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Electrochemical reduction of the corrosion products formed on copper surface in alkaline-sulphide solutions
Journal of Alloys and Compounds 432 (2007) 205–210
Electrochemical reduction of the corrosion products formed on copper surface in alkaline-sulphide solutions S.M. Abd El Haleem, E.E. Abd El Aal ∗ Chemistry Department, Faculty of Science, Zagazig University, Zagazig, Egypt Received 4 April 2006; accepted 27 May 2006 Available online 5 July 2006
Abstract Electroreduction of the preformed reaction products on copper in highly oxygenated NaOH solutions, on one hand, and NaOH solutions in presence of S2− , on the other hand, is followed galvanostatically. In sulphide free solutions, two cathodic steps corresponding to the reduction of Cu2 O and Cu(OH)2 , receptively, before H2 evolution commences were reported. The quantity of electricity consumed during the reduction of both species is found to depend on both the concentration of NaOH and the immersion time. The presence of S2− ions, at any concentration causes the complete elimination of the second step corresponding to the reduction of Cu(OH)2 and the partial disappearance of the step corresponding to the reduction of Cu2 O. The reduction of the preformed Cu2 S film occurs in a single reduction step. The quantity of electricity consumed along this step depends on the thickness of the sulphide film. © 2006 Elsevier B.V. All rights reserved. Keywords: Copper; Sodium hydroxide; Sulphide; Passivation; Corrosion; Electroreduction and galvanostatic polarization
1. Introduction
Cu2 S has been formed as an insoluble film on a Cu electrode in aqueous sulphide solutions according to [4]:
Several interesting studies have been concerned with the electrochemical behaviour of oxide and sulphide minerals from which metal sulphides have received the greatest interest because they are recognized as the main sources of nonferrous metals [1,2]. Many electrochemical studies on copper sulphide minerals have been carried out which aimed at an understanding their behaviour in the different stages of mineral processing [3,4]. Stoichiometric and non-stoichiometric copper sulphide films have been widely used in dielectric metallization [5], in solar cell technology [6] and as ion specific electrode [7]. When soluble sulphides are present in potable water or seawater, a thick black, poorly adherent scale forms on copper or brass surface [8]. This scale in composed mainly of Cu2 S although CuS, Cu2 O and non-stoichiometric copper sulphide species such as Cu1.8 S have also been reported [9].
2Cu + HS− + OH− Cu2 S (film) + H2 O + 2e
∗
Corresponding author. Tel.: +20 55 2329786; fax: +20 55 2345452. E-mail address: emad [email protected] (E.E. Abd El Aal).
0925-8388/$ – see front matter © 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.jallcom.2006.05.099
(1)
This reaction is expressed in terms of the HS− species that predominates in alkaline sulphide solutions of high pH [10]. In previous works, the open circuit potential of the Cu electrode was followed in NaOH solution free from and containing increasing concentrations of S2− . A mechanism was proposed for the formation of CuS film involved the formation of CuHS as an intermediate step [11]. The behaviour was further investigated using both the galvanostatic and cyclic voltammetry techniques [12]. In S2− free NaOH solutions, the anodic polarization reveals the formation of Cu2 O, CuO and soluble CuO2 2− , respectively. The first two oxides are reduced during the cathodic process. However, in presence of S2− ions, several Cu sulfide and oxide species are formed along the whole anodic polarization regions with a corresponding reduction peaks on cathodic polarization. Cu2 S and CuS films are assumed to form directly from metallic Cu or by a mechanism involving the adsorption of HS− and/or S2− on Cu surface. The present investigation aims to through new lights on the characteristics of the reaction products preformed on the Cu
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surface in NaOH solutions free from and containing S2− ions. This is followed through the electrochemical reduction of these species using the galvanostatic techniques. 2. Experimental Spectroscopically pure Cu electrode supplied by Johnson and Matthay (England) having the dimensions 3.92 cm × 3.80 cm × 0.2 cm. These were fixed to a borosilicate glass tube with epoxy resin so that the total exposed surface area was 32.72 cm2 . Electrical contact was achieved through a copper wire soldered to the end of the electrode not exposed to the solution. Before being used, the electrode surface was polished by 0-, 00- and 000-grades emery papers until it appeared free of scratches and other defects. Then, it was rinsed with acetone and finally washed with twice distilled water. The copper electrodes were first equilibrated with 100 ml oxygenated NaOH solutions of different concentrations devoid of and containing increasing concentrations of Na2 S, for different times of immersion. At the end of each time interval, the Cu electrode was quickly and carefully washed with twice distilled water, without affecting the adherent reaction products. The electrode was directly immersed in a polarization cell like that previously described [13,14], containing deareated NaOH solution of the same concentration as the test equilibration solution, under a constant cathodic current density of 0.03 mA/cm2 . The cell has a double wall jacket through which water at the adjusted temperature was circulated. The potential of the Cu electrode was then followed with respect to saturated calomel electrode (SCE) as a function of time till the H2 evolution reaction on a chart recorder type R-11 Shimadzu, Japan. Deareation of the alkali solution is achieved through bubbling highly purified N2 strongly through the solution 30 min before the immersion of the electrode and just before polarization started. Sodium hydroxide solutions of different concentrations were prepared by dilution of a BDH carbonate-free solution, kept in plastic bottles. The concentration of each solution was checked by titration against standard HCl solution, using methyl orange as indicator. Sodium hydroxide solutions containing different concentrations of S2− ions were similarly prepared from a stock solution of approximately 1 M Na2 S (Koch-Light Lab., England). Experiments were carried out at a constant temperature 25 ± 0.1 ◦ C, which was controlled using an ultra thermostat Type Polyscience (USA). Each run was carried out in a fresh test solution and with a newly polished electrode.
3. Results and discussion 3.1. Electrolytic reduction of the preformed oxygenated copper compounds Copper electrodes with specified surface area were previously equilibrated for different times of immersion in highly oxygenated NaOH solutions of two different concentrations. At the end of each time interval, the electrode was thoroughly washed with running distilled water and immersed directly in the electrolytic cell, containing deareated NaOH solution of the same concentration as the test electrolyte, under a constant cathodic current density of 0.03 mA/cm2 . The potential of the working electrode was then followed with time till the hydrogen evolution reaction. The curves of Figs. 1 and 2 represent such behaviour of the copper electrode in 0.01 and 0.10 M NaOH solutions, respectively, for different times ranged between 30 min and 6 h. The galvanostatic potential–time curves of Figs. 1 and 2 are characterized by the presence of two well-defined cathodic reduction steps before the hydrogen evolution reaction commences. In Table 1 the values of the starting potentials (eH ) of these steps are grouped.
Fig. 1. Galvanostatic potential–time curves for copper electrodes previously equilibrated for (1) 30, (2) 60, (3) 180, and (4) 360 min in oxygenated 0.01 M NaOH solutions under a cathodic current density 0.03 mA/cm2 .
Fig. 2. Galvanostatic potential–time curves for copper electrodes previously equilibrated for (1) 30, (2) 60, (3) 180, and (4) 360 min in oxygenated 0.1 M NaOH solutions under a cathodic current density 0.03 mA/cm2 .
These two cathodic polarization steps are expected to correspond to the reduction of Cu2 O and Cu(OH)2 , respectively [11,12,15]. Inspection of the potentials reported in Table 1 reveals, however, a considerable displacement of the potentials of these steps from their expected values towards more negative potentials. Similar behaviour has reported by different authors Table 1 Starting potentials (eH ) of the reduction steps observed during the cathodic reduction of the oxygenated copper compounds on the copper surface Immersion time (min)
0.01 M NaOH
0.10 M NaOH
First step (V)
Second step (V)
First step (V)
Second step (V)
30 60 180 360
−0.290 −0.270 −0.270 −0.290
−0.480 −0.460 −0.430 −0.460
−0.28 −0.24 −0.135 −0.160
−0.590 −0.560 −0.470 −0.510
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[16,17]. The reduction of both Cu2 O and Cu(OH)2 films takes place at potentials considerably far removed from their formation values according to: Cu2 O + H2 O + 2e → 2Cu + 2OH−
(2)
and Cu(OH)2 + 2e → Cu + 2OH−
(3)
The irreversibility of these reactions is intrinsic and related to the oxides themselves. Shams E1 Din and Abd El Wahab [15] have explained this behaviour in view of the lattice structure of both oxides. Cuprous oxide is regarded as having cation defects resulting in anion excess and positive holes. The latter is assumed to be localized with the formation of cupric ions. This arrangement is characteristic for p-type semiconductor oxides in which electrons are easily donated and difficulty accepted. Little is known regarding the type of electrical conductivity of crystalline Cu(OH)2 . However, CuO is known, to be an intrinsic semiconductor [18]. Adsorption of oxygen, dry and moist air, on the surface of cuprous oxides leads to the formation of CuO under certain conditions, increases considerably the p-conductivity of the oxide. Therefore, the electrochemical reduction of the oxides will not occur at their expected thermodynamic potentials, but at potential more negative than their formation values. The cathodic potential displacement depends on the potential to which the electrode was previously anodized [15], or on the time of equilibration of the electrode in solution, as shown in Table 1. Reduction of Cu2 O and CuO oxides does not proceed as expected according to: Cu(OH)2 → Cu2 O → Cu, but rather Cu2 O → Cu and Cu(OH)2 → Cu [11,12,15]. This, together with the off shooting of the potential towards negative values observed at the start of the second reduction step indicates that the cathodic process starts at the metal/oxide interface. When Cu2 O is totally reduced, the reduction of Cu(OH)2 proceeds in one single step at a potential more negative from the reversible Cu/Cu(OH)2 /OH− value. The reaction occurring during this stage is, however, complex since the system is unstable being directly transferred to the corresponding Cu2 O/Cu(OH)2 /OH− [16]. Comparison of the quantity of electricity consumed in the cathodic reduction of both oxides in 0.01 and 0.1 M NaOH solutions is of interest. In 0.01 M solution, the quantity of electricity consumed during the first cathodic step, Q1 , is markedly smaller in comparison with that consumed during the second cathodic step, Q2 . However, the reverse is true in the concentrated alkali solution, Q1 is much greater than that of Q2 . This could be attributed to the dissolution of part of the oxides in the concentrated solution. In the more dilute solution, the solubilities of the oxides seem to be at a minimum [17]. The relation between the quantities of electricity consumed during both cathodic steps with the immersion time, t, is best represented on a double logarithmic scale as is shown in Fig. 3A and B. Straight line relationships are obtained satisfying the equation: log Q = a1 + b1 log t
(4)
Fig. 3. Variation of quantity of electricity consumed during the reduction of (A) Cu2 O and (B) Cu(OH)2 in () 0.01 and () 0.10 M NaOH solutions as function of the immersion time.
where a1 and b1 are constants. The constant b1 is always the same. A behaviour and amount to which could be explained on the basis that the way by which both oxides are electroreduced is always the same. Electroreduction of both oxides involves a twoelectron transfer mechanism (Reactions (2) and (3)). The constant a1 represent, however the logarithm of the quantity of electricity (log Q) consumed at a time of immersion equals 1 min. The value of a1 is found to depend on the alkali concentration, on one hand, and on the type of oxide electroreduced, on the other. 3.2. Electrolytic reduction of the preformed cuprous sulphide Thin films of copper sulphide previously formed during the equilibration of the copper electrode of 32.72 cm2 surface area in oxygenated 0.01 and 0.10 M NaOH solutions, containing increasing concentrations of Na2 S, for different immersion times, are electroreduced galvanostatically in deoxygenated NaOH solutions of the same concentrations, at a current density of 0.03 mA/cm2 . The curves of Fig. 4 represents the galvanostatic E–time plots for the reduction of Cu2 S, previously formed in oxygenated 0.01 M NaOH in presence of 0.05 M Na2 S at different immersion times. The curves of Fig. 5 shows the same behaviour of the reduction of Cu2 S films formed in oxygenated 0.1 M NaOH in presence of 0.1 M Na2 S. Similar curves were also obtained in the two alkali concentrations in the presence of varying concentrations of Na2 S. Inspection of the galvanostatic potential–time curves of Figs. 4 and 5, reported, respectively, in 0.01 and 0.1 M NaOH solutions, in presence of increasing concentrations of Na2 S reveals some interesting features which deserve mentioning. Thus, the presence of S2− ions, at any time of immersion, causes the following: (i) The partial disappearance of the first reduction step reported in S2− free NaOH solution which corresponds to the reduction of Cu2 O.
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electricity consumed during this step depends markedly on the concentration of both NaOH and S2− ion.
Fig. 4. Galvanostatic potential–time curves for the reduction of copper sulphide films formed during the equilibration of the copper electrode in 0.01 M NaOH + 0.05 M Na2 S for different times: (1) 30, (2) 60, (3) 180, and (4) 360 min under a cathodic current density 0.03 mA/cm2 .
(ii) The complete disappearance of the second reduction step which corresponds to the reduction of Cu(OH)2 . (iii) The appearance of a new cathodic reduction step at a potential ∼−0.81 V (NHE), which is assumed to correspond to the reduction of the preformed Cu2 S film. The quantity of
Fig. 5. Galvanostatic potential–time curves for the reduction of copper sulphide films formed during the equilibration of the copper electrode in 0.1 M NaOH + 0.1 M Na2 S for different times: (1) 30, (2) 60, (3) 180, and (4) 360 min under a cathodic current density 0.03 mA/cm2 .
In one and the same solution composition, typical galvanostatic reduction curves for thin films (first-order colours), medium films (second-order colours) and thick films (third, fourth and fifth-order colour) are obtained with increasing time of equilibration. After a short initial period during which the initial potential falls rapidly to about −0.9 V (SCE), the sulphide reduction precedes over a potential range of about 0.05–0.15 V depending on the immersion time. The quantity of electricity consumed along this step depends on the thickness of the preformed Cu2 S film. During reduction, the cathode surface was observed visually through the side of the cell. First-order films showed interference colours corresponding to decreasing film thickness as the reduction proceeded. At the completion of reduction, the surface appeared almost identical to the original pre-immersion abraded surface on which the film had been formed. Thicker films showed, however, a different behaviour. Thus, in the early stages of reduction, the surface was gradually decreased in reflectivity, sometimes exhibiting patches of earlier interference colours. With further reduction, the surface took on a uniform pale brown appearance. When the circuit to the cathode was broken, this brown material gradually darkened to black. The greater part of it was loosely bound to the cathode, and could be swept into the solution. Hoar and Stockbridge [19] assumed that the overpotential for hydrogen ion reduction on reducing cuprous sulphide is the same as that on reduced copper. They reported that, at 0.03 mA/cm2 cuprous sulphide reductions at about 85% current efficiency will be followed by hydrogen ion-reduction at 100% efficiency. However, under the present experimental conditions, Figs. 4 and 5 and
Fig. 6. The relation between the quantity of electricity, Q, consumed during the reduction of copper sulphide and the immersion time, t, in 0.01 M NaOH solutions in the presence of increasing Na2 S concentrations: (1) 0.01, (2) 0.05, (3) 0.1, and (4) 0.5 M.
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209
catholyte or remain as isolated particles on the surface of the cathode. While these copper specks remain of colloidal size, interference colours, indicating diminishing thickness, would be shown by the remaining adherent film. When the film surface becomes flowed, by cracks or voids formed during the formation of large specks of copper, some hydrogen ion may be reduced within or at the base of cracks, the evolved gas detaching some unreduced film material from the cathode. Once out of electrical contact with the cathode, this film cannot be reduced. Thus, although the premature reduction of hydrogen ions lowers the current efficiency of sulphide reduction, the consequent increase in the reduction time is probably more than counterbalanced by failure of separated film material to reduce. The production of loose material experimentally observed is thus well interpreted as the cause of premature attainment of the end-point, and results in the observed under-estimation of thick sulphide films. 5. Conclusions Fig. 7. The relation between the quantity of electricity, Q, consumed during the reduction of copper sulphide and the immersion time, t, in 0.1 M NaOH solutions in the presence of increasing Na2 S concentrations: (1) 0.01, (2) 0.05, (3) 0.1, and (4) 0.5 M.
the like, the quantity of electricity (time in seconds till the beginning of H2 evolution) consumed during the reduction of cuprous sulphide films varies with the immersion time, t, Figs. 6 and 7, irrespective of the concentration of both NaOH and Na2 S according to: log Q = a2 + b2 log t
(5)
where a2 and b2 are constants, which vary irregularly with the composition of the solution. 4. Mechanism of cuprous sulphide reduction At room temperature, cuprous sulphide is a metal-deficit conductor, the electron conductivity being due to the presence of positive holes and the ion conductivity to vacant cation sites [14]. Although the bulk of the film is certainly non-stoichiometric cuprous sulphide [20], the metal deficit being the greater the more rapid the rate of its formation. The layer next to the metal is probably nearly stoichiometric [21]. Here, the electron and especially the ion conductivity are much lower than in the bulk. This will militate against cuprous ions being reduced at or near the metal/film interface, where ion movement will be especially difficult, although some reduction in this region may occur [22]. However, it is much more likely that electrons pass through the nearly stoichiometric layer, perhaps by tunnel effect, reduce cuprous ions in the non-stoichiometric outer layer. These can diffuse fairly readily, gradually forming specks of metal, by capturing further electrons coming through the film, encourage further cuprous ions to migrate to, and plate out, on them. As the specks grow, some of them lose physical and electrical contact with the film, and either move off into the
(1) Reduction of copper oxides preformed in solutions of oxygenated NaOH solution depends on both the electrolyte concentration and time of immersion. (2) The preformed oxides are reduced in two steps corresponding, respectively, to the reduction of Cu2 O and Cu(OH)2 to metallic Cu. (3) In presence of S2− ions, Cu2 S is formed on the expense of the absence of Cu(OH)2 and the partial formation of Cu2 O. (4) Cu2 S is reduced along one reduction step, with the quantity of electricity consumed, depends on the thickness of the preformed sulphide. (5) A mechanism is proposed for the electroreduction of Cu2 S. References [1] R.S. McMillan, D.J. Mackinnon, J.E. Dutrizac, J. Appl. Electrochem. 12 (1982) 743. [2] E.M. Arce, I. Gonz¯alez, Int. J. Miner. Process. 67 (2002) 17. [3] A.N. Buckly, R. Wood, J. Electroanal. Chem. 357 (1993) 387. [4] P. Velasquez, D. Leinen, J. Pascual, J.R. Ramos-Barrado, R. Cordova, H. Gomez, R. Schrebler, J. Electroanal. Chem. 510 (2001) 20. [5] L.I. Stepanova, T.V. Mozolevskaya, O.G. Purovskaya, Met. Finish. 101 (2003) 18. [6] M.I. Schimmel, N.R. De Tacconi, K. Rajeshwar, J. Electroanal. Chem. 453 (1998) 187. [7] A. Zirino, R. De Marco, I. Rivera, B. Pejcic, Electroanalysis 14 (2002) 493. [8] S. Jacobs, M. Edwards, Water Res. 34 (2000) 2798. [9] M.B. McNeil, A.L. Anos, T.L. Woods, Corrosion 49 (1993) 755. [10] A.M. Shams El Din, J.M. Abd El Kader, F.M. Abd El Wahab, T.M.H. Saber, A.A. El Azhary, A.A. El Waraky, Electrochim. Acta 30 (1985) 461. [11] S.M. Abd El Haleem, E.E. Abd El Aal, A.M. Atia, Corrosion 61 (2005) 838. [12] S.M. Abd El Haleem, E.E. Abd El Aal, Corrosion 62 (2006) 121. [13] E.E. Abd El Aal, Bull. Soc. Chim. Fr. 128 (1991) 351. [14] E.E. Abd El Aal, J. Power Sources 75 (1998) 36. [15] A.M. Shams El Din, F.M. Abd El Wahab, Electrochim. Acta 9 (1964) 113.
210 [16] [17] [18] [19]
S.M. Abd El Haleem, E.E. Abd El Aal / Journal of Alloys and Compounds 432 (2007) 205–210 S.E.S. El Wakkad, S.H. Emara, J. Chem. Soc. (1953) 3508. R. Glicksman, C.K. Morehouse, J. Electrochem. Soc. 104 (1957) 589. C.A. Hogorth, Z. Phys. Chem. 198 (1951) 30. T.P. Hoar, C.D. Stockbridge, Electrochim. Acta 3 (1960) 94.
[20] T.P. Hoar, A.J.P. Tucker, J. Inst. Met. 81 (1953) 665. [21] V.A. Kozheurov, Colloid J. 5 (1939) 45. [22] W. Jaenicke, E.M. Khairy, W. Schaefer, J. Phys. Chem. 70 (1966) 421.
Electrochemical reduction of the corrosion products formed on copper surface in alkaline-sulphide solutions S.M. Abd El Haleem, E.E. Abd El Aal ∗ Chemistry Department, Faculty of Science, Zagazig University, Zagazig, Egypt Received 4 April 2006; accepted 27 May 2006 Available online 5 July 2006
Abstract Electroreduction of the preformed reaction products on copper in highly oxygenated NaOH solutions, on one hand, and NaOH solutions in presence of S2− , on the other hand, is followed galvanostatically. In sulphide free solutions, two cathodic steps corresponding to the reduction of Cu2 O and Cu(OH)2 , receptively, before H2 evolution commences were reported. The quantity of electricity consumed during the reduction of both species is found to depend on both the concentration of NaOH and the immersion time. The presence of S2− ions, at any concentration causes the complete elimination of the second step corresponding to the reduction of Cu(OH)2 and the partial disappearance of the step corresponding to the reduction of Cu2 O. The reduction of the preformed Cu2 S film occurs in a single reduction step. The quantity of electricity consumed along this step depends on the thickness of the sulphide film. © 2006 Elsevier B.V. All rights reserved. Keywords: Copper; Sodium hydroxide; Sulphide; Passivation; Corrosion; Electroreduction and galvanostatic polarization
1. Introduction
Cu2 S has been formed as an insoluble film on a Cu electrode in aqueous sulphide solutions according to [4]:
Several interesting studies have been concerned with the electrochemical behaviour of oxide and sulphide minerals from which metal sulphides have received the greatest interest because they are recognized as the main sources of nonferrous metals [1,2]. Many electrochemical studies on copper sulphide minerals have been carried out which aimed at an understanding their behaviour in the different stages of mineral processing [3,4]. Stoichiometric and non-stoichiometric copper sulphide films have been widely used in dielectric metallization [5], in solar cell technology [6] and as ion specific electrode [7]. When soluble sulphides are present in potable water or seawater, a thick black, poorly adherent scale forms on copper or brass surface [8]. This scale in composed mainly of Cu2 S although CuS, Cu2 O and non-stoichiometric copper sulphide species such as Cu1.8 S have also been reported [9].
2Cu + HS− + OH− Cu2 S (film) + H2 O + 2e
∗
Corresponding author. Tel.: +20 55 2329786; fax: +20 55 2345452. E-mail address: emad [email protected] (E.E. Abd El Aal).
0925-8388/$ – see front matter © 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.jallcom.2006.05.099
(1)
This reaction is expressed in terms of the HS− species that predominates in alkaline sulphide solutions of high pH [10]. In previous works, the open circuit potential of the Cu electrode was followed in NaOH solution free from and containing increasing concentrations of S2− . A mechanism was proposed for the formation of CuS film involved the formation of CuHS as an intermediate step [11]. The behaviour was further investigated using both the galvanostatic and cyclic voltammetry techniques [12]. In S2− free NaOH solutions, the anodic polarization reveals the formation of Cu2 O, CuO and soluble CuO2 2− , respectively. The first two oxides are reduced during the cathodic process. However, in presence of S2− ions, several Cu sulfide and oxide species are formed along the whole anodic polarization regions with a corresponding reduction peaks on cathodic polarization. Cu2 S and CuS films are assumed to form directly from metallic Cu or by a mechanism involving the adsorption of HS− and/or S2− on Cu surface. The present investigation aims to through new lights on the characteristics of the reaction products preformed on the Cu
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surface in NaOH solutions free from and containing S2− ions. This is followed through the electrochemical reduction of these species using the galvanostatic techniques. 2. Experimental Spectroscopically pure Cu electrode supplied by Johnson and Matthay (England) having the dimensions 3.92 cm × 3.80 cm × 0.2 cm. These were fixed to a borosilicate glass tube with epoxy resin so that the total exposed surface area was 32.72 cm2 . Electrical contact was achieved through a copper wire soldered to the end of the electrode not exposed to the solution. Before being used, the electrode surface was polished by 0-, 00- and 000-grades emery papers until it appeared free of scratches and other defects. Then, it was rinsed with acetone and finally washed with twice distilled water. The copper electrodes were first equilibrated with 100 ml oxygenated NaOH solutions of different concentrations devoid of and containing increasing concentrations of Na2 S, for different times of immersion. At the end of each time interval, the Cu electrode was quickly and carefully washed with twice distilled water, without affecting the adherent reaction products. The electrode was directly immersed in a polarization cell like that previously described [13,14], containing deareated NaOH solution of the same concentration as the test equilibration solution, under a constant cathodic current density of 0.03 mA/cm2 . The cell has a double wall jacket through which water at the adjusted temperature was circulated. The potential of the Cu electrode was then followed with respect to saturated calomel electrode (SCE) as a function of time till the H2 evolution reaction on a chart recorder type R-11 Shimadzu, Japan. Deareation of the alkali solution is achieved through bubbling highly purified N2 strongly through the solution 30 min before the immersion of the electrode and just before polarization started. Sodium hydroxide solutions of different concentrations were prepared by dilution of a BDH carbonate-free solution, kept in plastic bottles. The concentration of each solution was checked by titration against standard HCl solution, using methyl orange as indicator. Sodium hydroxide solutions containing different concentrations of S2− ions were similarly prepared from a stock solution of approximately 1 M Na2 S (Koch-Light Lab., England). Experiments were carried out at a constant temperature 25 ± 0.1 ◦ C, which was controlled using an ultra thermostat Type Polyscience (USA). Each run was carried out in a fresh test solution and with a newly polished electrode.
3. Results and discussion 3.1. Electrolytic reduction of the preformed oxygenated copper compounds Copper electrodes with specified surface area were previously equilibrated for different times of immersion in highly oxygenated NaOH solutions of two different concentrations. At the end of each time interval, the electrode was thoroughly washed with running distilled water and immersed directly in the electrolytic cell, containing deareated NaOH solution of the same concentration as the test electrolyte, under a constant cathodic current density of 0.03 mA/cm2 . The potential of the working electrode was then followed with time till the hydrogen evolution reaction. The curves of Figs. 1 and 2 represent such behaviour of the copper electrode in 0.01 and 0.10 M NaOH solutions, respectively, for different times ranged between 30 min and 6 h. The galvanostatic potential–time curves of Figs. 1 and 2 are characterized by the presence of two well-defined cathodic reduction steps before the hydrogen evolution reaction commences. In Table 1 the values of the starting potentials (eH ) of these steps are grouped.
Fig. 1. Galvanostatic potential–time curves for copper electrodes previously equilibrated for (1) 30, (2) 60, (3) 180, and (4) 360 min in oxygenated 0.01 M NaOH solutions under a cathodic current density 0.03 mA/cm2 .
Fig. 2. Galvanostatic potential–time curves for copper electrodes previously equilibrated for (1) 30, (2) 60, (3) 180, and (4) 360 min in oxygenated 0.1 M NaOH solutions under a cathodic current density 0.03 mA/cm2 .
These two cathodic polarization steps are expected to correspond to the reduction of Cu2 O and Cu(OH)2 , respectively [11,12,15]. Inspection of the potentials reported in Table 1 reveals, however, a considerable displacement of the potentials of these steps from their expected values towards more negative potentials. Similar behaviour has reported by different authors Table 1 Starting potentials (eH ) of the reduction steps observed during the cathodic reduction of the oxygenated copper compounds on the copper surface Immersion time (min)
0.01 M NaOH
0.10 M NaOH
First step (V)
Second step (V)
First step (V)
Second step (V)
30 60 180 360
−0.290 −0.270 −0.270 −0.290
−0.480 −0.460 −0.430 −0.460
−0.28 −0.24 −0.135 −0.160
−0.590 −0.560 −0.470 −0.510
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[16,17]. The reduction of both Cu2 O and Cu(OH)2 films takes place at potentials considerably far removed from their formation values according to: Cu2 O + H2 O + 2e → 2Cu + 2OH−
(2)
and Cu(OH)2 + 2e → Cu + 2OH−
(3)
The irreversibility of these reactions is intrinsic and related to the oxides themselves. Shams E1 Din and Abd El Wahab [15] have explained this behaviour in view of the lattice structure of both oxides. Cuprous oxide is regarded as having cation defects resulting in anion excess and positive holes. The latter is assumed to be localized with the formation of cupric ions. This arrangement is characteristic for p-type semiconductor oxides in which electrons are easily donated and difficulty accepted. Little is known regarding the type of electrical conductivity of crystalline Cu(OH)2 . However, CuO is known, to be an intrinsic semiconductor [18]. Adsorption of oxygen, dry and moist air, on the surface of cuprous oxides leads to the formation of CuO under certain conditions, increases considerably the p-conductivity of the oxide. Therefore, the electrochemical reduction of the oxides will not occur at their expected thermodynamic potentials, but at potential more negative than their formation values. The cathodic potential displacement depends on the potential to which the electrode was previously anodized [15], or on the time of equilibration of the electrode in solution, as shown in Table 1. Reduction of Cu2 O and CuO oxides does not proceed as expected according to: Cu(OH)2 → Cu2 O → Cu, but rather Cu2 O → Cu and Cu(OH)2 → Cu [11,12,15]. This, together with the off shooting of the potential towards negative values observed at the start of the second reduction step indicates that the cathodic process starts at the metal/oxide interface. When Cu2 O is totally reduced, the reduction of Cu(OH)2 proceeds in one single step at a potential more negative from the reversible Cu/Cu(OH)2 /OH− value. The reaction occurring during this stage is, however, complex since the system is unstable being directly transferred to the corresponding Cu2 O/Cu(OH)2 /OH− [16]. Comparison of the quantity of electricity consumed in the cathodic reduction of both oxides in 0.01 and 0.1 M NaOH solutions is of interest. In 0.01 M solution, the quantity of electricity consumed during the first cathodic step, Q1 , is markedly smaller in comparison with that consumed during the second cathodic step, Q2 . However, the reverse is true in the concentrated alkali solution, Q1 is much greater than that of Q2 . This could be attributed to the dissolution of part of the oxides in the concentrated solution. In the more dilute solution, the solubilities of the oxides seem to be at a minimum [17]. The relation between the quantities of electricity consumed during both cathodic steps with the immersion time, t, is best represented on a double logarithmic scale as is shown in Fig. 3A and B. Straight line relationships are obtained satisfying the equation: log Q = a1 + b1 log t
(4)
Fig. 3. Variation of quantity of electricity consumed during the reduction of (A) Cu2 O and (B) Cu(OH)2 in () 0.01 and () 0.10 M NaOH solutions as function of the immersion time.
where a1 and b1 are constants. The constant b1 is always the same. A behaviour and amount to which could be explained on the basis that the way by which both oxides are electroreduced is always the same. Electroreduction of both oxides involves a twoelectron transfer mechanism (Reactions (2) and (3)). The constant a1 represent, however the logarithm of the quantity of electricity (log Q) consumed at a time of immersion equals 1 min. The value of a1 is found to depend on the alkali concentration, on one hand, and on the type of oxide electroreduced, on the other. 3.2. Electrolytic reduction of the preformed cuprous sulphide Thin films of copper sulphide previously formed during the equilibration of the copper electrode of 32.72 cm2 surface area in oxygenated 0.01 and 0.10 M NaOH solutions, containing increasing concentrations of Na2 S, for different immersion times, are electroreduced galvanostatically in deoxygenated NaOH solutions of the same concentrations, at a current density of 0.03 mA/cm2 . The curves of Fig. 4 represents the galvanostatic E–time plots for the reduction of Cu2 S, previously formed in oxygenated 0.01 M NaOH in presence of 0.05 M Na2 S at different immersion times. The curves of Fig. 5 shows the same behaviour of the reduction of Cu2 S films formed in oxygenated 0.1 M NaOH in presence of 0.1 M Na2 S. Similar curves were also obtained in the two alkali concentrations in the presence of varying concentrations of Na2 S. Inspection of the galvanostatic potential–time curves of Figs. 4 and 5, reported, respectively, in 0.01 and 0.1 M NaOH solutions, in presence of increasing concentrations of Na2 S reveals some interesting features which deserve mentioning. Thus, the presence of S2− ions, at any time of immersion, causes the following: (i) The partial disappearance of the first reduction step reported in S2− free NaOH solution which corresponds to the reduction of Cu2 O.
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electricity consumed during this step depends markedly on the concentration of both NaOH and S2− ion.
Fig. 4. Galvanostatic potential–time curves for the reduction of copper sulphide films formed during the equilibration of the copper electrode in 0.01 M NaOH + 0.05 M Na2 S for different times: (1) 30, (2) 60, (3) 180, and (4) 360 min under a cathodic current density 0.03 mA/cm2 .
(ii) The complete disappearance of the second reduction step which corresponds to the reduction of Cu(OH)2 . (iii) The appearance of a new cathodic reduction step at a potential ∼−0.81 V (NHE), which is assumed to correspond to the reduction of the preformed Cu2 S film. The quantity of
Fig. 5. Galvanostatic potential–time curves for the reduction of copper sulphide films formed during the equilibration of the copper electrode in 0.1 M NaOH + 0.1 M Na2 S for different times: (1) 30, (2) 60, (3) 180, and (4) 360 min under a cathodic current density 0.03 mA/cm2 .
In one and the same solution composition, typical galvanostatic reduction curves for thin films (first-order colours), medium films (second-order colours) and thick films (third, fourth and fifth-order colour) are obtained with increasing time of equilibration. After a short initial period during which the initial potential falls rapidly to about −0.9 V (SCE), the sulphide reduction precedes over a potential range of about 0.05–0.15 V depending on the immersion time. The quantity of electricity consumed along this step depends on the thickness of the preformed Cu2 S film. During reduction, the cathode surface was observed visually through the side of the cell. First-order films showed interference colours corresponding to decreasing film thickness as the reduction proceeded. At the completion of reduction, the surface appeared almost identical to the original pre-immersion abraded surface on which the film had been formed. Thicker films showed, however, a different behaviour. Thus, in the early stages of reduction, the surface was gradually decreased in reflectivity, sometimes exhibiting patches of earlier interference colours. With further reduction, the surface took on a uniform pale brown appearance. When the circuit to the cathode was broken, this brown material gradually darkened to black. The greater part of it was loosely bound to the cathode, and could be swept into the solution. Hoar and Stockbridge [19] assumed that the overpotential for hydrogen ion reduction on reducing cuprous sulphide is the same as that on reduced copper. They reported that, at 0.03 mA/cm2 cuprous sulphide reductions at about 85% current efficiency will be followed by hydrogen ion-reduction at 100% efficiency. However, under the present experimental conditions, Figs. 4 and 5 and
Fig. 6. The relation between the quantity of electricity, Q, consumed during the reduction of copper sulphide and the immersion time, t, in 0.01 M NaOH solutions in the presence of increasing Na2 S concentrations: (1) 0.01, (2) 0.05, (3) 0.1, and (4) 0.5 M.
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catholyte or remain as isolated particles on the surface of the cathode. While these copper specks remain of colloidal size, interference colours, indicating diminishing thickness, would be shown by the remaining adherent film. When the film surface becomes flowed, by cracks or voids formed during the formation of large specks of copper, some hydrogen ion may be reduced within or at the base of cracks, the evolved gas detaching some unreduced film material from the cathode. Once out of electrical contact with the cathode, this film cannot be reduced. Thus, although the premature reduction of hydrogen ions lowers the current efficiency of sulphide reduction, the consequent increase in the reduction time is probably more than counterbalanced by failure of separated film material to reduce. The production of loose material experimentally observed is thus well interpreted as the cause of premature attainment of the end-point, and results in the observed under-estimation of thick sulphide films. 5. Conclusions Fig. 7. The relation between the quantity of electricity, Q, consumed during the reduction of copper sulphide and the immersion time, t, in 0.1 M NaOH solutions in the presence of increasing Na2 S concentrations: (1) 0.01, (2) 0.05, (3) 0.1, and (4) 0.5 M.
the like, the quantity of electricity (time in seconds till the beginning of H2 evolution) consumed during the reduction of cuprous sulphide films varies with the immersion time, t, Figs. 6 and 7, irrespective of the concentration of both NaOH and Na2 S according to: log Q = a2 + b2 log t
(5)
where a2 and b2 are constants, which vary irregularly with the composition of the solution. 4. Mechanism of cuprous sulphide reduction At room temperature, cuprous sulphide is a metal-deficit conductor, the electron conductivity being due to the presence of positive holes and the ion conductivity to vacant cation sites [14]. Although the bulk of the film is certainly non-stoichiometric cuprous sulphide [20], the metal deficit being the greater the more rapid the rate of its formation. The layer next to the metal is probably nearly stoichiometric [21]. Here, the electron and especially the ion conductivity are much lower than in the bulk. This will militate against cuprous ions being reduced at or near the metal/film interface, where ion movement will be especially difficult, although some reduction in this region may occur [22]. However, it is much more likely that electrons pass through the nearly stoichiometric layer, perhaps by tunnel effect, reduce cuprous ions in the non-stoichiometric outer layer. These can diffuse fairly readily, gradually forming specks of metal, by capturing further electrons coming through the film, encourage further cuprous ions to migrate to, and plate out, on them. As the specks grow, some of them lose physical and electrical contact with the film, and either move off into the
(1) Reduction of copper oxides preformed in solutions of oxygenated NaOH solution depends on both the electrolyte concentration and time of immersion. (2) The preformed oxides are reduced in two steps corresponding, respectively, to the reduction of Cu2 O and Cu(OH)2 to metallic Cu. (3) In presence of S2− ions, Cu2 S is formed on the expense of the absence of Cu(OH)2 and the partial formation of Cu2 O. (4) Cu2 S is reduced along one reduction step, with the quantity of electricity consumed, depends on the thickness of the preformed sulphide. (5) A mechanism is proposed for the electroreduction of Cu2 S. References [1] R.S. McMillan, D.J. Mackinnon, J.E. Dutrizac, J. Appl. Electrochem. 12 (1982) 743. [2] E.M. Arce, I. Gonz¯alez, Int. J. Miner. Process. 67 (2002) 17. [3] A.N. Buckly, R. Wood, J. Electroanal. Chem. 357 (1993) 387. [4] P. Velasquez, D. Leinen, J. Pascual, J.R. Ramos-Barrado, R. Cordova, H. Gomez, R. Schrebler, J. Electroanal. Chem. 510 (2001) 20. [5] L.I. Stepanova, T.V. Mozolevskaya, O.G. Purovskaya, Met. Finish. 101 (2003) 18. [6] M.I. Schimmel, N.R. De Tacconi, K. Rajeshwar, J. Electroanal. Chem. 453 (1998) 187. [7] A. Zirino, R. De Marco, I. Rivera, B. Pejcic, Electroanalysis 14 (2002) 493. [8] S. Jacobs, M. Edwards, Water Res. 34 (2000) 2798. [9] M.B. McNeil, A.L. Anos, T.L. Woods, Corrosion 49 (1993) 755. [10] A.M. Shams El Din, J.M. Abd El Kader, F.M. Abd El Wahab, T.M.H. Saber, A.A. El Azhary, A.A. El Waraky, Electrochim. Acta 30 (1985) 461. [11] S.M. Abd El Haleem, E.E. Abd El Aal, A.M. Atia, Corrosion 61 (2005) 838. [12] S.M. Abd El Haleem, E.E. Abd El Aal, Corrosion 62 (2006) 121. [13] E.E. Abd El Aal, Bull. Soc. Chim. Fr. 128 (1991) 351. [14] E.E. Abd El Aal, J. Power Sources 75 (1998) 36. [15] A.M. Shams El Din, F.M. Abd El Wahab, Electrochim. Acta 9 (1964) 113.
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[20] T.P. Hoar, A.J.P. Tucker, J. Inst. Met. 81 (1953) 665. [21] V.A. Kozheurov, Colloid J. 5 (1939) 45. [22] W. Jaenicke, E.M. Khairy, W. Schaefer, J. Phys. Chem. 70 (1966) 421.