519–1524.
110
H.R. Watling / Hydrometallurgy 146 (2014) 96–110
Ruiz, M.C., Montes, K.S., Padilla, R., 2011. Chalcopyrite leaching in sulfate–chloride media
at ambient pressure. Hydrometallurgy 109, 37–42.
Safarov, J., Berndt, S., Millero, F., Feistel, R., Heintz, A., Hassel, E., 2012. (p, ρ, T) Properties
of seawater: extension to high salinities. Deep-Sea Res. I 65, 146–156.
Sawyer, H.D., Shaw, R.W., 1983. Hydrometallurgical recovery of copper. Australian Patent
Au-B-17181/83.
Saxena, N.N., Mandre, N.R., 1992. Mixed control of copper dissolution for copper ore using
ferric chloride. Hydrometallurgy 28, 111–117.
Schweitzer, F.W., Livingston, R., 1982. Duval's CLEAR hydrometallurgical process. In:
Parker, P.D. (Ed.), Chloride Electrometallurgy. TMS-AIME, New York, pp. 221–227.
Scott, K.M., 1987. Solid solution in, and classification of, gossan-derived members of the
alunite–jarosite family, northwest Queensland, Australia. Am. Mineral. 72, 178–187.
Senanayake, G., Muir, D.M., 2003. Chloride processing of metal sulphides: fundamentals
and applications. In: Young, C., Alfantazi, A., Anderson, C., James, A., Dreisinger, D.,
Harris, B. (Eds.), Hydrometallurgy 2003, vol. 1. TMS, Warrendale, pp. 517–531.
Skrobian, M., Havlik, T., Ukasik, M., 2005. Effect of NaCl concentration and particle size on
chalcopyrite leaching in cupric chloride solution. Hydrometallurgy 77, 109–114.
Slater, H.B., 1922. Leaching copper and like ores. U.S. Patent 1438869 (12 December 1922).
Spink, D.R., 1977. Energy requirements for the extractive metallurgy of copper. The
Metallurgical Society of CIM, Annual Volume 1977, pp. 167–175.
Taylor, A., Jansen, M.L., 1999. Hydrometallurgical treatment of copper sulphides—are we
on the brink? ALTA Copper Sulphides Symposium and Copper Hydrometallurgy
Forum. ALTA Metallurgical Services, Melbourne.
Tchoumou, M., Roynette, M., 2007. Leaching of complex sulphide concentrate in acidic
cupric chloride solutions. Trans. Nonferrous Metals Soc. China 17, 423–428.
Torres, C., Taboada, M.E., Graber, T., Galleguillos, H., Watling, H., 2013. Copper extraction
from oxide ore using seawater. In: Valenzuela, F., Young, C. (Eds.), Proceedings of
Hydroprocess 2013 (Santiago, Chile). Gecamin, Santiago, pp. 69–77.
Turkmen, Y., Kaya, E., Yavuzhan, O., 2012. Leaching of chalcopyrite flotation concentrate
with CuCl2–NaCl–HCl. In: Özdağ, H., Bozkurt, V., Ipek, H., Biler, K. (Eds.), Proceedings
of the XIII International Mineral Processing Symposium (Bodrum, Turkey), Arber,
Ankara, paper #361.
Turner, D.R., Whitfield, M., 1987. An equilibrium speciation model for copper in sea and
estuarine waters at 25 °C including Complexation with glycine. Geochim.
Cosmochim. Acta 51, 3231–3239.
Turner, D.R., Whitfield, M., Dickson, A.G., 1981. The equilibrium speciation of dissolved
components in fresh water and seawater at 25 °C and 1 atm pressure. Geochim.
Cosmochim. Acta 45, 855–881.
Velásquez-Yévenes, L., Nicol, M., Miki, H., 2010a. The dissolution of chalcopyrite in
chloride solutions. Part 1. The effect of solution potential. Hydrometallurgy 103,
108–113.
Velásquez-Yévenes, L., Miki, H., Nicol, M., 2010b. The dissolution of chalcopyrite in
chloride solutions. Part 2. Effect of various parameters on the rate. Hydrometallurgy
103, 80–85.
Wang, S., 2005. Copper leaching from chalcopyrite concentrates. J. Met. 57 (7), 48–51.
Wang, Y., Su, L., Zhang, L., Zeng, W., Wu, J., Wan, L., Giu, G., Chen, X., Zhou, H., 2012.
Bioleaching of chalcopyrite by defined mixed moderately thermophilic consortium including marine acidophilic halotolerant bacterium. Bioresour. Technol. 121, 348–354.
Watling, H.R., 2006. The bioleaching of sulphide minerals with emphasis on copper
sulphides—a review. Hydrometallurgy 84, 81–108.
Watling, H.R., 2013. Chalcopyrite hydrometallurgy at atmospheric pressure: 1. Review of
acidic sulfate, sulfate–chloride and sulfate–nitrate process options. Hydrometallurgy
140, 163–180.
Watling, H.R., Elliot, A.D., Maley, M., van Bronswijk, W., Hunter, C., 2009. Leaching of a
low-grade, copper–nickel sulfide ore. 1. Key parameters impacting on Cu recovery
during column bioleaching. Hydrometallurgy 97, 204–212.
Watling, H.R., Shiers, D.W., Li, J., Chapman, N.M., Douglas, G.B., 2014. Effect of water
quality on the leaching of a low-grade copper sulfide ore. Miner. Eng. 58, 39–51.
Welham, N.J., Johnston, G.M., Sutcliffe, M.L., 2012. Method for leaching of copper and
molybdenum. World Patent WO 2012/024744 A2 (1 March 2012).
Weston, J.M., Dreisinger, D.B., Hackl, R.P., King, J.A., 1995. Continuous biological leaching
of copper from a chalcopyrite ore and concentrate in a saline environment. In:
Cooper, W.C., Dreisinger, D.B., Dutrizac, J.E., Hein, H., Ugarte, G. (Eds.), Copper
1995–Cobre 1995, vol. 3. CIM, Ottawa, pp. 377–392.
Wiertz, J.V., 2009. When best water use efficiency is not enough, what can the mining
industry do? Proceedings Water in Mining. AusIMM, Melbourne, pp. 13–16.
Williams, T., 2006. BioHeapTM bacterial leaching of the Sherlock Bay nickel mine primary
nickel–copper sulphide ore in saline water. Proceedings of ALTA Nickel/Cobalt 2006.
ALTA Metallurgical Services, Melbourne.
Winand, R., 1991. Chloride hydrometallurgy. Hydrometallurgy 27, 285–316.
Xian, Y.J., Wen, S.M., Deng, J.S., Liu, J., Nie, Q., 2012. Leaching chalcopyrite with sodium
chlorate in hydrochloric acid solution. Can. Metall. Q. 52, 133–140.
Yoo, K., Kim, S., Lee, J., Ito, M., Tsunekawa, M., Hiroshi, N., 2010. Effect of chloride ions on
leaching rate of chalcopyrite. Miner. Eng. 23, 471–477.
Zammit, C.M., Mangold, S., rao Jonna, V., Mutch, L.A., Watling, H.R., Dopson, M., Watkin, E.L.J.,
2012. Bioleaching in brackish waters—effect of chloride ions on the acidophile population and proteomes of model species. Appl. Microbiol. Biotechnol. 93, 319–329.