Lixiviacion Con Fe3+ - Efecto en Potencial Redox

8/2/2019 Lixiviacion Con Fe3+ - Efecto en Potencial Redox

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Leaching of chalcopyrite with ferric ion. Part II: Effect of redox potential

E.M. Crdoba a, J.A. Muoz b, M.L. Blzquez b, F. Gonzlez b, A. Ballester b,a Escuela de Ingeniera Metalrgica y Ciencia de los Materiales, Facultad de Ingenieras Fsico-Qumicas, Universidad Industrial de Santander, Bucaramanga, Colombiab Departamento de Ciencia de Materiales e Ingeniera Metalrgica, Facultad de Ciencias Qumicas, Universidad Complutense de Madrid, 28040 Madrid, Spain

A B S T R A C TA R T I C L E I N F O

Available online 2 May 2008

Keywords:

ChalcopyriteFerric leachingRedox potential

This paper reports the effect of redox potential (or Fe3+/Fe2+ ratio) on chalcopyrite leaching. The relationshipbetween redox potential and other variables (iron concentration and temperature) is also evaluated. Leachingtests were performed in stirred Erlenmeyer flasks with 0.5 g of pure chalcopyrite and 100 mL of a Fe3+/Fe2+

sulphate solution. The redox potential ranged between 300 and 600 mV Ag/AgCl for the solution at a pH 1.8,180 rpm, with temperatures at 35 C or 68 C. The results show that although ferric ion is responsible for theoxidation of chalcopyrite, ferrous ion has an important role in that it controls precipitation and nucleation of

jarosites, which ultimately causes passivation of this sulphide. Chalcopyrite dissolves through the formationof an intermediary product (covellite, CuS) that is later oxidized by ferric ion, releasing Cu 2+ ions.

2008 Elsevier B.V. All rights reserved.

1. Introduction

Chalcopyrite, the most abundant mineral in copper ore bodies, isalso the most recalcitrant to hydrometallurgical processes (Dutrizac,1978), and for that reason copper is mainly extracted by pyrome-tallurgy. However, the depletion of ore deposits and declining mineralgrades has encouraged the development of hydrometallurgicalprocesses for the treatment of chalcopyrite.

Thedissolution of chalcopyrite by ferric ion is normally depicted bythe following chemical reactions (Dutrizac and MacDonald, 1974):

CuFeS2 4Fe3YCu25Fe2 2S- 1

CuFeS2 4Fe3 3O2 2H2OYCu2 5Fe2 2H2SO4: 2

The release of cuprous ion does not occur in a single step. APourbaix diagram (Garrels and Christ, 1965) for the CuFeS2H2Osystem would show the stability fields of several sulphides such asCu5FeS4,CuSorCu2S between thestability fields of chalcopyriteand ofits dissolution.

After almost a century of research into the mechanisms ofchalcopyrite leaching, no unanimous theory has been proposed toaccount for its slow kinetics. Nevertheless, there is consensus as to theformation of a passivating layer on the chalcopyrite surface that slowsoxidation. The main theories were discussed in part I. Most of thempoint to the formation of a diffusion layer surrounding thechalcopyrite during dissolution, consisting of: bimetallic sulphide

(Burkin, 1969), copper polysulphide with a deficit of iron with respectto chalcopyrite (Ammou-Chokroum et al., 1977 and Hackl et al., 1995),or elemental sulphur (reaction (1)) (Muoz et al., 1979; Majima et al.,1985 and Dutrizac, 1989).

In addition, several authors have concluded that the chalcopyritedissolution rate depends on the redox potential of the solution, thebest results being achieved under moderately oxidizing conditions(Kametani and Aoki, 1985; Hiroyoshi et al., 2001; Okamoto et al.,2003). Hiroyoshi et al. have proposed a model reaction involving theintermediatereduction of chalcopyrite by ferrous ion to Cu2S and lateroxidation by ferric ion to release cupric ions:

CuFeS2 3Cu2 3Fe2Y2Cu2S 4Fe

3: 3

Nicol and Lzaro (2003) have proposed a different mechanism toexplain the chalcopyrite dissolution at low redox potentials.

Kametani and Aoki identified a critical potential around 0.45 V vs.SCE, associated with the onset of pyrite oxidation. A relatively lowredox potential has also been reported to have a favorable effectduring bioleaching (Barr et al., 1992). These authors observed lowredox potentials (400 to 450 mV SCE) after addition of ferrous ion,which sped up the bioleaching kinetics of a copper concentrate. Thirdet al. (2000, 2002) concluded that high ferric ion concentrations orhigh redox potentials inhibited the bioleaching of chalcopyrite, butthey did not determine the cause of passivation.

Temperature is another factor which directly affects chalcopyriteleaching rate. The high values of activation energy found by differentauthors: 71 kJ/mol (Dutrizac et al., 1969), 84 kJ/mol (Muoz et al.,1979), 88 kJ/mol (Hirato et al., 1987); clearly demonstrate the need ofhigh temperatures to break down bonds in the chalcopyrite crystal

lattice. In this study, the relationship between temperature and

Hydrometallurgy 93 (2008) 8896

Corresponding author. Tel.: +34 91 3944339; fax: +34 91 3944357.

E-mail address: [email protected] (A. Ballester).

0304-386X/$ see front matter 2008 Elsevier B.V. All rights reserved.doi:10.1016/j.hydromet.2008.04.016

Contents lists available at ScienceDirect

Hydrometallurgy

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leaching kinetics is evaluated at 35 and 68 C, given that thesetemperatures are appropriate in bioleaching processes with meso-philic and thermophilic bacteria respectively.

This research was prompted by the debate as to the relativeinfluence of ferrous ion and of a low redox potential on chalcopyritedissolution. The specific aim was to study the role play by ferric and

ferrous ions, which control the potential, during leaching.

2. Materials and methods

2.1. Solids

All tests were performed with a chalcopyrite mineral (approximately80% CuFeS2) from Messina, Transvaal (South Africa). Table 1 shows thechemical composition of this mineral. The main impurities were pyrite(FeS2), siderite (FeCO3) and silica (SiO2), as determined by X-ray analysis(Fig.1).

The mineral was dry-ground using a ball mill. The particle sizedistribution was determined by laser pulse. The average particle sizewas around 70 m. Additionally a BET surface area of 0.07 m2/g was

determined.

2.2. Leaching solutions

The different redox potentials of solutions were obtained by mixingferric and ferrous sulphates while keeping total iron concentrationconstant at 5 g/L (in some tests 0.5 g/L) and pH at 1.8. Stock solutions offerric sulphate were prepared with 0 K (modified 9 K medium ofSilvermanand Lundgren,1959) and Norris(Norris and Barr,1985) nutrientmediums forthe tests at 35 C and68 C respectively, which are normallyused during bioleaching at these temperatures.

2.3. Leaching tests

All leaching tests were performed in an orbital shaker at 180 rpm and35 or 68 C using 250 mL Erlenmeyer flasks covered with hydrophobiccotton so as to admit oxygen while reducing water loss throughevaporation. A low pulp density of 0.5% (100 mL of leaching solutionand 0.5 g of mineral) was chosen to avoid sharp changes in the redoxpotential of the liquid medium during the onset of leaching.

Periodically, water evaporation wasrestored, pH adjusted when abovetheinitial value,redoxpotentialrecordedand 1 mLsamplesremovedfromthe liquid to obtain kinetic information on metal dissolution. Copper andtotal iron concentration were determined by atomic absorption spectro-scopy and ferrous ion concentration using a photocolorimetric methodbased on the formation of a reddish colored complex of Fe2+with ortho-phenantroline which was analyzed in an UVvis spectrophotometer at awavelength of 510 nm. Finally, solid residues were characterized by XRDand SEM-EDS.

3. Results and discussion

3.1. Influence of redox potential

The influence of redox potential on the chalcopyrite dissolution ratewas testedat 35C and 68C.Theredox potentialvalues studiedwere 300,400, 500 and 600 mV vs. Ag/AgCl.

At 35 C the effect of the initial redox potential was negligible, withvery low copper extractions (b2.5 %) in all cases (Fig. 2). The lowreactivity of the chalcopyrite surface is consistent with low consump-tion of the oxidizing agent (Fe3+).

Theevolution of theredox potential with time shows that Fe3+

/Fe2+

solutions tend towards equilibrium. This equilibrium potential, in arange of 400500 mV vs. Ag/AgCl, is presumably related to thestandard potential of the Fe3+/Fe2+ couple. The value normally cited inthe literature is 771 mV vs. SHE or 564 mV vs. Ag/AgCl (Dry andBryson, 1988; Lowson, 1982); however, Cabral and Ignatiadis (2001)obtained experimentally a lower value (644 mV vs. SHE or 437 mV vs.Ag/AgCl) while assuming that the ratio of activity coefficients of bothion species, Fe3+ and Fe2+, was equal to unity. This standard potentialvalue is within the equilibrium potential range observed in thepresent study.

Table 1

Chemical composition of the mineral tested

Element Content (%)

Copper 27.36Iron 29.65

Sulphur 34.30Zinc 0.31Lead 0.02

Fig.1. X-ray diffractogram patter of the starting mineral.

Fig. 2. Influence of redox potential on the chalcopyrite leaching at 35 C and[Fe]Total=5 g/L.

89E.M. Crdoba et al. / Hydrometallurgy 93 (2008) 8896
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SEM examination of leaching residues did not show apparentlypassivating layers on the chalcopyrite surface.The micrographs (Fig. 3)show very clean surfaces with small amounts of a precipitate that isinsufficient to affect the reactivity of chalcopyrite. EDS microanalysisof these residues are shown in Table 2. At low potential (400 mV), thecomposition of the attacked chalcopyrite surface was practically

identical to the unattacked surface except for the presence of a smalloxidized precipitate. At high potential (600 mV), an oxidized jarosite-type compound was detected on the mineral surface, presumablypotassium jarosite.

X-ray diffractograms of theseresidues (Fig. 4) show thepresence ofseveral products: elemental sulphur (probably formed by reaction(1)), covellite (formed possibly by chalcopyrite transformation, as willbe discussed later), goethite (only at low potential) and potassiumjarosite. The las two compounds are presumably formed by ferric ionhydrolysis, according to the following reactions:

Fe2 SO4 34H2OY2FeOOH 3H2SO4 4

K 3Fe3 2SO2

4 6H2OYKFe3 SO4

2OH

66H: 5

At high temperature (68 C), unlike at 35 C, the effect of the initialredox potential was very pronounced (Fig. 5). At lower initial redoxpotentials (Einitial400 mV), the chalcopyrite dissolution rate was veryfast during the first 5 days of leaching, and 80 or 90 % Cu extractionwas achieved (Fig. 5a). This successful leaching coincided withpotential values lower than approximately 450 mV (Fig. 5b) andvalues of the Fe3+/Fe2+ ratio lower than unity (Fig. 5c). Fortimeslongerthan 5 days, the potential rose to 500 mV and the Fe3+/Fe2+ ratioincreased appreciably. Under these new conditions in the leachingmedium, chalcopyrite dissolution stopped.

These observations tend to support the hypothesis of Hiroyoshiet al. (1997, 1999, 2000, 2001), according to which chalcopyriteleaching with ferric sulphate is catalysed by ferrous ion. Furthermore,

in agreement with Hiroyoshi's results, there is a critical redoxpotential value of 413 mV vs. Ag/AgCl above which chalcopyritedissolution slows down.

As noted earlier, various researchers have also observed betterchalcopyrite dissolution rates at low redox potentials at 30 C(Okamoto et al., 2003) and at 90 C (Kametani and Aoki, 1985).Kametani and Aoki determined a critical potential of 458 mV vs. Ag/AgCl, and detected by XRD the presence of CuS in the leachingresidues, working at a very low potential (338 mV vs. Ag/AgCl).

SEM micrographs of the leaching residue at 68 C and an initialredox potential of 400 mV (Fig. 6a) show chalcopyrite particlessurrounded by a precipitate identified by XRD as a mixture ofpotassium jarosite, elemental sulphur and goethite (Fig. 7a). Althoughthis compositionwas similar to the leaching tests at 35 C, the amountof precipitate was greaterat 68 C. EDS microanalyses (Table 2) of thatleaching residue indicate an enrichment in potassium jarosite.

Under more oxidizing conditions (Einitial500 mV), copper extrac-tions were lower than 35 C (Fig. 5a). The low-magnificationbackscattered electron micrograph of the residue at 600 mV(Fig. 6b), shows few bright areas corresponding to clean chalcopyritesurfaces and many dark areas corresponding to chalcopyrite covered

Fig. 3. SEM micrographs of the leaching residues at 35 C and [Fe]Total=5 g/L:a) Einitial=400 mV and b) Einitial=600 mV.

Table 2

EDS microanalysis of chalcopyrite leaching residues (results in weight percent)

Test Chalcopyrite surface S Fe Cu O K

Original 30.80 32.24 36.9635 C Attacked 30.50 32.96 36.54400 mV Precipitate 33.48 27.46 29.75 9.3135 C Attacked 34.17 29.18 31.49 5.16600 mV Precipitate 23.25 23.08 18.19 34.18 1.3068 C Attacked 30.81 27.44 29.14 12.32 0.29400 mV Precipitate 18.89 31.35 1.05 46.33 2.3868 C Attacked 33.30 27.34 29.31 9.84 0.21600 mV Precipitate 14.49 34.10 0.72 46.46 4.23

Fig. 4. X-ray diffractograms of the leaching residues at 35 C and [Fe] Total=5 g/L:a) Einitial=400 mV and b) Einitial=600 mV.

90 E.M. Crdoba et al. / Hydrometallurgy 93 (2008) 8896
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by a diffusion layer. That film formed on chalcopyrite particles isclearly observed at higher magnification in Fig. 6c. the low porosity ofthat layer would explain the poor results obtained at high redox

potentials. XRD analysis (Fig. 7b) and EDS microanalysis (Table 2)point to that thediffusion layer consisted mainly of potassium jarosite.

Those differences were also related to a higher iron precipitation,during the initial stages, at high than at low potentials (Fig. 5d).Assuming that nucleation of jarosites on mineral particles (hetero-geneous nucleation) was more important than nucleation out ofsolution (homogeneous nucleation, that fast iron precipitation at highpotentials caused the rapid passivation of chalcopyrite.

At low potentials, the fast leaching kinetics initial stage wasfollowed by another very slow (Fig. 5a) and related to an abundantironprecipitation(Fig. 5d). That would explain the presenceof jarosite

Fig. 5. Influence of redox potential on the chalcopyrite leaching at 68 C and [Fe]Total=5 g/L.

Fig.6. SEMmicrographsof theleachingresidues at68 Cand [Fe]Total=5g/L:a) Einitial=400mV,b) Einitial=600 mV and c) circled particle in (b) at higher magnification.

Fig. 7. X-ray diffractograms of the leaching residues at 68 C and [Fe]Total=5 g/L:a) Einitial=400 mV and b) Einitial=600 mV.

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and goethite in the leaching residues at low potential (Fig. 6a), like inexperiments at high potentials (Fig. 6b). X-ray diffractograms of theresidues also show the presence of sulphur, although in a smallerproportion than jarosite (Fig. 7).

The passivating nature of jarosites is in good agreement withthermodynamic data. For thedissolution of jarosite at 25 C, Baronand

Palmer (1996) determined:

KFe3 SO4 2 OH 66HYK

3Fe3 2SO24 6H2O 6

the solubility product constant, Ksp=1011, and the free energy of

jarosite formation,Gf, 3309.8 kJ/mol. These values areindicative ofthe very low solubility and high stability of jarosites.

Casas et al. (2000) noted that jarosite formation in ferric sulphatesolutions starts from iron hydroxides, like Fe(OH)3, formed byhydrolysis of ferric ion. The presence of goethite (FeOOH) in someresidues in the present study tends to confirm that hypothesis. Theaction of goethite as activating agent in the formation of jarosite maybe depicted as follows:

2 Fe3 K 2SO24 FeOOH s

4H2OYKFe3 SO4 2 OH 63H: 7

These results point to that passivation of chalcopyrite duringleaching is due to the formation of a layer of jarosite which preventstransportation of both electrons and ion species between the mineralsurface and the leaching medium.

Furthermore, the precipitation of iron, and hence the formation ofjarosites, is directly related to the redox potential of the solution. Thus,redox potentials higher than the critical value (between 400 and500 mV) favor Fe3+ precipitation as jarosite and the subsequentchalcopyrite passivation.

The above observation is supported by Bigham et al. (1996) in thatthe stability field of potassium jarosite at pH lower than 2 is located atpotentials higher than 563 mV vs. Ag/AgCl. Thus, the closer the redoxpotential of the leaching solution is to the beginning of the stabilityfield of the jarosite, the faster it will precipitate.

Asintestsat35C(Fig. 2), the kinetics curvesat 68 C (Fig. 5) showthat the Fe3+/Fe2+ ratio or the redox potential of the leaching solutiontends asymptotically towards an equilibrium value of approximately480 mV vs.Ag/AgCl. Only a fewresearchers have mentioned this trendfor leaching solutions of ferric sulphate. Of these, Dutrizac (1983)noted that whereas the amount of jarosite formed increased withferric ion concentration, the Feprecipitated/Fedissolved ratio did not,concluding that the solution tends toward equilibrium. Precipitationof iron from ferric sulphate solutions, then, is difficult to control sinceit occurs spontaneously until the system reaches equilibrium.

The leaching results at 68 C seem to demonstrate that precipita-tion and nucleation of jarosites on mineral particles would occur even

in tests with the fastest kinetics. Intermediate products were notdetected in the long-term tests, probably because of a jarosite filmcovering the particles. In order to generate residues to identifypossible intermediate products, new leaching tests were performed atshorter times (1 h, 5 h and 1 day).

Figs. 8 and 9 show micrographs and X-ray diffractograms ofleaching residues at 68 C and Einitial=400 mV with short attack-time.Chalcopyrite surface transformations began at times as short as 1 h(Fig. 8a). EDS microanalysis of precipitate shown in that figure(named precipitate in Table 3) indicates an impoverishment of ironand an enrichment of copper and sulphur. The XRD diffractogram ofthe residue (Fig. 9a) evidences the formation of three products:elemental sulphur, goethite and covellite. The last compound (CuS)would be related to the EDS analysis of zone named precipitate in

Fig. 8a.Micrographs of the 5 h leaching residue (Fig. 8b) show small

quantities of a precipitate over particles. EDS microanalyses (Table 3)

show again an enrichment of copper and sulphur on the chalcopyritesurface and the presence of an oxidized compound in the zone namedprecipitate, probably goethite. The diffractogram of that residue(Fig. 9b) reveals the presence of the same three products found at 1 h:elemental sulphur, goethite and covellite.

Finally, the growth of a product over the chalcopyrite surface

seems to start after 1 day of attack (Fig. 8c). EDS microanalyses(Table 3) of that product indicate the presence of oxidized compoundscontaining potassium and phosphorus in its composition. Twoadditional products were detectedby XRD (Fig. 9c): potassiumjarositeand ferric hydroxyphosphate.

The fact that goethite formed prior to the formation of jarositeconfirms the hypothesis that jarosite formation starts from Fe3+

hydro-oxidized compounds like goethite.Even although the proportion of ferric hydroxyphosphate detected

in the residue after 1 day of leaching is small, like jarosite it can bepassivating. In fact, its chemical composition (Fe4(PO4)3(OH)3) may berelated to the alunite group (AB3(XO4)2(OH)6), as in the case of jarosite(Lowson,1982), where A stands for cations such as Na+, K+, H3O

+, NH4+,

Pb2+ or Ag+; B stands for Fe3+ or Al3+, and XO4 usually for SO4, PO4 or

AsO4.Another noteworthy finding was an increase of the CuS/chalcopyr-ite ratio, at least during the first day of attack, as shown in thediffractogram.

Fig. 8. SEMmicrographsof theleachingresidues at68 C,[Fe]Total=5 g/L, Einitial=400mVand short times: a) 1 h, b) 5 h and c) 1 day.

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The presence of CuS in the leaching products at short timessuggests that chalcopyrite dissolves in two steps. First it oxidizes,forming CuS as an intermediary product:

CuFeS2 2Fe3YCuS 3Fe2 S-: 8

Then, covellite is oxidized by ferric sulphate, releasing Cu2+ ions:

CuS 2Fe3YCu2 2Fe2 S- 9

Also, we cannot rule out an initial reduction of chalcopyrite(reaction (3)), as proposed by Hiroyoshi et al. (2001), followed byoxidation of chalcocite to covellite by ferric sulphate through theformation of copper-deficient intermediate products, as reported byFerron (2003)in a study on theleachingof secondary copperminerals:

Cu2S Fe2 SO4 3YCuSO4 2FeSO4 CuS 10

Cu2Schalcocite

YCu1;931;96 Sdjujerite

YCu1;80Sdigenite

Y CuScovellite

: 11

Some researchers have detected the formation of covellite fromchalcopyrite under reducing conditions. Jang and Wadsworth, 1993,

concluded that covellite forms either aerobically or anaerobicallythrough the reaction:

CuFeS2 Cu2Y2CuS Fe2: 12

Moreover, this transformation speeds up with temperature.Kametani and Aoki, 1985, also found CuS in chalcopyrite leaching

residues carried out at a redox potential lower than 0.33 V vs. SCE.All this suggests, then, that CuS only forms at low redox potentials.Additionally, CuS formation can also occur during chalcopyrite

leaching by direct reaction between cupric ions and previously-formed elemental sulphur, as follows:

3Cu2 4H2O 4S-Y3CuS HSO4 7H: 13

Nevertheless, theauthors favor the idea that although this reactionis thermodynamically possible, it does not occur because of thehydrophobic nature of the elemental sulphur (Dutrizac and MacDo-nald, 1974).

Therefore, chemical reactions (8) and (12) become kinetically morefavorable.

3.2. Influence of iron concentration

The results reported in the previous section showed thatprecipitation and nucleation of jarosites on mineral particles led topassivation of chalcopyrite. In a first attempt to explain and controlpassivation, it was decided to study the effect of total ironconcentration.

Fig. 10 depicts the kinetics curves obtained during the leaching ofchalcopyrite at 68 C at two different iron concentrations: 0.5 and 5 gFeTotal/L. Total iron concentration plays an important role in theprocess: Reducing total iron concentration from 5 to 0.5 g Fe/L has anegative effect on copper dissolution at low redox potential (400 mV)and practically has no effect at high redox potential (600 mV).

The SEM study of the leaching residue at high potential and lowiron concentration showed chalcopyrite surfaces free of passivatinglayers (Fig. 11). The X-ray diffractogram of this residue (Fig. 12)confirmed the same compounds previously detected with 5 g/L of Fe(Fig. 7b): S, jarosite, goethite and ferric hydroxyphosphate, the lasttwo being more important in this case.

These results indicate that in the Fe3+/ Fe2+ couple the ion closelyrelated to chalcopyrite dissolution is Fe3+. The roleof Fe2+, then, wouldbe to achieve rapid equilibrium of the Fe3+/Fe2+ couple in solution,thus controlling hydrolysis of the ferric ion, which could be ultimatelyresponsible for the passivation of chalcopyrite.

Thepositive effectobserved in chalcopyrite leaching at 68 C wheniron concentration was increased from 0.5 to 5 g/L (or from 0.009 to0.09 M) suggests that the process is at least partially controlled by thediffusion of ions towards the chalcopyrite surface. That is consistentwith the findings of Hirato et al. (1987) that the dissolution rate at70 C increases with ferric ion concentration up to 0.1 M, while abovethat value the enhancement is negligible.

Table 3

EDS microanalysis of chalcopyrite leaching residues at 68 C, [Fe]Total=5 g/L,Einitial=400 mV and short times (results in weight percent)

Time Chalcopyrite surface S Fe Cu O K P

Original 30.80 32.24 36.961 h Attacked 30.16 31.99 36.33 1.52

Precipitate 34.06 8.87 48.76 8.315 h Attacked 37.32 29.64 33.04

Precipitate 28.99 28.76 30.38 11.871 day Attacked 32.84 29.74 33.05 4.37

Precipitate 12.39 23.91 4.45 54.68 3.92 0.65

Fig. 9. X-ray diffractograms of the leaching residues at 68 C, [Fe]Total=5 g/L,Einitial=400 mV and short times: a) 1 h, b) 5 h and c) 1 day.

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3.3. Influence of temperature

The effect of temperature was very pronounced, with negligiblecopper dissolution at 35 C after 13 days of leaching (b3%) vs. almostcomplete mineral dissolution at 68 C (Figs. 2 and 5). Thermalactivation therefore plays an important role in this process.

Theeffect of temperature in theleachingof chalcopyrite with ferricsulphate was evaluated by calculating the activation energy in theexperimental temperature range of 35 C to 68 C. Four assays wereperformed at 35, 46, 57 and 68 C, while all other experimentalconditions were constant, i.e.: 0.5% pulp density, Einitial=400 mV and[Fe]Total=5 g/L.

The activation energy was calculated from the kinetic constantsand the semilogarithmic plot of the Arrhenius equation.

First, the fraction of reacted chalcopyrite was plotted vs. time foreach temperature (Fig.13). Then, linear curves were plotted using thesimplified shrinking core model proposed by Sohn and Wadsworth(1979), according to the following expression:

1 1 X 1=3 k t 14

where Xis the fraction of reacted chalcopyrite, k the kinetic constantand t time. That equation is represented graphically in Fig. 14. Theslope of the final experimental data for the kinetic curves at 57 C and

68 C was zero (Fig.13), probably due to the precipitation of jarosites,and so they were removed from the linear regression.

According to Eq. (14), the slopes of the curves in Fig.14 correspond

to the following kinetic constant values: 0.0006, 0.0044, 0.0319 and0.0732 days1 for 35, 46, 57 and 68 C respectively.

Most chemical reactions obey the Arrhenius equation (Logan,2000; Levenspiel, 2004) and it was used to determine the effect oftemperature on the leaching of chalcopyrite:

k A eEa=RT 15

where Ea is the activation energy, A the pre-exponential factor withthe same units as the kinetic constant k, R the universal gas constant(8.314 J.K1.mol1) and T the absolute temperature (K). Then, thelogarithm of the kinetic constant is a linear function of the inverse oftemperature. Table 4 shows the values represented in the Arrheniusplot in Fig. 15.

The activation energy for chalcopyrite dissolution, deduced from

the slope of the straight line of the Arrhenius plot, was appreciablyhigher (130.7 kJ/mol) than that reported by other researchers (7188 kJ/mol) in the range of temperature between 50 and 94 C and insulphate medium (Dutrizac et al. (1969), Muoz et al. (1979), Hiratoet al. (1987). Those differences could be attributed to the temperaturerange assayed. The kinetic curves showthat chalcopyrite dissolution isnegligible at 35 C, andthat theenergy barrier responsiblefor theslowcopper dissolution rate could be overcome at 68 C. However, theenhancement of copper extraction is less pronounced above a certaintemperature (N50 C).

Based on the parameters established by Moore (1990) to elucidatereaction mechanisms in different chemical processes, the highactivation energy value registered in this study (130.7 kJ/mol)indicates that during the first stage of chalcopyrite leaching, when

the kinetics is approximately linear, the system is under chemicalcontrol. Therefore, the main drawback of chalcopyrite dissolution

Fig. 12. X-ray diffractogram of the leaching residue at 68 C, Einitial=600 mV and[Fe]Total=0.5 g/L.

Fig.13. Influence of temperature on the chalcopyrite leaching at Einitial=400 mV.Fig. 11. SEM micrograph of the leaching residue at 68 C, Einitial=600 mV and[Fe]Total=0.5 g/L.

Fig.10. Influence of iron concentration on the chalcopyrite leaching at 68 C.

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would be the strong chemical bonding of the crystal lattice, whichrequires a high input of energy to the system. These findings supportthe hypothesis of Hiskey (1993) that the transport of electronsthrough vacants is very poor due to the n-type semiconductivity ofchalcopyrite, and hence the first step during chalcopyrite oxidation isthe consumption of vacants to favor electron transportation throughthe crystal lattice. Thus, a great deal of the energy applied to thesystem in the form of heat is consumed by the displacement of ionsfrom thebulk of theparticleto thesurface, eliminating surface vacantsand favoring the transport of electrons through the chalcopyritesurface.

At the same time, increasing iron concentration from 0.5 to 5 g/Limproves the leaching rate at 68 C (Fig. 10), and so the diffusion offerric ion to the chalcopyrite surface may also control the process.

Researchers have not reached a consensus on the control stepsgoverning chalcopyrite leaching. Our findings are similar to thosereported by Linge (1976) and Hirato et al. (1987) with a linear branchin the copper kinetic curve. Nevertheless, many researchers (Dutrizacet al., 1969; Ferreira and Burkin,1975; Muoz et al., 1979; Parker et al.,

1981; Kametani and Aoki, 1985; Hackl et al., 1995, etc.) have obtainedparabolic kinetics with an initial linear branch, similar to our resultsbut with very varied conclusions. This variety in the hypotheses to befound in the literature reflects a lack of agreement on the mechanismsof chalcopyrite dissolution.

These results seem to indicate that hydrolysis and precipitation ofFe3+ has a main role in chalcopyrite passivation in ferric solutions bypreventing the contact between the mineral surface ant the oxidizingagent in solution.

Decreasing pH or removing some ions that favor iron precipitationcould be a way to prevent ferric ion hydrolysis. However, our attemptsin that direction (results not shown) have not solved the problem ofchalcopyrite passivation. Copper dissolution rate decreased when pHwas reduced from 0.5 to 2.0, perhaps because of that the species

responsible for the oxidation of chalcopyrite is not properly Fe3+ butprobably Fe(SO4)2

(Crdoba, 2005). Furthermore, the removal ofmonovalent cations (K+, Na+ or NH4

+) from solution neither preventshydrolysis and precipitation of iron as goethite and hydroniumjarosite that, like potassium jarosite, also tend to nucleate overchalcopyrite particles.

On the other hand, minority iron-bearing minerals in the startingmineral (siderite and pyrite) could have an important role in leachingkinetics. Siderite is quickly dissolved in acidic media, increasing theiron concentration in solution:

FeCO3 H2SO4YFeSO4 CO2 H2O: 16

In the case of pyrite, there are evidences that this sulphide onlydissolves when chalcopyrite is already passivated. Moreover, pyriterest potential is higher than that of chalcopyrite and, therefore, pyriteenhances dissolution chalcopyrite through galvanic contact.

Finally, the comparison between studies on chalcopyrite leachingis a difficult task because of differences in experimental conditionsused. However, unlike at high temperature, chalcopyrite dissolutionand ferric ion hydrolysis kinetics are very slow processes at lowtemperature (35 C).

4. Conclusions

1. The redox potential is a key factor in the leaching of chalcopyrite. A

high potential at theonsetof leaching provokes rapid passivation ofchalcopyrite.2. Ferric/ferrous sulphate leachingsolutions tendto reach equilibrium

when the activities of the two ions are equal, which is associatedwith a critical potential of approximately 450 mV. When the redoxpotential is very high initially, that tendency to equilibrium favorsrapid precipitation of ferric ion as jarosite and consequentlypassivation of chalcopyrite.

3. The activation energy during chalcopyrite leaching was 130.7 kJ/mol, which is a clear demonstration of the importance of thermalactivation in this process.

4. Increasing the iron concentration from 0.5 to 5 g/L had a positiveeffect in the chalcopyrite leaching at 68 C.

5. Chalcopyrite dissolves through the intermediate formation ofcovellite, CuS, which is later oxidized by ferric ion to release Cu2+

ions:

CuFeS2 2Fe3YCuS 3Fe2 S-

CuS 2Fe3YCu2 2Fe2 S-:

6. The elemental sulphur that forms during chalcopyrite leaching isporous and does not form a passivating layer on the chalcopyritesurface.

References

Ammou-Chokroum, M., Cambazoglu, M., Steinmez, D., 1977. Oxydation menage de lachalcopyrite en solution acide: analyses cintique de ractions. II. Modlesdiffusionales. Bulletin de la Societe francaise de mineralogie et decristallographie 100, 161177.

Fig. 15. Arrhenius plot.

Table 4

Values represented in the Arrhenius plot of Fig. 15

T (K) k (days1) k (s1) 1000 / T (K1) ln k (s1)

308 0.0006 6.9444109 3.25 18.79

319 0.0044 5.0926108 3.13 16.79330 0.0319 3.6921107 3.03 14.81341 00732 8.4722107 2.93 13.98

Fig.14. Variation of 1 (1X)1/3 over time.

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