The Influence of Halide Ions on the Kinetics of

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... an undesirably large overvoltage. It is known however, that this overvoltage may be reduced by. * Corresponding author. E-mail: doblhofer@fhi-berlin.mpg.de ...
Z. Phys. Chem. 217 (2003) 479–491  by Oldenbourg Wissenschaftsverlag, München

The Influence of Halide Ions on the Kinetics of Electrochemical Copper(II) Reduction By Karl Doblhofer 1 , ∗, Sabine Wasle 1 , David M. Soares 2 , Konrad G. Weil 3 , Gisela Weinberg 1 , and Gerhard Ertl 1 1

2 3

Fritz-Haber-Institut der Max-Planck-Gesellschaft, Faradayweg 4–6, D-14195 Berlin, Germany Present address: IFGW, UNICAMP; Cx. P. 6165, 13083-970 Campinas, S.P., Brasil Present address: Dept. of Materials Science and Engineering, Pennsylvania State University, 327A Steidle Building, University Park, PA 16802, USA

Dedicated to Prof. Dr. Dieter M. Kolb on the occasion of his 60 th birthday (Received July 18, 2002; accepted in revised form November 11, 2002)

Electrochemistry / Copper / Halides / Mechanism / Electroplating / Electrorefining The cathodic reduction of copper(II) in an electrolyte comparable to technical conditions (2.2 M H2 SO4 + 0.3 M CuSO4 ) is markedly affected by the presence of small concentrations of halide ions. Chloride ions accelerate the reaction, while it is slowed down by bromide. Experiments in which cyclic voltammetry is combined with an electrochemical quartz-crystal microbalance reveal the deposition and dissolution of crystalline CuCl or CuBr, respectively, on the copper surface. At the technically relevant more negative electrode potentials bulk CuCl and CuBr are unstable, however halide ions are adsorbed on the copper electrode. Although there is evidence for adsorption of the reacting Cu(II) species at the electrode surface, up to a concentration of 0.3 M CuSO4 in presence as well as in absence of adsorbed halogenide, there is no evidence for limiation of the reaction rate caused by a limited coverage of the surface with this species.

1. Introduction The reduction of copper(II) to metallic copper is a process of great technical importance [1–3]. Under typical conditions, i.e., at concentrations of both CuSO4 and H2 SO4 of the order of 1–2 M, the reaction needs an undesirably large overvoltage. It is known however, that this overvoltage may be reduced by * Corresponding author. E-mail: [email protected]

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the addition of about one mM Cl− to the electrolyte. The mechanism by which the chloride is catalytically active is still the subject of debate. In numerous cases of electron transfer reactions with redox ions the type of anions present has a significant effect on the reaction rate, see for example the dependence of both the homogeneous and heterogeneous Fe2+/3+ electron exchange rates on the presence of chloride [4, 5]. In the case of Cu(II) reduction it was suggested that an anion-bridged activated surface complex of the Cu−Cl− −Cu + type is responsible for the rate enhancement [6]. Under open-circuit conditions Cu2+ solutions in contact with Cu establish an equilibrium concentration of Cu + according to [7–9]: Cu 2+ + Cu = 2Cu +

pK = 6.2 .

(1)

Recently, we have shown [10] that this intermediate Cu+ reacts with chloride to form crystalline CuCl on the copper surface when the solubility product K s is exceeded: Cu + + Cl− = CuCl solid

K s = 1.72 × 10−7 M 2 .

(2)

In the present work the study was extended to the analysis of CuBr that forms according to Cu + + Br − = CuBr solid

K s = 5.9 × 10−9 M 2

(3)

when the electrolyte contains bromide. The discussion of the mechanism of CuCl and CuBr formation combined with the known adsorption behaviour of chloride [11–15] and bromide [16– 18] on copper leads to new insight into the role of halides in the kinetics of the copper deposition reaction.

2. Experimental The electrochemical experiments were conducted with conventional potentiostatic equipment. The saturated Hg/Hg2 SO4 reference electrode, “SMSE” has been used throughout the work, and all potentials are given with respect to this reference: E SMSE = E SHE − 0.69 V, where NHE is the normal hydrogen electrode. The electrolytes were prepared with analytical grade reagents and Milli-Q deionized water. The electrochemical quartz-crystal microbalance (EQCM) was described previously [19]. The 6 MHz quartzes were purchased from KVG, Germany. A central circular area of 6 mm diameter was gold plated on both sides of the crystal. On the gold face exposed to the electrolyte a copper layer was deposited electrochemically before each experiment (at −0.48 V). The film thickness was approx. 1 µm as determined by integrating the Faradaic current and from the reduction of the resonance frequency (EQCM sensitivity 12.8 ng/Hz cm 2 ). For the rotating-electrode experiments the circular front

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Fig. 1. Effect of chloride and bromide addition on the voltammetric current (I)/voltage (E) curves observed at a sweep rate of 20 mV s−1 . Base electrolyte: 0.3 M CuSO4 + 2.2 M H 2 SO 4 . Electrode: copper deposited on gold surface of the EQCM quartz, area 0.28 cm2 . The electrolyte was agitated by argon bubbling.

faces of copper (99.999%) cylinders were used, diameter 6 mm. Before each experiment they were polished with grain size down to 0.25 µm, and washed thoroughly in an ultrasonic bath with alkohol and Milli-Q water.

3. Results The experiments were conducted with aqueous electrolytes containing 2.2 M H2 SO4 + 0.3 M CuSO4 plus chloride ions (added as HCl) or bromide ions (added as KBr) of concentrations varying between 0 and 8 mM. Fig. 1 shows typical cyclic voltammograms that demonstrate that the reduction current is enhanced by chloride as discussed before [10], but is reduced considerably by the presence of bromide. Fig. 2 shows as solid lines the variation of the EQCM resonance frequency, reported as mass change, ∆m, for the experiments of Fig. 1. If the mass change

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Fig. 2. Comparison of the mass changes observed with the EQCM (——) with those obtained by integrating the current according to Eq. (4) (- - - - -). Experiments and system as in Fig. 1.

would be solely due to anodic dissolution/cathodic deposition processes it would be given according to Faraday’s law by: M M ∆q = − ∆m = −s∆ f = − nF nF

t2 Idt

(4)

t1

∆m s ∆f n M ∆q

is the mass change is the sensitivity of the EQCM in g Hz−1 change in resonance frequency number of electrons exchanged per metal atom atomic mass of the metal electric charge tranferred for Faradaic deposition/dissolution.

The dashed curves in Fig. 2 are calculated from the corresponding currents (Fig. 1) according to Eq. (4), with n = 2 and M = 63.5 for the Cu 2+ /Cu re-

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Fig. 3. Scanning electron micrograph showing the copper electrode surface after immersion for 3 minutes in the electrolyte 0.3 M CuSO4 + 2.2 M H 2 SO 4 + 4 mM KBr. The EDAX analysis reveals that the crystals consist of CuBr.

action. In absence of halogenide in the electrolyte the calculated curve agrees well with the EQCM result, indicating that under these conditions the Cu 2+ /Cu reaction is the only process leading to mass changes of the electrode. On the other hand, in the presence of chloride or bromide an additional mass gain is observed in the more positive potential region: in the case of chloride in the anodic scan at about −0.47 V, disappearing during the cathodic scan at about −0.5 V; in the case of bromide the extra mass starts to build up at E SMSE > −0.5 V, and disappears again at E SMSE < −0.57 V (outside the E-range of Fig. 2). When the electrode is kept at the open-circuit potential (E SMSE = −0.41 V) in presence of bromide at concentrations above ≈ 0.5 mM, the resonance frequency of the EQCM decreases at nearly constant rate (≈ −400 Hz/min in the case of 1.5 mM Br − with the solution unstirred). Fig. 3 is an electron

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Fig. 4. The dependence of ∆m, i.e., the maximum difference between measured and calculated mass changes as represented in Fig. 2, on the concentration of bromide in 0.3 M CuSO 4 + 2.2 M H 2 SO 4 .

Fig. 5. Tafel plots, i.e., charge-transfer current (I) vs. overvoltage (η), for the Cu(II)/Cu reaction in the electrolytes 0.3 M CuSO4 + 2.2 M H 2 SO 4 plus halogenides of indicated concentration. Electrode area 0.28 cm2 .

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Fig. 6. Dependence of the cathodic charge-transfer current at E SMSE = −0.51 V on the CuSO 4 concentration, [Cu++ ], in (a) 2.2 M H 2 SO 4 , (b) 2.2 M H 2 SO 4 + 2 mM HCl (c) 2.2 M H 2 SO 4 + 0.3 mM KBr. Electrode area 0.28 cm2 .

micrograph of the copper surface after contact with an electrolyte containing 4 mM KBr for 3 minutes, indicating that the mass increase is due to a crystalline deposit that forms without current flow. The EDAX analysis conducted as described before [10, 20] indicates unambiguously that the deposit is CuBr. It has been shown [10] that CuCl forms at open circuit at the rate of transport of chloride towards the electrode surface. In the present case of CuBr formation the situation is similar. In Fig. 4, the maximum excess mass in voltammetric sweeps as in Fig. 2, i.e., the maximum difference between the mass measured with the EQCM and that calculated from the current with n = 2 is plotted versus the bromide concentration. The linear relationship is indicative of rate control by bromide transport. The finite solubility of CuBr leads to the intercept on the x-axis. Finally, voltammetric currents (I) as in Fig. 1 were measured as a function of the rotation rate √ (ω) of the copper electrode. From an extrapolation of the plots 1/I vs. 1/ ω to ω → ∞, the charge-transfer (CT) limited currents were determined. In Fig. 5 the dependence of such CT currents on the overvoltage is shown (“Tafel plots”). Fig. 6 represent the CT currents for cathodic reduction in presence and absence of halides at E SMSE = −0.51 V, as a function of the concentration of copper(II) in the electrolyte.

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4. Discussion 4.1 The mechanism of CuX formation It was the aim of the work to analyze the system under conditions that are close to those of technical copper deposition. In particular, the reactions were studied at 2.2 M H2 SO4 and 0.3 M CuSO4 . Assuming that H2 SO4 is dissociated into H + and HSO4 − , and CuSO4 is fully dissociated, the ionic strength of the electrolyte is 2.8 M. Inspection of the tabulated values of activity coefficients for CuSO4 [21], along with the consideration that the single-ion activity coefficient of Cu ++ will not differ much from that of SO4 −− [22], the activity coefficient for Cu ++ in this electrolyte, γCu++ = 0.05 is obtained. This activity coefficient will be used for the following discussions. When a solution containing 0.3 M CuSO4 , i.e., with an activity a Cu++ = 0.015 M, is in contact with Cu, the equilibrium activity of Cu + will be a Cu+ = 0.10 mM (Eq. (1)). According to Eqs. (2) and (3), the minimum halide activities for reaching the solubility products would hence be a Cl− = 1.8 mM and a Br− = 0.06 mM, respectively. Considering that the activity coefficients of the halogenide ions are somewhat larger than unity [22], the values derived from the intercepts of the plots of measured mass change as a function of halide concentration (Fig. 4 for bromide and Ref. [10] for chloride) are in agreement with the predicted values. The formation of CuBr (or CuCl in presence of Cl − ) at the open circuit situation can thus be understood straightforwardly as a precipitation reaction between Cu + formed at a fast rate via Eq. (1) and the halide ions, X − . The rate of CuX formation is determined by the transport of X − towards the copper surface. Under the conditions of cyclic voltammetry the CuX deposit forms in the anodic scan as the potential approaches the equilibrium potential. Apparently, in this potential range the activity of Cu + exceeds the value defined by the solubility product with the halogenide. In the negative potential range the CuX deposit dissolves because the activity of Cu + falls below the value defined by the solubility product. To estimate these Cu + activities, consider the basic Butler–Volmer equations [23] for Cu ++ /Cu +         α I F E − E I0 (1 − α I )F E − E I0 a Cu+ a Cu++ 0 iI = iI exp exp − − 0 0 a Cu RT a Cu++ RT + (5) and for Cu + /Cu   i II = i

0 II

      α II F E − E II0 (1 − α II )F E − E II0 a Cu+ exp − 0 exp − RT a Cu+ RT (6)

where i I and i II are the charge-transfer controlled current densities of the two consecutive electrochemical reaction steps [24–28], i I0 and i II0 are the exchange

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0 0 0 0 current densities for a Cu + = a Cu ++ = 1 M, and E I and E II are the corresponding standard potentials:

Cu 2+ + e− = Cu + Cu + + e− = Cu metal

◦ E SMSE = −0.537 V ◦ E SMSE = −0.169 V .

(7) (8)

For deriving from Eqs. (5,6) the activity of Cu + , its transport away from the electrode is disregarded so that the stationary state of the system is characterized by i I = i II = i/2, where i is the observed current density. Additionally, the following assumptions are made: 1. activity of Cu++ is constant (a Cu++ ≈ const.) 2. all CT coefficients are equal (α I ≈ α II ≈ 0.5). The stationary activity of Cu+ is then:   0 F ( E−E I0 ) F ( E−E II i I0 a Cu ++ ) 0 exp − exp + i 0 II 2RT 2RT a ++ Cu 0   . (9) a Cu+ = a Cu + 0 F ( E−E I0 ) F ( E−E II 0 0 i I exp + i II exp − 2RT ) 2RT In Fig. 7, the Cu + activities obtained with Eq. (9) for three sets of values of the exchange current densities, i I0 and i II0 are plotted for the potential range employed in the experiments. The combination 0.6 A/cm 2 and 1.2 A/cm 2 for the i 0 values yields a current/potential plot, Fig. 8, that reflects the actual situation with sufficient accurracy for the present arguments. These exchange current densities may seem rather large. This is because frequently i 0 values refer to relatively small concentrations while here they are reported for the activities 0 0 a Cu + = a Cu ++ = 1 M. For a discussion of this point see for example the review in Bard’s Encyclopedia [27]. The rise of the calculated Cu+ concentration near the equilibrium potential, Fig. 7, agrees with that observed with the rotating ring-disk electrode [25]. It supports the proposed mechanism of CuX formation and dissolution. The CuX deposit forms by a precipitation mechanism when in the positive potential sweep the concentration of Cu + exceeds the value defined by the solubility product. In the reverse sweep the precipitate dissolves as [Cu + ] falls below this value. It is remarkable that the CuX formation is hardly noticeable in the experimental current/potential curves. This is because the formation of Cu + proceeds to a significant extent (at the equilibrium potential to 100%) via Eq. (1), i.e., the cathodic copper(II) reduction according to Eq. (5) and the anodic copper oxidation according to Eq. (6) proceed simultaneously.

4.2 Effect of the halides on the charge-transfer reaction rates In the potential range of the technical Cu deposition, i.e., at E SMSE < −0.5 V [1] the copper surface is free of bulk CuX. However, halide ions are adsorbed at

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Fig. 7. Stationary activities of Cu+ , aCu+ , obtained with Eq. (9) for aCu++ = 0.3 M and the indicated exchange current densities, i 0 . The curves for 0.5/1.0 A cm−2 and 0.6/ 1.2 A cm−2 are identical.

these potentials [11–18]. It is of interest to elucidate the effect of the adsorbed halide ions on the rate of copper(II) reduction. The Tafel plots obtained in presence and absence of halides (Fig. 5) show in all cases cathodic slopes of 120–135 mV/decade and anodic slopes of 50–60 mV/decade. This result is consistent [24] with the description of the currents by Eqs. (10a,b):

−(1 − α)FE i cath = −ka Cu++ exp (10a) RT

(1 + α)FE  . (10b) i anod = k exp RT They describe the sequence of two consecutive electrochemical charge-transfer steps, Cu ++ /Cu + and Cu + /Cu, in which the first step (Cu ++ /Cu + ) is rate determining. From the Tafel plots, Fig. 5, it may be concluded that this mechanism is not altered by the presence of adsorbed halide ions. It has been our experience that the CT current depends very strongly on the surface state of the copper electrode. This finding is consistent with the Tafel

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Fig. 8. Current densities obtained with Eqs. 7,8, using the exchange current densities, i 0 , of Fig. 7.

constants reported by Mattsson and Bockris [24], that demonstrate a variation of the current by an order of magnitude as a consequence of varying the method of electrode preparation. This points to specific interaction, i.e., adsorption of the reactant at the electrode surface. Thus, the charge-transfer rate might well be limited by full coverage according to

−(1 − α)FE i cath = −k θ Cu++ exp (11) RT where θ Cu++ is the fractional coverage of the electrode surface with the reacting Cu(II) species. To find out if θ Cu++ reaches saturation, the reduction rates were determined as a function of the Cu ++ concentration, Fig. 6. Since the activity coefficient may be considered constant, it is clear that up to a CuSO4 concentation of 0.3 M saturation is not noticeable, in presence as well as in absence of adsorbed halogenide. This means, up to a concentration of 0.3 M CuSO4 the adsorption of the reacting Cu(II) species does not reach a limiting value that would influence the reduction rate. At the ionic strength of 2.8 M, the electrostatic effect of the adsorbed halogenide ions is not expected to generate the changes in the reaction rates

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observed experimentally [29, 30]. Rather, one concludes that the free energy barrier of the activated complex forming on the copper surface in the course of the Cu ++ /Cu + transition increases when bromide is adsorbed, while it decreases slightly when the chloride adsorbate is present.

5. Conclusions Solutions containing CuSO4 + H2 SO4 + X− , where X − is Cl− or Br − , are studied with electrochemical methods and with the EQCM. It is found that: 1. In the electrochemical Cu(II)/Cu system both the reactions Cu ++ + e− = Cu + and Cu + + e− = Cu proceed according to their individual kinetics, establishing a potential dependent stationary concentration of Cu+ . At potentials, at which the product of the activities of Cu + and X − exceeds the value of the solubility product K s , crystalline deposits of CuCl or CuBr form on the copper surface. At more negative potentials, at which the activity of Cu + falls below the value required for the solubility product, the deposit dissolves again. 2. Technical copper deposition is conducted normally [1] at current densities not lower than 10 mA cm −2 , i.e., at E SMSE < −0.5 V. Under these conditions both the bulk compounds CuCl and CuBr are unstable. However, when chloride or bromide are present in the electrolyte, in the course of Cu(II) reduction the electrode surface is covered with a layer of adsorbed X − , that influences the rate of the charge-transfer reaction. 3. Up to a concentration of 0.3 M CuSO4 in presence as well as in absence of adsorbed halogenide there is no evidence for a limiation of the reaction rate caused by a limited coverage of the electrode surface with the reacting Cu(II) species.

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