Homogeneously Catalyzed Electroreduction of Carbon Dioxide

Review Cite This: Chem. Rev. XXXX, XXX, XXX−XXX

pubs.acs.org/CR

Homogeneously Catalyzed Electroreduction of Carbon DioxideMethods, Mechanisms, and Catalysts Robert Francke,*,† Benjamin Schille,† and Michael Roemelt‡,§ †

Institute of Chemistry, Rostock University, Albert-Einstein-Strasse 3a, 18059 Rostock, Germany Lehrstuhl für Theoretische Chemie, Ruhr-University Bochum, 44780 Bochum, Germany § Max-Planck Institut für Kohlenforschung, Kaiser-Wilhelm Platz 1, 45470 Mülheim an der Ruhr, Germany ‡

ABSTRACT: The utilization of CO2 via electrochemical reduction constitutes a promising approach toward production of value-added chemicals or fuels using intermittent renewable energy sources. For this purpose, molecular electrocatalysts are frequently studied and the recent progress both in tuning of the catalytic properties and in mechanistic understanding is truly remarkable. While in earlier years research efforts were focused on complexes with rare metal centers such as Re, Ru, and Pd, the focus has recently shifted toward earth-abundant transition metals such as Mn, Fe, Co, and Ni. By application of appropriate ligands, these metals have been rendered more than competitive for CO2 reduction compared to the heavier homologues. In addition, the important roles of the second and outer coordination spheres in the catalytic processes have become apparent, and metal−ligand cooperativity has recently become a well-established tool for further tuning of the catalytic behavior. Surprising advances have also been made with very simple organocatalysts, although the mechanisms behind their reactivity are not yet entirely understood. Herein, the developments of the last three decades in electrocatalytic CO2 reduction with homogeneous catalysts are reviewed. A discussion of the underlying mechanistic principles is included along with a treatment of the experimental and computational techniques for mechanistic studies and catalyst benchmarking. Important catalyst families are discussed in detail with regard to mechanistic aspects, and recent advances in the field are highlighted.

CONTENTS 1. Introduction 2. Fundamentals of Electrochemical CO2 Reduction 2.1. Background and Motivation 2.2. Possible Pathways 2.3. Comparison between Homogeneous and Heterogeneous Electrocatalysis 2.3.1. Heterogeneous Catalysts 2.3.2. Homogeneous Catalysts 2.4. Paired Processes 3. Molecular Catalysts and Electrochemical CO2 Reduction 3.1. General Mechanistic Guidelines 3.1.1. Electron Transfer between Electrode and Catalyst 3.1.2. Coordination Chemistry of CO2 3.1.3. Role of the Proton Source 3.2. Techniques for Benchmarking and Mechanistic Studies 3.2.1. Cyclic Voltammetry 3.2.2. Spectroelectrochemistry 3.2.3. Preparation of Catalytic Intermediates 3.2.4. Quantum Chemical Methods 3.3. Catalysts and Mechanisms 3.3.1. Mn and Re Complexes 3.3.2. Fe-Based Catalysts 3.3.3. Ru and Os Compounds 3.3.4. Co-Based Catalysts © XXXX American Chemical Society

3.3.5. Rh and Ir Complexes 3.3.6. Ni Complexes 3.3.7. Pd-Based Catalysts 3.3.8. Cu-Based Catalysts 3.3.9. Group 6 Metals 3.3.10. Organocatalysts 4. Tabular Overview 5. Conclusion and Outlook Author Information Corresponding Author ORCID Notes Biographies Acknowledgments Glossary Ligands References

A B B C C C D E F F F F H J J N Q Q V V Y AD AF

AI AJ AL AN AO AO AQ AQ BE BE BE BE BE BE BE BE BF

1. INTRODUCTION Among the numerous approaches for utilization of carbon dioxide as C1-building block,1−4 its electrocatalytic reduction Special Issue: Electrochemistry: Technology, Synthesis, Energy, and Materials Received: August 1, 2017

A

DOI: 10.1021/acs.chemrev.7b00459 Chem. Rev. XXXX, XXX, XXX−XXX

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represents a promising approach.5−7 However, an efficient and selective electroreduction of CO2 is not trivial due to the high overpotential of the cathodic process and a large number of possible reaction pathways. Therefore, it is still a subject of intensive studies after decades of research.8,9 One method to overcome the above-mentioned issues is to employ homogeneous redox catalysts in an indirect electrolysis, lowering kinetic barriers and often yielding an increased and/or entirely different selectivity compared to direct electroreduction.10,11 Accordingly, substantial progress has recently been made in this field, and the improvement of both the mechanistic understanding and the effectiveness of the catalysts is truly impressive. Since in the majority of the reported cases metal centers of complex compounds play a crucial role in CO2 activation, the advancements in homogeneously catalyzed CO2 electroreduction are closely related to the progress made in the field of metal−organic chemistry. While years ago the use of rare and precious metals such as Re and Pd played a major role,12−14 more abundant metals such as Fe, Mn, and Ni have recently attracted more attention.15−18 In order to render these metals more competitive with regard to catalytic performance, specific ligand design is often necessary.16,19,20 In fact, nowadays the role of the ligand in CO2 activation is well understood for many systems, and this knowledge recently triggered further intriguing innovations such as the use of Coulombic substituent effects and metal−ligand cooperativity for boosting the catalytic performance.18,21 The scope of this review is to present the most important fundamental aspects of electrochemical CO2 conversion with molecular electrocatalysts and to outline the development of this field in the last three decades. Not within the scope are processes where CO2 is reacted with electrogenerated nucleophiles (e.g., reduced aromatics,22,23 carbonyl compounds,24−26 or alkyl halogenides27−32) to give carboxylated organic compounds (“electrocarboxylation”), a broad field with numerous examples that has been reviewed elsewhere.33 Instead, we focus on such examples where CO2 is reduced to give C1 or C2 compounds such as CO, formate, methanol, or oxalate as potential sustainable fuels or synthetic building blocks. The heterogenization of molecular catalysts by attachment to the electrode surface has been reviewed elsewhere,34 and we will only touch on this matter briefly at several points in this review. With regard to the basic principles and the applied catalysts, the field of electrocatalytic CO2 reduction has a certain overlap with its photocatalytic counterpart. The latter topic, however, is not within the scope of this review, and readers with interest in photocatalytic or photoelectrocatalytic processes enabled by molecular catalysts are referred to other reviews.35−43 The present review is structured as follows: After a treatment of the fundamentals of electrochemical CO2 reduction (section 2) and the underlying mechanistic principles (section 3.1), we focus on electrochemical and nonelectrochemical methods which are used to study mechanisms (section 3.2). In this part of the review the most common toolscyclic voltammetry, spectroelectrochemistry, quantum chemical calculations, and the trapping/preparation of catalytic intermediateswill be introduced and discussed in the context of electrocatalytic CO2 reduction. Subsequently, important catalyst families are treated in detail with regard to mechanistic aspects and catalytic performance (section 3.3). Here we will focus on some of the most innovative and instructive examples instead of giving an exhaustive description where each catalyst is only briefly

touched. In order to provide a comprehensive overview, a summary of molecular catalysts which were reported in the last three decades to be active for CO2 reduction is given in form of a table in section 4. The reviewed field is shaped by researchers from various disciplines, and with this work we aim to address readers from all those communities, i.e., electrochemistry, coordination chemistry, spectroscopy, and theoretical chemistry. We hope that this review can provide a common knowledge base and can stimulate readers to extend their views to the neighboring fields. Since such an interdisciplinary work can naturally not cover every aspect in full depth, the reader will be referred to further literature at the appropriate points.

2. FUNDAMENTALS OF ELECTROCHEMICAL CO2 REDUCTION 2.1. Background and Motivation

The field of electrochemical CO2 reduction in a broader sense has a decade-long history, and in this time a rapidly growing number of studies has become available. Different subfields have been reviewed, among them engineering aspects and economic studies.44−48 With respect to electrodes, articles with focus on flat49,50 and nanostructured7,51 materials are available. Regarding product selectivity, methods for the selective generation of formic acid52,53 and methanol51,54 have been reviewed. Furthermore, the fundamental importance of electrocatalysis for the field is reflected by a number of reviews on different types of electrocatalysts.5,6,8,9,36,37,55−60 In light of all of these different aspects and developments, this review on homogeneous electrocatalysts covers only a relatively narrow segment of electrochemical CO2 reduction. However, we consider this segment as crucial, since studying molecular electrocatalysts with their well-defined active sites provides unique access to mechanistic information on a molecular level. These insights have also helped to obtain a better understanding of heterogeneous reactions up to the point that concepts from homogeneous catalysis are frequently adopted to tune heterogeneously catalyzed processes. Despite the long history of electrochemical CO2 reduction, the field still attracts much interest and the indicated motivations for the research are typically (1) the possibility to reduce the greenhouse effect by lowering the CO2 concentration in the atmosphere, (2) the generation of fuels from CO2, and (3) the transformation of CO2 into building blocks for the production of chemicals. These three options have previously been evaluated with respect to energy efficiency and economic aspects,44,61 and the outcome can be summarized as follows.44,61 (1) Reducing the greenhouse ef fect: Using electrochemistry for this purpose is unrealistic considering the low atmospheric CO 2 concentration (∼400 ppm, currently increasing by ∼2 ppm each year)62 and the problematic energy and CO2 balance of the overall process (including concentration of CO2, effluent treatment, stream recycling, and so on). Even if a process with an overall negative CO2 balance would be established (for instance, with the help of regenerative energy), the approach seems to be problematic considering the dimension of the problem (the worldwide emission of CO2 is in the order of 1010 t/a).62 (2) Electrochemical CO2-to-f uel conversion: For an economically feasible and ecofriendly process, replacing existing B


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carbonaceous fuels or for use in fuel cells, major improvements in terms of energy efficiency and overall CO2 balance would be required. With respect to the conversion of CO 2 into carbonaceous fuels for combustion engines, the market is simply too large and production via electrolysis of CO2 appears to be unrealistic.44 Regarding fuel cell technology, the use of CO2-derived methanol63,64 or formic acid65−67 may represent an alternative to H2 due to easier fuel transport and handling but still appears to be difficult considering the low efficiency of the state-of-the-art systems. (3) CO2 as C1 building block: A more promising approach is represented by the use of cathodic CO2 reduction for the synthesis of organic compounds. Compared to carbonaceous fuels, the production scales are small and the added value is considerably higher, which means that the overall energy and CO2 balance would not be a major issue, especially when a concentrated CO2 stream is available on site. Considering the fact that many industrially relevant chemical transformations involve an increase in the number of carbon atoms per molecule, CO2 represents an abundant and inexpensive source for C1 building blocks.2,68 For example, carbon monoxide represents such a building block which can be derived from CO2 and which is used, for instance, in carbonylation reactions or in the production of phosgene.69,70 Methanol and formic acid are further interesting CO2derived feedstocks: While the transformation of methanol into olefins (MTO process) is already a wellestablished process,71 formic acid represents a promising material for reversible chemical hydrogen storage and other applications.72 In summary, the use of CO2 as C1 feedstock for chemicals production represents the most promising option, while currently the large-scale production of fuels and the reduction of the greenhouse effect are rather unrealistic goals in the context of electrochemical CO2 reduction.

Table 1. Selected CO2 Reduction Processes and the Corresponding Standard Redox Potentials E00 for Aqueous Solutions8 E00 (vs SHE, V)

reduction process +

−

−0.25 −1.08 −0.11 −0.93 −0.07 −0.90 +0.02 −0.81 +0.17 −0.66 −0.50 −0.59 +0.06 +0.08

CO2(g) + 2H + 2e = HCOOH(l) CO2(g) + H2O(l) + 2e− = HCOO−(aq) + OH− CO2(g) + 2H+ + 2e− = CO(g) + H2O(l) CO2(g) + H2O(l) + 2e− = CO(g) + 2OH− CO2(g) + 4H+ + 4e− = CH2O(l) + H2O(l) CO2(g) + 3H2O(l) + 4e− = CH2O(l) + 4OH− CO2(g) + 6H+ + 6e− = CH3OH(l) + H2O(l) CO2(g) + 5H2O(l) + 6e− = CH3OH(l) + 6OH− CO2(g) + 8H+ + 8e− = CH4(g) + 2H2O(l) CO2(g) + 6H2O(l) + 8e− = CH4(g) + 8OH− 2CO2(g) + 2H+ + 2e− = H2C2O4(aq) 2CO2(g) + 2e− = C2O42−(aq) 2CO2(g) + 12H+ + 12e− = CH2CH2(g) + 4H2O(l) 2CO2(g) + 12H+ + 12e− = CH3CH2OH(l) + 3H2O(l)

Table 2. Selected CO2 Reduction Processes in Nonaqueous Electrolytes and the Corresponding Standard Potentials E00 process +

solvent −

CO2(g) + 2H (solv) + 2e = CO(g) + H2O(solv)

+

−

CO2(g) + 8H (solv) + 8e = CH4(g) + 2H2O(solv)

E00 (vs NHE, V)a

CH3CN

−0.65,20 −0.7573

DMF DMF + 2M H2O DMF + HBF4 CH3CN

−1.3673 −0.6920 −0.2620 −0.4873

DMF

−1.1173

a

Values from the literature converted to NHE using the conversion constants provided in ref 74.

overpotential required for catalysis (ηcat), is defined as the difference between the potential that is required to drive a reaction A → B at a specific rate and the standard potential E°(A/B) for the formation of product B from substrate A (ηcat is of fundamental importance for characterization of catalysts and will be discussed in more detail in sections 3.1.1 and 3.2.1). The selectivity of the CO2 reduction represents a further challenge. Since complicated reaction mechanisms and multiple possible pathways are often involved, a mixture of different products is typically obtained, whereby the composition strongly depends on the selected electrolysis parameters. The development and optimization of electrocatalysts, both of the homogeneous or heterogeneous type, is therefore essential for realizing sufficiently high rates and selectivity at moderate overpotentials.

2.2. Possible Pathways

The electroreduction of CO2 can proceed via one-, two-, four-, six-, and eight-electron reduction pathways. As reduction products, carbon monoxide (CO), formic acid or formate (HCOOH/HCOO−), oxalic acid or oxalate (C2O4H2/C2O42−), formaldehyde (CH2O), and methanol (CH3OH) are most frequently observed.8 Other products such as methane (CH4), ethylene (CH2CH2), and ethanol (CH3CH2OH) are far more difficult to generate and often occur only as byproducts. A selection of possible CO2 electroreduction processes along with the corresponding standard redox potentials E00 for aqueous solution is depicted in Table 1.8 For comparison, the E00 values for reduction processes in nonaqueous electrolytes are summarized in Table 2. It is important to emphasize that the E00 values in Table 1 have been calculated from the standard Gibbs energies of the reactants and merely represent the thermodynamics of the reaction, indicating the possibility for a particular pathway at a given electrode potential. Such proton-coupled multielectron reactions are relatively favorable, since they lead to products which are thermodynamically similar or more stable compared to CO2. However, in the practical case the kinetics of direct CO2 electroreduction are slow, which leads to a demand for potentials more negative than the thermodynamic requirement in order to obtain reasonable reaction rates. This additional driving force, the

2.3. Comparison between Homogeneous and Heterogeneous Electrocatalysis

The direct and uncatalyzed electrochemical reduction of CO2 generates a highly reactive CO2 radical anion via an outersphere single-electron transfer (SET). Due to the large reorganization energy between linear CO2 and bent radical anion, the equilibrium potential for the corresponding process is very negative (E00 = −1.90 V vs NHE).9 Consequently, the application of very negative potentials to inert (outer-sphere) electrodes such as carbon, lead, and mercury electrodes yields preferentially oxalate as C−C coupling product.50,75,76 2.3.1. Heterogeneous Catalysts. In contrast to the abovementioned outer-sphere electrodes, many materials provide the C


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oxide, metal−carbon, metal chalcogenide−carbon, etc.) can lead to cooperative and synergistic effects and may lead to accelerated kinetics, enhanced catalyst stability, and improved selectivity compared to the use of the individual components in their pure form.7,8 2.3.2. Homogeneous Catalysts. In parallel to the research on catalytically active electrode materials, molecularly defined compounds have been intensively studied as homogeneous catalysts for CO2 electroreduction over the last decades. The homogeneous catalyst is interposed as a shuttle between electrode and CO2, and the direct electron transfer between electrode and substrate is replaced by a homogeneous one in solution, a process which is also called indirect electrolysis (see Scheme 1, bottom).11 In this case, the reduction is carried out at the potential of the catalyst, not the one for direct CO2 reduction. For indirect conversion it is therefore important that the former potential is less negative than the latter one. In galvanostatic (constant current) electrolysis, the cathode potential automatically adjusts to the species with the least negative reduction potential, provided that the transport of the material through the diffusion layer to the electrode is fast enough to satisfy the applied current (if consumption of the redox-active species exceeds the transport rate, the cathode potential becomes more negative until it matches the potential of the next easiest species to reduce in solution).84 Since the compounds present in the reaction mixture are reduced sequentially in the order of increasingly negative redox potential, Cox can only be selectively reduced in the presence of CO2 as long as the catalyst is consumed at a rate that does not exceed diffusion of CO2 to the electrode surface. If the catalyst redox couple is reversible (allowing for multiple reaction cycles), the preconditions for a homogeneously catalyzed process are fulfilled. In order to render this indirect process efficient, both the heterogeneous electron transfer and the interaction between catalyst and CO2 must exhibit favorable kinetics. Analogously to the outer-sphere and inner-sphere electrodes discussed above, two types of indirect homogeneous electron transfer processes can be distinguished (see Scheme 2). In the

possibility for bond-forming interactions (chemisorption) prior to electron transfer (“inner-sphere electrodes”), thereby lowering the activation barrier and influencing the selectivity of the reaction (see Scheme 1, top). Typical examples in this Scheme 1. Electrochemical CO2 Reduction via Heterogeneous (top) and Homogeneous Catalysis (bottom)

context are Pt, Ag, or Cu, which adsorb CO2 prior to reduction, leading to lower values for ηcat and different reduction pathways. Such electrodes can therefore be considered as electrocatalysts, and consequently, the choice of the electrode material is one of the key factors determining the product distribution. However, generalizations and predictions for the selectivity of electrode materials are difficult, since the outcome depends on a number of additional parameters (solvent, additives, temperature, supporting electrolyte, etc.). Furthermore, the structure of the surface and the electrode pretreatment can play an important role. A few generalizations can still be made, for instance, that silver and gold in protic media give CO,50,77 while Cu cathodes in aqueous solutions give hydrocarbons, methanol, and formaldehyde.50,78−82 Enormous efforts have been undertaken in order to advance the development of catalytically active electrodes for the reduction of CO2. A vast number of different materials has been studied and reviewed.8,50,59 The materials can be categorized into four different groups: metals, metal alloys, transition metal oxides/chalcogenides, and metal organic frameworks. Comprehensive experimental and computational studies have been carried out in order to improve the mechanistic understanding of the reaction and the influence of the material, and it has been found that the binding energy between various reaction intermediates and the electrocatalyst is a key factor governing the catalytic activity and final distribution of products.6 In addition to the chemical composition of the catalyst, the influence of the surface morphology on the catalytic performance has been studied, and major differences between flat surfaces and their roughened counterparts have been found.8,50 A particular importance can be attributed to low-coordination sites such as edges, steps, and defects, and often these sites exhibit a much better intrinsic activity for CO2 reduction. In this regard, nanostructured catalysts are a frequent object of study, since they exhibit a larger surface area and a higher fraction of low-coordination sites, and in many cases it has been found that nanostructuring is notably improving the catalytic performance.6,83 A further promising direction is the use of composite cathodes for CO2 reduction. With numerous examples it has been demonstrated that combining two or more different materials to a composite (such as metal−metal

Scheme 2. Inner-Sphere and Outer-Sphere Electron Transfer between a Homogeneous Catalyst and CO2

first scenario, which is often referred to as redox catalysis,55 a nonbonded (outer-sphere) ET occurs between the reduced form of the catalyst couple (Cred) and CO2. In other words, Cred merely acts as an outer-sphere electron mediator, which shuttles electrons from the electrode to the substrate. In this case, the homogeneous electron transfer is associated with the same activation barriers and driving force limitations as the outer-sphere electron transfer at inert electrodes such as carbon or lead. The catalytic effect of such redox catalysts/outer-sphere mediators originates from two different phenomena: First, the D


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effluent treatment, and stream recycle in a flow cell are far more efficient when a heterogeneous catalyst is used. On the other hand, heterogeneous systems have a variety of different types of active sites, and therefore, a high selectivity for a certain product is difficult to achieve. In contrast, structure−reactivity relationships are much easier to study with molecularly defined catalysts, making optimization of the selectivity easier. Consequently, homogeneous catalysts with a FE of >95% with respect to a particular product are not unusual. A common challenge for both fields is represented by the degradation of the catalyst. While in the heterogeneous case electrode poisoning by either intermediates, byproducts, or impurities is limiting the lifetime, slow decomposition of the molecular catalyst due to undesired side reactions represents the biggest problem for homogeneous catalysis. In both fields the catalyst lifetime is typically far below 100 h, and therefore, substantial improvements are needed in order to establish this technology on an industrial scale.

availability of electrons is improved due to distribution over the same three-dimensional space as the substrate (instead of a two-dimensional dispersion on the electrode surface). Second, the oftentimes unfavorable position of the electron transfer equilibrium is moved in the desired direction if rapid and thermodynamically favorable chemical steps are following. With regard to a CO2 radical anion resulting from an outer-sphere electron transfer, protonation or radical coupling are good examples portraying rapid and thermodynamically favorable follow-up reactions. In the second scenario shown in Scheme 2, a binding interaction between C red and the substrate (chemical activation) occurs prior to the ET from Cred to the substrate, leading to transient adduct formation, before Cox is regenerated and the product formed. This process is equivalent to chemical catalysis, with the corresponding implications regarding driving force and activation barriers (a more detailed discussion of the binding between catalyst and CO2 is provided in section 3.1.2). For two reasons, chemical catalysis represents the superior option with regard to CO2 electroreduction. First, the generation of a CO2·− intermediate in an outer-sphere electron transfer is unfavorable and associated with a high reorganization energy (linear molecule vs bent radical anion).85−87 Second, in the presence of CO2, typical outer-sphere reagents such as radical anions of hydrocarbons with extended π-systems (generated in aprotic media) rather undergo carboxylation via electrophilic aromatic substitution than an outer-sphere electron transfer,22,23 whereby an exception from this trend is represented by radical anions of benzoic esters and benzonitriles (see section 3.3.10.1).88,89 Consequently, the CO2 molecule is in most cases activated with a chemical catalyst, whereas redox catalysts/outer-sphere mediators are more suitable for the electroconversion of larger organic molecules.90−94 Within the family of homogeneous CO2 electroreduction catalysts, a coordinated transition metal ion typically represents the active center and systems from nonelectrochemical catalysis (e.g., hydrogenation catalysts, photocatalysts, bioinspired systems, etc.) are frequently adapted, along with the corresponding ligand platforms. However, transforming a “conventional” catalyst into an electrochemical one is not always possible, and a set of rules and preconditions has to be taken into account. Generally, the reduction potential of the catalyst needs to be in the appropriate range, while the reduced form of the catalyst is sufficiently stable, and the chemical step is feasible under the desired electrolysis conditions (reaction medium, temperature, etc.). A good homogeneous electrocatalyst must be able to operate close to the thermodynamic potential of the reaction it is supposed to catalyze, while the kinetics of the chemical step have to be fast in order to facilitate a rapid turnover. For metal−organic compounds, these factors can be adjusted to the desired electrolysis conditions by choice of a suitable metal and optimized by ligand tuning. In the last few decades, the progress which has been made in this area is truly impressive, both in optimizing the electrocatalytic performance and in mechanistic understanding. An overview of the molecular-defined catalysts studied in the context of electrochemical CO2 reduction is given in section 4, while mechanistic aspects are discussed in section 3.3. At the present stage it is difficult to tell whether homogeneous or heterogeneous electrocatalytic CO2 reduction is the more promising technology, as both of them have their advantages and drawbacks. For instance, product separation,

2.4. Paired Processes

All electrochemical reactions proceed in pairs, which means that an anodic half-reaction is paired with a cathodic halfreaction. While most of the cases presented in this review represent half-cell studies with a focus on the cathodic reaction and the characterization of a specific catalyst, it should be noted that electrolyte degradation occurs simultaneously as sacrificial anodic half-reaction. Particularly in organic electrolytes, this process leads to an accumulation of degradation products and to a deterioration of the electrolysis conditions. For testing a CO2-reduction catalyst in a prolonged controlled potential electrolysis, the choice of an appropriate anodic half-reaction can be beneficial for avoiding the accumulation of degradation products and maintaining constant conditions in the cell. One example for such a benign anodic process is the Kolbe oxidation of acetate giving only CO2 and ethane, thereby avoiding acidification and contamination with degradation products.95 The choice of an appropriate anodic process is also important from the economic and ecologic point of view. The coupling of two electrochemical reactions, each of which gives a useful product or intermediate for chemical synthesis, constitutes a general strategy for sustainable production of chemicals (“paired electrosynthesis”).96 While the most prevalent paired process is chloralkali electrolysis, a great number of examples for paired electroorganic syntheses have been reported.96−98 Such an approach can be used to improve the energy efficiency and the atom economy in comparison to separately conducted processes. Since the passed electric charge is essentially used twice, a maximum Faradaic efficiency of 200% can be achieved by definition. In the case of CO2 reduction, the common approach for upscaled reactors is the coupling to water oxidation (oxygen evolution reaction, OER).99−103 In a homogeneous approach by Meyer and coworkers, this has been achieved by using the same Ru complex for both a cathodic CO2-to-CO conversion and an anodic OER (see Scheme 3).104 The organic catholyte (0.1 M NBu4PF6/ CH3CN) was separated from the aqueous anolyte by a Nafion membrane and the electrolysis carried out in a two-electrode arrangement using BDD as anode and cathode. While applying a cell voltage of 3 V led only to modest current densities, the proof of principle was achieved and a catalytic activity for the Ru complex demonstrated for both the anodic and the cathodic half-reaction with 76% Faradaic efficiency (FE) for CO and E


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and gave the benzimidazole in 65% isolated yield, while the FE for CO production was 100% after 2 h.

Scheme 3. Electrochemical Approach Toward CO2 splitting by Meyer et al. Using the Same Ru Catalyst for CO2 Reduction and Water Oxidation104

3. MOLECULAR CATALYSTS AND ELECTROCHEMICAL CO2 REDUCTION 3.1. General Mechanistic Guidelines

In general, the electrocatalytic CO2 reduction process is driven by three processes: the electron transfer from electrode to the catalyst or a catalytic intermediate, interaction between the active form of the catalyst and CO2 (either via binding between metal and CO2 or via hydride transfer from metal to CO2), and protonation of catalytic intermediates. Each of these three aspects is separately discussed in sections 3.1.1−3.1.3. 3.1.1. Electron Transfer between Electrode and Catalyst. In most cases, the catalytic reduction of CO2 is induced by the electrochemical reduction of the catalyst, creating an active species which in the presence of CO2 enters the catalytic cycle. Regarding electron transfer kinetics, catalysts which can be rapidly and reversibly reduced are preferred over those with slow electron transfer kinetics. With regard to the reduction potential of the catalyst E0cat, it is important to note that it has to be situated in an appropriate range. Toward increasingly negative potentials, this range is confined by the onset potential of direct CO2 reduction Eonset (see Figure 1),

60% FE for O2. The concept was later refined in a follow-up work by the same group,105 whereby a 0.1 M NaHCO3 buffer solution (pH 6.7) served both as anolyte and as catholyte in a cell divided by an anion exchange membrane. Applying a cell voltage of 2.8 V using the same catalyst as depicted in Scheme 3 (1 mM), syngas (H2/CO = 2:1) and O2 were produced in 53% energy efficiency (j = 1.4 mA cm−2). Although the large ηcat of the OER can be reduced by the choice of an appropriate electrocatalyst, large-scale O 2 production is not desirable in terms of product value. A first attempt to replace such a low-value half-reaction by a useful synthetic process has been recently reported by Kubiak, Moeller, and co-workers (see Scheme 4).106 In this example,

Figure 1. Prerequisite relative positions of Eonset, E0cat, and E0(CO2/P) for establishing a catalytic cycle.

Scheme 4. Schematic Illustration of a Paired Electrolysis Involving Cathodic CO2-to-CO Conversion Catalyzed with a Re Complex and Anodic Synthesis of a 2-Substituted Benzimidazole Mediated by Ceric Ammonium Nitrate106

since a concurrent direct process is typically not desired. At less negative potentials, the range suitable for E0cat is limited by the thermodynamic (equilibrium) potential E0 of the corresponding reaction couple CO2/P, since catalysts with E0cat > E0(CO2/ P) lack the required driving force. Although E0(CO2/P) does not represent a strict confinement for the applicability of a catalyst, useful rates for CO2 reduction are only achieved when a certain overpotential η with respect to E0(CO2/P) is applied. In fact, the best catalysts which are currently known operate at a potential a few 100 mV negative of E0(CO2/P). Within a series of catalysts, E0cat can be tuned by variation of the metal 0 center and/or the ligand, whereby a less negative E cat oftentimes also corresponds to a decreased rate of CO2 conversion. More details regarding redox potentials, reversibility, and catalyst benchmarking are discussed in section 3.2.1. In an alternative scenario the initial form of the catalyst interacts with CO2 forming an intermediate, which is further reduced to form the product and the initial form of the catalyst. In this case, the reduction potential of the CO2 adduct and not E0cat needs to be considered. Although this is a rather rare scenario, corresponding cases were reported, and one example is discussed in section 3.3.5. 3.1.2. Coordination Chemistry of CO2. The electrocatalytic reduction of CO2 is initiated by the interaction between the active form of the catalyst and the substrate molecule. For the discussion of different electrocatalytic pathways of CO2 reduction, it is worthwhile to start with a

CO2-to-CO electroreduction catalyzed by [LReCl(CO)3] (with L = 4,4′-di-tert-butyl-2,2′-bipyridine; for mechanistic details of this reaction see section 3.3.1) was coupled to an oxidative condensation of syringaldehyde and o-phenylenediamine giving the corresponding benzimidazole product. The electrolysis was carried out in an 8:3:1 CH3CN/THF/MeOH mixed solvent system containing 0.8 M Et4NBF4 as supporting electrolyte using glassy carbon electrodes. The anodic reaction was mediated by catalytic amounts of ceric ammonium nitrate F


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consideration of its molecular properties and coordination chemistry. CO2 is a 16e− molecule with a linear equilibrium structure in its electronic ground state, therefore belonging to the D∞h symmetry group. Despite the overall nonpolar nature of the molecule, some reactivity can be anticipated, resulting from the electrophilicity of the carbon atom, the lone pairs of electrons on the oxygen atoms, or the presence of the πelectron density of the C−O double bonds. The qualitative molecular orbitals (MO) energy diagram of CO2 is depicted in Figure 2. The MOs with the highest

Figure 3. CO2 coordination modes to a single metal center.

metals since a strong charge transfer between the metal center and the antibonding π* orbital of CO2 takes place. The first stable complex with a η1-C coordination mode was reported by Herskovitz and co-workers for [Rh(diars)2(Cl)(CO2)] (diars = o-phenylene-bis(dimethyl)arsine).113 In the η2(C,O) binding mode, the CO2 molecule is also bent, and coordinated through the carbon and one of the oxygen atoms. The first transition metal complex with a η2(C,O) coordination mode to be structurally characterized was [Ni(PCy3)2(CO2)] described by Aresta and co-workers.114 In rare cases, CO2 binds to the metal center as a Lewis base via its oxygen atoms. The η1(O) end-on coordination is preferred with electron poor metals, and the CO2 molecule can remain linear or be weakly bent. The first evidence for O-coordinated CO2 was reported only in 2004 by Meyer and co-workers for an adduct with a six-coordinate uranium(III) species, whereby it should be noted that this linear η1-O coordination of CO2 is likely enforced by bulky adamantyl substituents within the U(III) coordination sphere.115 3.1.2.2. Insertion of CO2 into M−H Bonds. As discussed above, one of the possible means for the activation of CO2 is the modification of the properties of the CO2 molecule by direct coordination to a metal center. Another option is represented by the insertion reactions of CO2 into M−X bonds (M = metal; X = C, O, and N), which are relevant in the catalytic formation of new bonds to other molecules as a key step in the conversion of CO2 into added-value organic compounds.116 With regard to the catalytic production of formic acid, the insertion of CO2 into M−H bonds is of particular importance for the reductive processes.117 Such an insertion reaction has been shown to follow two routes, as depicted in Scheme 5, affording either the M−OCHO group or the M−COOH moiety, both having an applicative interest. The two methods of insertion are known as normal (A) and abnormal insertion (B). The names of these insertion

Figure 2. Qualitative molecular orbital diagram of carbon dioxide.110,111

relevance to the reactivity are the 1πg and 2πu orbitals, which play the role of HOMO and LUMO, respectively. The doubly occupied nonbonding 1πg MOs are mainly localized at the terminal oxygen atoms, whereas the empty antibonding 2πu orbitals are mostly centered on the carbon atom. Therefore, CO2 can be considered as an amphoteric molecule, where the oxygen atoms display a Lewis basic character and the carbon atom behaves as a Lewis acid center. With a slightly negative electron affinity (Ea) of about −0.6 eV86 and a first ionization potential (IP) of about 13.8 eV,107 carbon dioxide is a better electron acceptor than donor and the reactivity of the molecule is dominated by the electrophilic character of the carbon atom rather than the weak nucleophilic properties of the oxygen atoms. When the LUMO orbitals of CO2 are filled via an electron transfer, the resulting lowest energy state corresponds to a bent geometry. As an example, the CO2•− radical anion is a bent molecule with an equilibrium angle of 134°.108 Also, any interaction of carbon dioxide with an electron-rich metal center will induce a loss of linearity.109 Despite the vast number of different reduction pathways, only two major modes for CO2 activation exist. In the first scenario, CO2 is activated via coordination to a vacant binding site of a metal complex, whereas insertion into a metal−hydride bond represents the second possibility (see discussion below). While the majority of the systems presented in section 3.3 follow this generalization, exceptions are found when reduction processes are promoted by organocatalysts or ionic liquids, which are treated in section 3.3.10. 3.1.2.1. Coordination of CO2 to Metal Centers. As depicted in Figure 3, there are three common modes of CO 2 coordination, the η1-C, η2(C,O), and η1-O binding modes.112 The η1-C coordination mode is preferred with electron-rich

Scheme 5. Mechanisms for CO2 Insertion into a M−H Bond

G


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the proton availability on the thermodynamics of a particular CO2 electroreduction process is given by the Nernst equation and is visualized by plotting the standard potential of the reactant−product couple to the pH (“Pourbaix diagram”, see Figure 4).

mechanisms are based on the concept that in the M−H bond the metal center usually bears the fraction of the positive charge while the negative is accumulated on H. Thus, it is expected that during the insertion of CO2 into a M−H bond, one of the nucleophilic oxygen atoms should interact with the metal atom and the electrophilic carbon atom with the hydride ligand. This description matches with the transition state TS-1, where the metallaformate species I and II are generated. An example from CO2 electroreduction chemistry where such an insertion mechanism has been proposed is represented by the [Ir(PCP)H2] complex where PCP is a pincer-type ligand (for details see section 3.3.5). Conversely, less is known about the abnormal insertion of CO2 via metallacarboxylic TS-2 (pathway B in Scheme 5), for which only a few claims exist.109 In contrast, the formation of a M−COOH moiety via reaction of a metal carbonyl species M−(CO) with a hydroxide anion or via protonation of a M−(η1-C)-CO2 complex represent the more common scenarios.118,119 The latter one has also been frequently observed in the case of electrochemical CO 2 reduction when, for example, Mn, Fe, or Ni complexes in low oxidation states were used (see sections 3.3.1, 3.3.2, and 3.3.6). It should be noted that the polarity of the M−H bond can be influenced or even inversed by both the metal itself and the ancillary ligands. Consequently, the H moiety bound to a metal can in principle be transferred as an H atom,120 a hydride,121 or a proton.122 Moreover, the same metal hydride may show a dual behavior (proton or hydride transfer) depending on the receiving substrate.123 3.1.3. Role of the Proton Source. Studies on electrocatalytic CO2 reduction with homogeneous catalysts are often carried out in organic solvents under rigorous exclusion of water. Such conditions provide two advantages over aqueous electrolytes, namely, a broader potential window for the electrocatalytic experiment and a higher solubility of CO2. While, for instance, the saturation concentration of CO2 in water is only 0.04 M, it is significantly higher in DMF (0.2 M) and CH3CN (0.28 M).20 Although the solubility of CO2 in aqueous electrolytes can be significantly improved by altering the pH, solvated CO2 is always in equilibrium with carbonic acid, bicarbonate, and carbonate, whereby bicarbonate and carbonate are much harder to reduce than CO2 and are not reduced under regular electrolysis conditions. In order to enable the electrocatalytic reaction, a suitable proton donor (e.g., water or alcohols) is typically added in small defined quantities. When the electrocatalytic reduction is carried out in aqueous electrolytes, the pH value has a pronounced influence on the outcome of the reaction. Generally, the proton source plays a central role in CO2 electroreduction chemistry and therefore deserves particular attention. In the following, we will discuss the influence of the proton source on the thermodynamics (section 3.1.3.1), the kinetics (section 3.1.3.2), and the selectivity of a studied CO2 electroreduction process (section 3.1.3.3). Interestingly, for some catalytic systems CO2 reduction has been reported to proceed in aprotic solvents and without protic additives. These cases will be discussed separately in section 3.1.3.4. 3.1.3.1. Influence on the Thermodynamics. Considering the possible reactions of CO2 shown in Table 1, it becomes clear that the availability of protons is essential for the formation of most reduction products. Exceptions are represented by oxalate formation and reductive disproportionation to CO and CO32− (see section 3.1.3.4). The influence of

Figure 4. Pourbaix diagram for the reduction of CO2 in aqueous electrolytes. Reprinted from ref 124. Copyright 2013 American Chemical Society.

3.1.3.2. Influence on the Kinetics. As pointed out in section 3.1.2, it can be distinguished between two modes of CO2 activation. In the first scenario, the CO2 molecule fills a free binding site of the catalyst, while protons are assisting in further activation the CO2 ligand. Such a synergistic action of electronrich metal and protons is exemplified in Scheme 6 (eq 1), Scheme 6. Possible Roles of Protons in CO2 Activation

where η1-C-bound CO2 requires activation by proton donors in order to facilitate C−O bond cleavage. The shown reaction follows a so-called push−pull mechanism,125 where the electronrich metal center pushes electron density into the CO2 ligand while the protons facilitate the electron transfer by pulling out electron density, ultimately leading to C−O bond cleavage and liberation of water. In the second scenario, CO2 interacts with a M−H species according to Scheme 5 via either TS-1 or TS-2 to form a M− OCHO or a M−COOH species. Along this pathway, the proton plays an indirect role in CO2 activation by protonation of the electrochemically reduced catalytic intermediate forming the active metal hydride species (Scheme 6, eq 2). Choosing the appropriate proton source is essential for developing an efficient system, whereby reactivity and selectivity of the catalyst need to be taken into account. A pronounced effect on the rate of the homogeneous reaction may be observed when the electron transfer between the reduced form of the catalyst and CO2 is concerted with a proton transfer (proton-coupled electron transfer, PCET). In this case, the coupled protonation has a dramatic effect on the energy profile of the reduction process, which is exemplified in Figure 5 for a single-electron reduction/protonation sequence. While the corresponding uncatalyzed reaction would be H


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most catalysts provide activity for CO2 reduction only in the presence of proton donors, excessive acidity may lead to catalyst degradation, for instance, by hydrogenation of reduced ligands or formation of metal hydrides that react otherwise. Regarding undesired side reaction promoted by proton donors, the hydrogen evolution reaction promoted by the catalyst (indirect HER) often represents the major challenge. Since in many catalytic scenarios HER is thermodynamically favored over CO2 reduction, a selective CO2 electroconversion in the presence of proton sources is often achieved through favorable kinetic influences. Consequently, the electrocatalysts can only efficiently operate within a certain acidity window, and the optimization of the procedure is therefore associated with a systematic variation of type and concentration of the proton source. Particular attention should also be paid to the choice of the cathode material in view of its activity for direct HER. Among the commonly used electrode materials, ηcat for direct HER varies within a wide range84 and also depends on the type of proton source.140 For instance, platinum is frequently used for CV studies. However, due to its exceptional activity for HER at low overpotentials in protic solutions, it represents a bad choice for the study of other electrocatalytic reductions. In contrast, Hg electrodes provide a broad potential window for electrocatalysis experiments in the presence of proton donors, and in the past, it has been the most commonly used electrode material in this kind of studies. However, safety and environmental aspects have led to a gradual replacement by glassy carbon (GC) electrodes. 3.1.3.4. Electrolysis without Protic Additives. It is important to note that in some instances CO2 reduction has been reported to proceed in aprotic electrolytes without added proton sources. These reports can be divided into three general cases. (1) Oxalate formation and reductive disproportionation: Some of the possible CO2 reduction processes such as reductive disproportionation or oxalate formation (see Scheme 7) do not involve protons and may occur when

Figure 5. Typical energy profile of a proton-coupled electron transfer (PCET). Note that the two displayed energy profiles correspond to different reaction coordinates.

associated with a charged high-energy intermediate (A−) and high activation energies, the catalyzed PCET process simply bypasses these unfavorable intermediates and barriers. In the context of homogeneously catalyzed CO2 electroreduction, many cases are reported wherein PCET steps have been proven or postulated as part of catalytic cycles (see, for instance, refs 124, 126, and 127). The realm of PCET reactions represents a broad and very active field, and readers with interest in the details are referred to refs 128−134. Additionally, the kinetics can be strongly influenced by cooperativity between the metal center and an H-bonddonating group positioned on the ligand. A typical example for such an interaction is represented by the group of Co− or Ni−N4 macrocycle catalysts bearing N−H functionalities in the second coordination sphere. For such catalysts, a stabilizing effect of the N−H group on the M−(η1-C)CO2 intermediate via H bonding has been proposed,135,136 as exemplified for the CO2 adduct of the Co−N4 macrocycle in Figure 6 (top left). A

Scheme 7. Reductive Disproportionation (1) and Oxalate Formation (2) as Proton-Free Reduction Pathways Figure 6. Catalysts with prepositioned H-bond donors in the coordination sphere.

similar effect was achieved by introduction of phenolic groups to Fe porphyrin catalysts (see Figure 6, right)20 and recently to Mn−bpy complexes (see example in Figure 6, bottom left).137−139 Such prepositioned acidic groups have a dramatic impact on the catalytic activity and may change both the mechanistic sequence and the product selectivity. The current understanding is that their benign effect on the reaction rate lies both in the stabilization of the M−(η1-C)CO2 intermediate via H-bond donation and the increase of the local proton concentration (a more detailed discussion of the catalysts shown in Figure 6 is provided in sections 3.3.1.1, 3.3.2.1, and 3.3.4.1). 3.1.3.3. Influence on the Selectivity. The proton source may promote undesired catalytic reactions or influence the stability of the catalyst itself. Regarding the catalyst stability in the presence of Brønsted acids, it should be noted that basic catalytic intermediates (e.g., metal hydrides or negatively charged species) are typically generated at the cathode. While

the electrolysis is carried out under aprotic conditions.12,141,142 The reductive disproportionation can also be stimulated when Lewis acids such as Mg2+ or other metal ions are introduced, by addition of stoichiometric quantities of well-soluble perchlorate or triflate salts and/ or by using a sacrificial anode.19,143 A disadvantage of this method is the formation of insoluble carbonate salts which may block the electrode and inhibit catalysis. The activation with Lewis acids will be discussed in greater detail in section 3.3.2.1 in the context of iron porphyrin catalysts. (2) Anodically generated acids: Even in aprotic electrolytes, the formation of H-containing products such as HCOO− or CO/H2O may be observed. A possible explanation is provided by the anodic half-reaction. When electrolyte degradation serves as a sacrificial anodic process, acids may be generated, which migrate to the cathodic half-cell I


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E0cat in the presence and absence of substrate. In the case of a rapid chemical follow-up reaction, a so-called catalytic current icat can be observed, which originates from turnover of CO2 by the catalyst (in some cases, the potential at which icat appears can also be shifted cathodically with respect to E0cat; for examples, see section 3.3). Typical voltammograms for the case of an indirect cathodic reduction are exemplified in Figure 7. As

and balance the consumption of protons by the cathodic reaction. A typical example for such a proton-liberating degradation process is acetonitrile oxidation in the presence of perchlorate salts.144 (3) Hof mann degradation: A further atypical proton source is represented by Hofmann degradation of alkylammonium-based supporting electrolytes. This process, although thus far only rarely observed in homogeneously catalyzed CO2 electroreductions,145−147 can be triggered by cathodic formation of a basic catalytic intermediate B− and leads to the formation of the corresponding alkylamine and alkene (Scheme 8). Scheme 8. Hofmann Degradation of NBu4+ as a Possible Proton Source for CO2 Reduction in Aprotic Electrolytes

3.2. Techniques for Benchmarking and Mechanistic Studies

In the following section, we briefly discuss different experimental and theoretical tools which are frequently used to study mechanisms of homogeneously catalyzed CO 2 electroreductions, to benchmark catalysts, and to identify side reactions. It is important to highlight that using only one of the presented methods usually renders only poor and incomplete results. While cyclic voltammetry (section 3.2.1) provides access to kinetic data, it is not suitable for obtaining structural information. The latter can only be obtained when one or several additional methods are used, e.g., spectroelectrochemistry (section 3.2.2), preparation of catalytic intermediates (section 3.2.3), and quantum chemical methods (section 3.2.4). In the best-case scenario, all these techniques are combined in order to make the picture as complete as possible. 3.2.1. Cyclic Voltammetry. In order to characterize a molecularly defined electrocatalyst, it is typically studied regarding its redox potential, heterogeneous electron transfer rate, and chemical kinetics. For such benchmarking purposes, cyclic voltammetry (CV) is the most frequently used electroanalytical method, and in addition, it can be a useful tool for studying mechanistic aspects. In this section, we briefly discuss the most important points regarding CV and catalyst benchmarking in the context of CO2 electroreduction, whereby the presented principles are also applicable to the electrochemical conversion of other molecules such as cathodic oxygen or proton reduction, the anodic oxidation of water or hydrogen, as well as the electrocatalytic conversion of organic molecules. For a more comprehensive treatment of this topic, the reader is referred to other resources available in the literature.10,55,148 A thorough description of the fundamentals of CV can be found, for instance, in refs 149−152. 3.2.1.1. Catalytic Current. For a preliminary assessment of the feasibility of an electrocatalytic process, CV is particularly helpful. First, it can be used to study the redox behavior in the absence of CO2, which allows one to determine whether a catalyst candidate can be reduced within the appropriate potential regime and at which potential a possible electrocatalytic reaction can proceed. Second, a possible catalytic turnover can be detected by observing changes of the shape of the CVs according to certain diagnostic criteria. Furthermore, the efficiency of the catalytic process can be evaluated by comparing the peak currents i at the potential of the catalyst

Figure 7. Schematic illustration of the voltammetric analysis of the electrocatalytic CO2 reduction. (1) Voltammetry of the catalyst under Ar atmosphere. (2) Catalyst, CO2 atmosphere. (3) Blank electrolyte, Ar atmosphere. (4) Blank electrolyte, CO2 atmosphere. We note that the definition of ηcat shown here is just one among several others (for details see text).

described above, the voltammetry experiments should first be carried out under Ar in order to determine the reduction potential E0cat and the cathodic peak current iP of the catalyst (curve 1). A reversible behavior as shown in the example can be beneficial, since it demonstrates that, in general, the catalyst is quickly and reversibly reduced. This reversibility is usually preferred but not a mandatory requirement: Since within the catalytic cycle the reoxidation of the catalyst proceeds chemically by CO2, compounds with irreversible voltammetric behavior under Ar can in principal also be catalytically active. However, it should be noted that in this case there is always a kinetic competition between the reaction making the redox couple irreversible and the reaction with CO2. When the potential has been estimated and found to be situated in the appropriate range, the solution can be saturated with CO2 and the measurement be repeated. When the conditions mentioned above are fulfilled, a catalytic current (icat > iP) is now observed (curve 2). In this case, the reversible shape of the CV (curve 1) often changes to the irreversible form indicated with curve 2. The peaked shape of this CV results from competition between consumption of CO2 by the rate-determining step of the catalytic cycle with diffusion of fresh substrate to the electrode. Despite depletion of the substrate, no reverse redox wave is observed since the entire amount of electrochemically reduced catalyst is reoxidized by catalytic turnover. It should be noted that depending on the ratio between concentrations of catalyst and CO2, the scan rate v, and the rate of the catalytic process, the shape of the voltammogram can vary significantly (vide infra), and for a correct interpretation, the experimenter should be familiar with the characteristic shapes and the parameters influencing them. A detailed treatment of this matter can be found in refs 10, 55, J


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converted per time and catalyst unit contained in solution. In contrast, TOF in the context of homogeneous electrocatalysis is referring only to the activity of catalyst molecules present in the reaction−diffusion layer close to the electrode surface and independent of the total amount of catalyst contained in the electrolyte. In general, TOF is a function of the applied potential and can be calculated when kapp is known (vide infra). Additionally, TOF0, which represents the TOF value for E = E0(CO2/P), has been proposed as a benchmark criterion. The peaked shape of the catalytic wave in Figure 7 (red line), although often observed, represents merely a special case where consumption of CO2 by the catalytic process is competing with diffusion of CO2 to the electrode, the latter process limiting the current response at higher potentials. Not surprisingly, the extraction of kinetic data from such CVs is problematic, and therefore, experimental conditions have to be found which lead to pure kinetic control over the entire potential range. In general, the determination of k app and TOF is not straightforward, and there are multiple approaches depending on the scenario. For the sake of simplicity, we restrict the following discussion to a simple catalytic EC process (Scheme 9) and show under which conditions the shown methods can

and 148. For the sake of brevity, this discussion is restricted to a few very common cases (vide inf ra). An overlap between direct and indirect CO2 reduction as well as with cathodic decomposition of electrolyte is not desirable, and in order to examine this possibility, the voltammetric response of the blank electrolyte under Ar (curve 3) and CO2 (curve 4) should also be known. CV can also be used for the benchmarking of a catalyst. The parameters which are often extracted from these measurements and used as benchmarking criteria are discussed below. (1) Overpotential required for catalysis (ηcat): An important feature for catalyst benchmarking is the overpotential ηcat required for a catalytic process with respect to E0(CO2/ P). This parameter ηcat represents the additional driving force required to drive a reaction at a specific rate. A widely adopted definition for ηcat is the difference between E0(CO2/P) and the half-wave potential of the catalytic wave E cat/2 . 153 Another frequently used definition for ηcat is the difference between E0(CO2/P) and the redox potential for activation of the catalyst E0cat, as shown in Figure 7. Instead of Ecat/2 or E0cat, other parameters such as peak potentials or onset potentials for the catalytic wave are sometimes used, and this lack of a standard definition can lead to deviations of 100 mV or more. Caution is therefore necessary when ηcat of a specific catalyst is compared to literature values. With regard to E0(CO2/P), it should be taken into account that this value strongly depends upon the solvent as well as on concentrations and pKa values of acids present in solution (we note that here E0(CO2/P) is defined as the concentration- and temperature-dependent equilibrium potential, whereas E00(CO2/P) is used for the standard potential). If the E0(CO2/P) value for the used electrolyte system is not available in the literature (as it is often the case especially for organic electrolytes), it has to be determined using a detailed thermodynamic cycle which also takes into account all relevant acid−base equilibria and free energy changes of the participating species between aqueous medium and the organic solvent of interest.154 We note that ηcat may not be confused with the overpotential η: While the former is intrinsic to a catalyst under certain reaction conditions, the latter is used to define the difference between the applied electrode potential and the thermodynamic potential E0(CO2/P) of the electrochemical generation of a product P from CO2 electroreduction. (2) Apparent rate constant (kapp): The observed or apparent rate constant kapp represents the overall rate of the homogeneously catalyzed reaction, and it can be considered as a composition of rate constants of the involved steps.10 In the case of a rate limitation by an individual partial reaction (e.g., addition of the catalyst to CO2 or protonation of an intermediate), kapp is essentially a rate expression for this rate-determining step. The term TOFmax is frequently used synonymously.154 When kapp values of different catalysts are compared to each other, one should be aware that the values may vary depending on the method used for their estimation (for more details see section 3.2.1.2). (3) Turnover f requency (TOF): In homogeneous catalysis, TOF is used to quantify the activity of a catalyst and reflects the number of product molecules which are

Scheme 9. EC Process Underlying the Voltammetric Profiles Depicted in Figure 8

be transferred to more complicated multielectron transfer mechanisms, which are underlying homogeneously catalyzed CO2 electroreductions. A theoretical treatment of the voltammetry of catalytic EC reactions is provided in refs 155−159. As pointed out above, the shape of the voltammetric profile depends strongly on the reaction kinetics and can under certain preconditions be influenced by varying the concentrations of CO2 (thereby influencing kapp) and the scan rate v (determining the time scale of the experiment). The influence of v on the shape of the CV is exemplified in Figure 8.

Figure 8. Influence of the scan rate v on the shape of the voltammetric response for the catalytic EC process depicted in Scheme 9 (CVs have been obtained by digital simulation using the DigiElch package). For simulation parameters see ref 160. (Inset) Plot of the catalytic peak currents iP vs v. K


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With increasing v, a transition from the peak-shaped profile (black curve in Figure 8) to an S-shaped response, where forward and reverse scans almost trace each other, can be observed (green curve in Figure 8). This S-shaped curve reflects a situation where the concentration of the substrate in bulk solution [A]0 is equal to its concentration on the surface [A] throughout the entire potential regime. In other words, substrate depletion on the electrode surface is suppressed by shortening the time scale of the CV experiment, which allows for observation of the voltammetric response under kinetic control without substrate consumption. Notably, the maximum current iP rises with increasing v and approaches a maximum value, which is determined by the rate of the coupled chemical reaction (see inset in Figure 8). Consequently, CV studies of electrocatalytic processes should always be carried out with a set of different scan rates in order to evaluate how the diagnostic criteria of the profiles change with v and to search for pure kinetic conditions. An increase of [A]0 has a similar effect and can be applied if solely the variation of v does not lead to pure kinetic control (for gaseous substrates such as CO2 this option does not apply, since the experiments are typically carried out in saturated solutions for the sake of reproducibility). 3.2.1.2. Extraction of Rate Constants. When several preconditions apply (i.e., rate-limiting chemical step with k = kr, [A]0 ≫ [Cox]0), the plateau current icat of the S-shaped wave is described by icat = zFS[Cox ]0 · kappD

shaped CVs as the one shown in Figure 8 (green curve). As discussed before, both v and [A]0 can be increased until the S shape is obtained and icat is independent from the scan rate (pure kinetic conditions). As pointed out above, the discussed method for determination of kapp has been developed for a simple catalytic EC mechanism described by Scheme 9, whereas the catalytic reactions of interest often proceed via complicated multielectron−multistep mechanisms. However, the waveforms presented in Figure 8 are relatively common, and in many cases, the described treatment can be applied with slight modifications to multielectron−multistep sequences which are typically found in electrocatalytic CO2 reduction. An important precondition for the applicability of the EC equations to more complex scenarios is that the rate of the regeneration of the oxidized form of the catalyst Cox is limited by a single ratedetermining step. Consequently, an extraction of kapp from voltammetry data with the same strategy can be successful, whereby eqs 1 and 4 have to be extended by the parameter n (the number of catalyst units required per turnover of a substrate molecule A), giving expressions 5 and 6. A selection of electrocatalytic two-step two-electron reductions where these equations apply is shown in Scheme 10.161 Principally, these cases are relevant for the electrocatalyzed reduction of CO2 to CO or to formic acid.

(1)

where z is the number of electrons transferred in the redox event, F the Faraday constant, S the electrode surface area, D the diffusion constant of the catalyst C (assuming no significant changes between species Cox and Cred), and kapp the observed (pseudo-first-order) rate constant.55 The latter one is independent from the applied electrode potential and therefore intrinsic to the homogeneous process. The observed reaction rate kapp is proportional to the rate constant of the chemical step kr as well as to [A]0, as described by eq 2. kapp = k r·[A]0

icat = zFS[Cox ]0 · nkappD

(5)

RT ·nkapp icat 1 = · iP 0.4463 zFv

(6)

Scheme 10. Selection of Electrocatalytic Two-Step TwoElectron Reductions where eqs 5 and 6 Are Applicable161

(2)

When [A]0, [Cox]0, and D are known, kapp can now be extracted from the CV data. Whereas [Cox]0 is typically known and D can be estimated via the Randles−Sevcik treatment (see eq 3) or other methods, the determination of the active electrode surface area S is not straightforward. Consequently, eq 1 is often divided by eq 3, the latter one describing the peak current of a reversible redox couple as a function of concentration and scan rate (without coupled chemical reactions) in order to obtain eq 4.10 iP = 0.4463·zFS[Cox ]0 · RT ·kapp icat 1 = · iP 0.4463 zFv

zFvD RT

(3)

(4)

For the application of eqs 3 and 4, the peak current iP is extracted from the voltammogram recorded in the same solution before addition of substrate (compare Figure 7). In contrast to eq 1, the use of eq 4 enables the determination of kapp without separate determination of D and S. It is worth highlighting once more that the use of eq 4 is restricted to S-

In the case of an ECEC mechanism (Scheme 10, bottom), eqs 5 and 6 are only valid if the second electron transfer is more favorable than the first one (E02 > E01 in the case of reduction). Obviously, a number of further scenarios is conceivable and not to all of them the shown equations are applicable. A detailed treatment of various mechanistic pathways including a L


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derivation of the mathematic expressions for icat is provided in ref 161. In addition to the estimation of kapp, an investigation of the dependency of icat from [Cox]0 and [A]0 by variation of the concentrations is recommended when a new catalyst is developed. By plotting icat vs [Cox]0 and the square root of [A]0, respectively, the applicability of eqs 1 and 2 can be confirmed for the studied catalyst. In the case of CO2 electroreduction, variation of [CO2]0 in the CV experiment can give direct proof of whether CO2 is involved in the ratedetermining step of the catalytic cycle. It should be noted that in many cases the achievement of the S-shaped voltammetric response required for the treatment shown above cannot be achieved. This problem occurs when side reactions such as catalyst deactivation or other redox events perturb the plateau current or when substrate consumption cannot be eliminated by variation of the experimental parameters. Achieving the no-substrate-consumption-scenario is particularly problematic for CO2 electroreduction due to the rather limited solubility of CO2 in most solvents. When the typical S shape cannot be obtained, an analysis of the onset of the catalytic current can be attempted (“foot-of-the-wave-analysis, FOWA”) under the premise that no other electrochemical process is superimposed in this potential regime.58,124,161 This treatment utilizes the fact that also for peak-shaped catalytic responses such as the one shown in Figure 8 (black line), the catalytic process is under kinetic control when the driving force for the catalytic process (the electrode potential) is low. For processes such as those depicted in Scheme 10, eq 7 gives the relationship between the current response and the applied electrode potential for the entire catalytic wave (eq 5 is actually derived from eq 7 for the limiting case E ≪ E0cat). Again, eq 7 can be divided by eq 3 in order to eliminate the diffusion coefficient, the electrode surface area, and the scan rate dependency, yielding eq 8. i=

Figure 9. Influence of substrate concentration [A]0 on the voltammetric responses for a catalytic reduction process (left, simulated using the DigiElch package) and linearization according to the foot of the wave treatment (right).162 Adapted with permission from ref 154. Copyright 2012 American Chemical Society.

3.2.1.3. Turnover Frequencies and Catalytic Tafel Plots. When the apparent rate constant has been successfully determined, the turnover frequency (TOF) can be calculated according to eq 9 as a function of the applied overpotential η with regard to the equilibrium potential of the conversion of a substrate A to a product P (η = E − E0A/P).154 TOF =

(7)

RT

2.24· zFv nkapp i = zF 0 ⎤ iP 1 + exp⎡⎣ RT (E − Ecat )⎦

(

F 0 (EA/P RT

F − Ecat/2)⎤⎦ ·exp − RT η

(

)

(9)

Here Ecat/2 refers to the potential where the catalytic wave reaches one-half of its maximum (catalytic half-wave potential). As pointed out above, this TOF value is referring only to the catalyst molecules present in the reaction−diffusion layer close to the electrode surface and is independent of the total amount of catalyst contained in the electrolyte. This TOF−overpotential relationship is often used for catalyst benchmarking in catalytic Tafel plots as illustrated in Figure 10 (for three hypothetic catalysts C1, C2, and C3). For the plots, it is assumed

zFS[Cox ]0 nkappD zF 0 ⎤ 1 + exp⎡⎣ RT (E − Ecat )⎦

kapp 1 + exp⎡⎣

(8)

For extraction of kapp from the CV data, i or i/iP (forward scan) is plotted vs 1/{exp[zF/RT(E − E0cat)]}, whereby the current at the foot of the wave behaves linearly (as long as no side reactions occur in the onset range). A linear extrapolation of this regime allows retrieval of the desired kinetic data. The procedure is exemplified in Figure 9 for the case of substrate consumption at different substrate concentrations [A]0. Whereas replotting of the S-shaped CV (green curve) gives a straight line (right), an increasing deviation from the linear behavior is obtained with growing substrate consumption. However, in each case a linear range is found at low potentials (the foot of the wave), which can be used for determination of kapp. Again, the current−potential relationship is not applicable to all mechanisms, and in many cases, the equations have to be adapted to the specific scenario in order to allow for an extraction of kapp from the voltammetry data. A detailed treatment of a number of scenarios for FOWA, including derivations of the mathematic expressions for icat and other parameters, is provided in refs 161 and163.

Figure 10. Plot of TOF vs the applied overpotential η for three hypothetic catalysts C1, C2, and C3 with different TOFmax and Ecat/2 values reducing the same starting material. Selected parameters: TOFmax(C1) = 1100 s−1, E0A/P − Ecat/2(C1) = 0.15 V, TOFmax(C2) = 500 s−1, E0A/P − Ecat/2(C2) = 0.3 V, TOFmax(C3) = 100 s−1, E0A/P − Ecat/2(C3) = 0.15 V. M


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3.2.1.4. Mechanistic Modeling Using Digital Simulations. Software packages for the simulation of electrochemical experiments such as DigiElch or DigiSim enable the user to calculate the current response on the basis of physicochemical parameters (e.g., standard potentials, diffusion constants) and experimental conditions (temperature, scan rate, concentrations, electrode surface area, etc.). When coupled chemical reactions are involved, the mechanistic sequence of all of the involved elemental steps must be proposed along with the appropriate rate constants in order to obtain accurate results. When the rate constants are to be determined for a new system, a fitting of simulated CVs to the experimental ones can be carried out, which leads to the desired values as long as the proposed sequence is in accordance with the actual mechanism. When the mechanism is not or only partially known, different scenarios can be fitted and the best fits of each proposed sequence are compared to each other. As for other forms of mechanistic analyses, it should be kept in mind that a good match between experimental and simulated voltammograms does not necessarily ensure a match between the proposed and the actual mechanism. However, a few measures can help in order to increase the certainty of the mechanistic proposal. (1) Inclusion of side phenomena: Since even toward minor influences the voltammetric profile is highly sensitive, side phenomena should be taken into account if possible. Of course, it is necessary to first identify these side processes, and in this regard, we highlight once more the importance of the combined use of CV with other analytical tools. Reactions which are not participating in the catalytic cycle (such as catalyst deactivation or formation of resting states) may also be included in the mechanistic model in order to obtain reliable kinetic data. Product adsorption on the electrode surface is a frequently encountered inhibition mode, which can be modeled using simulation packages such as DigiElch. (2) Input of the highest possible number of physicochemical parameters: Diffusion constants of the participating species and the redox potentials of the catalyst should be measured separately and fed to the simulation. For instance, E0cat can be measured using CV in the absence of substrate, whereas the diffusion constant can be determined using the Randles−Sevcik treatment (eq 3). If possible, the rate or equilibrium constants of individual chemical steps may also be determined in separate experiments (e.g., by using nonelectrochemical methods such as stopped-flow measurements). (3) Variation of the conditions: The kinetic model including the fitted parameters (e.g., potentials, rate constants) should be validated by variation of the conditions. The experiments should therefore be carried out at multiple concentrations and scan rates and the CV data compared to the corresponding simulations. The best global fit reflects the most probable mechanism. Some examples where digital simulation has been applied in order to elucidate mechanisms of CO2 electroreduction with homogeneous electrocatalysts are given in refs 125, 143, and 165−168. For more general aspects regarding the technique, the reader is referred to refs 169−172. 3.2.2. Spectroelectrochemistry. Electroanalytical techniques are used for studying the thermodynamic and kinetic aspects of electron transfer between the electrode and a compound of interest by applying potential sweep or step

that all three catalysts are reducing the same substrate A to product P (with E0A/P) while having different TOFmax and Ecat/2 values. As demonstrated with the shown example, such plots are often used for comparing different catalysts to each other and for evaluating which catalyst is most suitable for a certain potential regime. In the semilogarithmic plot, TOF increases linearly until it reaches a plateau value at elevated η, which corresponds to kapp. As mentioned in the beginning, the parameter kapp is often termed TOFmax (the maximum turnover frequency). The intersection with the vertical line at zero overpotential corresponds to the TOF0 value, which has been proposed as a metric to determine the intrinsic performance of a catalyst.154 In the shown scenario it is clear that C1 (black line) represents the best catalyst among the three examples throughout the entire potential range. However, between C2 (blue line) and C3 (red line) it cannot be clearly stated which system exhibits better performance, since their TOF functions exhibit an intersection (at η = 0.26 V). Catalyst C3 has the more favorable Ecat/2 value and the less favorable TOFmax which results in a better performance at η < 0.26 V, whereas C2 is the better catalyst when η > 0.26 V. In summary, it can be stated that CV represents a useful tool for catalyst benchmarking and the elucidation of mechanisms of CO2 reduction. Voltammetry using a rotating disc electrode (RDE), which is not discussed here, represents a further powerful method, and its application to homogeneous electrocatalysis is treated elsewhere.55 As discussed above, TOFmax, TOF0, and η are extremely useful for comparison between different catalysts, since these parameters reflect the intrinsic performance of a system, unperturbed by undesired phenomena such as catalyst deactivation or other side reactions. However, it should be highlighted that this data does not necessarily reflect the catalyst performance under practical conditions, i.e., controlled potential electrolysis (CPE) at elevated potentials (which are required in order to obtain useful current densities). It is exactly under these conditions where the side phenomena play an important role, and therefore, long-term studies of a catalyst under high strain (i.e., electrolysis at high current densities) are equally necessary in order to obtain information about long-term stability, product selectivity, and Faradaic efficiency. The limiting turnover number TONlim was introduced as a useful criterion for the ratio between catalyst effectiveness and long-time stability by Costentin et al. It is defined according to eq 10, whereby kdec expresses the rate constant of catalyst deactivation (a parameter which can be extracted from i−t profiles of potentiostatic bulk electrolyses).164 TONlim =

kapp kdec

(10)

With regard to the estimation of the intrinsic parameters TOFmax, TOF0, and η, it is important to highlight that the method of evaluation of the data needs to be carefully adapted to the present scenario when complex multistep−multielectron mechanisms are involved (which is typically the case for electrocatalytic CO2 reduction) or side phenomena are perturbing the ideal voltammetric responses. In any case, the combination of CV with other techniques such as spectroelectrochemistry (see section 3.2.2), synthesis and examination of catalytic intermediates (section 3.2.3), and/or DFT calculations (section 3.2.4) is highly recommended. N


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distance between the beam and the electrode and thus to focus either on the bulk electrolyte or the reaction−diffusion layer. However, the method is associated with long optical path lengths and is therefore only practical for UV−vis spectroscopy. Case B represents the more frequently used approach, where optically transparent electrodes (OTE) consisting of gold191−193 or platinum194,195 minigrids are used. The collected spectra contain information from bulk electrolyte, reaction− diffusion layer, and solid−liquid interface. In order to minimize the absorption by solvent and supporting electrolyte, an optically transparent thin-layer electrode (OTTLE) cell is used.196 SEC experiments based on this cell design have been reviewed in the literature, and interested readers are referred to refs 197 and 198. With regard to the reflectance mode, two subclasses can be distinguished. The first one is based on external (or specular) reflection of the working electrode (case C) that consists of a polished and reflecting material such as platinum, gold, or glassy carbon.199−203 Before the beam is reflected, it passes a transparent window and the electrolyte. Accordingly, the collected spectra contain information about the bulk electrolyte, reaction−diffusion layer, and interface. Analogously to OTTLE cells, a thin-layer cell has to be used in order to avoid total absorption by the electrolyte. The second type of reflectance mode is based on internal reflectance (or attenuated total reflectance, ATR) utilizing the evanescence effect, where the beam is reflected from the electrode−electrolyte interface in an angle greater than the critical angle (case D). Here, the beam passes only the electrode, and the choice of the electrode material is therefore restricted to optically transparent materials such as Ge crystals or indium tin oxide (ITO).204−207 Comparing the two reflectance modes, external reflection represents by far the more prominent method for studying redox processes in solution: First, the method is much more sensitive toward dissolved species, since the beam passes the electrolyte solution. In contrast, a strong interference from adsorbates can be expected in the case of internal reflection (which is in turn very useful for studying heterogeneous electrocatalysts).208 Second, the choice of electrode materials is much broader when the external reflection mode is applied, and most of the common materials including glassy carbon can be used. Reported SEC studies where external reflectance was employed have been reviewed in the literature, and interested readers are referred to refs 173, 198, and 209. Metal carbonyl complexes are frequently encountered in the electrochemical CO2 reduction, either as catalyst precursor or as intermediate, and IR SEC is therefore very often the method of choice. The carbonyl ligands of the catalytic intermediates as well as CO2 and its reduction products give rise to strong infrared absorptions. Whereas the detection of free CO as product is difficult due to its poor solubility in polar solvents, the consumption of CO2 in the course of the electrolysis can be monitored by following the intensity of its characteristic signal in the range between 2330 and 2350 cm−1 (the exact position is dependent on the solvent).209 Typical cell designs for IR-SEC in this kind of studies are shown in Figures 12 and 13. The three-electrode thin-layer cell reported by Hartl et al. (Figure 12) has been frequently used in the analysis of molecular electrocatalysts for CO2 reduction (see Table 3).196 In the gastight setup the three electrodes are melt sealed into a polyethylene spacer which is sandwiched between two optical windows, whereby the detection beam directly passes the through minigrid working electrode. Straightforward variation

methods. Furthermore, information on coupled chemical reactions can be obtained, which allows for investigation of homogeneous electrocatalysis as shown in section 3.2.1. However, despite the variety of available electroanalytical methods, there is no possibility to obtain any information on the structure of the species formed near the electrode upon electron transfer. The obvious first approach toward obtaining structural information is to carry out a preparative-scale electrolysis and to isolate the electrogenerated species for characterization with spectroscopic methods. A complementary method is to couple one or more spectroscopic techniques with the electrochemical experiment (spectroelectrochemistry, SEC), enabling the in situ characterization of products and intermediates formed in the course of the electrolysis. This integrative approach has been successfully applied to IR,173−175 UV−vis,176−178 fluorescence,179,180 resonance Raman,181 and EPR spectroscopy.182−184 Although experimentally more difficult to realize, progress has also been made recently in NMR-SEC.185−188 Despite the large number of coupled methods which became available in recent decades, UV−visSEC and IR-SEC represent the most frequently used techniques in the spectroelectrochemical investigations of redox-active compounds in solution. This popularity stems from the following reasons: First, the two methods cover a photon energy range between 1 and 6.5 eV, which covers the entire field of electronic and vibrational transitions within molecules and thus allows one to study their electronic structure as well as changes induced by electron transfer reactions. Second, UV−vis-SEC and IR-SEC are associated with a relatively low experimental effort, whereby simple cells can be used in combination with relatively inexpensive spectrometers. Third, relatively high sensitivity paired with a short time scale for data acquisition allow for time-resolved measurements during potential sweeps at reasonable scan rates. Regarding IR and UV−vis spectroscopy, SEC studies of processes in solution can be carried out in the transmission and in the reflection mode (see Figure 11). With regard to the

Figure 11. Different cell types for SEC experiments.

spatial range which is supposed to be investigated (electrode− electrolyte interface, reaction−diffusion layer and/or bulk electrolyte), one has to carefully choose the experimental setup among cases A−D shown in Figure 11. One should be aware that the thickness of a reaction−diffusion layer is typically smaller as for a “regular” Nernstian diffusion layer, typically on the order of D/2kapp .159 Hence, with increasing efficiency of the catalyst, absorbance issues become more likely. For realization of the transmission mode, the beam can be directed either in parallel (case A)189,190 or through the electrode (case B). Setup A principally allows one to vary the O


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Table 3. Overview of IR SEC Studies on the Electrochemical CO2 Reduction with Molecular Catalysts Reported in the Literature catalyst

Figure 12. Schematic illustration of an IR OTTLE cell with threeelectrode arrangement inside the thin layer (left, top view; right, exploded view). (1) Working electrode (minigrid). (2) Counter electrode (mesh). (3) Ag wire pseudo-reference electrode. (4) Steel pressure plate. (5) Electrolyte inlet. (6) PTFE gasket. (7) Front CaF2 window. (8) Polyethylene electrode spacer. (9) Back CaF2 window. (10) PTFE holder. (11) Back plate. Adapted with permission from ref 196. Copyright 1991 Elsevier.

of the windows and the electrode materials allows for an adaption to UV−vis as well as to Raman spectroscopy.210−212 Another setup which is frequently used has been published by Kubiak et al. (Figure 13 and Table 3)213 It allows the use of commercially available PEEK-jacketed disc electrodes in a concentric arrangement with ring-shaped counter and quasireference electrodes, which are embedded in the cell body (see Figure 13, middle right). The ring shapes of counter and quasireference can be achieved upon leading the respective wire from the backside of the cell body through a bore into a circular slot (resulting in a flush surface). The cell body including the electrodes is pressed on the optical window separated by a PTFE spacer, which determines the thickness of the electrolyte layer (typically between 200 and 300 μm). The arrangement is coupled to a mirror unit, which directs the IR beam through the optical window into the cell, where it is reflected by the electrode and ultimately directed to the detector unit (see Figure 13, bottom right). It should be noted that the SEC techniques discussed above, while extremely useful for the analysis of species accumulating in the electrolyte, is not capable of monitoring short-lived intermediates. For resolving of the kinetics on shorter time scales (typically in the range between 1 and 10−3 s), stopped-

cell typea b

electrode + solvent

ref

[M(CO)4(bpy)] (M = Cr, Mo, W) [Mn(bpy)(CO)3(CN)] [Mn(tBu-bpy)(CO)3Br] [Mn(Mes-bpy)(CO)3Br] [MnBr(CO)3(IP)] [Mn(tpy)CO2Br] [Mn(CO)3(iPr-DAB)]+ [M(CO)3(R1-DAB-R2)Br] (M = Mn, Re) [Mn(pd-bpy)(CO)3Br] [Mn(CO)3(CNAr)2] [Fe4N(CO)12]

A B B B Ab B Ab B

Au + THF GC + CH3CN GC + CH3CN GC + CH3CN Pt + CH3CN GC + CH3CN Pt + CH3CN GC + CH3CN

214 215 16 19, 203 216 217 218 219

A B A

137 220 17, 221

[Fe(dophen)(N-MeIm)2]+ Co(II)−pincer complexes [Ni2(μ-dppa)2(μ-CNR)(CNR)2] [Ni(cyclam]2+ [Mo(CO)2(η3-allyl)(bpy)(NCS)]

B B B B Ab

[Ru(phen)2(ptpb)]2+

Bb

[Re(bpy)(CO)3Cl] [Re(tBu-bpy)(CO)3Cl] [Re(MeCyIm)(CO)3Cl] Re(I)−NHC complexes [Ru(Mes-bpy)(CO)2Cl2] [Os(CO)(bpy)Cl3]

A B B B B Ab

Pt + CH3CN GC + CH3CN Au + CH3CN/H2O GC + DMSO Pt + CH3CN Pt + CH3CN GC + CH3CN Pt + THF or PrCN Pt + DMF/H2O Pt + CH3CN GC + CH3CN GC + CH3CN GC + CH3CN GC + CH3CN Pt + PrCN

222 223 224 225 226 227 228 202, 229 230 231 232 233

a

Abbreviation for the cell types: A corresponds to Figure 12, B corresponding to Figure 13. bThe same setup has been used for IR and UV−vis SEC.

flow rapid mixing coupled with spectroscopic monitoring represents a valuable method (provided that the reaction steps of can be carried out nonelectrochemically).148 Of course, rate constants of certain elemental steps independently retrieved

Figure 13. Schematic view of a thin-layer SEC cell for the external reflectance mode. (Left) Disassembled view. (Top right) Assembled view. (Middle right) Bottom view of the electrode configuration. (Bottom right) Cross section of the thin-layer cell. P


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then converted with CO2 in the presence of proton donors to give the Re−COOH intermediate, whereby the kinetics of the reaction with CO2 in the presence of MeOH as proton donor were investigated by means of stopped-flow mixing in tandem with rapid-scan IR spectroscopy. The identity of the product formed in the stopped-flow process was confirmed by separate preparation of the Re−COOH intermediate via conversion of [Re(bpy)(CO)4]OTf with KOH (Scheme 11, bottom). Another example which illustrates the value of such nonelectrochemical experiments has been recently reported by Beller and Francke et al. (see Scheme 12).18,242 In two

with this technique may also be used for validation of digital CV simulation. With regard to the kinetic analysis of molecular electrocatalysts in small molecule activation, stopped flow has proved to be particularly useful.234−237 For the analysis of transient intermediates on shorter time scales down to the picoand nanosecond regime, transient absorption spectroscopy has been used in combination with laser pulses that trigger the generation of excited-state reductants (flash photolysis) which replace the electrode for the reduction of the catalyst. While these photoreductants (typically Ru, Ir, or Re polypyridyl complexes) activate the catalyst, the interaction with CO2 can be observed with time-resolved spectroscopy. Additionally, photoacids can also be generated in situ to study protoncoupled reactions. Although this approach is, as stopped-flow rapid mixing, a nonelectrochemical technique, it allows one to analyze the (nonelectrochemical) elemental steps of processes catalyzed by molecular electrocatalysts134 and has also been used for analysis of elemental steps in CO2 electroreduction.238,239 3.2.3. Preparation of Catalytic Intermediates. The isolation and characterization of catalytically relevant species not only represent a key to obtain a better mechanistic understanding but are also essential for studying the relationship between catalyst structure and its activity. To elucidate the structure of the reduced form of a catalyst, chemical reduction with KC8 or Na-amalgam is frequently carried out, since isolation of a reduced species from the electrolyte is usually more difficult (such cases can be found for instance in refs 135, 237, 240, and 241). In addition to a structural analysis of the reduced intermediate, one can use the corresponding IR and/or UV−vis spectroscopic data for referencing in spectroelectrochemical experiments. Furthermore, the reactivity of the reduced species toward CO2 should be tested if applicable in the presence and absence of proton donors. An example for the preparation and analysis of a catalytic intermediate has been reported by Kubiak et al. (see Scheme 11).237 On the basis of several other previous reports, the

Scheme 12. Investigation of a Possible Catalytic Intermediate in the Electro- and Photochemical Reduction of CO218,242

studies on the photo- and electrocatalytic reduction of CO2 with iron cyclopentadienone complexes of the type [Fe(cpd)(CO)3], it was demonstrated that these compounds are competent catalysts for the generation of CO and water. While the photochemical reactions were carried out in NMP using an Ir photosensitizer and triethanolamine as electron and proton donor, the electrochemical reaction proceeded in acetonitrile in the absence of proton donor. Iron cyclopentadienones such as those shown in Scheme 12 are well known from hydrogenations of unsaturated organic compounds (e.g., carbonyls and imines) via the bifunctional catalytic intermediate [Fe(cpd−OH)(CO) 2 H] (see Scheme 12, right).243 Consequently, the latter species was considered as a possible active species in the photo- and electrocatalytic processes. For control reactions, it was prepared by consecutive addition of aq NaOH and H3PO4 to a solution of [Fe(cpd)(CO)3] in THF.18,242 The behavior of the resulting hydride toward CO2 was then tested, whereby generation of 1 equiv of CO was observed in NMP upon irradiation with visible light. Furthermore, [Fe(cpd−OH)(CO)2H] was detected in operando FTIR measurements, which allowed for a mechanistic proposal for the photocatalytic process which includes the hydride intermediate.242 In contrast, no interaction between CO2 and hydride species was observed in the dark, and [Fe(cpd−OH)(CO)2H] could therefore be ruled out as active species of the electrocatalytic process.18 3.2.4. Quantum Chemical Methods. In recent decades quantum chemistry has become an important pillar for the mechanistic investigation of numerous catalytic processes, including the electrocatalytic reduction of CO2. Beyond its capabilities to explore possible reaction pathways including possible intermediates and transition states it serves as a versatile tool that complements a number of experimental techniques.244−248 For example, modern quantum chemical

Scheme 11. Preparation of a Catalytic Intermediate from Electrochemical CO2 Reduction Catalyzed by [Re(bpy)(CO)3Cl]237

neutral complex [Re(bpy)(CO)3(COOH)] has been proposed as a part of the electrocatalytic CO2 reduction mechanism involving [Re(bpy)(CO)3Br] as catalyst precursor, but could not be directly observed. The Re−COOH species was thus separately prepared in two different ways. In order to mimic the electrochemical reduction, the catalyst precursor was chemically reduced using KC8 (Scheme 11, top), yielding the fivecoordinate intermediate [Re(bpy)(CO)3]−. This species was Q


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term EXC, the so-called exchange-correlation energy. The functional EXC[σ] comprises of the electronic exchange energy arising from the antisymmetry of the Slater determinant and the correlation energy. The latter corresponds to contributions from the correlated motion of interacting electrons. Additionally, EXC[ρ] contains the remaining parts of the kinetic energy that are not covered by the kinetic energy functional of a single Slater determinant. Minimization of the energy functional in eq 12 with respect to the shape of the KS-MOs leads to a set of pseudo-one-electron eigenvalue equations

methods can be used to calculate essential quantities such as redox potentials and pKa values of catalysts and catalytic intermediates.246 Moreover, they provide a means to predict a variety of spectroscopic properties, notably infrared (IR) and UV−vis absorption spectra.247 This offers the ability to interpret spectroscopic results that would otherwise be difficult if not impossible to understand and furthermore allows one to verify or falsify proposed intermediate structures. Through those links it is possible to relate theoretical results to experimental data, and it is our conviction that theoretical efforts to elucidate mechanistic details of electrocatalytic CO2 reduction are most efficient when they are combined with experimental studies. This section is devoted to providing a brief overview of the most widely used theoretical techniques in the field of electrocatalytic CO2 reduction. For this purpose, it is convenient to first highlight the challenges that theoretical studies of electrocatalytic processes face. A central concern of such studies is to adequately describe the electronic structure of the system under consideration as it is usually the electronic structure that governs the reactivity. With many of the molecularly defined catalysts discussed in this review containing transition metal centers, this goal is sometimes difficult to achieve. Beyond electronic effects it is essential to take into account thermochemical contributions arising from translational, vibrational, and rotational motions of the involved molecules if one is interested in reaction free energies and reaction barriers. Furthermore, as all discussed catalytic processes occur in solution a suitable treatment of solvent effects is of great importance. This aspect becomes particularly critical during calculation of redox potentials and pKa values. 3.2.4.1. Electronic Structure Methods. The following discussion of electronic structure methods to calculate the electronic energy of the ground state, Ee0 ({RA}), focuses mostly on density functional theory (DFT) since it is applied in the vast majority of theoretical studies reviewed in this work. Due to the limited scope of this work, only a brief outline of its concept and some technical aspects can be given here. For a more elaborate introduction to DFT, the reader is referred to dedicated works on this matter.244,247,249 The basic idea of DFT is to interpret the electronic energy as a functional of the electronic density, ρ, i.e. Ee0 = Ee0[ρ], instead of solving the electronic Schrödinger equation for a given wave function Ansatz. In its currently used form, DFT rests on the two theorems by Hohenberg and Kohn and utilizes the Kohn− Sham construction.244,250 Within this theoretical framework the electronic density of a given system is approximated by the density of a single Slater determinant consisting of a set of single-electron wave functions {ϕi}, the molecular Kohn−Sham orbitals (KS-MOs). Accordingly

KS

F ̂ ϕi = εiϕi

Solution of these so-called Kohn−Sham equations yields the optimal KS-MOs as the eigenfunctions of the operator F̂KS, and the corresponding eigenvalues {εi} can be interpreted as orbital energies. Since the operator F̂KS explicitly depends on the shape of the MOs, the Kohn−Sham equations can only be solved iteratively. Hence, the process of orbital optimization is referred to as self-consistent field procedure. If supplied with the exact EXC functional, the Kohn−Sham equations would yield the exact ground state energy as stated by the Hohenberg−Kohn theorems. However, the form of the exact EXC remains unknown, and therefore, nonexact functionals have to be used in reality. Over the last decades, numerous exchange correlation functionals have been introduced, usually labeled with acronyms that refer to the last names of their developers. According to their underlying physical model these nonexact functionals can be grouped in different classes. The common starting point for the derivation of many current exchange correlation energy functionals is the homogeneous electron gas. A number of different approximations to the exchange and correlation energy have been formulated that solely depend on the local density of this oversimplified system. Nevertheless, functionals that employ such a local density approximation (LDA) yield surprisingly good results for molecules and solids with probably the most successful LDA functionals being VWM and PW92.251,252 A straightforward way to improve the performance of exchange correlation functionals beyond LDA is to take into account the inhomogeneity of the electron distribution in molecules. Therefore, within the generalized gradient approximation (GGA) the exchange and correlation energies incorporate the gradient of the electron density. This approach has met great success in chemistry, and some GGA functionals such as PBE, BP86, and BLYP are still very popular.253−255 A natural development after GGA are functionals that include the second and higher derivatives of the electron density and are hence labeled meta-GGA. Prominent members of that family are TPSS and M06-L.256,257 Considerable improvement in terms of accuracy for many systems was achieved by the admixture of a fraction of nonlocal Hartree−Fock exchange to the exchange functional which has led to the class of hybrid functionals. Although the calculation of Hartree−Fock exchange comes at the cost of significantly increased computational demands, hybrid functionals such as B3LYP, PBE0, and TPSSh are among the most widely used methods in applied quantum chemistry.258−261 Eventually “double-hybrid” functionals such as B2PLYP and PWPB95-D3 also include a fraction of wave function-based correlation energy evaluated from second-order perturbation theory.262,263 Of course, the associated gains in accuracy are also connected to a considerable increase of the computational costs. At this point it is important to emphasize

N

ρ(r) ≈ ρSD (r) =

∑ ∫ |ϕi(x)|2 ds i=1

(13)

(11)

where the integration is over spin degrees of freedom. Analogous to the energy expectation value of a Slater determinant the Kohn−Sham energy functional reads E0e,KS[ρ] = Te[ρ] + Ven[ρ] + J[ρ] + E XC[ρ] + VNN (12)

In addition to the kinetic energy (Te) and the Coulombic interaction terms (VeN, J, and VNN) that are straightforwardly defined for a Slater determinant, eq 12 contains the compound R


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Here S is the overlap matrix with elements Sμv = ⟨φμ|φv⟩ and F is the matrix associated with the operator F̂KS in eq 13. Importantly, the solutions of eq 15 are only equivalent to the true molecular orbitals when the atomic basis {φμ} is complete in the mathematical sense. Nevertheless, since the error introduced by the expansion in eq 14 has been shown to rapidly decrease with the number of atomic basis functions and such a matrix equation can be handled efficiently by modern computers virtually all present-day DFT studies on molecular systems employ eq 15 instead of directly solving eq 13.245 Owing to their favorable properties during the numerical generation of one- and two-electron integrals, Gaussian-type orbitals (GTOs) attached to the positions of the nuclei are most widely used as atomic basis functions for the investigation of molecular systems. Alternative types of bases contain Slatertype orbitals or plane waves which become particularly important for calculations with periodic boundary conditions. Without going into greater detail, we just state that the type and size of the basis set required to obtain accurate results from molecular DFT calculations generally depends on the calculated property. For example, calculations of hyperfine couplings in electron paramagnetic resonance spectroscopy require a high degree of flexibility in the core region of the molecule.299 For electronic energy calculations we advise the reader to use a standard GTO split-valence basis set of at least triple-ζ quality with two sets of polarization functions and one set of diffuse function, although this matter is certainly subject to debate. As the computational demand of a DFT calculation strongly depends on the choice of functional and the basis set size a feasible trade-off between accuracy and computational cost must be found for every system under investigation. 3.2.4.2. Thermochemical Aspects. Electronic structure methods such as DFT provide an efficient way to calculate electronic energies of molecular systems. However, the relative stability of reaction intermediates and transition states is determined by the free energy G = H − TS, especially at temperatures T > 0 K. The required thermochemical contributions to the free energy can be evaluated by means of statistical mechanics. In particular, the partition function provides the required connection between the properties of individual molecules and a macroscopic ensemble. The microscopic partition function of a single molecule (or a group of molecules) is defined as a sum of exponentials involving all possible states

that despite the growing sophistication in the development of exchange-correlation functionals there is no rigorous way to systematically improve them toward the exact result. Instead, modern functionals are based on simple models that are refined by derivation, physical reasoning, and data fitting. This often results in rather complicated expressions that in many cases contain empirical parameters. Notably, the amount of Hartree− Fock exchange incorporated in hybrid functionals is an empirical parameter, and its actual size strongly influences the obtained results for total energies as well as many molecular properties. Therefore, DFT may be regarded as a semiempirical method, and one should compare results with experimental data whenever possible. In particular, the relative energies of different spin states and other spin-dependent properties in open-shell systems are usually highly dependent on the choice of functional and require careful calibration to obtain a reasonable estimate of the calculation error.264,265 An overview over the applied functionals in the reviewed works on electrocatalytic CO2 reduction is given in Table 4. Table 4. List of Used Exchange Correlation Functionals in the Reviewed Works on Electrocatalytic CO2 Reduction functional name

functional type

described in refs

PBE BP86 BLYP M06 B3LYP

GGA GGA GGA hybrid hybrid

253 254 254, 255 273 258, 259

TPSSh PBE0

hybrid hybrid

261 260

used in refs 266−269 15, 225, 269−272 15 238, 274, 275 18, 126, 127, 135, 240, 267−269, 271, 272, 274, 276−288 289 290

In contrast to DFT, wave function-based methods are systematically improvable toward the exact result, and chemical accuracy on the order of ∼1 mEh is reliably achieved (for closed-shell systems) with the coupled cluster theory with singles, doubles, and perturbative triples, CCSD(T), method.291 While calculations on the CCSD(T) level were for a long time restricted to small molecules due to their high computational cost, the implementation of modern local correlation schemes has made them affordable for larger and more relevant systems.292−298 As a result they have found their way in a number of recent studies on catalytic CO2 reduction as a means to obtain accurate electronic energies.280−282,285,286 However, it should be noted that geometry optimizations and frequency calculations (see below) were still performed on the DFT level in these studies. Nowadays the Kohn−Sham equations can efficiently be solved for systems with up to several hundred atoms by means of modern computer codes. While there are many technical aspects to computer-aided DFT calculations that cannot be discussed here due to the limited scope of this work, we will briefly discuss atomic basis sets as they are an integral matter to be dealt with during a DFT study. Expansion of the molecular orbitals in a set of NBS atomic basis functions {φμ} according to

all states

q=

qN

form Q = N ! . In most molecular DFT codes an ideal gas is assumed, and thus, the latter term enters the expressions for the enthalpy and entropy

(14)

transforms eq 13 into a matrix pseudoeigenvalue equation F(ci)ci = εi Sci

(16)

where kB is the Boltzmann constant and gK is the degeneracy of state K. For a system with N distinguishable molecules (as in a perfect crystal) the partition function becomes Q = qN, while for N indistinguishable particles (as in an ideal gas) it takes the

∑ cμiφμ μ

gK e−EK / kBT

K

NBS

ϕi =

∑

(15) S

⎛ ∂ ln Q ⎞ ⎛ ∂ ln Q ⎞ ⎟ + k TV ⎜ ⎟ H = kBT 2⎜ B ⎝ ∂T ⎠V ⎝ ∂V ⎠T

(17)

⎛ ∂ ln Q ⎞ ⎟ + k ln Q S = kBT ⎜ B ⎝ ∂T ⎠V

(18) DOI: 10.1021/acs.chemrev.7b00459 Chem. Rev. XXXX, XXX, XXX−XXX

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To calculate q (and Q), all possible eigenstates of the molecular Hamiltonian and their corresponding eigenvalues EI are needed. Within the Born−Oppenheimer approximation, the electronic Hamiltonian can be separated from the nuclear Hamiltonian which in turn can be further separated into terms corresponding to translational, rotational, and vibrational motion of the nuclei. Accordingly, the total energy of any state K becomes a sum of the corresponding contributions EK = vib EeK + Etrans + Erot K K + EK , and likewise, the partition function can be rewritten as q = qeqtransqrotqvib. Therefore H = H e + H trans + H rot + H vib

(19)

S = S e + Strans + S rot + S vib

(20)

produce meaningful results, the use of QM/MM methods is rather limited in the field of electrocatalytic CO2 reduction. Instead, the vast majority of theoretical studies reviewed in this work apply continuum solvation models. Within such models the large number of individual solvent molecules is removed, and the space formerly occupied by them is filled by a continuous medium that has properties consistent with the solvent. Before discussing the concepts of modern continuum solvation models, it is advantageous to consider the process of solvation itself. Although they cannot be measured separately, the solvation free energy of a given molecule (or a small group of molecules) can be divided into different contributions.245,246 First, the solute and solvent may directly interact with specific solvent molecules, for example, via hydrogen bonding, leading to preferred orientations of the solvent relative to the solute. Such interactions are usually restricted to solvent molecules of the first solvation shell. Likewise, van der Waals interactions which act to stabilize the solute have a relatively short range. Electrostatic interactions, on the other hand, also involve solvent molecules at larger distances from the solute. Eventually the generation of a vacuum cavity in the bulk solvent in which the solute is to be placed (Figure 14) is associated with an

Although the summations over states in all partition functions run to infinity, their numerical value is finite, since the energies in the numerator of the exponents become larger. Without going into detail, we state that in virtually all reviewed theoretical works the different contributions to H and S are evaluated by applying standard approximations such as the particle in a box and the harmonic oscillator. For a detailed discussion of this subject the reader is referred to standard textbooks.245,246 In the case of molecular vibrations, the obtained harmonic frequencies are furthermore used to verify the minimum or transition state character of a given molecular structure and to simulate vibrational spectra.244,247,300−304 The foregoing discussion of thermochemical properties is entirely based on approximate solutions of the time-independent Schrödinger equation for single molecules (or small groups of molecules) with static conformations. An alternative approach is in principle provided by molecular dynamics and Monte Carlo simulations.300 Although the former was successfully applied to compute thermodynamic properties of pure CO2,301 both approaches are unfeasible for the investigation of electrocatalytic CO2 reduction for various reasons. 3.2.4.3. Solvent Effects. All of the reviewed electrochemical CO2 reductions occur in polar solvents with DMF, CH3CN, and H2O being the most widely used. It is of great importance to take into account the occurring solvent effects during any theoretical study of the reduction process in order to obtain a realistic picture of the experimental observations. In the following we briefly discuss the physical background of the solvation models used in the reviewed works. The central quantity in this context is the solvation free energy, ΔG0sol(A), that refers to the free energy change of molecule A leaving the gas phase and entering the condensed phase in the limit of an ideal solution. Hereby, attention must be paid to the definition of the standard state concentrations of both phases. For example, most theoretical works use standard states of 1 M for both phases, whereas it is conventional to use a partial pressure of 1 atm to define the standard state in the experimental context. Such differences in the definition of the standard state give rise to corrections of the free solvation energy. Owing to the prohibitive computational costs, an explicit inclusion of a sufficient number of solvent molecules in quantum chemical calculation is unfeasible. In contrast, hybrid quantum mechanics/molecular mechanics (QM/MM) approaches address this problem in a practical fashion by describing only a few solvent molecules in the vicinity of the solute with quantum mechanical methods while the remaining solvent molecules are treated classically.302 However, as they are computationally demanding, still suffer from conceptual problems in some cases, and require considerable insight to

Figure 14. Reaction field concept: solute is placed in a cavity that consists of overlapping spheres attached to each of its nuclei, while solvent is modeled as a continuous medium with a dielectric constant ε.

energy penalty. While the specific solvent−solute interactions cannot be modeled realistically by a continuous medium, continuum solvation models are frequently used to describe the latter three contributions to the solvation free energy.246,248 Most modern continuum solvation models employ the concept of a “reaction field”, where the solvent is considered as a uniform medium with a dielectric constant ε and the solute is placed in a suitable cavity.245,246 While early models used a single sphere or ellipsoid as cavity, modern models generate more realistic cavities like overlapping spheres with scaled van der Waals radii attached to each nucleus, as exemplified in Figure 14. With such a cavity the van der Waals and cavitation contributions to the solvation free energy are usually approximated as a linear function of the cavity surface area.245 The remaining electrostatic contributions are calculated by using the Poisson−Boltzmann (PB) equation of classical electrostatics that describes the electrostatic potential of electrolytes as a function of the charge density, the dielectric constant, and the ionic strength of the solution.246 In cases without added supporting electrolyte, the PB equation becomes the simpler Poisson equation. Since the PB equation cannot be solved analytically for arbitrarily shaped cavities, numerical T


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Figure 15. Thermodynamic cycles that can be used to calculate the pKa value of AH (left) and the reduction potential of A+ (right).

features an electron in the gas phase.246 Thus, one does not have to worry about the solvation of a free electron, but only the formation free energy of an electron in the gas phase is needed. For the free electron, like the free proton, the electronic energy is zero and the PV term and translational free energy amount to 0.006 kcal mol−1 at 298 K and 1 atm. In practice, the reduction potential is given relative to a chosen standard whose potential has to be subtracted from the absolute obtained value. For example, to obtain the reduction potential of A+ versus the normal hydrogen electrode (NHE) one needs to subtract 4.36 V from E0(A+).246 An alternative approach to pKa values and reduction potentials involves isodesmic reactions.310−312 Within this approach the reduction potential of interest E(A/A−) is calculated utilizing a reference potential Eref(B/B−) of a chemically similar redox pair that has been studied experimentally. With such a reference system at hand one may formulate

integration schemes like the surface area boundary element method are employed which leads to the polarizable continuum model (PCM).303 Over the years, different types of PCMs were developed that use different numerical integration schemes (IEF-PCM),304 create the cavity using specifically parametrized radii (SMD),305 or assume an isosurface of the electron density as cavity surface (IPCM and SCIPCM).306 A frequently used variation of the reaction field concept is the conductor-like screening model (COSMO) in which the medium surrounding the solute has an infinite dielectric constant like in a conductor.307 This construction greatly simplifies the underlying electrostatic equations, and in order to describe the electrostatic potential of a solvent with a finite dielectric constant ε the obtained free solvation energy is scaled by a factor of 2(ε − 1)/2(ε + 1). Similar conductor-like variations have also been implemented in the PCM approach (CPCM).308 3.2.4.4. Calculation of pKa Values and Redox Potentials. The calculation of solvation free energies provides theoretical access to reduction potentials and pKa values of possible catalytic intermediates, which are of great importance in mechanistic studies on the electrocatalytic reduction of CO2. A straightforward way to calculation of these quantities is via the direct thermodynamic cycles shown in Figure 15. In this way, one may use high-level quantum chemical calculations to obtain the gas-phase free reaction energies ΔG0(g) for the deprotonation (DP) or reduction (R) and subsequently add solvent effects. The pKa value of compound A and the reduction potential of compound A+ are defined as pK a(AH) = E 0(A+) = −

B− → B + e−

FEref (B/B−)

A + B− → A− + B

ΔGr0(e−)

E(A/A−) = Eref (B/B−) −

ΔGr0(e−) F

(23)

Of course, analogous arguments can be made for the calculation of pKa values. A considerable advantage of isodesmic reactions is that this approach avoids the calculation of the free energy of the gas-phase electron or the free solvation energy of the proton entirely. Furthermore, isodesmic reactions account for systematic errors introduced through the computational model. Owing to the difficulties arising from the description of transition metal compounds in multiple oxidation states with DFT, calculated redox potentials exhibit mean errors on the order of a few hundred millivolts depending on the used methodology and the investigated redox couple.313−316 When experimental potentials are compared to computed values, particular attention should be paid to the reversibility of the observed electron transfer between electrode and catalyst or catalytic intermediate. In this regard, it is important to note that the thermodynamic potential E0 of a catalyst or catalytic intermediate is experimentally not accessible when the electron transfer between electrode and electroactive species is followed by a fast and irreversible chemical step.151,152 In such cases, peak potentials (EP) or half-wave potentials (E1/2) are often used to characterize a catalyst or an intermediate. These parameters do not necessarily reflect the thermodynamic situation, and therefore, one needs to be cautious when these values are compared to computed potentials.

(21)

0 ΔG(sol) (R)

F

−FE(A/A−)

and obtain the reduction potential as

0 ΔG(sol) (DP)

RT ln(10)

A + e− → A−

(22)

where F refers to the Faraday constant. Evaluation of pKa values using eq 21 and the corresponding thermodynamic cycle in Figure 2 involves the formation free energy as well as the solvation free energy of the proton. As the electronic energy of the proton is zero, the former reduces to contributions from a PV enthalpy term and translational free energy. At the 298 K and 1 atm standard state, the free energy of the proton is thus ca. −6.27 kcal mol−1. Reasonable estimates for the solvation free energy of the proton can be obtained from experimental data, but its uncertainty is of several kcal mol−1 and therefore a source of error. Together with the intrinsic errors in calculated solvation free energies and gas-phase deprotonation free energies, this may lead to unsatisfying errors in predicted pKa values of up to 5 units or more.246,309,310 Alternative approaches that involve different thermocycles may improve the results for certain cases.309,310 Interestingly, the definition of ΔG0(sol)(R) in the thermodynamic cycle in Figure 15 for the reduction of A+ U


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3.3. Catalysts and Mechanisms

This section is intended to provide a general overview over the most important developments in the field with a focus on mechanistic aspects. Instead of an exhaustive treatment of each catalyst family, we focus on the discussion of some instructive examples which highlight important mechanistic aspects. For a comprehensive overview of the catalysts which thus far have been reported as active for the electrochemical CO2 reduction, the reader is directed to section 4. 3.3.1. Mn and Re Complexes. The most frequently studied family of homogeneous Mn and Re catalysts for CO2 electroreduction is represented by complexes of the type fac[M(bpy)(CO)3X] (with X = Br, Cl).317 The fac-[Re(bpy)(CO)3Cl] complexes and bpy-substituted derivatives are known to be highly active for CO2-to-CO conversion since early reports published by Lehn et al.13 and Meyer et al.12 in the 1980s. However, due to the importance of more sustainable methods and materials for CO2 reduction, the focus has recently shifted from Re toward the bpy−Mn tricarbonyl congeners (since all of the Mn and Re catalyst precursors discussed in this section have three facially coordinated CO ligands, the fac label is henceforth omitted). The discussion in this section is restricted to bipyridyl complexes, while a comprehensive overview of further active Mn and Re complexes is provided in section 4. 3.3.1.1. Mn Bipyridyl Catalysts. Despite the fact that Mn is several orders of magnitude more abundant in the earth’s crust than Re,318 the first studies on the electrocatalytic activity of [Mn(bpy)(CO)3Br] complexes have been reported only recently. In the first report from 2011, Chardon-Noblat and Deronzier et al. described the activity of [Mn(bpy-R)(CO)3Br] with bpy-R = 4,4′-disubstiteded 2,2′-bipyridine (R = Me, H) toward CO2-to-CO reduction.319 Compared to the Re analogues, the reduction proceeded at a significantly lower overpotential (approximately 0.4 V less). In the CPE experiment carried out at a potential of the foot of the catalytic wave, the catalysts turned out to produce CO with exclusive selectivity and a stability over several hours (with moderate current densities in the range of 0.06 mA cm−2). It was also demonstrated that in contrast to their Re congeners, the use of Mn complexes requires the addition of a proton source such as water. Compared to [Mn(bpy)(CO)3Br], slightly better results with regard to selectivity and stability were obtained with the methyl-substituted version [Mn(bpy-Me)(CO)3Br]. The system was further tuned in a follow-up study by Kubiak et al, where the related [Mn(bpy-tBu)(CO)3Br] catalyst was investigated with respect to type and concentration of the proton source.16 TOFmax values were extracted from the CV data using the treatment discussed in section 3.2.1, and optimum rates were found in the presence of 1.4 M TFE (TOFmax = 340 s−1).16,203 Electrolysis at −2.2 V vs SCE using a 5 mM solution of [Mn(bpy-Me)(CO)3Br] in acetonitrile with added TFE yielded CO in 100% FE with current density up to 30 mA cm−2 and a stability over several hours. The mechanism of this process was studied by several groups, whereby CV,16,319 IR-SEC,16,203 pulsed EPR spectroscopy,271 pulse radiolysis,239 and DFT calculations285 were used among further techniques. The CV of [Mn(bpy-R)(CO)3Br] (for R = tBu, Me, H) under Ar atmosphere displays two irreversible one-electron reduction waves in the cathodic scan around −1.7 and −1.9 V vs Fc/Fc+, as shown exemplarily for [Mn(bpy)(CO)3Br] in Figure 16 (red line). The signals have been assigned to the processes shown in Scheme 13. The first

Figure 16. Cyclic voltammetry of 1 mM [Mn(bpy)(CO)3Br] (red line) and 1 mM [Mn(Mes-bpy)(CO)3Br] (black line) under N2 atmosphere in CH3CN + 0.1 M NBu4PF6 (v = 100 mV s−1, GC electrode). Adapted with permission from ref 203. Copyright 2014 American Chemical Society.

Scheme 13. Processes Associated with the Signals of the CV Shown in Figure 16

wave corresponds to reductive dehalogenation yielding a radical species, which rapidly dimerizes to form [Mn(bpy-R)(CO)3]2. The dimer is further reduced at −1.9 V, yielding the fivecoordinate monomer [Mn(bpy-R)(CO)3]−. The intermediates shown in Scheme 13 have also been detected in IR-SEC experiments using a thin-layer cell in the external reflectance mode (as the setup shown in Figure 13). The results for the case of [Mn(bpy-tBu)(CO)3Br] are shown in Figure 17 (plot A). The initial state of the catalyst exhibits three v(CO) stretches characteristic for fac-coordinated tricarbonyl complexes (2028, 1933, and 1923 cm−1, black line), which gradually diminish with increasingly negative potential. In the progression through the reductions at −1.7 and −1.9 V (red line), new bands appeared at 1973, 1928, 1878, and 1850 cm−1 (assigned to [Mn(bpy-R)(CO)3]2), as well as at 1907 and 1807 cm−1 (assigned to [Mn(bpytBu)(CO)3]−). Further progression ultimately led to complete consumption of [Mn(bpy-tBu)(CO)3]2, leaving [Mn(bpyR)(CO)3]− as the only species in solution (blue line). While the identity of [Mn(bpy-tBu)(CO)3]2 could be confirmed by comparison of the IR-SEC results to literature data, an authentic sample of [Mn(bpy-tBu)(CO)3]− was separately prepared by reduction of the parent complex with KC8.16 A further improvement of the system was later achieved by Kubiak and co-workers by using a bpy ligand with sterically demanding mesityl groups in positions 6 and 6′ (Mes-bpy).203 This modification leads to major differences in the redox behavior compared to the previously studied [Mn(bpyR)(CO)3Br] complexes where R = H, Me, tBu. As shown in Figure 16 (black line), the second wave at −1.9 V (which corresponds to the reduction of the dimer) disappears, and instead, the [Mn(Mes-bpy)(CO)3Br] complex exhibits a quasireversible two electron reduction wave at −1.7 V. The difference in the redox behavior originates from the bulky V


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Figure 18. Linear scans of [Mn(Mes-bpy)(CO)3Br] under CO2 atmosphere in CH3CN + 0.1 M NBu4PF6 with various amounts of added H2O (v = 100 mV s−1, GC electrode). Adapted with permission from ref 203. Copyright 2014 American Chemical Society. Figure 17. IR-SEC results obtained using a thin layer cell in the reflectance mode (see Figure 13) under N2 atmosphere in CH3CN for [Mn(bpy-tBu)(CO)3Br] (A) and [Mn(Mes-bpy)(CO)3Br] (B). In each case, the initial complex having three bands (black line) is ultimately converted to the doubly reduced species [Mn(L)(CO)3]− (blue line), whereby the reduction proceeds via the dimer [Mn(bpy)(CO)3]2 in case A and via the singly reduced monomer [Mn(bpy)(CO)3] in case B (both indicated with red line). Adapted with permission from refs 16 and 203. Copyright 2013 and 2014 American Chemical Society.

faster (TOFmax = 130 s−1 for 3.2 M MeOH and 340 s−1 for 1.4 M TFE). Moreover, it was shown in CPE electrolysis experiments that the Mes-bpy catalyst is highly selective (98% for CO generation) and stable over more than 7 h. While for the case of sterically less demanding bpy ligands it is assumed that the catalytic cycle can be initiated by addition both of the dimer [Mn(L)(CO)3]2 and of the anionic species [Mn(L(CO)3]− to CO2,271,319 a pathway which proceeds exclusively via the doubly reduced species [Mn(L)(CO)3]− was proposed by Kubiak et al. for the case of [Mn(Mesbpy)(CO)3Br] (see Scheme 14). The sequence is initiated by 2-fold reduction of the catalyst precursor, whereby the second electron transfer is easier than the first one. In this step the active species [Mn(L)(CO)3]− is generated at −1.6 V, which then binds to CO2 to give a Mn(I)−COOH intermediate. Before this species can be protonated to induce the dehydration step, a reduction step is necessary (“reduction first pathway”). This reduction step requires a potential which is about 400 mV more negative than the one where the catalyst precursor is reduced. After water is liberated, the resulting Mn(0) species is further reduced to release CO closing the catalytic cycle. Furthermore, it was shown in a later work by Kubiak et al. that catalysis works efficiently when [Mn(Mes-bpy)(CO)3Br] is used in combination with a Lewis acid such as Mg2+ (replacing the role of the Brønsted acid in the C−O bond cleavage step).220 In this Lewis-acid-promoted case, reductive disproportionation of CO2 occurs (2CO2 + 2e− → CO + CO32−), whereby ηcat could be significantly reduced compared to the Brønsted-acid-catalyzed procedure. However, in anhydrous electrolytes this approach leads to the formation of insoluble carbonate salts, which after prolonged electrolysis may block the electrode and inhibit the reaction. In this context, it has been recently shown that the use of [Zn(cyclam)]2+ represents a good alternative, since a similarly benign effect on the catalytic performance is obtained while the formation of insoluble salts is prevented.320 A further major advance for the Mn−bpy catalyst series was recently achieved by Rochford and co-workers using a [Mn(bpy-R)(CO)3]+ catalyst precursor. According to their report, the mechanism can be directed toward the more favorable “protonation first pathway” (see Scheme 14) by a small structural change of the bulky bpy ligand.238 By replacing the mesityl groups with 2,6-dimethoxyphenyl units, the steric

bpy ligands, which hinder the dimerization of the singly reduced Mn complex. Scan rate-dependent CV measurements and IR-SEC (Figure 17, spectrum B) indicate an ECE process, where the initial reduction is followed by rapid dehalogenation yielding the [Mn(Mes-bpy)(CO)3] intermediate, which is then further reduced to give [Mn(Mes-bpy)(CO)3]−. Again, the identity of these species observed in the IR-SEC experiment was confirmed by separate preparation and characterization of the intermediates (same setup as the one shown in Figure 13). The singly reduced intermediate [Mn(bpy-R)(CO)3] was prepared by reduction of the parent complex with 1.3 equiv. KC8 showed to be only stable in solution, whereby the characteristic IR bands were identical with those observed in the IR-SEC experiment (υ̃ = 1973, 1883, 1866 cm−1). With an increase of the amount of reduction agent (2.5 equiv), the doubly reduced species [Mn(bpy-R)(CO)3]− could be generated. In the presence of 18-crown-6, the compound was crystallized and fully characterized, whereby the observed IR bands (1909 and 1808 cm−1) also matched those from the IRSEC experiment. The voltammetric profiles of [Mn(Mes-bpy)(CO)3Br] under CO2 atmosphere in the presence of various amounts of H2O are shown in Figure 18. While in the absence of proton donor, the presence of CO2 does not alter the voltammetry significantly (black line), a catalytic wave appears in the presence of H2O with an onset at about −1.9 V vs Fc/Fc+. The peak current rises with increasing proton donor concentrations until reaching a plateau value. Similar results were obtained in analogous experiments using MeOH and TFE as proton donors. The TOFmax values extracted from these data were 2000 s−1 for 3.2 M MeOH and 5000 s−1 for 1.4 M TFE.203 Compared to [Mn(bpy-tBu)(CO)3Br], the second best catalyst of this series, CO2 reduction is more than 1 order of magnitude W


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reduction of the M(I)−COOH intermediate, which is reflected by the change of the voltammetric behavior of the catalyst (Figure 19, left). Whereas analogously to [Mn(Mes-bpy)(CO)3(CH3CN)]+, a reversible two-electron reduction associated with the same ECE process occurs at −1.6 V under Ar, a pronounced change is observed under CO2 in the presence of suitable proton donors such as TFE. In the latter case, a second catalytic wave with an onset at the potential where the active [Mn(L)(CO)3]− species is generated occurs. Consequently, the protonation first pathway is the exclusive one below 2.0 V, while above 2.0 V both processes proceed in parallel. With CPE experiments it was shown that by using the methoxysubstituted Mn−bpy catalyst, CO is selectively produced at reasonable rates at a potential which is 0.55 V more positive with respect to the reduction first process. Further support for the catalytic cycle shown in Scheme 14 was obtained by FTIR spectroscopy. While under Ar atmosphere and SEC conditions the active [Mn(L)(CO)3]− was observed, the postulated sixcoordinate [Mn(L)(CO)4]+ intermediate was detected after CO saturation of a solution of the [Mn(L)(CO)3(CH3CN)]+ precursor. Moreover, pulse radiolysis in combination with fast scan IR spectroscopy allowed for detection of the short-lived [Mn(L)(CO)3] intermediate generated after first reduction of the precursor. The catalytic behavior has been further tuned by introducing phenolic groups to the coordination sphere of the Mn bipyridyl complexes (see Figure 20, left). In an approach by Bocarsly and

Scheme 14. Proposed Catalytic Cycle for the Elctrochemical CO2 Reduction Using Mn Catalysts with Bulky 2,2′Bipyridine Ligands203,238

influence which blocks the Mn(0)−Mn(0) dimerization is maintained, while the methoxy units have a pronounced effect on the rate-determining dehydration step. On the basis of their electroanalytical results and DFT calculations, the authors proposed a mechanism which analogously to the one proposed by Kubiak et al. starts with the two-electron reduction of the catalyst precursor followed by generation of the metallocarboxylic acid species. Due to a combination of electronic substituent effects with a hydrogen-bonding interaction between methoxy substituents and the carboxyl proton (see Figure 19, right), the activation barrier for C−OH bond cleavage from the Mn(I)−COOH intermediate (assisted by an external proton source ROH) is significantly lowered, opening up the protonation first pathway. The reduction of the resulting cationic intermediate is much more favorable compared to the

Figure 20. Structures of the Mn bipyridyl complexes studied by Bocarsly et al. (left)139 and Nervi et al. (right)137 having a local proton source.

co-workers, improved rates and a decreased ηcat with respect to unsubstituted [Mn(bpy)(CO)3X] were reported when a local proton source was introduced to the outer coordination sphere of the Mn catalyst.139 The described [Mn(hp-bpy)(CO)3Br] species (with hp-bpy = 6-(2-hydroxyphenol)-2,2′-bipyridine) has a phenolic proton positioned closely to the CO2 binding site, facilitating the rate-determining proton-assisted C−O bond cleavage step. DFT studies indicated that the local proton source lowers the activation barrier of the C−O bond cleavage/dehydration significantly compared to the unsubstituted Mn−bpy catalyst. In a related work by Nervi and coworkers, the electrocatalytic behavior of the [Mn(pd-bpy)(CO)3Br] complex (with pd-bpy = 4-phenyl-6-(phenyl-2,6diol)-2,2′-bipyridine, see Figure 20, right) was investigated.137 Mechanistic studies using CV and UV−vis−IR-SEC indicated that Mn(0)−Mn(0) dimerization is suppressed and that the mechanistic pathway is altered compared to the one shown in Scheme 14, mainly due to reductive deprotonation of the phenolic groups. Interestingly, formate generation was also observed, which was presumably promoted by formation of small amounts of a Mn−H species. In addition to the efforts made toward ligand tuning, a number of studies on the heterogenization of Mn bipyridyl catalysts is available. In one approach, different [Mn(L)(CO)3Br] catalysts were cast together with carbon nanotubes

Figure 19. (Left) Linear potential scan for [Mn({MeO2Ph}bpy)(CO)3(CH3CN)]OTf under Ar (blue line) and under CO2 in the presence of PhOH (green line). (Right) Proposed transition state for the dehydration step of the catalytic cycle. Adapted with permission from ref 238. Copyright 2017 American Chemical Society. X


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electron reduction generates a ligand radical and does not alter the coordination number, the reduction of [Mn(bpy)(CO)3Br] is metal based and therefore leads to a loss of the bromide ligand followed by fast dimerization. Nevertheless, 2-fold reduction results in both cases in the active [M(bpy)(CO)3]− (M = Mn, Re) species that binds CO2, which is in agreement with the experimental evidence.16,237,239 For the Re-based catalyst the CO2 binding is thermodynamically favored, whereas the Mn catalyst requires subsequent protonation as driving force (see Figure 21). This finding explains the need for a weak Brønsted acid for the Mn-based catalyst to generate CO, which is not observed for the Re-based catalyst.

in Nafion membranes, which were used as catalytically active electrode coatings in aqueous solution for CO2 reduction.321,322 Further reports comprise the attachment of [Mn(L)(CO)3Br] species to glassy carbon,323 carbon nanotubes,324 TiO2,325 and silicon nanowires (including their implementation in photoelectrochemical cells used for CO2 reduction).326 3.3.1.2. Re Bipyridyl Catalysts. As pointed out above, the long history of Re bipyridyl catalysts in CO2 electroreduction started in the early 1980s with reports on the activity of [Re(bpy)(CO)3Cl] by Lehn et al.13 and Meyer et al.12 One of the main differences to the analogous Mn catalysts is that CO2 reduction also works in the absence of acids, following the reductive disproportionation pathway (2CO2 + 2e− → CO + CO32−). The influence of type and concentration of added Brønsted acid was studied, and it was found that acids with rather low pKa values such as TFE and PhOH have a positive effect on the rate and catalyst lifetime.327 It was also found that water, despite its comparable acidity, has a much weaker catalysis-enhancing effect, and it was concluded that it can effectively compete with CO2 for the free binding site on the Re catalyst intermediate. It was also reasoned that in the course of the electrolysis, the catalytic process is inevitably slowed down due to continuous formation of water as byproduct. Further tuning of the catalytic performance was achieved by Kubiak and co-workers through alteration of the substitution in position 4 and 4′ on the bpy ligand.202 The influence of proton-responsive bpy ligands was also studied in two independent works.272,288 Notably, considerable efforts have been made toward heterogenization of the Re−bpy system, whereby thiophene-,328 alkynyl-,329 and vinyl-330substituted Re−bpy complexes were used to generate catalytically active polymer films on the electrode surface. Promising results were also obtained by using graphite-conjugated Re−bpy catalysts331 and welldefined nanographene−Re−bpy complexes.332 The main mechanistic features of the CO2 electroreduction process catalyzed by homogeneous Re bipyridyl complexes are very similar to those shown in Scheme 14. The [Re(bpy)(CO)3]− intermediate, formed upon two successive singleelectron reductions of the catalyst precursor, is generally assumed to be the active species, and its formation and reaction with CO2 has been studied in detail with EPR spectroscopy,333 IR-SEC,202 and stopped-flow experiments using the separately prepared anionic species (see also section 3.2.3).237,241 Furthermore, detailed DFT studies have been carried out by the Carter group in order to explain the high rates and selectivity as well as the proton dependence of the reaction.283,287 These studies have shown that the high selectivity toward CO2 binding over the thermodynamically favored proton binding and subsequent H2 formation is governed by reaction kinetics. More precisely, the transfer of a proton from methanol to the Re center is associated with a high reaction barrier that substantially exceeds the barrier for binding of CO2. Only in the presence of a strong acid, protonation of the catalyst becomes competitive to CO2 binding. The remaining steps in the catalytic cycle are predicted to follow a sequence of protonation, reduction, and acidmediated C−O bond breaking. A more recent study by Carter and co-workers employed a combination of DFT and correlated wave function-based methods to highlight mechanistic similarities and differences of the reaction catalyzed by [Re(bpy)(CO)3Cl] and the analogous [Mn(bpy)(CO)3Br].285 In the course of this study, it was found that while for [Re(bpy)(CO)3Cl] the first single-

Figure 21. Potential energy surface for the addition of H+ and CO2 to the active catalysts [M(bpy)(CO)3]−. Catalytic intermediates are depicted with solid lines, while transition states are shown in dashed lines. Adapted with permission from ref 285. Copyright 2014 American Chemical Society.

Furthermore, microkinetics simulations indicated that at minimum operating potentials (−2.1 V vs Fc/Fc+ for Re and −1.8 V vs Fc/Fc+ for Mn) and with phenol as proton source [Re(bpy)(CO)3(CO2)]− proceeds via reduction, protonation, and acid-mediated CO bond breaking (analogously to the “reduction first pathway” in Scheme 14), although a reversed order of reduction and protonation is in principle feasible. In contrast, [Mn(bpy)(CO)3(CO2)] − is predicted to follow the protonation first pathway at −1.8 V. However, it should be noted that unless the bpy ligand is equipped with Lewis-basic groups in positions 6 and 6′ (see Figure 19), the protonation first pathway is very slow and does not lead to significant catalytic currents (as exemplified in Figure 18 for the Mes-bpy derivative). For reasonable reaction rates, the reduction first process has therefore to be enforced by increasing the electrode potential to values around −2.0 V. 3.3.2. Fe-Based Catalysts. Fe complexes belong to the more frequently studied catalysts for CO2 electroreduction, although the reported examples are restricted mostly to porphyrins,334,335 cyclopentadienone complexes18 and Fe carbonyl clusters17,336,337 (for a comprehensive overview see section 4). Among these Fe catalysts, the porphyrin complexes are the best-characterized systems, and the progress made in terms of ligand optimization is truly remarkable. 3.3.2.1. Porphyrin Iron Complexes. The long history of this class of CO2 electroreduction catalysts started in the 1980s with the pioneering work of Savéant and co-workers, who initiated a series of studies by close examination of the tetraphenyl porphyrin complex [(tpp)Fe III ]Cl (see Scheme 15, top).95,125,143,334,335 First, the catalyst was tested in the absence of proton donors in a NEt4ClO4/DMF electrolyte, whereby Y


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in the absence of CO2 shows three reversible redox couples, associated with the formation of [(tpp)FeII], [(tpp)FeII]−, and [(tpp)Fe0]2−. In the presence of CO2 and a Lewis acid, a catalytic wave appears at the potential where the Fe(0) species is formed, indicating that catalysis is triggered by the last electron transfer step. By analyzing the kinetics via concentration-dependent measurements, it was found that the reaction is second order in CO2. The reaction order in Lewis acid depends on the valence of the cation (one for divalent cations and two for monovalent cations). A catalytic cycle which is consistent with the findings described above and with other observations has been proposed (see Scheme 15).143 It starts with the formation of the Fe(0) complex, followed by a nucleophilic attack of the metal center on the electrophilic CO2 carbon, forming an intermediate adduct. Subsequently, the Fe(0)−CO2 adduct reacts under the assistance of Mg2+ with a further equivalent of CO2 (scenario 1 in Scheme 15), which gives rise to a Fe(II) carbonyl complex. The authors propose that in this step, ion pair formation between the negatively charged oxygen atoms in the Fe(II)− CO2 adduct and Mg2+, is responsible for the acceleration of the catalytic process (see Scheme 16). In other words, this ion pair

Scheme 15. Proposed Mechanism for the Electrocatalytic Reduction of CO2 Using [FeIII(tpp)]Cl (reduction potentials vs SCE) in the Presence of Mg2+ and Brønsted Acids125,143

Scheme 16. Proposed Synergistic Action between Fe and Mg2+ in the Breaking of the C−O Bond in CO2143

catalytic behavior was found under formation of carbon monoxide.334,335 However, the catalytic process under aprotic conditions is accompanied by rapid deactivation of the catalyst, presumably caused by carboxylation and hydrogenation of the ligand, and only few turnovers are therefore possible. Later it was found that the activity and catalyst stability is dramatically enhanced when a mono- or divalent Lewis acid (with a nonnucleophilic counterion such as perchlorate) is added to the electrolyte.143,335 Several Lewis acids were tested, and it was found that the activity has the order Mg2+ ≈ Ca2+ > Ba2+ > Li+ > Na+. Bulk electrolyses and product analysis revealed that carbon monoxide and carbonate are mainly formed, with 10− 30% formate as byproduct, whereby the exact product ratio depends on the type of Lewis acid. The mechanism of this catalytic process has been studied by means of cyclic voltammetry in the absence and presence of CO2 (see Figure 22). The scan of the negative potential regime

formation allows the Lewis acid to pull out an electron pair which has been pushed from Fe(0) to CO2 in the previous step (“push−pull mechanism”), thereby weakening the C−O bonds of CO2. Consequently, the added divalent cation can be considered as efficient Lewis acid synergist, whereby it should be noted that it is required in stoichiometric amounts (leading to formation of MgCO3) to drive the overall reaction. The C−O bond-breaking step results in the formation of [(tpp)FeII(CO)], which has been confirmed by thin-layer SEC and comparison to an authentic sample of the Fe(II) carbonyl complex.335 In the next step, [(tpp)FeII(CO)] is reduced to [(tpp)FeI(CO)]− in a homogeneous process by a second electrochemically generated Fe0 species. The appearance of the Fe(I) carbonyl complex is also indicated by the relatively weak signal 2′′ in the reverse scan of the CV (Figure 22, right), which has been confirmed by comparison to the voltammetry under CO atmosphere (feature 2′′ basically reflects the influence of CO on the Fe(II)/Fe(I) redox couple). As a final step, the catalytic cycle is closed by decarbonylation of the Fe(I) complex. While the formal oxidation states of the Fe center in the catalytic intermediates are clearly assigned throughout this discussion, it should be noted that, in particular, the electronic structure of the active [(tpp)Fe]2− species is currently under debate. One possibility is represented by the Fe(0) resonance form shown in Scheme 15, where the additional electrons are stored at the metal center. Furthermore, the Fe(I) and Fe(II) resonance formula are conceivable, where the additional electron density is shared between the metal center and the porphyrin unit or entirely located on the ligand, respectively. In fact, a recent combined spectroscopic and computational study by Ye and co-workers indicates that a significant portion of the

Figure 22. Cyclic voltammetry of 1 mM [(tpp)Fe]Cl in DMF + 0.1 M NEt4ClO4 (100 mV s−1) in the absence of CO2 (left) and in the presence of CO2 and Mg2+ (right). Adapted with permission from ref 143. Copyright 1996 American Chemical Society. Z


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negative charge is located on the ligand.338 Accordingly, the reduction in THF is described as mostly ligand based, resulting in a porphyrin-based biradical [(tpp••2−)Fe] that is antiferromagnetically coupled to the intermediate-spin Fe(II) center. However, the reactivity of the reduced Fe complex toward alkyl halides (alkylation proceeds exclusively on the metal center) indicates that the Fe(0) mesomeric structure must also play an important role.339 Generally, it appears that the electronic structure strongly depends on the used solvent. Although the study of the influence of Lewis acids on the electrocatalytic CO2 reduction represents a milestone with regard to mechanistic understanding, the practical use of the elaborated procedure is rather limited. This is due to the fact that MgCO3 formed upon CO2 reduction precipitates from the aprotic electrolyte, thus rapidly blocking the electrode surface and inhibiting further CO2 reduction. In order to being able to analyze the kinetics of the process, the authors restricted their analysis to the onset of the catalytic wave using the foot-of-thewave analysis described in section 3.2.1.2. For CPE, a mercury pool cathode had to be used, whereby a constant regeneration of active surface was assured by agitation with a magnetic stirrer. Analogously to their study on the Lewis-acid-assisted electrocatalytic reduction of CO2, Savéant and co-workers reported on the positive influence of weak Brønsted acids.95,125 It was found that addition of 1-propanol, 2-pyrrolidinone, or CF3CH2OH leads to a dramatic enhancement of the catalytic wave, whereby the effect becomes stronger with increasing acidity. Furthermore, an unexpected effect of the type of proton donor on the selectivity of the reaction was observed: Whereas the use of 1-propanol in bulk electrolysis leads to the generation of CO (FE ≈ 60%) along with considerable amounts of formate (FE ≈ 35%), the use of the more acidic CF3CH2OH gives CO almost exclusively (FE > 96%). On the basis of thorough analysis of the voltammetry, the authors proposed a catalytic cycle which is consistent with this selectivity and resembles the one proposed for the Lewisacid-promoted process (see Scheme 15). The voltammetric profile under CO2 atmosphere in the presence of Brønsted acid (HA) is similar to that observed in the presence of Lewis acids (compare Figure 22), with the catalytic wave appearing at the potential of the Fe(I)/Fe(0) couple. The proposed synergistic action of Brønsted acids is illustrated in Scheme 17 and consists

(push−pull mechanism), whereby ion pairing is replaced by hydrogen bonding. The dehydration step in Scheme 17 was later examined more closely by variation of acid type and concentration as well as by measuring the kinetic isotope effect.124 On the basis of the obtained data, a proton-coupled electron transfer bond cleavage (PCETBC) was proposed as rate-determining step at moderate [HA]. The intermediate adducts shown in Scheme 17 also explain the formation of formic acid when a weaker H-bond donor such as 1-propanol is used. In this case, the electron density is less efficiently pulled out of the ligand, resulting in a higher basicity of the Fe-bound C atom. These circumstances allow for protonation of the carbon, which leads to breaking of the Fe−C instead of the C−O bond and thereby to the generation of formic acid. In contrast to the use of divalent cation synergists, the use of Brønsted acids leads to the formation of water, and consequently, the electrode surface is not blocked by precipitated products. However, when high concentrations of CF3CH2OH are used, catalysis becomes so strong that product inhibition, namely reversible CO adsorption on the electrode surface, plays a significant role. In these cases, the kinetic analysis had to be carried out using the foot-of-the-wavetreatment, whereas for the weaker H-bond donors, kapp was directly extracted from S-shaped voltammograms (see section 3.2.1.2). A significant improvement of the catalytic performance of the [(tpp)Fe] system was later achieved by installing phenolic groups in ortho and ortho′ position of the arene substituents ([FeTDHPP], see Figure 23, top left).20 The catalysisenhancing effect of these local proton sources in [FeTDHPP] was demonstrated for the case of a CO2-saturated DMF/0.1 M NBu4PF6 electrolyte in the presence of 2 M H2O by comparison of the methoxy-substituted equivalent

Scheme 17. Proposed Synergistic Action of the Iron Metal Center and Brønsted Acids on the C−O Bond-Breaking Process

of hydrogen-bond formation between negatively charged CO2 and two HA molecules. This is in agreement with the observed reaction order, which is 2 with regard to HA (at high [HA], saturation kinetics are observed, whereby the formation of the Fe−CO2 complex becomes the only rate-determining step). The nature of the synergistic effect of weak Brønsted acids is therefore the same as that described above with Lewis acids

Figure 23. (Top) Structures of the modified FeTPP catalysts. (Bottom) Correlation between TOF and η for [FeTDHPP] and [FeTDMPP] in DMF containing 2 M water and 0.23 M CO2 (dashed line, catalytic Tafel plot; thick gray segment, data obtained from the FOWA treatment; star, data obtained from a bulk electrolysis). Adapted with permission from ref 20. Copyright 2012 AAAS. AA


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[FeTDMPP] (see Figure 23, top right): The catalytic Tafel plot (Figure 23, bottom) clearly indicates the superiority of [FeTDHPP] throughout the entire overpotential regime, including a TOF0 which is 9 orders of magnitude higher compared to that of [FeTDMPP]. In a follow-up study on the mechanistic consequences of the incorporation of phenolic groups into iron porphyrins, Costentin and co-workers demonstrated that this modification of the tpp ligand affects the consecutive order of elemental steps (see Scheme 18).290 As for the [(tpp)Fe] system shown Scheme 18. Alteration of the Mechanistic Pathway of CO2 Reduction by Fe Tetraphenylporphyrins in the Presence of Phenolic Groups on the Ligand

Figure 24. Influence of electronic substituent effects of iron porphyrins on the catalysis of the CO2 electroreduction. Electrolyte: DMF + 0.1 M NBu4PF6 + 0.1 M H2O + 1 M PhOH saturated with CO2. Reprinted with permission from ref 340. Copyright 2016 American Chemical Society.

CO both in DMF containing a proton donor and in pH-neutral aqueous electrolyte.341 The latter finding is particularly remarkable, since the byproduct H2 was found only in small quantities. On further studying the influence of ionic groups on the catalysis of CO2 electroreduction, a dramatic catalysisenhancing effect was achieved when trimethylammonium groups were installed in the ortho positions of all four phenyl rings (structure G in Figure 25). A significant decrease of ηcat

in Scheme 15, the catalytic cycle is initiated by generation of the Fe(0) species, followed by addition to CO2. The resulting Fe(II)−CO2 adduct is strongly stabilized by intramolecular H bonding with the pendant OH groups, which was confirmed by DFT calculations. A consequence of these stabilizing interactions is that catalysis requires the uptake of a second electron, whereby this second reduction is more difficult than the first one (which is indicated by a precatalytic wave in the CV). This behavior is in contrast to the case of [(tpp)Fe], where the second reduction is easier than the first one and occurs after the dehydration step. The second electron transfer in the case of [FeTDHPP] was suggested to be concerted with the cleavages of the Fe−C and of one of the C−O bonds as well as with proton transfer (see last step in Scheme 18). Moreover, the reprotonation of the catalyst by phenol is expected to occur simultaneously with the electron transfer step. It was concluded that the role of the phenolic groups on the tpp ligand and the catalysis-enhancing effect lies both in stabilizing the FeII−CO2 adduct and in generation of a high local proton concentration. Recently, the influence of electronic substituent effects on the iron porphyrin-catalyzed CO2-to-CO conversion was studied by successive replacement of the phenyl rings by one, two, and four perfluorophenyl groups (see Figure 24, species B−D).340 At a given proton-donor concentration (PhOH), two opposing trends with regard to catalyst performance have been observed upon increasing the number of electron-withdrawing substituents. First, catalysis is favorably affected with regard to η. The significant decrease of η corresponds to a positive shift of the redox potential of the Fe(I)/Fe(0) couple by the electronwithdrawing substituents. Second, the same inductive effects seem to lower the catalytic rate constant, possibly as a consequence of decreasing the electron density on the Fe0 center of the reduced form of the catalyst, thereby lowering its nucleophilicity and its capability for addition to CO2. A further improvement of the [(tpp)Fe] system was achieved by Costentin et al. upon substitution of the phenyl rings with ionic moieties. By introducing four trimethylammonium groups to the para positions, a water-soluble catalyst was generated which is able to efficiently catalyze the conversion of CO2 into

Figure 25. Cyclic voltammetry of [(tpp)Fe] (A), fluorosubstituted iron tetraphenylporphyrins (B−D, for color coding see Figure 24), and (tpp)Fe species containing ionic moieties (E−G). Electrolyte: DMF + 0.1 M NBu4PF6 + 0.1 M H2O + 3 M PhOH saturated with CO2; v = 100 mV s−1. Reprinted with permission from ref 21. Copyright 2016 American Chemical Society.

(to only 220 mV) was reported with a concomitant increase of TOFmax to values as high as 106 s−1 in DMF containing PhOH as proton donor.21 Since this finding is in contrast to the electronic substituent effect described above for the cases A−D, which leads to opposing tendencies with regard to η and TOF, it was concluded that a different effect must play a role in scenario G. Although this effect is not yet entirely understood, it was proposed that shifting the trimethylammonium groups from the para to the ortho positions allows for an efficient Coulombic stabilization of the initial Fe−CO2 adduct. To date, structure G represents the most efficient homogeneous catalyst for the CO2-to-CO electroconversion reported in the literature, exceeding other reported catalysts based on Mn, Ru, Re, and Pd in terms of TOF over the entire η-range of the catalytic Tafel plot, whereby the FE for CO production is practically 100%. The long-term stability of the catalyst was examined, and it was AB


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was established with CV under Ar atmosphere. The cathodic scan of [Fe4N(CO)12]− in an aprotic electrolyte reveals two reversible signals at E1/2 = −1.23 and −1.6 V vs SCE (corresponding to −0.98 and −1.35 V vs NHE),74 which have been assigned to the [Fe4N(CO)12]−/[Fe4N(CO)12]2− and [Fe4N(CO)12]2−/[Fe4N(CO)12]3− redox couples (see Figure 26, bottom left, black line).342 Upon addition of water, a new anodic peak appears at 0.5 V in the reverse scan (see Figure 26, bottom right, blue line), which corresponds to oxidation of a hydride species generated at the less negative reduction potential. The structure of both [Fe4N(CO)12]2− and [HFe4N(CO)12]− was determined by IR-SEC and X-ray analysis of separately prepared crystalline samples of the catalytic intermediates. Under CO2 atmosphere, the cathodic scan shows that the catalytic cycle is triggered by the reduction at the less negative potential (see Figure 26, bottom right, red line), indicating that either [Fe4N(CO)12]2− or [HFe4N(CO)12]− represent the active species. By successful conversion of CO2 into formate using [HFe4N(CO)12]− in a separate nonelectrochemical control experiment, the hydride could be identified as the active species in the catalytic cycle. Consequently, [HFe4N(CO)12]− reacts with CO2, whereby concentration-dependent CV measurements revealed that the reaction is first order with regard to catalyst and CO2. After liberation of formate, the catalytic cycle is closed by electrochemical regeneration of [Fe4N(CO)12]− from [Fe4N(CO)12]. Interestingly, both the selectivity over H2 generation and the rate of formate production improves upon increasing the amount of water in the electrolyte (the best results are obtained in pH neutral water). In order to explain this observation, additional thermochemical measurements were performed using IR-SEC in combination with a thermodynamic cycle originally proposed by DuBois et al.17,344 It was found that the hydricity (hydride donor ability) of [HFe4N(CO)12]− is strongly dependent on the reaction medium, leading to favorable thermodynamics in water (exergonic reaction) and to an unfavorable situation in CH3CN (endergonic reaction).17 In a follow-up study, the analogous [Fe4C(CO)12]2− complex (see Figure 27) was studied as catalyst in aqueous buffered

found that at least at low average current densities (in the range of 0.1 mA cm−2) the turnover and selectivity remain unaffected for 3.5 days. It should be noted that the unprecedented catalytic performance is associated with a substantial synthetic effort for the synthesis of catalyst structure G, which requires six preparative steps (including chromatographic purification of the modified porphyrin ligand) with 9% overall yield. 3.3.2.2. Iron Carbonyl Clusters. A series of iron carbonyl clusters containing carbide and nitride interstitial atoms was studied with regard to their electrochemical properties both in aqueous and in organic media by Berben et al.17,337,342,343 Among the tested systems, the butterfly-shaped cluster [Fe4N(CO)12]− (see Figure 26, employed as [(diglyme)2Na]+

Figure 26. (Top) Proposed catalytic cycle for the electroreduction of CO2 to formate using [Fe4N(CO)12]−. (Bottom left) Cyclic voltammetry of 1 mM [Fe4N(CO)12]− in 0.1 M NBu4PF6/CH3CN electrolyte in the absence of water (black line) and with 5% H2O (blue line). (Bottom right) Cyclic voltammetry of the same cluster compound in CH3CN/H2O (95:5) under Ar (black line) and under CO2 (red line). Adapted with permission from ref 17. Copyright 2015 American Chemical Society.

salt) exhibits the highest stability toward acids and the most promising properties with regard to electrocatalytic CO2 reduction. Noteworthy, this catalyst represents one of the few molecularly defined systems which allows for selective CO2-toformate electroconversion. In the absence of CO2, [Fe4N(CO)12]− catalyzes H2 generation, which illustrates that proton reduction is thermodynamically feasible and thus that formate generation must be kinetically favored.17,343 With regard to CO2 reduction, the catalyst is operative in the pH range between 5 and 13, whereby the optimum performance is obtained at pH = 7. Under optimum conditions, the catalyst is able to generate formate with 95% FE at a potential of −0.95 V vs NHE,17 which corresponds to an overpotential of only 440 mV. The mechanism of the CO2 reduction catalyzed by [Fe4N(CO)12]− was proposed based on results of CV studies and IR-SEC experiments (see Figure 26, top).17 The formation of a hydride intermediate as the first step

Figure 27. Modified iron carbonyl clusters studied by Berben et al.221

electrolytes.336 It was found that despite the higher hydricity of [HFe4C(CO)12]2− compared to [HFe4N(CO)12]−, only H2 is formed under CO2 atmosphere. The authors concluded that the lower hydride-donating ability of [HFe 4 N(CO) 12 ] − provides the required selectivity for formate generation and, consequently, that there is only a narrow range on the hydricity scale, where formate production is thermodynamically possible and no hydrogen is produced (“formate window”). In a further report by Berben et al., the electrocatalytic properties of the structural analogues [Fe4N(CO)11(PPh3)]− and [Fe4N(CO)11(PPh2(CH2)2OH)]− were described (see Figure 27).221 It was found that the phosphine ligands cause a negative shift of the reduction potential compared to AC


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[Fe4N(CO)12]− into a regime where direct proton reduction at the GC electrode is competing. Moreover, the proton relay in the latter compound changes the selectivity dramatically: Whereas the PPh3-substituted generates exclusively formate (H2 is formed as byproduct by background reduction at the electrode), only H2 generation was observed when the cluster containing the proton relay was used. Again, thermochemical measurements were performed using IR-SEC, and both the hydricity values of the hydride intermediates and the pKa value of the hydroxyl group in [Fe4N(CO)11(PPh2(CH2)2OH)]− were determined. It was found that the acidity of the proton relay is rather low and thus does not account for the high selectivity for H2 generation. Therefore, the authors concluded that the proximity of the proton donor to the reactive center plays the key role and that proton delivery kinetics are determining the selectivity in the presented case. 3.3.2.3. Fe Cyclopentadienones. Recently, a new system for the electrochemical CO2 reduction based on (cyclopentadienone)iron−tricarbonyl complexes (Figure 28)

does not proceed via an iron hydride intermediate (see also Scheme 12 in section 3.2.3 for more details).18 3.3.3. Ru and Os Compounds. Among the plethora of Ru catalysts used in the context of CO2 electroreduction (for a complete overview see section 4), those bearing bpy ligands have been most intensively studied. The first reports date back to the 1980s, when the catalytic properties of [(bpy)2Ru(CO)Cl]+, where the bpy ligands are located in cis positions to each other, were described by Tanaka and co-workers.349−351 The complex catalyzes CO2 reduction at −1.5 V vs SCE in buffered H2O/DMF (9:1 v/v) solutions, whereby a strong influence of the pH value on the product selectivity was found. While under slightly acidic conditions (pH 6) a mixture of CO and H2 was obtained, HCOO− and CO are produced in almost equimolar amounts along with H2 at pH 9.349 The voltammetric profile of [(bpy)2Ru(CO)Cl]+ exhibits a reversible wave at −1.2 V and an irreversible wave at −1.4 V, the latter one being assigned to the reductive dehalogenation affording five-coordinate species [(bpy)2Ru(CO)]. This neutral Ru complex represents the active species which initiates the catalytic cycle in the presence of CO2 by spontaneous formation of a η1-CO2 adduct which, according to the mechanistic proposal by Tanaka et al. (Scheme 19), is then Scheme 19. Proposed Catalytic Cycle for the Electrochemical CO2 Reduction Using [(bpy)2Ru(CO)Cl]+ or [(bpy)2Ru(CO)2]2+ 349

Figure 28. (Left) Structure of the (cyclopentadienone)iron-tricarbonyl complexes A−D used by Francke et al. for CO2-to-CO electroreduction. (Right) Species E as proposed intermediate of the catalytic cycle.18

was reported by Francke et al.18 A distinct advantage of this catalyst series compared to other systems is that both the Fe complex and the cyclopentadienone ligand are synthesized in a single straightforward step by cycloaddition of two alkynylsilane units and CO in the presence of Fe(CO)5.345,346 It was demonstrated that Fe cyclopentadienone derivatives A−D are highly selective toward the CO2-to-CO conversion, whereby the reaction rates were found to be in the order A > B > C > D. Using a 0.5 mM solution of species A, CPE at −1.65 V vs NHE renders current densities in the range of 3−4.5 mA cm−2 with FEs for CO generation typically above 95%. Notably, these high current densities were realized in a dry 0.1 M Bu4ClO4/CH3CN electrolyte in the absence of any Brønsted- or Lewis-acidic additives. Whereas usually reductive disproportionation of CO to CO/CO 32− is observed under such aprotic conditions,12,334,347,348 water was formed as a byproduct in stoichiometric quantities in the reported case. As an explanation of this unusual behavior, the authors proposed that anodic degradation of the electrolyte represents the terminal proton source enabling water formation. As a key to the efficiency of the process, the reductive transformation of the cyclopentadienone ligand to a hydroxycyclopentadienyl unit was suggested. This step creates a local proton source which stabilizes intermediate E via H-bond formation and facilitates the C−O bond cleavage step. A further interesting mechanistic feature is that in contrast to other previously reported nonelectrochemical applications of Fe cyclopentadienones such as hydrogenation of carbonyl compounds243 or photochemical CO2 reduction,242 the electrochemical CO2 reduction

protonated leading to the Ru formyl intermediate.349 Under acidic conditions, this formyl complex is further protonated and dehydrated to give [(bpy)2Ru(CO)2]2+, which is then reduced and decarbonylated closing the catalytic cycle. In neutral or slightly basic media, the formyl intermediate can also undergo a two-electron reduction under participation of a proton to generate formate and the pentacoordinate Ru(0) complex. In a control experiment it has been confirmed that the separately prepared intermediate [(bpy)2Ru(CO)2]2+ catalyzes CO2 reduction equally. Furthermore, it has been shown that both intermediates [(bpy)2Ru(CO)(COOH)]+ and [(bpy)2Ru(CO)(CO2)] can be formed from [(bpy)2Ru(CO)2]2+ by addition of 1 or 2 equiv of OH−, respectively.352 Using UV−vis spectroscopy at various pH values, it has also been established that all three species exist in equilibrium in aqueous solution AD


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(see Figure 29).349 The equilibrium distributions of the three intermediates shown in Figure 29 also explain why at pH 9 almost equimolar amounts of formate and CO are produced.

1a). The generated acetone is then subjected to carboxylation, possibly via abstraction of an α-proton by one of the catalytic intermediates and nucleophilic attack at CO2, whereby the liberated proton induces the generation of 1 equiv of HCOO− via catalytic reduction of a second equivalent of CO2 (Scheme 20, 1b). The same reactivity was used for the electrochemical carboxylation of acetophenone and hexan-1-one using [Ru(bpy)2(CO)2]2+ in a related study.353 An interesting outcome was also obtained upon use of Me2NH2Cl as proton donor in CH3CN electrolyte (Scheme 20, 2). In this case, DMF and water are produced (along with formate), whereby the carbamoyl species [Ru(bpy)2(CO)(CONMe2] has been identified as intermediate of the transformation using IR and NMR spectroscopy.354 In contrast to their bis(bpy)-coordinated congeners, Ru catalysts with only one chelating ligand such as [Ru(bpy)(CO)2Cl2], [Ru(bpy)(CO)2(CH3CN)2]2+, or [Ru(tpy)(CO)Cl2] tend to undergo cathodic formation of polymers of the type [Ru0(L)(CO)m]n having Ru−Ru bonds.355−357 Due to their poor solubility in CH3CN or aqueous electrolytes, these polymers tend to form films on the electrode surface which can be used as heterogeneous electrocatalysts for CO2 reduction, typically leading to CO and formate. Interestingly, it has been shown by Kubiak and co-workers that using a bulky bipyridine ligand (Mes-bpy = 6,6′-dimesityl-2,2′-bipyridine) can inhibit Ru−Ru bond formation and electrodeposition. Consequently, the studied [(Mes-bpy)Ru(CO)2Cl2] complex represents a homogeneous electrocatalyst, promoting the CO2-to-CO reduction with 95% FE in CH3CN in the presence of Brønsted acids.232 The CO2 reduction process via the Ru−carboxylate species shown in Scheme 19 is closely related to the same process using other catalysts such as Fe porphyrins (see section 3.3.2.1) or Ni cyclam (see section 3.3.6), where C−O bond cleavage occurs on a metallocarboxylate species generating a carbonyl complex and water. In each case, this C−O bond-breaking step is facilitated by addition of mild Brønsted acids, and in the case of Fe porphyrin, a dramatic enhancement of the catalytic activity was obtained upon use of phenolic ligands as internal proton donors (see Scheme 18 and Figure 23). However, it has been demonstrated by Fujita and co-workers that the installment of OH groups in the second coordination sphere does not necessarily have to lead to an enhancement of the activity. Testing their hypothesis that electron-donating OH groups in

Figure 29. Distribution of the different Ru species in aqueous solution at various pH values. Adapted with permission from ref 349. Copyright 1987 American Chemical Society.

An interesting change in the product selectivity was observed when Ru−bpy catalysts were used in combination with other proton sources and electrolytes (Scheme 20). Electrolysis of Scheme 20. C−C and C−N Bond-Forming Reactions upon CO2 Electroreduction with Ru Catalysts146,354

CO2 in a dry electrolyte consisting of CH3CN/DMSO (1:1) and 50 mM NMe4BF4 in the presence of catalytic amounts of [Ru(bpy)2(qu)(CO)]2+ (qu = quinoline) afforded acetone and small amounts of acetoacetic acid along with CO and HCOO−.146 In this reaction, 2 equiv of Me4N+ act as alkylating agent of the CO ligand of the reduced Ru complex (Scheme 20,

Scheme 21. Deactivation of [Ru(6-dhbpy)(tpy)(CH3CN)]2+ during Electrolysis in the Presence of CO2240

AE


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organic framework364 or grafted to a pyrolytic graphite electrode,363 these Co complexes efficiently catalyze the CO2to-CO conversion in aqueous electrolytes at relatively low overpotentials. Interestingly, the formation of small amounts of MeOH and CH4 as byproducts was observed at low pH when Co(tpp) was attached to the graphite electrode.363 A recent report compares the activity of Co(tpp) coated on carbon nanotubes to the same catalyst in solution: While the homogeneous reaction proceeds slowly and with considerable amounts of H2 as byproduct, the heterogeneous version shows dramatically enhanced rates and selectivity. 367 Such a preference for H+ reduction in the presence of CO2 is often encountered when homogeneous Co catalysts are used. Two examples where a preference for CO2 reduction was still achieved and the mechanism was thoroughly dissected are discussed in the following.369 3.3.4.1. Co N4-Macrocycles. Among the reported cases for homogeneous electrocatalysis with Co complexes, those with N4-macrocycic ligands play a prominent role.135,269,368−370 Although such complexes often favor H + over CO 2 reduction,371,372 their selectivity can be directed toward CO2 conversion as it has been demonstrated by Peters et al.135 In their combined synthetic/electrocatalytic study of the macrocycle complex [CoIIIN4H(Br)2]Br (see Figure 30, right), the

positions 4, 4′, 6, and 6′ of the bpy ligands of Ru complexes may lead to an increased nucleophilicity of the metal center, while the acidity of the OH group would further enhance the activity, the complex [Ru(6-dhbpy)(tpy)(CH3CN)]2+ (with 6dhbpy = 6,6′-dihydroxy-2,2′-bipyridine) was studied as electrocatalyst (see Scheme 21).240 Unexpectedly, the introduction of OH groups led to a poor catalytic activity and rapid deactivation of the catalyst. The deactivation pathway was thoroughly studied using cyclic voltammetry, acid−base titrations, IR-SEC, and DFT calculations. Under Ar atmosphere, the CV exhibits two irreversible reduction waves which have been assigned to reductive deprotonation of the catalyst under H2 formation (first step in Scheme 21). The two irreversible waves are followed by a third reversible wave, which can be assigned to (mostly tpy-based) one-electron reduction of the catalyst. While in the CV experiment a catalytic current was observed, the use of [Ru(6-dhbpy)(tpy)(CH3CN)]2+ in bulk electrolyses was associated with rapid current decrease and the formation of only small quantities of CO. The IR-SEC data for [Ru(6-dhbpy)(tpy)(CH3CN)]2+ shows that in the presence of CO2, a dicarboxylated species is generated, which is further reduced to form the stable Ru−CO complex (deactivation product, Scheme 21). CO cannot be released from the carbonyl complex, and catalysis is therefore inhibited. This finding is in sharp contrast to the related complex [Ru(bpy)(tpy)(CH3CN)]2+, which represents an active catalyst for CO2 electroreduction, and the inhibiting effect can therefore be unequivocally ascribed to the introduction of OH groups to the ligand. Although the attempt of the authors to create a more efficient catalyst was unsuccessful, the shown case illustrates the necessity for understanding deactivation pathways, providing important insights for a knowledge-based optimization of other systems. In contrast to its 4d and 5d neighbors Ru, Re, and Ir, Os has only rarely been used as a metal center for CO2 electroreduction catalysts.167,168,233,358 In these studies, bpy complexes were in the focus of investigation. Early reports by Meyers and co-workers in the 1980s indicated that [Os(CO)(bpy)2H]+ is active for CO generation under aprotic conditions, whereas the addition of small amounts of water leads to concurrent HCOO− formation.167,168 Analogously to Ru complexes bearing only one chelating ligand, a more recent study on the electrocatalytic properties of [Os(bpy)(CO)2Cl2] by Deronzier et al. showed that upon cathodic reduction, an Os−Os-bonded polymer of the type [Os(bpy)(CO)2]n is generated.358 In acetonitrile, the polymer is insoluble and forms a film on the electrode surface, which is active for CO2 reduction. Recently, it was shown by Hartl et al. that the film deposition upon use of Os−bpy complexes can be suppressed by choice of a suitable solvent, enabling their use as homogeneous electrocatalysts.233 3.3.4. Co-Based Catalysts. Compared to the neighboring 3d metals Fe and Ni there is only a small number of examples known, where Co complexes constitute efficient homogeneous catalysts for CO2 electroreduction. In several cases, immobilization of Co complexes to the electrode surface has proven to be an effective way to enhance the stability and activity compared to the corresponding molecular catalysts. These cases comprise Co chlorin and phthalocyanin complexes adsorbed on multiwalled carbon nanotubes359,360 or coated on carbon electrodes361 as well as Co terpyridines grafted to glassy carbon.362 Among the heterogenized Co complexes, Co tetraphenylporphyrins (tpp) play a particularly important role, as shown by a number of recent reports.363−367 Incorporated into a metal−

Figure 30. (Left) CV of 0.3 mM [CoIIIN4H(Br)2]+ in CH3CN/0.1 M NBu4PF6 under Ar (black line), after saturation with CO2 (red line), and after adding 10 M H2O to the same solution; v = 0.1 V s−1, WE = glassy carbon. (Right) Structure of the catalyst precursor [CoIIIN4H(Br)2]+. Adapted with permission from ref 135. Copyright 2014 American Chemical Society.

authors found that in CO2-saturated wet acetonitrile, CO can be generated at −2.0 V vs Ag/AgNO3 with a FE of up to 45% (along with 30% H2) using this catalyst.135 Notably, it was found that submonolayers of Co were deposited on the glassy carbon electrode in the course of the electrolysis and that this electrodeposited Co enhances neither CO2 nor H+ reduction at the given working electrode potential. The mechanism of this process has been thoroughly studied with cyclic voltammetry and with separate preparation and examination of catalytic intermediates. The CV of [CoIIIN4H(Br)2]+ recorded in CH3CN under Ar in the absence of H+ contains three reversible waves centered around −0.4, −0.9, and −1.9 V vs Fecp2, which correspond to the formal CoIII/CoII, CoII/CoI, and CoI/Co0 redox couples (see Figure 30, left). Successive addition of CO2 and H2O reveals that the catalytic cycle is initiated by generation of the Co0 species and that a proton source significantly enhances the catalytic activity (see Figure 30, red and blue lines). Additionally, UV−vis SEC experiments showed a [CoIIN4H]2+-type complex as the major AF


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Scheme 22. Probing the Mechanism of the Electrocatalytic CO2-to-CO Conversion with Co N4-Macrocycles with Chemical and Photochemical Experimentsa

a

Boxes in the upper row highlight isolated intermediates along with their X-ray crystal structures (H atoms except for NH groups and counterions left out for clarity). Reaction conditions: (a) Conversion of [CoIIN4H(Br)]Br with 2 equiv of KC8, THF, rt. (b) NaBPh4, MeCN/H2O (3:1), rt. (c) [H-DMF][OTf], MeCN, rt. (d) Photochemical reduction using NEt3 as a sacrificial electron donor in combination with a photosensitizer (for more details see text). Box in the lower row highlights the two-fold-reduced CO2 adduct, which has been detected in a photochemical experiment using rapid scan FT-IR spectroscopy and predicted by DFT calculations. Adapted with permission from refs 135 and 269. Copyright 2014 and 2015 American Chemical Society.

species indicate that the complex is best described as N4H•− radical anion which is antiferromagnetically coupled to a lowspin CoII ion.135 It was concluded that the stability of the N4H•− ligand and its ability for uptake of a second electron contributes to the preferential reduction of CO2 over H+ even at high H2O concentrations. Considering the reversibility of the redox couple at −1.9 V and the fact that [NiIN4H] is a known and stable species, the complex [Co0N4H] was initially proposed as the active 2e− species, although several attempts for its isolation and spectroscopic detection were unsuccessful.135,269 In a subsequent combined spectroscopic−theoretical study by Peters, Gordon, Frei, and co-workers it was found that CO2 and [CoIN4H(MeCN)]+ form a η1-C adduct where the CO2 ligand is only partially charged and slightly bent (see Scheme 22, bottom middle).269 Interestingly, this [CoIN4H(CO2)]+ complex was also selectively formed upon irradiation of [CoIIN4H(MeCN)]2+ with visible light in the presence of [Ru(bpy)3]2+ photosensitizer and an electron-donor reagent (NEt3). Furthermore, it was shown that in wet acetonitrile, [CoIIN4H(MeCN)]2+ reacts with CO2 and H2O to the bicarbonate adduct [CoIIN4H(HCO3)]+ (Scheme 22, bottom right). Taken together, these results suggest that in contrast to the initial assumption, [CoIN4H] is potentially not part of the electrocatalytic cycle and that the 2e− species could also be directly formed from [CoIN4H(CO2)]+. The 2e− intermediate was ultimately visualized upon photochemical treatment of [CoIIN4H(MeCN)]2+ with NEt3,

species present in solution, indicating such a complex as the probable resting state of the catalyst. In order to obtain an idea about possible catalytic intermediates of the catalytic cycle, several species were prepared separately from [CoIIN4H(Br)]Br (see Scheme 22) and characterized using single-crystal X-ray diffraction.135 Considering that the catalytic cycle is established at the formal CoI/Co0 couple (E0 = −1.9 V), a reduced Co species was prepared by conversion of the Co(II) precursor with 2 equiv of KC8 in THF, whereby only the singly reduced complex [CoIN4] was formed due to H2 generation (see Scheme 22, top). The structures of the CoI/CoII couple under electrochemical conditions were then established by further treatment: First, [CoN4] was dissolved in wet acetonitrile, followed by addition of NaBPh 4 resulting in precipitation of [CoIN4H(MeCN)][BPh4]. In the following step, protonation with [H-DMF][OTf] gave [CoIIN4H(MeCN)]2+ (along with 0.5 equiv of H2). Both resulting CoI and CoII species were studied with CV under Ar, and it was found that the voltammetric features associated with the CoII/CoI and CoI/ Co0 redox couples are consistent with those observed for [CoIIIN4H(Br)2]+ (compare Figure 30). The species resulting from electrochemical reduction under Ar at −0.4 and −0.9 V have thereby been established unequivocally.135 The situation in the presence of CO2 turned out to be more complicated. On the basis of the CV results obtained under Ar, [CoIN4H(MeCN)]+ was initially assumed as precatalyst for CO2 reduction. DFT calculations on the [CoIN4H(MeCN)]+ AG


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CO2, and [Ir(ppy)3] photosensitizer, which is known to have a stronger reducing power than the previously used [Ru(bpy)3]2+ photosensitizer.369 The catalytic cycle is closed under these conditions and CO is produced, analogously to the electrolysis at E = −2.0 V. Using rapid-scan FT-IR spectroscopy, [CoN4H(CO2)] was observed as 2e−-reduced intermediate, which dissociates spontaneously on the time scale of seconds to liberate CO and H2O. This result was also confirmed by an accompanying DFT electronic structure calculation, which predicted that this intermediate is a stable adduct of the CoI tetracycle, with substantial electron density on the carboxylate ligand (see also Scheme 22, bottom left). A detailed catalytic cycle for the electrochemical CO2 reduction process cannot be derived from the present results, since the electrolysis conditions can obviously not exactly be reproduced in the nonelectrochemical experiments, and more kinetic data would be necessary to obtain the consecutive order of the elemental steps. However, the studies presented above constitute an excellent example for how insights obtained from electrochemical and photochemical experiments can be combined in order to achieve a better understanding of possible intermediates involved in a CO2 reduction mechanism. The results clearly demonstrate the active role of the N4H ligand, which in cooperation with the metal center is involved in the activation of the CO2 molecule. A further important lesson with respect to steering the selectivity from H+ toward CO2 reduction is that the second electron injected into stable [CoIN4H(CO2)]+ complex creates the labile [CoIN4H(CO2)] species, which spontaneously dissociates to liberate CO and H2O under regeneration of the catalyst. This suggests that the ability to form an adduct with CO2 prior to the second electron transfer plays a crucial role for avoiding H+ reduction, since [CoIN4H(MeCN)]+ was shown to react spontaneously with H+ under electrochemical and photochemical conditions.135,369 Consequently, an optimal stability range for the CO2 adduct seems to be likely: while too strong interaction may result in an unreactive complex after reduction with the second electron, too little interaction would lead to insufficient lifetime for capturing a second electron and therefore to a preference for H+ reduction. 3.3.4.2. Cyclopentadienyl Co Complexes with Pendant Amines. A new type of CO2 reducing molecular catalyst based on Co and diphosphine ligands PR2NR′2 with two pendant amine moieties was recently reported by Artero et al. (see Figure 31, left).126 The system selectively generates formic acid in DMF/water mixtures with promising TOFmax numbers (up to 1000 s−1) and moderate ηcat values (500−700 mV, estimated using the half-wave potential of the catalytic wave Ecat/2). According to the benchmarking results, [(cp)CoIII(PR2NR′2)I]+ belongs to the most potent homogeneous catalysts for CO2-toformic acid electroconversion. Four different ligands with R1, R2 = Ph, Bn, Cy (A−D) were tested, and while all four catalysts display activity for CO2 reduction, it was found that the most electron-donating phosphine ligand (R2 = Cy) in combination with the most basic amine (R1 = Bn) exhibits the fastest kinetics (see Figure 31, right). In order to investigate the influence of the amine groups in the second coordination sphere, [(cp)Co(dppe)I] + with dppe = 1,3-bis(diphenylphosphino)propane was tested as analogous aminefree complex, whereby no catalytic activity was found. Mechanistic studies were carried out in combination with DFT calculations, and a catalytic cycle was proposed that matches the experimental results and the theoretical predictions

Figure 31. (Left) Structures of the Co complexes A−D with pendant amine groups studied by Artero et al. (Right) TOF-η relationship for a series of catalysts for the CO2-to-formate conversion including A−C, the iron carbonyl clusters discussed in section 3.3.2.2, the [LIr(PCP)H2] system (which will be discussed in section 3.3.5), and [Fe(N5)Cl2]+ (see section 4 and ref 373). Adapted with permission from ref 374. Copyright 2017 American Chemical Society.

(see Scheme 23). According to the proposed mechanism, [(cp)CoIII(PR2NR′2)I]+ (species I, Scheme 23) serves as Scheme 23. Catalytic Cycle Proposed by Artero et al. for the Electrochemical CO2-to-Formate Conversion Using Co Complexes with Pendant Amines126

precatalyst which is electrochemically activated by reductive deiodination. The resulting complex [(cp)CoII(PR2NR′2)]+ (II) is further reduced in two successive electron transfers to give Co(I) intermediate III and Co(II) hydride IV, whereby DFT calculations point toward a concerted PCET mechanism for the latter process (with intramolecular proton transfer from the ammonium group to the metal center). While Co(II) hydrides with similar coordination spheres have not been isolated before, they have been evidenced for a related Co PR2NR′2 complex by DuBois and co-workers in the context of electrocatalytic proton reduction.375 The interaction of IV with CO2 was studied with DFT calculations, and it was found that including water into the system is energetically favorable, owing to intramolecular AH


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hydrogen bonding between the bound CO2 molecule and one pendant amino group (see Scheme 23, species V). For the following step, a rate-determining hydride transfer from Co to CO2 via TS-V/VI was proposed, which is in contrast to other reports on Fe- and Co-based CO2 hydrogenation catalysts where a CO2 insertion into the metal−hydride bond was suggested.376,377 The proposed hydride transfer step rather resembles the mechanism followed by hydride donor catalysts based on the dihydropyridine unit, which is assisted by hydrogen-bonded H2O molecules.378,379 With regard to the conversion of V into VI, the hydride transfer mechanism is supported by measurement of the kinetic isotope effect (with fairly high H/D KIE values > 5, supporting a concomitant Co− H bond cleavage and C−H bond formation). The mechanism also provides a rationale for the higher activity of Co PR2NR′2 complexes containing benzylic amino groups (Figure 31, A and C) in comparison to complexes B and D, which are equipped with aniline units: In organic media, benzylamines units are significantly more basic and can therefore better contribute to the stabilization of transition state TS-V/VI via H bonding. 3.3.5. Rh and Ir Complexes. Compared to the neighboring 4d metals Ru and Pd, only a few cases of successful application of Rh complexes as electrocatalysts for CO2 reduction have been reported.145,147,380,381 In an early study by Meyer et al. it has been demonstrated that [Rh(bpy)2(TFMS)2]+ (TFMS = trifluoromethanesulfonate) represents a stable catalyst which promotes the CO2-to-formate electroreduction under aprotic conditions.145,147 Noteworthy, it has been shown that the NBu4+ cation of the supporting electrolyte acts as a proton source via Hofmann degradation (see Scheme 8). Later, the complex [(η5-C5Me5)Rh(bpy)Cl]+ has been studied by Deronzier et al., whereby activity for formate generation was found in the presence of water.382 In a case recently reported by Turro et al., formamidate-bridged dirhodium(II,II) complexes with chelating diamine ligands were successfully applied under protic conditions, whereby H2 and formate were found as major products.381 In contrast to Rh, considerably more reports on the use of Ir complexes for CO2 electroreduction are available. In a striking example from Brookhart, Meyer, and co-workers, a selective reduction of CO2 to formate was achieved by using the Ir PCP pincer dihydride complex [LIr(PCP)H2] (L = CH3CN, see Scheme 24).383 A CPE carried out at −1.45 V vs NHE in CH3CN/H2O (95:5) yielded formate in 85% FE and 40 turnovers per catalyst unit, whereby H2 (formed via nonspecific background reduction of water) was found as side product (15% FE). The mechanism was investigated with CV, NMR studies, and DFT calculations, and a catalytic cycle was proposed based on these results (see Scheme 24). In the CV of 1 mM [LIr(PCP)H2] recorded in CH3CN/H2O (95:5), the cathodic scan shows no signal within the stability range of the electrolyte (see Figure 32, left). However, under CO 2 atmosphere, a catalytic wave appears around −1.4 V vs NHE, which indicates that the employed [LIr(PCP)H2] complex represents the active species. In a control experiment, the reaction of [LIr(PCP)H2] with CO2 (steps 2 and 3 in the catalytic cycle) was studied with NMR spectroscopy. Whereas in dry CH3CN neither the formation of a CO2 adduct nor the generation of formate was observed, the addition of water leads rapidly to formation of [L 2Ir(PCP)H]+ and HCOO−. Interestingly, an equilibrium was observed in this reaction, with quantitative conversion at ≥4% H2O content, and the authors proposed that this equilibrium is driven by improved

Scheme 24. Catalytic Cycle for the CO2-to-Formate Conversion Catalyzed by [Ir(PCP)H2] Proposed by Brookhart et al.276

Figure 32. (Left) Cyclic voltammetry of 1 mM [LIr(PCP)H2] (L = CH3CN) in CH3CN/H2O (95:5) under Ar (black), under CO2 (red), and in THF/H2O (95:5) under CO2 (blue). (Right) CV of 1 mM [Ir(PCP)H]+ in CH3CN under Ar (black) and under CO2 (red). Supporting electrolyte in all experiments: 0.1 M NBu4PF6. Reprinted with permission from ref 276. Copyright 2012 American Chemical Society.

solvation of the generated formate anion in the presence of water. The κ1-formate complex [H(L)Ir(PCP)(OCHO)] (see Scheme 24, bottom) as transient intermediate of CO2 insertion into the M−H bond was not directly observed but postulated on the basis of the observation that other Ir pincer hydrides undergo formation of such κ1-formate complexes in the presence of CO2.276 To ensure that [L2 Ir(PCP)H]+ is also part of the electrocatalytic cycle, a separately prepared sample of this intermediate was studied with CV (see Figure 32, right). As expected, it exhibits a reduction signal under Ar at −1.4 V (black line), the potential where the catalytic wave occurs in the presence of CO2. The observed reduction signal corresponds to a 2e− process, and DFT calculations confirmed that it is the Ir(III) → Ir(I) (metal-centered) process corresponding to step 3 in Scheme 24. The irreversibility of the signal is explained by coupled hydride formation via protonation of the cathodically generated intermediate. In the reverse scan, the hydride species is reoxidized in two successive steps at 0.25 and 0.75 V (not shown).276 Under CO2 atmosphere (red line), a catalytic wave appears at the potential of the Ir(I)/Ir(III) couple, and the similarity to the voltammetric profile of [LIr(PCP)H2]/CO2 AI


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confirms the role of [L2Ir(PCP)H]+ as intermediate of the catalytic cycle. These interesting mechanistic features triggered several independent DFT studies.274,277,278 In the first report by Sun, Chen et al., the feasibility of the cycle proposed by Brookhart et. al was essentially confirmed.277 In this work, the CO2 insertion step was described by formation of a [H(L)(PCP)Ir− H···η1-C(CO2)] intermediate with a C−H bond length of 1.3 Å, followed by rearrangement to give [H(L)(PCP)Ir−H···η1O(CO2)]. Since the experimentally used solvent mixture (CH3CN/H2O = 95:5) is difficult to account for, separate calculations were carried out with solvent models for CH3CN and H2O. In agreement with the experimentally observed positive influence of water addition, a significant decrease of the activation barrier is resulting from using a H2O- instead of a CH3CN-based solvent model. In a later computational study by Nielsen and co-workers, several possible pathways for the reaction between catalyst and CO2 were investigated, and it was concluded that formation of [H(L)(PCP)Ir−H···η1-O(CO2)] followed by hydride transfer directly resulting in [(PCP)IrH]+ and HCOO− also represents a viable option.274 Furthermore, the authors studied proton reduction as a competitive pathway and found that protonolysis of the Ir−H bond is kinetically unfavorable under the given conditions (with activation barriers in excess of 25 kcal mol−1), which explains the experimentally observed selectivity toward formate generation over H2 evolution. On the basis of their computational results, a general guideline for the design of related formate-generating catalysts was proposed. In agreement with the conclusions by Berben et al. (see [Fe4N(CO)12] clusters in section 3.3.2),336 it was suggested that the hydricity of the M−H intermediate which reacts with CO2 must be moderate in order to prevent protonolysis but sufficiently high to provide driving force for hydride transfer (“formate window”). A further report by Ahlquist et al. confirms the feasibility of the hitherto proposed catalytic cycle but also addresses an interesting alternative pathway where the cathodically generated Ir(I) intermediate reacts directly with CO2 resulting in the [L(PCP)Ir−H···η1-C(CO2)]− intermediate instead of being protonated to give [L(PCP)IrH2].278 According to their calculations and in good agreement with a previous DFT study by Hazari et al.,279 this “Ir(I) hydride mechanism” is energetically more favorable and could therefore potentially compete with the mechanism outlined in Scheme 24. Further insights were obtained by Brookhart et al. upon examination of a water-soluble Ir PCP pincer complex (see Scheme 25) in an aqueous electrolyte containing 1% CH3CN.383 The triflate salts of WS-[L2(PCP)IrH]2+ and WS[(PCP)IrH(OTf)] + (Figure 33) were used as catalyst

Figure 33. Structures of the catalysts WS-[L2(PCP)IrH]2+ and WS[(PCP)IrH(OTf)]+ studied by Brookhart at al.383

precursors and the dihydride generated in situ via cathodic reduction and protonation. The catalysts were found to be stable under aqueous conditions, and formate was selectively formed at useful rates in each case. While in the presence of small quantities of CH3CN the mechanism was essentially found to follow the cycle proposed in Scheme 24, the key role of CH3CN became clear upon testing WS-[(PCP)Ir(OTf)H]+ in pure water. Interestingly, neither under Ar nor under CO2 the complex is reduced in the cathodic scan within the stability range of the electrolyte. In contrast, the addition of small amounts of CH3CN leads to the appearance of a reduction signal under Ar at −1.3 V and gives rise to a catalytic wave under CO2 at the same potential. Another experiment further highlights the importance of the CH3CN additive for displacing the formate ligand (see Scheme 25): By converting WS-[(PCP)Ir(OTf)H]+ with excess NaOOCH in water, the stable WS-[H(PCP)Ir(OCHO)]+ intermediate is formed as κ2-formate adduct (see Scheme 25), which is quantitatively converted to WS-[(PCP)IrH]2+ and formate upon addition of only 2 equiv of CH3CN. Thus, acetonitrile can be considered as a key ancillary ligand, both enabling the cathodic reduction of the Ir(III) species and sustaining the catalytic cycle by displacing formate from the catalyst. 3.3.6. Ni Complexes. The first successful application of Ni catalysts for CO2 electroreduction was reported in 1980 by Eisenberg et al., who studied a selection of complexes with N4macrocyclic ligands.370 Initiated by a study of Sauvage et al. a few years later,384 a number of reports on Ni(II) cyclam and derivatives appeared, and the progress in optimization and mechanistic understanding is still ongoing.127,225,385 While mechanistic aspects of this system are discussed below, a comprehensive overview of the Ni catalysts reported thus far is provided in section 4. Although the vast majority of the studied Ni catalysts are N4 macrocycles, it should be noted that other promising systems based on Pincer-type ligands275 and Ni cluster compounds have been investigated as well.224,386−388 As pointed out above, cyclam-based complexes constitute the most frequently studied group of Ni catalysts for CO2 electroreduction. On the basis of the pioneering work of Eisenberg et al.370 and Sauvage et al.384 in the 1980s, a number of studies on [Ni(cyclam)]2+ (see Scheme 26, top) with focus on mechanistic aspects,127,136,389−395 on optimization of catalyst structure and performance,225,385,396,397 and on improving the environmental compatibility15,398 appeared. The Ni cyclam system represents a remarkable case where the selectivity of CO2 reduction can be directed by choice of

Scheme 25. Studies on Potential Intermediates of the CO2to-Formate Conversion in Water Using Water-Soluble Ir PCP Pincer Complexes383

AJ


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complex adsorbed on the Hg electrode surface, whereby reduced Ni species in the bulk phase play no or only a minor role. In fact, it was found later that adsorption of the Ni(I) complex on the Hg surface proceeds already well below 1.0 V.389 The benign effect of catalyst adsorption is contrasted by a study where Ni cyclam has been used in combination with a glassy carbon electrode in a CH3CN/H2O (4:1) mixture.15 In this case it has been shown that the catalyst still exhibits a good activity with high selectivity for CO formation, although with a slower rate and at the potential of the Ni(II)/Ni(I) couple. A plausible mechanism based on the initial proposal by Sauvage et al.136 and later studies by Kubiak et al.225 and Ye et al.127 is depicted. The cycle is induced by single-electron reduction of [Ni(cyclam)]2+ concomitant with adsorption of the reduced complex on the electrode surface. The adsorbed Ni(I) species then forms an adduct with CO2, this step being supported by hydrogen bonding between Ni-bound CO2 and the adjacent N−H groups (see Scheme 26, right).127 The resulting species is then reduced and protonated, followed by dehydration and decarbonylation to close the catalytic cycle. As initially observed by Sauvage and co-workers with UV−vis and EPR spectroscopy, the active [Ni(cyclam)]+ readily forms the [Ni(cyclam)(CO)]+ under bulk electrolysis conditions, which is not part of the catalytic cycle and decreases the amount of available Ni(I) species for catalysis. Consequently, product inhibition limits the rate of the catalytic process.136 With regard to the shape of the catalytic wave in the CV, this phenomenon is expressed by the strong decrease in the catalytic current in the CV upon scanning beyond the peak potential of the catalytic wave (see Figure 34).393 In addition, it was also found that cathodic reduction of [Ni(cyclam)(CO)] + to its corresponding Ni(0) species represents a possible pathway to irreversible catalyst deactivation.225 To circumvent these issues, Kubiak et al. have shown an elegant possibility by using [Ni(TMC)]2+ (with TMC as the tetra-N-methylated version of the cyclam ligand) as a CO scavenger.225 The one-electronreduced [Ni(TMC)]+ exhibits a much higher binding constant for CO and a much lower activity for CO2 reduction compared to the parent cyclam complex and is therefore an ideal candidate for this purpose. Consequently, the catalytic current could be increased 10-fold in the presence of an excess amount of [Ni(TMC)]2+ when applied in a CH3CN/H2O (4:1) solvent mixture using a glassy carbon electrode. It has been known for a long time that among the five possible planar conformational isomers, [Ni(cyclam)]2+ predominantly exists in the trans-III and trans-I configuration (85% and 15%, respectively) in aqueous solution (see Scheme 26, top).400 Consequently, the influence of the configuration of the one-electron-reduced Ni(I) species has been investigated in several experimental and theoretical studies, whereby the outcome strongly depends on the choice of the electrode material. For the scenario where a carbon electrode is used and the catalytic intermediates are not adsorbed on the electrode, DFT studies indicate that the trans-I species represents the better catalyst (the binding between CO2 and trans-I cyclam was predicted to be more favorable compared to the trans-III species by several kcal mol−1).15,395 This finding could potentially be ascribed to the four N−H functionalities in trans-I cyclam which are available for hydrogen bonding with CO2, compared to only two available N−H groups in the transIII configuration.127 These results are contrasted by experimental and theoretical studies of the scenario where the catalytic intermediates are adsorbed on a Hg electrode. While

Scheme 26. Proposed Mechanism for the CO2 Electroreduction Catalyzed by [Ni(cyclam)]127,136

the electrolyte system: While CPE in water (pH 4 to 5) gives CO in high TON and FE, it leads preferentially to formic acid in aprotic solvents, with FEs of up to 75% and CO as byproduct.399 The mechanism of the Ni cyclam-mediated process has first been studied by Sauvage et al. using a Hg working electrode.136 The corresponding CVs recorded in water in the absence and presence of CO2 are depicted in Figure 34. Under N2

Figure 34. Cyclic voltammetry of 1 mM [Ni(cyclam)]Cl2 in the absence (curve a) and presence of CO2 (curve b). Scan rate: 0.1 V s−1, Electrolyte: 0.1 M KClO4 in H2O. WE: Hg. Reprinted with permission from ref 136. Copyright 1986 American Chemical Society.

atmosphere, a reversible wave centered around −1.3 V and assigned to the Ni(II)/Ni(I) couple is observed (curve a), which gives way to a quite impressive catalytic wave upon saturation of the electrolyte with CO2 (curve b). Several features of the voltammetric behavior under CO2 differ from other systems such as the Fe porphyrins discussed in section 3.3.2.1 (compare Figure 25) or the [CoN4H(Br)2]+ complex discussed in section 3.3.4.1 (compare Figure 30). The peak potential of the catalytic wave is shifted anodically by about 300 mV with regard to the Ni(II)/Ni(I) couple, and the peak current shows only a weak dependency on the catalyst concentration. This clearly points toward catalysis by a Ni AK


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early CV studies showed that a cathodically induced configurational change of the Ni(I) complex adsorbed on Hg occurs around E = −1.0 V vs NHE,389 recent DFT calculations carried out by Batista et al. indicate that at the potential where catalysis occurs, Hg promotes adsorption of the flat trans-III over the bowl-shaped trans-I conformer.401 Furthermore, it was found that the conformational change to trans-III accounts for the enhancement of the catalysis of the CO2-to-CO conversion on the Hg electrode: Since Ni−CO σ-bonding is weakened in the trans-III intermediate adsorbed on the surface, CO desorption kinetics are improved and a catalyst poisoning is disfavored.400,15,12,71,27 As mentioned above, the Ni cyclam complex produces formate instead of CO upon switching to aprotic electrolytes,399 a striking feature which has been studied by Ye et al. using DFT calculations (see Scheme 27).127 The selectivity for

Figure 35. Ni(cyclam) derivatives with improved catalytic performance.402

lower overpotentials compared to the parent cyclam complex, with an exclusive selectivity for CO formation. It has been proposed that the favorable geometry of these complexes is responsible for the higher catalytic activity. Interestingly, the related RSSR-[Ni(HTIM)]2+ shows a much lower activity compared to RRSS-[Ni(HTIM)]2+ and [Ni(cyclam)]2+, further highlighting the importance of stereochemistry for the catalytic performance. Further insights into the relationship between structure and reactivity of cyclam complexes were later obtained in a follow-up study.395 An additional interesting finding of these studies is that at pH < 2, H2 is predominantly produced, allowing for generation of syngas upon adjustment of the pH to the appropriate value.395 3.3.7. Pd-Based Catalysts. While recently very promising results for the electrocatalytic CO2 reduction were achieved with new bis-N-heterocyclic carbene pincer-type complexes,403−405 the focus of research was mostly on Pd phosphine complexes in the past. Among the Pd phosphine complexes, those having tridentate phosphine ligands (P3) play a prominent role (vide infra), whereas monodentate phosphine ligands have been less frequently studied in the context of CO2 electroreduction.406,407 A series of studies on the electrocatalytic activity of such Pd(II)−P3 catalysts was carried out by DuBois and co-workers (see Scheme 28).408−411 Cyclic voltammetry and CPE revealed that these compounds are catalytically active for the electroreduction of CO2 to CO and H2O in DMF or CH3CN containing small amounts of a strong Brønsted acid such as HBF4 or H3PO4. A water-soluble system

Scheme 27. Summary of the Results of a DFT Study by Ye et al. on the CO2 Reduction Catalyzed by Ni(cyclam)a

a

ΔG values in kcal mol−1, referring to H3O+ as proton source.127

CO production in aqueous medium was explained by formation of a η1-CO2 adduct with Ni(I), which is energetically favored by 14 kcal mol−1 compared to the η1-OCO version. The theoretical study also confirmed the earlier experimental finding that the Ni(I) species interacts with CO2 rather than the Ni(II) complex, since all attempts to locate a local minimum for an adduct between CO2 and [Ni(cyclam)]2+ resulted in CO2 dissociation. After reversible binding between Ni(I) and CO2, the resulting intermediate was predicted to be reduced in a concerted PCET, followed by an exergonic dehydration step. Since the C−O bond breaking was found to proceed barrier free upon assumption of H3O+ as proton donor, the conversion of the η1-CO2 adduct into the [LNi(CO)]2+ intermediate can be summarized as proton-coupled electron transfer/bond cleavage (PCETBC, comparable to the mechanism associated with the Fe porphyrins described in section 3.3.1). However, in the absence of suitable proton donors, the second electron transfer becomes more difficult, which explains why formate is preferentially formed under aprotic conditions via the η1-OCO intermediate. It should be noted that despite the common use of Ni(cyclam) as benchmark for other Ni N4-macrocycles, there have been other complexes with a better catalytic performance identified by Fujita et al. (see Figure 35).395,402 In aqueous electrolyte at the Hg electrode, RRSS-[Ni(HTIM)]2+ and [Ni(DMC)]2+ exhibit significantly higher catalytic currents at

Scheme 28. Catalytic Cycle Proposed by DuBois et al. for the CO2-to-CO Electroreduction Using Pd Complexes with Tridentate Phosphine Ligands14

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was also developed, which showed activity for CO formation when a phosphate buffer was used.412 Cyclic voltammetry under Ar reveals two separate electron transfers in the cathodic scan.409 This finding indicates the presence of a Pd(I) species which is stable enough to be trapped in the presence of CO2. Kinetic studies were carried out, and several general trends for this class of catalysts were revealed. While at both low and high [AH] the rate is first order with regard to both CO2 and catalyst, it is second order with regard to acid at low [AH] and reaches saturation kinetics at higher [AH].409,411,413 A catalytic cycle was proposed based on the CV results (see Scheme 28), where η1-C addition of the metal center to CO2 represents the rate-determining step at high acid concentrations. In view of the thermodynamically favored proton reduction, this addition step is crucial for the selectivity. Therefore, it is of utmost importance that the Pd(I) intermediate reacts preferentially with CO2, since Pd−hydride formation and subsequent protonolysis represent a competitive pathway. In fact, selectivities of up to 97% for CO generation were observed with the best catalysts of this series, even when strong proton donors such as HBF4 were used. As next steps of the catalytic cycle, further reduction and protonation of the [P3Pd(COO−)(CH3CN)]+ intermediate were proposed, eventually leading to the loss of the weakly bound CH3CN ligand. At this point, the important role of this weakly coordinating solvent molecule occupying the fourth coordination site becomes obvious: It keeps an additional vacant coordination site available and allows for the migration of H2O to the metal center, leading to the [P3Pd(CO)(H2O)]2+ intermediate. It was suggested that H2O migration to the vacant coordination site facilitates C−O bond cleavage (which is actually the ratedetermining step at low [AH]) and thereby improves the catalytic rate. Not surprisingly, it was found that the incorporation of monodentate phosphine ligands and the use of strongly coordinating solvents such as DMSO inhibit the reaction.413 A typical degradation pathway of [P3PdL]2+-type catalysts under the conditions of the electrocatalytic CO2 reduction is the formation of bimetallic complexes (see Scheme 29), which

carried out.14,410,413 The highest activity is obtained when P3 ligands with ethylene tethers forming two five-membered rings upon coordination to the Pd center are used. In contrast, catalysts containing one five-membered and one six-membered ring or two six-membered rings are considerably less active.414 Interestingly, the selectivity is exclusively directed toward proton reduction when P3 is replaced by a PNP or PCP pincer ligand.415 Consequently, the optimized f irst coordination sphere for optimal CO formation rates and selectivities consists of a P3 ligand with ethylene tethers coordinating to the Pd atom in combination with a weakly bound solvent molecule such as CH3CN or DMF. Within these optimized structural parameters, the rate can be further tuned by suitable modifications of the second coordination sphere. As pointed out above, the ratedetermining step under regular electrolysis conditions (1 atm CO2 and excess acid) is represented by the addition of the Pd(I) intermediate to CO2 (compare Scheme 28). When the substituent R′ on the central P atom of the P3 ligand is rather small, log(k) was found to be proportional to the potential of the Pd(II)/Pd(I) couple (see Figure 36, bottom, black diamonds).

Scheme 29. Degradation Pathway of [P3PdL]2+-Type Catalysts14

Figure 36. (Top) Steric interactions between R′ and coordinated CO2. (Bottom) Plot of the log of the second-order rate constants k for the reaction of [P3PdL]+ with CO2 versus the potentials for Pd(II) → Pd(I). Black diamonds: catalysts with small substituents R′ (Me or Ph) on the central P atom of the tridentate ligand. Red squares: catalysts with large substituents R′ on the central P atom. Green triangles: catalysts substituted on the central P atom with groups that may be able to stabilize the Pd(I)−CO2 adduct. Adapted with permission from ref 57. Copyright 2014 American Chemical Society.

results in low double-digit TONs.409 However, a significant improvement of the catalyst efficiency was achieved by incorporation of sterically demanding phosphine substituents. For instance, replacing the PhP(CH2CH2PEt2)2 ligand by the more bulky PhP(CH2CH2PCy2)2 leads to an increase in TON from 10 to 130, while the FE for CO formation is simultaneously improved from 65% to 85%.409 In a further attempt to suppress the dimerization pathway, the alkylene chain length between the P units in P3 was increased from C2 to C3.414 Although dimerization was prevented with this approach, the catalytic activity could not be improved, since Pd−H formation emerged as a competing pathway. With regard to the homogeneous reaction rate, a thorough investigation of the structure−reactivity relationship has been

However, the observed catalytic rate decreases by approximately 50% when a sterically demanding substituent such as tert-butyl or mesityl is attached to the central P atom (see red squares in Figure 36, bottom). This behavior was explained by hindrance of the coordination of CO2 by the substituent R′ (see Figure 36, top), and consequently, the second coordination sphere is inhibiting catalysis. In order to accelerate the catalytic process, P3 ligands with a modified second coordination sphere were developed with the intention to facilitate the binding to CO2 via metal−ligand cooperativity (see green triangles in Figure 36 and structures in Figure 37). In one example, the incorporation of a phosphonium moiety via an ethylene linker attached to the central P was tested, resulting in a 50% increase of the reaction rate (see Figure 37, top AM


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naphthaldimine.423 Promising results have also been obtained with dinuclear complexes bridged by phosphanylbipyridine ligands, whereby a selective formation of CO and CO32− has been observed in aprotic electrolytes.347,424 An outstanding case was reported by Bouwman et al. with the dinuclear complex [Cu2(L-L)]2+ (see Scheme 30), which was found to catalyze the CO2-to-oxalate conversion in almost quantitative FE. Remarkably, the catalytic cycle is established at +0.03 V vs NHE (almost 2 V more positive than E0(CO2/ CO2•−). In CH3CN under constant purging with CO2, 6 turnovers were achieved upon 7 h electrolysis. Interestingly, [Cu2(L-L)]2+ is capable of both capturing and reducing CO2 from air to form oxalate. A catalytic cycle was proposed on the basis of the analysis of catalytic intermediates and some CV data (see Scheme 30). Initially, attempts to prepare [Cu2(L-L)]2+ via conversion of [Cu(CH3CN)4]BF4 with the disulfide ligand (L-L) rendered the tetranuclear oxalate-bridged complex [Cu 2 (LL)2(C2O4)2]4+ upon exposure to air, whereby the incorporation of carbon dioxide was proven by using 13CO2 in a labeling experiment. Further evidence for the active species was obtained by cathodic reduction of separately prepared ([Cu2(L-L)Cl]2)4+ at 0.03 V vs NHE under Ar in a CH3CN/ NBu4PF6 electrolyte. After passing four charge equivalents, the identity of [Cu 2(L-L)] 2+ was confirmed with ESI-MS spectrometry and UV−vis spectroscopy. Furthermore, it was shown that this intermediate reacts spontaneously upon purging the solution with CO2 to give [Cu2(L-L)2(C2O4)2]4+, which was confirmed by comparison of the spectroscopic and voltammetric data to an authentic sample. Addition of LiClO4 to the same solution leads to precipitation of the Li oxalate and liberation of [Cu2(L-L)(CH3CN)2]2+, the latter one being confirmed by ESI-MS spectrometry. Consequently, the tetranuclear oxalate-bridged Cu(II) complex seems to be thermodynamically favored over the dinuclear Cu(II) complex, and the catalytic cycle can only be sustained by constant removal of oxalate through precipitation of the Li salt. A followup work where a chemically reduced Cu macrocycle dimer was used for CO2-to-oxalate conversion supports the idea that two Cu centers can perform C−C coupling between two CO2 molecules.425

Figure 37. Modification of the second coordination sphere in order to improve the catalytic rates of P3−Pd complexes.14

left).14,411 In contrast, the use of a 3-hydroxypropyl substituent showed no significant effect (see Figure 37, top right).14,412 This result is in sharp contrast to other catalysts such as the Fe porphyrin or Mn bipyridyl system, where the installment of phenolic groups in the vicinity of the metal center led to a dramatic increase of the catalytic performance (see Figures 20 and 23). The best results were achieved upon tethering a second [Pd(P3)(L)]2+ unit to the central P atom via a methylene linker (see Figure 37, bottom).14,416 This bimetallic complex showed an enhancement of 3 orders of magnitude for the second-order catalytic rate with respect to the expected value from the Pd(II)/Pd(I) potential, which was interpreted as a sign for metal−metal cooperativity. Other attempts to initiate metal−metal cooperativity with di-, tetra-, and pentanuclear Pd P3 complexes were not effective,416−418 and in this regard, the dinuclear methylene-bridged congener shown in Figure 37 remains the only successful example within this series of catalysts.416 3.3.8. Cu-Based Catalysts. While Cu (as bulk metal, in nanostructures or alloys) is a well-studied heterogeneous electrocatalyst for CO2 electroreduction,50,79−82,419 only a handful of successful applications of Cu catalysts in homogeneously catalyzed reactions have been reported. In several studies of mononuclear and dinuclear Cu complexes, promising catalytic currents were observed in the presence of CO2, although the products have not been identified. These cases comprise Cu(II) complexes with ligands based on cyclam,420 salicylaldimine,421 carbonate,422 and 2-hydroxy-1-

Scheme 30. Catalytic Cycle for the CO2-to-Oxalate Electroconversion Proposed by Bouwman et al.142

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3.3.9. Group 6 Metals. While historically, 4d and 5d transition metals from groups 7−10 such as Re, Ru, and Pd played an important role in CO2 electroreduction, the demand for more sustainable solutions based on less expensive elements has recently directed the attention of the community toward catalysts based on earth-abundant 3d metals such as Fe and Mn (see sections 3.3.1 and 3.3.2). In the ongoing quest for finding new and effective catalysts based on abundant elements, it has only recently been discovered that group 6 metals can also be applied for electrocatalytic CO2 reduction, of course provided that the appropriate ligands are used. In 2014, Kubiak et al. reported that [M(bpyR)(CO)4] (M = Mo, W, bpy-R = 4,4′ disubstituted 2,2′-bipyridine) constitute competent catalysts for CO2-to-CO electroreduction in acetonitrile in the presence of TFE as proton donor.426 While a full catalytic cycle has not yet been proposed, it was found that compared to the group 7 congeners (Mn and Re), the catalytic rates are lower (possibly due to a stronger back-donation to CO and slower product release). It was later reported by Hartl and co-workers that CO dissociation from reduced [M(bpyR)(CO)4] (M = Cr, Mo, W) can be accelerated by using the solvent NMP and that catalysis can be further enhanced by using a gold electrode.214 On the basis of in situ IR−vis vibrational sum frequency generation (VSFG) spectroscopy studies carried out by Cowan and coworkers, this benign effect was explained by an interaction between the Au electrode surface and a stable [M(bpyR)(CO)4]•− intermediate, which is critical for enabling a decarbonylation.427 In a recent study by Nervi et al., a series of Mo and W tetracarbonyl complexes coordinated by nonclassic diimines has been examined in aprotic media, whereby the 2,2′-dipyridylamine (dpa) complex clearly outperformed analogous catalysts with substituted 2,2′-bipyridine ligands.428 Furthermore, Grice et al. recently found evidence that in contrast to group 7 carbonyls M2(CO)10, their group 6 congeners M(CO)6 catalyze CO2-to-CO electroconversion in the presence and absence of proton sources.429 Although the mechanistic details of the presented cases are not yet entirely understood, it seems clear that there is a lot of potential for further developments in the field of group 6 catalysts. 3.3.10. Organocatalysts. While among the reported cases metal complexes represent the vast majority of catalysts for CO2 electroreduction, only a few successful applications of organic mediators are known. One explanation for the lack of suitable organocatalysts is that upon cathodic reduction most organic compounds form highly reactive intermediates which are difficult to control. While the reduction of most organic compounds induces irreversible follow-up chemistry, only a small number of systems can form anion radicals or anions which are stable enough for establishing a catalytic cycle. A further problem is that the few systems which can be reversibly reduced (e.g., unsaturated hydrocarbons with extended πsystems) tend to undergo carboxylation in the presence of CO2. While these reactions are interesting from the point of view of organic chemists, they render a catalytic conversion of CO2 impossible. However, some very interesting exceptions from this general trend have been reported over the years (vide infra), and the development of further metal-free systems for CO2 reduction may represent one of the future challenges in the field. 3.3.10.1. Benzoic Acid Esters and Benzonitriles. In the early 1980s, Vianello et al. reported on the electrocatalytic activity of benzonitrile in the reduction of CO2 in dry DMF, whereby oxalate is exclusively formed.88 It was shown that the reaction is

initiated by generation of the radical anion of benzonitrile, followed by electron transfer from the radical anion to CO2 and dimerization giving oxalate. At the time this finding was rather surprising since radical anions of aromatic compounds are typically carboxylated in the presence of CO2 via electrophilic aromatic substitution (see also discussion in section 2.3.2).22,23 Originally, an outer-sphere single electron transfer between the electrochemically generated radical anion of the catalyst (Cat•−) and CO2 was assumed.88 However, a later study carried out by Savéant, Vianello, and co-workers with several benzonitriles and alkyl benzoates revealed that the reaction between Cat•− and CO2 is rather of an inner-sphere nature.89 A selection of 14 catalysts of this series has been employed, and the rate constants for catalysis have been determined as a function of the catalyst standard potential. On the basis of the analysis of the resulting activation/driving force relationship, an inner-sphere mechanism was proposed, which is exemplified for the benzoate case in Scheme 31 (the sequence for the Scheme 31. Mechanism for the Electrochemical Generation of Oxalate Catalyzed by Alkyl Benzoates Proposed by Savéant et al.89

benzonitriles proceeds analogously). The reaction between Cat•− and CO2 consists of two successive steps. The first one is an addition of the oxygen or nitrogen of Cat•− to CO2 under intermediate adduct formation. This step can also be considered as an inner-sphere process, where electron transfer and bond formation are concerted. The second step is a homolytic cleavage of the C−O bond, whereby CO2•− is ultimately formed. 3.3.10.2. Pyridine. One of the most intriguing but also most controversially discussed systems for the electrochemical reduction of CO2 was introduced by Bocarsly and co-workers in the early 1990s.430 It was reported that pyridine (Pyr) as simple organic compound can facilitate the electrochemical reduction of CO2 to MeOH in aqueous electrolytes at acidic pH using a Pd electrode at −0.58 V vs SCE, which corresponds to an overpotential of only 200 mV with regard to E0(CO2/ MeOH) (compare Table 1). Up to 30% Faradaic yield for MeOH generation was reported with hydrogen evolution as the main competing reaction. Formic acid and formaldehyde were identified as intermediate products along the six-electron pathway to MeOH, and Pyr was found to play a role in the reduction of both intermediate products.430,431 These findings triggered a series of studies by different groups, whereby it became clear early that the process is highly sensitive toward the type of electrode material. While MeOH generation was also observed using other electrodes such as Pt,431 Sn,100 p-GaP (using a photoelectrochemical setup),432,433 and Pt/C-TiO2,434 no activity for CO2 reduction was found on carbon and Au.431,435 A later study based on CV and CPE questions the AO


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CO2 pressure. Second, monitoring the product distribution in the course of electrolysis revealed that the MeOH level stops to increase after a short induction period, whereas hydrogen is continuously produced. Importantly, the MeOH concentration did not increase after reactivation of the electrode, ruling out electrode passivation as an explanation and indicating that the process has a short-lived transient character. Two other recent studies by Xu et al.447 and Savéant et al.438 indicate that the formation of MeOH could not be reproduced when polycrystalline platinum electrodes were used. Since the major electrolysis product was hydrogen, it was reasoned that CO is formed in small quantities, which inhibits further CO2 reduction by electrode poisoning (but not the concurrent proton reduction).438 These results are contradicted by a study from Portenkirchner et al. comparing pyridine and pyridazine as catalysts, whereby in each case MeOH was detected as part of a product mixture.448 All of these recent reports further highlight that the process is still not sufficiently understood and that additional efforts are needed in order to obtain reasonable explanations for the observed effects. The long story of Pyr-assisted CO2 reduction with all its contradicting reports is instructive in several regards. First, it is a good example for a process which was for a long time thought to be a homogeneously catalyzed one but later turned out to be heterogeneous. We believe that this possibility should be taken into consideration for all catalysts! Second, the process represents a good example for a very delicate system with a strictly limited range of favorable electrolysis conditions. It appears that not only the electrode material but also the surface structure and the composition of different species competitively adsorbed on the electrode play a major role. Third, the case illustrates how such favorable conditions can be time dependent: While a short CV experiment may render promising results, the situation can dramatically change in the course of a prolonged bulk electrolysis.165,442,443,449,445 3.3.10.3. Imidazolium-Based Ionic Liquids. Since the discovery that imidazolium-based ionic liquids can promote the CO2-to-CO electroconversion in IL−water mixtures449 there has been considerable interest in the application of the imidazolium cation in the homogeneous and heterogeneous electrocatalytic CO2 reduction.284,450−452 It has been shown that this reaction is surface sensitive and that Ag, Sn, and Bi electrodes are particularly efficient.449,451,453,454 The fact that the efficiency of the IL-assisted CO2 reduction strongly depends on the electrode material suggests that the positive influence cannot simply be attributed to homogeneous electrocatalysis and that an intimate interaction of CO2, imidazolium, or both with the electrode surface (as in the pyridinium-promoted case, vide supra) plays a crucial role. While it has been shown that ηcat of CO2 reduction is lowered by several hundred millivolts in the presence of imidazolium cations,449,450 the exact mechanism is still under debate. On the basis of DFT calculations, Nakamura and coworkers proposed a mechanism where cathodic reduction of the investigated 1-ethyl-3-methylimidazolium ion ([EMIM]+) to EMIM• is followed by formation of a ([EMIM-COOH]•) key intermediate, whereby the hydrogen in position C-2 represents the proton source for formation of the formyl group.284 The finding that the EMIM radical is more likely the CO2-activating species than the EMIM carbene was later supported by Horne, Zhang, and co-workers, who identified the [EMIM-COO]− adduct resulting from interaction between the EMIM carbene and CO2 as deactivation product.452 However,

activity of Pyr on Pt electrodes, since earlier reports could not be reproduced.436 The Pyr-based system has created substantial controversy in the community with regard to the role of Pyr in the mechanism, and efforts in order to obtain a better understanding seem to be ongoing.166,437,438 The initially suggested mechanism involved the reduction of a pyridinium ion at the electrode to give the pyridinium radical PyrH•.431,439 It was proposed that PyrH• then reacts with CO2 to yield carbamate radical intermediate PyrCOOH•, followed by further singleelectron reductions and protonation steps, regenerating Pyr and releasing MeOH and water. However, subsequent quantum chemical studies carried out by Carter et al. indicated that the formation of PyrH• at such low potentials is unlikely, since the calculated potential for Pyr reduction was more than 1 V more negative than the −0.58 V vs SCE from the experimental study.280−282 As an explanation, two different scenarios were proposed based on computational studies: In the first scenario by Carter et al., dihydropyridine (DHP) as a result of a twoelectron two-proton reduction was considered as the active species.282 In the second approach by Musgrave et al., a preadsorption of PyrH+ on the electrode surface and a πorbital-based reduction was used to explain the discrepancy between experimental and calculated potentials, whereby the reaction was assumed to proceed via a PyrH• intermediate as initially proposed.440 While later the formation of DHP could not be experimentally confirmed,441 the involvement of PyrH• was disproved by further studies (refs 282, 440, 441, 442, 443, and 445). Concurrently, a new mechanism was suggested by Batista et al. based on quantum chemical free energy analysis of the process, explicitly taking into account the effect of the electrode surface using periodic boundary conditions (see Scheme 32, left).266 According to their proposal, hydride Scheme 32. Mechanism for Pyridine-Catalyzed CO2 Electroreduction Involving the Metal Electrode Surface Proposed by Batista et al.266,446 and Bocarsly et al.165

species on the electrode surface play a crucial part in the CO2 electroreduction process (see Scheme 32, right). Experimental support for the involvement of metal hydrides was later reported based on a number of different techniques, including isotopic probing (deuterium exchange).442,444 The key role of metal hydrides in the process was also confirmed by Bélanger et al. in a comparative study of the Pyr-based system with Ir, Au, and glassy carbon electrodes.445 In a more recent work by Rybchenko, Haywood, and coworkers, the electroreduction process was studied at elevated CO2 pressures (up to 55 bar).166 While the general feasibility of CO2 reduction to methanol and the necessity of Pyr therein was confirmed, severe limitations of the process became obvious. First, increasing the CO2 pressure (for a fixed Pyr concentration) leads to an increase of the cathodic peak current in the CV but not to an increase of the Faradaic yield for MeOH in bulk electrolysis, clearly showing that hydrogen production remains the dominant reaction even at elevated AP


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characterization including the extraction of kinetic data seems to have become the prevalent standard for reports on new molecular electrocatalysts. Although the relationship between catalyst structure and efficiency basically remains empirical, it has been well investigated for a number of individual systems. The importance of the second coordination sphere has been realized and pushed ligand design in new directions. In this context, developments such as metal−ligand cooperativity, ligand-based proton relays, and functional ionic groups that allow for catalysis-boosting Coulombic interactions represent milestones that helped to set new benchmarks and will trigger further developments. With regard to the metal center, the focus has shifted from precious elements of the second and third d-block series such as Ru, Pd, and Re to more abundant 3d metals such as Mn, Fe, Co, and Ni. It has been shown that these 3d metals can be more than competitive when the appropriate ligands are used. While a broad range of catalyst systems for CO and formate generation have become available, the efficient generation of the four, six, and eight electron reduction products formaldehyde, methanol, and methane with homogeneous electrocatalysts still remains as a major challenge. Only a few molecular electrocatalysts have been reported to be able to reduce CO2 beyond the two-electron reduction products (i.e., immobilized Co porphyrins363 and the controversially debated pyridine-related systems discussed in section 3.3.10.2). Recent results on the photochemical CO2-to-CH4 reduction using iron porphyrins514,515 indicate that this pathway may in principle be feasible with homogeneous electrocatalysts. Catalyst durability remains an issue that deserves much more attention in the future. In few cases, data on the long-term stability (more than a few hours) has been reported. While for each of the more common products of CO2 electroreduction, CO, formate, and oxalate, a number of optimized catalysts with exclusive selectivity, moderate overpotential, and excellent activity has been made available, the stability does often not exceed a few hours/days or has not been determined under realistic conditions (= high current densities). Moreover, the decomposition pathways including the parameters that determine the lifetime are poorly understood, as is the case for many other nonelectrochemical catalysts. Consequently, the development of more stable systems including a more systematic investigation of the factors that improve the catalyst lifetime represent a future challenge in this field. An interesting approach to overcome stability issues is represented by heterogenization of molecular catalysts through attachment to an appropriate cathode material. Such immobilization strategies lead to heterogeneous electrocatalysts with well-defined active sites, which can outperform their homogeneous counterparts in terms of durability and other heterogeneous catalysts with regard to selectivity. Promising approaches toward heterogenized molecular electrocatalysts have been reported, and this subfield will very likely continue to obtain increased attention. Future interest and plenty of room for exploration may lie in the combination of micro- and nanostructured electrodes with molecular electrocatalysis.516 As important as the development of new catalysts and the mechanistic investigations are, upscaling and the improvement of cost effectiveness will play an increasingly important role. In this regard, first implementations of molecular catalysts into medium-scale electrolyzers for CO2 conversion have been reported.99,100 While most of the reports discussed in this

the proposed mechanism does still not account for the surface sensitivity of the reaction. Furthermore, it was reported that the observed effect is not only restricted to imidazolium-based ILs. Other organic cations such as PR4+, NR4+, and SR3+ (R = alkyl) may exhibit similar performance-enhancing effects (depending on the experimental conditions), whereby in contrast to imidazolium, electrochemical double-layer effects are believed to be responsible for the promotion of catalysis.452,455

4. TABULAR OVERVIEW To provide the reader a comprehensive overview, we summarized the catalysts which were found to be active for CO2 electroreduction in Table 5. Along with the molecular structure, we included the reaction conditions and the formed products. We deliberately decided not to include ηcat values, Faradaic efficiencies, nor TOF/TON numbers. These numbers were determined under a variety of different experimental conditions using varying mathematic approaches, and showing benchmark values would therefore suggest a comparability which does not exist. Readers interested in greater detail are, of course, referred to the original literature. We note that only examples are shown, where catalytic activity has been demonstrated by CV and where a product analysis in bulk electrolyses shows a significant amount of CO2 reduction products. Not included are examples where merely CO2 activation and no catalytic turnover is achieved. 5. CONCLUSION AND OUTLOOK The existing challenges with regard to the environment and energy will very likely cause a continuing interest in catalyst development for CO2 electroreduction. Although many aspects are still unclear, recent progress in understanding the mechanisms of the conversion to CO, formate, and oxalate is truly impressive. While earlier mechanistic studies almost exclusively relied on electroanalytical techniques, additional methods such as spectroelectrochemistry and quantum chemical calculations are used with increasing frequency. This trend has led to a better mechanistic understanding and enabled a number of improvements and new developments. The pursuit for deeper mechanistic insights will therefore remain a cornerstone for the optimization of molecular catalysts, whereby the comprehensive approach that includes complementary techniques has proven to be much more effective for catalyst development than random trial and error. With regard to operando spectroscopy, FTIR-SEC and UV− vis-SEC have been most widely used to study CO2-electroreduction mechanisms. The techniques for coupling of electrochemistry to Raman,510,511 EPR,182 and X-ray spectroscopy512,513 are in principle available and provide further potential for future studies in the field of CO2 electroreduction. In parallel to the increasing number of available nonelectrochemical techniques for mechanistic investigations, progress has also been made in developing methods for the electroanalytical characterization and benchmarking of the catalysts. Theoretical tools for catalyst benchmarking such as the foot-of-the-wave analysis are now available that provide the possibility to correctly treat increasingly complex systems even in the presence of side phenomena. Although a direct comparison between different studies often remains difficult, these new tools definitively helped to improve the comparability of different systems. In fact, a thorough voltammetric AQ


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Table 5. Homogeneous Catalysts for the Electrochemical Reduction of CO2457−509

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a

Only such cases are shown in which significant catalytic currents in CV experiments are reported and where electroreduction products have been identified. bMinor products (typically less than 10% FE) are shown in parentheses. cIn ref 456 a supramolecular assembly of one Mn and one Re catalyst unit is used to study heterobimetallic cocatalysis.

review are half-cell studies (using a sacrificial anodic reaction), future efforts also need to be directed toward the coupling to

useful anodic processes, ideally leading to products with similar demands. Among the conceivable paired electrolyses, a BD


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GLOSSARY difference between the applied electrode potential and the thermodynamic potential of a reaction of interest (overpotential) η ηcat overpotential required for catalysis BDD boron-doped diamond CCSD(T) coupled cluster with singles, doubles, and perturbative triples CE counter electrode CPE controlled potential electrolysis CV cyclic voltammetry DFT density functional theory E electrode potential E0 equilibrium potential E00 standard equilibrium potential E1/2 half-wave potential EA electron affinity Ecat/2 half-wave potential of the catalytic wave EP peak potentials ET electron transfer EPR electron paramagnetic resonance FE Faradaic efficiency FOWA foot-of-the-wave analysis GC glassy carbon GGA generalized gradient approximation GTO Gaussian-type orbital i electric current iP peak current IP ionization potential IR infrared j current density jP peak current density kapp apparent rate constant (synonymous with TOFmax) KIE kinetic isotope effect LDA local density approximation NHE normal hydrogen electrode NMR nuclear magnetic resonance PB Poisson−Boltzmann PCET proton-coupled electron transfer QM/MM quantum mechanics/molecular mechanics RE reference electrode RVC reticulated vitreous carbon TFA 2,2,2-trifluoroacetic acid TFE 2,2,2-trifluoroethanol SCE saturated calomel electrode SEC spectroelectrochemistry TOF turnover frequency TOFmax maximum turnover frequency (synonymous apparent rate constant kapp) TOF0 turnover frequency at zero overpotential (η = 0) TON turnover number UV−vis ultraviolet−visible v scan rate WE working electrode

particularly interesting one would be a convergent process, where an intermediate derived from CO2 reduction and another one obtained by the coupled anodic oxidation react with each other to form a more complex structure.

AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected]. ORCID

Robert Francke: 0000-0002-4998-1829 Notes

The authors declare no competing financial interest. Biographies Robert Francke studied chemistry at Bonn University (Germany) and Alicante University (Spain). In 2008, he received his diploma degree (equivalent to a M.S. degree) from Bonn University while studying under the direction of Prof. S. R. Waldvogel. Then he relocated with the group of Prof. Waldvogel to Mainz University (Germany), where he completed his dissertation on fluorinated electrolytes for electrochemical energy storage devices in 2012. Equipped with a Feodor Lynen Fellowship (Alexander von Humboldt Foundation), he then joined the group of Prof. R. D. Little at the University of California, Santa Barbara, as a postdoctoral researcher. In 2014, he returned to Germany, where he started his independent career at Rostock University with financial support by the Fonds der Chemischen Industrie (Liebig Fellowship). The current focus of his research group is on electroorganic synthesis and on electrocatalytic CO2 reduction. Benjamin Schille was born in Berlin (Germany). He studied chemistry at Rostock University (Germany), where he received his M.S. degree in 2015. In the same year he joined the group of Robert Francke as a Ph.D. candidate with financial support by the FCI (Fonds der Chemischen Industrie). His research interests include the development of new electrocatalysts and sustainable electrolysis protocols. Michael Roemelt studied chemistry at the University of Bonn (Germany) and Stanford University (United States). In 2013, he obtained his Ph.D. degree in theoretical chemistry from the University of Bonn for his work under the supervision of Prof. Frank Neese on theoretical core electron spectroscopy of transition metal compounds. Afterwards he joined the group of Prof. Garnet K.-L. Chan at Princeton University (United States) for a postdoctoral stay to work on the development of electronic structure methods for large molecules with many strongly correlated electrons. Since 2015 he has been the head of an independent research group at the RuhrUniversity Bochum (Germany) and the Max-Planck Institute for Coal Research in Mülheim an der Ruhr (Germany). His current research interests primarily focus on the development and application of quantum chemical methods to study chemical processes involving complex molecular systems.

Ligands

BIAN dmbpy bpy DAB dpq diphos hp-bpy

ACKNOWLEDGMENTS We acknowledge financial support through a Liebig Fellowship (Fonds der Chemischen Industrie) and by the German Federal Ministry of Education and Research (BMBF, project no. 031A123). M.R. gratefully acknowledges funding from the Otto-Hahn award program of the Max-Planck Society. BE

bis(aryl)acenaphthenequinonediimine dimethyl-2,2′-bipyridine 2,2′-bipyridine 1,4-diazabutadiene dipyrido[3,2-f:2′,3′-h]quinoxaline 1,2-bis(diphenylphosphino)ethan 6-(2-hydroxyphenol)-2,2′-bipyridine DOI: 10.1021/acs.chemrev.7b00459 Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews pd-bpy phen py qu tpy

Review

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4-phenyl-6-(phenyl-2,6-diol)-2,2′-bipyridine phenanthroline pyridine quinoline 2,2′:6′,2″-terpyridine

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Homogeneously Catalyzed Electroreduction of Carbon Dioxide

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