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Acid leaching and dissolution of major sulphide ore minerals: processes and galvanic effects in complex systems

P. K. Abraitis^1,, R. A. D. Pattrick¹, G. H. Kelsall² and D. J. Vaughan^1,*

¹ Williamson Research Centre for Molecular Environmental Science, and Earth Sciences Department, University of Manchester, Manchester M13 9PL, UK
² Department of Chemical Engineering, Imperial College London, London SW7 2AZ, UK

^* E-mail: david.vaughan{at}man.ac.uk

	ABSTRACT

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
The kinetics and mechanisms of dissolution of the major base metal sulphide minerals, pyrite, chalcopyrite, galena and sphalerite in acidic (chloride) media have been investigated. Minerals were ground in air, then dissolved in air-equilibrated solutions at pH 2.5, while monitoring the redox potential. Solution samples were analysed by ICP-AES and HPLC, and surfaces of residual sulphides analysed using XPS. Dissolution of aerial oxidation products on pyrite particles in the first 15 min apparently led to a sulphur-rich surface, and was followed by slower dissolution of pyrite itself, driven by oxygen reduction, and resulting in net production of protons. Chalcopyrite dissolution resulted in a Cu, S-rich (near) surface layer, accompanied by net consumption of protons. Apparently incongruent dissolution of galena and sphalerite may reflect the formation of elemental S at the surface. The rates of dissolution of chalcopyrite, galena and sphalerite in the presence of pyrite were determined, respectively, as 18, 31 and 1.5 times more rapid than in single-mineral experiments. These data were consistent with galvanically-promoted mineral oxidation of the other sulphides in the presence of pyrite. In the case of galena, the experimental data suggested extensive release of Pb ions and development of a sulphur-rich surface during galvanically-promoted dissolution.

KEYWORDS: acid leaching, sulphides, pyrite, chalcopyrite, galena, sphalerite

	Introduction

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
THE sulphide ore minerals sphalerite, galena and chalcopyrite provide the major sources of the world’s base metals (Zn, Pb, Cu), whereas pyrite is virtually ubiquitous as a metalliferous mineral in sulphide ore deposits (Craig and Vaughan, 1994). Acid leaching of these minerals occurs in nature in the context of acid mine drainage (AMD) and acid rock drainage (Keith and Vaughan, 2000) and is of great importance in metal extraction using hydrometallurgical methods, including dump leaching of low-grade ores.

The present study was concerned with the overall mechanisms of dissolution of (air-ground) single-phase pyrite, chalcopyrite, galena and sphalerite particles in chloride media of pH 2.5 and with the complexities that may arise in systems involving two phases: pyrite + chalcopyrite, pyrite + galena, and pyrite + sphalerite.

	Experimental methods

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
Materials
Mineral powders were prepared by dry grinding of selected, hand-picked specimen-grade mineral samples. Grinding was accomplished using a tungsten carbide ball mill. The resultant mineral powders were size fractionated by dry sieving and the –45 + 150 µm powder size fraction was used in each experiment. Grinding and sieving were performed in air and the resultant powders were cleaned by ultrasonic washing in acetone and dried at room temperature. The identity and purity of the mineral powders was confirmed by X-ray diffraction (XRD).

Acid-leaching experiments
Experiments utilized a Mettler Toledo DL-67 auto-titration instrument, fitted with a Mettler Toledo DG 111-SC pH Electrode. The continuously stirred solution within the reaction vessel was maintained at 25±2°C. Prior to the addition of the mineral powder(s), a pre-titration was performed in which 0.1 M HCl (BDH, Convol) was added to 50 cm³ of air-equilibrated, high-purity water, resulting in a final solution pH of 2.5.

At the start of the experiments, 0.3 g of the sulphide powder was added to the solution. Leaching experiments were conducted for 3 h. During this time the pH of the resultant mineral pulp and the volume of 0.1 M HCl required to maintain a constant pH were recorded continuously. The solution redox potential (mV, SHE) was monitored using a Mettler Toledo InLab 501 combination Pt/Ag-AgCl redox electrode after 0, 30, 45, 60, 90, 120 and 180 min. 1 cm³ solution samples of the mineral pulp were removed at selected time intervals, filtered through 0.22 µm cellulose acetate membrane filters and acidified prior to analysis by inductively coupled plasma atomic emission spectroscopy (ICP-AES). A sample of the unacidified leachate was also analysed for dissolved sulphur species by anion chromatography. At the end of the experiment, the mineral powder was separated by filtration, dried at room temperature and stored under nitrogen prior to analysis by XPS.

Solution analysis
The concentrations of Fe, Cu, Pb, Zn and S were determined by ICP-AES using a VG Elemental Horizon instrument and conventional solution nebulization. Prior to ICP analysis, the solution samples were further acidified to 2% HNO₃. The instrument was calibrated using commercial multi-element standards.

Samples of the final leachates from each experiment were also analysed by high performance liquid chromatography (HPLC) which was conducted using a Dionex 4000i gradient ion chromatograph. Sulphur anion separation was accomplished using a combination of Dionex AG11 guard column (75 x 4 mm) and Dionex AS11 analytical column (250 x 4 mm) with sodium hydroxide eluant.

X-ray photoelectron spectroscopy (XPS) analysis
A Kratos XSAM 800 surface analysis instrument, employing a Mg-K X-ray source, was used to obtain XPS spectra of mineral fracture and powder surfaces. This instrument was calibrated using the Au 4f_7/2 83.98 eV, Cu 2p_3/2 932.67 eV and Ag 3d_5/2 368.27 eV peaks. All peaks were corrected for static charging effects using the 1s electron binding energy of carbon, which was taken as 284.6 eV. Due to instrumental limitations, no measures were taken to reduce potential evaporation of elemental sulphur or other volatile sulphur species within the vacuum chamber.

Fitting of the XPS data was accomplished using XPSPEAK41 software. Unless otherwise stated, all peaks have been assigned 80% Gaussian, 20% Lorentzian character. During fitting, all S(2p) components were assigned as doublets with an intensity ratio of 2:1 and a spin orbit splitting of 1.2 eV. Elemental atomic ratios based on the XPS data have been calculated using Kratos sensitivity factors.

	Results and discussion

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
Prior to mineral addition, the redox potential of the acid solutions was 300±10 mV (SHE). On addition of the mineral powder(s), this value changed rapidly, reaching the values shown in Fig. 1 at t_o. In most cases, the redox potential then decreased by a few tens to a hundred or more mV to reach a stable value, except for the single-phase sphalerite and the galena + pyrite mixture, for which the redox potential rose continuously throughout the experiment. The ranges of redox potentials at pH 2.5 recorded for each single mineral sample are plotted onto the potential-pH diagram for the S-H₂O system in Fig. 2.

View larger version (27K): [in this window] [in a new window]   FIG. 1. Time dependence of redox potential values in pH 2.5±0.1 pH-stat experiments at 25±2°C with sulphide mineral powders in air-saturated water.

View larger version (25K): [in this window] [in a new window]   FIG. 2. Potential-pH diagram for S-H2O system 298 K (ST = 10–4 M). Symbols indicate measured redox potential values in pH-stat single-sulphide mineral dissolution tests at 25°C, pH 2.5.

To these data on the redox potential changes and generation or consumption of protons can be added information on changing solution chemistry with time and surface analyses of the final mineral products for each of the systems studied.

Pyrite
In the case of pyrite, no acid addition was required during the experiment to maintain the solution pH at 2.5±0.1, because oxidation of the disulphide by dissolved oxygen at the experimental redox potential (36±35 mV) involves consumption of protons by oxygen reduction. That reaction was assumed to occur by a four-electron reduction (Biegler et al., 1970):

(1)

rather than forming H₂O₂ as an intermediate by a two-electron reduction. Reaction 1 was coupled to pyrite oxidation, assumed to produce S(VI), the thermodynamically most stable sulphide oxidation product under the redox potential conditions applied:

(2)

Hence, the overall process would involve the net generation of protons by the reaction:

(3)

However, if sulphoxy species with oxidation state +V (S₂O^2–₆) form, then no excess protons would be formed, whereas formation of S₄O^2–₆ ions would consume one mole of proton per mole of FeS₂. If elemental sulphur were the predominant sulphide oxidation product, then the reaction would consume two moles of protons per mole of FeS₂.

Figure 3 shows the time dependences of Fe and S concentrations dissolved oxidatively from pyrite. Linear regression analysis of these data indicate that the total dissolved sulphur concentration increases at twice the rate of Fe(II), consistent with stoichiometric dissolution of pyrite, according e.g. to reaction 3. However, the ratio of dissolved S to Fe concentrations increases from 1.38 to 1.5. The finite intercept given by extrapolating the regression line back to zero time, reflects the aerial oxidation of the pyrite, especially during grinding in air:

View larger version (13K): [in this window] [in a new window]   FIG. 3. Time dependences of dissolved Fe and S in a pH 2.5 (stat) experiment with powdered pyrite at 25°C.

(4)

Hence, dissolution during the initial 15 min was predominantly of the aerial oxidation products rather than the bulk pyrite, which exhibited slower, congruous dissolution thereafter. It is possible that a S-rich surface develops during the initial stages of dissolution, or that the oxidized powder surface is S-rich (relative to the S:Fe ratio in bulk pyrite).

The XPS data for pyrite are shown in Fig. 4. In each case, the spectrum for a freshly cleaved sample is shown at the bottom, the sample ground (in air) in the middle, and the final reaction product sample spectrum at the top. The dominant peak at 707.3±0.1 eV in the Fe 2p3/2 spectrum of the fracture sample is consistent with binding energies in pyrite. However, an additional peak is required at 709.6±0.1 eV to adequately fit the data. Nesbitt and Muir (1994) and Karthe et al. (1993) have attributed such a feature to surface defects. However, some hydroxyl oxygen was also identified at the pyrite fracture surface, so this feature could be attributable to ferrous hydroxide species, although the presence of surface defects cannot be ruled out (Bonnissel-Gissinger et al., 1998). The spectrum of the oxidized powder sample requires addition of further components at 711.6 eV and 714.3 eV attributable to Fe(III) in oxy(hydr)oxide and sulphate environments. The small component at higher binding energy is attributable to an Fe(III) multiplet structure. The Fe 2p_3/2 spectrum of the acid-leached sample was fitted in an identical manner. Peak areas show that Fe(III) comprises 30 and 36% of the total iron at the surface of the air-oxidized and acid-leached pyrite powder surfaces, respectively.

View larger version (29K): [in this window] [in a new window]   FIG. 4. Pyrite XPS: (a) Fe 2p, (b) S 2p,(c) O 1s; in each case, the bottom spectrum is for the fresh sample, middle for the air-ground sample and top the leached product.

Bonnissel-Gissinger et al. (1998) identified Fe and S oxidation products at the surfaces of pyrite samples ground in a nitrogen atmosphere. Therefore, minor oxidation of the pyrite fracture sample within the spectrometer chamber, or during fracturing under nitrogen, cannot be ruled out. Likely oxygen-bearing species in these systems include hydroxyl (OH^–), oxide (O^2–), adsorbed molecular water, sulphate species and additional sulphoxy anions (such as thiosulphate and sulphite). The O 1s data for the pyrite fracture sample have been closely simulated with a single peak having a binding energy of 531.1 eV. This is close to the reported binding energy of hydroxyl oxygen (as in Fe hydroxides). As noted above, the presence of ferrous hydroxyl species could explain the feature at 709 eV in the Fe 2p_3/2 spectrum of the pyrite fracture surface.

Two components were required to adequately simulate the O 1s data obtained for the air-oxidized pyrite sample. A peak centred at 531.1 eV is attributable to hydroxyl species and a second peak centred at 532 eV is attributed to sulphate. The O 1s spectrum of the acid leached pyrite powder was similarly fitted using two peaks at binding energies 531.1 eV and 532 eV. The greater relative area of the higher binding energy component in the O 1s spectrum of the acid leached sample is a reflection of the greater amount of sulphate present at the surface of this sample, consistent with the S 2p data.

The pyrite used in this study had a bulk Fe:S ratio very close to the ideal stoichiometry Fe:S = 0.5, whereas, based on the XPS data obtained for the pyrite fracture surface, Fe:S = 0.6. The net reaction stoichiometry for pyrite dissolution may be estimated from the XPS and solution chemistry data, the main reaction products being dissolved sulphate, ferrous ions and protons, leaving Fe-depleted sulphide and sulphate on the pyrite surface:

(5)

and

(6)

so overall:

(7)

driven by:

(8)

Hence, reactions 7 and 8 predict a net production of 2(y–x+xy) mol H⁺(mol FeS₂)^–1, (y+x–xy) mol Fe²⁺ (mol FeS₂)^–1 and 2y mol SO^2–₄(mol FeS₂) )^–1 However, in the experiment, 8.68 x 10^–5 mol H⁺ were produced and 4.70 x 10^–5 mol Fe²⁺ dissolved, giving a H⁺:Fe²⁺ molar ratio of 1.85, rather than the value of 2.0 predicted. Similarly, the SO4 ^2–:Fe²⁺ molar ratio is predicted as 2y:(y+x–xy), whereas, according to the data in Fig. 3, after the initial dissolution of air-formed products, the experimental value was 2, implying that x was very small or zero in the steady state.

	Sphalerite and galena

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
Figure 5 shows that the concentration of Zn(II) species exceeded that of dissolved sulphur species, the latter being essentially time independent, implying that the rate of the reaction 9 greatly exceeded that of reaction 10:

View larger version (16K): [in this window] [in a new window]   FIG. 5. Time dependence of dissolved Zn(II) and S concentrations in a pH 2.5 (stat) leaching experiment with powdered sphalerite at 25°C.

(9)

(10)

When coupled in electron balance with reaction 1, the overall chemical reaction:

(11)

results in a net proton consumption of 2(1 –y) mol H⁺ (mol ZnS)^–1. At and beyond 1.5 hours of leaching, the dissolved Zn(II) : S molar ratio (y^–1) was essentially constant with a value of ~10 (i.e. y = ~0.1), the time-dependence arising primarily from dissolution of products of aerial oxidation formed during grinding prior to the leaching experiments.

As shown in Fig. 6, dissolution of oxidation products formed during grinding galena produced dissolved Pb(II) and sulphur concentrations significantly higher than in the case of sphalerite leaching. However, concentrations of both species decreased over the first hour of the experiment, probably due to restricted solubility of PbCl₂ and PbSO₄, the latter phase having a particularly low solubility product K_sp(PbSO₄) = 10^–7.86

View larger version (15K): [in this window] [in a new window]   FIG. 6. Time dependence of dissolved Pb(II) and S concentrations in a pH 2.5 (stat) leaching experiment with powdered galena at 25°C.

(12)

(13)

(14)

Though the dissolved metal to sulphur ratios were not as high as in the case of sphalerite leaching, they increased from 1.3 (y = 0.7) after 19 min to 2.5 (y = 0.4) after 182 min. Though this may imply that the rate of reaction was greater than that of reactions such as 13, the insolubility of PbSO₄ would have restricted concentrations of both Pb(II) and dissolved sulphur, as implied in Fig. 6.

Based on the XPS data, no significant changes in S speciation occur at the surfaces of these minerals as a result of atmospheric oxidation and acid leaching. Elemental S has been reported at the surfaces of both air-oxidized and acid-leached sphalerite and galena, but no evidence for the presence of elemental S was obtained here, possibly due to sulphur desorption in the UHV chamber of the instrument, which had no low-temperature stage.

Chalcopyrite
Figure 7 shows that, after the initial dissolution, presumably of oxidation products formed mainly during grinding chalcopyrite in air, concentrations of dissolved S, Cu and Fe species decreased slightly, before increasing after the first hour and reaching time-independent values after 2 h. The redox potential (Fig. 1) decreased between 18 and 30 min and reached a relatively constant value of 472 mV. E⁰_{Cu2+/CuCl–2} (SHE)/V = 0.457, so that in 0.1 M HCl, the electrode potential for equal activities/concentrations of Cu²⁺ and CuCl^–₂ species is 0.339 V (SHE). Hence, the measured redox potentials correspond to conditions in which Cu²⁺ ions predominate.

View larger version (18K): [in this window] [in a new window]   FIG. 7. Concentrations of Fe, Cu and S as a function of time in a pH 2.5 (stat) experiment with powdered chalcopyrite at 25°C.

The S 2p spectra for the chalcopyrite powder sample and the acid-reacted powder clearly show the presence of sulphate at the powder surfaces (Fig. 8). In the case of the ground chalcopyrite sample, sulphate accounts for ~30% of the total surface S. In the case of the acid-reacted chalcopyrite powder, sulphate comprises ~3% of the total surface S. Spectra for the air-oxidized powder and the acid-reacted powder are also consistent with the presence of a ‘CuS’ phase, resulting in an additional component centred at ~162.5 eV in the S 2p spectra.

View larger version (23K): [in this window] [in a new window]   FIG. 8. S(2p) XPS spectra from; (a) a freshly cleaved surface of chalcopyrite (cleaved and transferred under nitrogen); (b) chalcopyrite powder prepared by dry grinding in air; and (c) powdered chalcopyrite following reaction with a pH 2.5 solution.

While Fe(II), Cu(II) and sulphate are the thermodynamically favoured aqueous reaction products, elemental S, non-stoichiometric Fe_(1–x)CuS₂ phases and covellite (CuS) have been reported as solid reaction products. Hence, a higher sulphur:sulphate molar ratio was expected than for pyrite. Based on mean values, the solution concentration ratios during the experiment were [Fe]:[Cu]:[S] = 1:0.3:0.8. This suggests that a Cu,S-rich surface alteration layer formed at the chalcopyrite surface. During the experiment, acid addition was required to maintain pH 2.5. Therefore, the mineral-solution reactions resulted in net proton consumption (cf. the pyrite experiment).

(15)

(16)

(17)

Overall:

(18)

(19)

Hence, in coupling reaction 19 to the oxygen reduction reaction (1), the net consumption of protons would be {8y –(2 + 6y + 2z)} = 2(y –z 31) mol H⁺ (mol FeCuS₂)^–1 and the predicted molar ratios [Fe(II)]:[Cu(II)]:[S(VI)] = 1:z:y. Values derived from experimental data in Fig. 7 were z = 0.32±0.01 and y = 0.73±0.06, omitting the first point at 18 min. The theoretically predicted value of –1.18 mol H⁺ (mol FeCuS₂)^–1 can be compared with the measured value of –1.59 mol H⁺ (mol FeCuS₂)^–1

Rates of leaching in single- and two-phase systems
Net mineral dissolution rates have been calculated on the basis of the temporal solution chemistry data (see Fig. 9). The surface areas of the –45 to +150 µm powder size fraction samples have been simply estimated by assuming that the powders comprise spherical mineral grains with an average uniform diameter of (45+150)/2 = 97.5 µm. Dissolution rates have then been estimated on the basis of the increase of the Fe (pyrite experiment), Cu (chalcopyrite experiment), Pb (galena experiment), Zn (sphalerite experiment) concentrations between the first and last sampling times in a given experiment.

View larger version (18K): [in this window] [in a new window]   FIG. 9. Dissolution rates (mol m–2 s–1) for single and two-phase systems.

The leaching rates of galena, chalcopyrite and sphalerite increased by factors of 31, 18 and 1.5, respectively, in the presence of pyrite, due to its superior catalytic properties (Biegler et al., 1970) for the oxygen reduction reaction 1. In mixed-sulphide mineral pulps and ore-bearing waste heaps, particles of different sulphides are in electronic contact and galvanic interactions must occur (Subrahmanyam and Forssberg, 1993).

In the galena + pyrite experiment, the concentration of Fe did not increase appreciably between the first and final sampling times, whilst the Pb concentration did increase significantly. Hence, galvanically-promoted dissolution of galena + pyrite decreases the pyrite electrode potential and its dissolution rate. Also in the galena + pyrite experiment, 75% of the total S in solution as measured by ICP-AES was not detectable by HPLC, which detects only anionic species; this could be due to the presence of colloidal elemental S. There was clear evidence of galvanically promoted chalcopyrite leaching in chalcopyrite + pyrite pulps; again in these experiments, 70% of the S in solution (measured by ICP-AES) was not detectable by HPLC, and so could be present largely as colloidal S rather than anionic sulphur species.

The rate of sphalerite leaching was increased only slightly in the presence of pyrite, though the mineral pulp pH decreased during the experiment. This suggests strongly that pyrite dissolution also occurred in this system, since pyrite dissolution alone decreased the solution pH in the single-mineral power experiments. However, the concentration of dissolved Fe did not increase significantly during the experiment, perhaps due to precipitation of Fe(III) species.

	Conclusions

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
In addition to the above specific conclusions regarding the rates and mechanisms of acid leaching of sulphide minerals under the particular conditions studied, a number of more general conclusions may be drawn: (1) the rates of acid dissolution of metal sulphides in mixed-mineral systems can be dramatically affected by galvanic effects, with rates increasing by factors as great as ~30 times in some cases; (2) the combination of conventional bulk leaching experiments with surface analysis techniques can lead to new insights into the mechanisms of the dissolution process through an understanding of reaction stoichiometry.

	Acknowledgements

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
The authors thank the UK Engineering and Physical Science Research Council (EPSRC) for a grant and Rio Tinto Technology Development Ltd for additional financial support. Drs K. England and P. Wincott are thanked for help with XPS experiments.

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 
Present address: BNFL plc, BTC (B170), Sellafield, Seascale, Cumbria, UK

Dedicated to the memory of Dr A. J. Criddle, Natural History Museum, London, who died in May 2002

[Manuscript received 9 July 2003: revised 13 November 2003]

	References

TOP ABSTRACT Introduction Experimental methods Results and discussion Sphalerite and galena Conclusions Acknowledgements References

 

Biegler, T., Rand, D.A.J. and Woods, R. (1970) Oxygen reduction on sulphide minerals. Part 1. Kinetics and mechanism at rotated pyrite electrodes. Electroanalytical Chemistry and Interfacial Electrochemistry, 60, 151–162.

Bonnissel-Gissinger, P., Alnot, M., Erhardt, J. and Behra, P. (1998) Surface oxidation of pyrite as a function of pH. Environmental Science and Technology, 32, 2839–2845.[CrossRef]

Craig, J.R. and Vaughan, D.J. (1994) Ore Microscopy and Ore Petrography (2^nd edition). Wiley Interscience, New York.

Karthe, S., Szargan, R. and Suoninen, E. (1993) Oxidation of pyrite surfaces: a photoelectron spectroscopic study. Applied Surface Science, 72, 157–170.[CrossRef]

Keith, C.N. and Vaughan, D.J. (2000) Mechanisms and rates of sulphide oxidation in relation to the problems of acid rock (mine) drainage. Pp. 197–219 in: Environmental Mineralogy: Microbial Interactions, Anthropogenic Influences, Contaminated Land and Waste Management (J.D. Cotter-Howells, L.S. Campbell, E. Valsami-Jones and M. Batchelder, editors). Mineralogical Society Series 9, Mineralogical Society, London.

Nesbitt, H.W. and Muir, I.J. (1994) X-ray photoelectron spectroscopic study of a pristine pyrite surface reacted with water vapour and air. Geochimica et Cosmochimica Acta, 59, 4667–4679.

Smith, E.E. and Schumate, K.S. (1970) Sulfide to Sulfate Reaction Mechanism. Federal Water Quality Administration Water Polution Control Research Series Report # 14010, FPS, 02/70.

Subrahmanyam, T.V. and Forssberg, K.S.E. (1993) Mineral–solution interface chemistry in minerals engineering. Minerals Engineering, 6, 439–454[CrossRef]

Vaughan, D.J., England, K.E.R., Kelsall, G.H. and Yin, Q. (1995) Electrochemical oxidation of chalcopyrite (CuFeS₂) and the related metal-enriched derivatives Cu₄Fe₅S₈, Cu₉Fe₉S₁₆ and Cu₉Fe₈S₁₆. American Mineralogist, 80, 725–731.[Abstract][ISI][GeoRef]

Williamson, M.A. and Rimstidt, J.D. (1994) The kinetics and electrochemical rate-determining step of aqucous pyrite oxidation. Geochimica et Cosmochimica Acta, 58, 5443–5454.[CrossRef][ISI][GeoRef]

Yin, Q., Kelsall, G.H., Vaughan, D.J. and England, K.E.R. (1995) Atmospheric and electrochemical oxidation of the surface of chalcopyrite (CuFeS₂). Geochimica et Cosmochimica Acta, 59, 1091–1100.[CrossRef][ISI][GeoRef]

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This Article Abstract Figures Only Full Text (PDF) Alert me when this article is cited Alert me if a correction is posted Services Email this article to a friend Similar articles in this journal Similar articles in ISI Web of Science Alert me to new issues of the journal Download to citation manager Citing Articles Citing Articles via HighWire Citing Articles via ISI Web of Science (13) Citing Articles via Google Scholar Google Scholar Articles by Abraitis, P. K. Articles by Vaughan, D. J. Search for Related Content GeoRef GeoRef Citation