Journal of The Electrochemical Society, 162 (1) F115-F122 (2015) F115
Mechanisms for Ethanol Electrooxidation on Pt(111) and Adsorption Bond Strengths Defining an Ideal Catalyst Haleema Aied Asiri and Alfred B. Anderson∗,z Chemistry Department, Case Western Reserve University, Cleveland, Ohio 44106-7078, USA
Ethanol electrooxidation on the Pt(111) electrode has been studied with computational theory. Using a solvation model and a modified Poison-Boltzmann theory for electrolyte polarization, standard reversible potentials for forming 17 reaction intermediates in solution were calculated with density functional theory. Reversible potentials for adsorbed intermediates were then determined by inputting calculated adsorption energies into a linear Gibbs energy relationship. A path to CO2 was found where surface potentials were low and close to the calculated 0.004 V reversible potential for the 12 electron oxidation of ethanol. An exception was the 0.49 V potential for forming the OH(ads) from H2O(l), this being required for oxidation of CO(ads) and RH(ads) intermediates. The surface potentials show that acetyl, OCCH3(ads) forms at small positive potentials and decomposes to CH(ads), CH3(ads), and CO(ads), which poison the surface at these potentials. Energy losses due to non-electron transfer reaction steps are small and cause a small shift in the reversible potential for the 12 electron oxidation. Values for adsorption bond strengths over a perfect catalyst were determined. It is concluded that on an ideal catalyst most intermediates will adsorb more weakly and OH more strongly than on Pt(111). © The Author(s) 2014. Published by ECS. This is an open access article distributed under the terms of the Creative Commons Attribution 4.0 License (CC BY, http://creativecommons.org/licenses/by/4.0/), which permits unrestricted reuse of the work in any medium, provided the original work is properly cited. [DOI: 10.1149/2.0781501jes] All rights reserved.
Manuscript submitted August 11, 2014; revised manuscript received October 10, 2014. Published November 21, 2014.
Unlike a fossil fuel, ethanol is renewable and can be produced from surements identified acetaldehyde, OCHCH3, forming at potentials biomass. Compared to methanol, ethanol is less toxic. The specific greater than about 0.3 V and identified CO2 forming at potentials energy density of ethanol is high (8.0 kWh/kg). It is liquid, which greater than about 0.5 V.2 Fourier transform infrared (FTIR) spec- makes it easy to store and transport. These advantages make the direct troscopy provided evidence for the functional groups in acetic acid, ethanol fuel cell (DEFC) a promising green energy source. However, HOOCCH3, acetaldehyde, and adsorbed methyl (CH3)ormethylene 2 commercialization of DEFC is hindered by the slow inefficient elec- (CH2) fragments. The Lamy group obtained functional group vibra- trooxidation reaction of ethanol on platinum. Platinum is active as tional signatures using subtractively normalized interfacial Fourier the anode electrocatalyst in hydrogen fuel cells but is much less ac- transform infrared spectroscopy (SNIFTIRS) and single potential al- tive for ethanol oxidation. Theoretical understanding of the failures of teration infrared spectroscopy (SPAIRS).3 They concluded that carbon platinum will establish the specific steps during the electrooxidation monoxide (CO) and acetyl (OCCH3) were intermediates, acetalde- which present the challenges. As shown in this paper, some of these hyde was an intermediate and final product, and acetic acid, carbon challenges can be overcome by using electrocatalysts where reaction dioxide, and methane were final products. intermediates have specific adsorption energies to the active site. De- In addition to the above, there have been numerous other investiga- signing materials with these properties is the goal for electrocatalyst tions that have given additional insight into the details and complexi- development. ties of ethanol electrooxidation on platinum5–19 and some of them are The complete electrochemical oxidation reaction which takes place comparative studies of platinum and platinum alloys,6,7,9,11,13,14,17–19 at the anode surface produces twelve electrons, twelve protons, and and a few study temperature dependencies of the electrooxidation.6,7,16 carbon dioxide: Findings in several of these studies will be discussed later in this paper. + − Theoretical analyses and predictions concerning ethanol electroox- HOCH CH (aq) + 3H O(l) → 2CO (g) + 12H (aq) + 12e [1] 2 3 2 2 idation are limited in number and all of them are applications of ex-situ The reversible potential, Uo, is 0.090 V in aqueous solution un- modeling of structures and energies using density functional theory der standard conditions, as calculated from standard thermodynamic (DFT) calculations of properties of the surface-vacuum interface.20–23 data.1 However, cyclic voltammetry studies reported in 2002 by Iw- The surface chemistry of ethanol at the vacuum interface on Pt(111) −1 24 asita indicated that under the conditions 1 mol L HOCH2CH3 has been examined this way, and the findings provide rough guid- −1 −1 + 0.1 mol L HClO4,50mVs and 298 K, ethanol oxi- ance to the electrode surface chemistry at the potential of zero charge dation does not occur on well-defined Pt(111) electrodes at po- (PZC), which is the potential of the uncharged electrode immersed in tentials below about 0.4 V on the reversible hydrogen electrode the electrolyte. However, the PZC can change by hundreds of milli- (RHE) scale.2 Reference 2 is of a review nature and contains volts between reactants and products and generally is different from references to earlier pioneering work of the Iwasita group. Sim- the applied potential.25 This makes it necessary for theoretical studies ilar results were seen in 2004 by the Lamy group in polar- to add the dimension of electrode potential to the calculations in or- ization curves for ethanol oxidation over carbon-supported plat- der to provide definitive predictions. Interactions with solvent water −1 inum electrodes under the conditions 1 mol L HOCH2CH3 molecules and electrolyte ions must be included for the modeling of −1 −1 3 + 0.1 mol L HClO4,5mVs and 293 K. Very recent experiments the electrochemical interface to be accurate and complete. by Cantane and Gonzalez for electrodes containing black platinum The purpose of this paper is to give results of our examination supported on gold show an onset potential for ethanol oxidation be- of the steps during the electrochemical oxidation of ethanol over −1 tween 0.4 V and 0.5 V under the conditions 1 mol L HOCH2CH3 + the Pt(111) electrode using the theory outlined below. We will first −1 −1 4 0.1 mol L HClO4, 100 mV s and 298 K. Thus, the overpotential obtain results for each oxidation reaction occurring with simultaneous for oxidation is about 300 mV- 400 mV under typically employed generation of H+(aq) + e− oxidation products. The results of our study conditions of study. of electrooxidation of ethanol also apply to CO2 reduction to ethanol, However, the oxidation is not complete, and several final oxida- acetic acid, acetaldehyde, and adsorbed intermediates. tion products have been observed over platinum electrodes. Iwasita’s Self-consistent field density functional theory, SCF-DFT, is used on-line differential electrochemical mass spectroscopy (DEMS) mea- to calculate standard reversible potentials, U0, for electron transfer steps between the electrode and reaction intermediates in solution. ∗Electrochemical Society Active Member. These potentials are used in the theory for predicting reversible po- zE-mail: [email protected] tentials for reactions on the catalyst surface and they are calculated
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because experimental values are not available for most of the elec- face, adsE, with the zero-point vibrational energies not included in E, tron transfer steps. In this theory, called the linear Gibbs energy are substituted for the adsG, and reversible potential predictions pre- relationship, LGER, the electrode surface is viewed as a perturbation dicted this way also have errors usually less than 0.2 V.This is because to an electron transfer process in solution. The perturbation consists changes in the solvation and zero-point vibrational energies upon ad- of the difference between the adsorption bond strengths between the sorption have been found to nearly cancel when calculated with ac- electrode surface and the reactants and products of the electron trans- curate theory.28 The conceptual model for this, discussed in refs. 25 fer reaction. Although the adsorption energies are calculated at the and 28, shows how on adsorption approximately equal amounts of PZC, the LGER approach has yielded predicted reversible potentials weakly-held solvation shells of the reactant and product are replaced for electron transfer reactions on Pt(111) surfaces in error by less than by bonds to the surface. rev 0.2 V in most of the cases where experimental data can be found for The equation used for Usurf predictions in this paper is comparison.25 rev ≈ 0 + − / Knowing the redox potentials for adsorbed intermediates can Usurf U [ ads E(Ox) ads E(R)] nF [3] greatly assist understanding the electrocatalysis. Modifying these po- To find the U0 values employed in eq. 3, we first calculated Gibbs tentials to approach the reversible potential for the overall process, energies of Ox and R in bulk solution and determined the energy for example 0.090 V for the 12 electron oxidation of ethanol, by find- of the electron, that is, the electrode potential U0 that satisfied the ing new catalytic materials and possibly new electrolytes is a goal equilibrium condition: for development of higher efficiency ethanol anodes for use in fuel + − 0 → cells. The electron transfer kinetics are expected to be generally fast Ox(aq) e (U ) ← R(aq) [4] at the reversible potentials. This is because past experience in this lab- The Gibbs energy balance in eq. 4 at equilibrium yields the oratory, found in reference 25 and references therein, indicates that equation activation energies for coupled electron plus proton transfer reactions 0 =−{ 0 − 0 }/ − ϕ/ at the reversible potentials on platinum surfaces are ∼ 0.1 eV. This U G sol(R) G sol(Ox) nF F [5] low value is compatible with fast reactions at the anode. At potentials In eq. 5, F is the Faraday constant, ϕ is the thermodynamic work higher than the reversible potential, the activation energies will be 29 25 function of the standard hydrogen electrode, calculated to be 4.43 eV, even less and at lower potentials they will increase. For reduction 0 and G sol is the standard Gibbs energy of the molecule in standard pH the inverse holds and activation energies increase as the electrode = 0 aqueous solution. The density functional code Interface was used potentials increase. 0 0 for determining G sol(R) and G sol(Ox) for electron transfer reactions The overpotential for current flow is around 300 mV to 400 mV, in solution and also when the reacting species are adsorbed on the but at these potentials the oxidation current density is much less than electrocatalyst surface.29,30 The code uses atomic orbitals and atomic would be obtained for the 12 electron oxidation. This is because at 0 pseudopotentials. In Interface, G sol is calculated as a sum of energy 300 mV overpotential the oxidation is only partial, generating two components by using the equation electrons and forming acetaldehyde. At higher overpotentials addi- 0 = + + − − + − tional products form, some of which desorb and are not oxidized Gsol E ss,nonel is,nonel TSe TSi Hvib TSvib [6] further, and little full oxidation to CO2 takes place. Therefore, the Here E is the molecule’s internal energy, ss,nonel and is,nonel are the electrical power obtained is a small fraction of the thermodynamic respective Gibbs energies from non-electrostatic solute-solvent and limit for complete oxidation. A goal for catalysis improvement is to ion-solute interactions with the molecule, determined using dielectric find catalytic materials that hold all intermediates on the electrode continuum for the solvent and a modified Poisson-Boltzmann equa- surface until CO2 has formed and is able to desorb rapidly. tion for the ion distribution, T is the temperature of the system, S There can be thermodynamically induced overpotentials caused by e and Si are entropies of the electrons and ions, Hvib - TSvib are the exergonic reactions on the electrode surface. These reactions waste vibrational contributions to the enthalpy and entropy of the molecule. Gibbs energy as heat and entropy, as happens at platinum and platinum The components of eq. 6 were calculated using the RPBE functional. alloy oxygen cathodes, resulting in effective reversible potentials that 30 25,26 Details of the computations are in ref. 29 and its follow-up paper. are less than the standard one for reduction reactions, and greater The same procedure was also used in refs. 25–28 and others. To reduce than the standard one for oxidation reactions as shown later in this computational effort, the Gaussian09 code31 was used to calculate the paper. An ideal catalyst will adsorb the intermediates with appropriate Hvib - TSvib components. adsorption bond strengths so as to make these reaction energies zero It might be wondered whether internal energies, E(Ox) and E(R), and shift the reversible potential for each electron transfer step to 0 could yield useful predictions when substituted for G sol(Ox) and the overall reversible potential. As shown later, an ideal catalyst for G0 (R) in eq. 5. The answer is yes in some cases, at least. As shown ethanol electrooxidation can do both. sol in ref. 27 for five different reduction reactions involving O2,H2O2, 0 OH, and H2O, there is a linear correlation between U calculated this Theoretical Details way and the measured values. The existence of the linear relationship is the result of the small differences of the small hydration energies The LGER theory got its start in 1999 for prediction bulk solu- 27 of the neutral reactant and product molecules. The calculated values tion reversible potentials and was subsequently refined up to the are on average 0.43 V less than the experimental values. Other series present date for predicting reversible potentials for reactions of ad- 26 involving groups of molecules possessing similar solvation energy sorbed species. This theory uses the standard reversible potentials differences might be expected to show trends too with an offset dif- for reactions in bulk solution and perturbs them with the difference in fering from 0.43 V, but there has not yet been exploration of the idea. adsorption Gibbs energies, adsG, on the electrode surface of prod- 0 0 To be safe, it is better to calculate and use G sol(R) and G sol(Ox), as ucts of reduction, the reduced species R, and the reactants that are is done for this paper. being oxidized, Ox. The reversible potential for the reaction between 0 rev With the predicted values for the U in hand, we proceeded to adsorbed species, Usurf ,isthen calculate adsE(Ox) and adsE(R) for use in eq. 3 to determine values rev = 0 + − / for U rev . Calculations of the adsorption internal energies were per- Usurf U [ adsG(Ox) adsG(R)] nF [2] surf formed with the two-dimensional periodic boundary condition option The adsG’s are dependent on the electrode potential, and using val- of the Interface code. Adsorption internal energies, with vibrational ues calculated at the PZC instead of at the reversible potential gives zero point energies omitted, were calculated as follows: reversible potential predictions for surface reactions that have errors usually less than 0.2 V for the aqueous interface when electrolyte and ads E(Ox) = E(Ox(ads)) − E(Surf) − E(Ox) [7] solvation models are used in good quality DFT.25 In the LGER theory, calculated internal energies of adsorption at the vacuum-surface inter- ads E(R) = E(R(ads)) − E(Surf) − E(R) [8]
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√ Figure 1. (a) top view of 1/12 ML HOCH2CH3(ads) showing 2 3×3trans- lational cell; and (b) side view.
In cases of intermediates with two carbon atoms, 1/12 ML cov- erage was used to minimize interactions among√ the adsorbates. A three-layer-thick slab was created using a 2 3×3 unit cell with 2- dimensional periodic boundary condition as shown in Figure 1.The unit cell was made of 36 atoms with 12 in the bottom layer frozen in bulk platinum sites and the positions of the remaining 24 were opti- mized for lowest energy. For intermediates with a single carbon atom, 1/6 ML coverage was used and a 3×2 unit cell was employed. The previously calculated lattice constant value√ of 4.03 Å was used.30 A 3×4×1 grid of k-points was used for the 2 3×3 unit cell and 3×6×1 grid was used for 3×2 unit cell to give a uniform Monkhorst-Pack sampling32 in the Brillouin-zone.
Results Adsorption internal energies.— Like methanol and water, ethanol adsorbs by electron lone-pair donation from oxygen to a Pt atom, bonding on the top site with the C-C bond nearly parallel to the surface. The calculated low-coverage, 1/12 ML, adsorption bond strength, shown to be 0.25 eV in Table I, is close to those of low-coverage, 1/6 ML, methanol and water, 0.28 and 0.23 eV, respectively, which we calculated in a study of methanol electrooxidation over Pt(111).33 Figure 2. Slightly off normal views of the optimized structures for ethanol and its oxidation intermediates on Pt(111) surface. (a) ethanol;(b) ethoxy; (c) Numbers from Table I are rounded to two figures after the decimal acetaldehyde; (d) 1-hydroxylethyl; (e) 1-hydroxylethylidene; (f) acetyl; (g) when used in the text. There are three possibilities for the dehydro- ketene; (h) ketenyl; (i) acetic acid; (j) acetate radical and anion; (k) methyl, (l) genation: loss of H from (i) the α-C of the ethyl group, (ii) the hydroxyl methane, (m) hydroxymethylene, (n) formyl, and (o) hydroxymethylidyne. group, and (iii) the methyl group; the last was not calculated. H loss from the α-C in ethanol, (i), gives 1-hydroxyethyl (HO*CHCH3), which bonds to a Pt atom through the α-C with acetyl. Loss of H from the β-C in acetyl gives ketene (O*C*CH2) with bond strength 2.03 eV. Note that each time the * notation occurs 1.80 eV adsorption bond strength and a short, 1.20 Å, C-O bond length. in the formulas, it means the atom following this symbol is bonded Loss of a second H from the β-C gives ketenyl (O*C*CH), which is to the surface in the variationally optimized structure. The OH and strongly adsorbed by 3.55 eV, and also has short 1.20 Å C-O bond CH3 groups bend away from the surface and the CO bond length is α length. 1.40 Å. By removing the second H from the -C, 1-hydroxylethylidene Table I includes 13 additional adsorption energies and C-C and (HO*CCH3) is formed and it remains on the top site. The adsorption C-O bond lengths. Some of the data are taken from our earlier study strength is increased to 3.23 eV and the CO bond length shortens to 33 of methanol electrooxidation and CO2 reduction. From these data 1.31 Å. Loss of H from the OH group now generates acetyl (O*CCH3), is seen that products of C-C bond breaking are strongly adsorbed, α which bonds through the -C, standing upright on a Pt atom with both bonding through the carbon atoms with strengths in the order, CO OandCH3 group pointing away from the surface. The adsorption < CH3 < CH. The last is strongly adsorbed with 6.13 eV bond bond strength is 2.36 eV and now the C-O bond length is shortened strength. Methane and CO2 are weakly adsorbed at 0.02 and 0.05 eV, to 1.20 Å. Structures for these and other intermediates considered in respectively. this study are shown in Figure 2. It is noted that the C-C distances for Other intermediates include strongly adsorbed hydroxylmethy- the adsorbed intermediates are relatively constant, lying in the range lene (HO*CH), possibly formed by combining OH(ads) and CH(ads). 1.54 Å for ethoxy (*OCH2CH3) to 1.48 Å for ketenyl (O*C*CH). Loss of one hydrogen atom from it gives strongly adsorbed hydrox- If the first H loss in ethanol is instead from the hydroxyl group, (ii), ymethylidyne (HO*C) or formyl (O*CH), and adsorbed CO is formed ethoxy is formed and it is bonded through oxygen, by 1.43 eV, with when the last H is removed. The strongly adsorbed hydrocarboxyl a shortened 2.01 Å Pt-O distance, compared to 2.41 Å for ethanol. radical (HOCO), and non-adsorbed acetic acid (HOOCCH3), are sug- Loss of H from the α-C of ethoxy gives acetaldehyde (*O*CHCH3), gested later to be produced by reaction between OH(ads) and re- which has a weak adsorption energy of 0.35 eV. The most stable spective adsorbed intermediates CO and acetyl. Hydrocarboxyl bonds site for acetaldehyde at 1/12 ML coverage is bridging, with oxygen through the carbon with the hydrogen pointing toward the surface. and carbon bonded to adjacent Pt atoms. The O-C bond is nearly The adsorption bond strength of acetic acid is negative, meaning parallel to the surface, and the methyl group is tilted away from it. If it that though a local energy minimum was found with the variational forms, stays on the surface, and loses the remaining α-H, it becomes structure optimization, the structure is metastable and acetic acid is
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Table I. Most stable adsorption sites, energies of reaction, and bond distances of C-C and C-O bonds of intermediates of ethanol oxidation over the Pt (111) surface. Atoms bonded to the surface are preceded with *. Structures are shown in Figure 1 except for *OH and *OH2.H*OCH, O*CH, and HO*C which are from ref. 33.
Species R(CC) R(CO) Dads Dads(ideal) Dads(ideal)-Dads
ethanol (H*OCH2CH3) 1.530 1.471 0.253 0.000 −0.253 1-hydroxylethyl radical (HO*CHCH3) 1.530 1.397 2.030 1.662 −0.368 1-hydroxylethylidene (HO*CCH3) 1.504 1.310 3.232 2.597 −0.635 ethoxy radical (*OCH2CH3) 1.542 1.426 1.431 2.155 0.724 acetaldehyde (*O*CHCH3) 1.525 1.347 0.351 0.570 0.219 acetyl radical (O*CCH3) 1.537 1.199 2.358 1.776 −0.582 ketene (O*C*CH2) 1.523 1.195 1.797 1.167 −0.630 ketenyl radical (O*C*CH) 1.477 1.196 3.551 3.056 −0.495 hydroxyl radical (*OH) - - 1.951 2.449 0.498 water (*OH2) 0.229 0.000 0.229 acetic acid (HOOCCH3) 1.543 1.209, 1.378 −0.173 − acetate anion (OOCCH3 ) 1.553 1.274, 1.274 0.304 acetate radical (*O*OCCH3) 1.534 1.279, 1.280 1.972 carbon monoxide (*CO) - 1.176 1.640 0.732 −0.908 methyl radical (*CH3) - - 2.051 methane (CH4) - - 0.023 0.000 −0.023 carbon dioxide (CO2) - 1.175, 1.175 0.048 0.000 −0.048 hydrocarboxyl radical (HO*CO) - 1.204, 1.351 2.320 2.097 −0.223 methylidyne radical (*CH) - - 6.126 5.372 −0.977 hydroxymethylene (HO*CH) - 1.385 3.105 3.255 0.150 formyl radical (O*CH) - 1.198 2.392 2.185 −0.207 hydroxymethylidyne radical (HO*C) - 1.342 4.442 3.691 −0.751
a adsG is replaced by adsE in LGER theory.
predicted to not chemisorb. On the other hand, the acetate radical, dard reversible potentials of the solution phase molecules. The re- (OOCCH3), is strongly bonded by 1.97 eV to the surface through versible potentials are best discussed in view of the reaction network in both oxygen atoms, and it stands upright over two adjacent Pt atoms Figure 3. Here it is seen that there are two initial reaction branches, the with the methyl group pointing away from the surface. In comparison, left-hand one for OH bond oxidation, with a positive reversible poten- − the acetate anion (OOCCH3 ), is weakly adsorbed by 0.30 eV. tial of 0.98 V and the right-hand one where H bonded to the α-C is ox- To summarize: all intermediates adsorb strongly to the surface idized, with a reversible potential of −0.35 V. The latter is more likely except acetaldehyde, acetic acid, acetate anion, and methane, all of to occur near the overall reversible potential of 0.090 V, being nega- which are observed partial oxidation products. The reactants ethanol tive of it. Both branches lead to OCCH3(ads), for which three possible and water and the product carbon dioxide also adsorb weakly. reactions are shown. There is also a crossover from HOCHCH3(ads) to acetaldehyde, OCHCH3(ads), in the left-hand branch which has predicted reversible potential 0.58 V. This is higher than, but close Predicted standard reversible potentials for reactions in aqueous to, the ∼0.35 V potential where acetaldehyde is first observed as the solution.— We calculated Gibbs energies for 18 electron transfer reac- potential of a platinum is increased.2,3 Staying with the right-hand tions in bulk aqueous solution phase and obtained standard reversible branch, HOCHCH (ads) is oxidized to HOCCH (ads) with a negative potentials by using them in eq. 5. The energies for oxidation reactions, 3 3 0 0 0 reversible potential and this is oxidized to OCCH3(ads) with a slightly G sol = G sol(R) - G sol(Ox) with 4.43 eV subtracted from them, are positive reversible potential. Acetaldehyde formation by reduction of given under the heading G0 /eV in Table II. The predicted U0 sol OCCH (ads) is prevented in the potential range of interest, near 0.09 values are in the second columns of Table II and their signs are nec- 3 V, because the reversible potential is −0.80 V. essarily the same as those calculated for G0 . The reaction Gibbs sol The OCCH (ads) has three pathways open as shown in Figure 3. energies would have the opposite sign had the reactions been written 3 The left-hand one: combines it with OH(ads) if the electrode potential for reductions. There is a wide range of calculated standard reversible is >∼0.5 V to form acetic acid in a 0.32 eV exothermic reaction; the potentials for the redox reactions considered, ranging from −3.00 V central one has it oxidized to OCCH with a slightly negative reversible to 2.44 V. 2 potential; and the right-hand has it undergoing exothermic C-C bond scission with a loss of 1.10 eV to form CO(ads) and CH3(ads). The Predicted reversible potentials for oxidations of adsorbed CO(ads) is oxidized to CO2 at potentials where OH(ads) is present molecules.— The LGER equation of eq. 3 wasusedtopredictre- and CH3(ads) is reduced to methane with a slightly negative reversible versible potentials for oxidation reactions of adsorbed molecules potential. We did not explore methyl oxidation. from the calculated standard reversible potentials and the calculated The left-hand branch, combining OH(ads) with the OCCH3(ads), internal energies of adsorption for the reactants and products. Be- forms acetic acid, which, because of the negative adsorption bond cause the Pt(111) surface is hydrophobic30 and the interaction with strength in Table I, does not bond to the surface. The reaction energy ethanol should be similar and the CO2 adsorption energy is <0.05 is calculated for this reaction is 0.32 eV, which is small enough for eV, H2O, HOCH2CH3 and CO2 adsorption energies were set to zero. reaction to take place. The solvated acetic acid molecule is calculated The calculated reaction Gibbs energies for the adsorbed molecules to be oxidized with reversible potential −0.08 V to adsorbed acetate are in the third column of Table II and the predicted reversible po- radical. It is also calculated to become ionized to acetate anion and tentials are in the fourth column. In the following discussion, num- H+(aq) with a 0.32 eV energy increase. bers taken from Table II will be rounded to two figures after the Continuing down the central path, OCCH2(ads) is oxidized with decimal. a small positive reversible potential to OCCH(ads), and OCCH(ads) All the predicted reversible potentials for the adsorbed molecules dissociates with 1.17 eV energy loss to CH(ads) + CO(ads). If the are closer to the overall ethanol oxidation potential than the stan- potential is high enough for it to form, the CO(ads) combines with
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