Molecular Catalysts for Water Oxidation - ACS Publications - American

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Molecular Catalysts for Water Oxidation James D. Blakemore,† Robert H. Crabtree,* and Gary W. Brudvig* Department of Chemistry and Energy Sciences Institute, Yale University, P.O. Box 208107, New Haven, Connecticut 06520-8107, United States these intermediates, resulting in a lower kinetic barrier and, consequently, faster rates of oxygen production. Here, we review the field of molecular catalysts for water oxidation. Inspired by the oxygen-evolving complex in Photosystem II, many researchers have endeavored to develop well-defined catalysts for water oxidation that operate in homogeneous solution. Such efforts have important consequences for understanding natural water oxidation and also CONTENTS possible applications in artificial photosynthesis, envisioned as a 1. Introduction A future source of clean energy for society. Thus, we briefly 1.1. Role of Water Oxidation in Energy Storage A review the motivations for studying water oxidation, as well as 1.2. Natural Photosynthetic Water Oxidation B the basic details of water oxidation in Photosystem II. We then 1.3. Water Electrolysis C turn our attention to artificial, molecular catalysts for water 1.4. Heterogeneous Oxide-Based Catalysts C oxidation. 2. Molecular Catalysts for Water Oxidation D The catalysts described in this review are organized by 2.1. Manganese Catalysts D elements, focusing first on manganese, ruthenium, and iridium. 2.2. Ruthenium Catalysts G Each of the elements in this privileged series shows good 2.3. Iridium Catalysts L activity in a variety of compounds and ligand environments for 2.4. Iron Catalysts R water oxidation, suggesting that the diagonal relationship 2.5. Cobalt Catalysts T between these elements in Groups 7, 8, and 9 is key to their 2.6. Other Catalysts T activity. We then review cobalt, iron, and other catalysts. 3. Outlook and Conclusions U Emphasis in this review is placed on synthetic, kinetic, and Author Information U electrochemical data, as they pertain to the mechanism of Corresponding Authors U action of each family of catalysts. Special attention is also paid Present Address U to distinguishing homogeneous, molecular catalysis from Notes U heterogeneous catalysis of metal oxide solids or particles that Biographies U can form in situ during studies of molecular precatalysts. Acknowledgments References

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1.1. Role of Water Oxidation in Energy Storage

Geopolitical concerns and the increasing threats of anthropogenic climate change encourage movement toward more stable, renewable, and environmentally benign sources of energy. Among the many renewable energy sources, the sun represents the most abundant source of renewable energy. The current world demand for energy is approximately 15 terawatts (15 × 1012 J s−1), and the sun can provide more than 50 terawatts of energy in the form of solar light irradiance.6 Since other forms of renewable energy are limited to total amounts less than either our current or our foreseeable future societal needs, the sun represents an ultimate source of energy. Existing photovoltaic technologies can convert light energy from the sun into an electrical potential. However, at best the sun shines only up to one-half the hours of each day, and the light is unevenly distributed across the Earth’s surface by shadowing. Additionally, the light energy is spread over the entire surface of the Earth, making it fairly dilute on a power per unit area basis. Thus, due to the intermittent and diffuse

1. INTRODUCTION The development of technologies for the production of chemical fuels or useful compounds with renewable energy (e.g., sunlight) and inexpensive feedstocks relies on an abundant supply of protons and electrons to form the reduced products.1−5 In natural oxygenic photosynthesis, a possible blueprint for artificial photosynthesis, the only suitably abundant source for the needed protons and electrons is water. Oxidation of water liberates protons and electrons and gives off oxygen gas 2H 2O → O2 + 4H+ + 4e−

(1)

However, the water-oxidation reaction is both thermodynamically and kinetically demanding, resulting in slow kinetics without the use of a catalyst. At an electrode, the slow uncatalyzed kinetics are observed in the form of high overpotentials. In both the chemical and the electrochemical cases, the sluggish kinetics are likely due to the many intermediates required to accomplish the complete 4H+/4e− oxidation of water to dioxygen. An effective catalyst stabilizes © XXXX American Chemical Society

Special Issue: Solar Energy Conversion Received: February 26, 2015

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nature of the sun’s energy, methods for concentrating and storing this energy are required. Water splitting, in which water is split into oxygen and hydrogen that can be recombined at the point of use in a fuel cell, represents one solution to the energy storage problem. 1.2. Natural Photosynthetic Water Oxidation7

In natural photosynthesis, light energy is used to oxidize water to dioxygen in the enzyme−cofactor complex Photosystem II (PSII).8−11 In oxidizing water to oxygen, four protons and four electrons are liberated; the protons liberated by PSII contribute to the membrane-spanning proton-motive force, and the electrons are used to reduce plastoquinones to plastoquinols. The primary electron-transfer pathway is well established (shown in Figure 1).12−14 The light-induced charge separation

Figure 2. Oxygen-evolving complex of Photosystem II, the site of catalytic water oxidation. The structure is from the recent 1.9 Å crystal structure of PSII obtained by Shen and co-workers.11,18 Bond lengths and atom labels are given according to ref 18. Colors: manganese, purple; calcium, yellow; bridging oxo groups, red; bound water molecules, orange. The site of O−O bond formation has been proposed to be between a manganese(IV)−oxyl species on Mn4 and a water W3 bound to calcium or between a manganese(IV)−oxyl species on Mn1 and O5. Reprinted with permission from ref 18. Copyright 2011 Elsevier B. V.

proposed the now well-accepted S-state cycle to explain this behavior.23 The P680+ cation radical is the most oxidizing redox cofactor known in biology, with an estimated reduction potential of 1.25 V vs NHE.24 This potential is sufficiently strong to drive the oxidation of water at pH 7 (E°′ = 820 mV) (where E°′ is the pH-adjusted thermodynamic potential for the reaction).25 However, oxygen evolution is observed with PSII at pH values as low as 5. Since E°′ = 935 mV at pH 5, this implies that the OEC operates at an “overpotential” below 300 mV.26,27 Oxygen production, which involves overcoming the known high kinetic barrier for this reaction, can proceed at a very fast rate with this overpotential. Interestingly, the rate-determining step in oxygen evolution by PSII is the 2e−/2H+ reduction of QB and its subsequent diffusion from its binding site into the membrane.28 Thus, it is the acceptor side of the electron-transport chain that dictates the turnover frequency of PSII, not any of the faster steps measured for oxidations of the OEC or release of dioxygen.20 These features suggest that the OEC is a highly optimized catalyst for water oxidation. Indeed, under conditions where electron acceptors and all supporting media are present in sufficient quantities, turnover frequencies in excess of 100 mol of O2 (mol OEC)−1 s−1 are observed.29 However, as pointed out by Dau and co-workers, the maximal turnover frequency based on the donor side kinetics could be as high as 400 mol of O2 (mol OEC)−1 s−1.26 Interestingly, the OEC and related cofactors in the D1 protein subunit of PSII are replaced as often as twice per hour.26 This seems consistent with the highly oxidizing chemistry taking place and the related consequences of side reactions. The need to constantly reassemble the OEC and related machinery to avoid the accumulation of photo-

Figure 1. Electron-transfer pathway in Photosystem II.12,13 The primary electron-transfer pathway is marked with the solid arrows, where electrons pass from water to P680 via the oxygen-evolving complex (OEC) and YZ. Secondary electron transfer14 that involves cytochrome b559 and accessory chlorophylls/carotenoids (a cyclic pathway) is also included with the dashed line.

from the reaction center chlorophylls (P680) to an adjacent pheophytin (PheoA) is followed by electron transfer to a PSIIbound plastoquinone (QA). This P680+/QA− state is a transiently stable charge-separated state; QA then passes the electron off to a second quinone (QB) that, upon two-electron/ two-proton reduction, diffuses away to continue the electrontransport chain. The electrons from water pass through the cytochrome b6 f complex and Photosystem I. Eventually they are used to reduce CO2 in the Calvin−Benson cycle.15 Biological water oxidation is catalyzed by the oxygen-evolving complex (OEC) in PSII,16 a tetramanganese cluster17 that also contains a required calcium atom (Figure 2).18,19 The demanding requirements of this reaction contribute to the observation that the OEC is totally conserved as the site of oxygen production from water in all known oxygenic phototrophs.12 Using the single-photon photochemistry of the P680 chlorophylls of the reaction center, the manganese cluster is progressively oxidized via a tyrosine residue (YZ) in one-electron steps. After four electrons are removed from the OEC, sufficient potential has been accumulated for oxygen evolution from water to occur.20 From the resulting reduced state, the cycle can begin again. Joliot first observed this periodfour behavior in single-turnover flash experiments,21,22 and Kok B

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An additional feature that should not be ignored is the possibility of formation of heterogeneous, catalytically active metal oxide nanoparticles, or related materials under catalytic conditions. Because the conditions required for driving water oxidation are harsh, the conversion of the homogeneous, molecular metal complexes into secondary materials is often possible. Starting with a molecular precatalyst, ligand oxidation/ modification to form new species in solution could eventually lead to complete loss of the initial ligand set. Complete loss of ligand under oxidizing conditions then gives a high probability of forming metal oxide particles. Thermodynamically speaking, the decomposition products of most transition-metal catalysts under oxidizing conditions would be the corresponding metal oxides.39 Thus, there are numerous challenges to be overcome in developing and characterizing the mechanistic details of water-oxidizing chemistry in homogeneous systems.

damaged centers presents a challenge for efforts to develop biohybrid approaches and related devices for artificial photosynthesis, although recent results have shown some promise.30−33 1.3. Water Electrolysis

The process of biological water splitting bears resemblance to water electrolysis, in which an electrical potential difference is used to split water at two electrodes. The two half reactions and the overall reaction at acidic pH values are as follows Anode: 2H 2O → O2 + 4H+ + 4e−

(2)

Cathode: 2H+ + 2e− → H 2

(3)

Overall: 2H 2O → O2 + 2H 2

(4)

The hydrogen gas fuel can then be collected, stored, and used at a later time. The overall reaction of water splitting was accomplished as early as 1789 by van Troojstwijk and Deiman34 and with Volta’s pile (i.e., battery) in 1800 by Nicholson and Carlisle.35−37 Because the H2/H+ equilibrium is essentially reversible at a platinized platinum electrode,38 the equilibrium potential of this couple (the so-called normal or standard hydrogen electrode (NHE)) is arbitrarily set at 0 V. On this scale, the standard electrode potential (E°) for the water-oxidation reaction (eq 2) is 1.23 V vs NHE.25 Because this is a four-electron reaction, this gives a quite endothermic free-energy change of 359 kJ mol−1.12 Reaction kinetics are an important feature of oxygen evolution. In an aqueous cyclic voltammety experiment, up to 1.23 V vs NHE, essentially no current flows though the cell. This is due to the high kinetic barrier for oxygen evolution, which introduces the requirement for large overpotentials to observe significant conversion of water to oxygen. Here, we define the term overpotential as the additional potential required above the thermodynamic minimum to drive an electrode reaction at a given rate; for water oxidation with heterogeneous catalysts, this rate is typically normalized by electrode geometric surface area, for example, 0.5 mA cm−2. The need for a high overpotential (typically, more than 400 mV) in electrochemical water oxidation arises from the multiple intermediates required to couple two oxygen atoms while simultaneously liberating four protons and four electrons. Further increasing the difficulty of electrochemical water splitting, the required oxidizing equivalents must be properly coupled to yield dioxygen, rather than partially oxidized side products (e.g., HO•, H2O2).1 Ligand oxidation and modification are also of concern in molecular catalysts that rely on sensitive organic-based ligands. These difficulties hamper the development of synthetic water-oxidation electrocatalysts, especially catalysts that can operate under the required harsh oxidizing conditions in a sustained fashion with minimal deactivation. Homogeneous, molecular catalysts can be especially susceptible to deactivation and unproductive pathways, since they can undergo side reactions involving ligand oxidation and degradation. Heterogeneous metal oxide catalysts, on the other hand, have been in use as catalysts for water electrolysis for many years. As a general rule, the robustness (or lack thereof) of a homogeneous catalyst is an important characteristic that is worth evaluating in all cases but most particularly for water oxidation where harsh conditions are required and where the ligand incorporates a potentially oxidation-labile group. In terms of ligand design, oxidative robustness is a primary consideration in the field.

1.4. Heterogeneous Oxide-Based Catalysts

A dependence of the electrochemical overpotential for a given reaction on the nature of the electrode material has been recognized since the dawn of the 20th century.40−43 Transitionmetal oxides (e.g., ruthenium, iridium) were observed early to have good catalytic properties for water oxidation.44−46 Among these, iridium materials have proven to be among the most active and stable for oxygen evolution.47 Among all materials for water oxidation, transition-metal oxides of Groups 7, 8, and 9 occupy a privileged position. Manganese, as the archetypal element for water-oxidizing chemistry with its known role in PSII, anchors the group of known elements for oxygen evolution. Ruthenium and iridium, located diagonally in the Periodic Table relative to manganese, are the other elements also strongly associated with wateroxidation chemistry, because they are currently found in commercial proton-exchange membrane (PEM) electrolyzers.48−50 Ruthenium and iridium oxides benefit from stability at low pH values, making them suitable for the low pH values found in the PEM systems. The electrochemical behavior of Ru and Ir heterogeneous water-oxidation catalysts has recently been reviewed.51 Not surprisingly, though, due to their low cost and relevance to PSII, manganese oxides have received increased attention in recent years. Work from Kurz et al. with manganese(III) oxides52 has yielded higher activity manganese materials than previously available. In their preparations, solution-phase methods give crystalline but high surface area materials which are catalysts for oxygen evolution, even with the one-electron oxidant cerium(IV) (see below for information on this oxidant problem) or in a light-triggered system with [Ru(bpy)3]2+ (bpy = 2,2′-bipyridine) as a photoinitiator and a cobalt(III) complex as the electron acceptor. Perhaps the most interesting aspect of the Kurz work has been the high activity obtained from manganese(III) oxides containing calcium; their high surface area preparations of the naturally occurring mineral Marokite (CaMn2O4) seem to be more active for water oxidation than the calcium-free analogs. However, it is difficult to ascribe this to a difference in intrinsic activity, as addition of calcium could also change the morphology (surface area or porosity) of these materials in unpredictable ways. Jaramillo and co-workers recently revisited manganese oxides for water oxidation and oxygen reduction in a bifunctionalelectrode approach. In their preparation,53,54 manganese acetate was used for electrodeposition of manganese oxide, which was then calcined at 480 °C for 10 h. The activity obtained with this C

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In 1999, our group published the structure and further characterized the water-oxidizing chemistry of the related Mn2(III,IV) mixed-valence terpyridine dimer of manganese, the so-called “terpy dimer.”87,88 This complex, isolated as [Mn(terpy)(H2O)O]2(NO3)3 (Figure 3), is a functional model for

material was impressive, but as previously reported, high pH values were required for high activity.25 Interestingly, in their work, iridium on carbon was found to be the superior catalyst, and ruthenium and platinum were less highly performing.53 Jaramillo and co-workers also investigated the solar photochemistry of crystalline, semiconducting Birnessite-type MnO2,55 which undergoes direct band-gap excitation to produce oxygen with an applied bias voltage. Unfortunately, however, this material is a poor performer, giving low external quantum efficiencies for conversion of light energy. Other base metal catalysts, most based on cobalt, nickel, and nickel−iron oxides,56−59 are also of current interest for wateroxidizing chemistry at alkaline pH values. Kanan and Nocera’s report60 of water oxidation with amorphous cobalt oxide containing phosphate has motivated much research toward the use of abundant elements for water-oxidizing chemistry.61 This has produced much information about the mechanism(s) of cobalt oxide-catalyzed water oxidation62−68 and includes efforts aimed at light-driven water oxidation in heterogeneous systems.69−73 A recent study74 examined water oxidation catalyzed by Co(II) adsorbed on silica nanoparticles; notably, differences between simple cobalt (oxide) materials and this hybrid system suggest future catalyst improvements may come from such supported or hybrid approaches.

2. MOLECULAR CATALYSTS FOR WATER OXIDATION Like their heterogeneous analogues, perhaps the bestcharacterized homogeneous catalysts rely on the three elements mentioned above as privileged catalysts for water oxidation: manganese, ruthenium, and iridium.75 The first homogeneous catalyst for water oxidation, Meyer’s so-called “Blue Dimer”, relies on ruthenium.76 The use of ruthenium in the first case is likely a consequence of the relative abundance of synthetically accessible complexes. Relatively slower ligand exchange rates also make observation of intermediates more likely. In the case of iridium, it was known to be the most active and stable oxide catalyst for water oxidation for many years but waited for work in the direction of homogeneous catalysis until late in the decade 2000−2010. There is also considerable interest in the use of cheap, base metals as catalysts, and this topic has recently been reviewed.77

Figure 3. Molecular structure of [Mn(terpy)(H2O)O]2(NO3)3.87 Nitrate counterions and hydrogen atoms on the bound water molecules have been removed for clarity. Colors: manganese, green; nitrogen, blue; oxygen, red. Thermal ellipsoids are shown at 30% probability.

O−O bond formation by the OEC in PSII, since high-valent manganese species are involved in formation of dioxygen. Oxygen evolution with this complex continues over hours with oxone (HSO5−) or hypochlorite (ClO−) as oxidants. The proposed mechanism of action for the complex with oxone as primary oxidant is shown in Figure 4. In the proposed mechanism of oxygen evolution,89 oxone first binds to the (III,IV) dimer. Because there are two manganese centers, one Mn(III) and one Mn(IV), there are already two paths here. We imagine that when oxone binds to Mn(IV), there can be no oxidation since a two-electron oxidation would give a Mn(VI), which is inaccessible in this ligand environment even with the high-potential oxone. When it binds to the Mn(III), however, productive chemistry does result.90 The departure of sulfate (SO42−) results in the twoelectron oxidation of manganese(III) to manganese(V) and the formation of the key high-valent intermediate responsible for O−O bond formation. This oxidation of manganese(III) to what would be formally manganese(V) appears to be the ratedetermining step in the catalytic cycle, as no intermediates besides the (III,IV) dimer can be detected. The high activity of this manganese(V)−oxo or manganese(IV)−oxyl91 intermediate may stem, in part, from its trans position to one of the bridging oxo ligands which are expected to have a larger trans influence. Under the catalytic conditions, there is a large amount of catalytically irrelevant (IV,IV) dimer in solution, observable by X-ray absorption spectroscopy.90,92 It likely forms by comproportionation of the (III,IV) dimer and transient (V,IV) dimer formed by oxidation with oxone; the (IV,IV) dimer is not sufficiently oxidizing to catalyze water oxidation.92

2.1. Manganese Catalysts

Many groups, including our own, have long been interested in the basic coordination chemistry of manganese due to the relevance that synthetic manganese complexes can have for modeling various aspects of the oxygen-evolving complex in Photosystem II.78−83 Here, we focus on work with functional manganese catalysts for water oxidation, but it should be noted that there are other systems which have been investigated as models of the OEC in PSII.84,85 In 1997, our group found that solutions containing manganese(II) dipicolinate (dipic) or 2,2′:6′,2″-terpyridine (terpy) complexes produced oxygen upon treatment with the high potential two-electron, oxo-donating oxidant potassium persulfate (KHSO5, oxone; E° = 1.85 V).86 In the case of the [ONO] X2L-type dipic ligand, oxygen evolution was quickly lost as permanganate built up; this was presumably occurring by loss of the dipic ligand and subsequent overoxidation of manganese. However, the analogous [NNN] L3-type terpyridine (terpy) ligand gave sustained oxygen evolution with little deactivation even after 1 h, as well as no permanganate formation. D

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none have been found to function better than the parent terpyridine dimer.93,94 In some cases, this likely arises from solubility effects, because the terpyridine ligands having ethoxy or phenyl substitutions are less soluble in water. Manganese can be lost from the terpyridine ligand during the formation of lowvalent intermediates that result from reduction of the complex during dioxygen formation; thus, ligand loss would result in irreversible precipitation of the ligand and loss of the ability to reform the relevant manganese complexes. Second, the donor effects of the substituted manganese complexes may be sufficiently different than the parent terpyridine to destabilize or adversely affect formation of the various intermediates required for oxygen evolution to take place. Along these lines, it has been of interest to move work with this system to one-electron chemical oxidants, such as those used in ruthenium-catalyzed water oxidation. Unfortunately, cerium(IV), the most common high-potential one-electron oxidant used for oxygen-evolution experiments, is only stable at very low pH values (i.e., below pH 1.5).44 When a solution of the [Mn(terpy)(H2O)O]23+ complex is added to a solution containing cerium(IV), oxygen is evolved.95 However, only up to ca. 0.5 equiv of dioxygen per dimanganese dimer can be measured. A calculation based upon the four-electron conversion of a dimanganese(IV,IV) dimer to manganese(III) and 1 equiv of dioxygen gives a predicted 48% yield of dioxygen in this experiment. On the basis of the electrochemistry of the complex, cerium(IV) has sufficient potential (E° = 1.7 V vs NHE)96,97 to oxidize the Mn−terpy dimer from the starting (III,IV) state to its known (IV,IV) state, but the nature of the further reaction is unclear. If the dimer solution is adjusted to pH 1 before addition to the cerium(IV) solution (already at pH 1), no oxygen is evolved. This has been shown by related experiments to be due to loss of the dimeric structure by protonation of the susceptible oxo bridges and subsequent formation of a Mn dimer of dimers.98 These results suggest that stabilization of the pH-susceptible oxo bridges would restore catalytic chemistry with this one-electron oxidant. However, attempts to stabilize the dimeric structure with specialized, “strapped” chelate ligands have not been successful.99 Extensive control experiments have thoroughly ruled out a role for manganese oxide as the catalytic species in this system.92,93,95 Interestingly, however, colloidal manganese oxide solutions with terpyridine added do produce oxygen upon treatment with oxone, likely due to in situ assembly of the manganese terpyridine dimer.92 Manganese oxide has been reported to have some activity for water oxidation with manganese(IV) pyrophosphate and [Ru(bpy)3]2+ as a photoinitiator;100 however, this reaction is likely a noncatalytic disproportionation reaction of the manganese(IV) complexes, because no primary oxidant was used. The involvement of the ruthenium species in the catalysis cannot be discounted either, considering the high activity of ruthenium complexes and ruthenium oxide for water oxidation (see below). The electrochemistry of the Mn−terpy complex has been of great interest, since it is relevant to advancing the catalyst in one-electron steps.101 Collomb et al.102 found that one-electron oxidation of the starting [Mn(terpy)(H2O)O]23+ complex in aqueous solution (the mixed-valence (III,IV) dimer) occurs at ca. 1.3 V vs NHE (Figure 5). This oxidation causes conversion of the oxidized dimer to a tetrameric Mn4(IV,IV,IV,IV) complex in which two (IV,IV) dimers are linked via a monoμ-oxo bridge.98 However, the first oxidized intermediate, the Mn2(IV,IV) dimer, can be detected by its reduction back to the

Figure 4. Proposed mechanism for the reaction between [Mn(terpy)(H2O)O]23+ and two-electron, oxo-donating reagents, XO.89 (Left pathway) Attack of XO on the terminal oxo ligand gives rise to oxone−oxone coupling (no incorporation of oxygen atoms from oxidation of water). (Right pathway) Attack of solvent water on oxidant-deposited oxo ligand gives rise to water oxidation and consequent incorporation of atoms from substrate water into product dioxygen. Reprinted with permission from ref 89. Copyright 2001 American Chemical Society.

Despite this high known concentration of (IV,IV) dimer under the catalytic conditions, impressive rates with respect to total initial dimer loading of up to 2420 mol of O2 (mol dimer)−1 h−1 have been reported, which gives a turnover frequency of 0.67 s−1.89 This is an impressive apparent turnover frequency that is likely an underestimate of the real turnover frequency of the small portion of the solution that is active for the catalytic reaction. 18 O isotope experiments have been key in elucidating detailed aspects of this catalytic system.89,92 In these experiments, the heavy isotope is added in the form of H218O that is catalytically converted to 18O-labeled 34O2 or 36O2. Because oxone is always responsible for delivery of one oxygen atom in the product dioxygen, the chemistry taking place can be considered formally as one part of the oxygen-evolution half reaction of the overall water-splitting process. In the absence of oxo exchange, only up to 50% of the oxygen atoms in the product could arise from oxidation of solvent water resulting in formation of 34O2. 18O-labeling experiments reveal the formation of both 34O2 and 36O2, indicating that some oxo exchange can occur. Because oxone is slow to exchange with water on the time scale of these experiments, these experiments validate the ability of the high-valent manganese(V)−oxo intermediate to oxidize water. Along this line, it is interesting to note that at high oxone loadings, very little catalytic water oxidation is observed; instead, oxone−oxone coupling is seen. In this pathway, the high-valent Mn(V)−oxo attacks a second oxone molecule, giving rise to dioxygen in which both oxygen atoms come from oxone. This highlights the importance of carefully monitoring the conditions of any catalytic system to include all conditions and reagents present. Following these initial investigations, further work has continued elucidating the details of this complex catalytic system. Substituted derivatives of terpyridine have been investigated for use as related homogeneous catalysts, but E

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Similarly, Naruta et al. studied a manganese bis(porphyrin) system for oxygen evolution.108,109 In this system, mchloroperbenzoic acid drives formation of a bis-Mn(V)−oxo intermediate that stoichiometrically decomposes to form dioxygen upon treatment with triflic acid. Thus, this “catalyst” can achieve nearly a single turnover by sequential chemical treatments. Similarly, a single-turnover system based on a salen ligand derivative (N,N-3,5-dichlorobis(salicylidene)-1,2-diaminoethane) and a related system based on a dinuclear analogue have been reported.110−113 McKenzie and co-workers studied a dinuclear manganese complex as a catalyst for water oxidation driven by tert-butyl hydroperoxide.114,115 Isotope incorporation experiments with H218O confirmed incorporation of 18O into O2 produced by this system. However, we note that use of such a peroxide reagent opens the possibility of the involvement of radical species in oxygen evolution. A number of other manganese compounds have recently been reported by Anderlund and others to function as catalysts for water oxidation, driven by oxone or cerium(IV).116,117 Stoichiometric incorporation of added 18O from H218O was observed in one prominent case.118 A related systematic screening of manganese catalysts with a variety of chemical oxidants revealed several interesting trends in activity.85 More generally, the various chemical oxidants used in assaying wateroxidation catalysts have also recently been reviewed.119 Oxone, cerium(IV) ammonium nitrate, [Ru(bpy)3]3+, and Co(III) compounds are among the commonly used oxidants. Work on heterogenized systems of [Mn(terpy)(H2O)O]23+ absorbed on insoluble clay minerals such as kaolinite, montmorillonite, and mica shows activity for oxygen evolution with one-electron chemical oxidants such as cerium(IV) and [Ru(bpy)3]2+/persulfate.120,121 A tetrameric analogue has also been explored in a heterogenized system.122 Other dimeric manganese complexes have been found to evolve oxygen when heterogenized into clay.123 It should be noted, however, that determining the catalytic turnover frequency or mechanism is made quite difficult by formation of the heterogeneous systems. The total turnover attained in these systems was less than 20 mol of O2 (mol dimer)−1 and rates no greater than 0.0002 turnovers s−1.124,125 Moreover, X-ray absorption spectroscopy suggests layered manganese oxides are formed in these systems, complicating the search for atomically precise catalytic motifs.126 In a particularly intriguing recent report from Yagi et al.,125 the rate constant for oxygen production was found to be correlated with the potential of the Mn2(III,IV)/Mn2(IV,IV) couple found in cyclic voltammetry experiments. In their study, various derivatives of the [Mn(terpy)(H2O)O]23+ dimer were absorbed on mica, where the parent terpyridine dimer gave the best results, with the other derivatives having decreasing activity as the donor power of the ligand increased, leading to lower values for the Mn2(III,IV)/Mn2(IV,IV) potential. Analogous work aimed at functionalizing electrode surfaces with water-oxidation catalysts has yielded similar results. Li et al. found that nearly amorphous titanium dioxide could stabilize the Mn2(III,IV) dimer of terpyridine, based on EPR spectroscopy.127 Some oxygen was evolved from this system when treated with cerium(IV), encouraging further development of similar systems. In a related report, one group has found128 that the [Mn(terpy)(H2O)O]23+ dimer can be polymerized onto an electrode by extensive potential pulsing, although these results have yet to be corroborated, and may be inconsistent with

Figure 5. Electrochemistry of [Mn(terpy)(H2O)O]23+ in water.102 Pa1 corresponds to oxidation of the dimer from the Mn2(III,IV) to the Mn2(IV,IV). Pc2 likely corresponds to the paired reduction of the Mn2(IV,IV) dimer to Mn2(III,IV). Pc1′ corresponds to reduction of the tetrameric Mn4(IV,IV,IV,IV) complex which forms upon dimerization of Mn2(IV,IV) following its electrochemical generation in Pa1. Conditions: ca. 5 mM [Mn] in pH 4 aqueous solution with 0.1 M KCF3SO3 as supporting electrolyte; scan rate 5 mV/s; working electrode was a vitreous carbon disk. (Ag/AgCl vs NHE: +197 mV.) Adapted with permission from ref 102. Copyright 2005 American Chemical Society.

(III,IV) dimer at 1.2 V. Additionally, reduction of the tetrameric species is observed at 1.1 V. Notably, no catalytic waves were observed in the electrochemistry; Collomb et al. ascribed this lack of electrocatalysis to the formation of the catalytically inactive tetramer, at least at the potentials examined. Additionally, it is likely that the Mn2(IV,IV) to Mn2(V,IV) oxidation is inaccessible under these conditions; in contrast, the Mn2(III,IV) to Mn2(V,IV) oxidation is carried out by the reaction of the dimer with oxone. This work has been extended to pH-dependent electrochemical studies.103 Cady et al. found that the one-electron oxidation of the Mn2(III,IV) dimer was pH dependent. This is a result of the greatly lowered pKa of a water molecule bound to manganese in the Mn2(IV,IV) oxidation state. Consistent with expectations from electrochemical theory,104 within the range of stable pH values for the dimer, the one-electron oxidation of the dimer is pH dependent, showing a 59 mV/pH-unit decrease in the potential for oxidation as the pH value is increased. The potential for oxidation of the dimer shifts from near 1.35 V at pH 2 to near 1.10 V at pH 5.5. Furthermore, Cady et al. found that under acetate-buffered conditions, only one acetate is bound to the Mn2(III,IV) dimer, presumably at the manganese(IV) site. This leaves a single bound water for participation in pH-dependent electrochemistry and is consistent with the mechanistic work showing the dissymmetric chemistry of the dimer under catalytic conditions with oxone as primary oxidant.90 All current work thus suggests that catalysis with the [Mn(terpy)(H2O)O]23+ complex cannot be driven with oneelectron oxidants in aqueous solution. This includes work with related systems in molecular constructs with [Ru(bpy)3]2+ which can form the Mn2(IV,IV) state but not a form of relevance to oxygen evolution.105 This is also the case in systems involving photooxidation of mononuclear manganese complexes.106 A related recent study of a mononuclear catalyst with an anionic carboxamido N-donor ligand suggests that among similar complexes, an anionic ligand lowers the potential required for oxidation of the metal complex and consequently enables water oxidation in certain cases.107 F

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A landmark report came from Moyer and Meyer in 1981;147 in their paper, the redox properties of [RuII(bpy)2(py)H2O]2+ are described (shown schematically in Figure 6). They found

preparative chemistry of tetrameric manganese(IV) terpyridine complexes carried out in our own group.90,98 Similarly, Hou and co-workers investigated thermal decomposition of the [Mn(terpy)(H2 O)O]2 3+ dimer129 to give formation of heterogeneous material that can function as a water-oxidation catalyst. This work follows previous studies in which nanostructured manganese oxide materials could be formed from dimeric high-valent manganese complexes.130 In all these systems, though, solid evidence of the activity of the molecular decomposition products versus simple amorphous manganese oxides remains elusive. The question of catalyst decomposition and its prevention is a topic of current interest,131 but the harsh oxidative conditions of water oxidation put a particular strain on the stability of the ligand set, as emphasized in a recent review on that aspect of water-oxidation catalysis.132 Along these lines, Spiccia and Dismukes reported133,134 a system in which a coordinatively saturated tetramanganese complex was doped into a Nafion-coated electrode, thereby heterogenizing the catalyst. The electrode assembly oxidizes water upon illumination with UV light and application of a significant 1.2 V vs NHE bias voltage. In other work, the bias voltage could be eliminated by inclusion of a ruthenium− polypyridyl complex as photosensitizer.135 Results from Spiccia et al.,136 as well as related work from Young et al. in our group, 137 have now demonstrated that the molecular manganese complex in the electrode assembly is not the true catalyst for oxygen evolution. Rather, manganese oxide materials formed in situ are the true catalyst. This was exemplified in studies by Young et al. on the infrared spectra of [MnO(tacn)]44+-doped Nafion electrode assemblies where tacn is 1,4,7-triazacyclononane.137 In the IR spectra, a characteristic 730 cm−1 breathing mode of the adamantane core of the manganese tetramer disappears after subjecting the system to the relevant catalytic conditions, showing conversion of the heterogenized molecular complex to secondary material. Various other systems involving nominally molecular species that evolve oxygen have been described but defy exact characterization. Assembly of a device often makes catalyst identification difficult.138 In other cases, ligands prone to side reactions may be involved in promoting heterogeneous chemistry.139−141 Prior work from the Dismukes group demonstrated that the manganese tetramer used to prepare some of the electrode assemblies135 is not a water-oxidation catalyst under the relevant homogeneous conditions.142,143 A recent electrochemical study of a homogeneous complex and its heterogenized analogue showed that electrocatalysis does not occur with the molecular species in solution.144

Figure 6. Electrochemical studies from Moyer and Meyer147 on [RuII(bpy)2(py)H2O]2+. (Top) Cyclic voltammetry of the metal complex at pH 2. (Middle) Processes taking place in electrochemistry at pH 0. (Bottom) Processes taking place in electrochemistry at pH 7. All potentials are reported vs SSCE (SSCE vs NHE +236 mV). Conditions: gold-wire working electrode; 20 mV/s; 0.1 M LiClO4 as electrolyte. Adapted with permission from ref 147. Copyright 1981 American Chemical Society.

that at pH 0, oxidation of the complex to the RuIII−OH2 state occurs at 1.02 V, followed by oxidation to RuIVO at 1.24 V. However, at pH values above 0.85 (the pKa value of bound water in the RuIII−OH2 ion), the first oxidation is proton coupled, resulting in formation of the RuIII−OH complex rather than RuIII−OH2. Thus, the redox chemistry at pH 7 behaves as follows: oxidation of RuII−OH2 to RuIII−OH occurs at 0.66 V, followed by oxidation to RuIVO at 0.77 V. This is consistent with the expected 59 mV/pH-unit shift expected for a 1H+/1e− process.43,104 In this same report,147 the possibility of comproportionation of [Ru I I (bpy) 2 (py)H 2 O] 2+ and [RuIV(bpy)2(py)O]2+ to form a dimeric species postulated to be [{RuIII(bpy)2(py)}2O]4+ was explored. In fact, the species of interest here for water-oxidizing chemistry is the dimeric complex with two open sites for binding waters. Thus, the initial Meyer work on single-site catalysts did not find any water-oxidation behavior; they did find, however, that oxidation of organic substrates was possible with the mononuclear complexes.148 Meyer’s blue dimer (BD), the first reported homogeneous water-oxidation catalyst, is the [{RuIII(bpy)2(H2O)}2O]4+ ion shown in Figure 7.76 In the initial communication, they found a catalytic oxidation wave for BD by cyclic voltammetry in 0.1 M sulfuric acid at around 1.54 V vs NHE (Figure 8). In this complex, the protonation state of the bound water molecules in solution is unclear, as well as the nature of the involvement of the oxo bridge in the water-oxidizing chemistry. However, the complex is clearly a catalyst both in chemical and in electrochemical experiments: when the dimer is added to a

2.2. Ruthenium Catalysts

The first homogeneous water-oxidation catalyst, the ruthenium polypyridyl complex reported by Meyer and co-workers in 1982,76 has emphasized the role of proton-coupled electrontransfer (PCET) processes for water activation,145 as outlined in a recent review.146 In such processes, oxidation of the metal center causes a pKa shift of water ligands bound at the metal; the net result is activation of bound water upon metal oxidation, giving hydroxo or oxo complexes upon oxidation of the ruthenium center. Enhancing the donor power of the ligand set in each PCET step, by going from aqua to hydroxo to oxo, makes the next oxidation state advance easier to achieve, thus promoting the multielectron oxidation normally considered necessary for water oxidation. G

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minimum for onset of catalysis to occur; at pH 7, this additional voltage is 630 mV (E°′ at pH 7 = 820 mV).

Figure 7. Molecular structure of [{RuIII(bpy)2(H2O)}2O]4+ (BD).149 Counterions and hydrogens of the bound water molecules are omitted for clarity. Colors: ruthenium, orange; nitrogen, blue; oxygen, red. Thermal ellipsoids are shown at 30% probability.

Figure 9. Pourbaix (potential vs pH) diagram for BD.149 The regions of the diagram are labeled with the expected oxidation state of the ruthenium centers in the dimer under the conditions. The pHindependent oxidation to form the key (V,V) intermediate is shown as the horizontal line above pH 2 of about 1.44 V vs NHE (1.2 V vs SCE). (SCE vs NHE +241 mV.) Adapted with permission from ref 149. Copyright 1985 American Chemical Society.

The mechanism of oxygen evolution with BD has been the focus of much research over the past 30+ years. Mechanistic work on water oxidation by ruthenium catalysts has recently been thoroughly reviewed by Hurst.150 The key mechanistic question, much as in the case of the manganese catalysts, has centered on the source of the O atoms incorporated into product dioxygen. Isotope-labeling studies, described in detail in ref 150 and carried out in both the Hurst and the Meyer laboratories, has validated two primary pathways for O−O bond formation. The primary pathway involves a nucleophilic attack of an uncoordinated water molecule on a high-valent Ru−oxo species resulting in formation of a peroxidic intermediate. Notably, this general mechanism has been supported as operative in numerous molecular water-oxidation catalysts, including those based on manganese and iridium that were developed following the early work on BD. Classic work151 included spectroscopic identification and measurement of the kinetics of formation of the peroxidic intermediates resulting from O−O bond formation via nucleophilic attack of water. The isotopic-labeling work has conversely ruled out direct coupling of two adjacent [Ru]-bound oxo groups from an oxidized form of BD, as shown in eq 5 (where O* is isotopically tagged, often as 18O)

Figure 8. Electrochemical studies from Meyer et al. on BD at pH 1.76 A catalytic oxidation wave onsets at ca. 1.54 V vs NHE following reversible oxidation of the dimer below 1.2 V vs NHE. Conditions: glassy-carbon working electrode; 100 mV/s; 0.1 M H2SO4 as electrolyte. Potentials given versus SSCE. (SSCE vs NHE +236 mV.) Adapted with permission from ref 76. Copyright 1982 American Chemical Society.

solution containing 50 or 100 equiv of cerium(IV), oxygen was evolved cleanly and in the proper proportions based on the redox equivalents involved (4 equiv of CeIII were produced along with 1 equiv of O2). In the second report from Meyer on this chemistry,149 the pH-dependent speciation and electrochemical properties of BD were described. Notably, formation of a final Ru2(V,V) oxidation state (with two oxo ligands) is thought to be required for oxygen evolution to occur. On the basis of the number of protons in the initial bis-aqua dimer, the final oxidation of the complex to form the required high-valent intermediate is not proton coupled. This results in a pHindependent potential for water oxidation above pH 2 of about 1.44 V, as shown in the Pourbaix diagram of Figure 9. Notably, this is a relatively large potential above the thermodynamic

2Ru V O* + 2H 2O → 2Ru III−OH 2 + O*O*

(5)

This pathway is conceivably possible with BD, but nucleophilic attack of a second water molecule is preferred. This intramolecular oxo-oxo coupling does not seem possible in part due to steric considerations with BD. A second operative pathway leading to oxygen evolution with BD has been suggested to involve “direct” oxygen evolution from two water molecules. This pathway, though, is more H

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A mechanistic study on the BD with UV−visible, EPR, and XAS methods gave information on processes occurring on the 100 ms time scale. Under the conditions employed, the (III,IV) form was identified as the resting state and its oxidation by Ce(IV) was considered the turnover-limiting step. In addition, a (IV,V) state was seen as a transient that immediately precedes O−O bond formation.162 In related work, an EPR study probed a Ru(V)O intermediate of the BD labeled with 17O, where hyperfine splittings consistent with high radical character on oxygen were seen. This outcome was ascribed to destabilization of Ru(V)O π bonding, a feature that also enhanced the reactivity of this intermediate.163 Another EPR study looked at the (V,V) intermediate of the BD with the conclusion that this species has an S = 1/2 configuration with Ru oxyl character.164 More recent work from the Meyer group has examined methods by which to improve catalysis with BD, since it is slow and requires high potentials. In some reports, the complex was absorbed onto electrode surfaces,165 but in that case the exact coordination sphere of the ruthenium is unclear. Heterogeneous chemistry could even be occurring. Elsewhere, Meyer examined mediator-assisted water oxidation with BD, which has given good mechanistic insights.166 In 2004, Llobet and co-workers reported a new ruthenium dimer for water oxidation,167 in which the specialized Hbpp ligand (where Hbpp is 3,5-bis(2-pyridyl)pyrazole) was used to enforce a cis geometry of the two ruthenium centers. Addition of the metal complex to cerium(IV) evolves oxygen slowly at a rate of 0.014 turnovers s−1. Since the turnover or disproportionation of various intermediates is slow in the case of this complex, mechanistic studies have examined and validated168 with a good level of detail that the oxo−oxo coupling pathway leads to oxygen evolution. This is in contrast to the work discussed above with BD. In some later work with a related analogue, an intermolecular pathway was preferred over an intramolecular one.169 The coordination chemistry of the Hbpp ligand has been explored in great detail for ruthenium.170,171 In a binucleating version of the ligand, shown in Figure 10, the properties of the

difficult to visualize and deserves further investigation. Work from Hurst and co-workers152−155 centered on the possibility that ligand oxidation and hydration was involved in this observed pathway; this hypothesis has not so far gained a wide following and is not discussed here in depth. Observations have been made regarding noninnocence of the cerium(IV) ammonium nitrate often used to drive catalysis with BD,156 which further complicates work in this area. In the unexpected twist, the nitrate counterion of the Ce(IV) oxidant, usually assumed to be innocent, was shown by UV−visible, resonance Raman, and Raman data to have an influence on the dynamics of the BD, although no O derived from nitrate appeared in the O2 formed.156 The central position of the BD is further emphasized by the many computational studies that have been carried out on it. For example, Bianco et al.157,158 looked at the O−O bondforming step by DFT with an explicit solvent treatment. They introduced 80 water molecules of which the inner 4 were described quantum mechanically. A key proposal in this work was that a proton relay links adjacent active sites and mediates proton transfers via a chain of two hydrogen-bonded water molecules. Nucleophilic attack of a water on a RuO was accompanied by a proton transfer from that water via the water chain. The calculated barrier for this process was consistent with the experimental value. Thus, the water nucleophilic attack (WNA) mechanism appears to be the predominant but not exclusive path for WO, as demonstrated in the experimental work. Meyer and colleagues carried out a comprehensive combined experimental−theoretical study to quantify the extensive electronic coupling that links the two metals via the μ-oxo group in the (III,III) state. The DFT analysis provides assignments for the near-IR and visible bands; RuORu-based dπ → dπ* and interconfiguration transitions also appear in the near-IR, while MLCT and LMCT transitions are in the visible. Bending of the Ru−O−Ru unit was described to arise from a Jahn−Teller distortion.159 The oxidation state changes leading to oxygen evolution have also been discussed extensively, and a 2004 article from Hurst160 is notable for its insights. In concentrated triflic acid, electrochemical studies161 show well-resolved features for the various oxidations of the complex from the (II,II) state up to the key (V,V) intermediate that gives rise to oxygen evolution. The formation of this (V,V) state occurs at 1.6 V vs NHE under these conditions where the (V,V) form of the dimer is relatively stable in solution, slowly giving rise to oxygen production. Thus, Hurst was able to independently measure H/D kinetic isotope effects with the chemically prepared (V,V) form of the dimer, as well as by addition of BD to a solution containing excess cerium(IV) as primary oxidant. In both cases, the H/D KIE was found to be 1.6−1.7. This is consistent with cleavage of O−H bonds during catalysis and strongly suggests that the (V,V) state of the dimer or a closely related species is the key intermediate in this chemistry. These studies demonstrate that catalysis of oxygen evolution with BD is slow. In fact, depending on the conditions, the rate of oxygen evolution is a modest 0.0053 turnovers s−1. These low turnover frequencies are a benefit for mechanistic understanding, since the key intermediate leading to oxygen evolution can be prepared and characterized. The rate-determining step in the chemistry is considered to be the O−O bond-forming step which requires the (V,V) oxidation state of the dimer. This mechanistic proposal is consistent with the measured H/D KIE as well as the isotopic-labeling studies.

Figure 10. Binucleating Hbpp-type ligand used for preparing tetranuclear Ru complexes. Reprinted with permission from ref 172. Copyright 2011 American Chemical Society.

tetraruthenium catalyst depended on the ortho, meta, or para orientation of the linker, the O2/CO2 ratio of evolved gases being highest for the ortho case. The CO2 formation was ascribed to ligand oxidation.172 This area has now been comprehensively reviewed.173 Other related complexes come from Thummel et al.174−178 and Sun et al.179,180 In the Sun systems, dimeric ruthenium complexes contain anionic carboxylate donors in addition to the usual pyridyl groups. This was intended to lower the I

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modulate the electron-donor power of the bpy. An O−H···F hydrogen bond was proposed on the basis of a crystal structure of the aqua complex.193 In another recent study, the activity of the Ru(terpy)(bpy) system with Ce(IV) was shown to be enhanced by the presence of a redox mediator of the [Ru(bpy)3]2+ type. Spectroscopic and electrochemical data support the presence of an interaction between the catalyst and the mediator that affects the water oxidation.194 An unusual series of catalysts is based on the parent structure shown in Figure 12, containing a dipyridyl functionalized with

potential required to reach the key +5 oxidation state and, thus, lower the potential required for oxygen evolution. Indeed, impressive rates of up to 300 turnovers s−1 have been observed.181 Thus, under certain conditions and for certain specialized catalysts, the oxo coupling pathway (OCP) may be operative over the more common water nucleophilic attack (WNA) path. Sun, Privalov, and co-workers found this to be true in a 2010 report.182 In a rather exotic system, the OCP was traced out by DFT studies in the context of an electrocatalyst183,184 featuring two Ru centers with quinone-like ligands separated by a bis(terpyridine)-substituted anthracene linker.185−187 Interestingly, the quinone ligands are thought to be noninnocent by acting as a temporary store of redox equivalents. More recently, mononuclear ruthenium complexes have been found to be active catalysts. In 2005, Thummel reported oxygen evolution catalyzed by a mononuclear ruthenium precursor174 (T1) with the bis-naphthyridyl-pyridine ligand shown in Figure 11 and containing pendant base sites adjacent to the active site

Figure 12. Distorted catalyst (left) facilitating attainment of the 7coordinate oxo structure shown (right). Reprinted from ref 195. Copyright 2012 National Academy of Sciences.

carboxylate groups in the 2 and 9 positions so they can bind to the metal in an O,N,N,O-donor arrangement. The complexes have a strongly distorted octahedral coordination that is believed to facilitate attainment of the Re(V)O intermediate, now 7 coordinate. The preferred ligands for O2 formation from Ce(IV) were electron withdrawing and hydrophobic.195,196 In one case, a turnover number of >100 000 was achieved with a Ce(IV) oxidant.197 Computational work has provided insight into the competition between the WNA and the OCP mechanisms,198 and improved performance has been achieved by incorporating the catalysts into a nanofibrous coordination polymer.199 A three-component light-driven water-oxidation system composed of a conventional photosensitizer, persulfate as a sacrificial electron acceptor, and the same type of wateroxidation catalyst (WOC) as catalyst produced oxygen with an excellent quantum yield of 17%.200 Following Thummel’s report, Meyer reported201 extensive electrochemical and mechanistic work from 2008 showing oxygen evolution occurring with [Ru(terpy)(bpm)H2O]2+, where bpm = 2,2′-bipyrimidine. The mechanism described here is shown in Figure 13. Starting with the RuII−OH2 complex, electrochemical oxidation proceeds with a 2H+/2e− oxidation to form the RuIVO species (pKa (RuII−OH2/RuII− OH) = 9.7). Meyer speculated that the potential for oxidation of Ru(II) to Ru(III) is higher than the potential for Ru(III) to Ru(IV) due to stabilization of the low-valent state by the chelate bipyrimidine and, conversely, stabilization of the higher valent Ru(IV) by σ donation. At around 1.6 V vs NHE, a pHindependent, 1e− wave appears in the voltammogram as a shoulder to the onset of a catalytic wave corresponding to oxygen evolution. As shown by related experiments with cerium(IV), the formation of the ruthenium(V)−oxo intermediate gives catalytic water oxidation. A species assigned as a side-on peroxide intermediate can be observed in steady-state UV−visible absorption. This is presumed to form following attack of nucleophilic water on the electrophilic RuVO

Figure 11. Molecular structure of Thummel’s mononuclear ruthenium water-oxidation catalyst174 (T1) supported by the bis-naphthyridylpyridine ligand. Counterions and hydrogen atoms on the organic ligands are omitted for clarity. Colors: ruthenium, orange; nitrogen, blue; oxygen, red; protons on bound water, green. Reprinted with permission from ref 174. Copyright 2005 American Chemical Society.

of the catalyst. This complex evolved oxygen with a turnover frequency of 0.00035 s−1 in initial experiments and could reach 600 turnovers overall. The rate was later reported to be 0.014 s−1 under different conditions.188 A light-driven version proved possible in this case using a conventional photosensitizer and persulfate as sacrificial oxidant.189 In a closely related complex to that shown in Figure 11, an additional benzo group is fused to the wingtips of the pincer ligand. This change is advantageous in that attainment of the Ru(V)O state is no longer required (see below), as in the Ru(terpy)(bpy) series; in neutral and alkaline solutions, the mechanism instead passes through the Ru(IV)O via PCET steps. The reason for the success of the ligand modification is not yet clarified.190 Other successful implementations of the pendant base ligand design for Ru-catalyzed water oxidation have also been reported.191 The Ru(terpy)(bpy) series has also been modified by replacing bpy with 2,2′-bypyrimidine, a somewhat less electron-donating analogue,192 and with fluoro groups at the 5,5′ positions, also to J

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Figure 14. Pourbaix (potential vs pH) diagram for T1.204 The regions of the diagram are labeled with the expected species under the conditions. The pH-independent oxidation to form the key Ru(V) intermediate is shown as the horizontal line above pH 1 of about 1.4 V vs NHE. The two-electron oxidation from Ru(II) to Ru(IV) is a 2H+/ 2e− process above pH 2.9 but a 1 e− process below this pH. The result is a change in the slope from 59 mV/pH-unit to 118 mV/pH-unit in the respective regions, as shown in the blue lines and text labels. Adapted with permission from ref 204. Copyright 2011 American Chemical Society.

Figure 13. Meyer mechanism for single-site water oxidation with the [Ru(terpy)(bpm)H2O]2+ complex.201 A rapid oxidation from Ru(II) to form the key high-valent Ru(V) intermediate is shown on the right of the scheme. Decay of the peroxidic intermediate RuIV−OO to yield dioxygen is the rate-determining step; the rate of oxygen evolution was measured to be 0.00075 turnovers s−1. Reprinted with permission from ref 201. Copyright 2008 American Chemical Society.

species. This species then decomposes to form dioxygen and return the starting RuII−OH2 complex. Decay of the peroxidic intermediate to yield dioxygen is the rate-determining step; the rate of oxygen evolution is 0.00075 turnovers s−1. Under basic conditions, Meyer and co-workers found that the Ru(terpy)(bpy) electrocatalyst operates via single-electron activation by a combination of base-assisted atom proton transfer (APT) and direct reaction with OH−. APT is a concerted process in which a base removes a proton from water in concert with O···O bond formation, in this case, to form a Ru(III) hydroperoxide intermediate that can go on to form O2 with an overpotential of only 180 mV and a TOF of 5.4 s−1.202 Related kinetic data have been obtained on this class of catalysts by Sakai and coworkers.203 Another notable mechanistic study was carried out by Polyansky and Fujita204 in collaboration with Thummel’s group. In their study, the electrochemical and chemical oxidation of Thummel’s mononuclear water-oxidation T1 catalyst bearing the bis-naphthyridyl-pyridine ligand (Figure 11) was thoroughly investigated. Polyansky et al. found that above pH 2.9, oxidation of the T1 RuII−OH2 complex proceeds via a 2H+/2e− couple to form the RuIVO complex. This was consistent with complete pH-dependent electrochemical studies, which found a 59 mV/pH-unit slope in the Pourbaix diagram, as expected for a 2H+/2e− couple (Figure 14). However, below the pKa of 2.9, oxidation of the T1 RuII−OH2 complex gives RuIII−OH2. Formation of the RuVO was found to be pH independent above pH 1, similar to the case of Meyer’s [Ru(terpy)(bpm)] complex. Additionally, pulse radiolysis experiments were able to characterize the intermediate RuIII−OH complex at higher pH values, a species which typically would disproportionate rapidly or undergo further oxidation in electrochemical experiments.204 They generated this complex using the carbonate radical, which was generated in situ during the pulse radiolysis experiment and has a reduction potential of 1.59 V vs NHE,205 which is similar to the standard potential of the cerium(III/IV) couple at 1.6−1.7 V vs NHE.97 The UV− visible spectrum of the intermediate and the kinetics of its disproportionation to form RuII−OH2 and RuIVO were

obtained. These observations are, thus, consistent with the complete model that suggests that the Ru(II)/Ru(III) couple lies at a higher midpoint potential than the subsequent oxidation of Ru(III)/Ru(IV). In fact, this had been observed previously in an unrelated complex.206 Polyansky et al. also considered204 the oxidation of water by the lower valent RuIVO, whereas in the Meyer system, only the RuVO had previously been considered201 for O−O bondforming chemistry. Owing to the participation of a proton in the oxidation to form the Ru(IV) species, at higher pH values the participation of the Ru(IV) in O−O bond formation becomes progressively more favorable. This pH dependence of the Ru(III)/Ru(IV) couple causes the potentials for the Ru(III)/Ru(IV) and Ru(IV)/Ru(V) couples to approach each other at low pH values (ca. pH 1).207 Thus, the RuIVO and RuVO species were found in related experiments to be in equilibrium, which can contribute to mechanistic ambiguities due to disproportionation reactions under those conditions. Thus, mechanistic “branching” is possible in this system and depends on the conditions used for water-oxidizing chemistry with these ruthenium complexes. Multiple pathways seem especially likely and easily accessible in the cases where strong chemical oxidants such as cerium(IV) are used to drive oxygen evolution. Taken together, these results emphasize the point that in-depth knowledge of all conditions in a catalytic system are necessary for understanding the catalytic chemistry taking place.208 Subsequent work with mononuclear ruthenium wateroxidation catalysts has been extensive. A number of valuable reviews have appeared on the topic of mononuclear catalysts.209−211 Continuing work from the Meyer group has provided mechanistic details on other single-site catalysts,212 including significant work on electrocatalysis. Paul, Papish, and Grotjahn and their co-workers much improved the activity of the mononuclear Ru(terpy)(bpy) catalysts by installing pendant electron-donor ligands such as −OH and −OMe on the bpy unit. An X-ray structure suggests a possible origin of K

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the effect in that a pendant −OH group was found to hydrogen bond to an aqua ligand on the metal.213 One of the main applications of these catalysts lies in photoelectrochemical cells (PECs),214 where the goal is to achieve light-driven water oxidation, although challenges still remain.215 For example, Sun and co-workers216 describe a 3component system with persulfate as electron acceptor, a RuII(bpy)3 derivative as sensitizer, and a series of mononuclear Ru complexes of the Ru(NNN-pincer)(4-picoline)2 type as catalysts. A complete PEC would have a WOC and a photosensitizer at the photoanode and a water-reduction catalyst to make H2 at the cathode. An example using a mononuclear molecular Ru complex as WOC has been reported by Sun and co-workers. 217 With a RuII(bpy)3 derivative as the sensitizer, an initial current density of 1.1 mA/cm2 was achieved. A full review up to 2011 of the field of visible light-driven water oxidation in photoelectrochemical cells is available.218 In an important observation by Meyer and co-workers, chloride was shown to interfere with electrochemical water oxidation by a mononuclear Ru catalyst of the Ru(N,N,Npincer)(bpy) type. An intermediate Ru(V)O was shown to oxidize chloride to HOCl. Once formed, HOCl is further oxidized to O2, potentially complicating interpretation of the data if the chloride effect were to go unnoticed.219 Ru complexes have been widely adopted in the design of sensitizer−catalyst dyads for photodriven water oxidation.220 Meyer and co-workers suggested using a Ru complex in both roles, with a [Ru(bpy)3]2+ unit as the sensitizer and a [Ru(terpy)(bpy)]2+ unit as catalyst. The link is made by introducing an amide between the preformed complexes, the link having a methylene unit to partially isolate the two electronically.221 An unusual [ONO]-pincer ligand, shown in Figure 15, has been used by Åkermark and co-workers to obtain an active Ru-

Some unexpected organometallic ligand sets have proved viable in providing useful catalysts. For example, a carbene Ru complex proved effective in a Ru(terpy) system and showed enhanced reactivity compared to the polypyridyl analogues.212 Similar carbene catalysts also proved useful in a chromophore− catalyst assembly where the presence of the chromophore assisted oxidation of water even when driven by Ce(IV).224 The same catalysts were also successfully attached to a nano-TiO2 surface in a rotating ring-disk electrode (RRDE) electrochemical experiment that confirmed water-oxidation catalysis.225 Another carbene catalyst incorporates the monodentate dimethylimidazolium-2-ylidene ligand with a 2,2′-bipyridine6,6′-dicarboxylic acid coligand for which a water-nucleophilic attack (WNA) mechanism was proposed on the basis of DFT and kinetic data.226 Even “abnormal” NHCs,227 often easily cleaved from the metal, were found by Llobet, Albrecht, and coworkers to be effective. They have shown that a pyridinefunctionalized abnormal triazolylidene ligand bound to Ru makes an effective catalyst for WO. CO2 production was monitored as an indicator to see if there had been any significant degree of ligand decomposition.228 An NHC, normal this time, was also effective for iridium in allowing access to a stable Ir(IV) state in a complex analogous to a Cp*Ir catalyst precursor discussed in the next section.229 Another unexpected ligand set has been reported as an efficient WOC with TOFs up to 77 s−1. In this case, the catalyst is a RuRu-containing complex, [Ru2(OAc)4]. A WNA mechanism is proposed from DFT calculations.230 Oxidative robustness of the ligand set, a key property in water-oxidation catalysis, has been examined by Lau and coworkers with the unexpected result that a Ru quaterpyridine catalyst undergoes oxidation of the pair of wingtip pyridyl groups without loss of water-oxidation catalytic activity. In this case, the N,N,N,N-donor ligand, quaterpyridine, is converted into an O,N,N,O donor ligand, quaterpyridine bis-N-oxide under the reaction conditions.231 Ligand modification under reaction conditions must, therefore, continue to be considered as a possible pathway in water-oxidation catalysis. 2.3. Iridium Catalysts

The use of iridium in molecular catalysts for water oxidation originated with the work of Bernhard and co-workers. In 2008, this group showed that certain iridium(III) complexes, bearing unsubstituted or substituted 2-phenylpyridine (ppy) ligands, were suitable catalyst precursors.232 The compounds have two open sites in the first coordination sphere for binding substrate water to the iridium center; a control experiment with [Ir(ppy)(bpy)]+ showed no activity, confirming a key role for these open coordination sites in binding and activation of substrate water. Oxygen-evolution experiments confirmed essentially stoichiometric consumption of cerium(IV) coupled to generation of oxygen gas that was confirmed by gas chromatography. Analysis of the kinetics of oxygen evolution with these catalysts revealed two regimes. First, a linear region lasting ∼30 min and then, second, a parabolic region indicating progressively more rapid oxygen evolution lasting hours until total consumption of the cerium(IV) oxidant. Notably, the reaction order with respect to iridium precursor in this initial stage was found to be greater than 1, suggesting a role for multiple metal sites in oxygen evolution. Bernhard et al. also noted that the two-phase behavior indicates an in situ alteration of the catalyst and, thus, the possibility of multiple catalytically

Figure 15. [ONO] ligand architecture from Åkermark and co-workers bearing a phenolate donor and a carboxylate donor. Reprinted with permission from ref 222. Copyright 2012 WILEY-VCH Verlag GmbH & Co. KGaA.

based catalyst with 4000 turnovers after 15 min.222 Catalyst deactivation may be associated with the presence of a phenol in the ligand, this being an easily oxidized functional group. This emphasizes the important general point that ligands for water oxidation should be designed to be as oxidatively robust as possible. Another report that includes a phenol unit in the overall ligand design involves a conventional phenylenediamine Schiff base adduct with salicylaldehyde.223 The water-oxidation activity of the Ru complex with Ce(IV) proved satisfactory at a turnover frequency of 0.65 s−1, and activity was maintained for 10 h. The same group also successfully drove a catalyst of the Sun-type electrochemically. More than 15 000 turnovers of water oxidation were seen during 10 h of electrolysis at 1.42 V vs NHE. The low overpotential permitted construction of a light-driven version with a conventional photosensitizer and persulfate as sacrificial oxidant.223 L

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proposed sequential oxidation of the iridium(III) precatalysts to form a key iridium(V) species that is sufficiently electron deficient to undergo nucleophilic attack by substrate water and form the O−O bond, shown schematically in Figure 18.

Figure 16. General structure of the Bernhard-type iridium precatalyst for water oxidation.232 Reprinted with permission from ref 232. Copyright 2008 American Chemical Society.

active species under these conditions. Rate values of ca. 0.1 turnovers min−1 were obtained during this first hour of catalysis. In 2009, our group published the use of pentamethylcyclopentadienyl (Cp*) complexes as molecular precatalysts for water oxidation driven by cerium(IV) as sacrificial oxidant.233 In this work, catalyst precursors were described that bear the Cp* ligand along with a LX-type bidentate chelate ligand, either 2-phenylpyridine or 2-phenylpyrimidine (Figure 17). Notably, rates of oxygen evolution at short times were found on the order of 10 turnovers min−1, significantly faster than other systems known at that time. On the basis of a first-order dependence of oxygen evolution on the concentration of iridium precatalyst, it was noted that single-site catalysis was possible. Along this line, calculations from Eisenstein described the electronic structure of an iridium(V) oxo complex as a possible intermediate, although this species was not observed experimentally. Related [Cp*Ir] complexes were also applied to C−H activation, and stereoretentive alkane hydroxylation activity was observed in chemically driven experiments with cerium(IV).234,235 Our 2010 report,236 as well as a parallel report from Macchioni,237 expanded the number of half-sandwich iridium complexes known to be active precatalysts (Figure 17). Specifically, coordination by Cp* along with an L2-type bidentate chelate (e.g., 2,2′-bipyridine, bpy) was found to give an active catalyst operating at a turnover frequency of around 10 min−1, much as in the case of the LX chelates already studied. Kinetic studies from Macchioni237 showed a first-order dependence of oxygen evolution on iridium concentration in systems with a bidentate chelate in addition to Cp*, in agreement with our own results. On the basis of the observed kinetics of oxygen evolution, as well as further modeling, we

Figure 18. Proposed pathway for oxygen evolution with iridium complexes serving as the catalyst. Reprinted with permission from ref 236. Copyright 2010 American Chemical Society.

We also studied a number of cycopentadienyl (Cp) analogues for comparison and found similar initial rate data under standard conditions with cerium(IV). However, an unusual feature was encountered in the oxygen-evolution rate data obtained for complexes of the type [CpIr(bpy)X]+: at high precatalyst concentrations, the system undergoes a deactivation process and the rate of catalysis tends toward a zero-order behavior. This suggested that the less bulky Cp ligand permits formation of a coordinatively saturated mono-μ-oxo complex at high concentrations of [CpIr(bpy)]. Notably, this behavior contrasts with the [Cp*Ir] analogues, which show a clear firstorder dependence of oxygen evolution on precatalyst concentration. Studies of a Cp*Ir complex with a strongly donating NHCbased [CC] chelating ligand provided initial evidence for the presence of metastable iridium(IV) intermediates with these complexes.229 A reversible oxidation from Ir(III) to Ir(IV) was observed at +210 mV vs Fc+/0 in acetonitrile solvent.

Figure 17. [Cp*Ir] and [CpIr] complexes studied as molecular precatalysts for water oxidation. Reprinted with permission from ref 236. Copyright 2010, American Chemical Society. M

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Complementary studies with [Ru(bpy)3]3+ as chemical oxidant revealed a rhombic signal in EPR spectra at 8.5 K. A closely related example was reported in 2012 using the Co-containing tripodal Kläui ligand, [Co(η5-C5H5){P(O)(OEt)2}3]−.238 Electrochemically reversible oxidation from Ir(III) to Ir(IV) was observed for a complex with three chloride ligands in addition to the tripodal ligand. An axial EPR spectrum was obtained for the isolated material, which was also crystallographically characterized. Consistent with the relative stability of this complex, the iridium(III/IV) midpoint potential is at approximately 0 V vs Fc+/0, considerably less oxidizing than the Ir NHC complex mentioned above.229 Bernhard and Albrecht studied239 bidentate [CN] L2-type and [CC] LX-type carbene-containing complexes of iridium that showed good activity as precatalysts for water oxidation driven with cerium(IV) as the sacrificial oxidant. Analogous work from Macchioni has examined strongly donating [CC] and [NN] NHC-based chelating ligands to iridium.240 Biphenyl complexes of [Cp*Ir] have also been synthesized and studied under oxidative conditions. These complexes showed reversible iridium(III/IV) couples; the formally iridium(IV) complexes could be observed by EPR spectroscopy in some cases.241 A notable precatalyst in the 2010 reports236,237 is the [Cp*Ir(H2O)3]2+ complex, which shows behavior distinct from the many other Cp* or Cp complexes studied. This complex is a highly water-soluble organometallic aqua cation, first isolated by Kölle, and a member of a family of related species with multiple aqua or hydroxo ligands.242 Under the standard assay conditions used (5 μM catalyst, 78 mM cerium(IV)), this trisaqua complex showed a lower rate of oxygen evolution, ∼5 turnovers min−1, than the related complexes with bidentate chelates. However, kinetic studies showed that oxygen evolution on an iridium basis increases at higher precatalyst concentrations. This suggested to us that dimerization or polymerization might contribute to catalysis. Finally, nonlinearity in the conventional log−log plot suggested that equilibrium speciation processes affect the rate of oxygen evolution. At the higher concentrations, turnover numbers of over 20 min−1 were recorded. Results on an iridium basis for [(Cp*Ir)2(OH)3]OH and [Cp*IrCl2]2 are similar, consistent with the expected speciation of these complexes to form [Cp*Ir(H2O)3]2+ under the acidic aqueous conditions used for the assay. The oxidative electrochemical response of [Cp*Ir(H2O)3]2+ is complex and unique243 and indicates deposition of heterogeneous material referred to as “Blue Layer” (BL) upon oxidation of the parent complex (see Figure 19). Thus, BL is the true catalyst in electrochemically driven water oxidation starting with the [Cp*Ir(H2O) 3]2+ precursor complex. This represents a relatively unusual and clear case in which a well-defined coordination complex gives rise to a heterogeneous material that functions as the true catalyst for water oxidation. We, thus, investigated this process in detail and the resulting heterogeneous material in the context of developing molecular water-oxidation catalysts and working toward a better understanding of the interplay between homogeneous and heterogeneous catalysis. Anodic cycling of a solution containing [Cp*Ir(H2O)3]2+ results in the appearance of a new feature centered near 0.9 V vs NHE, which grows in intensity upon excursions in potential beyond ca. 1.2 V. The new feature is preserved in voltammograms collected after rinsing the electrode and transferring to a blank solution of electrolyte (free from the precursor

Figure 19. Cyclic voltammetry of BL (blue line) and basal-plane graphite electrode background (black line) in 0.1 M KNO3 at pH 6. The catalytic wave corresponding to oxygen evolution onsets at ca. 1.1 V. (Inset) Cyclic voltammogram (successive scans) showing deposition of BL from a solution containing [Cp*Ir(H2O)3]2+. The quasi-reversible peak is centered at 0.88 V (ΔEp = 15 mV). Deposition conditions: 2.3 mM [Ir], 0.1 M KNO3, pH 2.9; scan rate 50 mV s−1. Reproduced with permission from ref 243. Copyright 2011 The Royal Society of Chemistry.

[Cp*Ir(H2O)3]2+ complex). The peak currents of this redox feature show a linear dependence on scan rate, consistent with a nondiffusional species that is bound to the surface of the electrode. Consistent with this observation, more extensive oxidative electrolysis on a transparent conducting electrode gives visible deposition of a blue layer of heterogeneous material on the electrode (BL). This deposited material is an exceptionally robust and highly active heterogeneous catalyst that shows virtually quantitative oxygen evolution based on charge transferred. It has been applied to a hematite electrode for catalysis in light-driven water oxidation.244 The deposited material is amorphous, showing only a broad, low-angle feature in X-ray diffraction data. Surprisingly, BL contains a significant amount of carbon (∼9% by mass following deposition) which has its origin in the Cp* ligand of the precursor complex.245 The Cp* ligand is necessary for deposition, as other complexes bearing chelating ligands in addition to Cp* do not show deposition behavior. IR spectra include features consistent with CO and C−O functionalities not present in the starting material, suggesting that the incorporated carbon is highly oxidized. X-ray scattering data collected for BL suggest that small di-μ-oxo-bridged clusters containing ∼5 iridium atoms dominate the amorphous structure.246 Consistent with all these data, characteristic features in the X-ray scattering data for [Cp*Ir(H2O)3]2+ complex arising from scattering between the iridium center and the Cp* methyl groups are totally absent in analogous data collected for BL.245 Consistent with this observation, the activity of BL for water oxidation is stable over many hours of electrolysis, despite the almost total loss of carbon content from the catalyst material over this same time period.245 Thus, Cp* is extensively modified and lost from the sites that are active for water oxidation in BL. N

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tions.253 A possible involvement of cerium oxides in promoting formation of iridium oxide was discussed, and formation of nanoparticulate oxides by TEM was reported in the case of some precatalyst derivatives. Formation of characteristic visible absorption bands for many catalysts in the range of 550−650 nm was ascribed to formation of iridium oxide, although subsequent studies have shown that these features can also arise from molecular species (see below). Consistent with a range of active species in catalysis, sequestration of molecular iridium species in solution has been suggested to explain the observed inhibition of catalysis upon addition of encapsulating cucurbit[n]urils.254 Beller studied255 the nature of the true catalyst beginning with a number of [Cp*Ir] precatalysts and implicated molecular species as catalysts. In situ X-ray absorption spectroscopy, along with other techniques, also suggested that iridium oxide particles were present in reaction mixtures under certain conditions and at longer times. Under the chosen conditions, which relied on cerium(IV) as chemical oxidant, an immobilized form of iridium dioxide on MCM-41 mesoporous silicate did not show improved activity over molecular precursors. In parallel with these studies on soluble, molecular precatalysts, Lin developed metal−organic frameworks,256,257 one of which is shown in Figure 21, that contain immobilized

Macchioni et al. investigated the possibility of ligand degradation of a number of Cp*Ir complexes in the presence of cerium(IV) as chemical oxidant247 or hydrogen peroxide as a chemical oxidant.248 In situ NMR experiments showed that treatment with cerium(IV) resulted in oxidative degradation of the Cp* ligand and formation of acetic and formic acids. Oxidation of a quaternary carbon in the Cp* ring system, as well as a Cp* methyl group, were implicated in these studies. In the analogous work with hydrogen peroxide as chemical oxidant, initial attack on a quaternary carbon was followed by oxidation of a methyl group, as studied by NMR spectroscopy. Gas-evolution experiments showed onset of oxygen evolution along with a small amount of carbon dioxide in some cases, consistent with ligand oxidation. In one notable study249 examining a complex bearing the 2-benzoylpyridine (bzpy) ligand, Macchioni’s group was able to isolate and crystallize a Cp*-modified analogue of the precursor (Figure 20). This

Figure 20. Molecular structure of Macchioni’s Cp*-activated iridium complex.249 For clarity, hydrogen atoms are shown only on the alcohol moieties that result from oxidative activation of the Cp* ring. Colors: iridium, magenta; nitrogen, blue; oxygen, red. Thermal ellipsoids are shown at 10% probability. Figure generated from CCDC entry number 971168. Figure 21. Lin’s MOF-immobilized [Cp*Ir] complex (left) based on a modified 2-phenylpyridine complex (right). Reprinted with permission from ref 256. Copyright 2011 American Chemical Society.

complex has undergone double functionalization of the Cp* ring, resulting in a pseudo-octahedral iridium(III) complex. Considering all their observations, the authors concluded that no modification of the bzpy ligand occurred, but modification of the ligation of Cp* resulted in complexes competent for water oxidation driven by chemical oxidants. They also noted that less oxidatively robust chelate ligands would be more likely to undergo modification and possibly result in heterogeneous oxide-based catalysis. In contrast to Macchioni’s work showing modification of Cp*, low-temperature studies of a [Cp*Ir] complex of 2phenylpyridine with oxygen atom-transfer reagents suggested that ligand oxidation readily occurs under the chosen conditions.250 O-atom insertions into Ir−C or Ir−N bonds of a 2-phenylpyridine ligand were also observed under some conditions with [Cp*Ir] complexes.251 Similarly, O-atom insertion into the Rh−C bond of coordinated 2-phenylpyridine was found in studies with a [Cp*Rh] complex.252 Grotjahn also studied a number of [Cp*Ir] complexes and implicated involvement of catalytically active iridium oxides in the observed oxygen-evolution activity under some condi-

analogues of the previously investigated Cp*Ir(2phenylpyridine)Cl precatalyst for water oxidation.236 Initial results on these MOF-immobilized species showed that oxygen was evolved with cerium(IV) as the sacrificial oxidant, although activity was less than in solution due to steric difficulty associated with cerium(IV) accessing the active sites in the MOF. In their later study,257 the immobilized derivative of the [Cp*Ir(bpy)Cl]Cl was studied in greater detail. Following treatment with cerium(IV), generation of oxygen was confirmed, and a new species, [Ir(bpy)(X)(Cl)(OH2)2], was implicated in the observed catalysis by a variety of spectroscopic techniques. Decomposition of the Cp* ligand to acetate and formate, among other products, was observed. This work on [Cp*Ir] complexes, taken together, highlights the observation that molecular water-oxidation catalysts are susceptible to degradation over time due to side reactions that modify or destroy their ligands. However, an additional concern O

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Figure 22. Comparison of current and mass responses of solutions containing the [Cp*Ir] complexes shown as a function of time. (Top) Gray solid lines represent applied potential; black solid lines represent current response recorded by the potentiostat. Dashed lines are background current. (Bottom) Solid lines represent the mass response in each experiment, and dashed lines are the corresponding catalyst-free backgrounds. Conditions: 0.1 M KNO3 in air-saturated deionized water; 3 mM 1; ca. 2.5 mM 3. Scan rate: 50 mV/s. Reprinted with permission from ref 258. Copyright 2011 American Chemical Society.

(∼150 ng/cycle). In the case of the pyridine alkoxide complex, no such change in mass was detected, consistent with an origin of catalysis solely arising from soluble iridium species. Thus, the EQCN is a suitable tool for distinguishing between heterogeneous film catalysis and homogeneous catalysis at short times.260 However, as noted in the original report, the EQCN cannot provide data regarding the formation of modified catalytic molecules or particles that are soluble and resistant to electrodeposition. With the EQCN method, we also compared the electrodeposition characteristics of [Cp*Ir(H2O)3]2+ to its cyclopentadienyl analogue, [CpIr(H2O)3]2+.261 Additionally, the electrodeposition properties of small, iridium dioxide nanoparticles262 and [Ir(OH)6]n+ were surveyed263 as common precursors to electrodeposited iridium oxide materials. Notably, the CpIr complex deposited twice as much heterogeneous material on each cycle of voltammetry as the Cp*Ir analogue. In addition, the nanoparticles were not irreversibly deposited but rather were removed from the electrode upon voltage excursion below +600 V vs NHE at pH 14. Finally, annealing the amorphous electrodeposited materials to crystalline IrO2 resulted in loss of all quasi-reversible redox features arising from the deposited layer and a marked increase in the overpotential from 240 to 490 mV. These results highlight the remarkable activity of amorphous iridium oxides and also the difficulty of

is the generation of secondary materials that are responsible for observed oxygen-evolution activity. This is of special concern in the area of iridium catalysts, because iridium oxides are among the best catalysts known for oxygen evolution, especially under acidic conditions.51 As an additional complication, techniques such as microscopy and catalyst isolation are hampered by the dilute aqueous conditions and high salt concentrations often used for studies of water oxidation. In 2011, our group described use of an electrochemical quartz crystal nanobalance (EQCN) to address this challenge.258 After observation of the anodic electrodeposition of heterogeneous and highly active catalyst material from solutions containing the [Cp*Ir(H2O)3]2+ complex, we were interested in distinguishing between heterogeneous surface catalysis and homogeneous, solution-phase catalysis. The EQCN is an ideal instrument for this purpose as it provides an accurate real-time mass of the working electrode during electrochemical studies (i.e., it is a piezoelectric gravimetric method).259 In our initial study with the EQCN, we compared the electrochemical properties of two water-soluble iridium complexes, the [Cp*Ir(H2O)3]2+ complex mentioned before, and a [Cp*Ir] complex bearing 2-(2′-pyridyl)-2-propanolate ligand as an oxidation-resistant, LX-type bidentate chelate. The results are summarized in Figure 22. In the case of the tris-aqua complex, excursion to oxidizing potentials results in a large pseudocatalytic current and a clear increase in mass of the electrode P

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been found to bind strongly to metal oxide surfaces and maintain high water-oxidation activity without loss or degradation of the bidentate (2′-pyridyl)-2-propanolate ligand.272 Thus, observations from studies of water-oxidation driven either electrochemically or chemically strongly suggest that the presence of an oxidatively stable chelate ligand in addition to Cp* is key for molecular catalytic activity in this family of catalysts. Numerous iridium complexes have been synthesized and tested for catalytic activity since the first reports of molecular iridium precatalysts. Macchioni’s group studied [Cp*Ir] derivatives with carboxylate ligation to the iridium center, including complexes of ethylenediamine tetraacetic acid (EDTA)273 and several bidentate ligands.274 Crabtree’s group investigated a number of monomeric and dimeric complexes to investigate the possibility of oxo-coupling pathways in oxygen evolution.275,276 Notably, the doubly chelating bispyridine− pyrazolide ligand was used to closely space two iridium centers, but no enhancement of catalysis was observed versus the appropriate mononuclear analogue as apparent in the electrochemical data shown in Figure 23. A dinuclear iridium(I)

observing even small amounts of crystalline iridium oxide by electrochemical methods. In the case of water oxidation or C−H oxidation264 driven with chemical oxidants, the possibility of particle formation from molecular precatalysts was also investigated.265 Dynamic light scattering (DLS) can detect particles down to nanometer scales in aqueous solutions and is an especially attractive method for such investigations because it can be used on working catalyst solutions. In the case of oxidations driven with sodium periodate, particle formation was not observed from [Cp*Ir] complexes with oxidation-resistant bidentate chelate ligands 2,2′-bipyridine (bpy), 2-phenylpyridine, or 2-(2′pyridyl)-2-propanolate. In the cases of bpy and the pyridine alkoxide ligands, available data from EQCN studies confirmed no electrodeposition upon electrochemical oxidation of these complexes. Hintermair et al. also concluded that the blue color of catalyst solutions (band at λmax ≈ 580 nm) after catalysis has ceased is not related to particle formation. Macchioni also obtained blue solutions following catalysis and implicated molecular species.266 A notable result confirming that highly colored species are not necessarily particulate in nature comes from a recently isolated complex that shows interconversion of Ir−OH and Ir−O−Ir units.267 Studies by advanced mass spectroscopic techniques have also investigated the species present when using [Cp*Ir] complexes as precatalysts. An initial report from Johnson examined spectral features of [Cp*Ir] complexes ligated by 2-phenylpyridine or 2-(2′-pyridyl)-2-propanolate.268 Mass spectroscopic injection of species from acetonitrile showed spectral features consistent with Cp*-containing compounds. A later report from Crabtree, Johnson, and Zare in 2014,269 however, implicated several activation modes of the Cp* ligand. The aqueous reaction mixture obtained upon treatment of the Cp*Ir complex bearing 2-(2′-pyridyl)-2-propanolate with periodate as chemical oxidant was investigated from milliseconds to seconds using desorption electrospray ionization, electrosonic spray ionization, and cryogenic ion vibrational predissociation spectroscopy. The data obtained implicate initial electrophilic C−H hydroxylation of a Cp* methyl group. In recent work, compelling evidence has been provided that Cp* is a sacrificial placeholder ligand in the family of wateroxidation catalysts with chelate ligands in addition to Cp*.270 In accord with Macchioni’s prior observations under related conditions, NMR and mass spectrometry showed a rapid and irreversible loss of Cp* in chemically driven water oxidation. When oxidatively stable chelate ligands are bound to iridium, solutions with a characteristic blue color remain active for water oxidation (or C−H oxidation) without an induction period. Electrophoresis results suggested the presence of well-defined iridium cations, and a variety of additional spectroscopic techniques (including 17O NMR) implicate the structure of the catalyst to a di-μ-oxo-bridged dimer of iridium(IV). Consistent with the proposal that Cp* is a sacrificial placeholder ligand, studies of cyclooctadiene (cod) complexes of iridium(I) showed very similar kinetic profiles for water oxidation in comparison with their Cp*IrIII analogues.270 Moreover, studies of extensive electrolysis of solutions containing [Cp*Ir] precursor complexes show that blue solutions of activated catalyst are produced, and these blue solutions show higher yields of oxygen in electrochemically driven water oxidation versus their unactivated [Cp*Ir] analogues.271 The activated catalyst in the blue solution formed from the Cp*Ir 2-(2′-pyridyl)-2-propanolate complex has also

Figure 23. Electrochemical response of mononuclear and dinuclear iridium complexes used to interrogate the viability of an oxo-coupling pathway in oxygen evolution.276 Conditions: 1 mM [Ir] in 0.1 M aqueous KNO3 electrolyte (pH 4) measured at 100 mV/s at room temperature. Reprinted with permission from ref 276. Copyright 2013 American Chemical Society.

complex has also been investigated,277 but mononuclear analogues were not compared, and a pronounced lag phase in catalysis under certain conditions suggests the possibility of extensive ligand modification. Related Me2NHC complexes have been investigated278 for their activity in both electrochemically and chemically driven water oxidation. High activity was obtained using periodate as the chemical oxidant,279 although there has been debate about the possibility of O-atom transfer reactivity from periodate.280 A Me2NHC system adsorbed on a polycrystalline gold surface was studied by a number of techniques and showed electrochemical activity.281 Albrecht and Bernhard also studied [Cp*Ir] precatalysts bearing monodentate carbene ligands for use in light-driven water oxidation.282 Dicarbene complexes have also been studied, showing EPR signals from iridium(IV) in reaction mixtures under certain conditions.283 Carbene iridium complexes have been modeled Q

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in computational studies of possible mechanistic pathways.284 Carbene iridium complexes were also appended to scaffolds containing a distant terpyridine ligand as a strategy for building up more complex supramolecular constructs.285 Iridium complexes with [N]- or [C]-donating carbene ligands have been investigated for C−H oxidation activity.286 A large group of substituted 2,2′-bipyridyl complexes has been studied. Papish’s group studied a number of hydroxysubstituted bipyridine complexes and found that deprotonation of this ligand resulted in more rapid oxygen evolution.287 Fukuzumi implicated288 particle formation in catalysis with hydroxy-appended bipyridine complexes, and related studies were carried out by Fujita and Himeda on a family of related hydroxy-substituted bipyridine complexes.289 A study comparing a number of bipyridine derivatives with electron-donating or -withdrawing groups as ligands suggested a role for heterogeneous iridium material in the observed catalysis in some cases.290 Bipyridine complexes of iridium have also recently been modeled by Siegbahn291 to computationally compare a number of oxidation states and their competency for oxygen evolution. Specifically, oxo species of Ir(V), Ir(VI), and Ir(VII) were modeled on iridium centers that do not bear the Cp* ligand. Cp*Ir complexes have also been applied for modification of electrodes and photoelectrodes for water-oxidation catalysis. A [Cp*Ir] complex was codeposited with a porphyrin dye for use in light-driven water oxidation,292 but as in other cases, the nature of the active species is not clear.293,294 Additionally, these surface-attached systems show only modest performance. Covalently linked chromophore−catalyst electron-transfer dyads have been synthesized with a [Cp*Ir] fragment and perylene-3,4:9,10-bis(dicarboximide) (PDI) photooxidant. In one case,295 excitation of the PDI unit resulted in formation of multiple populations of the PDI•−−Ir(IV) ion pair, and biexponential charge recombination largely to the PDI−Ir(III) ground state was measured. In a later study, a more elaborate molecular triad system was prepared296 with a [Cp*Ir] fragment, perylene-3,4-dicarboximide chromophore, and a naphthalene-1,8:4,5-bis(dicarboximide) (NDI) electron acceptor. In this system, longer-lived charge separation was achieved by formation of the NDI•−−PMI−Ir(IV) state with a longer ion pair distance versus the dyad system studied initially.295 A notable time-resolved X-ray absorption experiment carried out on this triad compound at the Ir LIII X-ray absorption edge demonstrated light-driven photooxidation of the iridium center by monitoring a shift of the white line signal from 11 212.5 to 11 215.9 eV at 100 ps after excitation of the chromophore (see Figure 24).

Figure 24. (Left axis) XANES spectra at the Ir LIII edge of the NDI− PMI−Ir catalyst triad ground state (black) and photoexcited state at 100 ps following photoexcitation (red). (Right axis) Difference between the photoexcited-state and the ground-state spectra (blue). Arrows indicate changes between ground and photoexcited states in the absorption edge and the white line at 11212.5 and 11215.9 eV, respectively. The NDI−PMI−Ir catalyst compound is shown below the data. Reproduced with permission from ref 296. Copyright 2013 The Royal Society of Chemistry.

the primary oxidant. The least oxidatively stable TAML catalyst was ineffective, presumably because of ligand oxidation, but four other variants proved successful. Improved catalyst performance was associated with increasingly electron-withdrawing substituents on the TAML. The best catalyst gave an initial turnover frequency of more than 1.3 s−1, although a slower phase takes over after a few tens of seconds, as seen in Figure 25, which then continues for hours. Early work by Elizarova and co-workers on iron tetrapyrrole catalysts300 gave similar activities, but the catalyst lifetimes were very short, no doubt as a result of ligand degradation. In an electrocatalytic version of the TAML water-oxidation reaction,

2.4. Iron Catalysts

Iron-based catalysts for water oxidation have recently been reviewed in detail by Singh and Spiccia.297 The story has its origin with Collins’s work from 1980 on the development of robust tetraamido macrocyclic ligands (TAMLs) for iron.298 These were initially developed for organic oxidation catalysis with H2O2 as the primary oxidant. The key points emerging from this work were the use of very strongly donor deprotonated amide ligands to allow access to high oxidation state iron and the careful, stepwise elimination of oxidatively unstable substructures in the TAML itself, culminating in a series of extremely robust macrocycles. Their application to water-oxidation catalysis came in a 2010 publication in collaboration with Bernhard,299 describing work with CAN as

Figure 25. (A) FeIII−TAMLs used for water oxidation: 1, X1 = X2 = H, R = CH3; 2, X1 = X2 = H, R = (CH2)2; 3, X1 = X2 = H, R = F; 4, X1 = NO2, X2 = H, R = F; 5, X1 = X2 = Cl, R = F. Y = H2O. (B) Plots of O2 evolution with time from 1 to 5 (0.6 μmol) upon addition of cerium(IV) (145.7 μmol) in unbuffered water (total volume 0.8 mL). Reprinted with permission from ref 299. Copyright 2010 American Chemical Society. R

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case, simple Fe salts proved to be even more active than the complex, and the real catalyst for both the salt and the complex was thought to be α-Fe2O3 nanoparticles, formed under the reaction conditions. DLS, energy-dispersive X-ray spectroscopy, and X-ray photoelectron spectroscopy helped identify the nature of the nanoparticles. In addition, authentic α-Fe2O3 nanoparticles gave similar water-oxidation activity under the photochemical conditions. Llobet, Fukuzumi, and co-workers312 looked at related Ce(IV)-driven WOCs in which the tetradentate ligand is equipped with oxidatively stable 8-quinolyl groups that also provide enhanced rigidity (Figure 27). The aliphatic backbone

an immobilized TAML system (X1 = X2 = H; R = Me; Y = OH2) on glassy carbon produces O2 with much higher TONs than the homogeneous predecessors.301 In this case, a FeIII/IV wave was seen at 720 mV (vs Ag/AgCl) prior to the onset of dioxygen evolution at ∼1 V. The Faradaic efficiency was ∼45%, no doubt due to parasitic oxidation of the carbon in the carbon paper electrode. Theoretical analysis of the TAML catalysts by Cramer and co-workers302 with density functional theory and multireference second-order perturbation theory suggests the intermediacy of a Fe(V) oxo with a 1e −-oxidized macrocycle. Waternucleophilic attack was identified as the O−O bond-forming step. In more recent work, Liao and colleagues come to a similar conclusion but add nitrate attack on the oxo as a pathway having a reasonable barrier.303 They also identify opening of the arene ring as the most probable pathway for ligand degradation. Gupta, Dahr, and co-workers304 modified the TAML system to allow photochemical oxidation with persulfate as the sacrificial oxidant and [Ru(bpy)3]2+ as the photosensitizer. The maximum TON was 220 with a TOF of 0.76 s−1 and a Fe(V)−oxo intermediate was proposed. Fillol, Costas, and co-workers tested a series of tetraaza complexes305 of Mn, Fe, Co, and Ni for water-oxidation catalysis either with Ce(IV) or NaIO4 as primary oxidants. Good activity required both Fe being the metal and the ligand being compatible with the presence of two adjacent labile sites. In the case of one compound with mixed aryl and alkyl L-type [N] donor ligands, for example, NaIO4 proved to be the best primary oxidant, giving TONs of >1050 and initial TOFs of 222 h−1. Dynamic light scattering (DLS) and nanoparticletracking analysis (NTA) studies eliminated metal oxide particles as a possible alternative thermal catalyst. A waternucleophilic attack mechanism was proposed. Theoretical work has appeared306,307 on one of the Fillol− Costas catalysts, shown in Figure 26. These reports agree on

Figure 27. Iron-based catalyst structures from Llobet and Fukuzumi. Reprinted with permission from ref 312. Copyright 2013 American Chemical Society.

was still vulnerable to oxidation, however, leading to limited catalyst lifetimes with TONs up to 80. Photochemical water oxidation was attempted with [Ru(bpy)3]2+ and persulfate, but in this case, the true catalyst was ∼100 nm iron hydroxide nanoparticles. A recent review by Fukuzumi and Hong313 examines the evidence for homogeneous versus heterogeneous catalysis in water-oxidation catalysis; another recent review covers methods that have been used to make this distinction for catalysts in general.314 An iron complex, studied by Najafpour and co-workers,315 was based on a different tetradentate N-donor ligand, tpa (tpa = tris(2-pyridylmethyl)amine). In this case, the water-oxidation catalyst precursor was the known mu-oxo Fe(III) dinuclear complex, [tpa(H2O)FeOFe(H2O)tpa](ClO4)4, successfully driven with Ce(IV) as primary oxidant. To examine the possibility that the Fe(VI) ferrate ion, [FeO4]2−, was the true catalyst, formed after loss of tpa, a ferrate solution was added to the Ce(IV) oxidant. Immediate O2 evolution took place as a result of the low pH of the Ce(IV) solution, but this was only stoichiometric, not catalytic, so ferrate is less likely to be the true catalyst, at least under these conditions. Parent, Sakai, and co-workers316 have shown that the tpa catalyst and the related 1-(bis(2-methylpyridyl)amino)-2-methyl-2-propanoate catalyst can be driven by NaIO4. The latter complex showed simpler kinetics, while studies on the former were complicated by the presence of more than one active species in the catalytic solutions. Li, Sun, and co-workers317 screened a diverse series of 11 isolated iron complexes and a further 9 complexes generated in situ and found two complexes with wateroxidation activity comparable to the best prior catalysts. Meyer’s group also examined one iron complex in an electrochemically driven system.318 Yang and co-workers319 looked into the effect of introducing hydrogen-bonding proton acceptors adjacent to the presumed oxo-binding site in tetraaza Fe complexes. Trials were run with

Figure 26. Fillol−Costas iron-based catalyst for water oxidation. Reprinted with permission from ref 307. Copyright 2012 American Chemical Society.

the viability of the water-nucleophilic attack mechanism, although disagreeing on the details of the intermediates involved. The latest work in this area308 posits access to a Fe(V)−oxo via a PCET pathway with an OH ligand located cis to the oxo site acting as an internal base to help deprotonate the attacking water molecule. Close agreement between the theoretical (18.9 kcal/mol) and the experimental (17.6 ± 1.6 kcal/mol) barriers was achieved. Some authors have also argued from theoretical work that an oxo-coupling pathway for O2 formation can be accessible with iron complexes309 or from ferrate ion.310 Lau and co-workers311 re-examined the Fillol−Costas cyclohexane diamine catalyst and found comparable results for the Ce(IV)-driven process at low pH but came to different conclusions on the photochemical version at high pH. In this S

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CAN at pH 0.7 and sodium periodate at pH 4.7. Some complexes had significant water-oxidation activity (TOFs up to 140 h−1), but others gave problems.

Berlinguette has done careful studies in the case of a related mononuclear cobalt complex that electrocatalytically evolves oxygen from water.346,347 pH-dependent electrochemical data in this system were assembled into a Pourbaix diagram, suggesting molecular catalysis. Further interrogation of the electrochemical data revealed a concerted proton-coupled electron transfer during oxidation of this complex.348 Tetranuclear Co4O4 cubane compounds, such as the one shown in Figure 28, have recently been discussed as molecular

2.5. Cobalt Catalysts

There has been great interest in developing molecular cobalt catalysts for water oxidation. For example, a cobalt phthalocyanine320 and a fluorinated cobalt corrole321 have been suggested as heterogenized homogeneous catalysts. Stahl et al. reported a dinuclear cobalt complex as a catalyst that is functional at pH values normally inaccessible to cobalt oxides.322 Cobalt porphyrins have also been studied as catalysts.323,324 However, because suspensions containing simple cobalt salts, phosphate buffer, and primary oxidants are competent for oxygen evolution,325−327 distinguishing between true homogeneous catalysis and heterogeneous catalyst formation followed by oxygen evolution is difficult. Hill et al. reported a fast, soluble, and homogeneous cobalt water-oxidation catalyst supported by polyoxotungstate (POM) ligands.328 The catalyst can be triggered with light in the conventional Ru(bpy)3−persulfate assay.329,330 This was an encouraging result, because the older ruthenium-based POM complexes from Hill et al.331,332 and Bonchio et al.333,334 and an iridium-based POM system from Hill et al.335 for water oxidation were less active. The reported turnover frequencies for the Co−POM complex were impressive, reaching above 5 turnovers s−1 at pH 8. A more recent paper reports even higher numbers.336 Building on these results, other POM-based cobalt catalysts have been reported 337 and used in diverse conditions.338 The POM catalyst area has now been reviewed in detail.339 However, at the time of the original publication, little mechanistic detail was available, and the electrochemistry of the Co−POM complex showed only a strong catalytic wave in electrochemical studies with no metastable intermediates that could be characterized. In response, Finke and co-workers reported340 that the Co−POM catalyst is not stable under specific electrochemical conditions and catalysis results from a small amount of cobalt(II) ions in solution. Once these ions are oxidized, a nearly undetectable but nonetheless catalytically viable amount of cobalt oxide is formed. In the Finke study, only 4.3% of the UV−visible absorbance of the Co−POM was lost over 3 h of continuous monitoring before catalysis was initiated. This aging process gave sufficient free ions, however, to quantitatively account for all of the observed catalytic water oxidation within the 12% experimental error of their measurements. After sufficient electrolysis, cobalt oxide is visible on the electrode surface by microscopy. Since the surface layer is completely free of tungsten, and in concert with the related experiments detailed above, Finke et al. concluded that the Co−POM is merely a precatalyst which forms cobalt oxide in situ.340 Follow-up reports provide strong evidence that, under certain other conditions in which chemical oxidants are used, the observed oxygen evolution arises from the solution phase, rather than heterogeneous particles.341 The same group has made suggestions regarding interpretation of kinetics data in chemically driven cases.342 Further work from Finke and coworkers has examined the same system under various conditions, finding evidence for either homogeneous or heterogeneous catalysis, depending upon the chosen conditions.343−345

Figure 28. Molecular structure of [Co4O4] complex investigated as a water-oxidation precatalyst. Reprinted with permission from ref 355. Copyright 2014 American Chemical Society.

catalysts for water oxidation. Dismukes et al. reported349 catalytic water oxidation from a coordinatively saturated cobalt cubane in solution with chemical oxidants, while Britt et al. found350 that the same complex gave no yield of dioxygen when oxidized electrochemically. Oxygen evolution from other molecular Co4O4 systems has been published.351−354 More recently, Nocera et al. found that catalysis in this system arises from cobalt(II) impurities.355 A variety of techniques, including NMR, EPR, and electrochemistry, support this assignment. 2.6. Other Catalysts

Copper coordination complexes have recently attracted attention as catalysts for water oxidation. Mayer and coworkers described the first catalyst based on copper as a bipyridine-based system.356 The catalyst was operated electrochemically in basic media, and although a relatively large applied overpotential (∼750 mV) was needed to achieve catalysis, a high turnover frequency of ∼100 s−1 was estimated. The available data also strongly suggested that the catalyst is a soluble species ligated by bpy and hydroxide. Hydroxysubstituted bipyridine ligands have also been used with copper to generate analogous catalysts for water oxidation.357,358 A lower overpotential for catalysis was estimated for the hydroxysubstituted analogues versus the simple bipyridine-based system. Meyer’s group explored a number of copper systems for water oxidation. A copper(II) complex with a polypeptidebased ligand has been reported as a catalyst,359 although simple copper salts have been found to be catalytically active.360,361 Coordination of buffer base was postulated to prevent T

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Biographies

precipitation of heterogeneous catalysts, and rate enhancements have been reported with certain bases.362 Other copper compounds have also been shown to be active precatalysts for water oxidation.363,364 Aqueous copper has been studied computationally in bicarbonate solutions as a catalyst for water oxidation.365 Importantly, copper oxide has been shown to serve as a heterogeneous catalyst,366,367 suggesting that further mechanistic work is required to better understand the implicated molecular activity of other systems. In other work, Glusac’s group studied metal-free, organic compounds as catalysts for water oxidation.368,369 Specifically, certain flavin derivatives have been found to be capable of mediating oxygen evolution at glassy carbon and platinum electrodes but not on conducting fluorine-doped tin oxide electrodes. Catalysis thus involves heterogeneous intermediates that rely the electrode itself to form. Very high potentials are needed for catalysis (+1.9 V vs NHE), and relatively low turnover numbers of oxygen (∼13) have been shown to date.

James Blakemore was an undergraduate student at Wichita State University and completed his B.S. degree in Chemistry and B.A. degree in Spanish in 2007. He then enrolled at Yale University, where he completed his Ph.D. degree in Chemistry on the topic of wateroxidation catalysis in 2012 under the direction of Gary Brudvig and

3. OUTLOOK AND CONCLUSIONS

Robert Crabtree. In recognition of his doctoral work, he was awarded

This review summarizes work on synthetic and mechanistic inorganic chemistry aimed at achieving well-defined systems for oxidation of water to dioxygen. Manganese and ruthenium complexes are perhaps the most well-known molecular systems for water oxidation, but relatively new efforts have targeted development of iridium-, iron-, cobalt-, and copper-based systems. In all these cases, it is interesting to note that heterogeneous metal oxides based on these elements are known to show activity for water oxidation, perhaps prefiguring their use in molecular catalysts. Due to the ability of the analogous metal oxides to catalyze water oxidation, catalysts of manganese, iridium, and cobalt have recently attracted major scrutiny in a search for evidence of true molecular reactivity. The high activities observed in some systems are fundamentally exciting but also encourage further work in improved mechanistic understanding. With this understanding in hand, the community can advance toward more practical, higher efficiency devices for fuel generation from renewable sources.

the Richard Wolfgang Prize. James then moved to Caltech to take up a position as a postdoctoral scholar in Harry Gray’s group and a Fellow of the NSF Center for Chemical Innovation in Solar Fuels. In 2014, he was named a Prize Postdoctoral Fellow of the Resnick Sustainability Institute at Caltech.

Bob Crabtree, educated at New College Oxford with Malcolm Green, did his Ph.D. work with Joseph Chatt at Sussex University and spent 4

AUTHOR INFORMATION

years in Paris in Hugh Felkin’s lab at the CNRS Natural Products

Corresponding Authors

Institute, then headed by Derek Barton. In 1977 he became Assistant

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

Professor at Yale, where he is now Whitehead Professor of Chemistry. He has received the ACS and Royal Chemical Society prizes for

Present Address

organometallic chemistry and is a Fellow of the RSC, ACS, and

†

James D. Blakemore: Beckman Institute, Division of Chemistry and Chemical Engineering, and Resnick Sustainability Institute, California Institute of Technology, MC 139-74, Pasadena, California 91125, United States.

American Academy. He has also been Dow Lecturer (Berkeley), Williams Lecturer (Oxford), Centenary Lecturer (RSC), and Osborn Lecturer (Strasbourg). He is the author of The Organometallic

Notes

Chemistry of the Transition Metals and has long been involved in

The authors declare no competing financial interest.

collaboration with Gary Brudvig in water-oxidation catalysis. U

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(12) Brudvig, G. W. Water Oxidation Chemistry of Photosystem II. Philos. Trans. R. Soc. London, Ser. B: Biol. Sci. 2008, 363, 1211−8 (discussion 1218−9). (13) Moore, G. F.; Brudvig, G. W. Energy Conversion in Photosynthesis: A Paradigm for Solar Fuel Production. Annu. Rev. Condens. Matter Phys. 2010, 2, 303−327. (14) Shinopoulos, K. E.; Brudvig, G. W. Cytochrome b559 and Cyclic Electron Transfer within Photosystem II. Biochim. Biophys. Acta, Bioenerg. 2012, 1817, 66−75. (15) Berg, J. M.; Tymoczko, J. L.; Stryer, L. Biochemistry, 6th ed.; W. H. Freeman and Co.: New York, 2007. (16) McEvoy, J. P.; Brudvig, G. W. Water-Splitting Chemistry of Photosystem II. Chem. Rev. 2006, 106, 4455−4483. (17) Armstrong, F. A. Why Did Nature Choose Manganese to Make Oxygen? Philos. Trans. R. Soc. London, Ser. B: Biol. Sci. 2008, 363, 1263−1270. (18) Kawakami, K.; Umena, Y.; Kamiya, N.; Shen, J.-R. Structure of the Catalytic, Inorganic Core of Oxygen-Evolving Photosystem II at 1.9 Å Resolution. J. Photochem. Photobiol. B: Biol. 2011, 104, 9−18. (19) Suga, M.; Akita, F.; Hirata, K.; Ueno, G.; Murakami, H.; Nakajima, Y.; Shimizu, T.; Yamashita, K.; Yamamoto, M.; Ago, H.; Shen, J.-R. Native Structure of Photosystem II at 1.95 Å Resolution Viewed by Femtosecond X-ray Pulses. Nature 2015, 517, 99−103. (20) Haumann, M.; Liebisch, P.; Mueller, C.; Barra, M.; Grabolle, M.; Dau, H. Photosynthetic O2 Formation Tracked by Time-Resolved Xray Experiments. Science 2005, 310, 1019−1021. (21) Joliot, P. Reaction Kinetics of Coupled Photosynthetic Oxygen Evolution. Biochim. Biophys. Acta, Biophys. Incl. Photosynth. 1965, 102, 116−134. (22) Joliot, P. Kinetics of Photosynthetic Fluorescence Induction in Relation to Oxygen Evolution. Biochim. Biophys. Acta, Biophys. Incl. Photosynth. 1965, 102, 135−48. (23) Kok, B.; Forbush, B.; McGloin, M. Cooperation of Charges in Photosynthetic Oxygen Evolution. I. A Linear Four Step Mechanism. Photochem. Photobiol. 1970, 11, 457−475. (24) Rappaport, F.; Guergova-Kuras, M.; Nixon, P. J.; Diner, B. A.; Lavergne, J. R. M. Kinetics and Pathways of Charge Recombination in Photosystem II. Biochemistry 2002, 41, 8518−8527. (25) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions; Pergamon Press: Oxford, 1966. (26) Dau, H.; Limberg, C.; Reier, T.; Risch, M.; Roggan, S.; Strasser, P. The Mechanism of Water Oxidation: from Electrolysis via Homogeneous to Biological Catalysis. ChemCatChem 2010, 2, 724− 761. (27) This calculation assumes that hydrogen production from the resulting protons is the coupled half-reaction (i.e., the pH-adjusted potential for the H+/H2 couple). (28) de Wijn, R.; van Gorkom, H. J. Kinetics of Electron Transfer from QA to QB in Photosystem II. Biochemistry 2001, 40, 11912− 11922. (29) Vinyard, D. J.; Ananyev, G. M.; Dismukes, G. C. Photosystem II: The Reaction Center of Oxygenic Photosynthesis. Annu. Rev. Biochem. 2013, 82, 577−606. (30) Esper, B.; Badura, A.; Rögner, M. Photosynthesis as a Power Supply for (Bio-)Hydrogen Production. Trends Plant Sci. 2006, 11, 543−549. (31) Terasaki, N.; Iwai, M.; Yamamoto, N.; Hiraga, T.; Yamada, S.; Inoue, Y. Photocurrent Generation Properties of Histag-Photosystem II Immobilized on Nanostructured Gold Electrode. Thin Solid Films 2008, 516, 2553−2557. (32) Larom, S.; Salama, F.; Schuster, G.; Adir, N. Engineering of an Alternative Electron Transfer Path in Photosystem II. Proc. Natl. Acad. Sci. U.S.A. 2010, 107, 9650−9655. (33) Lubner, C. E.; Applegate, A. M.; Knörzer, P.; Ganago, A.; Bryant, D. A.; Happe, T.; Golbeck, J. H. Solar Hydrogen-Producing Bionanodevice Outperforms Natural Photosynthesis. Proc. Natl. Acad. Sci. U.S.A. 2011, 108, 20988−20991. (34) de Levie, R. The Electrolysis of Water. J. Electroanal. Chem. 1999, 476, 92−93 and references cited.

Gary Brudvig received his B.S. degree (1976) from the University of Minnesota and Ph.D. degree (1981) from Caltech working with Sunney Chan, and he was a Miller Postdoctoral Fellow with Ken Sauer at the University of California, Berkeley from 1980 to 1982. He has been on the faculty at Yale since 1982, where he is currently the Benjamin Silliman Professor of Chemistry, Professor of Molecular Biophysics & Biochemistry, and Director of the Yale Energy Sciences Institute. His research involves study of the chemistry of water oxidation in natural and artificial photosynthesis and work to develop artificial bioinspired systems for solar fuel production.

ACKNOWLEDGMENTS The U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences is gratefully acknowledged for support of our studies of iridium-based water-oxidation catalysts as part of the Argonne-Northwestern Solar Energy Research (ANSER) Center, an Energy Frontier Research Center funded under Award Number DE-SC0001059, and oxomanganese wateroxidation complexes (DE-FG02-07ER15909). J.D.B. thanks the Resnick Sustainability Institute at Caltech for support during preparation of this manuscript. REFERENCES (1) Bockris, J. O. M. Energy: The Solar-Hydrogen Alternative; Hogbin and Poole: Redfern, Australia, 1975. (2) Bard, A. J.; Fox, M. A. Artificial Photosynthesis: Solar Splitting of Water to Hydrogen and Oxygen. Acc. Chem. Res. 1995, 28, 141−145. (3) Meyer, T. J. Chemical Approaches to Artificial Photosynthesis. Acc. Chem. Res. 1989, 22, 163−170. (4) Lewis, N. S.; Nocera, D. G. Powering the Planet: Chemical Challenges in Solar Energy Utilization. Proc. Nat. Acad. Sci. U.S.A. 2006, 103, 15729−15735. (5) Gray, H. B. Powering the Planet with Solar Fuel. Nat. Chem. 2009, 1, 7. (6) Cho, A. Energy’s Tricky Tradeoffs. Science 2010, 329, 786−787. (7) The mechanism of photosynthetic water oxidation has been the subject of considerable work, and is beyond the scope of this review. A brief overview is given here. (8) Zouni, A.; Witt, H.-T.; Kern, J.; Fromme, P.; Krauss, N.; Saenger, W.; Orth, P. Crystal Structure of Photosystem II from Synechococcus elongatus at 3.8 Å Resolution. Nature 2001, 409, 739−743. (9) Kamiya, N.; Shen, J.-R. Crystal Structure of Oxygen-Evolving Photosystem II from Thermosynechococcus vulcanus at 3.7-Å Resolution. Proc. Natl. Acad. Sci. U.S.A. 2003, 100, 98−103. (10) Ferreira, K. N.; Iverson, T. M.; Maghlaoui, K.; Barber, J.; Iwata, S. Architecture of the Photosynthetic Oxygen-Evolving Center. Science 2004, 303, 1831−1838. (11) Umena, Y.; Kawakami, K.; Shen, J.-R.; Kamiya, N. Crystal Structure of Oxygen-Evolving Photosystem II at a Resolution of 1.9 Å. Nature 2011, 473, 55−60. V

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