Anodic dissolution of copper sulphides in sulphuric acid solution

HydrometaUurgy, 11 (1983) 195--206 Elsevier Science Publishers B.V., Amsterdam -- Printed in The Netherlands

195

ANODIC DISSOLUTION O F COPPER SULPHIDES IN SULPHURIC ACID SOLUTION. II. T H E ANODIC DECOMPOSITION O F CuS

E I L H A R D HILLRICHS and R O L F B E R T R A M InstitutfiirPhysikalischeund Theoretische Chemie, T U Braunschweig, Hans-Sommer-Str. 10, D-3300 Braunschweig (West Germany) (Received July 12, 1982; accepted in revisedform April 19, 1983)

ABSTRACT

Hillrichs, E. and Bertram, R., 1983. Anodic dissolution of copper sulphides in sulphuric acid solution. II. The anodic decomposition of CuS. HydrometaUurgy, 11: 195--206. The anodic decomposition of CuS in sulphuric acid solutions was investigated by means of cyclic voltammetry. The pH of the electrolyte and the scan rate were varied. Our results are interpreted using the experience of semiconductor electrochemistry and of passivation of pure metals. At potentials E < Enp (Enp = potential of pitting nucleation), the current/voltage curves were interpreted by the formation of metastable nonstoichiometric copper oxide or hydroxide. After the dissolution of these layers at the potential Enp , total decomposition of CuS starts. These results were confirmed by rest potential measurements and by scanning electron microscopy. INTRODUCTION Th e current/voltage characteristics o f CuS (covellite) differ f r o m those o f C u 2 - x S electrodes. The apparently passive behaviour o f CuS surfaces has n o t been discussed in detail. Cole (1972) and Mackinnon (1976) f o u n d no anodic d e c o m p o s i t i o n o f CuS in their fluidized bed cells. Wright (1971), Etienne (1967) and Peters (1977) r e p o r t e d a decom posi t i on at current densities i ~ I m A / m 2. Higher c ur r e nt densities lead t o an apparent passivation. Kato et al. (1972) and Bertram et al. (1980) supposed from their current/ voltage spectrum t h a t metastable CuO layers blocked o f f the anodic decomposition. In this paper we present detailed investigations o f the electrochemical behaviour o f CuS. We obtained the main results by cyclic v o l t a m m e t r y experiments. METHODS, TECHNIQUES AND MATERIALS The e q u i p m e n t f or cyclic v o l t a m m e t r y experiments are described in Part I (Hillrichs et al., 1983). Our 3.5 M calomel reference electrode (E) has a potential o f +0.252 V versus NHE. 0304-386X/83/$03.00

© 1983 Elsevier Science Publishers B.V.

196

The solid electrodes are prepared from pellets of natural covellite from Butte, Montana, USA. The pellets are m o u n t e d on a brass holder with silver and e m b e d d e d in an e p o x y resin (Scan Dia, Hagen, West Germany). The plane surface was polished with sandpaper of 1000 grade. The rest potential of the natural probes was in the range o f 0.26--0.30 V vs. 3.5 M CE. This indicates an equilibrium m C u : _ x S / n CuS with m ~ n. After a preanodization at E = 0.3 V the rest potential became constant within 10 minutes. We suppose that the surface composition was changed to an equilibrium CuS/S during this procedure. Synthetic CuS pellets prepared from analytically pure Cu and S show potentials in a range E = 0.30--0.325 V. These values are in good agreem e n t with those published by Koch et al. (1976). These rest potentials m a y be explained b y the phase equilibrium m CuS/n S with m >~ n. For kinetic investigations we used "natural" CuS, because the high porosity of synthetic CuS gave poor reproducibility o f the current/voltage curves. All experiments were carried o u t in sulphuric acid solutions at 20 -+ 1°C. RESULTS

Cyclic voltammograms of CuS (Fig. 1) are changed in structure significantly when the scan rate increases from 1 mV/s to 10 mVfs. In 1 M H2SO4 peak i

I

I

I

(c)

i / A m -2

30 2O 10 OI

I

J

1.0

2.0

3.0

I

i

I

E/V t

(b)

1.0 E I I"

% .~ 0.75

~tiI

0.5

, 0

0

1.0

s

2.0

2

, I

3.0

E/V

Fig. 1. Cyclic voltammograms of CuS; T = 293 K. (a) 1 M H2SO4; v = 1 mV/s. ( b ) - : I M H , S O , ; - - - :l--yMH2SO4/yMNa2SO4, pH3.25;vffi10mV/s.

197

C splitsinto two maxima (v = 10 mV/s). Peak D is sharp and the peak potential is nearly constant (Ep ~ 2.3 V). W h e n v increases peak E becomes broader. At scan rates v ~ 20 mV/s we recorded a broad unstructured wave at E ~ 1 V and analysis does not seem to be successful. The CuS surface changes its appearance (colour, topography) after anodization (Fig. 2, Table 1). The change of the colour was observed by a microscope after an interruption of the electrode process at characteristic potentials, which are given by the i--E curve (Fig. 3b). The variation of the potential E m a ( m a x i m u m anodic potential) in 3.5 M H2SO4 is shown in Fig. 3a. W e identified typical potentials according to the potential of pitting nucleation, Enp , and to the potential of criticalpitting nucleation, Ecp , (Szklarska-Smialowska, 1971; Sato, 1982). At potentials E ~ Enp S formation was observed in the apparently passive areas (Fig. 2).

Fig. 2. CuS surface after a n o d i z a t i o n in 1 M H2~O 4. (a) Ema = 0.8 V; v = 4 mV/s (potential ramp). (b) Ema = 2 V ( E m a ) Enp); v = 4 mV/s (potential ramp).

TABLE 1 Appearance of CuS surface after interruption of anodization at potential E E (V)

Appearance

ER 1 1.4 1.9 2.1

blue dark blue with brown areas black (porous) surface greenish-grey with yellow sulphur

198 (a)

.

.

.

.

.

i

E0.4 % 0.3 ¸

0.2

01

0.5

0.3

1.0

1.5

2.0

2.5

(b)

E/V

fl

/

'E

% 0.2

Enp ,,

0

0.5 '

10 ,.

1.5 '

2.0 ~

2~.5

E/V

Fig, 3. Variation of the maximum anode potential Ema. (a) CUS/3.5 M H~SO+; v = 4 mV/s. (b) CuS/1 M H2SO,; v : 4 mV/s.

The SEM photos are similar to those from anodized Cu2S surfaces. In 1 M H~SO+ the current density increases less when the potential returns from a maximum value Ema > Enp (Fig. 3b). The different structure of the i--E curve at reverse sweeps may be influenced by the consistency o f the product layer: in 1 M H2SO4 a solid yellow-grey product layer was found, while in 3.5 M H2SO+ a loose sulphur layer was formed. However, the current--voltege structure in 3.5 M H2SO4 at reverse sweeps is not understood in any detail.

199

The current efficiency was calculated for the reaction CuS -*

C u 2÷ +

S

+

2 e-

by analysin~ the Cu 2+ concentration in the solution by atomic absorption spectroscopy (AAS) (Hillrichs, 1976). We observed no gas evolution and we assumed no SO:`- formation, because in HC1 solutions no SO:,- was found. In 1 M H2SO4 the current efficiency decreases at potentials E < 1.8 V to around 80%. Therefore the whole loss of the current efficiency (20%) may be contributed to the formation of a copper oxide layer. Rest potential measurements

We analysed the open circuit potentials after anodization for characterization of the surface product layers. At low potentials (E < 1 V) the layer thickness limits the indication of the products by other techniques (X-ray reflection, electron beam microanalysis). The product layer, which was built up at higher potentials, was too inhomogeneous and porous, and the reproducibility of our results was poor. Typical E R --t curves of anodized CuS in 1 M H2SO4/0.1 M CuSO4 are given in Fig. 4. In Table 2 rest potentials are compared with standard potentials of suggested equilibria of the system Cu--S--H:O. The measured rest potentials of CuS in 1 M CuSO4/y M H2SO, (0.001 ~< y ~< 0.1) are nearly independent of y. For correlation of rest potentials to standard potential measurements in 1 M H2SO4 electrolytes (pH 0.1) are useful. 2.0 ERIV I

15 --CuO/Cu203

~.0

0.5-

--Cu20/Cu (OH)2 --Cu20 /CuO CuS/$

"--'~()

8"0

1½0

180

' 240

t/s

Fig. 4. ER--t decay of CuS in 1 M H~SO4/0.1 M CuSO 4 after anodization (holding time T~ 3h).

200 TABLE 2 E q u i l i b r i u m p o t e n t i a l s o f t h e s y s t e m C u - - S - - H 2 0 w i t h r e s p e c t t o r e a c t i o n s w i t h E ° > 0.3 V. R e s t p o t e n t i a l s a f t e r a n o d i z a t i o n : p o t e n t i a l range I: E R (t = 60 s) = 0.41 + 0.1 V (0.4 < E (V) < 0.8); p o t e n t i a l range II: E R (t ffi 60 s) = 0.59 + 0.03 V (0.9 < E (V) < 1.8) Equations

Equilibrium potential E ° (V vs. 3.5 M CE)

CuS .~' S + Cu :÷ + 2 eCu=O + H20 ~ 2 CuO + 2 I-I+ + 2 eCu20 + 3 H20 .~ 2 Cu(OH)= + 2 H + + 2 eCuS + H20 ,~ S + CuO + 2 I-I* + 2 eCuS + 2 H20 ~ S + CuO + 2 H ÷ + 2 e-

0.325 0.417 0.495 0.581 0.637

C u O ~ C u 2÷ + lh O 2 + 2 e2 C u ( O H ) 2 ~ C u 2 0 s + H 2 0 + 2 H + + 2 e2 C u O + H 2 0 ~ C u ~ O j + 2 H ÷ + 2 e-

0.744 1.326 1.396

The rest potentials, E R , of unhomogeneous CuS surfaces may be understood as mixed potentials. Under stationary conditions these mixed potentials may be compared to equilibrium potentials of a dominant reaction. The E R values at t = 60 s determined over the anodization potential range I (0.4 ~ E (V) ~ 0.8) are in good agreement with the Cu20/CuO normal potential. A theoretically expected pH dependence of 59 mV/pH (T = 298 K) of the Cu20/CuO equilibrium potential cannot be verified. We found circa 40 + 10 mV/pH in (1 - - y ) M H:SO4/y M Na2SO4 (0 ~ pH ~ 4). The ER values at an anodization potential range II (0.9 ~ E (V) ~ 1.8) agree with CuS/CuO and CuS/Cu(OH)2 standard potentials and no definite change with pH can be evaluated. This seems not to be surprising considering mixed potential theory and the constitution of the electrode surface as described before. At present an exact analysis of the copper oxide composition by rest potential measurements is impossible, because the rest potentials of non-stoichiometric copper oxide solid/solid equilibria are not known. The standard potential for the equilibrium n Cu20/m CuO with m = n = 1 may be expressed as CuOl-d with d = 0.33. Copper hydroxide layers must be included, but the difference between their equilibrium potentials and the oxide equilibrium potentials is too small to allow a distinction. Cu203 may be formed at the surface if the anodization potential is higher than the equilibrium potential E ° (CuO/Cu203) = 1.396 V and any potential losses across the product layer are neglected. After Bickl (1966), Cu203 at the surface may not be determined by rest potential analysis, because the exchange current density of the reaction CuO/Cu203 (E ° = 1.396 V) is much lower than that of the reaction Cu20/CuO (E° = 0.417 V). Therefore from rest potential measurements and microscopic surface observations we may assume a formation of copper-rich brownish oxide layer CuO,--d (d = 0.33) (anodization potential range 0.4 ~ E (V) ~ 0.8), which

201

m a y be oxidized to black CuO (anodization potential range 0.8 < E (V) < 1.8).

Cyclic voltammetry investigations As described above the CuS decomposition starts at 0.3 V and visual evidence of sulphur formation is found after anodization at 2.1 V. The decomposition o f the CuS bulk electrode m a y be hindered by oxide layers. The cyclic voltammograms o f CuS in (1 - - y ) M H2SO4/y M Na2SO4 solutions are influenced b y pH (Fig. l b ) , as expected. At low potentials the current density is lowered when the pH increases. The ascent to peak C (Ep = 1.5--1.9 V) shifts to lower potentials on increasing pH (dE/dpH = --0.03 V), whilst the potential of pitting nucleation, Enp , (dE/dpH = 0.04 V), the potential of critical pitting nucleation, Ecp, (dE/dpH = 0.03 V) and the peak potential of peak D (dEp/dpH = 0.03 V) shift to higher potentials.

Potential range A The influence of pH on the polarization curve at a range E < 0.8 V, which is nearly identical with range A from Fig. 1, is shown in Fig. 5. At v/> 100 mV/s the anodic cycle can be divided into two broad maxima (Ep(1) = 0.3(a)

i / A m -2

20

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~15 "._-

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,..!

(l~)

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/ / - , < ~ ~:.:............ ,. -..-y // . . . . ].~; /

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pH 3.25~ - ~ "

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pH o.87

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--pH 0.1

-15I

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0.'25 0'.5 0'.75 E/V 0 0.25 0.5 0.75 ElY Fig. 5. Cyclic voltammograms of CuS. (a) In 1 M H~SO4: . . . . : start in anodic d i r e c t i o n ; - : start in cathodic direction; (1): first cycle, (2): second cycle (identical); v = 4 mV/s. (b) In 1 --y M Na2SO,/y M H2SO4;second cycle; v = 10 mV/s. 0

202 0.4 V; Ep(2) > 0.5 V). The slope of the function ip(V 1/2) decreases with increasing pH for both maxima. Control of the current density by transport of holes in the bulk electrode may be excluded, because this is unusual for anodic decomposition of psemiconductors with high hole concentrations (~ 1022 holes per cubic centimetre in natural CuS (Shuey, 1975)). In agreement with Geriscber et al. (1968), the breaking of Cu--S bonds may be slow in comparison with transport processes. But a solid state reaction without oxide formations may not explain the pH influence. The rate determinant step of these reactions should not be a transport process in solution, because rotating disc experiments have no significant effect on current density. The decrease of the current density with increasing pH may be explained by the following model of parallel reactions: (a) a pH-independent decomposition starts: CuS + H20 ~ Cu2+(aq) + S + 2 e-

E(CuS/S) ~ 0.3 V

(1)

(b) in a concurrent reaction, formation of non-stoichiometric metastable oxide (hydroxide) layers may partially block off the decomposition reaction

(I): CuS + (1 - - d ) H20-~ C u O , - - d + S + 2(1 - - d ) H" + 2(1 - - d ) e-

(2a)

(d = 0.33: product composition is equal to Cu20/CuO equilibrium; E ° = 0.581 V) Cu8 + 2 H:O -~ Cu(OH)2 + 2 H ÷ + 2 e-

(2b)

(E ° -- 0.637 V) The chemical dissolution rate of the oxide (hydroxide) layers at the CuS anode surfaces, according to the equations CuO~- d + 2(1 - - d ) H ÷ ~- Cu2÷ + (1 - - d ) H20

(3a)

Cu(OH)2 + 2 H ÷ ~ Cu 2÷ + 2 H20

(3b)

depends on pH and should be low (Bickl, 1966). The thickness of a CuO,-d product layer (d = 0.33) which is formed by reaction (2a), is estimated from the integrated i--E curve (Fig. 5) to be circa 1.0 nm (supposing a current efficiency of about 50%). Progress of the overall reaction should lead to a closed coverage with oxide and further reactions for the CuS bulk electrode may be blocked. Only when a partial current density according to reaction (2a) or the whole current density is low enough and the rate of formation of CuO,--d becomes low compared to the rate of chemical dissolution (3a), will decomposition of CuS proceed. This may be in accordance with Etienne (1967), Wright (1971) and Peters (1977), who reported a non-passivated CuS decomposition at very low current densities only (i < I mA/m 2) and a sharp potential rise at higher current densities in galvanostatic experiments.

203 Potential range B and C A separate study of peak B was n o t successful. Peak C involves B as a poorly defined prewave at rates v/> 10 mV/s. At potentials in the range o f peak C the surface colour changes from brownish to black and rest potential increases to 0.62 V. This indicates an oxidation of a CuOl--d surface layer to CuO. Peak C itself (Fig. l b ) can be divided into two peaks (Ep(1) = 1.5 V; Ep(2) = 1.8 V) in 1 M H2SO4; in solutions with a higher pH we observed only a prewave. Assignment of these peaks to oxide or hydroxide formation seems to be impossible at the present. Peak C is separated after potentiostatic pre-anodization at a potential E = 1.2 V, where the formation of a closed CuOl--d surface layer is completed. Based on thermodynamic considerations (Pourbaix, 1976) and considering a nearly pH-independent stationary current density at pre-anodization for a holding time greater than 120 s, we assume for stationary and potentiostatic conditions that the surface oxide composition is mainly determined by potential and that pH has less influence. The change of the separated current voltage curve with pH of the solution is illustrated in Fig. 6. The i--E response changes significantly: in solutions with pH 3.25 we obtained a peak, whereas in 1 M H2SO4 (pH 0.1) we obtained a limited current. The same change of the wave form is discussed by Casadio (1976) in the case of metal dissolution with adsorption of a passive layer. E [.V]

//A\ l\OrnVis

1.9 E

1.2

/

<~

\~

/s

//

o.o,-

;/ \ \ ..........//

/ ..;/'!

I

I

1.3

1,5

/

I

1.7 E / V

1

19

Fig. 6. Ourvei--EofCuSin(1--y)MNa2SO4;v=lOmV/s; '- : p H 3 . 2 5 ; . . . . : pH . . . . : pH 0.87 ; . . . . : pH 0.1 ; holding potential E~ = 1.2 V; holding time r ffi 180 s.

2.4;

204

The linear functions ip(V) (pH 3.25) and iE=l.9 V(v) (pH 3.25) (Fig. 7) are a diagnostic criterion for an adsorption-controlled reaction. We assume copper oxide (CuO) formation: CuOl--d + d H:O ~ CuOl--d + d OH s + d H ÷

(4a)

CuO,_d + d OHs ~ CuO + d H ÷ + 2d e-

(4b)

CuO,_ d + d H 2 0 ~ C u O + 2 d H

E <~ 3-

"--

x ip pH 3.25 • i (=1.9V) pH0.1 o i (=1.9V) pHO.9

(4a) + (4b)

÷+2de-

/• " -

-

/

2-

200

3[)0 I 400 v / m V s -1

Fig. 7. Current density at peak maximum (or at limited current) in dependence of the scan rate, v (see Fig. 6).

Reaction (4b) may determine the rate. The exact composition of the copper oxide is n o t clear. The assumption of the existence of an oxygen concentration gradient within this layer is realistic. The peak potential of peak C (Ep = 1.8 V in 1 M H:SO,, v = 10 mV/s) is higher than the equilibrium potential E(CuO/Cu203) = 1.396 V. If the potential drop in the product layer is small, formation of a Cu203 surface layer m a y be possible. Another reason for the difficulty of a kinetic interpretation of this part of the current/voltage curve is the overlapping of this peak with an anodic shift on decreasing pH and the following anodic oxide dissolution (peak D) with a cathodic shift on decreasing pH.

Peak D, Enp and Ecp The anodic dissolution o f the CuO layer begins at a potential Enp (Enp = 1.975 V in 1 M H~SO4). The total decomposition of the bulk electrode can start when the surface is partially free o f oxide. The activation is indicated by several results: (a) the current/voltage curve shows a course which is typical for the activation o f passivated surfaces (Sato, 1982). Enp and peak D (Fig. l b ) shift to lower potential when the pH of the solution decreases.

205

(b) the rest potential after anodization is lower than in the passive region (Fig. 4). (c) SEM photos show sulphur formation at the whole surface; the influence of pH on the CuO dissolution peak D may be understood by a reaction mechanism given by Bickl (1966) for CuO bulk electrodes: 20G2- + 2 H ÷

20Hs

(5a)

20Hs + 2 p* -+ 20Had

(5b)

20Had + 2 p÷ ~ O2 + 2 H÷

(5c)

2 CUG2÷ + H20 ~ 2 Cu2+ H20

(5d)

~

2CuG2 + + 2 0 6 2 - + H 2 0 + 4 p ÷ ~ 2 C u2+H20+O2

(5)

The resultingeqn. (5) can be simplifiedto CuO -+ Cu2÷ + lh O2 + 2 e-

(6)

Bickl (1966) explained that the forward step of eqn. (5a) determines the rate of this process. A theoretical shift of the start of this reaction (59 mV/pH) may be calculated from a combination of eqns. (5a) and (5b); we found circa 30 mV/pH. These discrepancies are cited by Vetter (1960) for O2 evolution reactions, too, and are accepted for electrocatalytic mechanisms. Holes (p÷) participate in the mechanism, because CuO is a p-semiconducting material. This active state remains at potentials E > Ecp. At potentials E < Ecp a growth of oxide again stops the decomposition. Therefore the shift dEcp is about 30 mV/pH (Fig. lb). CONCLUSIONS

The anodic decomposition may be divided into several sections. One kind of consideration is given by characteristic potential regions (A, B, C, D and E in Fig. la), which are analysed from cyclic voltammograms. However, the reaction according to range B is n o t clear. At potentials E < 0.8 V (range A) the total decomposition of CuS is stopped by a thin layer of copper oxide with a proposed composition CuOt--d (d = 0.33); high current densities and increasing pH favour the formation of this layer. The rest potentials after anodization (E < 0.8 V) are compared with the equilibrium potential E(Cu20/CuO) = 0.417 V. Brown areas at the electrode surface after anodization at E = 1.4 V confirm the existence of copper-rich CuOt--d product layers. In a second potential region oxidation of the CuOt--d surface to CuO proceeds. The i--E curve (peak C) shifts with increasing pH in cathodic direction and is determined by adsorption kinetics, which is shown by a linear function ip(V). The rest potential increases up to 0.62 V and the product layer shows a black colour. At a characteristic potential Enp the anodic dissolution of the CuO layer (peak D) starts. From an anodic shift of Enp and Ep(D) with increasing pH

206 we deduced an electrocatalytic mechanism as k n o w n in the case of CuO bulk electrodes. The covering o f the anode with sulphur indicates t h a t the total decomposition o f CuS has started in a subsequent process. The Cu 2÷ flux is mainly hindered by adherent sulphur. A further activation occurs when surface-active anions such as CI-, C104or NO3- are added to the solution (Bertram et al., 1981). The reaction model presented includes some uncertainties, which depend on investigation m e t h o d (rest potential measurements) and on the kind of reaction (thin metastable product layers, current efficiencies of about 80%). A difficulty in characterizing the products is the change in composition after the anodization potential is switched off. In situ analyses are needed for exact determination o f product layer composition. ACKNOWLEDGEMENT These investigations were supported by the DFG, Bonn--Bad Godesberg, and by the Land Niedersachsen, West Germany. REFERENCES Bertram, R., Greulich, H., Hillrichs, E. and M/iller, R., 1980. Erzmetall, 33(2): 88. Bertram, R. and HiUrichs, E., 1981. J. Electroanal, Chem., 129: 365. Bickl, H., 1966. Thesis, TH Miinchen. Biegler, T. and Swift, D.A., 1976/77. Hydrometallurgy, 2: 335. Casadio, S., 1976. J. Electroanal. Chem., 72: 243. Cole, S.H., 1972. Thesis, University of Columbia. Etienne, A., 1967. Thesis, University of British Columbia. Gerischer, H., 1968. Electrochim. Acta, 13: 1329. Hillriehs, E., 1976. Diplomarbeit, TU Braunschweig, W. Germany. Hillrichs, E., 1981. Thesis, TU Braunschweig. Hillrichs, E. and Bertram, R., 1983. Hydrometallurgy, 11: 181. Kato, T. and Oki, T., 1972. Denki Kagaku, 40(9): 670. Koch, D.F.A. and McIntyre, R.J., 1976. J. Electroanal. Chem., 71: 285. Macdonald, D.D., 1977. Transient Techniques in Electrochemistry. Plenum Press, New York. Mackinnon, D.J., 1976. Hydrometallurgy, 2: 65. Peters, E., 1977. In: Bockris, J.O.M., Rand, D.A.F. and Welch, R.J. (Eds.), Trends in Electrochemistry. Plenum Press, New York, p. 267. Pourbaix, M., 1973. Lectures on Electrochemical Corrosion. Plenum Press, New York. Sato, N., 1982. J. Electrochem. Soc., 129(2): 255. Shuey, R.T., 1975. Semiconductor Ore Minerals, Elsevier, Amsterdam. Szklarska-Smialowska, Z. and Janik-Czachor, H., 1971. Corros. Sci., 11: 901. Vetter, J.J., 1960. Elektrochemische Kinetik, Springer Verlag, Berlin. Wright, J., 1971. Bull. Aust. Mineral. Develop. Lab., 12: 47.