J. Phys. Chem. B 1997, 101, 2411-2414
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Electrocatalytic Reduction of Peroxodisulfate in 0.5 M NaOH at Copper Electrodes. A Combined Quartz Microbalance and Rotating Ring/Disc Electrode Investigation S. Ha1 rtinger,† J. Rosenmund, E. Savinova,‡ S. Wasle, and K. Doblhofer* Fritz-Haber-Institut der Max-Planck-Gesellschaft, Faradayweg 4-6, D-14195 Berlin-Dahlem, Germany ReceiVed: September 6, 1996; In Final Form: December 4, 1996X
From an electrochemical investigation by means of an electrochemical quartz microbalance, a rotating disc electrode, and a ring/disc electrode, two mechanisms for the reduction of S2O82- became apparent. Besides the well-known outer-sphere cathodic reduction, a catalytic mechanism of S2O82- reduction operates in a potential range between the surface oxide region (≈-0.5 V/SCE) and -1.0 V/SCE. It involves the chemical oxidation of the copper surface to a soluble Cu(I) species. The catalytic mechanism is concluded to result from the specific interaction between S2O82- and the Cu surface modified by the presence of subsurface oxygen.
Introduction The electrochemical reduction of peroxodisulfate anions (S2O82-) is one of the classical model systems for the study of the correlation between the electric double-layer structure and the rate of the charge-transfer step across the electrode/ electrolyte interface. Since the first reports of Frumkin in 1933,1 there has been a growing interest in the variety of dynamic features that lead under certain conditions to a negative faradaic impedance, current oscillations, spatiotemporal patterns, and other nonlinear phenomena.2-9 The reaction has been conducted on a large number of electrode materials, which were generally assumed to behave inertly in the course of the irreversible electrode reaction
S2O82- + 2e- f 2SO42-
(1)
However, on some metals the nonlinear phenomena observed during the cathodic reduction of S2O82- are so complex that they cannot be explained completely on the basis of the above model. This is true in particular for copper electrodes. To elucidate the nature of such complicated processes, the present study was conducted. The peroxodisulfate ion is a strong oxidant (E° ) 1.76 V/SCE10). In an early investigation Levie et al.11 studied the action of peroxodisulfate on metals of group IB, IIB, and VIIIB. Of all the metals investigated, only Pt and Au behave inertly, but the peroxodisulfate ion catalytically decomposes water in contact with these metals. A second group of metals (Ag, Fe, Hg, Zn, Cd) readily dissolve at varying rates in the presence of peroxodisulfate, and some of them form insoluble sulfates. Finally, there are metals (Cu, Ni, Co) that react under chemical (oxidative) dissolution with simultaneous formation of metal oxides and sulfate species. The chemical dissolution rate of copper specimens has additionally been determined from an exsitu gravimetric analysis. From the data obtained by Bond et al.,12 a dissolution rate of 1.6 × 10-6 g (Cu) cm-2 s-1 in 0.1 M K2S2O8 at 19 °C can be derived, showing that the chemical dissolution reaches the order of monolayers per second. It appears important to quantify the potential-dependent rate of copper dissolution that takes place during cathodic S2O82† Present address: Instituto de Quimica, Universidade Estadual de Campinas, CP 6154, 13081 970 Campinas, SP, Brazil. ‡ Permanent address: Boreskov Institute of Catalysis, Lavrentjeva 5, 630090 Novosibirsk, Russia. * To whom correspondence should be addressed. X Abstract published in AdVance ACS Abstracts, March 1, 1997.
S1089-5647(96)02727-7 CCC: $14.00
reduction. This is done in the present work with the technique of the electrochemical quartz microbalance (EQMB). The rotating ring/disc electrode (RRDE) technique is used for identifying the nature of the formed soluble species. The formation of surface oxides over the course of peroxodisulfate reduction has been investigated for Au and Pt electrodes, by means of standard electrochemical methods13-16 and surface-enhanced Raman spectroscopy (SERS).17 No such studies are available on metals such as Fe, Ni, and Cu, which are known to yield particularly stable current oscillations and rich dynamic features.18 For the present paper we therefore investigated the state of the reactive copper-electrode surface. The EQMB proved again very useful. In addition, we used the results of a recent SERS investigation of single-crystal copper electrodes19 for interpreting surface processes and the mechanism of the S2O82- reduction. Experimental Section A computer-controlled EQMB setup, with a circular, goldcoated AT-cut 6 MHz quartz resonator, was used in the present investigation. The properties of the driver circuit and its special features have been described in recent communications.20,21 The registration of the resonance frequency f and its changes (∆f) was done via a frequency counter (Hewlett-Packard HP 53181A), which allowed sample reading and storage on the PC up to a maximum rate of 40 s-1. A data acquisition time of 80 ms was used for the following experiments, resulting in a frequency resolution of (0.5 Hz. The gold/copper coating on one side of the quartz was in contact with the electrolyte. It was grounded and acted as the working electrode of a conventional three-electrode configuration, with a saturated calomel electrode (SCE) as the reference and a platinum sheet as counter electrode. By use of two digital multimeters (Hewlett-Packard HP 34401A), the current/potential (I/E) curve was recorded simultaneously with f. Since changes in the damping resistance of the quartz20,21 were never observed, the ∆f readings were converted straightforwardly into mass changes by the proportionality factor SEQMB.22 For the 6 MHz quartz crystal used in this work, a value of SEQMB ) 12.88 ng/ (Hz cm2) was found in a calibration procedure based on the electrochemical deposition and dissolution of copper and silver. For the present investigation all potentials are referred to the SCE. The electrochemical active area of the quartz resonator was 0.38 cm2. Before each measurement the gold-coated quartz was freshly plated with copper by galvanostatic deposition of © 1997 American Chemical Society
2412 J. Phys. Chem. B, Vol. 101, No. 14, 1997
Ha¨rtinger et al.
Figure 2. Current (I, lower trace) and EQMB frequency change (∆f, upper trace) of the EQMB-copper electrode in a stirred solution of 10-3 M Na2S2O8 and 0.5 M NaOH on a continuous time scale. The electrode potential varies linearly (sweep rate (20 mV/s) between the values -1.2 V (at positions 0, 2, and 4) and -0.55 V (at positions 1, 3, and 5). Figure 1. Cyclic voltammograms of the EQMB-copper electrode in a solution of 0.5 M NaOH without Na2S2O8 (- - -) and with 10-3 M Na2S2O8 (s) at 20 mV/s. The electrolytes were deaerated and stirred with a magnetic stirrer. The arrow marked A indicates the potential E ) -0.55 V, which is of relevance in the following experiments.
approximately 2000 monolayers from a 0.5 M CuSO4 solution at pH 1. After flushing with triply distilled water, the copper/ gold EQMB working electrode was mounted in the Teflon holder, transferred to the nitrogen-flushed measurement cell, and held for 10 min at -1.3 V in order to reduce air-formed surface oxides. During the measurement the electrolyte was agitated by means of a magnetic stirrer at a constant rotation rate of 720 rpm. Measurements at the RRDE were conducted using a PAR 366 bipotentiostat, and the electrode was rotated at a rate of 1800 rpm. Both the ring and the disc were constructed of platinum. The disc was copperized using the same procedure as used for the quartz resonator. From the geometric dimensions of the electrode (disc radius 0.3005 cm, inner ring radius 0.3605 cm, and outer ring radius 0.4606 cm), the theoretical collection efficiency (N) was calculated to be N ) 0.362.23 From an independent RRDE experiment for the reduction of K3[Fe(CN)6] in 0.01 M KCl, an experimental value for Nexp ) 0.368 was determined. This value was used to quantify the RRDE results in this work. All measurements were conducted in 0.5 M NaOH as base electrolyte with 1 mM Na2S2O8. A fresh solution was daily prepared from Na2S2O8 (Merck, p.a.), NaOH (Merck, suprapur), and triply distilled water. Results Voltammetric Response of the Cu Electrode in Alkaline Electrolytes. Figure 1 shows an anodic and the following cathodic sweep from a continuous cyclic voltammogram of the EQMB-Cu electrode in the 0.5 M NaOH electrolyte with and without S2O82-. At potentials more positive than -0.55 V (marked A), significant anodic copper(I) oxide formation starts. It leads to the observed current peak at -0.5 V24-26 with an apparent anodic charge (Qa) of 1.5 mC cm-2. On the following negative sweep the cathodic reduction of Cu2O gives rise to the current peak near -0.7 V. From the comparison of the electrochemical charges, in absence of S2O82- approximately 83% of Cu2O is converted back to copper in the cathodic peak region. The corresponding S2O82- reduction currents
during the anodic and cathodic potential sweeps exhibit a marked hysteresis. Behavior of the Cu Electrode in the Potential Range -1.2 < E < -0.55 V. In Figure 2, lower curve, the current corresponding to a sequence of positive and negative potential scans between -1.2 and -0.55 V applied to the EQMB-copper electrode is represented on a continuous time axis. The upper trace shows the EQMB results obtained simultaneously. The electrolyte solution contains 0.5 M NaOH and 1 mM peroxodisulfate. Significant S2O82- reduction currents are observed only at potentials more negative than -0.5 V. Since E° ) 1.76 V/SCE, one finds that the overpotential exceeds -2 V! The anodic and cathodic I/E curves coincide practically. At the negative end of the investigated potential range a current plateau is formed, which results from transport limitations of the peroxodisulfate to the electrode. Note that in the absence of the reducible species on the I scale of this figure no voltammetric features would be noticeable.19 The EQMB frequency (upper curve) shows a change of ∆f ≈ 2 Hz after each complete potential cycle. This corresponds to a mass loss of ≈ 2.8 × 10-8 g (Cu) per cycle or a dissolution of half a monolayer of copper atoms from the electrode surface. The rate of metal dissolution, which follows from the slope of the graph (d(∆f)/dt), is significant in the potential range where the peroxodisulfate reduction proceeds under charge-transfer control, i.e., where there is S2O82- present near the electrode surface at a finite concentration. The maximum Cu dissolution rate is 2 × 10-9 g cm-2 s-1. Assuming proportionality with the S2O82- surface concentration, this dissolution rate corresponds to the same order of magnitude as the one reported by Bond.12 The dissolution process ceases when the electrode potential is held in the region of the diffusion-current plateau. EQMB Investigation in the Extended Sweep Range -1.2 to -0.4 V. The EQMB response changes significantly when the positive potential limit is extended into the copper(I) oxide region. The result represented in Figure 3 shows an anodic and the following cathodic sweep obtained again from a cyclic voltammogram. In each anodic sweep, between -0.7 and -0.5 V, ∆f increases by 19.5 Hz to a maximum. This is equivalent to a mass loss of 2.5 × 10-7 g (Cu). The maximum dissolution rate (from d(∆f)/dt) is 1.5 × 10-7 g cm-2 s-1. It is observed at a potential 50 mV negative to the anodic current peak. After
Electrocatalytic Reduction of Peroxodisulfate
J. Phys. Chem. B, Vol. 101, No. 14, 1997 2413
Figure 5. RRDE current vs potential curves in 10-3 M Na2S2O8 and 0.5 M NaOH at 20 mV/s; fRRDE ) 1800 rpm. Copper disc current (upper trace) and platinum ring current (lower trace) vs disc potential ED: -1.3 < ED < -0.38 V. The ring potential is held constant at ER ) 0.6 V. Figure 3. A positive and the following negative potential sweep taken from a cyclic voltammogram as in Figure 2, but with the potential range extended to -1.3 < E < -041 V. I is the measured current, and ∆f is the EQMB frequency change observed simultaneously.
TABLE 1: Possible Surface Reactions of Copper That Might Proceed over the Course of Electrochemical Peroxodisulfate Reduction at pH 13.7 anodic electrode reaction e-
(1) Cu f Cu(I)sol + (2) Cu f Cu(II)sol + 2e(3) 2Cu + 2 OH- f Cu2O + H2O + 2e(4) Cu + H2O f CuOH + H+ + e(5) Cu2O + 2OH- f 2CuO + H2O + 2e(6) Cu2O + 2OH- + H2O f 2Cu(OH)2 + 2e-
K × 105,a A s2 cm-2 1.953 3.905 -15.510 -7.295 -15.510 -4.769
a K is the constant arising from eq 2 (positive sign of K means an increase of f, i.e., a decrease of electrode mass associated with copper oxidation).
Figure 4. Curve 1: EQMB frequency change with time (t)/electrode potential (E) obtained in the presence of Na2S2O8 in the alkaline electrolyte (data taken from the experiment of Figure 3, curve 1). Curve 2: the corresponding EQMB signal obtained in the absence of Na2S2O8.
the maximum, the resonance frequency decreases by 9.5 Hz, corresponding to a mass gain. On the following negative potential sweep the electrode mass remains at first almost unchanged. As E approaches the potential of the cathodic Cu2O reduction peak, the rate of mass loss increases to ≈5 × 10-8 g cm-2 s-1 and remains at a high value (≈3 × 10-8 g cm-2 s-1) until it decreases at E ≈ -1.0 V. Thus, under these conditions, after every completed potential excursion to -0.41 V the resonance frequency has increased by more than 30 Hz relative to the value at the end of the previous cycle; i.e., the EQMB-copper electrode loses a mass 0.4 µg/cm2 per cycle! In absence of S2O82-, i.e., in the pure 0.5 M NaOH electrolyte, the EQMB signal shows only the increase (-∆f) and decrease (+∆f) of electrode mass associated with the formation and reduction of Cu2O on the surface (Figure 4, curve 2). Investigation with the Rotating Ring/Disc Electrode (RRDE). The following RRDE experiments were conducted to obtain information on the oxidation state of the soluble copper species formed in the course of electroreduction of S2O82- in alkaline electrolyte. For the experiment reported in Figure 5 the potential of the platinum ring electrode was set to +0.6 V. As no electroreduction of peroxodisulfate occurs on platinum at this potential,
any current detected on the ring electrode is expected to be due to the oxidation of Cu(I) to Cu(II) species. From the collection curve at the ring electrode (lower part of the Figure 5) the formation of Cu(I) species (probably a hydroxy complex) is strongly indicated in the positive potential region. Thus, on the positive potential sweep the ring current peak shows the dissolution of copper as Cu(I) species, accompanying Cu2O film formation. Unfortunately, the ring current decreases quickly, apparently due to the buildup of a passivating oxide/hydroxide layer of Cu(II).26,27 When the electrolyte does not contain S2O82-, no measurable ring current flows. The same is true when the electrolyte contains S2O82-, but the anodic potential limit is kept at values E e -0.55 V. Discussion Analysis of the Potential-Dependent Surface Processes. On the basis of Faraday’s law, an electrochemical current flow I corresponds to the time derivative of the EQMB ∆f/t curves (Figures 2-4), if it leads to a mass (m) change of the electrode:
I)
d(∆f) zF dm zF SEQMB d(∆f) ) )K M dt M dt dt
(2)
where M is the molecular mass of the electrochemically converted species. K is a constant with units of A s2 cm-2, with positive or negative sign, depending on the particular surface reaction. In Table 1 possible surface reactions and the corresponding constants K are collected. By using the value of K ) -15.51 (reaction 3 of Table 1), one obtains indeed the experimental ∆f curve 2 of Figure 4 from the current integral of the corresponding voltammogram of Figure 1. Consider, further, that the “normal” corrosion, as
2414 J. Phys. Chem. B, Vol. 101, No. 14, 1997
Ha¨rtinger et al. encountered phenomenon which leads to various oxydehydration, hydrogen abstraction, and electrophilic reactions.19,29-34 These oxygen centers embedded in the electrode surface strongly polarize the copper/oxygen bonds.34 In the presence of chemisorbed peroxodisulfate ions the electron deficiency at the metal centers is further decreased, thus promoting the electron uptake through the S2O82-. The transient Cu(I) species may be oxidized by the sulfate radical, finally yielding two SO42- and one Cu(II) per catalytic cycle. The proposed course of the catalytic S2O82- reduction is sketched in Figure 6. Conclusion
Figure 6. Mechanism for the chemical reaction of S2O82- with the copper electrode in the presence of the copper/subsurface oxygen entity. The Lewis base character (step 1) of the embedded oxygen and the presence of the electron-withdrawing peroxodisulfate at the interface (step 2) facilitate the copper oxidation (step 3).
described in the literature12 and observed in Figure 2, amounts to a mass loss corresponding to approximately 2 Hz per potential cycle. Thus, if the S2O82- reduction reaction on Cu electrodes would proceed “normally”, one would expect that the two effects of surface oxide formation and corrosion would superimpose; i.e., in the experiment of Figure 3 the frequency change ∆f on the E (or t) scale should resemble curve 2 of Figure 4, with a rise of 2 Hz superimposed. However, what is observed is curve 1 (Figure 4)! This means that when the electrode is subjected to such positive electrode potential that surface Cu2O forms, the electrode surface can transform into a state in which it corrodes strongly over the course of S2O82- reduction. The process that leads to the sharp mass decrease in the anodic sweep of Figure 3 starts at a potential where the current is still cathodic. Furthermore, in the negative scan, at -0.7 < E < -1.0 V Cu dissolution proceeds at a rate of approximately 5 × 10-8 g/(cm2 s), which would correspond to an anodic current of approximately 0.03 mA (eq 2, dissolution as Cu(I)) if it were a straightforward faradaic process. Since the observed electrochemical current flow in this potential region is cathodic, one must postulate a chemical (corrosion) reaction between S2O82- and the copper surface that proceeds in parallel to the electrochemical S2O82- reduction. In our recent SERS investigation on the surface state of a single-crystal Cu(111) electrode in 0.5 M NaOH,19 it was concluded that a subsurface oxygen species is present in this potential range below the Cu2O surface film. Figure 5 of the authors’ previous work19 shows that the subsurface oxygen (Cu*(O)ad) concentration exhibits the same hysteresis between the anodic and cathodic potential sweep as it is observed for the catalytic activity of the copper surface for the corrosive reduction of S2O82- in the present paper. Thus, we attribute the observed high catalytic activity for the reaction of S2O82in the cited regions (anodic sweep -0.6 < Ea < -0.5 V and cathodic sweep -1.1 < Ec < -0.75 V) to the fact that the Cuelectrode surface is transformed into a catalytically active state by the presence of this subsurface oxygen species. The catalytic activity ceases, as the subsurface oxygen is reduced at very negative potentials. The surface is reactivated over the course of formation and reduction of the surface Cu2O film. Mechanism of S2O82- Catalysis. The high reactivity of low coordinated copper/oxygen species, like chemisorbed, absorbed, or subsurface oxygen, toward surface reactions is a frequently
The electrochemical reduction of peroxodisulfate at copper electrodes is associated with chemical oxidation (corrosion) of the electrode surface. A pure voltammetric investigation does not yield the complete information on these surface reactions. The application of the electrochemical quartz microbalance and the rotating ring/disc electrode technique proved to be most valuable tools for separating the chemical and electrochemical effects. The corrosive parallel pathway for the reduction of S2O82- proceeds on a copper/subsurface oxygen electrode surface. References and Notes (1) Frumkin, A. N. Z. Phys. Chem. (Munich) 1933, 164, 121. (2) Wolf, W.; Ye, J.; Purgand, M.; Eiswirth, M.; Doblhofer, K. Ber. Bunsen-Ges. Phys. Chem. 1992, 96, 1797. (3) Fla¨tgen, G.; Krischer, K.; Ertl, G. Z. Naturforsch. 1995, 50a, 1097. (4) Fla¨tgen, G.; Krischer, K. J. Chem. Phys. 1995, 103, 5428. (5) Fla¨tgen, G.; Krischer, K.; Pettinger, B.; Doblhofer, K.; Junkes, H.; Ertl, G. Science 1995, 269, 668. (6) Fla¨tgen, G.; Krischer, K. Phys. ReV. E 1995, 51, 3997. (7) Fla¨tgen, G.; Krischer, K.; Ertl, G. J. Electroanal. Chem., in press. (8) Koper, M. T. M. Ber. Bunsen-Ges. Phys. Chem. 1996, 100, 497. (9) Damaskin, B. B.; Safonov, V. A.; Fedorovich, N. V. J. Electroanal. Chem. 1993, 349, 1. (10) Latimer, W. M. The Oxidation States of the Elements and Their Potentials in Aqueous Solutions; Prentice-Hall: New York, 1952. (11) Levie, M. G.; Migliorini, E.; Ercolini, G. Gazz. Chim. Ital. 1908, 38, 583. (12) Bond, G. C.; Hill, G. M.; Tennison, R. J. Chem. Soc. 1959, 33. (13) Mu¨ller, L. J. Electroanal. Chem. 1967, 13, 275. (14) Mu¨ller, L.; Wetzel, R.; Otto, H. J. Electroanal. Chem. 1970, 24, 175. (15) Mark, H. B.; Anson, F. C. J. Electroanal. Chem. 1963, 6, 251. (16) Burke, L. D.; O’Sullivan, J. F.; O’Dwyer, K. J.; Scannell, R. A.; Ahern, M. J. G.; McCarthy, M. M. J. Electrochem. Soc. 1990, 137, 2476. (17) Desilvestro, J.; Weaver, M. J. J. Electroanal. Chem. 1987, 234, 237. (18) Ye, J.; Wolf, W.; Doblhofer, K.; Eiswirth, M. Unpublished work. Wolf, W. Ph.D. Thesis, Frei Universita¨t Berlin, Berlin, 1994. (19) Ha¨rtinger, S.; Pettinger, B.; Doblhofer, K. J. Electroanal. Chem. 1995, 397, 335. (20) Soares, D. M. Meas. Sci. Technol. 1993, 4, 549. (21) Frubo¨se, C.; Doblhofer, K.; Soares, D. M. Ber. Bunsen-Ges. Phys. Chem. 1993, 97, 475. (22) Sauerbrey, G. Z. Phys. 1959, 155, 206. (23) Albery, W. J.; Bruckenstein, S. Trans. Faraday Soc. 1966, 62, 1920. (24) Ikemiya, N.; Kubo, T.; Hara, S. Surf. Sci. 1995, 323, 81. (25) Strehblow, H. H.; Titze, B. Electrochim. Acta 1980, 25, 839. (26) Shirkhanzadeh, M.; Thompson, G. E.; Ashworth, V. Corr. Sci. 1990, 31, 293. (27) Miller, B. J. Electrochem. Soc. 1969, 116, 1675. (28) Tindall, G. W.; Bruckenstein, S. J. Electroanal. Chem. 1969, 22, 367. (29) Schubert, H.; Tegetmeyer, U.; Schlo¨gl, R. Cat. Lett. 1994, 28, 383. (30) Schedl-Niedrig, T.; Bao, X.; Muhler, M.; Neisius, T.; Schlo¨gl, R. Annual Fachbeirat of Fritz-Haber-Institut of the Max-Planck-Gesellschaft, Abstract AC12, Berlin, 1995. (31) Werner, H.; Demuth, D.; Schubert, H.; Weinberg, G.; Braun, T.; Herein, D.; Schlo¨gl, R. Annual Fachbeirat of Fritz-Haber-Institut of the Max-Planck-Gesellschaft, Abstract AC13, Berlin, 1995. (32) Polak, M. Surf. Sci. 1994, 321, 249. (33) Carley, A. F.; Davies, P. R.; Roberts, M. W.; Vicent, D. Top. Catal. 1994, 1, 35. (34) Davies, P. R.; Roberts, M. W.; Shukla, N.; Vicent, D. Surf. Sci. 1995, 325, 50.
