Activity Descriptors for CO2

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Activity Descriptors for CO2 Electroreduction to Methane on Transition-Metal Catalysts Andrew A. Peterson†,§ and Jens K. Nørskov*,†,‡ †

SUNCAT Center for Interface Science and Catalysis, Department of Chemical Engineering, Stanford University, Stanford, California, United States ‡ SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, Menlo Park, California, United States § Center for Atomic-scale Materials Design, Department of Physics, Technical University of Denmark, DK-2800 Lyngby, Denmark S Supporting Information *

ABSTRACT: The electrochemical reduction of CO2 into hydrocarbons and alcohols would allow renewable energy sources to be converted into fuels and chemicals. However, no electrode catalysts have been developed that can perform this transformation with a low overpotential at reasonable current densities. In this work, we compare trends in binding energies for the intermediates in CO2 electrochemical reduction and present an activity “volcano” based on this analysis. This analysis describes the experimentally observed variations in transition-metal catalysts, including why copper is the best-known metal electrocatalyst. The protonation of adsorbed CO is singled out as the most important step dictating the overpotential. New strategies are presented for the discovery of catalysts that can operate with a reduced overpotential.

U

nlike renewable electricity, which can come from several sources (hydro, solar, wind, geothermal, and others), renewable carbon-based fuels are currently only produced from biomass. The only other possibility to produce such fuels is by chemically reducing carbon dioxide; if CO2 can be electrochemically reduced to hydrocarbons, then any of the renewable electricity sources can also be used to create fuels, simultaneously addressing storage issues for these often intermittent resources.1 Over the past 3 decades, many heterogeneous and homogeneous electrocatalysts have been screened for their effectiveness in this reaction and the related photoelectrochemical process (artificial photosynthesis),2−18 and none have been found to be more effective than pure metallic copper for its repeatable demonstration to produce high (>50%) Faradaic yields of hydrocarbons at reasonable (∼5 mA cm−2) current densities. The high selectivity of copper for hydrocarbons in CO2 electroreduction has sparked much interest in uncovering its mechanism.11−16,19,20 However, the overpotential necessary to achieve this reaction on copper electrodes is prohibitively high, on the order of 1 V,2 which dwarfs the overpotential of the oxygen evolution reaction,21,22 the typical counter electrode reaction for these systems. From a purely thermodynamical point of view, CO2 reduction to the simplest hydrocarbon, CH4, should be possible at a standard potential of +0.17 V versus the reversible hydrogen electrode, or RHE. Because, by definition, hydrogen evolution becomes thermodynamically possible at 0 V versus RHE (at any pH), CO2 reduction will be in competition with the hydrogen evolution reaction (HER)23 at all negative potentials, and a key criterion for selective catalysts in CO2 reduction will © 2012 American Chemical Society

Beyond the criterion of selectivity over the hydrogen evolution reaction, no systematic criteria for effective CO2 reduction catalysts have been proposed. It is the purpose of the present Perspective to develop such criteria. be comparatively poor HER activity in the presence of CO2. However, beyond this criterion of selectivity over the HER, no systematic criteria for effective CO2 reduction catalysts have been proposed. It is the purpose of the present Perspective to develop such criteria. Many of the trends in CO2 reduction by heterogeneous electrocatalysts can be observed from the pioneering experimental work of Hori and co-workers.2,10 In Table 1, we have highlighted a subset of these electrocatalysts. This table captures many of the major trends in CO2 reduction, including selectivity toward hydrocarbons (Cu and, to a lesser extent, Ni Received: November 4, 2011 Accepted: January 5, 2012 Published: January 5, 2012 251

dx.doi.org/10.1021/jz201461p | J. Phys. Chem. Lett. 2012, 3, 251−258

The Journal of Physical Chemistry Letters

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shown in Figure 1, which is bound by stoichiometry to involve the transfer of eight proton−electron pairs, resulting in seven adsorbed intermediates. In both the previous analysis and in the current work, we focus on the limiting potential (UL) at which each elementary step of a reaction becomes exergonic (or downhill in free energy) as a simple measure of the potential dependence of the electrochemical rate. For an example elementary hydrogenation reaction, A* + H+ + e− → B* (where A* and B* are generic adsorbed species), the limiting potential can be found as a function of the chemical potentials (μ) of the bound species

Table 1. Faradaic Yields in CO2 Reduction on FaceCentered Cubic (fcc) Metal Electrodes, As Reported by Hori et al.,10 for Experiments at 5 mA cm−2 Current Density in a 0.1 M KHCO3 Buffer at 18.5°C Faradaic yield, % electrode

V vs RHE

hydrocarbons/ organicsa

CO

HCOOH

H2

total

Ni Cu Pd Ag Pt Au

−1.09 −1.05 −0.81 −0.98 −0.68 −0.75

2.1 72.3 2.9 0.0 0.0 0.0

0.0 1.3 28.3 81.5 0.0 87.1

1.4 9.4 2.8 0.8 0.1 0.7

88.9 20.5 26.2 12.4 95.7 10.2

92.4 103.5 60.2 94.6 95.8 98.0

μ[B*] − μ[A*] − μ[H+ + e−](U = 0 VRHE) e 0V −ΔG elem = e

UL = −

a Predominantly CH4, but also including C2H4, C2H5OH, C3H7OH, C2H6, C3H5OH, CH3CHO, and C2H5CHO.

and Pd), the production of CO (Au and Ag), and the lack of strong catalytic activity in CO2 reduction, evolving H2 instead (Pt and Ni). These electrode materials are all late transition metals that can be represented in the face-centered cubic (fcc) crystal structure; this uniformity allows us to perform a systematic analysis of the factors determining the electrochemical activity and selectivity toward hydrocarbons. We focus here on the simplest hydrocarbon methane, but the analysis is quite general. Our analysis is based on an extensive set of density functional theory (DFT) calculations of the key adsorbates in CO2 reduction on these surfaces, conducted for this study. In previous work, we proposed19,20 a mechanism by which copper electrocatalysts reduce CO2 to methane; this route is

(1)

where e is the (positive) charge of an electron and the chemical potential of the proton−electron pair at 0 V can be found from the computational hydrogen electrode model.19,24 This is equivalent to the negative Gibbs free-energy change of the elementary reaction divided by the electronic charge. The chemical potentials of the bound species A* and B* can be found from standard statistical mechanics treatment of our calculated binding energies (EB) of these adsorbed species. It has been found that such an analysis correctly predicts the onset potential for the formation of different products in CO2 reduction over Cu surfaces19,20 and that it provides a good description of trends in reactivity of different metals for the oxygen reduction reaction.25

Figure 1. Adsorption energy scaling. The proposed pathway from CO2 to CH4 on copper surfaces is shown at the top of the figure, and the calculated adsorption energies of the key bound intermediates on fcc (211) facets are shown in the two lower figures. (More tightly bound adsorbates correspond to more negative binding energies.) The adsorption energies of those adsorbates binding to the surface through carbon can be correlated and are plotted against the binding energy of CO in the left plot. Similarly, the adsorption energies of those adsorbates binding through oxygen can be correlated and are plotted versus the binding energy of OH in the right plot. Ir and Rh are included as additional fcc metals to those shown in Table 1 in order to provide more points for the scaling relations. 252



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Figure 2. Limiting potentials (UL) for elementary proton-transfer steps in the mechanism of Figure 1. Each line is the calculated potential at which the indicated elementary reaction step is neutral with respect to free energy, as a function of the electrocatalysts’ carbon or oxygen affinity, EB[CO] or EB[OH]. Each elementary reaction involves the addition of a proton−electron pair (H+ + e−) to the reactant. The equilibrium potential for the overall electrochemical reduction of CO2 to CH4 is +0.17 V versus RHE, which is indicated on the figure. Thermodynamics limits any real catalyst to operate at potentials more negative than the equilibrium potential; therefore, the theoretical overpotential as a function of EB[CO] or EB[OH] can be represented by the distance between the equilibrium line and the most-negative limiting potential line (highlighted in gray). The elementary steps in this figure are limited to the steps in the mechanism of Figure 1 and exclude CH2O* → OCH3*, which has dependence on both EB[CO] and EB[OH]. The CH2O* → OCH3* transformation can be found in the Supporting Information. The EB[CO] and EB[OH] values for each metal are shown on the plot, along with their overall predicted value of UL based on these scaling relations.

CO2 → COOH*, are the most negative of the reaction steps shown in Figure 1; thus, for materials that follow this same reaction network, they would be expected to dictate the overpotential requirement of the reaction. Copper is seen to sit near the top of this “volcano”-type relation. A similar plot is shown on the right-hand side of Figure 2 for the steps that involve oxygen-bound intermediates. The elementary UL lines on this plot are much less negative; it is only for materials that have very high affinities for OH that the potential of clearing OH from the surface becomes comparable to the limitations set by the CO* → CHO* reaction. This does indicate that for materials that exhibit strong OH binding (more negative EB[OH]), OH removal will likely limit the catalysts’ effectiveness; however, for most materials, the potential-limiting step will be set by the CO* → CHO* transformation discussed above. A striking feature of Figure 2 is the severity of the limiting potential needed to protonate CO* to CHO*. This UL is not only the strongest limitation, it is also the least sensitive to changes in EB[CO]; that is, the line is nearly horizontal.

Given that there are seven intermediates in the reduction of CO2 to CH4, such an analysis involves an understanding of the problem in seven-dimensional space of adsorbate−surface interaction energies. However, as seen in Figure 1, roughly the first half of the intermediates interact with the catalyst surface through a carbon atom, and the second half interact through an oxygen atom. Scaling relations associated with the d-band theory of adsorption suggest there may be strong correlations within the carbon-bound and oxygen-bound species.26 Using this principle, we are able to successfully correlate the adsorption energies of the carbonbound species to EB[CO] and those of the oxygen-bound species to EB[OH], as shown in Figure 1. For example, the binding energy of CHO can be found from the linear equation in Figure 1

EB[CHO] = 0.88EB[CO] + 2.03 eV

(2)

These scaled affinity relations allow us to reduce the dimensionality of this reaction network from seven to two, thus making the search for trends within these transition-metal catalysts tractable. In this way, the limiting potentials for each of the elementary steps can be estimated for each surface as a function of that surface’s affinity for CO and OH. Full mathematical details are available in the Supporting Information. The elementary limiting potentials that scale with EB[CO] are shown in the left pane of Figure 2. Each solid line on this figure shows the limiting potential (UL) of an elementary protonation reaction for the reaction network of Figure 1. Because UL is the potential at which the elementary reaction becomes exergonic, each limiting potential gives a first-order indication of the electrical potential at which that elementary step begins to have an appreciable rate. The equilibrium potential is also plotted in Figure 2. The difference between each UL and the equilibrium potential gives a first-order estimate of the overpotential requirement for that elementary step; thus, the most negative UL line at any CO adsorption energy dictates the theoretical overpotential. This is indicated as the gray area of the plot. The two bottom-most UL lines, CO* → CHO* and

The magnitude and insensitivity of the energetics of this reaction to the electronic properties of the catalyst material can explain both the reason for the large overpotential requirement and the dearth of materials that have been found that can catalyze this reaction with lower overpotentials than Cu. 253



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that produce CHO*; the most limiting of these steps is shown, which is the protonation of CH* to form CH2*. For materials that form CHO*, other downstream pathways are possible; these later steps will be important to selectivity15 but should not affect the overall theoretical overpotential. The most significant is the subsequent protonation of CHO* to CH2O* or CHOH*. In both cases, this reaction appears to require less overpotential than the protonation of CO*. Over most of the region of the plot, the CHO* route is favored over the COH* route. However, at low values of EB[CO], the route to COH* begins to become possible. This upturn suggests the possibility that materials with a high affinity for carbon may have reduced overpotentials; however, some attributes of materials in this region may prevent this from being realized. Materials that bind carbon strongly will lead to high coverages when in actual electrochemical operation; this will weaken the binding strength of each incremental adsorbate, thus effectively shifting the binding strength of these materials to the right on the plot. Additionally, many of the materials that bind carbon strongly also bind oxygen strongly; therefore, removing OH from the surface may prove to be problematic, as seen for materials at the left edge of the EB[OH] plot in Figure 2. The experimentally observed trends in CO2 reduction can be deduced from Figure 3. Au and Ag sit to the right of the CO desorption line, agreeing with the experimental observations in Table 1, as well as numerous studies showing CO formation from CO2 on Au10,34−41 and Ag10,34,42,43 electrodes. Copper sits atop the volcano but still requires an extreme overpotential in order to overcome the CO to CHO hydrogenation step. The theoretically predicted overpotential is in agreement with experiment, as shown in earlier studies.19,20 Experimentally, the use of CO as a reactant has been found to produce a similar product spectrum as CO2, and the potential requirement to reach a current density of 5 mA cm−2 in 0.1 M KHCO3 buffer is nearly identical to that of CO2 (although canceling effects of an increased solution pH and decreased mass transport likely play a role),10,33 indicating that the potential-limiting step comes after the adsorption of CO on the catalyst surface. The materials in Figure 3 that are known to have significant activity in hydrogen evolution are marked with open circles; the HER would be expected to dominate over any CO2 reduction in these materials. However, in these materials, particularly Pt and Ni, the onset of H2 evolution is significantly higher than that in the CO2-free system. This can easily be understood as a poisoning of the catalyst surface by adsorbed CO. The CO is bound very tightly to these materials, and the only means of removing it under these conditions is to react it off, which requires significant potentials. Some of the CO can be seen to react off as CH4 in the case of Ni in Table 1, which was brought to potentials slightly more negative than that of Cu, in agreement with the more negative potentials needed by the volcano relation in Figure 3. In the case of Pt, the hydrogen evolution current density in Table 1 precluded the potential from being raised more negative than −0.68 V versus RHE; however, in experiments in which Pt was taken to more negative potentials, some hydrocarbons were observed.42 This is in agreement with a large body of experimental evidence, in which CO has been observed to poison the electrocatalyst surface in CO2 electroreduction experiments on Pt44−48 and Ni49,50 surfaces. Under these highly reducing conditions, Pd would be expected to form a hydride due to the high effective H fugacity,51

The magnitude and insensitivity of the energetics of this reaction to the electronic properties of the catalyst material can explain both the reason for the large overpotential requirement and the dearth of materials that have been found that can catalyze this reaction with lower overpotentials than Cu. The root cause of this insensitivity can be seen in Figure 1, in which the CO and CHO binding energies scale with a slope of 0.88, meaning that a surface that stabilizes CHO will stabilize CO by a similar amount. Thus, the binding energies are coupled, and the relative binding energies between the two change only slightly with catalyst material. Of course, electrocatalysts are not bound to follow the reaction mechanism shown in Figure 1. Other mechanistic possibilities are included in Figure 3, focusing on the mechanisms that

Figure 3. Limiting potentials (UL), allowing for competitive reaction mechanisms. Free-energy neutrality lines are plotted as in Figure 1, but multiple pathways are possible. Competitive pathways are shown as lines of the same color, and the more favorable route is shown as a solid line, while the less favorable route is shown dotted. Also shown is the binding energy at which adsorbed CO is in equilibrium with gaseous CO (at a partial pressure of 1%); the CO would be predicted to desorb at binding energies weaker than this line. As in Figure 1, the EB[CO] values are shown for each metal. The open circles correspond to surfaces that have an experimental exchange current density in the hydrogen evolution reaction greater than 10−4 A cm−2;27−29 in these materials, the hydrogen evolution reaction would be expected to dominate over CO2 reduction.

proceed through a CO* intermediate. (This is based on the experimental11,15,30−33 and theoretical19 findings that CO, rather than HCOOH, is an intermediate in hydrocarbon production on Cu; however, this does neglect the possibility that a formate intermediate could produce more reduced products than formic acid.) A vertical equilibrium line shows the binding energy at which gas-phase CO (at a partial pressure of 1%) is more stable than adsorbed CO*; materials to the right of this line would be expected to liberate any produced CO rather than further react it to hydrocarbons. CO* can be protonated to form COH* rather than CHO*, which may change the reaction pathway for the strongest binding materials. Lines for both of these reactions are shown in the same color, with the most favorable reaction at any value of EB[CO] shown as a solid line and the less favorable shown as a dotted line. For materials that protonate CO* to form COH*, the downstream pathway will be different than that for those 254



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which will likely weaken its affinity for carbon52 and shift it to the right (weaker CO binding) in Figure 3. Experimental evidence of this occurring can be seen in Table 1, in which a Faradaic balance was not closed. This behavior has been observed and exploited numerous times in electrochemical CO2 reduction.54−59 Nevertheless, this study captures the trends that Pd weakly produces CH4 but is in major competition with the hydrogen evolution reaction and the formation of a hydride.

Crucially, effective catalysts must be capable of efficiently catalyzing the protonation of adsorbed CO to adsorbed CHO or COH and exhibit simultaneous poor activity for the competitive hydrogen evolution reaction. Crucially, effective catalysts must be capable of efficiently catalyzing the protonation of adsorbed CO to adsorbed CHO or COH and exhibit simultaneous poor activity for the competitive hydrogen evolution reaction. Regardless of the downstream pathway after the formation of CHO or COH, products that are more reduced than CO, such as hydrocarbons and alcohols, will be produced as long as the subsequent products desorb from the electrocatalyst surface. We observe that Cu exhibits a slightly betterthan-average ability to perform the hydrogenation of adsorbed CO, but this is only in conjunction with its relatively poor activity at the hydrogen evolution reaction that this works, as exhibited by it having among the most negative potentials in Table 1. For an electrocatalytic system to be more effective at reducing CO2 to CH4, a means to protonate CO at a less negative potential must be found; in other words, the binding energy EB[CHO] must be strengthened (made more negative) relative to EB[CO]. For materials at the right side of the volcano of Figure 3, a possible strategy is to increase the CO(g) partial pressure, either by letting product gases build up or by employing a series of electrochemical reactors. However, for materials over the broadest range of carbon-binding energies, CO and CHO will both be bound to the surface, and their binding energies must be decoupled. We can use our understanding of surface reactivity to suggest general means in which the binding energetics of CO and CHO may be decoupled, possibly leading to electrocatalysts with substantially improved overpotentials relative to that of Cu. Several strategies are listed below and shown in Figure 4. • Alloying with metals with higher oxygen af f inity. CO tends to bind to surfaces in an upright geometry with respect to the carbon atom. In contrast, CHO tends to bind in a planar geometry with respect to the carbon atom, with the oxygen not extending directly away from the surface as is the case with CO*. (See also Figure 1.) Catalyst materials have varying degrees of affinity for oxygen and for carbon; alloying an element that has high oxygen affinity into an electrocatalyst may allow CHO* to bind to the surface through both the carbon and oxygen atoms, thus increasing its stability without affecting the stability of

Figure 4. Suggested decoupling strategies exploiting the geometric differences between adsorbed CHO and adsorbed CO.

CO*. However, the element with high oxygen affinity may change the potential-determining step, such that clearing of OH* from those sites may limit UL. Thus, a balance must be found between these two effects. • Ligand stabilization. In an analogous concept to alloying, an electrophilic chemical ligand can be applied in homogeneous catalysis, in which a ligand can be constructed that is geometrically capable of interacting with the planar geometry of CHO* but not that of the linear CO*. This partial bond formation with CHO* will act to stabilize the complex, thus lowering the energy of CHO* relative to that of CO*. Many other homogeneous catalysis techniques may prove beneficial in performing this reaction more effectively.60−62 • Tethering. Homogeneous and heterogeneous approaches can be combined through tethering of a ligand to an electrode surface, such that the adsorbates (CO* and CHO*) bind to the electrode surface, but only the planar CHO* can interact with the tethered ligand. The reverse is also possible, in which the adsorbates bind to the ligand and the CHO* interacts with the surface. • Addition of promoters. Promoters can change the relative binding strength of adsorbates through electronic effects 255



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(ligand effects, induced fields) and through structural effects (geometry disruption, binding through promoter atoms) or through combinations of these effects. It is likely that many promoted catalysts will not follow the same correlation of binding energies of CO and CHO, as is observed on the pure transition metals in this study. The addition of promoters that are resistant to electrochemical reduction may lead to new possibilities, and one such possibility for geometrical stabilization is shown in Figure 4, in which the promoter atom acts to make the adsorbed promoter− adsorbate complex adsorb in a bidentate manner. • Hydrogen bond stabilization/solvent ef fects. Due to the differences in binding geometries, hydrogen bond donors near the surface may interact with CO* and CHO* in different fashions. Also, a hydrogen bond acceptor may work to preferentially stabilize COH* (with a strong hydrogen-bonding group, OH) over CO*. The hydrogen binding difference may be achieved by changes in the electrolyte or by the presence of hydrogen-bonding ligands attached to the catalyst surface.

ASSOCIATED CONTENT

S Supporting Information *

The computational methodolgy, the additional scaling relations necessary to create Figure 3, a comparison to direct hydrogenation reactions, and a two-dimensional volcano plot illustrating transformation of CH2O* to OCH3*. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Biographies Andrew A. Peterson is an Assistant Professor at Brown University. He conducted this work while a postdoctoral scholar at Stanford. Andrew also worked with Jens Nørskov as a postdoc at the Technical University of Denmark and received his Ph.D. and bachelors from MIT and the University of Minnesota, respectively. Jens K. Nørskov is the Leland T. Edwards Professor in Chemical Engineering at Stanford University and is the Director of the SUNCAT Center for Interface Science and Catalysis at Stanford and the SLAC National Accelerator Laboratory. Previously, Jens directed the Center for Atomic-scale Materials Design at the Technical University of Denmark.

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REFERENCES

(1) Lewis, N. S.; Nocera, D. G. Powering the Planet: Chemical Challenges in Solar Energy Utilization. Proc. Natl. Acad. Sci. U.S.A. 2006, 103, 15729−15735. (2) Hori, Y. Electrochemical CO2 Reduction on Metal Electrodes. In Modern Aspects of Electrochemistry; Springer: New York, 2008; Vol. 42, Chapter 3, pp 89−189. (3) Benson, E. E.; Kubiak, C. P.; Sathrum, A. J.; Smieja, J. M. Electrocatalytic and Homogeneous Approaches to Conversion of CO2 to Liquid Fuels. Chem. Soc. Rev. 2009, 38, 89−99. (4) Roy, S. C.; Varghese, O. K.; Paulose, M.; Grimes, C. A. Toward Solar Fuels: Photocatalytic Conversion of Carbon Dioxide to Hydrocarbons. ACS Nano 2010, 4, 1259−1278. (5) Darensbourg, D. J. Chemistry of Carbon Dioxide Relevant to Its Utilization: A Personal Perspective. Inorg. Chem. 2010, 49, 10765− 10780. (6) Barton Cole, E.; Lakkaraju, P. S.; Rampulla, D. M.; Morris, A. J.; Abelev, E.; Bocarsly, A. B. Using a One-Electron Shuttle for the Multielectron Reduction of CO2 to Methanol: Kinetic, Mechanistic, and Structural Insights. J. Am. Chem. Soc. 2010, 132, 11539−11551. (7) Le, M.; Ren, M.; Zhang, Z.; Sprunger, P. T.; Kurtz, R. L.; Flake, J. C. Electrochemical Reduction of CO2 to CH3OH at Copper Oxide Surfaces. J. Electrochem. Soc. 2011, 158, E45−E49. (8) Rakowski Dubois, M.; Dubois, D. L. Development of Molecular Electrocatalysts for CO2 Reduction and H2 Production/Oxidation. Acc. Chem. Res. 2009, 42, 1974−1982. (9) Savéant, J. M. Molecular Catalysis of Electrochemical Reactions. Mechanistic Aspects. Chem. Rev. 2008, 108, 2348−2378. (10) Hori, Y.; Wakebe, H.; Tsukamoto, T.; Koga, O. Electrocatalytic Process of CO Selectivity in Electrochemical Reduction of CO2 at Metal Electrodes in Aqueous Media. Electrochim. Acta 1994, 39, 1833−1839. (11) DeWulf, D. W.; Jin, T.; Bard, A. J. Electrochemical and Surface Studies of Carbon Dioxide Reduction to Methane and Ethylene at Copper Electrodes in Aqueous Solutions. J. Electrochem. Soc. 1989, 136, 1686−1691. (12) Frese, K., Jr. Electrochemical Reduction of CO2 at Solid Electrodes. In Electrochemical and Electrocatalytic Reactions of Carbon Dioxide; Sullivan, B.; ; Krist, K., Guard, H., Eds.; Elsevier: New York, 1993; Chapter 6, pp 145−216. (13) Gattrell, M.; Gupta, N.; Co, A. A Review of the Aqueous Electrochemical Reduction of CO2 to Hydrocarbons at Copper. J. Electroanal. Chem. 2006, 594, 1−19. (14) Hori, Y.; Takahashi, R.; Yoshinami, Y.; Murata, A. Electrochemical Reduction of CO at a Copper Electrode. J. Phys. Chem. B 1997, 101, 7075−7081. (15) Schouten, K. J. P.; Kwon, Y.; van der Ham, C. J. M.; Qin, Z.; Koper, M. T. M. A New Mechanism for the Selectivity to C1 and C2 Species in the Electrochemical Reduction of Carbon Dioxide on Copper Electrodes. Chem. Sci. 2011, 2, 1902−1909. (16) Chaplin, R.; Wragg, A. Effects of Process Conditions and Electrode Material on Reaction Pathways for Carbon Dioxide Electroreduction with Particular Reference to Formate Formation. J. Appl. Electrochem. 2003, 33, 1107−1123. (17) Delacourt, C.; Ridgway, P. L.; Kerr, J. B.; Newman, J. Design of an Electrochemical Cell Making Syngas (CO + H2) from CO2 and H2O Reduction at Room Temperature. J. Electrochem. Soc. 2008, 155, B42−B49. (18) Whipple, D. T.; Kenis, P. J. A. Prospects of CO2 Utilization via Direct Heterogeneous Electrochemical Reduction. J. Phys. Chem. Lett. 2010, 1, 3451−3458. (19) Peterson, A. A.; Abild-Pedersen, F.; Studt, F.; Rossmeisl, J.; Nørskov, J. K. How Copper Catalyzes the Electroreduction of Carbon Dioxide into Hydrocarbon Fuels. Energ. Environ. Sci. 2010, 3, 1311− 1315. (20) Durand, W. J.; Peterson, A. A.; Studt, F.; Abild-Pedersen, F.; Nørskov, J. K. Structure Effects on the Energetics of the Electrochemical Reduction of CO2 by Copper Surfaces. Surf. Sci. 2011, 605, 1354−1359.

METHODS Electronic structure calculations were undertaken using the DACAPO plane wave implementation of density functional theory, using the (211) facet of a periodic fcc crystal and the RPBE exchange−correlation functional.63 Full details are available in the Supporting Information.

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ACKNOWLEDGMENTS

We appreciate insights from Professor Clifford Kubiak and Dr. Kyle A. Grice of the University of California at San Diego. This material is based on work supported by the Air Force Office of Scientific Research through the MURI program under AFOSR Award No. FA9550-10-1-0572. The authors also acknowledge support from the Catalysis for Sustainable Energy (CASE) initiative, which is funded by the Danish Ministry of Science, Technology, and Innovation. 256



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