Perspective pubs.acs.org/JPCL
Catalysts and Reaction Pathways for the Electrochemical Reduction of Carbon Dioxide Ruud Kortlever, Jing Shen, Klaas Jan P. Schouten, Federico Calle-Vallejo, and Marc T. M. Koper* Leiden Institute of Chemistry, Leiden University, PO Box 9502, 2300 RA Leiden, The Netherlands ABSTRACT: The electrochemical reduction of CO2 has gained significant interest recently as it has the potential to trigger a sustainable solar-fuel-based economy. In this Perspective, we highlight several heterogeneous and molecular electrocatalysts for the reduction of CO2 and discuss the reaction pathways through which they form various products. Among those, copper is a unique catalyst as it yields hydrocarbon products, mostly methane, ethylene, and ethanol, with acceptable efficiencies. As a result, substantial effort has been invested to determine the special catalytic properties of copper and to elucidate the mechanism through which hydrocarbons are formed. These mechanistic insights, together with mechanistic insights of CO2 reduction on other metals and molecular complexes, can provide crucial guidelines for the design of future catalyst materials able to efficiently and selectively reduce CO2 to useful products.
T
he electrocatalytic reduction of carbon dioxide has attracted the interest of electrochemists and inorganic chemists for decades as it can facilitate a sustainable lowtemperature redox cycle for energy storage and conversion.1,2 However, major issues still need to be resolved before the reduction of CO2 to fuels becomes appealing for technological applications. The main problems holding back electrocatalytic CO2 reduction are the high overpotentials needed and the poor product selectivity and faradaic efficiency. The high overpotentials and poor product selectivities are the result of inappropriate adsorption energies of key reaction intermediates.3−5 The low faradaic efficiencies are due to the competition with the hydrogen evolution reaction (HER), which takes place in the same range of potentials as CO2 reduction.6,7 Thus, new catalysts need to be developed that increase the product selectivity and efficiency of electrocatalytic CO2 reduction while simultaneously lowering the overpotentials. In recent years, numerous research articles on electrochemical CO2 reduction have focused on the determination of the reaction mechanism, both experimental8,9 and theoretical,3−5,10 as more mechanistic insight is expected to lead to better, tailor-made catalytic systems. In this Perspective, we will discuss the current trends in understanding and designing electrocatalysts for CO2 reduction and the pathways through which various products are formed, focusing on heterogeneous electrocatalysts and (immobilized) metal complexes. We will also outline the implications of these recent mechanistic insights for the design of selective and efficient catalysts for CO2 reduction. Theoretical Considerations on the Electroreduction of CO2. The electrochemical reduction of carbon dioxide can be viewed as a multiple proton−electron reaction leading to various products “P” and water © 2015 American Chemical Society
In this Perspective, we will discuss the current trends in understanding and designing electrocatalysts for CO2 reduction and the pathways through which various products are formed, focusing on heterogeneous electrocatalysts and (immobilized) metal complexes. kCO2 + n(H+ + e−) ⇄ P + mH 2O
(1)
The most typical products “P” in aqueous media, the coefficients k, n, and m, and the equilibrium potentials are shown in Table 1. Oxalate (C2O42−, k = 2, n = 2, m = 0, E0 = −0.59 V versus NHE) has also been detected in significant amounts, though primarily in aprotic solvents, for which the quoted redox potential does not apply. Note that for all reactions in which an equal number of protons and electrons is transferred (which corresponds to the vast majority of the reactions in aqueous media), the reversible hydrogen electrode (RHE) is the most sensible reference scale, and it will be used throughout this Perspective. The thermodynamic theory of multiple proton−electron transfer predicts that for all reactions in which only 2 electrons (that is n = 2) are transferred, catalysts will exist that are able to Received: July 20, 2015 Accepted: September 24, 2015 Published: September 24, 2015 4073
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Table 1. Main Products of the Electrochemical Reduction of CO2a
a
product name and formula
k
n
m
E0 (V versus RHE)
carbon monoxide, CO formic acid, HCOOH formaldehyde, HCHO methanol, CH3OH methane, CH4 ethanol, CH3CH2OH ethylene, C2H4
1 1 1 1 1 2 2
2 2 4 6 8 12 12
1 0 1 1 2 3 4
−0.10 −0.20 (for pH < 4); −0.20 + 0.059[pH-4] (for pH > 4) −0.07 0.02 0.17 0.09 0.08
The coefficients k, n, and m in eq 1 are provided in each case together with the standard equilibrium potentials.
passing that Cu surfaces with square symmetry, namely, (100) facets, are able to break certain scaling relations during CO2 reduction to C2 products due to ensemble effects.23 An additional consideration that needs to be taken into account in the design of suitable electrocatalysts for CO2 reduction is the role of proton-coupled electron transfer. Most theoretical/computational studies on heterogeneous electrocatalysts3−5,22 make use of the computational hydrogen electrode (CHE) approach to model proton-coupled electron transfers.24 Within this simple and convenient model, it is assumed that at every step in the mechanism, concerted (simultaneous) proton−electron transfer takes place. Importantly, pH and potential effects are not included directly in the simulations; instead, they are added externally in a linear fashion. This implies that changes in pH do not lead to actual changes in activity on the RHE scale as the energies of all steps are shifted proportionally. In practice, however, many protoncoupled electron-transfer reactions, such as eq 1, show significant pH dependence on the RHE scale. In recent theoretical work, we have shown that this is typically due to the decoupling of proton and electron transfer at some stage in the reaction mechanism, for instance, in an elementary step that includes only electron transfer but not proton transfer.25 Whenever this occurs, the overall reaction rate becomes pHdependent on the thermodynamically relevant RHE scale. Below, we will show that these considerations are highly relevant for electrocatalytic CO2 reduction, not only in optimizing reactivity but also in optimizing selectivity, for instance, with respect to H2 evolution, which is normally an important reaction but undesired in the context of CO2 reduction. Carbon Dioxide Reduction on Copper. Ever since Hori made his landmark discovery in 1985 that copper has the unique ability to electrochemically reduce CO2 to hydrocarbons such as methane and ethylene with good faradaic efficiencies in comparison to other catalysts,26 substantial effort has been invested to understand the special reactivity of copper for this reaction.27 It was shown in early work that CO is a key intermediate in the formation of hydrocarbons from the reduction of CO2 on copper,28 which is now widely accepted in the literature. However, proposing a conclusive mechanism for the reduction of CO2 on copper is challenging, as illustrated by the recent observation of 16 different products formed from CO2.29 Besides methane and ethylene, these products include a broad mix of aldehydes, ketones, carboxylic acids, and alcohols, out of which 12 are C2 or C3 species, showing the complexity of this reaction. In Figure 1 the potential dependence of these products is shown, as published recently by Jaramillo et al.29 So far, detailed mechanistic pathways have been suggested for the formation of C1 and C2 products, to be discussed below. The origin of the many other products remains largely unclear,
catalyze the corresponding redox reactions in both directions, namely, both reduction and oxidation, with negligible overpotentials.11 We will refer to this situation as “reversible catalysis”. For instance, hydrogen oxidation and hydrogen evolution catalyzed on platinum electrodes are good examples of reversible catalysis. It is worth noting that reversible catalysts satisfy the Sabatier principle, which means that they bind key reaction intermediates in an optimal fashion. Optimality is defined in this context as a compromise in binding strength; too strong binding leads to catalyst poisoning, while too weak binding prohibits the commencement of the reaction. For twoelectron transfer reactions, there is typically only a single intermediate, and hence, the identification of the optimal catalyst becomes a “simple” one-dimensional optimization problem with the ideal catalyst being on top of the so-called activity volcano. These ideas go back to Parsons12 and Gerischer13 but have been revived recently14 owing to the ability to calculate accurate binding energies of presumed catalytic intermediates by first-principles quantum chemical calculations, typically employing the density functional theory (DFT) formalism. Detailed discussions of the use of such calculations for the design of (electro)catalysts can be found in the literature.15,16 Redox reactions involving the transfer of more than two electrons (and protons) will take place through more than one catalytic intermediate, which ultimately causes their catalytic irreversibility.11 This is because the binding energies of certain adsorbed intermediates follow linear energetic scaling relationships,17−19 which often locate the top of activity volcano plots far from the equilibrium potentials. The best-known example is the universal energy difference of ∼3.2 eV between the *OOH and *OH intermediates during oxygen reduction and oxygen evolution, which is considerably larger than the “ideal” value of 2.46 eV, required for reversible catalysis.11 In the multielectron transfer reduction of CO2, a similar universal energetic scaling between *CO and *HCO intermediates leads to an overpotential for the six- or eight-electron reduction of CO2 to methanol or methane.20 Universal scaling relations are the result of the existence of similar chemical bonds between various adsorbed species and different catalytic surfaces, for instance, through a single bond pairing a single valence electron.18 This is the case for oxygenates (*OR, where R can be H, OH, CH3, CH2CH3, etc.),19 which imposes substantial intrinsic overpotentials for the reduction of CO2: the energetic scaling of *OH with *OCH3 in the pathways to CH4 or CH3OH;4,5 the scaling of *OH and *OCHCH2 for the reduction to C2H4;3 and the scaling between *OCHCH2, *OCHCH3, and *OCH2CH3 for the reduction to CH3CH2OH.3 Therefore, materials where different bonding modes exist need to be designed and synthesized if scaling relations are to be broken.20−22 Note in 4074
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toward formic acid is thus separate from the hydrocarbon pathway, which must go through carbon monoxide.30,31 On the other hand, a carbene species (*CH2) on the surface, formed from *CO, was proposed to be a common intermediate for the production of methane and ethylene formation. Methane is then formed by double proton−electron transfer to the carbene intermediate, while ethylene is formed by either dimerization of *CH2 species or CO insertion in a Fischer−Tropsch-like step, which has also been suggested to be the pathway for the formation of alcohols.32 In contrast with this “carbene” mechanism, Peterson et al. have performed DFT calculations of the reduction of CO2 to methane on Cu(211) surfaces (which contain short (111) terraces separated by (100)-like steps), suggesting that in the thermodynamically most favorable pathway, the second C−O bond is broken only at a late stage of the mechanism (see Figure 2a).4 After the initial formation of *CO, there is subsequent hydrogenation to *HCO, *H2CO, and *H3CO (methoxy), and this methoxy intermediate is reduced to CH4 and *O, which is finally reduced to H2O. Ethylene is formed by dimerization of HxCO species and subsequent deoxygenation.10 Note that this mechanism includes *H2CO and *H3CO, which cannot explain the fact that formaldehyde (CH2O) is reduced only to methanol (CH3OH) and that methanol cannot be reduced to methane.8 The inclusion of kinetic barriers in the theoretical analysis made by Nie et al. on Cu(111)5 leads to a substantially different mechanism for the formation of CH4, which is in better agreement with various experimental observations, namely, the deactivation of the catalyst while producing methane presumably by coking30,31,33 and the reduction of formaldehyde to methanol and not to methane.8 Note that within their analysis, Nie et al. predict the formation of C2H4 to occur via coupling of *CH2 moieties. It is important to note that despite their different features, none of the above mechanisms are able to explain three important experimental observations. The first observation was made by Hori et al.,32 namely, that the formation of methane from CO shows a different pH dependence from the formation of ethylene. The second observation is that during CO2 reduction, often ethylene formation takes place at less negative potentials without any simultaneous formation of methane. The third observation links these two observations to the copper surface structure, based on our own experiments with copper single-crystal electrodes;34,35 on Cu(111), the reduction of CO
Figure 1. Current efficiencies of the products of CO2 reduction on Cu electrodes in 0.1 M KHCO3 (pH 6.8) as a function of potential for major (top), intermediate (middle), and minor (bottom) products. Reproduced from ref 29 with permission from The Royal Society of Chemistry.
though we emphasize that they occur in such small amounts that they are only detectable with NMR spectroscopy. Formic acid is one of the products of CO2 reduction on Cu. Early mechanistic studies found that formic acid cannot be reduced to other products, suggesting that the mechanistic pathway
Figure 2. Pathways for the electrochemical production of methane from CO2 on Cu electrodes from (a) a thermodynamic analysis, adapted from ref 4, and (b) a combined thermodynamic and kinetic analysis from ref 5. Species in black are adsorbates, while those in red are reactants or products in solution. The two mechanisms are identical up to the second proton−electron transfer, that is, until *CO is formed, after which they form (a) *CHO and (b) *COH. 4075
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Figure 3. Possible reaction pathways for the electrocatalytic reduction of CO2 to products on transition metals and molecular catalysts: (a) pathways from CO2 to CO, CH4 (blue arrows), CH3OH (black arrows), and HCOO− (orange arrows); (b) pathways from CO2 to ethylene (gray arrows) and ethanol (green arrows); (c) pathway of CO2 insertion into a metal−H bond yielding formate (purple arrows). Species in black are adsorbates, while those in red are reactants or products in solution. Potentials are reported versus RHE, while RDS indicates rate-determining steps and (H+ + e−) indicates steps in which either concerted or separated proton−electron transfer takes place.
to both methane and ethylene is observed to take place simultaneously. On the other hand, on Cu(100), a second pathway is observed that forms ethylene at lower potentials without methane formation, especially at high (alkaline) pH, in agreement with Hori’s observation on the pH dependence of ethylene formation.32 Earlier work by Hori and co-workers also showed a strong structural dependence for the reduction of CO2 and CO on Cu; methane is preferentially formed on Cu(111), while ethylene is the main product on Cu(100).36 Therefore, the combined evidence suggests the existence of structure-sensitive and pH-dependent pathways on Cu. On the basis of this conclusion and taking into account other experimental observations,8 as well as our own DFT results on the reduction of CO on Cu(100),3 we propose the comprehensive mechanism for the reduction of CO2 on copper shown schematically in Figure 3. In this mechanism, a distinction is made between the pathway leading to methane (C1 pathway) and the pathway leading to ethylene (C2 pathway). In the C1 pathway, the CO intermediate is first reduced to a formyl species (*CHO) or a *COH species, which is further reduced to methane. The mechanism assumes an early breaking of the second C−O bond (as was shown earlier in Figure 2b), though we emphasize that a late breaking
of this C−O bond may be more thermodynamically favorable according to the DFT calculations of Peterson et al. (see Figure 2a).4 Dimerization of the intermediates in this C1 pathway may also yield ethylene at high applied overpotentials.10 It should be noted that this pathway toward ethylene is believed to be the pathway that produces the bulk of the amount of ethylene once current densities reach 10 mA cm−2 and takes place on both Cu(100) and Cu(111). In the C2 pathway, the key C−C bond-making step at low overpotentials is a CO dimerization step mediated by electron transfer rendering a *C2O2− intermediate. Proton transfer happens only af ter the formation of the negatively charged adsorbed CO dimer.3,37 Note that this dimerization is the ratedetermining step for the reduction of CO (RDS2 in Figure 3). Such a decoupling resulting in first electron then proton transfer explains why the CO reduction prefers alkaline media. The C−C bond making by the reductive coupling of CO is somewhat similar to the reductive carbonyl coupling to olefins as originally suggested by McMurry.38 Our DFT calculations have shown that the formation of the negatively charged CO dimer is most stable on square arrangements of four surface atoms,23 explaining the observed preferential formation of ethylene on Cu(100) at low potentials. At this point, it is 4076
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Carbon Dioxide Reduction on Other Metals. Recently, Peterson and Nørskov have outlined the requirements for optimal catalysts for the reduction of CO2 to methane.20 According to their work, it is essential that a catalyst efficiently catalyzes the protonation of *CO to *COH or *CHO while simultaneously showing poor activity for the HER, in order to optimize methane production. Metals that bind *CO weakly will not produce methane as CO will desorb before further reduction can take place, while metals that bind *CO strongly still require large overpotentials to reduce it to methane because of the highly unfavorable thermodynamic conditions needed for the formation of *CHO or *COH. Copper is near the top of the volcano, indicating that the *CO adsorption strength on copper is suitable for the production of methane. The binding energies of *CO on gold and silver are too weak; therefore, they mostly produce CO and are not able to produce any hydrocarbons. Metals such as Pt, Pd, and Ni bind CO too strongly. Consequently, they will be poisoned by *CO, which cannot be removed from the surface. However, small amounts of methane have been observed experimentally on Pd and Ni electrodes.1,49 In earlier reports, the production of hydrocarbons and alcohols was mostly assumed to be a unique property of Cu that was absent for other transition metals. However, Jaramillo et al. recently showed that the production of methane and methanol is more general than initially thought. They have shown the production of methane on Fe, methane and methanol on Au, Zn, Ni, and Pt, and methane and methanol (and ethanol) on Ag, albeit in small amounts.9,50 Their proposed mechanism for the formation reduction of CO2 to methane and methanol on Ag shows similarities to the mechanism that we proposed for the formation of the same products on Cu in Figure 3.8,9 Adsorbed CO is first formed, and then it is either desorbed or further transformed into a formyl (*CHO) species or a *COH species on the surface. These intermediates then react to form methane and/or methanol. Transition metals mainly produce carbon monoxide but also formate.51 It is believed that the formation of CO or formic acid depends on the initial binding mode of the first intermediate of CO2 reduction, as is illustrated in Figure 3a. The precursor that leads to the formation of CO binds to the catalyst via the C atom, that is, the carboxyl intermediate (*COOH). In DFT simulations and in the mechanism shown in Figure 3a, *COOH is assumed to be formed through a concerted proton−electron transfer to CO2. However, in some of the older experimental literature, the formation of *COOH was considered to take place via the formation of a CO2− radical, which was suggested to adsorb at copper and gold electrode surfaces.49,52 The existence of this anionic intermediate implies the decoupling of proton and electron transfer (see Figure 3a) and is, therefore, expected to show a different pH dependence from the concerted pathway. A potential example will be discussed below in the section on molecular catalysts. On the other hand, the catalytic intermediate that leads to formate or formic acid is expected to bind to the catalyst through (one of) the oxygen atoms, either in a monodentate or bidentate fashion, such that the C atom is available for hydrogenation (see Figure 3a). It is not clear whether this intermediate forms through a reaction with *H via a CO2 insertion reaction into the metal− hydrogen bond or if the intermediate can be formed through direct protonation with H+ from solution.
important to note that Bard and co-workers were able to reduce CO in liquid ammonia to C2O22− using Pt, Ni, and Hg electrodes, which suggests that C−C coupling indeed involves electron transfer without any proton transfer.39 It is also assumed that this CO dimerization is the key step in the recently observed selective reduction of CO to ethanol on oxide-derived copper electrodes.40 Nanoparticulate copper electrodes have recently gained significant attention due to reports that oxide-derived copper nanoparticles can reduce the onset potentials for both formic acid and CO compared to polycrystalline copper electrodes.41,42 Furthermore, it was shown that on a roughened copper-nanoparticle-covered electrode, prepared by electrodeposition of copper, the relative selectivity toward ethylene over methane could be increased.43,44 Both the oxide-derived nanoparticles and the electrodeposited nanoparticles showed an increased stability in comparison to polycrystalline copper, which generally shows fast deactivation. These observations are attributed to the surface structure of the nanoparticles as these tend to contain more defect sites such as kinks and steps. For instance, a recent study by Verdaguer−Casadevall et al. has linked the activity change for oxide-derived copper catalysts for the reduction of CO to the number of metastable sites that bind CO strongly and can exist at grain boundaries.45 However, we have shown that the initial crystal orientation of Cu2O films before reduction has only a minor effect on the product selectivity and deactivation, while the oxide layer thickness has a significant effect.46 Therefore, we hypothesize that the selectivity change is not only associated with the surface structure but also (or perhaps even rather) with the increase of the effective surface area, which in turn has a substantial effect on the local pH in the boundary layer near the electrode under reductive conditions.33,46 To support this hypothesis, we have shown that a low buffer capacity of the electrolyte favors ethylene formation on copper nanoparticles, while increasing the buffer capacity leads to an increase in the selectivity toward methane.44 This was recently confirmed by Varela et al., who demonstrated that the activity and selectivity of a smooth copper electrode can also be tuned by changing the bicarbonate concentration of the electrolyte.47 Following the reaction mechanism in Figure 3, this can be explained by the different pH gradients near the electrode surface, with a smaller pH gradient near the electrode with increasing buffer strength. A low buffering capacity leads to a high local pH near the electrode, yielding an enhanced production of ethylene at such low buffer capacities, in agreement with the earlier observation that CO can selectively be reduced in the C2 pathway to ethylene in alkaline media, separate from the C1 pathway producing both methane and ethylene. Interestingly, when the selectivity toward methane is increased, a deactivation effect is observed that is not present when ethylene is the major product.44 This implies that large amounts of poisoning intermediates form on copper nanoparticles in the C1 pathway, and such intermediates do not take part in the C2 pathway. The effect of Cu nanoparticle size for the reduction of CO2 was recently studied, showing that the reduction of CO2 on Cu nanoparticles of 2 nm and smaller is significantly enhanced.48 This increase in activity however is due to an increase in H2 and CO production, while the efficiency toward hydrocarbons vanishes. These catalytic properties are believed to be due to an increase in undercoordinated sites, which enhance both the HER and the reduction of CO2 to CO. 4077
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transfer reaction, a feature that had previously been demonstrated only for immobilized electroactive enzymes.67 Role of the Electrolyte. The nature of the electrolyte plays an important role in electrocatalytic systems for CO2 reduction. Solvent and pH control the (relative) concentrations of the reactants, CO2 and H+ or *H. Furthermore, solvent, pH, and the presence of certain cations or anions can stabilize reaction intermediates or inhibit their formation and are speculated to aid directly in the choice of certain reaction pathways.1 Therefore, the use of suitable electrolytes at the right pH values can be an important step toward more active and selective electrocatalytic processes. One important parameter to control proton availability is the pH of the electrolyte. Given that CO2 forms bicarbonate and carbonate when it is purged through alkaline solutions, electrolytes for CO2 reduction are limited to the neutral and acidic pH range. However, because both the CO2 reduction and the HER consume protons (or generate OH−, alternatively) at the electrode surface, the local pH near the electrode is typically higher than that in the bulk. As a result, CO2 near the electrode can react with OH−, forming bicarbonate or carbonate, depending on the local pH in the proximity of the electrode. Besides, intermediates of the CO2 reduction such as CO find a local pH at the electrode that differs from the bulk pH, which according to Figure 3 may have significant consequences on the selectivity of the reaction. This is because the C1 and C2 pathways have distinctly different pH dependence. Although most studies use buffered solutions as electrolytes (either bicarbonate buffers or phosphate buffers) to counteract the overall changes in pH, near the electrode, there might be significant pH differences with respect to the bulk electrolyte. Gupta et al. have illustrated this phenomenon with some simple calculations.1,68 To a first approximation, the local pH at the electrode is determined by the bulk pH of the electrolyte, its buffer strength, as well as the local current density flowing. In turn, the local current density depends on the local surface roughness. We believe that this effect is often overlooked when explaining observed changes in reaction selectivity with changes in surface structure.46 Several studies have also probed the role of the anions and cations in the electrolyte.69,70 It has been hypothesized that cations play a role in electrochemical CO2 reduction either by specifically adsorbing on the electrode, thus influencing the potential at the outer Helmholtz plane, or by delivering water molecules from their solvation shell to the electrode, thus influencing the hydrogen coverage on the electrode. Additionally, one could imagine that different cations may have different interactions with (negatively charged) intermediates of the reaction. On Cu electrodes, an increase in the selectivity for the production of ethylene over methane has been observed with an increase in the Stokes radii of the hydrated cations, Li+ > Na+ > K+ > Cs+.69,70 Note that Hori and Murata also reported hindrance of the HER by using different alkaline cations and obtained the highest selectivity toward hydrocarbons with K+.70 Furthermore, Kyriacou et al. have recently shown that the rate of reduction can be influenced by using multivalent cations as supporting electrolytes.69 A consistent understanding of the role of cations remains elusive, but they are appealing to tune the catalyst’s reactivity and selectivity. Most studies in electrochemical CO2 reduction have focused on aqueous electrolytes. However, CO2 dissolves poorly in water (0.034 M), leading to low concentrations of CO2 in a saturated aqueous electrolyte. Besides, HER is a prominent
On p-block metals such as In, Sn, Hg, and Pb, formic acid and formate are the main products.51,53 The mechanistic pathway by which formate and formic acid are formed on these metals has been speculated to proceed via a (weakly adsorbed) CO2− radical that reacts with water to form formate or formic acid.53 Formic acid and formate can be formed with high selectivity on these metals because they are poor catalysts for the competing HER. However, in order to reach these high selectivities, high overpotentials are needed as the redox potential for the formation of the CO2− radical is −1.9 V (versus NHE).51 It should be noted that recently some improvements have been reported, showing high selectivities for CO2 reduction at lower overpotentials on nanostructured Pb and Sn electrodes.54,55 However, we expect that very low (near-zero) overpotentials for the reduction of CO2 to formate/formic acid require the stabilization of an adsorbed formate intermediate as illustrated in the top part of Figure 3. As was the case for Cu, some studies have probed the effect of nanoparticles size on the reduction of CO2 on other metals.56,57 A decrease in the size of Au nanoparticles caused an increase in the observed current density and a decrease in the faradaic efficiency toward CO. This is believed to be caused by an increase of undercoordinated sites, as was also hypothesized for Cu nanoparticles.56 Tuning the size of Pd nanoparticles showed that an increase in faradaic efficiency toward CO can be achieved by decreasing the particle size. In this case, the increase in CO production was ascribed to an increase in corner and edge sites, which display a higher activity toward CO2 reduction to CO.57 A promising area of research is the design of bimetallic catalysts for the reduction of CO2. Early work showed that the selectivity of a metal electrode for CO or formic acid can be altered by the addition of adatoms on the surface.49 Sakata et al. studied the effects of alloying Cu with Ni, Sn, Pb, Zn, Cd, and Ag.58 They showed that alloying can have a (beneficial) effect on the onset potentials for the formation of products from CO2 and that some alloys are able to make products that both metals separately cannot produce in detectable amounts. A Cu/Ni alloy, for instance, produces methanol at low overpotentials with a faradaic efficiency of 5%, while Cu and Ni separately produce methanol only in negligible amounts.58 Moreover, theoretical studies are also enriching the state of the art in the design of more energetically efficient and chemically selective bimetallic catalysts for the reduction of CO2. Approaches based on copper alloys and the so-called “isolated active sites” aim at retaining or mimicking the unique reducing features of Cu while trying to either hinder the HER or reduce the onset potential.59,60 Cu−Au alloys and Cu overlayers on Pt have also been investigated for the production of CO and hydrocarbons, respectively.61−64 For Cu−Au alloys, it was shown that the activity and selectivity toward CO production can be tuned by controlling the electronic and geometric properties,61,62 while for Cu overlayers on Pt, differences in the surface strain allowed tuning the selectivity toward methane or ethylene.64 The potential of bimetallic catalysts is further illustrated by our own recent work on a Pt−Pd alloy.65,66 This catalyst was shown to produce formic acid from CO2 with an onset potential very close to the corresponding equilibrium potential, with high faradaic efficiencies at high current density and very moderate overpotentials, and with good stability. Furthermore, this alloy is able to reversibly reduce CO2 to formic acid and oxidize formic acid to CO2, as expected for a two-electron4078
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metal centers, either HCOOH is produced as an end product without consecutive reaction or CO is formed as a precursor to the more reduced products CH3OH and CH4. M−COOH was assumed as the first and only key intermediate. For the activation of CO2, their calculations suggest a competition between *H and *COOH adsorption, with *H being more strongly bound. In order to overcome *H poisoning, they suggest to reduce CO instead of CO2, and their DFT results predict Rh centers to be the most active catalysts. Nielsen et al. have also utilized DFT calculations to investigate the potential intermediates and the reaction mechanism of the electrochemical reduction of CO2 to CO on a cobalt porphyrin in water.85 They suggest that CO2 adsorption on cobalt centers is likely to take place simultaneously with electron transfer to form a [CoIPCO2]2− intermediate (CoP denoting the Coporphyrin ring), which will be subsequently protonated by a neighboring H3O+ and the motion of a water molecule toward the CO 2 − group to form the next intermediate [CoIIPCOOH]−. Recent work by Varela et al. on a metaldoped nitrogenated carbon, which is a solid-state analogue of macrocyclic complexes, showed an independence on the nature of the metal dopant for the activity of CO2 reduction toward CO.86 Therefore, it is believed that the production of CO on this material occurs on the nitrogen functionalities. Methane production is believed to be dependent on the strength of the interaction between CO and the metal center, which determines if CO can be protonated before desorption. Cobalt complexes have been found to be the most effective immobilized molecular electrocatalysts for CO2 reduction. Kapusta and Hackerman found formate as the main product on a Co phtalocyanine deposited on a carbon electrode, with small amounts of methanol.87 From Tafel plots, they concluded that the transfer of the first electron to CO2 is the rate-determining step. However, other papers using Co macrocycles report CO as the main product of electrochemical CO2 reduction in water.81,83 Recently, our group has obtained novel insights into the electrochemical reduction of CO2 catalyzed by an immobilized Co protoporphyrin by combining cyclic voltammetry with online electrochemical mass spectrometry at different values of pH. The pH-dependent formation of CO indeed suggests the formation of a CO2•− anion bound to the Co macrocycle, with Co presumably being in the CoI state.88 This CO2•− anion is then protonated by water, rather than by H+ as suggested by Nielsen and Leung,85,89 explaining why at less acidic conditions (pH = 3), CO formation can reach up to 60% faradaic efficiency compared to
