Mechanism of chalcopyrite formation from iron monosulphides in aqueous solutions (< 100°C, pH 2–4.5)

Chemical Geology, 78 {1989) 325-341 Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands

325

Mechanism of chalcopyrite formation from iron monosulphides in aqueous solutions (< 1O0 ° C, pH 2-4.5)* MARK COWPER and DAVID RICKARD Department of Geology, University of Wales, Cardiff CFI, 3 YE, Wales (Great Britain) (Accepted for publication August 24, 1989)

Abstract Cowper, M. and Rickard, D., 1989. Mechanism of chalcopyrite formation from iron monosulphides in aqueous solutions ( < 100°C, pH 2-4.5). In: J. Schott and A.C. Lasaga {Editors), Kinetic Geochemistry. Chem. Geol., 78: 325341. In low-temperature aqueous solutions ( < 100 ° C, pH 2-4.5), chalcopyrite (CuFeS2) does not form through direct precipitation from solution. The pathway is exclusively via precursor iron sulphides and dissolved Cu salts. The reaction of dissolved Cu (II) salts with natural hexagonal pyrrhotite (Feo.9S) is diffusion controlled. The initial stage has an apparent activation energy of 11.4 + 1.8 kJ mol- 1and the rate ( in units of mol d m - 3 s - 1cm- 2) is independent of the solid reactant surface area. The reaction proceeds through a series of metastable Cu-Fe-sulphide intermediaries. These phases form a series of ephemeral layers penetrating into the pyrrhotite surface. The first phase formed has the stoichiometry Cuo.lFeo.gS.No Fe is released into the solution during its formation and this, together with the extremely low apparent activation energy and the stoichiometry, suggest that it is formed by stuffing of electron holes in the pyrrhotite structure with Cu ions. The transformation from the hexagonal close-packed arrangement of the pyrrhotite structure to the essentially cubic packing in chalcopyrite proceeds through a series of intermediaries, approximating in composition to members of the cubanite group. The rate of formation of these phases is controlled by the coupled diffusion of Fe(II) , Fe(III), Cu(I) and Cu(II) species through the surface reaction zone, although the process as a whole can be approximated by steady-state diffusion of total Cu into a semi-infinite medium. Experiments with metastable precursor iron monosulphide phases, includingamorphous FeS and synthetic mackinawite indicate similar reaction pathways. The results suggest that chalcopyrite formation in low-temperature natural systems may be significantly constrained by kinetic factors. Chalcopyrite is, at least, a diagenetic mineral since its formation requires the prior formation of iron sulphides. However, at ambient temperatures its formation is probably limited to very early diagenesis.

1. I n t r o d u c t i o n Chalcopyrite is the most common Cu ore mineral. However, the mechanisms and conditions of chalcopyrite formation from aqueous

solution, particularly at low ( < 200 ° C) temperature, are poorly understood. Previous workers have defined three possible mechanisms: 1.1. Direct nucleation from solution

*Paper presented at the International Congress of Geochemistry and Cosmochemistry, Paris, France, August 29September2, 1988.

0009-2541/89/$03.50

Roberts (1963) identified chalcopyrite as the only product when equal molarities of Fe (II)

© 1989 Elsevier Science Publishers B.V.

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and Cu (II) solutions were added to a slight excess of sodium sulphide. The immediate precipitate gave a broad X-ray diffraction (XRD) pattern of chalcopyrite. A more defined pattern was obtained when the precipitate was aged for 1 week in distilled water at room temperature and 150 ° C. Roberts (1963) proposed that chalcopyrite could be formed by direct nucleation from solution, 1.2. Solid state diffusion Young and Moore (1916) and Roberts ( 1961, 1963 ) proposed that chalcopyrite forms within days at room temperature when polished faces of pyrrhotite ( F e l _ x S ) a n d c h a l c o c i t e (Cu2S) were held in contact. It occurred as a brass-yellow deposit that replaced chalcocite at the contact whilst the surface of the pyrrhotite face was etched. The reaction was thought to proceed by the solid-state diffusion of Fe ( I I ) f r o m pyrrhotite into chalcocite, which was driven by a galvanic couple [measured as + 0.51 V by Roberts (1961) ] between the blocks. The rate of the reaction was faster in water (5 days compared to 1 month under vacuum at room temperature). However, Young and Moore (1916) only observed chalcopyrite after 2 months at the contact between Fe-bearing chalcocite (but not pure chalcocite) and pyrrhotite when the blocks were immersed in a solution saturated with H2S at room temperature, Roberts (1961) observed no reaction when the experiments were repeated using pyritechalcocite (FeS2-Cu2S) and pyrrhotite-covellite (FeI_xS-CuS) blocks. However, Young and Moore (1916) found that chalcopyrite formed rapidly when magnetite was added to pyrrhotite and covellite blocks immersed in H2S-saturated solutions, 1.3. Heterogeneous reactions Zies et al. (1916) found that pyrrhotite was altered to chalcopyrite in 5% copper (II) sulphate solutions between 40 ° and 200 ° C. If Cu

M. COWPER AND D. RICKARD

was in excess, chalcopyrite was replaced sequentially by bornite (CusFeS4), covellite (CuS), chalcocite (Cu2S), native Cu and Cuoxides. This work was repeated by Schouten (1934) and by us during this study. Under the same conditions, Zies et al. (1916) found that pyrite (FeS2) altered to the copper sulphides covellite and chalcocite with ferrous sulphate and sulphuric acid left in solution. Chalcopyrite was thought to be an intermediate phase during this alteration. This reaction was very slow below 100 ° C. Schouten (1934) found that chalcocite, covellite and bornite could be altered to chalcopyrite in concentrated iron (II) and iron (III) sulphate solutions at 120°C. Roberts (1961) observed no reaction between chalcocite blocks and solutions containing Fe (II) and both iron (II) and copper (II) salts after 10 days at room temperature. In this study, we have failed to observe any reaction between pure copper sulphides and dissolved Fe species but such reactions are well documented under oxidizing, acidic (pH < 2) conditions. These are analogous to supergene oxidation environments. However, in these environments chalcopyrite is unstable and tends to be oxidized to idaite (Cu3FeS4), covellite and blaubleibender covellite (Cul +xS) (Sillitoe and Clark, 1969). Dutrizac et al. (1970) observed chalcopyrite as an intermediate phase during the dissolution of bornite in acidified iron (II) sulphate solutions. It was found to persist metastably due to the formation of a sulphur coating that inhibited further dissolution. Walsh and Rimstidt (1986) produced a mixture of chalcopyrite and blaubleibender covellite by treating bornite with nitric acid. Our experiments demonstrate that chalcopyrite only forms at low temperatures from aqueous solutions at pH > 2 from the reaction between solid iron sulphides and dissolved Cu species. It does not form through direct precipitation nor through the reaction between solid copper sulphides and dissolved Fe species. In this paper, we report the results of studies of the

MECHANISM OF CHALCOPYRITE FORMATION FROM Fe-MONOSULPHIDES

reaction between dissolved Cu salts and iron monosulphides at pH 2-4.5. 2. E x p e r i m e n t a l m e t h o d s The reactant iron sulphide phases used were polished pyrrhotite crystals and freshly precipitated metastable iron monosulphides. Pyrrhotite was obtained as macroscopic aggregates from an unspecified location in Mexico. It was in the form of hexagonal crystals. The pyrrhotite was predominantly pure but common microscopic inclusions were pyrite, galena, silicate gangue, and rare arsenopyrite, pentlandite and primary chalcopyrite. Only polished sections with very few or no inclusions were used in the experimentation. Electron microprobe analysis (EPMA) showed it to be hexagonal pyrrhotite of the nC type, Feo.s9_o.91S(Scott and Kissin, 1982 ). Metastable iron (II) monosulphides were prepared by adding equal molarities of sodium sulphide to a m m o n i u m iron (II) sulphate in phosphate-based buffered solutions at pH 7-8 at room temperature. All precipitates were prepared, washed and filtered under a nitrogen atmosphere. In all experiments, the nitrogen was first bubbled through alkaline pyrogallol solutions to remove any traces of oxygen. Products were identified using XRD data obtained from a Philips ® 1820PW X-ray diffractometer. The initial phase formed was amorphous iron sulphide (Morse et al., 1987). In some cases the initial precipitate was aged in solution under nitrogen and the product was identified as mackinawite (Fel.,S). Both these phases are important in that they are the initial products in sedimentary pyrite formation (Rickard, 1975; Berner, 1984; Morse et al., 1987 ) formed by the reaction between detrital Fe minerals and sulphide produced by bacterial sulphate reduction in anaerobic marine sediments, Polished sections and precipitates were initially washed with 0.1 M hydrochloric acid (to remove any oxide coating) and then washed with distilled water before being sealed under a

327

stream of nitrogen into 150-ml round-bottom glass flasks containing 100 ml of the preheated reactant solution. Nitrogen had been previously bubbled into the solution before the solid reactant was added in order to displace any dissolved oxygen. Copper (II) sulphate solutions containing 200-5000 ppm Cu (0.003-0.08 M C u S Q ) were used. These solutions were buffered between pH 3 and 4.5 by using hydrogen phthalate (C6H4 (CO2)2H-) and acetate ( C H 3 C O 0 - ) solutions. All pH readings given in this paper were measured at room temperature. The ionic strength of buffered solutions varied between 0.1 and 0.5, whilst unbuffered solutions were usually an order of magnitude lower (0.025-0.075). In experiments where the temperature was <50°C, the flasks were placed into a shaking waterbath whilst between 50 ° and 100°C, the flasks were rotated in ovens. Both kept temperature to an accuracy of + 1 ° C. The longest runs were 2 months. Rate data were obtained by removing 1 ml of solution at defined intervals and analysing for Cu and Fe using standard atomic absorption spectrophotometry (AAS) techniques. The rate was determined using initial rate theory as outlined by Rickard ( 1974, 1982 ). The advantages of using polished sections were three-fold: (1) the geometric area of the reacting pyrrhotite face was controlled; (2) the section could be removed from solution at any time so that any reaction products could be examined under the microscope; and (3) it provided a convenient surface for determining concentration-depthprofiles by ball-cratering. This method involves grinding a circular hole into the polished reacted face with a steel ball coated in diamond paste. Knowing the radius of the ball and diameter of the crater, it is simple to determine the depth at any point in the crater (see the Appendix). The gradient of the crater is so shallow (e.g., the depth of a 1-mm crater made by a 10-mm radius ball is 12.5 # m deep) that the composition at any point within the crater can be determined using the line-scan facility of the electron microprobe. Thus, ac-

328

curate copper concentration-depths profiles could be obtained which enabled us to determine diffusion rates. The analyses were done using the Cambridge ® Microscan III probe at Cardiff and the Camdca ® Camdbax Dual WDS probe in the Department of Geology, University of Manchester. In the reactions between precipitates or powders and Cu solutions, the resultant solid product (s) were filtered and washed under nitrogen, and freeze dried under vacuum. The dried product (s) was identified by XRD. 3. R e s u l t s We carried out a series of experiments in order to evaluate the mechanisms of chalcopyrite formation outlined in the introduction. These showed that chalcopyrite is only formed at low temperatures through the reaction of solid iron sulphides and Cu solutions at pH > 2.9.

3.1. Direct nucleation from solution We repeated the experiment of Roberts (1963) by adding a mixed solution of 0.01 M Cu 2+-Fee + to an equal quantity of a 0.0205 M NaeS solution under a stream of nitrogen, such that sulphide was in slight excess after mixing, A black solid was precipitated at pH 7.14 and was immediately filtered and washed with distilled water under nitrogen. The dried precipitate gave a broad XRD pattern that matched a mixture of amorphous iron monosulphide Fel+xS (a broad peak at 5 ,£.) and the poorly crystalline blaubleibender covellite, Cul+xS, described by Shea and Helz (1989). However, this pattern could easily be mistaken for poorly crystalline chalcopyrite, particularly when interpreting Debye-Scherrer XRD patterns which were used by Roberts (1963). To prove the precipitate was not chalcopyrite, we obtained a semi-quantitative energy-dispersive spectrum (EDS) of the dried powder using an analytical Philips ® 400T scanning and transmission electron microscope. This showed the

M. COWPER AND D. RICKARD

precipitate not to be a single phase but a series of phases with different Cu/Fe ratios, consistent with an admixture of Fel+xS and CUl+xS. Furthermore, when we added equal quantities of a 0.01 M sulphide solution to a solution containing 0.01 M Cu e+ and Fe 2+, the bluish-black precipitate that formed at pH 4.46 was identified as poorly crystalline blaubleibender covellite only. When the solution was analysed by AAS, it contained 92% of the original Fe and no Cu. This implies that the rate of reaction between Cu and sulphide is not only faster than any reaction between Fe and sulphide, but also much faster than the reported rate of apparent direct chalcopyrite nucleation. Thus, it appears that chalcopyrite does not nucleate directly from solution, a view consistent with the inferences of a recent experimental study on Cu speciation at 25 ° C by Shea and Helz (1988).

3.2. Solid-state diffusion experiments It is difficult to evaluate the results of Roberts (1961) because experimental conditions such as pH and Eh were not specified. However, we have been unable to repeat these experiments using pyrrhotite and chalcocite blocks in buffered solutions at pH 4-7. In all cases, the pyrrhotite face became rapidly coated in orange hydrated ferric oxides. The possibility of the reaction being kinetically slow was examined by using powdered reactions at elevated temperatures. In no case was chalcopyrite detected.

3.3. Reaction between solid copper sulphides and iron solutions We carried out a series of experiments between 20 ° and 100 ° C to evaluate whether chalcopyrite could be formed by the reaction between solid copper sulphides and ferrous iron solutions. Polished sections of covellite and chalcocite were used as the solid reactants. Iron (II) ammonium sulphate solutions, buffered between pH 3 and 4.5, were used as reactant solutions. The longest runs lasted for 3 months

MECHANISM OF CHALCOPYRITE FORMATION FROM Fe-MONOSULPHIDES

at 75 ° C. In all experiments we failed to observe any reaction. Reactions between copper sulphides and ferric solutions are well documented (e.g., Walsh and Rimstidt, 1986, and references therein) and do not result in chalcopyrite formation, 3.4. Solid iron sulphides and copper solutions Our experiments show that chalcopyrite is formed below 100 ° C at pH > 2.9 in aqueous solutions by the reaction between solid iron sulphides and Cu solutions. We have failed to reproduce reported syntheses using other pathways and conclude that the reaction between solid iron sulphides and dissolved Cu species is the predominant, if not exclusive, route for chalcopyrite formation under these conditions. Below we present an experimental investigation of the mechanisms of the reaction involving iron monosulphides as reactants. The experimental conditions are outlined in Table III on p. 333. 3.5. Reaction between pyrrhotite and copper (II)solutions Pyrrhotite is readily altered to chalcopyrite in copper (II) sulphate solutions at low temperatures. However, ball-cratering through the reacted polished surfaces after various times showed that this alteration was not direct but proceeded through a series of intermediate compositions that form discrete layers between chalcopyrite at the surface and unreacted pyrrhotite at the base of the crater (Figs. 1 and 2). The two lowermost layers have no immediate natural analogue and we have called them phases A and B. The third layer is analogous to natural cubanite and synthetic chalcopyrite appears at the surface. These phases consistently supersede each other in the order (from the deepest within the pyrrhotite surface): (1) phase A; (2) phase B; (3) cubanite; and (4) chalcopyrite. This sequence is time dependent and related to the concentration of Cu in solu-

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tion, so that partial sequences can be observed. Electron microprobe analyses of the phases observed are listed in Table I. All results are the average of ten separate analyses for each run. The uncertainties are due in part to the fact that natural hexagonal pyrrhotite was used as a reactant. For comparison, the compositions of known natural phases in the Cu-Fe-S system are given in Table II. 3.5.1. Phase A. Phase A is the initial phase to form at the surface of the pyrrhotite. For all runs its composition ranges around 5.6 + 1.8% Cu (Table I) but is commonly observed with a maximum composition of 7.4% Cu. This corresponds to CUo.lFeo.gS.It occurs as a slight paleyellow coating at the pyrrhotite surface within hours at pH 4-4.5, but is only observed within the first 100 hr. The disappearance of phase A suggests that it is a transitional product, converting to more stable phases such as phase B and cubanite with time. Phase A is not yet an established phase in the literature. It is similar in composition and paragenesis to a cuprian mackinawite, Cuo.12Feo.94S, which was reported by Clark (1970) from the Y15jarvi Cu-W mine in SW Finland (Table II). It occurred as rods and lamellae associated with orthorhombic cubanite and chalcopyrite. Birks et al. (1959) had observed a similar Cu-bearing phase in the Mackinaw Mine, Washington, U.S.A. It occurred as small grains up to 30/~m in size which were enclosed in chalcopyrite. The phase contained 05% Cu and 51-58% Fe based on uncorrected probe data. Birks et al. inferred a composition approximating to "FeS" from these results. 3.5.2. Phase B. Phase B, in contact with phase A and apparently replacing it, has a distinct composition around 11.5_+ 1.5% Cu in all runs (Table I). Its stoichiometry varies between Cuo.15Feo.s2Sand Cuo.19Feo.soS,which appears to be significantly different from natural cubanite (Table II) and we have not been able to trace

330

M. COWPER AND D. RICKARD

Fig. 1. Reflectedlight micrographshowingan intermediarylayer (cub) of copper-ironsulphideswith cubanite-typecompositions (Fig. 2) betweenchalcopyrite(cpy) and pyrrhotite (pyr) in a crater bored into sample4A-7.

" CHALCOPYRITE LAYER

28)L~--J m

UNREACTED PYRRHOTITE Fig. 2. Trace of line scan across a crater bored in sample dA-]2 (reacted for 23 ] hr. ). The shaded areas reflect small pits in the surface produced during the crat~ring procedure.

any low-temperature, natural analogue of phase B in the literature, Phase B appears to be kinetically more stable than phase A in that it has been observed in contact with pyrrhotite in runs of > 200 hr. at 75°C. For example, in runs 4A-12 and 4A-13 (Table III), phase A was observed in contact with pyrrhotite in both samples after 50-hr. reaction, whilst phase B occurred at the surface, The partially reacted pyrrhotites were then

placed in the original solutions and reacted for another 181 hr. After this time, phase A was no longer observed and phase B was the phase in contact with pyrrhotite (run 4A-13). It is likely that a layer of phase A still exists between pyrrhotite and phase B but that it is too thin for detection by EPMA. However, the relative abundance of phase A in the products was substantially reduced. Likewise in run 4A-12, both phases A and B had been substantially replaced

MECHANISMOF CHALCOPYRITEFORMATIONFROMFe-MONOSULPHIDES

331

TABLE I

Electron microprobe analysis of phases observed during this study Fe (wt.%)

S (wt.%)

Ni (wt.%)

Total

Formula

61.09 _+0.2

39.09 _+0.2

0.03

100.21

Feo.s96S

4.02 _+0.5 5.35 + 0.6 7.41 7.62 + 0.5

57.46 + 0.5 58.31 + 0.7 56.74 56.83 _+0.5

38.26 -+ 0.5 36.87 + 0.6 35.74 35.57 _+0.3

99.74 100.47 99.91 100.02

Cuo.o6Feo.s6S Cuo.oTFeo.9,S CUo.lFeo.91S CUo.llFeo.91S

4B-2 3A-1

10.82 _+0.1 11.80_+1.0 12.20 12.84 _+0.2

52.27 _+0.4 51.50_+0.7 51.20 50.88 _+0.6

36.31 + 0.3 36.17_+0.4 36.85 36.81 _+0.3

99.40 99.47 100.26 100.59

Cuo.l~Feo.s4S CUo.lsFeo.s3S Cuo.,TFeo.soS Cuo.lsFeo.soS

Cubanite- type phases: 4B-2 4A-1 3A-2 4B-4

16.42 18.81_+0.4 22.53 _+0.7 23.51

47.39 46.10_+0.4 43.14 _+0.4 40.82

36.02 34.91_+0.3 34.58 _+0.3 34.51

0.01

0.02

99.85 99.82 100.25 99.86

Cuo.TFe2.3S3 Cuo.s3Fe2.3S3 CuFe2jsS3 Cul.o3Fe2.o3S3

Chalcopyrite: 4A-1 4A-12

34.08 33.14 _+0.6

30.93 32.30 Jr 0.5

34.46 34.95 _+0.4

-

99.47 100.39

CuFel.o2S2 Cuo.95Fel.o~S2

R u n No.

Cu (wt.%)

Unreacted pyrrhotite

Phase A: 4B-1 2F-D

4B-2 3A-1

0.02

Phase B: 4B- 1 2F-A

by cubanite, implying that it is kinetically more stable than both these intermediate phases,

3.5.3. Cubanite and related compositions. With time, phase B is replaced by a cubanite-type phase at the surface. This occurs after ~ 50 hr. at pH 4 and 75 ° C. Its composition varies around 19 + 3% Cu (Table I) which corresponds to a range of stoichiometries between Cuo.TFe2.3S3 and CuFe2S3 (naturally occurring cubanite is approximately CuFe2S3). It appears as a pale yellow-brown phase between pyrrhotite and chalcopyrite in longer runs (runs 4A-12 and 4A13, Table III).

3.5.4. Chalcopyrite and other phases. Chalcopyrite, CuFeS2, replaces the cubanite-type phase after ~ 70 hr. at pH 4 and 75 ° C (run 4B4, Table III ). It tends to be slightly Fe rich compared to stoichiometric chalcopyrite. Chalco-

0.01

pyrite was not observed in runs where the pH was < 2.9. If excess Cu is present in solution, the yellow chalcopyrite at the surface is tarnished blue [as was observed by Zies et al. (1916) in most of their runs ] . This is due to the formation of a thin coat of bornite at the surface. This is accompanied by a sharp drop in pH of the solution in unbuffered runs. In our experiments, if Cu is in excess, chalcopyrite is sequentially replaced by bornite (Cu~FeS4), digenite ((Cu,Fe)9Ss) and Cu metal at the surface. Sillitoe and Clark (1969)showed that the phases idaite (Cu3FeS4) and covellite are oxidation products of chalcopyrite and these were not observed in this study. Rickard (1973) showed that the formation of covellite was controlled by the amount of free elemental sulphur available. The stuffed chalcopyrite derivatives, talnakhite, mooihoekite and haycockite (Table II) were not observed in this study.

332

M. COWPER AND D. RICKARD

T A B L E II Phases in the C u - F e - S system at low temperature ( < 250°C) Name

Composition

Pyrite (marcasite)

FeS2 FeS Feo.s9_o.91S

Mackinawite

FeTS8 FegSs

Cuprian

Cuo.12Feo.94S

Troilite nC pyrrhotite

4C pyrrhotite

Cu (wt.%)

Fe (wt.%)

S (wt.%)

-

46.55 63.53 60.79-61.32

53.45 36.47 39.21-38.68

-

60.39 66.21

39.61 33.79

8.08

57.06

33.95

Note

NiAs structure hexagonal supercells based on NiAs vacancy

structure monoclinic can obtain large amounts of Co, Ni, Cu

mackinawite Cubanite

CuFe2S3

23.39

41.25

35.56

orthorhombic wurtzite-like

Chalcopyrite

CuFeS2

34.60

30.52

34.88

orthorhombic sphalerite-like structure

Haycockite

Cu4Fe~Ss CugFegS16 Cu9 (Fe,Ni)gS16* Cu3FeS4 CusFeS6 CusFeS4

32.18 36.03 37.20 50.88 56.15 63.30

35.35 31.66 29.87 14.95 9.90 11.17

32.47 32.31 32.40 34.17 33.95 25.53

CusFeS4+~

63.07

11.00

25.93

(Cu,Fe)gSs CuS CugSs Cu~9S28 CU1.96S Cu2S

77.13 66.49 69.04 73.41 79.52 79.87

1.00 -

21.87 33.51 30.96 26.59 20.48 20.13

structure

Mooihoekite Talnakhite Idaite Nukundamite Bornite "Anomalous bornite" Digenite Covellite Yarrowite Spionkopite Djurleite Chalcocite

] (

both have "stuffed" chalcopyrite structures oxidation product of bornite

complete solid solution with digenite above 83 ° C metastable oxidation product of bornite

] ~

blaubleibender covellites; metastable phases

*Contains 0.63 wt.% Ni.

4. R a t e d a t a 4.1. P h a s e - A kinetics

The rate of reaction between pyrrhotite and Cu(II) solutions was followed by measuring changes in metal concentration in solutions. A typical concentration-time curve is shown in Fig. 3. As phase A forms at the surface, there is an initial rapid loss of Cu from solution, but no increase of Fe into the solution. The initial rate (which applies only to phase A formation; Fig. 3 ) is directly proportional to the geometric area

of the polished pyrrhotite, suggesting that the rate can be expressed in rate per unit geometric surface area. A plot of initial rate per unit geometric area vs. Cu concentration is a straight line passing through the origin (Fig. 4), indicating that the rate of phase A formation is pseudo first-order with respect to Cu concentration in solution. The initial rate of Cu loss from solution is also temperature dependent. This relationship is defined by the Arrhenius equation: In k' = ( - E a / R ) ( 1 / T ) + In A

(1)

where k' is the pseudo first-order rate constant

pyrrhotite block pyrrhotite block pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrrhotite face pyrite face pyrite face mackinawite mackinawite mackinawite amorphous FeS amorphous FeS amorphous FeS

1-D 1-E 2A-4 2A-2 2F-A 2F-D 3A-1 3A-1 3A-2 4A-1 4A-1 4A-7 4A-10 4A-12 4A-12 4A-13 4A-13 4A-16 4A-17 4B-1 4B-2 4B-3 4B-4 2F-E 4B-5 6A 6A 6A 6B 6D 6F

11.71 5.31 5.02 6.83 1.10 0.45 1.17 1.17 1.05 1.40 1.40 0.85 1.38 0.96 0.96 1.10 1.10 1.20 1.38 0.45 1.00 0.90 1.00 1.00 1.25 0.5 g ppt 0.5 g ppt 0.5 g ppt 6.00 g ppt 0.05 g ppt 0.50 g ppt

Surface area (cm 2) 600 600 500 500 2,500 1,000 600 600 600 600 600 1,000 1,000 2,000 2,000 600 600 600 1,000 600 600 600 600 1,000 1,000 1,000 1,000 1,000 1,000 1,000 1,940

Reactant solution (ppm CuS04 ) 4.24-3.80 4.24-4.01 4.21-3.37 4.21-3.20 3.00 (B) 3.00-3.38 3.54 (B) 3.57 (B) 3.60 (B) 3.99 (B) 3.97 (B) 4.45 (B) 4.45-4.24 4.51-4.20 4.51-4.20 4.56 (B) 4.56 (B) 4.45 (B) 4.53-4.16 4.02 (B) 4.02 (B) 3.95 (B) 3.97 (B) 3.30-3.05 4.03 (B) 4.05 (B) 4.05 (B) 4.05 (B) 4.21-3.80 4.58-2.41 3.95 (B)

pH (B = buffered)

25 50 75 70 75 75 75 75 75 75 75 75 75 75 75 75 75 75 75 75 75 75 75 75 75 18 50 75 75 75 75

Temperature (°C)

187 187 214 214 50 212 100 200 24 50 100 231 231 50 231 50 231 168 100 50 50 24 65 200 212 168 168 168 47 96 48

Duration (hr.)

4.78_+0.24 7.07 _+0.40 9.12_+0.33 8.73 + 0.32 7.30 _+0.60 9.91 _+1.00 26.0 _+1.50 21.4 _+3.50 8.00+0.80 12.4 _+1.0 710 _+70 1,100 _+100 1,780 _+150 all Cu reacted all FeS reacted all Cu reacted

Initial rate (10 -9 tool d i n - 3 s - 1 cm - 2) Cpy Cpy Bn, Cpy Bn, Cpy Cub Cub, B, A B, A Cub, B, A A Cub, B, A Bn, Cpy Cpy, Cub, A Bn, Cpy, B Cpy, Cub, B, A Cpy, Cub B, A Cpy, Cub, B Cpy, Cub, B Cub, B, A Cub, B, A Cub, B, A B, A Cpy, Cub, B, A no reaction no reaction BBCov, Dig BBCov, Dig BBCov, Dig Cpy, mackinawite Bn, Dig, Chal Cpy, Cub

Products*

*Key: Cpy = chalcopyrite; Cub = cubanite; Bn = bornite; Dig = digenite; Chal = chalcocite; B = phase B; A = phase A; BBcov = blaubleibender covellite.

Solid reactant

Run No.

Experimental conditions

TABLE III

12.0 _+0.6 9.0_+0.8 12.0_+ 1.5 5.2 _+0.7 12.0 _+1.0 22.0 _+1.0 15.0 _+1.1 19.7 -+ 0.5 11.0 + 0.3 29.0 _+1.2 11.7-+0.5 17.7-+0.6 15.5 _+1.5 12.0 + 1.0 13.5 +0.9 9.0_+0.6 15.8 _+1.0

Layer thickness (/~m)

>

¢D ¢.O ¢.~

~

oZ

¢~

C)

>

o

z

334

M. C O W P E R A N D D. R I C K A R D 600

z

--0400 12: Z300

Phase A forming

z

I I I

°200

II

Phase B forming

'~' 100

0

20

40

60

REACTION TIME

80

100

(hr.)

Fig. 3. Plot showing the change in bulk metal content of the solution with time for run 3A-1 ([] = Cu; A = Fe). Error bars (vertical lines in the symbols) reflect uncertainties in AAS values. The changes in Cu and Fe concentrations shown in this figure are typical for the initial stages of the reaction between pyrrhotite and copper (II) sulphate under our experimental conditions (Table I I I ) . { 1 ppm Cu = 1.5737-10 -5 mol dm-3; 1 ppm F e = 1.7906" 10 -5 mol dm -3. )

~" 2

o 1 ~ -0 000~ 001

0.015 002

0.025 003

o.035

COPPER CONCENTRATION (mol dm -3)

Fig. 4. Plot showing the variation between initial rate of Cu loss per unit geometric area of pyrrhotite ( = rate of phase A formation) vs. initial Cu concentration at 75°C and pH = 4.35 4- 0.2. The linear relationship indicates "pseudo" first-order dependence of phase A formation on Cu concentration,

in s-1; Ea is the apparent activation energy (or strictly the critical increment of energy for a heterogeneous reaction) in J mol-1; R is the universal gas constant in J m o l - 1 K - 1; T is the absolute temperature in kelvins; and A is the pre-exponential function, Thus, a plot of I n k ' vs. 1/T yields a straight

line plot of gradient (-Ea/R). Generally, diffusion-controlled reactions in solution have low activation energies in the range 10-21 kJ m o l - 1 (Rickard, 1975; Lasaga, 1981). Activation energies for chemically-controlled reactions are usually an order of magnitude higher, typically between 40 and 90 kJ mol-1 (Lasaga, 1981). An Arrhenius plot [ln (initial rate, k' ) vs. 1/ T] in the temperature range 20-80°C is shown in Fig. ha. This plot yields an apparent activation energy of 11.4+1.8 kJ mo1-1. This extremely low activation energy suggests that the initial stage of the heterogeneous reaction between pyrrhotite and Cu (II) solutions (phase A formation) is transport controlled. This is consistent with the independence of the rate per unit area of the surface area of the sample. A1though it is possible t h a t the number of chemically reactive sites on the pyrrhotite surface is directly proportional to the surface area for an individual specimen, it is improbable t h a t this fortunate relationship pertains over a series of samples. Normally, therefore independence of rate per unit area of sample surface area dependence is a feature of transport-controlled reactions. We also observed that the products of this reaction form an even coating over the whole

MECHANISM OF CHALCOPYRITE FORMATION FROM Fe-MONOSULPHIDES TEMPERATURE (°C)

~0

80

40

20

-185s o18.65

187~ ~c -18.85 -~ 18.95 -19,o5

335

As more Cu diffuses into the pyrrhotite lattice, phase B is formed at the surface. The rate of loss of Cu from solution decreases and Fe (II) (identified by a red colour when dimethylgloxime is added) is released into solution (Fig. 3). Fe (II) continues to be released into solution as cubanite and chalcopyrite forms. This suggests that later stages of the reaction between pyrrhotite and Cu (II) solutions are due to both Cu and Fe diffusion.

-19.15 -19.25

t 2.9

2.8

, 3.0

I 3.1

I 3.2

L 3.3

4.2. Diffusion rates: thickness versus time 3.4

1] TEMPERATURE(103K-1)

-,3, ~,33

,35 _~-~37 ~39 -,4, -,~3 2.8

2.9

3.0

31

32

I/TEbAPERATURE(103

33

34

35

K-1)

Fig. 5. a. Arrhenius plot for the initial stage of the reaction

between pyrrhotite and copper (II) sulphate, the formation of phase A, between 20° and 80 ° C, with initial p H 4.24.The gradient of the line drawn yields an "apparent" activation energy (see t e x t ) o f 11.37_ 1.8 kJ tool-1. (Error bars represent uncertainties in the measurement of the initial rate, k'. ) b. Arrhenius plot for the initial stage of the reaction between mackinawite, Fel.IS, and copper (II) sulphate between 18 ° and 75 oC at p H 4.00 ___0.08. The gradient of the line drawn represents an "apparent" activation energy of

13.4_+2 kJ tool-1.

surface of the sample and are not concentrated at potential high-energy surface sites such as fractures and crystal boundaries. This is also consistent with the reaction being transport controlled rather t h a n surface reaction controlled,

The use of the ball-cratering to produce craters with shallow gradients allowed us to use E P M A to accurately determine concentrationdepth profiles and diffusion rates of Cu in the pyrrhotite lattice. A pair of typical concentration-depth profiles for various experimental conditions of dissolved Cu concentrations, pH and temperature are shown in Fig. 6a and b. The depth profiles display a series of layers with different Cu contents. The Cu contents of the layers decrease downwards into the surface but the composition of individual layers is approximately constant and the Cu concentration gradients within each layer are so small as not to be measurable with our experimental technique. The gradients within each layer are t h e n far less t h a n the gradients between layers as shown schematically in Fig. 6a and b. T h e s e q u e n c e of the layers is dynamic, m o v ing downward into the pyrrhotite crystal with time. The total thickness of Cu penetration is defined as the point at which the Cu content drops below the detection limit of the electron microprobe. Plots of depth of Cu penetration vs. time are shown in Fig. 7. This plot shows that there is an increase in the depth of Cu penetration with time but the rate of Cu penetration decreases with time. Furthermore, there is a pH effect. The depth of Cu penetration increases up to p H 4, but begins to decrease at pH > 4. The pH dependence is thought to be due to speciation of metal ions in solution.

336

M. COWPER AND D. RICKARD

36 " 33-

~

A

LCO PYRITE

(a)

3027" 24-

H=4,0

~ LU E 12

"CU

°°~w s

- . . . . . . . . .

6 3

PHASE A ~ i

o

2

4 6 8 10 12 14 DEPTH INTO CRATER [,,'U,m)

16

~_~ 4 18

(b)

20 ~--]"CUBANITE"] "CUBANITE" 18- ' ' ~ 16i----~--- - i 14. . . . .

~J ,~ 12-

8 o 6

~ ~

, 2

,

,

,

4

6

8

"~k 10

12

i

14

I 50

I I 100 150 REACTION TIME (hr.)

I 200

250

Fig. 7. Plot showing variation between depth of Cu penetration (total thickness of the layers) and time for pyrrhotite-CuSO4 runs at 75°C and an initial Cu concentration of 600 ppm (0.00944M). The pH variation ( A = 3.5; [] = 4; V = 4.5 ) is thought to be due to changes in metal speciation in solution. P r o d u c t s were identified b y X R D . It was n o t possible to e v a l u a t e w h e t h e r p h a s e A or p h a s e B was f o r m e d as X R D d a t a are n o t available for t h e s e phases. C h a l c o p y r i t e is readily f o r m e d b y t h e reac-

P.~Sse-&-~ . . . . . ......... . . . . . . ~,

4

0 0

24-

~o P.~SEs

=

o

I

22-

20

/

~t~ "~'u.16

+ . . . . . . "7

1~

~

24

-- ~ 20 ~

"~ 2118.

£L

! z)

16

18

DEPTHINTOCRATER(A.km) Fig. 6. Schematic plot showing the variation of Cu content (fromEPMA)withdepthfromthesurface (depth=0/Lm) to unreactedpyrrhotite (0wt.% Cu) for: (a) 4B-4 and (b) 4B-2 (reaction conditions are given in Table III). Overall, there is an inverse relationship between Cu content and depth (thick line) but the actual profile is stepwise (dashed line ) due to the formation of layers of intermediary products (phase A, phase B and cubanite) between unreacted pyrrhotite and the surface. In each layer, O[Cu]/O (depth) is approximately zero, and is much less than the overall gradient between the layers (thick line),

4.3. Reaction between metastable iron (II) monosulphides and Cu solutions F r e s h l y p r e c i p i t a t e d a m o r p h o u s iron sulp h i d e (Fel + xS ) a n d m a c k i n a w i t e (Fel.IS) were r e a c t e d in b u f f e r e d a n d u n b u f f e r e d c o p p e r (II) s u l p h a t e solutions ( p H 2-4.5 ) a t t e m p e r a t u r e s r a n g i n g f r o m 20 ° to 100 ° C. T h e r a t e o f r e a c t i o n was d e t e r m i n e d b y following c h a n g e s in m e t a l c o n c e n t r a t i o n in s o l u t i o n w i t h t i m e (by A A S ) .

tion

between

copper

(II) sulphate

and

meta-

stable iron (II) m o n o s u l p h i d e s below 100°C. T h e r e a c t i o n results in Cu loss from, a n d Fe release into, t h e solution ( T a b l e III). T h i s implies t h a t t h e same d i f f u s i o n - c o n t r o l l e d mecha n i s m is a c t i n g for b o t h p y r r h o t i t e a n d t h e m e t a s t a b l e iron (II) m o n o s u l p h i d e s d u r i n g rea c t i o n s w i t h c o p p e r solutions. T h i s was f u r t h e r s u p p o r t e d b y t h e results o f e x p e r i m e n t 6A (Wable III), in w h i c h t h e initial rate o f Cu r e m o v a l f r o m solution was d e t e r m i n e d at 18 ° , 50 ° a n d 75 ° C. A plot o f l n (initial r a t e ) vs. 1 / T (Fig. 5b) yields a n a p p a r e n t a c t i v a t i o n e n e r g y of 13.4 _+2 k J m o l - 1 for t h e initial r e a c t i o n b e t w e e n a m o r p h o u s iron m o n o s u l p h i d e a n d c o p p e r (II) sulp h a t e . T h i s is w i t h i n t h e same range as t h e app a r e n t a c t i v a t i o n e n e r g y of 11.4 _+ 1.8 k J m o l - 1 o b t a i n e d for p h a s e A f o r m a t i o n d u r i n g t h e initial reaction b e t w e e n p y r r h o t i t e a n d copper (II) sulphate. In r u n 6B, w h e r e F e S was in excess, all the Cu h a d r e a c t e d w i t h i n 2 days at 75°C. T h e

337

M E C H A N I S M OF C H A L C O P Y R I T E FORMATION FROM Fe-MONOSULPHIDES

products were identified by a well-defined Xray pattern of chalcopyrite and mackinawite. A duplicate experiment interrupted after 1 day gave an X-ray pattern of mackinawite, and an apparent mixture of cubanite and chalcopyrite. Analysing this precipitate using the EDS facility of an SEM, revealed two distinct Cu-Fe-S phases; however, the stoichiometry could not be determined. When Cu was in excess (run 6D), all the FeS had been converted to digenite and chalcocite. This was accompanied by a sharp drop in the pH of the solution, 5. Discussion Since the total thickness of the pyrrhotite crystal is very large compared with the reaction layer, the reaction system can be approximated by considering it analogous to Cu diffusion through a semi-infinite medium, where the Cu concentration at the boundary is kept at a constant concentration, Co, and the initial Cu concentration is zero at depth in the medium. This requires a solution to Fick's second law of diffusion for the one-dimensional case (x increasing into the surface): (2)

8c/Ot=D'O2c/Ox 2

ness of the layers) should be proportional to the square root of time. This is shown in Fig. 8. The regression lines drawn have correlation coefficients of 0.995 (n--4) at pH 3.5 and 0.992 (n = 5 ) at pH 4 with errors well within the experimental error. The relationship is approximate because eq. 2 does not accurately define the system. In reality, an extra term relating to the chemical reactions of Cu is required. However, if the reactions are fast relative to the diffusion rate, the effect will be to reduce D in eq. 2. In the simplest case where S, the concentration of Cu fixed by chemical reaction, is directly proportional to the amount of Cu free to diffuse, c, i.e: S = rc (5) where r is a function characteristic of the reaction, then: Oc/c)t---D. c)2c/Ox 2 - OS/Ot (6) and solving this equation (from Crank, 1956): Oc/Ot= ( D / r + 1 ). 02c/0x 2 (7)

24

z

~E Z ~ 16 ~ ~ ~ 12 ~

C=Co,

0¢o

t>0

pH:40

i ~,

'

~

~

4 °~ pH :3.5

u-z

and c=O,

J

~ ~ 20

where D is the diffusion coefficient for Cu in pyrrhotite; t is time; and x is depth in the crystal. With the boundary conditions: X:0,

~

o_~

o 0

x>0,

5~4

t=O

o_

the solution to eq. 2 (Crank, 1956) is: C=Co erfc[x/{2(Dt)l/2}]

00 (3)

where erfc is the reciprocal error function integral. This relationship involves only the single dimensionless parameter: x / { 2 ( D t ) ~/2}

(4)

If this model is correct, the depth of Cu penetration into the pyrrhotite surface (the thick-

2

4

6

8

10

12

~4

SQUARE ROOT OF TIME

Fig. 8. Linear plots between depth of Cu penetration (total thickness of the layers) and square root of time for pyrrhotite-CuS04 runs at 75 oC, initial Cu concentration of 600 ppm and pH = 3.5 ( A; n = 4; correlation coefficient--0.995) and p H - - 4 {r-l; n=5; correlation coefficient--0.992). These imply that overall the reaction between pyrrhotite and copper (II) sulphate can be modelled by steady-state diffusion of total Cu into a semi-infinite

medium(seetext).

338

M. COWPER AND D. RICKARD

The effect is to reduce D in eq. 2 to D/r + 1 in eq. 7. This underlies the fact that D in eq. 2 is not the simple diffusion coefficient. Because of the variation in composition of the medium with time and depth, the diffusion coefficient through the experiment must vary. However, this will not effect the relationship between concentration and square root of time because Boltzmann (1894) showed that the square root relationship is a boundary condition for the process, Not only is diffusion the rate-limiting step for the reaction, but the chemical potential gradient between the solution and the solid interior can be shown to be the driving force. In the case of a single layer of thickness l, with the boundary conditions: x = 0, and

c = Co

dc/dt=O

x = l, c= 0 Fick's second law in one dimension (eq. 2) reduces to:

Oc/Ot=0

(8)

Plots of Cu concentration vs. depth (Fig. 6a and b ) show a series of steps, as different layers with different Cu concentrations are developed. However, overall the plots of Cu concentration vs. depth are straight lines (Fig.6a and b), indicating that eq. 8 is approximately held. This steady-state approximation, which is valid at any instant for the reaction, was first used by Stefan (1890) for the case of diffusion with a moving boundary. In the present study, the fastest diffusion rate measured was < 10 -s cm s - 1. Therefore, the steady-state model could be used to roughly approximate the experimental product, Although the chemical potential gradient between pyrrhotite and the surface drives the overall process, it does not function in detail. Since the Cu concentration-depth profile within any individual layer is less than for the

whole system, the rate of diffusion of Cu within the individual layers should be less than the rate of diffusion through the reaction zone as a whole. This is emphasized by the apparent large Cu concentration discontinuities at the boundaries of the individual layers. The layers appear to act as theoretical bottle-necks to the diffusion of Cu ions. However, since the diffusion is both ion specific and coupled it is possible that the rate within each layer is controlled by the coupled diffusion of Cu and Fe ionic species. ThatisthatCu(II) andCu(I),andFe(III)and Fe (II) diffuse at different rates whilst maintaining a charge balance at any point in the layer and relatively constant total Cu and total Fe concentrations. This idea is encouraged by the fact that although the Cu in solution is in the divalent state, the Cu in the copper-iron sulphides, although complex, appears to be mainly monovalent (Vaughan and Craig, 1978). Therefore, redox reactions must be occurring within the reaction zone, even though these are not directly observed. However, the question still remains as to why pyrrhotite is not altered directly to chalcopyrite but goes through a series of intermediary phases (phase A, phase B and cubanite). The reactant pyrrhotite, Feo.sg_o.91S, is non-stoichiometric and is characterized by an electron defect structure based upon the hexagonal NiAs structure (Vaughan and Craig, 1978). For electronic neutrality, up to 20% of the Fe must be Fe(III) leaving a series of electron holes. Phase A forms when Cu ions diffuse into the pyrrhotite lattice and fill these defects, "stuffing" the pyrrhotite lattice with Cu atoms. This interpretation is supported by the fact that the change from Feo.gS to Cu0.,Feo.gS by infilling electron holes requires ~ 7% Cu, coincident with the maxim u m composition of phase A. It is further supported by the extremely low activation energy of 11.4 + 1.8 kJ m o l - ' for this process (Fig. 5 ) and by the experimental observation that no Fe (II) was released into solution as phase A was formed (Fig. 3). The increase of the Cu proportion above 7% results in a structural rearrange-

MECHANISM OF CHALCOPYRITE FORMATION FROM Fe-MONOSULPHIDES

ment, causing the first discontinuity, Phase B and subsequent phases are formed by both Cu and Fe diffusion. The change from pyrrhotite to chalcopyrite and beyond involves a rearrangement of the sulphur lattice from hexagonal close packing (h.c.p.) in pyrrhotite to cubic close packing (c.c.p.) in chalcopyrite, which has an orthorhombic superstructure based on the unit cell of sphalerite (Hall and Stewart, 1973 ). Since it is probable that phase A retains the structure of the pyrrhotite, the major change appears to centre around the development of cubanite. This adopts an orthorhombic (pseudo-hexagonal) structure (Szymanski, 1974) which Buerger (1947) characterized as "slabs of wurtzite-likestructure parallel to ( 0 1 0 ) " . The wurtzite and sphalerite structures are closely related: both can be regarded as polytypes which differ only in the stacking of a layer of tetrahedra-sharing vertices - a c.c.p, sequence in sphalerite and a h.c.p, sequence in wurtzite (Wuensch, 1974). Therefore, phase B and Fe-rich cubanite compositions (Table I) represent transitional structural states between the pure h.c.p, sulphur array in pyrrhotite and the h.c.p, sulphur array in the wurtzite-type structure of cubanite, 6. C o n c l u s i o n s Chalcopyrite is formed by the reaction between precursor iron sulphides and dissolved Cu. Below 200 ° C, the reaction between pyrrhotite or other iron (II) monosulphides and Cu solutions (pH 3-4.5) is fast, and is geologically significant at ambient surface temperatures, The pyrrhotite to chalcopyrite reaction is diffusion controlled and goes through a series of metastable intermediaries. The first phase to form represents simple Cu diffusion into the pyrrhotite lattice. The activation energy of this process is 11.4 + 1.8 kJ mol-1. Continued addition of Cu involves both Fe and Cu diffusion and results in a series of cubanite-type compo-

339

sition. These represent transitional structural states between hexagonal sulphur packing in pyrrhotite and cubic sulphur packing in chalcopyrite. It appears that phase A, phase B and possibly, the cubanite-like phase observed have only a limited lifetime and may not be abundant naturally. However, it is possible that more detailed studies ofthe contacts between natural iron (II) monosulphides and chalcopyrite will reveal thin layers of these intermediaries. The rate of the reaction is a function of total dissolved Cu concentration, geometric pyrrhotite area and pH. The pH effect is unclear but is thought to be due to changes in aqueous Cu speciation. The reaction between metastable iron (II) monosulphides and copper sulphate appears to proceed via a similar diffusion-controlled mechanism. The apparent activation energy for the initial stage of this reaction is 13.4 _+2 kJ m o l - 1, which is similar to that recorded for pyrrhotite under the same conditions (Table III). The reaction proceeds by both Fe and Cu diffusion. Like pyrrhotite, mackinawite, FellS, has a vacancy structure based on a c.c.p, array of S atoms with Fe occupying some of the tetrahedral interstices but vacancies in larger octahedral interstices (Morse et al., 1987). These vacancies provide a pathway for Cu diffusion into the lattice. The structures differ in that whilst pyrrhotite is Fe-deficient, mackinawite is Sdeficient. Preliminary results from the reaction between pyrite and dissolved Cu indicate that this reaction is characterized by high activation energy and is strongly temperature dependent. At < 200 °C it is almost exclusively localised to high-energy surface sites, such as defects, crystal edges and fractures. The reaction rate is thought to be controlled by a surface chemical reaction. Chalcopyrite was not observed to form by direct nucleation from solution or by reaction between copper sulphides and Fe solutions. This implies that chalcopyrite in low-temperature synsedimentary Cu deposits is at least an early

340

diagenetic mineral: that is a precursor iron sulp h i d e m u s t be p r e c i p i t a t e d first w h i c h subseq u e n t l y reacts w i t h Cu s o l u t i o n s to f o r m chalcopyrite. I f t h e iron sulphide p r e c u r s o r is a

monosulphide, chalcopyrite formation may be r a p i d at a m b i e n t surface t e m p e r a t u r e s . H o w ever, t h e r e a c t i o n o f p y r i t e w i t h Cu s o l u t i o n s at t h e s e t e m p e r a t u r e s is relatively slow. T h i s m e a n s that synsedimentary chalcopyrite either f o r m s by: (1) r e a c t i o n o f iron ( I I ) m o n o s u l -

phides (e.g., mackinawite, amorphous iron sulp h i d e ) w i t h dissolved Cu d u r i n g early diagenesis before pyritization, or (2) reaction of pyrite with C u - r i c h s o l u t i o n s d u r i n g a n early hight e m p e r a t u r e process, or (3) d u r i n g a late stage

of basin development when higher temperatures may be attained intrinsically. Acknowledgements T h e s t u d y was s u p p o r t e d b y t h e N a t u r a l E n -

vironmental Research Council grant GR3/5634. We t h a n k P e t e r Fisher, K e v i n O ' F a r r e l l y a n d

Anthony Oldroyd for help and advice, and Tim H o p k i n s a n d D a v i d P l a n t at M a n c h e s t e r U n i v e r s i t y for t h e i r help w h i l s t u s i n g t h e C a m d c a ®

electron microprobe. We would also like to thank Martin Schoonan and L.S. Sj/hberg for their constructive criticisms and n u m e r o u s helpful c o m m e n t s d u r i n g t h e i r reviews o f this

manuscript. Appendix In a two layer system (A and B), ball-cratering through the top layer (A) will reveal two circles when the surface is viewed from above. If the bottom circle has diameter d2, and the top circle has a diameter dl, the thickness of the top layer can be simply calculated by: (thickness of A ) = [ (dl) 2_ (d2) 2]/8R (A- 1 ) where R is the radius of the steel ball that made the crater,

References Berner, R.A., 1984. Sedimentary pyrite formation: An update. Geochim. Cosmochim. Acta, 48: 605-615. Birks, L.S., Brooks, E.J., Adler, I. and Milton, C., 1959.

M.COWPERANDD.RICKARD Electron probe analysis of minute inclusions of a copper-iron mineral. Am. Mineral., 44: 974-978. Boltzmann, V.L., 1894. Zur Integration der Diffusionsgleichung bei variabeln Diffusions-koeffizienten. Ann. Phys. (Leipzig),53: 959-964. Buerger, M.J., 1947. The crystal structure of cubanite. Am. Mineral., 32: 415-425. Clark, A.H., 1970. An unusual copper-iron sulphide, Cuo.12Feo.94Sl.o0, from the Ylhjarvi deposit, Finland. Econ. Geol., 65: 590-591. Crank, J., 1956. The Mathematics of Diffusion, Oxford University Press, London, 1st ed., 348 pp. Dutrizac, J.E., Mcdonald, R.J.C. and Ingraham, T.R., 1970. The kinetics of dissolution of bornite in acidified ferric sulfate solutions. Metall. Trans., 1: 225-231. Hall, S.R. and Stewart, J.M., 1973. The crystal structure refinement ofchalcopyrite. ActaCrystallogr.,B31:2105_ 2112. Lasaga, A.C., 1981. Rate laws of chemical reactions. In: A.C. Lasaga and R.J. Kirkpatrick (Editors), Kinetics of Geochemical Processes. Reviews in Mineralogy, Vol. 8. Mineral. Soc. Am., Washington, D.C., pp. 1-68. Morse, J.W., Millero, F.J., Cornwell, J.C. and Rickard, D.T., 1987. The chemistry of the hydrogen sulphide and iron sulphide systems in natural waters. Earth-Sci. Rev., 24: 1-42. Rickard, D.T., 1973. Copper sulphide formation chemistry at low temperatures. Tschermaks Mineral. Petrogr. Mitt., 19: 60-76. Rickard, D.T., 1974. Kinetics and mechanism of the sulphidation ofgoethite. Am. J. Sci., 274: 941-952. Rickard, D.T., 1975. Kinetics and mechanisms of pyrite formation at low temperatures. Am. J. Sci., 275: 636652. Rickard, D.T., 1982. Reaction kinetics in low temperature ore formation: Principles and applications. Stockholm Contrib. Geol., 37: 215-236. Roberts, W.M.B., 1961. Formation of chalcopyrite by reaction between chalcocite and pyrrhotite in cold solution. Nature (London), 191: 560-562. Roberts, W.M.B., 1963. The low temperature synthesis in aqueous solution of chalcopyrite and bornite. Econ. Geol., 58: 52-61. Schouten, C., 1934. Structure and textures of synthetic replacements in "open space". Econ. Geol., 29: 611-658. Scott, S.D. and Kissin, S.A., 1982. Phase relations involving pyrrhotite below 350 °C. Econ. Geol., 77: 1739-1754. Shea, D. and Helz, G.R., 1988. The solubility of copper in sulphidic waters: Sulphide and polysulphide complexes in equilibrium with covellite. Geochim. Cosmochim. Acta, 52: 1815-1825. Shea, D. and Helz, G.R., 1989. Solubility product constants of covellite and a poorly crystalline copper sulphide precipitate at 298 K. Geochim. Cosmochim. Acta, 53: 229236. Sillitoe, R.H. and Clark, A.H., 1969. Copper and copperiron sulphides as the initial products of supergene oxi-

MECHANISMOF CHALCOPYRITEFORMATIONFROMFe-MONOSULPHIDES dation, Copiapo mining district, northern Chile. Am. Mineral., 54: 1684-1710. Stefan, J., 1890.0ber die Theorie der Eisbildung, insbesondere fiber die Eisbildung in Polarmeere. Sitzungsber. Akad. Wiss. Wien, Math.-Naturwiss. Kl., Abt. 1, 98: 965983. Szymanski, J.T., 1974. A refinement of the crystal structure of cubanite, CuFe2S3. Z. Kristallogr., 140: 218-239. Vaughan, D.J. and Craig, J.R., 1978. Mineral Chemistry of Metal Sulphides. Cambridge University Press, London, 493 pp. Walsh, C.A. and Rimstidt, J.D., 1986. Rates of reaction of

341

covellite and blaubleibender covellite with ferric iron at pH 2.0. Can. Mineral., 24: 35-44. Wuensch, B.J., 1974. Sulphide crystal chemistry. In: P.H. Ribbe (Editor), Sulphide Mineralogy. Mineral. Soc. Am. Short Course Notes, 1: W21-W27. Young, S.W. and Moore, N.P., 1916. Laboratory studies in sulphide ore enrichment, II. The formation of chalcopyrite by artificial replacement. Econ. Geol., 11: 574581. Zies, E.G., Allen, E.T. and Merwin, H.E., 1916. Some reactions involved in secondary copper sulphide enrichment. Econ. Geol., 11: 407-503.