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O2 reduction and CO oxidation at the Pt-electrolyte interface. The role of H2O and OH adsorption bond strengths
Electrochimica Acta 47 (2002) 3759 /3763 www.elsevier.com/locate/electacta
O2 reduction and CO oxidation at the Pt-electrolyte interface. The role of H2O and OH adsorption bond strengths Alfred B. Anderson * Chemistry Department, Case Western Reserve University, 10900 Euclid Avenue, Cleveland, OH 44106-7078, USA Received 8 November 2001; received in revised form 1 March 2002
Abstract The nature of surface sites on which OH(ads) forms in acid and basic electrolytes is discussed based on the analysis of experimental results in the literature and the results of quantum calculations. Theoretical evidence is given for OH(ads) forming in acid solution on Pt surfaces by the reaction PtOH2 OH2 (OH)2 l PtOH H OH2 (OH2 )2 e (U) (i) and an analysis of experimental results suggests that in base it forms by the reaction PtOH2 OH (HOH)2 l PtOH HOH(HOH)2 e (U) (ii) The reversible potential for reaction (i) is calculated to be 0.62 V, which is essentially the onset potential for OH(ads) formation in weak acid electrolytes. It is suggested that OH(ads) forms at potentials as low as /0.17 V in weak basic electrolyte by reaction (ii) where H2O molecules bond by lone-pair donation to unblocked Ptd sites of the hydrided electrode surface, and that as the potential is increased to the double layer region, beginning at 0.4 V, H2O no longer bonds with the surface to participate in this reaction. This behavior would explain the /0.3 V prewave in CO(ads) oxidation by OH(ads). At the potential of zero charge, /0.6 V, H2O again adsorbs and the reaction resumes. It is concluded based on experimental and theoretical evidence that four-electron O2 reduction on Pt electrodes at low overpotentials requires a special site characterized by a relatively small ratio of OH to H2O adsorption bond strengths and or a higher activation energy for OH(ads) formation by reaction (i), as well as a weak ability to adsorb anions. # 2002 Elsevier Science Ltd. All rights reserved. Keywords: OH(ads) formation; CO oxidation; O2 reduction; Theory
1. Introduction At the solid /liquid interface the O2 reduction reaction of electrochemical interest takes O2 to water by the route of combining it with four protons from the electrolyte and four electrons from the cathode. This topic has been reviewed frequently over the years [1 /4]. Platinum and platinum alloys are still, despite efforts to identify replacements, the best known electrocatalysts for this reaction. The standard reversible potential for O2 reduction is 1.23 V on the hydrogen scale, but due to kinetic effects, the oxygen cathode in a fuel cell has a working potential of around 0.8 V so that the overpotential is around 400 mV.
* Tel.: /1-216-368-5044; fax: /1-216-368-3006 E-mail address: [email protected] (A.B. Anderson).
The overpotential at which the oxygen cathode operates is believed to be influenced by adsorbed molecules that are blocking the approach of O2 molecules to the surface sites where they can undergo reduction. Much of the evidence has been reviewed [5]. It is presently thought that in weak perchloric acid OH(ads), formed on Pt(111) from the oxidation of water: H2 O(l) 0 OH(ads)H (aq)e (U)
(1)
blocks the surface and is the cause of the reduction in surface activity at the positive end of the 0.4 /0.6 V double layer region. The OH(ads) formation current begins to flow at 0.6 V and the rate of increase slows at 0.7 V, and then the current rises to a sharp peak at 0.8 V, and it finally drops to a low value as the potential is increased further. There is also vibrational spectroscopic evidence, visible first at 0.7 V as the potential is
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increased, that OH(ads) is the product of the oxidation reaction [6]. The low index (100) and (110) platinum surfaces show effects of OH(ads) on O2 reduction being similar to observations of the (111) surface, though a noted difference for these surfaces is that the formation of OH(ads) shows less reversible behavior in the cyclic voltammograms compared with the (111) surface and there is no sharp peak at 0.7 V [4,7]. Other platinum electrode surface-blocking adsorbates include OH (aq) in basic electrolyte and sulfate and phosphate anions in the corresponding acid electrolytes [4,7]. Cyclic voltammograms in basic solution show profiles in the 0.6 /0.8 V region that are broadly similar to those obtained in perchloric acid for the three low index faces, and so it is presumed that OH (aq) is oxidized at the surfaces to OH(ads). The (111) surface is particularly strongly blocked by bisulfate which has the proper structure to bond through three oxygen atoms, and therefore, strongly, to three platinum atoms. This results in a much smaller current for OH(ads) formation from 0.6 to 0.8 V and about a 100 mV increase in overpotential for O2 reduction on the (111) surface compared with the other two surfaces [4]. Oxygen reduction activity over the other two surfaces in sulfuric acid shows a fairly small inhibition compared with carrying it out in perchloric acid [4], and from this it seems that OH(ads) becomes the predominant blocking agent in weak sulfuric acid electrolytes on these surfaces too. It is believed by some that one mechanism involving one type of active surface site accounts for the fourelectron reduction of O2 on the various surfaces in these electrolytes. A compelling reason for this supposition is the finding from temperature-dependent Tafel plots the same 42 kJ mol 1 activation energy for the three lowindexsurfaces in dilute sulfuric acid [4]. Model calculations of the potential dependence of O2 reduction show it to be easily reduced to H2O2 when initially bonded end-on to a platinum atom but H2O2 has a high barrier to reduction at the same site [8]. This appears to be consistent with the observed reduction to hydrogen peroxide in the hydrogen upd potential region on Pt(111) where the surface is at least partially blocked by hydrogen atoms. Recent further theoretical work [9] with the same theoretical approach shows that for bridge-bonded O2 the step with the highest activation energy is the first step: O2 (ads)H (aq)e (U) 0 HOO(ads)
(2)
The adsorbed HOO is calculated to dissociate with a small activation energy to OH(ads) and O(ads) [8] and these adsorbates are reduced to adsorbed water molecules over activation barriers that are much smaller at all electrode potentials than for Eq. (2) [8,9]. It is of interest to understand the potential dependence of OH(ads) formation and the relationship to
calculable and measurable bond strengths. We believe the adsorption bond energies of H2O and OH are key, rather than the potential dependent activation energies for forming OH(ads) by Eq. (1), because recent calculations show that the activation energy barrier for this reaction is small near the reversible potential [10]. This means that OH(ads) formation is under thermodynamic instead of kinetic control. The purpose of this paper is to provide an understanding of the potential dependence of OH(ads) formation and its relationship to the electrooxidation of CO(ads) and to the electroreduction of O2. Reference will be made to calculated and experimental surface/gas interface properties and electrochemical surface/aqueous electrolyte interface properties.
2. Results and discussion 2.1. Possible relationship between cyclic voltammograms for OH(ads) and calculated coverage-dependent OH adsorption energies at the vacuum interface The shape of the butterfly region of cyclic voltammograms of Pt(111) electrodes in dilute perchloric acid indicates that there is an onset potential for OH(ads) formation at 0.6 V and as the potential increases at the rate of 50 mV s 1 (a typically used rate) the current begins to level out but then rises sharply to a spike at 0.8 V and then drops to nearly zero [4,7]. The spike behavior is characteristic of a phase transition and it has been proposed that adsorption of Cl that is present in perchloric acid may be involved [5]. The coverage by OH at the peak has been determined to be about 1/2 monolayer (ML) [7]. The onset potential is the reversible potential in the low-coverage limit. The width of the current signal is influenced by the scan rate, though slow, but also by surface heterogeneity. Interestingly, recent density functional theory band calculations of OH structures and adsorption energies at various ordered coverages of 1/9, 1/6, 1/4, and 1/3 ML show the 1-fold top sites to be 0.02 /0.05 eV more stable than 2-fold bridge sites [11]. The adsorption energies per OH were, respectively, 2.25, 2.29, 2.27 and 2.31 eV. These numbers are fairly constant and at 1/2 ML coverage the adsorption energy per OH increases to 2.49 eV because of hydrogen bonding between the adsorbed OH molecules. At 2/3, 3/4, and 1 ML the respective energies are 2.53, 2.44, and 2.50 eV and chains of hydrogen bonded OH(ads) are the most stable structures to 3/4 ML. If the sudden increase in the average OH adsorption bond strength at 1/2 ML coverage also takes place at the electrochemical interface, then this can tentatively be considered as a contributor to the sharp peak at 0.8 V in the cyclic volatmmograms. However, the strength of the hydrogen bonds between the adsorbed OH is low, at 0.2
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eV, and this stability might not be enough to overcome the energetic cost of displacing hydrogen bonds to water molecules in the double layer, as would be expected to take place during their formation. If adsorbed Cl plays a role, it might be by participating in the hydrogen bonding. At this potential Cl , if oxidized and adsorbed, would still be in the form of Cl (ads), having gained, just like OH(ads), its negative charge from the Pt atoms to which it is bonded. Adsorbed bisulfate, present during OH(ads) formation in weak sulfuric acid electrolytes, reduces the oxidation current by blocking surface sites, and no OH(ads) forming current spikes are observed over the three low index platinum surfaces, which implies that the transition is blocked. However, voltammogram signals for OH(ads) formation on Pt(111) in 0.1 M potassium hydroxide and in 0.1 M perchloric acid have similar width and peak positions [4], indicating that the same transition is occurring and that chloride is not an issue. Monte Carlo simulations have indicated that second-order order /disorder phase transitions of OH(ads) can lead to the observed sharply peaked voltammograms, and Cl was absent from the model [12]. Therefore, Cl may be perturbing the spike in current but not creating it.
may be unavailable for hydrogen bonding to the OH(ads) that forms at high coverage of OH(ads) alone, and may also weaken the adsorption bond of OH that is bonded to adjacent atop sites. The /1.0 eV underestimate in the experimental determination of the Pt /OH bond energy for Eq. (4) is proposed by this author to be due to this equation not corresponding to the process that is actually taking place at the electrode surface. OH(ads) begins to form at 0.6 V is where the surface charge is becoming positive (the potential of zero charge is about 0.6 V, which is at the high end of the double layer potential range [5,10]). Water molecules should be bonded to the surface through oxygen at this potential and at the 0.8 V current peak potential. Furthermore, from available data one can estimate for the solution reaction
2.2. Model for reversible OH(ads) formation in acid electrolyte by Eq. (1)
Comparing Eqns. (4) and (6), it is evident that the isosteric heat of adsorption of OH can be used to define the Pt /OH adsorption bond strength for Eq. (4) only after the Pt/OH2 adsorption bond strength is added to the 1.4 eV that was determined. The Pt/OH2 bond strength has been determined by thermal desorption studies to be 0.65 eV [14], and so the final Pt/OH(ads) bond strength should be /2.0 eV, which is within 0.3 and 0.5 eV of the calculated low and high coverage values reported in Ref. [11] and within 0.5 eV of the measured low-coverage value for the vacuum interface in Ref. [13]. There may be several causes for the remaining 0.3 /0.5 eV differences and it is beyond the scope of this work to try to evaluate them now. This lab has recently studied the OH(ads) formation reaction according to Eq. (6), determining an approximate reversible potential for it [8] and approximate electrode potential dependent activation energies for it [10]. The ab initio MP2 approach was used along with a single platinum atom to represent the 1-fold adsorption site. The reversible potential 0.57 V was obtained based on the reaction energy for Eq. (6) and an empiricallybased 0.49 eV correction corresponding to enthalpic, entropic, counterion, and overall electrolyte contributions to the Gibbs free energy of reaction [15]. The oxidation study was based on a model for Eq. (6) as follows:
The isosteric heat of adsorption of OH, DHu (OH), has been determined using the formula DHu (OH) @(DGad (OH)=T)u =@T 1
(3)
at u $/1/2 ML coverage on Pt(111), that is, at the potential of the 0.8 V voltammogram peak in alkaline solution [7]. The authors of Ref. [7], noting that the current peaks in alkaline and dilute perchloric acid solutions are at the same potential, proposed that the isosteric heats are the same in the two electrolytes but used measurements for the basic electrolyte out of concern over possible interference from impurity chloride ions that are present in perchloric acid. Using the obtained value for DHu(OH) of /200 kJ mol 1 with the standard enthalpy of formation of OH(g), the Pt / OH bond enthalpy of /136 kJ mol 1 or 1.4 eV was calculated for the reaction H2 OPt(s) 0 Pt OH(ads)H (aq)e (U)
(4)
The Pt/OH bond energy should be close to the bond enthalpy but the value 1.4 eV is a full 1.1 eV less than the calculated values of Ref. [11] for this coverage and 0.9 eV less than the low-coverage value. It is noted that the bond strength of OH to a Pt wire has been found to decrease with increasing p (O2)/p (H2) pressure ratio from about 2.5 to 1.5 eV [13], but it is an increasing O(ads) coverage that is evidently responsible for the decrease: O(ads), sitting in high-coordinate surface hollow sites,
H2 O(aq) 0 OH(aq)H (aq)e (U)
(5)
a reversible potential of /3 V, which means that oxidation at 0.8 V must take place with water bonded to the surface so that OH is adsorbed upon electron transfer. The actual reaction that is taking place should then be PtOH2 0 PtOHH (aq)e (U)
(6)
PtOH2 OH2 (OH2 )2 l PtOH H OH2 (OH2 )2 e (U)
(7)
where electron transfer takes place at a transition state
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which has an intermediate structure, Pt / OH H OH2(OH2)2 for which the ionization potential is equal to 4.6 eV/(U/V) eV. As discussed in Ref. [10], the counterion and electrolyte Madelung fields were approximated by placing a /1/2 point charge ˚ from the O. The resulting along the H /O axis 10 A calculated reaction energy for Eq. (7), with the electron at zero on the vacuum scale, is 5.22 eV. Taking this as the reaction Gibbs free energy, which is a reasonable approximation based on the conserved bond order and small change in structure on going the reduced and oxidized states in Eq. (7), the predicted reversible potential is 0.62 V, and this is essentially the same as the potential of zero charge and as the onset potential for OH(ads) formation and is 0.05 V greater than the predicted value mentioned above. Proton migration is probably by a vacancy diffusion mechanism, the vacancies being formed at the cathode and migrating to the anode to consume the H . The diffusion process is rapid and so the activation barrier for the proton leaving the double layer that is represented on the right-hand side of Eq. (7) should be small so that the reversible potential determined above for this equation will be close to that which would be calculated for the overall process of presenting the proton to the bulk solution phase of the electrolyte. It is noted that in the reaction model of Eq. (7) the hydrogen-bonded water trimer on the left and the partially solvated hydronium ion on the right are part of the double layer. The double layer is a reactant and is included in the model calculations. 2.3. Model for reversible OH(ads) formation in basic electrolyte For OH(ads) to form in basic electrolyte with the same onset potential as reaction (7) according to the reaction OH (aq)Pt(s) 0 PtOH(ads)e (U)
(8)
suggested in Ref. [7] would appear to be a coincidence, and it might be thought that adsorbed water would be the species actually oxidized. Using the standard reversible potential of 2.02 V for aqueous OH reduction to aqueous OH and the Pt/OH bond strength value of 2.50 eV, the reversible potential for Eq. (8) would be / 0.48 V, which is deep in the potential range of hydrogen evolution and far from the observed 0.6 V onset potential. However, the same correction for H2O displacement that was made to the oxidation reaction in acid electrolyte might be assumed to apply in this case too. The rationale for this would be that in the hydrogen upd range H(ads) polarizes the surface, forming Hd and Ptd at the surface. This will result in lone-pair bonding of H2O to these Pt atoms. Thus the appropriate reaction in base would then be
PtOH2 OH (aq) 0 PtOHH2 Oe (U)
(9)
When the water adsorption bond strength of 0.65 eV is included, the predicted reversible potential becomes 0.17 V. At electrode potentials positive of this, then, it would be thermodynamically possible to form OH(ads) in basic solution by discharge if OH (aq). Relative solvation energies of adsorbed H2O and OH could change this reversible potential prediction as could changes in the relative adsorption energies. Mechanisms for this reaction have not yet been explored theoretically. Two candidates come immediately to mind, displacement of adsorbed H2O by OH with concomitant or subsequent electron transfer or H abstraction with simultaneous electron transfer. The former would require displacement of strongly held solvating water molecules from the small OH anion, which might require too much energy, and the latter could be modeled in much the same way was done for oxidation of adsorbed H2O in Eq. (7): PtOH2 OH (HOH)2 l PtOH HOH(HOH)2 e (U)
(10)
where two water molecules are hydrogen bonded to the two lone pairs of OH and the product water molecule. Work on this mechanism has been initiated in the author’s lab. Might OH (aq) actually discharge at /0.17 V to form OH(ads)? Recent studies of CO electrooxidation, where in OH(ads) is believed, from experimental analysis [7,16] and from quantum mechanical model calculations [17] to be the oxidant for CO(ads), show evidence for some OH(ads) being formed at potentials as low as 0.2 V in base [7,16]. The oxidation current density for CO in solution in this ‘prewave’ potential region, which has a weak peak at 0.3 V and weakens at 0.4 V, the edge of the H upd range, is much smaller than the current densities that are associated with 0.6 V OH(ads) formation onset potential discussed earlier. This adsorbed OH is in low surface concentration and is believed to be bonded to surface defects on Pt(111) and to step sites on the (100) and (110) surfaces [7,16]. In acid electrolytes, by comparison, these defect and step sites are believed to be blocked by adsorbed anions. If OH bonds relatively strongly compared with H2O at defect sites, then, by Eq. (10), the reversible potential for its formation will be smaller than when it bonds relatively weakly. The drop in CO oxidation current at 0.4 V can be associated with the absence of upd hydrogen at this potential and the fact that this potential is /0.2 V less than the pzc, with the result that water molecules will not form strong lone-pair donation bonds to the surface and will therefore have difficulty in being activated for generating OH(ads). The higher onset potential for forming OH(ads) on non-defect and nonstep sites in the basic electrolyte would be the conse-
A.B. Anderson / Electrochimica Acta 47 (2002) 3759 /3763
quence of a smaller difference between OH and H2O adsorption bond strengths at these sites. 2.4. The adsorption bond energy of O2 and its relationship to reduction In the previously discussed density functional calculations of O2 reduction, O2 was calculated to bond to a Pt2 site with a bond strength of 0.43 eV when upright on a 1fold site and 0.94 eV when lying down in the di-s structure [9]. Density functional band calculations have yielded 0.72 eV for the di-s site on the (111) surface [18] and other density functional cluster calculations produce energies in the range 0.53 /0.83 eV for O2 in this configuration [19]. Experimental determinations for the adsorption bond strength of O2 on Pt(111) are 0.4 [20] and 0.5 eV [21], but a competing dissociation reaction forming adsorbed O atoms complicated the analysis. In the electrode potential range from 0.5 to 0.9 V, O2 reduction is a four-electron process in acid electrolyte, with curved Tafel plots [4]. The observed curvature could in part be caused by the potential dependence of the activation energy for the first step, according to theory discussed above [9]. Kinetic models incorporating a dependence on coverage of OH or O site blocking species have accounted for the limiting Tafel slopes in the high overpotential and low overpotential regions, 2RT /F and RT /F , respectively [22]. The relative contributions of these two effects to the observed change in Tafel slopes is an unresolved issue. Despite the formation of OH(ads) from oxidation of H2O(ads) by Eq. (6), and despite the adsorption of anions, sites evidently exist for O2 coordination, reduction to OOH(ads), and dissociation to easily reduced O(ads) and OH(ads). The nature of these sites would be very interesting to elucidate. They will be characterized by either a relatively weak OH adsorption bond strength, so that the reversible potential for OH(ads) formation by Eq. (7) is pushed to /0.9 V or higher (this assumes the formation is under thermodynamic control), or by an increase in the activation energy for OH(ads) formation at these sites (this would bring about kinetic control). It is also possible for a combination of these two effects to be responsible for the absence of OH adsorption on these sites. These sites will also be required to adsorb anions weakly so that O2 can displace them.
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Acknowledgements I thank my coworkers, Dr. T.V. Albu, Dr. R.A. Sidik, Dr. N.M. Neshev, and Dr. P. Shiller for their many calculations and discussions that led up to the work referenced herein, some of which is published and some of which is submitted for publication. This work is supported by the National Science Foundation, Grant No. CHE-9982179, and by the Army Research Office, Grant No. DAAD 19-99-1-0253.
References [1] E. Yeager, Electrochim. Acta 29 (1984) 1527. [2] J.O’M. Bockris, S.U.M. Khan, Surface Electrochemistry (Section 3.23), Plenum Press, New York, 1993. [3] R. Adzic, in: J. Lipkowski, P.N. Ross (Eds.), Electrocatalysis, Wiley, New York, 1998, pp. 197 /242. [4] N.M. Markovic, P.N. Ross, in: A. Wieckowski (Ed.), Interfacial Electrochemistry. Theory, Experiment and Applications, Marcel Dekker, New York, 1999, pp. 821 /841. [5] N. Markovic, P.N. Ross, J. Electroanal. Chem. 330 (1992) 499. [6] T. Iwasita, X. Xia, J. Electroanal. Chem. 411 (1995) 95. [7] N.M. Markovic, T.J. Schmidt, B.N. Grgur, H.A. Gasteiger, R.J. Behm, P.N. Ross, J. Phys. Chem. B103 (1999) 8568. [8] A.B. Anderson, T.V. Albu, J. Electrochem. Soc. 147 (2000) 4229. [9] R.A. Sidik, A.B. Anderson, Density functional study of O2 electroreduction when bonded to a Pt dual site, J. Electroanal. Chem. 528 (2002) 69. [10] A.B. Anderson, N.M. Neshev, R.A. Sidik, P. Shiller, Mechanism for the electrooxidation of water to OH(ads) and O(ads) on platinum: quantum chemical theory, Electrochim. Acta 47 (2002) 2999. [11] A. Michaelides, P. Hu, J. Chem. Phys. 114 (2001) 513. [12] M.T.M. Koper, J.J. Lukkien, J. Electroanal. Chem. 485 (2000) 161. [13] C.E. Mooney, L.C. Anderson, J.H. Lunsford, J. Phys. Chem. 97 (1993) 2505. [14] G.B. Fisher, J.L. Gland, Surf. Sci. 94 (1980) 446. [15] A.B. Anderson, T.V. Albu, J. Am. Chem. Soc. 121 (1999) 11855. [16] T.J Schmidt, P.N. Ross, N.M. Markovic, Temperature-dependent surface electrochemistry on Pt single crystals in alkaline electrolyte: Part 1: CO oxidation, J. Phys. Chem. B 105 (2002) 12082. [17] A.B. Anderson, N.M. Neshev, Mechanism for the electrooxidation of carbon monoxide on platinum, including electrode potential dependence. Theoretical determination, J. Electrochem. Soc., in press. [18] A. Eichler, J. Hafner, Phys. Rev. Lett. 79 (1997) 4481. [19] T. Li, P.B. Balbuena, J. Phys. Chem. B 105 (2001) 9943. [20] J.L. Gland, B.A. Sexton, G.B. Fisher, Surf. Sci. 95 (1980) 587. [21] H. Steininger, H. Ibach, Surf. Sci. 123 (1982) 1. [22] A. Damjanovic, V. Brusic, Electrochem. Acta 12 (1967) 615.
O2 reduction and CO oxidation at the Pt-electrolyte interface. The role of H2O and OH adsorption bond strengths Alfred B. Anderson * Chemistry Department, Case Western Reserve University, 10900 Euclid Avenue, Cleveland, OH 44106-7078, USA Received 8 November 2001; received in revised form 1 March 2002
Abstract The nature of surface sites on which OH(ads) forms in acid and basic electrolytes is discussed based on the analysis of experimental results in the literature and the results of quantum calculations. Theoretical evidence is given for OH(ads) forming in acid solution on Pt surfaces by the reaction PtOH2 OH2 (OH)2 l PtOH H OH2 (OH2 )2 e (U) (i) and an analysis of experimental results suggests that in base it forms by the reaction PtOH2 OH (HOH)2 l PtOH HOH(HOH)2 e (U) (ii) The reversible potential for reaction (i) is calculated to be 0.62 V, which is essentially the onset potential for OH(ads) formation in weak acid electrolytes. It is suggested that OH(ads) forms at potentials as low as /0.17 V in weak basic electrolyte by reaction (ii) where H2O molecules bond by lone-pair donation to unblocked Ptd sites of the hydrided electrode surface, and that as the potential is increased to the double layer region, beginning at 0.4 V, H2O no longer bonds with the surface to participate in this reaction. This behavior would explain the /0.3 V prewave in CO(ads) oxidation by OH(ads). At the potential of zero charge, /0.6 V, H2O again adsorbs and the reaction resumes. It is concluded based on experimental and theoretical evidence that four-electron O2 reduction on Pt electrodes at low overpotentials requires a special site characterized by a relatively small ratio of OH to H2O adsorption bond strengths and or a higher activation energy for OH(ads) formation by reaction (i), as well as a weak ability to adsorb anions. # 2002 Elsevier Science Ltd. All rights reserved. Keywords: OH(ads) formation; CO oxidation; O2 reduction; Theory
1. Introduction At the solid /liquid interface the O2 reduction reaction of electrochemical interest takes O2 to water by the route of combining it with four protons from the electrolyte and four electrons from the cathode. This topic has been reviewed frequently over the years [1 /4]. Platinum and platinum alloys are still, despite efforts to identify replacements, the best known electrocatalysts for this reaction. The standard reversible potential for O2 reduction is 1.23 V on the hydrogen scale, but due to kinetic effects, the oxygen cathode in a fuel cell has a working potential of around 0.8 V so that the overpotential is around 400 mV.
* Tel.: /1-216-368-5044; fax: /1-216-368-3006 E-mail address: [email protected] (A.B. Anderson).
The overpotential at which the oxygen cathode operates is believed to be influenced by adsorbed molecules that are blocking the approach of O2 molecules to the surface sites where they can undergo reduction. Much of the evidence has been reviewed [5]. It is presently thought that in weak perchloric acid OH(ads), formed on Pt(111) from the oxidation of water: H2 O(l) 0 OH(ads)H (aq)e (U)
(1)
blocks the surface and is the cause of the reduction in surface activity at the positive end of the 0.4 /0.6 V double layer region. The OH(ads) formation current begins to flow at 0.6 V and the rate of increase slows at 0.7 V, and then the current rises to a sharp peak at 0.8 V, and it finally drops to a low value as the potential is increased further. There is also vibrational spectroscopic evidence, visible first at 0.7 V as the potential is
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increased, that OH(ads) is the product of the oxidation reaction [6]. The low index (100) and (110) platinum surfaces show effects of OH(ads) on O2 reduction being similar to observations of the (111) surface, though a noted difference for these surfaces is that the formation of OH(ads) shows less reversible behavior in the cyclic voltammograms compared with the (111) surface and there is no sharp peak at 0.7 V [4,7]. Other platinum electrode surface-blocking adsorbates include OH (aq) in basic electrolyte and sulfate and phosphate anions in the corresponding acid electrolytes [4,7]. Cyclic voltammograms in basic solution show profiles in the 0.6 /0.8 V region that are broadly similar to those obtained in perchloric acid for the three low index faces, and so it is presumed that OH (aq) is oxidized at the surfaces to OH(ads). The (111) surface is particularly strongly blocked by bisulfate which has the proper structure to bond through three oxygen atoms, and therefore, strongly, to three platinum atoms. This results in a much smaller current for OH(ads) formation from 0.6 to 0.8 V and about a 100 mV increase in overpotential for O2 reduction on the (111) surface compared with the other two surfaces [4]. Oxygen reduction activity over the other two surfaces in sulfuric acid shows a fairly small inhibition compared with carrying it out in perchloric acid [4], and from this it seems that OH(ads) becomes the predominant blocking agent in weak sulfuric acid electrolytes on these surfaces too. It is believed by some that one mechanism involving one type of active surface site accounts for the fourelectron reduction of O2 on the various surfaces in these electrolytes. A compelling reason for this supposition is the finding from temperature-dependent Tafel plots the same 42 kJ mol 1 activation energy for the three lowindexsurfaces in dilute sulfuric acid [4]. Model calculations of the potential dependence of O2 reduction show it to be easily reduced to H2O2 when initially bonded end-on to a platinum atom but H2O2 has a high barrier to reduction at the same site [8]. This appears to be consistent with the observed reduction to hydrogen peroxide in the hydrogen upd potential region on Pt(111) where the surface is at least partially blocked by hydrogen atoms. Recent further theoretical work [9] with the same theoretical approach shows that for bridge-bonded O2 the step with the highest activation energy is the first step: O2 (ads)H (aq)e (U) 0 HOO(ads)
(2)
The adsorbed HOO is calculated to dissociate with a small activation energy to OH(ads) and O(ads) [8] and these adsorbates are reduced to adsorbed water molecules over activation barriers that are much smaller at all electrode potentials than for Eq. (2) [8,9]. It is of interest to understand the potential dependence of OH(ads) formation and the relationship to
calculable and measurable bond strengths. We believe the adsorption bond energies of H2O and OH are key, rather than the potential dependent activation energies for forming OH(ads) by Eq. (1), because recent calculations show that the activation energy barrier for this reaction is small near the reversible potential [10]. This means that OH(ads) formation is under thermodynamic instead of kinetic control. The purpose of this paper is to provide an understanding of the potential dependence of OH(ads) formation and its relationship to the electrooxidation of CO(ads) and to the electroreduction of O2. Reference will be made to calculated and experimental surface/gas interface properties and electrochemical surface/aqueous electrolyte interface properties.
2. Results and discussion 2.1. Possible relationship between cyclic voltammograms for OH(ads) and calculated coverage-dependent OH adsorption energies at the vacuum interface The shape of the butterfly region of cyclic voltammograms of Pt(111) electrodes in dilute perchloric acid indicates that there is an onset potential for OH(ads) formation at 0.6 V and as the potential increases at the rate of 50 mV s 1 (a typically used rate) the current begins to level out but then rises sharply to a spike at 0.8 V and then drops to nearly zero [4,7]. The spike behavior is characteristic of a phase transition and it has been proposed that adsorption of Cl that is present in perchloric acid may be involved [5]. The coverage by OH at the peak has been determined to be about 1/2 monolayer (ML) [7]. The onset potential is the reversible potential in the low-coverage limit. The width of the current signal is influenced by the scan rate, though slow, but also by surface heterogeneity. Interestingly, recent density functional theory band calculations of OH structures and adsorption energies at various ordered coverages of 1/9, 1/6, 1/4, and 1/3 ML show the 1-fold top sites to be 0.02 /0.05 eV more stable than 2-fold bridge sites [11]. The adsorption energies per OH were, respectively, 2.25, 2.29, 2.27 and 2.31 eV. These numbers are fairly constant and at 1/2 ML coverage the adsorption energy per OH increases to 2.49 eV because of hydrogen bonding between the adsorbed OH molecules. At 2/3, 3/4, and 1 ML the respective energies are 2.53, 2.44, and 2.50 eV and chains of hydrogen bonded OH(ads) are the most stable structures to 3/4 ML. If the sudden increase in the average OH adsorption bond strength at 1/2 ML coverage also takes place at the electrochemical interface, then this can tentatively be considered as a contributor to the sharp peak at 0.8 V in the cyclic volatmmograms. However, the strength of the hydrogen bonds between the adsorbed OH is low, at 0.2
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eV, and this stability might not be enough to overcome the energetic cost of displacing hydrogen bonds to water molecules in the double layer, as would be expected to take place during their formation. If adsorbed Cl plays a role, it might be by participating in the hydrogen bonding. At this potential Cl , if oxidized and adsorbed, would still be in the form of Cl (ads), having gained, just like OH(ads), its negative charge from the Pt atoms to which it is bonded. Adsorbed bisulfate, present during OH(ads) formation in weak sulfuric acid electrolytes, reduces the oxidation current by blocking surface sites, and no OH(ads) forming current spikes are observed over the three low index platinum surfaces, which implies that the transition is blocked. However, voltammogram signals for OH(ads) formation on Pt(111) in 0.1 M potassium hydroxide and in 0.1 M perchloric acid have similar width and peak positions [4], indicating that the same transition is occurring and that chloride is not an issue. Monte Carlo simulations have indicated that second-order order /disorder phase transitions of OH(ads) can lead to the observed sharply peaked voltammograms, and Cl was absent from the model [12]. Therefore, Cl may be perturbing the spike in current but not creating it.
may be unavailable for hydrogen bonding to the OH(ads) that forms at high coverage of OH(ads) alone, and may also weaken the adsorption bond of OH that is bonded to adjacent atop sites. The /1.0 eV underestimate in the experimental determination of the Pt /OH bond energy for Eq. (4) is proposed by this author to be due to this equation not corresponding to the process that is actually taking place at the electrode surface. OH(ads) begins to form at 0.6 V is where the surface charge is becoming positive (the potential of zero charge is about 0.6 V, which is at the high end of the double layer potential range [5,10]). Water molecules should be bonded to the surface through oxygen at this potential and at the 0.8 V current peak potential. Furthermore, from available data one can estimate for the solution reaction
2.2. Model for reversible OH(ads) formation in acid electrolyte by Eq. (1)
Comparing Eqns. (4) and (6), it is evident that the isosteric heat of adsorption of OH can be used to define the Pt /OH adsorption bond strength for Eq. (4) only after the Pt/OH2 adsorption bond strength is added to the 1.4 eV that was determined. The Pt/OH2 bond strength has been determined by thermal desorption studies to be 0.65 eV [14], and so the final Pt/OH(ads) bond strength should be /2.0 eV, which is within 0.3 and 0.5 eV of the calculated low and high coverage values reported in Ref. [11] and within 0.5 eV of the measured low-coverage value for the vacuum interface in Ref. [13]. There may be several causes for the remaining 0.3 /0.5 eV differences and it is beyond the scope of this work to try to evaluate them now. This lab has recently studied the OH(ads) formation reaction according to Eq. (6), determining an approximate reversible potential for it [8] and approximate electrode potential dependent activation energies for it [10]. The ab initio MP2 approach was used along with a single platinum atom to represent the 1-fold adsorption site. The reversible potential 0.57 V was obtained based on the reaction energy for Eq. (6) and an empiricallybased 0.49 eV correction corresponding to enthalpic, entropic, counterion, and overall electrolyte contributions to the Gibbs free energy of reaction [15]. The oxidation study was based on a model for Eq. (6) as follows:
The isosteric heat of adsorption of OH, DHu (OH), has been determined using the formula DHu (OH) @(DGad (OH)=T)u =@T 1
(3)
at u $/1/2 ML coverage on Pt(111), that is, at the potential of the 0.8 V voltammogram peak in alkaline solution [7]. The authors of Ref. [7], noting that the current peaks in alkaline and dilute perchloric acid solutions are at the same potential, proposed that the isosteric heats are the same in the two electrolytes but used measurements for the basic electrolyte out of concern over possible interference from impurity chloride ions that are present in perchloric acid. Using the obtained value for DHu(OH) of /200 kJ mol 1 with the standard enthalpy of formation of OH(g), the Pt / OH bond enthalpy of /136 kJ mol 1 or 1.4 eV was calculated for the reaction H2 OPt(s) 0 Pt OH(ads)H (aq)e (U)
(4)
The Pt/OH bond energy should be close to the bond enthalpy but the value 1.4 eV is a full 1.1 eV less than the calculated values of Ref. [11] for this coverage and 0.9 eV less than the low-coverage value. It is noted that the bond strength of OH to a Pt wire has been found to decrease with increasing p (O2)/p (H2) pressure ratio from about 2.5 to 1.5 eV [13], but it is an increasing O(ads) coverage that is evidently responsible for the decrease: O(ads), sitting in high-coordinate surface hollow sites,
H2 O(aq) 0 OH(aq)H (aq)e (U)
(5)
a reversible potential of /3 V, which means that oxidation at 0.8 V must take place with water bonded to the surface so that OH is adsorbed upon electron transfer. The actual reaction that is taking place should then be PtOH2 0 PtOHH (aq)e (U)
(6)
PtOH2 OH2 (OH2 )2 l PtOH H OH2 (OH2 )2 e (U)
(7)
where electron transfer takes place at a transition state
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which has an intermediate structure, Pt / OH H OH2(OH2)2 for which the ionization potential is equal to 4.6 eV/(U/V) eV. As discussed in Ref. [10], the counterion and electrolyte Madelung fields were approximated by placing a /1/2 point charge ˚ from the O. The resulting along the H /O axis 10 A calculated reaction energy for Eq. (7), with the electron at zero on the vacuum scale, is 5.22 eV. Taking this as the reaction Gibbs free energy, which is a reasonable approximation based on the conserved bond order and small change in structure on going the reduced and oxidized states in Eq. (7), the predicted reversible potential is 0.62 V, and this is essentially the same as the potential of zero charge and as the onset potential for OH(ads) formation and is 0.05 V greater than the predicted value mentioned above. Proton migration is probably by a vacancy diffusion mechanism, the vacancies being formed at the cathode and migrating to the anode to consume the H . The diffusion process is rapid and so the activation barrier for the proton leaving the double layer that is represented on the right-hand side of Eq. (7) should be small so that the reversible potential determined above for this equation will be close to that which would be calculated for the overall process of presenting the proton to the bulk solution phase of the electrolyte. It is noted that in the reaction model of Eq. (7) the hydrogen-bonded water trimer on the left and the partially solvated hydronium ion on the right are part of the double layer. The double layer is a reactant and is included in the model calculations. 2.3. Model for reversible OH(ads) formation in basic electrolyte For OH(ads) to form in basic electrolyte with the same onset potential as reaction (7) according to the reaction OH (aq)Pt(s) 0 PtOH(ads)e (U)
(8)
suggested in Ref. [7] would appear to be a coincidence, and it might be thought that adsorbed water would be the species actually oxidized. Using the standard reversible potential of 2.02 V for aqueous OH reduction to aqueous OH and the Pt/OH bond strength value of 2.50 eV, the reversible potential for Eq. (8) would be / 0.48 V, which is deep in the potential range of hydrogen evolution and far from the observed 0.6 V onset potential. However, the same correction for H2O displacement that was made to the oxidation reaction in acid electrolyte might be assumed to apply in this case too. The rationale for this would be that in the hydrogen upd range H(ads) polarizes the surface, forming Hd and Ptd at the surface. This will result in lone-pair bonding of H2O to these Pt atoms. Thus the appropriate reaction in base would then be
PtOH2 OH (aq) 0 PtOHH2 Oe (U)
(9)
When the water adsorption bond strength of 0.65 eV is included, the predicted reversible potential becomes 0.17 V. At electrode potentials positive of this, then, it would be thermodynamically possible to form OH(ads) in basic solution by discharge if OH (aq). Relative solvation energies of adsorbed H2O and OH could change this reversible potential prediction as could changes in the relative adsorption energies. Mechanisms for this reaction have not yet been explored theoretically. Two candidates come immediately to mind, displacement of adsorbed H2O by OH with concomitant or subsequent electron transfer or H abstraction with simultaneous electron transfer. The former would require displacement of strongly held solvating water molecules from the small OH anion, which might require too much energy, and the latter could be modeled in much the same way was done for oxidation of adsorbed H2O in Eq. (7): PtOH2 OH (HOH)2 l PtOH HOH(HOH)2 e (U)
(10)
where two water molecules are hydrogen bonded to the two lone pairs of OH and the product water molecule. Work on this mechanism has been initiated in the author’s lab. Might OH (aq) actually discharge at /0.17 V to form OH(ads)? Recent studies of CO electrooxidation, where in OH(ads) is believed, from experimental analysis [7,16] and from quantum mechanical model calculations [17] to be the oxidant for CO(ads), show evidence for some OH(ads) being formed at potentials as low as 0.2 V in base [7,16]. The oxidation current density for CO in solution in this ‘prewave’ potential region, which has a weak peak at 0.3 V and weakens at 0.4 V, the edge of the H upd range, is much smaller than the current densities that are associated with 0.6 V OH(ads) formation onset potential discussed earlier. This adsorbed OH is in low surface concentration and is believed to be bonded to surface defects on Pt(111) and to step sites on the (100) and (110) surfaces [7,16]. In acid electrolytes, by comparison, these defect and step sites are believed to be blocked by adsorbed anions. If OH bonds relatively strongly compared with H2O at defect sites, then, by Eq. (10), the reversible potential for its formation will be smaller than when it bonds relatively weakly. The drop in CO oxidation current at 0.4 V can be associated with the absence of upd hydrogen at this potential and the fact that this potential is /0.2 V less than the pzc, with the result that water molecules will not form strong lone-pair donation bonds to the surface and will therefore have difficulty in being activated for generating OH(ads). The higher onset potential for forming OH(ads) on non-defect and nonstep sites in the basic electrolyte would be the conse-
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quence of a smaller difference between OH and H2O adsorption bond strengths at these sites. 2.4. The adsorption bond energy of O2 and its relationship to reduction In the previously discussed density functional calculations of O2 reduction, O2 was calculated to bond to a Pt2 site with a bond strength of 0.43 eV when upright on a 1fold site and 0.94 eV when lying down in the di-s structure [9]. Density functional band calculations have yielded 0.72 eV for the di-s site on the (111) surface [18] and other density functional cluster calculations produce energies in the range 0.53 /0.83 eV for O2 in this configuration [19]. Experimental determinations for the adsorption bond strength of O2 on Pt(111) are 0.4 [20] and 0.5 eV [21], but a competing dissociation reaction forming adsorbed O atoms complicated the analysis. In the electrode potential range from 0.5 to 0.9 V, O2 reduction is a four-electron process in acid electrolyte, with curved Tafel plots [4]. The observed curvature could in part be caused by the potential dependence of the activation energy for the first step, according to theory discussed above [9]. Kinetic models incorporating a dependence on coverage of OH or O site blocking species have accounted for the limiting Tafel slopes in the high overpotential and low overpotential regions, 2RT /F and RT /F , respectively [22]. The relative contributions of these two effects to the observed change in Tafel slopes is an unresolved issue. Despite the formation of OH(ads) from oxidation of H2O(ads) by Eq. (6), and despite the adsorption of anions, sites evidently exist for O2 coordination, reduction to OOH(ads), and dissociation to easily reduced O(ads) and OH(ads). The nature of these sites would be very interesting to elucidate. They will be characterized by either a relatively weak OH adsorption bond strength, so that the reversible potential for OH(ads) formation by Eq. (7) is pushed to /0.9 V or higher (this assumes the formation is under thermodynamic control), or by an increase in the activation energy for OH(ads) formation at these sites (this would bring about kinetic control). It is also possible for a combination of these two effects to be responsible for the absence of OH adsorption on these sites. These sites will also be required to adsorb anions weakly so that O2 can displace them.
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Acknowledgements I thank my coworkers, Dr. T.V. Albu, Dr. R.A. Sidik, Dr. N.M. Neshev, and Dr. P. Shiller for their many calculations and discussions that led up to the work referenced herein, some of which is published and some of which is submitted for publication. This work is supported by the National Science Foundation, Grant No. CHE-9982179, and by the Army Research Office, Grant No. DAAD 19-99-1-0253.
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