Electrochemical Carbon Monoxide Reduction on Polycrystalline

Jul 18, 2018 - A 0.1 M solution of KOH (pH 13.0) prepared with 18.2 MΩ cm ... of the reduction of surface species from Cu2+ to Cu1+ and Cu1+ to Cu0 (...
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Research Article Cite This: ACS Catal. 2018, 8, 7445−7454

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Electrochemical Carbon Monoxide Reduction on Polycrystalline Copper: Effects of Potential, Pressure, and pH on Selectivity toward Multicarbon and Oxygenated Products Lei Wang,† Stephanie A. Nitopi,† Erlend Bertheussen,§ Marat Orazov,† Carlos G. Morales-Guio,† Xinyan Liu,† Drew C. Higgins,†,‡ Karen Chan,†,‡ Jens K. Nørskov,†,‡ Christopher Hahn,*,†,‡ and Thomas F. Jaramillo*,†,‡ Downloaded via TUFTS UNIV on July 19, 2018 at 06:50:48 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

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SUNCAT Center for Interface Science and Catalysis, Department of Chemical Engineering, Stanford University, Stanford, California 94305, United States § Department of Physics, Technical University of Denmark, 2800 Kgs Lyngby, Denmark ‡ SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, Menlo Park, California 94025, United States S Supporting Information *

ABSTRACT: Understanding the surface reactivity of CO, which is a key intermediate during electrochemical CO2 reduction, is crucial for the development of catalysts that selectively target desired products for the conversion of CO2 to fuels and chemicals. In this study, a customdesigned electrochemical cell is utilized to investigate planar polycrystalline copper as an electrocatalyst for CO reduction under alkaline conditions. Seven major CO reduction products have been observed including various hydrocarbons and oxygenates which are also common CO2 reduction products, strongly indicating that CO is a key reaction intermediate for these further-reduced products. A comparison of CO and CO2 reduction demonstrates that there is a large decrease in the overpotential for C−C coupled products under CO reduction conditions. The effects of CO partial pressure and electrolyte pH are investigated; we conclude that the aforementioned large potential shift is primarily a pH effect. Thus, alkaline conditions can be used to increase the energy efficiency of CO and CO2 reduction to C−C coupled products, when these cathode reactions are coupled to the oxygen evolution reaction at the anode. Further analysis of the reaction products reveals common trends in selectivity that indicate both the production of oxygenates and C−C coupled products are favored at lower overpotentials. These selectivity trends are generalized by comparing the results on planar Cu to current state-of-the-art high-surface-area Cu catalysts, which are able to achieve high oxygenate selectivity by operating at the same geometric current density at lower overpotentials. Combined, these findings outline key principles for designing CO and CO2 electrolyzers that are able to produce valuable C−C coupled products with high energy efficiency. KEYWORDS: carbon monoxide reduction, copper, selectivity, carbon monoxide, pH effect, overpotential

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The global demand for the further reduced (>2e−) products from CO2R, such as multicarbon hydrocarbons and/or alcohols, is much greater than the 2e− products CO and HCOO−.4 Of the metallic catalyst systems,5−7 Cu-containing materials are the only ones that can catalyze aqueous CO2R toward >2e− products in significant amounts.7−11 However, Cu-containing electrocatalysts require large overpotentials and are not particularly selective toward any single >2e− product, leading to poor energy efficiency for the reaction and possibly energy-intensive product separations. On the pathway to >2e−

INTRODUCTION Utilization of renewable energy sources to convert CO2 into value-added chemicals and fuels is an attractive strategy to not only mitigate the excess CO2 levels in the atmosphere, but also to convert renewable electricity into energy-dense liquid forms that are convenient for storage and transportation.1−3 Electrochemical processes in aqueous solution are promising for catalytic CO2 reduction (CO2R), since water is a sustainable electron and proton source.2 However, developing catalysts with high energy efficiency and product selectivity is still a major challenge, because of the complexity of the reaction networks involved. Thus, there is motivation to further understand the reaction mechanisms in order to guide the selectivity toward desired products. © XXXX American Chemical Society

Received: March 26, 2018 Revised: June 8, 2018

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ACS Catalysis products, CO2 is first reduced to CO and/or HCOO− through processes that require 2e−. A recent study determined the key thermodynamic descriptors to better understand the selectivity between these two initial pathways.12 While CO and HCOO− are both 2e− products, CO is the intermediate that has been shown to be subsequently reduced to >2e− products, while HCOO− appears to be a terminal pathway.13,14 Therefore, direct investigation of electrochemical CO reduction (COR) simplifies the reaction network that leads to the formation of further reduced CO2R products for a given catalyst. These factors suggest that a deeper understanding of COR on copper may reveal insights on how to develop more active and selective electrocatalysts for CO 2 R to high-value >2e − products. The discussion above indicates that there is a need for detailed studies of COR on Cu catalysts. The major challenge of COR is the low solubility of CO in aqueous electrolytes.15 This poor solubility leads to low mass-transport-limited current densities, complicating product detection and limiting the potential window for investigating the intrinsic reaction kinetics on a given electrocatalyst. COR on polycrystalline Cu (pc-Cu) has been previously conducted under very negative applied potentials and high total current densities, which suggests that the data may have been collected in a mass-transport-limited regime.16,17 While several other studies focused on single-crystal-based Cu electrodes for COR,18−20 comprehensive product quantification was not feasible, because of the limited electrode sizes. Oxide-derived Cu (OD-Cu) has recently drawn substantial attention, because it favors oxygenate production under COR conditions at relatively low overpotentials.21−23 In order to deconvolute the effects of intrinsic catalytic activity and mass transport in these highly complex, porous materials, understanding the behavior of planar copper under similar conditions is necessary. Thus, there is a need for more comprehensive and quantitative COR studies on pc-Cu in order to enhance our understanding of reaction mechanisms for COR and CO2R. In a previous study,24 we designed and implemented an electrochemical cell with a high ratio of electrode surface area to electrolyte volume, demonstrating that this parameter improves the overall product detection sensitivity. This enhanced analytical capability enabled the study of a broader range of potentials, and provided new mechanistic insights on CO2R on pc-Cu by revealing previously unquantified products. Here, we improve mass transport in this electrochemical cell by adding a glass frit to increase gas dispersion.25 We have used this modified version of our electrochemical cell to carefully examine COR on planar pc-Cu electrodes in 0.1 M KOH (pH 13.0) electrolyte. We chose this alkaline electrolyte, because it has been used elsewhere to ensure that the local pH on the electrode surface remains relatively stable during COR.21,26 In addition, COR studies in different pH conditions have indicated that alkaline conditions can promote the formation of C2+ (multicarbon-based compounds) products.17,19 Similarly, recent advancements implementing vapor-fed electrolyzers have enabled CO2R studies in alkaline electrolytes, and have further demonstrated that alkaline conditions can improve the energy efficiency for CO2 conversion to C2+ products.27,28 With the aforementioned measurement platform, we are able to provide the most comprehensive data to date of COR on pc-Cu. Our results show that pc-Cu reduces CO to oxygenates and hydrocarbons with Faradaic efficiencies of up to 65% at a low overpotential of −0.63 V vs RHE. A

comparison of COR and CO2R on pc-Cu reveals a large decrease in the overpotential for C2+ products for COR. By varying the CO partial pressure and the electrolyte pH, we determine that the major factor in the potential shift is due to the pH. More importantly, two distinct selectivity trends common to both COR and CO2R were uncovered through examining our own results in context with high-surface-area catalysts from the literature: lower overpotentials result in higher selectivity for C2+ products and oxygenates. Based on these results, we establish a deeper understanding of how the electrode potential and pH can be used to guide COR and CO2R selectivity to C−C coupled oxygenates.

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EXPERIMENTAL SECTION Materials. Potassium hydroxide (semiconductor grade, 99.99%) was purchased from Sigma−Aldrich; ortho-phosphoric acid (85% in water) was purchased from Fisher Scientific; carbon monoxide (99.999%) and customized CO/ Ar gas mixtures were purchased from Advanced Specialty Gases; copper foil (99.9999%, 0.1 mm thick) was purchased from Alfa Aesar. All chemicals were used without further purification. Electrolyte solutions were prepared with deionized water that was purified using a Thermo Scientific Barnstead Nanopure water purification system (resistivity of 18.2 MΩ cm). Electrochemical Measurements. All electrochemical measurements were carried out on a Bio-Logic VMP3 multichannel potentiostat/galvanostat with a built-in EIS analyzer under ambient pressure. A previously described polycarbonate electrochemical cell24 was modified with a fritted glass gas dispersion tube in order to enhance CO dispersion.25 A three-electrode system was introduced into the aforementioned modified cell: (1) polycrystalline Cu foils were used as working electrodes after mechanical polishing and then electropolishing (85% ortho-phosphoric acid, 2.1 V vs graphite), with a 5.9 cm−2 geometric area exposed to the electrolyte; (2) a platinum foil was used as the counter electrode with the same geometric area exposed to the electrolyte; (3) an Ag/AgCl electrode (Accumet) was introduced in the working electrode compartment as the reference electrode through a Luggin capillary. An anion exchange membrane (Selemion AMV, AGC, Inc.) was used between the working electrode and counter electrode compartments. The distance between the working and reference electrodes was kept small (0.5 cm) to reduce solution resistance. After the cell assembly, humidified CO was flowed through the working electrode compartment, with the flow rate regulated by a mass flow controller (πMFC, MKS Instruments) at 20 sccm, and then vented directly into the gas sampling loop of a gas chromatograph (SRI Instruments, Model 8610C); argon gas with the same flow rate was flowed through the counter electrode compartment and then vented to exhaust directly. A 0.1 M solution of KOH (pH 13.0) prepared with 18.2 MΩ cm deionized water was used as the electrolyte. Each compartment contained 10 mL of 0.1 M KOH electrolyte, and the solution in the cathodic compartment was purged with CO for 20 min before the start of electrolysis. A 0.1 M boric acid solution (pH adjusted by 0.1 M KOH) was used for the pH 7 COR experiments. A bipolar membrane (Fumasep FBM, FuelCellStore) was used for these experiments to maintain a stable bulk pH in the electrolyte. 7446

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ACS Catalysis Product Quantification. Product identification was conducted according to protocols described in detail in our previous work.24 For each measured cathode potential, 1 mL of reactor exhaust was injected into a GC (SRI 8610C) during the chronoamperometry (CA) measurements at times of 5, 23, 41, and 59 min, to analyze the concentration of products in the gas phase. The average current density for the 1 min window before each GC injection was used for analysis. Average current efficiencies and partial current densities over the course of 1 h were obtained by averaging the results from individual injections. After CA measurements, the electrolyte was purged with Ar for at least 15 min, and then the liquid phase products were analyzed by using water suppression 1D 1H NMR (600 MHz, Varian Inova). A 700 μL aliquot of post-reaction electrolyte was mixed with 35 μL of internal standard (10 mM DMSO and 50 mM phenol in D2O solution) for quantification. Samples were shipped to quantify aldehydes by static headspace-gas chromatography (HS-GC). We expect the aldehyde concentrations to be an underestimation, because of the aging of the samples.23 The entire protocol was repeated three times for each potential to establish statistical significance of the data (see Table S1 in the Supporting Information).

contaminants; thus, if any contaiminants are present, they are below the detection limits (see Figures S3−S5 and Note 1 in the Supporting Information). This strongly suggests that the performance of Cu is obtained in a clean environment, and reflects the catalytic behavior of Cu instead of contaminants. To assess the COR selectivity and activity of pc-Cu, the gasand liquid-phase products were quantified by using a combination of gas chromatography (GC), and nuclear magnetic resonance (NMR) spectroscopy, respectively. The aldehydes were quantified using static headspace−gas chromatography (HS-GC).32 Faradaic efficiencies (FE) for all detected products show the selectivity of pc-Cu for COR (see Figure 1a). Methane, ethylene, ethanol, acetate, n-

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RESULTS AND DISCUSSION pc-Cu COR in 0.1 M KOH. Cyclic voltammetry (CV) (see Figure S1 in the Supporting Information) was used to assess the overall current−potential profiles of pc-Cu electrodes across a wide range of potentials in both CO- and Ar-purged 0.1 M KOH electrolyte. Freshly prepared Cu electrodes in both purge environments show reduction peaks at approximately 0.1 V and −0.1 V (vs RHE), which are characteristic of the reduction of surface species from Cu2+ to Cu1+ and Cu1+ to Cu0 (see Figure S1).29 These features are not present in subsequent potential sweeps between −0.05 V and −0.8 V (vs RHE), indicating that the surface oxide is reduced during the first sweep. The oxide was likely formed after electropolishing from air exposure during the cell assembly process. In contrast to previous studies of oxide-derived Cu,21 these oxide layers are expected to be thin (2e− CO2R products.8,16,17 The lower FEs of aldehydes that are observed here, relative to other studies done under similar conditions, can be attributed to spontaneous reactions under alkaline conditions from aging of the electrolyte.33 A small portion of the Faradaic efficiency is unaccounted for at low overpotentials, which is likely due to the reduction of CuO and/ or product concentrations that are below the detection limits of NMR/GC. Nevertheless, pc-Cu in our system exhibits a maximum FE of 65% toward COR at −0.63 V, which is the highest selectivity for COR among the reported pc-Cu catalysts.10,17 This high selectivity toward COR could be due to the improved hydrodynamics in our electrochemical cell, since the glass frit improves CO mass transport to the electrode by dispersing small CO bubbles.25 Partial current densities, as a function of applied potential, reveal trends in activity toward the different COR products (see Figure 1b). Ethylene and ethanol are the major products at potentials more negative than −0.55 V, with their partial 7447

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ACS Catalysis current densities plateauing at −0.63 V, where the previously discussed CV analyses suggest that there is a mass transport limitation of CO (see Figure S1). At more-negative potentials where the supply of CO is mass-transport-limited, the methane activity increases and there is a corresponding decrease in the activity to C2+ products. Note that more-positive onset potentials for product detection are observed for all of the C2+ products than for the C1 product methane, suggesting that the C2+ and C1 pathways likely diverge at an early stage in COR. These trends are corroborated by previous theoretical work that suggests C−C coupling steps are favored over CO hydrogenation at low overpotentials.34,35 However, methane production occurs only at potentials where CO is masstransport-limited. Therefore, it is difficult to attribute the selectivity and activity changes to the intrinsic reaction kinetics or a CO mass-transport phenomenon. Among the C2+ products, aldehydes are detected at more positive potentials than the corresponding alcohols with the same number of carbons. Partial current densities for the aldehydes increase with more negative potentials until approximately −0.55 V, where alcohols are detected. This trend suggests that the aldehydes could be intermediates for the corresponding alcohols.17 In addition to the aldehydes and alcohols, acetate was also detected as a minor oxygenate product. The acetate formation mechanism is currently unclear; however, one possibility may be through homogeneous chemical reactions in the alkaline electrolyte, such as Cannizzaro-type disproportionation.36 Overall, these trends demonstrate that the C2+ products are favored over the C1 product methane at morepositive electrode potentials, highlighting the mechanistic differences for the C1 and C2+ pathways. Comparison of COR and CO2R on pc-Cu. Additional mechanistic insights are revealed by comparing COR in 0.1 M KOH and CO2R 0.1 M KHCO3 on pc-Cu. When comparing different reaction conditions, it is critical to identify the potential ranges where the activity is limited by the intrinsic activity of the catalyst or the mass transport of reactant species to the electrode surface. To this end, the total molar CO/CO2 reduction rates were evaluated (see Figure S6 in the Supporting Information). Based on this analysis of COR, potentials more negative than −0.63 V are clearly limited by mass transport of CO, corroborating the aforementioned conclusions from CV measurements. Thus, the potential range of −0.40 V to −0.63 V is identified to be the region where the COR activity is more influenced by the intrinsic reaction kinetics on pc-Cu. A similar analysis for CO2R data shows that the corresponding region for CO2R is between −0.65 V and −1.10 V. To compare activity, the partial current densities of the major products of COR were plotted together with previous CO2R data (see Figure 2).24 Across a wide range of potentials, the partial current densities for these products exhibit strikingly similar trends in these two different studies, indicating shared mechanistic pathways for these two reactions and further demonstrating that CO is a primary intermediate for >2e− products in CO2 reduction.37 Despite the similarities in reaction products, a major difference between COR and CO2R on pc-Cu is the significantly more-positive onset potentials for COR. The onset potentials are shifted ∼0.4 V more positive for C2+ products and ∼0.15 V for the C1 product methane, indicating that there is a major improvement in energy efficiency under COR conditions, if these cathode reactions are coupled to oxygen evolution reaction (OER) at the anode. We hypothesize that the primary factors that lead to

Figure 2. Partial current densities of individual products: CO2 reduction (solid symbols) and CO reduction (open symbols), both on pc-Cu.

the significant onset potential shift are differences in the CO coverage on the catalyst surface and/or the pH of the electrolyte. Pressure-Dependent COR on pc-Cu. In order to determine if the CO coverage is a major contributor to the aforementioned shift in onset potential, the reaction kinetics were further investigated by using different partial pressures of CO with CO/Ar gas mixtures. As expected, cyclic voltammograms show that a decrease in the CO partial pressure correspondingly shrinks the previously described COR shoulder at −0.63 V (see Figure S7 in the Supporting Information). As the CO partial pressure is decreased to 0.1 and 0.01 bar, the contribution of COR to the total current density becomes minimal, compared to the HER, and the shoulder is not observed. This trend is also corroborated by 1 h electrolysis experiments, which show a corresponding decrease in the COR rates with the CO partial pressure (see Figure S8 in the Supporting Information). In addition, the COR rates appear to plateau at more-positive potentials with decreasing CO partial pressure, because of increased mass-transport limitations. A limited amount of information on COR reaction kinetics has been obtained from previous studies, because these mass-transport limitations have led to intrinsic challenges in product identification and quantification.7,17−22,38−40 In this study, we have been able to detect methane, ethylene, and ethanol as major products at all investigated partial pressures, because of the high detection sensitivity enabled by our custom electrochemical cell (see Figure 3). While there is a major decrease in partial current densities of the detected products

Figure 3. Partial current densities for methane, ethylene, and ethanol at various CO partial pressures. 7448

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ACS Catalysis with the decrease of CO partial pressure, the potentialdependent trends are similar, indicating that the primary factor in the aforementioned shift in onset potential is unlikely to be from a difference in CO coverage on the electrode surface. Further examination of the 1 h electrolysis data (Figure 3) reveals mechanistic insights on the reaction kinetics for the COR pathways. While it is difficult to extract meaningful information on the kinetics for the C1 pathway, since methane formation only occurs at potentials within a CO masstransport-limited regime, the C2+ products are clearly more sensitive to the CO partial pressure than the C1 product methane. This experimental evidence may support previous theory studies that suggest a second-order dependence of the reaction rate to C2+ products on the coverage of adsorbed CO.35,41 Quantitative analysis of the changes in partial current density that are due to pressure indicates that ethylene follows second-order reaction kinetics when the CO partial pressure is decreased from 1.0 bar to 0.5 bar but switches to first-order reaction kinetics at lower pressures (0.1 and 0.01 bar). In addition, a comparison of the partial current densities for ethylene and ethanol demonstrates that ethanol is less sensitive to the decrease in CO pressure. A possible explanation for these trends could be that the heterogeneity of surface sites on pc-Cu leads to a multicomponent Langmuir isotherm, since step and kink sites will bind CO stronger than terrace sites. Competitive adsorption leads to the CO coverage saturating on the undercoordinated sites at lower partial pressures than on the terrace sites. Thus, the partial current densities for COR to C−C coupled products are less sensitive to changes within the low end of the CO partial pressure range chosen for this study. These trends are consistent with previous work that suggests Cu step sites have COR activities that are several orders of magnitude higher than that of terrace sites.41 To an extent, the combined observations support the previous hypothesis that step sites have an integral role on the kinetics of C−C coupling and oxygenate formation on multisite catalysts such as pc-Cu,42,43 motivating further work to provide direct evidence of the difference in the catalytic behavior of step and terrace sites. The pH Effect. Another hypothesis for the shift in onset potential between CO2R and COR is a pH effect, since the CO2R and COR data were collected at pH 7 and 13, respectively. To investigate whether there is a pH effect, Figure 2 is replotted using the standard hydrogen electrode (SHE) scale (see Figure 4, as well as Figure S9 in the Supporting Information). On the SHE scale, the overpotentials for C2+ products in COR are similar to that of CO2R, indicating that the rate-limiting step for C2+ formation is pH-independent of an absolute potential scale. This invariance in the overpotential with pH will henceforth be described as a pH independence, that is based on a given reference scale (SHE or RHE). Note that acetate is an outlier that still shows a positive shift in onset potential on the SHE scale for COR, indicating that the ratedetermining step for acetate may be different from the other C2+ products. In addition, higher partial current densities for acetate under COR conditions suggest that alkaline conditions are more favorable for producing acetate. In comparison to the C2+ products, methane shows a smaller shift of ∼0.15 V on the RHE scale, indicating that the C1 pathway has a more complex pH dependence (vide infra). We hypothesize that the overall data are representative of the pH dependence of COR. In order to validate this hypothesis, we also investigated COR in a pH 7 borate solution using the same pc-Cu. As shown in

Figure 4. Partial current densities of individual products: (solid) CO2 reduction and (dashed) CO reduction, on a SHE scale.

Figure S10 in the Supporting Information, it is clear that the potential dependence of C1 and C2+ products are very similar for COR and CO2R in near-neutral electrolytes, confirming the aforementioned pH effects. In addition, previous work suggested that the CO2 to CO steps are not rate-limiting on Cu,44 which is also confirmed by the fact that similar onset potentials were obtained for CO2R and COR in the same electrolyte (see Figure S10). Overall, we conclude that the significant shift in onset for C2+ products for CO2R, compared to COR on the RHE scale, is primarily due to a pH effect. A pH independence for the C2+ pathway on the SHE scale can be rationalized through several possible rate-limiting processes, including the following: slow electron transfer rates in comparison to that of proton transfer, field stabilization of a chemical reaction step (e.g., for CO dimerization), or proton−electron transfer from water. While it is common for electron transfers to be rate-limiting for reactions in solution, processes within close vicinity of a metal surface typically have very high electron transfer rates, which are therefore not ratelimiting.45 Thus, it is unlikely that a rate-limiting electron transfer process is leading to the aforementioned pH effect. Instead, one possible origin of this pH independence could be a rate-limiting chemical reaction step, such as C−C coupling through CO dimerization, that is strongly promoted by field stabilization from electrolyte cations in the Helmholtz plane. A previous study hypothesized this would lead to a pH independence on the SHE scale,34 since the potential of zero charge and charging of the electrical double layer of a metal is dependent on the absolute potential. Another possible explanation for the pH independence on the SHE scale is a proton−electron transfer from water. The potential and pH dependence of proton−electron transfers from protons and water can be derived from the chemical potentials of electrons and the ions involved. As an example, consider the Volmer reaction with both protons and water as the proton source: H+ + e− + * → H* H 2O + e− + * → H* + OH−

By calculating the free energy of the species in the above reactions (see Note 2 in the Supporting Information), the impact of pH on proton-coupled reaction barriers on an RHE and SHE scale can be illustrated in Figure 5. Generally, reaction energies (see eqs S5−S8 in the Supporting 7449

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chemical reaction and proton−electron transfer from water during the rate-limiting step are plausible explanations for the experimentally measured pH effect. A simple kinetic analysis, using a single electrochemical ratelimiting step, can elucidate how early in the reaction this step occurs for the C1 and C2+ products. Assuming a Boltzmann distribution for the adsorbate coverages leading up to the ratelimiting step, coverages resulting from the n previous proton− electron transfer steps will be pH-independent of the RHE scale, while the barriers will remain pH-independent of the SHE scale. For a rate-limiting step involving proton−electron transfer to adsorbate X, the rate has the following form: ij G 0 + eβUSHE yz zz R = PθX* expjjj− a zz j kT k {

where P is the reaction pre-exponential factor, θX* is the coverage of adsorbate X, USHE is the potential with respect to the SHE, G0a is the activation energy under standard conditions (where USHE = 0 V), k is the Boltzmann constant, T is the reaction temperature, and β is the symmetry factor (which is the amount of charge transferred to the adsorbate at the transition state and gives the potential dependence of the barrier for the rate-limiting step).48,49 We assume that X follows a Boltzmann distribution:

Figure 5. Free energy barriers of acidic (solid) and alkaline (dashed) proton−electron transfer barriers on SHE and RHE scales. In SHE scale, an increased pH leads to an increased acidic barrier, while the alkaline barrier is unaffected by pH. While reaction energies by definition are invariant with the potential vs RHE, an increased pH, but constant potential vs RHE will lead to an increase in acidic barriers and a decrease in barriers from water.

i enURHE zy zz θX* = θA* expjjj− kT { k

Information) show a pH independence on the RHE scale, since shifts in pH are, by definition, balanced by shifts in absolute potential (i.e., 2.3kTΔpH = −eΔU). However, activation energies are affected differently since the configurational entropy of the H+ and OH− species in bulk solution do not play a role in the transition states, which are affected only by the absolute potential.45 At a fixed potential vs RHE, an increase in pH will either increase or decrease the activation energy when the proton transfer is from H+ or water, respectively (see Figure 5, as well as eqs S10 and 12 in the Supporting Information). At a fixed potential vs SHE see (Figure 5, as well as eqs S9 and S11 in the Supporting Information), activation energies for proton transfer from H+ increase with pH, while those from water are unaffected, since a pH shift affects only the free energy of the final state for a proton transfer from water. Under alkaline conditions, water serves as the primary source for proton-coupled electrochemical reactions, because of the low proton concentration.46,47 Thus, we conclude that both field stabilization of a

where θA* represents the coverage of a preceding intermediate that shows no potential dependence (e.g., *CO), and n is the number proton−electron transfer steps from A to X. Therefore, the rate can be written as a function of URHE and pH as follows: ÄÅ ÉÑ ÅÅÅ Ga0 + (n + β)eURHE ÑÑ R = θA*P expÅÅÅ− + 2.3β pHÑÑÑÑ ÅÅÅ ÑÑÑ kT Ç Ö Therefore, the corresponding shift in overpotential with a β change in pH is −2.3kT ΔpH n + β on a RHE scale. Using a simple case, where water is the proton source and the ratelimiting step is the first proton−electron transfer (n = 0), a reaction pathway will show a pH independence on the SHE scale, since the shift in overpotential on the RHE scale will simplify to −2.3kTΔpH. Thus, the experimentally observed

Figure 6. Selectivity of Cu toward CO2R and COR: (a) oxygenate/hydrocarbon ratios of the reduction products as a function of potential, and (b) Faradaic efficiencies for C1, C2, and C3 products, each as a function of potential. 7450

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dimerization, when compared to *CO hydrogenation, could lead to this selectivity enhancement. We also note that the selectivity of pc-Cu for C2+ products is greatly improved under COR conditions, which is most likely due to the aforementioned combination of a higher pH and CO coverage. Thus, the combined results indicate that low overpotentials favor C2+ products that are oxygenated, highlighting an effective design principle for both COR and CO2R systems to yield liquid-phase products. Normalized Activity for COR on Cu. In order to gain a deeper understanding of the intrinsic reaction kinetics on different Cu-based materials, we have compared the results from this work with other published data for COR and CO2R on Cu-based electrodes by normalizing the geometric current densities with the reported roughness factor (RF) for each electrode surface (see Figure 7, as well as Figure S12 in the

pH independence for C2+ products on the SHE scale suggests that if the rate-limiting step involves an electrochemical reaction with water as the proton donor, the reaction is limited by the first proton−electron transfer step. On the other hand, a pathway that is limited by a later proton−electron transfer (n > 0) will exhibit a lower Tafel slope and a morecomplex pH dependence.49 Therefore, the later the ratelimiting proton−electron transfer step, the lower the Tafel slope and the larger overpotential shift on the RHE scale. A previous study reported this catalytic behavior for COR to CH4, and proposed that the rate-limiting step for the C1 pathway arises from the second or third proton electron transfer.49 Qualitatively, the same trends are observed in our study as both the Tafel slope and the overpotential shift on the RHE scale are smaller for the C1 pathway, in comparison to those of the C2+ pathways. As previously discussed, a more quantitative investigation of the C1 pathway will require improved CO transport to the electrode, since mass-transport limitations convolute analysis of the intrinsic reaction kinetics. Overall, the pH effect highlighted in this section strongly motivates the utilization of alkaline conditions for both COR and CO2R, since increasing the pH is a simple and effective way to increase the overall energy efficiency to C−C coupled products. Potential-Dependent Selectivity to Oxygenates and C2+ Products. Additional mechanistic insights are gained by comparing the COR and CO2R selectivity of pc-Cu for >2e− CO2R products (see Figure 6). It is useful to group the >2e− CO2R products into the categories of oxygenates and hydrocarbons (see Figure 6a, as well as Figure S11 in the Supporting Information), since hydrocarbons generally require additional C−O bond scission and H+/e− transfer steps. A comparison of the oxygenate/hydrocarbon ratios for COR and CO2R reveals a common potential-dependent trend, where a higher selectivity for oxygenates is observed at lower overpotentials. This trend could be due to reaction intermediates being less polarized at more-positive potentials, resulting in a smaller driving force for C−O bond scission and thus a higher selectivity for oxygenates. In addition, significantly higher oxygenate to hydrocarbon ratios are observed for COR, when compared to CO2R. Previous work suggested that selectivity toward oxygenates and hydrocarbons is very likely related to the surface coverage of CO* and H*, with high H* coverages favoring hydrocarbons and low H* favoring oxygenates.42 Therefore, the higher oxygenate to hydrocarbon ratios in the case of COR could be due to a higher coverage of CO*, leading to fewer active sites for H* under these conditions. However, we do not exclude the possibility of base-catalyzed chemical reactions having an impact on the oxygenate/hydrocarbon ratios due to the alkaline conditions used for COR. Nevertheless, the common potential-dependent trend for both CO 2 R and COR demonstrates that low overpotentials are a design principle for improving not only energy efficiency but also selectivity to valuable oxygenated products. In order to investigate the C−C coupling selectivity of CO2R and COR, the Faradaic efficiencies for >2e− products are grouped by the number of carbons within a given product into the categories C1, C2, and C3 (Figure 6b). Similar to the trend for oxygenate/hydrocarbon ratios, selectivity for C−C coupled (C2+) products is enhanced over the C1 product methane at more-positive potentials for both COR and CO2R. As previously discussed, a lower barrier for surface *CO

Figure 7. Normalized partial current densities for Cu, CO2R, and Cu COR: [1] ref 24; [2] this work; [3,4] refs 21 and 22; [5,6] refs 10 and 16; and [7] ref 50. Solid symbols represent CO2R data while hollow symbols represent COR data; blue symbols represent partial activities for >2e− C1 products, while red symbols represent partial activities for >2e− C2+ products. All of the current densities have been normalized based on the reported roughness factor of the corresponding electrodes. The black dashed line indicates the similar Tafel slopes for different Cu materials.

Supporting Information, on the RHE and SHE scales, respectively). When comparing the potential ranges where the intrinsic reaction kinetics is expected to dominate the current density (i.e., before the COR or CO2R current plateaus on the Tafel plots due to mass-transport limitations) for high surface area Cu (e.g., OD-Cu21,22,50) and planar pc-Cu,10,16,24 the similar slopes indicate that both low- and high-RF materials exhibit comparable intrinsic activities for COR and the formation of >2e− products in CO2R. This suggests that the active site densities are very similar between high surface area and planar Cu catalysts, regardless of what types of active sites are contributing to the COR activity.41,42,51−54 While increasing the RF does not have a large impact on the intrinsic activity of Cu catalysts, major differences in selectivity toward oxygenates are reported between high surface area Cu and planar pc-Cu.21,24,42 As discussed in the previous section, higher oxygenate/hydrocarbon ratios are favored at lower overpotentials. Therefore, it is reasonable to assume that a large part of the selectivity difference between high surface area Cu and planar pc-Cu is due to the two electrocatalysts operating at different applied potentials. Our conclusions also 7451

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agree with a very recent work that has uncovered similar surface area effects on CO2R.55 In order to confirm this hypothesis, the potential dependence of oxygenate to hydrocarbon ratios on both metallic and flat polycrystalline Cu foils and OD-Cu have been compared in Figure S13 in the Supporting Information. We note that the potential-dependent trends are closely matched to each other, further indicating again that the absolute applied potential is a major factor in the selectivity between oxygenates and hydrocarbons on Cu-based catalysts. These results suggest that, while planar pc-Cu and high surface area Cu have a similar intrinsic activity toward CO2R and COR, high-surface-area Cu enables operation with the same geometric current density at lower overpotentials where Cu is more selective for oxygenated products. Thus, the conclusions from the previous38,42 and present work indicate that high undercoordinated site density, alkaline electrochemical conditions, and operating at more positive potentials are all design principles to steer selectivity to C2+ oxygenated products. This motivates the rational design of new nanostructured hierarchical electrode morphologies that are able to combine all of these guidelines to ultimately enable highenergy-efficiency CO2 and CO electrolyzers.

Research Article

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acscatal.8b01200. Supplemental figures (Figures S1−S13) and data (Table S1); supplemental discussions regarding the physical and chemical integrity of the Cu surface during electrocatalysis (Note 1) and the free-energy analysis of the Volmer reaction with both protons and water as the proton source (Note 2) (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected] (C. Hahn). *E-mail: [email protected] (T. F. Jaramillo). ORCID

Carlos G. Morales-Guio: 0000-0002-5840-5591 Drew C. Higgins: 0000-0002-0585-2670 Karen Chan: 0000-0002-6897-1108 Christopher Hahn: 0000-0002-2772-6341 Thomas F. Jaramillo: 0000-0001-9900-0622

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CONCLUSIONS Using a custom-designed electrochemical cell, polycrystalline Cu has been carefully studied as a catalyst for electrochemical CO reduction under alkaline condition. High Faradaic efficiencies for COR were obtained due to the improved hydrodynamics in this cell, e.g., over 65% at −0.63 V, and the major products were determined to be ethylene and ethanol. A comparison between COR and CO2R data for pc-Cu showed a substantial positive shift in overpotential for COR to C2+ products. After investigating the hypotheses of a CO partial pressure or a pH effect, it was determined that the shift in overpotential is primarily due to the different operating pH conditions. While the CO partial pressure does not have a strong influence on the overpotential for COR, the relative insensitivity to CO partial pressure for ethanol, in comparison to ethylene, corroborates previous reports that indicate step sites are selective for oxygenates. Further comparison of the product distributions for COR and CO2R on a SHE scale reveals that the rate-determining steps of C2+ pathways are independent of the pH on an absolute potential scale. These findings suggest alkaline pH conditions as a design principle for increasing the overall energy efficiency to C−C coupled products for future CO2R/COR applications, when coupled to OER. In addition to the pH effect, common potentialdependent trends are revealed for both CO2R and COR on pc-Cu, demonstrating that C2+ products and oxygenates are preferentially formed at low overpotentials. By normalizing to the RF of the catalyst, we show that pc-Cu has a similar normalized activity as previously reported high-surface-area Cu-based catalysts, indicating that the surface fraction of active sites on these catalysts are similar regardless of the morphology. While all Cu catalysts exhibit similar normalized activities, high-surface-area catalysts can operate at higher geometric current densities in a potential window where the formation of valuable oxygenated products is favorable due to the aforementioned trend in increased oxygenate selectivity at lower overpotentials. Overall, our work outlines several rational design strategies to enable the development of CO2 and CO electrolyzers that can generate multicarbon liquid products with high energy efficiency.

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This material is based upon work performed by the Joint Center for Artificial Photosynthesis, a DOE Energy Innovation Hub, supported through the Office of Science of the U.S. Department of Energy under Award No. DE-SC0004993. Part of this work was performed at the Stanford Nano Shared Facilities (SNSF), supported by the National Science Foundation, under Award No. ECCS-1542152. Additional thanks go to the Stanford NMR Facility. L.W. thanks the Knut & Alice Wallenberg Foundation for the financial support.

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REFERENCES

(1) Costentin, C.; Robert, M.; Saveant, J.-M. Catalysis of the Electrochemical Reduction of Carbon Dioxide. Chem. Soc. Rev. 2013, 42, 2423−2436. (2) Gattrell, M.; Gupta, N.; Co, A. Electrochemical Reduction of CO2 to Hydrocarbons to Store Renewable Electrical Energy and Upgrade Biogas. Energy Convers. Manage. 2007, 48, 1255−1265. (3) Appel, A. M.; Bercaw, J. E.; Bocarsly, A. B.; Dobbek, H.; DuBois, D. L.; Dupuis, M.; Ferry, J. G.; Fujita, E.; Hille, R.; Kenis, P. J. A.; Kerfeld, C. A.; Morris, R. H.; Peden, C. H. F.; Portis, A. R.; Ragsdale, S. W.; Rauchfuss, T. B.; Reek, J. N. H.; Seefeldt, L. C.; Thauer, R. K.; Waldrop, G. L. Frontiers, Opportunities, and Challenges in Biochemical and Chemical Catalysis of CO2 Fixation. Chem. Rev. 2013, 113, 6621−6658. (4) Shaner, M. R.; Atwater, H. A.; Lewis, N. S.; McFarland, E. W. A Comparative Technoeconomic Analysis of Renewable Hydrogen Production Using Solar Energy. Energy Environ. Sci. 2016, 9, 2354− 2371. (5) Hahn, C.; Abram, D. N.; Hansen, H. A.; Hatsukade, T.; Jackson, A.; Johnson, N. C.; Hellstern, T. R.; Kuhl, K. P.; Cave, E. R.; Feaster, J. T.; Jaramillo, T. F. Synthesis of Thin Film AuPd Alloys and Their Investigation for Electrocatalytic CO2 Reduction. J. Mater. Chem. A 2015, 3, 20185−20194. (6) Torelli, D. A.; Francis, S. A.; Crompton, J. C.; Javier, A.; Thompson, J. R.; Brunschwig, B. S.; Soriaga, M. P.; Lewis, N. S. Nickel−Gallium-Catalyzed Electrochemical Reduction of CO2 to Highly Reduced Products at Low Overpotentials. ACS Catal. 2016, 6, 2100−2104. 7452

DOI: 10.1021/acscatal.8b01200 ACS Catal. 2018, 8, 7445−7454

Research Article

ACS Catalysis (7) Jovanov, Z. P.; Hansen, H. A.; Varela, A. S.; Malacrida, P.; Peterson, A. A.; Nørskov, J. K.; Stephens, I. E. L.; Chorkendorff, I. Opportunities and Challenges in the Electrocatalysis of CO2 and CO Reduction Using Bifunctional Surfaces: A Theoretical and Experimental Study of Au−Cd Alloys. J. Catal. 2016, 343, 215−231. (8) Hori, Y. In Modern Aspects of Electrochemistry; Vayenas, C. G., White, R. E., Gamboa-Aldeco, M. E., Eds.; Springer: New York, 2008; pp 89−189. (9) Peterson, A. A.; Abild-Pedersen, F.; Studt, F.; Rossmeisl, J.; Norskov, J. K. How Copper Catalyzes the Electroreduction of Carbon Dioxide into Hydrocarbon Fuels. Energy Environ. Sci. 2010, 3, 1311− 1315. (10) Hori, Y.; Kikuchi, K.; Suzuki, S. Production of CO and CH4 in Electrochemical Reduction of CO2 at Metal Electrodes in Aqueous Hydrogencarbonate Solution. Chem. Lett. 1985, 14, 1695−1698. (11) 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. (12) Feaster, J. T.; Shi, C.; Cave, E. R.; Hatsukade, T.; Abram, D. N.; Kuhl, K. P.; Hahn, C.; Nørskov, J. K.; Jaramillo, T. F. Understanding Selectivity for the Electrochemical Reduction of Carbon Dioxide to Formic Acid and Carbon Monoxide on Metal Electrodes. ACS Catal. 2017, 7, 4822−4827. (13) Kortlever, R.; Shen, J.; Schouten, K. J. P.; Calle-Vallejo, F.; Koper, M. T. M. Catalysts and Reaction Pathways for the Electrochemical Reduction of Carbon Dioxide. J. Phys. Chem. Lett. 2015, 6, 4073−4082. (14) Hori, Y. In Solar to Chemical Energy Conversion; Springer: Cham, Switzerland, 2016; pp 191−211. (15) Haynes, W. M.; Lide, D. R.; Bruno, T. J. CRC Handbook of Chemistry and Physics: A Ready-Reference Book of Chemical and Physical Data; CRC Press: Boca Raton, FL, 2015. (16) Hori, Y.; Murata, A.; Takahashi, R.; Suzuki, S. Electroreduction of Carbon Monoxide to Methane and Ethylene at A Copper Electrode in Aqueous Solutions at Ambient Temperature and Pressure. J. Am. Chem. Soc. 1987, 109, 5022−5023. (17) Hori, Y.; Takahashi, R.; Yoshinami, Y.; Murata, A. Electrochemical Reduction of CO at a Copper Electrode. J. Phys. Chem. B 1997, 101, 7075−7081. (18) Schouten, K. J. P.; Pérez Gallent, E.; Koper, M. T. M. Structure Sensitivity of the Electrochemical Reduction of Carbon Monoxide on Copper Single Crystals. ACS Catal. 2013, 3, 1292−1295. (19) Schouten, K. J. P.; Qin, Z.; Perez Gallent, E.; Koper, M. T. M. Two Pathways for the Formation of Ethylene in CO Reduction on Single-Crystal Copper Electrodes. J. Am. Chem. Soc. 2012, 134, 9864− 9867. (20) Javier, A.; Chmielowiec, B.; Sanabria-Chinchilla, J.; Kim, Y.-G.; Baricuatro, J. H.; Soriaga, M. P. A DEMS Study of the Reduction of CO2, CO, and HCHO Pre-Adsorbed on Cu Electrodes: Empirical Inferences on the CO2RR Mechanism. Electrocatalysis 2015, 6, 127− 131. (21) Li, C. W.; Ciston, J.; Kanan, M. W. Electroreduction of Carbon Monoxide to Liquid Fuel on Oxide-derived Nanocrystalline Copper. Nature 2014, 508, 504−507. (22) Verdaguer-Casadevall, A.; Li, C. W.; Johansson, T. P.; Scott, S. B.; McKeown, J. T.; Kumar, M.; Stephens, I. E. L.; Kanan, M. W.; Chorkendorff, I. Probing the Active Surface Sites for CO Reduction on Oxide-Derived Copper Electrocatalysts. J. Am. Chem. Soc. 2015, 137, 9808−9811. (23) Bertheussen, E.; Verdaguer-Casadevall, A.; Ravasio, D.; Montoya, J. H.; Trimarco, D. B.; Roy, C.; Meier, S.; Wendland, J.; Nørskov, J. K.; Stephens, I. E. L.; Chorkendorff, I. Acetaldehyde as an Intermediate in the Electroreduction of Carbon Monoxide to Ethanol on Oxide-Derived Copper. Angew. Chem., Int. Ed. 2016, 55, 1450− 1454. (24) Kuhl, K. P.; Cave, E. R.; Abram, D. N.; Jaramillo, T. F. New Insights into the Electrochemical Reduction of Carbon Dioxide on Metallic Copper Surfaces. Energy Environ. Sci. 2012, 5, 7050−7059.

(25) Lobaccaro, P.; Singh, M. R.; Clark, E. L.; Kwon, Y.; Bell, A. T.; Ager, J. W. Effects of Temperature and Gas-liquid Mass Transfer on the Operation of Small Electrochemical Cells for the Quantitative Evaluation of CO2 Reduction Electrocatalysts. Phys. Chem. Chem. Phys. 2016, 18, 26777−26785. (26) Pérez-Gallent, E.; Marcandalli, G.; Figueiredo, M. C.; CalleVallejo, F.; Koper, M. T. M. Structure- and Potential-Dependent Cation Effects on CO Reduction at Copper Single-Crystal Electrodes. J. Am. Chem. Soc. 2017, 139, 16412−16419. (27) Ma, S.; Sadakiyo, M.; Luo, R.; Heima, M.; Yamauchi, M.; Kenis, P. J. A. One-step Electrosynthesis of Ethylene and Ethanol from CO2 in an Alkaline Electrolyzer. J. Power Sources 2016, 301, 219−228. (28) Dinh, C.-T.; Burdyny, T.; Kibria, M. G.; Seifitokaldani, A.; Gabardo, C. M.; García de Arquer, F. P.; Kiani, A.; Edwards, J. P.; De Luna, P.; Bushuyev, O. S.; Zou, C.; Quintero-Bermudez, R.; Pang, Y.; Sinton, D.; Sargent, E. H. CO2 electroreduction to ethylene via hydroxide-mediated copper catalysis at an abrupt interface. Science 2018, 360, 783−787. (29) Strehblow, H.-H.; Maurice, V.; Marcus, P. Initial and Later Stages of Anodic Oxide Formation on Cu, Chemical Aspects, Structure and Electronic Properties. Electrochim. Acta 2001, 46, 3755−3766. (30) Zhang, Y.-J.; Sethuraman, V.; Michalsky, R.; Peterson, A. A. Competition between CO2 Reduction and H2 Evolution on Transition-Metal Electrocatalysts. ACS Catal. 2014, 4, 3742−3748. (31) Hori, Y.; Konishi, H.; Futamura, T.; Murata, A.; Koga, O.; Sakurai, H.; Oguma, K. Deactivation of Copper Electrode” in electrochemical reduction of CO2. Electrochim. Acta 2005, 50, 5354−5369. (32) Bertheussen, E.; Abghoui, Y.; Jovanov, Z. P.; Varela, A.-S.; Stephens, I. E. L.; Chorkendorff, I. Quantification of Liquid Products from the Electroreduction of CO2 and CO Using Static HeadspaceGas Chromatography and Nuclear Magnetic Resonance Spectroscopy. Catal. Today 2017, 288, 54−62. (33) Bertheussen, E.; Hogg, T. V.; Abghoui, Y.; Engstfeld, A. K.; Chorkendorff, I.; Stephens, I. E. L. Electroreduction of CO on Polycrystalline Copper at Low Overpotentials. ACS Energy Lett. 2018, 3, 634−640. (34) Montoya, J. H.; Shi, C.; Chan, K.; Nørskov, J. K. Theoretical Insights into a CO Dimerization Mechanism in CO2 Electroreduction. J. Phys. Chem. Lett. 2015, 6, 2032−2037. (35) Calle-Vallejo, F.; Koper, M. T. M. Theoretical Considerations on the Electroreduction of CO to C2 Species on Cu(100) Electrodes. Angew. Chem., Int. Ed. 2013, 52, 7282−7285. (36) Birdja, Y. Y.; Koper, M. T. M. The Importance of CannizzaroType Reactions during Electrocatalytic Reduction of Carbon Dioxide. J. Am. Chem. Soc. 2017, 139, 2030−2034. (37) 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. (38) Kim, Y.-G.; Javier, A.; Baricuatro, J. H.; Torelli, D.; Cummins, K. D.; Tsang, C. F.; Hemminger, J. C.; Soriaga, M. P. Surface Reconstruction of Pure-Cu Single-crystal Electrodes under COreduction Potentials in Alkaline Solutions: A Study by Seriatim ECSTM-DEMS. J. Electroanal. Chem. 2016, 780, 290−295. (39) Pérez-Gallent, E.; Figueiredo, M. C.; Calle-Vallejo, F.; Koper, M. T. M. Spectroscopic Observation of a Hydrogenated CO Dimer Intermediate During CO Reduction on Cu(100) Electrodes. Angew. Chem., Int. Ed. 2017, 56, 3621−3624. (40) Schouten, K. J. P.; Pérez Gallent, E.; Koper, M. T. M. The Influence of pH on the Reduction of CO and to Hydrocarbons on Copper Electrodes. J. Electroanal. Chem. 2014, 716, 53−57. (41) Liu, X.; Xiao, J.; Peng, H.; Hong, X.; Chan, K.; Nørskov, J. K. Understanding Trends in Electrochemical Carbon Dioxide Reduction Rates. Nat. Commun. 2017, 8, 15438. (42) Hahn, C.; Hatsukade, T.; Kim, Y.-G.; Vailionis, A.; Baricuatro, J. H.; Higgins, D. C.; Nitopi, S. A.; Soriaga, M. P.; Jaramillo, T. F. 7453

DOI: 10.1021/acscatal.8b01200 ACS Catal. 2018, 8, 7445−7454

Research Article

ACS Catalysis Engineering Cu Surfaces for the Electrocatalytic Conversion of CO2: Controlling Selectivity toward Oxygenates and Hydrocarbons. Proc. Natl. Acad. Sci. U. S. A. 2017, 114, 5918−5923. (43) Tang, W.; Peterson, A. A.; Varela, A. S.; Jovanov, Z. P.; Bech, L.; Durand, W. J.; Dahl, S.; Norskov, J. K.; Chorkendorff, I. The Importance of Surface Morphology in Controlling the Selectivity of Polycrystalline Copper for CO2 Electroreduction. Phys. Chem. Chem. Phys. 2012, 14, 76−81. (44) Peterson, A. A.; Abild-Pedersen, F.; Studt, F.; Rossmeisl, J.; Nørskov, J. K. How Copper Catalyzes the Electroreduction of Carbon Dioxide into Hydrocarbon Fuels. Energy Environ. Sci. 2010, 3, 1311− 5. (45) Nørskov, J. K.; Studt, F.; Abild-Pedersen, F.; Bligaard, T. In Fundamental Concepts in Heterogeneous Catalysis; John Wiley & Sons: Hoboken, NJ, 2014; pp 155−174. (46) Strmcnik, D.; Uchimura, M.; Wang, C.; Subbaraman, R.; Danilovic, N.; van der Vliet, D.; Paulikas, A. P.; Stamenkovic, V. R.; Markovic, N. M. Improving the Hydrogen Oxidation Reaction Rate by Promotion of Hydroxyl Adsorption. Nat. Chem. 2013, 5, 300−306. (47) Ooka, H.; Figueiredo, M. C.; Koper, M. T. M. Competition between Hydrogen Evolution and Carbon Dioxide Reduction on Copper Electrodes in Mildly Acidic Media. Langmuir 2017, 33, 9307−9313. (48) Chan, K.; Nørskov, J. K. Potential Dependence of Electrochemical Barriers from ab Initio Calculations. J. Phys. Chem. Lett. 2016, 7, 1686−1690. (49) Hori, Y.; Takahashi, R.; Yoshinami, Y.; Murata, A. Electrochemical Reduction of CO at a Copper Electrode. J. Phys. Chem. B 1997, 101, 7075−7081. (50) Raciti, D.; Cao, L.; Livi, K. J. T.; Rottmann, P. F.; Tang, X.; Li, C.; Hicks, Z.; Bowen, K. H.; Hemker, K. J.; Mueller, T.; Wang, C. Low-Overpotential Electroreduction of Carbon Monoxide Using Copper Nanowires. ACS Catal. 2017, 7, 4467−4472. (51) Lee, S.; Kim, D.; Lee, J. Electrocatalytic Production of C3−C4 Compounds by Conversion of CO2 on a Chloride-Induced Bi-Phasic Cu2O-Cu Catalyst. Angew. Chem. 2015, 127, 14914−14918. (52) Eilert, A.; Cavalca, F.; Roberts, F. S.; Osterwalder, J.; Liu, C.; Favaro, M.; Crumlin, E. J.; Ogasawara, H.; Friebel, D.; Pettersson, L. G. M.; Nilsson, A. Subsurface Oxygen in Oxide-Derived Copper Electrocatalysts for Carbon Dioxide Reduction. J. Phys. Chem. Lett. 2017, 8, 285−290. (53) Gao, D.; Zegkinoglou, I.; Divins, N. J.; Scholten, F.; Sinev, I.; Grosse, P.; Roldan Cuenya, B. Plasma-Activated Copper Nanocube Catalysts for Efficient Carbon Dioxide Electroreduction to Hydrocarbons and Alcohols. ACS Nano 2017, 11, 4825−4831. (54) Garza, A. J.; Bell, A. T.; Head-Gordon, M. Is Subsurface Oxygen Necessary for the Electrochemical Reduction of CO2 on Copper? J. Phys. Chem. Lett. 2018, 9, 601−606. (55) Ren, D.; Fong, J.; Yeo, B. S. The effects of currents and potentials on the selectivities of copper toward carbon dioxide electroreduction. Nat. Commun. 2018, 9, 925.

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