The Mechanism of Water Oxidation

DOI: 10.1002/cctc.201000126 The Mechanism of Water Oxidation: From Electrolysis via Homogeneous to Biological Catalysis Holger Dau,*[a] Christian Limberg,*[b] Tobias Reier,[c] Marcel Risch,[a] Stefan Roggan,[b] and Peter Strasser*[c]

724 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761 Striving for new solar fuels, the water oxidation reaction cur- catalytic manganese–calcium complex of photosynthetic water rently is considered to be a bottleneck, hampering progress in oxidation. Recent developments in homogeneous water-oxida- the development of applicable technologies for the conversion tion catalysis are outlined with a focus on the discovery of of light into storable fuels. This review compares and unifies mononuclear ruthenium (and non-ruthenium) complexes that viewpoints on water oxidation from various fields of catalysis efficiently mediate O2 evolution from water. Water oxidation in research. The first part deals with the thermodynamic efficien- photosynthesis is the subject of a concise presentation of cy and mechanisms of electrochemical water splitting by metal structure and function of the natural paragon—the manga- oxides on electrode surfaces, explaining the recent concept of nese–calcium complex in photosystem II—for which ideas con- the potential-determining step. Subsequently, novel cobalt cerning redox-potential leveling, proton removal, and OO oxide-based catalysts for heterogeneous (electro)catalysis are bond formation mechanisms are discussed. The last part high- discussed. These may share structural and functional properties lights common themes and unifying concepts. with surface oxides, multinuclear molecular catalysts and the

Introduction tensified, as it can serve as an important inspiration and benchmark for the development of artificial water-oxidation In recent years, the increased interest in direct conversion of catalysts. solar energy into storable fuels has led to intensified research Herein we will discuss and compare concepts and recent efforts with respect to artificial photosynthesis.[1] A successful finding on the mechanism of water oxidation obtained in the imitation of the natural energy-storing process requires the following separate research communities: combination of several distinct processes, including light har- * Heterogeneous electrocatalysis at electrode surfaces (sec- vesting, charge separation, electron transfer, and water oxidation 1), * tion (O2 evolution), as well as the fuel-forming reaction. Where- heterogeneous catalysis by oxidic bulk materials in colloi- as a close adaption to the biological process of photosynthesis dal form or deposited on electrodes (section 2), would involve the transformation of carbon dioxide into a * homogeneous catalysis by transition metal complexes chemical compound of higher energy content, at present the (section 3), * generation of molecular hydrogen (H2) by means of proton re- biocatalytic water oxidation in photosynthesis (section 4). duction seems to be a more realistic proposition for storing Rather than approaching the impossible, that is, a fully com- solar energy by artificial photosynthesis. prehensive review, we focus on recent progress and strive to The electrons and protons required for fuel generation address common themes and potentially unifying concepts. should be obtained from oxidation of water to oxygen, thereby converting the whole process into an overall water-splitting scheme, where hydrogen is the targeted fuel. Hydrogen itself 1. Electrochemical Water Splitting can be utilized as a fuel. Moreover, solar hydrogen will be the key for the realization of a number of related renewable fuel Electrolysis, the electrocatalytic splitting of liquid water into hy- production processes, such as hydrodeoxygenation of lignocel- drogen and oxygen using electricity, is the longest-studied cat- lulosic biomass into fuels and chemicals, or the reduction of alytic way to oxidize water, predating even the development of the concept of chemical catalysis. The overall reaction, cata- CO2 into methanol or other carbonaceous fuels. However, not only the generation of molecular hydrogen from water re- lyzed at the electrified solid–liquid–gas three-phase interface quires water oxidation. For all substances that today are typi- of the anode and the cathode, is given by cally considered as fuels, it holds that utilization of the chemi- 1 cally stored energy involves the reduction of molecular H2OðlÞ ! H2ðgÞþ =2 O2ðgÞ ð1Þ oxygen, that is, the transfer of electrons from the fuel to oxygen. For carbonaceous fuels, this process is coupled to CO [a] Prof. Dr. H. Dau, M. Risch 2 Freie Universitt Berlin, Physics Dept., Molecular Biophysics formation and, in most cases, also to H2O formation. In fuel Arnimallee 14, 14195 Berlin (Germany) generation, a source for these electrons is needed, and water Fax: (+ 49)30-838-56299 is the one and only truly attractive candidate; every scheme E-mail: [email protected] for sustainable, large-scale production of non-fossil fuels will [b] Prof. Dr. C. Limberg, Dr. S. Roggan eventually involve utilization of water as a raw material. Humboldt Universitt Berlin, Department of Chemistry Brook-Taylor-Strasse 2, 12489 Berlin (Germany) Water oxidation is currently considered a major bottleneck, Fax: (+ 49)30-2093-6966 hampering progress in the development of applicable devices E-mail: [email protected] for the conversion of light into storable fuels. Intense research [c] T. Reier, Prof. Dr. P. Strasser efforts are being pursued to develop both heterogeneous and Technische Universitt Berlin, Chemistry Dept. molecular catalysts that enable water oxidation at potentials The Electrochemical Energy, Catalysis and Materials Science Laboratory Strasse des 17.Juni 124, 10623 Berlin (Germany) close to the thermodynamic limit. Research into the biological Fax: (+ 49)30-314-22261 process of photosynthetic water oxidation has also recently in- E-mail: [email protected]

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Holger Dau received his physics diplo- Marcel Risch received his Bachelor of ma in 1985 and doctoral degree in Science (physics) at the Technical Uni- 1989 in Kiel, Germany, working with versity of Darmstadt, Germany, in 2006 U.-P. Hansen and at the Weizmann In- and a Master of Science (physics) at stitut (Rehovot, Israel, winter 1987/88). the University of Saskatchewan, After postdoctoral work with K. Sauer Canada, in 2008. In 2009, he joined the at the UC Berkeley, USA, and H. Senger Berlin International Graduate School of in Marburg, Germany, he received his Natural Sciences and Engineering (BIG- habilitation degree at the Biology De- NSE) and he is currently studying for a partment of Philipps University Mar- PhD in the group of Prof. Dau at the burg in 1994. Besides photosynthesis Free University of Berlin. The topic of research in Marburg, he developed bi- his thesis is the characterization of an otest applications at bbe Moldaenke GmbH (1997–1999). Since electrochemical cobalt catalyst. 2000, he has been a full Professor at the Physics Department of the Free University in Berlin, Germany, where he investigates biological and synthetic metal sites with X-ray spectroscopy and complementary methods. His current focus lies on light-driven water Stefan Roggan graduated in chemistry oxidation and H2 formation in biological and biomimetic systems. in 2003 at the Ruprecht-Karls Universi- ty of Heidelberg in Germany. He then moved together with his supervisor Prof. Christian Limberg to the Hum- Christian Limberg was born in 1965 in boldt-Universitt zu Berlin, where he Essen, Germany, and studied chemistry performed his doctoral studies. In from 1985 until 1990 in Bochum, Ger- 2007, he received his PhD in chemistry many. He earned his doctorate with A. and then joined Prof. T. Don Tilley’s Haas in 1992 and concluded postdoc- group at the UC Berkeley, USA, as a toral work together with A. J. Downs postdoctoral fellow. After his return at Oxford University, UK, with a further from the US in 2009, he continued re- PhD thesis (D. Phil.). Having performed search in Christian Limberg’s group before starting to work as a his habilitation in Heidelberg (1995– scientist for Bayer Technology Services GmbH in Leverkusen, Ger- 1999) he moved to the TU Munich in many. Germany to lead the inorganic chemistry chair on behalf of W. A. Herrmann. In 2002, he accepted an offer to become full professor at the Hum- boldt-Universitt zu Berlin, Germany. His research is mainly con- Peter Strasser is a full Professor at the cerned with oxo metal complexes, oxidation reactions and their Chemical Engineering Division of the mechanisms (surfaces of metal–oxo catalysts as inspiration for mo- Department of Chemistry at the Tech- lecular models; biomimetic O2 activation and oxidation) as well as nical University of Berlin. His research with the activation of small substrates utilizing dinuclear com- interests focus on materials and cataly- plexes. sis for electrochemical (energy) technologies. Prior to his appointment, he was an Assistant Professor at the De- partment of Chemical and Biomolecu- lar Engineering at the University of Tobias Reier received his Master of Houston, USA. From 2001 to 2004, he Science (Dipl.-Ing.) in 2010 at the Tech- served as senior member of staff at nical University Berlin, Germany, in Symyx Technologies, Inc., Santa Clara, CA, developing and utilizing Chemical Engineering. He is currently combinatorial and high throughput screening methodologies in working on his doctoral thesis at the heterogeneous and electrocatalysis. In 1999, Peter Strasser ob- Technical University of Berlin in the tained his doctoral degree in Physical Chemistry and Electrochem- group of Prof. Strasser. His research is istry from the Fritz-Haber-Institute of the Max-Planck-Society, focused on a fundamental understand- Berlin, under the direction of the 2007 Chemistry Noble Laureate, ing of mechanistic aspects and struc- Professor Gerhard Ertl. He studied chemistry at Stanford University, tural effects of electrochemical water USA, the University of Tbingen, Germany, and the University of splitting as well as the design of im- Pisa, Italy, and received his diploma degree (MS) in chemistry in proved water-splitting electrocatalysts. 1995.

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In the future, water electrolysis could play a key role in sus- Today, water electrolysis is frequently cited as a clean corner- tainable energy conversion and storage infrastructure.[2] First stone production technology in a largely hydrogen-based sus- observed in 1789 in a transient discharge experiment by van tainable energy infrastructure (hydrogen economy[13]). Electrol- Troostwijk,[3] then possibly observed by Volta using one of his ysis could serve to liberate hydrogen from a range of chemi- piles[4] and finally made known to a wider audience in 1800 by cals, not only from water. Also, being a scalable technology, Nicolson and Carlisle,[5] it largely remained a scientific curiosity electrolysis would allow the initial deployment of smaller-scale throughout the 19th century. During the early decades of the electrolysis units at comparable efficiencies to large-scale elec- 20th century, water electrolysis was industrialized and served to trolyzers.[13b–d] The vision of a broader hydrogen-based energy produce very pure hydrogen for industrial uses, such as ammo- economy, however, still faces a large number of serious techni- nia production or early petroleum refining.[6] Later in the cal and nontechnical road blocks.[13d] Among the scientific and 20th century, materials and system improvements resulted in technical challenges, the need for advanced materials is critical significant increases in net efficiencies of electrolyzers above and will likely require a large amount of additional basic 50%. Alkaline water electrolysis, where hydroxide anions are energy materials and catalysis research. the major charge carrier in the liquid ion-conducting electro- Hydrogen is an important industrial feedstock for petroleum lyte, became a technology of choice for production of oxygen refining and fertilizer industry. However, electrolysis accounts and hydrogen for life-support systems in anaerobic environ- only for about 4% of global hydrogen production[9b,13b–d,14]. ments, such as submarines. The use of separator materials (dia- The recent surge of interest and investments in renewable phragms) aided in the design of more-compact electrolyzer electricity through solar–thermal, solar–electric (photovoltaic), cells with improved separation of hydrogen and oxygen gas.[7] or wind farms, combined with the emerging interest in electro- Later on, high-pressure alkaline electrolyzers were designed mobility concepts for urban centers, has raised a substantial and commercialized. Work by Beer resulted in the technologi- amount of awareness about pressurized electrolytic hydrogen cally important dimensionally stable anode (DSA),[8] which re- for storage of excess renewable electricity. As a results, hydro- mains the basis of electrolyzer anodes to date. In the DSA, cat- gen–wind hybrid power plants are currently being commercial- alysts are supported directly on metallic titanium as a thin film. ized at a number of locations worldwide.[15] On a smaller scale, Space exploration drove the development of polymer electro- more and more consumer wind and photovoltaic units have lyte membrane (PEM) electrolyzers for acidic media that been deployed in accordance with energy policies that offer showed very similar system architectures as hydrogen PEM generous subsidies for renewable electricity fed back into the fuel cells (PEMFCs).[9] Membrane materials such as perfluorosul- grid. Once these subsidies are cut back and grid electricity fonic acid (Nafion) allow the conduction of protons across a prices continue to rise, a growing interest in home energy stor- thin polymer electrolyte layer. Advantages of proton-conduct- age units, possibly based on water electrolysis and compressed ing PEM devices over liquid alkaline systems include higher hydrogen, will result. safety and reliability, because no caustic soda is circulated in The overall electrochemical water splitting is divided into the cell, greater energy efficiency, sustainability to high differ- two half-cell redox reactions. The reduction process at the ential gas pressures, higher hydrogen production rates, and a cathode (hydrogen evolution reaction, HER) proceeds accord- more-compact cell design.[9a] Drawbacks relate to the use of ing to more expensive noble metal anode catalyst materials, such as Ru or Ir. Current developments in the field of electrolysis in- þ clude the coupling of water electrolysis with photovoltaic solar 2H þ2e ! H2 ð2Þ energy conversion[10] into an integrated direct photoelectroca- talytic (PEC) water-splitting technology.[11] PEC devices promise more-compact electrode and possibly device architectures, yet while the oxidation process (oxygen evolution reaction, OER) require multifunctional hybrid electrode materials. Low solar at the anode of the electrolyzer is[16] conversion efficiencies combined with electrocatalytic overpotential losses largely limit the efficiencies of today’s PEC devi- 1 þ ces to below 5%. Another current development centers on the H2O ! =2 O2þ2H þ2e ð3Þ synthesis of practical alkaline polymer electrolyte membranes. This technology would combine the advantages of the PEM technology with the benefits of a mature alkaline electrolyzer In the total process, two electrons are transferred per formula industry, in particular with regards to the use of less expensive or a total charge of 192928 C (= 2 F) crosses the electrocatalyt- non-noble metal electrode materials. Finally, the integration of ic interface per mole of water converted. electrolyzer and fuel-cell functionalities into one single unitized We will focus herein on the mechanistic aspects of water regenerative fuel-cell device is the subject of current research splitting (electrolysis) itself. Photoelectrochemical water split- and development efforts.[12] A regenerative fuel cell/electrolyzer ting, whereby a photon-absorbing component is directly cou- system offers the prospect of a two-way conversion of chemi- pled with a gas-evolution electrocatalyst, is beyond the scope cal and electrical energy using bifunctional electrodes with the of this review. A recent article has highlighted photoelectro- advantage of further increased system simplicity and compact- chemical water splitting in the context of photoelectrochemi- ness. cal solar fuels.[11]

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1.1. The thermodynamic efficiency of electrochemical water where Ho is the molar enthalpy of water at standard pres- H2OðlÞ ,T splitting sure and temperature T. As a result, the voltage range of endothermic water splitting broadens.[18] The equilibrium thermodynamics of galvanic and electrolysis The formalism above assumed the electrolytic production of cells extends the set of thermodynamic variables of a chemical perfectly dry hydrogen and oxygen gas. Taking into account process that does not involve free charges, an electric poten- the presence of a water-vapor partial pressure pw at a total tial E, or the difference of two potentials, the voltage V.[17] Cell pressure p of hydrogen and oxygen, a total energy for splitting voltages at constant pressure P and temperature T are related of one mole of water into humidified gases can be derived, to thermal energy quantities by the amount of electricity trans- which is associated with a thermo-neutral voltage VTN accord- ferred nF, where n denotes the number of electrons transferred ing to[18–19] and F is the Faraday constant. p Water splitting (ws) according to Equation (1) at standard DH ¼DH þ 1:5 w H Ho TN;T;p HHV;T;ðppwÞ H2OðlÞ ;T;pw H2OðlÞ;298 temperature and pressure (STP) is associated with a molar en- ðÞp pw ð8Þ o thalpy change of reaction, herein denoted as DHws,298, which is ¼ 2 FVTN;T;p equal to DHo , the heat of formation of one mole of f,H2 OðlÞ,298 liquid water at temperature 298 K. The absolute value of these The first term of the central expression is the higher heating 1 quantities is 286 kJmol (water) or 2.96 eV. The quantity value of the electrolysis process at temperature T, and gas o o DHws,298 is also referred to as the higher heating value DHHHV,298 pressure (ppw), while the second term accounts for the stoi- of one mole hydrogen. The change in Gibbs free energy of re- chiometric evaporation of water from STP to operating condi- action (1) Go is +237 kJmol1 (water) or 2.46 eV. Note: Dif- D ws,298 tions. ferent energy units are used in the reviewed fields and we do Owing to a variety of energy and efficiency losses, cell vol- not use consistently the same units but rather those most con- tages in practical electrolysis cells[9b] are generally above both venient in the specific context. For a rough transformation of [18] the thermoneutral voltage VTN and the higher heating value energy units: 1 eV100 kJmol1 25 kcalmol1. voltage VHHV. That is, excess thermal energy generally leads to At equilibrium, the minimum electrical energy, 2FVrev, re- heating of the electrolysis cell. quired to split one mole of water under STP is related to the The efficiency of an isothermal electrolysis cell operated at a Gibbs energy as o cell voltage Vop,T relates its chemical energy output to its electrical energy input, 2FVo . The energy output may be consid- DGo ¼ 2 FVo op,T ws,298 rev,298 ð4Þ ered in two separate ways. Firstly, the maximum reversible work (Gibbs free energy) DG from reconverting hydrogen and o Considering only absolute values of cell voltages, Vrev,298 = oxygen, for instance in a fuel cell, may be of interest and [9b] o 1.23 V at STP and is referred to as the reversible electrolysis cell hence a faradaic (voltage) efficiency of electrolysis, h298;Faradaic, voltage at STP. can be written as The total energy required to split one mole of water at STP is given by o o o DGws;298 Vrev;298 h298;Faradaic ¼ ¼ o ð9Þ 2 FVop Vop;298 DHo ¼ DHo ¼DHo ¼ DGo þ T DSo ws,298 HHV,298 f,H2OðlÞ,298 ws,298 ws,298 ð5Þ This quantity essentially represents the reciprocal value of 1 [20] where the last term at STP is 49 kJmol and corresponds to the voltage efficiency of a fuel cell with Vop replaced by the the heat required from the surroundings. Thus, a higher-heat- experimentally observed output cell voltage of a fuel cell. The [18] ing-value voltage, VHHV, or an equivalent amount of electrici- faradaic efficiency can never be larger than one. Secondly, hy- ty, 2FVHHV, can be related to the water splitting energy accord- drogen and oxygen may be viewed in terms of their heating ing to value and hence a thermal efficiency of an electrolysis cell, ho , is written as[18] o o 298;thermal DHHHV,298 ¼2 FVHHV,298 ð6Þ o o o DHHHV;298 VHHV;298 o o h298;thermal ¼ ¼ o ð10Þ where VHHV,298 = 1.48 V. The voltage difference between Vrev,298 2 FVop Vop;298 o and VHHV,298 implies that water electrolysis below 1.48 V and above 1.23 V is an endothermic process at STP. Hence, neglect- This efficiency has become widely accepted even though it ing ohmic and other loss processes, an electrolysis cell operat- would never be possible to recover the entire higher heating ed below 1.48 V would cool down during electrolysis or would value of hydrogen in an energy application. In principle, require a heat supply to operate isothermally. o [6,18–19] h298;thermal can be larger than unity. It is the reciprocal o At elevated temperatures, Vrev,T drops based on Equations (4) value of the reversible thermodynamic efficiency of a fuel and (5), whereas DHo and Vo increase according to[18–19] [20] o HHV,T rev,T cell. Typical values of hT;thermalof industrial electrolyzers range between 60 % to 80 % at a given current density (typically o o o o o DH ¼>¼DH þðH Þ¼2 FV 2 HHV,T f,H2 OðlÞ,T H2OðlÞ ,T f,H2 OðlÞ ,298 HHV,T ð7Þ 1Acm ) and at 363 K (908C) and 0.1 MPa.

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Equations (9) and (10) provide a measure of the efficiency of turnover frequency of an electrochemical half-cell process. On overall water splitting. They are generally valid, regardless of Pt electrodes (Figure 1), for instance, HER shows values of j0 of whether it is an electrocatalytic process at an inorganic elec- up to several hundred A cm2,[21] resulting in a very reversible trode, a biocatalytic process, or a catalytic pathway involving half-cell electrode with high forward and backward reaction homogeneous organometallic catalysts. In a ‘translation’ of the electrochemical terminology, anode and cathode potentials correspond to the midpoint redox potentials of the oxidant (anode potential) and reductant (cathode potential) that are used to drive the OER or HER. Similarly, in electrochemical, homogeneous, and biological catalysis, research typically consid- ers only half-cell reactions, as is also the case in the mechanistic discussions of water oxidation (section 1.2). For half-cell reactions, however, Equations (9) and (10) are not directly applicable. Equation (9) may be reformulated as Figure 1. Catalytic reactions associated with electrochemical water electrolysis. An external voltage, symbolized by the battery symbol, drives the evolu- o o Vrev;298 tion of hydrogen at the cathode (HER) and the evolution of oxygen at the h298;Faradaic ¼ o ð11Þ Vrev;298 þ OPanode þ OPcathode anode (OER). Protons are migrating from anode to cathode. Both mechanisms are depicted assuming a sequence of proton-coupled single-electron redox reactions. While HER shows one adsorbed intermediate, OER has where OPanode and OPcathode are the absolute values of the over- three. The spheres symbolize a metal electrode, for example, Pt, covered potentials of the OER (anodic reaction) and HER (cathodic reac- with an oxide layer. The numbers in parentheses refer to chemical equations tion), respectively. These overpotentials can be calculated in given in the text. homogeneous and biocatalysis as the difference between the midpoint potential of the oxidant (or reductant) and the OER (or HER) equilibrium potential. To quantify the efficiency of the rates. However, OER on Pt, the catalytically most active ele- half-cell reaction, we now assume for OER (or HER) that ment for this reaction, exhibits values of j0 on the order of 9 2 OPcathode (or OPanode) equals zero. The thereby-obtained value of 10 Acm , which is why much recent work has shifted o o toward OER kinetics. Therefore, the OER is our focus here and hOER (hHER) represents a meaningful measure of the energetic efficiency of the catalysis of water oxidation (hydrogen evolu- also in upcoming sections. tion) in electrochemical, homogeneous, and biocatalysis. For a In electrocatalysis, a fundamental concept in characterizing combination of two half-cell reactions, calculation of the over- catalytic processes on surfaces is the overpotential.[22] Overpo- o all efficiency h from the two half-cell efficiencies is straightfor- tential is the deviation of the observed output cell voltage Vop o o o (galvanic cell) or the applied cell voltage V (electrolysis) used ward. However, h is not equal to the product of hOER and hHER op o 1 o 1 1 in Equations (9) and (10) from the thermodynamic reversible but to [(hOER) +(hHER) 1] . cell voltage Vrev [Equation (4)]. The overpotential is generally considered a kinetic phenomenon and splits into various con- tributions such as the ohmic voltage loss (ohmic overpotential) 1.2. Reaction mechanisms of electrochemical water splitting or mass-transport limitations (transport overpotential) during As described above, the electrochemical splitting of water in- the flow of charge.[22] In the context of catalysis, “reaction over- volves two concurrent catalytic half-cell reactions according potentials” relate to a kinetic hindrance of individual elementa- Equation (2) (HER) and Equation (3) (OER). Also shown are ry reaction steps due to activation barriers. On that basis, reac- mechanistic schemes for either reaction process, which have tion overpotential is a function of the flowing current and is been proposed in literature in the past. Clearly, oxygen evolu- thought to be linked to the slowest of the elementary reaction tion being a 4 electron process is more complex as compared steps, the rate-determining step (RDS), with the highest kinetic to the evolution of hydrogen and involves several surface ad- activation barrier. sorbed intermediates. As either process may contribute to volt- Despite much work on the OER, mechanistic models and hy- age and efficiency losses, a discussion of the electrocatalysis of potheses of the sequence of elementary steps based on water splitting should in principle touch upon both the HER atomic-scale experiments have remained scarce. Over the and OER. Work on biocatalytic water splitting has focused course of five decades, proposed mechanisms, conclusions as either on photosystem II (oxygen evolution), as discussed in to the RDS, and experimental approaches have changed only section 4, or hydrogenase enzymes (hydrogen evolution), little. A severe hurdle to more direct atomic-scale insight into which are not discussed herein. Similarly, research concerning OER is the fact that reactive intermediates have largely eluded homogeneous catalysis of water splitting has involved systems direct spectroscopic observation. Most early mechanistic work liberating either hydrogen or oxygen (or, in rare cases, both; relies on classical current–potential–time analyses and the de- see section 3.2 for comparison). In electrochemical environ- rived Tafel slopes. Tafel slopes are derived from semi-logarith- ments, HER on selected noble metals shows extremely high ex- mic potential current plots and have been frequently used to change current densities, j0, which is a measure of the intrinsic obtain information on the mechanism and the rate-determin-

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 729 H. Dau, C. Limberg, P Strasser et al. ing step of the overall half-cell reaction. An early mechanistic de1.[23d,27c,d] Similar to the single crystal work, reaction (16) was scheme of electrochemical water splitting based on experi- proposed for these films as a second elementary step. Low- ments on RuO2-based electrodes in acidic environments is out- temperature Ru oxide films showed high BET surface area lined by Equations (12)—(15):[23] values combined with higher values of q*. The OER activity was high for these films and the Tafel slope ranged in the þ 1 SþH2O ! SOHþH þe ð12Þ 40 mV decade range. Hence, an electrochemical reaction (13) was now proposed as the rate-determining step and the full SOH ! SOþHþþe ð13Þ mechanistic sequence thus becomes (12)–(13)–(15). SOHþSOH ! SOþSþH2O ð14Þ Binary and ternary mixed oxides containing Ir, Sn, and Ru showed their own distinct complex Tafel slope patterns with SOþSO ! O2þ2 S ð15Þ no clear trend or controlling parameter.[25,27c,d, g,29] A number of recent studies focused on the effect of particle size on the OER [30] In this mechanistic hypothesis, S denotes a catalytically mechanism for RuO2 and Ru1xCoxO2 mixed oxides. The au- active surface site. Equations (12), (13), and (15), on the one thors used differentially pumped mass spectrometry (DEMS) to hand, and (12), (14), and (15), on the other, represent parallel directly correlate the evolution of molecular oxygen and its far- reaction pathways. This mechanism and very similar schemes adaic current profile. They found that the OER reaction mecha- were proposed for both well-defined Ru(110) single-crystal nism was independent of catalyst particle size but, however, [23d,24] [23d] [30b] oxide surfaces and compact RuO2 films. RuO2 single was affected by Co doping. The second metal lowered the crystals were reported with a near 60 mV decade1 and a Tafel slope from values near 120 mVdecade1 to 40–50 mVde- 120 mVdecade1 Tafel slope at low and high overpotentials, re- cade1, consistent with a chemical step becoming rate-deter- spectively.[23d,24] On that basis, reaction (12) was thought to be mining. the RDS of the sequence (12)–(13)–(15) at high electrode po- In alkaline solution, non-noble metal oxides of Ni, Mn, Co, or tentials. An additional second chemical step, such as the hypo- Pb with rutile-type,[27a] perovskite-type,[27a, 31] and spinel- thetical rearrangement of the SOH species according to type[27a,32] structures were investigated in terms of their reac- Equation (16):[24] tion mechanism for OER. An initial discharge of hydroxide ions was proposed as the first step in the reaction sequence:[33] SOH ! SOH* ð16Þ SþOH ! SOHþe ð17Þ was proposed to account for the smaller experimental Tafel slopes at low overpotentials, with the reaction sequence now To facilitate comparison of mechanisms in heterogeneous being (12)–(16)–(14). electrocatalysis, homogeneous, and biological catalysis, a tran-

Ti-supported polycrystalline RuO2 and Ru–MO2 mixed oxide sition from the physical terminology involving ‘adsorption’ and films were studied extensively, because they compositionally ‘hydroxide discharge’ to a description in terms of chemical pro- resemble the industrial DSA used in virtually all commercial cesses is required. The hydroxide ‘discharge’, by binding to the electrolyzers.[25] A very similar reaction mechanism to Equa- electrode surface, could relate either to formation of a bound [26] tions (12)–(15) was suggested for IrO2. A large number of hydroxide with radical character (OH·) or, more likely, to liga- studies considered the influence of the preparation conditions tion of the hydroxide ion to a metal of the surface oxide such applied for Ru oxide and Ru–Ir bimetallic oxide films on the that the oxidation state of this specific metal ion formally in- OER activity and mechanism.[26] Synthetic parameters studied creases by one unit. included the metal precursor and synthesis route (e.g. sol–gel The step described by Equation (17) was thought to be fol- versus precipitation), annealing temperature of the catalyst lowed by either a deprotonation and subsequent dis- film, the resulting BET surface, the integral voltammetric charge[33–34] charge q*, and the point of zero charge.[23d,27] Due to the lack of a single parameter controlling trends in OER activity or SOHþOH ! SO þH2O ð18Þ mechanism, many studies resorted to correlations of synthesis SO ! SOþe ð19Þ parameters,[27a,28] voltammetric charge q*,[27a] OER activity (Tafel slopes), macroscopic structure, and morphology, often intro- 2 SO ! 2 SþO2 ð20Þ ducing somewhat arbitrary descriptors for catalyst morphology such as cracked, smooth, compact, or rough.[27a] Sol–gel-type or by a slow breaking of the MOH bond to form a peroxide Ru oxide films followed the Tafel slope pattern of single crys- species,[31a,c] tals,[28] whereas other synthetic methods resulted in deviating slopes. SOHþOH ! SþH2O2þe ð21Þ High annealing temperatures resulted in polycrystalline Ru oxide films with low BET surface areas and small voltammetric that subsequently decomposes into O2. Similar to studies in charges q*, which was interpreted in terms of a low density of acid environments, synthesis conditions and Tafel slopes varied active sites.[23d,24,27a,d] The OER activity of these films is generally widely. Bockris and Otagawa presented a detailed study[31b] low and their Tafel slope showed values around 60 mVdeca- where they correlated the OER activity with a large number of

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microscopic and material specific parameters. Their findings activity (higher overpotential at given current density). The IrO2 suggested that the catalytic OER activity correlated inversely content in Ru–Ir mixed oxides, in contrast, correlated inversely with the surface bond energy of OH, from which he concluded with the PZC. The PZC of solid surfaces, even though the that the desorption of OH or oxygenated species may be the result of a complex convolution of factors, can be viewed in rate determining step in the OER under his conditions. Tafel the context of the Brønsted acid behavior of m-oxo bridges, as slopes on doped perovskites varied from 43 to 200 mVdeca- discussed for the {Mn4Ca} center in photosynthetic water oxi- de1[31b] and hence provided limited insight into the reaction dation by photosystem II, where proton dissociation is crucial mechanism. In yet further mechanistic studies in alkaline in redox-potential leveling (see Section 4). media,[35] the formation of higher-valent metal oxides was A seminal concept introduced by Ruetschi and Delahay,[38] stressed as also concluded in studies on water oxidation in later extended by Parsons,[39] Bockris[31] and Trasatti,[27a] involves acidic media.[27f] An early molecular orbital theory study of the the hypothesis that a single microscopic parameter, such as OER mechanism showed that chemical recombination of ad- the oxygen chemisorption energy or the molar enthalpy of the sorbed surface oxygen[33] has a low barrier [Equation (15)] and oxidation of metal oxides, may be a key controlling factor in the study also proposed the formation of an adsorbed perox- the OER and possibly also in the chlorine evolution reaction ide species (SOOH). (ClER). Along these lines, OER can be viewed as the formation The specific adsorption of ions from the electrolyte at the and subsequent decomposition of high-valent metal surface inner portion of the electrochemical double layer (inner Helm- oxides. To investigate this hypothesis, Bockris[31b] and Trasat- holtz layer) can have a tremendous impact on the available ti[27a] both reported volcano plots of OER activity (overpoten- number of active electrocatalytic sites. Specific adsorption can tial) versus the metalOH bond energy and versus the enthal- block active sites and hence seriously affect the catalytic activi- py of formation of higher oxides, building on Sabatier’s princi- ty. Knowledge and control of such ion effects is therefore key ple of balanced intermediate adsorption in catalysis[40] (see sec- to fundamental understanding and design of electrocatalytic tion 1.3). In Trasatti’s study in acidic environments, RuO2 was interfaces. Anion effects in the oxygen reduction reaction located near the top of the activity volcano curve owing to its (ORR) have been studied in much detail over recent decades.[36] balanced energetics in the formation and decomposition of However, little thorough work has been done in the area of co- higher-valence oxides. This could potentially be the reason valent and noncovalent interactions of OER catalysts with spec- that a majority of mononuclear complexes and coordination tator species..[37] Lodi et al.[23d] reported no influence of the compounds with OER activity contain Ru ions. nature of the electrolyte anion (sulfate or perchlorate) on the In summary, considering the large number of partially con-

OER catalysis on RuO2, in stark contrast to the effect of anion flicting reports and the variety of synthetic routes and condi- adsorption in the oxygen reduction reaction.[36b,g] Lodi’s result tions Tafel slope-based mechanistic work should be viewed suggests that active OER sites were not blocked by sulfate with caution and care where reliable conclusions as to the rate anion adsorption. As a consequence, catalytic and kinetic im- determining step of the OER are concerned. Correlation of ac- provements by selective blocking of adsorbing anions may not tivity with microscopic parameters related to the surface chem- be an option in the search for improved OER catalysts.[36h] ical bond is important for an atomic level insight into the elec- Babak and co-workers[37b] reported a weak effect of the sup- trocatalytic processes. Only experimental in situ spectroscopy porting electrolyte anion (sulfate, nitrate, or perchlorate) on of surface intermediates combined with atomic-level analysis the OER activity in neutral conditions on PbO2 electrodes only and a reliable theoretical computational framework will ulti- at high overpotentials. Amadelli et al.[37a] reported a blocking mately be able to resolve kinetic activation barriers, elementary influence of fluoride anions for the OER on the PbO2 electro- steps, and possibly the nature of the active site of the OER on des. Clearly, more work has to be done on understanding spec- oxidic surfaces. tator effects in OER catalysis, as it may represent an alternative There are a number of developments towards in situ atomic- route to improved surface electrocatalysis.[36h] scale understanding of electrochemical water splitting. In situ Trasatti[27a] highlighted the relation between preparation electrochemical Raman spectroscopy[41] or mass spectrometry, conditions and the point of zero charge (PZC) of an oxide sur- for instance, have become valuable tools for insight into elec- face. The PZC is the pH at which the surface carries no net trocatalytic mechanisms. The incorporation of atomic oxygen electrical charges. The PZC controls the acid–base properties originating from reactant water into the surface and subsur- of surface coordination complexes such as SOH. A low PZC, face of the RuO2 and IrO2 lattice were confirmed by using for instance, can imply that the SO bond is strong and the in situ differentially pumped mass spectrometry studies (in situ surface hydroxide groups can behave like a weak Brønsted DEMS).[26,42] The formation of higher-valent Ru compounds acid: such as RuO4 was also reported, which corroborates the degradation mechanisms of Ru oxides during water splitting due to þ SOH ! SO þH ð22Þ volatile higher-valent oxides. Novel in situ tools for the investigation of solid–liquid inter- Thus, indirectly, the PZC may provide information on the faces include modern synchrotron-based X-ray methodologies, bond strength or chemisorption energy of S and O. The an- such as ambient-pressure X-ray photoemission spectroscopy nealing temperature of pure RuO2 films was found to correlate (XPS), in situ X-ray absorption spectroscopy (XAS), or X-ray well with the measured PZC. Higher PZC resulted in lower OER emission spectroscopy (XES), which allows the simultaneous

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 731 H. Dau, C. Limberg, P Strasser et al. detection of a variety of different surface-adsorbed oxygen species.[43] In the future, these methods are likely to advance our understanding of the interaction of water with solid surfaces and mechanistic aspects of water splitting.

1.3. The thermochemical overpotential of water splitting The activation energy of the slowest elementary electrochemical step is generally thought to control the relation between electrode overpotential and the catalytic rate of a reaction process. This purely kinetic notion tacitly assumes that all elementary steps along the reaction pathway are thermodynamically Figure 2. Plot of Gibbs free energies of reactive species and Gibbs free ener- downhill. Studies into the relationship between activity and gies of chemisorptions of intermediates (horizontal lines) versus the reaction thermochemical quantities date back to reports by the groups coordinate of the hydrogen evolution reaction (HER). Blue lines and red of Parsons,[39] Bockris,[31b] and Trasatti.[27a] lines indicate energetics of a typical (real) catalyst and ideal catalyst at three Building on those early concepts, recent theoretical compu- different electrode potentials (E1, E2, E3), respectively. Dashed lines indicate energetics at the electrode potential where all thermochemical barriers dis- tational studies[44] of electrocatalytic reactions highlighted the appear. DGi denotes the free reaction energy of the two elementary reaction thermochemical aspects of the overpotential of electrochemi- steps. The red case corresponds to an overpotential-free catalyst. Eo denotes cal reactions using density functional theory (DFT). In their the- the standard reversible electrode potential of the HER, h is the overpoten- oretical framework, Rossmeisl and co-workers calculated Gibbs tial. free adsorption (chemisorption) energies DGi of the adsorbed surface intermediates as function of the electrode potential Figure 2 shows the Gibbs free adsorption energies of the in- and combined these with a thermodynamic model in which dividual species of Equations (23) and (24) as horizontal lines. pH and electrode potential are represented as control parame- The Gibbs free energies of adsorption can be derived from the ters.[44b,d] The authors related the Gibbs chemisorption energies total energies of adsorption, DE“H+,e”, DEH , and DEH , of the of the intermediates to the normal hydrogen electrode (NHE) (ad) 2 reacting proton and electrons, the intermediate H , and the + (ad) by assuming equilibrium conditions for the H2/H redox [44d] product H2, respectively. The x axis represents the reaction system,[45] which rendered the calculation of the chemical po- coordinate. A typical electrocatalyst (Figure 2, blue traces) ex- tential of an aqueous proton unnecessary. The difference be- hibits an imbalance between DG1 and DG2; in this case, for in- tween the chemisorption energies of two subsequent inter- stance, DG1 >DG2. At electrode potentials E1 above the reversi- mediates was equal to the Gibbs free reaction energy, DGrxn,of o o ble potential E (positive overpotentials: E1E = h> 0; the elementary mechanistic step at a given electrode potential. Figure 2, lower traces) all elementary reactions are uphill and The differences in chemisorption energy caused each individu- thermodynamically unfavorable. As the electrode potential is al G to turn negative at different electrode potentials. D rxn lowered, all Gibbs energies of charged species shift more posi- Under the premise that each elementary step must be thermo- tively by neE, where n is the number of electrons in Equa- chemically downhill (i.e. each G must be negative) for the D rxn tions (23) and (24) and e the elementary charge. At the reversi- overall reaction to occur at an appreciable rate, the authors o ble electrode potential E2 (E2E =h= 0; Figure 2, center could predict the minimum electrode (over)potential for the traces), the elementary reaction in Equation (23) has a positive overall reaction to occur. Koper and Heering reported that the Gibbs free energy and is therefore the potential-determining last elementary step to attain a negative DGrxn value is the ‘po- step. At the electrode potential E3 (negative overpotential; tential-determining step’ in analogy to the kinetic concept of a Figure 2, dashed blue top trace), the elementary reaction in rate-determining step.[46] We note that the ‘potential-determin- Equation (23) becomes reversible. There are no more thermo- ing step’ and the rate-determining step are not necessarily dynamic barriers along the reaction pathway, and hence, the identical. o potential (E3E ) represents the thermochemical overpotential Figures 2 and 3 illustrate the idea of the thermochemical of the hydrogen evolution of the real catalyst (Figure 2, blue overpotential for a mechanism of the hydrogen evolution reac- line). This value represents the overpotential at which the over- tion (HER).[47] The reaction is assumed to follow the simple 2- all reaction rate becomes independent of the applied electrode electron Volmer–Heyrovsky reaction sequence: potential and may reach a certain prescribed magnitude. Kinet- ic barriers are not considered in this framework; they may DG 2Hþþ2e 1 H þHþþe ð23Þ keep the actual reaction rate low and introduce a further elec- ! ðadÞ trode-potential dependence of the reaction rate. DG The energetics of the ideal catalyst (Figure 2, red trace) is H þHþþe 2 H ð24Þ ðadÞ ! 2 characterized by equidistant adsorption energy levels of the reactive species, implying identical Gibbs reaction energies of where DGi denotes the elementary Gibbs reaction energies of the two elementary steps, DG1 =DG2 at any given electrode Equations (23) and (24). potential. As a result, the Gibbs free energies of all species and

732 www.chemcatchem.org 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761 The Mechanism of Water Oxidation intermediates are also equal at the equilibrium potential (Figure 2, red dashed trace), and hence no thermodynamic overpotential is expected. Figure 3 shows the relation between the negative thermochemical overpotential, hthermo, of each reaction step as function of the energy of adsorption DEH(ad) of the H(ad) intermediate.[44a,48] The thermochemical overpotential is taken as a mea-

Figure 4. Plot of Gibbs free energy of reactive species and intermediates (horizontal lines) of the oxygen evolution reaction (OER) versus the reaction coordinate. Blue lines and red lines indicate energetics of a real (typical) catalyst and an ideal catalyst, respectively, at three different electrode potentials. Dashed lines indicate energetics at the electrode potential where all

thermochemical barriers disappear (“thermochemical overpotential”); DGi denotes the free reaction energy of the two elementary reaction steps. The red ideal case corresponds to a catalyst with vanishing overpotential.

Figure 3. One-dimensional thermochemical volcano plot of the HER. The negative thermodynamic overpotential hthermo is plotted against the total DG 1 þ 25 hydrogen chemisorption energy DE . The left and right side of the volcano 2H2O OHðadÞþH þe þH2O ð Þ H(ad) ! are associated with a “potential-determining” reaction step given in the plot. The volcano curve provides the thermodynamic overpotential, at which the DG OH þHþþeþH O 2 O þ2Hþþ2eþH O ð26Þ HER reaction occurs at an appreciable rate (potential independent rate). The ðadÞ 2 ! ðadÞ 2 circles correspond to the real and the overpotential-free ideal case from Figure 2, respectively. In this two-electron process, tuning the hydrogen ad- þ DG3 þ OðadÞþ2H þ2e þH2O OOHðadÞþ3H þ3e ð27Þ sorption energy can result in the ideal case. ! DG OOH þ3Hþþ3e 4 O þ4Hþþ4e ð28Þ ðadÞ ! 2 o sure of activity and is defined as hthermo =DGi /e evaluated at the reversible electrode potential Eo of the process. The plot The four-step mechanism of Equations (25)–(28) differs from has a volcano-type shape with a single maximum. At more the classical mechanisms for acidic and alkaline solutions de- negative DEH(ad) (left side of volcano) hydrogen is adsorbed scribed by Equations (12)–(15) and (18)–(21), respectively. strongly and hence the desorption of atomic hydrogen DG2 As for HER, we compare and contrast the energetics of an controls the thermodynamic overpotential. On the opposite ideal (red) and a real (blue) catalyst in Figures 4 and 5. The inside of the volcano, the adsorption of atomic hydrogen limits dividual chemisorption energies of the intermediates reacting the overall reaction. As “DEH(ad) (real)” of the real catalyst on the real catalyst are assumed to be associated with the

(Figure 3) is more positive than the optimum (red dot), reac- Gibbs reaction energy order DG3 >DG1 =DG2 > DG4; the for- tion (23) with Gibbs energy DG1 controls the overall rate. The mation of the peroxide intermediate (DG3) is the thermochemi- ideal catalyst (Figure 3) has equidistant energy levels and cally least favorable step for the real catalyst. Figure 4 indicates hence “DEH(ad) (ideal)” is at the top of the volcano. At the equi- that the OOH(ad) species is bound a little too weakly to the cat- librium potential, all thermodynamic barriers disappear and alyst. As for HER, the ideal catalyst shows equally spaced the reaction is ideally reversible without any thermochemical chemisorption energies of the intermediate and hence equal overpotential. This simple analysis shows that for two-electron Gibbs reaction energies for each elementary step; DG3 =DG1 = processes with one adsorbed intermediate I, there generally DG4 = DG2. For instance, at an overpotential of h=1.23 V, exists an ideal chemisorption energy “DEI(ad) (ideal)” of the inter- corresponding to an electrode potential of 0 V vs NHE (“zero mediate such that the thermochemical overpotential vanishes. bias” at pH 0), the overall change in free energy of the four- Based on that, the search for the ideal catalyst material would electron water splitting reaction amounts to 4.92 eV; in this sit- [49] require tuning DEI(ad) to the ideal value. uation, the ideal catalyst would exhibit a Gibbs reaction The anodic electrocatalytic OER[44b,d] (Figures 4 and 5) in- energy of 1.23 eV for each step. Taking water as the zero point volves three adsorbed intermediates; OH(ad),O(ad), and OOH(ad) of the energy scale, the Gibbs energies of adsorption of the (Figure 1). Similar to the HER mechanism above, the OER ideal intermediates would be 1.23 eV, 2.46 eV, and 3.69 eV for mechanism underlying this analytical framework consists exclu- OH(ad),O(ad), and OOH(ad), respectively. sively of single-electron charge-transfer steps[44d] according to Figure 4 illustrates the energetics of the real and ideal cata-

Equations (25)–(28): lyst at various electrode potentials. At electrode potential E1, all steps are uphill for either catalyst and hence OER cannot

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 733 H. Dau, C. Limberg, P Strasser et al.

given structural family, such as rutile-type oxides, (111) metal, or nitrides. These findings have important implications. Firstly,

given the scaling relations, tuning of DEOOH(ad) to smaller values

will also change DEi for all other intermediates, so that energetic equidistance is in principle unachievable. Secondly, based on a given set of linear scaling relations, the reaction free ener-

gies DGi (i=1–4) can be expressed by using a one-dimensional volcano plot with a single controlling chemisorption parame- [44–45] ter, such as DEO(ad). The solid lines in Figure 5 illustrate the linear relations be-

tween the negative thermodynamic overpotential, hthermo = o DGi /e, evaluated at the reversible potential, and the parame-

ter DEO(ad) for all four elementary steps on a rutile-type oxide [44d] Figure 5. Relation between the thermodynamic overpotential, hthermo, associ- surface. For instance, the scaling relation between the ad- ated with elementary steps of the OER [Equations (25)–(28)] and the total sorption energies of OOH(ad) and O(ad) on rutile surfaces was cal- energy of adsorption of ‘nonmolecular’ oxygen DEOad at the reversible elec- [44d,51] trode potential on a non-ideal (real) and ideal catalyst surface (solid lines culated in detail as and dashed lines, respectively). Energetics for the two catalysts are assumed to be consistent with considerations in Figure 4. Solid lines follow theoretical DEOOHðadÞ ¼ 0:64 DEOðadÞ þ 2:03 eV ð29Þ studies on rutile (110) surfaces in reference.[44d] For the real catalyst, overpotentials associated with Equations (28), that is, DG3, and (27), that is, DG2, are dominant over most adsorption energies. The linear relations of DG3 (solid Taking zero point energies and entropies into consideration, green) and DG4 (solid brown) control the overall overpotential of the OER this equation becomes on the real catalyst and are the “potential-determining steps” at low and [44d] high DEO(ad). The blue circle represents the situation of the real catalyst in DGOOH ¼ 0:64 DGO þ 2:40 eV 30 ðadÞ ðadÞ ð Þ Figure 4. Tuning of DEO(ad) to the transition between the green and brown relation optimizes the catalyst for the given scaling relations. However, the overpotential cannot be completely eliminated for this set of scaling rela- The scaling relations between OH and O were calculat- tions. Dashed lines represent an ideal set of scaling relations where appro- (ad) (ad) ed as priate tuning of DEO(ad) results in an overpotential-free catalyst and OER reaction (red circle).

DEOH ¼ 0:61 DEO 0:90 eV 31 ðadÞ ðadÞ ð Þ

and proceed. At the reversible potential E2 (Figure 4, center, solid blue and dashed red traces), the real and ideal catalysts DGOH ¼ 0:61 DGO 0:58 eV ð32Þ behave distinctly differently. For the real catalyst, DG1 and DG2 ðadÞ ðadÞ vanish, DG4 is negative, and DG3 remains positive, hindering the OER. Thus, Equation (27) describes the potential-limiting Koper pointed out that the difference, DGOOH(ad)DGOH(ad), re- step. The ideal catalyst shows no more uphill energetics at the sults in an approximately constant energy gap of about reversible electrode potential and has therefore no overpoten- 3eV,[51] which is clearly above the ideal gap of 2.46 eV required tial, provided the kinetic limitations are negligible. The elec- in an ideal case. The difference in the constant energy term is trode potential E3 (Figure 4, bottom) is required in order to believed to originate from the difference in the chemical make all steps downhill for the real catalyst, at which OER can nature and bonding of the two intermediates and is likely not occur at a prescribed rate. to depend greatly on the nature of the catalytic surface. Simple graphical inspection of Figure 4 suggests that a Herein lies the fundamental challenge in improving the water theory-guided improvement of the real OER catalyst (blue) oxidation reaction rate. would need to shift the chemisorption free energy of the The scaling relations [Equations (30) and (32)] of the inter- peroxo species OOH(ad), DGOOH(ad), to more negative values, char- mediates further imply that OH(ad) binds too strongly compared acterizing a stronger surface bonding until equally spaced re- to the ideal case, whereas OOH(ad) binds at close to the ideal action free energies are obtained (red). Rossmeisl et al.,[44b,d] case. The energetic position of the near 3 eV gap relative to however, demonstrated that the independent change of a the ideal case depends on the nature of the surface; on metal- single chemisorption energy to achieve optimum relative ad- lic close-packed (111) surfaces, OH(ad) was found to bind close sorption energetics is not feasible, because the chemisorption to ideality, whereas OOH(ad) binds much too weakly. energies of the three individual surface intermediates of the From Equations (29)–(32), the following Gibbs reaction OER mechanism in Equations (25)–(28) on solid surfaces are energy relations were derived, with E denoting the electrode strongly correlated.[44a,c,50] The authors’ calculations suggest potential versus NHE:[44d] that the individual chemisorption energies of the intermediates are linked by linear scaling relations. Their work further provid- DG2 ¼ 0:39 DEOðadÞ þ 0:60 eV eE ð33Þ ed evidence that these relations appear to be independent of G 0:36 E 2:38 eV eE the detailed composition of the solid catalyst surface for a D 3 ¼ D OðadÞ þ ð34Þ

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The slope was related to the making or breaking of surface imply limits in the ability to design a low-overpotential electro- bonds, with positive and negative slopes indicating the forma- catalyst for reactions with multiple adsorbed intermediates, tion and breaking of a bond, respectively. Values close to 0.5 such as the electrochemical evolution of molecular oxygen. were calculated for the slope on metal(111) surfaces, whereas Direct experimental support of generalized scaling relations the slopes on the rutile surfaces were slightly off a value of 0.5. between reaction intermediates has not been reported to o o A graphical representation of DG2 and DG3 at the reversible date, yet is a current subject of much interest. Experimental electrode potential E= 1.23 V is given by the brown and green methods to determine surface binding energies of oxygen- solid lines in Figure 5. These two reactions control the thermo- containing intermediates in electrochemical surfaces, such as dynamic overpotential over a large portion of the oxygen calorimetry and in situ ambient pressure photoemission spec- chemisorption energy range shown. troscopy, will likely play a critical role in such future studie- [43a,b,52] The blue circle in Figure 5 and its associated -hreal value indi- s. cate the thermodynamic overpotential of the real catalyst from The origin of the generalized scaling relations for different

Figure 4, given its oxygen chemisorption energy DEOðadÞ (real). surface adsorbates, such as CHx,NHx,OHx,orSHx, appears to As confirmed in Figure 5, the Gibbs energy of formation of be related to the number of unsaturated bonds of the inter-

OOH, DG3 (Figure 5, green), controls the thermochemical over- mediates, the nature of the binding site, and the resulting potential hreal for the real catalyst (blue). Following Equa- bond order of the intermediate. Rutile(110) and metal(111) sur- tion (34), weaker oxygen chemisorption (more positive DEOðadÞ ) faces exhibit distinctly different sets of scaling relations for the lowers the thermodynamic overpotential until a minimum is oxygenated surface;[44d] some other structural classes of solid reached at the transition where the formation of adsorbed surfaces may exhibit scaling relations that more closely match atomic oxygen [Equation (27)] becomes energetically least fa- the desirable equidistant Gibbs free energy levels of the inter- vorable (Figure 5). At even weaker oxygen chemisorption, DG2 mediates. However, for an overpotential-free catalyst, both the and reaction (27) control the overpotential. It is clear that scaling relations and the adsorption energies have to be ideal tuning the chemisorption energies of the intermediate will not (Figure 5). A surface with ideal scaling relations, yet unfavora- result in an overpotential-free ideal catalyst (Figure 4, red). ble adsorption energies can exhibit a lower catalytic activity To achieve the ideal case, novel catalysts with a distinct set than a surface with unfavorable scaling yet optimal binding of ideal scaling relations need to be found, allowing for equally energies. Hence, optimizing a water splitting electrocatalyst spaced reaction free energies (Figure 5; dashed lines). The red may possibly involve tuning both the slope and y-intercept of ball indicates the combination of ideal scaling relations with the scaling relations in combination with tuning surface ad- ideal oxygen chemisorption energy (DEOðadÞ 2.46 eV, DGOðadÞ = sorption energies - an ambitious task that will require much 2.46 eV) resulting in vanishing overpotential. Each bond-form- more fundamental and practical knowledge on how to experi- ing and bond-breaking linear relation falls on top of each mentally tune surface-bond energies. The slope of scaling rela- other and intersect exactly in the red ball at the ideal oxygen tions appears to be loosely related to the bond order of the in- chemisorption energy. The slopes of the ideal scalings are not termediate species, whereas the y intercept could be related to exactly known, yet an absolute value around 0.5 is assumed, in the oxygen atom. line with findings for metal(111) surfaces. Even for ideal scaling Experimental catalysis research is faced with a huge range of relations, the actual oxygen chemisorption energy may be sub- structural and compositional materials parameters to be ex- optimal, suggesting that the experimentally observed OER re- plored in the quest for a reversible electrocatalyst. Computa- action rate of a material with ideal scaling yet suboptimal tional predictions of suitable surface structures with favorable oxygen chemisorption may be inferior to those for the current- scaling and binding energies, possibly combined with combi- ly known materials, and offering a practical challenge in the natorial and high-throughput screening approaches, in con- identification of ideal water-splitting catalysts. junction with in situ spectroscopy will be critical tools for de- On the surface of a real catalyst surface, such as a supported signing and developing reversible electrocatalysts for water metal oxide nanoparticle, the multiplicity of structurally distinct splitting.[53] facets leads to a range of surface sites with distinctively different coordination and chemisorption properties. Generally, low- 2. Water Oxidation by Bulk Metal Oxides coordinated sites bind adsorbates more strongly and hence, in case of the real catalyst in Figure 5, there may be a small frac- In 2008, Nocera and co-workers reported efficient water oxidation of surface sites with a more favorable adsorption thermo- tion by a cobalt oxide film formed by electrodeposition from a chemistry. The overall overpotential will then be a convolution solution containing cobalt, potassium, and phosphate ions.[54] of individual overpotentials of different sites. Under these con- Hereafter, we will denote the cobalt catalyst film as CoCF.[55] ditions, it is feasible that different surface sites are thermo- The report of Nocera and co-workers has attracted much at- chemically limited by different elementary reactions. tention and may have initiated a revival of bulk metal oxides In conclusion, a simple thermochemical framework[44a,d] has as catalysts for water oxidation (we use the term ‘bulk metal provided insights into qualitative trends of the OER overpoten- oxides’ to differentiate these from the ‘surface oxides’ formed tial and reactivity on well-defined metal oxide surfaces. The on metal electrodes which are exposed to a voltage sufficiently framework predicts linear scaling relations between the chemi- positive for water oxidation). Without attempting a compre- sorption energies of intermediates. These scaling relations hensive review, we will discuss selected cornerstone studies on

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 735 H. Dau, C. Limberg, P Strasser et al. bulk metal oxides as catalysts for water oxidation, with a focus 2.1. Properties of an electrodeposited cobalt catalyst on cobalt oxides. In 1950, El Wakkad and Hickling[56] reported the performance The aforementioned catalyst developed by Kanan and Nocera of various CoO/OH films deposited at 5 mAcm2 from was first produced by electrodeposition from an aqueous solu-

CoSO4·7H2O, H3BO3, and NH4Cl, on Pt wire electrodes. These tion of KH2PO4 and K2HPO4 (KPi) at pH 7 containing m m films were tested in a variety of electrolytes, among which was [Co(OH2)6(NO3)2], with concentrations of 0.1 (KPi) and 0.5 m m 2+ a solution of KH2PO4 and NaH2PO4 at pH 6.8 (0.2 each). They (Co ) on an indium tin oxide (ITO) substrate. More recently, found that the presence of phosphate stopped the dissolution the following modifications have been discussed: 1) Deposition of Co observed with other alkaline and acidic electrolytes, and on other semiconductor and metal substrates, such as glassy that an initial dissolution was reverted in phosphate buffer, carbon, carbon felt, fluorinated tin oxide (FTO), and nickel which indicates that the films had a self-repair mechanism in metal,[63] and especially importantly, deposition and operation [64] [65] phosphate buffer. The ability of the films to split water was not on a-Fe2O3 and ZnO photoanodes; 2) exchange of the discussed. Benson, Briggs, and Wynne-Jones[57] deposited films electrolyte, specifically, substitution of potassium for 2 [54,66] [66b] from 0.1m Co(NO3)2 at 5–10 mAcm on platinum and nickel sodium and phosphate for phosphinate or borate. substrates. Deposition in KOH yielded a black film (like the The significance of point (2) is that neither potassium nor [54] CoCF ) with the composition K0.08CoO1.51·1.03H2O and suffi- phosphate itself is essential for catalytic activity. Likely, they are cient crystallinity to resolve the unit-cell parameters (hexago- not part of the catalytic unit itself. For discussion of the mech- nal; a =6.75 , c=5.36 ).[57a] Oxygen evolution is reported, anism, it may not be required to extend the Cooxo units (Fig- but the origin of the oxygen is not discussed.[57b] Better docu- ures 7 and 8) by a distinct cation, such as potassium, or anionic mentation of the oxygen evolution of an early CoO/OH film ligand, such as phosphate. Thus, we refer to this catalyst may be found in the patent application of Osamu et al.[58] Their herein as the cobalt catalyst film (CoCF), instead of cobalt crystalline films have the formula CoOm·nH2O(m=1.4–1.7, n= phosphate (Co–Pi) catalyst used elsewhere.

0.1–1.0). O2 evolution was described for operation in various The CoCF forms on a variety of conducting substrates, such nonbuffering electrolytes. as those discussed above. Surface morphology and film thick- Harriman, Pickering and Thomas compared the oxygen evo- ness depend on the deposition time, the composition of the lution of various metal oxide powders.[59] They used spinel- electrolyte, and the applied voltage.[63] The first layers of the II/III [64b] type Co3 O4, but IrO2 showed the highest activity. The rate of deposited films conform to the topology of the substrate.

O2 evolution for Co3O4 was given as 25.5 In thicker films, it is possible that only the ragged surface is 106 moldm3 min1.[59] At the beginning of the 1990s, Tseung electrochemically active.[64b] Thinner films, with deposition and co-workers published an extensive series of reports on re- times of 15 min, gave more stable performance.[64a] CoCF de- active deposition of cobalt oxides from 0.25 m CoCl2 at posited at voltages in the range of a prewave in the cyclic vol- 20 mAcm2.[60] The surface morphology of their electrodes tammogram assigned to Co oxidation are microscopically showed the same nodules as the CoCF described by Nocera smooth, whereas deposition at voltages which promote water and co-workers, and the reported overpotential of 0.39 V for oxidation produces CoCF with nodules on the surface.[54,63] The operation in 7m KOH was virtually identical to that of the overpotential to split water at pH 7 (for a current of m 2 [54] CoCF in 0.5 PO4 from KH2PO4 and K2HPO4 (pH 7). 1mAcm ) in the presence of KPi was reported as 0.41 V. Much of the more recent research has focused on crystalline Current densities as high as 100 mAcm2 were reported for un- [61] Co3O4. Singh et al. prepared spinel-type Co3O4 by micro- specified plastic composite electrolyzer cells operated at wave-assisted synthesis on a Ni support. The catalyst shows an 608C.[63b] For a thin film, we estimate the TOF to be as high as 2 1 apparent current density of 100 mAcm in 1m KOH at room 0.2 s per O2 molecule for the benign conditions reported in temperature with an overpotential of 0.24 V. They showed that Ref. [54] at 1.45 V vs NHE.[64c] the overpotential could be lowered to 0.22 V by doping with Both methylphosphonate and borate electrolytes foster cata- lanthanum. Jiao and Frei reported nanostructured Co3O4 in a lyst growth and support catalytic activity. These anions behave mesoporous silica support,[62] which has a turnover frequency in the same manner as phosphate but clearly differently from 1 2 4 [66b] (TOF) of 0.01 s per surface Co atom for an overpotential of nonbuffering electrolytes such as SO4 ,NO3 , and ClO . 0.35 V, at pH 5.8 and room temperature. The latter catalyst is Catalyst operation and formation does not require deionized III photochemically driven by a [Ru (bpy)3] species with visible reagent-grade water. Nocera and co-workers demonstrated op- light (476 nm, 240 mW). Frei estimated that a stack of 100 eration in brine and river water,[63,66b] suggesting that other

Co3O4 nanorod bundles would be needed for a TOF of ions, and especially chloride ions, do not inhibit O2 evolution. 100 s1 nm2 required to keep up with the solar flux.[62b] Potassium and sodium as components of the electrolyte were The cobalt catalyst for electrochemical water oxidation re- found to be interchangeable. Films grown with either alkali ported by Kanan and Nocera[54] has attracted much interest be- metal cation were catalytically active after cation exchange.[66a] cause of its efficiency at neutral pH, self-assembly from low- Leeching of both the alkali metal cation and the phosphate cost materials, and for its self-repair mechanism; it is discussed anion is faster than that of cobalt, as shown by radioactive la- in the following sections. beling experiments.[66a] This observation supports a cobalt-only framework for the catalyst, as proposed based on X-ray absorption spectroscopy (XAS)[55] and discussed below (sec-

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[67] [70] tion 2.2). Recently, McKlintock and Blackman proposed a long-range order. The FT of crystalline LiCoO2 is shown for new bidentate bonding mode of phosphate for the termina- comparison. In summary, films deposited at voltages lower tion of the Cooxo units proposed by Risch et al.[55] The dis- than the catalytic wave are microscopically smooth[63b] and tance between Co and P atoms is only 2.52 . The X-ray ab- atomically more ordered than films deposited under the same sorption spectroscopy (XAS) data in ref. [55, 68] does not pro- conditions but at voltages high enough to support catalytic ac- vide support for such a CoP distance, but phosphate coordi- tivity. The latter films show more disorder, not only on the mi- nation to a minor population of Co ions can not be ruled out croscopic scale (nodules) but also at atomic level (Figure 6). entirely.[55] The difference in the order at the atomic level suggests that Hill and co-workers recently reported an efficient water oxi- CoCF cannot be viewed as an aggregate of a multinuclear dation catalyst consisting of a soluble, carbon-free heteronu- cobalt complex of a distinct size. Either size and nuclearity of 10 clear molecule, [Co4(H2O)2(a-PW9O34)2] , in which a planar the molecular unit depends on the growth conditions or, more

[Co4(m-O)6O10] unit may represent the catalytically active likely, CoCF is a heterogeneous assembly of interconnected [69] part. It will be interesting to elucidate whether the mecha- {Cox(m2/3-O)y} units, as discussed in more detail in the following. nism of water oxidation in the molecular Co4W9 oxide and in XAS simulations on the cobalt catalyst films (CoCF) suggest CoCF is the same. that the central structural unit is a cluster of interconnected complete or incomplete CoIIIoxo cubanes (Figure 7).[55,68] The 2.2. Atomic structure of the catalytic cobalt oxide For the highly amorphous CoCF, diffraction techniques are not applicable. Thus, XAS is especially well suited to analyze the local structure of the cobalt metal site of thin disordered films. Figure 6 shows the Fourier transform (FT) of the extended X- ray absorption fine structure (EXAFS) extracted from XAS measurements at 20 K. Average bond lengths between the absorbing cobalt atoms and its neighboring ‘shells’ of atoms relate to peaks of the FT. The x axis shows reduced distances which are about 0.3 lower than the true nucleus–nucleus distance, which can be determined by EXAFS simulations. The amplitude of the peaks in the FT is a rough measure of the average number of the atoms at the respective distance. The peaks assigned to CoCo vectors of a CoCF deposited at 1.15 V vs NHE (Figure 6, red circles) are higher than for a film deposited at 1.35 V (Figure 6, black squares), indicating more extended

Figure 7. Possible structural motif, deduced from XAS data, for the bulk of the cobalt catalyst film (Co: dark grey, O: light grey). The catalytic film may be composed of a mixture of complete and incomplete cubanes. Presently, we cannot exclude that the CoCF material contains exclusively interconnect-

ed complete {Co4(m-O)4} cubanes or exclusively incomplete {Co3(m-O)4}cubanes. For an example of the latter, see Figure 8. The protonation states of the bridging oxides and terminal water species remain unknown.

octahedral geometry of the Co sites and the oxidation state found by XAS was confirmed recently by EPR spectroscopy.[71] In the case of incomplete cubanes, the extent of edge-sharing needs to be higher to comply with the simulation results. We modified the illustration from ref. [55] to emphasize that the catalytic unit most likely is not a distinct molecule. Figure 7 shows the proposed bulk structure of the CoCF catalyst, illus- Figure 6. Fourier-transform (FT) of an EXAFS spectrum of the cobalt catalyst trating the possible presence of complete and/or incomplete film (CoCF) for growth at a voltage not supporting water oxidation (1.15 V versus NHE; circles) and in the catalytic regime (1.35 V; squares) compared cubanes. Further interconnections of the building blocks may to crystalline LiCoO2 (triangles). LiCoO2 and CoCF are characterized by the result in an extended but overall disordered network of com- same octahedral Co building blocks, but in LiCoO2 those blocks are arranged plete and/or incomplete cubanes. III in an extended layer of side-sharing {Co O6} octahedra, which can be Here, we also consider another motif agreeing with the XAS viewed also as a layer of side-sharing, incomplete Co–oxo cubanes {Co (m- 3 data (Figure 8).[55,68] It is a tile-shaped Co O unit composed O)4}. Simulations of the data are shown as lines. All FT peaks shown exceed 10 32 the noise level; for further experimental details, see ref. [68]. of edge-sharing octahedra, which are exclusively incomplete

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gives them an advantage over heterogeneous systems and allows for a more systematic approach in the design of new generations of water oxidation catalysts with increased performance. However, the prediction of molecular structures that are able to catalyze the oxidation of water to oxygen at low overpotentials is rather difficult. Despite of their importance in the context of artificial photosynthesis purposes, only a relatively small number of molecular water oxidation catalysts is hitherto known, with the majority of them having been discovered only in the last few years. Based on the multinuclear structure of the oxygen evolving complex in photosystem II (PS II; see section 4) and the discovery of the nowadays well-known “ruthenium blue dimer” water 4+ oxidation catalyst [(bpy)2Ru(H2O)(m-O)(H2O)Ru(bpy)2] in the early 1980s (Figure 9),[73] it was initially assumed that, in molec-

Figure 8. Structure of m-oxo-bridged cobalt atoms, which is compatible with [55, 68] the EXAFS data (Co: drak grey, O: light grey). This {Co10O32} unit has an average of 3.8 CoCo vectors at 2.8 and 1.6 CoCo vectors at 5.6 (per Co atom) and thus satisfies the constraints resulting from simulations of the XAS data.[55,68] It is conceivable that several of these ‘tiles’ are connected to form an extended network or porous sheet. Layers of water molecules and cations may separate the Com-oxo sheets, in analogy to layered Mn or Co dioxides.[70,72] Figure 9. Examples of known dinuclear water oxidation catalysts.

Co-oxo cubanes. The macroscopic CoCF would be composed of a large number of these tiles; the space between them is ular complex catalysts, the presence of at least two metal cen- + 2 likely filled with cations (e.g. K ), anions (e.g. HPO4 ), and ters was essential to accomplish the mechanistically demand- water, in analogy to more-ordered layered structures, such as ing oxidative coupling of two water molecules to dioxygen. [72] LiCoO2 or Co-containing buserites. We emphasize that it Both the accumulation of the required amounts of oxidation would not be in conflict with the EXAFS data if some of the equivalents and a beneficial geometrical orientation of the two tiles were loosely interconnected. water molecules (or rather oxo ligands in later mechanistic The metal oxide/hydroxide materials that catalyze water oxi- steps) prior to the formation of the OO bond were believed dation frequently were operated electrochemically in an elec- to be more difficult to accomplish in mononuclear systems. Ac- trolysis experiment. However, it is unclear to what extent the cordingly, since the beginning of the 21st century, a number of concepts described in section 1 for surface oxides can be ap- dinuclear manganese and ruthenium complexes have been plied to the amorphous bulk metal oxides. Clearly, the CoCF shown to be more or less effective catalysts for water oxida- lacks the long-range order of crystalline solid-state materials. It tion, which seemed to confirm the prerequisite of at least two is an interesting conceptual question whether the CoCF can be metal centers in active water oxidation catalysts. Consistently, viewed best as an extended solid-state material or rather an mechanistic studies on Llobet’s dinuclear Ru–Hbpp water oxi- aggregate of multi-nuclear cobalt–oxo complexes with molec- dation catalyst (bpp= bis(2-pyridyl)-3,5-pyrazolate, Figure 9) ular properties. Water content and the analogy with layered have evidenced that oxygen–oxygen bond formation proceeds transition metal oxides suggests that water molecules fill the intramolecularly, benefitting from the favorable ligand matrix gaps between structures (Figures 7 and 8). Thus, it is conceiva- that preorientates the metal centers and bound H2O ligands of ble that not only the apparent surface catalyzes water oxida- this dinuclear complex for the coupling reaction.[74] However, tion but the bulk material can also contribute. In conclusion, water oxidation catalyzed by the blue dimer is nowadays be- the bulk metal oxides represent a catalytic material that may lieved to occur at a single Ru site within the dinuclear com- share properties with both heterogeneous electrocatalysts and plex, which raised the idea that water oxidation catalysis might homogeneous molecular catalysts. also be possible with mononuclear metal complexes.[75] In fact, within the last couple of years the discovery of several mono- 3. Molecular Catalysts for Water Oxidation nuclear water oxidation catalysts has been reported, which represents one of the major recent breakthroughs in the field In recent years, molecular catalysts for water oxidation have of water oxidation catalysis both from a mechanistic and a syn- also received a considerable amount of attention. The relative thetic point of view. However, the development of new and ef- ease with which molecular systems can be studied mechanisti- fective molecular water oxidation catalysts is not the only im- cally and can be tailored both structurally and electronically portant target to be pursued; the incorporation of the water

738 www.chemcatchem.org 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761 The Mechanism of Water Oxidation oxidation catalysts into integrated systems for solar water split- during the water oxidation cycle, given that, for steric reasons, ting has to be achieved, and also in this respect mononuclear an intermediate formation of dimeric or higher nuclear species complexes bear a lot of potential. Owing to their ease of modi- should be inhibited.[78b] fication and adaptation to the given needs, they are consid- To investigate the electronic influence of the ligand environ- ered as very promising candidates for surface grafting. ment on the catalytic activity of mononuclear ruthenium water oxidation catalysts, Thummel et al. then studied a wide range of related ruthenium complexes.[78b,79] One class of the investi- 3.1. Water oxidation catalysis by mononuclear metal com- gated compounds contained ruthenium centers coordinatively plexes saturated exclusively with N-donor ligands, mostly of the pyri- As a consequence of the increased interest in artificial photo- dine type, of varying denticity (Figure 10); a second series was synthesis in general and water oxidation catalysis in particular, characterized by the general formula [Ru(tpy)(NN)Cl](PF6) several review articles with emphasis on different aspects of (NN=bidentate N-donor ligand; Scheme 2). this reaction have been published recently, and these reviews have also taken molecular compounds into account.[76] Never- theless, the growing number of publications dealing with water oxidation certainly justifies an update of the field. In the past two years, several studies with respect to mononuclear water oxidation catalysts have been reported while only few novel dinuclear systems have appeared and most new findings concerning dinuclear systems were obtained in the course of continuing studies on already established systems.[74b,77] In this chapter we will therefore concentrate on mononuclear water oxidation catalysts (mostly ruthenium complexes) and mention dinuclear systems only briefly.

3.2. Mononuclear ruthenium complexes as water oxidation catalysts There are some early reports about water oxidation catalyzed by mononuclear ruthenium complexes, but their reliability, however, has been doubted.[76d] Not until 2005, Thummel et al. showed the mononuclear ruthenium aqua complexes [Ru(bina- 2+ py)(4-Rpy)2(H2O)] (Scheme 1) to be active water oxidation

Figure 10. Mononuclear water oxidation catalysts containing ruthenium centers coordinatively saturated with pyridine-type ligands. Turnover numbers are given in parentheses.

In contrast to 1–4, for which the electronic absorption, the redox properties, and the catalytic activities were strongly dependent on the nature of the axial 4-Rpy ligands,[78a] substitution at the equatorial ligands in 5–14 (Figure 10) did not have a strong or consistent influence on these properties; no correlation between the catalytic activity in water oxidation and the ligand set was revealed. At pH 1, TONs between 146 (10) and Scheme 1. Mononuclear water oxidation catalysts bearing the binapy ligand. Turnover numbers are given in parentheses. 416 (5) were determined for these catalysts in the presence of a 5000-fold excess of [Ce(NH4)2(NO3)6] as a stoichiometric oxidant.[79] catalysts in the presence of excess Ce4+ oxidant in aqueous For complexes 15–21 (Scheme 2), the dependency of the acidic media.[78] The catalytic activity of these complexes electronic absorption and redox properties on the ligand set is strongly depends on the nature of the 4-pyridyl substituent R more distinct. With increasing electron-withdrawing character giving turnover numbers (TON) between 20 (R=NMe2, that is, of the 4,4’-substituents at the bipyridyl ligand (bpy, 15–19), the + NHMe2 in acidic media) and 260 (R= CH3). From the observa- characteristic metal-to-ligand charge transfer (MLCT) band is tion that the sterically encumbered derivative 4 (Scheme 1) is shifted to longer wavelengths. At the same time, the oxidation also an active water oxidation catalyst (TON 260), it was con- potential of the RuII/III couple, as measured in acetonitrile, included that the mononuclear constitution of 4 is preserved creases. However, no direct correlation between the redox

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susceptible to coordination of a water molecule as an additional ligand. Subsequent loss of two protons and two electrons (proton-coupled electron transfer) generates a high-valent [O= VI 2+ Ru L6] species. Attack of a second water molecule on the electrophilic oxo group then leads to the formation of a IV + [Ru L6OOH] hydroperoxide intermediate. The loss of a second proton initiates the evolution of oxygen, which regen- erates the catalyst and thus closes the catalytic cycle. The hypothesis that the intermediate formation of a seven- coordinate Ru species during catalytic water oxidation is a viable mechanism, was recently supported by the identification of the seven-coordinate ruthenium aqua complex [m-(HO- IV 3+ HOH)(Ru {6,6’-(COO)2bpy}(pic)2)2] ·2H2O 23 (pic =4-picoline; Figure 11) as an intermediate of water oxidation catalyzed by

Scheme 2. Mononuclear water oxidation catalysts of the type [Ru(t- py)(NN)Cl]+. Turnover numbers are given in parentheses.

properties and the catalytic activity of these complexes could be drawn: Both the presence of electron-donating (e.g. CH3 in

16) or electron-withdrawing (e.g. NO2 in 18) 4,4’-substituents led to lower TONs as compared to the parent system 15. A remarkably high activity (TON= 1170) was found for 20, although the reason for the distinct activity in this special case remained unclear. Exemplarily, mechanistic studies were performed for

15. The initial rate of O2 evolution was found to be first order in catalyst concentration, indicating that dimerization is not a Figure 11. Seven-coordinate intermediate of water oxidation catalyzed by prerequisite for these complexes to be active in water oxida- II [Ru {6,6’(COO)2bpy}(pic)2] 22. tion. Additional evidence for a mechanism in which the origi- nal monomeric nature of the catalyst is conserved was ob- III 2+ II tained by the isolation of [Ru (tpy)(bpy)Cl] as an intermedi- [Ru {6,6’-(COO)2bpy}(pic)2] 22; accordingly 23 itself is catalyti- ate from a reaction mixture, which had undergone approxi- cally active, too.[80] It should be noted, however, that in this mately 10 water oxidation turnovers. This complex was identi- case (consistent with the dimeric structure of 23) the kinetic fied as such by mass spectrometry and also chemically by data and also the results of calculations[81] suggest that dimeri- showing that addition of ascorbic acid as a reductant leads zation of 22 plays an important role in the course of the reac- back to the RuII complex 15, as evidenced by 1H NMR spectros- tion (see below). copy. On the basis of these findings and supporting DFT calcu- Assuming that none of the N-donor functions temporarily lations on 5, a mechanism for catalysis of water oxidation by dissociates from the Ru centers to create an empty site for complexes 5–21 was proposed, which involves the formation water coordination, a mechanism like that outlined in of seven-coordinate ruthenium centers during the course of Scheme 3 might indeed be followed for catalysts 5–14. For the 4-electron oxidation of water (Scheme 3).[78b] complexes of the type 15–21 ([Ru(tpy)(NN)Cl]), however, II Accordingly, removal of two electrons from [Ru L6] generates recent findings of Sakai et al. and especially of Meyer et al. in- IV 2+ a 16-electron [Ru L6] species (Scheme 3), which should be dicate a different mechanism of water oxidation. Sakai et al. independently investigated the related mononuclear ruthenium complexes [Ru(tpy)(bpy)Cl]+, 15, and [Ru(t- 2+ py)(bpy)(H2O)] , 26, as well as their dinuclear congeners [Ru(t- 2+ 4+ py)(Cl)L(Cl)(tpy)Ru] and [Ru(tpy)(H2O)L(H2O)(tpy)Ru] (L= bis[5-(5’-methyl-2,2’-bipyridinyl)]ethane; Scheme 4).[82] The potential of these complexes as water oxidation catalysts was investigated in water/acetonitrile mixtures at pH 0.4

using [Ce(NH4)2(NO3)6] as the oxidant, monitoring oxygen evolution with a Clarke-type electrode in solution. Interestingly, the dinuclear ruthenium complexes were found to be less active catalysts than the related mononuclear ones. The aqua Scheme 3. Water oxidation mechanism proposed for mononuclear rutheni- complexes 24 and 26 proved superior to the corresponding um complexes exhibiting Ru centers coordinated by N-donor ligands. chloro complexes 25 and 15, which was the main finding of

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Scheme 4. Mono- and dinuclear complexes of the [Ru(NNN)(NN)X]n +-type. Turnover numbers are given in parentheses. this study. A 2–3 h induction period for the onset of water oxidation was observed when the chloro complexes 25 and 15 were employed as the catalysts. This observation may be ex- plained by an equilibrium between the chloro and aqua complexes in aqueous solution, with the aqua form of [Ru(tpy)(b- py)X]n+ being the active species for water oxidation. Accord- ingly, the addition of NaCl to an aqueous solution containing [Ru(tpy)(bpy)Cl]+ completely inhibited the formation of oxygen.[82] II n+ An extensive study of the [Ru (NNN)(LL)(H2O)] aqua system was performed by Meyer et al., who investigated several related water oxidation catalysts, differing in the nature of the tridentate NNN and bidentate LL ligands (Figure 12).[75,83] From electrochemical and kinetic studies on the complexes 2+ 2+ [Ru(tpy)(bmp)(H2O)] 27 and [Ru(tpy)(bpz)(H2O)] 28 c, a detailed mechanism for the oxidation of water, with Ce4+ as the Figure 12. Meyer’s mononuclear ruthenium water oxidation catalysts of the [83c] type [RuII(NNN)(LL)(H O)]n +. No maximum turnover numbers were reported oxidant, was derived (Scheme 5). In diluted acidic media 2 for these catalysts. 28a=26. DMAP= 4-dimethylaminopyridine. m II 2+ IV 2+ (0.1 HNO3) 27 ([Ru OH2] ) is rapidly oxidized to [Ru =O] by two equivalents of Ce4+ . Subsequent reaction with further Ce4+ leads to the formation of a high-valent [RuV=O]3+ species, Meyer’s investigations clearly evidence that water oxidation which reacts with water to give the terminal hydroperoxide at single metal sites is not only possible with Ir complexes (see [RuIIIOOH]2+ . The latter is further oxidized to the peroxo com- below), but also at ruthenium centers. The easier synthetic IV 2+ plex [Ru (O2)] , whose decomposition, releasing molecular di- access to mononuclear Ru compounds of the type shown in oxygen, is the rate-limiting step of water oxidation in diluted Figure 12 (as compared to the known dinuclear systems) dis- 4+ HNO3. As a result of the stronger oxidizing power of Ce at burdens the precise tailoring of the ligand sphere for the link- IV 2+ V 3+ lower pH, oxidation of [Ru (O2)] to [Ru OO] competes age to surfaces, for the synthesis of molecular assemblies, or with its decomposition in 1m HNO3. Oxygen evolution from for the systematic investigation of electronic factors governing the RuV-peroxide is followed by subsequent oxidation of [RuIII- the catalytic activity. For instance, the complex cation 29a 2+ IV 2+ 2+ OH] to [Ru =O] , which closes the catalytic cycle. The plau- PO3H2 [Ru(Mebimpy){4,4’-(H2O3PCH2)2bpy}(H2O)] (Scheme 5) sibility of the proposed intermediates was supported by DFT has been successfully grafted on FTO, ITO or TiO2 nanoparticle calculations, which suggested that the initial intermediate surfaces to result in materials with good electrocatalytic activi- [RuIIIOOH]2+ is best described as a terminal peroxide, whereas ties.[84] the peroxide ligand of [RuIVOO]2+ is coordinated in a side-on Recently, Meyer et al. found that the presence of redox me- 2+ fashion. diators such as [Ru(bpy)2(LL)] (LL=bpy, bpm, bpz) provokes Similar mechanistic investigations for some other [Ru(N- a significant rate enhancement in water oxidation, when the 2+ NN)(LL)(H2O)] derivatives suggested that the mechanism de- ruthenium blue dimer and [Ce(NH4)(NO3)6] are used as the cat- picted in Scheme 5 should be generally valid for catalysts be- alyst and stoichiometric oxidant, respectively.[73c] Consequently, longing to this class of complexes, albeit with changes in the the effect of such mediators was studied for the mononuclear individual reaction rates and rate limiting steps.[83b] systems as well.[85] Molecular catalyst–mediator assemblies of

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to impressively high turnover numbers up to 28 000 and turnover rates up to 0.6 s1 in water oxidation experiments for the latter systems without any signs of activity loss after longer peri- ods of time.[85] The presence of the redox mediator in the heterogeneous assembly allowed water oxidation catalysis to be performed at 1.43 V, which is close to the thermodynamic po-

tential for the O2/H2O couple at pH 0. Notably, under these conditions, no electrocatalytic current was detected for solutions of the related monomeric catalyst Scheme 5. Mechanism for water oxidation catalyzed by Ru–aqua complexes [Ru(NNN)(LL)(H O)]n+, as proposed by 2+ 2 [Ru(tpy)(bpm)(H2O)] . Meyer’s Meyer et al. recent findings doubtlessly represent major progress towards the design of integrated water- II II 4+ the type [(bpy)2Ru (bpm)Ru (LLL)(OH2)] (LLL= tpy: 30; LLL= splitting devices based on artificial photosynthesis, for which Mebimpy: 31), in which a mediator and a water oxidation cata- the availability of robust electrocatalysts is an essential prereq- lyst are combined within a single molecule, were obtained uisite. II from the reaction of [Ru (bpy)2Cl2]·2H2O with Besides the mononuclear ruthenium complexes stemming [RuII(LLL)(bpm)Cl] + (bpm =2,2’-bipyrimidine) in ethanol/water from the extensive studies performed by the groups of Thum- mixtures (Scheme 6). In this case, the bipyrimidine ligand, be- mel and Meyer, a few other examples have been shown to be longing to the redox mediator as well as to the catalyst, plays active as water oxidation catalysts, too. In addition to the the important role of bridging the two functional units within aforementioned investigations on 15 and 24–26, Sakai et al. re- the dinuclear ruthenium complex. ported ruthenium aqua complexes of the type [Ru(NNN)(4,4’- 2+ Cyclic voltammetry (CV) analysis of 30 at pH 1 revealed mul- R2bipy)(H2O)] containing the facially coordinating 1,4,7-tri- tiple sequential one-electron oxidation steps until formation of methyl-1,4,7-triazacyclononane (tmtacn) ligand rather than the III V 6+ the species [(bpy)2Ru (bpm)Ru (tpy)(O)] , which triggers cata- meridional tpy or Mebimpy derivatives, which were employed lytic water oxidation. In these systems the reaction of by Thummel and Meyer (Scheme 7).[86] III V 6+ [(bpy)2Ru (bpm)Ru (tpy)(O)] with water was found to be the As a consequence of the lacking p acceptor and the more key OO bond formation step, leading to the hydroperoxo pronounced s-donating properties of the tmtacn ligand, the III III 5+ II III species [(bpy)2Ru (bpm)Ru (tpy)(OOH)] . The latter then un- electrochemical potentials of the Ru (H2O)/Ru (OH) and dergoes further oxidations by Ce4+ until oxygen is released, RuIII(OH)/RuIV=O couples of 32–34, as determined in aqueous presumably from the peroxidic species acidic solution, are strongly shifted to lower potentials, for ex- III V 6+ 2+ [(bpy)2Ru (bpm)Ru (tpy)OO)] , resulting in the formation of ample, in comparison to [Ru(tpy)(bpy)(H2O)] 28a. On the II IV 4+ [(bpy)2Ru (bpm)(tpy)Ru =O] , which closes the catalytic cycle. other hand, from these reduced oxidation potentials, no lower- Anchoring the closely related redox assemblies [{(4,4’- ing of the overpotential for electrocatalytical water oxidation II II 4+ (H2O3PCH2)2bpy}2Ru (bpm)Ru (NNN)(OH2)] (NNN=tpy, Me- followed. For example, for 34, bearing the bpy ligand with the bimpy) to ITO or TiO2/FTO (nanoparticle TiO2 films on FTO) led most strongly s-donating substituent, the potential that had to be applied to initiate water oxidation was found to be of a similar magnitude to that required by the tpy derivative 28 a. For the derivatives 32 and 33, even higher potentials were necessary. Water oxidation experiments were also performed in

diluted aqueous solutions with [Ce(NH4)2(NO3)6] as a chemical oxidant, and oxygen evolution was monitored with a Clarke- type electrode. From the total amounts of oxygen that had evolved after a 40 h period, turnover numbers amounting to 148 (32), 173 (33), and 251 (34) were determined, revealing an increasing catalytic activity with increasing electron-donating capabilities of the 4,4’-bpy substituents. However, the activities Scheme 6. Catalyst-mediator assemblies for water oxidation catalysis. were again lower than that for the tpy derivative 28a; for this

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water oxidation with 22 as the catalyst is based on a binuclear process, which very recently has also been found feasible by calculations:[81] Dioxygen was proposed to be formed by coupling of two RuIV–oxyl radicals. Mass spectroscopic investigations with the aim of identifying intermediates of the catalytic IV + cycle revealed a peak for [Ru {6,6’-(COO)2bpy}(pic)2(OH)] , and consistently a compound containing two of these cationic complex units, that is, the abovementioned seven-coordinated ruthenium complex [m-(HOHOH)(RuIV{6,6’- 2 + Scheme 7. Water oxidation catalysts of the [Ru(NNN)(NN)(H2O)] type, con- (COO)2bpy}(pic)2)2](PF6)3·2H2O, 23(PF6)3, (Figure 11, Figure 13) taining the facially chelating tmtacn ligand. Turnover numbers are given in parentheses.

complex, the same authors determined a TON of 310 under analogous conditions. Accordingly, the replacement of the tpy ligand by tmtacn leads to a lowering of the Ru-based oxidation potentials in 32–34, although they were found to be generally less active in water oxidation than 28 a. To incorporate the water oxidizing process into a light driven integrated system for water splitting, the generally applied cerium(IV) oxidant has to be replaced by a photosensitizer, which takes over the responsibility of providing the necessary amounts of oxidation equivalents required by the water oxidation catalyst. For this purpose, a lowering of the oxidation potentials of the catalysts to a certain degree can be advanta- geous and, as can be seen from the investigations on complexes 32–34 this goal can be achieved by introduction of stronger s-donating and less p-accepting ligands. A different Figure 13. Molecular structure of [m-(HOHOH)(RuIV{6,6’- strategy for lowering the oxidation potentials of mononuclear (COO)2bpy}(pic)2)2](PF6)3·2H2O, 23(PF6)3, a possible intermediate of water oxi- Ru-based water oxidation catalysts was followed by Sun et al., II dation catalyzed by [Ru {6,6’-(COO)2bpy}(pic)2] 22. The hydrogen atoms of who, as mentioned before, studied the catalytic properties of the organic ligands and the (PF6) counterions are omitted for clarity. II [80a,b] the complex [Ru {6,6’-(COO)2bpy}(pic)2] 22 (Scheme 8).

was isolated after addition of NH4PF6 to the reaction mixture. IV + In 23, two [Ru {6,6’-(COO)2bpy}(pic)2(OH)] units are bridged by a proton sitting in the middle between the two OH ligands. Two water molecules are bonded to the two ends of the internal [HOHOH] moiety by bridging hydrogen atoms, and additional hydrogen bonds between these water molecules and the 6- and 6’-carboxylate groups of adjacent complex cations further stabilize the overall hydrogen bonding network as well as the dinuclear structure of 23.[80a] The active participation of 23 during water oxidation was not only suggested by the second order kinetic dependence of II Scheme 8. Synthesis of the mononuclear water oxidation catalyst [Ru {6,6’- the decay of Ce4+ in catalyst 22 but further evidenced by (COO) bpy}(pic) ] 22. Turnover number is given in parentheses. 2 2 proving the ability of 23 to catalyze water oxidation itself: Oxygen evolution was observed in an aqueous acid solution

As expected, the presence of two negatively charged car- when [Ce(NH4)2(NO3)6] was added as an oxidant. Although boxylate functions as part of the bpy-type ligand in 22 signifi- 23(PF6)3 does not represent the first example of a structurally cantly lowers the oxidation potentials of the ruthenium center, characterized seven-coordinate ruthenium complex, its suc- as compared, for example, to 28a. Water oxidation was in- cessful isolation and identification as an intermediate in water duced by addition of 22 to an aqueous solution containing oxidation, catalyzed by 22, certainly contributes valuable infor- Ce4+. However, only a rather modest TON of approximately mation with respect to possible reaction pathways for the con-

120 was determined by monitoring oxygen evolution with a version of H2OtoO2 at transition metal sites. However, it Clark-type oxygen electrode in solution. Kinetic analysis re- should be borne in mind that all other mononuclear rutheni- vealed a second order decay of Ce4+ for 22 indicating that um water oxidation catalysts that have been discussed to date

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 743 H. Dau, C. Limberg, P Strasser et al. obey first-order kinetics of oxygen evolution with respect to the reaction mixture for a certain number of times. Other deac- III 2 the corresponding catalysts, rendering a participation of dimer- tivation pathways, valid for the Co and for the S2O8 system, ic {Ru2} species during the catalytic process unlikely for those include the oxidative decomposition both of the photosensitiz- 2+ systems. Therefore, water oxidation starting out from mononu- er and the catalyst. Notably, in the absence of [Ru(bpy)3] ,no clear ruthenium catalysts, but proceeding via intermediate for- photodegradation of 22 occurred under similar reaction condi- mation of dinuclear species, rather represents a special case. tions. Regarding 22, dimerization certainly is favored by the presence A very different and novel reaction scheme for water oxida- of the 6,6’-carboxylate groups in the ligand backbone, which tion was discovered by Milstein et al., who showed that the contribute to a stabilization of the dinuclear Ru–hydroxo struc- mononuclear PNN–ruthenium pincer-type complex 34 is capa- ture by offering base functions for hydrogen-bonding interac- ble of performing the entire water splitting reaction tions. (Scheme 10).[87] Remarkably, 34 mediates the consecutive gen- Under neutral aqueous conditions, cyclovoltammetric analysis of solutions containing 22 as an electrocatalyst showed a catalytic water oxidation curve beginning at 0.98 V (vs NHE). This relatively low potential for water oxidation allowed the III 3+ utilization of [Ru (4,4’-R2bpy)3] (R= H, CH3) as photosensitiz- ers to realize photochemical water oxidation. The latter was performed in a three-component system composed of catalyst 2+ 22, [Ru(4,4’-R2bpy)3] , and [Co(NH3)5Cl]Cl2 or Na2S2O8 as sacrificial electron acceptors at pH 7.0.[80b] In the absence of light, no formation of oxygen could be detected. Irradiation, however, immediately triggered oxygen evolution, which rapidly ceased when the light was turned off (Scheme 9). Scheme 10. Water splitting at a single ruthenium center.

eration of hydrogen and oxygen, probably via a H2O2 intermediate, from water at a single metal center via thermally and photolytically induced steps. A central role within these trans- formations is played by the cooperative participation of the PNN ligand, which undergoes facile deprotonation and proto-

nation at the methylene unit of the Py’CH2PR2-arm. Milstein’s system has found considerable attention and recent highlight articles reflect the great potential attributed to this novel water splitting scheme,[88] although the overall reaction is ac- Scheme 9. Photochemical water oxidation in a three-component system tually stoichiometric and slow with rather low yields in gener- consisting of catalyst 22, photosensitizer [Ru(R bpy) ]2+ and a sacrificial elec- 2 3 ated oxygen and hydrogen.[87] tron acceptor. Redox potentials for the onset of oxygen evolution and for 2+/3 + Quantum chemical calculations have further elucidated the the [Ru(R2bpy)3] couple are given vs NHE at pH 7. mechanism of this new system.[89] The release of hydrogen,

which is induced by proton transfer from the (Py’CH2PR2) arm Initial TOFs of greater than 550 h1 and greater than of the aromatic PNN ligand in 34I to the ruthenium-bound hy- 1 III 2 1250 h were determined for the Co and S2O8 systems, re- dride, was found to be the rate-limiting step of the reaction. spectively. However, both systems suffered from rapid deacti- The presence of water reduces the activation barrier for this vation, and oxygen evolution stopped after only 7 min for CoIII step from 37.6 to 33.6 kcalmol1 (1 kcal =23.9 kJ) by the for- 2 or 50 s for S2O8 , under the diluted conditions required for mation of a CH···OH2···HRu bridge between the methylene oxygen detection in the liquid phase. However, when water proton and the {RuH} unit. Nevertheless, this relatively high oxidation was performed in solutions containing higher activation barrier might be the reason for the slow rate of H2 amounts of catalyst, oxygen evolution, as monitored by GC, generation and the low yield of 34II (45%) even after 3 days persisted for about 2 h and a TON of 100 was determined for at 1008C. Time-dependent DFT calculations on 34II suggested III the Co system. The fairly quick loss in activity on employment that the photochemical reductive elimination of H2O2, as the of Na2S2O8 as the sacrificial electron acceptor could be traced ultimate source of O2 (see above), from 34II, for which two back to the continuous acidification of the solution in the strong RuOH bonds need to be broken, occurs from a disso- course of water oxidation [Equation (3) and Scheme 9], which ciative triplet state via a singlet–triplet crossing. does not occur with [Co(NH3)5Cl]Cl2, which simultaneously acts as a buffer. Consequently, it is possible to regenerate the pho- tocatalytic activity of the persulfate system by neutralization of

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3.3. Non-ruthenium systems

In contrast to the rapidly developing field of ruthenium-catalyzed water oxidation, examples of catalysts containing other metal centers remain scarce, both for multinuclear and mononuclear systems. Hence, most of the known water oxidation systems still rely on highly expensive and/or rare transition metals, and this reliance can be considered one of the major drawbacks regarding future application of artificial photosynthesis for power generation. Nevertheless, some examples of Figure 14. Mononuclear iridium-based water oxidation catalysts. water oxidation or water splitting promoted by mononuclear complexes of metals other than ruthenium have been reported recently and are discussed in this section. tron-donating pentamethylcyclopentadiene (Cp*) ligand Purely photochemical water splitting has been carried out (Figure 14).[92] with osmocene as a photocatalyst (Scheme 11).[90] Photolysis of Oxygen evolution experiments were performed in diluted solutions using a Clark-type oxygen electrode, which revealed complexes 37–39 to be much more active catalysts than 36. Initial rates amounting to 0.91 and 0.28 turnovers per second were determined for 37 and 39, respectively, which are significantly higher than that which the same authors determined 1 for 36 (R1 =R2 =H; 0.005 s ) under analogous conditions. However, the turnover numbers attained by these new iridium catalysts (1500 for 37, 213 for 39) were lower. Oxygen evolution was found to be first order in the Ir catalyst, which renders it plausible that a similar mechanism for water oxidation is fol- Scheme 11. Hypothetical catalytic cycle of water splitting using osmocene lowed as in the case of the earlier discussed mononuclear as a photocatalyst. II n+ ruthenium complexes [Ru (NNN)(LL)(H2O)] 27–30 (see above). DFT calculations on the potential IrV intermediate [Cp*IrV(ppy)(O)]+ revealed the presence of a low lying p* (Ir=O) IV + an acidic aqueous solution containing [Cp2Os H] , 35 I, with orbital, which includes a significant contribution from the O white light induced the evolution of hydrogen with concomi- atom. Formation of an OO bond might therefore result from IV 2+ tant formation of the osmocene(IV) complex [Cp2Os (H2O)] . nucleophilic attack of a water molecule on this high-valent iri- It was assumed that, on photolysis, hydrogen is released from dium–oxo unit.[92] IV IV 2+ dimeric [Cp2Os H···HOs Cp2] units, leading initially to The use of defect polyoxometalate (POM) derivatives with III III 2+ [Cp2Os Os Cp2] , which then undergoes photodisproportio- vacant metal sites as ligands for transition metals is an interest- IV 2+ nation to give [Cp2Os (H2O)] . The latter was found to be ing approach for the design of water oxidation catalysts with photostable under acidic conditions, but when an aqueous so- enhanced stability, because decomposition routes based on lution of the closely related hydroxo complex oxidative degradation of organic ligands is not possible. There- IV [Cp2Os (OH)](PF6), 35 II(PF6), was photolyzed, a nearly quantita- fore, the recent discovery that the purely inorganic Ru poly- II 10 tive conversion of 35 II to oxygen and [Cp2Os ] resulted. Photo- oxometalate [{Ru4O4(OH)2(H2O)4}(g-SiW10O36)2] represents an oxidation of 35 II probably involves the intermediate genera- active water oxidation catalyst, has stimulated ongoing re- III III [69,77a,b,93] tion of the peroxo complex [Cp2Os (O2)Os Cp2]. The results of search efforts on this class of catalysts. As a mononu- these two separately studied photochemical processes may be clear (with respect to the catalytic site) representative, the iridi- formally combined to an interesting cyclic process um polytungstate K14[(IrCl4)KP2W20O72]·23H2O, 40, was synthe- 9 (Scheme 11). Difficulties associated with the pH-dependent for- sized by the reaction of IrCl3 with [PW9O34] in the presence of [94] mation of oligomeric (m-OH), (m-O), and (m-Cp) osmocene spe- KOH in water. Low yields of O2 (30 %) were detected by GC IV 2+ 3+ cies in aqueous solutions of [Cp2Os (H2O)] , however, have so when [Ru(bpy)3] and 40 were mixed at neutral pH. However, far impeded the establishment of a real catalytic photochemi- 40 was found to be unstable in solution, dissociating into [90] 13 cal water-splitting system. [IrCl4(H2O)2] and [KP2W20O72] . Although this dissociation was Bernhard et al. recently discovered a new family of iridium- found to proceed much slower than the oxidation of water in III + based water oxidation catalysts [Ir {2-(4’-R’Ph)(5-Rpy)}2(H2O)2] , the presence of 40, it could not be ruled out that, as a conse-

36, for which impressively high TONs of up to 2760 (R1 =Me, quence of dissociation, small amounts of catalytically active [76a,91] R2 =F) were reported (Figure 14). However, the rates of IrO2 are responsible for the formation of the detected oxygen. oxygen evolution were relatively low. Related mononuclear Ir By using ligands containing xanthene and corrole units complexes 37–39 were obtained by formally replacing one 2- linked to each other by amide bonds, kermark et al. synthe- (2-pyridyl)phenyl (ppy) and one water ligand in 36 by the elec- sized the mono- and dinuclear manganese complexes 41 and 42 (Scheme 12).[95] By means of cyclic voltammetry on basic

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which remain very scarce, in spite of the importance of catalysts based on cheap and abundant materials.

3.4. Conclusions In recent years, increased research efforts on the development of new molecular water oxidation catalysts exhibiting enhanced performance and stability has led to a notable progress in this area, resulting in the discovery of some new families of catalysts, along with a deeper insight into the mechanisms that can lead to the generation of oxygen from two water molecules. The realization that dimer- Scheme 12. Mn–corrole complexes active in water oxidation. ic or higher-order assemblies of multiple metal centers are not

CH2Cl2/CH3CN solutions, both complexes were shown to act as essential to accomplish water oxidation catalysis certainly rep- water oxidation catalysts, with 42 being more active than 41. resents a major advance. The easier derivatization associated The activity of the mononuclear manganese complex 41 indi- with structurally simpler mononuclear systems allows for a cated that water oxidation might proceed—similarly to what more rational synthesis of tailored systems to study the struc- has been considered a possible OO bonding step in natural ture–activity relationships that might govern water oxidation. photosynthesis—via a single-site mechanism involving nucleo- Similarly, an introduction of surface-binding groups can facili- philic addition of water or OH to a high-valent mangane- tate the strong attachment of molecular catalysts to electrode se(V)–oxo unit. DFT calculations supported this single-site surfaces or enable the synthesis of molecular assemblies suited mechanism as the preferred pathway with both 41 and 42.[96] for coupling light-harvesting and/or redox mediation to water Recently, the same authors reported further experimental evi- oxidation. As the final goal is to use the electrons resulting dence for such a water oxidation mechanism, demonstrating from the oxidation of water to drive fuel-forming reactions, that rapid oxygen generation resulted from the addition of growing efforts are dedicated to the coupling of the different V Bu4NOH to a solution containing [Corrole–Mn =O] 43a (Cor- processes involved in artificial photosynthesis. In this context, role=5,10,15-tris(4-nitrophenyl)corrole) in acetonitrile light-driven water oxidation mediated by suitable photosensi- (Scheme 12).[97] Ultimately, however, only very small amounts tizers plays a central role, as does the attachment of molecular of dioxygen, corresponding to 0.01 turnovers of 43, were de- catalysts to conducting electrode surfaces. Inspection of the tected. By using 18O-labeled water, it could be proved that one new generation of water oxidation catalysts, however, reveals of the oxygen atoms in the evolved oxygen originated from them to be mostly based on expensive ruthenium metal cen- water. On the basis of the results of these 18O-labeling studies ters. Therefore, the discovery of effective and long-lived water in combination with UV/Vis spectroscopy and high-resolution oxidation catalysts derived from abundant and low-priced first- mass spectrometry investigations, it was proposed that, after row transition metals remains a major objective for future re- an initial oxidation of the [MnIII–Corrole] complex 43 by search. Advances concerning this aspect might be considered tBuOOH to give 43 a, attack of a hydroxide ion at the electro- indispensable in terms of a significant contribution of artificial philic MnV=O function leads to the MnIII–hydroperoxide species photosynthesis applications for sustainable power generation 43b. The latter undergoes further oxidation, presumably by in the near future. 43a, yielding the MnIV–peroxide 43c before the elimination of oxygen occurs. Although the efficiency of oxygen generation 4. Biological Water Oxidation in the Mn–corrole/tBuOOH system is only very low, the finding that water oxidation at manganese centers can proceed via Plants, algae, and cyanobacteria use solar energy to oxidize nucleophilic addition of hydroxide to high-valent manganese– water and reduce plastoquinone molecules. The light-driven oxo groups, that is, at single sites, in a similar manner to the water–plastoquinone oxidoreductase is denoted as photo- mononuclear Ru and Ir systems, is an important finding not system II (PSII).[98] This cofactor–protein complex of impressive only regarding the debate about the actual mechanism of nat- size and complexity is embedded in a lipid bilayer membrane ural water oxidation in PSII; it might also stimulate the devel- (thylakoid membrane), which separates the aqueous phases of opment of more manganese-based water oxidation catalysts, the lumen and stroma compartment (Figure 15). Crystallo-

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Hereafter, the states on the left and right side are denoted ·+ ·+ as YZ and YZ , respectively. YZ formation is followed by elec-

tron transfer from a pentanuclear {Mn4Ca} complex, which represents the catalytic center of the photosynthetic water oxidation. Water oxidation involves four sequential steps of one- ·+ electron oxidation of the catalytic complex by YZ , resulting in the following overall equation:

þ þ 4YZ þ2H2O ! 4YZþ4H þO2 ð36Þ

·+ The maximal value of the redox potential of the YZ /YZ pair is around +1.2 V vs NHE.[104] It resembles the anodic potential in electrochemical water oxidation. The lowest pH for efficient water oxidation by PSII is around 5.5,[105] which corresponds to

an equilibrium potential of the H2O/O2 couple of about 0.9 V

Figure 15. Redox factors in PSII. The D1 and D2 protein subunits carry all es- (at pO2 =1 bar). The potential difference of about 0.3 V (= sential redox factors. Besides the D1 protein, the CP43 subunit also provides 1.2 V0.9 V) can be viewed as the ‘overpotential’ of photosyn- ligands to the {Mn4Ca} complex. The proteins of the PSII core complex carry- thetic water oxidation. The corresponding energetic efficiency ing antenna pigments (CP43, CP47) and peripheral antenna proteins are not calculated for the half-cell reaction [Equation (11) in section 1.1] shown. PSII is part of the thylakoid membrane, a dense array of proteins and lipids separating the aqueous compartments at the PSII acceptor side is 80 % when the second half-cell reaction is assumed to be (lumen) and donor side (stroma). proton reduction, and around 75 % when referenced against plastoquinone reduction by PS II. These figures are impressive, especially when taking into account that water oxidation in graphic models of the core complex of PS II have been deter- PSII is so fast that it efficiently outcompetes recombination- mined[99] and, more recently, a resolution of 2.9 was ach- loss pathways.[105b] ieved.[100] The PSII core complex is most likely evolutionarily The turnover frequency (TOF) of the complete PSII may be 1 fully conserved with respect to every site of direct functional as high as 100 s per O2 molecule; it is determined by rate- importance. In intact organisms, additional pigment-carrying limiting reactions at the acceptor side. For catalytic processes 1 proteins are found, which increase the light-harvesting capabil- at the donor side, a maximal TOF of about 400 s per O2 mol- ity of the core complex. Composition of these peripheral an- ecule can be calculated from the rate constants of the elec- tenna proteins is species-dependent and relates to differences tron-transfer steps. The maximal turnover number (TON) of the in color observed when comparing taxonomic groups of pho- catalyst is determined by the so-called PS II repair cycle (see tosynthetic organisms (e.g. green, red, brown, gold, or blue section 4.2) and may be around 105. For a review of the ener- alga).[98, 101] getics and solar energy conversion efficiency of PSII, see Light absorption by one of the numerous chlorophyll ref. [104c]. The catalytic process is discussed in the following groups (Chl) of PSII is followed by ultra-fast excitation energy section.[106] transfer and excitation of a special pair or group of chlorophylls denoted as P680.[102] The excited-state energy of P680* 4.1. Structure of the catalyst drives a sequence of electron-transfer (ET) steps as well as the proton-coupled redox reactions at the electron acceptor and In 2001, Zouni and co-workers presented the first crystallo- donor sides of PS II. Firstly, by ultrafast ET, a pheophytin group graphic model (at 3.8 resolution) of a PSII complex in which + (Phe) is reduced, resulting in formation of the P680 Phe radi- the position of the catalytic site, the {Mn4Ca} complex, was cal pair. At the acceptor side, this process is followed by rapid uniquely determined.[99a] Higher resolution resulted in structur- reduction of QA (ca. 300 ps), a firmly bound plastoquinone (Q) al models for the Mn complex and its ligand environment, followed by slower ET to QB, a second plastoquinone molecule showing that the protein-derived ligands are carboxylates sup- (ca. 1 ms; Figure 15). A second photon drives formation of a plemented by a single histidine residue.[99b–d,100,107] However, all doubly reduced and protonated QH2 molecule, which leaves crystallographic data related to a non-native arrangement of the QB site; a third quinone in PSII, QC, may also be in- the Mn ions. The high X-ray intensity needed for data collec- [100,103] III,III,IV,IV volved. tion results in rapid X-ray photoreduction of the {Mn4 }to At the donor-side, P680 + is reduced by electron transfer a level where all Mn ions are present in the MnII state.[108] from a specific redox-active tyrosine, denoted as YZ, within less Whereas Mn reduction and the loss of the internal m-oxo struc- than 1 ms. This ET step is assumed to be coupled to the move- ture of the Mn complex by X-ray irradiation during crystallo- ment of the phenolic proton within a hydrogen bond towards graphic data collection is well documented,[108c] it is largely un- a nearby histidine sidechain: clear to what extent the coordinates of the Mn ligands are modified. One Ca ion is likely connected by two or more bridging P680þþY OH NHis ! P6800þY O Hþ NHis ð35Þ Z Z oxides to Mn ions, as suggested by EXAFS spectroscopy[109]

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 747 H. Dau, C. Limberg, P Strasser et al. and crystallography.[99b,d] After substituting chloride by bro- mide, the question of halide ligation was addressed by EXAFS at the Br K-edge using an approach previously established for a bromoperoxidase and synthetic model complexes.[110] The results show that the halide is not a first-sphere ligand of manganese and is possibly at a distance of about 5 to either a Mn or a Ca ion.[111] However, recent results obtained by protein crystallography suggested a metal–halide distance of about 6.5 .[100, 112] Mostly on the basis of the EXAFS results, the group of Klein,

Sauer, and Yachandra suggested the presence of two di-m2- oxo-bridged Mn dimers (2MnMn distances of 2.7 ), which are connected by a mono-m-oxo, bis-carboxylato bridge (Mn Mn distance of 3.3 ), such that a C-shaped structure resulted.[115] Simulations of EPR data[116] and the crystallographic results[99] are difficult to reconcile with a C-shaped dimer-of- dimers model. In 2004, Barber and co-workers presented a crystallographic model involving a {Mn3Ca(m3-O)4} cubane connected to the fourth Mn ion by a bridging oxide.[99b] This Figure 16. Model of the catalytic site for water oxidation in PSII. The Mn [117] ligand environment was obtained by molecular-mechanics modeling, start- model was used as a starting point for DFT calculations and ing with the crystallographic coordinates of Loll et al.[99d] The model of the [106b,117–118] [113] mechanistic considerations, but was difficult to recon- {Mn4Ca(m-O)n} core is based on EXAFS results. The shown protein-derived cile with EXAFS results.[108a,119] Moderate modifications resulted groups are from the D1 subunit, unless labeled as being CP43 side chains. in a model which may be compatible with the EXAFS data.[120] The structure in the S1 state is shown. The m2-oxo group between Mn3 and Mn4 is proposed to be protonated in the S0 state (m2-OH). In the S2 !S3 tran- Recently, XAS data were collected on single crystals of sition, a hydroxide bridging between Mn1 and Mn2 may become deproton- [119] PSII, thereby providing further constraints for structural ated and transformed into a sixth ligand of Mn1.[114] A cavity close to Mn4 models of the Mn complex, which complement earlier results could harbor several water molecules (dotted line). The Asp61 is close to obtained by XAS on isotropic[114c,121] and unidirectionally orient- Mn4 and has been proposed to represent the entrance of a path for proton [108a,121b,122] translocation to the aqueous lumen. Accumulation of oxidizing equivalents ed PSII samples. Comparison of the single-crystal ·+ by the Mn complex involves electron transfer to YZ , the oxidized form of ·+ EXAFS with a set of hypothetical structures led to a hypothesis the Tyr160/161 (YZ). Due to the vicinity to the Mn complex, YZ formation could affect pK values of groups in the Mn ligand environment, electrostati- on the structure and orientation of the {Mn4Ca(m-O)n} core of the Mn complex; the protein-derived ligand environment was cally and through hydrogen-bond networks. We emphasize that several aspects of the described model are merely hypothetical. not modeled.[119] An alternative working model of the Mn complex and its ligand environment is shown in Figure 16;[113] similar structures have been invoked in computational studies.[123] ligand environment may relate to a metal–(m-O) core that re- The following structural features of the Mn complex of PSII sembles an intrinsically stable motif also found in mineralic Mn [127] have been identified unequivocally: The {Mn4Ca} complex is oxides; the Mn complex of PSII may resemble a minuscule bound to several carboxylate oxygens as well as to a single N- piece of rock. Nonetheless, its oxidative assembly and water- histidine; the metal ions are internally interconnected by bridg- oxidation function clearly require the presence of most of the ing oxides. There are 2 or 3 di-m-oxo bridges between Mn ions ligands provided by the (apo)proteins. The striking structural involving m2-O and presumably also m3-O bridging. The Ca ion similarities have inspired the recent investigations on water ox- could cap the bridging oxides (Figure 16). idation by transition metal oxides which resemble minerals (see below). Neither the Mn complex nor the PS II protein complex are as 4.2. Formation of the catalytic metal complex stable as the rock. The so-called PS II repair cycle involves ex-

Assembly of the {Mn4Ca} complex of PSII does not require any change of the D1 protein against a fresh copy; the D1-turnover chaperone-like enzymes but is driven by light and thus denot- frequency may be as high as 2 per hour for organisms exposed ed as photoactivation (of PSII).[124] The mechanism involves the to high light intensities.[128] The D1 protein is likely exchanged 2+ ·+ oxidation of Mn ions by YZ formed upon illumination and after oxidative damage to chlorophylls, quinines, or amino acid [129] formation of the {Mn4Ca(m-O)n} core within the ligand system residues. It is unlikely that the costly exchange of the D1 provided by the protein. The affinity of the apoprotein to Mn protein is aiming to repair the Mn complex. Repair of the ions is so high that stoichiometric amounts are sufficient for {Mn4Ca(m-O)n} core is likely achieved in a different manner; complete assembly (4 Mn ions per PSII); a {Mn2(m-O)2}complex when exposed to light, the photoactivation mechanism contin- has been identified as a likely intermediate of the assembly uously counteracts the occasional loss of Mn ions from the process.[125] Some amino acid residues that are assumed to complex, representing a cost-efficient and inconspicuous ligate Mn ions can be exchanged against nonligating residues repair mechanism of the metal complex itself. In conclusion, and the Mn complex is assembled nonetheless, albeit with the {Mn4Ca(m-O)n} core of the catalyst is formed after oxidation lower stability.[126] This relative insensitivity to details of the of Mn2+ ions by self-assembly at the ligand system provided

748 www.chemcatchem.org 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761 The Mechanism of Water Oxidation by the apoprotein. The oxidative self-assembly also may repre- an especially comprehensive investigation by Haumann et al., sent an efficient self-repair route. both Mn oxidation-state changes and structural changes were addressed for all transitions between semi-stable S-states, on the basis of data collected at liquid helium temperatures and 4.3. Kok’s basic reaction cycle: Accumulation of oxidizing at room temperature.[114b] The authors concluded that the oxi- equivalents dation-state changes of the Mn complex are accompanied by

Four stable or semistable intermediate states of the reaction deprotonation of a bridging hydroxides in S0!S1 and by for- cycle of the photosynthetic Mn complex can be populated by mation of additional di-m-oxo bridges in S2 !S3 (Sche- saturating laser flashes of visible light (Scheme 13). These are me 14).[106a,114b,121b]

Scheme 13. S-state cycle of photosynthetic water oxidation (Kok cycle).

Starting in the dark-stable S1 state, absorption of a photon causes formation of Y ·+ within less than 1 ms. Reduction of Y ·+ by ET from the Mn complex Z Z Scheme 14. Structural changes of the Mn complex of PSII. Only three of the results in the S !S transition; typical time constants of the ET step are in- i i+1 four Mn ions are shown because this is sufficient to illustrate the changes in dicated. A plausible set of oxidation-state combinations of the four Mn ions the bridging type in the course of the S-state cycle. The indicated distances is shown. The presence of MnII in S and of a ligand radical in S (as well as 0 3 have been determined by EXAFS spectroscopy. Reprinted from ref. [114b] delocalized valencies) also have been proposed. with permission of the American Chemical Society. the S-states of Kok’s basic reaction cycle: S (dark-stable), S ,S, 1 2 3 4.4. Redox-potential leveling and S0; the subscripts indicate the number of accumulated oxi- [130] dizing equivalents. In the dark-stable S1 state, the likely oxi- In each Si !Si+1 transition, the Mn complex is oxidized by the III IV [114c] ·+ dation state of the manganese ions is {Mn2 Mn2 }, but con- same oxidant, namely YZ . The requirement of a low overpo- flicting opinions have been expressed.[131] Currently, it is mostly tential thus implies that several single-electron oxidations pro- III IV assumed that each Si !Si+1 transition involves Mn !Mn oxi- ceed at about the same redox potential of the Mn complex. If dation (at least up to the S3 state). For the S2 !S3 transition, it the four Mn ions were electronically isolated, we might ob- is disputed whether Mn-centered[114b,c,132] or ligand-centered[133] serve four essentially independent MnIII !MnIV oxidations at oxidation takes place. Favoring the former mechanism, we the same potential. However, four isolated Mn ions are an un- arrive at the oxidation-state combination shown in Scheme 13. likely water oxidant and, moreover, the short MnMn distances This oxidation-state cycle (Scheme 13) bears an important found by EXAFS[121b,c,136] and EPR results[116,131b,137] suggest that mechanistic consequence; before the onset of water oxidation, the four Mn ions are strongly coupled, most likely via bridging the catalytic metal complex accumulates oxidizing equivalents oxygen atoms (see above). Consequently any oxidation of one by increasing the oxidation state of the Mn ions. There is no Mn ion can be predicted to increase the oxidation potential for evidence for partial oxidation of water at an early state in the the next single-electron oxidation of the complex significantly, reaction cycle (no indications for OH·, terminal oxyl or peroxide as reported for numerous synthetic m-oxo bridged com- [138,139] creation before S4-formation). plexes. The structure of the Mn complex has been investigated by For example, a first oxidation of the Mn complex will in- X-ray spectroscopy for all four semistable S-states, mostly by S- crease the oxidation potential for the next step by 0.5– state population at room temperature followed by rapid freez- 1.5 V,[138] unless prevented by an appropriate charge-compen- ing and XAS measurements at 10–20 K.[114b,121b,134] By laser-flash sating change in the structure of the Mn complex. The redox- ·+ illumination during the X-ray experiment, data collection at potential difference between YZ /YZ and the Si/Si+1 redox cou- room temperature became possible. The room-temperature re- ples is probably lower than 100 mV,[140] so an increase by 0.5 V sults exclude that the structural data collected at 20 K were af- or more would prevent the Si+1!Si+ 2 transition, suggesting fected by a temperature-dependent redox isomerism.[114b,135] In that ‘redox-potential leveling’ by appropriate coupling of Mn

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p d oxidation-state changes to charge-compensating deprotona- The difference of Em and Em (Scheme 15) was chosen to be MeCom tions or more complex structural changes is essential for the 0.6 V (e DEm = 0.6 eV) and consequently the pK values for catalytic metal complex to function.[141] Investigations on syn- deprotonation in the reduced and oxidized state differ by 10 thetic complexes suggest that Mn oxidation m-OH deprotona- units (DpKred/ox =10). A similar DpKred/ox was determined for a [142] [142e] tion or other structural changes leading to an additional m- specific [Mn2(m-O)(m-OH)] complex. [139] oxo bridge represent potential-lowering steps that can effi- In PSII, the S0!S1 transition may involve m-OH deprotona- [114b,121b,134a] ciently counteract the increase in the redox potential for a sub- tion (Scheme 14) whereas the S1!S2 transition is sequent Mn oxidation step. probably not associated with any deprotonation of the metal Scheme 15 illustrates the relation between the redox poten- complex.[143] Thus, for the (a)-set of pK values in Scheme 15, III tial and pKA values for oxidation of a [Mn 2(m-O)(m-OH)] com- the transition from A via C to D corresponds to the S0!S1 plex. In Scheme 15, each rectangle of states represents a ther- transition in PSII, whereas the transition from D to F models [144] S1!S2. Concerted electron–proton transfer (CEPT) can be ruled out for this set of pK values. Only for an alternative set of pK values (denoted as (b) in Scheme 15), the transition from C to F could, at neutral or acidic pH, involve CEPT because the paths via D or E are energetically strongly disfavored (uphill to D and E). At this point, a clarifying comment on the nomenclature may be appropriate. Distinctions are frequently made between (1) electron transfer followed by (de)protonation (ET/PT), (2) proton transfer (PT) preceding the ET step (PT/ET) and concerted transfer of an electron and a proton by simultaneous tunneling of the two particles through energy barriers (ETPT or, preferably, CEPT).[144] The acronym PCET (proton-coupled electron transfer) comprises all ET/PT, PT/ET, and CEPT. These three modes of coupling electron and proton transfer are often experimentally not easily distinguishable and theoretically not mutually exclusive. Therefore, the use of the term PCET is appropriate and convenient whenever a clear-cut distinction cannot be made. Scheme 15. Relation between deprotonation and successive oxidation steps The above considerations suggest how redox-potential level- for a hypothetical Mn complex. Two sets of pK values for the protonated m- ing might be realized in PS II. The interrelation of ET and depro- oxo bridge are indicated, which are denoted as (a) and (b). For (a), the tran- n + tonation steps is clearly crucial. The concerted transfer of an sitions A!C!D!F model the following transitions in PSII: S0 !S1 ! n + electron and a proton (CEPT) may be involved in redox-poten- S1 !S2 . Concerted electron–proton transfer (CEPT; dotted gray arrow from C to F) represents an option for the second set of pK values (b), but is irrele- tial leveling, but sequential oxidation and deprotonation steps vant for the first one at neutral pH (a). (ET/PT or PT/ET) could perform the same role. Further insights into the ways and modes of redox-potential leveling, resulting from investigations on synthetic complexes modynamic cycle. For example, for the cycle made up of states or computational chemistry, would be desirable. The role of F, E, D, and C, the free-energy change for going from C to F via the nuclearity of the metal complex is largely unknown (typi- D must be identical to that when going via E. The free-energy cally binuclear complexes have been studied). It is also still un- change ET resulting in oxidation of the metal complex (vertical clear whether deprotonation of a terminally coordinated water transitions) is given by Equation (37): ligand is sufficient for complete redox-potential leveling. In PSII water oxidation, m-OH deprotonation may facilitate redox- o ox MeCom DG e ðEm Em Þ ð37Þ potential leveling in the S0!S1 transition; deprotonation coupled to formation of new m-oxo bridges may serve the same purpose in the S !S transition (see above). where Eox and EMeCom are the midpoint potential (in volts) of 2 3 m m A mere oxidation step results in a positive charge of the the oxidant and of the metal complex, respectively, and e is metal site, whereas all redox-potential lowering steps are asso- the elementary charge. The free-energy change of the depro- ciated with ‘de-charging’ typically achieved by removal of a tonation step (horizontal transitions) is given by Equation (38): proton. Anion binding or cation release could also facilitate

þ charge compensation, but are unlikely to play a role in PSII DGH 0:06 eV ðpHpK Þ ð38Þ i i water oxidation. Even though redox-potential leveling clearly relates to charge compensation, a discussion solely in terms of Thus, for each rectangle in Scheme 15 holds: the energetic penalty to be paid for solvation of a positive charge in a low-dielectric protein medium would be mislead- MeCom red=ox: e DEm ¼ 0:06 eV DpK ð39Þ ing. The solvation penalty may contribute, but is probably not

750 www.chemcatchem.org 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761 The Mechanism of Water Oxidation the major determinant of the redox-potential increase resulting from a mere oxidation step. An uncompensated one-electron oxidation will always result in an increase of the potential for the second oxidation, irrespective of the dielectric properties of the medium surrounding the metal complex.

4.5. ‘Smart’ proton removal Concerted transfer of an electron and a proton has been discussed in the context of redox-potential leveling.[144] Oxidation ·+ of the Mn complex by ET to YZ is most likely directly coupled to the shift of a proton from His190 to Tyr160 (YZ ; Figure 16) and possibly also to a similar shift of a second proton within the catalytic site.[145] However, a discussion in terms of a concerted Scheme 16. Extended S-state cycle model. Kok’s classical S-state cycle multisite transfer of one electron and two protons is insuffi- model[130a] is extended by including not only 4 oxidation but also 4 deprotonation steps.[130b,147] Four of the nine intermediates are stable for tens of sec- cient in providing a complete picture of the relation between n + + n onds or fully dark-stable (S1 ,S2 ,S3 , and S0 ). They correspond to the electron and proton transfer in biological water oxidation. The states S1,,S2,S3, and S0 of Kok’s reaction cycle. Electrons and protons are re-

Mn complex of PSII is embedded deeply in the PSII protein. moved alternately from a Mn complex which comprises the {Mn4Ca(m-O)n} Long-distance inner-protein proton transfer typically requires core, the protein environment and bound water molecules. We emphasize tens or hundreds of microseconds. Consequently, ET by tunnel- that proton removal from the catalytic site does not necessarily imply deprotonation of water species. The energetics of the oxidation and deprotona- ·+ ing of the electron from the Mn complex to YZ and the tion steps are described by Equations (37) and (38), respectively. The nine in- +/n proton translocation from the Mn complex to the aqueous termediate states of the Mn complex are denoted as Si , where the sub- phase can not occur concertedly, but need to proceed sequen- script gives the number of accumulated oxidizing equivalents and the super- script indicates the relative charge: positive (+) or neutral (n) relative to the tially. For the transition S !S !S +O , time-resolved X-ray ab- 3 4 0 2 dark-stable S state.[106a] The shown reaction cycle has been discussed pre- [135d] 1 sorption spectroscopy and complementary fin- [130b,147] + + viously using a nomenclature where the nine states from S0 to S4 [130b,143,146] dings have provided evidence for the following se- were numbered consecutively and denoted as I0–8. quence of events: Absorption of a photon results in formation ·+ of YZ within less than 1 ms. The electric field emanating from ·+ the positive charge at YZ affects pK values and drives several 4.7. Water deprotonation and oxidation: Four basic routes proton hopping steps and eventually results in a transfer of a proton from the Mn complex to the aqueous phase, within Three crucial and interrelated aspects of water oxidation have + n n [135d] about 200 ms(S3 !S3 ;S3 also has been denoted as S4 ). so far been discussed herein: 1) accumulation of oxidizing The potential-lowering deprotonation facilitates the subse- equivalents by transition metal oxidation; 2) redox-potential ·+ n n n [135d] quent ET to YZ (S3 !S4 ;S4 previously denoted as S4’ ) leveling; 3) ‘smart’ proton removal prior to the onset of the O followed by rapid onset of water oxidation. The reaction se- O bond-formation step. All three points relate to deprotona- quence is concluded by the release of a second proton. tion at the catalytic site. In this section, we address the relation

The discovery of a proton-first ET in the S3!S4!S0 transi- of the four deprotonation steps to the process of forming di- tion led us to an extended working hypothesis on the se- oxygen from the two molecules which often are denoted as quence of electron and proton transfer steps (Scheme 16).[147] ‘substrate water’. Four prototypical routes towards OO bond We emphasize that the proton removal steps relate to the mul- formation are shown in Scheme 17. tistep transfer from the catalytic site to the aqueous phase, Route A (Scheme 17) involves partial water oxidation by per- which can not be coupled directly to the ET step. However, it oxide formation at an early stage, that is, after accumulation of + n is conceivable, for example in the S2 !S2 transition, that a three oxidizing equivalents in the S3 state. This route was dis- proton acceptor B is ‘prepared,’ which renders possible the cussed with regards to water oxidation at mono-nuclear metal ·+ concerted transfer of an electron to YZ and the shift of a complexes (see sections 3.2 and 3.3) and is also considered to proton to B . be of high relevance in electrochemical water oxidation [Equa- Concerted electron–proton transfer may also be favorable in tions (25)–(28) in Section 1.3]. Early peroxide formation has water oxidation by artificial catalysts. However, oxidation of also been discussed in the context of PSII water oxida- [145,148] the metal complex directly coupled to a proton transfer to a tion. However, peroxide formation in the S3 state of the water molecule is an unlikely event in the acidic and neutral PSII manganese complex, in conflict with the results obtained pH regime because a water molecule is an unfavorable proton by X-ray absorption spectroscopy,[114b,121b,132,133b] implies Mn re- acceptor (formal pK of H2O/H3O couple is about 1.7). This im- duction in the S2 !S3 transition: plies that, also in nonbiological water oxidation, ‘smart’ removal of protons from the catalytic site may be an issue when fMn III,IV,IV,IVðO2Þ gþY þ !fMn III,III,IV,IVðO1Þ gþY ð40Þ aiming at fast and efficient water oxidation. 4 2 Z 4 2 Z

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formation step. It relates to the ‘alkaline mechanism’ in the older electrochemical literature, in contrast to the ‘acidic mechanism’ of Route B, and typically is not considered a relevant pathway in the context of photosynthetic water oxidation. In route D (Scheme 17), accumulation of oxidizing equivalents is not coupled to substrate-water deprotonation. In- stead the catalyst is deprotonated, for example, by deprotonation of bridging oxides. Thereby ‘acceptor bases’ are created, which facilitate a direct coupling of water oxidation and water deprotonation in the OO bond- formation step. The acceptor- base hypothesis was put forward to explain water oxidation in PSII.[106a,121b,141] A combination of

Scheme 17. Four routes towards water oxidation. The catalyst is symbolized by the grey shape. In Si, the number early water deprotonation of accumulated oxidizing equivalents always is equal to i. However, accumulation of oxidizing equivalents (by oxi- (route B) and acceptor base ac- dation of metal ions) is explicitly indicated only in D in form of the encircled ‘+’ symbols. A–D differ in the mini- cumulation (route D) has also mal number of water and proton binding sites (n). been considered.[117,121b] A more detailed discussion is provided below. In route B (Scheme 17), water oxidation is facilitated by a Experimental results relating to the mode of substrate-water regular pattern of water-deprotonation steps, eventually result- coordination and deprotonation are rare. Presently, the most ing in two oxides terminally coordinated to two high-valent important results may be the water exchange rates determined 16 18 [152] metal ions. The oxo group may have radical character and thus in elegant H2O /H2O exchange experiments, which have might be better described as an oxyl radical (e.g. MnVO! revealed the presence of two different binding sites for the MnIVO·). Route B is appealingly straightforward and has been two substrate-water molecules with distinguishable exchange invoked frequently in mechanistic models of water oxidation rates. Typical results for the S-state dependence of the ex- at electrode surfaces by metal complexes in solution (homoge- change rates are summarized in Table 1.[152f] neous catalysis) and also in PSII. Also, the influential ‘hydrogen atom abstraction’ model of PSII water oxidation, suggested by Babcock and co-workers,[149] involves route B. However, termi- Table 1. Rate constants for exchange of substrate-water molecules (de- [152f] nal oxo formation is predicted to strongly enhance the so- tected for thylakoid membranes from spinach, T= 108C). The kfast de- called pre-edge peak in the X-ray absorption spectra,[150] but crease in S2 !S3, by a factor of 3, corresponds to marginal changes in binding energy. this is not observed.[114b,132b,133b] Consequently, route B is unlike- 1 ly to be followed by PSII. The unique aspect of Babcock’s hy- kslow/fast of water exchange [s ]S0 S1 S2 S3

substrate pothesis is the concerted transfer of one electron (Mn oxida- Slow HnO ca. 10 ca. 0.02 ca. 2 ca. 2 · substrate tion) and one proton (water deprotonation) to YZ amounting Fast HnO – >120 ca. 120 ca. 40 to a hydrogen atom transfer in each Si !Si+1 transition. In this case, the oxidized tyrosine residue is not a radical quasi-cation ·+ (not YZ with the phenolic proton staying in close to the phe- The exchange rates of water bound to a pentanuclear metal nolic oxygen), unlike in Equation (35). According to Babcock’s complex buried inside a protein are not well understood,[152b,f] hypotheses, the phenolic proton is released to the bulk phase so any seemingly straightforward interpretation might also be before hydrogen atom transfer can take place, transforming misleading. Therefore, the following discussion is tentative the initial radical cation into a neutral radical. This hydrogen only; some aspects of the results in Table 1 have been inter- · [99a,b] [152b,f] atom abstraction by YZ is in serious conflict with structural preted differently. and spectroscopic data.[143,150–151] To relate the results summarized in Table 1 to mechanistic In route C (Scheme 17), water oxidation involves bound or proposals on photosynthetic water oxidation, we will use the free hydroxides that act as proton acceptors in the OO bond- following guidelines:

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G1) Oxo groups in bridging position between high-valent Assuming that a bridging hydroxide is deprotonated in S0 ! 2 1 III,IV [114b,121b,134a] Mn ions exchange extremely slowly (< 10 s in {Mn2 } and S1 (Scheme 14), G1 (with the results summarized in 6 1 IV,IV [153] <10 s in {Mn2 }). Table 1) implies that, rather than substrate-water deprotona- G2) Deprotonation of any water bound to a metal ion tion, an unprotonated m-oxo group is formed that can serve as should increase the binding energy and thus decrease its ex- a proton acceptor in S4!S0+O2 (formation of an acceptor change rate strongly. base). Similarly, in S2 !S3 another acceptor base is likely

G3) Partial water oxidation by peroxide or radical formation formed, whereas in the S1!S2 transition there is no deprotona- (O·) should affect the substrate water exchange rate. tion event (Scheme 16). The available data is compatible with G4) Oxidation of Mn(i)III!IV should increase the binding the following notion: In the three Mn-oxidizing S-state transi- strength of any water species coordinated to Mn(i) and thus tions (from S0 to S3) no substrate water is deprotonated but decrease the exchange rate strongly (i.e. by several orders of rather two bases are created that can accept a proton from magnitude). water in the process of OO bond formation. The employment

The decrease in kslow in the S0 !S1 transition might indicate of several bases as proton acceptors in the OO bond forma- water deprotonation, but the subsequent acceleration is in- tion reaction could be crucial (acceptor-base hypothesis; compatible with persistent deprotonation of the substrate route D in Scheme 17). water molecule. In explaining these variations in kslow, any Most likely, OH ions cannot diffuse within the PSII protein simple model will fail. A complex and largely hypothetical ex- to the catalytic site. However, H2O molecules certainly move to planation follows (Scheme 18): The deprotonation of a m-oxo the catalytic site effectively. A water molecule may bind to the

Mn complex immediately after completion of the O2-formation step to a vacant coordination site. The pK of the bound water changes drastically resulting in deprotonation and formation + of a bound hydroxide (MnOH2 !MnOH+H ). Transfer of the proton to the aqueous bulk phase completes the process. The net effect of this hypothetical sequence of events would be hydroxide binding to the Mn complex. Such an indirect hydroxide-binding step has not been supported by experimental + n findings, but cannot be excluded for the S0 !S0 transition (Scheme 16). Similarly, it represents an open question whether + n the proton removed from the Mn complex in the S3 !S3 transition stems from a substrate-water deprotonation.

4.8. Mechanism of OO bond formation Scheme 18. Hypothesis on the structural changes underlying the peculiar variations in the water exchange rate of the slowly exchanging water mole- Numerous hypotheses on dioxygen formation in PSII have n n cule (H2Os). The m-hydroxo deprotonation in the S0 !S1 transition facilitates been put forward; see ref. [154] for a summary of 34 early sug- formation of an additional hydrogen bond; ‘charging’ of the complex by oxi- gestions and ref. [106b] for a discussion of more recent hy- n + dation, in S1 !S2 , weakens this H-bond. potheses. It is mostly assumed that, in PSII, O2 formation proceeds via transient formation of a peroxidic intermediate, possibly in a two-electron step.[155] Scheme 19 shows prominent [106a,121b,134a] bridge in the S0 !S1 transition facilitates the forma- hypotheses on the mode of peroxide formation in the S4!S0 tion of a new hydrogen bond, thereby binding the substrate transition of PSII. water molecule more strongly. The MnIII!IV oxidation in the Since the late 1970s, peroxide formation between two bridg-

S1!S2 transition is not coupled a deprotonation, so the charge ing oxygen atoms (Scheme 19 A) has been suggested in nu- of the metal complex will increase. Consequently, the pK value merous studies.[156] Analogous chemistry has been demonstrat- of a bridging oxide connected to the oxidized Mn ion will ed for bis(m-oxo)dicopper(III) complexes and their transforma- change strongly[142e] (Scheme 15) and render the m-oxygen a tion to the (m-h2:h2-peroxo)dicopper(II) form.[157] Oxygen radi- weak hydrogen-bond acceptor, thereby causing an increase in cals bridging between Mn ions have been invoked.[121a] Mecha- the exchange rate of the water molecule. nistic models centered around a cuboid/butterfly In conjunction with the results summarized in Table 1, G1 transformation involve OO bond formation between oxygen implies that the substrate water molecules are not in a bridg- atoms bridging between Mn ions.[158] This bond-formation ing position between Mn ions. G2 is in conflict with deproto- mode is considered only rarely in the context of photosynthet- nation of the slowly exchanging water in S0 !S1 or S1!S2, and ic water oxidation, as, among other reasons, it is in conflict substrate excludes deprotonation in S2 !S3 for both HnO groups. with G1 and also difficult to reconcile with G2 and G4. G3 provides an argument against peroxide or O·substrate forma- The nucleophilic attack of a hydroxide upon a terminal tion in the transitions for which water deprotonation is exclud- oxide (MnV=O) or oxyl radical (MnIVO·; Scheme 19 B), has been ed by G2. G4 suggests that substrate water molecules are not proposed repeatedly and is supported by molecular models [106b] bound to the Mn ions that are oxidized in S2!S3. (see sections 3.2 and 3.3). The attacking hydroxide may be

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an unconventional mode of peroxide formation (Sche- [123b] me 19 E). In this model, the S2 !S3 transition involves a ‘re- construction’ of the complex, which is associated with binding of the second substrate-water molecule. It is unclear whether this and other aspects of the model can be reconciled with the water-exchange rates (Table 1). As early as the late 1980s, Krishtalik argued that a two-electron step resulting in peroxide formation and possibly coupled to the transfer of a proton to an acceptor base that is part of the Mn complex represents a favorable path towards dioxygen formation.[155,159] It presumably relates to conceptual problems that such two-electron/n-proton transitions are not typically considered in explicit mechanistic models of water oxidation. Further development in the theory of such multiparticle reactions may be required before complex multielectron/multipro- ton steps can be addressed in computational studies. The new mechanistic proposal in Scheme 20 is inspired by the hydrogen-bonding pattern found in a dimeric unit of Ru complexes (Figure 13).[80a,81] The starting points for peroxide

Scheme 19. Proposed modes of peroxide formation in the S4 !S0 transition. The upper structure in each panel (and the only one in C) refers to a state formed in the S3 !S4 transition. This state is formed by oxidation of the Mn complex and possibly also by the relocation of 1 or 2 protons to 1 or 2 bases (not shown). The lower structure shows the peroxidic state, which is assumed to form transiently. For peroxide formation between two oxo or oxyl groups, each terminally coordinated to a MnV/IV ion, see route B in Scheme 17.

a ligand bound to Mn or to Ca, or an outer-sphere water. In ref. [123d] (Scheme 19C), it is suggested that the slowly exchanging water is terminally coordinated to the Ca ion of a · {CaMn3(m-O)4} cube, whereas the terminal O is formed in the

S3 !S4 transition by deprotonation of a hydroxide, terminally Scheme 20. Hypothesis on the coupling of electron transfer and shifts of coordinated to Mn4 in Figure 16. In this and a related protons in hydrogen-bond networks for peroxide formation in the S !S [117] 4 0 model, the Arg357 of the CP43-protein (not shown in transition. (s and f on oxygen represents slow and fast, respectively) Figure 16) is crucial in ‘managing’ the proton removal from the catalytic site. Scheme 19D illustrates a specific proposal involving an formation are two fully protonated water molecules. Together [121b] outer-sphere water species attacking a terminal oxyl, which with Mna, these molecules form a motif (H2OHOHMn) has been formed in the S3 !S4 transition by oxidation and de- which resembles Zundel’s hydronium ion (H2O-H-OH2). In this protonation of a terminally coordinated water species (not quasi-Zundel ion, the OO distance may be as small as 2.4 , shown). A unique aspect of this proposal is the coupling of so that the transition state for formation of a peroxidic OO peroxide formation to the transfer of a proton to a bridging bond is reached more readily. The bond formation process is oxide, possibly in form of a concerted ‘removal’ of one electron assumed to involve the transfer of HB2 and two electrons, re- and one proton from the outer-sphere water (CEPT), which for- sulting in m-O protonation and Mn reduction (Scheme 20), and mally amounts to abstraction of a hydrogen atom. Specifically, can thus be viewed as a hydride abstraction by the Mn com- the inner-sphere electron transfer from oxygen to a MnIV ion is plex. Moreover, this OO bond formation by hydride abstrac- coupled to proton transfer to its m-O ligand (Scheme 19D). tion is assumed to be either directly coupled (concerted pro- Owing to the relation between the pK value of the m-oxo cess), or preceded or followed by the transfer of HB1 and HB3. B4 group and the oxidation potential of the metal complex The subsequent transfer of H is required for completion of O2 (Scheme 15), a concerted mechanism could be especially ad- formation, but is not indicated in Scheme 20. This proposed vantageous in facilitating the critical peroxide-formation mechanism is speculative and is presented to illustrate two step.[106a,121b,141,155,159] points: Recent DFT calculations resulted in especially low energetic a) Two fully protonated water molecules could represent the barriers in OO bond formation for a mechanistic model with starting point for OO bond formation.

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b) The relocation of several protons from water to acceptor differ from synthetic paths that have been followed for the bases likely involves a specific hydrogen-bond network; the preparation of molecular metal complexes (although the latter coupling of proton and electron transfer in the OO bond may can also be achieved employing high oxidation state precur- be crucial in avoiding high energetic barriers. sors already). We note that the strong coupling between the pK of m-oxo The oxidation of metal surfaces and the formation of bulk bridges and the midpoint potential for reduction of the metal metal oxides proceed in the absence of any elaborate ligand ions (Scheme 15) can render unprotonated oxo groups suitable system. The oxide formation can be viewed as an oxidative acceptor bases in a multielectron/proton transfer reaction. self-assembly and implies a self-repair mechanism (see sec- The OO bond formation chemistry remains insufficiently tion 2.1). In PS II, an analogous self-repair could occur but is understood. We tentatively conclude that the {Mn4Ca} complex obscured by the intricate and costly exchange of the protein not only positions the substrate-water molecules by hydrogen subunit that provides most ligands to the {Mn4Ca} complex bonding and terminal ligation to a metal ion, but also partici- (see section 4.2). In homogeneous catalysis, loss and rebinding pates actively in formation of the OO bond by alteration of of metal ions to the ligand system might occur. However, pres- its metal–oxo core, specifically by bridging-type changes. The ently this question is of little relevance because, most likely, ox- OO bond formation step itself may involve the coupled trans- idative modification/destruction of the ligand system limits the fer of electrons and protons from water to the catalysts. maximal turnover number much more severely than loss of metal ions.

5. Common Themes and Unifying Concepts 5.1. Nuclearity and structure of the catalyst: obligatory 5.3. Potential-determining step: Redox-potential leveling m-oxo-bridging? The thermochemical analysis of electrocatalytic water splitting The structure of the photosynthetic manganese complex is on solid surfaces reveals that an ideal reversible catalyst should characterized by m-oxo bridges between metal ions. Hereafter, exhibit four proton-coupled electron transfer steps with identi- we do not discriminate between m-O, m-OH, and m-OH2 bridg- cal DGi for each of the four steps (see section 1.3) so that the ing. The putative mechanistic role of m-oxo bridges in redox- same minimal electrical potential suffices to drive each of the potential leveling and as proton acceptors in OO bond forma- elementary steps. In real catalysts, deviations from the equal tion is discussed in section 4. In electrochemical water oxida- DGi distribution result in a potential-determining step; the re- tion at metal electrodes, all metal surfaces are transformed action with the highest DGi determines the overpotential. into metal oxides at the positive potentials needed for water Each of the four discussed steps [Equations (25)–(28)] repre- oxidation; it is likely that extensive m-oxo bridging is prevalent. sents a deprotonation-coupled oxidation, either of the active How deep the oxide layer penetrates the metal electrode is site at the electrode surface or of a water species itself. Identi- largely unknown and probably depends on both the metal cal DGi in an ideal electrocatalyst implies that the values of the used and the operation conditions. For mineral, amorphous or individual (electrical) reversal potentials, Vrev(i), are identical for microcrystalline metal oxides, the m-oxo bridging involves not the four individual oxidation steps, that is, Vrev(1) =Vrev(2) = only the surface atoms but also the bulk material (see sec- Vrev(3) = Vrev(4). The Vrev(i) in electrocatalysis corresponds to the tion 2). On these grounds, we conclude that bridging oxides midpoint redox potentials for four sequential single-electron are a candidate for a common structural motif in all water oxi- oxidations of a molecular catalyst. The problem to find an elec- dation catalysts. trocatalyst with approximately the same Vrev(i) for each oxida- However, there are a growing number of mononuclear tion step is thus analogous to the problem of redox-potential water-oxidation catalysts that seemingly facilitate water oxida- leveling in homogeneous and biological catalysis, (see section in the absence of any m-oxo bridging (see section 3). The tion 4.4). mere existence of these functional mononuclear catalysts indi- The redox-potential leveling, that is the maintenance of an cates that m-oxo bridging is not an essential prerequisite for approximately constant redox potential for four sequential oxi- water oxidation. However, it should not be overlooked that the dation steps, is a chemically demanding task that, to our turnover frequencies of these complexes are generally low and knowledge, has not yet been achieved for any synthetic mole- that the redox-potential leveling is inadequate. It remains pos- cule. However, two sequential oxidations steps at the same po- sible that efficient water oxidation generally requires m-oxo tential have been realized as discussed for binuclear Mn com- bridging between metal ions. plexes (see section 4.4). In PSII, formation of unprotonated m- oxo bridges appears to be crucial (Scheme 14). Amorphous metal oxide catalysts may carry flexible m-hydroxo bridges with 5.2. Self-assembly and self-repair of catalytic metal centers suitable pK values, related to the PZC, which could facilitate The proteins of PS II represent an elaborate ligand system. The the redox-potential leveling by coupling of m-OH deprotona- assembly of the catalytic complex remains insufficiently under- tion to the oxidation step, possibly in close analogy to process- stood but probably involves oxidation of Mn2+ ions and forma- es in PSII. This hypothesis may explain why amorphous oxides tion of {Mn4Ca(m-O)n} within the framework of the protein li- are OER-active at relatively low electrochemical overpotentials. gands. Conceptually, this oxidative self-assembly does not Experiments addressing changes in m-oxo bridging for surface

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 755 H. Dau, C. Limberg, P Strasser et al. and bulk oxides represent a major challenge but the results oxidation (see section 3) and it has recently also been favored could be of exceptional importance. for electrocatalytic water oxidation (see section 1.3). It should In conclusion, redox-potential leveling in homogeneous and be noted that route A does not involve accumulation of four biological water oxidation, on the one hand, and the require- oxidation equivalents before the onset of OO bond forma- ment for the absence of a potential-determining step in an tion—as opposed to routes B, C, and D—but rather of only ideal electrocatalyst, on the other hand, are different facets of two oxidation equivalents, as suggested for some reaction the same problem. We propose that the quality of redox-po- cycles of mononuclear complexes (see section 3). Therefore, tential leveling generally determines the maximal energetic ef- route A can be ruled out as the main path for water oxidation ficiency of a catalyst. Mode and quality of redox-potential lev- in photosynthesis. eling may also determine the kinetic competence of the cata- To our knowledge, route D has only been considered for lyst, which is quantified in terms of rate constants or turnover water oxidation at the {Mn4Ca} complex of photosynthesis. The frequencies. Notably in this context, in biological water oxida- possibility of concerted inner-complex electron–proton move- tion the OO bond formation step itself is not rate-determin- ments in the OO bond formation step represents a unique ing but rather it is the oxidation of the catalyst in the individu- aspect of route D (see section 4.8). In the field of photosynthet- al Si !Si+1 transitions. For each Si !Si+1 transitions, the driving ic water oxidation, it represents a working hypothesis present- [106a] force of the ET step (DGi of oxidation of the Mn complex) will ed before and detailed herein (section 4). Route D lacks the be maximal for optimal redox-potential leveling. A greater driv- simplicity of routes A and B because not only water-binding ing force typically implies an increase in the respective rate sites but also four acceptor bases are proposed to be part of constant. Similarly in electrocatalysis, the concept of the poten- the catalyst, thereby ruling out route D for mononuclear cata- tial-determining steps involves, at least implicitly, that this step lytic sites. is also the overpotential-determining and thus the overall rate- The investigation of mononuclear complexes revealed the determining step. We conclude that the redox-potential level- feasibility of a hydroxide/water attack at terminal oxo/oxyl li- ing determines generally the maximal energetic efficiency and gands, which has recently also been favored for electrochemi- is also a key determinant of the kinetic competence of each cal water oxidation (see section 1.3) and for the biological par- water oxidation catalyst. A generic mode of optimal redox-po- agon (Scheme 19B–D). In all fields it has remained difficult to tential leveling may involve coupling of metal oxidation to obtain experimental proof, but specifically the electrophilic ligand deprotonation and changes of the oxo bridging be- attack of hydroxide at a terminal oxyl appears to be the mostly tween metal ions (see section 4). This conjecture, however, is favored mode of OO bond formation. still at the level of a working hypothesis. When weighing the evidence in favor and against a distinct mechanism, none of the mechanistic proposals is supported experimentally up to a level that the term ‘proof’ would be ap- 5.4. Mechanism of OO bond formation plicable. Simplicity and seeming plausibility still have to substi- The requirements for OO bond formation have been dis- tute for the lack of experimental evidence, in spite of tremen- cussed at the level of molecular orbitals repeatedly, either in dous efforts and recent progress by experimentalists and theo- form of plausibility considerations or based on quantum chem- retical chemists. The largest body of detailed experimental re- ical calculations. Instead of attempting a discussion of electron- sults has been obtained for photosynthetic water oxidation. ic aspects, we focus on basic considerations and reaction For these reasons, route A, B, and C can essentially be ruled models. out for the biocatalytic reaction. Whether this conclusion is of Routes A, B, and C in Scheme 17 represent the prominent high relevance for electrochemical or homogeneous water oxi- types of mechanisms that have been considered in the field of dation is still unclear. electrochemical water oxidation. In route B, proton-coupled electron transfer results in accumulation of four oxidizing 5.5. Concluding Remarks equivalents by the catalyst and formation of two terminal metal–oxo (Me=O) or metal–oxyl (MeO·) groups at neighbor- We have reviewed and compared recent findings on water oxi- ing sites before OO bond formation takes place. Route B has dation in three research fields: heterogeneous (electro)catalysis indeed been considered in all of the three fields; electrochemi- by surface and bulk oxides (sections 1 and 2), homogeneous cal, homogeneous, and biological catalysis of water oxidation. catalysis by metal complexes (section 3), and biocatalytic water Route B has never been proven experimentally in electrochem- oxidation in photosynthesis (section 4). Each field has devel- ical water oxidation and essentially has been disproven for bio- oped its own terminology, methods, concepts, and, conse- logical water oxidation, whereas recent O16/O18 isotope experi- quently, its own strengths. ments support its prevalence for a binuclear ruthenium complex.[74a] We conclude that route B is appealingly simple and Heterogeneous electrocatalytic water oxidation plausible, but only rarely supported experimentally. There is also an appealingly simple aspect to route A; bind- Stringent approaches towards the energetics and efficiencies ing, deprotonation and oxidation of water take place at a of water splitting have been developed only for electrochemi- single metal site. The existence of mononuclear catalysts cal water-splitting (section 1.1). Traditionally the ‘macroscopic proves that a mono-site mechanism can be operative in water energetics’ were considered with a focus on system parame-

756 www.chemcatchem.org 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761 The Mechanism of Water Oxidation ters of relevance in chemical engineering. At the atomic level, ble to the noncrystalline bulk metal oxides described in sec- however, electrolysis is not well understood. Traditional mecha- tion 2, represents a task for the future. nistic proposals [Equations (12)–(21), routes A–C in Scheme 17] may be plausible but largely are supported only circumstantial- Homogeneous catalysis by transition metal complexes ly by macroscopic kinetic parameters, such as the experimental Tafel slope (see section 1.2). Metal complexes synthesized as homogeneous water-oxidation A major step forward in our understanding of qualitative catalysts are typically characterized crystallographically, so that trends in electrocatalytic activity has resulted from thermo- their structure is known at atomic resolution. Moreover syn- chemical analysis of so-called volcano plots relating thermody- thetic variations of the metal complex facilitate investigation namic overpotential or catalytic activity with microscopic pa- of structure–function relations. These two experimental advan- rameters accessible through modern quantum-chemical calcu- tages at hand, synthetic chemistry has established the founda- lations. Thereby, suboptimal scaling and chemisorption ener- tions for understanding crucial mechanistic aspects in water gies are currently found to limit the achievable catalyst activity oxidation, such as the need for and modes of redox-potential (section 1.3). This work may represent a first step towards an leveling (see section 4.4). atomic-scale rational design of electrocatalysts. However, even Research in recent years has revealed that dimeric or higher- for the basic reaction scheme of the reviewed computational order assemblies of multiple metal centers are not essential for approach [Equation (25)–(28), route A in Scheme 17], the sup- water-oxidation catalysis; mononuclear systems are sufficient port by experimental findings is, at present, insufficient. We and this finding represents a major breakthrough (section 3). also lack strategies and methods to independently manipulate The easier derivatization associated with structurally simpler undesired scaling relations and chemisorption energies. Com- mononuclear systems allows for a more rational synthesis of bination of electrochemical methods with structure-sensitive tailored systems, structure–function analyses, and mechanistic spectroscopies represents a promising avenue towards under- studies. Future studies should evaluate more model systems standing at the atomic level. This endeavor is seriously compli- with respect to the redox-potential leveling and also include cated by the active-site problem of heterogeneous catalysis; complexes with nuclearities higher than two. Although it neither the effective nuclearity (monosite versus multisite seems clear that such systems will never reach the efficiency of mechanisms in Scheme 17) nor the local structure of the cata- nature, they might be designed for satisfactory artificial devi- lyst at the active site are well known, the latter because devia- ces, once ligand decomposition is under control. Immobiliza- tions from crystalline order (defects) probably play a crucial tion of molecular complexes on surfaces may increase drasti- role. cally the catalytic performance and lifetime,[160] possibly be- In the past, electrochemical water oxidation was mostly dis- cause bimolecular decomposition routes are suppressed. Stabi- cussed in physical terms, such as adsorption, discharge, or lizing surfaces or matrices in conjunction with self-assembly Fermi level, which tend to obscure the chemical character of and self-repair of the catalyst will likely be of high importance the process. However, at the atomic level, equivalent reaction in the development of technological catalysts based on mono- schemes and considerations are applicable in electrocatalytic, or multinuclear transition metal complexes. homogeneous, and biological water oxidation (section 5.4). The conceptual gap between electrocatalytic water oxidation, Biocatalytic water oxidation on the one hand, and homogeneous or biocatalytic water oxidation, on the other hand, can often be bridged by a mere The {Mn4Ca} complex of PS II can be driven synchronously translation, for example, substitution of ‘chemisorption’ by the through its reaction cycle (S-state cycle) by four light flashes, term ‘binding’ or ‘ligation.’ In other cases, deeper conceptual thus rendering time-resolved experiments on the four S-state questions may arise. For example, hydroxide binding (or chem- transitions possible (determination of reaction sequences and isorption) to one of the many metal ions of a surface oxide rate constants). Moreover, the four reaction intermediates can may change the oxidation state of this specific metal ion ac- be stabilized by freezing and studied conveniently in sophisti- cording to Equation (41): cated low-temperature experiments, specifically advanced EPR spectroscopy and X-ray methods. These exceptional experimental advantages, in conjunction with high biological impor- MnþþOH ! Mðnþ1Þþ ðOHÞðÞþe ð41Þ tance, have motivated a large number of investigators, specifically spectroscopists, to focus their research on elucidation of This point of view, with roots in the concept of formal oxida- the mechanism of photosynthetic water oxidation. The final tion states (or ‘valencies’) in coordination chemistry, is difficult goal of this truly interdisciplinary endeavor has not yet been to reconcile with the band-gap model often used in solid-state reached. Nonetheless, an exceptionally deep understanding of chemistry and physics. The difference in concepts—distinct lo- mechanistic details and pertinent questions has been obtained. calized oxidation-state changes versus modification of the Many key concepts of likely relevance also for nonbiological band structure close to the surface—is clearly related to the water oxidation have been established, namely accumulation discussion in coordination and bioinorganic chemistry on local- of four oxidizing equivalents, redox-potential leveling, and cou- ized versus delocalized valencies (or oxidation states). The de- pling of electron transfer to deprotonation of the catalyst and velopment of a unifying concept or theory that is also applica- removal of protons from the catalytic site. Specifically the

ChemCatChem 2010, 2, 724 – 761 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 757 H. Dau, C. Limberg, P Strasser et al. latter aspect, that is the transfer of protons from the catalyst [11] G. Centi, S. Perathoner, ChemSusChem 2010, 3, 195. to the bulk medium, has been almost completely neglected in [12] a) L. Jorissen, J. Power Sources 2006, 155, 23; b) K. A. Burke, IEEE Aero- space And Electronic Systems Magazine 1999, 14, 23; c) F. Mitlitsky, B. research on nonbiological water oxidation, but could be crucial Myers, A. H. Weisberg, Energy Fuels 1998, 12, 56; d) S. D. Song, H. M. in the developments of catalysts with satisfactory turnover fre- Zhang, X. P. Ma, Y. N. Zhang, B. L. Yi, Prog. Chem. 2006, 18, 1375; e) A. quencies. In conclusion, research on photosynthetic water oxi- Verma, S. Basu, J. Power Sources 2004, 135, 62. dation will continue to be a source of inspiration, concepts, [13] a) Department of Energy: http://www.sc.doe.gov/bes/reports/list.html, Washington, 2003; b) R. L. Busby, Hydrogen and Fuel Cells, Pennwell and experimental technologies for investigations on nonbio- Corporation, Tulsa, 2005; c) J. Rifkin, The Hydrogen Economy, Tarcher logical catalysts of water oxidation. Penguin, New York, 2002; d) J. J. Romm, The Hype about Hydrogen: We believe that, in the strive for knowledge-based develop- Fact and Fiction in the Race to Save the Climate, Island Press, Washing- ment of efficient and eventually technology-relevant catalysts ton, 2004. [14] S. Dunn, Int. J. Hydrogen Energy 2002, 27, 235. of water oxidation, the fields of heterogeneous, homogeneous, [15] a) Enertrag Corporation, 2010, press release: http://www.enertrag.- and biological catalysis will grow increasingly together. Unify- com/de/hybridkraftwerk.html; b) Hydrogen Wind Hybrid Power Plant: ing concepts will not only play a role on a theoretical, concep- http://en.wikipedia.org/wiki/Wind-hydrogen_hybrid_power_system tual level—as discussed herein—but also in terms of new tech- 2010. [16] J. P. Hoare, The Electrochemistry of Oxygen, Wiley, New York, 1968. nological systems that combine approaches from homogene- [17] a) H. B. Callen, Thermodynamics and an Introduction to Thermostatistics, ous and heterogeneous catalysis with surface and nanoscien- Wiley, New York, 1985;b)Y.A.C¸ engel, M. A. Boles, Thermodynamics: ces, for example in terms of immobilized molecular systems, An Engineering Approach,4th ed., McGraw–Hill, Boston, 2002;c)B.G. nd bulk metal oxides in nanostructured materials, or purely inor- Kyle, Chemical and Process Thermodynamics,2 ed., Prentice Hall, Eng- lewood Cliffs, 1992. ganic catalysts with molecular properties on surfaces.[62b] [18] R. L. LeRoy, C. T. Bowen, D. J. LeRoy, J. Electrochem. Soc. 1980, 127, 1954. [19] R. L. LeRoy, J. Electrochem. Soc. 1983, 130, 2158. Acknowledgements [20] R. O’Hayre, S.-W. Cha, W. Colella, F. B. Prinz, Fuel Cell Fundamentals, Wiley, New York, 2006. [21] K. C. Neyerlin, W. B. Gu, J. Jorne, H. A. Gasteiger, J. Electrochem. Soc. Joint scientific work in the framework provided by the UniCat 2007, 154, B631. cluster of excellence (Unifying Concepts in Catalysis, Berlin) has [22] a) J. O. Bockris, A. K. N. Reddy, Modern Electrochemistry 2B: Electrodics inspired this Review article. Financial support by the UniCat clus- in Chemistry, Engineering, Biology, and Environmental Science, Vol. 2B, ter is gratefully acknowledged. We also thank the European Kluwer Academic, New York, 2000; b) V. S. Bagotsky, Fundamentals of Electrochemistry, 2nd ed., The Electrochemical Society, Penningt, New Union (7th framework program, SOLAR-H2, #516510) and the Jersey, 2006; c) C. H. Hamann, A. Hamnett, W. 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