the mechanism of water oxidation: from electrolysis via homogeneous to biological ... ·...

DOI: 10.1002/cctc.201000126

The Mechanism of Water Oxidation: From Electrolysis viaHomogeneous to Biological CatalysisHolger Dau,*[a] Christian Limberg,*[b] Tobias Reier,[c] Marcel Risch,[a] Stefan Roggan,[b] andPeter Strasser*[c]

724 � 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761

Introduction

In recent years, the increased interest in direct conversion ofsolar energy into storable fuels has led to intensified researchefforts with respect to artificial photosynthesis.[1] A successfulimitation of the natural energy-storing process requires thecombination of several distinct processes, including light har-vesting, charge separation, electron transfer, and water oxida-tion (O2 evolution), as well as the fuel-forming reaction. Where-as a close adaption to the biological process of photosynthesiswould involve the transformation of carbon dioxide into achemical compound of higher energy content, at present thegeneration of molecular hydrogen (H2) by means of proton re-duction seems to be a more realistic proposition for storingsolar energy by artificial photosynthesis.

The electrons and protons required for fuel generationshould be obtained from oxidation of water to oxygen, there-by converting the whole process into an overall water-splittingscheme, where hydrogen is the targeted fuel. Hydrogen itselfcan be utilized as a fuel. Moreover, solar hydrogen will be thekey for the realization of a number of related renewable fuelproduction processes, such as hydrodeoxygenation of lignocel-lulosic biomass into fuels and chemicals, or the reduction ofCO2 into methanol or other carbonaceous fuels. However, notonly the generation of molecular hydrogen from water re-quires water oxidation. For all substances that today are typi-cally considered as fuels, it holds that utilization of the chemi-cally stored energy involves the reduction of molecularoxygen, that is, the transfer of electrons from the fuel tooxygen. For carbonaceous fuels, this process is coupled to CO2

formation and, in most cases, also to H2O formation. In fuelgeneration, a source for these electrons is needed, and wateris the one and only truly attractive candidate; every schemefor sustainable, large-scale production of non-fossil fuels willeventually involve utilization of water as a raw material.

Water oxidation is currently considered a major bottleneck,hampering progress in the development of applicable devicesfor the conversion of light into storable fuels. Intense researchefforts are being pursued to develop both heterogeneous andmolecular catalysts that enable water oxidation at potentialsclose to the thermodynamic limit. Research into the biologicalprocess of photosynthetic water oxidation has also recently in-

tensified, as it can serve as an important inspiration andbenchmark for the development of artificial water-oxidationcatalysts.

Herein we will discuss and compare concepts and recentfinding on the mechanism of water oxidation obtained in thefollowing separate research communities:

* Heterogeneous electrocatalysis at electrode surfaces (sec-tion 1),

* heterogeneous catalysis by oxidic bulk materials in colloi-dal form or deposited on electrodes (section 2),

* homogeneous catalysis by transition metal complexes(section 3),

* biocatalytic water oxidation in photosynthesis (section 4).Rather than approaching the impossible, that is, a fully com-

prehensive review, we focus on recent progress and strive toaddress common themes and potentially unifying concepts.

1. Electrochemical Water Splitting

Electrolysis, the electrocatalytic splitting of liquid water into hy-drogen and oxygen using electricity, is the longest-studied cat-alytic way to oxidize water, predating even the developmentof the concept of chemical catalysis. The overall reaction, cata-lyzed at the electrified solid–liquid–gas three-phase interfaceof the anode and the cathode, is given by

H2OðlÞ ! H2ðgÞþ1=2 O2ðgÞ ð1Þ

Striving for new solar fuels, the water oxidation reaction cur-rently is considered to be a bottleneck, hampering progress inthe development of applicable technologies for the conversionof light into storable fuels. This review compares and unifiesviewpoints on water oxidation from various fields of catalysisresearch. The first part deals with the thermodynamic efficien-cy and mechanisms of electrochemical water splitting by metaloxides on electrode surfaces, explaining the recent concept ofthe potential-determining step. Subsequently, novel cobaltoxide-based catalysts for heterogeneous (electro)catalysis arediscussed. These may share structural and functional propertieswith surface oxides, multinuclear molecular catalysts and the

catalytic manganese–calcium complex of photosynthetic wateroxidation. Recent developments in homogeneous water-oxida-tion catalysis are outlined with a focus on the discovery ofmononuclear ruthenium (and non-ruthenium) complexes thatefficiently mediate O2 evolution from water. Water oxidation inphotosynthesis is the subject of a concise presentation ofstructure and function of the natural paragon—the manga-nese–calcium complex in photosystem II—for which ideas con-cerning redox-potential leveling, proton removal, and O�Obond formation mechanisms are discussed. The last part high-lights common themes and unifying concepts.

[a] Prof. Dr. H. Dau, M. RischFreie Universit�t Berlin, Physics Dept. , Molecular BiophysicsArnimallee 14, 14195 Berlin (Germany)Fax: (+ 49) 30-838-56299E-mail : [email protected]

[b] Prof. Dr. C. Limberg, Dr. S. RogganHumboldt Universit�t Berlin, Department of ChemistryBrook-Taylor-Strasse 2, 12489 Berlin (Germany)Fax: (+ 49) 30-2093-6966E-mail : [email protected]

[c] T. Reier, Prof. Dr. P. StrasserTechnische Universit�t Berlin, Chemistry Dept.The Electrochemical Energy, Catalysis and Materials Science LaboratoryStrasse des 17.Juni 124, 10623 Berlin (Germany)Fax: (+ 49) 30-314-22261E-mail : [email protected]

ChemCatChem 2010, 2, 724 – 761 � 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim www.chemcatchem.org 725

www.chemcatchem.org

Holger Dau received his physics diplo-

ma in 1985 and doctoral degree in

1989 in Kiel, Germany, working with

U.-P. Hansen and at the Weizmann In-

stitut (Rehovot, Israel, winter 1987/88).

After postdoctoral work with K. Sauer

at the UC Berkeley, USA, and H. Senger

in Marburg, Germany, he received his

habilitation degree at the Biology De-

partment of Philipps University Mar-

burg in 1994. Besides photosynthesis

research in Marburg, he developed bi-

otest applications at bbe Moldaenke GmbH (1997–1999). Since

2000, he has been a full Professor at the Physics Department of

the Free University in Berlin, Germany, where he investigates bio-

logical and synthetic metal sites with X-ray spectroscopy and com-

plementary methods. His current focus lies on light-driven water

oxidation and H2 formation in biological and biomimetic systems.

Christian Limberg was born in 1965 in

Essen, Germany, and studied chemistry

from 1985 until 1990 in Bochum, Ger-

many. He earned his doctorate with A.

Haas in 1992 and concluded postdoc-

toral work together with A. J. Downs

at Oxford University, UK, with a further

PhD thesis (D. Phil.). Having performed

his habilitation in Heidelberg (1995–

1999) he moved to the TU Munich in

Germany to lead the inorganic chemis-

try chair on behalf of W. A. Herrmann.

In 2002, he accepted an offer to become full professor at the Hum-

boldt-Universit�t zu Berlin, Germany. His research is mainly con-

cerned with oxo metal complexes, oxidation reactions and their

mechanisms (surfaces of metal–oxo catalysts as inspiration for mo-

lecular models; biomimetic O2 activation and oxidation) as well as

with the activation of small substrates utilizing dinuclear com-

plexes.

Tobias Reier received his Master of

Science (Dipl.-Ing.) in 2010 at the Tech-

nical University Berlin, Germany, in

Chemical Engineering. He is currently

working on his doctoral thesis at the

Technical University of Berlin in the

group of Prof. Strasser. His research is

focused on a fundamental understand-

ing of mechanistic aspects and struc-

tural effects of electrochemical water

splitting as well as the design of im-

proved water-splitting electrocatalysts.

Marcel Risch received his Bachelor of

Science (physics) at the Technical Uni-

versity of Darmstadt, Germany, in 2006

and a Master of Science (physics) at

the University of Saskatchewan,

Canada, in 2008. In 2009, he joined the

Berlin International Graduate School of

Natural Sciences and Engineering (BIG-

NSE) and he is currently studying for a

PhD in the group of Prof. Dau at the

Free University of Berlin. The topic of

his thesis is the characterization of an

electrochemical cobalt catalyst.

Stefan Roggan graduated in chemistry

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-

boldt-Universit�t zu Berlin, where he

performed his doctoral studies. In

2007, he received his PhD in chemistry

and then joined Prof. T. Don Tilley’s

group at the UC Berkeley, USA, as a

postdoctoral fellow. After his return

from the US in 2009, he continued re-

search in Christian Limberg’s group before starting to work as a

scientist for Bayer Technology Services GmbH in Leverkusen, Ger-

many.

Peter Strasser is a full Professor at the

Chemical Engineering Division of the

Department of Chemistry at the Tech-

nical University of Berlin. His research

interests focus on materials and cataly-

sis for electrochemical (energy) tech-

nologies. Prior to his appointment, he

was an Assistant Professor at the De-

partment of Chemical and Biomolecu-

lar Engineering at the University of

Houston, USA. From 2001 to 2004, he

served as senior member of staff at

Symyx Technologies, Inc. , Santa Clara, CA, developing and utilizing

combinatorial and high throughput screening methodologies in

heterogeneous and electrocatalysis. In 1999, Peter Strasser ob-

tained his doctoral degree in Physical Chemistry and Electrochem-

istry from the Fritz-Haber-Institute of the Max-Planck-Society,

Berlin, under the direction of the 2007 Chemistry Noble Laureate,

Professor Gerhard Ertl. He studied chemistry at Stanford University,

USA, the University of T�bingen, Germany, and the University of

Pisa, Italy, and received his diploma degree (MS) in chemistry in

1995.

726 www.chemcatchem.org � 2010 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim ChemCatChem 2010, 2, 724 – 761

H. Dau, C. Limberg, P Strasser et al.

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In the future, water electrolysis could play a key role in sus-tainable energy conversion and storage infrastructure.[2] Firstobserved in 1789 in a transient discharge experiment by vanTroostwijk,[3] then possibly observed by Volta using one of hispiles[4] and finally made known to a wider audience in 1800 byNicolson and Carlisle,[5] it largely remained a scientific curiositythroughout the 19th century. During the early decades of the20th century, water electrolysis was industrialized and served toproduce very pure hydrogen for industrial uses, such as ammo-nia production or early petroleum refining.[6] Later in the20th century, materials and system improvements resulted insignificant increases in net efficiencies of electrolyzers above50 %. Alkaline water electrolysis, where hydroxide anions arethe major charge carrier in the liquid ion-conducting electro-lyte, became a technology of choice for production of oxygenand hydrogen for life-support systems in anaerobic environ-ments, such as submarines. The use of separator materials (dia-phragms) aided in the design of more-compact electrolyzercells with improved separation of hydrogen and oxygen gas.[7]

Later on, high-pressure alkaline electrolyzers were designedand commercialized. Work by Beer resulted in the technologi-cally important dimensionally stable anode (DSA),[8] which re-mains the basis of electrolyzer anodes to date. In the DSA, cat-alysts are supported directly on metallic titanium as a thin film.Space exploration drove the development of polymer electro-lyte membrane (PEM) electrolyzers for acidic media thatshowed very similar system architectures as hydrogen PEMfuel cells (PEMFCs).[9] Membrane materials such as perfluorosul-fonic acid (Nafion�) allow the conduction of protons across athin polymer electrolyte layer. Advantages of proton-conduct-ing PEM devices over liquid alkaline systems include highersafety and reliability, because no caustic soda is circulated inthe cell, greater energy efficiency, sustainability to high differ-ential gas pressures, higher hydrogen production rates, and amore-compact cell design.[9a] Drawbacks relate to the use ofmore expensive noble metal anode catalyst materials, such asRu or Ir. Current developments in the field of electrolysis in-clude the coupling of water electrolysis with photovoltaic solarenergy conversion[10] into an integrated direct photoelectroca-talytic (PEC) water-splitting technology.[11] PEC devices promisemore-compact electrode and possibly device architectures, yetrequire multifunctional hybrid electrode materials. Low solarconversion efficiencies combined with electrocatalytic overpo-tential losses largely limit the efficiencies of today’s PEC devi-ces to below 5 %. Another current development centers on thesynthesis of practical alkaline polymer electrolyte membranes.This technology would combine the advantages of the PEMtechnology with the benefits of a mature alkaline electrolyzerindustry, in particular with regards to the use of less expensivenon-noble metal electrode materials. Finally, the integration ofelectrolyzer and fuel-cell functionalities into one single unitizedregenerative fuel-cell device is the subject of current researchand development efforts.[12] A regenerative fuel cell/electrolyzersystem offers the prospect of a two-way conversion of chemi-cal and electrical energy using bifunctional electrodes with theadvantage of further increased system simplicity and compact-ness.

Today, water electrolysis is frequently cited as a clean corner-stone production technology in a largely hydrogen-based sus-tainable energy infrastructure (hydrogen economy[13]). Electrol-ysis could serve to liberate hydrogen from a range of chemi-cals, not only from water. Also, being a scalable technology,electrolysis would allow the initial deployment of smaller-scaleelectrolysis units at comparable efficiencies to large-scale elec-trolyzers.[13b–d] The vision of a broader hydrogen-based energyeconomy, however, still faces a large number of serious techni-cal and nontechnical road blocks.[13d] Among the scientific andtechnical challenges, the need for advanced materials is criticaland will likely require a large amount of additional basicenergy materials and catalysis research.

Hydrogen is an important industrial feedstock for petroleumrefining and fertilizer industry. However, electrolysis accountsonly for about 4 % of global hydrogen production[9b, 13b–d, 14] .The recent surge of interest and investments in renewableelectricity through solar–thermal, solar–electric (photovoltaic),or wind farms, combined with the emerging interest in electro-mobility concepts for urban centers, has raised a substantialamount of awareness about pressurized electrolytic hydrogenfor storage of excess renewable electricity. As a results, hydro-gen–wind hybrid power plants are currently being commercial-ized at a number of locations worldwide.[15] On a smaller scale,more and more consumer wind and photovoltaic units havebeen deployed in accordance with energy policies that offergenerous subsidies for renewable electricity fed back into thegrid. Once these subsidies are cut back and grid electricityprices continue to rise, a growing interest in home energy stor-age units, possibly based on water electrolysis and compressedhydrogen, will result.

The overall electrochemical water splitting is divided intotwo half-cell redox reactions. The reduction process at thecathode (hydrogen evolution reaction, HER) proceeds accord-ing to

2 Hþþ2 e� ! H2 ð2Þ

while the oxidation process (oxygen evolution reaction, OER)at the anode of the electrolyzer is[16]

H2O! 1=2 O2þ2 Hþþ2 e� ð3Þ

In the total process, two electrons are transferred per formulaor a total charge of 192 928 C (= 2 F) crosses the electrocatalyt-ic interface per mole of water converted.

We will focus herein on the mechanistic aspects of watersplitting (electrolysis) itself. Photoelectrochemical water split-ting, whereby a photon-absorbing component is directly cou-pled with a gas-evolution electrocatalyst, is beyond the scopeof this review. A recent article has highlighted photoelectro-chemical water splitting in the context of photoelectrochemi-cal solar fuels.[11]


The Mechanism of Water Oxidation

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1.1. The thermodynamic efficiency of electrochemical watersplitting

The equilibrium thermodynamics of galvanic and electrolysiscells extends the set of thermodynamic variables of a chemicalprocess that does not involve free charges, an electric poten-tial E, or the difference of two potentials, the voltage V.[17] Cellvoltages at constant pressure P and temperature T are relatedto thermal energy quantities by the amount of electricity trans-ferred nF, where n denotes the number of electrons transferredand F is the Faraday constant.

Water splitting (ws) according to Equation (1) at standardtemperature and pressure (STP) is associated with a molar en-thalpy change of reaction, herein denoted as DHo

ws,298, which isequal to �DHo

f,H2 OðlÞ ,298, the heat of formation of one mole ofliquid water at temperature 298 K. The absolute value of thesequantities is 286 kJ mol�1 (water) or 2.96 eV. The quantityDHo

ws,298 is also referred to as the higher heating value DHoHHV,298

of one mole hydrogen. The change in Gibbs free energy of re-action (1) DGo

ws,298 is +237 kJ mol�1 (water) or 2.46 eV. Note: Dif-ferent energy units are used in the reviewed fields and we donot use consistently the same units but rather those most con-venient in the specific context. For a rough transformation ofenergy units : 1 eV�100 kJ mol�1�25 kcal mol�1.

At equilibrium, the minimum electrical energy, 2 F Vrev, re-quired to split one mole of water under STP is related to theGibbs energy as

DGows,298 ¼ 2 F Vo

rev,298 ð4Þ

Considering only absolute values of cell voltages, Vorev,298 =

1.23 V at STP and is referred to as the reversible electrolysis cellvoltage at STP.

The total energy required to split one mole of water at STPis given by

DHows,298 ¼ DHo

HHV,298 ¼ �DHof,H2OðlÞ ,298 ¼ DGo

ws,298þ T DSows,298 ð5Þ

where the last term at STP is 49 kJ mol�1 and corresponds tothe heat required from the surroundings. Thus, a higher-heat-ing-value voltage, VHHV,

[18] or an equivalent amount of electrici-ty, 2 F VHHV, can be related to the water splitting energy accord-ing to

DHoHHV,298 ¼ �2 F V o

HHV,298 ð6Þ

where VoHHV,298 = 1.48 V. The voltage difference between Vo

rev,298

and VoHHV,298 implies that water electrolysis below 1.48 V and

above 1.23 V is an endothermic process at STP. Hence, neglect-ing ohmic and other loss processes, an electrolysis cell operat-ed below 1.48 V would cool down during electrolysis or wouldrequire a heat supply to operate isothermally.

At elevated temperatures, Vorev,T drops based on Equations (4)

and (5), whereas DHoHHV,T and Vo

rev,T increase according to[18–19]

DHoHHV,T ¼>¼ �DHo

f,H2 OðlÞ ,Tþ ðHo

H2OðlÞ ,T�o

f,H2 OðlÞ ,298Þ ¼ 2 F VoHHV,T ð7Þ

where HoH2OðlÞ ,T

is the molar enthalpy of water at standard pres-sure and temperature T. As a result, the voltage range of endo-thermic water splitting broadens.[18]

The formalism above assumed the electrolytic production ofperfectly dry hydrogen and oxygen gas. Taking into accountthe presence of a water-vapor partial pressure pw at a totalpressure p of hydrogen and oxygen, a total energy for splittingof one mole of water into humidified gases can be derived,which is associated with a thermo-neutral voltage VTN accord-ing to[18–19]

DHTN;T;p ¼DHHHV;T ;ðp�pwÞ þ 1:5pw

p� pwð Þ HH2OðlÞ ;T ;pw� Ho

H2OðlÞ ;298

� �

¼ 2 F VTN;T;p

ð8Þ

The first term of the central expression is the higher heatingvalue of the electrolysis process at temperature T, and gaspressure (p�pw), while the second term accounts for the stoi-chiometric evaporation of water from STP to operating condi-tions.

Owing to a variety of energy and efficiency losses, cell vol-tages in practical electrolysis cells[9b] are generally above boththe thermoneutral voltage VTN

[18] and the higher heating valuevoltage VHHV. That is, excess thermal energy generally leads toheating of the electrolysis cell.

The efficiency of an isothermal electrolysis cell operated at acell voltage Vo

op,T relates its chemical energy output to its elec-trical energy input, 2 F Vo

op,T. The energy output may be consid-ered in two separate ways. Firstly, the maximum reversiblework (Gibbs free energy) DG from reconverting hydrogen andoxygen, for instance in a fuel cell, may be of interest andhence a faradaic (voltage) efficiency[9b] of electrolysis, ho

298;Faradaic,can be written as

ho298;Faradaic ¼

DGows;298

2 F Vop¼

Vorev;298

Voop;298

ð9Þ

This quantity essentially represents the reciprocal value ofthe voltage efficiency of a fuel cell[20] with Vop replaced by theexperimentally observed output cell voltage of a fuel cell. Thefaradaic efficiency can never be larger than one. Secondly, hy-drogen and oxygen may be viewed in terms of their heatingvalue and hence a thermal efficiency of an electrolysis cell,ho

298;thermal, is written as[18]

ho298;thermal ¼

DHoHHV;298

2 F Vop¼

V oHHV;298

V oop;298

ð10Þ

This efficiency has become widely accepted even though itwould never be possible to recover the entire higher heatingvalue of hydrogen in an energy application. In principle,ho

298;thermal can be larger than unity.[6, 18–19] It is the reciprocalvalue of the reversible thermodynamic efficiency of a fuelcell.[20] Typical values of ho

T ;thermalof industrial electrolyzers rangebetween 60 % to 80 % at a given current density (typically1 A cm�2) and at 363 K (90 8C) and 0.1 MPa.



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Equations (9) and (10) provide a measure of the efficiency ofoverall water splitting. They are generally valid, regardless ofwhether it is an electrocatalytic process at an inorganic elec-trode, a biocatalytic process, or a catalytic pathway involvinghomogeneous organometallic catalysts. In a ‘translation’ of theelectrochemical terminology, anode and cathode potentialscorrespond to the midpoint redox potentials of the oxidant(anode potential) and reductant (cathode potential) that areused to drive the OER or HER. Similarly, in electrochemical, ho-mogeneous, and biological catalysis, research typically consid-ers only half-cell reactions, as is also the case in the mechanis-tic discussions of water oxidation (section 1.2). For half-cell re-actions, however, Equations (9) and (10) are not directly appli-cable.

Equation (9) may be reformulated as

ho298;Faradaic ¼

Vorev;298

Vorev;298 þ OPanode þ OPcathode

ð11Þ

where OPanode and OPcathode are the absolute values of the over-potentials of the OER (anodic reaction) and HER (cathodic reac-tion), respectively. These overpotentials can be calculated inhomogeneous and biocatalysis as the difference between themidpoint potential of the oxidant (or reductant) and the OER(or HER) equilibrium potential. To quantify the efficiency of thehalf-cell reaction, we now assume for OER (or HER) thatOPcathode (or OPanode) equals zero. The thereby-obtained value ofho

OER (hoHER) represents a meaningful measure of the energetic

efficiency of the catalysis of water oxidation (hydrogen evolu-tion) in electrochemical, homogeneous, and biocatalysis. For acombination of two half-cell reactions, calculation of the over-all efficiency ho from the two half-cell efficiencies is straightfor-ward. However, ho is not equal to the product of ho

OER and hoHER

but to [(hoOER)�1+(ho

HER)�1�1]�1.

1.2. Reaction mechanisms of electrochemical water splitting

As described above, the electrochemical splitting of water in-volves two concurrent catalytic half-cell reactions accordingEquation (2) (HER) and Equation (3) (OER). Also shown aremechanistic schemes for either reaction process, which havebeen proposed in literature in the past. Clearly, oxygen evolu-tion being a 4 electron process is more complex as comparedto the evolution of hydrogen and involves several surface ad-sorbed intermediates. As either process may contribute to volt-age and efficiency losses, a discussion of the electrocatalysis ofwater splitting should in principle touch upon both the HERand OER. Work on biocatalytic water splitting has focusedeither on photosystem II (oxygen evolution), as discussed insection 4, or hydrogenase enzymes (hydrogen evolution),which are not discussed herein. Similarly, research concerninghomogeneous catalysis of water splitting has involved systemsliberating either hydrogen or oxygen (or, in rare cases, both;see section 3.2 for comparison). In electrochemical environ-ments, HER on selected noble metals shows extremely high ex-change current densities, j0, which is a measure of the intrinsic

turnover frequency of an electrochemical half-cell process. OnPt electrodes (Figure 1), for instance, HER shows values of j0 ofup to several hundred A cm�2,[21] resulting in a very reversiblehalf-cell electrode with high forward and backward reaction

rates. However, OER on Pt, the catalytically most active ele-ment for this reaction, exhibits values of j0 on the order of10�9 A cm�2, which is why much recent work has shiftedtoward OER kinetics. Therefore, the OER is our focus here andalso in upcoming sections.

In electrocatalysis, a fundamental concept in characterizingcatalytic processes on surfaces is the overpotential.[22] Overpo-tential is the deviation of the observed output cell voltage Vop

(galvanic cell) or the applied cell voltage Vop (electrolysis) usedin Equations (9) and (10) from the thermodynamic reversiblecell voltage Vrev [Equation (4)] . The overpotential is generallyconsidered a kinetic phenomenon and splits into various con-tributions such as the ohmic voltage loss (ohmic overpotential)or mass-transport limitations (transport overpotential) duringthe flow of charge.[22] In the context of catalysis, “reaction over-potentials” relate to a kinetic hindrance of individual elementa-ry reaction steps due to activation barriers. On that basis, reac-tion overpotential is a function of the flowing current and isthought to be linked to the slowest of the elementary reactionsteps, the rate-determining step (RDS), with the highest kineticactivation barrier.

Despite much work on the OER, mechanistic models and hy-potheses of the sequence of elementary steps based onatomic-scale experiments have remained scarce. Over thecourse of five decades, proposed mechanisms, conclusions asto the RDS, and experimental approaches have changed onlylittle. A severe hurdle to more direct atomic-scale insight intoOER is the fact that reactive intermediates have largely eludeddirect spectroscopic observation. Most early mechanistic workrelies on classical current–potential–time analyses and the de-rived Tafel slopes. Tafel slopes are derived from semi-logarith-mic potential current plots and have been frequently used toobtain information on the mechanism and the rate-determin-

Figure 1. Catalytic reactions associated with electrochemical water electroly-sis. An external voltage, symbolized by the battery symbol, drives the evolu-tion of hydrogen at the cathode (HER) and the evolution of oxygen at theanode (OER). Protons are migrating from anode to cathode. Both mecha-nisms are depicted assuming a sequence of proton-coupled single-electronredox reactions. While HER shows one adsorbed intermediate, OER hasthree. The spheres symbolize a metal electrode, for example, Pt, coveredwith an oxide layer. The numbers in parentheses refer to chemical equationsgiven in the text.



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ing step of the overall half-cell reaction. An early mechanisticscheme of electrochemical water splitting based on experi-ments on RuO2-based electrodes in acidic environments is out-lined by Equations (12)—(15):[23]

SþH2O! S�OHþHþþe� ð12Þ

S�OH! S�OþHþþe� ð13Þ

S�OHþS�OH! S�OþSþH2O ð14Þ

S�OþS�O! O2þ2 S ð15Þ

In this mechanistic hypothesis, S denotes a catalyticallyactive surface site. Equations (12), (13), and (15), on the onehand, and (12), (14), and (15), on the other, represent parallelreaction pathways. This mechanism and very similar schemeswere proposed for both well-defined Ru(110) single-crystaloxide surfaces[23d, 24] and compact RuO2 films.[23d] RuO2 singlecrystals were reported with a near 60 mV decade�1 and a120 mV decade�1 Tafel slope at low and high overpotentials, re-spectively.[23d, 24] On that basis, reaction (12) was thought to bethe RDS of the sequence (12)–(13)–(15) at high electrode po-tentials. An additional second chemical step, such as the hypo-thetical rearrangement of the S�OH species according toEquation (16):[24]

S�OH! S�OH* ð16Þ

was proposed to account for the smaller experimental Tafelslopes at low overpotentials, with the reaction sequence nowbeing (12)–(16)–(14).

Ti-supported polycrystalline RuO2 and Ru–MO2 mixed oxidefilms were studied extensively, because they compositionallyresemble the industrial DSA used in virtually all commercialelectrolyzers.[25] A very similar reaction mechanism to Equa-tions (12)–(15) was suggested for IrO2.

[26] A large number ofstudies considered the influence of the preparation conditionsapplied for Ru oxide and Ru–Ir bimetallic oxide films on theOER activity and mechanism.[26] Synthetic parameters studiedincluded the metal precursor and synthesis route (e.g. sol–gelversus precipitation), annealing temperature of the catalystfilm, the resulting BET surface, the integral voltammetriccharge q*, and the point of zero charge.[23d, 27] Due to the lackof a single parameter controlling trends in OER activity ormechanism, many studies resorted to correlations of synthesisparameters,[27a, 28] voltammetric charge q*,[27a] OER activity (Tafelslopes), macroscopic structure, and morphology, often intro-ducing somewhat arbitrary descriptors for catalyst morphologysuch as cracked, smooth, compact, or rough.[27a] Sol–gel-typeRu oxide films followed the Tafel slope pattern of single crys-tals,[28] whereas other synthetic methods resulted in deviatingslopes.

High annealing temperatures resulted in polycrystalline Ruoxide films with low BET surface areas and small voltammetriccharges q*, which was interpreted in terms of a low density ofactive sites.[23d, 24, 27a, d] The OER activity of these films is generallylow and their Tafel slope showed values around 60 mV deca-

de�1.[23d, 27c, d] Similar to the single crystal work, reaction (16) wasproposed for these films as a second elementary step. Low-temperature Ru oxide films showed high BET surface areavalues combined with higher values of q*. The OER activitywas high for these films and the Tafel slope ranged in the40 mV decade�1 range. Hence, an electrochemical reaction (13)was now proposed as the rate-determining step and the fullmechanistic sequence thus becomes (12)–(13)–(15).

Binary and ternary mixed oxides containing Ir, Sn, and Rushowed their own distinct complex Tafel slope patterns withno clear trend or controlling parameter.[25, 27c, d, g, 29] A number ofrecent studies focused on the effect of particle size on the OERmechanism for RuO2 and Ru1�xCoxO2 mixed oxides.[30] The au-thors used differentially pumped mass spectrometry (DEMS) todirectly correlate the evolution of molecular oxygen and its far-adaic current profile. They found that the OER reaction mecha-nism was independent of catalyst particle size but, however,was affected by Co doping.[30b] The second metal lowered theTafel slope from values near 120 mV decade�1 to 40–50 mV de-cade�1, consistent with a chemical step becoming rate-deter-mining.

In alkaline solution, non-noble metal oxides of Ni, Mn, Co, orPb with rutile-type,[27a] perovskite-type,[27a, 31] and spinel-type[27a, 32] structures were investigated in terms of their reac-tion mechanism for OER. An initial discharge of hydroxide ionswas proposed as the first step in the reaction sequence:[33]

SþOH� ! S�OHþe� ð17Þ

To facilitate comparison of mechanisms in heterogeneouselectrocatalysis, homogeneous, and biological catalysis, a tran-sition from the physical terminology involving ‘adsorption’ and‘hydroxide discharge’ to a description in terms of chemical pro-cesses is required. The hydroxide ‘discharge’, by binding to theelectrode surface, could relate either to formation of a boundhydroxide with radical character (OH·) or, more likely, to liga-tion of the hydroxide ion to a metal of the surface oxide suchthat the oxidation state of this specific metal ion formally in-creases by one unit.

The step described by Equation (17) was thought to be fol-lowed by either a deprotonation and subsequent dis-charge[33–34]

S�OHþOH� ! S�O�þH2O ð18Þ

S�O� ! S�Oþe� ð19Þ

2 S�O! 2 SþO2 ð20Þ

or by a slow breaking of the M�OH bond to form a peroxidespecies,[31a, c]

S�OHþOH� ! SþH2O2þe� ð21Þ

that subsequently decomposes into O2. Similar to studies inacid environments, synthesis conditions and Tafel slopes variedwidely. Bockris and Otagawa presented a detailed study[31b]

where they correlated the OER activity with a large number of



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microscopic and material specific parameters. Their findingssuggested that the catalytic OER activity correlated inverselywith the surface bond energy of OH, from which he concludedthat the desorption of OH or oxygenated species may be therate determining step in the OER under his conditions. Tafelslopes on doped perovskites varied from 43 to 200 mV deca-de�1 [31b] and hence provided limited insight into the reactionmechanism. In yet further mechanistic studies in alkalinemedia,[35] the formation of higher-valent metal oxides wasstressed as also concluded in studies on water oxidation inacidic media.[27f] An early molecular orbital theory study of theOER mechanism showed that chemical recombination of ad-sorbed surface oxygen[33] has a low barrier [Equation (15)] andthe study also proposed the formation of an adsorbed perox-ide species (S�OOH).

The specific adsorption of ions from the electrolyte at theinner portion of the electrochemical double layer (inner Helm-holtz layer) can have a tremendous impact on the availablenumber of active electrocatalytic sites. Specific adsorption canblock active sites and hence seriously affect the catalytic activi-ty. Knowledge and control of such ion effects is therefore keyto fundamental understanding and design of electrocatalyticinterfaces. Anion effects in the oxygen reduction reaction(ORR) have been studied in much detail over recent decades.[36]

However, little thorough work has been done in the area of co-valent and noncovalent interactions of OER catalysts with spec-tator species. .[37] Lodi et al.[23d] reported no influence of thenature of the electrolyte anion (sulfate or perchlorate) on theOER catalysis on RuO2, in stark contrast to the effect of anionadsorption in the oxygen reduction reaction.[36b, g] Lodi’s resultsuggests that active OER sites were not blocked by sulfateanion adsorption. As a consequence, catalytic and kinetic im-provements by selective blocking of adsorbing anions may notbe an option in the search for improved OER catalysts.[36h]

Babak and co-workers[37b] reported a weak effect of the sup-porting electrolyte anion (sulfate, nitrate, or perchlorate) onthe OER activity in neutral conditions on PbO2 electrodes onlyat high overpotentials. Amadelli et al.[37a] reported a blockinginfluence of fluoride anions for the OER on the PbO2 electro-des. Clearly, more work has to be done on understanding spec-tator effects in OER catalysis, as it may represent an alternativeroute to improved surface electrocatalysis.[36h]

Trasatti[27a] highlighted the relation between preparationconditions and the point of zero charge (PZC) of an oxide sur-face. The PZC is the pH at which the surface carries no netelectrical charges. The PZC controls the acid–base propertiesof surface coordination complexes such as S�OH. A low PZC,for instance, can imply that the S�O bond is strong and thesurface hydroxide groups can behave like a weak Brønstedacid:

S�OH! S�O�þHþ ð22Þ

Thus, indirectly, the PZC may provide information on thebond strength or chemisorption energy of S and O. The an-nealing temperature of pure RuO2 films was found to correlatewell with the measured PZC. Higher PZC resulted in lower OER

activity (higher overpotential at given current density). The IrO2

content in Ru–Ir mixed oxides, in contrast, correlated inverselywith the PZC. The PZC of solid surfaces, even though theresult of a complex convolution of factors, can be viewed inthe context of the Brønsted acid behavior of m-oxo bridges, asdiscussed for the {Mn4Ca} center in photosynthetic water oxi-dation by photosystem II, where proton dissociation is crucialin redox-potential leveling (see Section 4).

A seminal concept introduced by Ruetschi and Delahay,[38]

later extended by Parsons,[39] Bockris[31] and Trasatti,[27a] involvesthe hypothesis that a single microscopic parameter, such asthe oxygen chemisorption energy or the molar enthalpy of theoxidation of metal oxides, may be a key controlling factor inthe OER and possibly also in the chlorine evolution reaction(ClER). Along these lines, OER can be viewed as the formationand subsequent decomposition of high-valent metal surfaceoxides. To investigate this hypothesis, Bockris[31b] and Trasat-ti[27a] both reported volcano plots of OER activity (overpoten-tial) versus the metal�OH bond energy and versus the enthal-py of formation of higher oxides, building on Sabatier’s princi-ple of balanced intermediate adsorption in catalysis[40] (see sec-tion 1.3). In Trasatti’s study in acidic environments, RuO2 waslocated near the top of the activity volcano curve owing to itsbalanced energetics in the formation and decomposition ofhigher-valence oxides. This could potentially be the reasonthat a majority of mononuclear complexes and coordinationcompounds with OER activity contain Ru ions.

In summary, considering the large number of partially con-flicting reports and the variety of synthetic routes and condi-tions Tafel slope-based mechanistic work should be viewedwith caution and care where reliable conclusions as to the ratedetermining step of the OER are concerned. Correlation of ac-tivity with microscopic parameters related to the surface chem-ical bond is important for an atomic level insight into the elec-trocatalytic processes. Only experimental in situ spectroscopyof surface intermediates combined with atomic-level analysisand a reliable theoretical computational framework will ulti-mately be able to resolve kinetic activation barriers, elementarysteps, and possibly the nature of the active site of the OER onoxidic surfaces.

There are a number of developments towards in situ atomic-scale understanding of electrochemical water splitting. In situelectrochemical Raman spectroscopy[41] or mass spectrometry,for instance, have become valuable tools for insight into elec-trocatalytic mechanisms. The incorporation of atomic oxygenoriginating from reactant water into the surface and subsur-face of the RuO2 and IrO2 lattice were confirmed by usingin situ differentially pumped mass spectrometry studies (in situDEMS).[26, 42] The formation of higher-valent Ru compoundssuch as RuO4 was also reported, which corroborates the degra-dation mechanisms of Ru oxides during water splitting due tovolatile higher-valent oxides.

Novel in situ tools for the investigation of solid–liquid inter-faces include modern synchrotron-based X-ray methodologies,such as ambient-pressure X-ray photoemission spectroscopy(XPS), in situ X-ray absorption spectroscopy (XAS), or X-rayemission spectroscopy (XES), which allows the simultaneous



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detection of a variety of different surface-adsorbed oxygenspecies.[43] In the future, these methods are likely to advanceour understanding of the interaction of water with solid surfa-ces and mechanistic aspects of water splitting.

1.3. The thermochemical overpotential of water splitting

The activation energy of the slowest elementary electrochemi-cal step is generally thought to control the relation betweenelectrode overpotential and the catalytic rate of a reaction pro-cess. This purely kinetic notion tacitly assumes that all elemen-tary steps along the reaction pathway are thermodynamicallydownhill. Studies into the relationship between activity andthermochemical quantities date back to reports by the groupsof Parsons,[39] Bockris,[31b] and Trasatti.[27a]

Building on those early concepts, recent theoretical compu-tational studies[44] of electrocatalytic reactions highlighted thethermochemical aspects of the overpotential of electrochemi-cal reactions using density functional theory (DFT). In their the-oretical framework, Rossmeisl and co-workers calculated Gibbsfree adsorption (chemisorption) energies DGi of the adsorbedsurface intermediates as function of the electrode potentialand combined these with a thermodynamic model in whichpH and electrode potential are represented as control parame-ters.[44b,d] The authors related the Gibbs chemisorption energiesof the intermediates to the normal hydrogen electrode (NHE)by assuming equilibrium conditions for the H2/H+ redoxsystem,[45] which rendered the calculation of the chemical po-tential of an aqueous proton unnecessary. The difference be-tween the chemisorption energies of two subsequent inter-mediates was equal to the Gibbs free reaction energy, DGrxn, ofthe elementary mechanistic step at a given electrode potential.The differences in chemisorption energy caused each individu-al DGrxn to turn negative at different electrode potentials.Under the premise that each elementary step must be thermo-chemically downhill (i.e. each DGrxn must be negative) for theoverall reaction to occur at an appreciable rate, the authorscould predict the minimum electrode (over)potential for theoverall reaction to occur. Koper and Heering reported that thelast elementary step to attain a negative DGrxn value is the ‘po-tential-determining step’ in analogy to the kinetic concept of arate-determining step.[46] We note that the ‘potential-determin-ing step’ and the rate-determining step are not necessarilyidentical.

Figures 2 and 3 illustrate the idea of the thermochemicaloverpotential for a mechanism of the hydrogen evolution reac-tion (HER).[47] The reaction is assumed to follow the simple 2-electron Volmer–Heyrovsky reaction sequence:

2 Hþþ2 e�DG1

! HðadÞþHþþe� ð23Þ

HðadÞþHþþe�DG2

! H2ð24Þ

where DGi denotes the elementary Gibbs reaction energies ofEquations (23) and (24).

Figure 2 shows the Gibbs free adsorption energies of the in-dividual species of Equations (23) and (24) as horizontal lines.The Gibbs free energies of adsorption can be derived from thetotal energies of adsorption, DE“H+ ,e�”, DEH(ad)

, and DEH2, of the

reacting proton and electrons, the intermediate H(ad), and theproduct H2, respectively.[44d] The x axis represents the reactioncoordinate. A typical electrocatalyst (Figure 2, blue traces) ex-hibits an imbalance between DG1 and DG2 ; in this case, for in-stance, DG1>DG2. At electrode potentials E1 above the reversi-ble potential Eo (positive overpotentials : E1�Eo = h>0;Figure 2, lower traces) all elementary reactions are uphill andthermodynamically unfavorable. As the electrode potential islowered, all Gibbs energies of charged species shift more posi-tively by n e E, where n is the number of electrons in Equa-tions (23) and (24) and e the elementary charge. At the reversi-ble electrode potential E2 (E2�Eo =h= 0; Figure 2, centertraces), the elementary reaction in Equation (23) has a positiveGibbs free energy and is therefore the potential-determiningstep. At the electrode potential E3 (negative overpotential ;Figure 2, dashed blue top trace), the elementary reaction inEquation (23) becomes reversible. There are no more thermo-dynamic barriers along the reaction pathway, and hence, thepotential (E3�Eo) represents the thermochemical overpotentialof the hydrogen evolution of the real catalyst (Figure 2, blueline). This value represents the overpotential at which the over-all reaction rate becomes independent of the applied electrodepotential and may reach a certain prescribed magnitude. Kinet-ic barriers are not considered in this framework; they maykeep the actual reaction rate low and introduce a further elec-trode-potential dependence of the reaction rate.

The energetics of the ideal catalyst (Figure 2, red trace) ischaracterized by equidistant adsorption energy levels of the re-active species, implying identical Gibbs reaction energies ofthe two elementary steps, DG1 =DG2 at any given electrodepotential. As a result, the Gibbs free energies of all species and

Figure 2. Plot of Gibbs free energies of reactive species and Gibbs free ener-gies of chemisorptions of intermediates (horizontal lines) versus the reactioncoordinate of the hydrogen evolution reaction (HER). Blue lines and redlines indicate energetics of a typical (real) catalyst and ideal catalyst at threedifferent electrode potentials (E1, E2, E3), respectively. Dashed lines indicateenergetics at the electrode potential where all thermochemical barriers dis-appear. DGi denotes the free reaction energy of the two elementary reactionsteps. The red case corresponds to an overpotential-free catalyst. Eo denotesthe standard reversible electrode potential of the HER, h is the overpoten-tial.



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intermediates are also equal at the equilibrium potential(Figure 2, red dashed trace), and hence no thermodynamicoverpotential is expected.

Figure 3 shows the relation between the negative thermo-chemical overpotential, hthermo, of each reaction step as func-tion of the energy of adsorption DEH(ad)

of the H(ad) intermedia-te.[44a, 48] The thermochemical overpotential is taken as a mea-

sure of activity and is defined as hthermo =DGoi /e evaluated at

the reversible electrode potential Eo of the process. The plothas a volcano-type shape with a single maximum. At morenegative DEH(ad)

(left side of volcano) hydrogen is adsorbedstrongly and hence the desorption of atomic hydrogen DG2

controls the thermodynamic overpotential. On the oppositeside of the volcano, the adsorption of atomic hydrogen limitsthe overall reaction. As “DEH(ad)

(real)” of the real catalyst(Figure 3) is more positive than the optimum (red dot), reac-tion (23) with Gibbs energy DG1 controls the overall rate. Theideal catalyst (Figure 3) has equidistant energy levels andhence “DEH(ad)

(ideal)” is at the top of the volcano. At the equi-librium potential, all thermodynamic barriers disappear andthe reaction is ideally reversible without any thermochemicaloverpotential. This simple analysis shows that for two-electronprocesses with one adsorbed intermediate I, there generallyexists an ideal chemisorption energy “DEI(ad)

(ideal)” of the inter-mediate such that the thermochemical overpotential vanishes.Based on that, the search for the ideal catalyst material wouldrequire tuning DEI(ad)

to the ideal value.[49]

The anodic electrocatalytic OER[44b, d] (Figures 4 and 5) in-volves three adsorbed intermediates; OH(ad), O(ad), and OOH(ad)

(Figure 1). Similar to the HER mechanism above, the OERmechanism underlying this analytical framework consists exclu-sively of single-electron charge-transfer steps[44d] according toEquations (25)–(28):

2 H2ODG1

! OHðadÞþHþþe�þH2O ð25Þ

OHðadÞþHþþe�þH2ODG2

! OðadÞþ2 Hþþ2e�þH2O ð26Þ

OðadÞþ2 Hþþ2e�þH2ODG3

! OOHðadÞþ3 Hþþ3e� ð27Þ

OOHðadÞþ3 Hþþ3e�DG4

! O2þ4 Hþþ4 e� ð28Þ

The four-step mechanism of Equations (25)–(28) differs fromthe classical mechanisms for acidic and alkaline solutions de-scribed by Equations (12)–(15) and (18)–(21), respectively.

As for HER, we compare and contrast the energetics of anideal (red) and a real (blue) catalyst in Figures 4 and 5. The in-dividual chemisorption energies of the intermediates reactingon the real catalyst are assumed to be associated with theGibbs reaction energy order DG3>DG1 =DG2>DG4 ; the for-mation of the peroxide intermediate (DG3) is the thermochemi-cally least favorable step for the real catalyst. Figure 4 indicatesthat the OOH(ad) species is bound a little too weakly to the cat-alyst. As for HER, the ideal catalyst shows equally spacedchemisorption energies of the intermediate and hence equalGibbs reaction energies for each elementary step; DG3 =DG1 =

DG4 = DG2. For instance, at an overpotential of h=�1.23 V,corresponding to an electrode potential of 0 V vs NHE (“zerobias” at pH 0), the overall change in free energy of the four-electron water splitting reaction amounts to 4.92 eV; in this sit-uation, the ideal catalyst would exhibit a Gibbs reactionenergy of 1.23 eV for each step. Taking water as the zero pointof the energy scale, the Gibbs energies of adsorption of theideal intermediates would be 1.23 eV, 2.46 eV, and 3.69 eV forOH(ad), O(ad), and OOH(ad), respectively.

Figure 4 illustrates the energetics of the real and ideal cata-lyst at various electrode potentials. At electrode potential E1,all steps are uphill for either catalyst and hence OER cannot

Figure 3. One-dimensional thermochemical volcano plot of the HER. Thenegative thermodynamic overpotential �hthermo is plotted against the totalhydrogen chemisorption energy DEH(ad)

. The left and right side of the volcanoare associated with a “potential-determining” reaction step given in the plot.The volcano curve provides the thermodynamic overpotential, at which theHER reaction occurs at an appreciable rate (potential independent rate). Thecircles correspond to the real and the overpotential-free ideal case fromFigure 2, respectively. In this two-electron process, tuning the hydrogen ad-sorption energy can result in the ideal case.

Figure 4. Plot of Gibbs free energy of reactive species and intermediates(horizontal lines) of the oxygen evolution reaction (OER) versus the reactioncoordinate. Blue lines and red lines indicate energetics of a real (typical) cat-alyst and an ideal catalyst, respectively, at three different electrode poten-tials. Dashed lines indicate energetics at the electrode potential where allthermochemical barriers disappear (“thermochemical overpotential”) ; DGi

denotes the free reaction energy of the two elementary reaction steps. Thered ideal case corresponds to a catalyst with vanishing overpotential.



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proceed. At the reversible potential E2 (Figure 4, center, solidblue and dashed red traces), the real and ideal catalystsbehave distinctly differently. For the real catalyst, DG1 and DG2

vanish, DG4 is negative, and DG3 remains positive, hinderingthe OER. Thus, Equation (27) describes the potential-limitingstep. The ideal catalyst shows no more uphill energetics at thereversible electrode potential and has therefore no overpoten-tial, provided the kinetic limitations are negligible. The elec-trode potential E3 (Figure 4, bottom) is required in order tomake all steps downhill for the real catalyst, at which OER canoccur at a prescribed rate.

Simple graphical inspection of Figure 4 suggests that atheory-guided improvement of the real OER catalyst (blue)would need to shift the chemisorption free energy of theperoxo species OOH(ad), DGOOH(ad)

, to more negative values, char-acterizing a stronger surface bonding until equally spaced re-action free energies are obtained (red). Rossmeisl et al. ,[44b, d]

however, demonstrated that the independent change of asingle chemisorption energy to achieve optimum relative ad-sorption energetics is not feasible, because the chemisorptionenergies of the three individual surface intermediates of theOER mechanism in Equations (25)–(28) on solid surfaces arestrongly correlated.[44a,c, 50] The authors’ calculations suggestthat the individual chemisorption energies of the intermediatesare linked by linear scaling relations. Their work further provid-ed evidence that these relations appear to be independent ofthe detailed composition of the solid catalyst surface for a

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 valueswill also change DEi for all other intermediates, so that ener-getic equidistance is in principle unachievable. Secondly, basedon a given set of linear scaling relations, the reaction free ener-gies DGi (i = 1–4) can be expressed by using a one-dimensionalvolcano plot with a single controlling chemisorption parame-ter,[44–45] such as DEO(ad)

.The solid lines in Figure 5 illustrate the linear relations be-

tween the negative thermodynamic overpotential, �hthermo =

�DGoi /e, evaluated at the reversible potential, and the parame-

ter DEO(ad)for all four elementary steps on a rutile-type oxide

surface.[44d] For instance, the scaling relation between the ad-sorption energies of OOH(ad) and O(ad) on rutile surfaces was cal-culated in detail[44d, 51] as

DEOOHðadÞ ¼ 0:64 DEOðadÞ þ 2:03 eV ð29Þ

Taking zero point energies and entropies into consideration,this equation becomes

DGOOHðadÞ¼ 0:64 DGOðadÞ þ 2:40 eV ð30Þ

The scaling relations between OH(ad) and O(ad) were calculat-ed as

DEOHðadÞ ¼ 0:61 DEOðadÞ� 0:90 eV ð31Þ

and

DGOHðadÞ ¼ 0:61 DGOðadÞ � 0:58 eV ð32Þ

Koper pointed out that the difference, DGOOH(ad)�DGOH(ad)

, re-sults in an approximately constant energy gap of about3 eV,[51] which is clearly above the ideal gap of 2.46 eV requiredin an ideal case. The difference in the constant energy term isbelieved to originate from the difference in the chemicalnature and bonding of the two intermediates and is likely notto depend greatly on the nature of the catalytic surface.Herein lies the fundamental challenge in improving the wateroxidation reaction rate.

The scaling relations [Equations (30) and (32)] of the inter-mediates further imply that OH(ad) binds too strongly comparedto the ideal case, whereas OOH(ad) binds at close to the idealcase. The energetic position of the near 3 eV gap relative tothe ideal case depends on the nature of the surface; on metal-lic close-packed (111) surfaces, OH(ad) was found to bind closeto ideality, whereas OOH(ad) binds much too weakly.

From Equations (29)–(32), the following Gibbs reactionenergy relations were derived, with E denoting the electrodepotential versus NHE:[44d]

DG2 ¼ 0:39 DEOðadÞ þ 0:60 eV� e E ð33Þ

DG3 ¼ �0:36 DEOðadÞ þ 2:38 eV� e E ð34Þ

Figure 5. Relation between the thermodynamic overpotential, hthermo, associ-ated with elementary steps of the OER [Equations (25)–(28)] and the totalenergy of adsorption of ‘nonmolecular’ oxygen DEOad at the reversible elec-trode potential on a non-ideal (real) and ideal catalyst surface (solid linesand dashed lines, respectively). Energetics for the two catalysts are assumedto be consistent with considerations in Figure 4. Solid lines follow theoreticalstudies on rutile (110) surfaces in reference.[44d] For the real catalyst, overpo-tentials associated with Equations (28), that is, DG3, and (27), that is, DG2, aredominant over most adsorption energies. The linear relations of DG3 (solidgreen) and DG4 (solid brown) control the overall overpotential of the OERon the real catalyst and are the “potential-determining steps” at low andhigh DEO(ad)

.[44d] The blue circle represents the situation of the real catalyst inFigure 4. Tuning of DEO(ad)

to the transition between the green and brown re-lation optimizes the catalyst for the given scaling relations. However, theoverpotential cannot be completely eliminated for this set of scaling rela-tions. Dashed lines represent an ideal set of scaling relations where appro-priate tuning of DEO(ad)

results in an overpotential-free catalyst and OER reac-tion (red circle).



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The slope was related to the making or breaking of surfacebonds, with positive and negative slopes indicating the forma-tion and breaking of a bond, respectively. Values close to 0.5were calculated for the slope on metal(111) surfaces, whereasthe slopes on the rutile surfaces were slightly off a value of 0.5.

A graphical representation of DGo2 and DGo

3 at the reversibleelectrode potential E = 1.23 V is given by the brown and greensolid lines in Figure 5. These two reactions control the thermo-dynamic overpotential over a large portion of the oxygenchemisorption energy range shown.

The blue circle in Figure 5 and its associated -hreal value indi-cate the thermodynamic overpotential of the real catalyst fromFigure 4, given its oxygen chemisorption energy DEOðadÞ (real).As confirmed in Figure 5, the Gibbs energy of formation ofOOH, DG3 (Figure 5, green), controls the thermochemical over-potential hreal for the real catalyst (blue). Following Equa-tion (34), weaker oxygen chemisorption (more positive DEOðadÞ )lowers the thermodynamic overpotential until a minimum isreached at the transition where the formation of adsorbedatomic oxygen [Equation (27)] becomes energetically least fa-vorable (Figure 5). At even weaker oxygen chemisorption, DG2

and reaction (27) control the overpotential. It is clear thattuning the chemisorption energies of the intermediate will notresult in an overpotential-free ideal catalyst (Figure 4, red).

To achieve the ideal case, novel catalysts with a distinct setof ideal scaling relations need to be found, allowing for equallyspaced reaction free energies (Figure 5; dashed lines). The redball indicates the combination of ideal scaling relations withideal oxygen chemisorption energy (DEOðadÞ �2.46 eV, DGOðadÞ =

2.46 eV) resulting in vanishing overpotential. Each bond-form-ing and bond-breaking linear relation falls on top of eachother and intersect exactly in the red ball at the ideal oxygenchemisorption energy. The slopes of the ideal scalings are notexactly known, yet an absolute value around 0.5 is assumed, inline with findings for metal(111) surfaces. Even for ideal scalingrelations, the actual oxygen chemisorption energy may be sub-optimal, suggesting that the experimentally observed OER re-action rate of a material with ideal scaling yet suboptimaloxygen chemisorption may be inferior to those for the current-ly known materials, and offering a practical challenge in theidentification of ideal water-splitting catalysts.

On the surface of a real catalyst surface, such as a supportedmetal oxide nanoparticle, the multiplicity of structurally distinctfacets leads to a range of surface sites with distinctively differ-ent coordination and chemisorption properties. Generally, low-coordinated sites bind adsorbates more strongly and hence, incase of the real catalyst in Figure 5, there may be a small frac-tion of surface sites with a more favorable adsorption thermo-chemistry. The overall overpotential will then be a convolutionof individual overpotentials of different sites. Under these con-ditions, it is feasible that different surface sites are thermo-chemically limited by different elementary reactions.

In conclusion, a simple thermochemical framework[44a,d] hasprovided insights into qualitative trends of the OER overpoten-tial and reactivity on well-defined metal oxide surfaces. Theframework predicts linear scaling relations between the chemi-sorption energies of intermediates. These scaling relations

imply limits in the ability to design a low-overpotential electro-catalyst for reactions with multiple adsorbed intermediates,such as the electrochemical evolution of molecular oxygen.Direct experimental support of generalized scaling relationsbetween reaction intermediates has not been reported todate, yet is a current subject of much interest. Experimentalmethods to determine surface binding energies of oxygen-containing intermediates in electrochemical surfaces, such ascalorimetry and in situ ambient pressure photoemission spec-troscopy, will likely play a critical role in such future studie-s.[43a,b, 52]

The origin of the generalized scaling relations for differentsurface adsorbates, such as CHx, NHx, OHx, or SHx, appears tobe related to the number of unsaturated bonds of the inter-mediates, the nature of the binding site, and the resultingbond order of the intermediate. Rutile(110) and metal(111) sur-faces exhibit distinctly different sets of scaling relations for theoxygenated surface;[44d] some other structural classes of solidsurfaces may exhibit scaling relations that more closely matchthe desirable equidistant Gibbs free energy levels of the inter-mediates. However, for an overpotential-free catalyst, both thescaling relations and the adsorption energies have to be ideal(Figure 5). A surface with ideal scaling relations, yet unfavora-ble adsorption energies can exhibit a lower catalytic activitythan a surface with unfavorable scaling yet optimal bindingenergies. Hence, optimizing a water splitting electrocatalystmay possibly involve tuning both the slope and y-intercept ofthe scaling relations in combination with tuning surface ad-sorption energies - an ambitious task that will require muchmore fundamental and practical knowledge on how to experi-mentally tune surface-bond energies. The slope of scaling rela-tions appears to be loosely related to the bond order of the in-termediate species, whereas the y intercept could be related tothe oxygen atom.

Experimental catalysis research is faced with a huge range ofstructural and compositional materials parameters to be ex-plored in the quest for a reversible electrocatalyst. Computa-tional predictions of suitable surface structures with favorablescaling and binding energies, possibly combined with combi-natorial and high-throughput screening approaches, in con-junction with in situ spectroscopy will be critical tools for de-signing and developing reversible electrocatalysts for watersplitting.[53]

2. Water Oxidation by Bulk Metal Oxides

In 2008, Nocera and co-workers reported efficient water oxida-tion by a cobalt oxide film formed by electrodeposition from asolution containing cobalt, potassium, and phosphate ions.[54]

Hereafter, we will denote the cobalt catalyst film as CoCF.[55]

The report of Nocera and co-workers has attracted much at-tention and may have initiated a revival of bulk metal oxidesas catalysts for water oxidation (we use the term ‘bulk metaloxides’ to differentiate these from the ‘surface oxides’ formedon metal electrodes which are exposed to a voltage sufficientlypositive for water oxidation). Without attempting a compre-hensive review, we will discuss selected cornerstone studies on



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bulk metal oxides as catalysts for water oxidation, with a focuson cobalt oxides.

In 1950, El Wakkad and Hickling[56] reported the performanceof various CoO/OH films deposited at 5 mA cm�2 fromCoSO4·7 H2O, H3BO3, and NH4Cl, on Pt wire electrodes. Thesefilms were tested in a variety of electrolytes, among which wasa solution of KH2PO4 and NaH2PO4 at pH 6.8 (0.2 m each). Theyfound that the presence of phosphate stopped the dissolutionof Co observed with other alkaline and acidic electrolytes, andthat an initial dissolution was reverted in phosphate buffer,which indicates that the films had a self-repair mechanism inphosphate buffer. The ability of the films to split water was notdiscussed. Benson, Briggs, and Wynne-Jones[57] deposited filmsfrom 0.1 m Co(NO3)2 at 5–10 mA cm�2 on platinum and nickelsubstrates. Deposition in KOH yielded a black film (like theCoCF[54]) with the composition K0.08CoO1.51·1.03 H2O and suffi-cient crystallinity to resolve the unit-cell parameters (hexago-nal ; a = 6.75 �, c = 5.36 �).[57a] Oxygen evolution is reported,but the origin of the oxygen is not discussed.[57b] Better docu-mentation of the oxygen evolution of an early CoO/OH filmmay be found in the patent application of Osamu et al.[58] Theircrystalline films have the formula CoOm·n H2O (m = 1.4–1.7, n =

0.1–1.0). O2 evolution was described for operation in variousnonbuffering electrolytes.

Harriman, Pickering and Thomas compared the oxygen evo-lution of various metal oxide powders.[59] They used spinel-type Co3

II/IIIO4, but IrO2 showed the highest activity. The rate ofO2 evolution for Co3O4 was given as 25.5 �106 mol dm�3 min�1.[59] At the beginning of the 1990s, Tseungand co-workers published an extensive series of reports on re-active deposition of cobalt oxides from 0.25 m CoCl2 at20 mA cm�2.[60] The surface morphology of their electrodesshowed the same nodules as the CoCF described by Noceraand co-workers, and the reported overpotential of 0.39 V foroperation in 7 m KOH was virtually identical to that of theCoCF in 0.5 m PO4 from KH2PO4 and K2HPO4 (pH 7).

Much of the more recent research has focused on crystallineCo3O4. Singh et al.[61] prepared spinel-type Co3O4 by micro-wave-assisted synthesis on a Ni support. The catalyst shows anapparent current density of 100 mA cm�2 in 1 m KOH at roomtemperature with an overpotential of 0.24 V. They showed thatthe overpotential could be lowered to 0.22 V by doping withlanthanum. Jiao and Frei reported nanostructured Co3O4 in amesoporous silica support,[62] which has a turnover frequency(TOF) of 0.01 s�1 per surface Co atom for an overpotential of0.35 V, at pH 5.8 and room temperature. The latter catalyst isphotochemically driven by a [RuIII(bpy)3] species with visiblelight (476 nm, 240 mW). Frei estimated that a stack of 100Co3O4 nanorod bundles would be needed for a TOF of100 s�1 nm2 required to keep up with the solar flux.[62b]

The cobalt catalyst for electrochemical water oxidation re-ported by Kanan and Nocera[54] has attracted much interest be-cause of its efficiency at neutral pH, self-assembly from low-cost materials, and for its self-repair mechanism; it is discussedin the following sections.

2.1. Properties of an electrodeposited cobalt catalyst

The aforementioned catalyst developed by Kanan and Nocerawas first produced by electrodeposition from an aqueous solu-tion of KH2PO4 and K2HPO4 (KPi) at pH 7 containing[Co(OH2)6(NO3)2] , with concentrations of 0.1 m (KPi) and 0.5 mm

(Co2+) on an indium tin oxide (ITO) substrate. More recently,the following modifications have been discussed: 1) Depositionon other semiconductor and metal substrates, such as glassycarbon, carbon felt, fluorinated tin oxide (FTO), and nickelmetal,[63] and especially importantly, deposition and operationon a-Fe2O3

[64] and ZnO photoanodes;[65] 2) exchange of theelectrolyte, specifically, substitution of potassium forsodium[54, 66] and phosphate for phosphinate or borate.[66b]

The significance of point (2) is that neither potassium norphosphate itself is essential for catalytic activity. Likely, they arenot part of the catalytic unit itself. For discussion of the mech-anism, it may not be required to extend the Co�oxo units (Fig-ures 7 and 8) by a distinct cation, such as potassium, or anionicligand, such as phosphate. Thus, we refer to this catalystherein as the cobalt catalyst film (CoCF), instead of cobaltphosphate (Co–Pi) catalyst used elsewhere.

The CoCF forms on a variety of conducting substrates, suchas those discussed above. Surface morphology and film thick-ness depend on the deposition time, the composition of theelectrolyte, and the applied voltage.[63] The first layers of thedeposited films conform to the topology of the substrate.[64b]

In thicker films, it is possible that only the ragged surface iselectrochemically active.[64b] Thinner films, with depositiontimes of 15 min, gave more stable performance.[64a] CoCF de-posited at voltages in the range of a prewave in the cyclic vol-tammogram assigned to Co oxidation are microscopicallysmooth, whereas deposition at voltages which promote wateroxidation produces CoCF with nodules on the surface.[54, 63] Theoverpotential to split water at pH 7 (for a current of1 mA cm�2) in the presence of KPi was reported as 0.41 V.[54]

Current densities as high as 100 mA cm�2 were reported for un-specified plastic composite electrolyzer cells operated at60 8C.[63b] For a thin film, we estimate the TOF to be as high as0.2 s�1 per O2 molecule for the benign conditions reported inRef. [54] at 1.45 V vs NHE.[64c]

Both methylphosphonate and borate electrolytes foster cata-lyst growth and support catalytic activity. These anions behavein the same manner as phosphate but clearly differently fromnonbuffering electrolytes such as SO4

2�, NO3� , and ClO4�.[66b]

Catalyst operation and formation does not require deionizedreagent-grade water. Nocera and co-workers demonstrated op-eration in brine and river water,[63, 66b] suggesting that otherions, and especially chloride ions, do not inhibit O2 evolution.Potassium and sodium as components of the electrolyte werefound to be interchangeable. Films grown with either alkalimetal cation were catalytically active after cation exchange.[66a]

Leeching of both the alkali metal cation and the phosphateanion is faster than that of cobalt, as shown by radioactive la-beling experiments.[66a] This observation supports a cobalt-onlyframework for the catalyst, as proposed based on X-ray ab-sorption spectroscopy (XAS)[55] and discussed below (sec-



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tion 2.2). Recently, McKlintock and Blackman[67] proposed anew bidentate bonding mode of phosphate for the termina-tion of the Co�oxo units proposed by Risch et al.[55] The dis-tance between Co and P atoms is only 2.52 �. The X-ray ab-sorption spectroscopy (XAS) data in ref. [55, 68] does not pro-vide support for such a Co�P distance, but phosphate coordi-nation to a minor population of Co ions can not be ruled outentirely.[55]

Hill and co-workers recently reported an efficient water oxi-dation catalyst consisting of a soluble, carbon-free heteronu-clear molecule, [Co4(H2O)2(a-PW9O34)2]10�, in which a planar[Co4(m-O)6O10] unit may represent the catalytically activepart.[69] It will be interesting to elucidate whether the mecha-nism of water oxidation in the molecular Co4W9 oxide and inCoCF is the same.

2.2. Atomic structure of the catalytic cobalt oxide

For the highly amorphous CoCF, diffraction techniques are notapplicable. Thus, XAS is especially well suited to analyze thelocal 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 meas-urements at 20 K. Average bond lengths between the absorb-ing cobalt atoms and its neighboring ‘shells’ of atoms relate topeaks of the FT. The x axis shows reduced distances which areabout 0.3 � lower than the true nucleus–nucleus distance,which can be determined by EXAFS simulations. The amplitudeof the peaks in the FT is a rough measure of the averagenumber of the atoms at the respective distance. The peaks as-signed to Co�Co vectors of a CoCF deposited at 1.15 V vs NHE(Figure 6, red circles) are higher than for a film deposited at1.35 V (Figure 6, black squares), indicating more extended

long-range order. The FT of crystalline LiCoO2[70] is shown for

comparison. In summary, films deposited at voltages lowerthan the catalytic wave are microscopically smooth[63b] andatomically more ordered than films deposited under the sameconditions but at voltages high enough to support catalytic ac-tivity. The latter films show more disorder, not only on the mi-croscopic scale (nodules) but also at atomic level (Figure 6).The difference in the order at the atomic level suggests thatCoCF cannot be viewed as an aggregate of a multinuclearcobalt complex of a distinct size. Either size and nuclearity ofthe molecular unit depends on the growth conditions or, morelikely, CoCF is a heterogeneous assembly of interconnected{Cox(m2/3-O)y} units, as discussed in more detail in the following.

XAS simulations on the cobalt catalyst films (CoCF) suggestthat the central structural unit is a cluster of interconnectedcomplete or incomplete CoIII�oxo cubanes (Figure 7).[55, 68] The

octahedral geometry of the Co sites and the oxidation statefound by XAS was confirmed recently by EPR spectroscopy.[71]

In the case of incomplete cubanes, the extent of edge-sharingneeds to be higher to comply with the simulation results. Wemodified the illustration from ref. [55] to emphasize that thecatalytic unit most likely is not a distinct molecule. Figure 7shows the proposed bulk structure of the CoCF catalyst, illus-trating the possible presence of complete and/or incompletecubanes. Further interconnections of the building blocks mayresult in an extended but overall disordered network of com-plete and/or incomplete cubanes.

Here, we also consider another motif agreeing with the XASdata (Figure 8).[55, 68] It is a tile-shaped Co10O32 unit composedof edge-sharing octahedra, which are exclusively incomplete

Figure 6. Fourier-transform (FT) of an EXAFS spectrum of the cobalt catalystfilm (CoCF) for growth at a voltage not supporting water oxidation (1.15 Vversus NHE; circles) and in the catalytic regime (1.35 V; squares) comparedto crystalline LiCoO2 (triangles). LiCoO2 and CoCF are characterized by thesame octahedral Co building blocks, but in LiCoO2 those blocks are arrangedin an extended layer of side-sharing {CoIIIO6} octahedra, which can beviewed also as a layer of side-sharing, incomplete Co–oxo cubanes {Co3(m-O)4}. Simulations of the data are shown as lines. All FT peaks shown exceedthe noise level; for further experimental details, see ref. [68] .

Figure 7. Possible structural motif, deduced from XAS data, for the bulk ofthe cobalt catalyst film (Co: dark grey, O: light grey). The catalytic film maybe 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} cu-banes. For an example of the latter, see Figure 8. The protonation states ofthe bridging oxides and terminal water species remain unknown.



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Co-oxo cubanes. The macroscopic CoCF would be composedof a large number of these tiles; the space between them islikely filled with cations (e.g. K+), anions (e.g. HPO4

2�), andwater, in analogy to more-ordered layered structures, such asLiCoO2 or Co-containing buserites.[72] We emphasize that itwould not be in conflict with the EXAFS data if some of thetiles were loosely interconnected.

The metal oxide/hydroxide materials that catalyze water oxi-dation frequently were operated electrochemically in an elec-trolysis experiment. However, it is unclear to what extent theconcepts described in section 1 for surface oxides can be ap-plied to the amorphous bulk metal oxides. Clearly, the CoCFlacks the long-range order of crystalline solid-state materials. Itis an interesting conceptual question whether the CoCF can beviewed best as an extended solid-state material or rather anaggregate of multi-nuclear cobalt–oxo complexes with molec-ular properties. Water content and the analogy with layeredtransition metal oxides suggests that water molecules fill thegaps between structures (Figures 7 and 8). Thus, it is conceiva-ble that not only the apparent surface catalyzes water oxida-tion but the bulk material can also contribute. In conclusion,the bulk metal oxides represent a catalytic material that mayshare properties with both heterogeneous electrocatalysts andhomogeneous molecular catalysts.

3. Molecular Catalysts for Water Oxidation

In recent years, molecular catalysts for water oxidation havealso received a considerable amount of attention. The relativeease with which molecular systems can be studied mechanisti-cally and can be tailored both structurally and electronically

gives them an advantage over heterogeneous systems andallows for a more systematic approach in the design of newgenerations of water oxidation catalysts with increased perfor-mance. However, the prediction of molecular structures thatare able to catalyze the oxidation of water to oxygen at lowoverpotentials is rather difficult. Despite of their importance inthe context of artificial photosynthesis purposes, only a rela-tively small number of molecular water oxidation catalysts ishitherto known, with the majority of them having been discov-ered only in the last few years.

Based on the multinuclear structure of the oxygen evolvingcomplex in photosystem II (PS II ; see section 4) and the discov-ery of the nowadays well-known “ruthenium blue dimer” wateroxidation catalyst [(bpy)2Ru(H2O)(m-O)(H2O)Ru(bpy)2]4 + in theearly 1980s (Figure 9),[73] it was initially assumed that, in molec-

ular complex catalysts, the presence of at least two metal cen-ters was essential to accomplish the mechanistically demand-ing oxidative coupling of two water molecules to dioxygen.Both the accumulation of the required amounts of oxidationequivalents and a beneficial geometrical orientation of the twowater molecules (or rather oxo ligands in later mechanisticsteps) prior to the formation of the O�O bond were believedto be more difficult to accomplish in mononuclear systems. Ac-cordingly, since the beginning of the 21st century, a number ofdinuclear manganese and ruthenium complexes have beenshown to be more or less effective catalysts for water oxida-tion, which seemed to confirm the prerequisite of at least twometal centers in active water oxidation catalysts. Consistently,mechanistic studies on Llobet’s dinuclear Ru–Hbpp water oxi-dation catalyst (bpp = bis(2-pyridyl)-3,5-pyrazolate, Figure 9)have evidenced that oxygen–oxygen bond formation proceedsintramolecularly, benefitting from the favorable ligand matrixthat preorientates the metal centers and bound H2O ligands ofthis dinuclear complex for the coupling reaction.[74] However,water oxidation catalyzed by the blue dimer is nowadays be-lieved to occur at a single Ru site within the dinuclear com-plex, which raised the idea that water oxidation catalysis mightalso be possible with mononuclear metal complexes.[75] In fact,within the last couple of years the discovery of several mono-nuclear water oxidation catalysts has been reported, whichrepresents one of the major recent breakthroughs in the fieldof water oxidation catalysis both from a mechanistic and a syn-thetic point of view. However, the development of new and ef-fective molecular water oxidation catalysts is not the only im-portant target to be pursued; the incorporation of the water

Figure 8. Structure of m-oxo-bridged cobalt atoms, which is compatible withthe EXAFS data (Co: drak grey, O: light grey).[55, 68] This {Co10O32} unit has anaverage of 3.8 Co�Co vectors at 2.8 � and 1.6 Co�Co vectors at 5.6 � (perCo atom) and thus satisfies the constraints resulting from simulations of theXAS data.[55, 68] It is conceivable that several of these ‘tiles’ are connected toform an extended network or porous sheet. Layers of water molecules andcations may separate the Co�m-oxo sheets, in analogy to layered Mn or Codioxides.[70, 72]

Figure 9. Examples of known dinuclear water oxidation catalysts.



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oxidation catalysts into integrated systems for solar water split-ting has to be achieved, and also in this respect mononuclearcomplexes bear a lot of potential. Owing to their ease of modi-fication and adaptation to the given needs, they are consid-ered as very promising candidates for surface grafting.

3.1. Water oxidation catalysis by mononuclear metal com-plexes

As a consequence of the increased interest in artificial photo-synthesis in general and water oxidation catalysis in particular,several review articles with emphasis on different aspects ofthis reaction have been published recently, and these reviewshave also taken molecular compounds into account.[76] Never-theless, the growing number of publications dealing withwater oxidation certainly justifies an update of the field. In thepast two years, several studies with respect to mononuclearwater oxidation catalysts have been reported while only fewnovel dinuclear systems have appeared and most new findingsconcerning dinuclear systems were obtained in the course ofcontinuing studies on already established systems.[74b, 77] In thischapter we will therefore concentrate on mononuclear wateroxidation catalysts (mostly ruthenium complexes) and mentiondinuclear systems only briefly.

3.2. Mononuclear ruthenium complexes as water oxidationcatalysts

There are some early reports about water oxidation catalyzedby 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-py)(4-Rpy)2(H2O)]2 + (Scheme 1) to be active water oxidation

catalysts in the presence of excess Ce4 + oxidant in aqueousacidic media.[78] The catalytic activity of these complexesstrongly depends on the nature of the 4-pyridyl substituent Rgiving turnover numbers (TON) between 20 (R = NMe2, that is,NHMe2

+ in acidic media) and 260 (R = CH3). From the observa-tion that the sterically encumbered derivative 4 (Scheme 1) isalso an active water oxidation catalyst (TON 260), it was con-cluded that the mononuclear constitution of 4 is preserved

during the water oxidation cycle, given that, for steric reasons,an intermediate formation of dimeric or higher nuclear speciesshould be inhibited.[78b]

To investigate the electronic influence of the ligand environ-ment on the catalytic activity of mononuclear ruthenium wateroxidation catalysts, Thummel et al. then studied a wide rangeof related ruthenium complexes.[78b, 79] One class of the investi-gated compounds contained ruthenium centers coordinativelysaturated exclusively with N-donor ligands, mostly of the pyri-dine type, of varying denticity (Figure 10); a second series wascharacterized by the general formula [Ru(tpy)(NN)Cl](PF6)(NN = bidentate N-donor ligand; Scheme 2).

In contrast to 1–4, for which the electronic absorption, theredox properties, and the catalytic activities were strongly de-pendent on the nature of the axial 4-Rpy ligands,[78a] substitu-tion at the equatorial ligands in 5–14 (Figure 10) did not havea strong or consistent influence on these properties; no corre-lation between the catalytic activity in water oxidation and theligand set was revealed. At pH 1, TONs between 146 (10) and416 (5) were determined for these catalysts in the presence ofa 5000-fold excess of [Ce(NH4)2(NO3)6] as a stoichiometric oxi-dant.[79]

For complexes 15–21 (Scheme 2), the dependency of theelectronic absorption and redox properties on the ligand set ismore distinct. With increasing electron-withdrawing characterof the 4,4’-substituents at the bipyridyl ligand (bpy, 15–19), thecharacteristic metal-to-ligand charge transfer (MLCT) band isshifted to longer wavelengths. At the same time, the oxidationpotential of the RuII/III couple, as measured in acetonitrile, in-creases. However, no direct correlation between the redox

Scheme 1. Mononuclear water oxidation catalysts bearing the binapy ligand.Turnover numbers are given in parentheses.

Figure 10. Mononuclear water oxidation catalysts containing ruthenium cen-ters coordinatively saturated with pyridine-type ligands. Turnover numbersare given in parentheses.



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properties and the catalytic activity of these complexes couldbe drawn: Both the presence of electron-donating (e.g. CH3 in16) or electron-withdrawing (e.g. NO2 in 18) 4,4’-substituentsled to lower TONs as compared to the parent system 15. A re-markably high activity (TON= 1170) was found for 20, althoughthe reason for the distinct activity in this special case remainedunclear. Exemplarily, mechanistic studies were performed for15. The initial rate of O2 evolution was found to be first orderin catalyst concentration, indicating that dimerization is not aprerequisite for these complexes to be active in water oxida-tion. Additional evidence for a mechanism in which the origi-nal monomeric nature of the catalyst is conserved was ob-tained by the isolation of [RuIII(tpy)(bpy)Cl]2 + as an intermedi-ate from a reaction mixture, which had undergone approxi-mately 10 water oxidation turnovers. This complex was identi-fied as such by mass spectrometry and also chemically byshowing that addition of ascorbic acid as a reductant leadsback to the RuII complex 15, as evidenced by 1H NMR spectros-copy. On the basis of these findings and supporting DFT calcu-lations on 5, a mechanism for catalysis of water oxidation bycomplexes 5–21 was proposed, which involves the formationof seven-coordinate ruthenium centers during the course ofthe 4-electron oxidation of water (Scheme 3).[78b]

Accordingly, removal of two electrons from [RuIIL6] generatesa 16-electron [RuIVL6]2 + species (Scheme 3), which should be

susceptible to coordination of a water molecule as an addition-al ligand. Subsequent loss of two protons and two electrons(proton-coupled electron transfer) generates a high-valent [O=

RuVIL6]2 + species. Attack of a second water molecule on theelectrophilic oxo group then leads to the formation of a[RuIVL6OOH]+ hydroperoxide intermediate. The loss of asecond 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 aviable mechanism, was recently supported by the identificationof the seven-coordinate ruthenium aqua complex [m-(HO-HOH)(RuIV{6,6’-(COO)2bpy}(pic)2)2]3 + ·2 H2O 23 (pic = 4-picoline;Figure 11) as an intermediate of water oxidation catalyzed by

[RuII{6,6’-(COO)2bpy}(pic)2] 22 ; accordingly 23 itself is catalyti-cally active, too.[80] It should be noted, however, that in thiscase (consistent with the dimeric structure of 23) the kineticdata and also the results of calculations[81] suggest that dimeri-zation of 22 plays an important role in the course of the reac-tion (see below).

Assuming that none of the N-donor functions temporarilydissociates from the Ru centers to create an empty site forwater coordination, a mechanism like that outlined inScheme 3 might indeed be followed for catalysts 5–14. Forcomplexes of the type 15–21 ([Ru(tpy)(NN)Cl]), however,recent findings of Sakai et al. and especially of Meyer et al. in-dicate a different mechanism of water oxidation.

Sakai et al. independently investigated the related mononu-clear ruthenium complexes [Ru(tpy)(bpy)Cl]+ , 15, and [Ru(t-py)(bpy)(H2O)]2+ , 26, as well as their dinuclear congeners [Ru(t-py)(Cl)L(Cl)(tpy)Ru]2 + and [Ru(tpy)(H2O)L(H2O)(tpy)Ru]4 + (L =

bis[5-(5’-methyl-2,2’-bipyridinyl)]ethane; Scheme 4).[82]

The potential of these complexes as water oxidation cata-lysts was investigated in water/acetonitrile mixtures at pH 0.4using [Ce(NH4)2(NO3)6] as the oxidant, monitoring oxygen evo-lution with a Clarke-type electrode in solution. Interestingly,the dinuclear ruthenium complexes were found to be lessactive catalysts than the related mononuclear ones. The aquacomplexes 24 and 26 proved superior to the correspondingchloro complexes 25 and 15, which was the main finding of

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

Scheme 3. Water oxidation mechanism proposed for mononuclear rutheni-um complexes exhibiting Ru centers coordinated by N-donor ligands.

Figure 11. Seven-coordinate intermediate of water oxidation catalyzed by[RuII{6,6’(COO)2bpy}(pic)2] 22.



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this study. A 2–3 h induction period for the onset of water oxi-dation was observed when the chloro complexes 25 and 15were employed as the catalysts. This observation may be ex-plained by an equilibrium between the chloro and aqua com-plexes 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 ofoxygen.[82]

An extensive study of the [RuII(NNN)(LL)(H2O)]n + aquasystem was performed by Meyer et al. , who investigated sever-al related water oxidation catalysts, differing in the nature ofthe tridentate NNN and bidentate LL ligands (Figure 12).[75, 83]

From electrochemical and kinetic studies on the complexes[Ru(tpy)(bmp)(H2O)]2 + 27 and [Ru(tpy)(bpz)(H2O)]2 + 28 c, a de-tailed mechanism for the oxidation of water, with Ce4 + as theoxidant, was derived (Scheme 5).[83c] In diluted acidic media(0.1 m HNO3) 27 ([RuII�OH2]2+) is rapidly oxidized to [RuIV=O]2 +

by two equivalents of Ce4+ . Subsequent reaction with furtherCe4 + leads to the formation of a high-valent [RuV=O]3+ species,which reacts with water to give the terminal hydroperoxide[RuIII�OOH]2+ . The latter is further oxidized to the peroxo com-plex [RuIV(O2)]2+ , whose decomposition, releasing molecular di-oxygen, is the rate-limiting step of water oxidation in dilutedHNO3. As a result of the stronger oxidizing power of Ce4 + atlower pH, oxidation of [RuIV(O2)]2 + to [RuV�OO]3+ competeswith its decomposition in 1 m HNO3. Oxygen evolution fromthe RuV-peroxide is followed by subsequent oxidation of [RuIII-OH]2 + to [RuIV=O]2+ , which closes the catalytic cycle. The plau-sibility of the proposed intermediates was supported by DFTcalculations, which suggested that the initial intermediate[RuIII�OOH]2+ is best described as a terminal peroxide, whereasthe peroxide ligand of [RuIV�OO]2 + is coordinated in a side-onfashion.

Similar mechanistic investigations for some other [Ru(N-NN)(LL)(H2O)]2+ derivatives suggested that the mechanism de-picted in Scheme 5 should be generally valid for catalysts be-longing to this class of complexes, albeit with changes in theindividual reaction rates and rate limiting steps.[83b]

Meyer’s investigations clearly evidence that water oxidationat single metal sites is not only possible with Ir complexes (seebelow), but also at ruthenium centers. The easier syntheticaccess to mononuclear Ru compounds of the type shown inFigure 12 (as compared to the known dinuclear systems) dis-burdens the precise tailoring of the ligand sphere for the link-age to surfaces, for the synthesis of molecular assemblies, orfor the systematic investigation of electronic factors governingthe catalytic activity. For instance, the complex cation 29 a�PO3H2 [Ru(Mebimpy){4,4’-(H2O3PCH2)2bpy}(H2O)]2+ (Scheme 5)has been successfully grafted on FTO, ITO or TiO2 nanoparticlesurfaces to result in materials with good electrocatalytic activi-ties.[84]

Recently, Meyer et al. found that the presence of redox me-diators such as [Ru(bpy)2(LL)]2 + (LL = bpy, bpm, bpz) provokesa significant rate enhancement in water oxidation, when theruthenium blue dimer and [Ce(NH4)(NO3)6] are used as the cat-alyst and stoichiometric oxidant, respectively.[73c] Consequently,the effect of such mediators was studied for the mononuclearsystems as well.[85] Molecular catalyst–mediator assemblies of

Scheme 4. Mono- and dinuclear complexes of the [Ru(NNN)(NN)X]n +-type.Turnover numbers are given in parentheses.

Figure 12. Meyer’s mononuclear ruthenium water oxidation catalysts of thetype [RuII(NNN)(LL)(H2O)]n + . No maximum turnover numbers were reportedfor these catalysts. 28 a = 26. DMAP = 4-dimethylaminopyridine.



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the type [(bpy)2RuII(bpm)RuII(LLL)(OH2)]4 + (LLL = tpy: 30 ; LLL =

Mebimpy: 31), in which a mediator and a water oxidation cata-lyst are combined within a single molecule, were obtainedfrom the reaction of [RuII(bpy)2Cl2]·2 H2O with[RuII(LLL)(bpm)Cl]+ (bpm = 2,2’-bipyrimidine) in ethanol/watermixtures (Scheme 6). In this case, the bipyrimidine ligand, be-longing to the redox mediator as well as to the catalyst, playsthe important role of bridging the two functional units withinthe dinuclear ruthenium complex.

Cyclic voltammetry (CV) analysis of 30 at pH 1 revealed mul-tiple sequential one-electron oxidation steps until formation ofthe species [(bpy)2RuIII(bpm)RuV(tpy)(O)]6+ , which triggers cata-lytic water oxidation. In these systems the reaction of[(bpy)2RuIII(bpm)RuV(tpy)(O)]6 + with water was found to be thekey O�O bond formation step, leading to the hydroperoxospecies [(bpy)2RuIII(bpm)RuIII(tpy)(OOH)]5+ . The latter then un-dergoes further oxidations by Ce4+ until oxygen is released,presumably from the peroxidic species[(bpy)2RuIII(bpm)RuV(tpy)OO)]6 + , resulting in the formation of[(bpy)2RuII(bpm)(tpy)RuIV=O]4+ , which closes the catalytic cycle.Anchoring the closely related redox assemblies [{(4,4’-(H2O3PCH2)2bpy}2RuII(bpm)RuII(NNN)(OH2)]4 + (NNN = tpy, Me-bimpy) to ITO or TiO2/FTO (nanoparticle TiO2 films on FTO) led

to impressively high turnovernumbers up to 28 000 and turn-over rates up to 0.6 s�1 in wateroxidation experiments for thelatter systems without any signsof activity loss after longer peri-ods of time.[85] The presence ofthe redox mediator in the heter-ogeneous assembly allowedwater oxidation catalysis to beperformed at 1.43 V, which isclose to the thermodynamic po-tential for the O2/H2O couple atpH 0. Notably, under these con-ditions, no electrocatalytic cur-rent was detected for solutionsof the related monomeric cata-lyst[Ru(tpy)(bpm)(H2O)]2 + . Meyer’srecent findings doubtlessly rep-resent major progress towardsthe design of integrated water-

splitting devices based on artificial photosynthesis, for whichthe availability of robust electrocatalysts is an essential prereq-uisite.

Besides the mononuclear ruthenium complexes stemmingfrom the extensive studies performed by the groups of Thum-mel and Meyer, a few other examples have been shown to beactive as water oxidation catalysts, too. In addition to theaforementioned investigations on 15 and 24–26, Sakai et al. re-ported ruthenium aqua complexes of the type [Ru(NNN)(4,4’-R2bipy)(H2O)]2+ containing the facially coordinating 1,4,7-tri-methyl-1,4,7-triazacyclononane (tmtacn) ligand rather than themeridional tpy or Mebimpy derivatives, which were employedby Thummel and Meyer (Scheme 7).[86]

As a consequence of the lacking p acceptor and the morepronounced s-donating properties of the tmtacn ligand, theelectrochemical potentials of the RuII(H2O)/RuIII(OH) andRuIII(OH)/RuIV=O couples of 32–34, as determined in aqueousacidic solution, are strongly shifted to lower potentials, for ex-ample, in comparison to [Ru(tpy)(bpy)(H2O)]2 + 28 a. On theother hand, from these reduced oxidation potentials, no lower-ing of the overpotential for electrocatalytical water oxidationfollowed. For example, for 34, bearing the bpy ligand with themost strongly s-donating substituent, the potential that hadto be applied to initiate water oxidation was found to be of asimilar magnitude to that required by the tpy derivative 28 a.For the derivatives 32 and 33, even higher potentials werenecessary. Water oxidation experiments were also performed indiluted aqueous solutions with [Ce(NH4)2(NO3)6] as a chemicaloxidant, and oxygen evolution was monitored with a Clarke-type electrode. From the total amounts of oxygen that hadevolved after a 40 h period, turnover numbers amounting to148 (32), 173 (33), and 251 (34) were determined, revealing anincreasing catalytic activity with increasing electron-donatingcapabilities of the 4,4’-bpy substituents. However, the activitieswere again lower than that for the tpy derivative 28 a ; for this

Scheme 5. Mechanism for water oxidation catalyzed by Ru–aqua complexes [Ru(NNN)(LL)(H2O)]n + , as proposed byMeyer et al.

Scheme 6. Catalyst-mediator assemblies for water oxidation catalysis.



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complex, the same authors determined a TON of 310 underanalogous conditions. Accordingly, the replacement of the tpyligand by tmtacn leads to a lowering of the Ru-based oxidationpotentials in 32–34, although they were found to be generallyless active in water oxidation than 28 a.

To incorporate the water oxidizing process into a lightdriven integrated system for water splitting, the generally ap-plied cerium(IV) oxidant has to be replaced by a photosensitiz-er, which takes over the responsibility of providing the neces-sary amounts of oxidation equivalents required by the wateroxidation catalyst. For this purpose, a lowering of the oxidationpotentials of the catalysts to a certain degree can be advanta-geous and, as can be seen from the investigations on com-plexes 32–34 this goal can be achieved by introduction ofstronger s-donating and less p-accepting ligands. A differentstrategy for lowering the oxidation potentials of mononuclearRu-based water oxidation catalysts was followed by Sun et al. ,who, as mentioned before, studied the catalytic properties ofthe complex [RuII{6,6’-(COO)2bpy}(pic)2] 22 (Scheme 8).[80a,b]

As expected, the presence of two negatively charged car-boxylate functions as part of the bpy-type ligand in 22 signifi-cantly lowers the oxidation potentials of the ruthenium center,as compared, for example, to 28 a. Water oxidation was in-duced by addition of 22 to an aqueous solution containingCe4 + . However, only a rather modest TON of approximately120 was determined by monitoring oxygen evolution with aClark-type oxygen electrode in solution. Kinetic analysis re-vealed a second order decay of Ce4 + for 22 indicating that

water oxidation with 22 as the catalyst is based on a binuclearprocess, which very recently has also been found feasible bycalculations:[81] Dioxygen was proposed to be formed by cou-pling of two RuIV–oxyl radicals. Mass spectroscopic investiga-tions with the aim of identifying intermediates of the catalyticcycle revealed a peak for [RuIV{6,6’-(COO)2bpy}(pic)2(OH)]+ , andconsistently a compound containing two of these cationiccomplex units, that is, the abovementioned seven-coordinatedruthenium complex [m-(HOHOH)(RuIV{6,6’-(COO)2bpy}(pic)2)2](PF6)3·2 H2O, 23(PF6)3, (Figure 11, Figure 13)

was isolated after addition of NH4PF6 to the reaction mixture.In 23, two [RuIV{6,6’-(COO)2bpy}(pic)2(OH)]+ units are bridgedby a proton sitting in the middle between the two OH� li-gands. Two water molecules are bonded to the two ends ofthe internal [HOHOH] moiety by bridging hydrogen atoms,and additional hydrogen bonds between these water mole-cules and the 6- and 6’-carboxylate groups of adjacent com-plex cations further stabilize the overall hydrogen bonding net-work as well as the dinuclear structure of 23.[80a]

The active participation of 23 during water oxidation wasnot only suggested by the second order kinetic dependence ofthe decay of Ce4+ in catalyst 22 but further evidenced byproving the ability of 23 to catalyze water oxidation itself :Oxygen evolution was observed in an aqueous acid solutionwhen [Ce(NH4)2(NO3)6] was added as an oxidant. Although23(PF6)3 does not represent the first example of a structurallycharacterized seven-coordinate ruthenium complex, its suc-cessful isolation and identification as an intermediate in wateroxidation, catalyzed by 22, certainly contributes valuable infor-mation with respect to possible reaction pathways for the con-version of H2O to O2 at transition metal sites. However, itshould be borne in mind that all other mononuclear rutheni-um water oxidation catalysts that have been discussed to date

Scheme 7. Water oxidation catalysts of the [Ru(NNN)(NN)(H2O)]2 + type, con-taining the facially chelating tmtacn ligand. Turnover numbers are given inparentheses.

Scheme 8. Synthesis of the mononuclear water oxidation catalyst [RuII{6,6’-(COO)2bpy}(pic)2] 22. Turnover number is given in parentheses.

Figure 13. Molecular structure of [m-(HOHOH)(RuIV{6,6’-(COO)2bpy}(pic)2)2](PF6)3·2 H2O, 23(PF6)3, a possible intermediate of water oxi-dation catalyzed by [RuII{6,6’-(COO)2bpy}(pic)2] 22. The hydrogen atoms ofthe organic ligands and the (PF6)� counterions are omitted for clarity.



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obey first-order kinetics of oxygen evolution with respect tothe corresponding catalysts, rendering a participation of dimer-ic {Ru2} species during the catalytic process unlikely for thosesystems. Therefore, water oxidation starting out from mononu-clear ruthenium catalysts, but proceeding via intermediate for-mation of dinuclear species, rather represents a special case.Regarding 22, dimerization certainly is favored by the presenceof the 6,6’-carboxylate groups in the ligand backbone, whichcontribute to a stabilization of the dinuclear Ru–hydroxo struc-ture by offering base functions for hydrogen-bonding interac-tions.

Under neutral aqueous conditions, cyclovoltammetric analy-sis of solutions containing 22 as an electrocatalyst showed acatalytic water oxidation curve beginning at 0.98 V (vs NHE).This relatively low potential for water oxidation allowed theutilization of [RuIII(4,4’-R2bpy)3]3 + (R = H, CH3) as photosensitiz-ers to realize photochemical water oxidation. The latter wasperformed in a three-component system composed of catalyst22, [Ru(4,4’-R2bpy)3]2+ , and [Co(NH3)5Cl]Cl2 or Na2S2O8 as sacrifi-cial electron acceptors at pH 7.0.[80b] In the absence of light, noformation of oxygen could be detected. Irradiation, however,immediately triggered oxygen evolution, which rapidly ceasedwhen the light was turned off (Scheme 9).

Initial TOFs of greater than 550 h�1 and greater than1250 h�1 were determined for the CoIII and S2O8

2� systems, re-spectively. However, both systems suffered from rapid deacti-vation, and oxygen evolution stopped after only 7 min for CoIII

or 50 s for S2O82�, under the diluted conditions required for

oxygen detection in the liquid phase. However, when wateroxidation was performed in solutions containing higheramounts of catalyst, oxygen evolution, as monitored by GC,persisted for about 2 h and a TON of 100 was determined forthe CoIII system. The fairly quick loss in activity on employmentof Na2S2O8 as the sacrificial electron acceptor could be tracedback to the continuous acidification of the solution in thecourse of water oxidation [Equation (3) and Scheme 9], whichdoes not occur with [Co(NH3)5Cl]Cl2, which simultaneously actsas a buffer. Consequently, it is possible to regenerate the pho-tocatalytic activity of the persulfate system by neutralization of

the reaction mixture for a certain number of times. Other deac-tivation pathways, valid for the CoIII and for the S2O8

2� system,include the oxidative decomposition both of the photosensitiz-er and the catalyst. Notably, in the absence of [Ru(bpy)3]2 + , nophotodegradation of 22 occurred under similar reaction condi-tions.

A very different and novel reaction scheme for water oxida-tion was discovered by Milstein et al. , who showed that themononuclear PNN–ruthenium pincer-type complex 34 is capa-ble of performing the entire water splitting reaction(Scheme 10).[87] Remarkably, 34 mediates the consecutive gen-

eration of hydrogen and oxygen, probably via a H2O2 inter-mediate, from water at a single metal center via thermally andphotolytically induced steps. A central role within these trans-formations is played by the cooperative participation of thePNN ligand, which undergoes facile deprotonation and proto-nation at the methylene unit of the Py’CH2PR2-arm. Milstein’ssystem has found considerable attention and recent highlightarticles reflect the great potential attributed to this novelwater splitting scheme,[88] although the overall reaction is ac-tually stoichiometric and slow with rather low yields in gener-ated oxygen and hydrogen.[87]

Quantum chemical calculations have further elucidated themechanism of this new system.[89] The release of hydrogen,which is induced by proton transfer from the (Py’CH2PR2) armof the aromatic PNN ligand in 34 I to the ruthenium-bound hy-dride, was found to be the rate-limiting step of the reaction.The presence of water reduces the activation barrier for thisstep from 37.6 to 33.6 kcal mol�1 (1 kcal = 23.9 kJ) by the for-mation of a C�H···OH2···H�Ru bridge between the methyleneproton and the {Ru�H} unit. Nevertheless, this relatively highactivation barrier might be the reason for the slow rate of H2

generation and the low yield of 34 II (45 %) even after 3 daysat 100 8C. Time-dependent DFT calculations on 34 II suggestedthat the photochemical reductive elimination of H2O2, as theultimate source of O2 (see above), from 34 II, for which twostrong Ru�OH bonds need to be broken, occurs from a disso-ciative triplet state via a singlet–triplet crossing.

Scheme 9. Photochemical water oxidation in a three-component systemconsisting of catalyst 22, photosensitizer [Ru(R2bpy)3]2+ and a sacrificial elec-tron acceptor. Redox potentials for the onset of oxygen evolution and forthe [Ru(R2bpy)3]2+ /3 + couple are given vs NHE at pH 7.

Scheme 10. Water splitting at a single ruthenium center.



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3.3. Non-ruthenium systems

In contrast to the rapidly developing field of ruthenium-cata-lyzed water oxidation, examples of catalysts containing othermetal centers remain scarce, both for multinuclear and mono-nuclear systems. Hence, most of the known water oxidationsystems still rely on highly expensive and/or rare transitionmetals, and this reliance can be considered one of the majordrawbacks regarding future application of artificial photosyn-thesis for power generation. Nevertheless, some examples ofwater oxidation or water splitting promoted by mononuclearcomplexes of metals other than ruthenium have been reportedrecently and are discussed in this section.

Purely photochemical water splitting has been carried outwith osmocene as a photocatalyst (Scheme 11).[90] Photolysis of

an acidic aqueous solution containing [Cp2OsIVH]+ , 35 I, withwhite light induced the evolution of hydrogen with concomi-tant formation of the osmocene(IV) complex [Cp2OsIV(H2O)]2+ .It was assumed that, on photolysis, hydrogen is released fromdimeric [Cp2OsIVH···HOsIVCp2]2 + units, leading initially to[Cp2OsIII�OsIIICp2]2+ , which then undergoes photodisproportio-nation to give [Cp2OsIV(H2O)]2 +

. The latter was found to bephotostable under acidic conditions, but when an aqueous so-lution of the closely related hydroxo complex[Cp2OsIV(OH)](PF6), 35 II(PF6), was photolyzed, a nearly quantita-tive conversion of 35 II to oxygen and [Cp2OsII] resulted. Photo-oxidation of 35 II probably involves the intermediate genera-tion of the peroxo complex [Cp2OsIII(O2)OsIIICp2]. The results ofthese two separately studied photochemical processes may beformally combined to an interesting cyclic process(Scheme 11). Difficulties associated with the pH-dependent for-mation of oligomeric (m-OH), (m-O), and (m-Cp) osmocene spe-cies in aqueous solutions of [Cp2OsIV(H2O)]2 + , however, have sofar impeded the establishment of a real catalytic photochemi-cal water-splitting system.[90]

Bernhard et al. recently discovered a new family of iridium-based water oxidation catalysts [IrIII{2-(4’-R’Ph)(5-Rpy)}2(H2O)2]+ ,36, for which impressively high TONs of up to 2760 (R1 = Me,R2 = F) were reported (Figure 14).[76a, 91] However, the rates ofoxygen evolution were relatively low. Related mononuclear Ircomplexes 37–39 were obtained by formally replacing one 2-(2-pyridyl)phenyl (ppy) and one water ligand in 36 by the elec-

tron-donating pentamethylcyclopentadiene (Cp*) ligand(Figure 14).[92]

Oxygen evolution experiments were performed in dilutedsolutions using a Clark-type oxygen electrode, which revealedcomplexes 37–39 to be much more active catalysts than 36.Initial rates amounting to 0.91 and 0.28 turnovers per secondwere determined for 37 and 39, respectively, which are signifi-cantly higher than that which the same authors determinedfor 36 (R1 = R2 = H; 0.005 s�1) under analogous conditions.However, the turnover numbers attained by these new iridiumcatalysts (1500 for 37, 213 for 39) were lower. Oxygen evolu-tion was found to be first order in the Ir catalyst, which rendersit plausible that a similar mechanism for water oxidation is fol-lowed as in the case of the earlier discussed mononuclearruthenium complexes [RuII(NNN)(LL)(H2O)]n + 27–30 (seeabove). DFT calculations on the potential IrV intermediate[Cp*IrV(ppy)(O)]+ revealed the presence of a low lying p* (Ir=O)orbital, which includes a significant contribution from the Oatom. Formation of an O�O bond might therefore result fromnucleophilic attack of a water molecule on this high-valent iri-dium–oxo unit.[92]

The use of defect polyoxometalate (POM) derivatives withvacant metal sites as ligands for transition metals is an interest-ing approach for the design of water oxidation catalysts withenhanced stability, because decomposition routes based onoxidative degradation of organic ligands is not possible. There-fore, the recent discovery that the purely inorganic Ru poly-oxometalate [{Ru4O4(OH)2(H2O)4}(g-SiW10O36)2]10� represents anactive water oxidation catalyst, has stimulated ongoing re-search efforts on this class of catalysts.[69, 77a, b, 93] As a mononu-clear (with respect to the catalytic site) representative, the iridi-um polytungstate K14[(IrCl4)KP2W20O72]·23 H2O, 40, was synthe-sized by the reaction of IrCl3 with [PW9O34]9� in the presence ofKOH in water.[94] Low yields of O2 (30 %) were detected by GCwhen [Ru(bpy)3]3 + and 40 were mixed at neutral pH. However,40 was found to be unstable in solution, dissociating into[IrCl4(H2O)2]� and [KP2W20O72]13�. Although this dissociation wasfound to proceed much slower than the oxidation of water inthe presence of 40, it could not be ruled out that, as a conse-quence of dissociation, small amounts of catalytically activeIrO2 are responsible for the formation of the detected oxygen.

By using ligands containing xanthene and corrole unitslinked to each other by amide bonds, �kermark et al. synthe-sized the mono- and dinuclear manganese complexes 41 and42 (Scheme 12).[95] By means of cyclic voltammetry on basic

Scheme 11. Hypothetical catalytic cycle of water splitting using osmoceneas a photocatalyst.

Figure 14. Mononuclear iridium-based water oxidation catalysts.



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CH2Cl2/CH3CN solutions, both complexes were shown to act aswater oxidation catalysts, with 42 being more active than 41.The activity of the mononuclear manganese complex 41 indi-cated that water oxidation might proceed—similarly to whathas been considered a possible O�O bonding step in naturalphotosynthesis—via a single-site mechanism involving nucleo-philic addition of water or OH� to a high-valent mangane-se(V)–oxo unit. DFT calculations supported this single-sitemechanism as the preferred pathway with both 41 and 42.[96]

Recently, the same authors reported further experimental evi-dence for such a water oxidation mechanism, demonstratingthat rapid oxygen generation resulted from the addition ofBu4NOH to a solution containing [Corrole–MnV=O] 43 a (Cor-role = 5,10,15-tris(4-nitrophenyl)corrole) in acetonitrile(Scheme 12).[97] Ultimately, however, only very small amountsof dioxygen, corresponding to 0.01 turnovers of 43, were de-tected. By using 18O-labeled water, it could be proved that oneof the oxygen atoms in the evolved oxygen originated fromwater. On the basis of the results of these 18O-labeling studiesin combination with UV/Vis spectroscopy and high-resolutionmass spectrometry investigations, it was proposed that, afteran initial oxidation of the [MnIII–Corrole] complex 43 bytBuOOH to give 43 a, attack of a hydroxide ion at the electro-philic MnV=O function leads to the MnIII–hydroperoxide species43 b. The latter undergoes further oxidation, presumably by43 a, yielding the MnIV–peroxide 43 c before the elimination ofoxygen occurs. Although the efficiency of oxygen generationin the Mn–corrole/tBuOOH system is only very low, the findingthat water oxidation at manganese centers can proceed vianucleophilic addition of hydroxide to high-valent manganese–oxo groups, that is, at single sites, in a similar manner to themononuclear Ru and Ir systems, is an important finding notonly regarding the debate about the actual mechanism of nat-ural water oxidation in PS II ; it might also stimulate the devel-opment of more manganese-based water oxidation catalysts,

which remain very scarce, inspite of the importance of cata-lysts based on cheap and abun-dant materials.

3.4. Conclusions

In recent years, increased re-search efforts on the develop-ment of new molecular wateroxidation catalysts exhibiting en-hanced performance and stabili-ty has led to a notable progressin this area, resulting in the dis-covery of some new families ofcatalysts, along with a deeper in-sight into the mechanisms thatcan lead to the generation ofoxygen from two water mole-cules. The realization that dimer-ic or higher-order assemblies ofmultiple metal centers are not

essential to accomplish water oxidation catalysis certainly rep-resents a major advance. The easier derivatization associatedwith structurally simpler mononuclear systems allows for amore rational synthesis of tailored systems to study the struc-ture–activity relationships that might govern water oxidation.Similarly, an introduction of surface-binding groups can facili-tate the strong attachment of molecular catalysts to electrodesurfaces or enable the synthesis of molecular assemblies suitedfor coupling light-harvesting and/or redox mediation to wateroxidation. As the final goal is to use the electrons resultingfrom the oxidation of water to drive fuel-forming reactions,growing efforts are dedicated to the coupling of the differentprocesses involved in artificial photosynthesis. In this context,light-driven water oxidation mediated by suitable photosensi-tizers plays a central role, as does the attachment of molecularcatalysts to conducting electrode surfaces. Inspection of thenew generation of water oxidation catalysts, however, revealsthem to be mostly based on expensive ruthenium metal cen-ters. Therefore, the discovery of effective and long-lived wateroxidation catalysts derived from abundant and low-priced first-row transition metals remains a major objective for future re-search. Advances concerning this aspect might be consideredindispensable in terms of a significant contribution of artificialphotosynthesis applications for sustainable power generationin the near future.

4. Biological Water Oxidation

Plants, algae, and cyanobacteria use solar energy to oxidizewater and reduce plastoquinone molecules. The light-drivenwater–plastoquinone oxidoreductase is denoted as photo-system II (PS II).[98] This cofactor–protein complex of impressivesize and complexity is embedded in a lipid bilayer membrane(thylakoid membrane), which separates the aqueous phases ofthe lumen and stroma compartment (Figure 15). Crystallo-

Scheme 12. Mn–corrole complexes active in water oxidation.



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graphic models of the core complex of PS II have been deter-mined[99] and, more recently, a resolution of 2.9 � was ach-ieved.[100] The PS II core complex is most likely evolutionarilyfully conserved with respect to every site of direct functionalimportance. In intact organisms, additional pigment-carryingproteins are found, which increase the light-harvesting capabil-ity of the core complex. Composition of these peripheral an-tenna proteins is species-dependent and relates to differencesin color observed when comparing taxonomic groups of pho-tosynthetic organisms (e.g. green, red, brown, gold, or bluealga).[98, 101]

Light absorption by one of the numerous chlorophyllgroups (Chl) of PS II is followed by ultra-fast excitation energytransfer and excitation of a special pair or group of chloro-phylls denoted as P680.[102] The excited-state energy of P680*drives a sequence of electron-transfer (ET) steps as well as theproton-coupled redox reactions at the electron acceptor anddonor sides of PS II. Firstly, by ultrafast ET, a pheophytin group(Phe) is reduced, resulting in formation of the P680+ Phe� radi-cal pair. At the acceptor side, this process is followed by rapidreduction of QA (ca. 300 ps), a firmly bound plastoquinone (Q)followed by slower ET to QB, a second plastoquinone molecule(ca. 1 ms; Figure 15). A second photon drives formation of adoubly reduced and protonated QH2 molecule, which leavesthe QB site; a third quinone in PS II, QC, may also be in-volved.[100, 103]

At the donor-side, P680+ is reduced by electron transferfrom a specific redox-active tyrosine, denoted as YZ, within lessthan 1 ms. This ET step is assumed to be coupled to the move-ment of the phenolic proton within a hydrogen bond towardsa nearby histidine sidechain:

P680þþYZ�OH � � � N�His! P6800þYZ�O� � � � Hþ N�His ð35Þ

Hereafter, the states on the left and right side are denotedas YZ and YZ

· + , respectively. YZ· + formation is followed by elec-

tron transfer from a pentanuclear {Mn4Ca} complex, which rep-resents the catalytic center of the photosynthetic water oxida-tion. Water oxidation involves four sequential steps of one-electron oxidation of the catalytic complex by YZ

· + , resulting inthe following overall equation:

4 YZ�þþ2 H2O! 4 YZþ4 HþþO2 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 potentialin electrochemical water oxidation. The lowest pH for efficientwater oxidation by PS II is around 5.5,[105] which corresponds toan equilibrium potential of the H2O/O2 couple of about 0.9 V(at pO2

= 1 bar). The potential difference of about 0.3 V (=1.2 V�0.9 V) can be viewed as the ‘overpotential’ of photosyn-thetic water oxidation. The corresponding energetic efficiencycalculated for the half-cell reaction [Equation (11) in section 1.1]is 80 % when the second half-cell reaction is assumed to beproton reduction, and around 75 % when referenced againstplastoquinone reduction by PS II. These figures are impressive,especially when taking into account that water oxidation inPS II is so fast that it efficiently outcompetes recombination-loss pathways.[105b]

The turnover frequency (TOF) of the complete PS II may beas high as 100 s�1 per O2 molecule; it is determined by rate-limiting reactions at the acceptor side. For catalytic processesat the donor side, a maximal TOF of about 400 s�1 per O2 mol-ecule can be calculated from the rate constants of the elec-tron-transfer steps. The maximal turnover number (TON) of thecatalyst is determined by the so-called PS II repair cycle (seesection 4.2) and may be around 105. For a review of the ener-getics and solar energy conversion efficiency of PS II, seeref. [104c] . The catalytic process is discussed in the followingsection.[106]

4.1. Structure of the catalyst

In 2001, Zouni and co-workers presented the first crystallo-graphic model (at 3.8 � resolution) of a PS II complex in whichthe position of the catalytic site, the {Mn4Ca} complex, wasuniquely determined.[99a] Higher resolution resulted in structur-al models for the Mn complex and its ligand environment,showing that the protein-derived ligands are carboxylates sup-plemented by a single histidine residue.[99b–d, 100, 107] However, allcrystallographic data related to a non-native arrangement ofthe Mn ions. The high X-ray intensity needed for data collec-tion results in rapid X-ray photoreduction of the {Mn4

III,III,IV,IV} toa level where all Mn ions are present in the MnII state.[108]

Whereas Mn reduction and the loss of the internal m-oxo struc-ture of the Mn complex by X-ray irradiation during crystallo-graphic data collection is well documented,[108c] it is largely un-clear to what extent the coordinates of the Mn ligands aremodified.

One Ca ion is likely connected by two or more bridgingoxides to Mn ions, as suggested by EXAFS spectroscopy[109]

Figure 15. Redox factors in PS II. The D1 and D2 protein subunits carry all es-sential redox factors. Besides the D1 protein, the CP43 subunit also providesligands to the {Mn4Ca} complex. The proteins of the PS II core complex carry-ing antenna pigments (CP43, CP47) and peripheral antenna proteins are notshown. PS II is part of the thylakoid membrane, a dense array of proteinsand lipids separating the aqueous compartments at the PS II acceptor side(lumen) and donor side (stroma).



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and crystallography.[99b, d] After substituting chloride by bro-mide, the question of halide ligation was addressed by EXAFSat the Br K-edge using an approach previously established fora bromoperoxidase and synthetic model complexes.[110] The re-sults show that the halide is not a first-sphere ligand of man-ganese and is possibly at a distance of about 5 � to either aMn or a Ca ion.[111] However, recent results obtained by proteincrystallography suggested a metal–halide distance of about6.5 �.[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 (2 � Mn�Mn distances of 2.7 �), whichare connected by a mono-m-oxo, bis-carboxylato bridge (Mn�Mn distance of 3.3 �), such that a C-shaped structure result-ed.[115] Simulations of EPR data[116] and the crystallographic re-sults[99] are difficult to reconcile with a C-shaped dimer-of-dimers model. In 2004, Barber and co-workers presented acrystallographic model involving a {Mn3Ca(m3-O)4} cubane con-nected to the fourth Mn ion by a bridging oxide.[99b] Thismodel was used as a starting point for DFT calculations[117] andmechanistic considerations,[106b, 117–118] but was difficult to recon-cile with EXAFS results.[108a, 119] Moderate modifications resultedin a model which may be compatible with the EXAFS data.[120]

Recently, XAS data were collected on single crystals ofPS II,[119] thereby providing further constraints for structuralmodels of the Mn complex, which complement earlier resultsobtained by XAS on isotropic[114c, 121] and unidirectionally orient-ed PS II samples.[108a, 121b, 122] Comparison of the single-crystalEXAFS with a set of hypothetical structures led to a hypothesison the structure and orientation of the {Mn4Ca(m-O)n} core ofthe Mn complex; the protein-derived ligand environment wasnot modeled.[119] An alternative working model of the Mn com-plex and its ligand environment is shown in Figure 16;[113] simi-lar structures have been invoked in computational studies.[123]

The following structural features of the Mn complex of PS IIhave been identified unequivocally: The {Mn4Ca} complex isbound to several carboxylate oxygens as well as to a single N-histidine; the metal ions are internally interconnected by bridg-ing oxides. There are 2 or 3 di-m-oxo bridges between Mn ionsinvolving m2-O and presumably also m3-O bridging. The Ca ioncould cap the bridging oxides (Figure 16).

4.2. Formation of the catalytic metal complex

Assembly of the {Mn4Ca} complex of PS II does not require anychaperone-like enzymes but is driven by light and thus denot-ed as photoactivation (of PS II).[124] The mechanism involves theoxidation of Mn2+ ions by YZ

· + formed upon illumination andformation of the {Mn4Ca(m-O)n} core within the ligand systemprovided by the protein. The affinity of the apoprotein to Mnions is so high that stoichiometric amounts are sufficient forcomplete assembly (4 Mn ions per PS II) ; a {Mn2(m-O)2}complexhas been identified as a likely intermediate of the assemblyprocess.[125] Some amino acid residues that are assumed toligate Mn ions can be exchanged against nonligating residuesand the Mn complex is assembled nonetheless, albeit withlower stability.[126] This relative insensitivity to details of the

ligand environment may relate to a metal–(m-O) core that re-sembles an intrinsically stable motif also found in mineralic Mnoxides;[127] the Mn complex of PS II may resemble a minusculepiece of rock. Nonetheless, its oxidative assembly and water-oxidation function clearly require the presence of most of theligands provided by the (apo)proteins. The striking structuralsimilarities have inspired the recent investigations on water ox-idation by transition metal oxides which resemble minerals(see below).

Neither the Mn complex nor the PS II protein complex are asstable as the rock. The so-called PS II repair cycle involves ex-change of the D1 protein against a fresh copy; the D1-turnoverfrequency may be as high as 2 per hour for organisms exposedto high light intensities.[128] The D1 protein is likely exchangedafter oxidative damage to chlorophylls, quinines, or amino acidresidues.[129] It is unlikely that the costly exchange of the D1protein is aiming to repair the Mn complex. Repair of the{Mn4Ca(m-O)n} core is likely achieved in a different manner;when exposed to light, the photoactivation mechanism contin-uously counteracts the occasional loss of Mn ions from thecomplex, representing a cost-efficient and inconspicuousrepair mechanism of the metal complex itself. In conclusion,the {Mn4Ca(m-O)n} core of the catalyst is formed after oxidationof Mn2 + ions by self-assembly at the ligand system provided

Figure 16. Model of the catalytic site for water oxidation in PS II. The Mnligand environment was obtained by molecular-mechanics modeling, start-ing with the crystallographic coordinates of Loll et al.[99d] The model of the{Mn4Ca(m-O)n} core is based on EXAFS results.[113] The shown protein-derivedgroups are from the D1 subunit, unless labeled as being CP43 side chains.The structure in the S1 state is shown. The m2-oxo group between Mn3 andMn4 is proposed to be protonated in the S0 state (m2-OH). In the S2!S3 tran-sition, a hydroxide bridging between Mn1 and Mn2 may become deproton-ated and transformed into a sixth ligand of Mn1.[114] A cavity close to Mn4could harbor several water molecules (dotted line). The Asp61 is close toMn4 and has been proposed to represent the entrance of a path for protontranslocation to the aqueous lumen. Accumulation of oxidizing equivalentsby the Mn complex involves electron transfer to YZ

· + , the oxidized form ofthe Tyr160/161 (YZ). Due to the vicinity to the Mn complex, YZ

· + formationcould affect pK values of groups in the Mn ligand environment, electrostati-cally and through hydrogen-bond networks. We emphasize that several as-pects of the described model are merely hypothetical.



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by the apoprotein. The oxidative self-assembly also may repre-sent an efficient self-repair route.

4.3. Kok’s basic reaction cycle: Accumulation of oxidizingequivalents

Four stable or semistable intermediate states of the reactioncycle of the photosynthetic Mn complex can be populated bysaturating laser flashes of visible light (Scheme 13). These are

the S-states of Kok’s basic reaction cycle: S1 (dark-stable), S2, S3,and S0; the subscripts indicate the number of accumulated oxi-dizing equivalents.[130] In the dark-stable S1 state, the likely oxi-dation state of the manganese ions is {Mn2

IIIMn2IV},[114c] but con-

flicting opinions have been expressed.[131] Currently, it is mostlyassumed that each Si!Si + 1 transition involves MnIII!MnIV oxi-dation (at least up to the S3 state). For the S2!S3 transition, itis disputed whether Mn-centered[114b, c, 132] or ligand-centered[133]

oxidation takes place. Favoring the former mechanism, wearrive at the oxidation-state combination shown in Scheme 13.

This oxidation-state cycle (Scheme 13) bears an importantmechanistic consequence; before the onset of water oxidation,the catalytic metal complex accumulates oxidizing equivalentsby increasing the oxidation state of the Mn ions. There is noevidence for partial oxidation of water at an early state in thereaction cycle (no indications for OH·, terminal oxyl or peroxidecreation before S4-formation).

The structure of the Mn complex has been investigated byX-ray spectroscopy for all four semistable S-states, mostly by S-state population at room temperature followed by rapid freez-ing and XAS measurements at 10–20 K.[114b, 121b, 134] By laser-flashillumination during the X-ray experiment, data collection atroom temperature became possible. The room-temperature re-sults exclude that the structural data collected at 20 K were af-fected by a temperature-dependent redox isomerism.[114b, 135] In

an especially comprehensive investigation by Haumann et al. ,both Mn oxidation-state changes and structural changes wereaddressed for all transitions between semi-stable S-states, onthe basis of data collected at liquid helium temperatures andat room temperature.[114b] The authors concluded that the oxi-dation-state changes of the Mn complex are accompanied bydeprotonation of a bridging hydroxides in S0!S1 and by for-mation of additional di-m-oxo bridges in S2!S3 (Sche-me 14).[106a, 114b, 121b]

4.4. Redox-potential leveling

In each Si!Si + 1 transition, the Mn complex is oxidized by thesame oxidant, namely YZ

· + . The requirement of a low overpo-tential thus implies that several single-electron oxidations pro-ceed at about the same redox potential of the Mn complex. Ifthe four Mn ions were electronically isolated, we might ob-serve four essentially independent MnIII!MnIV oxidations atthe same potential. However, four isolated Mn ions are an un-likely water oxidant and, moreover, the short Mn�Mn distancesfound by EXAFS[121b, c, 136] and EPR results[116, 131b, 137] suggest thatthe four Mn ions are strongly coupled, most likely via bridgingoxygen atoms (see above). Consequently any oxidation of oneMn ion can be predicted to increase the oxidation potential forthe next single-electron oxidation of the complex significantly,as reported for numerous synthetic m-oxo bridged com-plexes.[138, 139]

For example, a first oxidation of the Mn complex will in-crease the oxidation potential for the next step by 0.5–1.5 V,[138] unless prevented by an appropriate charge-compen-sating change in the structure of the Mn complex. The redox-potential difference between YZ

· +/YZ and the Si/Si + 1 redox cou-ples is probably lower than 100 mV,[140] so an increase by 0.5 Vor more would prevent the Si + 1!Si + 2 transition, suggestingthat ‘redox-potential leveling’ by appropriate coupling of Mn

Scheme 13. S-state cycle of photosynthetic water oxidation (Kok cycle).Starting in the dark-stable S1 state, absorption of a photon causes formationof YZ

· + within less than 1 ms. Reduction of YZ· + by ET from the Mn complex

results in the Si!Si + 1 transition; typical time constants of the ET step are in-dicated. A plausible set of oxidation-state combinations of the four Mn ionsis shown. The presence of MnII in S0 and of a ligand radical in S3 (as well asdelocalized valencies) also have been proposed.

Scheme 14. Structural changes of the Mn complex of PS II. Only three of thefour Mn ions are shown because this is sufficient to illustrate the changes inthe bridging type in the course of the S-state cycle. The indicated distanceshave been determined by EXAFS spectroscopy. Reprinted from ref. [114b]with permission of the American Chemical Society.



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oxidation-state changes to charge-compensating deprotona-tions or more complex structural changes is essential for thecatalytic metal complex to function.[141] Investigations on syn-thetic complexes suggest that Mn oxidation m-OH deprotona-tion[142] or other structural changes leading to an additional m-oxo bridge[139] represent potential-lowering steps that can effi-ciently counteract the increase in the redox potential for a sub-sequent Mn oxidation step.

Scheme 15 illustrates the relation between the redox poten-tial and pKA values for oxidation of a [MnIII

2(m-O)(m-OH)] com-plex. In Scheme 15, each rectangle of states represents a ther-

modynamic cycle. For example, for the cycle made up of statesF, E, D, and C, the free-energy change for going from C to F viaD must be identical to that when going via E. The free-energychange ET resulting in oxidation of the metal complex (verticaltransitions) is given by Equation (37):

DGo � �e ðEoxm�EMeCom

m Þ ð37Þ

where Eoxm and EMeCom

m are the midpoint potential (in volts) ofthe oxidant and of the metal complex, respectively, and e isthe elementary charge. The free-energy change of the depro-tonation step (horizontal transitions) is given by Equation (38):

DGHþ

i � �0:06 eV ðpH�pK iÞ ð38Þ

Thus, for each rectangle in Scheme 15 holds:

e DEMeComm ¼ 0:06 eV DpK red=ox: ð39Þ

The difference of Epm and Ed

m (Scheme 15) was chosen to be0.6 V (e DEMeCom

m = 0.6 eV) and consequently the pK values fordeprotonation in the reduced and oxidized state differ by 10units (DpKred/ox = 10). A similar DpKred/ox was determined for aspecific [Mn2(m-O)(m-OH)] complex.[142e]

In PS II, the S0!S1 transition may involve m-OH deprotona-tion[114b, 121b, 134a] (Scheme 14) whereas the S1!S2 transition isprobably not associated with any deprotonation of the metalcomplex.[143] Thus, for the (a)-set of pK values in Scheme 15,the transition from A via C to D corresponds to the S0!S1

transition in PS II, whereas the transition from D to F modelsS1!S2. Concerted electron–proton transfer (CEPT)[144] can beruled out for this set of pK values. Only for an alternative set ofpK values (denoted as (b) in Scheme 15), the transition from Cto F could, at neutral or acidic pH, involve CEPT because thepaths via D or E are energetically strongly disfavored (uphill toD and E).

At this point, a clarifying comment on the nomenclaturemay 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 con-certed transfer of an electron and a proton by simultaneoustunneling of the two particles through energy barriers (ETPT or,preferably, CEPT).[144] The acronym PCET (proton-coupled elec-tron transfer) comprises all ET/PT, PT/ET, and CEPT. These threemodes of coupling electron and proton transfer are often ex-perimentally not easily distinguishable and theoretically notmutually exclusive. Therefore, the use of the term PCET is ap-propriate and convenient whenever a clear-cut distinctioncannot be made.

The above considerations suggest how redox-potential level-ing might be realized in PS II. The interrelation of ET and depro-tonation steps is clearly crucial. The concerted transfer of anelectron and a proton (CEPT) may be involved in redox-poten-tial leveling, but sequential oxidation and deprotonation steps(ET/PT or PT/ET) could perform the same role.

Further insights into the ways and modes of redox-potentialleveling, resulting from investigations on synthetic complexesor computational chemistry, would be desirable. The role ofthe nuclearity of the metal complex is largely unknown (typi-cally binuclear complexes have been studied). It is also still un-clear whether deprotonation of a terminally coordinated waterligand is sufficient for complete redox-potential leveling. InPS II water oxidation, m-OH deprotonation may facilitate redox-potential leveling in the S0!S1 transition; deprotonation cou-pled to formation of new m-oxo bridges may serve the samepurpose in the S2!S3 transition (see above).

A mere oxidation step results in a positive charge of themetal site, whereas all redox-potential lowering steps are asso-ciated with ‘de-charging’ typically achieved by removal of aproton. Anion binding or cation release could also facilitatecharge compensation, but are unlikely to play a role in PS IIwater oxidation. Even though redox-potential leveling clearlyrelates to charge compensation, a discussion solely in terms ofthe energetic penalty to be paid for solvation of a positivecharge in a low-dielectric protein medium would be mislead-ing. The solvation penalty may contribute, but is probably not

Scheme 15. Relation between deprotonation and successive oxidation stepsfor a hypothetical Mn complex. Two sets of pK values for the protonated m-oxo bridge are indicated, which are denoted as (a) and (b). For (a), the tran-sitions A!C!D!F model the following transitions in PS II : S0

n!S1+!

S1n!S2

+ . Concerted electron–proton transfer (CEPT; dotted gray arrow fromC to F) represents an option for the second set of pK values (b), but is irrele-vant for the first one at neutral pH (a).



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the major determinant of the redox-potential increase resultingfrom a mere oxidation step. An uncompensated one-electronoxidation will always result in an increase of the potential forthe second oxidation, irrespective of the dielectric propertiesof the medium surrounding the metal complex.

4.5. ‘Smart’ proton removal

Concerted transfer of an electron and a proton has been dis-cussed in the context of redox-potential leveling.[144] Oxidationof the Mn complex by ET to YZ

· + is most likely directly coupledto the shift of a proton from His190 to Tyr160 (YZ ; Figure 16) andpossibly also to a similar shift of a second proton within thecatalytic site.[145] However, a discussion in terms of a concertedmultisite transfer of one electron and two protons is insuffi-cient in providing a complete picture of the relation betweenelectron and proton transfer in biological water oxidation. TheMn complex of PS II is embedded deeply in the PS II protein.Long-distance inner-protein proton transfer typically requirestens or hundreds of microseconds. Consequently, ET by tunnel-ing of the electron from the Mn complex to YZ

· + and theproton translocation from the Mn complex to the aqueousphase can not occur concertedly, but need to proceed sequen-tially. For the transition S3!S4!S0+O2, time-resolved X-ray ab-sorption spectroscopy[135d] and complementary fin-dings[130b, 143, 146] have provided evidence for the following se-quence of events: Absorption of a photon results in formationof YZ

· + within less than 1 ms. The electric field emanating fromthe positive charge at YZ

· + affects pK values and drives severalproton hopping steps and eventually results in a transfer of aproton from the Mn complex to the aqueous phase, withinabout 200 ms (S3

+!S3n ; S3

n also has been denoted as S4[135d]).

The potential-lowering deprotonation facilitates the subse-quent ET to YZ

· + (S3n!S4

n ; S4n previously denoted as S4’

[135d])followed by rapid onset of water oxidation. The reaction se-quence is concluded by the release of a second proton.

The discovery of a proton-first ET in the S3!S4!S0 transi-tion led us to an extended working hypothesis on the se-quence of electron and proton transfer steps (Scheme 16).[147]

We emphasize that the proton removal steps relate to the mul-tistep transfer from the catalytic site to the aqueous phase,which can not be coupled directly to the ET step. However, itis conceivable, for example in the S2

+!S2n transition, that a

proton acceptor B� is ‘prepared,’ which renders possible theconcerted transfer of an electron to YZ

· + and the shift of aproton to B� .

Concerted electron–proton transfer may also be favorable inwater oxidation by artificial catalysts. However, oxidation ofthe metal complex directly coupled to a proton transfer to awater molecule is an unlikely event in the acidic and neutralpH regime because a water molecule is an unfavorable protonacceptor (formal pK of H2O/H3O couple is about �1.7). This im-plies that, also in nonbiological water oxidation, ‘smart’ remov-al of protons from the catalytic site may be an issue whenaiming at fast and efficient water oxidation.

4.7. Water deprotonation and oxidation: Four basic routes

Three crucial and interrelated aspects of water oxidation haveso far been discussed herein: 1) accumulation of oxidizingequivalents by transition metal oxidation; 2) redox-potentialleveling; 3) ‘smart’ proton removal prior to the onset of the O�O bond-formation step. All three points relate to deprotona-tion at the catalytic site. In this section, we address the relationof the four deprotonation steps to the process of forming di-oxygen from the two molecules which often are denoted as‘substrate water’. Four prototypical routes towards O�O bondformation are shown in Scheme 17.

Route A (Scheme 17) involves partial water oxidation by per-oxide formation at an early stage, that is, after accumulation ofthree oxidizing equivalents in the S3 state. This route was dis-cussed with regards to water oxidation at mono-nuclear metalcomplexes (see sections 3.2 and 3.3) and is also considered tobe of high relevance in electrochemical water oxidation [Equa-tions (25)–(28) in Section 1.3] . Early peroxide formation hasalso been discussed in the context of PS II water oxida-tion.[145, 148] However, peroxide formation in the S3 state of thePS II manganese complex, in conflict with the results obtainedby X-ray absorption spectroscopy,[114b, 121b, 132, 133b] implies Mn re-duction in the S2!S3 transition:

fMn4III,IV,IV,IVðO2�Þ2gþYZ

�þ ! fMn4III,III,IV,IVðO1�Þ2gþYZ ð40Þ

Scheme 16. Extended S-state cycle model. Kok’s classical S-state cyclemodel[130a] is extended by including not only 4 oxidation but also 4 deproto-nation steps.[130b, 147] Four of the nine intermediates are stable for tens of sec-onds or fully dark-stable (S1

n, S2+ , S3

+ , and S0n). They correspond to the

states S1,, S2, S3, and S0 of Kok’s reaction cycle. Electrons and protons are re-moved alternately from a Mn complex which comprises the {Mn4Ca(m-O)n}core, the protein environment and bound water molecules. We emphasizethat proton removal from the catalytic site does not necessarily imply depro-tonation of water species. The energetics of the oxidation and deprotona-tion steps are described by Equations (37) and (38), respectively. The nine in-termediate states of the Mn complex are denoted as Si

+ /n, where the sub-script gives the number of accumulated oxidizing equivalents and the super-script indicates the relative charge: positive (+) or neutral (n) relative to thedark-stable S1 state.[106a] The shown reaction cycle has been discussed pre-viously[130b, 147] using a nomenclature where the nine states from S0

+ to S4+

were numbered consecutively and denoted as I0–8.



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In route B (Scheme 17), water oxidation is facilitated by aregular pattern of water-deprotonation steps, eventually result-ing in two oxides terminally coordinated to two high-valentmetal ions. The oxo group may have radical character and thusmight be better described as an oxyl radical (e.g. MnVO!MnIVO·). Route B is appealingly straightforward and has beeninvoked frequently in mechanistic models of water oxidationat electrode surfaces by metal complexes in solution (homoge-neous catalysis) and also in PS II. Also, the influential ‘hydrogenatom abstraction’ model of PS II water oxidation, suggested byBabcock and co-workers,[149] involves route B. However, termi-nal oxo formation is predicted to strongly enhance the so-called pre-edge peak in the X-ray absorption spectra,[150] butthis is not observed.[114b, 132b, 133b] Consequently, route B is unlike-ly to be followed by PS II. The unique aspect of Babcock’s hy-pothesis is the concerted transfer of one electron (Mn oxida-tion) and one proton (water deprotonation) to YZ

· amountingto a hydrogen atom transfer in each Si!Si + 1 transition. In thiscase, the oxidized tyrosine residue is not a radical quasi-cation(not YZ

· + with the phenolic proton staying in close to the phe-nolic oxygen), unlike in Equation (35). According to Babcock’shypotheses, the phenolic proton is released to the bulk phasebefore hydrogen atom transfer can take place, transformingthe initial radical cation into a neutral radical. This hydrogenatom abstraction by YZ

· is in serious conflict with structural[99a, b]

and spectroscopic data.[143, 150–151]

In route C (Scheme 17), water oxidation involves bound orfree hydroxides that act as proton acceptors in the O�O bond-

formation step. It relates to the‘alkaline mechanism’ in the olderelectrochemical literature, incontrast to the ‘acidic mecha-nism’ of Route B, and typically isnot considered a relevant path-way in the context of photosyn-thetic water oxidation.

In route D (Scheme 17), accu-mulation of oxidizing equiva-lents is not coupled to sub-strate-water deprotonation. In-stead the catalyst is deprotonat-ed, for example, by deprotona-tion of bridging oxides. Thereby‘acceptor bases’ are created,which facilitate a direct couplingof water oxidation and water de-protonation in the O�O bond-formation step. The acceptor-base hypothesis was put forwardto explain water oxidation inPS II.[106a, 121b, 141] A combination ofearly water deprotonation(route B) and acceptor base ac-cumulation (route D) has alsobeen considered.[117, 121b] A moredetailed discussion is providedbelow.

Experimental results relating to the mode of substrate-watercoordination and deprotonation are rare. Presently, the mostimportant results may be the water exchange rates determinedin elegant H2O16/H2O18 exchange experiments,[152] which haverevealed the presence of two different binding sites for thetwo substrate-water molecules with distinguishable exchangerates. Typical results for the S-state dependence of the ex-change rates are summarized in Table 1.[152f]

The exchange rates of water bound to a pentanuclear metalcomplex buried inside a protein are not well understood,[152b, f]

so any seemingly straightforward interpretation might also bemisleading. Therefore, the following discussion is tentativeonly; some aspects of the results in Table 1 have been inter-preted differently.[152b, f]

To relate the results summarized in Table 1 to mechanisticproposals on photosynthetic water oxidation, we will use thefollowing guidelines:

Scheme 17. Four routes towards water oxidation. The catalyst is symbolized by the grey shape. In Si, the numberof accumulated oxidizing equivalents always is equal to i. However, accumulation of oxidizing equivalents (by oxi-dation of metal ions) is explicitly indicated only in D in form of the encircled ‘+’ symbols. A–D differ in the mini-mal number of water and proton binding sites (n).

Table 1. Rate constants for exchange of substrate-water molecules (de-tected for thylakoid membranes from spinach, T = 10 8C).[152f] The kfast de-crease in S2!S3, by a factor of 3, corresponds to marginal changes inbinding energy.

kslow/fast of water exchange [s�1] S0 S1 S2 S3

Slow HnOsubstrate ca. 10 ca. 0.02 ca. 2 ca. 2Fast HnOsubstrate – >120 ca. 120 ca. 40



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G1) Oxo groups in bridging position between high-valentMn ions exchange extremely slowly (<10�2 s�1 in {Mn2

III,IV} and<10�6 s�1 in {Mn2

IV,IV}).[153]

G2) Deprotonation of any water bound to a metal ionshould increase the binding energy and thus decrease its ex-change rate strongly.

G3) Partial water oxidation by peroxide or radical formation(O·) should affect the substrate water exchange rate.

G4) Oxidation of Mn(i)III!IV should increase the bindingstrength of any water species coordinated to Mn(i) and thusdecrease the exchange rate strongly (i.e. by several orders ofmagnitude).

The decrease in kslow in the S0!S1 transition might indicatewater deprotonation, but the subsequent acceleration is in-compatible with persistent deprotonation of the substratewater molecule. In explaining these variations in kslow, anysimple model will fail. A complex and largely hypothetical ex-planation follows (Scheme 18): The deprotonation of a m-oxo

bridge in the S0!S1 transition[106a, 121b, 134a] facilitates the forma-tion of a new hydrogen bond, thereby binding the substratewater molecule more strongly. The MnIII!IV oxidation in theS1!S2 transition is not coupled a deprotonation, so the chargeof the metal complex will increase. Consequently, the pK valueof a bridging oxide connected to the oxidized Mn ion willchange strongly[142e] (Scheme 15) and render the m-oxygen aweak hydrogen-bond acceptor, thereby causing an increase inthe exchange rate of the water molecule.

In conjunction with the results summarized in Table 1, G1implies that the substrate water molecules are not in a bridg-ing position between Mn ions. G2 is in conflict with deproto-nation of the slowly exchanging water in S0!S1 or S1!S2, andexcludes deprotonation in S2!S3 for both HnOsubstrate groups.G3 provides an argument against peroxide or O·substrate forma-tion in the transitions for which water deprotonation is exclud-ed by G2. G4 suggests that substrate water molecules are notbound to the Mn ions that are oxidized in S2!S3.

Assuming that a bridging hydroxide is deprotonated in S0!S1 (Scheme 14),[114b, 121b, 134a] G1 (with the results summarized inTable 1) implies that, rather than substrate-water deprotona-tion, an unprotonated m-oxo group is formed that can serve asa proton acceptor in S4!S0+O2 (formation of an acceptorbase). Similarly, in S2!S3 another acceptor base is likelyformed, whereas in the S1!S2 transition there is no deprotona-tion event (Scheme 16). The available data is compatible withthe following notion: In the three Mn-oxidizing S-state transi-tions (from S0 to S3) no substrate water is deprotonated butrather two bases are created that can accept a proton fromwater in the process of O�O bond formation. The employmentof several bases as proton acceptors in the O�O bond forma-tion reaction could be crucial (acceptor-base hypothesis;route D in Scheme 17).

Most likely, OH� ions cannot diffuse within the PS II proteinto the catalytic site. However, H2O molecules certainly move tothe catalytic site effectively. A water molecule may bind to theMn complex immediately after completion of the O2-formationstep to a vacant coordination site. The pK of the bound waterchanges drastically resulting in deprotonation and formationof a bound hydroxide (Mn�OH2!Mn�OH+H+). Transfer of theproton to the aqueous bulk phase completes the process. Thenet effect of this hypothetical sequence of events would behydroxide binding to the Mn complex. Such an indirect hy-droxide-binding step has not been supported by experimentalfindings, but cannot be excluded for the S0

+!S0n transition

(Scheme 16). Similarly, it represents an open question whetherthe proton removed from the Mn complex in the S3

+!S3n

transition stems from a substrate-water deprotonation.

4.8. Mechanism of O�O bond formation

Numerous hypotheses on dioxygen formation in PS II havebeen put forward; see ref. [154] for a summary of 34 early sug-gestions and ref. [106b] for a discussion of more recent hy-potheses. It is mostly assumed that, in PS II, O2 formation pro-ceeds via transient formation of a peroxidic intermediate, pos-sibly in a two-electron step.[155] Scheme 19 shows prominenthypotheses on the mode of peroxide formation in the S4!S0

transition of PS II.Since the late 1970s, peroxide formation between two bridg-

ing oxygen atoms (Scheme 19 A) has been suggested in nu-merous studies.[156] Analogous chemistry has been demonstrat-ed for bis(m-oxo)dicopper(III) complexes and their transforma-tion to the (m-h2 :h2-peroxo)dicopper(II) form.[157] Oxygen radi-cals bridging between Mn ions have been invoked.[121a] Mecha-nistic models centered around a cuboid/butterflytransformation involve O�O bond formation between oxygenatoms bridging between Mn ions.[158] This bond-formationmode is considered only rarely in the context of photosynthet-ic water oxidation, as, among other reasons, it is in conflictwith G1 and also difficult to reconcile with G2 and G4.

The nucleophilic attack of a hydroxide upon a terminaloxide (MnV=O) or oxyl radical (MnIV�O·; Scheme 19 B), has beenproposed repeatedly and is supported by molecular models(see sections 3.2 and 3.3).[106b] The attacking hydroxide may be

Scheme 18. Hypothesis on the structural changes underlying the peculiarvariations in the water exchange rate of the slowly exchanging water mole-cule (H2Os). The m-hydroxo deprotonation in the S0

n!S1n transition facilitates

formation of an additional hydrogen bond; ‘charging’ of the complex by oxi-dation, in S1

n!S2+ , weakens this H-bond.



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a ligand bound to Mn or to Ca, or an outer-sphere water. Inref. [123d] (Scheme 19 C), it is suggested that the slowly ex-changing water is terminally coordinated to the Ca ion of a{CaMn3(m-O)4} cube, whereas the terminal O· is formed in theS3!S4 transition by deprotonation of a hydroxide, terminallycoordinated to Mn4 in Figure 16. In this and a relatedmodel,[117] the Arg357 of the CP43-protein (not shown inFigure 16) is crucial in ‘managing’ the proton removal from thecatalytic site.

Scheme 19 D illustrates a specific proposal involving anouter-sphere water species attacking a terminal oxyl,[121b] whichhas been formed in the S3!S4 transition by oxidation and de-protonation of a terminally coordinated water species (notshown). A unique aspect of this proposal is the coupling ofperoxide formation to the transfer of a proton to a bridgingoxide, possibly in form of a concerted ‘removal’ of one electronand one proton from the outer-sphere water (CEPT), which for-mally amounts to abstraction of a hydrogen atom. Specifically,the inner-sphere electron transfer from oxygen to a MnIV ion iscoupled to proton transfer to its m-O ligand (Scheme 19 D).Owing to the relation between the pK value of the m-oxogroup and the oxidation potential of the metal complex(Scheme 15), a concerted mechanism could be especially ad-vantageous in facilitating the critical peroxide-formationstep.[106a, 121b, 141, 155, 159]

Recent DFT calculations resulted in especially low energeticbarriers in O�O bond formation for a mechanistic model with

an unconventional mode of peroxide formation (Sche-me 19 E).[123b] In this model, the S2!S3 transition involves a ‘re-construction’ of the complex, which is associated with bindingof the second substrate-water molecule. It is unclear whetherthis and other aspects of the model can be reconciled with thewater-exchange rates (Table 1).

As early as the late 1980s, Krishtalik argued that a two-elec-tron step resulting in peroxide formation and possibly coupledto the transfer of a proton to an acceptor base that is part ofthe Mn complex represents a favorable path towards dioxygenformation.[155, 159] It presumably relates to conceptual problemsthat such two-electron/n-proton transitions are not typicallyconsidered in explicit mechanistic models of water oxidation.Further development in the theory of such multiparticle reac-tions may be required before complex multielectron/multipro-ton steps can be addressed in computational studies.

The new mechanistic proposal in Scheme 20 is inspired bythe hydrogen-bonding pattern found in a dimeric unit of Rucomplexes (Figure 13).[80a, 81] The starting points for peroxide

formation are two fully protonated water molecules. Togetherwith Mna, these molecules form a motif (H2O�H�OH�Mn)which resembles Zundel’s hydronium ion (H2O-H-OH2). In thisquasi-Zundel ion, the O�O distance may be as small as 2.4 �,so that the transition state for formation of a peroxidic O�Obond is reached more readily. The bond formation process isassumed to involve the transfer of HB2 and two electrons, re-sulting in m-O protonation and Mn reduction (Scheme 20), andcan thus be viewed as a hydride abstraction by the Mn com-plex. Moreover, this O�O bond formation by hydride abstrac-tion is assumed to be either directly coupled (concerted pro-cess), or preceded or followed by the transfer of HB1 and HB3.The subsequent transfer of HB4 is required for completion of O2

formation, but is not indicated in Scheme 20. This proposedmechanism is speculative and is presented to illustrate twopoints:

a) Two fully protonated water molecules could represent thestarting point for O�O bond formation.

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 stateformed in the S3!S4 transition. This state is formed by oxidation of the Mncomplex and possibly also by the relocation of 1 or 2 protons to 1 or 2bases (not shown). The lower structure shows the peroxidic state, which isassumed to form transiently. For peroxide formation between two oxo oroxyl groups, each terminally coordinated to a MnV/IV ion, see route B inScheme 17.

Scheme 20. Hypothesis on the coupling of electron transfer and shifts ofprotons in hydrogen-bond networks for peroxide formation in the S4!S0

transition. (s and f on oxygen represents slow and fast, respectively)



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b) The relocation of several protons from water to acceptorbases likely involves a specific hydrogen-bond network; thecoupling of proton and electron transfer in the O�O bond maybe crucial in avoiding high energetic barriers.

We note that the strong coupling between the pK of m-oxobridges and the midpoint potential for reduction of the metalions (Scheme 15) can render unprotonated oxo groups suitableacceptor bases in a multielectron/proton transfer reaction.

The O�O bond formation chemistry remains insufficientlyunderstood. We tentatively conclude that the {Mn4Ca} complexnot only positions the substrate-water molecules by hydrogenbonding and terminal ligation to a metal ion, but also partici-pates actively in formation of the O�O bond by alteration ofits metal–oxo core, specifically by bridging-type changes. TheO�O bond formation step itself may involve the coupled trans-fer of electrons and protons from water to the catalysts.

5. Common Themes and Unifying Concepts

5.1. Nuclearity and structure of the catalyst: obligatorym-oxo-bridging?

The structure of the photosynthetic manganese complex ischaracterized by m-oxo bridges between metal ions. Hereafter,we do not discriminate between m-O, m-OH, and m-OH2 bridg-ing. The putative mechanistic role of m-oxo bridges in redox-potential leveling and as proton acceptors in O�O bond forma-tion is discussed in section 4. In electrochemical water oxida-tion at metal electrodes, all metal surfaces are transformedinto metal oxides at the positive potentials needed for wateroxidation; it is likely that extensive m-oxo bridging is prevalent.How deep the oxide layer penetrates the metal electrode islargely unknown and probably depends on both the metalused and the operation conditions. For mineral, amorphous ormicrocrystalline metal oxides, the m-oxo bridging involves notonly the surface atoms but also the bulk material (see sec-tion 2). On these grounds, we conclude that bridging oxidesare a candidate for a common structural motif in all water oxi-dation catalysts.

However, there are a growing number of mononuclearwater-oxidation catalysts that seemingly facilitate water oxida-tion in the absence of any m-oxo bridging (see section 3). Themere existence of these functional mononuclear catalysts indi-cates that m-oxo bridging is not an essential prerequisite forwater oxidation. However, it should not be overlooked that theturnover frequencies of these complexes are generally low andthat the redox-potential leveling is inadequate. It remains pos-sible that efficient water oxidation generally requires m-oxobridging between metal ions.

5.2. Self-assembly and self-repair of catalytic metal centers

The proteins of PS II represent an elaborate ligand system. Theassembly of the catalytic complex remains insufficiently under-stood but probably involves oxidation of Mn2 + ions and forma-tion of {Mn4Ca(m-O)n} within the framework of the protein li-gands. Conceptually, this oxidative self-assembly does not

differ from synthetic paths that have been followed for thepreparation of molecular metal complexes (although the lattercan also be achieved employing high oxidation state precur-sors already).

The oxidation of metal surfaces and the formation of bulkmetal oxides proceed in the absence of any elaborate ligandsystem. The oxide formation can be viewed as an oxidativeself-assembly and implies a self-repair mechanism (see sec-tion 2.1). In PS II, an analogous self-repair could occur but isobscured by the intricate and costly exchange of the proteinsubunit that provides most ligands to the {Mn4Ca} complex(see section 4.2). In homogeneous catalysis, loss and rebindingof metal ions to the ligand system might occur. However, pres-ently this question is of little relevance because, most likely, ox-idative modification/destruction of the ligand system limits themaximal turnover number much more severely than loss ofmetal ions.

5.3. Potential-determining step: Redox-potential leveling

The thermochemical analysis of electrocatalytic water splittingon solid surfaces reveals that an ideal reversible catalyst shouldexhibit four proton-coupled electron transfer steps with identi-cal DGi for each of the four steps (see section 1.3) so that thesame minimal electrical potential suffices to drive each of theelementary steps. In real catalysts, deviations from the equalDGi distribution result in a potential-determining step; the re-action with the highest DGi determines the overpotential.

Each of the four discussed steps [Equations (25)–(28)] repre-sents a deprotonation-coupled oxidation, either of the activesite at the electrode surface or of a water species itself. Identi-cal DGi in an ideal electrocatalyst implies that the values of theindividual (electrical) reversal potentials, Vrev(i), are identical forthe four individual oxidation steps, that is, Vrev(1) = Vrev(2) =

Vrev(3) = Vrev(4). The Vrev(i) in electrocatalysis corresponds to themidpoint redox potentials for four sequential single-electronoxidations of a molecular catalyst. The problem to find an elec-trocatalyst with approximately the same Vrev(i) for each oxida-tion step is thus analogous to the problem of redox-potentialleveling in homogeneous and biological catalysis, (see sec-tion 4.4).

The redox-potential leveling, that is the maintenance of anapproximately constant redox potential for four sequential oxi-dation steps, is a chemically demanding task that, to ourknowledge, has not yet been achieved for any synthetic mole-cule. However, two sequential oxidations steps at the same po-tential have been realized as discussed for binuclear Mn com-plexes (see section 4.4). In PS II, formation of unprotonated m-oxo bridges appears to be crucial (Scheme 14). Amorphousmetal oxide catalysts may carry flexible m-hydroxo bridges withsuitable pK values, related to the PZC, which could facilitatethe redox-potential leveling by coupling of m-OH deprotona-tion to the oxidation step, possibly in close analogy to process-es in PS II. This hypothesis may explain why amorphous oxidesare OER-active at relatively low electrochemical overpotentials.Experiments addressing changes in m-oxo bridging for surface



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and bulk oxides represent a major challenge but the resultscould be of exceptional importance.

In conclusion, redox-potential leveling in homogeneous andbiological water oxidation, on the one hand, and the require-ment for the absence of a potential-determining step in anideal electrocatalyst, on the other hand, are different facets ofthe same problem. We propose that the quality of redox-po-tential leveling generally determines the maximal energetic ef-ficiency of a catalyst. Mode and quality of redox-potential lev-eling may also determine the kinetic competence of the cata-lyst, which is quantified in terms of rate constants or turnoverfrequencies. Notably in this context, in biological water oxida-tion the O�O bond formation step itself is not rate-determin-ing but rather it is the oxidation of the catalyst in the individu-al Si!Si + 1 transitions. For each Si!Si + 1 transitions, the drivingforce of the ET step (�DGi of oxidation of the Mn complex) willbe maximal for optimal redox-potential leveling. A greater driv-ing force typically implies an increase in the respective rateconstant. Similarly in electrocatalysis, the concept of the poten-tial-determining steps involves, at least implicitly, that this stepis also the overpotential-determining and thus the overall rate-determining step. We conclude that the redox-potential level-ing determines generally the maximal energetic efficiency andis also a key determinant of the kinetic competence of eachwater oxidation catalyst. A generic mode of optimal redox-po-tential leveling may involve coupling of metal oxidation toligand deprotonation and changes of the oxo bridging be-tween metal ions (see section 4). This conjecture, however, isstill at the level of a working hypothesis.

5.4. Mechanism of O�O bond formation

The requirements for O�O bond formation have been dis-cussed at the level of molecular orbitals repeatedly, either inform of plausibility considerations or based on quantum chem-ical calculations. Instead of attempting a discussion of electron-ic aspects, we focus on basic considerations and reactionmodels.

Routes A, B, and C in Scheme 17 represent the prominenttypes of mechanisms that have been considered in the field ofelectrochemical water oxidation. In route B, proton-coupledelectron transfer results in accumulation of four oxidizingequivalents by the catalyst and formation of two terminalmetal–oxo (Me=O) or metal–oxyl (Me�O·) groups at neighbor-ing sites before O�O bond formation takes place. Route B hasindeed been considered in all of the three fields; electrochemi-cal, homogeneous, and biological catalysis of water oxidation.Route B has never been proven experimentally in electrochem-ical water oxidation and essentially has been disproven for bio-logical water oxidation, whereas recent O16/O18 isotope experi-ments support its prevalence for a binuclear ruthenium com-plex.[74a] We conclude that route B is appealingly simple andplausible, but only rarely supported experimentally.

There is also an appealingly simple aspect to route A; bind-ing, deprotonation and oxidation of water take place at asingle metal site. The existence of mononuclear catalystsproves that a mono-site mechanism can be operative in water

oxidation (see section 3) and it has recently also been favoredfor electrocatalytic water oxidation (see section 1.3). It shouldbe noted that route A does not involve accumulation of fouroxidation equivalents before the onset of O�O bond forma-tion—as opposed to routes B, C, and D—but rather of onlytwo oxidation equivalents, as suggested for some reactioncycles of mononuclear complexes (see section 3). Therefore,route A can be ruled out as the main path for water oxidationin photosynthesis.

To our knowledge, route D has only been considered forwater oxidation at the {Mn4Ca} complex of photosynthesis. Thepossibility of concerted inner-complex electron–proton move-ments in the O�O bond formation step represents a uniqueaspect of route D (see section 4.8). In the field of photosynthet-ic water oxidation, it represents a working hypothesis present-ed before[106a] and detailed herein (section 4). Route D lacks thesimplicity of routes A and B because not only water-bindingsites but also four acceptor bases are proposed to be part ofthe catalyst, thereby ruling out route D for mononuclear cata-lytic sites.

The investigation of mononuclear complexes revealed thefeasibility of a hydroxide/water attack at terminal oxo/oxyl li-gands, which has recently also been favored for electrochemi-cal water oxidation (see section 1.3) and for the biological par-agon (Scheme 19 B–D). In all fields it has remained difficult toobtain experimental proof, but specifically the electrophilicattack of hydroxide at a terminal oxyl appears to be the mostlyfavored mode of O�O bond formation.

When weighing the evidence in favor and against a distinctmechanism, none of the mechanistic proposals is supportedexperimentally up to a level that the term ‘proof’ would be ap-plicable. Simplicity and seeming plausibility still have to substi-tute for the lack of experimental evidence, in spite of tremen-dous efforts and recent progress by experimentalists and theo-retical chemists. The largest body of detailed experimental re-sults has been obtained for photosynthetic water oxidation.For these reasons, route A, B, and C can essentially be ruledout for the biocatalytic reaction. Whether this conclusion is ofhigh relevance for electrochemical or homogeneous water oxi-dation is still unclear.

5.5. Concluding Remarks

We have reviewed and compared recent findings on water oxi-dation in three research fields: heterogeneous (electro)catalysisby surface and bulk oxides (sections 1 and 2), homogeneouscatalysis by metal complexes (section 3), and biocatalytic wateroxidation in photosynthesis (section 4). Each field has devel-oped its own terminology, methods, concepts, and, conse-quently, its own strengths.

Heterogeneous electrocatalytic water oxidation

Stringent approaches towards the energetics and efficienciesof water splitting have been developed only for electrochemi-cal water-splitting (section 1.1). Traditionally the ‘macroscopicenergetics’ were considered with a focus on system parame-



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ters of relevance in chemical engineering. At the atomic level,however, electrolysis is not well understood. Traditional mecha-nistic proposals [Equations (12)–(21), routes A–C in Scheme 17]may be plausible but largely are supported only circumstantial-ly by macroscopic kinetic parameters, such as the experimentalTafel slope (see section 1.2).

A major step forward in our understanding of qualitativetrends in electrocatalytic activity has resulted from thermo-chemical analysis of so-called volcano plots relating thermody-namic overpotential or catalytic activity with microscopic pa-rameters accessible through modern quantum-chemical calcu-lations. Thereby, suboptimal scaling and chemisorption ener-gies are currently found to limit the achievable catalyst activity(section 1.3). This work may represent a first step towards anatomic-scale rational design of electrocatalysts. However, evenfor the basic reaction scheme of the reviewed computationalapproach [Equation (25)–(28), route A in Scheme 17], the sup-port by experimental findings is, at present, insufficient. Wealso lack strategies and methods to independently manipulateundesired scaling relations and chemisorption energies. Com-bination of electrochemical methods with structure-sensitivespectroscopies represents a promising avenue towards under-standing at the atomic level. This endeavor is seriously compli-cated by the active-site problem of heterogeneous catalysis ;neither the effective nuclearity (monosite versus multisitemechanisms in Scheme 17) nor the local structure of the cata-lyst at the active site are well known, the latter because devia-tions from crystalline order (defects) probably play a crucialrole.

In the past, electrochemical water oxidation was mostly dis-cussed in physical terms, such as adsorption, discharge, orFermi level, which tend to obscure the chemical character ofthe process. However, at the atomic level, equivalent reactionschemes and considerations are applicable in electrocatalytic,homogeneous, and biological water oxidation (section 5.4).The conceptual gap between electrocatalytic water oxidation,on the one hand, and homogeneous or biocatalytic water oxi-dation, on the other hand, can often be bridged by a meretranslation, for example, substitution of ‘chemisorption’ by theterm ‘binding’ or ‘ligation.’ In other cases, deeper conceptualquestions may arise. For example, hydroxide binding (or chem-isorption) to one of the many metal ions of a surface oxidemay change the oxidation state of this specific metal ion ac-cording to Equation (41):

MnþþOH� ! Mðnþ1Þþ ðOHÞð�Þþe� ð41Þ

This point of view, with roots in the concept of formal oxida-tion states (or ‘valencies’) in coordination chemistry, is difficultto reconcile with the band-gap model often used in solid-statechemistry and physics. The difference in concepts—distinct lo-calized oxidation-state changes versus modification of theband structure close to the surface—is clearly related to thediscussion in coordination and bioinorganic chemistry on local-ized versus delocalized valencies (or oxidation states). The de-velopment of a unifying concept or theory that is also applica-

ble to the noncrystalline bulk metal oxides described in sec-tion 2, represents a task for the future.

Homogeneous catalysis by transition metal complexes

Metal complexes synthesized as homogeneous water-oxidationcatalysts are typically characterized crystallographically, so thattheir structure is known at atomic resolution. Moreover syn-thetic variations of the metal complex facilitate investigationof structure–function relations. These two experimental advan-tages at hand, synthetic chemistry has established the founda-tions for understanding crucial mechanistic aspects in wateroxidation, such as the need for and modes of redox-potentialleveling (see section 4.4).

Research in recent years has revealed that dimeric or higher-order assemblies of multiple metal centers are not essential forwater-oxidation catalysis ; mononuclear systems are sufficientand this finding represents a major breakthrough (section 3).The easier derivatization associated with structurally simplermononuclear systems allows for a more rational synthesis oftailored systems, structure–function analyses, and mechanisticstudies. Future studies should evaluate more model systemswith respect to the redox-potential leveling and also includecomplexes with nuclearities higher than two. Although itseems clear that such systems will never reach the efficiency ofnature, they might be designed for satisfactory artificial devi-ces, once ligand decomposition is under control. Immobiliza-tion of molecular complexes on surfaces may increase drasti-cally the catalytic performance and lifetime,[160] possibly be-cause bimolecular decomposition routes are suppressed. Stabi-lizing surfaces or matrices in conjunction with self-assemblyand self-repair of the catalyst will likely be of high importancein the development of technological catalysts based on mono-or multinuclear transition metal complexes.

Biocatalytic water oxidation

The {Mn4Ca} complex of PS II can be driven synchronouslythrough its reaction cycle (S-state cycle) by four light flashes,thus rendering time-resolved experiments on the four S-statetransitions possible (determination of reaction sequences andrate constants). Moreover, the four reaction intermediates canbe stabilized by freezing and studied conveniently in sophisti-cated low-temperature experiments, specifically advanced EPRspectroscopy and X-ray methods. These exceptional experi-mental advantages, in conjunction with high biological impor-tance, have motivated a large number of investigators, specifi-cally spectroscopists, to focus their research on elucidation ofthe mechanism of photosynthetic water oxidation. The finalgoal of this truly interdisciplinary endeavor has not yet beenreached. Nonetheless, an exceptionally deep understanding ofmechanistic details and pertinent questions has been obtained.Many key concepts of likely relevance also for nonbiologicalwater oxidation have been established, namely accumulationof four oxidizing equivalents, redox-potential leveling, and cou-pling of electron transfer to deprotonation of the catalyst andremoval of protons from the catalytic site. Specifically the



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latter aspect, that is the transfer of protons from the catalystto the bulk medium, has been almost completely neglected inresearch on nonbiological water oxidation, but could be crucialin the developments of catalysts with satisfactory turnover fre-quencies. In conclusion, research on photosynthetic water oxi-dation will continue to be a source of inspiration, concepts,and experimental technologies for investigations on nonbio-logical catalysts of water oxidation.

We believe that, in the strive for knowledge-based develop-ment of efficient and eventually technology-relevant catalystsof water oxidation, the fields of heterogeneous, homogeneous,and biological catalysis will grow increasingly together. Unify-ing concepts will not only play a role on a theoretical, concep-tual level—as discussed herein—but also in terms of new tech-nological systems that combine approaches from homogene-ous and heterogeneous catalysis with surface and nanoscien-ces, for example in terms of immobilized molecular systems,bulk metal oxides in nanostructured materials, or purely inor-ganic catalysts with molecular properties on surfaces.[62b]

Acknowledgements

Joint scientific work in the framework provided by the UniCatcluster of excellence (Unifying Concepts in Catalysis, Berlin) hasinspired this Review article. Financial support by the UniCat clus-ter is gratefully acknowledged. We also thank the EuropeanUnion (7th framework program, SOLAR-H2, #516510) and theGerman Federal Ministry of Education and Research (BMBF, H2Design Cell, #03SF0355D) for financial support (to HD). PS grate-fully acknowledges financial assistance by the BMBF (grant033R018A-G) in meeting the publication costs of this article. Thepresented EXAFS data (Figure 6) was collected in cooperationwith Dr. F. Sch�fers at beamline KMC1 of the Berlin synchrotron(BESSY, Helmholtz Zentrum Berlin) ; F. Ringleb, Dr. I. Zaharieva, Dr.P. Chernev, and N. Leidel contributed to sample characterization(MR) and data collection. Discussions (of PS) with M. Koper(Leiden, The Netherlands), J. Norskov, and J. Rossmeisl (bothLyngby, Denmark) are gratefully acknowledged.

Keywords: oxygen evolution · photosynthesis · solar fuels ·transition metals · water splitting

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