chalcopyrita.pdf

7/24/2019 chalcopyrita.pdf

1/18

Review

Chalcopyrite hydrometallurgy at atmospheric pressure: 1. Review ofacidic sulfate, sulfatechloride and sulfatenitrate process options

H.R. WatlingCSIRO Minerals Down Under, CSIRO Process Science and Engineering, P.O. Box 7229, Karawara, WA 6152, Australia

a b s t r a c ta r t i c l e i n f o

Article history:

Received 13 May 2013

Received in revised form 19 August 2013Accepted 22 September 2013Available online 10 October 2013

Keywords:

ChalcopyriteLeachingDissolutionPassivationSuldes

The need to process low-grade and/or complex chalcopyrite-containing ores that cannot be concentrated is themaindriver for the development of hydrometallurgicalprocesses. The ferric sulfatesulfuricacid system, with orwithoutthe assistance of microorganisms, hasbeen studiedextensivelybecause it comprises themost promising,low-cost processroute. Alternative oxidants to ferric ionare known but,as yet,their superior oxidation strengthshavenot beenexploited other thanat laboratoryscale, probablydue to their higher costs.Hybridsulfatechlorideand sulfatenitrate systems were included because they may offer specic advantages in some instances. Theaims of this review were to summarise current knowledge in respect of these systems and highlight potentiallyrewarding areas for future research.

2013 Published by Elsevier B.V.

Contents

1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1641.1. Process options for chalcopyrite concentrates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1641.2. Process options for low-grade chalcopyrite ores . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1651.3. Scope of this review . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 166

2. Sulfuric acidferric sulfate systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1662.1. Chemistry of leaching . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1662.2. Chalcopyrite surface overlayers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1672.3. Chalcopyrite surface structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1672.4. Redox control in ferric sulfate systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1692.5. Microorganisms as catalysts in ferric sulfate systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 170

2.5.1. Ambient- to moderate-temperature bioleaching . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1702.5.2. Bioleaching at moderate to high temperature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1702.5.3. Reduced sulfur additives to increase extraction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1712.5.4. Separation of biological and chemical processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 171

2.6. Cations as catalysts in ferric sulfate systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1713. Sulfuric acid alternative oxidants . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 172

3.1. Sulfuric aciddichromate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1723.2. Sulfuric acidchlorate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1733.3. Sulfuric acidpermanganate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1733.4. Sulfuric acidhydrogen peroxidehydroxyl radical . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1733.5. Sulfuric acidozone . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1743.6. Sulfuric acidperoxodisulfate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 174

4. Hybrid sulfatechloride systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1754.1. H2SO4NaClO2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1754.2. H2SO4Fe2(SO4)3NaCl. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1754.3. H2SO4Fe2(SO4)3LiCl . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 176

Hydrometallurgy 140 (2013) 163180

E-mail address:[email protected].

0304-386X/$ see front matter 2013 Published by Elsevier B.V.

http://dx.doi.org/10.1016/j.hydromet.2013.09.013

Contents lists available atScienceDirect

Hydrometallurgy

j o u r n a l h o m e p a g e : w w w . e l s e v i e r . c o m / l o c a t e / h y d r o m e t
http://dx.doi.org/10.1016/j.hydromet.2013.09.013http://dx.doi.org/10.1016/j.hydromet.2013.09.013http://dx.doi.org/10.1016/j.hydromet.2013.09.013mailto:[email protected]:[email protected]://dx.doi.org/10.1016/j.hydromet.2013.09.013http://www.sciencedirect.com/science/journal/0304386Xhttp://www.sciencedirect.com/science/journal/0304386Xhttp://dx.doi.org/10.1016/j.hydromet.2013.09.013mailto:[email protected]://dx.doi.org/10.1016/j.hydromet.2013.09.013http://crossmark.crossref.org/dialog/?doi=10.1016/j.hydromet.2013.09.013&domain=f


2/18

5. Hybrid sulfatenitrate systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1765.1. Sulfuric acidnitric acid . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1765.2. Sulfuric acidsodium nitrate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 177

6. Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 177Acknowledgements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 178References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 178

1. Introduction

There is currently an imbalance between copper supply and worlddemand (Fig. 1). The lack of new, large, high-grade deposits to replacethose that are nearing the ends of their lives, together with theincreased delays between discovery and production due to higherindustry standards and extensive permitting requirements, meansthat the current global imbalance between supply and demand couldcontinue for several years.

At the same time, there has been a decline in copper grades, oftenremarked upon as a future challenge to the copper industry. As anexample, at Escondida Mine, Chile, the world's largest copper mineand producer of 9.5% of the global copper supply, the average coppercontent of mined ore fell from 1.65% in 20072008 to 1.14% in20112012, with measured resource, indicated resource and inferredresource all substantially less than 1% grade (Basto, 2012). In his reviewof historical trends in Australian mining, Mudd (2010)presented datafor copper grades mined in the last 150 years and concluded thatthe decline in ore grades would continue and, in addition, that oremineralogy would become more complex and make the low-gradeores more difcult to process. Strategies to reduce the imbalancebetween supply and demand could include the processing of complexores, the recycling of metals from electronic and other copper-containing waste materials and the development of processes to extractcopper from dirty concentrates containing penalty elements or fromlow grade ores and mine waste materials.

These strategies are not new. Heap, dump, in situ and vatleaching ofwhole ores are the preferred technologies for the processing of low-

grade ores. Data from 2010 showed that proven concentration andpyrometallurgical technologies accounted for 80% of world copperproduction while hydrometallurgical processing of low-grade orescontaining copper oxides or secondary copper suldes contributedabout 20% of annual copper production (Index Mundi, 2013). Aschalcopyrite (CuFeS2) is the most abundant but also themost refractoryof the copper suldes, and with the current extensive exploitation of

low-grade oxide and secondary suldes, it is clear that the proportionof low-grade ores containing chalcopyrite will increase in the future.Therefore, there are imperatives to improve hydrometallurgical tech-nologies for the extraction of copper from chalcopyrite in ores of suchlow grades that they are uneconomic to concentrate, and to augmentthe supply by extracting copper from polymetallic ores of complexmineralogy that cannot be concentrated and from chalcopyrite con-centrates that contain undesirable impurities such as arsenic andcannotbe smelted.

1.1. Process options for chalcopyrite concentrates

In recent years, many processes for the extraction of copper fromchalcopyrite concentrates have been developed, some of them atatmospheric pressure. In these processes, high copper extractionscould be achieved, soluble copper could be separated and puriedusing well-established technologies such as solvent extraction andelectrowinning, and pure, high-quality metal products were recov-erable (Dreisinger, 2006). Using examples from acidic sulfate-basedchemical systems, four strategies employed to maximise copper ex-traction particularly from chalcopyrite were: (i) as high a temperatureprocess as could be reasonably managed at atmospheric pressure (e.g.,BioCOP Batty and Rorke, 2006); (ii) ne grinding to increase chal-copyrite reactivity and overcome passivation (e.g., BacTech/Mintek van Staden, 1998); (iii) the use of additives (e.g., Galvanox Dixonet al., 2008); and (iv) exploitation of innovative combinations ofprocessing technologies, such as heap leaching of concentrates (e.g., the

Geocoat process Harvey and Bath, 2007).Most of the processes were taken to pilot scale and some were dem-

onstrated at a larger scale. However, there are still few commercial-scale operations and energy-intensive processes are unlikely to beeconomic for concentrates, except in circumstances where thecompeting pyrometallurgical technologies cannot be employed. Basedon a survey of process options for copper concentrates assisted by theapplication of a qualitative ranking technique, Lunt et al. (1997)concluded that two of the best options for the hydrometallurgicalprocessing of concentrates were sulfationroastleach and bioleach.They also noted that pyrometallurgical process routes offered somedistinct advantages, not least low cost and industry acceptance andthat for any given project there were a limited number of suitableprocess options. Similar points were made by Peacey et al. (2004),

who outlined a potential opportunity as being the hydrometallurgicaltreatment of dirty copper concentrates with subsequent discharge ofthe acidic leachate to a heap leach/SX/EW operation in circumstanceswhere copper concentrate was readily available but ore for heapleaching was diminishing.

Dreisinger (2006) summarised nine reasons why hydromet-allurgical processes for the treatment of copper sulde concentrateshad failed to achieve sustained commercial production. They can bedistilled into: (i) Production: incomplete copper and precious metalrecovery and/or poor product quality; (ii) Wastes: difculties withtreatment and/or disposal of insoluble residues with particularemphasis on sulfur and (iii) Technoeconomics: costs not competitivewith pyrometallurgy and higher perceived risks associated withimplementing new processes. Peacey et al. (2004) and Dreisinger

(2006) noted that successful processes tended to represent niche

-6

-4

-2

0

2

4

6

2000

2001

2002

2003

2004

2005

2006

2007

2008

2009

2010

2011

[2012]

[2013]

copperproductionsurplusor

deficit[as%o

fproduction]

Fig. 1.Imbalance between copper production and copper demand. Annual surpluses ordecits (data fromInternational Copper Study Group, 2012). Square brackets indicate

predictions.

164 H.R. Watling / Hydrometallurgy 140 (2013) 163180


3/18

opportunities and predicted that theeld would continue to advancebyway of necessity or unique opportunity. Other criteria for successfulprocesses could include (iv) Universality: applicable to all copperconcentrates and able to deal with their various impurities, for example,Zn, Pb, As, Sb, Bi, Se, Hg and F and (v) Simplicity: robust, reliabletechnology suitable for remote locations (Jones, 1996).

1.2. Process options for low-grade chalcopyrite ores

Heap and dump technologies were developed to overcome thechallenges of low copper grade in vast quantities of ores. Thesetechnologies may also be applied to less-difcult ores, not necessarilyof the lowest grade, for (i) smaller-sized ore deposits, (ii) deposits inremote locations lacking the necessary infrastructure for, or access to,pyrometallurgical process routes, or (iii) when pyrometallurgicalprocessing is marginal or uneconomic. Copper recovery rates duringleaching in acidic systems at atmospheric pressure and ambienttemperature vary greatly for different mineral phases (e.g., Table 1).Rates are inuenced by one or more of the physico-chemical propertiesof the ores:

The copper minerals present (Fig. 2). As a general rule M2S is morereadily dissolved than MS and impurities also inuence dissolutionrates.

The mineral associations and copper mineral liberation (Fig. 3).For the example ore, QEMSCAN analysis indicated that only the0.85 mm size fraction contained substantial liberated chalcopyrite;b30% of the chalcopyrite in the larger size fractions had surfaceexpression and, overall, only about 40% of the ore copper contentcould be leached. Particle structure analysis using X ray tomog-raphy failed to reveal particle fracturing that might have assisted inexposing further chalcopyrite grains to the leachate.

The chalcopyrite grain size. For concentrates, the need to ne grindchalcopyrite concentrates to obtain a reasonable (economic) copperextraction rate during leaching is widely reported and constitutes akey parameter in the Albion process (Hourn and Halbe, 1999; Hourn

et al., 1999) and the Mintek/BacTech process (Gericke et al., 2009;van Staden, 1998; Wang, 2005). The same effect applies to large-grained chalcopyrite in ores, assuming that the grains are exposedto the leachate (e.g.,Naderi et al., 2011).

Gangue mineral dissolution consumes acid during leaching orbioleaching and may form amorphous silica gel, thus increasing theviscosity of leachates. Gangue mineral dissolution may increase ironconcentrations in leachates and therefore promote jarosite formation,and may release potentially toxic elements to bioleaching solutions(Dopson et al., 2009; Halinen et al., 2009; Watling et al., 2009).

Heap leaching of copper oxide ores and bioleaching of secondarycopper sulde (chalcocite) ores is widely practised (Domic, 2007).However, heap leaching of chalcopyrite has yet to be implemented atcommercial scale and the extraction of copper from dumps of low-grade chalcopyrite ores or tailings is a practical option only becausethe slow and low copper recoveries are proportionate to the lowprocessing costs (Schnell, 1997). In their reviews,Watling (2006)andPradhan et al. (2008)focused on the potential of heap leaching as aprocess route for low-grade chalcopyrite ores. However, among currentand past heap or dump leaching operations, only one chalcopyrite heapleach operation (Straits Resources' Girilambone Copper Company,NSW) was identied and two large-scale test heaps of chalcopyriteore were described or discussed(Schlitt, 2006; van Staden et al., 2005;Watling, 2006). Watling(2006) expressed the opinion that the efcientheap leaching of chalcopyrite would require greater management andcontrol than was thus far required in oxide and secondary suldeheap leaching andvan Staden et al. (2005) described some of those

Table 1

Comparative dissolution rates in laboratory-scale tests using acidic sulfate systems fordifferent copper-bearing minerals found in heaps and dumps of low-grade ores.

Required Mineral Ideal formula

Hours to days Atacamite Cu2Cl(OH)3Chrysocolla CuSiO32H2ONeotocite (Cu,Mn)2H2Si2O5(OH)4nH2OTenorite CuOMalachite Cu2(CO3)(OH)2Azurite Cu3(CO3)2(OH)2Antlerite CuSO4.2Cu(OH)2Brochantite CuSO4.3Cu(OH)2

Days to months Native Cu CuCuprite Cu2OChalcocite Cu2S

Months to years Bornite Cu5FeS4Covellite CuSEnargite Cu3AsS4Chalcopyrite CuFeS2

0

10

20

30

40

50

60

70

80

90

0 5 10 15 20 25 30 35

Cuextracted(%)

Time (days)

Djurleite

Bornite

Covellite

pyritic chalcopyrite

porphyry chalcopyrite

chalcocite (+digenite)

chalcopyrite (+ 1500g/t Ag)

Fig. 2. Copper extraction from copper sulde minerals(bioleaching, pH1.82,30C, 2wt.%mineral).Data fromFu et al. (2012),Johnson et al. (2008), andRuan et al. (2010).

Quartz

Chalcopyrite

Pyrite

Biotite

Feldspar

K-Feldspar

Titanite

Quartz

K-Feldspar

Fig. 3.Ore particle (N7 mm, 0.7% Cu). QEMSCAN mineral association analysis of multipleblocks of each size fraction of this ore revealed that chalcopyrite was primarily associated

with pyrite and quartz/feldspar.

165H.R. Watling / Hydrometallurgy 140 (2013) 163180


4/18

requirements and how they were being achieved in a pilot heap ofchalcopyrite ore. Vat leaching, largely superseded by heap leaching, isapplied at Mantos Blancos, Chile (Schlesinger et al., 2011) for theextraction of copper from rapid-leaching minerals such as copperoxides or carbonates and may enjoy a revival and wider applicationwith the development of continuous vat technologies (Mackie andTrask, 2009; Schlitt and Johnston, 2010).

The Geocoat bioleaching technology is a hybrid technology. It

combines the high recoveries associated with reactor leaching ofchalcopyrite concentrate with the lower capital cost of heap leaching.In this process, copper concentrate is coated as a thickened slurry ontohost rock particles (625 mm). The coated host rock particles arethen stacked and irrigated as for heap leaching (Harvey and Bath,2007). Sulde dissolution is catalysed by acidophilic microorganismsappropriate to the heap temperature. The complementary Geoleachtechnology comprised a control strategy of aeration and irrigationrates to maximise microbial activity and heat generation and/orconservation within heaps, intended to promote faster leaching,particularly of chalcopyrite. Though widely tested and publicised, thetechnology has not been commercialised.

1.3. Scope of this review

The purposes of this review were to describe the chemistry ofchalcopyrite leaching at atmospheric pressure, in sulfate media withdifferent oxidants or reductants and/or other additives and treatmentsand, where possible, to compare copper extractions. Not surprisingly,most research on chalcopyrite leaching has been undertaken usingchalcopyrite concentrates, whether or not the target application wasthe processing of concentrate or of low-grade ore. Such tests utilised asimple matrix, largely free of gangue minerals that might obscureimportant relationships between starting and product materials.However, one of the difculties of making direct comparisons betweenthe results of published studies has been the failure of manyresearchersto include sufcient detail in methods, or test-mineral properties.Where data are presented in different formats, or disguised for reasonsof condentiality, comparisons between studies become even more

difcult.The review was limited to those sulfate-based systems operated at

ambient pressure within the temperature range constrained by thatcondition. Chloride, nitrate and other acidic leaching systems, and pre-and co-treatments associated with chalcopyrite hydrometallurgy willbe reviewed separately. Sulfate systems operating at above ambientpressure and higher temperatures were comprehensively describedand discussed byDreisinger (2006)andMcDonald and Muir (2007a,b)and references therein, and are not discussed further. In the presentreview, no account was taken of the possible economics of processing,

but rather the aim was to inform researchers, metallurgists and plantoperators about the wide variety of chemical systems that might beapplied in the future when copper demand is higher, ore gradesare lower and new technologies have been developed. While theadvantages or disadvantages of current technologies might be referredto in the contexts of reported results or applications of specic systems,a detailed account of the engineering of such technologies, theirmanagement and/or control are outside the scope of the review.

2. Sulfuric acidferric sulfate systems

The most commonly employed hydrometallurgical process for theoxidation of chalcopyrite and extraction of copper is the sulfuric acidferric sulfate system. It is the system of choice for bioleaching processesat atmospheric pressure, including stirred tank technologies forconcentrates (Batty and Rorke, 2006) and heap or dump technologiesfor low-grade ores (Watling, 2006). However, whether chemical orbio-assisted chemical leaching is undertaken, the oxidation of chal-copyrite is slow and incomplete, possibly the consequence of themineral crystalline structure and changes therein (de Oliveira et al.,2012; Klauber, 2003) but also inhibited by insoluble secondary reactionproducts (elemental sulfur, ferric hydroxides or hydroxysulfates)forming what are termed passivation or overlayers on chalcopyritesurfaces. There is value in examining the chemistry of ferric sulfateand acid leaching systems, some of which are summarised inTable 2,because they comprise the basic chemical systems against which theuse of other chemical systems would be assessed during thedevelopment of alternative atmospheric-pressure leaching processes.

2.1. Chemistry of leaching

The generally accepted reactions for the extraction of copper fromchalcopyrite via oxidation (oxygen, ferric ions) or acid (H2SO4) leachingare:

CuFeS2 2Fe2SO43 CuSO45FeSO4 2S0

1

CuFeS2 O2 2H2SO4 CuSO4FeSO42S0

2H2O 2

4FeSO42H2SO4O2 2Fe2SO43 2H2O: 3

Dutrizac (1981)reviewed the literature on ferric ion leaching ofchalcopyrite (reaction (1)) and, with data from ancillary studies,summarised existing knowledge at that time as: (i) leaching wasindependent of acid above that required to keep iron in solution (andiron neednot exceed 0.1M Fe2(SO4)3 Parker et al., 1981); (ii) leaching

Table 2

Acidic sulfate processes developed for chalcopyrite or other copper sulde concentrates or ores and operated at atmospheric pressure.

Process Temperature(C)

Size(m)

Differentiating conditions Scale

Heap leaching (e.g., Quebrada Blanca, Chile) (Domic, 2007) Ambient Crushed ore, stacked, irrigated with dilute H2SO4solution CROM dump leaching (e.g., La Escondida, Chile) (Domic, 2007) Ambient ROM ore stacked and irrigated with dilute H2SO4 CGeocoat (Harvey and Bath, 2007) Ambient n.d. Sulde concentrate supported on host rock particles; sulfate

bioleaching in heapsD

BacTech/Mintek (Gericke et al., 2009; van Staden, 1998; Wang, 2005) 3550 510 Sulfate bioleach;ne grind DBioCOP (Batty and Rorke, 2006) 6580 37 Sulfate leach at higher temperature using archaea as catalysts DSepona (Baxter et al., 2003) 80 100 Ferric sulfate leach of chalcocite rimming pyrite (FeS2) CGalvanox (Dixon et al., 2008) 80 5375 Ferric sulfate leach; pyrite to chalcopyrite in 2:1 ratio PAlbion (Hourn and Halbe, 1999; Hourn et al., 1999) 85 510 Ferric sulfate leach;nely ground concentrate PCobre Las Cruces Project (Fleury et al., 2010) 90 150 Ferric sulfate generated from pyrite oxidation with O2; mainly Cu9S5

and Cu2S (6.2 wt.%), minor CuFeS2

C

Cuprochlor process (Espejo et al., 2001; Herreros et al., 2006) Ambient Sulfatechloride leach with CaCl2agglomeration CNitric acid route(Bjorling et al., 1976) 90 n.d. Sulfatenitric leach of chalcopyrite concentrate L

a In theSepon process, residual pyrite isoated andthen oxidised under pressure to generate acidand ferricions foruse in thesecondaryheap leach. C commercial;D demonstration;L

laboratory P pilot; n.d. not disclosed.

http://-/?-http://-/?-


5/18

rates were proportional to chalcopyrite surface area, only slightlydependent on ferric ion concentrations and increased with increasedtemperature; and (iii) increased sulfate ion concentration contributedto slower leaching rates. Hiroyoshi et al. (1997) proposed thatchalcopyrite was oxidised by dissolved oxygen (reaction(2)) and thatthe generated ferrous ions were oxidised to ferric ions (reaction(3)).Their evidence for these reactions was that the pH increased withtime when ferrous sulfate was added to acid medium containing chal-

copyrite, indicating the consumption of protons. Subsequent studiesby the authors were incorporated into a two-stage reaction model toexplain ferrous-promoted chalcopyrite dissolution (Hiroyoshi et al.,2000and references therein) in which (i) chalcopyrite was reducedby ferrous ions in the presence of cupric ions to form chalcocite and(ii) chalcocite was oxidised (more readily than chalcopyrite) bydissolved oxygen and/or ferric ions to form cupric ions and elementalsulfur insoluble product. From thermodynamic calculations, chalcociteformation only occurred when the redox potential of the solutionwas lower than the critical potential (a function of the ferrous andcupric ion activities); and the optimum oxidationreduction potential(ORP) for chalcopyrite leaching increased with increased copper(II)concentrations. Conversion of the normalised redox potential to thesolution redox potential showed the optimum ORP for chalcopyriteleaching to be a function of cupric and ferrous ion concentrations(Hiroyoshi et al., 2008).Sandstrm et al. (2005)examined the reactionproducts formed during the chemical leaching of chalcopyrite andfound Cu(I) species in all samples, consistent with theHiroyoshi et al.(2000)two-stage model.

Nicol and Lzaro (2003)showed that dissolution of chalcopyritecould occur in the absence of any oxidising reagent at potentialslower than 0.4 V versus Standard Hydrogen Electrode (SHE), with theformation of a detectable soluble sulfur species, such as hydrogensulde. They presented data indicating that reaction(4)was the mostproton-consuming reaction and hypothesised that it was governed bytwo steps: (i) rapid dissolution to establish the equilibrium betweensoluble species at the chalcopyrite surface and the bulk solid and(ii) rate-determining diffusion of the soluble species away from thesurface. Nicol and Lzaro (2003) examined the thermodynamic

feasibility of reaction (4) at 25 C and higher temperatures usingOutukumpu HSC software and the CrissCobble technique.

CuFeS2 4HCu2 Fe2 2H2S: 4

2.2. Chalcopyrite surface overlayers

There is a general view that theslow dissolution rates of chalcopyritein ferric sulfate leaching media are, in part, a consequence of theformation of secondary reaction product on the mineral surface. Thesesecondary reaction products, passivation or overlayers, may be ofsuch depth and sufciently compact (Fig. 4) as to hinder the diffusionof reagents to the chalcopyrite surface and/or the diffusion of ionsaway from the surface, or they may have low electrical conductivity.

Reaction products thought to be implicated in chalcopyrite passivationincluded:

Complex lms of suldes, polysuldes and/or elemental sulfur(reaction (1)), or an unreactive, copper-rich layer on partially-oxidised chalcopyrite surfaces (Fu et al., 2012; Hackl et al., 1995;Linge, 1976; Nicol and Lzaro, 2003; Warren et al., 1982);

Iron oxides, hydroxides, iron hydroxysulfates and/or jarosite(reaction((5))(Crdoba et al., 2008a,b; Gmez et al., 1996; Parkeret al., 2003; Sandstrm et al., 2005; Stott et al., 2001).

3Fe3 2SO42

6H2O MMFe3 SO4 2 OH 6 6H

5

where M = K+

, Na+

, NH4+

or H+

.

In a critical review of the surface science investigations of whatmight be responsible for hindered chalcopyrite dissolution in ferric

sulfate systems,Klauber (2008) clari

ed some of the discrepanciesbetween studies.Klauber (2008)concluded that polysuldes could berejected as passivation candidates because they were too reactiveand oxidised to elemental sulfur on exposure to air, especially in thepresence of moisture and metal-decient suldes and that the presenceof metal decient suldes was questionable. Overlayers of elementalsulfur formed on chalcopyrite surfaces could hinder dissolution butmight subsequentlybe peeledoff,allowing furtherdissolution. Overlayersof iron(III)-insoluble compounds such as jarosite were an inevitableconsequence of long leach times, triggered by the formation of anintermediate ferric sulfatecompound at the chalcopyrite surface as anintegral part of the ferric ion leaching mechanism.

More recently,Debernardi and Carlesi (2012)presented a review ofchalcopyrite passivation in which they outlined the capabilities andlimitations of electrochemical and surface analysis techniques inresolving the nature of chalcopyrite passivation. As part of their review,in which they summarised many of the papers discussed byKlauber(2008),Debernardi and Carlesi (2012)questioned whether the currenttechniques in electrochemistry and surface analysis could providereliable data on surface phenomena and whether those data werebeing interpreted correctly. They identied potential research topicsthat might assist in the future interpretation of data obtained usingsophisticated instrumentation.

2.3. Chalcopyrite surface structure

In their initial study, using X-ray photoelectron spectroscopy (XPS),Klauber et al. (2001) proposed that a disulde phase was among

the reaction products on leached chalcopyrite surfaces (Fig. 5). The

a

b

CuFeS2

CuFeS2

Fig. 4.Examples of reaction products on chalcopyrite surfaces: (a) ferric hydroxysulfatelayer after 96 h exposure to ferric sulfate solution (pH 1.8); (b) elemental sulfur after 2 hexposure to HOCl (pH 4).

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


6/18

synthesis of modelpolysuldes and theirsubsequentexamination using

XPS indicated that polysuldes did not play a role in inhibitingchalcopyrite dissolution, leading to the conclusion that chalcopyritedissolution occurred via the oxidation of the disulde phase (Parkeret al., 2003). In a subsequent detailed study of a freshly-cleavedchalcopyrite surface prepared in an inert atmosphere, Klauber (2003)again identied a disulde phase and presented a well-constructedargument for the formation of a two-layer surface phase containingpyritic FeS2at 50% of the density of bulk pyrite. He proposed a modelthat described thephysical andredox mechanisms of thereconstructionand concluded that it would lead to a lower-energy chalcopyritesurface.

de Oliveira and Duarte (2010) used density functional theory withinplane wave framework to analyse the reconstruction of the chalcopyrite(001) surface and its reactivity, at the molecular level. The calculations

predicted that the sulfur-terminated (001) surface underwent areconstruction in which disulde groups formed on the surface,consistentwith the XPS results ofKlauber (2003), with the concomitantreduction of iron(III) to iron(II) oxidation state.

Subsequentlyde Oliveira et al. (2012)applied plane wave densityfunctional calculations to a comparative study of the (001), (100),(111), (112), (101) and (110) chalcopyrite surfaces including bothsulfur- and metal-terminated cleavages of the rst three (Fig. 6).Three different surface reconstructions emerged from the calculationsof de Oliveira et al. (2012)which they described as follows: (i) onsurfaces which have relatively close sulfur atoms in the rst atomic

layer with lower coordination number than in the bulk chalcopyrite

(the (001)-S, (100)-S, and (112) surfaces), the formation of disuldegroups via an oxidative process with concomitant reduction of Fe(III)to Fe(II) or, for the sulfur-terminated (111)-S face, the formation of theS42 group; (ii) on surfaces terminated in metal atoms (the (001)-M,

(100)-M and (111)-M surfaces), the formation of metalmetal bondsresulting in alloy-like structures beneath a layer of exposed sulfuratoms, and (iii) the relaxation of metal atoms with downward migration(the (101), (110) and part of the (112) surfaces). The experimentaldetection of the disulde groups has been reported (e.g., Klauber,2003) but the predicted formation of the proposed metalmetal, alloy-like groups and/or metal migration requires greater investigation inrespect of elucidating the slow dissolution kinetics of chalcopyrite.

de Lima et al. (2011)directed their study towards water adsorptionon the reconstructed (001)-S and (001)-M chalcopyrite surfaces. Their

calculations predicted that (i) an iron atom was the most favourablesite for a water molecule to adsorbon thereconstructed (001)-Ssurface,(ii) dissociation of the water molecule was not favoured, and (iii) wateradsorption on the (001)-M surface was also unfavourable, suggestingthat the (001)-M surface would exhibit hydrophobic characteristicsthat might inuence surface reactivity. In an extension of that study,the calculations ofde Lima et al. (2012) revealed additional insightsrelevant to chalcopyrite leaching. Their calculations predicted that(iv) water molecules and chloride ions, with their similar adsorptionenergies, competed for iron sites on (001)-S chalcopyrite surfaces,(v) that sulfate and bisulfate ions were more strongly bound to the

Fig. 5.X-ray photoelectron spectrum of a chalcopyrite surface after treatment with 0.1M ferric sulfate solution at pH 1.8. Deconvolution of the spectrum to reveal the underlying sulde,and the disulde, elemental sulfur and sulfate species on the surface (Klauber et al., 2001).

Fig. 6.Chalcopyrite surface reconstructions which may partially impact on leaching kinetics.Reprinted with permission from deOliveira, C., de Lima G.F., de Abreu, H.A. and Duarte, H.A. (2012), Journal of Physical ChemistryC 116, 63576366. Copyright 2012 American Chemical

Society.

http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B6%80


7/18

surfacethan water or chloride ion,that binding wasbidentate, replacingtwo water molecules, and that bisulfate was bound more strongly thansulfate, and (vi) that high acidity would lead to the protonation of thedisuldes on the surface, probably weakening the disulde bond. It isanticipated that future contributions from this research group shouldadvanceour understanding of the slow kinetics of chalcopyrite leaching,especially where predicted surface changes and interactions lendthemselves to experimental validation.

Recently,Li et al. (2013)conducted a comprehensive review of theliterature on the leaching of chalcopyrite, specically the nature androles of secondary product overlayers, the reconstruction of the bulksulde crystal when fractured or otherwise exposed as a surface andthe consequences of reconstruction in different leaching systems, andthe participation of redox reactions (both oxidation and reduction)during leaching. While they reviewed many of the publicationsdiscussed by Klauber (2008) andDebernardi and Carlesi (2012), Liet al. (2013)adopted a broader approach, taking account of researchusing oxidants other than ferric ion and a consideration of the abovemodelling predictions.Li et al. (2013)concluded that three importantquestions remained to be answered in order to explain the chalcopyriteleaching mechanism(s), relating to: (i) the interaction betweenchalcopyrite surfaces and oxidants; (ii) the relationships betweenoxidants and secondary surface reaction products and (iii) elucidationof the key factor that determines whether a chalcopyrite surface willbecome passivated.

2.4. Redox control in ferric sulfate systems

The acceleration of copper extraction from chalcopyrite in acidicferric/ferrous sulfate media at low ORP in the temperature range of2050 C (e.g.,Fig. 7), is well established. As a consequence, differentstrategies for ORP control were implemented in fundamental studiesand process developments. Bruynestein et al. (1986)established aslurry potential of 540660 mV versus SHE via the initial addition ofthiosulfate and cupric ions in a process that was specically designedto promote the extraction of copper from chalcopyrite and the partial

oxidation of the sul

de to elemental sulfur (not sulfate). Kametaniand Aoki (1985)maintained a bioleaching suspension at the relativelylow ORP of 635 mV (versus SHE) by controlled additions of potassiumpermanganate solution, to maximise copper extraction from achalcopyrite concentrate.Sandstrm et al. (2005)maintained an ORPof 635mV (versusSHE) during the chemical leaching of a chalcopyriteconcentrate using potassium permanganate, but reduced the airow(oxygen) and added sodium sulte to maintain a similarly low ORP

during bioleaching.Third et al. (2002)andGericke et al. (2010)alsoused oxygen limitation to control ORP during bioleaching, the formerby temporarily arresting the air supply to the reactor when the ORPwas greater than a designated set point, and the latter by adjustingimpeller speed to control oxygen mass transfer.Ahmadi et al. (2010)exercised electrochemical control of ORP by applying a direct currentinto the ore suspension with/without bacteria.

Viramontes-Gamboa et al. (2007)studied the oxidative dissolution

of chalcopyrite in acidi

ed ferric/ferrous sulfate media in the tem-perature range of 2580 C and concluded that chalcopyrite displayedthe activepassive behaviour of passivating metals. They showed thatthe predicted electrochemical passivation potentials (E-pp) obtainedusing electrochemical techniques were consistent with the resultsfrom leaching experiments and reported that E-pp increased from680 mV (versus SHE) at 25 C to 755 mV at 80 C. In their systematicstudy of the effect of acidity in the range of 2100 g H2SO4L

1, theyalso showed that E-pp were insensitive to acidity in the temperatureranges of 2540 C and 6080 C but that between 40 and 60 C, anacid-dependent transition of E-pp occurred from 680 to 755mV (versusSHE). Subsequently, Viramontes-Gamboa et al. (2010) collated datafrom studies in which chalcopyrite leaching kinetics were investigatedunder conditions of controlled ORP and found that chalcopyritepassivation was reported to occur at one of two potentials, eithergreater than about 680 mV or greater than about 750 mV (versus SHE)for samples from different locations and with different impurities.They then demonstrated electrochemically that transitions from activeto passive and vice versa did not take place at the same potential; attemperatures of 50 C or higher and conditions of increasing potential,chalcopyrite passivated at about 750755 mV (versus SHE) but if thechalcopyrite was already passivated and the potential was decreased,the chalcopyrite only became active at about 680 mV (versus SHE).Viramontes-Gamboa et al. (2010) noted that this result offered apossible explanation for the apparent discrepancies between publishedcontrolled-ORP leaching studies but did not elucidate the factors thatdetermined whether chalcopyrite would be active or passive duringleaching in the potential range of 680755 mV (versus SHE).

The benets of enhanced chalcopyrite dissolution under conditions

of controlled ORP were sufciently important to have prompted patentapplications relating to stirred tank and heap technologies. The patentofBruynestein et al. (1986)was noted above. In the patent ofDixonand Tshilombo (2005), the chemical leaching of chalcopyrite in acidicsulfate medium, with added pyrite as catalyst, was conducted underconditions whereby the pyrite is not materially oxidised, for exampleby maintaining the operating solution potential at a suitable level.Redox-control during chalcopyrite bioleaching was also a key claim inpatents byvan der Merwe et al. (1998),Pinches et al. (2001a, b)andLastra and Budden (2002).Pinches et al. (2001a)proposed to bioleachchalcopyrite in stirred tanks by controlling air, oxygen and/or carbondioxide concentrations to limit the efciency of iron(II)-oxidation bythe microbial community, specically to control the ORP in the range575675 mV (versus SHE) thereby to promote the rate and extent of

chalcopyrite oxidation.Pinches et al. (2001b)claimed that the abovemethodology could also be applied in a heap leaching process inwhich the leach solution was conditioned to provide the said surfacepotential. This latter claim was supported by the results of comparativecolumn leaching tests in which the iron(III)H2SO4feed solutions wereproduced in a separate bacterial ferric-ion-generator and could thus becontrolled at high redox (775 mV versus SHE) or low redox(637 mV), as required. The authors noted that in a normalheap leachsystem, the presence of the natural bacterial community would makecontrol of ORP extremely difcult.Lastra and Budden (2002)proposedtwo methodsof controlling the ORP in heap bioleaching of chalcopyrite.In their invention, they described (i) the direct method: the im-provement ofmaintaining the reduction potential at below or aroundEH b 550 mV (reference electrode not specied but probably SCE or

Ag/AgCl), a condition in which it was proposed could be achieved by

0

20

40

60

80

100

0 5 10 15 20 25

copperex

trac

tion

(%)

leach duration (days)

chemical

chemical ORP control

bacterial

bacterial ORP control

Fig. 7.Enhanced copper extraction at 50 C from chalcopyrite with bacteria and redoxcontrol (re-drawn from Ahmadi et al., 2010). Electrochemical control of ORP at

400

450 mV versus Ag/AgCl reference electrode.

http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80http://localhost/var/www/apps/conversion/tmp/scratch_4/image%20of%20Fig.%E0%B7%80


8/18

limiting the concentrations of oxygen in a heap operated at N50C; and(ii) the indirect method, in which they proposed that biologicaloxidation of ferrous ions would be achieved in a separate reactor andthat oxygen would be excluded from the heap, so as to prevent ferrousion oxidation in situ. Key claims in the patent ofHunter (2006)werethat a chalcopyrite heap would be inoculated with asulde-oxidisingbacterial culture that was either inefcient at oxidising iron(II) or didnot have that specic ability, and that the process water being fed to

the heap would be maintained at ORP lower than 695 mV (versusSHE), such that the prevailing chemical conditions are conduciveto leaching chalcopyrite while being non-conducive to surfacepassivation. As heaps are open systems, the maintenance of low-ORPwould be difcult should adventitious colonisation of the ore by nativeiron(II)-oxidising microorganisms occur.

2.5. Microorganisms as catalysts in ferric sulfate systems

Bioleaching can be considered as a chemical system catalysed bymicroorganisms. The bioleaching of sulde minerals with emphasis oncopper sulde leaching using low-temperature heap technology wasreviewed (Watling, 2006 and references therein). Briey, suldeminerals in pristine (but exposed) deposits, mine waste dumpsand tailings, and/or managed heaps represent complex habitats foracidophilic microorganisms but very few of the microorganisms havebeen cultured, isolated or identied.

2.5.1. Ambient- to moderate-temperature bioleaching

Microorganisms participate in the dissolution of sulde minerals intwo main ways: (i) by catalysing the oxidation of ferrous ions to ferricions (reaction(3)) more rapidly than occurs in chemical systems and(ii) by oxidising sulfur to sulfate, helping to remove sulfur overlayersand generating additional acid (reaction (6)). The costs associatedwith the maintenance of sulde bioleaching microorganisms areminimal because they gain energy from the redox reactions(3) and(6) (oxygen is usually supplied by air), acquire carbon for growthfrom the carbon dioxide in air and obtain phosphorus, nitrogen,potassium and micronutrients from the ore environment. Microbialattachment to mineral surfaces is mediated by the properties of theextracellular polymeric substances produced by the cells andattachment is mineral and site specic and may change the propertiesof the underlying mineral surface (Watling, 2006 and referencestherein).In some cases theaddition of organic compounds mayenhancecopper extraction (e.g.,Hiroyoshi et al., 1995; Li and Li, 2010; Mallickand Dasgupta, 1997) but, more generally, the presence of organiccompounds is deleterious to microbial growth, iron(II)- and/or sulfur

oxidation (e.g.,Mazuelos et al., 1999; Okibe and Johnson, 2002; Tormaand Itzkovitch, 1976).

2S0 3O22H2O 2H2SO4: 6

2.5.2. Bioleaching at moderate to high temperature

Given the strong temperature dependence of chalcopyrite dis-solution (Fig. 8), the bioleaching of chalcopyrite at temperatures

60

90 C is an attractive R&D goal (e.g.,Crdoba et al., 2008d; Gerickeet al., 2001; Rodrguez et al., 2003) but not yet a full-scale, commercialprocess (e.g.,Batty and Rorke, 2006). The iron(II)-(reaction(3)) andsulfur-oxidising (reaction (6)) capabilities of archaeal members ofAcidianus,MetallosphaeraandSulfolobusand possibly other genera areexploited in higher-temperature bioleaching. Some recent keyndings,most often arising from studies using individual species but applicableto other archaea, are summarised: Archaeal species contribute to stableleach environments in continuous pilot plants operated at 1220%solids loadings (Gericke and Pinches, 1999; Sandstrm and Petersson,1997). Limitations resulting from reduced gas solubilities in tanksoperated at high temperatures (6080 C) can be minimised bysupplying oxygen and/or carbon dioxide to the reactors (Batty andRorke, 2006; Dew et al., 2001). Microbial sensitivity to shear from

agitators results in an upper limit for solids loading of about 12.5% atdemonstration scale (Batty and Rorke, 2006). Archaea exhibit differentsensitivities to soluble metals arising from the leaching of impureconcentrates or complex ores, but in the case of copper, archaealcultures tolerate up to 35 g Cu L1 (Batty and Rorke, 2006). Whileredox-controlled bioleaching of chalcopyrite at temperatures 2050Cresults in enhanced copper extraction, redox-control is less benecialin the temperature range of 6080C(Gericke et al., 2010). The additionof iron(II) (0.5 g L1) promotes cell growth, high ORP and increasedchalcopyrite dissolution rates but the presence of iron at N5 g L1

promotes jarosite formation on chalcopyrite surfaces that hindersfurther dissolution (Crdoba et al., 2008d; Sandstrm et al., 2005).Thus operating conditions should be optimised to minimise theformation of iron-rich overlayers, possibly requiring strategies for ironcontrol. Cell attachment to surfaces is rapid (within 20 min) ( Konishiet al., 1999, 2001). As long as a fraction of the population can attachto the chalcopyrite surface, copper extraction is efcient, but thepopulation may not grow if cells are prevented from approaching thechalcopyrite surface (Gautier et al., 2008). Attached cells oxidiseelemental sulfur on particle surfaces (Liang et al., 2012) and planktoniccells oxidise the soluble reduced sulfur or iron(II) species that accu-mulate in solution (Jordan et al., 2006). The joint activity of attachedcells and planktonic cells has been termed cooperative leaching

0

20

40

60

80

100

120

25 30 35 45 50 55 60 65 70 75 82

copperex

trac

ted

(%)

in96hours

temperature (degrees Celsius)

inoculated tests abiotic tests

mesophiles

moderatethermophiles

archaea

Fig. 8.Temperature dependence of chalcopyrite dissolution. Copper extraction from a chalcopyrite concentrate of composition 64% chalcopyrite, 6.6% pyrite, 3.3% pyrrhotite and 25%quartz,with meanparticlesize 27m, P8066m andsurface area 0.36m

2g1 (5-pointBET analysis). Tests (96h) inoculatedwith single speciesor mixed cultures; 2535C, mesophiles;

45

55 C, moderate thermophiles; and 60

82 C, thermophiles (archaea). Abiotic tests were not inoculated.

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


9/18

(Rodrguez et al., 2003). Archaeal activity does not change theunderlying mechanism of chalcopyrite dissolution by ferric ions.However, surface-attached archaea may increase the corrosionpotential and corrosion current density of chalcopyrite and may alsoreduce the polarisation resistance of chalcopyrite during bioleaching(Li and Huang, 2011).

2.5.3. Reduced sulfur additives to increase extraction

The addition of elemental sulfur to sul

de bioleach operations is astrategy that can be used to meet three needs: (i) an alternativemeans of providing the acid required to extract the target metals(reaction(6)); (ii) the generation of heat to increase copper extractionrates; and (iii) a means of promoting microbial colonisation in a low-grade sulde ore.

In the usual method of heap leaching, sulde oxidation occurs inacidic ferric sulfate solutions, so the presence of acid is essential.However, for heaps of low-grade ore, greater than 99% of the bulk oreis comprised of gangue minerals which also consume acid, to a greateror lesser extent, with time. This acid consumption during ganguemineral dissolution constitutes a reagent loss because it does notcontribute to the target-metal extraction. For the majority of oresprocessed using heap technology, acid consumption is a majoroperating cost. It is not surprising, therefore, that the biologicalproduction of sulfuric acid through the oxidation of suldes orelemental sulfur has been described. Young et al. (2003)patented amethod for microbiological production of acid in which twoapplications (processes) were described, one having many of thecharacteristics of a heap leach of sulde minerals and the otheremploying reactor technology, such as an agitated tank. In both, theaim was to oxidise elemental sulfur or mineral suldes to producesulfuric acid using sulfur-oxidising acidophilic microorganisms thattolerate the high sulfate and metal concentrations occurring in processsolutions. In the patentdescribedby Duyvesteyn et al. (2002), a primaryclaim wasthat theprocess described would lessenor eliminate theneedto supply commercially-purchased or chemically-produced sulfuricacid (such as from an on-site acid production plant). The proposedtechnology/application was broad; the sulfur-amended ore could be in

the form of a slurry, a heap, or a charge in a vat, with or without prioragglomeration. Bouffard et al. (2009) listed the benets of sulfuraddition directly to a bioleaching heap as being reduced acid con-sumption by gangue minerals, faster mineral leaching kinetics, andheat generation. However, they also noted the disadvantages of addingcommercially-available sulfur, which is sold and transported in prill,slate or pellet form. In these commercial products, particle sizesrange from 1 to 13 mm and, as such, have low surface areas. Sulfur isvirtually insoluble in water, the solubility of rhombic sulfur being1.90.6108molS8kg

1 (Boulegue, 1978). Therefore the thrust oftheBouffard et al. (2009)study was the addition of lignosulfonate toimprove the wettability of sulfur particles during a wet-grindingprocess. In laboratory tests, Salo-Zieman et al. (2006) studied thebioleaching of a pyrrhotite-rich, acid-consuming nickelcopper ore

and reported that the addition of elemental sulfur reduced the needfor acid supplementation during bioleaching with mesophiles.

Kohr et al. (2004)described a process for heating a heap rapidly totemperatures at least 50C by augmenting thefuelcontent of the oreto at least 10 kg of exposed sulde per tonne of ore, via the addition ofsulde minerals. They further dened the process as the additionpreferably of small particle sizes of sulfur-containing materials thatwould generate a large amount of heat when oxidised. The details ofthe process were as already described in a family of patents on theGEOCOAT process (Harvey and Bath, 2007), in which host rock wascoated with ne particles of material rich in sulde, and leached in aheap. Essentially the same process was described in their more recentpatent (Kohr et al., 2011). Of thepatents examined thus far, theprocessdescribed by Hunter and Williams (2006) wastheonlyone inwhich the

addition of elemental sulfur was specically proposed, compared with

others in which sulfur addition was subsequently dened to meanthe addition of mineral suldes. This patent differed also in that theprimary purpose of the addition was to achieve a sufcient sulfurcontent (between 2 and 20% w/w) to promote a rapid increase in theinternal temperature of a chalcopyrite heap, such that thermophilicmicroorganisms could grow and function. The concomitant productionof acid in the heap from microbial sulfur oxidation was noted asbenecial.

In a laboratory study, augmentation of a chalcopyrite concentratewith elemental sulfur was intended to selectively promote the growthof sulfur-oxidising microorganisms in an iron- and sulfur-oxidising,mixed-mesophilic culture during bioleaching, the rationale being thatsulfur overlayers formed on chalcopyrite surfaces could become ratelimiting if not removed efciently. However,Xia et al. (2012) foundthat the strategy conferred an advantage onAcidithiobacillus thiooxidansat the expense of Leptospirillum ferriphilum to such an extent thatiron-oxidation was inhibited and copper extraction was reduced.Optimisation of the conditions should correct the imbalance betweensulfur- and iron-oxidation. This is essentially the underlying conceptof the Bioheap process (Hunter, 2002a,b) in which a sulfur-oxidisingculture inefcient at iron oxidation was used to inoculate a heap ofchalcopyrite-containing ore. The process also provided for iron controlby passage of process solution through a second heap of largely barrenand inert ore, thus reducing the possible formation of passivatingiron(III) oxides/hydroxides/ hydroxysulfates on chalcopyrite surfaces.Uhrie et al. (2012)proposed the use of elemental sulfur to promotemicrobiological growth in a separate bioreactor for the purpose ofgenerating an inoculum for a heap leaching process.

2.5.4. Separation of biological and chemical processes

The benets of implementing an indirect bioleaching process are:(i) the independent optimisation of the biological and chemicalprocesses, (ii) restricted access of microorganisms to sulfur reactionproducts, minimising sulfur oxidation to sulfate (reaction(6)) and, insome reactor congurations, (iii) lower oxygen requirement duringleaching, (iv) a degree of acid and iron control through iron hydrolysisand jarosite formation (reaction ((5)) and (v) lower exposure of

microorganisms to potentially toxic components in concentrates orores.

The separation of biological and chemical processes duringbioleaching was briey mentioned above in the context of ORP controland is generally referred to in the literature as indirect bioleaching. Itdepends upon the propensity of acidophilic microorganisms to becomeattached to inert substrates and form biolms, within which bacterialoxidation of ferrous ion occurs (reaction (3)) and from which thecirculating ferric-ion rich medium is fed to a chemical leaching reactor.Ferrous ion oxidation can be achieved with different types of reactorsfurnished with different inert support materials for the biolm. Mostapplications apply packed beds, uidised beds or rotating biologicalcontactors furnished with a variety of support materials (e.g., resin,glass or alginate beads, activated carbon, ground silica, jarosite, poly-

ethylene, polyvinyl chloride, polystyrene). Relatively small (economic)reactors can be used because ferric ion productivity is generally high. Anumber of two-stage processes have been developed and implemented,most providing suitable growth conditions forAt. ferrooxidans, the speciesthought to be most prevalent in bioleaching at the time of those studies.Acronyms used to describe three examples are BACFOX (bacterial lmoxidation Karavaiko, 1985), BFIG (bacterial ferric ion generator VanStaden, 1998) and IBES (indirect bioleaching with effects separation Carranza et al., 1993).

2.6. Cations as catalysts in ferric sulfate systems

Several studies have been conducted in which selected cations wereadded to sulfuric acidferric sulfateleach systems and the enhancement

(or not) of copper extraction from chalcopyrite reported. Barriga

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


10/18

Mateos et al. (1987)compared a number of elements and showed thatthe precipitation of silver sulde onto fresh chalcopyrite surfaces, oronto already-passivated surfaces, reactivated chalcopyrite. In othercomparative tests, silver ions were the most efcient at promotingchalcopyrite dissolution (Ballester et al., 1990, 1992; Escudero et al.,1993; Muoz et al., 2007a). Johnson et al. (2008) reported thatconcentrates containing high silver contents were amenable tobioleaching at low temperatures withmesophiles but that concentrates

with lower silver contents required higher temperatures and moderatethermophiles.The benecial action of silver(Ag+) wasthought to proceed through

reaction with the chalcopyrite surface with the formation of silversulde (Ag2S) on the sulde mineral surface (reaction(7)) followedby the oxidation of Ag2S by ferric ions (reaction (8)) (Miller andPortillo, 1979). However, this simple mechanism does not account forall the observations made then and subsequently. For example, it wasshown that ferric ion oxidation of Ag2S (reaction (8)) requiredtemperatures greater than 100C (Dutrizac, 1994) and elemental sulfurwas not detected on a chalcopyrite surface post silver-catalysedleaching (reactions (8)and (9)) (Parker et al., 2003). Hypothesesabout the formation of a small amount of silver (Muoz et al., 1998;Price and Warren, 1986) included reactions(9)((11)](Kolodziej andAdamski, 1984; Price and Warren, 1986). In addition, the formation ofAg2SO4in the presence of a large excess of Ag+ (Warren et al., 1984)was attributed to high oxidation potentials and excess Ag+ near theAg2S surface (Price and Warren, 1986), a reaction that would depleteAg+ concentrations and result in lower (catalysed) copper extraction,as was also reported by Crdoba et al. (2008c, 2009). In the reactionmechanism proposed byHiroyoshi et al., 2002for systems operated atlow solution potential, chalcopyrite wasreduced to chalcocite and silverions reacted with hydrogen sulde (overall reaction (12)) andchalcocite was oxidised by ferric ions more readily than chalcopyrite(reaction(13)).Hiroyoshi et al. (2007)reported that Ag+ caused thecritical potential (that potential below which chalcopyrite leachingoccurs faster) to rise.

CuFeS24Ag

2Ag2Son chalcopyrite surfaceCu2

Fe2 7

2Fe3 Ag2S2Ag

regenerated catalyst S0

2Fe2 8

CuFeS24Ag Cu2 Fe2 4Ag0 2S0 9

Ag Fe2 Ag0 Fe3 10

2AgCl Ag0 Cl2photochemical decomposition 11

2CuFeS2 4Ag Cu2S2Fe

22Ag2SS

012

Cu2S4Fe3

2Cu2

4Fe2

S0

:

13

In tests using ore particles in columns, simulating heap leaching, thebenet of silver catalysis was less clear. Silver addition resulted in asharp but transient increased copper extraction rate from chalcopyritecontained within in a complex ore (Ahonen and Tuovinen, 1990) butresulted in improved copper recoveries from low-grade chalcopyriteore in 300-day tests (Muoz et al., 2007b). In those process develop-ments where silver catalysis was shown to be effective, there remaineda concern about the economics of adding silver to full scale operations.Hu et al. (2002)suggested the use of silver-bearing concentrates asa cost-effective strategy. In the IBES process (Carranza et al., 1997;Palencia et al., 1998) and the BRISA process (Romero et al., 2003), theow sheets included a methodology to recover and recycle silver from

the leached residues.

The benets of pre-treating pyrite with silver and then using thetreated pyrite in the Galvanox process were described by Nazariet al. (2012). In that process, in the absence of silver-treated pyrite,the surface layer of sulfur that forms on chalcopyrite particles is non-conductive and limits electrical contact between the pyrite andchalcopyrite particles. However copper extraction in the Galvanoxprocess is enhanced when silver-treated pyrite is used. Nazari et al.(2012)proposed that the mechanism of silver-enhanced extraction

involved a small part of the added silver reacting with the sulfur layerto increase its conductivity and facilitate the transfer of electronsbetween chalcopyrite andpyrite particles.The pyrite(plus added silver)is recycled in the Galvanox process.

Bismuth has received less attention than silver as a copper leachingcatalyst.Hiroyoshi et al. (2007)ranked bismuth second, andEscuderoet al. (1993) ranked bismuth fourth most-effective catalyst of themetal ions tested. In contrast, Muoz et al. (2007a) reported thatbismuth did not enhance the leaching of copper from a sulde ore andGmez et al. (1999)reported only a small positive effect. Mier et al.(1994) suggested that the mechanism involved the reaction ofbismuth with the phosphate added as part of the bioleaching mediumand precipitation of bismuth phosphate, preventing ferric ions fromcomplexing with the phosphate and thus increasing the oxidisingpotential of the Fe(III)/Fe(II) couple.

3. Sulfuric acidalternative oxidants

The sulfuric acidferric sulfate system for the oxidative leaching ofchalcopyrite has been studied widely, and the redox chemistry, reactionproducts, and consequences of adopting certain controlled conditionsare well understood. The reagents are inexpensive and, with biologicalassistance, can be regenerated during processing. However, theoxidising potential of ferric ion is not particularly high and severalstronger oxidants have been tested for their ability to extract copperfrom chalcopyrite. Half reactions with standard reduction potentialsare summarised for selected oxidants (Table 3) and copper extractiondata compared inFig. 9. It should be noted that the cupric/cuprouscouple does not play a role in sulfate systems as concentrated as those

used in oxidative hydrometallurgy, as sulfate anions do not complexand thus stabilise the cuprous ion, which is rapidly oxidised inoxygenated or ferric ion systems.

3.1. Sulfuric aciddichromate

The dissolution of a chalcopyrite concentrate in medium containing0.2 M dichromate and 0.5 M sulfuric acid showed strong temperaturedependence in the range of 3080 C (Antonijevic et al., 1994). Inthose tests, the consumption of dichromate was greater than wouldbe expected if elementalsulfurwas the sole product of sulde oxidation(reaction (14)) indicating that some of the sulde was oxidised to

Table 3

Selected oxidants, half reactions and standard reduction potentials.Handbook of Chemistry and Physics, 89th edition, CRC Press, 2008.

Oxidant Half reaction Standard reductionpotential(E0) [V]

Cupric ion Cu2+ + eCu+ 0.153Ferric ion Fe3+ + e Fe2+ 0.771Dichromate Cr2O7

2+14H+ + 6e2Cr3+ + 7H2O 1.232Chlorate ClO3

+ 6H+ + 6e3H2O +C l 1.451

Permanganate MnO4+ 8H+ + 5eMn2+ + 4H2O 1.507

Hydrogenperoxide H2O2 + 2H+ + 2e2H2O 1.776

Ozone O3 + 2H+ + 2eO2 + H2O 2.076

Peroxodisulfate S2O82+ 2H+ + 2e 2HSO4

2.123S2O8

2+ initiator2SO4 2.6

http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-http://-/?-


11/18

sulfate (reaction (15)). The results were consistent with those of earlierstudies (Murr and Hiskey, 1981; Shantz and Morris, 1974). Theconcentration of dichromate did not affect the dissolution rate and it

was concluded that the chalcopyrite surface was covered by Cr(VI)(Antonijevic et al., 1994; Murr and Hiskey, 1981). Dichromate activityincreased with increasing acid concentration in the range 0.252.0 MH2SO4 but there was minimal additional benet when acid con-centrations N2M were used. The presence of chloride ions in the systemdecreased copper recovery, an effect that was attributed to thedisplacement of Cr(VI) ions with Cl ions on the chalcopyrite surface.Concentrations of ferric ions (up to 0.09 M) and copper ions (up to0.08 M) had negligible effects on reaction rates.

6CuFeS25K2Cr2O7 35H2SO4 6CuSO4 3Fe2SO43 5K2SO4

12S0 5Cr2SO43 35H2O 14

6CuFeS217K2Cr2O7 71H2SO4 6CuSO43Fe2SO4317K2SO4

17Cr2SO43 71H2O: 15

In the study byAydogan et al. (2006), dissolution of a chalcopyriteconcentrate in H2SO4 (0.10.5 M) and K2Cr2O7 (0.010.15 M) mediawas temperature dependent, the greatest extraction being achieved at97 C, the top temperature tested. At 70 C, about 70% copper wasextracted in 150 min in medium containing 0.4 M H2SO4 and 0.1 Mdichromate. While dissolution with dichromate was faster than withferric ions, it was nevertheless controlled by diffusion through a poroussulfur layer (reaction(14)).

3.2. Sulfuric acidchlorate

Chlorate ion is a strong oxidising agent (Table 3) that has been usedto enhance chalcopyrite leaching in both hydrochloric acid and sulfuricacid systems.Kariuki et al. (2009) studied chalcopyrite oxidation bymixing 2 g concentrate with up to 3 g sodium chlorate and 30 mLof 10 g L1 H2SO4 and heating in sealed Teon-lined vessels in thetemperature range of 45200 C. The results showed that chalcopyriteoxidation (reaction((16)) increased with increased temperature up toapproximately 165 C.

6CuFeS217NaClO33H2SO43Fe2SO43 6CuSO4

17NaCl3H2O:

16

Sodium chlorate concentrations up to 1 M were tested at 45 C withstirring in batch reactors (Xian et al., 2012). In these tests, the leachingrate accelerated with increased acid up to 1.5 M. At 0.5 M and 1 M HCl,

respectively, 45% and 65% Cu were extracted in 300min.

3.3. Sulfuric acidpermanganate

Permanganate ion has been used to maintain solution ORP in the

acid leaching of chalcopyrite and pyrite but not as the primary oxidant.Kametani and Aoki (1985)and Sandstrm et al. (2005)used potassiumpermanganate to control ORP during ferric sulfate chemical leaching ofnely-ground chalcopyrite concentrates at 90 and 65 C, respectively.From their results, Sandstrm et al. (2005) concluded that thepermanganate oxidised ferrous ionsto ferric ions(reaction (17)), ratherthan oxidising the chalcopyrite directly. They noted thata disadvantageof using potassium permanganate was the loss of ferric ions fromsolution via reaction with potassium and sulfate ions to form jarosite(reaction ((5)) resultingin lower copper extraction. Potassium perman-ganate was also a suitable oxidant for controlling ORP during the ferricsulfate leaching of pyrite, and the mass of potassium permanganatesolution required to maintain an ORP set point was a reliable indicatorof leaching progress (Bouffard et al., 2006).

5Fe2 MnO4

8H 5Fe3 Mn2 4H2O: 17

3.4. Sulfuric acidhydrogen peroxidehydroxyl radical

Hydrogen peroxide is a strong oxidising agent in acidic medium(Adebayo et al., 2003). It has been used as a leaching agent for uraniumores(Eary and Cathles, 1983) and various studies on theuse of peroxidewith zinclead concentrate, pyrite and sphalerite have also beenconducted (e.g., Antonijevic et al., 1994; Olubambi et al., 2006). Asuggested oxidation mechanism in dilute aqueous solutions ofhydrogen peroxide involved dissociation to form the hydroxyl radical(reaction (18), catalysed by ferrous ions) that oxidised the suldemoiety of minerals (reaction (19)) (Adebayo et al., 2003; Lin andLuong, 2004).

H2O2 Fe2HO OH Fe3 18

2HO 2S2 s 2S0

s H2O 0:5O2: 19

Both Adebayo et al. (2003) and Antonijevic et al. (2004) usedchalcopyrite concentrates in their studies, while Olubambi andPotgieter (2009)conducted an electrochemical study focused on themechanism of chalcopyrite oxidation with acidic peroxide solution.Olubambi and Potgieter (2009)noted that chalcopyrite leaching withsulfuric acid alone was slow (reaction(2)) but that, when hydrogen

peroxide was added to the system, copper extraction was rapid

0

10

20

30

40

50

60

70

80

90

0 10 20 30 40 50 60

Cu

(%e

xtrac

ted)

time (hours)

0.1 M Fe(III) in H2SO4; 80C

0.1 M Fe(III); 0.18 M H2SO4; 90C

1 M Fe(III); 0.25 M H2SO4; 90C

0.2 M Cr2O7; 0.5 M H2SO4; 70C

0.1 M Cr2O7; 0.4 M H2SO4; 70C

0.5 M ClO3; 1 M HCl; 65C

20% H2O2; 0.1 M H2SO4; 70C

2 M H2O2; 3 M H2SO4; 40C

3 M H2O2; 1 M H2SO4; 25C

2.5% O3; 0.5 M H2SO4; 22C

0.2 M S2O8; 0.01 M Fe(III); 0.1 M H2SO4; 23C

Fig. 9. Examplecopperextraction data forchalcopyrite concentrates of similar particle sizes: ferricsulfate systems (Ferreira andBurkin, 1976; Jones andPeters,1976;Munoz et al.,1979),dichromate (Antonijevic et al., 1994; Aydogan et al., 2006), sodium chlorate (Xian et al., 2012), hydrogen peroxide (Adebayo et al., 2003; Antonijevic et al., 2004; Olubambi et al., 2006),ozone (Havlik and Skrobian, 1990); peroxodisulfate (Dakubo et al., 2012).



12/18

and that the sulde was oxidised to elemental sulfur and sulfate(reactions(20) and (21)).

2CuFeS2 5H2O2 5H2SO4 2CuSO4 Fe2SO43 4S0

10H2O

20

2CuFeS2 17H2O2H2SO4 2CuSO4 Fe2SO43 18H2O: 21

Leaching rates were strongly dependent on temperature in a rangeof 3080 C; with maximum copper extractions of 5% to 60% after60 min. However, rates slowed after 6080 min at temperatures6080 C, attributed to increased peroxide at higher temperatures(Adebayo et al., 2003). These authors also reported that best resultswere obtained using 30% H2O2v/v with H2SO4and moderate agitation,and suggested that high peroxide concentrations and strong agitationduring leaching both caused accelerated hydrogen peroxide decom-position.Antonijevic et al. (2004)reported increased leaching rates inthe range 0.15 M H2O2in 2 M H2SO4at 40 C. In both studies, highersulfuric acid concentrations resulted in higher dissolution rates (testsconducted in the range 0.16 M H2SO4). More recently, Adebayo(2006)showed that copper extraction from a low-grade chalcopyriteore was faster when leached in a solution (pH 4.3) containing ammo-

nium sulfate (72 g L1

) with 45% v/v H2O2.The addition of ethylene glycol to batch reactors containing nelyground chalcopyrite in 1 M H2SO40.26 M H2O2 medium at 65 Chindered peroxide decomposition, resulting in improved copperextraction from 20% to 60% in 240min (Mahajan et al., 2007). Similarly,copper extraction of 65% was achieved in 60 min when ethylene glycolwasaddedtoa1.25MH2SO42.1MH2O2 solution butno further copperwas extracted up to 300 min (Solis-Marcal and Lapidus, 2013).

Olubambi et al. (2006) reported copper extraction of 74% in 180minfrom a ground, complex ore leached in 1 M H2O21 M H2SO4solutionand attributed the increased dissolution rate with high peroxideconcentration to the formation of Caro's acid (reaction((22)).

H2O2 H2SO4 H2SO5 H2O: 22

Olubambi et al. (2006) suggested that optimisation of a system with1 M H2SO4 and peroxide concentrations up to 1 M could yield aneconomical process because the reagent costs were low and the overallprocess would not be particularly corrosive in respect of materials ofconstruction. While the test was encouraging in that the peroxide wasactive when an ore was used, this ore was atypical, comprisingmassive chalcopyrite (66%) with minor pyrite (11%) and silica (7%). Inan investigation of the kinetics of peroxide decomposition in thepresence of ground pyrite crystals in hydrochloric acid,Chiri (2007)reported that peroxide decomposition was catalysed by both the ironsites on pyrite surfaces and by aqueous ferric ions. Increases in pyritesurface area, temperature, peroxide concentration and acid con-centration all resulted in increased peroxide decomposition rates. The

highest losses of peroxide by decomposition into molecular oxygenand water were observed at moderate temperature (45 C) and initialsolution pH 2. In the presence of ferric ion ligands, only the mechanismof decomposition catalysed by the iron sites on the pyrite surfacecontributed and thus overall decomposition was lower. The dissolutionof chalcopyrite involves the production of ferrous ions which wouldbe oxidised to ferric ions in leach solutions (reactions(20) and (21)).The additional ferric ions entering the solution would increase thedecomposition rate of hydrogen peroxide (Garten, 1962; Kolthoffet al., 1973), constituting a further reagent loss.

3.5. Sulfuric acidozone

Potential advantages of the use of ozone as an oxidising agent for

chalcopyrite are that it does not introduce additional ions into the

solution and it has a high oxidising potential (Table 3).Havlik andSkrobian (1990)studied the oxidation of a chalcopyrite concentrateusing an isothermal reactor containing 3 g sample and 1 L of 0.5 MH2SO4 into which ozone was introduced using a gas impeller. Themaximum ozone concentration was 3 vol.% in the gaseous streambeing sparged into the leachate. Copper extraction was efcient at20 C but decreased at higher temperatures, an effect attributed to thelower ozone solubility. Because no elemental sulfur was detected in

any experiment and the ratio of Fe:Cu was 1:1, Havlik and Skrobian(1990) concluded that the reaction, comprising several possibleintermediate steps, proceeded according to overall reaction (23).Using results obtained at different temperatures and visualisation ofleached surfaces, they proposed that the rate limiting step was thediffusion of ozone from the acidic medium to the chalcopyrite surface.They also noted the possibility of using powerful ozonisers, which areused to prepare drinking water,in futuremineral processingowsheets.In subsequent research,Havlik et al. (1999)examined ways of makingthe process more efcient, particularly in respect of the contact timesbetween gas bubbles and the mineral particles.

3CuFeS2 8O3 3CuSO43FeSO4: 23

Carillo-Pedroza et al. (2010)studied the use of ozone as a strongoxidant in an acidic ferric sulfate system applied to copper extractionfrom an ore. They conducted isothermal agitated tests in which thesolution was heated prior to ore addition (300 g in 1 L solution), andthe oxygen/ozone mixture was injected into the base of the reactor(O30.5 or 1 g h

1). The tests were carried out at 25 C, using a low-grade, groundore (0.7%Cu) split into different size fractions andleachedin solutions of sulfuric acid (0.10.5 M) and ferric ions (00.5 M). Notsurprisingly, best copper extractions were obtained using the smallestsize fraction and highest ferric and acid concentrations, with 0.5 g h1

O3inow. Overall, the results showed that ozone reduced the leachingtime and allowed the use of lower ferric ion and acid concentrationswhen compared with a typical ferric sulfate leach system. It is alsoworth noting that gangue minerals in the low grade ore apparently

did not consume ozone. However, the enhanced chalcopyrite leachingwas less effective with larger particle sizes, from which result it wasconcluded that the presence of a single oxidant was sufcient topromote copper extraction from samples with low chalcopyrite surfaceareas. In their subsequent, broader study,Carillo-Pedroza et al. (2012)summarised their results on the use of ozone as an alternative oxidantfor sulde ores as follows: Results show that the extraction of goldand silver is increased by at least 15%, with lower cyanide consumption;the extraction of copper increased by 16% and in less time; and in thecase of coal, sulfur is removed above 70%. They concluded that theoxidation using ozone was a promising clean technology for processingof sulde ores.

3.6. Sulfuric acidperoxodisulfate

The advantages of peroxodisulfate (S2O82) as an oxidising agent

were reported to be that it could be made from sulfuric acid using anelectrochemical cell (Canizares et al., 2005; Radimer and McCarthy,1979) at low energy cost ($0.20 kg1) (Kimizuka et al., 2001).Peroxodisulfate undergoes decomposition or hydrolysis, dependingupon the prevalent conditions (reactions(24) and (25)) but may alsoform reactive species and radicals (Table 3). In the leaching of chal-copyrite using peroxodisulfate, the main activator for the formation ofsulfate radicals is expected to be the ferrous ion released from themineral during dissolution (Dakubo et al., 2012). In contaminatedaquifer remediation using ferrous ion-activated peroxodisulfate, thedecomposition rate of peroxodisulfate was 1020% per week (Brown,

2003). Once the peroxodisulfate has been converted to sulfate ions,



13/18

it can be regenerated electrochemically, minimising sulfuric acidconsumption.

2Na2S2O8 2H2O O2 2H2SO42Na2SO4 24

Na2S2O8 H2ONaHSO4 NaHSO5: 25

Dakubo et al. (2012) leached chalcopyrite concentrate or ore instirred reactors or small columns at 23 C. When ground massivechalcopyrite samples were leached in columns with solutions (pH 2)containing sulfuric acid or peroxodisulfate (10 g L1) in sulfuric acid,copper extraction with peroxodisulfate was 500 times greater thanwithout (% extraction not shown but roughly calculated to beapproximately 10% extraction in 75 h). In tests using stirred batchreactors containing either 1.5 g of chalcopyrite concentrate or 20 g ofground, low-grade ore in 1 L H2SO4(pH2), dissolved oxygen enhancedcopper extraction but the pre-addition of ferrous ions to the reactorshad little effect. Up to 30% of the contained copper was extracted in90 h when the ground ore was leached with 10g L1 Na2S2O8at pH 2.Up to 50% Cu was recovered from concentrate at 45 C using 50g L1

Na2S2O8at pH 2. Particle size inuenced extraction; extractions were70% (25 m particle size) and 25% (125 m particle size), under similarleaching conditions.Dakubo et al. (2012)estimated a molar ratio of

S2O82 added to Cu extracted as 8.8:1, for a 96-hour test with 50 g L1Na2S2O8(i.e., not reagent limited).

4. Hybrid sulfatechloride systems

4.1. H2SO4NaClO2

The advantage of using sulfuric acid and sodium chloride to create apseudo-chloride lixiviant is that these reagents are cheaper than ferricchloride or cupric chloride. The addition of chloride to a sulfate systemradically changes the solution speciation, as both iron and copper ionscan form complex species with chloride ions which may assist leaching.The oxidant in this system is oxygen. There should be no need to addferric ions to the system initially; ferrous ions produced during

chalcopyrite leaching will be oxidised to ferric ions (reactions(2) and(3)) and will subsequently contribute to overall copper extraction.With careful selection of leaching conditions, most of the iron andsome sulfate can be precipitated as sodium jarosite (reaction((5)).

The addition of chloride to a sulfate system changes the speciation ofboth copper and iron through the formation of Cu- and Fe-complexchloride ions in concentrations reecting the overall solution compo-sition. However, there is as yet no consensus on which complex speciesare benecial to copper extraction under different conditions. Li et al.(2010) examined the extraction rates of copper from a chalcopyriteconcentrate for a suite of leach systems and concluded that Fe3+ activitywas a key parameter. In tests conducted at 75 C, they measured fasterkinetics for the NaCl (0.25M)H2SO4system at pH2 (97% Cu extractionin 170 h) than at pH 1 (58% Cu extraction in 170 h). The slower rate at

pH 1 was attributed to the increased solubility of iron-containingsecondary minerals due to the formation of FeCl2+ solution speciesand the consequent lower Fe3+ activities in the presence of chloride.Ruiz et al. (2011)reported that leaching of a nely-ground chalcopyriteconcentrate in sulfatechloride solutions was rapid, 90% of the copperbeing extracted in 180 min at 100 C. The presence of 0.5 M chlorideions (29 g L1 NaCl) enhanced the leaching rate signicantly but theaddition of 3 g L1 Fe3+ caused both an increase in the ORP and alarge decrease in the leaching rate.Muoz-Ribadeneira and Gomberg(1970, 1971) attributed the increased copper extraction fromchalcopyrite when as little as 0.1 M HCl was added to a sulfuric acidleach to the in situ generation of the oxidant CuCl4

2 from the mixedacid. Subramanian and Ferrajuolo (1976) leached a complex suldeore in sulfuric acid with O2over pressure (100 C) and reported that

thepresence of chloride ion(0.5

10gL1

) resulted in slightly enhanced

copper extraction from chalcopyrite, reduced the amount of sulfateformation and reduced the concentration of iron in solution. Similarly,Deng et al. (2001) reported that the addition of small amounts ofchloride to a sulfuric acidO2leach system enhanced copper extractionfrom the copper residue produced from the oxidative leaching of nickelmatte.

In 1983 the Broken Hill Associated Smelters patented a process forthe recovery of copper from a copperlead matte using oxygenated

acidic sulfate

chloride solutions (Sawyer and Shaw, 1983). AccordingtoLu et al. (2000), this was possibly the only low-pressure suldeleach process to have been commercialised. Lu et al. (2000)appliedthe copper matte process to the dissolution of nely-ground chal-copyrite concentrate in solutions of pH b 0.8 (0.8 M H2SO4) containing1 M NaCl in the temperature range of 6095 C. They achieved up to97% copper extraction in 9-hour tests. Based on their results showingthat chloride concentrations N 0.5 M did not enhance the leaching rateof chalcopyrite (consistent with the results ofPalmer et al., 1981),Luet al. (2000) concluded that it was important only that there weresufcient chloride ions present rather than an excess. This nding hascommercial implications for geographical areas lacking in freshwaterbut with a supply of brackish water or seawater; seawater containsapproximately0.5M chloride ions. From an examination of theresidues,Lu et al. (2000)concluded that the sulfur reaction product obtained inthe p

Documents

chalcopyrita.pdf