First-Principles Mechanistic Analysis of Dimethyl Ether Electro-Oxidation on Monometallic Single-Crystal SurfacesJeffrey A. Herron, Peter Ferrin, and Manos Mavrikakis*

Department of Chemical and Biological Engineering, University of WisconsinMadison, Madison, Wisconsin 53706, United States

*S Supporting Information

ABSTRACT: Dimethyl ether is an attractive alternative to petroleum fuels due to itsphysical properties, comparable energy density to methanol and ethanol, and minimaldeleterious environmental/toxicological effects. For direct fuel cells, it has a number ofadvantages over other prominent fuels, including easier storage with respect to hydrogen,lower toxicity and crossover when compared to methanol, and more facile completeoxidation as compared to ethanol (which includes a relatively difficult to break C−C bond).However, the dimethyl ether electro-oxidation reaction is poorly understood, hindering thedevelopment of improved electrocatalysts. Using periodic, self-consistent (PW91-GGA)density functional theory calculations, we evaluate the thermochemistry of dimethyl ether(DME) electro-oxidation, at the elementary step level, on 12 model, closed-packed facets ofpure transition metals: Au, Ag, Cu, Pt, Pd, Ni, Ir, Rh, Co, Os, Ru, and Re. From thecalculated thermochemistry, we determine the most probable reaction paths on each ofthese surfaces, focusing on Pt as a model system. Our results predict two key electro-oxidation peaks. At lower potentials, there is a peak corresponding to partial oxidation of DME to CO (and other surfacepoisoning species) or complete oxidation to CO2 via formic acid as a key intermediate. A second, higher-potential peak is due tocomplete oxidation of adsorbed CO (and other surface poisoning species) to CO2. Assuming the catalysts remain in theirmetallic state during the DME electro-oxidation process, our results suggest that the onset potential of the surfaces increases inthe order Cu < Ni < Os < Rh < Ir < Co < Ru < Pt < Ag < Pd < Re < Au. Using our results, we construct a theoretical phasediagram showing predicted catalyst activity based on two key reactivity descriptors, the free energies of adsorbed CO and OH.We compare all results to methanol electro-oxidation to understand key mechanistic differences and their impacts on optimalcatalyst design for direct DME fuel cells.

1. INTRODUCTION

In recent years, there has been a drive to find alternatives topetroleum-based fuels that can be used efficiently, transportedeasily, and utilized with reduced environmental impact.Dimethyl ether (DME) is a promising alternative fuel that isphysically similar to liquefied petroleum gases (LPGs), allowingfor easy transport and storage.1 DME is quite environmentallybenign, in terms of both toxicity as well as its global warmingpotential. It is industrially produced over two steps: (1)methanol synthesis from syngas and (2) dehydration ofmethanol to dimethyl ether.1,2 As a fuel, DME has acomparable lower heating value (28.62 kJ/g) to methanol(19.99 kJ/g) and ethanol (26.87 kJ/g).1 Direct fuel cellsutilizing DME as a feed have an equilibrium voltage of 1.20 Vand a high theoretical specific energy of 8377 Wh/kg (forcomparison, methanol has an equilibrium potential of 1.21 Vand theoretical specific energy of 6073 Wh/kg, respectively).3

Furthermore, compared to methanol, the lower dipole momentof DME has been shown to limit DME crossover to thecathode,4 which is a significant problem in methanol fuel cells.5

Due to these reasons, there has been significant interest indeveloping direct fuel cells utilizing dimethyl ether as a fuel.6−14

In order to improve the efficiency3 and activity of direct fuelcells, improved anode electrocatalysts must be developed.

Toward this goal, a fundamental understanding of the reactionmechanism is essential in designing alloy catalysts.15,16

Although there have been some theoretical studies of DMEvapor-phase catalysis,17 there has been a lack of exploration intothe DME electro-oxidation reaction. On the other hand,experimental studies have provided some key mechanisticinsights. Of note, the analysis of reaction byproducts has shedsome light on competing reaction pathways. Tsutsumi et al.18

detected formic acid in the exhaust of a direct polymerelectrolyte membrane fuel cell, in agreement with themechanism proposed by Kerangueven.19 Yet, Muller et al.9

found methanol but no formic acid. Finally, Mizutani and co-workers4 detected both methanol and methyl formate(HCOOCH3), but in negligible quantities.Kerangueven and co-workers19 studied DME electro-

oxidation on polycrystalline Pt electrodes. The authorsidentified two oxidation peaks, which suggested that therewere either multiple active sites or parallel reactionmechanisms. The first peak occurred between 0.3 and 0.65 Vversus the reversible hydrogen electrode (RHE), while the

Received: June 13, 2014Revised: August 19, 2014Published: August 27, 2014

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second peak began around 0.65 V with maximum currentoccurring at 0.78 V. (Note, hereafter we quote all potentialsversus the reversible hydrogen electrode and drop thedistinction for brevity.) They suggested that the first peakcorresponded to incomplete oxidation of DME to either formicacid (HCOOH) or CO. The mechanism postulated involvesDME hydrolysis to hydroxymethyl (CH2OH), followed byoxidation to a CHO intermediate which reacts with water toform HCOOH or is oxidized to CO. The second peak wasattributed to complete oxidation of CO through activation ofwater. Using in situ infrared reflectance spectroscopy theseauthors were able to identify adsorbed CO, adsorbed carboxyl(COOH), and CO2.Lu and co-workers performed extensive studies on Pt

electrodes, uncovering remarkable structure sensitivity for thisreaction.7,8,11−13,20,21 Using Pt(111)7 and (100)8 single crystals,they demonstrated a much higher activity for the (100) facetthan that for the (111) facet. The authors proposed thefollowing reaction mechanism

+ * → * + ++ −CH OCH CH OCH H e3 3 3 2 (1)

* → * + ++ −CH OCH CH OC 2H 2e3 2 3 (2)

* + → * + + ++ −CH OC H O CO CH OH H e3 2 3 (3)

* + * → + * + ++ −CO OH CO 2 H e2 (4)

+ + *

→ * + + ++ −

CH OCH H O

CH O CH OH H e3 3 2

3 3 (5)

* + → + * + ++ −CH O H O CO 5H 5e3 2 2 (6)

where reactions 1 through 4 correspond to indirect oxidationoccurring on both crystal facets and reactions 5 through 6correspond to direct oxidation through a key methoxyintermediate. These latter reactions 5 and 6 were suggestedto occur exclusively on the (100) facet where methoxy ispreferentially stabilized. On the (111) surface, the authorssuggest that DME is initially decomposed to CO, with a peak at0.35 V, and then CO is oxidized between 0.6 and 0.8 V inagreement with the results of Kerangueven and co-workers.19 Liand co-workers, in a combined experimental and theoreticalstudy, concluded that the (100) facet is more active due tofavorable stabilization of a CHOC intermediate, which is theprecursor to C−O bond breaking.22

There have been limited efforts in the design andoptimization of the electrocatalyst for DME electro-oxidation.Kerangueven et al. studied Pt/C, PtRu/C, and PtSn/Celectrocatalysts based on their effectiveness in methanolelectro-oxidation.6 They found that alloying Pt with Ru or Snreduces the onset potential from 0.5 to 0.4 V. Though theelectro-oxidation occurs at lower potentials, the peak currentdensity for these alloys is reduced as compared to pure Pt.Furthermore, though the Ru and Sn provide bifunctionalitythrough activation of water, they do not affect the rate of DMEadsorption. Liu et al. compared the performance of PtMe/Ccatalysts (where Me = Ru, Sn, Mo, Cr, Ni, Co, W) to Pt/C.14

They found that in the low overpotential range (<0.55 V)PtRu/C showed the best activity and CO tolerance, whileabove this potential Pt/C was the best catalyst. Furthermore, ascompared to methanol electro-oxidation, DME electro-oxidation has a much stronger temperature dependence for

the PtRu/C electrocatalyst. Other studies of PtRu catalysts23,24

have demonstrated their superior properties while also showingthat pure Pd is inactive.First-principles methods have become invaluable toward

understanding reaction mechanisms and in the design ofimproved alloy catalysts.15,16,25−32 In particular, the utility ofthese methods has been demonstrated for a number of reactionsystems, including oxygen reduction,30,32 methanol electro-oxidation,33 ethylene epoxidation,34 and acetylene hydro-genation,26 to name a few examples. Toward developingimproved electrocatalysts for dimethyl ether electro-oxidation,we seek to understand the fundamental electro-oxidationmechanism and the energetics of competing elementary stepson monometallic surfaces. We begin our study by examiningthe reactivity trends of model close-packed facets of thefollowing transition metals: Au, Ag, Cu, Pt, Pd, Ni, Ir, Rh, Co,Os, Ru, and Re. From this data, we construct theoreticalelectro-oxidation activity phase diagrams, based on simplescaling relationships, using the free energies of adsorbed COand OH as reactivity descriptors. From the phase diagrams, weidentify regions of the descriptors’ parameter space wherepromising alloy surfaces may be found. Finally, we compare theelectro-oxidation of DME to that of methanol35 to understandfundamental differences in the mechanisms and how these mayimpact catalyst design for these two fuel molecules.

2. METHODSThe free energies of reaction species are calculated using spin-polarized density functional theory, as implemented inDACAPO,25,36 a planewave total energy code. The surfacesare represented using a periodic 3 × 3 unit cell with three layersof metal atoms and at least five equivalent layers of vacuumbetween successive slabs. The metal atoms in the slab are fixedat their optimized bulk positions. Previous calculations haveshown that relaxation effects on these close-packed surfaces areminimal.37−40 Calculations are performed on the (111) facet offace-centered cubic (fcc) metals (here: Rh, Ir, Ni, Pd, Pt, Cu,Ag, and Au) and the (0001) facet of hexagonal close-packed(hcp) metals (here: Re, Ru, Os, and Co). (Note we drop thefacet distinction (111) versus (0001) hereafter for brevity.) Thelattice constants for these metals were optimized usingDACAPO. The optimized bulk lattice constants are (exper-imental value41 in parentheses): Ag 4.14 Å (4.09 Å), Au 4.18 Å(4.08 Å), Co 2.50 Å (2.51 Å), Cu 3.67 Å (3.62 Å), Ir 3.86 Å(3.83 Å), Ni 3.52 Å (3.52 Å), Os 2.74 Å (2.73 Å), Pd 3.99 Å(3.89 Å), Pt 4.00 Å (3.92 Å), Re 2.76 Å (2.76 Å), Rh 3.83 Å(3.80 Å), and Ru 2.74 Å (2.71 Å). For Co, Os, Ru, and Re, weused a c/a ratio of 1.63. Adsorption is permitted on only one ofthe two exposed surfaces of each slab, and the dipole moment isadjusted accordingly.42,43 The ionic cores are described usingultrasoft Vanderbilt pseudopotentials.44 The Kohn−Sham one-electron states are expanded in a series of plane waves with anenergy cutoff of 25 Ry. The surface Brillouin zone is sampledwith 18 special Chadi−Cohen45 k-points based on convergenceof the total energy. The exchange-correlation energy andpotential are calculated self-consistently using the PW91generalized-gradient approximation.46 The electron density isdetermined by iterative diagonalization of the Kohn−ShamHamiltonian, Fermi population of the Kohn−Sham states (kBT= 0.1 eV), and the Pulay mixing of the resulting electronicdensity. All total energies are extrapolated to kBT = 0 eV.The vibrational frequencies of all species are calculated by

numerical differentiation of forces using a second-order finite

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difference approach with a step size of 0.015 Å.47 The Hessianmatrix is mass-weighted and diagonalized to yield thefrequencies and normal modes of the adsorbed species. Forsurface adsorbates, the vibrational frequencies were onlyevaluated on Pt(111) and were assumed (as a firstapproximation) to be the same on other surfaces due tosimilar adsorption geometries. The entropy of all surfacespecies is calculated including translation, vibration, androtation. The zero-point energy (ZPE) is included in thecalculation of the free energy for all surface adsorbates. Thezero-point energies are calculated assuming a quantumharmonic oscillator possessing the calculated vibrationalfrequencies.Calculations of free energy of gas-phase species are done

relative to DFT-derived energies of H2O(g), CO2(g), andH2(g). For example, the free energy of dimethyl ether iscalculated from the reaction: 2CO2(g) + 6H2(g) → 3H2O(l) +CH3OCH3(g) which leads to

Δ = − +

+ − + − −

+ − − +

G E TS Z

E TS Z E TS

Z E TS Z

( )

3( ) 2(

) 6( )

CH OCH CH OCH CH OCH CH OCH

H O H O H O CO CO

CO H H H

3 3 3 3 3 3 3 3

2 2 2 2 2

2 2 2 2

where Ex is the total energy of species x calculated from DFT; Tis the standard temperature (298 K); Sx is the calculatedentropy for species x, and Zx is the calculated zero-point energyfor species x. The calculation of free energies of other species isdone similarly. When calculating the free energy of surfacespecies, the total energy of the species is taken relative to theclean surface, while the entropy and zero-point energy arecalculated for the adsorbed state. Using this approach, wecalculate free energies of formation for gas-phase CH3OCH3,CH3OH, and HCOOH as 0.39, 0.11, and 0.44 eV, respectively.The corresponding experimental values are −0.09, 0.04, and0.34 eV, respectively.41 We note that there are well-knowndifficulties for the GGA-PW91 functional for describing thesegas-phase energetics, though we do not expect thesediscrepancies to affect the overall results of our analysis.To correct the calculated free energy for the electrochemical

potential, we employ the computational hydrogen electrodeapproach, which has previously been applied to studies ofoxygen evolution,48,49 oxygen reduction,50,51 and methanoloxidation37,38 at electrode surfaces. First, we choose thereversible hydrogen electrode as a reference. At standardconditions, hydrogen gas is in equilibrium with protons andelectrons, at a defined potential of 0 V. A change in theelectrode potential by U will shift the free energy of theelectrons exergonically by |e|U, where |e| is the absolute chargeof an electron. Therefore, if we consider the first dehydrogen-ation step in dimethyl ether, CH3OCH3(g) + * → CH2OCH3*+ H+ + e−, a change in the potential by U will adjust the freeenergy change of this reaction by −|e|U.Activation energy barriers are not calculated for any of the

reaction steps studied. Because previous studies of Brønsted−Evans−Polanyi relationships suggest a strong correlationbetween the thermochemistry of reaction steps and theirkinetics, we rely on the reaction thermochemistry alone in ouranalysis.52,53 The existence of C−O bond breaking events doescomplicate the analysis though as these steps are notelectrochemical in nature, and thus the barriers are not afunction of the electrode potential. Furthermore, thethermochemistry is calculated at 1/9 ML surface coverage.

The surface coverage under electrochemical oxidation con-ditions may be different, and this could contribute todifferences between experiments and the present calcula-tions.47,54−58 Also, we neglect any contributions to bindingenergies from solvation effects.59−62 We note that past DFTstudies have calculated that solvation stabilizes OH, which isadsorbed on Pt(111), by ∼0.5 eV, and this may affect theoverall conclusions.51,63,64 However, it is likely that otheradsorbates in the reaction, perhaps especially those with an−OH group, may be stabilized to a similar extent. Therefore, aswe (or others in the literature) have not performed systematicstudies of solvation on all of the adsorbates, we omit thoseeffects from our analysis so as not to unfairly bias the results.Keeping in mind these limitations to our analysis, we proceedwith presenting its results.

3. RESULTS AND DISCUSSIONThe overall anode, electro-oxidation reaction is CH3OCH3(g)+ 3H2O(l) → 2CO2(g) + 12(H+ + e−). Evaluation of theelectro-oxidation mechanism is a complicated task. Inconstructing the reaction network, we consider proton−electron transfer and C−O bond breaking steps. Further, weconsider Heyrovsky-type reactions65 of the form: X + H2O(l)→ X−OH + (H+ + e−). We consider only direct (non-electrochemical) C−O bond breaking events, which do notinvolve the simultaneous release of a proton/electron. There-fore, we exclude from our study concerted steps involving C−Obond breaking and a Heyrovsky-type of reaction: an example ofsuch a reaction is CH2OCH3* + H2O → CH2O* + CH3OH +(H+ + e−). In our study, we treat this as two separate reactionsteps, C−O bond breaking and a Heyrovsky step (withinterchangeable order of events). We note that C−H or O−Hbond breaking could also occur via nonelectrochemical means.However, though these processes might be competitive withthe electrochemical processes at low (or negative) potentials, athigher potentials (i.e., at the calculated onset potential) theywill not be competitive. Therefore, we exclude nonelectro-chemical C−H and O−H bond breaking reactions from ouranalysis. The reaction network studied, shown in Scheme 1,includes 29 species providing an enormous set of possiblepathways for complete oxidation.In order to deconstruct the reaction network into tractable

pieces, we attempt to analyze the pathways in a systematic way.First, we determine a reaction pathway that follows the mostexergonic transformation possible at each step along thereaction coordinate (in the forward direction). Following thispathway will serve to identify species that are possiblypoisoning the surface and the difficulties in their removal.Second, we compare these results to the “lowest-free-energyreaction path”, as described in previous studies of methanolelectro-oxidation.37,38 This pathway is found by first consider-ing all possible pathways, j, and calculating the free energychange of each electrochemical elementary step, i, along thatpathway (ΔGij). From each pathway, we find the mostenergetically demanding step (i.e., maxi ΔGij). The lowest-free-energy reaction path is the pathway which minimizes thedifficulty of the most energetically demanding step (i.e., minj(maxi ΔGij)). This can be considered a pathway which adheresto the Sabatier principle, therefore having the minimum “mostendergonic” reaction step of all possible pathways. As theremoval of adsorbed surface poisoning species is usually highlyendergonic, this pathway, in practice, avoids producing these. Inthese two pathway analyses, we omit the possibility of

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Heyrovsky reactions occurring prior to C−O bond breaking.Finally, we briefly comment on additional pathways whichmanifest when the C−O−C backbone is oxygenated viaHeyrovsky reactions prior to breaking the C−O bond.Experimental evidence4 showing methyl formate as a byproducton Pt suggests that these reactions may be important towardthe overall observed reactivity. The additional intermediatesstudied within this final pathway are not included in the originalreaction network diagram (Scheme 1).In the following analysis, we first consider Pt, in detail, due to

its ubiquitous usage in fuel-cell electrocatalysis and thenconsider differences observed on the 11 other late transitionmetals studied: Au, Ag, Cu, Pd, Ni, Ir, Rh, Co, Os, Ru, and Re.3.1. Most Exergonic Pathway. For the most exergonic

pathway, we follow the reaction pathway that arises if the mostexergonic reactions occur consecutively. In general, the resultsof this analysis are a function of the cell potential, so we beginassuming that the potential is 0 V. First we analyze thepathways that break the DME C−O−C backbone through C−O bond scission (see Scheme 2 for energetics on Pt(111)).3.1.1. Pathways to C−O Bond Scission. The electro-

oxidation of DME (CH3OCH3) on Pt(111) begins with eitherproton/electron removal or C−O bond breaking. On Pt,proton/electron transfer (to yield CH2OCH3) is mildlyendergonic by 0.20 eV, whereas C−O bond breaking (toCH3 and OCH3) is endergonic by 0.50 eV. The difficulty of C−O bond breaking is due to the instability of the methoxy(OCH3) intermediate that would be formed. Since proton/electron transfer to yield CH2OCH3 is less endergonic than C−O bond breaking, the first reaction will be proton/electrontransfer forming CH2OCH3. Studies by Osawa and co-workers7,8 suggested that the higher activity of Pt(100) versusPt(111) is due to the activation of a direct mechanism featuringa stabilized methoxy intermediate. The results of our study on

Pt(111) provide evidence that pathways featuring the formationof methoxy are indeed unfavorable on that surface.On Pt, the second step in the reaction can be proton/

electron removal from either the CH2 or CH3 functional groupsor C−O bond breaking. Removal of a hydrogen from the CH2group to form CHOCH3 is endergonic by 0.25 eV, whileremoving a hydrogen from CH3 is exergonic by 0.03 eV. C−Obond breaking can lead to the formation of two sets ofintermediates: formaldehyde (CH2O) and methyl (CH3) or toCH2 and OCH3. The former reaction is exergonic by 0.13 eV,and the latter is endergonic by 0.64 eV. Since C−O bondbreaking to form CH2O and CH3 is the most exergonic stepamong all possible options at this stage, we choose this as thesecond step in the pathway.On the basis of these results, it seems likely that the C−O

bond breaking will occur early in the mechanism on Pt(111).However, mechanisms that have been suggested based onexperimental findings7 feature further dehydrogenation of theC−O−C backbone to COCH3 prior to C−O bond breaking.Our calculations at 0 V show that while COCH3 is more stablerelative to CH3OCH3, intermediate transformations betweenthese two species are more difficult than C−O bond breaking.Therefore, it is unlikely that COCH3 is an active intermediateat 0 V. However, if the operating potential is increased,selectivity toward proton/electron transfer steps would beincreased relative to C−O bond breaking because the ΔG ofthe reaction for the electrochemical steps (i.e., non C−O bondbreaking) is a function of the potential.If we return to the second step in the reaction, C−O bond

breaking is exergonic by 0.13 eV, while proton/electron transferof CH2OCH3 to CH2OCH2 is exergonic by 0.03 eV. If weoperate at a potential of 0.35 V, the potential that Lu and co-workers7 (and within the range suggested by Kerangueven19)found the formation of CO, the proton/electron transfer wouldbe exergonic by 0.38 eV, a much more favorable reaction thanC−O bond breaking. (Note, in Scheme 3 we show theenergetics at 0.35 V on Pt(111).) The subsequent proton/electron extraction to form CHOCH2 is exergonic by 0.57 eV(at 0.35 V), while C−O bond breaking to form formaldehydeand CH2 is endergonic by 0.24 eV. Then, further extractionleads to either COCH2 (endergonic by 0.31 eV) or CHOCH(endergonic by 0.37 eV). Breaking the C−O bond in CHOCH2forms either CH and formaldehyde (endergonic by 0.15 eV) or

Scheme 1. Reaction Network for DME Electro-Oxidationa

a(a) Proton/electron transfer from the COC backbone and C−Obond breaking. (b) Electro-oxidation of DME C−O bond breakingintermediates. Proton/electron transfer steps shown with black arrows;C−O bond breaking steps shown with blue arrows; and Heyrovskyreactions shown with dotted red arrows. All species are adsorbed onthe catalyst surface except for CH3OCH3, CH3OH, HCOOH, CH2O,and CO2, which are gas-phase species.

Scheme 2. Free Energies of Reaction for Proton/ElectronTransfer Steps and C−O Bond Scission of the DMEBackbone on Pt(111) at 0 Va

aIn black font and black arrows are proton/electron transfer steps, andin blue font and blue arrows are C−O bond scission steps. For theequivalent results on the additional 11 metals studied, see SI. Thereaction energetics is calculated with all species adsorbed on thecatalyst surface except for CH3OCH3 and CH2O, which are gas-phasespecies. All reaction energies are given in eV.

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CHO and CH2 (endergonic by 0.12 eV). This would suggestthat at 0.35 V we would anticipate the formation of CHOCH2,which would then dissociate to CHO and CH2. However, thereis an additional possibility which has not been considered:CHOCH2 can isomerize to COCH3. Isomerization is exergonicby 0.20 eV, which is significantly favored relative to eitherproton/electron transfer or C−O bond breaking. Therefore, at0.35 V, there exists a thermodynamically favorable pathway toform COCH3 from DME, supporting the experimentallypostulated mechanism.Theoretically, the formation of COCH3 becomes preferred

to dissociation of CH2OCH3 at potentials above 0.10 V. Abovethis potential, it becomes more favorable for CH2OCH3 todehydrogenate to CH2OCH2 rather than dissociate to CH2O +CH3 (both reactions are exergonic by 0.13 eV at 0.10 V).However, it is not until 0.20 V that the initial proton/electronextraction (to form CH2OCH3 from DME) becomes exergonic.Therefore, we would expect formation of COCH3 (andtherefore CO through direct C−O bond breaking) to beginabove 0.20 V, somewhat lower than the experimental findings.If we increase the potential even higher than 0.35 V, there willnot be significant changes in the reaction pathway. At very highpotentials (>0.86 V) formation of COCH2 from CHOCH2(ΔG < −0.20 eV) will be competitive with isomerization ofCHOCH2 to COCH3 (ΔG = −0.20 eV), but this intermediate(COCH2) will be quickly dissociated to CO and CH2. Asummary of the most exergonic reaction pathways leading toC−O bond breaking on Pt at 0 and 0.35 V is shown in Scheme4.Comparing the results on Pt with the other 11 metals studied

reveals some interesting differences. Results, equivalent toScheme 2, for these metals are provided in the SI. At 0 V, Ag,Cu, Ni, Ir, Rh, Co, Os, Ru, and Re favor direct dissociation of

DME to CH3 and OCH3, while Au and Pd (like Pt) favorproton/electron extraction to CH2OCH3. On Au and Pd, thesubsequent favored reaction is C−O bond scission, similarly toPt. For the rest of the metals studied to favor initial proton/electron extraction over C−O bond scission, the operatingpotential must be increased. The required potentials for thesemetals are Ag (0.05 V), Cu (0.70 V), Ni (0.98 V), Ir (0.12 V),Rh (0.56 V), Co (1.26 V), Os (0.65 V), Ru (0.99 V), and Re(1.63 V). In Scheme 5, we show the thermodynamically favored

pathway at 0 V and at two other characteristic operatingpotentials of Pt, 0.35 and 0.60 V (experimental7,19 partialoxidation and complete oxidation potentials on Pt, respec-tively). The results show that as the potential increases C−Obond scission becomes delayed as proton/electron transfersteps become more favorable. As the potential continues toincrease, all metals studied eventually form COCH3 which willeither be dissociated to CO and CH3 (at moderate potentials)or further dehydrogenated (at very high potentials). In general,CO and CH3 are very stable intermediates, and dissociation toCO and CH3 will be favored except at unrealistically highpotentials.

3.1.2. Oxidation of C−O Scission Products. In the previoussection, we showed that changing the operating potential willalter the products of C−O bond scission. If we again considerScheme 1, the operating potential will change the COC-typeintermediate broken (part a of the scheme) and will thereforeaffect “where” in part b of the scheme we will “begin” oursubsequent analysis. This can have some consequences for theoverall reaction mechanism. For a particular set of COC-typeintermediates broken, only a subset of the reactions in theoverall network will be relevant. Therefore, one might postulatethat for a particular catalyst there exists an operating potential(or range of potentials) that selectively dissociates a specificCOC-type species such that surface poisons do not formdownstream. To try and answer this hypothesis, we continueour analysis of electro-oxidation on Pt and compare it to therest of the metals studied.The analysis of the remainder of the electro-oxidation

mechanism is somewhat complicated by the fact that we aretracking two separate fragments. When the C−O bond in DMEis cleaved, we are left with an oxygenated product (CHxO) anda hydrocarbon product (CHy). For any given fragment, we havethe possibility of two separate reactions: we can activate watervia a Heyrovsky-type reaction, or we can remove a proton/electron pair bound to either O or C. For Pt, at 0 V we expect

Scheme 3. Free Energies of Reaction for Proton/ElectronTransfer Steps and C−O Bond Scission of the DMEBackbone on Pt(111) at 0.35 Va

aIn black font and black arrows are proton/electron transfer steps, andin blue font and blue arrows are C−O bond scission steps.Isomerization of CHOCH2 to COCH3 is shown in pink. The reactionenergetics are calculated with all species adsorbed on the catalystsurface except for CH3OCH3 and CH2O, which are gas-phase species.All reaction energies are given in eV.

Scheme 4. C−O Bond Breaking Pathways for DME onPt(111) at Different Operating Potentialsa

aValues shown are free energies of reaction (eV); stoichiometry isbalanced with H2O, H

+, and e−. The reaction energetics are calculatedwith all species adsorbed on the catalyst surface except for CH3OCH3and CH2O, which are gas-phase species.

Scheme 5. Pathways Leading to Breaking of the C−O Bondon All Metals Studied at Selected Operating Potentials (0.00,0.35, and 0.60 V)a

aStoichiometry is balanced with H2O, H+, and e−.

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CH2O and CH3 as C−O bond scission products, while above0.10 V it is favorable to form CO and CH3 via COCH3

dissociation. In Scheme 6 we present the entire possibility oftransformations for C−O bond scission products on Pt.

Since CH3 is a common dissociation product at both 0 and0.35 V, we will start with the analysis of its oxidation. Note,here we quote all thermochemistry at 0 V. For CH3, we canactivate water to form methanol or remove a proton/electronpair to form CH2. The former reaction is endergonic by 0.36eV, while the latter is endergonic by 0.34 eV. Since CH2formation is favored, we assume that CH2 is the dominantproduct. From CH2 we can form CH2OH, which is endergonicby 0.33 eV, or CH, which is exergonic by 0.31 eV. CH is clearlyfavored and can be removed by forming carbon, which is 0.96eV endergonic, or formation of CHOH which is endergonic by0.74 eV. CHOH is the preferred product, yet neither of thesereactions is facile. Therefore, we might expect poisoning byCH. CHOH reacts to form either CHO, exergonic by 0.15 eV,or COH, exergonic by 0.46 eV. COH is preferred, which thenreleases a proton/electron pair to form strongly bound CO.CO is difficult to remove, and the preferred method is by

Heyrovsky-type reaction to form COOH, which is endergonicby 0.81 eV. This is the most difficult step within the describedreaction mechanism. Finally, COOH is oxidized to CO2 whichis 0.39 eV exergonic. The prediction of COOH as a keyintermediate is in agreement with an in situ IR reflectancespectroscopy19 study which observed a stretch at 1700 cm−1,which was assigned to the C−O stretch in COOH or CHO.We have calculated the C−O stretch in COOH at 1715 cm−1

and the C−O stretch in CHO at 1786 cm−1.The identity of the other C−O bond scission product varies

with operating potential (CH3 fragment is the preferred CHyspecies, in general). Assuming low potential, the relevantintermediate species is CH2O (formaldehyde). Formaldehydecan react to form an H2COOH intermediate or lose a proton/electron pair from its carbon atom. The former species is quiteunstable requiring 0.77 eV to form, while CHO is more stablethan formaldehyde with an exergonic reaction energy, −0.34eV. This CHO intermediate was suggested in the mechanismproposed by Kerangueven.19 Then, CHO is oxidized to CO,which is quite exergonic, −0.80 eV. As the operating potentialincreases (to realistic operational levels), we have shown thatthe DME C−O bond is cleaved later in the mechanism, leadingto direct formation of CO, bypassing formaldehyde and CHO.This may explain why formaldehyde is not detectedexperimentally.If we perform the same analysis for the remainder of the

metals studied, we uncover some differences. (Note, in theSupporting Information we provide equivalent energetics asshown in Scheme 6 for the other metals studied.) Assumingoperation at low potential, by tracking the oxygenated fragmentmade by the C−O bond scission event, we see that every metalfeatures formaldehyde as an active intermediate. Formaldehydeis oxidized to an H2COOH intermediate for Ag and Cu,whereas it is dehydrogenated to CHO for the remaining metals.Au, Ag, and Cu then form HCOOH which dehydrogenates toformate (HCOO) and then CO2. For the metals promotingCHO formation, CHO is dehydrogenated to CO, which reactswith an adsorbed hydroxyl to form CO2 directly (Pd, Ni, Rh,Co, Os, Ru, and Re) or is oxidized to stable carboxyl (Pt, Ir)before forming CO2. Tracking the hydrocarbon fragment madeby the C−O bond scission event, we find that the morestrongly binding metals, Pt, Ir, Rh, Co, Os, Ru, and Re, preferdehydrogenation of CH3 to CH (Pt, Ir, and Os) or C (Rh, Co,Ru, and Re) followed by oxidation to CHOH or COH,

Scheme 6. Electro-Oxidation Network for C−O BondScission Products, Produced from DME Electro-Oxidationon Pt(111)a

aValues presented are free energies of reaction (eV) at 0 V.Stoichiometry is balanced with H2O, H+, and e−. In dotted redarrows are Heyrovsky-type reactions, and in black arrows are proton/electron transfer reactions. In pink arrows, the most exergonic pathwayis outlined. Similar data for the rest of the metals studied are providedin SI. The reaction energetics is calculated with all species adsorbed onthe catalyst surface except for CH3OH, HCOOH, CH2O, and CO2,which are gas-phase species.

Figure 1. Onset potential for DME electro-oxidation on a variety of close-packed monometallic surfaces through the (a) most exergonic pathwayand (b) lowest-free-energy pathway. Color code above the graphs gives the most difficult electrochemical transformations within the reactionmechanisms studied.

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respectively. CHOH dehydrogenates to COH, which dehydro-genates to form CO. For Au, Ag, Cu, Pd, and Ni, CH3 isoxidized to methanol which will then dehydrogenate to eitherCH2OH (Au and Pd) or methoxy (Ag, Cu, and Ni). Eitherway, both paths converge on formaldehyde which reacts asdiscussed previously.3.1.3. Onset Potentials for Most Exergonic Pathways. In

analyzing the pathways (pre and post C−O bond scission) forPt (at 0 V), we find that the most difficult electrochemicalreaction (following the most exergonic route) is CO removalthrough formation of carboxyl. On Pt, this reaction is 0.81 eVendergonic. For this reaction to become thermoneutral, thepotential must be raised to 0.81 V. We refer to this potential(the point at which all free energies of reaction are nonpositive)as the onset potential for the specific metal surface. (Note, sinceC−O bond scission events are not potential dependent, weexclude the energetics of these steps from this analysis, thoughthe pathway may depend on their energetics.) This calculatedonset potential is on the high end of the experimentallyobserved range of approximately 0.6−0.8 eV.7,19

The onset potential is a quantity we can use to judge theefficacy of each catalyst. Catalysts with low onset potentialshave reaction coordinates which do not contain highlyendergonic transformations; i.e., they are not poisoned byreactive species. Furthermore, a low onset potential maximizesthe chemical energy available for producing useful work,improving the energy efficiency of the fuel cell. We calculate theonset potential for each of the surfaces studied and describe themost difficult proton/electron transfer step within theirreaction network that gives rise to this potential. The resultsof this analysis are shown in Figure 1a.If we operate at a potential more positive than 0 V, the C−O

bond scission products formed may differ, thereby requiring aniterative approach to converge the onset potential. For Pt, byanalyzing the entirety of Scheme 6, no matter the initial C−Obond scission products, the reaction will be channeled throughCO as a necessary intermediate. Therefore, there does not existan optimal potential, for Pt, which provides favorable C−Oscission products which can avoid the formation of CO duringcomplete electro-oxidation. Of the other metals studied, onlyNi, Re, Ru, Co, and Os might benefit from different C−O bondscission products, as the most difficult proton/electronextraction is from OCH3. However, when observing theenergetics on these metals, the potential required to delay theC−O bond breaking event (thereby avoiding direct OCH3formation) is much higher than the potential required tofacilitate proton/electron extraction from OCH3. Thus, for allmetal surfaces studied, there is not an optimal potential; ratherthe reaction rate should monotonically increase with appliedpotential.The results shown in Figure 1a suggest that we can rank the

metals in terms of their onset potential, from lowest to highest(lowest being the best): Cu < Ni < Os < Rh < Ir < Co < Ru <Pt < Ag < Pd < Re < Au. Importantly, these rankings assumethat the active surface remains in the metallic state and do not,at all, account for the stability of these materials under workingconditions. Furthermore, since these rankings are based on amechanism that follows the most exergonic transformations,these rankings can generally serve to illustrate the resistance ofeach catalyst to surface poisoning. The high onset potential forAu and Ag, however, is required to activate. Though, for Au andAg, the high onset potential is required to activate closed-shelled species, not for removal of surface poisons. Finally,

these rankings only include proton/electron extraction eventsand Heyvrosky-type steps; slow C−O bond scission kineticsmay alter the overall observed reactivity.Observing the potential-determining steps provides some

interesting information that can help in the design of better,alloy catalysts. Of the steps, we see that CO removal is the mostdifficult reaction for only Pt and Ir (also Pd, as the mostdifficult step is H2O activation, which is required for COremoval). CO poisoning is a common problem in methanolelectro-oxidation, but here we see a new poison challengingDME electro-oxidation: CH removal is the most difficultreaction step for Rh. These two species, CO and CH, can beconsidered poisons to the catalyst as a consequence of strongaffinity for C-binding. For metals which have a strong affinityfor O-binding (Re, Ru, Os, Co, and Ni), OCH3 removalbecomes the most difficult reaction step. For this argument, wenote that the binding energies of C and O on the surfacesstudied are not strongly correlated (see Supporting Informa-tion). The remaining metals, Au, Ag, Cu, and Pd, have the mostdifficult steps that correspond to activation of stable, closed-shell species, i.e., methanol, formaldehyde, and water (seeFigure 1a). Therefore, we can conclude that these metals are“on the other side of the volcano”, binding species too weaklyto activate the reaction.

3.2. Lowest-Free-Energy Pathways. Next we sought toinvestigate the “lowest-free-energy pathways”.37,38 These path-ways are constructed by exploring all possible pathways andfinding the pathway with the least difficult, most difficult step(when considering only electrochemical elementary steps). Asconstructed, the onset potential for this pathway is theminimum possible. (Note, based on this formulation, aparticular metal surface can have multiple possible pathwayswhich have the same most difficult step, i.e., the same onsetpotential. In those cases, we present the pathway which followsmost exergonic steps for the ambiguous sections. See theSupporting Information for more details.) At start up, and lowpotentials, these pathways may lead to complete oxidation ofDME to CO2. However, over time, surface poisons willaccumulate on the surface, thereby reducing activity. Here, wepresent the lowest-free-energy pathway for Pt(111) and thendiscuss trends across metal surfaces.In Scheme 7 we illustrate (one of) the lowest-free-energy

pathway(s) on Pt(111). The most difficult electrochemical stepis C(OH)2 → COOH + H+ + e− (ΔG = 0.25 eV), and thecorresponding onset potential is 0.25 V. This is a directoxidation pathway since CO is not formed. Furthermore, CH isavoided. In reality, these surface poisons would form throughthe most exergonic pathways, which operate in parallel to thelowest-free-energy mechanism. Therefore, they will eventuallylimit catalytic activity.We can verify that this is the lowest-free-energy path using a

simple argument. Upon C−O bond scission of CHxOCHy, CHxand OCHy are formed. For CHx, if CH3 were formed, then itcan be oxidized to either CH3OH (ΔG = 0.36 eV) or CH2 (ΔG= 0.34 eV). Both of these are more difficult than the step wehave found (see paragraph above). If CH2 were formed, it iseasiest to remove via a Heyrovsky-type reaction, formingCH2OH (ΔG = 0.33 eV). Again, this is more difficult than thepathway illustrated in Scheme 7. Finally, if CH were formed,then it is very difficult to oxidize, requiring a potential of 0.74 Vto form CHOH. Therefore, for the CHx fragment, the lowest-free-energy pathway exists for C, as the C−O bond scissionproduct. The oxidation pathway of one possible OCHx

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fragment, OCH3, has its most difficult step being CHO →HCOOH (ΔG = 0.05 eV). So, there exists at least one pathwaywhich is easier than the most difficult oxidation for the CHxfragment. Two other possible OCHx fragments (i.e., CH2O andCHO) can be oxidized in the same pathway, with the samemost difficult step. If, however, CO is a fragment, it is quitedifficult to remove. So, based on this analysis, the lowest-free-energy pathway to oxidize DME to CO2 must have C andOCHx as C−O bond scission products (x = 1, 2, or 3).However, we need to also consider the potential to form such aCOCHx species. By inspection, we can see that COCH3 is theCOCHx species which can be formed at the most negativepotential. The most difficult step in its formation is CH2OCH3→ CHOCH3 (ΔG = 0.25 eV), the same difficulty as Coxidation to CO2. If there did not exist such a pathway, wewould need to return to our analysis of the CHx fragments tofind the next preferred pathway. (We note that we alreadyexplored an alternative pathway via isomerization of CHOCH2to COCH3, which presents a slightly easier pathway, though itwill not change the overall most difficult transformation in theentire pathway).This calculated potential of 0.25 V is close to experimentally

found oxidation peaks. Kerangueven19 found an initialoxidation peak beginning around 0.3 V, which was proposedas partial oxidation of DME to CO or HCOOH. Similarly, Lu7

found CO formation at 0.35 V.7 We showed, previously, thatCO can be produced by breaking of COCH3 at potentials aslow as 0.10 V, though the onset is at 0.20 V. However, it is notuntil the potential is raised to 0.25 V when it does becomefavorable to oxidize the C−O bond scission fragments (i.e., C)to CO2. Furthermore, HCOOH can be formed at this potentialvia the other C−O scission fragment (i.e., OCH3). AnyHCOOH that forms should readily oxidize to CO2 at thispotential. This means that at 0.25 V we would expect both

partial oxidation of DME to CO and complete oxidation toCO2 (through HCOOH as a key reactive intermediate).Experimental studies of formic acid electro-oxidation on Ptreport steady increases in current density with appliedpotential, with onset as low as 0.1 V,66 supporting our findings.We caution that this analysis is based only on the

electrochemical steps, without accounting for difficulties inC−O bond scission. This is especially important given themuch more endergonic C−O bond scission energy for COCH3→ C + OCH3 (ΔG = 1.73 eV) versus COCH3 → CO + CH3(ΔG = −0.81 eV). If the latter reaction were to occur instead,then the onset potential for CH3 oxidation to CO2 is 0.34 V(via HCOOH), quite close to the experimental partialoxidation peak.

3.2.1. Onset Potentials for Lowest-Free-Energy Pathwaysand Comparisons to Most Exergonic Pathways. Thecalculated onset potentials for the lowest-free-energy pathwaysfor the other metal surfaces studied are shown in Figure 1b.(Note, the pathways for other surfaces are provided in theSupporting Information). We see that for Ag, Cu, and Ni thereis no change in the predicted onset potential, and the calculatedpathways are the same as in the most exergonic formulation ofthe mechanism. As a result, we would expect to only see asingle oxidation peak on these surfaces, corresponding tocomplete DME electro-oxidation. For the remaining metals, thepathways are quite different from those previously discussed,leading to reduced onset potentials. In particular, thehydrocarbon fragment of C−O bond scission is C (on Pt, Ir,Co, Os, Ru, and Re), CH2 (on Au), and CH3 (on Ag, Cu, Pd,Ni, and Rh), while all surfaces featured CH3 as an intermediatein the most exergonic pathway. For Pt, Ir, Os, and Ru, C isoxidized to COH, C(OH)2, COOH, and then CO2, while onCo and Re, CO is still an active intermediate. The CH2 andCH3 species are oxidized via Heyrovsky-type steps to CH2OHand CH3OH, respectively, rather than dehydrogenation to CH(which poisons the surface). The OCHx fragment of C−Obond scission is oxidized through HCOOH on Au, Ag, Cu, Pt,Pd, Ir, Rh, Os, and Ru and through CO on Ni, Co, and Re.A comparison between the calculated onset potentials for

these two approaches (most exergonic, lowest-free-energychange) can lead to predictions for experimental studies. ForAg, Cu, and Ni, where both pathways have the same onsetpotential, we would expect only a single oxidation peak,corresponding to complete oxidation of DME. For theremaining metals, the onset potential is reduced in this lattermechanistic analysis, which may manifest as a low potential,transient, oxidation peak. The peak at lower potentials wouldcorrespond to partial oxidation to CO, CH, OCH3, orHCOOH and complete oxidation of HCOOH to CO2. Thesecond, higher-potential, peak would be due to the removal ofsurface poisoning species: CO (Pt, Pd, and Ir), CH (Rh), andOCH3 (Co, Os, Ru, and Re) which is significantly moredifficult.

3.3. Oxidation of the DME Backbone Prior to C−OBond Breaking. Experimental results have shown theexistence of methyl formate (in negligible quantities, however)on Pt catalysts during DME electro-oxidation.4 To rationalizethis result, we investigated the thermochemistry of intermedi-ates which can be derived by oxygenating the DME backboneprior to breaking the C−O bond. Our results show that whilemethyl formate is relatively stable (the calculated free energy ofgas-phase methyl formate is 0.07 eV less stable than gas-phase

Scheme 7. Lowest-Free-Energy Pathway on Pt(111)a

aValues presented are free energies of reaction (eV) at 0 V.Stoichiometry is balanced with H2O, H+, and e−. In dotted redarrows are Heyrovsky-type reactions; in black arrows are proton/electron transfer reactions; and C−O bond breaking steps are shownwith blue arrows. In green arrows, the lowest-free-energy pathway isoutlined. The reaction energetics is calculated with all species adsorbedon the catalyst surface except for CH3OCH3, CH3OH, HCOOH,CH2O, and CO2, which are gas-phase species.

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DME) the pathways leading to its formation are energeticallyunfavorable (see Scheme 8 for results on Pt).

For Pt, at 0.35 V, all elementary steps leading to theformation of methyl formate are exergonic (along at least onepathway), but these elementary steps are not the mostexergonically selective. Methyl formate may form via wateractivation/OH addition to CH2OCH3, forming an HO-CH2OCH3 intermediate, but this is significantly less favorable(by 0.38 eV) than dehydrogenation to CH2OCH2. Therefore,at operating potentials we might expect some formation ofmethyl formate via these routes, but the selectivity should below.For other metals studied, there are some differences in

behavior. For Au, Ag, and Cu, which, in general, bind speciesmore weakly than Pt, formation of HOCH2OCH3 is favorablerelative to other elementary steps, and subsequent trans-formations leading to methyl formate are also selective. Thisdoes not mean that we would expect high activity towardmethyl formate formation, as the energetics are still quiteunfavorable (in comparison to other catalysts), but rather thatthese pathways are more favorable than the alternatives onthese surfaces. Pd, which binds species similarly to Pt, isselective toward methyl formate. More reactive metals (Ni, Ir,Rh, Co, Os, Ru, and Re) are unselective toward methyl formatedue to selective C−O bond breaking (Ni, Co, Os, Ru, and Re)or selectivity toward direct dehydrogenation to COCH3 (Ir andRh). Overall, based on the results of this study, the surfaceswith the greatest probability of forming methyl formate are Pd,Pt, and Cu.If these oxygenated species were to form, would they

contribute to any differences in the observed activity? Likelynot. If the species is only oxygenated once, when the C−Obond is broken it will form either a OCO-type species and a C-type species or two CO-type species. For example, methylformate (HCOOCH3) will dissociate into HCOO and CH3 orHCO and OCH3. If the OCO species (i.e., HCOO) were toform, it will be oxidized readily to CO2, but the remaining Cspecies will have the same difficulties as previously described.Similarly, if two CO-type species are formed, the reaction willproceed as described before. There may be some benefit byreducing CH poisoning, but catalysts which bind stronglyenough to suffer from poisoning will not be active towardforming these oxygenated species (selective toward early C−Obond breaking). If we attempt to add additional OH/O groupsto the C−O−C backbone, the species obtained are veryunstable.

In general, pathways which may lead to the formation ofmethyl formate do not provide significant thermochemicaladvantages. Surfaces which bind strongly are selective towardC−O bond breaking, never forming these oxygenated species.Catalysts which are weakly binding may be selective towardthese species, but the energetics are too unfavorable overall.

3.4. Optimal Catalyst Design/Comparisons withMethanol. To determine the optimal binding characteristicsfor a DME electro-oxidation catalyst, we construct a theoreticalphase diagram showing the free energy of reaction for the mostdifficult reaction (including proton/electron transfer and C−Obond breaking) by following the most-exergonic reactionpathway as a function of two reactivity descriptors, the freeenergies of adsorbed OH and CO. We choose to examine onlythe most exergonic pathway as this analysis seems to bestrepresent observed experimental behavior.Utilizing our database of DFT-calculated free energies, we

write the free energy of adsorbed reaction intermediates as alinear function of these two descriptors depending on if theintermediate binds through oxygen (linear in OH) or carbon(linear in CO).67,68 This procedure has been applied previouslyto methanol electro-oxidation catalysis.37,38 (The linearcorrelations are provided in the Supporting Information.)Thus, we can envision a catalyst with particular binding for COand OH and calculate the energetics of the entire reactionnetwork. For each value of GCO and GOH, we calculate the mostendergonic reaction step (i.e., the proposed “rate-determining”step, ΔGRDS) within the most exergonic reaction formulation asa measure for the overall activity of a catalyst possessing thesebinding characteristics. The results at 0.0 and 0.60 V are shownin Figure 2 (an analogous plot for results at 0.35 V is shown inthe Supporting Information, Figure S9).Analyzing the results shows that the behavior of DME

electro-oxidation is quite similar to our previous investigationinto methanol electro-oxidation.37,38 Though we do not choosethe mechanism using analogous indirect (CO as intermediate)or direct (CO not an intermediate) distinctions, the calculatedphase space for DME electro-oxidation is most similar to themethanol electro-oxidation through an indirect mechanism.Similar to methanol electro-oxidation, with very strong CObinding, the reaction rate is limited by CO poisoning. CO isremoved as carboxyl at weak OH binding and by reacting withcoadsorbed OH at moderate to strong OH binding. At verystrong OH binding, we might consider OH and CO to both bepoisoning the surface. Also, at moderate to strong OH bindingand moderate to weak CO binding, we find that the reactionrate is limited by the formation of formaldehyde from adsorbedmethoxy. When both CO and OH binding are weak, the mostdifficult step is in activation of closed-shell species. Formethanol this is methanol and formaldehyde, while for DMEwe have activation of DME (both direct dissociation to CH3and OCH3 and proton/electron extraction to form CH2OCH3)and activation of methanol (which is formed during thereaction).When we move away from the strong and weak binding

asymptotes, we note important differences from methanolelectro-oxidation. First, the optimal binding properties of theDME electro-oxidation catalyst change as a function of theoperating potential, while the peak is invariant for methanolelectro-oxidation. The peak moves for DME electro-oxidationbecause there is competition between C−O bond breaking(whose energetics are independent of operating potential) andproton/electron transfer. As the potential increases from 0 to

Scheme 8. Pathways to Formation of Methyl Formate(HCOOCH3) on Pt(111)a

aValues presented are free energies of reaction (eV) at 0 V on Pt(111).Stoichiometry is balanced with H2O, H

+, and e−. In dotted red arrowsare Heyrovsky-type reactions, and in black arrows are proton/electrontransfer reactions. All species are adsorbed on the surface except forCH3OCH3 and HCOOCH3 which are gas-phase species.

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0.60 V, we see significant changes in the shape of the regionscorresponding to direct dissociation of DME, proton/electronremoval from DME, and proton/electron removal frommethanol. As the potential increases, release of the proton/electron pairs becomes more facile. Methanol dehydrogenationto methoxy becomes easier than C−O bond breaking, butsimultaneously proton/electron extraction from DME becomesfavorable as compared to C−O bond breaking. We note that, ingeneral, as CHxOCHy is dehydrogenated, the free energy of C−O bond dissociation becomes more exergonic (see Scheme 2for energetics on Pt). Therefore, as the potential is increased,CHxOCHy is dehydrogenated, which reduces the apparentdifficulty of C−O bond breaking. In the range of potentialsstudied, CH3OCH3 can dehydrogenate to COCH3 which canbe dissociated to strongly bound CO and CH3.Another interesting result for CH3OCH3 electro-oxidation is

that CH presents a possible poisoning species. When the free

energy of adsorbed CO is more stable than −0.40 eV, CH3 ispreferentially oxidized to CH2 rather than being oxidized tomethanol. If methanol is formed, it will be oxidized to CO,which is relatively easily removed by reaction with coadsorbedOH. However, CH2 is oxidized to CH which is rather difficultto remove, thereby poisoning the surface. CH is removed eitheras CHOH (at weaker CO binding) or by forming carbon(stronger CO binding). Therefore, at moderate to strong CObinding, and moderate OH binding, CH becomes an importantspecies that limits the rate of DME electro-oxidation.As a result, the DME electro-oxidation phase diagram shows

two peaks. The two peaks exist at similar GOH but differ in GCO.The two peaks are formed due to the CH surface poisoningthat creates a local minima in the phase space, splitting thesingle, broad volcano peak that is observed in methanol electro-oxidation.37,38 Note, the more active peak is found at weakerCO binding than the less active peak. If we consider themonometallic surfaces studied, most bind CO around thisminimum (Pt, Pd, Ni, Ir, Rh, Co, Os, Ru, Re), and a numberare situated within the minima (Ni, Ir, Rh, Co, Os, Ru),meaning they have close to optimal OH binding but too strongCO binding. This result will have a significant impact on alloydesign.For methanol electrocatalysts,35,69 Pt is alloyed with another

metal (Ru for example) which is responsible for activatingwater, leading to a bifunctional catalyst. Pt and Ru have similarbinding for CO, and the alloying is targeted toward OHbinding. In the case of DME, we see that Pt has both too weakOH and suboptimal CO binding. The CO binding is too strongsuch that CH is stabilized and cannot be removed as CHOH,but it is not strong enough to form carbon. Ru and many of theother metals have similar CO binding but stronger OH binding.They could enhance the catalyst by providing the bifunction-ality of water activation, but the CO binding would likely besuboptimal. Therefore, it might be better to target the alloyingto either reduce or enhance CO binding. In general, it would bepreferable to reduce CO binding to reach the more active peak.(We also note that the CO surface coverage would likely belower at weaker binding, which is not accounted for in ourmodel but could only help with reactivity.) To this goal, it maybe best to remove Pt from the alloy search as it has bothunfavorable CO and OH binding. Instead, alloying a metalwithin the minima (Ni, Ir, Rh, Co, Os, and Ru) with a metalwhich binds CO weakly (perhaps Cu) might provide the bestcatalyst. However, it is also essential to consider the stability ofthe catalyst under reaction conditions, which we omit from thispresent analysis.Finally, one must also consider the effect of the operating

potential on optimal catalyst design. Since C−O bond breakingis not affected by the potential, as the potential increases therate of the reaction will inevitably be limited by C−O bondbreaking. Therefore, as the operational potential increases, themaximum activity shifts to stronger binding, facilitating C−Obond breaking. In order to achieve a careful balance betweenenergy efficiency and catalyst activity, the binding properties ofthe catalyst should be chosen accordingly.

4. CONCLUSIONUsing periodic, self-consistent (DFT-GGA) calculations, wehave analyzed the electro-oxidation mechanism of dimethylether on a series of monometallic close-packed transition metalsurfaces. Following two different mechanisms, we discover theexistence of two main oxidation peaks. From a mechanism that

Figure 2. Theoretical phase diagram showing free energy of reactionfor the most energetically difficult step (ΔGRDS) within the mostexergonic pathway formulation at (a) 0 V and (b) 0.60 V. Shown areregions with the same most endergonic reactions (note: proton/electron pairs and water are not written explicitly in the reactions).Monometallic catalysts are plotted for comparison/optimization, butmost endergonic reaction identities (and ΔGRDS) may differ frommore rigorous DFT results presented in the text. We note that GOH forAu is 1.71 eV, and therefore is not shown in the phase diagram. Erroris introduced due to the approximate nature of the scaling relationsused to construct this diagram. The blue color is the region of highestactivity (circled in red), while red is the least active region. For asimilar plot at 0.35 V, see Figure S9 (Supporting Information).

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follows the most exergonic transformations at each point alongthe reaction coordinate, we find strongly bound surfacepoisons, most notably CO, CH, and OCH3. The mainoxidation peak is due to the complete oxidation of thesestrongly bound intermediates at high potentials. Following asecond mechanism which is derived by minimizing the difficultyof the most difficult step, we find the existence of a lowerpotential peak due to complete oxidation through formic acidas a key reactive intermediate and partial oxidation to producestrongly adsorbed CH and CO. The existence of this stronglybound CH intermediate leads to a different optimal catalystformulation than is used for methanol electro-oxidation. Inparticular, the optimal catalyst for DME electro-oxidation hassimilar OH binding characteristics as an optimal methanolelectro-oxidation catalyst but has a lower affinity for CO.

■ ASSOCIATED CONTENT

*S Supporting Information

Calculated free energies, zero-point energy corrections, andentropies of all reaction species; most exergonic and lowest-free-energy pathways on all surfaces studied; linear free energycorrelations for all adsorbates; and a theoretical phase diagramcalculated at 0.35 V. This material is available free of charge viathe Internet at http://pubs.acs.org.

■ AUTHOR INFORMATION

Corresponding Author*E-mail: manos@engr.wisc.edu.

Notes

The authors declare no competing financial interest.

Biographies

Jeffrey A. Herron received his B.Sc. (2008) in chemical engineering at

the Pennsylvania State University. In 2013, he received his Ph.D. from

the Department of Chemical & Biological Engineering at the

University of WisconsinMadison under the supervision of Professor

Manos Mavrikakis. Since 2013, he is a postdoctoral researcher at the

University of WisconsinMadison under joint supervision by

Professor Christos T. Maravelias and Professor Manos Mavrikakis.

His research interests include the design of improved heterogeneous

catalysis from first-principles as well as in kinetic modeling.

Peter Ferrin received a B.S. in Chemical Engineering from BrighamYoung University (2004) and a Ph.D. in Chemical Engineering fromthe University of WisconsinMadison (2009). His research interestsin graduate school were in molecular simulation of catalytictransformations of hydrogen and hydrocarbons on transition metalsand their alloys. He is currently employed as a research engineer withHexcel Corporation.

Manos Mavrikakis is the Paul A. Elfers Professor of ChemicalEngineering at UWMadison. His main research interests include theelucidation of detailed reaction mechanisms for heterogeneouslycatalyzed reactions and the identification of improved catalysts fromfirst-principles. According to Thomson-Reuters Prof. Mavrikakis wasone of the “Top 100 chemists for the 2000−2010 decade”, was electedto APS Fellow in 2013, and is the recipient of the 2009 Paul H.Emmett award in Fundamental Catalysis from the North AmericanCatalysis Society. He is the Editor-in-Chief of Surface Science andserves on the editorial board of five more journals.

■ ACKNOWLEDGMENTS

Work at UW-Madison was supported by DOE-BES, Office ofChemical Sciences. JAH thanks Air Products & Chemicals, Inc.for partial support through a graduate fellowship. We thankLuke T. Roling for his comments on the manuscript.Computational work was performed in part using super-computing resources from the following institutions: EMSL, aNational scientific user facility at Pacific Northwest NationalLaboratory (PNNL); the Center for Nanoscale Materials atArgonne National Laboratory (ANL); and the National EnergyResearch Scientific Computing Center (NERSC). EMSL issponsored by the Department of Energy’s Office of Biologicaland Environmental Research located at PNNL. CNM andNERSC are supported by the U.S. Department of Energy,

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Office of Science, under contracts DE-AC02-06CH11357 andDE-AC02-05CH11231, respectively.

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