Theoretical Investigation of Methane Oxidation on Pd(111) and

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Cite This: J. Phys. Chem. C XXXX, XXX, XXX−XXX

Theoretical Investigation of Methane Oxidation on Pd(111) and Other Metallic Surfaces Jong Suk Yoo,*,†,∥ Julia Schumann,† Felix Studt,†,‡,§ Frank Abild-Pedersen,*,† and Jens K. Nørskov†

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†

SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory & Department of Chemical Engineering, Stanford University, Stanford, California 94305, United States ‡ Institute of Catalysis Research and Technology, Karlsruhe Institute of Technology, Hermann-von-Helmholtz-Platz 1, 76344 Eggenstein-Leopoldshafen, Germany § Institute for Chemical Technology and Polymer Chemistry, Karlsruhe Institute of Technology, Engesserstr. 18, 76131 Karlsruhe, Germany S Supporting Information *

ABSTRACT: Density functional theory and microkinetic modeling are employed to investigate CH4 oxidation to CO, CO2, CH2O, and CH3OH on Pd(111) under mildly oxidizing conditions. Although our energetic analysis indicates that the metallic site on Pd(111) is more active than O* and OH* on the Pd surface for C−H bond activation, our microkinetic analysis indicates that the metallic site produces mostly CO, whereas the O* site produces mostly CH2O. In addition, we show that the product selectivity can change significantly depending on the pressures of the products (CO, CO2, CH2O, and CH3OH). Increasing the product pressures leads to the promotion of CO2 production, because CO oxidation becomes more active than CH4 oxidation. We then extend the study to other FCC(111) surfaces by incorporating the linear scaling relations in the mean-field microkinetic model. We find that most transition-metal surfaces cannot effectively activate CH4 under the reaction conditions employed. We find that the kinetics of CH4 oxidation to CO, CO2, CH2O, and CH3OH can be described generally as a function of two descriptors, enabling identification of the most promising catalyst surface for selective production of the desired product.

1. INTRODUCTION Natural gas, which primarily consists of CH4, has become an important source in global energy consumption due to the increase in the production of shale gas.1−3 However, our society has not been efficient in utilizing natural gas, because our CH4 industry is mainly based on combustion (CH4 + 2O2 → CO2 + 2H2O)4,5 and steam reforming (CH4 + H2O → CO + 3H2),6,7 which all require high operating temperature and pressure. It would be desirable if we can directly convert CH4 to more versatile and value-added chemicals such as CH2O and CH3OH via catalytic partial oxidation.8,9 This will not only diversify our utilization of CH4, but also provide a potential solution to the storage issue, because partial oxygenates such as CH2O and CH3OH are energy-dense liquids that can be transported easily with the existing infrastructure.9,10 According to the calculated reaction thermodynamics (eq 1−4), complete oxidation of CH4 to CO2 is significantly more favorable than partial oxidation of CH4 to CO, CH2O, or CH3OH. This means that we need to lower the reaction temperature and the oxygen partial pressure to kinetically promote partial oxidation over complete oxidation. Under mildly oxidizing conditions, metallic catalysts are generally more stable than oxide catalysts because many oxides can be © XXXX American Chemical Society

reduced by, e.g., CO*, OH*, or H* reacting with lattice oxygen on the catalyst surface.11−13 Nevertheless, high CH3OH (or CH2O) selectivity has been achieved, mostly by oxide catalysts such as MoO3 and V2O3.14−17 This is because, O-insertion into CH3* (or CH2*) to form CH3O* (or CH2O*) and then reacting further to CH3OH (or CH2O) is likely to be more selective on oxides than metallic catalysts.18 The coexistence of both metal and lattice-oxygen sites on oxide catalysts plays an important role in preventing severe dehydrogenations of adsorbed hydrocarbon fragments. CH4 + 0.5O2 → CH3OH

ΔH ° = −1.3 eV

CH4 + 1.0O2 → CH 2O + H 2O CH4 + 1.5O2 → CO + 2H 2O CH4 + 2.0O2 → CO2 + 2H 2O

ΔH ° = −2.9 eV ΔH ° = − 5.4 eV ΔH ° = − 8.3 eV

(1) (2) (3) (4)

Received: March 3, 2018 Revised: June 25, 2018 Published: June 26, 2018 A

DOI: 10.1021/acs.jpcc.8b02142 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C

chemisorption and van der Waals interaction on metal surfaces reasonably well.32 The slab models for FCC(111) surfaces are created based on 3 × 3 × 4 supercells, which are separated by more than 13 Å of vacuum in the direction perpendicular to the surface. Thus, all adsorption energies (see Table S1 in the Supporting Information (SI) for values) are obtained at the fixed coverage of 1/9 ML. Transition-state geometries are calculated with the CI-NEB33 method with an initial path generated by stretching the relevant bond from the most stable associated state. The Vanderbilt ultrasoft pseudopotentials34 are used to describe the ionic cores, and the electronic energy calculations are carried out using a 4 × 4 × 1 Monkhorst−Pack k-point mesh35 with a plane-wave cutoff of 500 eV, a density cutoff of 5000 eV, and a Fermi-level smearing of 0.1 eV. The top two atomic layers of the slab models as well as all the adsorbates are relaxed until forces below 0.03 eV/Å per atom is reached. The convergence of adsorption energy with respect to the slab thickness, k-point sampling, and cutoff energies has been confirmed elsewhere.36,37 The vibrational frequencies of adsorbates (see Table S1 in the SI for values) are obtained based on a finite difference approximation to the Hessian and subsequent diagonalization to find the normal modes, as implemented in ASE.30 They are used to obtain freeenergy corrections for adsorbates (see Table S2 in SI for values) by assuming that all degrees of freedom are harmonic vibrational modes. Note that the free-energy corrections are obtained from calculations of adsorbates on Pd(111) and Cu(111) only. For surfaces other than Pd(111), the values obtained for Cu(111) are used (see the caption of Table S2 in SI for justification). Free-energy corrections for gas-phase molecules are calculated using the Shomate equation with parameters obtained by fitting from experimental data. The electronic energies of gas-phase O2 and H2 are corrected by +0.7 and +0.09 eV, respectively, in order to accurately describe the reaction enthalpies of overall reactions, as discussed elsewhere.38 Note that the energy correction is particularly large for gas-phase O2 because of the well-known limitation of the generalized gradient approximation (GGA) in describing the O−O bond. This brings some uncertainties to calculations of oxygen coverages as the initial state of O2 dissociative adsorption is corrected by +0.7 eV, but the transition state (as well as the final state) is not corrected as it is very final-state like. The electronic energy of gas-phase CO2 is also corrected by +0.41 eV in order to address the systematic error that originates from inaccurate descriptions of carbon− oxygen double bonds, as shown elsewhere.39,40 The electronic energy of adsorbed CO* is corrected by +0.25 eV in order to account for the overestimation of the interaction between the 2π* of CO* and metal d-electrons, as shown elsewhere.41 2.2. Microkinetic Modeling. The microkinetic modeling is performed using the Catalysis Microkinetic Analysis Package (CatMAP).42 Nondissociative adsorption steps are treated as if gas-phase molecules are in thermodynamic equilibrium with their environment. Therefore, adsorption prefactors are obtained purely from the entropic losses based on the transition state theory. The calculated prefactors of e.g. CH4 and O2 adsorptions are generally in agreement with the typical experimental pseudo first-order prefactor of slow adsorption, i.e., 102−105 s−1. Site coverages are modeled using the meanfield approach and the steady-state approximation. The FCC(111) surface is modeled using two surface sites: a “hydrogen reservoir” site and a site for all other intermediates.36,43 This means that both sites can have a coverage of one

In this context, a promising approach to exploiting the stability of metallic catalysts under mildly oxidizing conditions while pursuing high selectivities to partial oxygenates is to utilize the oxygen species chemisorbed on metallic surfaces as active sites for CH4 oxidation. Previous studies have shown that chemisorbed oxygen species can play an important role depending on the identity of the metallic surface.19−22 For example, density functional theory (DFT) studies were recently conducted to show that O* and OH* hinder C−H bond activation on Pt-group metals whereas they significantly promote the bond activation on coinage metals.21,22 However, it is still elusive what the kinetic contributions of chemisorbed oxygen species are on different metallic surfaces during CH4 oxidation under mildly oxidizing conditions. Herein, we employ periodic density functional theory and microkinetic modeling to investigate the reaction energetics and kinetics of CH4 oxidation to CO, CO2, CH2O, and CH3OH on various FCC(111) surfaces under mildly oxidizing conditions. The FCC(111) facet is chosen since oxygen species chemisorb relatively weakly on this facet, thus they are more reactive compared to those on other facets. Scheme 1 shows the considered CH4 oxidation pathways that are generally in agreement with previous studies.23−25 Scheme 1. Reaction Pathways Considered for CH4 Oxidation (Top) and Oxidant Activation (Bottom)a

a

Black, red, blue, and green arrows represent dehydrogenation, Oinsertion, OH-insertion, and desorption reactions, respectively. The molecular adsorption (physisorption) steps as well as the desorption steps of CH4, CH3OH, and CO2 are excluded as they are assumed to be nearly barrier-less.

The effects of chemisorbed oxygen species during CH4 oxidation will be addressed by considering three different C−H bond activation pathways for all the dehydrogenation steps shown in Scheme 1, i.e., the direct (e.g., CH4 → CH3* + H*), O-assisted (e.g., CH4 + O* → CH3* + OH*), and OHassisted (e.g., CH4 + OH* → CH3* + H2O) mechanisms.22−25 We will first focus on discussing our calculation results obtained for Pd(111), since Pd based catalysts are widely used in hydrocarbon oxidation.25−28 However, we will later develop a broader understanding of the performance of other FCC(111) surfaces by combining energy scaling relations with our microkinetic model to map the kinetics of CH4 oxidation as a function of only two descriptors, i.e., carbon and oxygen adsorption energies.

2. COMPUTATIONAL METHODS 2.1. DFT Calculation. The Quantum ESPRESSO29 code interfaced with the Atomic Simulation Environment (ASE)30 is used to perform periodic DFT calculations using the BEEFvdW functional.31 This functional is known to describe both B


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ultimately leading to the most stable state where only CO* and O* are present on the surface. At this point in the reaction, there is a close competition between CO* desorption (CO* → CO) and CO2 formation (CO* + O* → CO2). However, it is difficult to judge merely based on the reaction energetics which reaction is dominant as the free-energy barrier for CO2 formation is comparable to the free energy required to desorb CO*. Kinetic information that takes into account the steadystate coverages of CO* and O* on the Pd surface is required to accurately determine the product selectivity of Pd(111) during CH4 oxidation. Figure 1 also shows that all the dehydrogenation steps involved in the minimum energy pathway occur via the direct C−H bond activation mechanism. This indicates that the metallic site on Pd(111) is more active than O* or OH* on the surface during CH4 oxidation.21,22 However, it is important to note here that both O* and OH* on the Pd surface play an important role in the reaction. They contribute to the catalytic activity and selectivity via different reaction pathways, since the calculated activation barrier for O-assisted (or OH-assisted) CH4 activation is found to be only ∼0.4 eV (or ∼0.2 eV) higher than that for direct CH4 activation (see Table S1 in the SI).21 A more quantitative analysis of all possible reaction pathways via microkinetic modeling is necessary to precisely determine the role of adsorbed O* and OH* on the catalytic activity and selectivity of Pd(111). 3.2. Microkinetic Analysis of CH4 Oxidation on Pd(111). Figure 2a shows the calculated production rates of CO2, CO, CH2O, and CH3OH on Pd(111) as functions of temperature (see the caption for the reaction conditions). We find that all rates generally increase with increasing temperature. However, CO and CH2O productions are found to be substantially more active than CO2 and CH3OH productions under all temperature conditions. More importantly, we observe that the calculated production rates are very low even at high temperature (TOF < ∼1 s−1 at T < 900 K). This is in agreement with the previous experimental finding that the formation of metallic Pd on the catalyst surface generally decreases the conversion rate during CH4 combustion on Pd oxides.47,48 As discussed in the first paragraph of section 3.1, metallic Pd is not very efficient in activating CH4 compared to e.g. PdO(101),45,46 although it appears to selectively promote the desired partial oxidation reactions (CO and CH2O productions) under the reaction conditions studied. Figure 2b shows the calculated product selectivity of Pd(111) as a function of temperature, as obtained from the data in Figure 2a. It can be seen that the CO selectivity is always greater than 0.5 and generally increases with increasing temperature, whereas the CH2O selectivity is always less than 0.4 with maximum at T = 550 K. We identify from the free energy diagram in Figure 1 (or see Figure S1b (section 3) in the SI for the free energy diagram obtained with considering lateral adsorbate−adsorbate interactions) that CO is produced via a nonassisted reaction path, whereas CH2O is produced primarily via the O-assisted pathway that involves CH3* and CH3O* as intermediates (i.e., CH4 → CH3* + H*, CH3* + O* → CH3O*, and CH3O* + O* → CH2O + OH*). This suggests that the selectivity toward CH2O can be enhanced when increasing the O* coverage on the Pd surface. However, this will occur at the expense of decreasing the CH4 conversion rate since the metallic sites are more active than adsorbed O* on the Pd surface for CH4 activation.

simultaneously, which is unphysical, but can be justified as H* coverage is found to be extremely low in all cases considered in this study. As mentioned above, the effects of chemisorbed oxygen species during CH4 oxidation is taken into account by considering the direct, O-assisted, and OH-assisted mechanisms for all of the dehydrogenation steps shown in Scheme 1. In addition, lateral adsorbate−adsorbate interactions between relevant intermediate species such as CO*, O*, C*, CH*, COH*, and CHO* are included in the microkinetic model (see section 2 in the SI for details). They are obtained using a second-order expansion in the coverage for the integral adsorption energy from which the differential adsorption energy can be derived (see section 2 in the SI, and refs 43 and 44 for details). More details about our microkinetic model are provided in section 6 in the SI.

3. RESULTS AND DISCUSSION 3.1. Energetic Analysis of CH4 Oxidation on Pd(111). Figure 1 shows the calculated minimum energy pathway for

Figure 1. Calculated minimum energy pathway (black) for CH4 oxidation on Pd(111) at 523 K and PCH4 = 0.91 bar, PO2 = 0.11 bar, PH2O = 0.02 bar, and PH2, PCH3OH, PCH2O, PCO, PCO2 = 10−39 bar. Note that the product pressures are extremely low such that the reaction is modeled at nearly zero percent approach to equilibrium. Red, green, and blue show some other competing pathways of interests, such as those that lead to CH2O, CH3OH, and CO productions, respectively. Note that adsorbate−adsorbate interactions were not considered here. See Figure S1 (section 3) in the SI for the free-energy diagram on Pd(111) obtained with adsorbate−adsorbate interactions considered.

CH4 oxidation on Pd(111) at T = 523 K. We observe that O2 dissociation (O2 → 2O*) is exothermic and has the lowest free-energy barrier of ∼0.2 eV, whereas CH4 activation (CH4 → CH3* + H*) is endothermic with the highest free-energy barrier of ∼1.6 eV. This is in agreement with the previous theoretical finding that metallic Pd is not very efficient in activating CH4 compared to Pd oxides (e.g., the CH4 activation barrier is found to be lower on PdO(101) than Pd(111), Pd(211), and Pd(100)).28,45,46 Herein, we focus on reactions on metallic Pd as we are more interested in partial oxidation of CH4 under mildly oxidizing conditions where pure metals are generally more stable than oxides. Interestingly, the CH3* fragment formed from CH4 does not combine with O* (or OH*) to form CH 3 O* (or CH 3 OH). It rather dehydrogenates to CH* that then combines with OH* to form CHOH* on the Pd(111) surface. The CHOH* species then decomposes to CO* via either CHO* or COH*, C


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Figure 2. Temperature dependent (a) production rates, (b) product selectivities, (c) steady-state coverages, and (d) CH4 activation rates on Pd(111) at PCH4 = 0.91 bar, PO2 = 0.11 bar, PH2O = 0.02 bar, and PH2, PCH3OH, PCH2O, PCO, PCO2 = 10−39 bar. The oxygen deficient reaction condition is specifically chosen such that metallic surfaces are thermodynamically stable and no surface oxides are formed. Note that the product pressures are extremely low such that the reaction is modeled at nearly zero percent approach to equilibrium. Adsorbate−adsorbate interactions were considered here. See Figure S2 (section 4) in the SI for the microkinetics on Pd(111) obtained without considering adsorbate−adsorbate interactions.

function of temperature. We find that CH4 activation at low temperature occurs primarily via the direct mechanism (i.e., on the metallic site), in agreement with previous theoretical studies.21,22 At higher temperatures the O*-assisted pathway is nearly as active as the direct pathway on the metallic Pd site. We can also observe that the sum of all CH4 activation rates shown in Figure 2d matches with the sum of all production rates shown in Figure 2a, indicating that CH4 activation is the rate-limiting step for CH4 oxidation on Pd(111) under these reaction conditions. Although Figure 2a shows that CH3OH production is insignificant on Pd(111), we still examine the reaction pathway and identify that CH3OH is produced mainly via OH-assisted hydrogenation of CH3O* (CH3O* + OH* → CH3OH + O*) that directly competes with O-assisted dehydrogenation of CH3O* to form CH2O (CH3O* + O* → CH2O + OH*). Thus, a promising approach to achieving high CH3OH selectivity on the Pd catalyst is to increase the relative coverage of OH* to O* during CH4 oxidation, which can be done by either cofeeding H2 with O2 to promote OH* formation (O* + H* → OH*) or switching to hydrogen peroxide (HOOH) as oxidant that decomposes to OH* species on Pt-group metals,49 if we assume that OH* formed from HOOH does not further decompose to O* and H*. This theoretical prediction is in agreement with previous experimental finding that Au−Pd/TiO2 catalyst can selectively oxidize CH4 to CH3OH (∼50% selectivity) using HOOH as oxidant at T = 323−363 K.50 Again, the increased CH3OH selectivity on Au−Pd/TiO2 is a consequence of the catalyst being able to form and maintain OH* species on the surface

Figure 2c shows the calculated steady-state coverages of adsorbates on Pd(111) as a function of temperature. We find that O* is the only reaction intermediate partially covering the Pd surface under all temperature conditions. This shows that OH-assisted dehydrogenation pathways are irrelevant on the Pd(111) surface under these reaction conditions. In addition, we can observe that the CO* coverage is nearly zero even at room temperature, which is somewhat surprising given that metallic Pd catalysts are generally prone to CO*-poisoning particularly at low temperature. However, this can be explained by the low CO partial pressure, making CO desorption highly favorable. We also need to consider that CH4 conversion (i.e., CO* production) is much slower than O2 activation at all temperatures. As a result, the absence of CO* on the Pd surface suppresses CO2 formation (CO* + O* → CO2) more than CO* desorption (CO* → CO), in agreement with the kinetic results shown in Figure 2a. We also note here that our calculated steady-state C* coverage on Pd(111) is very low (nearly zero) compared to the previous theoretical study by Grönbeck et al.25 This is because C* + OH* → COH* is significantly promoted by considering adsorbate−adsorbate interactions for CO*, O*, C*, CH*, CHO*, and COH* (see section 2 in the SI for more details) as can be seen by the comparison between Figure S1a and S1b in the SI. These species are specifically chosen as our microkinetic modeling with no adsorbate−adsorbate interactions indicate that a transition-metal surface can only be partially covered with combinations of these species. Figure 2d shows the calculated CH4 activation rates via different C−H bond activation mechanisms on Pd(111) as a D


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Figure 3. Temperature dependent (a) production rates, (b) product selectivities, (c) steady-state adsorbate coverages, and (d) consumption rates on Pd(111) at PCH4 = 0.91 bar, PO2 = 0.11 bar, PH2O = 0.02 bar, and PH2, PCH3OH, PCH2O, PCO, and PCO2 = 0.01 bar. Note that the product pressures are increased from 10−39 to 0.01 bar. The inset in panel d is a magnified view of the figure. Note that adsorbate−adsorbate interactions were considered here.

at T < 900 K). This strongly indicates that the produced CO2 and CH2O comes from carbon sources other than the main reactant CH4, such as CO, CH2O, and CH3OH. This will be discussed in more detail below when we analyze the different consumption rates of different gas-phase species. We also find that CO and CH3OH are produced at very low rates at low temperature (TOF < ∼10−6 s−1 at T < 400 K). Figure 3b shows the temperature dependent product selectivity of Pd(111), as obtained from the data in Figure 3a. We find that the CO selectivity (0.85−1) is higher than CO2 selectivity (0−0.15) only at temperatures lower than 350 K. However, as the temperature increases above 350 K, the CO selectivity dramatically decreases to nearly zero, which is accompanied by a dramatic increase in the CO2 selectivity. At temperatures higher than 650 K, the CH2O selectivity slightly increases but remains relatively insignificant with maximum of ∼0.1 at T = 700 K. Finally, the CH3OH selectivity is found to be close to zero at all temperatures. These indicate that the product selectivity of Pd(111) can change significantly depending on the product pressures, and it is important to keep them very low to obtain high selectivities toward partial oxidation products. Figure 3c shows the temperature dependent steady-state coverages of adsorbates on Pd(111) with increased product pressures. Comparison between Figure 2c and 3c reveals that the increased CO pressure in the reaction feedstock leads to the accumulation of CO* on the Pd surface particularly at low temperature, which is, however, compensated by the decrease in the O* coverage, thus effectively preserving the amount of free (metallic) sites required for the direct dehydrogenation of hydrocarbons. The coexistence of both CO* and O* on the Pd

during CH4 oxidation. Reported DFT calculations on a similar catalyst (Au0.89Pd0.11(221)) show that HOOH → 2OH* is exothermic (ΔH = −1.63 eV), whereas OH* → O* + H* is endothermic (ΔH = 1.66 eV) for the synthesis of HOOH from H2 and O2.51 3.3. Effect of Increasing the Product Pressures during CH4 Oxidation on Pd(111). Another significant problem with catalytic partial oxidation of CH4 is that products such as CO, CH2O, and CH3OH are likely to be more reactive than the main reactant CH4. This means that, although a catalyst may be initially selective to the partial oxidation products, the product selectivity can change significantly over catalyst contact time due to the increased product pressures. Thus, we discuss below how the activity and selectivity of the Pd catalyst change when we arbitrarily increase all the product pressures (i.e., PCO, PCO2, PCH2O, and PCH3OH) from 10−39 to 10−1 bar. This will lead to a better understanding of which partial oxidation product is the most reactive on the Pd surface, and how it affects the catalytic properties during CH 4 oxidation. Figure 3a shows the temperature dependent production rates of CO, CO2, CH2O, and CH3OH on Pd(111) obtained with increased product pressures (see the caption for the reaction conditions). It is now observed that the Pd catalyst produces mostly CO2 under all temperature conditions with substantially enhanced production rates (TOF < ∼105 s−1 at T < 900 K) compared to those shown in Figure 2a (TOF < ∼10−5 s−1 at T < 900 K). In addition, CH2O production now becomes active only at high temperature with substantially enhanced production rates (TOF < ∼103 s−1 at T = 650−900 K) compared to those shown in Figure 2a (TOF < ∼10−1 s−1 E


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Figure 4. Logarithms of the calculated production rates (s−1 per site) of (a) CO2, (b) CO, (c) CH2O, and (d) CH3OH during CH4 oxidation at T = 523 K, PCH4 = 0.91 bar; PO2 = 0.11 bar; PH2O = 0.02 bar; and PH2, PCH3OH, PCH2O, PCO, and PCO2 = 10−39 bar. Pressures are identical to those specified in the caption of Figure 2 (nearly zero percent approach to equilibrium). ΔEC and ΔEO are taken relative to gas-phase CH4, H2O, and H2. The expected errors for ΔEC and ΔEO on transition metals are ±0.2 eV. Note that adsorbate−adsorbate interactions were considered here. The chemical equilibrium (constant temperature and volume) obtained with the reaction condition specified above as the initial state is 93.64% of CH4, 0.01% of CO, 6.35% of CO2, and nearly 0% of CH2O and CH3OH.

Figure 5. Calculated (a) CO2, (b) CO, (c) CH2O, and (d) CH3OH selectivities during CH4 oxidation at T = 523 K, PCH4 = 0.91 bar; PO2 = 0.11 bar; PH2O = 0.02 bar; and PH2, PCH3OH, PCH2O, PCO, and PCO2 = 10−39 bar. White circles in panels c and d are added based on calculated ΔEC and ΔEO values for the respective surfaces. Pd3Au and Au3Pd represent the FCC(111) surfaces of homogeneously mixed binary alloys with different atomic compositions, whereas Pd@Au, Pd3Au@Au, and PdAu@Au represent the FCC(111) surfaces of shell@core structures with only one atomic layer as the shell. ΔEC and ΔEO are taken relative to gas-phase CH4, H2O, and H2. The expected errors for ΔEC and ΔEO on transition metals are ±0.2 eV. Note that adsorbate−adsorbate interactions were considered here.

Figure 3d shows the temperature dependent consumption rates of different gas-phase species with increased product pressures. Negative consumption rates mean that the gas-phase

surface is a factor that promotes CO2 formation (CO* + O* → CO2) over CO* desorption (CO* → CO), corroborating the results shown in Figure 3a. F


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CH2O, Pd(111) produces both species in similar amounts, whereas Rh(111) produces significantly more CH2O than CO. Interestingly, the differences in the CH2O/CO production ratio on the Pt-group metal surfaces can be correlated to the differences in the steady-state coverage of O* shown in Figure S4 (section 7) in the SI. The more the surface is covered with O*, the more CH2O is produced relative to CO. This indicates that O* on Pt-group metal surfaces leads to the desired CH2O production via the same pathway identified for Pd(111) (i.e., CH4 → CH3* + H*, CH3* + O* → CH3O*, and CH3O* + O* → CH2O + OH*), whereas the metallic site on Pt-group metal surfaces leads to the undesired CO production via the same pathway identified for Pd(111) (i.e., the minimum energy pathway shown in Figure 1). In Figure 5a−d, we observe that coinage metal surfaces such as Ag, Au, and Cu(111) tend to produce mainly partial oxygenates such as CH2O and/or CH3OH compared to other products. However, Ag(111) produces both CH2O and CH3OH in similar amounts, Au(111) produces significantly more CH3OH than CH2O, whereas Cu(111) produces significantly more CH2O than CH3OH. We find that all dehydrogenation steps occur via the O-assisted mechanism on coinage metals. For example, CH2O is produced mainly via CH4 + 3O* → CH3* + OH* + 2O* → CH3O* + OH* + O* → CH2O* + 2OH* → CH2O + O* + H2O on Ag(111) and Cu(111), whereas CH3OH is produced mainly via CH4 + O* → CH3* + OH* → CH3OH on Ag(111) and Au(111). The CH2O vs CH3OH production on coinage metals is determined by the competition between CH3* + O* → CH3O* (that ultimately leads to CH2O production) vs CH3* + OH* → CH3OH. This indicates that the presence of OH* on the catalyst surface could promote the selective formation of CH3OH during CH4 oxidation on coinage metals. Finally, Figure 5a−d shows that no metallic surfaces can oxidize CH4 to CO2 with high selectivities under the considered reaction conditions, indicating the importance of preserving the product pressures low to promote partial oxidation reactions (particularly, CH2O and CH3OH productions) over CO2 production on all metallic surfaces. For example, see section 8 and 9 in the SI for our discussion about the effects of increasing product pressures (10−39 → 10−2 bar) and reaction temperature (523 → 823 K) on the activity and selectivity trends. 3.5. Comparison with Previous Experiments. It is generally difficult to compare calculated results with previous experiments directly due to the following reasons. (1) The calculated production rates shown in Figure 4a−d are low (TOF < ∼10−3 s−1) for most transition-metal surfaces, and experiments are usually conducted at high conversion rates. (2) The experimental operating conditions including reaction temperatures, pressures, catalyst contact times, and the ratio between CH4 and O2 (CH4/O2) vary significantly depending on the catalyst as well as the targeted product. For example, high temperature (T = 800−1300 K) is generally used to produce synthesis gas from CH4,54 whereas low temperature (T = 300−850 K) is most often used to produce partial oxygenates (e.g., CH2O and CH3OH) from CH4.50,55 (3) It is not clear, whether the active phase of the employed catalyst surface in experiments is metallic or in the form of an oxide. (4) Synthesized catalysts are often supported on oxides and effects due to support interactions can be significant particularly for reducible oxides. We stress that metal−support interactions are not addressed in this study.

species are produced rather than consumed. Interestingly, CO consumption rates shown in Figure 3d match with CO2 production rates shown in Figure 3a, indicating that CO2 is produced mainly from the introduced pressure of CO rather than CH4. In addition, CH3OH consumption rates shown in Figure 3d match with CH2O production rates shown in Figure 3a, indicating that CH2O is produced mainly from CH3OH and not from CH4. We note that CH4 is neither consumed nor produced in significant amounts under all temperature conditions, again indicating that Figure 3a−c shows the results of CO and CH3OH oxidations rather than CH4 oxidation. This indicates that CO and CH3OH are more easily oxidized than CH4; hence, the key to achieving high selectivities toward the less oxidized products on Pd based catalysts is to preserve low product pressures during the reaction. 3.4. Microkinetic Analysis of CH4 Oxidation as Functions of Descriptors. We now extend our study to include other FCC(111) surfaces using linear scaling relations between calculated adsorption and transition state energies of all of the reaction intermediates on Ag, Cu, Pd, Pt, and Rh(111) and the carbon and oxygen adsorption energies (ΔEC and ΔEO). These descriptors are chosen, specifically, since all of the reaction intermediates involved in CH4 oxidation, except H*, bind to the catalyst surface through either a carbon and/or an oxygen atom.52 Hydrogen binding can, to a good approximation, be linearly correlated with the carbon binding as shown elsewhere.37,53 Figure S3 (section 5) in the SI shows the obtained linear scaling relations used in our mean-field microkinetic model. This approach allows for the description of the kinetics of CH4 oxidation as functions of the two descriptors covering a wide range of the chemical space. Figure 4a−d shows the calculated production rates of CO2, CO, CH2O, and CH3OH as a function of the two descriptors (see the caption for the reaction conditions). We observe that the modeled production rates are relatively low (TOF < ∼10−3 s−1 at T = 523 K) for transition-metals surfaces. The most active production region (the darkest red region shown in Figure 4a,b) that activates CH4 via the direct mechanism requires moderate oxygen binding energies to avoid O poisoning and very strong carbon binding energies enabling CH4 activation, and no transition metal surfaces are found in that region of the phase space. Figure 4a−d also shows that both Pt-group metal surfaces such as Pd, Pt, and Rh(111) as well as noble metal surfaces such as Ag(111) have similar activities. This is because Pd, Pt, and Rh(111) have carbon binding energies that promote direct CH4 activation, whereas Ag(111) is in a region where oxygen interaction with the surface is strong enough to promote both O2 dissociation (strong oxygen binding preferred) and Oassisted CH4 activation (weak oxygen binding preferred). In the case of Au(111), all production rates are low (TOF < ∼10−6 s−1 at T = 523 K). The Au surface has very weak oxygen binding, resulting in an oxygen surface coverage close to zero. (see Figure S4 (section 7) in the SI for the steady-state coverages of adsorbates on different metallic surfaces including Au(111)). We find that O* is the main active site for CH4 activation on noble metals. Figure 5a−d shows the CO2, CO, CH2O, and CH3OH selectivities calculated from the data in Figure 4a−d as a function of the two descriptors. We observe that Pt-group metal surfaces such as Pd, Pt, and Rh(111) tend to produce mainly CO and/or CH2O compared to other products. However, Pt(111) produces significantly more CO than G


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CH3O* (CH3O* + OH* → CH3OH + O*) that directly competes with O-assisted dehydrogenation of CH 3O* (CH3O* + O* → CH2O + OH*). Thus, a possible approach to achieve higher CH3OH selectivity on the Pd catalyst is to increase the relative coverage of OH* to O* during CH4 oxidation. Furthermore, we have also examined the effect of increasing the product pressures during CH4 oxidation on Pd(111). We find that it is important to keep the product pressures low to achieve high selectivities toward the less oxidized products, since CO and CH3OH oxidations to CO2 and CH2O, respectively, are found to be significantly more active than CH4 oxidation. With increased product pressures, Pd(111) now produces mostly CO2 with high selectivity (>0.9) at a temperature higher than 400 K. We have then extended the study to other FCC(111) surfaces by incorporating the linear scaling relations into our microkinetic model such that the catalytic activity and selectivity are described as functions of carbon and oxygen adsorption energies. Interestingly, most transition-metal surfaces such as Ag, Au, Cu, Pd, Pt, and Rh(111) are not highly active for CH4 oxidation under the reaction conditions. However, Pt-group metal surfaces such as Pd, Pt, and Rh(111) are selective to CO and/or CH2O productions, whereas coinage metal surfaces such as Ag, Au, and Cu(111) surfaces are selective to CH2O and/or CH3OH productions. In the case of the Pt-group metal catalysts, the metallic site leads to CO production, whereas the O* site leads to CH2O production. In the case of the coinage metal catalysts, all dehydrogenation steps occur via the O-assisted mechanism, and it is the competition between CH3* + O* → CH3O* (that ultimately leads to CH2O production) and CH3* + OH* → CH3OH, which determines the CH2O vs CH3OH selectivity. Thus, forming and maintaining OH* on the catalyst surface to promote the latter reaction is the key to achieving high CH3OH selectivity during CH4 oxidation on coinage metals. The activity and selectivity trends obtained in this study are in agreement with previous experiments. The theoretical insights obtained in this study should provide guiding principles for the development of new and more efficient metallic catalysts for CH4 partial oxidation under mildly oxidizing conditions.

Nevertheless, some agreements between our calculations and previous experiments can be found. For example, experiments have found that Pt-group metal catalysts (i.e., Pd, Pt, and Rh nanoparticles supported on alumina) are suitable for CH4 oxidation to synthesis gas, achieving high CO selectivities (>90%) at high temperatures (T = 650−1050 K, CH4/O2 = 1−2).56−58 This is in agreement with our study showing that increasing reaction temperature from 523 to 823 K (at CH4/O2 = 8.27) leads to enhanced CO production rates and CO selectivities on Pd, Pt, and Rh(111), as shown in Figure S9b and S10b (section 9) in the SI. It is also widely known that coke formation can be a problem for Pt-group metal catalysts during CH4 oxidation to synthesis gas particularly at high temperatures.57,58 This is also in agreement with our study showing that increasing reaction temperature from 523 to 823 K dramatically increases the C* coverage on strongly carbon-binding surfaces, as shown in Figure S11b (section 9) in the SI. A further increase in reaction temperature above 823 K can severely deactivate Pt-group metal surfaces via C*-poisoning. Previous experiments have also found that the CuOx/SBA15 catalyst shows the highest activity (TOF = 3.1−5.6 s−1) and selectivity (58−71%) for partial oxidation of CH4 to CH2O (at T = 898 K, CH4/O2 = 1) among various M/SBA-15 (M = V, Cr, Mn, Fe, Co, Ni, Cu, Mo, and W) catalysts.59 Although the active site of the Cu catalyst is not well-defined, this is partially in agreement with our study showing that Cu is one of the most suitable elemental catalysts for active and selective production of CH2O from CH4 (at T = 823 K, CH4/O2 = 8.27), as shown in Figure S9c and S10c (section 9) in the SI, respectively. In addition, previous experiments have found that Au−Pd/TiO2 catalyst shows high CH3OH selectivity (∼50% with TOF = 0.0002 s−1) using HOOH as an oxidant at low temperature (T = 323 K).50 This is also in agreement with our study showing that different compositions of Au−Pd binary alloys or core−shell structures can be highly selective for CH2O or CH3OH production, as shown in Figure 5c,d, although O2 is still used as an oxidant at slightly higher temperature (T = 523 K). Pd-rich surfaces are generally preferred for CH2O production, whereas Au-rich surfaces are generally preferred for CH3OH production.

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4. CONCLUSIONS Based on periodic DFT and microkinetic modeling, we have examined the reaction energetics and kinetics of CH4 oxidation to CO, CO2, CH2O, and CH3OH on a partially oxygen covered Pd(111) surface. Our energetic analysis indicates that the metallic site is more active than O* on the Pd surface for C−H bond activation. Our microkinetic analysis indicates that even though O* on the Pd surface does not participate directly in the activation of CH4, it is important for the overall reaction mechanism. We have identified that the metallic site on Pd(111) produces mostly CO via the pathway that involves CHOH*, whereas the O* on the Pd surface produces mostly CH2O via O-assisted dehydrogenation of CH3O* (CH3O* + O* → CH2O + OH*). As a result, the CO selectivity is found to be always greater than 0.5 and generally increases with increasing temperature, whereas the CH2O selectivity is found to be always less than 0.4 with maximum at T = 550 K. Although CH3OH production on Pd(111) is found to be insignificant during CH4 oxidation under all temperature conditions, we have examined the pathway and found that CH3OH is produced mainly via OH-assisted hydrogenation of

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.8b02142.

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Additional DFT data including adsorption energies, vibrational frequencies, details about adsorbate−adsorbate interactions, and microkinetic modeling under different reaction conditions (PDF)

AUTHOR INFORMATION

Corresponding Authors

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

Jong Suk Yoo: 0000-0001-6472-7004 Julia Schumann: 0000-0002-4041-0165 Frank Abild-Pedersen: 0000-0002-1911-074X H


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The Journal of Physical Chemistry C Present Address

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Department of Mechanical Engineering, Massachusetts Institute of Technology, Cambridge, Massachusetts 02139, United States. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We gratefully acknowledge the support from the U.S. Department of Energy, Office of Sciences, Office of Basic Energy Sciences to the SUNCAT Center for Interface Science and Catalysis. J.S.Y. gives special thanks to the U.S. Department of State for supporting his Ph.D. studies through the International Fulbright Science & Technology Award program. J.S. would like to acknowledge the support from the German Academic Exchange Service (DAAD) through a postdoctoral fellowship.

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Theoretical Investigation of Methane Oxidation on Pd(111) and

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