Electrons, Photons, Protons and Earth-Abundant Metal Complexes for

Review pubs.acs.org/acscatalysis

Electrons, Photons, Protons and Earth-Abundant Metal Complexes for Molecular Catalysis of CO2 Reduction Hiroyuki Takeda,† Claudio Cometto,‡ Osamu Ishitani,*,† and Marc Robert*,‡ †

Department of Chemistry, Faculty of Science, Tokyo Institute of Technology, 2-12-1, NE-1 O-okayama, Meguro-ku, Tokyo, 152-8550, Japan ‡ Université Paris Diderot, Sorbonne Paris Cité, Laboratoire d’Electrochimie Moléculaire, Unité Mixte de Recherche Université−CNRS no. 7591, Bâtiment Lavoisier, 15 rue Jean de Baïf, 75205 CEDEX 13 Paris, France ABSTRACT: Electrochemical and photochemical reduction of CO2, or a smart combination of both, are appealing approaches for the storage of renewable, intermittent energies and may lead to the production of fuels and of value-added chemicals. By using only earth-abundant metal (Cu, Ni, Co, Mn, Fe) complexes, cheap electrodes and/or cheap sacrificial electron donors and visible light sensitizers, systems functioning with molecular catalysts have been recently designed, showing promising results, in particular, for the two-electron reduction of the carbon dioxide. By combining experimental and mechanistic studies, key parameters controlling the catalysis efficiency have been deciphered, opening the way to the design of future, more efficient and durable catalysts, as well as to the development of electrochemical or photoelectrochemical cells, all being key steps for the emergence of applied devices. The most recent advances related to these issues are discussed in this review. KEYWORDS: CO2 reduction, earth-abundant metals, molecular catalysis, electrochemical catalysis, photochemical catalysis

1. INTRODUCTION The conversion of CO2 into fuels or commodity chemicals by means of solar energy remains a chemist’s dream. It would open the way to an economy based on renewable energy, and it would potentially give access to energy to everyone on earth.1 One may directly use light to perform the reactions or to preliminarily convert the light into electricity in photovoltaic cells and then use the electrons to reduce the carbon dioxide. Direct hydrogenation of CO2 is another approach. Catalytic strategies are necessary to achieve the transformations due to the inertness of the substrate, but developing selective, efficient, and stable catalysts currently remains a highly challenging task. The selectivity issue is partly due to the reaction product’s close thermodynamics (the apparent E0 in water at pH 7 range from −0.24 V vs SHE for CO2/CH4 to −0.61 V vs SHE for CO2/HCOOH). Another facet of the selectivity, notably when doing experiments in aqueous media, is to avoid competitive hydrogen evolution. Molecular approaches, using electrochemistry or photochemistry, may be interesting in that connection because the fine-tuning of the catalyst structures may lead to a good selectivity. Molecular catalysts are, however, considered as less durable than solid catalysts because few of them can convert CO2 in aqueous conditions, and there are few examples of molecular catalysts able to reduce CO2 beyond the two-electron reduction products (CO, HCOOH). To add one hurdle more, catalysts based on earth-abundant metals or materials need to be designed, notably in view of their integration into devices for economically viable © XXXX American Chemical Society

applications. Molecular catalysts have involved, in particular, reduced states of transition-metal complexes, and we will exclusively focus on metal complexes in this review, including the most abundant metals, namely, Cu, Co, Ni, Mn, and Fe. Thus, this review describes recent advances regarding the catalytic reduction of CO2 with molecular catalysts involving earth-abundant metals, using electrodes or photons as input energy. The first section is devoted to the molecular electrochemical reduction of CO2, whereas the second one is dedicated to the photochemical approaches.

2. RESULTS AND DISCUSSION 2.1. Electrochemical Reduction of CO2. The use of electrochemical techniques in particular cyclic voltammetry, has proven particularly fruitful to get mechanistic insights in the catalytic processes as well as to benchmark the intrinsic properties of the catalysts with the help of catalytic Tafel plots (turnover frequency (TOF) as a function of the overpotential; see below for an illustration for the CO2-to-CO conversion in aprotic solvents). Such mechanistic studies are complementary of spectroscopic studies, including in situ and operando experiments. Mechanistic studies need to be vigorously Received: August 1, 2016 Revised: October 22, 2016 Published: October 28, 2016 70

DOI: 10.1021/acscatal.6b02181 ACS Catal. 2017, 7, 70−88

Review

ACS Catalysis Table 1. Electrochemical CO2 Reduction Using Abundant-Metal Complexes entry

catalyst

conditions (solvent, electrode)

potential vs SCE (V)

products (FE%)

ref

1 2 3

1 2 3

CH3CN, Hg pool CH3CN, glassy carbon DMF, Hg pool

−2.15 −0.27 −1.4

6 7 10

20

H2O/CH3CN (1:4), glassy carbon H2O (pH 4.1), Hg pool H2O (pH 2), Hg/Au amalgam

−1.45 −1.24 −1.23

C2O42− (90%) C2O42− (96%) HCOO− (75%) CO (24%) CO (90%) CO (96%) CO (66%) H2 (15%) CO (22%) CO (20−30%) CO (45%) H2 (30%) CO (98%) CO (85%) H2 (15%) CO (51%) H2 (24%) CO (100%) CO (98%) CO (70%) HCOO− (22%) CO (76%) CO (96%) CO (94%) CO (60%) HCOO− (35%) CO (100%) CO (90%) CO (97%) CO (80%) CO (60%) HCOO− (51%) H2 (13%) CO (93%) H2 (4%) CO (95%) CO (67%) H2 (12%) H2CO (39%) H2CO (28%) CO (89%) H2 (5%) CO (41%) H2 (60%) CO (91%) CO (41%) HCOO− (74%) CO (13%) C2O42− (7%) HCOO− (75%) HCOO− (94%) HCOO− (96%)

4 5 6 7 8 9

4 5

CH3CN, glassy carbon CH3CN, pyrolytic graphite CH3CN (+ 10 M H2O), glassy carbon

−1.8 −1.54 −1.76

10 11

6 7

DMF (+ 1.2 M TFE), glassy carbon CH3CN (+ 5% H2O), glassy carbon

−2.35 −1.40

H2O (pH 7.0), glassy carbon, cat. supported in Nafion membrane

−1.55

12 13 14 15

8 9 10

CH3CN (+ 1.4 M TFE), glassy carbon CH3CN (+ 0.3 M TFE), glassy carbon CH3CN, glassy carbon

−2.2 −1.8 −1.8

16 17 18 19

11 12 13

CH3CN (+ 5% H2O), glassy carbon CH3CN, glassy carbon DMF (+ 1 M phenol), Hg pool DMF (+ 6.7 M 1-propanol), Hg pool

−1.5 −1.92 −1.70 −1.7

20 21 22 23 24 25

17 16 21 22 23 24

DMF (+ 3 M phenol) H2O (pH 7.2), glassy carbon H2O (+ 0.5 M KHCO3), carbon (gas diffusion electrode), 20 atm CO2 H2O (pH 4.6), glassy carbon, cat. supported on multiwalled carbon nanotubes H2O (pH 3.0), pyrolytic graphite, 10 atm CO2 H2O (pH 7.0), pyrolytic graphite, cat. supported in Nafion membrane

−1.2 −1.1 −1.0 −1.34 −1.02 −1.10

26

25

−1.27

27 28

26 27

solvent: H2O (pH 7.3) electrode: glassy carbon, multiwalled carbon nanotubes, Nafion membrane H2O (pH 7.3), glassy carbon, cat. supported multiwalled carbon nanotubes CH3CN/H2O (19/1), mesoporous TiO2

29 30 31

28 29 30

H2O, glassy carbon H2O, glassy carbon H2O (pH 4.7), graphite

−1.15 −1.12 −1.25

32

31

CH3CN (+ 1 M TFE), FTO

−1.54

33 34 35

32 33 36

H2O (pH 7.3), carbon H2O (pH 7.3), carbon DMSO, glassy carbon

−1.35 −1.35 −1.76

36 37 38

37 38

DMF, glassy carbon CH3CN (+ 5% H2O), glassy carbon H2O (pH 7), glassy carbon

−1.25 −1.2 −1.2

−1.3 −1.3

11 9 36a 14 16 17 19 20 36c 21a 21b 21d 23 24 25 54 33 34 37d 38 37e 39 40 44 45 46 46 48 50 49b 49b 55

56 57 57

macrocycles) and metals (Ni, Co, Mn, Fe), leading to efficient homogeneous electrochemical CO2 reduction. It could be noted that many of the molecular catalysts currently studied and published are similar or even identical to those that have been thoroughly investigated in the 1980s and 1990s, during the f irst wave of the CO2 reduction studies. We will in the following focus on a selection of characteristic examples, including the most

developed so as to get a deeper understanding of the catalytic mechanisms that will in turn lead to new, more efficient catalysts. Transition-metal complex catalysts of CO2 electrochemical reduction based on earth-abundant metals have been carefully reviewed some time ago,2 with the description of the main classes of ligands (polypyridyl, phosphine, cyclam, and aza-macrocyclic ligands, as well as porphyrins, phtalocyanines, and related 71


Review

ACS Catalysis recent advances, notably those related to the use of Mn and Fe complexes, because these two metals are the two top most abundant metals in Earth’s crust and provide, like in photocatalytic systems, the most promising candidates for the CO2 reduction. The data discussed below for the electrochemical reduction of the carbon dioxide are summarized in Table 1. 2.1.1. Homogeneous One-Electron Reduction of CO2. Oxalate could be obtained in good yields from direct reduction of CO2 in aprotic solvent at inert electrodes, by applying negative potential close to the standard potential of the CO2/CO2•− redox couple (−2.21 V vs SCE in DMF + 0.1 M Bu4NClO4).3 For example, at a mercury electrode in DMF as solvent, a faradaic efficiency as high as 89% was obtained (10% CO was then identified as byproduct). In these cases, it has been demonstrated that oxalate is issued from the dimerization of the radical anion of CO2.4 Radical anions of aromatic esters and nitriles have also been used to catalyze the electrochemical CO2 reduction into oxalate in homogeneous conditions.5 For example, with methyl benzoate MB (E0(MB/MB•−) = −2.202 V vs SCE in DMF) as catalyst, 100% faradaic yield for oxalate was measured in DMF at a mercury pool and 1.6 mA/cm2 current density. The catalytic process is not simply an outer-sphere electron transfer from the radical anion of the donor to CO2 followed by dimerization, but rather, it involves a nucleophilic addition of the catalyst radical anion on CO2 forming an oxygen−carbon bond, which further breaks homolytically, generating the parent ester and CO2•−, which would then dimerize. A similar process occurred when using electrogenerated macrocyclic Ni(I) chelate complexes (1, Chart 1), leading, upon electrolysis at a potential close to −2.15

Chart 2

mercury electrode (see details below in this section), while the CO2 reduction in an aprotic solvent (DMF) under 1 atm CO2 led, upon 5 h electrolysis, to the production of formic acid mainly with 75% faradaic efficiency and a TON of 3.2 (and 15% of CO) at 460 mV overpotential (entry 3 in Table 1).10 It is a unique example for a molecular system, and it was suggested that in DMF the NiI reduced active species binds to CO2 and that after further reduction and protonation a NiII−OCHO intermediate was formed rather than a NiII−COOH intermediate, which would have furnished CO and H2O as products. Note that in this mechanism, no Ni hydride is postulated as an intermediate toward HCOOH production. Replacing the mercury electrode used in the above-mentioned studies by an inert glassy carbon electrode led also to good catalytic activity for CO production in 1:4 water/acetonitrile mixture as solvent, although with a slower rate. It was shown that 1 h of electrolysis at −1.21 V vs NHE (at 470 mV overpotential) and a current density of 1.8 mA/cm2 gave 90% faradaic efficiency (entry 4 in Table 1).11 As initially identified by Sauvage et al.,9b it was recognized that the NiI complex strongly binds to CO to give a deactivated species NiI(cyclam) (CO)+ and that catalysis is kinetically controlled by the CO loss from this NiI−CO adduct. An elegant way to remediate this deactivation reaction consists of using a CO scavenger, and the tetramethylated NiII(TMC)2+ cyclam bearing a methyl group at each nitrogen atom was used in the purpose.12 The one-electron reduced species NiI(TMC)+ has both a very large binding constant for CO (K = 1.2 × 105 M−1 s−1) and a lower reactivity toward CO2 than the parent, unsubstituted cyclam. When a 1:20 [NiII(cyclam)2+]:[NiII(TMC)2+] ratio was prepared in 1:4 water/acetonitrile solvent, the catalytic current for CO2 reduction was increased by a factor 10, indicating a larger number of catalytic turnovers. These studies also point to a degradation pathway due to the formation of Ni(0) species formed upon reduction of the NiI(cyclam)(CO)+ species. Recent mechanistic insights have been gained from density functional theory (DFT) calculations.13 That reactivity on Hg surface can be sustained for longer times than on carbon surfaces may arise from the fact that the catalysis takes place at more positive potentials on Hg, slowing down the formation of the deleterious Ni(0) species.12 N-heterocyclic carbine-isoquinoline complexes (4, Chart 2) have been recently proposed for both photocatalysis (see the photochemical section of this review) and electrochemical catalysis of CO2 into CO.14 It was shown that 8 h of

Chart 1

V vs SCE in an acetonitrile solution saturated with CO2, to oxalate production with faradaic yield equal or larger than 90% and a maximum TON of 750 (entry 1, Table 1).6 Note that in that case, the electrode was a mercury pool, and the counter electrode (not separated from the cathodic compartment) was an aluminum wire. A recent striking example of molecular catalyst for efficient CO2-to-oxalate conversion was given by a dinuclear Cu complex (2, Chart 1), able to produce oxalate in quantitative yield at very low overpotential at a glassy carbon electrode (6 TON were obtained upon 7 h electrolysis at −0.03 V vs SHE upon continuous purging of CO2, entry 2 in Table 1).7 Remarkably, it was possible to both capture and reduce CO2 from the air with the complex to make the C2 product. 2.1.2. CO2-to-CO Conversion in Aprotic Solvents. Ni Complexes. The use of Ni macrocycle complexes has been pioneered by Eisenberg et al.8 and Sauvage et al.9 The Ni(cyclam)2+ (3, Chart 2) provides a remarkable example for which the CO2-to-CO conversion occurs very efficiently and selectively (high TON and high FE) in water (pH 4 to 5) at a 72


Review

ACS Catalysis Chart 3

Table 1). Importantly, they gave evidence that the real catalyst is a molecular species and not metallic nanoparticles that may arise from ligand reduction and loss. More insights in the catalytic mechanism were obtained from combined DFT calculations and FT-IR spectroscopy, as detailed in the photochemical section, leading notably to the experimental identification of the doubly reduced complex bound to CO2, a key intermediate before cleavage of the C−O occurs.17,18 Such an integrated approach, combining electrochemical and photochemical studies along with quantum calculations and spectroscopic investigation will certainly provide more insights in the fine details of the catalytic mechanisms and further help the rational design of new efficient catalysts. Another example of an efficient Co catalyst for the CO2-to-CO conversion was discovered with a macrocycle including four aminopyridine as ligands, linked by pendant amine groups (6, Chart 2).19 Highly selective CO formation (98%) could be sustained for a couple of hours in a DMF solution saturated with CO2 and containing 1.2 M of trifluoroethanol as a

electrolysis at 840 mV overpotential affords CO with 90% faradaic yield during the first 30 min, while the yield drops to 22% for the 8 h period (entry 7 in Table 1). The electrolysis, performed in a 0.02 mM of catalyst at a glassy carbon electrode, furnishes less than 1 μmol of CO. Clearly, there is a need for new, highly active catalysts using Ni as metal center. Co Complexes. The same considerations apply to Co as metal catalyst. A representative example concerns the use of tetraazamacrocycle (5, Chart 2) that was first investigated by Tinnemans et al.15 and Che. et al.16 in the 1980s. For example, Che et al. have shown that electrolysis in acetonitrile at a carbon electrode produces CO in 20−30% faradaic yield at 940 mV overpotential, with no H2 as byproduct (entry 8, Table 1). The active catalytic species is the CoI ligand radical anion species (i.e., the doubly reduced species issued from the CoII neutral). In recent studies, Peters et al. have found that in the presence of 10 M water, the faradaic efficiency could be increased to 45% but along with non-negligible amount of H2 (30% FE, entry 9 in 73


Review

ACS Catalysis

acetonitrile solution without addition of a weak acid, with a TOF similar to what was obtained with 7.21d Remarkably, in addition to CO formation (70% faradaic yield after 4 h electrolysis at −1.8 V vs SCE), HCOOH was also produced with a 22% faradaic yield (entry 15, Table 1), suggesting an alternative reaction pathway. This pathway may involve the formation of a hydride intermediate issued from intramolecular protonation in the reduced catalyst anion, that would further insert CO2 and evolve formic acid. Recently, a derivative of fac-Mn(bpy)(CO)3Br bearing an hydroxyphenyl substituent appended to the bipyridine ligand (11, Chart 3) was synthesized and shown to produce CO efficiently in good yield (76% in 4 h electrolysis; no indication was given if formic acid has been or not obtained as a byproduct) at 540 mV overpotential in ACN/H2O (5%) mixture (entry 16, Table 1).23 In the search of new catalysts, MnI complexes with a planar coordinating tridentate PNP ligand (12, Chart 3) and bearing a positive charge was recently synthesized and studied, showing excellent efficiency for CO production (close to 100%) in dry acetonitrile during more than 5 h electrolysis at 960 mV overpotential (entry 17, Table 1).24 The presence of carbonate points to a reductive disproportionation of the CO2. Surprisingly, catalysis efficiency drops upon addition of 5% water to the solvent with concurrent H2 evolution (41% faradaic efficiency upon 1.7 h electrolysis), suggesting a competitive pathway to CO2 reduction that has not been observed within the abovedescribed family of Mn(L)(CO)3Br complexes. Iron Complexes. Iron porphyrins (Chart 3), when electrochemically reduced to the Fe0 species, are among the most efficient molecular catalysts for the CO2-to-CO conversion in aprotic solvent (DMF, ACN), in terms of both efficiency, catalytic rate and robustness.2a,25 14, 15, 16, and 17 (Chart 3) are the most active compounds within this class of catalysts. Recently, cofacial Fe porphyrin dimers have also shown excellent catalytic performances for CO production.26 In all of these cases, weak Brönsted acids such as water, trifluoroethanol, and phenol strongly enhance catalysis, and Lewis acids exert the same effect. For example, with 13 as catalyst and phenol as acid, CO yield varies from 100% to 94% when the acid concentration increases from 0.1 to 1 M, for electrolysis at −1.46 V vs SHE (mercury pool, DMF as solvent) corresponding to 600 mV overpotential (entry 18, Table 1). From cyclic voltammetry studies, the rate constants kcat for catalysis were obtained systematically as a function of both the CO2 concentration and the acid concentration.25c,27 The reaction order toward CO2 is 1, and the reaction order toward phenol is 2 at low acid concentration and 1 at large concentration, leading to the mechanism shown in Scheme 1. Catalysis starts with the formation of a 2−Fe0-CO2 adduct (DFT calculations suggest that −FeI−CO2•− is a predominant resonance form), which is stabilized by H-bonding with one PhOH molecule. Protonation from a second PhOH molecule leads to C−O bond cleavage and CO is released after an homogeneous one electron reduction of the FeII−CO adduct by the Fe0 species. Recent studies by resonance Raman spectroscopy at low temperature further confirm these mechanistic studies and provide new insights in the catalytic mechanism.28 By installing acid functionalities directly on the phenyl groups of the iron porphyrin (14 and 15, Chart 3), strong enhancement of the catalysis was achieved while the CO formation remains highly selective. TOF in the range of 102 s−1 (at 450 mV overpotential) and of 104 s−1 (at large overpotential) were measured while electrolysis could be pursued for several hours

weak acid to help the catalysis, at a glassy carbon electrode and 680 mV overpotential. In total, 6.2 TON of CO was formed, and it was suggested that the pendant NH groups helps stabilize the CoI−CO2 adduct, similarly to what has been shown with pendant OH groups in Fe porphyrins (see below) and with some other related Co macrocyles in the photochemical reduction of CO2 (see section 2.1.1). Mn Complexes. A new class of highly active Mn catalysts was introduced in 2011 by Chardon-Noblat and Deronzier et al. They used Mn(L) (CO)3Br complexes with L being either a 2,2′bipyridine (7, Chart 3) or a 4,4′-dimethyl-2,2′-bipyridine ligand.20 A 4 h electrolysis with 7 in ACN + 5% H2O at −1.40 versus SCE (420 mV overpotential) and 0.2 mA/cm2 converts CO2 into CO with unity faradaic efficiency. After 22 h, CO is still produced with 85% yield with 15% of H2 as byproduct, showing the remarkable stability of the catalyst, even if the current has then dropped to 0.07 mA/cm2 (entry 11, Table 1). With L = 4,4′dimethyl-2,2′-bipyridine, CO remains selectively produced for more than 18 h (TON 34) under the same conditions. Further pulsed-EPR studies complemented by DFT studies have shown that the catalytic mechanism involves the formation a Mn0−Mn0 dimer after one electron reduction of the MnI starting complex. Oxidative addition of CO2 and H+ then leads to a MnII− carboxylic acid intermediate, which finally gets reduced with a second electron and protonated to evolve a CO and a H2O molecules. Kubiak et al. further tune the substituents appended onto the L ligand while adding various weak Brönsted or Lewis acids to the solvent, thus getting excellent performances.21a,b,22 With L = 4,4′-ditertbutyl-2,2′-bipyridine (8, Chart 3), a maximum TOF of 340 s−1 in the presence of 1.4 M trifluoroethanol (approximately 300 times more than with 7) was calculated from analysis of the voltammograms, and electrolysis at −2.2 V vs SCE led to current density as high as 30 mA/cm2 with 100% CO faradaic efficiency for several hours (entry 13 in Table 1). Finally, with L = 6,6′-dimesityl-2,2′bipyridine (9, Chart 3), a maximum TOF of 5 × 103 s−1 was reached in ACN + 1.4 M trifluoroethanol, with again high faradaic efficiency for CO evolution at several mA/cm2 and during 7 h of electrolysis (CO remains selectively produced for about 24 h, although the faradaic efficiency declined, entry 14 in Table 1). Mechanistic and spectroscopic joint studies have shown that the complex is successively reduced with two electrons (the second reduction being easier than the first one), producing the active species Mn0(L)(CO)3−, while the bulky bipyridine substituent prevents dimerization to occur. The Mn0(L)(CO)3− was independently prepared by chemical reduction of the starting complex and characterized by NMR, X-ray crystallography, and FTIR spectroscopy. The doubly reduced, negatively charge complex then binds to CO2 and is protonated, leading to a MnI−CO2H intermediate. This later species needs to be reduced to release CO, necessitating to apply a 400 mV more negative potential as compared to the potential at which Mn0(L)(CO)3− is produced. By replacing the trifluoroethanol by a Lewis acid (Mg2+, 0.1 M), catalysis goes through a reductive disproportionation (2CO2 + 2e ↔ CO + CO32−), a second molecule of CO2 itself playing the role an acid in the C− O bond cleavage in the MnI−CO2−Mg intermediate adduct. A TOF of 630 s−1 for CO could be obtained at an overpotential as small as ca. 350 mV, showing the boosting effect of the Lewis acid on the catalytic process. With a proton source close to the metal introducing L = 4phenyl-6-(1,3-dihydroxybenzen-2-yl)-2,2′-bipyridine (10, Chart 3), the catalyst remains highly active toward CO production in an 74


Review

ACS Catalysis Scheme 1. Mechanism for the CO2-to-CO Conversion with 13 in the Presence of Phenol (AH = PhOH)

TOF = 1+

kcat[CO2 ] F 0 0 ⎤ exp⎡⎣ RT (ECO )⎦exp − Ecat 2 /CO

(− RTF η)

(1)

0

where E cat is the standard redox potential of the catalyst. Scheme 2 provides TOF = f(η) for the most active catalysts that can be found in the literature, illustrating the exceptional activity of iron porphyrins, in particular, that of 17. Scheme 2. Catalysts Benchmarking for the CO2-to-CO Conversion [log TOF = f(η = E0CO2/CO − E)]a

with little catalyst decomposition with 14 (note that an external acid, for example, water or phenol, was added to the solvent in all of these experiments so as to efficiently reprotonate the catalyst itself during the course of the reaction). The prepositioned −OH groups both play the role of H-bonding substituents by stabilizing the Fe0-CO2 adduct and of proton donors (or proton relays) for the protonation and then the cleavage of the protonated intermediate. Note that in this case and in contrast to what occurs with 13, the second electron transfer takes place before C−O bond cleavage occurs because of the strong stabilization of the intermediate protonated adduct.58 Introduction of 10 fluorine atoms on 2 of the phenyl rings of the porphyrin (15, Chart 3) also lead to high TOF and catalytic efficiency at even more positive potentials as a result of a 70 mV anodic shift of the FeI/Fe0 redox standard potential as compared to 14.29b By introducing four charged trimethylanilinium groups in para- or ortho-position of the phenyls (16 and 17, respectively, Chart 3) and adding a large concentration of phenol to assist the cleavage of the C−O bond of CO2, unprecedented catalytic efficiency has been obtained. With 17 as catalyst, catalysis led to selective formation of CO with a TOF of 106 s−1 at an overpotential of 220 mV (TOF at zero overpotential is in the range of 300 s−1). Remarkably, the catalyst is highly stable in catalytic conditions, showing no significant alteration and no decrease in CO selectivity after more than 80 hours electrolysis at 220 mV overpotential in DMF with phenol 3 M (entry 20, Table 1).33 The reasons for such activity likely stand from the stabilization of the initial Fe0-CO2 adduct by the interaction between the negative charge borne by the oxygen atoms of CO2 in this adduct and the positive charges borne by the trimethylanilinium substituents on the porphyrin phenyls (the closer distance between the trimethylanilinium groups and CO2 in 17 may also explain its even higher activity as compared to 16). The catalytic properties of the homogeneous catalysts for the CO2-to-CO conversion could be compared by plotting the graphs relating the TOF to the overpotential η (the difference between the standard potential of the substrate, E0CO2/CO, and the applied potential E). It allows benchmarking the catalysts independently of the characteristics of the electrolytic cell in use, and it necessitates deciphering the reduction mechanism for each catalyst complex. In the case where the electron transfer from the electrode to the catalyst is fast and the mechanism consists of a single catalytic step (or an apparent single step) with a rate kcat, the log TOF vs η is given by eq 1:

a

1: 17, 2: 16, black: 13, dark blue: 14, brown: 15, red: 9b,21b cyan: Re(bpy) (CO)3py,30 yellow: m-(triphos)2Pd2,31 dark green: 8,21a dotted brown: Ru(tpy) (Mebim-py),32 dotted pink: Ru(tpy) (bpy).32 [py: pyridine, tpy: 2,2′:6′,2″-terpyridine, Mebimpy: 2,6-bis(1-methyl benzimidazol-2-yl)pyridine]

2.1.3. CO2-to-CO Conversion in Water. By appending positively charged substituent on each para position of the four phenyls, a water-soluble catalyst was obtained (Chart 3, 16).34 While being highly active toward CO2 reduction in DMF (as described just above), it appears that 16 is also remarkably selective for CO production in unbuffered aqueous conditions at neutral pH (7.2). Selectivity in CO of 90% (the main by-product being hydrogen) was obtained at current densities on the order of 0.1 mA cm−2 and 450 mV overpotential (Eelectrolysis = −0.86 V vs SHE) for more than 1 day of electrolysis (entry 21, Table 1). One major reason for the selective reduction of CO2 stands from the fact that carbonic acid H2CO3 formation is thermodynamically uphill and kinetically slow, favoring the nucleophilic attack of the Fe(0) species onto CO2 and making it able to resist the fast H2 evolution pathway. 3 (NiII(cyclam)2+, Chart 2) is another example of molecular catalyst able to selectively reduce CO2 into CO in water.9 In a CO2 saturated solution without any added buffer (pH 4.1), 4 h of electrolysis at −1.00 V vs SHE, corresponding to an overpotential of 640 mV, led to quantitative formation of carbon monoxide with faradaic efficiency of 96% and a turnover number 75


Review

ACS Catalysis

Ti1−xZrO2), although in that case, no CO2 catalysis has been reported.36b Electrodes, in particular carbon electrodes, have been also film coated with molecular insoluble catalysts, leading to CO2 catalytic activity in water, notably with porphyrins and phtalocyanines.37 The tetraphenyl Co porphyrin (21, Chart 5) when deposited at carbon black gas diffusion electrode produces CO with 97% efficiency under 20 atm of CO2 and at −0.76 V vs SHE (corresponding to an overpotential of 230 mV) in a 0.5 M KHCO3 solution (entry 22, Table 1). More recently, a Co chlorin complex (22, Chart 5) adsorbed onto carbon nanotubes was used to evolve CO in 80% yield at pH 4.6, with TOF of 140 h−1 and a maximum TON of 1100 at 700 mV overpotential (entry 23, Table 1).38 Remarkably, a Co protoporphyrin (23, Chart 5) was adsorbed onto pyrolytic graphite and was able to produce CO even at very acidic pH (3) under 10 atm of CO2 with a 60% faradaic efficiency and 500 mV overpotential, while at pH 1, 2.5% of CH4 were obtained, showing that CO2 could be reduced with molecular catalyst beyond the two electron reduction products (entry 24, Table 1).37e Another notable example involves the use of bis(2,2′:6′,2″-terpyridine)CoII complex incorporated into a Nafion membrane (24). In a CO2-saturated solution at pH 7, formic acid was produced with 51% faradaic efficiency in a 4 h electrolysis experiment at 250 mV overpotential, with a TON of 11 (entry 25, Table 1).39 Although H2 was produced as byproduct with 13% faradaic efficiency and the catalysis was globally slow, this is a rare case of immobilized molecular catalyst able to produce formic acid with good selectivity. Mn and Fe catalysts can also be appended onto carbon surfaces by noncovalent immobilization techniques. fac-Mn(bpy)(CO)3Br (7, Chart 3) was casted in a Nafion membrane at pH 7, leading to CO:H2 ratio of 2:1 at an overpotential close to 750 mV for more than 4 h of electrolysis, and with a TON of 460 (entry 12 in Table 1).36c Recently, a pyrene-derivatized iron triphenylporphyrin (25, Chart 5)

of 116 (entry 5 in Table 1). The catalyst was operated at a mercury electrode, with favorable specific interactions between the catalytic species and the mercury surface, making uncertain the determination of the active catalytic species. Variously substituted Ni cyclams at that the nitrogen atoms have been recently investigated, showing in all cases that CO is produced in high yields (84%−92% FE) at pH 5 and 550 mV overpotential (1h electrolysis). In very acidic solutions (pH 2), the selectivity was shifted toward hydrogen evolution, with a CO:H2 ratio in between 0.04 and 0.4 as a function of the exact nature of the catalyst. Note that at pH 5, the two catalysts 18 and 19 (Chart 4) Chart 4

show higher current densities than the parent NiII(cyclam)2+, which was ascribed to favorable surface geometries for adsorption onto the mercury working electrode as well as to favorable environment for accommodating and reducing CO2 at the metal center.35 Recently, a carboxylic acid appended derivative of 3 shows excellent catalytic selectivity for CO formation at a Hg/Au amalgam in acidic water at pH 2, with a 4:1 CO:H2 ratio and a total faradaic efficiency of 81% during 1 h electrolysis at −0.99 V versus SHE, corresponding however to an overpotential of 740 mV (20, Chart 4, entry 6 in Table 1).36a Catalytic activity could be maintained over several hours. The catalyst may also be attached to semiconductive surface (TiO2, Chart 5

76


Review

ACS Catalysis

in the MOF likely helps access of both reactant and solvent to catalytic sites, the reaction was suggested to be limited by slow charge diffusion within the material. Moreover, the use of a weak acid (trifluoroethanol, TFE) to boost the catalysis led to catalyst chemical degradation after a couple of hours. In the case of cobalt porphyrin-derived 2D covalent organic frameworks (32), the material was deposited onto conductive carbon fabric and led to the selective conversion of CO2 into CO in aqueous solutions (pH 7.3 bicarbonate buffer). Optimal results were obtained at 550 mV overpotential (Eelectrolysis = −1.11 V vs SHE) with a catalytic activity that could be maintained over a 24 h period (entry 33, Table 1). A TON in CO of 3901 and 91% faradaic efficiency was measured when the porphyrin units were linked by biphenyl-4,4′-dicarboxaldehyde struts. With bimetallic COF (Co, Cu) and a 1:100 Co:Cu ratio (33), a TON of 24 000 for carbon monoxide was obtained, with 41% faradaic efficiency (5.8 days electrolysis, entry 34 in Table 1), illustrating the high potential of these structures.49b Integration of molecular catalyst able to transform CO2 into products in an integrated electrochemical cell, able to split, for example, CO2 and H2O, into CO and O2 along the following equations:

derived from 14 (Chart 3) was adsorbed on multiwalled carbon nanotubes via noncovalent interactions, and the modified carbon material proved to be a selective, stable, and fast catalyst for CO2to-CO conversion at low overpotential (480 mV) in neutral pH unbuffered water.40 After optimization of the system and upon 12 h electrolysis, CO was formed with a very high selectivity (92:8 CO:H2 ratio) and an excellent total faradaic yield (97%), with a TOF of 164 h−1 and a maximum TON in CO of 1574 (entry 26, Table 1). These examples illustrate that, starting from a highly efficient catalyst, the catalytic activity could be transferred in water and onto surfaces. Covalent attachment of a molecular catalyst onto surfaces is another appealing option that may lead to the design of new, versatile catalytic materials. It is worth noting that although redox molecules could be attached to a variety of conductive surfaces,41 few examples leading to CO2 catalysis have been reported so far. The covalent attachment of a Co-alkyne-modified porphyrin to an azide-functionalized diamond surface42 and the electrochemical grafting of a Co-tpy (tpy = terpyridine) complex43 led in both cases to the production of a small amount of CO, in the former case at 950 mV overpotential in ACN (CO was not quantified but only qualitatively identified from FTIR experiment) and at 610 mV overpotential in DMF in the latter case. Finally, a derivative of 14 with one phenyl ring being substituted by a carboxylic acid group at the para position was also covalently grafted to carbon nanotubes (26). After deposition on glassy carbon, the modified electrode shows high activity toward CO2 with CO formation (90% selectivity) in water pH 7 at a 510 mV overpotential, with good stability over several hours of electrolysis and a TOF of 178 h−1 (532 catalytic cycles per immobilized electroactive catalyst were obtained in 3 h, with a maximum TON of 750, entry 27, Table 1).44 Mesoporous TiO2 has also been modified with a Mn complex (fac-[MnBr(4,4′-bis(phosphonic acid)-2,2′-bipyridine) (CO) 3 ], 27). Anchoring of the complex through the phosphonate groups led to good selectivity for CO in CH3CN/H2O 19/1 mixture upon 2 h electrolysis, with a 67% faradaic efficiency at an overpotential of approximately 420 mV (entry 28, Table 1).45 Electropolymerization of a suitable group appended to a molecular catalyst is another way to functionalize the electrode surface. Again, few examples have proven to be successful using earth-abundant metals. An early one concerns the CO2 reduction into formaldehyde (a very rare case of CO2 product using molecular catalysts) with electropolymerized films of vinylterpyridine complexes of Co and Fe (28 and 29, respectively). In CO2-saturated aqueous solutions containing only 0.1 M NaClO4 at ca. 530 mV overpotential, faradaic yield for HCHO was 39% with the former metal and 28% with the later, and with very high TON (1.1−1.5 × 103, entries 29 and 30, Table 1).46 Among more recent examples (see ref 47a for a review as well as ref 47b), a coordination polymer of poly-4-vinylpyridine with Co phtalocyanine (30) led to the generation of CO from CO2 at pH 4.7, with a very good selectivity (89% faradaic yield) at 615 mV overpotential, for a period of 2 h of electrolysis (2 mA/ cm2).48 In that case, only 5% of H2 was produced as a byproduct (entry 31, Table 1). Recently, thin films of nanosized metal−organic frameworks (MOF) incorporating Co49 or Fe porphyrins50 deposited onto carbon or FTO electrode were shown to catalyze the CO2-to-CO conversion. In the case of the Fe-based MOF-film electrocatalysts (31), electrolysis for 3.2 h in ACN produces CO in moderate yields (41%, 1520 TON) with the remaining electrical charge being used for H2 evolution (60%), at 580 mV overpotential (entry 32, Table 1). Even if the nanoscale porosity

CO2 + 2H+ + 2e− ⇄ CO + H 2O H 2O ⇄ 1/2O2 + 2H+ + 2e−

CO2 ⇄ CO + 1/2O2

(E 0CO2 /CO = − 0.11V)

(E 0O2 /H2O = 1.23V)

(ΔG 0 = 2.68eV)

(potentials referred to SHE) remains a challenging task, especially when selecting friendly conditions including ambient temperature, neutral pH as well as the use of cheap and abundant materials at both the cathode and the anode. Even if one remains far from any practical devices that would lead delivering several A/cm2, two recent examples are worth noting. The first one involves the design of a simple two-compartment cell (H-type cell with a glass frit between the two compartments) functioning at pH close to neutral, using perfluorinated cobalt phtalocyanine (34, Chart 5) immobilized on carbon cloth as both catalyst for CO2 reduction and H2O oxidation.51 Under 1 atm CO2 and with catalyst loading in the range of 10−8 mol/cm2 at pH 7.2, 90% CO faradaic efficiency was maintained for 11 h with a full cell voltage between 2.5 and 3 V and current densities varying from 1 to 6 mA/cm2, corresponding to a global energy efficiency in the range of 70%−50%. The second example concerns the use of an iron porphyrin bearing trimethylammonio groups as substituents at all para positions of the phenyl rings (16, Chart 3).52 The complex was deposited on carbon paper (cathode) and a 200 nm thick phosphate cobalt oxide film was grown on a steel gauze as an oxygen evolution catalyst (anode). In a two-compartment electrolyzer with a Nafion membrane as separator, current densities of 1 mA/cm2 could be sustained for more than 30 h of electrolysis at a 2.5 V cell voltage at ambient temperature, corresponding to 50% energy efficiency for the splitting of CO2 and H2O into CO and O2. Such low voltage allows the use of solar panels so as to power the cell. By improvement of the cell design, optimization of the catalyst and its loading, use of gas diffusion electrodes and higher CO2 pressure, cells based on CO2RR molecular catalysts with much higher current densities may likely be developed in the near future. 2.1.4. CO2-to-HCOOH Conversion. Hydrogenation of CO2 to formate in the presence of a base under relatively mild temperature and pressure conditions has been achieved recently with Cu-, Co-, and Fe-based molecular complexes.53 Regarding 77


Review

ACS Catalysis

dioxide are gathered in Table 2. Photocatalytic systems for CO2 reduction typically comprise a catalyst (CAT) for CO2 reduction and a redox photosensitizer (PS) that can photochemically mediate the transfer of one electron from a sacrificial donor (SD) to the catalyst (CAT). RuII complexes, such as [RuII(N^N)3]2+, where N^N is an α-diimine such as 2,2′-bipyridine (bpy) and its derivatives (39 and 40 in Chart 7), are common PSs, and can absorb visible light up to 600 nm.58 The lowest excited state of the PS (PS*, Scheme 3a, step (1)), which is the triplet metal-toligand charge transfer (3MLCT) excited state in the case of [RuII(N^N)3]2+, is a strong enough oxidant to accept an electron from the SD. In other words, emission from PS* (Scheme 3a, step (2)) is reductively quenched to the corresponding oneelectron-reduced species of the PS (PS-OERS, Scheme 3a, step (3)). Because this species is a stronger reductant than both PS and PS*, the added electron can be donated to an electron acceptor (EA), that is, the CAT or an electron mediator, providing that the redox potential for EA reduction is more positive compared to that of the PS-OERS. Then, PS-OERS returns to its original state (Scheme 3a, step (4)). A second possibility involves oxidative quenching of PS* (Scheme 3b). In that case, the reaction between PS* and EA is favorable, and EA accepts an electron from PS* directly (Scheme 3b, step (3)). The one-electron oxidized species of PS (PSOEOS) is thus produced, and it returns back to its original state by one-electron supplying from SD (Scheme 3b, step (4)). Favorable conditions for oxidative quenching include long lifetime and high reducing power of PS* as well as high concentration of EA because the initial electron transfer is a bimolecular reaction between PS* and EA. The CAT should gain at least two electrons to activate CO2 and produce stable compounds (with the exception of oxalate that may arise from CO2•− dimerization, see the Electrochemical section).59 Consequently, the CAT promotes the multielectron reduction of CO2 using the photoredox cycle of the PS (Scheme 3). As mentioned above, electrochemical catalysts have often been used in photochemical systems as well. Generally, the reduced CAT releases a ligand and accepts a CO2 molecule onto the central metal forming the corresponding CO2 adduct. These CO2 reduction processes are basically the same as those in electrocatalytic reduction systems. However, there are some significant differences when comparing photochemical and electrochemical systems. In electrochemical conditions, many highly reducing electrons (with the same potential) are available from the electrode interface to reduce CAT and the CO2 adduct. In photochemical systems, the PS-OERS supplies only one electron to the CAT through a bimolecular reaction. Therefore, a second electron must be supplied by another source, such as another PS-OERS molecule, the deprotonation product of the oxidized SD (which is usually a strong reductant), or another CAT that would have been reduced through disproportionation. Globally, the photochemical process is likely to be less efficient. 2.2.1. CoII and NiII Cyclams, Polypyridines, and Their Derivatives. CoII and NiII complexes, as shown in Chart 8, have been used as CATs in photocatalytic CO2 reduction. The first photocatalytic systems using CoII and NiII cyclams (51, 5, and 53) with the Ru complex (39) as the PS, ascorbic acid as the SD, and water as the solvent were reported by Tinnemans et al.15 in 1984, following studies in which they were used as electrochemical catalysts by Eisenberg et al. (1980)8 and Sauvage et al. (1984).9 In these systems, CO was photocatalytically produced, but the main product was H2, resulting in low selectivity (95% and >2400 based on CAT in similar reaction conditions (entry 17, Table 2).67 By modifying 82


Review

ACS Catalysis Chart 10. Co and Fe Phthalocyanines and Corroles

following, was investigated in order to overcome the abovementioned problems.75 The standard reduction potential of PSOERS of 45 was negative enough (−2.45 V vs SCE in dimethylamine) to reduce the FeIIP species (−1.05 V vs SCE for 13, − 1.02 V vs SCE for 60, and −1.00 V vs SCE for 61) and FeIP (−1.66 V vs SCE for 13, − 1.61 V vs SCE for 60, and −1.55 V vs SCE for 61) so as to form the corresponding Fe0P species. This system exhibited 10 times higher photocatalytic efficiencies than those measured in the absence of 45. The Co porphyrin (62) can also be used to produce CO, and its photocatalytic efficiency was about 1.5 times higher compared to that of the corresponding Fe porphyrin (63, entries 27−29, Table 2).75 Fe and Co phthalocyanines (64 and 65, Chart 10) and corroles (66−68, Chart 10) also exhibit photocatalytic activity for CO2 reduction. In the case of phtalocyanines, the TONs (50−90) were not higher from that of the porphyrin analogues even in the presence of 45 as PS (entries 30−32, Table 2).76 In the case of corroles, the photocatalytic activities (TONs = 40−100) were comparable to those using 13 and 21 under the same conditions, with 45 as PS (entries 33−35, Table 2).77 The photocatalytic reaction using porphyrin 14 with two OH groups on the o-positions of the four phenyl groups of the phenyl rings was investigated in a CH3CN solution containing the catalyst and TEA as the SD. Irradiation (λex > 280 nm) of the solution gave CO with higher selectivity (ca. 85%) compared to that obtained using 13 (entries 36−38, Table 2).78 The introduction of OH groups stabilizes the CO2 adducts owing to the hydrogen bonding between the OH groups and the CO2 bound to the metal center. It was also shown that the OH groups accelerate cleavage of the C−O bond (see the section 2.1 for details). In the photocatalytic CO2 reduction, the TONCO was ca. 30 (with the production of a small amount of H2), although in electrochemical conditions, the selective CO2 reduction continued with little degradation, exhibiting a high faradaic yield of CO (>95%). The reason for the lower durability of the photocatalytic system as compared to the electrochemical one may be related to the hydrogenation of the porphyrin ring through photo-Birch reduction. Bonin and Robert et al. further improved this photocatalytic CO2 reduction by adding 41 as PS. Visible-light irradiation at >420 nm in the presence of TEA as SD led to CO evolution with a higher efficiency and selectivity (ca. 93%).79 The TON was also improved to 140 (entry 39, Table 2), which was significantly higher than in the absence of PS. An organic dye (47, Chart 7) was also used as PS instead of the Ir complex, and the new system exhibited visible-light-driven CO2 reduction, with a high

(sc) CO2, which forms a two-phase solution, may improve the selectivity for CO2 reduction. Irradiation of a water/scCO2 solution containing 3 as the CAT, 39 as the PS, and ascorbic acid as the SD produced CO with a CO/H2 selectivity of 7.1, 7 times higher than that using normal aqueous media with dissolved CO2 (entry 24, Table 2).72 One reason that could explain such improvement is the adsorption of the reduced-state CAT at the interface between the two phases, which induces preferential reaction with CO2 from the scCO2 phase. Metal−organic frameworks (MOFs) containing CoII complexes have also been used as CATs. In one study, a MOF containing benzimidazolate (bIm) as a ligand for CoII complexes, which was an insoluble zeolitic solid with micropores (58, Chart 8), was used.73 Irradiation of a 4:1 CH3CN/H2O (v/v) suspension containing the powder of 58 as the CAT in the presence of 39 as the PS and TEOA as the SD, produced CO and H2 with a ΦCO of 1.48% at λex = 420 nm (entry 25, Table 2). Repeated use of the CAT powders (5 times) gave a total TONCO of 450. This powder also exhibited CO2 capturing ability originating from the bIm moieties in the porous framework, which enhanced the photocatalysis. 2.2.2. Co and Fe Porphyrins and Their Derivatives. As described in section 2.1 (Electrochemical Reduction of CO2), metal porphyrins (Chart 9), especially Fe porphyrins, have been actively investigated as electrochemical CATs for CO2 reduction. Iron porphyrins were first described in photocatalytic systems for CO2 reduction by Neta et al. in 1997.74 Photoexcitation of the ligand-to-metal charge transfer (LMCT) absorption band at 360 nm of a DMF/TEA (5%) solution containing 13 with an axial chloride ligand caused one-electron reduction of the central metal from FeIII to FeII, releasing the chloride ligand. The FeII species could be further reduced to FeI by TEA. Disproportionation of two FeI molecules produces the active catalytic Fe0 species, which reacts to CO2. The product of CO2 reduction was CO with TONCO ∼ 70 and H2 as a minor product (entry 26, Table 2). However, photo-Birch reduction of the porphyrin ring, converting it to the corresponding chlorin structure, followed by further photochemical decomposition rapidly occurs. The cationic Fe porphyrin 59 has also been reported to function as a photocatalyst, generating CO from CO2 by photoexcitation in an aqueous solution (pH 8.8) containing TEA as SD and NaHCO3. Because the efficiencies of the photochemical reduction of the Fe porphyrins were very low, the CO2 photoreduction proceeds with extremely low efficiencies. Addition of p-terphenyl (45, Chart 8) as a PS to photocatalytic systems using Fe porphyrins, which are abbreviated to FeP in the 83


Review

ACS Catalysis Chart 11. Fe Polypyridines and Carbonyl Complexes

Chart 12. CuI Complexes Used as PS in the Photocatalytic Reduction of CO2

selectivity for CO formation. The PS-OERS of 47 exhibits strong reducing power (E0 = −1.58 V vs SCE), high enough to reduce the Fe porphyrins to Fe0 species. However, the durability still needs to be improved (TONCO = ca. 60, entry 40, Table 1). 2.2.3. Fe Polypyridines and Carbonyl Complexes. CH3CN solution containing Fe2+ ions, TEA or TEOA as SD, and 45 as PS was irradiated at λex > 300 nm under a CO2 atmosphere giving CO.80 The generation of CO was half that when 63 was used, and the selectivity for CO generation was much lower ( 400 nm), exhibit good durability, efficiency, and selectivity for CO2 reduction remain rare probably because of the lack of suitable photosensitizers that do not contain a precious metal. Emissive CuI complexes have been attracting attention as PSs in photocatalytic systems for H2 evolution82 and organic synthesis.83 Recently, CuI complexes have also been employed as photosensitizers for CO2 reduction, and 71 (Chart 11) was used as CAT.84,86 Heteroleptic CuI complexes [Cu(dmp)(P)2]2+ (dmp = 2,9-dimethyl-1,10-phenanthroline, P = phosphine ligand: Chart 12), were used as PS because their oxidation ability from the excited state is stronger than those of the corresponding homoleptic complexes Cu(dmp)2+.85 In the presence of BIH (10 mM) as SD, visible-light irradiation of a 5:1 CH3CN/TEOA (v/v) solution containing the 75 (0.25 mM) and 71 (0.05 mM) generated CO with a maximum selectivity of 78%, along with H2 as the minor product.86 When the Cu dimer

75 was used as PS, the efficiency and durability of the CO formation were high, with ΦCO = 6.7% and TONCO > 270 based on 71 used (TONCO > 54 based on 75 used, entries 47−49, Table 2). The catalytic process was initiated by reductive quenching of the excited state of the Cu complex by BIH. The tetradentate ligand, which contained both a dmp moiety and two phosphine moieties, coordinates with CuI ions to produce dimeric structures where each CuI center is surrounded by the dmb moiety, by one phosphine moiety of one tetradentate ligand, and by one phosphate moiety of another tetradentate ligand, as shown in Chart 12. This tetradentate structure of the ligand leads to improved stability for the CuI complexes, especially in their one-electron reduced state. As a consequence, the photocatalytic activity of the system using 75 was higher than that using 76 (ΦCO = 2.6%) A photocatalytic system using the quaterpyridine FeII complex (72, Chart 11, 0.05 mM) as CAT, 39 as PS (0.2 mM), and BIH (0.1 M) as SD in CH3CN/TEOA (4:1 v/v) generated CO with a TONCO of 1879, a selectivity of 97%, and a ΦCO of 8.8% (entries 50 and 51, Table 2).71 The organic photosensitizer 48 (Chart 7) could be used instead of 39. A DMF solution consisting of 48 (2 mM), 72 (0.005 mM), and BIH (0.1 M) converted CO2 to CO selectively with a TONCO of 1365 based on 72 used and a quantum yield ΦCO = 1.1%, which is remarkable for a system using only Fe as a metal (entries 52 and 53, Table 2).71 An interesting case where thermal CATs developed for organic synthesis were applied as efficient CATs for selective photoconversion of CO2 to CO is provided by Fe cyclopentadienone complexes (73 and 74, Chart 11).87 Fe cyclopentadienones are known to be effective for the reduction of ketones via the corresponding hydrido complex as an active intermediate, and the carbonyl moiety of the cyclopentadienone ligand also acts as a proton-donor/acceptor site. Visible-light irradiation of an NMP/TEOA (5:1 v/v) solution containing 73 (0.13 mM, Chart 11) and 43 as PS (1.67 mM, Chart 7) generated CO (TONCO = 84


Review

ACS Catalysis 421) and HCOOH (TONHCOOH < 40). The quantum yield of the CO formation was 68% at λex = 440 nm (entries 54−56, Table 2). In this reaction, TEOA works not only as SD but also as a proton donor.59 The corresponding hydrido complex was detected in the solution during the photocatalytic reaction. 2.2.4. Mn Diimine Complexes. 7 (Chart 3) was recently used as a CAT for photocatalytic CO2 reduction. A DMF and TEOA (4:1 v/v) solution containing the Mn complex (0.05 mM), 40 (0.05 mM) as PS, and 1-benzyl-1,4-dihydronicotinamide (BNAH) as SD was irradiated at λex = 480 nm under a CO2 atmosphere, producing HCOOH after a short induction period with a ΦHCOOH of 5.3%. A TONHCOOH of 149 was obtained after 12 h irradiation (entry 57, Table 2).88 Such performance is comparable to the photocatalytic system using fac-Re(dmb) (CO)3Cl, under similar reaction conditions (ΦCO = 6.2%).89 Note that the reaction product was HCOOH and not CO, although CO was the main product in the electrochemical CO2 reduction using the same Mn catalyst.20,21 During the induction period of HCOOH formation, a Mn dimer ([M(bpy)(CO)3]2) with a Mn−Mn bond20b was produced via one-electron reduction of the original Mn complex. During this period, CO was the main product. This photocatalytic reaction was initiated by reductive quenching of the 3MLCT excited state of 40 to give the corresponding PS-OERS, a reductant strong enough (E1/2red = −1.82 V vs Ag/AgNO3) to transfer one electron to 7 (Ep = −1.65 V vs Ag/AgNO3). As described above, the reduced Mn complex was converted to the Mn dimer. Since the reduction potential of the Mn dimer is comparable (Ep = −1.84 V vs Ag/AgNO3) with that of 40, reduction of the Mn dimer might proceed giving [Mn(bpy)(CO)3]−, along a slow process. Use of 39 (E1/2red = −1.72 V vs Ag/AgNO3), instead of 40 did not affect the photocatalytic activity of the system (TONHCOOH = 157 after 12 h irradiation). After the Mn dimer was no longer detected in the reaction solution, production of HCOOH proceeded more rapidly, and formation of CO just stopped. This suggests that the Mn dimer does not lead to HCOOH formation but might produce CO. A similar photocatalytic system consisting of 77 (Chart 13), which is photochemically more stable than 7,21e 40 as the PS, and

reduction in a DMF/TEOA (4:1 v/v) mixed solution in the presence of 40 as PS and BNAH as SD.91 This system produced HCOOH as reduction product with a high selectivity of 96% and a ΦHCOOH of 13.8% at λex = 470 nm. Because the TONHCOOH (110 for 18 h irradiation) was larger than that of the corresponding homogeneous systems (70 at 18 h, ΦHCOOH = 9.6% with 7 as CAT) under similar reaction conditions (entries 62 and 63, Table 2), the MOF was suggested to protect the Mn complexes from decomposition (by avoiding the Mn dimer formation due to the isolation of each Mn site in the MOF). Because this system was heterogeneous, the Mn-MOF could be recovered by decantation from the dispersed solution after the photocatalytic reaction, and could be reused as CAT in another experiment. The total TONHCOOH was 170 after 4 recycling steps. The photocatalytic activity of the system gradually decreased because of partial decomposition of the fac-Mn(CO)3 structure, even though the Mn content itself did not decrease in the MOF. In all the systems using Mn complexes described above, Ru complexes were used as PSs. Very recently, Bian et al. reported a photocatalytic system using 49 (0.5 mM) as PS with 79 (2 mM, Chart 13) as CAT and TEA as SD.92 In an aqueous CH3CN solution (CH3CN/H2O (20:1 v/v)), CO was mainly produced (TONCO = 119 based on the Mn complex used) with also some HCOOH (TONHCOOH = 19, entry 64, Table 2). In summary, various catalysts using abundant metals (i.e., Co, Ni, Fe, and Mn) have been developed. Despite some encouraging and even remarkable examples, notably those discovered recently with Mn and Fe metal complexes, photosensitizers using only abundant elements remain limited, and the reaction mechanisms of the photocatalytic reactions, especially those involving the above-mentioned Fe and Mn complexes, have not been fully elucidated. Most of the proposed mechanisms were investigated on the basis of data primarily issued from electrochemical experiments; however, the product distribution is sometimes very different between photochemical and electrochemical systems, even when the same CAT is used. Future direction also includes the design of new, efficient hybrid systems by association of a molecular catalyst to a semiconductive surface that could act as an efficient photosensitizer so as to generate the active molecular catalytic species.93 Because these fields have become increasingly important in recent times, new photocatalytic systems will be developed, and hopefully, the above-mentioned difficulties and challenges will find solutions in the near future.

Chart 13. Mn(I) Complexes

3. CONCLUSIONS Over the last 5 years or so, efficient molecular catalysts based on earth-abundant metals for the two electrons electrochemical and photochemical reduction of CO2 have appeared, showing high selectivity, good turnover number, and fast catalytic rate. Most of them function in aprotic solvent. Catalysts working efficiently in water and over long period of time (several days) are still rare. Even if mechanistic studies have made important progress, more studies including spectroscopic detection and analysis of intermediates are needed so as to not only design better catalysts but also to better understand the intrinsic structural factors that may lead to selectivity control during catalysis. Benchmark of the catalyst intrinsic properties based on the establishment of catalytic Tafel plot relating the turnover frequency to the overpotential of the reaction has been proposed for the electrochemical CO2-to-CO conversion. More generally, benchmark of catalysts independently from a special set of

BNAH as SD, also produced HCOOH as the main product with a ΦHCOOH of 3.9% and a TONHCOOH of 130 in a mixed DMF/ TEOA (4:1 v/v) solution.90 The product distribution is affected by the nature of the solvent, with HCOOH/CO being equal to 18 in DMF/TEOA and 0.42 in CH3CN/TEOA (entries 58−61, Table 2). The two-electron reduction products, i.e., [Mn(bpy)(CO)3]− produced by disproportionation of the OERS [Mn(bpy) (CO)3(CN)]•−,21f was proposed as an active species for the CO2 reduction,21e,f and the Mn dimer was not observed during the photocatalytic reaction. A MOF consisting of 7 as unit and Zn(IV) bridges (78, Chart 13) was used as CAT in a photocatalytic system for CO2 85


Review

ACS Catalysis

(16) Che, C.-M.; Mak, S.-T.; Lee, W.-O.; Fung, K.-W.; Mak, T. C. W. J. Chem. Soc., Dalton Trans. 1988, 2153−2159. (17) Zhang, M.; El-Roz, M.; Frei, H.; Mendoza-Cortes, J. L.; HeadGordon, M.; Lacy, D. C.; Peters, J. C. J. Phys. Chem. C 2015, 119, 4645− 4654. (18) Sheng, H.; Frei, H. J. Am. Chem. Soc. 2016, 138, 9959−9967. (19) Chapovetsky, A.; Do, T. A.; Haiges, R.; Takase, M. K.; Marinescu, S. C. J. Am. Chem. Soc. 2016, 138, 5765−5768. (20) (a) Bourrez, M.; Molton, F.; Chardon-Noblat, S.; Deronzier, A. Angew. Chem., Int. Ed. 2011, 50, 9903−9906. (b) Bourrez, M.; Orio, M.; Molton, F.; Vezin, H.; Duboc, C.; Deronzier, A.; Chardon-Noblat, S. Angew. Chem., Int. Ed. 2014, 53, 240−243. (21) (a) Smieja, J. M.; Sampson, M. D.; Grice, K. A.; Benson, E. E.; Froehlich, J. D.; Kubiak, C. P. Inorg. Chem. 2013, 52, 2484−2491. (b) Sampson, M. D.; Nguyen, A. D.; Grice, K. A.; Moore, C. E.; Rheingold, A. L.; Kubiak, C. P. J. Am. Chem. Soc. 2014, 136, 5460−5471. (c) Zeng, Q.; Tory, J.; Hartl, F. Organometallics 2014, 33, 5002−5008. (d) Franco, F.; Cometto, C.; Ferrero Vallana, F.; Sordello, F.; Priola, E.; Minero, C.; Nervi, C.; Gobetto, R. Chem. Commun. 2014, 50, 14670− 14673. (e) Agarwal, J.; Stanton, C. J., III; Shaw, T. W.; Vandezande, J. E.; Majetich, G. F.; Bocarsly, A. B.; Schaefer, H. F., III Dalton Trans. 2015, 44, 2122−2131. (f) Machan, C. W.; Stanton, C. J.; Vandezande, J. E.; Majetich, G. F.; Schaefer, H. F.; Kubiak, C. P.; Agarwal, J. Inorg. Chem. 2015, 54, 8849−8856. (g) Mukhopadhyay, T. K.; MacLean, N. L.; Gan, L.; Ashley, D. C.; Groy, T. L.; Baik, M.-H.; Jones, A. K.; Trovitch, R. J. Inorg. Chem. 2015, 54, 4475−4482. (22) Sampson, M. D.; Kubiak, C. P. J. Am. Chem. Soc. 2016, 138, 1386− 1393. (23) Agarwal, J.; Shaw, T. W.; Schaefer, H. F., III; Bocarsly, A. B. Inorg. Chem. 2015, 54, 5285−5294. (24) Rao, G. K.; Pell, W.; Korobkov, I.; Richeson, D. Chem. Commun. 2016, 52, 8010−8013. (25) (a) Costentin, C.; Drouet, S.; Robert, M.; Savéant, J.-M. Science 2012, 338, 90−94. (b) Costentin, C.; Robert, M.; Savéant, J.-M. Chem. Soc. Rev. 2013, 42, 2423−2436. (c) Costentin, C.; Robert, M.; Savéant, J.-M. Acc. Chem. Res. 2015, 48, 2996−3006 and references therein.. (26) Zahran, Z. N.; Mohamed, E. A.; Naruta, Y. Sci. Rep. 2016, 6, 24533. (27) Costentin, C.; Drouet, S.; Passard, G; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2013, 135, 9023−9031. (28) Mondal, B.; Rana, A.; Sen, P.; Dey, A. J. Am. Chem. Soc. 2015, 137, 11214−11217. (29) (a) Costentin, C.; Passard, G.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2014, 136, 11821−11829. (b) Costentin, C.; Passard, G.; Robert, M.; Savéant, J.-M. Proc. Natl. Acad. Sci. U. S. A. 2014, 111, 14990−14994. (30) Wong, K.-Y.; Chung, W.-H.; Lau, C.-P. J. Electroanal. Chem. 1998, 453, 161−169. (31) Raebiger, J. W.; Turner, J. W.; Noll, B. C.; Curtis, C. J.; Miedaner, A.; Cox, B.; DuBois, D. L. Organometallics 2006, 25, 3345. (32) Chen, Z.; Chen, C.; Weinberg, D. R.; Kang, P.; Concepcion, J. J.; Harrison, D. P.; Brookhart, M. S.; Meyer, T. J. Chem. Commun. 2011, 47, 12607−12609. (33) Azcarate, I.; Costentin, C.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc., DOI: 10.1021/jacs.6b07014. (34) Costentin, C.; Robert, M.; Savéant, J.-M.; Tatin, A. Proc. Natl. Acad. Sci. U. S. A. 2015, 112, 6882−6886. (35) Schneider, J.; Jia, H.; Kobiro, K.; Cabelli, D. E.; Muckerman, J. T.; Fujita, E. Energy Environ. Sci. 2012, 5, 9502−9510. (36) (a) Neri, G.; Aldous, I. M.; Walsh, J. J.; Hardwick, L. J.; Cowan, A. J. Chemical Science 2016, 7, 1521−1527. (b) Neri, G.; Walsh, J. J.; Wilson, C.; Reynal, A.; Lim, J. Y. C.; Li, X.; White, A. J. P.; Long, N. J.; Durrant, J. R.; Cowan, A. J. Phys. Chem. Chem. Phys. 2015, 17, 1562− 1566. (c) Walsh, J. J.; Neri, G.; Smith, C. L.; Cowan, A. J. Chem. Commun. 2014, 50, 12698−12701. (37) (a) Meshitsuka, S.; Ichikawa, M.; Tamaru, K. J. Chem. Soc., Chem. Commun. 1974, 158−159. (b) Furuya, N.; Matsui, K. J. Electroanal. Chem. Interfacial Electrochem. 1989, 271, 181−191. (c) Atoguchi, T.; Aramata, A.; Kazusaka, A.; Enyo, M. J. Chem. Soc., Chem. Commun. 1991,

experimental conditions is an important challenge to advance the field. Catalytic CO2 reduction could also be performed in photochemical conditions, with visible light, appropriate sensitizers and sacrificial electron donors, but much remains to be done to get more durable and efficient systems. Until now, almost no cheap molecular catalyst is able to reduce carbon dioxide beyond the two electron reduction products, with the noticeable recent exceptions of a Cu-based porphyrin94 and carbene-supported Ni complexes.95 By taking advantage of the selectivity offered by molecular catalysts, association to conductive or semiconductive materials based on earth-abundant elements (cheap metal oxides or various forms of carbon derivatives like, e.g., graphene, graphene oxides or carbon nitride) appears as a promising way for getting new photocatalytic or electro-photocatalytic systems functioning in aqueous conditions. It will likely open new pathway for catalytic systems able to reduce efficiently and selectively CO2 with more than two electrons so as to meet the promise of making fuels from the carbon dioxide.

■

AUTHOR INFORMATION

Corresponding Authors

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

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS Support from JSPS (Japan Society for the Promotion of Science, short-term research fellowship S15138) to M. R., as well as from CREST, JST to H. T. and O. I. is gratefully acknowledged. PhD fellowship to C. C. from Université Sorbonne Paris Cité is also warmly acknowledged.

■

REFERENCES

(1) (a) Robert, M. ACS Energy Lett. 2016, 1, 281−282. (b) Tatin, A.; Bonin, J.; Robert, M. ACS Energy Lett. 2016, 1, 1062−1064. (2) (a) Savéant, J.-M. Chem. Rev. 2008, 108, 2348−2378. (b) Benson, E. E.; Kubiak, C. P.; Sathrum, A. J.; Smieja, J. M. Chem. Soc. Rev. 2009, 38, 89−99. (3) Lamy, E.; Nadjo, L.; Savéant, J.-M. J. Electroanal. Chem. Interfacial Electrochem. 1977, 78, 403−407. (4) Gennaro, A.; Isse, A. A.; Severin, M.-G.; Vianello, E.; Bhugun, I.; Savéant, J.-M. J. Chem. Soc., Faraday Trans. 1996, 92, 3963−3968. (5) Gennaro, A.; Isse, A. A.; Savéant, J.-M.; Severin, M.-G.; Vianello, E. J. Am. Chem. Soc. 1996, 118, 7190−7196. (6) Rudolph, M.; Dautz, S.; Jäger, E.-G. J. Am. Chem. Soc. 2000, 122, 10821−10830. (7) Angamuthu, R.; Byers, P.; Lutz, M.; Spek, A. L.; Bouwman, E. Science 2010, 327, 313−315. (8) Fisher, B.; Eisenberg, R. J. Am. Chem. Soc. 1980, 102, 7361−7363. (9) (a) Beley, M.; Collin, J.-P.; Ruppert, R.; Sauvage, J.-P. J. Chem. Soc., Chem. Commun. 1984, 2, 1315−1316. (b) Beley, M.; Collin, J. P.; Ruppert, R.; Sauvage, J. P. J. Am. Chem. Soc. 1986, 108, 7461−7467. (10) Collin, J.-P.; Jouaiti, A.; Sauvage, J.-P. Inorg. Chem. 1988, 27, 1986−1990. (11) Froehlich, J. D.; Kubiak, C. P. Inorg. Chem. 2012, 51, 3932−3934. (12) Froehlich, J. D.; Kubiak, C. P. J. Am. Chem. Soc. 2015, 137, 3565− 3573. (13) Song, J.; Klein, E. L.; Neese, F.; Ye, S. Inorg. Chem. 2014, 53, 7500−7507. (14) Thoi, V. S.; Kornienko, N.; Margarit, C. G.; Yang, P.; Chang, C. J. J. Am. Chem. Soc. 2013, 135, 14413−14424. (15) Tinnemans, A. H. A.; Koster, T. P. M.; Thewissen, D. H. M. W.; Mackor, A. Recl. Trav. Chim. Pays-Bas 1984, 103, 288−295. 86


Review

ACS Catalysis 156−157. (d) Sonoyama, N.; Kirii, M.; Sakata, T. Electrochem. Commun. 1999, 1, 213−216. (e) Shen, J.; Kortlever, R.; Kas, R.; Birdja, Y. Y.; DiazMorales, O.; Kwon, Y.; Ledezma-Yanez, I.; Schouten, K. J. P.; Mul, G.; Koper, M. T. M. Nat. Commun. 2015, 6, 8177. (38) Aoi, S.; Mase, K.; Ohkubo, K.; Fukuzumi, S. Chem. Commun. 2015, 51, 10226−10228. (39) Yoshida, T.; Iida, T.; Shirasagi, T.; Lin, R.-J.; Kaneko, M. J. Electroanal. Chem. 1993, 344, 355−362. (40) Maurin, A.; Robert, M. J. Am. Chem. Soc. 2016, 138, 2492−2495. (41) Bélanger, D.; Pinson, J. Chem. Soc. Rev. 2011, 40, 3995−4048. (42) Yao, S. A.; Ruther, R. E.; Zhang, L.; Franking, R. A.; Hamers, R. J.; Berry, J. F. J. Am. Chem. Soc. 2012, 134, 15632−15635. (43) Elgrishi, N.; Griveau, S.; Chambers, M. B.; Bedioui, F.; Fontecave, M. Chem. Commun. 2015, 51, 2995−2998. (44) Maurin, A.; Robert, M. Chem. Commun. 2016, 52, 12084−12087. (45) Rosser, T. E.; Windle, C. D.; Reisner, E. Angew. Chem., Int. Ed. 2016, 55, 7388−7392. (46) Ramos Sende, J. A.; Arana, C. R.; Hernandez, L.; Potts, K. T.; Keshevarz-K, M.; Abruna, H. D. Inorg. Chem. 1995, 34, 3339−3348. (47) (a) Sun, C.; Gobetto, R.; Nervi, C. New J. Chem. 2016, 40, 5656− 5661. (b) Hu, X.-M.; Salmi, Z.; Lillethorup, M.; Pedersen, E. B.; Robert, M.; Pedersen, S. U.; Skrydstrup, T.; Daasbjerg, K. Chem. Commun. 2016, 52, 5864−5867. (48) Kramer, W. W.; McCrory, C. C. L. Chem. Sci. 2016, 7, 2506−2515 and references therein.. (49) (a) Kornienko, N.; Zhao, Y.; Kley, C. S.; Zhu, C.; Kim, D.; Lin, S.; Chang, C. J.; Yaghi, O. M.; Yang, P. J. Am. Chem. Soc. 2015, 137, 14129− 14135. (b) Lin, S.; Diercks, C. S.; Zhang, Y.-B.; Kornienko, N.; Nichols, E. M.; Zhao, Y.; Paris, A. R.; Kim, D.; Yang, P.; Yaghi, O. M.; Chang, C. J. Science 2015, 349, 1208−1213. (50) Hod, I.; Sampson, M. D.; Deria, P.; Kubiak, C. P.; Farha, O.; Hupp, J. T. ACS Catal. 2015, 5, 6302−6309. (51) Morlanés, N.; Takanabe, K.; Rodionov, V. ACS Catal. 2016, 6, 3092−3095. (52) Tatin, A.; Comminges, C.; Kokoh, B.; Costentin, C.; Robert, M.; Savéant, J.-M. Proc. Natl. Acad. Sci. U. S. A. 2016, 113, 5526−5529. (53) (a) Zall, C. M.; Linehan, J. C.; Appel, A. M. ACS Catal. 2015, 5, 5301−5305. (b) Badiei, Y. M.; Wang, W.-H.; Hull, J. F.; Szalda, D. J.; Muckerman, J. T.; Himeda, Y.; Fujita, E. Inorg. Chem. 2013, 52, 12576− 12586. (c) Zhang, Y.; MacIntosh, A. D.; Wong, J. L.; Bielinski, E. A.; Williard, P. G.; Mercado, B. Q.; Hazari, N.; Bernskoetter, W. H. Chem. Sci. 2015, 6, 4291−4299. (d) Rivada-Wheelaghan, O.; Dauth, A.; Leitus, G.; Diskin-Posner, Y.; Milstein, D. Inorg. Chem. 2015, 54, 4526−4538. (e) Bertini, F.; Gorgas, N.; Stöger, B.; Peruzzini, M.; Veiros, L. F.; Kirchner, K.; Gonsalvi, L. ACS Catal. 2016, 6, 2889−2893. (54) Bhugun, I.; Lexa, D.; Savéant, J.-M. J. Am. Chem. Soc. 1996, 118, 1769−1776. (55) Pun, S.-N.; Chung, W.-H.; Lam, K.-M.; Guo, P.; Chan, P.-H.; Wong, K.-Y.; Che, C.-M.; Chen, T.-Y.; Peng, S.-M. J. Chem. Soc., Dalton Trans. 2002, 575−583. (56) Chen, L.; Guo, Z.; Wei, X.-G.; Gallenkamp, C.; Bonin, J.; Anxolabéhère-Mallart, E.; Lau, K.-C.; Lau, T.-C.; Robert, M. J. Am. Chem. Soc. 2015, 137, 10918−10921. (57) (a) Rail, M. D.; Berben, L. J. Am. Chem. Soc. 2011, 133, 18577− 18579. (b) Taheri, A.; Thompson, E. J.; Fettinger, J. C.; Berben, L. ACS Catal. 2015, 5, 7140−7151. (c) Taheri, A.; Berben, L. Inorg. Chem. 2016, 55, 378−385. (d) Taheri, A.; Berben, L. Chem. Commun. 2016, 52, 1768−1777. (58) Juris, A.; Balzani, V.; Barigelletti, F.; Campagna, S.; Belser, P.; von Zelewsky, A. Coord. Chem. Rev. 1988, 84, 85−277. (59) Yamazaki, Y.; Takeda, H.; Ishitani, O. J. Photochem. Photobiol., C 2015, 25, 106−137. (60) (a) Grant, J. L.; Goswami, K.; Spreer, L. O.; Otvos, J. W.; Calvin, M. J. Chem. Soc., Dalton Trans. 1987, 2105−2109. (b) Craig, C. A.; Spreer, L. O.; Otvos, J. W.; Calvin, M. J. Phys. Chem. 1990, 94, 7957− 7960. (61) Mochizuki, K.; Manaka, S.; Takeda, I.; Kondo, T. Inorg. Chem. 1996, 35, 5132−5136.

(62) (a) Lehn, J. M.; Ziessel, R. Proc. Natl. Acad. Sci. U. S. A. 1982, 79, 701−704. (b) Hawecker, J.; Lehn, J.-M.; Ziessel, R. J. Chem. Soc., Chem. Commun. 1983, 536−538. (c) Ziessel, R.; Hawecker, J.; Lehn, J.-M. Helv. Chim. Acta 1986, 69, 1065−1084. (63) (a) Matsuoka, S.; Yamamoto, K.; Pac, C.; Yanagida, S. Chem. Lett. 1991, 20, 2099−2100. (b) Matsuoka, S.; Yamamoto, K.; Ogata, T.; Kusaba, M.; Nakashima, N.; Fujita, E.; Yanagida, S. J. Am. Chem. Soc. 1993, 115, 601−609. (64) Ogata, T.; Yamamoto, Y.; Wada, Y.; Murakoshi, K.; Kusaba, M.; Nakashima, N.; Ishida, A.; Takamuku, S.; Yanagida, S. J. Phys. Chem. 1995, 99, 11916−11922. (65) (a) Fujita, E.; Szalda, D. J.; Creutz, C.; Sutin, N. J. Am. Chem. Soc. 1988, 110, 4870−4871. (b) Fujita, E.; Creutz, C.; Sutin, N.; Brunschwig, B. S. Inorg. Chem. 1993, 32, 2657−2662. (c) Ogata, T.; Yanagida, S.; Brunschwig, B. S.; Fujita, E. J. Am. Chem. Soc. 1995, 117, 6708−6716. (d) Ogata, T.; Yanagida, S.; Brunschwig, B. S.; Fujita, E. Energy Convers. Manage. 1995, 36, 669−672. (e) Fujita, E.; Furenlid, L. R.; Renner, M. W. J. Am. Chem. Soc. 1997, 119, 4549−4550. (66) Chan, S. L.-F.; Lam, T. L.; Yang, C.; Yan, S.-C.; Cheng, N. M. Chem. Commun. 2015, 51, 7799−7801. (67) Yang, C.; Mehmood, F.; Lam, T. L.; Chan, S. L.-F.; Wu, Y.; Yeung, C.-S.; Guan, X.; Li, K.; Chung, C. Y.-S.; Zhou, C.; Zou, T.; Che, C.-M. Chem. Sci. 2016, 7, 3123−3136. (68) Wang, F.; Cao, B.; To, W.-P.; Tse, C.-W.; Li, K.; Chang, X.-Y.; Zang, C.; Chan, S. L.-F.; Che, C.-M. Catal. Sci. Technol. 2016, 6, 7408− 7420. (69) Lam, K.-M.; Wong, K.-Y.; Yang, S.-M.; Che, C.-M. J. Chem. Soc., Dalton Trans. 1995, 7, 1103−1107. (70) Tamaki, Y.; Koike, K.; Morimoto, T.; Ishitani, O. J. Catal. 2013, 304, 22−28. (71) Guo, Z.; Cheng, S.; Cometto, C.; Anxolabéhère-Mallart, E.; Ng, S.-M.; Ko, C.-C.; Liu, G.; Chen, L.; Robert, M.; Lau, T.-C. J. Am. Chem. Soc. 2016, 138, 9413−9416. (72) Méndez, M. A.; Voyame, P.; Girault, H. H. Angew. Chem., Int. Ed. 2011, 50, 7391−7394. (73) Wang, S.; Yao, W.; Lin, J.; Ding, Z.; Wang, X. Angew. Chem., Int. Ed. 2014, 53, 1034−1038. (74) Grodkowski, J.; Behar, D.; Neta, P.; Hambright, P. J. Phys. Chem. A 1997, 101, 248−254. (75) Dhanasekaran, T.; Grodkowski, J.; Neta, P.; Hambright, P.; Fujita, E. J. Phys. Chem. A 1999, 103, 7742−7748. (76) Grodkowski, J.; Dhanasekaran, T.; Neta, P.; Hambright, P.; Brunschwig, B. S.; Shinozaki, K.; Fujita, E. J. Phys. Chem. A 2000, 104, 11332−11339. (77) Grodkowski, J.; Neta, P.; Fujita, E.; Mahammed, A.; Simkhovich, L.; Gross, Z. J. Phys. Chem. A 2002, 106, 4772−4778. (78) Bonin, J.; Chaussemier, M.; Robert, M.; Routier, M. ChemCatChem 2014, 6, 3200−3207. (79) Bonin, J.; Robert, M.; Routier, M. J. Am. Chem. Soc. 2014, 136, 16768−16771. (80) Grodkowski, J.; Neta, P. J. Phys. Chem. A 2000, 104, 4475−4479. (81) Alsabeh, P. G.; Rosas-Hernández, A.; Barsch, E.; Junge, H.; Ludwig, R.; Beller, M. Catal. Sci. Technol. 2016, 6, 3623−3630. (82) (a) Mejía, E.; Luo, S.-P.; Karnahl, M.; Friedrich, A.; Tschierlei, S.; Surkus, A.-E.; Junge, H.; Gladiali, S.; Lochbrunner, S.; Beller, M. Chem. Eur. J. 2013, 19, 15972−15978. (b) Karnahl, M.; Mejía, E.; Rockstroh, N.; Tschierlei, S.; Luo, S.-P.; Grabow, K.; Kruth, A.; Brüser, V.; Junge, H.; Lochbrunner, S.; Beller, M. ChemCatChem 2014, 6, 82−86. (c) Luo, S.; Mejía, E.; Friedrich, A.; Pazidis, A.; Junge, H.; Surkus, A.-E.; Jackstell, R.; Denurra, S.; Gladiali, S.; Lochbrunner, S.; Beller, M. Angew. Chem., Int. Ed. 2013, 52, 419−423. (d) Tschierlei, S.; Karnahl, M.; Rockstroh, N.; Junge, H.; Beller, M.; Lochbrunner, S. ChemPhysChem 2014, 15, 3709−3713. (e) Khnayzer, R. S.; McCusker, C. E.; Olaiya, B. S.; Castellano, F. N. J. Am. Chem. Soc. 2013, 135, 14068−14070. (f) Lazorski, M. S.; Castellano, F. N. Polyhedron 2014, 82, 57−70. (83) Knorn, M.; Rawner, T.; Czerwieniec, R.; Reiser, O. ACS Catal. 2015, 5, 5186−5193. (84) König, E.; Ritter, G.; Madeja, K.; Kobetić, R.; Gembarovski, D.; Baranović, G.; Gabelica, V. J. Inorg. Nucl. Chem. 1981, 43, 2273−2280. 87


Review

ACS Catalysis (85) Sakaki, S.; Mizutani, H.; Kase, Y.; Inokuchi, K.; Arai, T.; Hamada, T. J. Chem. Soc., Dalton Trans. 1996, 1909−1914. (86) Takeda, H.; Ohashi, K.; Sekine, A.; Ishitani, O. J. Am. Chem. Soc. 2016, 138, 4354−4357. (87) Rosas, A.; Alsabeh, P. G.; Barsch, E.; Junge, H.; Ludwig, R.; Beller, M. Chem. Commun. 2016, 52, 8393−8396. (88) Takeda, H.; Koizumi, H.; Okamoto, K.; Ishitani, O. Chem. Commun. 2014, 50, 1491−1493. (89) Gholamkhass, B.; Mametsuka, H.; Koike, K.; Tanabe, T.; Furue, M.; Ishitani, O. Inorg. Chem. 2005, 44, 2326−2336. (90) Cheung, P. L.; Machan, C. W.; Malkhasian, A. Y. S.; Agarwal, J.; Kubiak, C. P. Inorg. Chem. 2016, 55, 3192−3198. (91) Fei, H.; Sampson, M. D.; Lee, Y.; Kubiak, C. P.; Cohen, S. M. Inorg. Chem. 2015, 54, 6821−6828. (92) Zhang, J.-X.; Hu, C.-Y.; Wang, W.; Wang, H.; Bian, Z.-Y. Appl. Catal., A 2016, 522, 145−151. (93) Jin, T.; Liu, C.; Li, G. Chem. Commun. 2014, 50, 6221−6224. (94) Weng, Z.; Jiang, J.; Wu, Y.; Wu, Z.; Guo, X.; Materna, K. L.; Liu, L.; Batista, V. S.; Brudvig, G. W.; Wang, H. J. Am. Chem. Soc. 2016, 138, 8076−8079. (95) Lu, Z.; Williams, T. J. ACS Catal. 2016, 6, 6670−6673.

88


Electrons, Photons, Protons and Earth-Abundant Metal Complexes for

Recommend Documents