Efficient water oxidation with organometallic iridium complexes as precatalysts.

PCCP View Article Online

Published on 30 January 2014. Downloaded by Universidad de Oviedo on 30/10/2014 09:18:51.

PAPER

Cite this: Phys. Chem. Chem. Phys., 2014, 16, 11976

View Journal | View Issue

Efficient water oxidation with organometallic iridium complexes as precatalysts† Anna Lewandowska-Andralojc,‡a Dmitry E. Polyansky,a Chiu-Hui Wang,a Wan-Hui Wang,bc Yuichiro Himedabc and Etsuko Fujita*a Catalytic water oxidation has been investigated using five iridium complexes as precatalysts and NaIO4 as an oxidant at various pH conditions. An increase in the activity of all complexes was observed with increasing pH. A detailed analysis of spectroscopic data together with O2-evolution experiments using Cp*Ir(6,6 0 -dihydroxy-2,2 0 -bipyridine)(OH2)2+ as a precatalyst indicate that the high catalytic activity is closely connected with transient species (A) that exhibits an absorption band at lmax 590 nm. The formation of this active form is strongly dependent on reaction conditions, and the species was distinctly observed using a small excess of periodate. However, another species absorbing at 600 nm (B), which seems to be a

Received 3rd December 2013, Accepted 30th January 2014 DOI: 10.1039/c3cp55101f

less active catalyst, was also observed and was more prominent at high oxidant concentration. Dynamic light scattering analysis and transmission electron microscopy have identified species B as 120 nm nanoparticles. The ultrafiltration method has revealed that species A can be attributed to particles with size in the range of 0.5–2 nm, possibly small IrOx clusters similar to those described previously

www.rsc.org/pccp

by Harriman and co-workers (J. Phys. Chem., 1991, 95, 616–621).

Introduction Nature utilizes and stores solar energy in the photosynthetic process which converts water and carbon dioxide to oxygen and carbohydrates. In photosynthesis, oxidation of H2O to O2 is the key reaction, producing electrons and protons which are ultimately used to reduce carbon dioxide and produce biomass. In pursuit of clean and sustainable energy sources, increasing efforts are being made to achieve efficient solar energy conversion into fuels utilizing lightdriven water splitting. The overall water splitting process consists of two half-reactions: (i) proton reduction (4H+ + 4e - 2H2) and (ii) water oxidation (2H2O - O2 + 4H+ + 4e) with the latter usually considered as the bottleneck of the overall process. Water oxidation is a challenging reaction in which the transfer of four protons and four electrons as well as the formation of an O–O bond has to be accomplished while avoiding the formation of high energy intermediates. In addition, water oxidation is very energy demanding, with an equilibrium potential of 0.82 V vs. NHE at pH 7. In natural a

Chemistry Department, Brookhaven National Laboratory, Upton, New York 11973-5000, USA. E-mail: [email protected] b National Institute of Advanced Industrial Science and Technology, Tsukuba Central 5-2, 1-1-1 Higashi, Tsukuba, Ibaraki, 305-8565 Japan c Japan Science and Technology Agency, ACT-C, 4-1-8 Honcho, Kawaguchi, Saitama, 332-0012 Japan † Electronic supplementary information (ESI) available. See DOI: 10.1039/ c3cp55101f ‡ On leave from the Faculty of Chemistry, Adam Mickiewicz University, Umultowska 89b, Poznan, Poland.

11976 | Phys. Chem. Chem. Phys., 2014, 16, 11976--11987

systems, the water oxidation reaction occurs at the oxygen evolving complex (OEC), whose structure has been recently clarified at a resolution of 1.9 Å:1 this OEC is made of a cluster of four manganese and one calcium ions, held together by oxygen bridges. Inspired by the function of the OEC, tremendous efforts have been made in developing catalysts able to carry out water oxidation reactions (WORs) at high rates for sustained periods. The discovery of efficient and robust catalysts capable of driving water oxidation at relatively low overpotential is a key challenge for efficient WORs. Extensive efforts have been devoted to developing water oxidation catalysts using transition metal complexes.2–19 In 2008, Bernhard and co-workers reported the first mononuclear iridium water oxidation catalyst.20 Since then the use of iridium complexes as water-oxidation catalysts has attracted significant attention because of their remarkable performance in terms of turnover frequencies (TOFs) and turnover numbers (TONs). Crabtree21–25 and others26–33 have been exploring the chemistry of a number of iridium complexes for use as water-oxidation catalysts. Most of the iridium catalysts reported to date contain the Cp* (Z5-pentamethylcyclopentadienyl) ligand. The Cp* ligand is believed to stabilize reaction intermediates with high-valent oxidation states. In addition to Cp*, the Ir complexes have other ligands such as phenyl pyridines, benzoylpyridine, bipyridines, or phenanthrolines. On the other hand, iridium oxides are well-known, robust and highly active heterogeneous water oxidation catalysts (WOCs),10,34–39 which leads to ambiguity regarding the true

This journal is © the Owner Societies 2014

View Article Online


Paper

PCCP

identity of the catalytic species. It has been demonstrated that coordination metal complexes, and in particular iridium compounds, can produce heterogeneous metal oxide catalysts under oxidative conditions.23,32,40,41 Distinguishing homogeneous from heterogeneous catalysis based on the experimental data may not be straightforward since only a small (undetectable) fraction of decomposed species can account for the observed catalytic activity. Possible degradation of the Ir complexes is especially pronounced if highly oxidizing cerium(IV) ammonium nitrate (CAN) is used as a sacrificial electron acceptor (E1/2 = 1.6 V at 1 M of HNO3).42 Several groups have reported oxidative degradation of ligands of Ir complexes leading to the formation of iridium oxide nanoparticles during WORs using CAN as a sacrificial oxidant.31,32 Recent studies based on the combination of catalytic experiments, UV-vis, NMR, in situ X-ray absorption, XRD and STEM measurements have shown that two regimes can possibly be identified: a short initial time period when molecular complexes are believed to be the active catalytic species, followed by longer times when oxide nanoparticles mainly contribute to the water oxidation reactions.31–33 In this context, alternative ways of driving WORs can be attractive in order to minimize ligand degradation and enable mechanistic studies of molecular catalysts. For instance, the replacement of CAN by a milder sacrificial oxidant such as NaIO4 may help to reduce ligand degradation as well as consequent nanoparticle (NP) formation. However, redox properties of

Table 1

periodates go beyond simple electron transfer. For example, it can act as an oxygen atom donor and thus activating pathways where produced O2 at least in part is not derived from water.43 It is difficult, however, to verify experimentally whether O2 production results from catalytic water oxidation or decomposition of periodate. In addition to both reactions having the same overall stoichiometry, 18O-labeling experiments are usually inconclusive owing to fast equilibrium of IO4 with water to produce H4IO6.43 Despite its ambiguity as a sacrificial electron acceptor, periodate is frequently used as a chemical oxidant in catalytic water oxidation (Table 1).23,24,29,43–45 Recent studies based on time-resolved dynamic light scattering (DLS) showed that water oxidation using NaIO4 with [Cp*Ir(OH2)3]2+ complexes may lead to the formation of amorphous oxide nanoparticles, but nanoparticle formation was not observed when an additional chelating ligand, such as 2,2 0 -bipyridine, 2-phenylpyridine or 2-(20 -pyridyl)-2-propanolate, was used.23 The effects of electron-donating substituents of Cp*Ir(bpy)Cl2 (bpy = 2,2 0 -bipyridine) on their water oxidation activity have been recently investigated.45,46 Papish and coworkers have studied [Cp*Ir(4DHBP)Cl]Cl (4DHBP = 4,40 -dihydroxy-2,2 0 -bipyridine) and [Cp*Ir(6DHBP)Cl]Cl (6DHBP = 6,6 0 -dihydroxy-2,2 0 -bipyridine) complexes for catalytic water oxidation under relatively mild conditions (near-neutral pH and periodate as oxidizing reagent) and reported that [Cp*Ir(4DHBP)Cl]Cl is more active than [Cp*Ir(6DHBP)Cl]Cl.45

Turnover frequencies and turnover numbers for some Ir catalysts

Catalysta

Oxidant

TOF (s1)

TON

Ref.

1 (5 mM) 2 (5 mM) 3 (5 mM) 4 (5 mM) 5 (5 mM) Cp*Ir(bpy)Cl2 (5 mM) Cp*Ir(bpy)SO4 (0.47 mM) Cp*Ir(bpy)Cl2 (41.5 mM) Cp*Ir(bpy)Cl2 (21.3 mM) Cp*Ir(pyalc)Cl (0.1 mM) Cp*Ir(bpm)Cl2 (5 mM) Cp*Ir(ppy)Cl (10 mM) Cp*Ir(bzpy)NO3 (1.3–3 mM) Cp*Ir(bzpy)Cl2 (38.5 mM) [Cp*Ir(4DHBP)H2O]SO4 (5 mM) Cp*Ir(4DHBP)Cl2 (5 mM) Cp*Ir(6DHBP)Cl2 (5 mM) Cp*Ir(k-2-N,O)Cl (0.5 mM) Cp*Ir(k-2-N,O)NO3 (1 mM) Cp*Ir(k-2,6-N,O)Cl (5 mM) Cp*Ir(NHC)Cl2 (70 mM) Cp*Ir(NHC)Cl2 (10 mM) Cp*IrCl2(Me2NHC) Cp*Ir(OH)2(Me2NHC) Ir(acac)3 (110 mM) IrCl3xH2O (90 mM) [Cp*Ir(OH2)3](NO3)2 (1–10 mM) [Cp*Ir(OH2)3](NO3)2 (23.6 mM) IrO2 (117 mM) IrO2 (0.31 mM)

NaIO4 (50 mM, pH 7.2) NaIO4 (50 mM, pH 7.2) NaIO4 (50 mM, pH 7.2) NaIO4 (50 mM, pH 7.2) NaIO4 (50 mM, pH 7.2) CAN (78 mM) NaIO4 (10 mM, pH 5.5) NaIO4 (23.2 mM, pH 5.4) NaIO4 (14.4 mM, pH 5.4) NaIO4 (10 mM, pH 5.5) CAN (78 mM) CAN (78 mM) CAN (10 mM) NaIO4 (14.8 mM, pH 5.4) CAN (10 mM) NaIO4 (20 mM, pH 5.6) NaIO4 (20 mM, pH 5.6) CAN (20 mM) CAN (40 mM) CAN (20 mM) CAN (125 mM) NaIO4 (500 mM) CAN (333 mM) CAN (333 mM) CAN (170 mM) CAN (170 mM) CAN (10 mM) NaIO4 (15.8 mM, pH 5.4) CAN (170 mM) NaIO4 (10 mM, pH 5.2)

3.75 2.2 3.5 5.2 0.28 0.24 0.042 0.048 0.04 2.2 0.07 0.9 0.21 0.11 0.75 0.2 0.17 4.78 4.62 0.17 0.36 0.31 0.5 1.5 0.33 0.08 0.08–0.63 0.36 0.097 0.38

6700 6330 6720 6700 7200 — — 200 225 — — — — 168 — 120b 100b — — 750 400 424 000 120 c2000 — — — 305 — —

This This This This This 21 43 29 29 43 21 25 27 29 32 45 45 26 26 26 44 44 66 66 67 67 27 29 67 68

work work work work work

a Note that TOFs are very dependent on catalyst and oxidant concentrations, TOFs are reported per Ir center; k-2-N,O = 2-pyridinecarboxylic acid; k-2,6-N,O = 2,6-pyridinedicarboxylic acid; NHC = 3-methyl-1-(1-phenylethyl)-imidazoline-2-ylidene; Me2-NHC = N-dimethylimidazolin-2-ylidene; pyalc = 2-(2 0 -pyridyl)-2- propanolate; bzpy = 2-benzoylpyridine; ppy = 2-pyridyl-2 0 -phenyl. b TON after 10 min.


Phys. Chem. Chem. Phys., 2014, 16, 11976--11987 | 11977

View Article Online


PCCP

They found that the –OH groups of 4DHBP improve the catalytic activity by changing the electronic properties of the complexes upon deprotonation. It should be noted that these catalysts change to the corresponding aqua complexes upon dissolving in water. Importantly, based on dynamic light scattering experiments, the authors concluded that the catalysis appears to be homogeneous. Lin and coworkers reported also that these complexes are oxidized electrochemically or with Ru(bpy)33+ without producing oxygen, but generate oxygen with NaIO4 and CAN as oxidants.46 Fukuzumi and coworkers have shown that the same catalyst [Cp*Ir(4DHBP)(H2O)]2+ forms nanoparticles composed of iridium hydroxide with a small amount of carbonaceous residue when the stronger oxidant CAN is used.32 Interestingly, Himeda previously demonstrated that this complex and the related 4,7-dihydroxy-1,10phenanthroline complex exhibit enhanced CO2 hydrogenation catalysis in aqueous basic solution due to the proton-responsive nature of the ligands which have pKa values around 5.47–50 Himeda’s group as well as ours have further improved the catalytic activity of CO2 hydrogenation using related complexes.51–53 The high catalytic activity of iridium(III) complexes was attributed to the strong electron-donating abilities of the ligands containing oxyanion groups generated from deprotonation of –OH in basic solutions. We have carried out detailed mechanistic studies of water oxidation driven by periodate and using water-soluble Cp*Ir mononuclear and dinuclear complexes with electron-donating hydroxyl groups on the bipyridine ligand: [Cp*Ir(6DHBP)(H2O)]2+ (1); [Cp*Ir(th4bpm)(H2O)]2+, th4bpm = 2,2 0 ,6,6 0 -tetrahydoxy-4,40 bipyrimidine (2); and [{Cp*Ir(H2O)}2(thbpm)]4+, thbpm = 4,4 0 ,6,60 tetrahydroxy-2,2 0 -bipyrimidine (3), as shown in Chart 1. While all these complexes were reported as efficient catalysts for CO2 hydrogenation, we have compared their reactivity as water oxidation catalysts with the corresponding complexes with unsubstituted ligands (4–5) in order to examine the effect of the hydroxyl groups on the catalytic performance. We also provide insight into the nature of the catalytically active species based on UV-vis spectroscopy, 1H NMR, DLS, transmission electron microscopy (TEM), and kinetic analyses. Here we report

Chart 1 Iridium(III) complexes used for water oxidation with periodate as the oxidant.


Paper

that the collected data based on a combination of catalytic experiments and spectroscopic measurements suggest that a species with an absorption at lmax 590 nm (A) formed at low oxidant concentration is the active species for water oxidation, and is composed of small iridium clusters with size o2 nm. The formation of larger nanoparticles that have lmax 600 nm (B) is identified at a high oxidant/catalyst ratio, but these seem to be less active than the small iridium clusters.

Experimental section Materials Tris(2,20 -bipyridine)ruthenium(II)chloride hexahydrate ([Ru(bpy)3]Cl26H2O) was obtained from Strem Chemicals Inc.; potassium persulfate and sodium periodate were supplied by Sigma-Aldrich. All chemicals commercially available were used without further purification. Aqueous solutions were prepared with distilled water that had been passed through a Millipore ultra-purification system. Complexes 1, 2, 3, and 4 were prepared using previously published methods.51–53 Spectroscopic measurements and characterization of intermediates UV-vis spectra were measured on an Agilent 8453 diode-array spectrophotometer. 1H NMR spectra were measured on a Bruker UltraShield 400 MHz spectrometer in D2O solutions. An Applied Photophysics stopped-flow system configured for two-syringe mixing was used to carry out stopped flow experiments. Spectral changes were detected within the 300–1100 nm range. In a typical experiment, one syringe was charged with a sodium periodate solution and the other syringe contained a solution of the iridium complex. The spectral change was recorded after equal amounts of both solutions were mixed in a 1 cm stopped-flow cell. The ultrafiltration technique was used to identify the presence of small particles of size o2 nm in the solution containing 1 (100 mM) and periodate (5 mM) in acetate buffer (50 mM, pH 5.5). Millipore Ultracel Ultrafiltration Discs (Molecular Weight Cut Off MWCO 1 kDa, Stokes radius o 0.5 nm) and 10 kDa (MWCO 10 kDa, Stokes radius o 2 nm) were used in an Amicon stirred cell under a constant pressure of 50 psi. The UV-vis spectrum of the solution after filtration was compared with that of the unfiltered solution. The procedure for DLS analysis was as follows. A catalyst solution (2 mM) was added to a solution of sodium periodate (50–200 mM) and allowed to react. The final catalyst concentration was 1 mM. When the formation of bubbles stopped, the sample was transferred to a clear cell and then measured on a Precision Detectors Model P2000DLS/Batch Light Scattering System equipped with 100 mW 800 nm laser with scattered light being detected at 901. Each solution was analyzed at least two times. TEM images of nanoparticles, which were mounted on a copper microgrid coated with elastic carbon, were observed using a JEOL JEM3000F TEM equipped with a field emission gun and a Gatan imaging filter (GIF). The instrument was operated at 300 kV.


View Article Online

Paper


Catalytic oxygen evolution Chemical water oxidation was performed as follows. Oxygen measurements were performed using a calibrated O2 probe (Ocean Optics HIOXY probe with a factory calibrated multipoint calibration). A three-neck round bottom flask with two side arms equipped with T bore Teflon plugs was used to purge the flask and facilitate oxidant injection. The O2 probe was inserted through the third neck and was sealed using an UltraTorr fitting. The total working volume was 35 mL. The gas-tight vessel was equipped with a stir bar and placed in a temperature controlled water bath (24 1C). In a typical experiment, an aqueous solution of sodium periodate containing acetate buffer (50 mM, pH 5.5) or phosphate buffer (50 mM, pH 7.2) in a glass vessel was purged with argon for ca. 20 min until a stable reading was obtained. Then a degassed aqueous solution of the iridium complex was injected using a gas tight syringe (Hamilton). In most experiments the total volume of the solution after mixing catalyst and oxidant solutions was 10 mL except the experiments presented in Fig. 2 where the total volume was 15 mL. Alternatively, Ru(bpy)33+ was used as the chemical oxidant. The Ru(bpy)33+ was prepared by bulk electrolysis (applied potential 1400 mV vs. NHE) of a solution containing 2 mM of Ru(bpy)32+ in 0.02 M triflic acid. Bulk electrolysis was conducted with a BAS 100b potentiostat from Bioanalytical Systems. Platinum gauze was used as a working electrode, a platinum wire as a counter electrode, and Ag/AgCl as a reference electrode. Water oxidation by Ru(bpy)33+ was examined in the presence of 8 mM of catalyst 1. Just before adding the catalyst to the acidic Ru(bpy)33+ solution, sodium acetate was added to obtain a pH of 5.5. Photocatalytic water oxidation was performed as follows. Light-induced water oxidation was performed in a cuvette (1 cm 1 cm) with a total volume of 6.2 mL sealed with a vacuum tight valve. The vessel was filled with 2.9 mL of solution with 50 mM of 1, 1 mM [Ru(bpy)3]Cl2 and 10 mM K2S2O8 dissolved in an aqueous solution with sodium phosphate buffer (50 mM, pH 7.2) or acetate buffer (50 mM, pH 5.5). The solutions were prepared in a glove box under an Ar atmosphere. All procedures were performed with a minimum exposure to ambient light. The reactions were started by irradiation with monochromatic light of lexc = 476 nm from an Ar/Kr ion laser (Coherent Innova 70 C). The light absorption of the catalyst at 476 nm is negligible in comparison to absorption by the photosensitizer. The laser beam was expanded to a diameter B1 cm. The solution was irradiated under vigorous stirring and cooled with a water cooling system to 20 1C. The kinetics of oxygen evolution in the gas phase were monitored by an Ocean Optics oxygen sensor (HIOXY). Determination of CO2 produced by oxidation of ligands Gases produced during catalytic water oxidation were analyzed using a QMS 300 Gas Analyzer (Stanford Research Systems). The reaction vessel was connected to a vacuum manifold which in turn was connected to the inlet of the gas analyzer. Solutions of an oxidant and a metal complex were carefully loaded into the two-compartment reaction vessel without mixing. Both solutions


PCCP

were degassed by intermittently opening a valve to the vacuum manifold equipped with a LN2 trap. Degassing was repeated until no pressure reading was detected on the vacuum gauge under static vacuum. Then the cold trap was switched to a LN2–ethanol mixture (30 1C) in order to trap water, but not CO2 gas. Deaerated solutions of chemical oxidant and metal complex were rapidly mixed and allowed to stand under stirring for a period of 5 minutes, similar to the observed induction time in oxygen evolution. The reaction headspace was then transferred into the vacuum manifold and was allowed to equilibrate to ensure that water vapor was trapped in a cold trap prior to sampling. The analyzer was set up to monitor m/z signals which corresponded to oxygen, nitrogen, and carbon dioxide.

Results and discussion Oxygen evolution studies in the presence of Ir complexes and periodate Catalytic water oxidation by periodate was examined using a series of Ir complexes, which may act as precatalysts, in aqueous solutions at three different pH values: 4.0, 5.5, and 7.2 (Fig. 1). Upon the addition of catalysts 1–3 to a solution with a large excess of sodium periodate (i.e., 2000 times), oxygen begins to evolve with an increasing rate after an induction time of B2 min, reaching a maximum after B6 min. This maximum rate was used to analyze the kinetic behavior of the system. The apparent TOF was observed at an almost constant value during the course of the oxygen evolution in all of the studied reactions. At the end of the reaction, the apparent TOF value decreased due to consumption of NaIO4. The stoichiometry for the oxygen produced was half of the employed periodate, which suggests that all periodate was consumed as a two-electron oxidant. A change in the activity of all complexes 1–5 (Fig. 1) was observed with increasing pH. Our findings are in agreement with recent work on complex 1 by Papish et al.45 The pKa values for 1 and 3 are 4.645 and 3.854 respectively. Although the UV-vis changes for 2 were small, the pKa value for 2 was estimated to be about 6.7 (Fig. S1, ESI†). While the increased activity for 1–3 can be attributed to the electronic effect based on the acid–base equilibrium of the phenolic hydroxyl group attached to the ligands, the rate of water oxidation can also increase with an increase of pH if the rate determining step involves a protoncoupled electron-transfer (PCET) reaction. In fact, a rate increase associated with increasing pH was observed with complexes 4–5, which do not have any OH group(s) on the bpm or bpy ligand. It should be noted that complexes 1–5 are all precatalysts in these experiments, and the active catalyst seems to be small nanoclusters in the range of 0.5–2 nm, especially in the use of precatalyst 1, which will be discussed in a later section. Complex 2 bearing four hydroxyl groups showed lower activity than 1, which has only two hydroxyl substituents. Complex 5 showed very limited catalytic activity, and only a slight increase in TOF was detected upon changing the pH (Fig. 1e). However, unexpected behavior was observed for catalyst 4. At pH 5.5 the catalyst was found to have low activity


View Article Online


PCCP

Paper

Fig. 1 pH-dependent water oxidation as measured by oxygen evolution with precatalyst 1–5 (5–6 mM) and NaIO4 (10 mM). Color code: red – pH 4.0, blue – pH 5.5, black – pH 7.2; TOFs are maximum rate/[Ir center].

Fig. 2 Water oxidation as measured by oxygen evolution with precatalysts 1–5 (5 mM) and NaIO4 (50 mM) in phosphate buffer (50 mM, pH 7.2); TOFs are maximum rate/[Ir center].

toward water oxidation with an induction period of 30 min and a TOF of 2.5 min1 (Fig. 1d), very similar to the TOF obtained for 5. At pH 7.2 the induction period was reduced to 3 min and the TOF increased to 41.6 min1 under the same experimental conditions ([cat] = 6 mM, [IO4] = 10 mM). To explain this enhanced catalytic activity with a pH change from 5.5 to 7.2, the dependence of the UV-vis spectrum on the pH was measured (Fig. S2, ESI†). Based on the analysis of the UV-vis changes with


the pH, the pKa of the complex was determined to be 5.9. The observed pKa is attributed to an equilibrium between the aqua and hydroxo complex 4. The observed pKa of the aqua ligand is usually higher54 but Fukuzumi et al.,55 using 1H NMR, have determined that the pKa value of the aqua ligand in [Cp*Ir(bpy)(OH2)]2+ (pKa = 7.7) is lower by 1.5 unit than that of the one with the more electron-donating 4,4 0 -OMe2-bpy ligand (pKa = 9.2). The obtained results show that catalytic activity may be tuned significantly with a change in the pH not only for hydroxyl substituted ligands (1–3) but also for the aqua complex 4 if the pKa of the aqua ligand is lower than the pH used in the water oxidation experiment. Effects of these local OH group(s) on the decomposition of the precatalysts or the structure on the formed active catalyst remain unclear. Under optimized conditions (properly regulated stirring, higher oxidant concentration), the maximum TOF for mononuclear precatalysts was found to be 5.2 s1 and 3.8 s1 for 4 and 1, respectively (Fig. 2). They afforded a turnover number of 6700 solely limited by the total consumption of the sacrificial oxidant. For dimeric Ir complex 3 the maximum TOF was found to be 3.5 s1 (TOF per Ir center). These values render precatalysts 1, 3 and 4 as the most active Ir complexes at room temperature for water oxidation (Table 1). Macchioni et al. reported an initial TOF of 4.8 s1 for oxidative water splitting to O2 driven by CAN


View Article Online


Paper

Fig. 3 O2 evolution experiment for 1 (0.5 mmol) performed with sequential addition of 25 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5).

with [Cp*Ir(k2-N,O)Cl] (k2-N,O = 2,4-pyridinedicarboxylic acid), which was the highest value reported until now for an Ir catalyst (Table 1).26 The highest TOF value reported for a catalyst under milder conditions (higher pH, NaIO4 as oxidant) was 2.2 s1 for Cp*Ir(pyalc)Cl (pyalc = 2-(2 0 -pyridyl)-2-propanolate).43 In order to provide better mechanistic insight into water oxidation catalyzed by complexes 1–5 we studied kinetic parameters of the catalytic reaction in more detail. Further studies were focused on complex 1, because results could be directly compared to detailed studies by Crabtree et al. for complex 5 with unsubstituted bpy.23,24,43 To understand whether the degradation of the complex occurs, and if so, if it is parallel to or precedes water oxidation, the experiments were performed with a small excess of the sacrificial oxidant. First of all, it was noticed that oxygen evolution is observed only when the ratio of [NaIO4]/[1] 4 20 (Fig. S3, ESI†). It was found that when using 25 eq. of NaIO4 oxygen formation began 120 s after the addition of the oxidant (Fig. 3). Since the water oxidation process requires 4 one-electron oxidation steps to generate 1 equivalent of oxygen, it is conceivable that modification of the precatalyst to form the catalytically active species occurs during the induction period before the onset of oxygen evolution. Once the oxygen evolution finished, a second 25 eq. of sacrificial oxidant was added to the solution. The maximum TOF was 3 times higher than that observed after the first addition of periodate (Fig. 3). The maximum TOF values further increased after the addition of the third aliquot. Subsequently, the TOFs remained constant at least for another 4 additions.

PCCP

Moreover, the induction time was reduced to 80 s after the second addition of the oxidant and to 20 s after the third addition of the oxidant. An induction time of 20 s is similar to the response time of the O2 probe. This observation confirms that the Ir complex 1 acts indeed as a precatalyst which is converted into a catalytically active species during the induction period. The rates of water oxidation remained almost unchanged during the remaining course of the reaction involving multiple NaIO4 additions after the initial transformation during the first two additions. Initial rates of O2 formation have a first-order dependence on precatalyst at low catalyst concentration (2.5–10 mM) (Fig. 4a). The order of the reaction is reduced significantly at higher iridium complex concentrations, suggesting an aggregation of the iridium species to form dimers or more complex species. The concentration effect of precatalyst 1 on the O2 yields was also examined (Fig. 4b). The amount of evolved O2 increased with an increase in the concentration of the precatalyst 1 up to 5 mM with a maximum O2 yield of 96%. Then the O2 yield gradually decreased upon further increases in concentration of precatalyst 1. This suggests that at higher catalyst concentrations the iridium species may aggregate to a less reactive form. Comparison of water oxidation by precatalyst 1 using periodate and photochemically produced Ru(bpy)33+ The role of periodate in water oxidation catalyzed by 1 was examined by replacing it with another sacrificial oxidant, Ru(bpy)33+, with one-electron reduction potential (1.27 V vs. NHE)56,57 significantly higher than that of IO4 (0.3–0.45 V vs. NHE at pH 7).58 Interestingly, no O2 formation was observed when Ru(bpy)33+ was applied as a chemical oxidant at pH 5.5 and pH 7.2. In addition, we attempted to achieve photo-induced water oxidation using Ru(bpy)32+ as a photosensitizer and Na2S2O8 as a sacrificial electron acceptor and 1 as a precatalyst. The Ru(bpy)32+/ Na2S2O8 is a well-known couple, often used in photo-induced reactions to generate Ru(bpy)33+.59–63 The experiments performed at pH 5.5 and 7.2 did not lead to any measurable O2 formation. Even after precatalyst 1 was activated with 30 eq. of NaIO4 prior to the light-driven water oxidation experiment, O2 formation could not be detected. These results suggest that the action of periodate might be more complex than just pure electron transfer, possibly involving an oxygen atom transfer to the catalyst.64 In this case the

Fig. 4 (a) Plot of the oxygen evolution rates vs. initial catalyst amount in the solution; (b) total oxygen formation for different catalyst concentrations for solution with 100 mmol of NaIO4 in acetate buffer (50 mM, pH 5.5).



View Article Online


PCCP

Paper

formed oxygen molecule would contain one atom derived from water and one atom from periodate. It was shown previously that in the reaction of dinuclear manganese complexes with oxone, one of the oxygen atoms of the O2 produced does not originate from bulk water but most probably from the oxidant.65 In addition, Hetterscheid and Reek have suggested based on DFT calculations and in situ mass spectrometry experiments that in water oxidation with [Cp*Ir(Me2NHC)(OH)2] (Me2NHC = N-dimethylimidazolin-2ylidene) by periodate water may not necessarily be the oxygen source. A low-lying pathway may exist in which O2 production occurs without direct involvement of water. The coordination of periodate to the relatively electron-rich intermediates involved into the catalytic cycle and O-atom transfer from coordinated periodate to iridium were found based on DFT calculations to be very facile.64 Unfortunately, 18O-labeling experiments will be inconclusive owing to the fast exchange of oxygen in water with periodate. The drastically different behaviors of precatalysts under the same water oxidation conditions and their pH dependence indicate that 1–4 are more prone to oxidation, consistent with the fact that the electron donating nature of the OH facilitates oxidation of the catalyst.

all peaks corresponding to 1 and rapid formation of new signals corresponding to acetic acid and formic acid (time for sample preparation and recording 1H NMR B6 min). The acetic and formic acid peaks in the 1H NMR spectrum were confirmed by spiking the solution with authentic samples. Macchioni et al.27 observed similar behavior when a Cp*-functionalized iridium complex underwent oxidation by CAN. Methyl groups bound to the aromatic moiety are oxidized to form alcohols, aldehyde, or acid derivatives that finally break apart the Cp* structure to form CH3CO2H or HCO2H. The 1 H NMR study suggests a quick oxidative degradation of the iridium complex through oxidation of the Cp* ligand. Further evidence of ligand oxidation was obtained from oxygen and carbon dioxide detection after mixing the catalyst with the oxidant using MS analysis. Upon addition of the solution of 1 (150 mM) to the solution of CAN (15 mM) at pH 1 both oxygen and carbon dioxide were detected after 5 minutes with the ratio of O2/CO2 = 0.8, showing ligand degradation to carbon dioxide. The addition of 1 (150 mM) to a solution of NaIO4 (15 mM) at pH 7.2 resulted in the formation of both oxygen and carbon dioxide after 5 minutes with the ratio of O2/CO2 B 4. This demonstrates that ligand oxidation of 1 with periodate also takes place but is less efficient than with CAN.

NMR investigation of chemical composition of the reaction mixture

UV-vis spectral changes during water oxidation and kinetic analysis

In order to explore the possibility of oxidative degradation of molecular Ir complexes and, in particular, the oxidation of the organic ligands upon reaction with NaIO4, the following experiments were conducted. 1H NMR measurements were performed in D2O solutions containing 1 (1 mM) and periodate (0–10 mM) to detect soluble species after the catalytic water oxidation by periodate. In Fig. 5, the red spectrum corresponds to the complex 1 before oxidation. Subsequent additions of periodate led to complete disappearance of peaks characteristic of the initial catalyst. The addition of only 10 eq. of periodate to the catalyst solution resulted in the complete disappearance of

The UV-vis spectrum of a solution of complex 1 (100 mM) in acetate buffer (pH 5.5) shows important changes during the first 10 min of reaction after the addition of NaIO4 (50 eq.). The addition of NaIO4 into a solution of 1 leads (in less than 1 second) to the disappearance of the complex peak at 375 nm (Fig. 6a, black spectrum) and the formation of a new feature in the UV-vis spectrum with a maximum at 345 nm (Fig. 6a, blue spectrum). This band quickly (B60 s) disappears (Fig. 6b) and a new band with maximum at 590 nm is formed. However formation of the band at 590 nm is slower than the disappearance of the band at 345 nm (Fig. 6b). Accordingly, the solution changes color from yellow to blue. The species that exhibits a band at lmax 590 nm (defined as species A) decayed on a longer time scale in solution with the rate depending on the initial periodate/catalyst ratio. The monitored UV-vis changes for complex 1 are similar to those reported earlier for complex 5.24 The initial induction time in the catalytic process can be tentatively attributed to the transformation of the iridium complex to form the true active species. The absorption band at 590 nm is strongly correlated with O2 formation. The formation of the 590 nm band (i.e., formation of species A) as well as O2 formation was not observed if less than 20 eq. of NaIO4 was added to the catalyst solution. Similarly, for precatalyst 4 a very long induction time for O2 formation was observed in the experiment at pH 5.5, but was significantly reduced at higher pH (Fig. 1d). This behavior was reflected in the UV-vis changes for precatalyst 4 at different pH values. After addition of 25 eq. of NaIO4 to the solution of 4 (230 mM) in acetate buffer (pH 5.5) the formation of the 590 nm band was slow and the maximum intensity was reached after 2700 s

Fig. 5 1H NMR spectra of 1 (1 mM) in the absence of periodate (red) and the presence of 10 eq. of NaIO4 (blue).



View Article Online


Paper

PCCP

Fig. 6 (a) UV-vis changes after mixing 1 (100 mM) with 50 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5) and (b) kinetics of the 590 nm and 345 nm absorption band.

Fig. 7 Intensity of the band at 590 nm vs. O2 evolution of the same reaction at 100 mM of 1 and 10 mM NaIO4 in acetate buffer (pH 5.5). Inset: whole UV-vis kinetics of the reaction at 590 nm.

(Fig. S4, ESI†). When the analogous experiment was performed at pH 7.2 the band at 590 nm was formed within 500 s (Fig. S5, ESI†) as observed for precatalyst 1. Parallel O2-monitoring and UV-vis analysis of the catalytic reactions permitted observation of interesting correlations between spectroscopic features and the water oxidation activity. In two separate experiments, equal

volumes of NaIO4 (30 mM) and catalyst 1 (200 mM) solutions in acetate buffer were mixed and O2 formation or UV-vis changes were monitored. A comparison between O2 production and the growth of the 590 nm band is shown in Fig. 8 under the same experimental conditions. The O2 evolution starts only when the concentration of species A builds up. The absorbance at 590 nm remains almost at a constant level during the O2 evolution and can be attributed to some intermediate being in a steady state. After the O2 evolution stops, the intensity at 590 nm slightly increases for a short period and then decays completely, resulting in the yellow solution (inset of Fig. 7). When a 1 mM solution of 1 was mixed with 50 eq. of NaIO4 at pH 5.5 the yellow solution turns blue within 10 minutes. This corresponds to the formation of the 590 nm band (Fig. 8), which disappears within one hour (Fig. 8 and 10). When the analogous experiment was performed with 200 eq. of NaIO4 the solution still developed its characteristic blue color with an intense absorption at 590 nm within minutes (Fig. 9, green line). Then the band increased in intensity and shifted slightly to 600 nm (Fig. 9, red line) (formation of species B). The solution remained blue (Fig. 10) and when the reaction mixture was allowed to stand overnight, a blue-black precipitate was observed. Formation of the blue-black precipitate

Fig. 8 UV-vis changes after mixing 1 (1 mM) with 50 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5): (a) initial changes from 0–500 s; (b) changes from 500–3850 s. Inset: whole UV-vis kinetics of the reaction at 590 nm.



View Article Online


PCCP

Fig. 9 UV-vis changes of 1 (1 mM) (black) after mixing with 200 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5): green – 500 s, red – 1250 s, purple – 3850 s, blue – overnight.

Fig. 10 Different colors of the acetate buffer (50 mM, pH 5.5) solutions containing 1 mM 1 before and after addition of 50 and 200 equivalents of periodate.

was also observed when 1 mM solution of 1 was mixed with 50 eq. of NaIO4 in water instead of buffer solution at pH 5.5. Identity of reactive species To determine whether the blue species formed after addition of only 50 eq. of NaIO4 to a 1 mM buffer solution of 1 is molecular or nanoparticles, a large excess of ethanol (100 mL) was added to the solution when the maximum absorbance at 590 nm was reached. The addition of ethanol causes acceleration of the

Paper

590 nm band disappearance (Fig. 11a); however UV-vis bands of the starting materials were not recovered. In an analogous experiment performed on the solution to which 200 eq. of periodate was added, the 600 nm band persisted with only a small decrease in intensity after the addition of ethanol (Fig. 11b). Thus, the blue species A formed upon oxidation of 1 with 50 eq. NaIO4 at pH 5.5 is distinct from the species B formed upon oxidation of 1 with 200 eq. NaIO4. The reactivity of species B with ethanol is similar to that of IrOx nanoparticles reported previously.23,69,70 DLS experiments were performed to ascertain whether nanoparticles were formed during oxidation of 1 with different NaIO4 concentrations. DLS experiments with 1 (1 mM) with 50 eq. of NaIO4 in acetate buffer did not show evidence for nanoparticle formation during a 16 h period. The DLS analysis indicates that if nanoparticles are formed, the concentration must be below the detection limit of the instrument and/or the sizes must be smaller than 1 nm. When the solution contained 1 mM of 1 and 200 eq. of NaIO4, particle formation was observed 50 minutes after mixing the catalyst with the oxidant (Fig. S6 inset, ESI†). A particle size of 120 nm was detected by DLS measurements as shown in Fig. S6 (ESI†). The species produced by oxidation of 1 with 50 eq. (sample X) and 200 eq. of NaIO4 (sample Y) in acetate buffer (pH 5.5) and with 50 eq. of NaIO4 in water (sample Z) were examined by transmission electron microscopy. TEM of samples Y and Z revealed the presence of nanoparticles which are shown in Fig. S7 (ESI†). For sample Z nanoparticles of the average size B2 nm were observed (Fig. S7b, ESI†) and for sample Y only conglomerates of sizes 500–1000 nm were found. No nanoparticles were detected for sample X. Different results obtained for samples X and Z indicate that complex behavior is highly sensitive to water-oxidation conditions. Based on the results obtained from DLS, TEM and the absence of reactivity with ethanol, species B formed after oxidation of 1 with high (200 eq.) concentration of periodate can be identified as nanoparticles containing IrOx and possibly other elements such as carbon. Unambiguous demonstration of the homogeneity of the active species A at low catalyst/ oxidant ratio is not straightforward. On the one hand, lack of observation of nanoparticles in DLS or TEM together with the

Fig. 11 UV-vis spectra of a solution of (a) 1 (1 mM) with 50 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5); (b) (1 mM) with 200 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5), after addition of ethanol.



View Article Online


Paper

PCCP

Fig. 12 UV-vis spectra of the solution of 1 (100 mM) with 50 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5): black line – without filtration, red line – after filtration. (a) Ultrafiltration disc (MWCO 1 kDa, Stokes radius o 0.5 nm), (b) disc (MWCO 10 kDa, Stokes radius o 2 nm).

acceleration of the 590 nm band disappearance after addition of EtOH suggest that the active catalytic species A might be molecular in nature. On the other hand, the unexplained induction time before the catalysis begins, and the formation of oxygen only when [oxidant]/[catalyst] 420 might be an indication of heterogeneity. Moreover, 1H NMR and gas analysis indicated that ligands (Cp* and DHBP) undergo oxidation to form acetic acid and CO2. In order to explore the possibility of the formation of small IrOx clusters (o1 nm), the ultrafiltration technique was employed. A solution of 1 (100 mM) with 50 eq. of NaIO4 in acetate buffer (50 mM, pH 5.5) was prepared. Half of the solution was filtered through an ultrafiltration disc (MWCO 1 kDa, Stokes radius o0.5 nm). Afterwards, the UV-vis spectra of the two solutions were compared (Fig. 12a). Interestingly, the 590 nm band that is attributed to the catalytically active species completely disappeared after filtration. An analogous experiment was performed using an ultrafiltration disc with a larger Stokes radius (MWCO 10 kDa, Stokes radius o2 nm). The solution after filtration retained most of the 590 nm band (Fig. 12b), indicating that the active species absorbing around 590 nm are unlikely to be molecular in nature, but rather are small particles with size in the range 0.5–2 nm. Iridium oxide clusters of 4–5 Ir atoms were reported earlier to have a broad absorption band around 580 nm, consistent with our observations.36 The lack of an observed signal in DLS experiments could be attributed to low sensitivity of this technique to very small particles (o1 nm). Our observation of different precatalyst behavior in terms of nanoparticle formation depending on the oxidant equivalents added to the solution is supported by the recent findings of Crabtree et al.23 They found that nanoparticle formation and size crucially depend on the exact reaction conditions including [oxidant]/[catalyst] ratio, catalyst concentration or presence of a salt.23

Conclusion The activity of five iridium complexes in the periodate-driven oxidation of water to oxygen has been compared at different pH values. Application of sodium periodate rather than CAN as a sacrificial oxidant for water oxidation allowed us to study water oxidation at neutral pH. The catalytic activity of 1–3 was found


to depend remarkably on pH. The OH groups are deprotonated at higher pH which facilitates the formation of the active form of the catalyst. A strikingly higher catalytic activity of [Cp*Ir(bpm)(H2O)]2+ over [Cp*Ir(bpy)(H2O)]2+ by a factor of 18 was found at neutral pH. The catalytic activation of precatalyst 4 above pH 6 is attributed to the deprotonation of the aqua ligand. At neutral pH, the activities of species formed from precatalysts 1–4 are among the highest ever reported for iridium catalysts (5.2 s1 for 4, 3.8 s1 for 1 and 3.5 s1 for 3). However, the effects of OH groups attached to the bpy or bpm ligand or a coordinated OH ion toward the formation of active catalyst A (or the geometric and electronic structure of A) are not clear at this stage. A detailed analysis of spectroscopic data together with O2-evolution experiments indicate that catalytic water oxidation promoted by complex 1 is mostly associated with transient species A that exhibits a band at lmax 590 nm. The formation of the active form was found to be strongly dependent on reaction conditions. While species A was distinctly observed using a small excess of periodate, under high oxidant concentration the formation of species B absorbing at 600 nm was more prominent. DLS analysis and TEM have identified species B as 120 nm nanoparticles. The ultrafiltration method has revealed that species A can be assigned to particles with size in the range of 0.5–2 nm, possibly small IrOx clusters similar to those described previously, e.g., by Harriman and co-workers.36 The activity of the active species formed from precatalysts 1–4 was found to be substantially higher compared to complex 5. Complex 5 might be more stable toward oxidative decomposition as proposed by Brudvig, Crabtree and co-workers.23 They suggested instead that the active form of 5 is a bis-m-oxo diiridium(IV) dimer complex with two waters and one chelate ligand bound to each other.24 Based on a combination of multiple techniques, the active species in the presently described Ir-catalyzed NaIO4-driven water oxidation appear to be heterogeneous in nature even at low oxidant concentration. The exact structure and number of Ir centers in the active species remain to be clarified.

Acknowledgements We thank Dr James T. Muckerman for a careful reading of this manuscript. Dr Yimei Zhu is acknowledged for providing access


View Article Online

PCCP


to a JEOL JEM3000F TEM. We thank Dr Sergei Lymar for help with calculations of one-electron reduction potentials of periodate. The work at BNL was carried out under contract DE-AC0298CH10886 with the U.S. Department of Energy and supported by its Division of Chemical Sciences, Geosciences, & Biosciences, Office of Basic Energy Sciences. Y. H. and W.-H. W. thank the Japan Science and Technology Agency (JST), ACT-C for financial support.

References 1 N. Kamiya, K. Kawakami, J.-R. Shen and Y. Umena, Nature, 2011, 473, 55–60. 2 S. Berardi, G. La Ganga, M. Natali, I. Bazzan, F. Puntoriero, A. Sartorel, F. Scandola, S. Campagna and M. Bonchio, J. Am. Chem. Soc., 2012, 134, 11104–11107. 3 Z. Chen, J. J. Concepcion, J. F. Hull, P. G. Hoertz and T. J. Meyer, Dalton Trans., 2010, 39, 6950–6952. 4 P. Comte, M. K. Nazeeruddin, F. P. Rotzinger, A. J. Frank and M. Gratzel, J. Mol. Catal., 1989, 52, 63–84. 5 J. J. Concepcion, J. W. Jurss, J. L. Templeton and T. J. Meyer, J. Am. Chem. Soc., 2008, 130, 16462–16463. 6 J. J. Concepcion, J. W. Jurss, M. R. Norris, Z. Chen, J. L. Templeton and T. J. Meyer, Inorg. Chem., 2010, 49, 1277–1279. 7 D. E. Polyansky, J. T. Muckerman, J. Rochford, R. Zong, R. P. Thummel and E. Fujita, J. Am. Chem. Soc., 2011, 133, 14649–14665. 8 L. Duan, A. Fischer, Y. Xu and L. Sun, J. Am. Chem. Soc., 2009, 131, 10397–10399. ¨gerler, D. A. Hillesheim, D. G. Musaev 9 Y. V. Geletii, B. Botar, P. Ko and C. L. Hill, Angew. Chem., 2008, 120, 3960–3963. 10 M. Hara, C. C. Waraksa, J. T. Lean, B. A. Lewis and T. E. Mallouk, J. Phys. Chem. A, 2000, 104, 5275–5280. 11 D. Hong, J. Jung, J. Park, Y. Yamada, T. Suenobu, Y.-M. Lee, W. Nam and S. Fukuzumi, Energy Environ. Sci., 2012, 5, 7606–7616. ¨rka ¨s, T. Åkermark, E. V. Johnston, S. R. Karim, 12 M. D. Ka T. M. Laine, B.-L. Lee, T. Åkermark, T. Privalov and B. Åkermark, Angew. Chem., Int. Ed., 2012, 51, 11589–11593. 13 N. Kaveevivitchai, R. Chitta, R. Zong, M. El Ojaimi and R. P. Thummel, J. Am. Chem. Soc., 2012, 134, 10721–10724. 14 N. D. McDaniel, F. J. Coughlin, L. L. Tinker and S. Bernhard, J. Am. Chem. Soc., 2007, 130, 210–217. 15 J. T. Muckerman, D. E. Polyansky, T. Wada, K. Tanaka and E. Fujita, Inorg. Chem., 2008, 47, 1787–1802. 16 L. Tong, Y. Wang, L. Duan, Y. Xu, X. Cheng, A. Fischer, M. S. G. Ahlquist and L. Sun, Inorg. Chem., 2012, 51, 3388–3398. 17 H.-W. Tseng, R. Zong, J. T. Muckerman and R. Thummel, Inorg. Chem., 2008, 47, 11763–11773. 18 Y. Xu, A. Fischer, L. Duan, L. Tong, E. Gabrielsson, B. Åkermark and L. Sun, Angew. Chem., 2010, 122, 9118–9121. 19 Q. Yin, J. F. Tan, C. Besson, Y. V. Geletii, D. G. Musaev, A. E. Kuznetsov, Z. Luo, K. I. Hardcastle and C. L. Hill, Science, 2010, 328, 342–345.


Paper

20 N. D. McDaniel, F. J. Coughlin, L. L. Tinker and S. Bernhard, J. Am. Chem. Soc., 2008, 130, 210–217. 21 J. D. Blakemore, N. D. Schley, D. Balcells, J. F. Hull, G. W. Olack, C. D. Incarvito, O. Eisenstein, G. W. Brudvig and R. H. Crabtree, J. Am. Chem. Soc., 2010, 132, 16017–16029. 22 T. P. Brewster, J. D. Blakemore, N. D. Schley, C. D. Incarvito, N. Hazari, G. W. Brudvig and R. H. Crabtree, Organometallics, 2011, 30, 965–973. 23 U. Hintermair, S. M. Hashmi, M. Elimelech and R. H. Crabtree, J. Am. Chem. Soc., 2012, 134, 9785–9795. 24 U. Hintermair, S. W. Sheehan, A. R. Parent, D. H. Ess, D. T. Richens, P. H. Vaccaro, G. W. Brudvig and R. H. Crabtree, J. Am. Chem. Soc., 2013, 135, 10837–10851. 25 J. F. Hull, D. Balcells, J. D. Blakemore, C. D. Incarvito, O. Eisenstein, G. W. Brudvig and R. H. Crabtree, J. Am. Chem. Soc., 2009, 131, 8730–8731. 26 A. Bucci, A. Savini, L. Rocchigiani, C. Zuccaccia, S. Rizzato, A. Albinati, A. Llobet and A. Macchioni, Organometallics, 2012, 31, 8071–8074. 27 A. Savini, P. Belanzoni, G. Bellachioma, C. Zuccaccia, D. Zuccaccia and A. Macchioni, Green Chem., 2011, 13, 3360–3374. 28 A. Savini, G. Bellachioma, G. Ciancaleoni, C. Zuccaccia, D. Zuccaccia and A. Macchioni, Chem. Commun., 2010, 46, 9218–9219. 29 A. Savini, A. Bucci, G. Bellachioma, L. Rocchigiani, C. Zuccaccia, A. Llobet and A. Macchioni, Eur. J. Inorg. Chem., 2014, 690–697. õ, L. Rocchigiani, A. Savini 30 C. Zuccaccia, G. Bellachioma, S. Bolan and A. Macchioni, Eur. J. Inorg. Chem., 2012, 1462–1468. 31 D. B. Grotjahn, D. B. Brown, J. K. Martin, D. C. Marelius, M.-C. Abadjian, H. N. Tran, G. Kalyuzhny, K. S. Vecchio, Z. G. Specht, S. A. Cortes-Llamas, V. Miranda-Soto, C. van Niekerk, C. E. Moore and A. L. Rheingold, J. Am. Chem. Soc., 2011, 133, 19024–19027. 32 D. Hong, M. Murakami, Y. Yamada and S. Fukuzumi, Energy Environ. Sci., 2012, 5, 5708–5716. 33 H. Junge, N. Marquet, A. Kammer, S. Denurra, M. Bauer, ¨rtner, M.-M. Pohl, A. Spannenberg, S. Wohlrab, F. Ga S. Gladiali and M. Beller, Chem.–Eur. J., 2012, 18, 12749–12758. 34 A. Harriman, I. J. Pickering, J. M. Thomas and P. A. Christensen, J. Chem. Soc., Faraday Trans., 1988, 84, 2795–2806. 35 S. J. Kwon, F.-R. F. Fan and A. J. Bard, J. Am. Chem. Soc., 2010, 132, 13165–13167. 36 G. S. Nahor, P. Hapiot, P. Neta and A. Harriman, J. Phys. Chem., 1991, 95, 616–621. 37 T. Nakagawa, N. S. Bjorge and R. W. Murray, J. Am. Chem. Soc., 2009, 131, 15578–15579. 38 R. Nakamura and H. Frei, J. Am. Chem. Soc., 2006, 128, 10668–10669. 39 Y. Zhao, E. A. Hernandez-Pagan, N. M. Vargas-Barbosa, J. L. Dysart and T. E. Mallouk, J. Phys. Chem. Lett., 2011, 2, 402–406. 40 J. D. Blakemore, N. D. Schley, M. N. Kushner-Lenhoff, A. M. Winter, F. D’Souza, R. H. Crabtree and G. W. Brudvig, Inorg. Chem., 2012, 51, 7749–7763.


View Article Online


Paper

41 N. D. Schley, J. D. Blakemore, N. K. Subbaiyan, C. D. Incarvito, F. D’Souza, R. H. Crabtree and G. W. Brudvig, J. Am. Chem. Soc., 2011, 133, 10473–10481. 42 E. Wadsworth, F. R. Duke and C. A. Goetz, Anal. Chem., 1957, 29, 1824–1825. 43 A. R. Parent, T. P. Brewster, W. De Wolf, R. H. Crabtree and G. W. Brudvig, Inorg. Chem., 2012, 51, 6147–6152. `, J. M. S. Cardoso, B. Royo, M. Costas and J. Lloret44 Z. Codola Fillol, Chem.–Eur. J., 2013, 19, 7203–7213. 45 J. DePasquale, I. Nieto, L. E. Reuther, C. J. Herbst-Gervasoni, J. J. Paul, V. Mochalin, M. Zeller, C. M. Thomas, A. W. Addison and E. T. Papish, Inorg. Chem., 2013, 52, 9175–9183. 46 T. Zhang, K. E. deKrafft, J.-L. Wang, C. Wang and W. Lin, Eur. J. Inorg. Chem., 2014, 698–707. 47 Y. Himeda, Eur. J. Inorg. Chem., 2007, 3927–3941. 48 Y. Himeda, N. Onozawa-Komatsuzaki, H. Sugihara, H. Arakawa and K. Kasuga, Organometallics, 2004, 23, 1480–1483. 49 Y. Himeda, N. Onozawa-Komatsuzaki, H. Sugihara and K. Kasuga, J. Am. Chem. Soc., 2005, 127, 13118–13119. 50 Y. Himeda, N. Onozawa-Komatsuzaki, H. Sugihara and K. Kasuga, J. Photochem. Photobiol., A, 2006, 182, 306–309. 51 W.-H. Wang, J. F. Hull, J. T. Muckerman, E. Fujita and Y. Himeda, Energy Environ. Sci., 2012, 5, 7923–7926. 52 W.-H. Wang, J. T. Muckerman, E. Fujita and Y. Himeda, New J. Chem., 2013, 37, 1860–1866. 53 W.-H. Wang, J. T. Muckerman, E. Fujita and Y. Himeda, ACS Catal., 2013, 3, 856–860. 54 J. F. Hull, Y. Himeda, W.-H. Wang, B. Hashiguchi, R. Periana, D. J. Szalda, J. T. Muckerman and E. Fujita, Nat. Chem., 2012, 4, 383–388. 55 S. Ogo, R. Kabe, H. Hayashi, R. Harada and S. Fukuzumi, Dalton Trans., 2006, 4657–4663.


PCCP

56 M. R. DeFelippis, C. P. Murthy, M. Faraggi and M. H. Klapper, Biochemistry, 1989, 28, 4847–4853. 57 R. Szentrimay, P. Yeh and T. Kuwana, Electrochemical Studies of Biological Systems, American Chemical Society, Washington, D. C, 1977, vol. 38, pp. 143–169. ¨ning, J. R. Byberg and K. Sehested, J. Phys. Chem. A, 58 U. K. Kla 1999, 103, 5055–5060. From that ref. one-electron reduction potential of IO4 was estimated to be 0.3–0.45 V vs. NHE at pH 7. 59 L. Duan, Y. Xu, L. Tong and L. Sun, ChemSusChem, 2011, 4, 238–244. 60 L. Duan, Y. Xu, P. Zhang, M. Wang and L. Sun, Inorg. Chem., 2010, 49, 209–215. 61 Y. V. Geletii, Z. Huang, Y. Hou, D. G. Musaev, T. Lian and C. L. Hill, J. Am. Chem. Soc., 2009, 131, 7522–7523. 62 A. Lewandowska-Andralojc and D. E. Polyansky, J. Phys. Chem. A, 2013, 117, 10311–10319. 63 A. Lewandowska-Andralojc, D. E. Polyansky, R. Zong, R. P. Thummel and E. Fujita, Phys. Chem. Chem. Phys., 2013, 15, 14058–14068. 64 D. G. H. Hetterscheid and J. N. H. Reek, Eur. J. Inorg. Chem., 2014, 742–749. 65 K. Beckmann, H. Uchtenhagen, G. Berggren, M. F. Anderlund, A. Thapper, J. Messinger, S. Styring and P. Kurz, Energy Environ. Sci., 2008, 1, 668–676. 66 D. G. H. Hetterscheid and J. N. H. Reek, Chem. Commun., 2011, 47, 2712–2714. ¨rtner, S. Losse, M.-M. Pohl, H. Junge and 67 N. Marquet, F. Ga M. Beller, ChemSusChem, 2011, 4, 1598–1600. 68 A. R. Parent, J. D. Blakemore, G. W. Brudvig and R. H. Crabtree, Chem. Commun., 2011, 47, 11745–11747. 69 S. E. Castillo-Blum, D. T. Richens and A. G. Sykes, Inorg. Chem., 1989, 28, 954–960. 70 A. R. Parent, R. H. Crabtree and G. W. Brudvig, Chem. Soc. Rev., 2013, 42, 2247–2252.


Modes of activation of organometallic iridium complexes for catalytic water and C-H oxidation.

Iridium-based double perovskites for efficient water oxidation in acid media.

Photochemical water oxidation by cyclometalated iridium(iii) complexes: a mechanistic insight.

Air- and water-tolerant rare earth guanidinium BINOLate complexes as practical precatalysts in multifunctional asymmetric catalysis.

Iridium oxide-polymer nanocomposite electrode materials for water oxidation.

Iridium(III) complexes with phenyl-tetrazoles as cyclometalating ligands.

Efficient water oxidation catalyzed by mononuclear ruthenium(II) complexes incorporating Schiff base ligands.

Highly efficient and robust molecular water oxidation catalysts based on ruthenium complexes.

Benchmarking Water Oxidation Catalysts Based on Iridium Complexes: Clues and Doubts on the Nature of Active Species.

Phase transformation fabrication of a Cu2S nanoplate as an efficient catalyst for water oxidation with glycine.

Underestimated potential of organometallic rhenium complexes as anticancer agents.

Highly efficient electrochemiluminescence from iridium(III) complexes with 2-phenylquinoline ligand.

Luminogenic iridium azide complexes.

Cyclometalated iridium(III) complexes with deoxyribose substituents.

Chemiluminescence detection with water-soluble iridium(III) complexes containing a sulfonate-functionalised ancillary ligand.

Carboxylate-rich hybrid ligands in Mn(II) complexes as precursors for water oxidation reactions.

Efficient water oxidation using nanostructured α-nickel-hydroxide as an electrocatalyst.

Coordination tuning of cobalt phosphates towards efficient water oxidation catalyst.

CoO x nanoparticle anchored on sulfonated-graphite as efficient water oxidation catalyst.

Mononuclear ruthenium polypyridine complexes that catalyze water oxidation.

Nickel-vanadium monolayer double hydroxide for efficient electrochemical water oxidation.

Photochemical water oxidation and origin of nonaqueous uranyl peroxide complexes.

Thin-Layer Fe2TiO5 on Hematite for Efficient Solar Water Oxidation.

CoSe2 Nanobelt composite for electrochemical water oxidation.