, Volume 8, Issue 6, pp 577–586 | Cite as

Methanol Oxidation on Pt(111) from First-Principles in Heterogeneous and Electrocatalysis

  • Sung Sakong
  • Axel Groß
Original Article


The catalytic oxidation of methanol on Pt(111) has been addressed based on first-principles electronic structure calculations. The chemical environment corresponding to the conditions in heterogeneous and electro-catalysis has been taken into account in a grand-canonical approach. Furthermore, the aqueous electrolyte in electrocatalysis has been described in an implicit solvent model. Thus, we find characteristic differences between the methanol oxidation paths in heterogeneous and electro-catalysis. The presence of the aqueous electrolyte stabilizes reaction intermediates containing hydrophilic groups thus also influencing the selectivity in the methanol oxidation. In addition, adsorbed hydrogen on Pt(111) is shown to render the electro-oxidation of methanol less efficient.

Graphical Abstract

The difference between methanol oxidation in heterogeneous and electro-catalysis has been studied theoretically from first principles employing a grand-canonical approach.


Electrocatalysis Heterogeneous catalysis Methanol oxidation Pt(111) Quantum Chemistry Density functional theory Theoretical electrochemistry 


Reactions at electrochemical interfaces play an increasingly important role in devices for chemical energy conversion and storage [1]. This is reflected in a growing number of theoretical studies addressing electrode/electrolyte interfaces from first-principles [2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, 19, 20, 21, 22, 23] on the basis of periodic density functional theory (DFT) calculations. Electrocatalytic processes occur at the electrode/electrolyte interface. For a proper theoretical description of these reactions, it is critical to take into account the explicit structure of this interface [16, 19] which is governed by the species present in the electrolyte. Furthermore, the liquid nature of the electrolyte requires a numerically demanding statistical averaging over many possible configurations [2, 24] in order to derive free energies rather than total energies. However, because of limited computer power, first-principles calculations are still computationally too expensive to perform such statistical averages on a routine basis. As an attractive alternative, the influence of the electrolyte on structures and processes at electrochemical interfaces might be described in a grand-canonical approach with the species present in the electrolyte represented through their (electro-)chemical potential.

In the field of heterogeneous catalysis where reactions typically occur at the gas/solid-interface, this grand-canonical approach has been coined ab initio thermodynamics [25, 26]. Usually, the gas phase is not explicitly modeled but treated as a thermodynamic reservoir. The thermodynamic conditions just enter through the dependence of the chemical potentials of the atomic species on temperature, pressure, concentration, etc. Thus, e.g., the structure of metal catalysts in oxidation catalysis as a function of CO and oxygen partial pressure and temperature could be determined from first principles [27].

The concept of combining atomistic first-principles calculations with a grand-canonical approach to describe the environment was extended by Nørskov et al. to address the electrochemical environment characterized by the electrode potential and the concentrations of the species solvated in the electrolyte [28, 29, 30, 31]. This method is now called the “computational hydrogen electrode”.

Chemical reactions in heterogeneous catalysis occur at solid/gas interfaces, whereas the electrochemical reactions in electrocatalysis occur at solid/liquid interfaces. In heterogeneous catalysis, therefore, it is natural for reaction intermediates to refer their adsorption energy to free gas molecules. However, in electrocatalysis, the reference species correspond to charged particles or ions in solution. To compute the energetics of these states again in principle require to calculate solvation energies on the basis of numerically very demanding thermodynamic integration schemes [32]. However, these expensive calculations can be avoided realizing the fact that typically solvated species are related to corresponding gaseous species given by standard electrode potentials. For example, the hydrogen electrochemical potential for a given electrode potential U can be expressed as
$$ \tilde{\mu}_{\mathrm{H}^{+}} + \tilde{\mu}_{\mathrm{e}^{-}} = \frac{1}{2} \mu_{\mathrm{H}_{2}} - eU_{\text{SHE}} -k_{B}T \ln(10)\text{pH} \ , $$
where one has used the fact that at standard conditions defining the standard hydrogen electrode (SHE) potential U SHE the solvated proton is in equilibrium with the H 2 molecule in the gas phase. The entropic contribution of varying proton concentration \(-k_{\mathrm {B}}T\ln (10) \text {pH}\) corresponds to a shift of about 59 meV in the electrochemical potential at room temperature if pH is changed by one. Thus, the computational hydrogen electrode method offers a convenient way to replace the reference energy of solvated ionic species by the reference energy of gaseous molecules.

In studies based on the concept of the computational hydrogen electrode, often the influence of the presence of the electrolyte on the energetics of the adsorbed system is neglected [15, 19, 29] which can be justified by the relatively weak interaction of water with metal electrodes and simple adsorbates [14, 33, 34]. However, the concept of the computational hydrogen electrode allows to take the effect of the environment on adsorption energies into account. Still, as mentioned above, considering the presence of an aqueous electrolyte explicitly is prohibitively computationally expensive.

As a numerically attractive alternative, the electrolyte can be described as a dielectric continuum, i.e., in an implicit solvent model [35, 36, 37, 38]. Recently, we have used this approach to derive the equilibrium coverage of Pt(111) with hydrogen and OH as a function of the electrode potential [14]. Furthermore, we have addressed the methanol electro-oxidation on Pt(111) which is the crucial process occurring on the anode of the direct methanol fuel cell (DMFC) [39, 40]. We demonstrated that the presence of the aqueous electrolyte has a decisive influence on the selectivity within the reaction scheme [16] at an electrode potential of U = 0.6 V, at which Pt(111) is not covered by neither hydrogen nor OH and at which CO 2 formation becomes thermodynamically stable.

In this study, we will extend this previous study [16] by establishing a comparable scheme valid for the methanol oxidation in heterogeneous catalysis. By performing a detailed comparison of the reaction scheme of methanol oxidation in heterogeneous and electro-catalysis, we will gain a better understanding of the role of the electrochemical environment on electro-catalytic processes. Furthermore, we will also consider methanol electro-oxidation at lower electrode potentials at which Pt(111) electrode is covered by protons, thus assessing the influence of adsorbates on the reaction energetics in methanol oxidation.

Computational Details

Periodic DFT calculations have been performed using the software package vasp [41]. The exchange-correlation energies are evaluated within the generalized gradient approximation as suggested by Hammer and Nørskov, known as a revised version of the Perdew-Burke-Ernzerhof (RPBE) functional [42]. Dispersion effects have been considered within the semi-empirical D3 dispersion correction scheme of Grimme using the zero damping function [43, 44, 45]. The RPBE-D3 approach correctly predicts properties of liquid water [17, 46, 47]. In addition, it reliably describes the interaction of organic molecules [16, 48] and water [17, 46, 47] with metal surfaces. The liquid electrolyte is modeled by a polarizable dielectric continuum as implemented into the vasp code by Mathew and Hennig [14, 49]. Thus, the thermally averaged contribution of the electrolyte is effectively included in the DFT calculations yielding solvation energies and the local potential [35, 36, 37, 38].

The wave functions were expanded up to a cutoff energy of 700 eV using a plane wave basis set to describe the solvation cavity accurately. The electronic cores are described by the projector augmented wave method [50]. The Pt lattice parameter of a = 3.99 Å is optimized using a fine k-point grid of 21 × 21 × 21. The Pt catalyst is modeled by a Pt(111) slab consisting of five atomic layers. The top three layers of the slabs are fully relaxed, while the lowest 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 cutoff radius for the pair interactions of the D3 correction has been chosen to be 10 Å. However, the screening of the van der Waals interactions in bulk metals is not correctly described in dispersion correction schemes [46, 51]. Therefore we exclude the dispersion correction for all metal atoms below the first layer.

The properties of liquid water are described with a dielectric constant of 𝜖 b = 80 and a cutoff charge density of ρ cut = 0.0025 Å −3. The cavitation energies are calculated using a surface tension parameter of 0.525 meV/Å 2 [52] as described in Ref. [53, 54]. The energies and forces of stable configurations satisfy the convergence criteria of 10 −6 eV and 0.01 eV/Å.

In the calculations, we will treat the chemical potential as a free parameter. The dependence of the reaction energetics on the thermodynamic conditions then enter through the corresponding dependence of the chemical potentials. Thus, the energetics of a given atomic adsorption configuration is given by the formation enthalpy
$$ {\Delta} H = E_{tot}-\sum\limits_{i} n_{i} \mu_{i} $$
with the total energy E t o t , chemical potential μ i , and the number n i of atomic species i per unit cell.

Chemical and Electrochemical Environments

As the first step in a grand-canonical simulation, the proper thermodynamic reservoirs of stable species have to be defined. The elements that are involved in the chemical reactions are carbon, hydrogen and oxygen. The specific reservoirs employed in this study are specified in Table 1. Let us first concentrate on the approach appropriate for the conditions in heterogeneous catalysis. Reactions in heterogeneous catalysis typically occur at the solid/gas interfaces. Therefore, the reference of the considered chemical species corresponds to gas-phase molecules. Let us consider the total oxidation of methanol on Pt(111) which can be expressed as
$$ 2\mathrm{CH_{3}OH (g)}+3\mathrm{O_{2}(g)} \rightarrow 2\mathrm{CO_{2}(g)} + 4\mathrm{H_{2}O(g)} \ . $$
Specifically, for the heterogeneous catalysis situation, we refer C and H to gas-phase methanol and water, respectively.
Table 1

The selection of (electro-)chemical potentials in methanol oxidation heterogeneous catalysis and electrocatalysis


Heterogeneous catalysis

chemical potential μ


electrochemical potential \(\tilde {\mu }\)


\(E_{\mathrm {CH_{3}OH(g)}}- 2 E_{\mathrm {H_{2}O(g)}} + \mu _{\mathrm {O}}\)

\(E_{\mathrm {CH_{3}OH(aq)}}-E_{\mathrm {H_{2}O(aq)}}-2\tilde {\mu }_{\mathrm {H}^{+}}\)


\((E_{\mathrm {H_{2}O(g)}}-\mu _{\mathrm {O}})/2\)

\( \tilde {\mu }_{\mathrm {H}^{+}}\) (1)


\(\mu _{\mathrm {O}}[p_{\mathrm {O_{2}(g)}},T]\)

\(E_{\mathrm {H_{2}O(aq)}}-2\tilde {\mu }_{\mathrm {H}^{+}}\)

In practice, we will treat the chemical potential of O species μ O as the control parameter. The chemical potential is related to the partial pressure and temperature of an O 2 gas. Then, the formation enthalpy calculated by Eq. 2 becomes a function of μ O as shown in Fig. 1a. Note that we relate the hydrogen chemical potential to gas phase water so that according to Table 1 the chemical potential of hydrogen becomes a function of the chemical potential of oxygen.
Fig. 1

a Formation enthalpies of selected molecules and adsorbates in heterogeneous catalysis are displayed as a function of oxygen chemical potential at the chemical equilibrium between methanol and water molecules. Hydrogen and oxygen adsorption energies have been determined at a coverage of 1/9. b Formation enthalpies of intrinsic H and OH species in electrocatalysis at the electrode/electrolyte interface are plotted as a function of electrode potential with respect to the standard hydrogen electrode, aqueous methanol, and aqueous water

Now we specify a range of oxygen chemical potentials that is beneficial for methanol oxidation. First of all, we want to avoid any adsorbed hydrogen on Pt(111) in order to prevent any back reaction towards methanol. Thus under operating conditions, according to Fig. 1a the chemical potential of oxygen should be larger than − 1.4 eV. Second, the adsorption of oxygen on the surface leading to a passivation and a possible oxidation of the catalyst should also be avoided, leading to an upper bound of − 1.1 eV for the oxygen chemical potential. The favorable range of − 1.4 < μ O < −1.1 eV is illustrated by the shaded area in Fig. 1a.

The chemical potential can be adjusted by changing the temperature and/or the partial pressure of dioxygen gas [55, 56, 57]. Assuming conditions of an ideal gas [58, 59, 60], at 700 K, the operating conditions correspond to a range of oxygen partial pressures of \(10<p_{\mathrm {O_{2}(g)}}<10^{5}\) μPa.

In contrast to heterogeneous catalysis, in electrocatalysis reactions occur at the solid electrode/liquid electrolyte interfaces. The reference states of the considered chemical species usually correspond to ions solvated in the electrolyte. The relevant overall reaction in methanol electro-oxidation can be expressed as
$$ \mathrm{CH_{3}OH (aq)}+\mathrm{H_{2}O(aq)} \leftrightharpoons \mathrm{CO_{2}(aq)} + 6\mathrm{H^{+}} +6e \ . $$

In order to establish the electrochemical environment, we select as the stable species in the grand-canonical reservoir solvated methanol and protons and aqueous water according to Table 1. The crucial parameter electrode potential enters the formalism through the electrochemical potential of solvated protons as shown in Eq. 1.

In Fig. 1b, the formation enthalpies per surface area of the important surface species hydrogen and hydroxyl are plotted. As recently discussed [16], in the potential range 0.5 < U < 0.75 V, the so-called double-layer region, clean Pt(111) is thermodynamically stable, i.e., the Pt(111) electrode is not covered by any specifically adsorbed species. This means that hydrogen atoms on the electrode created through dehydrogenation reactions will readily be transfered into the electrolyte, thus prohibiting backward hydrogenation reactions. Aqueous water molecules serve as the source for oxygen via the water dissociation into adsorbed hydrogen and hydroxyl. As surface oxygen is energetically less favorable than adsorbed hydroxyl at the Pt(111)/water interface [14], surface hydroxyl is supposed to take part in the reactions.

Energetics of Reaction Intermediates

As the elementary reaction steps, we take into account C-H and O-H bond breaking and C-O bond formation. All possible reaction intermediates based on this assumption are listed in Table 2 and plotted in Fig. 2 with respect to the selected chemical and electrochemical environments presented in Table 1. For both environments, gas phase and water, CO 2 is not the energetically most favorable product at the lower limit of the chemical potentials, but it only becomes the most stable adsorbed species at μ O = −1.34 eV with respect to oxygen gas in heterogeneous catalysis and at U = 0.6 V with respect to the standard hydrogen electrode in electrocatalysis. Note that the reported values for the standard electrode potential of the total electro-oxidation of methanol are 0.02 V [61] and 0.032 V [62], whereas we find a value of 0.24 V in water described as an implicit solvent [16]. Note furthermore that the gas phase chemical potential is a function of the partial pressure and temperature, i.e., any given chemical potential corresponds to a range of combinations of these two quantities. A chemical potential of μ O = −1.34 eV can for example been obtained by having a temperature of 700 K and a partial pressure of 25 μPa.
Fig. 2

Formation enthalpy of selected reaction intermediates in a heterogeneous catalysis and b electrocatalysis. The vertical dashed lines represent the chemical potential where CO 2 becomes the most favored species

Table 2

Formation enthalpies of reaction intermediates of methanol oxidation in heterogeneous catalysis and in electrocatalysis on clean Pt(111)


Heterogeneous catalysis


potential (eV):

μ O = −1.38

μ O = −1.10

U = 0.50

U = 0.75


− 0.57

− 0.57

− 0.55

− 0.55


− 0.12

− 0.26


− 0.23


− 0.96

− 1.10

− 0.91

− 1.16

H2C(OH) 2

− 1.10

− 1.38

− 0.83

− 1.33

HC(OH) 2

− 1.51

− 1.93

− 1.24

− 1.99


− 0.57

− 0.85

− 0.44

− 0.94


− 1.63

− 2.05

− 1.68

− 2.43


− 1.37

− 1.79

− 1.28

− 2.03


− 1.15

− 1.43

− 1.05

− 1.55


− 0.22

− 0.78


− 0.80


− 0.72

− 1.14

− 0.47

− 1.22


− 1.70

− 2.26

− 1.34

− 2.34


− 1.88

− 2.58

− 1.49

− 2.74


− 2.31

− 3.01

− 1.97

− 3.22

C(OH) 2

− 2.03

− 2.59

− 1.86

− 2.86

CO f c c

− 2.22

− 2.78

− 2.07

− 3.07

CO t o p

− 2.24

− 2.80

− 2.07

− 3.07


− 2.29

− 3.13

− 1.88

− 3.38

The presence of the aqueous environment does not change the adsorption energies dramatically. However, when we compare the formation enthalpies of reaction enthalpies of isomers such as methoxy (CH 3O) and hydroxymethyl (H 2COH) that have a common reference energy in the reservoir, we find that species containing hydrophilic groups such as the hydroxyl group in H 2COH become more energetically stabilized in the aqueous environment which can be quantisized by evaluating
$$\begin{array}{@{}rcl@{}} {\Delta} H_{\text{OH}} &=& ({\Delta} H^{\mathrm{l}} (\mathrm{H}_{2}\text{COH}) - {\Delta} H^{\mathrm{l}} (\text{CH}_{3}\mathrm{O})) \\ &&- ({\Delta} H^{\mathrm{g}} (\mathrm{H}_{2}\text{COH}) - {\Delta} H^{\mathrm{g}} (\text{CH}_{3}\mathrm{O})) \ , \end{array} $$
where ΔH l is the formation enthalpy according to Eq. 2 in the liquid phase appropriate for electrocatalysis and ΔH g in the gas phase appropriate for heterogeneous catalysis. In fact, this significantly modifies the selectivity in the electrochemical environment compared to the gas-phase environment [16] and is in good agreement with a joint experimental-theoretical study of the methanol decomposition on platinum [67]. This additional stabilization through the presence of the hydroxyl group is also observed when comparing COH with HCO, HCOOH with H 2COO, and COOH with HCOO, as illustrated in Fig. 3 where the stabilisation ΔH OH of reaction intermediates in the liquid phase is plotted for the four isomeric pairs mentioned above.
Fig. 3

Stabilization ΔH OH of reaction intermediates in an aqueous electrolyte containing hydrophilic OH groups with respect to their isomers without OH groups determined according to Eq. 5

The reaction paths connecting the intermediates in heterogeneous catalysis and electrocatalysis are illustrated in Fig. 4. The main difference is that OH is involved in C-O bond formation in electrocatalysis and atomic oxygen in heterogeneous catalysis. Consequently, an additional O-H bond breaking is required in electro-oxidation. Furthermore, the diol molecules having two hydroxyl groups can only be formed readily in electrocatalysis.
Fig. 4

Reaction paths of methanol oxidation a in heterogeneous catalysis and b in electrocatalysis

Under the conditions favorable for total oxidation described above, a higher oxidation number of the carbon species is associated with a larger formation enthalpy (see Table 2). As there is a range of chemical potentials at which CO 2 formation on clean Pt(111) is thermodynamically favorable, previous theoretical studies addressing the methanol oxidation in heterogeneous [63, 64, 65, 66] and electrocatalysis [67, 68, 69, 70, 71] correctly captured the essentials of methanol total oxidation. However, methanol oxidation on clean Pt(111) does not correspond to the optimum situation. The oxidation will be more efficient in oxygen rich condition in which the catalyst becomes oxygen- or hydroxyl-covered [57].

We stress that a meaningful comparison of the results of first-principles calculations with experiment can only be done when the appropriate chemical environment is taken into account in the calculations. For example, the mass spectroscopy experiments by Reichert et al. [40, 72] demonstrated that the CO 2 formation from methanol and the decrease of the amount of adsorbed CO start at 0.6 V (see Fig. 6 of Ref. [40]). This corresponds to the regime in which the clean Pt(111) electrode is the most stable surface termination. Hence, the formation enthalpies shown in Fig. 2b are suitable to describe the experimental situation. When we assume that the energetically mostly stable species exists dominantly on the electrode, the crossover of the largest formation enthalpy from CO to CO 2 in Fig. 2b is directly comparable to the experimental findings. To illustrate the importance of the proper consideration of the chemical environment, let us use the formation enthalpies of the reaction intermediates derived from the DFT calculations in vacuum for the electrocatalytic situation. Then, CO 2 becomes the most favorable reaction product already at an electrode potential of 0.52 V which corresponds to a down-shift of about 0.1 V compared to the calculations taking the aqueous solvent into account, thus also reducing the agreement with the experiment.

Figure 4 summarizes the energetics of the reaction intermediates in methanol oxidation in heterogeneous and electrocatalysis with respect to gas phase methanol and solvated methanol, respectively. The oxygen gas phase reservoir is characterized by an oxygen chemical potential of μ O = −1.34 eV, whereas the energies of the intermediates in electrocatalysis are derived for an electrode potential of U = 0.6 V. The green and blue arrows represent C-H and O-H bond breaking reactions, respectively, while the brown and red arrows denote C-O bond formation by O and OH associations. The solid and dashed arrows indicate exothermic and endothermic reactions, respectively. In vacuum, the oxidation through HCOH proceeds via a sequence of exothermic reactions (see Fig. 4a). An analysis of the energetics indicates that CH 3O and CH 2O formation are not favorable. However, methoxy formation is kinetically accessible because of the comparable activation barriers of C-H and O-H bond breaking in vacuum [16, 65]. Thus, CO 2 can be formed either in the so-called indirect mechanism involving adsorbed CO or in a direct mechanism from formic acid (HCOOH) [64, 65].

In water, the electro-oxidation paths through HCOH consist of consecutive exothermic reaction steps (see Fig. 4b). However, due to the presence of the aqueous electrolyte methoxy formation becomes kinetically hindered, as we demonstrated recently [16]. This result agrees well with the experimental finding of a predominant H 2COH formation at the water/Pt(111) interface [73].

The reaction schemes depicted in Fig. 4 rely on an efficient supply channel of oxygen species, but the mechanism has not been explicitly discussed yet. In heterogeneous catalysis, oxygen atoms are provided by the dissociative adsorption of O 2 molecules which is operative above 150 K [74, 75, 76].

However, in the reaction in the aqueous environment, oxygen species are supplied by the dissociation of water molecules. Since the water dissociation on Pt(111) is hindered by a relatively large activation barrier of 0.8 eV [77, 78, 79], the hydroxyl production will be a rate limiting step in the whole oxidation process. Without sufficient OH production, CO adsorbed on the electrode will be the dominant reaction product. In the experiments by Reichert et al. [40, 72], below 0.6 V, CO is the only detectable surface species and no CO 2 formation occurs. This indicates that there is a need for an efficient water dissociation mechanism in order to supply hydroxyl, for example at step edges of Pt or small Ru particle dispersed on the electrode.

Hydrogen-Covered Pt(111) Electrode

Finally, we discuss the influence of the electrochemical environment beyond optimal operating conditions. As discussed in the “Chemical and Electrochemical Environments” section, below 0.5 V the total oxidation of methanol is thermodynamically not favorable. In addition, the Pt(111) electrode is covered by H atoms that are supplied by the protons in the electrolyte. Hence, we focus on a partial oxidation path to CO formation. Note that in heterogeneous catalysis adsorbed hydrogen originates from gas-phase H 2 molecules which typically adsorb dissociatively on metal surfaces [80].

First, we address the initial methanol adsorption on Pt(111) as a function of hydrogen coverage. At the saturation coverage of 𝜃 H = 1, methanol is placed on top of the hydrogen layer, whereas for smaller hydrogen coverages methanol is put at one of the free Pt sites.

In Fig. 5a, the formation enthalpies of hydrogen and methanol co-adsorption on Pt(111) are plotted as a function of the electrode potential. Methanol adsorption on the fully H-saturated electrode (red line) is only stable below 0.36 V, whereas methanol replaces an adsorbed hydrogen atom above 0.36 V.
Fig. 5

a Formation enthalpy per area of methanol on H-covered Pt(111) electrodes for various H coverages and b formation enthalpy per area of the reaction intermediates on an H-covered electrode with 𝜃 H = 8/9. The concentration of the reaction intermediates on the electrode is 1/9 within a 3 × 3 surface unit cell

The calculations show that a methanol molecule can replace a surface H atom only above 0.36 V. Below 0.36 V, on the H saturated Pt electrode, because of the weaker interaction the adsorbed methanol molecule is located about 1 Å further away from the surface than on the clean Pt electrode. Then, it becomes difficult to catalyze the initial step of the methanol electro-oxidation directly by the Pt electrode, and CO formation from methanol will be slow and limited as observed in mass spectroscopy experiments [40].

In Fig. 5b, the formation enthalpies of the considered reaction intermediates are displayed for a hydrogen coverage of 𝜃 H = 8/9 (solid lines) and without any adsorbed hydrogen atoms (dashed lines). The adsorption of the reaction intermediates becomes more stable on clean Pt(111) than on H-covered Pt(111) at potentials above 0.4 V. Our calculations indicate that in the range 0.4 < U < 0.5 V, the reaction intermediates in the methanol oxidation replace the adsorbed hydrogen atoms in a competitive fashion, similar to what has be observed in the co-adsorption of halides and hydrogen [19, 81]. Thus, the presence of methanol in the aqueous electrolyte even reduces the potential region in which so-called underpotential deposited (upd) hydrogen is stable on Pt(111).

Below 0.4 V, the reaction intermediates will be surrounded by adsorbed H atoms. Thus, hydrogenation reactions reversing the oxidation become possible. Specifically, the formation of H 2COH (solid green line) becomes more favorable than HCOH formation (solid brown line) below 0.4 V which means that the hydrogenation of HCOH becomes exothermic. Then, this oxidation step will hinder the further progress of the partial oxidation of methanol significantly and will be responsible for a low efficiency towards CO formation.


Using a grand-canonical approach, we have addressed methanol oxidation on Pt(111), both in a situation corresponding to heterogeneous catalysis at the solid/gas interface as well as for a electrocatalytic environment at the solid/water interface. The aqueous electrolyte has been taken into account within an implicit solvent model. Thus, we have identified characteristic differences between the methanol oxidation paths in heterogeneous and electro-catalysis. The presence of the aqueous electrolyte leads to the stabilization of reaction intermediates that contain hydrophilic groups such as hydroxyl when isomers are compared. This also modifies the selectivitiy of the methanol oxidation in heterogeneous compared to electro-catalysis.

Furthermore, we have addressed the partial oxidation of methanol on H-covered Pt electrodes corresponding to a situation of low electrode potentials. The presence of adsorbed hydrogen atoms on Pt(111) is shown to be detrimental for methanol oxidation.



This research has been supported by the German Research Foundation (DFG) through contract GR 1503/21-2. The authors acknowledge the computer time supported by the state of Baden-Württemberg through the bwHPC project and the Germany Research Foundation (DFG) through grant number INST 40/467-1 FUGG.


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© Springer Science+Business Media New York 2017

Authors and Affiliations

  1. 1.Institute of Theoretical ChemistryUlm UniversityUlmGermany

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