The Importance of the Electrochemical Environment in the Electro

Research Article pubs.acs.org/acscatalysis

The Importance of the Electrochemical Environment in the ElectroOxidation of Methanol on Pt(111) Sung Sakong*,† and Axel Groß*,†,‡ †

Institute of Theoretical Chemistry, Ulm University, 89069 Ulm, Germany Electrochemical Energy Storage, Helmholtz Institute Ulm (HIU), 89069 Ulm, Germany

‡

ABSTRACT: The electro-oxidation of methanol on Pt(111) is studied based on periodic density functional theory calculations. The aqueous electrolyte is taken into account using an implicit solvent model, and the dependence of the reaction energetics on the electrode potential is derived using the concept of the computational hydrogen electrode. The total oxidation of methanol becomes thermodynamically preferred at electrode potentials above U = 0.6 V relative to the standard hydrogen electrode. We propose a most favorable reaction path involving surface carboxyl as the last reaction intermediate before CO2 formation, which can either be formed in a indirect mechanism from adsorbed CO or in a direct mechanism from formic acid. The presence of the aqueous electrolyte significantly stabilizes reaction intermediates that contain hydrophilic groups. This also leads to a selectivity for the initial C− H bond breaking process with respect to the initial O−H bond breaking of methanol that is increased by 3 orders of magnitude at room temperature when solvent effects are considered. KEYWORDS: electrocatalysis, direct methanol fuel cell, electro-oxidation of methanol, atomistic modeling, first-principles method

1. INTRODUCTION Recently, there has been an increased interest in the theoretical description of processes at electrochemical electrode−electrolyte interfaces.1−6 A reliable modeling of elementary electrochemical processes at electrode−electrolyte interfaces could be a key process in the better understanding of the basics of energy storage and energy conversion mechanism in batteries and fuel cells.6−10 Thus, systematic computational modeling could contribute to improving the design of batteries and fuel cells systematically. In practice, however, there are still severe limitations due to the complexity of electrochemical electrode−electrolyte interfaces. Typically, systems addressed in density functional theory (DFT) calculations consist of up to several hundreds of atoms because of limited computational resources. This is often not sufficient for an appropriate description of the electrode− electrolyte interface. A computationally efficient and successful method to take the environment in the modeling of catalytic reactions into account is to describe this environment as a reservoir in a grand canonical approach.11−13 Usually the environment is not explicitly considered, the thermodynamic conditions just enter through the dependence of the chemical potentials of the reacting species on temperature, pressure, concentration, and so forth. With regard to the modeling of electrochemical systems, there are two main obstacles: First, within electronic structure theory, there is hardly any practical realization of the grand canonical ensemble of electrons in terms of the electrode © XXXX American Chemical Society

potential, as modern DFT codes typically rely on a variational method for the energy minimization within a microcanonical ensemble of the electrons.2,9,14 Second, the liquid nature of the electrolyte requires a numerically demanding statistical sampling.15 Recently, we have suggested to combine the concept of the computational hydrogen electrode13 with an implicit solvent method16,17 to make electrochemical interfaces manageable within a conventional DFT framework. Using this approach, we were able to correctly predict the dependence of the hydrogen and hydroxyl coverage on a Pt(111) electrode as a function of the electrode potential.17 Here, we will extend this approach to address a technologically important electrocatalytic reaction, namely, the electro-oxidation of methanol. This corresponds to the crucial process occurring on the anode of the direct methanol fuel cell (DMFC). The DMFC is considered to be an efficient device for the electrochemical energy conversion from simple alcohols. Therefore, the electro-oxidation of methanol has been intensively studied both experimentally18−30 as well as theoretically.3,30−43 In heterogeneous catalysis, methanol oxidation and synthesis are also important technological processes that are typically performed on Cu/ZnO/Al2O3 industrial catalysts.44 The overall reaction in the total oxidation of methanol is given by Received: March 31, 2016 Revised: July 5, 2016

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DOI: 10.1021/acscatal.6b00931 ACS Catal. 2016, 6, 5575−5586

Research Article

ACS Catalysis 2CH3OH(g) + 3O2 (g) → 2CO2 (g) + 4H 2O(g)

reaction intermediates as a function of the electrode potential and thus propose the most favorable reaction scheme. In addition, we derive the reaction barriers that are the crucial input for any kinetic modeling. We particularly focus on the role of the solvent in the methanol electro-oxidation. We will show that there is a significant influence of the solvent on the energetics when some of the intermediates contain hydrophilic groups. This can influence the selectivity of certain reaction steps by 3 orders of magnitude at room temperature. Our results show that the combination of the concept of the computational hydrogen electrode with an implicit solvent method represents a reliable tool to analyze electrocatalytic processes from first-principles.

(1)

In contrast, the electro-oxidation of methanol can be expressed as the electrochemical equilibrium between methanol and carbon-dioxide CH3OH(aq) + H 2O(aq) ⇋ CO2 (aq) + 6H+ + 6e

(2)

which occurs at the anode of a DMFC at a half cell potential of 0.02 V.24 Still, in both heterogeneous and electro-catalysis, the methanol oxidation requires hydrogen-poor and oxygen-rich conditions at the catalyst surface. In heterogeneous catalysis, methanol oxidation is thus controlled by the partial pressures of hydrogen and oxygen,45−47 whereas in the elecro-oxidation, the electrode potential is the crucial control parameter. In electrocatalysis, oxygen is typically coming from the water molecules in the electrolyte. However, this requires the dissociation of water, and Pt electrodes are rather inactive as far as water dissociation is concerned.17 Therefore, an efficient catalyst for methanol electro-oxidation needs active sites for water dissociation. These sites might be provided by Ru atoms in a PtRu bimetallic alloy, as for example suggested by Shubina and Koper.48 Still, it is not only important to supply oxygen, hydrogen also needs to be removed from the catalyst because it, for example, passivates electrodes.49 In heterogeneous catalysis, according to reaction scheme 1, the formation of water molecules reduces the concentration of hydrogens at the catalyst surface.46 In electro-catalysis, protons need to be solvated, as indicated in reaction scheme 2. This is in principle energetically favorable as long as the protons are not too strongly chemisorbed on the electrode.50 On Pt(111), hydrogen is strongly bound14,42 such that Pt(111) is typically hydrogen-covered at low electrode potentials, and it needs an electrode potential larger than 0.5 V relative to the standard hydrogen electrode (SHE) to remove hydrogen from the Pt electrode.17,51 This demonstrates that in the theoretical modeling it is important to consider the appropriate electrochemical conditions under which a desired electrocatalytic process occurs. Taking the electrode potential into account via the concept of the computational hydrogen electrode13 and modeling the presence of the electrolyte in an implicit solvent approach,16,17 we have addressed possible reaction schemes for methanol electro-oxidation on Pt(111) using periodic DFT calculations. In most of the previous first-principles studies addressing the oxidation of methanol on Pt(111), the influence of the electrochemical environment has been neglected. There have been DFT studies37,38 based on the double-reference method,52 but in these studies, the role of the solvent was not explicitly discussed and only a limited set of possible reaction intermediates was considered. The initial dehydrogenation step in the methanol oxidation43 and the formic acid oxidation40 was addressed using a periodic continuum solvation method. In addition, the activation barriers involved in these selected reactions were derived in a constant-charge framework for different charge states of the electrode. Still, to the best of our knowledge the electrochemical environment has not yet been taken into account in any theoretical study addressing the complete methanol electro-oxidation scheme on Pt(111) including the energetics of reaction intermediates and activation barriers. Here, we first determine the equilibrium coverage of the Pt electrode in the presence of an aqueous electrolyte as a function of the electrode potential. We evaluate the energies of possible

2. COMPUTATIONAL DETAILS Periodic DFT calculations have been performed using the software package VASP.53 The wave functions have been expanded up to 700 eV using a plane wave basis set. The electronic cores are described by the projector augmented wave method,54 and the exchange-correlation energies are evaluated within the generalized gradient approximation (GGA) as suggested by Hammer and Nørskov, known as a revised version of the Perdew−Burke−Ernzerhof (RPBE) functional.55 The lattice parameters of fcc Pt of a = 3.99 Å and hcp Ru of a = 2.73 and c = 4.30 Å are optimized using a fine k-point grid of 21 × 21 × 21. The Pt electrode is modeled by a Pt(111) slab consisting of five atomic layers. Calculations for Ru have been performed using a five-layer slab of Ru(0001). The top three layers of the slabs are fully relaxed, whereas the lower two layers are fixed at their bulk positions. The slabs are separated by a vacuum of 15 Å to avoid any interaction between the periodic images. The energetics of the reaction intermediates and barriers have been determined within a 3 × 3 surface super cell employing a 5 × 5 × 1 k-point grid to integrate over the first Brillouin zone. The consideration of dispersion effects is crucial for a reliable first-principles description of the interaction of organic molecules56 and water57,58 with metal surfaces. Here, we use the semiempirical D3 dispersion correction scheme of Grimme with zero damping function,59,60 which together with the RPBE functional reliably predicts properties of water−metal interfaces.57,58,61 The cutoff radius for the pair interactions has been chosen to be 10 Å. However, the screening of the van der Waals (vdW) interactions in bulk metals is not correctly described in dispersion correction schemes.57,62 Therefore, we exclude the dispersion correction for all metal atoms below the first layer. The electrolyte is implicitly modeled by a polarizable dielectric continuum as implemented into the VASP code by Mathew and Hennig.16 Thus, the computationally demanding statistical sampling of the liquid water degrees of freedom can be avoided.17 Hence, the averaged contribution of the electrolyte is effectively included in the DFT calculations yielding solvation energies and the local potential as described in the joint density functional theory framework.63−66 The properties of water are described with a dielectric constant of ϵb = 80 and a cutoff charge density of ρcut = 0.0025 Å−3. The cavitation energies are calculated using a surface tension parameter of 0.525 meV/Å267 as described in refs 68 and 69. We note that the current implementation of the implicit solvent method in the VASP code may lead to an erratic surface dipole in a static calculation; thus, the local potential should be selfconsistently corrected by a compensating dipole field to remove spurious interactions of a surface dipole with its periodic 5576


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highest occupied and the lowest unoccupied molecular states in CO, the so-called HOMO−LUMO gap, leads to a strong backdonation into the CO π* orbital, and thus, overly stabilized CO prefers high-coordination sites.79,80 It is known that this problem can be corrected by using a hybrid functional,81 which, however, sacrifices the correct description of bulk metal properties.82 As Lazić and Blügel have demonstrated, the problem is also related to the appropriate consideration of correlation energy.83 Despite the overestimated CO adsorption energy, we note that the RPBE-D3/zero functional still represents the optimal compromise in the choice of the exchange-correlation functionals as there is no other functional that reproduces the metal and water properties and the metal− molecule interaction of most of the considered reaction intermediates equally well at the same computational cost. Next, to analyze the energetics of reaction intermediates (carbon species), we evaluate their formation enthalpy with respect to methanol in the aqueous phase. As worked out by Reuter and Scheffler, the reference energies of reacting species in catalysis in their appropriate reservoirs can be expressed through their chemical potentials, which are functions of the reaction conditions.11,12 Nørskov extended this approach to electrochemical systems13,84 within the concept of the computational hydrogen electrode. Effectively, the chemical potentials will be treated as free parameters, but the dependence of the reaction energetics on the thermodynamic conditions enter through the corresponding dependence of the chemical potentials. For example, the hydrogen chemical potential for a given electrode potential U can be expressed as

images. The energies and forces of stable configurations satisfy the convergence criteria of 10−6 eV and 0.01 eV/Å, respectively. We assume that the processes in methanol oxidation are electronically adiabatic both in vacuum and in water. Therefore, the transition state search of each reaction step has been performed on the multidimensional Born−Oppenheimer ground-state potential energy surface. The minimum energy paths of all reaction steps are determined using the nudged elastic band (NEB) method within variational transition state theory.70 We used eight images for the calculations. The exact energies of the transition states are determined by an interpolation scheme or by using the climbing NEB method.71 In the identification and validation of transitions states by the climbing NEB method, the forces on each atom are required to be smaller than 0.01 eV/Å.

3. ENERGETICS OF REACTION INTERMEDIATES As a first step, we validate our computational setup by comparing the calculated energetics of selected reaction intermediates with values measured by microcalorimetry or temperature-programmed desorption experiments under ultra high vacuum conditions. As shown in Table 1, the calculated Table 1. Adsorption Energies of Selected Molecules on Clean Pt(111) in Vacuuma adsorbate

RPBE-D3/zero

experiment

CH3OH CH3O HCOOH COfcc COtop H2O (H-up) H2O (H-down)

0.57 1.87 0.59 1.83 1.85 0.57 0.61

0.66b 1.93b 0.74c

μH ̃ (U ) = μ H ̃ +(aq) − e(U + USHE) =

1.37d

1 E H (g) − eU 2 2

(3)

where one has used the fact that, at standard conditions defining the standard hydrogen electrode potential USHE, the solvated proton is in equilibrium with the H2 molecule in the gas phase. This approach avoids the computationally demanding evaluation of solvation energies of ionic species.85 The entropic contribution of varying proton concentration, kBT ln(10)pH, can be added to eq 3, which at room temperature corresponds to a shift of approximately 59 meV in the chemical potential if the pH is changed by one. The chemical potentials of carbon and oxygen are specified by referring them to solvated methanol and water according to μO = EH2O(aq)−2 μ̃ H(U) and μC = ECH3OH(aq)−μO−4 μ̃H(U), respectively. These chemical potentials can be easily derived from the appropriate DFT total energy calculations of the neutral molecules either in the gas phase or in their solvated form employing the implicit solvent model. Thus, the formation enthalpy of a reaction intermediate is defined by

0.7e

a

The energy gain of an adsorbate is calculated in comparison to the molecule in gas phase. The energies are given in eV. bFrom ref 72. c From ref 73. dFrom ref 74. eFrom ref 75.

adsorption energies are in general in good agreement when compared with experimental measurements of selected molecules on clean Pt(111). We note that the calculated adsorption energies are in general in good agreement with experimental values72−75 with discrepancies of less than 0.1 eV, except for the case of CO adsorption. Note that the deviation between theory and experiment is reduced with respect to previous DFT studies in which dispersion corrections were not considered.72,73,75 We stress that the improvement in the evaluation of adsorption energies by using the RPBE-D3/zero functional is not restricted to weakly bound physisorbed species. Also, the adsorption energy of the methoxy radical is in better agreement with the experiment compared to calculations without dispersion interactions.72 More interestingly, we note that the RPBE-D3/zero functional furthermore leads to the correct adsorption site assignment of carbon monoxide on the Pt surface, thus not following the so-called CO/Pt(111) puzzle.76,77 According to spectroscopy studies of CO vibrations on metal surfaces, CO adsorbs at the top site at low CO coverages,74,78 whereas most GGA functionals in first-principles calculations erroneously predict fcc hollow as the most favorable adsorption site.76,77 Still, the CO adsorption energy is strongly overestimated by the RPBE-D3/zero functional. It has been suggested that the underestimated gap between the

ΔH(U ) = Etot − Eslab −

∑ niμi i

(4)

where Etot and Eslab represent the total energies with and without the adsorbed reaction intermediate, respectively, and ni is the number of species i characterized by the chemical potential μi. Note that the formation enthalpies of the carboncontaining species are given with respect to solvated methanol, liquid water, and protons in solution. For a given concentration of the species in the solution, the formation energies can then be expressed as a function of the electrode potential U. Entropic contributions are neglected as the change in the vibrational entropy typically only yields a small contribution, 5577


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layer region is rather favorable as protons created during electro-oxidation processes will be immediately removed from the surface. Interestingly enough, this double layer region corresponds to the range of electrode potentials in which the methanol electro-oxidation is most effective.18 We are now addressing the influence of the presence of the solvent on the formation energies of the reaction intermediates on Pt(111). Note that in previous studies considering the methanol oxidation on Pt(111), typically the influence of the solvent on the adsorption energies has been neglected.33,34 In Table 2, the formation enthalpies of reaction intermediates on a

and the configurational entropy of the liquid is taken into account in the implicit solvent model.17 First, we determine the equilibrium coverage of the electrode surfaces in the presence of an aqueous electrolyte. In Figure 1,

Table 2. Formation Enthalpies ΔH of Reaction Intermediates in the Electro-Oxidation of Methanol on Pt(111) in Vacuum and Solvated in Implicit Water According to ref 4, i.e., as a Reference in All Cases Solvated Methanol, Liquid Water, and Solvated Protons were Useda ΔH(U) (eV) on Pt(111) in water

on Pt(111) in vacuum

adsorbate

U = 0.5 V

0.75 V

U = 0.5 V

0.75 V

Esolv (eV)

CH3OH CH3O H2COH CH2O HCOH HCO COH COfcc COtop H2COOH HCOOH H2COO HCOO COOH C(OH)2 CO2

−0.55 0.02 −0.91 −0.44 −1.05 −1.28 −1.68 −2.07 −2.07 −0.47 −1.34 0.20 −1.49 −1.97 −1.86 −1.88

−0.55 −0.23 −1.16 −0.94 −1.55 −2.03 −2.43 −3.07 −3.07 −1.22 −2.34 −0.80 −2.74 −3.22 −2.86 −3.38

−0.38 0.07 −0.76 −0.38 −0.96 −1.18 −1.44 −2.03 −2.05 −0.22 −1.20 0.27 −1.39 −1.81 −1.54 −1.80

−0.38 −0.18 −1.01 −0.88 −1.46 −1.93 −2.19 −3.03 −3.05 −0.97 −2.20 −0.73 −2.64 −3.06 −2.54 −3.30

−0.17 −0.05 −0.15 −0.06 −0.09 −0.10 −0.24 −0.04 −0.02 −0.25 −0.14 −0.07 −0.10 −0.16 −0.32 −0.08

The enthalpies have been determined for one molecule per 3 × 3 surface unit cell. Esolv is the difference in the formation enthalpies in implicit water and in vacuum. a

clean Pt surface are listed for the electrode potential at U = 0.5 and 0.75 V, respectively, in the absence and presence of the implicit solvent; the difference is denoted as Esolv, which is the difference in the formation enthalpies in implicit water and in vacuum. We use the same reference in all cases, namely, solvated methanol, liquid water, and solvated protons. Thus, Esolv corresponds to the additional gain in adsorption energy upon considering the presence of the solvent at the surface. All considered species are water-soluble, including the solvent, which leads to an additional energy gain; however, the exact amount depends on the particular species. It is relatively large when the reaction intermediates contain a hydrophilic hydroxyl group, such as CH3OH, H2COH, COH, H2COOH, HCOOH, C(OH)2, and COOH. The maximal stabilization of 0.32 eV occurs for the adsorption of C(OH)2, which has two hydroxyl groups. The variation of 0.3 eV in Esolv shows that the consideration of solvation effects can significantly contribute to the ordering in the stability of adsorbates. Thus, we demonstrate that, in the determination of the stability of the reaction intermediates, the presence of the solvent indeed needs to be taken into account. Still, it is important to realize that, upon adsorption from the liquid phase, the solvation shell

Figure 1. Formation enthalpies per surface area of H and OH adsorption on (a) Pt(111), (b) Ru(0001), and (c) the Ru and Pt sites of a Pt2Ru1/Pt(111) surface alloy denoted by Ru@Pt(111) using the concept of the computational hydrogen electrode. The shaded area represents the double layer region in which Pt(111) is not covered by hydrogen or hydroxyl.

the formation enthalpies per surface area of H and OH adsorption are plotted as a function of the electrode potential. The curve with the lowest specific formation energy corresponds to the thermodynamic stable phase. As already found in our previous study, 17 our approach almost quantitatively reproduces the experimentally well-known fact51 that at low potentials Pt(111) is covered by hydrogen, followed by the so-called double layer region between U = 0.5 and 0.75 V where no adsorbate is present; for more positive potentials, OH adsorption begins. As discussed in the Introduction for the methanol electro-oxidation, the double 5578


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energy is strongly overestimated using the RBPE-D3 functional, whereas the adsorption energies of the other considered reaction intermediates are reproduced semiquantitatively. To test the influence of the inaccuracy in the CO adsorption energy on the reaction scheme, we have considered a reduced formation enthalpy of CO in Figure 2 taking the error of approximately 0.5 eV in the adsorption energy according to Table 1 into account. Thus, one obtains the red dashed line in Figure 2. For the corrected CO formation enthalpy, CO2 becomes more stable on Pt(111) than CO already at U = 0.25 V. However, the onset of methanol total oxidation is still at U = 0.6 V, because upon applying this correction to the CO formation enthalpy, surface carboxyl (COOH) becomes the most stable species below 0.6 V. Thus, our findings with respect to the onset of methanol total oxidation are not influenced by the inaccuracy in the CO formation enthalpy, and the uncorrected energy will be used in all further discussions. As far as the possible reaction mechanisms in the methanol electro-oxidation are concerned, on the basis of previous computational studies,3,30,32,39,41 we focus on only three elementary bond breaking and making processes: C−H bond breaking, O−H bond breaking, and C−OH bond creation. Thus, we disregard C−O bond breaking and methane formation. Furthermore, we have chosen an electrode potential of U = 0.6 V at which the total oxidation of methanol becomes stable. In Figure 3, we have constructed possible reaction paths

of the adsorbates becomes broken, whereas a corresponding phenomenon does not occur upon adsorption from the gas phase at the solid−vacuum interface. Hence, the stabilization of the adsorbates at solid−liquid interfaces with respect to the solvated species is in general smaller than the stabilization of adsorbates at solid−vacuum interfaces with respect to the species in the gas phase. Formation enthalpies reflect the chemical stability of reaction intermediates with respect to a reference chemical environment. As mentioned above, here we take the solvated proton via the concept of the computational hydrogen electrode, liquid water, and solvated methanol as the reference species. In thermal equilibrium, the oxidation reaction will end with the reaction product that has the most favorable formation enthalpies. Figure 2 shows the formation enthalpy per molecule

Figure 2. Formation enthalpies per molecule of selected reaction intermediates on clean Pt(111) as a function of electrode potential U in implicit water calculated within a 3 × 3 surface unit cell.

of some selected reaction intermediates on clean Pt(111) determined within a 3 × 3 surface unit cell in implicit water as a function of the electrode potential. Note that the assumption of a clean Pt(111) electrode is only justified in the shaded area in Figure 2 illustrating the double layer region; however, this is the region we are mainly concerned with as this corresponds to the conditions at which methanol electro-oxidation occurs. As demonstrated in Figure 2, CO2 becomes the most stable reaction product only for potentials above U = 0.6 V. Note that the calculated redox potential of methanol in aqueous solution using the implicit solvent model and the concept of the computational hydrogen electrode is 0.24 eV for the RPBE-D3 functional, which compares well with a previously calculated value obtained with a slightly different flavor of a dispersioncorrected functional.17 This means that entirely based on the thermodynamics of the reaction intermediates we predict an overpotential of (0.60 − 0.24) V = 0.36 V for the methanol electro-oxidation on Pt(111). Below U = 0.6 V, the partial oxidation of methanol will be dominant. According to Figure 2, adsorbed CO is the most stable reaction product at these potentials. This would predict that the Pt(111) electrode is blocked by adsorbed CO at low potentials. However, as shown in Table 1, the CO adsorption

Figure 3. Schematic representation of the reaction scheme in methanol electro-oxidation on Pt(111). H removal is denoted by blue arrows, and OH addition is denoted by red arrows. The energies of the reaction intermediates are given in eV relative to aqueous methanol at an electrode potential of U = 0.6 V.

similarly to what has been suggested before.39,41 Hydrogen removal is denoted by blue arrows, and OH addition is denoted by red arrows. Endothermic reaction steps are marked by dashed arrows. The energies given in Figure 3 are relative to aqueous methanol at a potential of U = 0.6 V. Adsorption of methanol from the aqueous phase is associated with an energy gain of 0.55 eV. As the first oxidation step, there 5579


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presence of implicit water on Ru(0001) and on the Ru site of a Pt2Ru1/Pt(111) surface alloy. Indeed, water dissociation is exothermic on Ru(0001) and on Pt2Ru1/Pt(111) by 0.52 and 0.04 eV per water molecule, respectively. The strong interaction of OH with Ru leads to much earlier onset of OH adsorption on Ru(0001) and Pt2Ru1/Pt(111) than on Pt(111) as a function of the electrode potential, as demonstrated in Figure 1b and c. Panels b and c of Figure 1 also show that in the double layer region 0.5 V < U < 0.75 V, hydrogen adsorption is not stable on Ru(0001) and Pt2Ru1/Pt(111). Thus, after water dissociation, protons will be transferred into the electrolyte. Hence on a PtRu alloy, a bifunctional mechanism can be operative: Ru sites provide hydroxyl groups, whereas on Pt sites the OH addition to the carbon species is concerned. As our calculations show, on extended Ru areas, the OH adsorption is rather strong, and OH diffusion is thus not very likely. On single Ru sites in a PtRu surface alloy, OH binding is much weaker because of effective compressive strain effects on the Ru atoms embedded in a matrix of larger Pt atoms.79,90 This means that finely dispersed Ru atoms in a Pt catalyst might be rather beneficial for the methanol electro-oxidation as they can supply hydroxyl groups without any diffusion limitations, as recently demonstrated in a joint experimental and theoretical work addressing CO oxidation on two-dimensional RuPt core-edge nanoclusters.91

are two competing dehydrogenation channels, either to CH3O via an O−H bond breaking process or to H2COH associated with a C−H bond breaking process. In fact, H2COH formation is exothermic with respect to adsorbed methanol, whereas CH3O formation is endothermic, which means that the methanol decay to H2COH is favored on a solvated Pt electrode at U = 0.6 V. Interestingly enough, methanol C−H and O−H breaking are competitive at the Pt(111)−vacuum interface because of the similar height of their activation barriers.33,34 In an electrochemical environment, rather high electrode potentials of approximately U = 1.4 V are needed to make the methoxy formation exothermic. This shows that it is crucial to take the electrochemical environment properly into account in the modeling of electrocatalytic processes. As the second oxidation step, hydroxymethyl (H2COH) can decay to formaldehyde (CH 2 O) or hydroxymethylene (HCOH). The formation of CH2O is slightly endothermic; therefore, we assume that the dehydrogenation step from H2COH to HCOH is favored. As the third oxidation step, both formyl (HCO) and hydroxymethylidyne (COH) can be formed associated with a gain in energy, so their formation might be competitive. As the next step, both removal of H and addition of OH can occur. Hydrogen abstraction converts both COH and HCO exothermically to CO. Methanol oxidation involving a CO intermediate has been called an indirect mechanism.39,41 The direct mechanism avoids the formation of CO and proceeds either from COH to C(OH)2 or from HCO to formic acid (HCOOH) through hydroxyl addition. All these processes are exothermic. As the fifth oxidation step, a carboxyl group (−COOH) is created from C(OH)2, CO, or HCOOH. Formic acid can also be converted to formate (HCOO). Finally, as the sixth reaction step, CO2 is formed either from COOH or HCOO, and the electro-oxidation process is completed. Note that the CO2 desorption and transfer into the electrolyte still requires approximately 0.3 eV according to Figure 3. However, this is no electrochemical step as no charge transfer is involved, and thus the overpotential in the methanol electro-oxidation is not influenced by it. The total oxidation of methanol to carbon dioxide requires the addition of an oxygen atom to some carbon species. In methanol electro-oxidation, this is done through hydroxyl addition, as highlighted by the red arrows in Figure 3, and the hydroxyl needs to be created from water dissociation. Hence, we discuss in the next section how OH can be supplied in the electro-oxidation of methanol.

5. REACTION SCHEME OF METHANOL ELECTROCATALYSIS Up to now we have only considered the energetics of possible reaction intermediates in the methanol electro-oxidation. However, the kinetics of catalytic processes are typically dominated by reaction barriers. Hence, we have determined the activation barriers separating the reaction intermediates in the methanol electro-oxidation using the nudged elastic band (NEB) method.70 As the bond-making and -breaking processes are catalyzed by the electrode surface, in the NEB calculations all reactants and products are taken as adsorbed species. Note, however, that in the reaction scheme shown in Figure 3, we have assumed that the reactants and products are in thermal equilibrium with the electrochemical reservoir in the electrolyte. Thus, for example, all produced hydrogen atoms are not supposed to stay adsorbed on the surface, but they are rather transferred as protons into the electrolyte. Similarly, OH is assumed to be provided by the electrolyte. This should be taken into account when discussing the barrier heights. In the determination of the barrier heights, we assume an adiabatic charge transfer occurs during the electrocatalytic reactions.92−96 In the following, we will address how we estimate the dependence of the reaction barriers on the electrode potential. 5.1. Electrode Potential Dependence of the Reaction Barriers. The electrocatalytic reaction steps of dehydrogenation through H+ formation and OH− association change the oxidation state of the reaction intermediates. In each step, one electron is transferred to the electrode. H+ and OH− ions are not explicitly treated in the present DFT calculations, their energetics are rather referred to the corresponding gas phase species within the concept of the computational hydrogen electrode. In the present formulation of this concept, we assume that the adsorption energies of the reacting species do not depend on the electrode potential, which is a reasonable approximation for species chemisorbed on metal electrodes.13 However, along the reaction path of the reaction step involving electron transfer, the oxidation state of the reacting species and

4. EFFICIENT HYDROXYL SUPPLY CHANNEL IN THE METHANOL ELECTRO-OXIDATION As just mentioned, the supply of OH through water activation is crucial for the methanol electro-oxidation. However, on Pt(111), the dissociation of water into H and OH is unfavorable, being endothermic by approximately 0.4 eV per H2O molecule. Experimentally, Reichert et al.28,29 demonstrated that CO2 formation from HCOOH, which does not need the presence of OH, starts at smaller overpotentials on Pt than CO2 formation from CH3OH. This indicates that the CO2 formation from methanol on Pt requires additional sites that can more easily activate water. These might be provided by Ru sites of a PtRu alloy.48 In general, PtRu alloys are intensively studied86−89 because of their favorable performance as electrocatalysts. To address the role of Ru sites in the water activation, we have determined the energy of H and OH adsorbates in the 5580


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even become spontaneous. It has recently be suggested that the transfer coefficient λ⧧ might be derived from a charge analysis scheme at the transition state.95,96 However, the value of the charge transfer depends on the particular partition scheme used in the calculations. In the following, we rather use the transfer coefficients λ⧧ = 1/2 for the dehydrogenation and λ⧧ = −1/2 for the OH association, which corresponds to the typical value assumed for outer-sphere electron transfer reactions,92,101,102 which is also supported by experimental findings.92 This means that for all calculated barrier configurations an energy correction of −0.5eU (eV) will be applied. Furthermore, we also extend this potential correction scheme to yield an interpolation along the reaction path. Our basic assumption is that the electron transfer rate is linearly related to the length d of the corresponding bond that becomes broken or formed, respectively, in the particular reaction step. Thus, we define the transfer coefficient as a function of the reaction coordinate q according to the NEB calculations as

hence also the dependence of the energetics on the electrode potential changes. Hence, to describe an electrocatalytic reaction at a given electrode potential U, the energies along the reaction path need to be corrected, as has been noted before (e.g., refs 95−101). The dependence of reaction barriers on the electrode potential can, for example, be related to the change of the surface dipole moment along the reaction path;97 it can be derived from a series of calculations with different constant charges,99 or it can be treated in a Butler−Volmer-like formalism invoking a reaction symmetry factor or transfer coefficient.100,101 Here, we will apply a simplified scheme in the spirit of the computational hydrogen electrode, also based on the concept of a transfer coefficient. We assume an adiabatic charge transfer along the reaction path92,93 characterized by a transfer coefficient λ that corresponds to the partial charge being transferred. The resulting correction scheme, in the spirit of the computational hydrogen electrode, is illustrated in Figure 4. At every point

λ(q) = λ⧧

along the reaction path, an electron transfer coefficient λ from or to the electrode is specified. Every DFT energy is then corrected according to the partial charge transfer λ, i.e., by −λeU. The transfer coefficient λ is positive for the electron transfer from H to the electrode, whereas λ is negative for the electron transfer from the electrode to OH. In the dehydrogenation steps, as shown in Figure 4a, the final state energy is corrected by −eU to account for the energy change associated with the transfer of one electron from H to the electrode (λ = 1). In the OH− association, on the other hand, the initial state energy of OH on the Pt electrode is corrected by eU as shown in Figure 4b to reflect the transfer of one electron transfer from the electrode to OH (λ = −1). As a consequence, the energy of the transition state will be modified accordingly. When we denote the electron transfer coefficient at the transition state by λ⧧, the activation barriers of the dehydrogenation steps become (5)

⧧

with a positive λ . In contrast, the activation barrier for the OH association at an electrode potential U is Ea(U ) = Ea − e(λ⧧ + 1)U

d ⧧ − d0

(7)

where d⧧ and d0 are the bond lengths at the transition state and the initial configuration, respectively. The bond length between the NEB images is linearly interpolated; furthermore, |λ| is assumed to be smaller than or equal to 1. Note that in this scheme a molecular reorientation along the reaction path without any change of the relevant bond length keeps the transfer coefficient constant. Note that this scheme relies on the fact that the energetics of the adsorbed species and the barriers only weakly depend on the potential, or equivalently, on the charge state of the Pt(111) electrode or any varying applied electric field. Then, it does not matter whether the work function changes along the reaction path, as for example observed in the oxygen dissociation on Pt(111).99 Such a dependence only vanishes in the limit of an infinite cell size.103 In fact, on metal surfaces this assumption is obviously to a large degree justified due to the good screening properties of metals, which is the reason underlying the success of the computational hydrogen electrode. Adsorption enthalpies on Pt(111) only weakly depend on the strength of an applied electric field.104 Furthermore, it has been demonstrated that the energetics of the barriers in the formic acid degradation on Pt(111) is hardly influenced by varying the surface excess charge.40 Recently, Fang and Liu addressed the dependence of activation barriers on the electrode potential in the initial reaction step of methanol electro-oxidation on Pt(111).43 They used a periodic continuum solvation method in a constantcharge framework for different charge states of the electrode and an a posteriori derivation of the corresponding electrode potential through the determination of the work function. Indeed, they found a linear relationship between the calculated activation barriers and the electrode potential, however, generally not with a slope corresponding to a transfer coefficient of 1/2. For example, for the initial C−H bonding breaking step in the methanol oxidation on Pt(111), they obtained λ⧧ = 0.36. This shows that the concept of outer-sphere reactions is not necessarily applicable to elementary electrocatalytic steps on metal electrodes. However, the consideration of the results from the computationally rather demanding explicit calculation of the potential dependence of the activation barriers43 compared to our simple assumption of an outer-

Figure 4. Schematic illustration of the electrode potential correction applied to the potential energy curves for (a) dehydrogenation and (b) OH association reactions.

Ea(U ) = Ea − eλ⧧U

d(q) − d0

(6)

⧧

with a negative λ For both reaction processes, the barriers become effectively lowered at positive electrode potentials U. At sufficiently high electrode potentials, activated processes can 5581


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exp[−Ea/kBT] with the same prefactor for both reactions, the C−H bond breaking rate then becomes approximately 7,000 times higher than the O−H bond breaking rate at 298 K, compared to a factor of 7 derived from the calculations in vacuum. Thus, our calculations yield a strong preference for an initial C−H bond breaking process in the methanol electrooxidation, in agreement with the experimental data.18 Consideration of the electrochemical environment changes the selectivity between the two initial dehydrogenation channels by 3 orders of magnitude, emphasizing the importance of a proper modeling of the electrocatalytic environment. With regard to the mechanism leading to the change of the selectivity in the implicit solvent, we already discussed that the presence of a hydrophilic hydroxyl group stabilizes the reaction intermediates in aqueous solution. In the C−H bond breaking process from methanol to hydroxymethyl, the O−H group is not modified along the reaction path. Hence, the shift between the potential curves calculated without and with solvent is almost constant as a function of the reaction coordinate, there is only a small reduction in the barrier height. In the methoxy formation, on the other hand, the O−H bond is broken. Hence, the stabilization in the solvent due to the presence of the hydrophilic hydroxyl group is lost along the reaction path, leading to a sizable increase in the barrier height. This selectivity is in good agreement with the calculations by Cao et al., who determined potential-dependent reaction energies in the methanol electro-oxidation on electrified Pt(111) and also considering one specific explicit water layer configuration.30 In contrast, DFT calculations addressing the oxidation of ethanol and isopropanol found that the presence of one single adsorbed water molecule favors the O−H breaking.105 However, first of all, it is not clear whether results with respect to ethanol and isoproponal oxidation can directly be related to methanol electro-oxidation. Second, calculations with just one single water molecule in optimized geometries do not capture the statistical nature of the presence of an aqueous electrolyte. There is certainly a need to validate implicit solvent models, but this requires prohibitively expensive first-principles simulations. 5.3. Kinetics of Methanol Electro-Oxidation. We will now discuss the mechanism of methanol electro-oxidation taking the calculated barrier heights between the reaction intermediates shown in Figure 3 into account. In the discussion, we will exclude CH3O as a reaction intermediate because of the fact that H2COH formation is much more favored, as just shown. The effect of the continuous energy correction scheme according to eq 7 for various electrode potentials on the energies along the reaction path determined by the NEB calculations is illustrated in Figure 6. The dehydrogenation of H2COH is shown in panel (a). In principle, upon applying the correction scheme, the position of the activation barrier can change. In the case of the decomposition of H2COH, the C−H bond stretches steadily along the reaction path. Hence, the barrier position is hardly changed, and the barrier height is lowered by roughly 0.5eU for all considered electrode potentials U (see the inset of Figure 6a). In contrast, the HCOO dehydrogenation starts with a reorientation of the molecule, which costs approximately 0.9 eV. As the C−H bond length hardly changes during this reorientation, according to our scheme, the energetics along the initial part of the reaction path are not altered as a function of the electrode potential (see Figure 6b). Hence, the barrier can only be lowered by up to 0.15 eV as a function of the electrode

sphere mechanism does not modify any of the conclusions drawn from our work. 5.2. Initial Methanol Dehydrogenation Step. Before discussing the whole reaction scheme, we will first focus on the initial methanol dehydrogenation step and, particularly, address the role of the electrolyte with regard to the barrier height. On the basis of spectroscopy experiments, Franaszczuk et al. could show18 that, on platinum in contact with an electrochemical environment, methanol is predominantly converted to hydroxymethyl (H2COH), which is associated with the breaking of one of the C−H bonds of the methyl group of methanol. In contrast, at the Pt−vacuum interface, methanol also decays to methoxy (CH3O) involving the breaking of the O−H bond. It has been speculated that, apart from electric field effects, this difference is due to additional hydrophilic− hydrophobic interactions of methanol in solution.18 The energetics along the initial methanol dehydrogenation steps according to our NEB calculations are illustrated in Figure 5. Note that here we have not applied any energy correction

Figure 5. Potential curves of the initial methanol dehydrogenation on Pt(111) in vacuum (dashed lines) and in implicit water (solid lines) as a function of the reaction coordinate, which corresponds to the images used in the NEB calculations. The potential curves in implicit water correspond to the reactions at U = 0 V. The blue lines correspond to the formation of H2COH involving the breaking of a C−H bond, and the red lines represent the creation of CH3O involving the breaking of an O−H bond.

scheme, i.e., the energetics correspond to U = 0.0 V vs SHE. However, using the correction scheme described in the previous section does not change the difference between two reaction barriers as the same correction is applied to all barrier heights. In vacuum, the barriers for the C−H and O−H bond breaking processes (dashed lines in Figure 5) are very similar at 0.70 and 0.75 eV, respectively. This is in good agreement with the results of previous DFT calculations, which suggested that the initial C−H and O−H bond breaking processes are competitive at the Pt−vacuum interface.34 Interestingly enough, upon introducing the aqueous electrolyte within an implicit water model, the C−H bond breaking barrier is lowered to 0.65 eV, whereas the O−H bond breaking barrier is increased to 0.88 eV (solid lines in Figure 5) at U = 0.0 V. These values are in good agreement with recent calculations by Fang and Liu.43 Plugging these numbers into a simple reaction rate estimate using the Arrhenius relation Γ ≈ 5582


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a C−H bond, respectively. Note that the formation enthalpy of COH is lower than the one of HCO by 0.4 eV. However, in contrast to the initial methanol oxidation step, the O−H bond breaking barrier (0.18 eV at U = 0.0 V) is 0.29 eV smaller than the C−H bond breaking barrier. Applying the correction scheme according to eq 7 even makes some intermediates unstable, i.e., these intermediates are then no longer a local minimum in the multidimensional potential energy surface. Thus, the HCOH decomposition to HCO becomes spontaneous for electrode potentials U > 0.54 V. Hence, HCOH decays predominantly to HCO rather than to COH despite the fact that COH formation is energetically more favorable. Therefore, we neglect the COH formation path based on kinetic reasons. Note that COH has been suggested to be part of the minimum energy path toward methanol oxidation on Pt according to periodic DFT calculations;39 however, in this study, no activation barriers were determined. Now, formyl (HCO) can be converted to carbon monoxide (CO) or formic acid (HCOOH). The dehydrogenation of HCO to CO is hindered by a small activation barrier of 0.16 eV. In the presence of hydroxyl, formic acid (HCOOH) can be created, which also has a rather small activation barrier of 0.19 eV. CO and HCOOH formation are thus competing processes. When CO is created, in the absence of OH, it will stay on the surface because of its large adsorption energy of CO and poison the Pt electrocatalyst. Under the assumption that OH is supplied in a bifunctional mechanism, CO reacts further and forms surface carboxyl (COOH) associated with a gain in free energy at U > 0.6 V. The activation barrier at U = 0.6 V is 0.10 eV. COOH can then decay to CO2 hindered by an activation barrier of 0.49 eV, thus completing the electro-oxidation. In the direct mechanism, HCOOH can be either dehydrogenated or desorbed into the electrolyte. Because of the relatively small adsorption energy of HCOOH of 0.4 eV with respect to solvated HCOOH, HCOOH can be replaced by CH3OH on the electrode surface. The partial oxidation channel ending up with HCOOH formation in solution reduces the efficiency of the DMFC. However, at sufficiently high overpotential, HCOOH dehydrogenation becomes kinetically favored. At U = 0.6 V, formic acid is dehydrogenated to either COOH (activation barrier of 0.14 eV) or HCOO (activation barrier of 0.13 eV). Because of the similar barrier heights, these are certainly competing processes. As Wang and Liu already demonstrated,40 HCOO formation results from a HCOOH configuration with the O−H bond oriented toward the electrode, whereas COOH formation occurs when HCOOH lies flat on the electrode. Upon COOH formation, we are back to the reaction scheme of the indirect mechanism discussed above where CO2 formation proceeds through a barrier of 0.49 eV. HCOO, on the other hand, is rather strongly bound, and its decay to CO2 is hence hindered by a larger barrier of 0.94 eV. Summarizing the discussion, at an electrode potential of U = 0.6 V corresponding to an overpotential of 0.36 V, the indirect mechanism of methanol electro-oxidation involving a CO intermediate and subsequently COOH becomes thermodynamically favorable as the energies of all reaction intermediates along the reaction path are downhill in energy, and the largest activation barrier associated with the last oxidation step is 0.49 eV. The same is true for the direct mechanism involving HCOOH and COOH as reaction intermediates. The reaction intermediates in the direct mechanism going through HCOO as a reaction intermediate are downhill in energy even at lower

Figure 6. Activation barriers of the C−H bond breaking from (a) H2COH and (b) HCOO at U = 0.00, 0.44, 0.60, and 0.75 V. The contributions of the adiabatic electron transfer is approximated by the bond length of the breaking or forming bonds. In insets, the activation barriers are displayed with a function of the electrode potential. A barrier change of −0.5eU is displayed in dashed lines as a guide to the eye.

potential; for potentials larger than 0.3 V, the barrier height does not change any more, as indicated in the inset of Figure 6b. The calculated activation barriers at U = 0.6 V are listed in Table 3, where the configuration at the transition states is also characterized by the distances between the atoms involved in the corresponding bond-breaking or -forming process and the nearest Pt surface atom. Typically, at the transition states, the distance between the Pt surface and the bonding partners of the corresponding reaction is in the range of 1.9−2.4 Å. The most favorable reaction paths for the indirect and the direct mechanism are illustrated in Figure 7, where the reaction energies and the activation barriers are given for U = 0.44, 0.60, and 0.75 V. In the following, we will base the discussion on the energetics at U = 0.6 V as this corresponds to the electrode potential at which CO2 formation on Pt becomes thermodynamically favored. After the first oxidation step from methanol to hydroxymethyl, there is a further decay to hydroxymethylene (HCOH) hindered by an activation barrier of 0.31 eV. Hydroxymethylene can be converted to formyl (HCO) or hydroxymethylidyne (COH) upon breaking either an O−H or 5583


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Table 3. Activation Barriers (Ea) of Reaction Steps in Methanol Electro-Oxidation Calculated Using the Implicit Water Model at U = 0.6 and 0.0 Va Ea(U) (eV) reaction CH3OH CH3OH H2COH HCOH HCOH HCO+OH− HCO COH HCOOH HCOOH CO+OH− HCOO COOH

→ → → → → → → → → → → → →

CH3O+H+ H2COH+H+ HCOH+H+ HCO+H+ COH+H+ HCOOH CO+H+ CO+H+ COOH+H+ HCOO+H+ COOH CO2+H+ CO2+H+

U = 0.6 V

U = 0.0 V

Ea(0.6) − Ea(0.0) (eV)

0.59 0.37 0.31 spontaneous 0.22 0.19 0.16 0.83 0.14 0.13 0.10 0.94 0.49

0.88 0.65 0.58 0.18 0.47 0.32 0.34 1.13 0.44 0.41 0.31 1.08 0.75

−0.29 −0.28 −0.27 −0.18 −0.25 −0.13 −0.18 −0.30 −0.30 −0.28 −0.21 −0.14 −0.26

transition state configurations (Å) at U = 0.0 V dPt−O dPt−C dPt−C dPt−C dPt−C dPt−C dPt−C dPt−C dPt−C dPt−O dPt−C dPt−O dPt−C

= = = = = = = = = = = = =

2.09 2.36 2.00 1.94 2.03 2.29 1.96 2.01 2.38 3.10 1.92 2.35 2.12

dPt−H dPt−H dPt−H dPt−H dPt−H dPt−OH dPt−H dPt−H dPt−H dPt−H dPt−OH dPt−H dPt−H

= = = = = = = = = = = = =

1.64 1.64 1.72 1.76 1.68 2.23 1.90 1.82 2.56 1.58 2.13 1.95 1.65

dO−H dC−H dC−H dO−H dC−H dC−O dC−H dO−H dC−H dO−H dC−O dC−H dO−H

= = = = = = = = = = = = =

1.59 1.51 1.42 1.29 1.34 2.03 1.27 1.30 1.74 1.70 1.88 1.44 1.59

The charge transfer coefficients at the transition state λ⧧ are selected to be 1/2 and −1/2 for the dehydrogenation and the OH association steps, respectively. The intermediates are bound to Pt(111) either through a carbon or oxygen atom. The transition state configurations at U = 0.0 V are characterized by the distance between the bonding oxygen or carbon atom and the nearest Pt atom (dPt−C/O), by the distance between the reaction partner H or OH and the nearest Pt atom (dPt−H/OH), and by the distance between bonding oxygen or the carbon atom and the reaction partner H or OH (dC/O−H/OH). a

Figure 7. Methanol electro-oxidation scheme for the indirect mechanism involving a CO intermediate (starting from the left-hand side) and the direct mechanism through a HCOO intermediate (starting at HCO from the right-hand side) at U = 0.44, 0.60, and 0.75 eV. The reaction starts with CH3OH and OH− being present on the Pt electrode. All reaction energies given in eV in green are taken with respect to methanol in aqueous solution. The activation barriers at U = 0 V (in violet) are derived from NEB calculations. Spontaneous reactions are denoted by *.

6. CONCLUSIONS

overpotentials. However, in this mechanism, the last oxidation step from formate to CO2 is hindered by a relatively large activation barrier, which amounts to 0.94 eV at U = 0.6 V, whereas for the other two paths, the last oxidation step is only hindered by a barrier of 0.49 eV. This means that formate can be formed relatively easily, but it might be hard to remove it so that it could act as a poison for the methanol electro-oxidation by blocking active sites of the Pt electrode. Indeed, adsorbed formate was observed in the electro-oxidation of methanol on Pt.26,29 However, there are experimental indications that adsorbed formate created during methanol electro-oxidation would be replaced by the more strongly bound CO.29

We have derived a reaction scheme of methanol electrooxidation on a Pt electrode as a function of the electrode potential based on DFT calculations representing the electrolyte by an implicit solvation method and using the concept of the computational hydrogen electrode. Thus, we have taken the electrochemical environment into account in the determination of the reaction scheme. It turns out that it is crucial to consider the aqueous electrolyte in the calculations as reaction intermediates containing hydrophilic groups become significantly stabilized in the presence of water. We predict that the total electro-oxidation of methanol to CO2 on Pt becomes thermodynamically stable at electrode potentials above U = 0.6 V vs SHE, in agreement with the experimental data. 5584


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(13) Nørskov, J. K.; Rossmeisl, J.; Logadottir, A.; Lindqvist, L.; Kitchin, J. R.; Bligaard, T.; Jónsson, H. J. Phys. Chem. B 2004, 108, 17886−17892. (14) Schnur, S.; Groß, A. Catal. Today 2011, 165, 129−137. (15) Hartnig, C.; Koper, M. T. M. J. Am. Chem. Soc. 2003, 125, 9840. (16) Mathew, K.; Sundararaman, R.; Letchworth-Weaver, K.; Arias, T. A.; Hennig, R. G. J. Chem. Phys. 2014, 140, 084106. (17) Sakong, S.; Naderian, M.; Mathew, K.; Hennig, R. G.; Groß, A. J. Chem. Phys. 2015, 142, 234107. (18) Franaszczuk, K.; Herrero, E.; Zelenay, P.; Wieckowski, A.; Wang, J.; Masel, R. I. J. Phys. Chem. 1992, 96, 8509−8516. (19) Gasteiger, H. A.; Markovic, N.; Ross, P. N.; Cairns, E. J. J. Phys. Chem. 1993, 97, 12020−12029. (20) Gasteiger, H. A.; Marković, N.; Ross, P. N.; Cairns, E. J. J. Electrochem. Soc. 1994, 141, 1795−1803. (21) Kua, J.; Goddard, W. A. J. Am. Chem. Soc. 1999, 121, 10928− 10941. (22) Dinh, H. N.; Ren, X.; Garzon, F. H.; Zelenay, P.; Gottesfeld, S. J. Electroanal. Chem. 2000, 491, 222−233. (23) Hoster, H. E.; Iwasita, T.; Baumgärtner; Vielstich, W. Phys. Chem. Chem. Phys. 2001, 3, 337. (24) Iwasita, T. Electrochim. Acta 2002, 47, 3663−3674. (25) Dubau, L.; Coutanceau, C.; Garnier, E.; Léger, J.-M.; Lamy, C. J. Appl. Electrochem. 2003, 33, 419−429. (26) Chen, Y. X.; Miki, A.; Ye, S.; Sakai, H.; Osawa, M. J. Am. Chem. Soc. 2003, 125, 3680−3681. (27) Zhao, X.; Yin, M.; Ma, L.; Liang, L.; Liu, C.; Liao, J.; Lu, T.; Xing, W. Energy Environ. Sci. 2011, 4, 2736−2753. (28) Reichert, R.; Schnaidt, J.; Jusys, Z.; Behm, R. J. ChemPhysChem 2013, 14, 3678−3681. (29) Reichert, R.; Schnaidt, J.; Jusys, Z.; Behm, R. J. Phys. Chem. Chem. Phys. 2014, 16, 13780−13799. (30) Cao, D.; Lu, G.-Q.; Wieckowski, A.; Wasileski, S. A.; Neurock, M. J. Phys. Chem. B 2005, 109, 11622−11633. (31) Ishikawa, Y.; Liao, M.-S.; Cabrera, C. R. Surf. Sci. 2000, 463, 66− 80. (32) Desai, S. K.; Neurock, M.; Kourtakis, K. J. Phys. Chem. B 2002, 106, 2559−2568. (33) Greeley, J.; Mavrikakis, M. J. Am. Chem. Soc. 2002, 124, 7193− 7201. (34) Greeley, J.; Mavrikakis, M. J. Am. Chem. Soc. 2004, 126, 3910− 3919. (35) Hartnig, C.; Spohr, E. Chem. Phys. 2005, 319, 185. (36) Kandoi, S.; Greeley, J.; Sanchez-Castillo, M.; Evans, S.; Gokhale, A.; Dumesic, J.; Mavrikakis, M. Top. Catal. 2006, 37, 17−28. (37) Janik, M. J.; Taylor, C. D.; Neurock, M. Top. Catal. 2007, 46, 306−319. (38) Neurock, M.; Janik, M.; Wieckowski, A. Faraday Discuss. 2009, 140, 363−378. (39) Ferrin, P.; Nilekar, A. U.; Greeley, J.; Mavrikakis, M.; Rossmeisl, J. Surf. Sci. 2008, 602, 3424−3431. (40) Wang, H.-F.; Liu, Z.-P. J. Phys. Chem. C 2009, 113, 17502− 17508. (41) Ferrin, P.; Mavrikakis, M. J. Am. Chem. Soc. 2009, 131, 14381− 14389. (42) Ferrin, P.; Kandoi, S.; Nilekar, A. U.; Mavrikakis, M. Surf. Sci. 2012, 606, 679−689. (43) Fang, Y.-H.; Liu, Z.-P. Surf. Sci. 2015, 631, 42−47. (44) Behrens, M.; Studt, F.; Kasatkin, I.; Kühl, S.; Hävecker, M.; Abild-Pedersen, F.; Zander, S.; Girgsdies, F.; Kurr, P.; Kniep, B.-L.; Tovar, M.; Fischer, R. W.; Nørskov, J. K.; Schlögl, R. Science 2012, 336, 893−897. (45) Günther, S.; Zhou, L.; Hävecker, M.; Knop-Gericke, A.; Kleimenov, E.; Schlögl, R.; Imbihl, R. J. Chem. Phys. 2006, 125, 114709. (46) Sakong, S.; Groß, A. J. Catal. 2005, 231, 420−429. (47) Sakong, S.; Groß, A. J. Phys. Chem. A 2007, 111, 8814−8822. (48) Shubina, T. E.; Koper, M. T. Electrochem. Commun. 2006, 8, 703−706.

However, not only thermochemistry but also kinetics are important to derive the most probable reaction schemes. In the first dehydrogenation step, the formation of hydroxymethyl associated with the breaking of a C−H bond and the formation of methoxy involving the breaking of the O−H bond are hindered by similar activation barriers at the Pt−vacuum interface. In the presence of the aqueous electrolyte, the selectivity toward hydroxymethyl formation is increased by 3 orders of magnitude at room temperature. The reaction scheme involving COH as an intermediate is disregarded because of kinetic reasons. As the last reaction intermediate before CO2 formation, we identify adsorbed surface carboxyl (COOH), which might be formed either in the indirect mechanism involving adsorbed CO as a reaction intermediate or from formic acid (HCOOH) in the direct mechanism. Our findings are all consistent with the known experimental data on methanol electro-oxidation on Pt electrodes. This study indicates that it is important to appropriately take the electrochemical environment into account in the first-principles modeling of electrocatalytic reactions.

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This research has been supported by the German Research Foundation (DFG) through the research unit FOR 1376 (DFG contract GR 1503/21-2) and by the Baden-Württemberg Foundation within the Network of Competence “Functional Nanostructures” through project CT-04 ORR-Scale of the CleanTech program. Computer time on the JUSTUS cluster at Ulm University has been provided by the bwHPC initiative and the bwHPC-C5 project, which are funded by the BadenWürttemberg government (MWK) and the German Research Foundation (DFG).

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