Bioinspired Photocatalytic Water Reduction and Oxidation with Earth

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Bioinspired Photocatalytic Water Reduction and Oxidation with Earth-Abundant Metal Catalysts Shunichi Fukuzumi, Dachao Hong, and Yusuke Yamada J. Phys. Chem. Lett., Just Accepted Manuscript • DOI: 10.1021/jz401560x • Publication Date (Web): 26 Sep 2013 Downloaded from http://pubs.acs.org on October 1, 2013

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The Journal of Physical Chemistry Letters

Bioinspired Photocatalytic Water Reduction and Oxidation with Earth-Abundant Metal Catalysts Shunichi Fukuzumi,*,†,‡ Dachao Hong,† and Yusuke Yamada† †

Department of Material and Life Science, Graduate School of Engineering, Osaka

University, ALCA, Japan Science and Technology Agency, Suita, Osaka 565-0871, Japan ‡

Department of Bioinspired Science, Ewha Womans University, Seoul 120-750, Korea

E-mail: [email protected]

1 ACS Paragon Plus Environment


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ABSTRACT: The development of efficient water reduction and oxidation catalysts is the key issue for achieving solar energy conversion to obtain sustainable energy sources in replace of fossil fuels. Platinum group metal (PGM) catalysts have been recognized as active catalysts for both water reduction and oxidation. However, it is highly desired to replace precious and scarce PGM catalysts by earth-abundant metal catalysts for water reduction and oxidation. In the past five years, there has been significant progress in the development of water reduction and oxidation catalysts based on earth-abundant metals such as iron, nickel, copper and manganese, which have been combined with organic photocatalysts. This work describes the state of the art and future challenges in bioinspired photocatalytic water reduction and oxidation with earth-abundant metals. TOC graphic:

Keywords: artificial photosynthesis, ruthenium, nickel, cobalt, iron, copper


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Clean and renewable energy is becoming increasingly more important as the world faces the threat of global warming caused by rapid growth of carbon dioxide (CO2) mission in combustion of fossil fuels. Thus, it is highly desired to make a fuel using sunlight through artificial photosynthesis.1-4 Molecular hydrogen (H2) can be clean and sustainable energy of the next generation, because combustion of H2 produces only water and H2 can be produced by electrolysis of water using solar cells.5-8 There have also been extensive studies on photocatalytic water splitting to produce H2 from water using solar energy, although the efficiency has yet to be much improved for any practical applications.8-11 Platinum group metals (PGM) have been reported to exhibit high catalytic activity for both water reduction and oxidation.12-14 However, the high cost and scarcity of PGM have precluded the large scale production of H2 by electrolysis of water.12-14 Thus, it is quite important to develop earth-abundant metal catalysts for water reduction and oxidation. Nature does not use precious metals such as platinum for the catalytic water reduction and oxidation. Hydrogenases that catalyze the reduction of protons of water as well as the oxidation of H2 are composed of iron and nickel, both of which are earth-abundant.15-18 On the other hand, the oxygen evolving complex (OEC) in photosystem II, which catalyze four-electron oxidation of water to evolve oxygen (O2), are composed of oxo-clusters of earth-abundant manganese and calcium.19,20 Herein we describe rational design for the development of efficient water reduction and oxidation catalysts using earth-abundant metals, inspired by the enzymatic reduction of water with hydrogenases and the oxidation of water with OEC in photosystem II, respectively.

Photocatalytic Water Reduction with Earth-Abundant Metals. In natural systems, catalytic proton reduction is exclusively based on the enzymatic activity of hydrogenases whose active sites contain Fe and/or Ni complexes tethering an Fe-S



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cluster.15-18 Various types of metal complexes of Fe and/or Ni have been synthesized as structural models and investigated as functional models of hydrogenases.21-25 As functional models, not only Fe and/or Ni complexes but metal complexes containing Pt, Ir, Pd, Rh, Ru, Co, etc. have been reported to exhibit catalytic activity for H2 evolution.26-35 For a prominent example, a Ru-Ni heterodinuclear complex has been reported as a functional and structural model of Fe-Ni hydrogenases and the µ-hydride complex as a reactive intermediate was isolated and structurally characterized by the single crystal X-ray analysis.33 Inspired by such a structural and functional model complex of hydrogenases, nanoparticles of Ru, Fe and Ni have been examined as hydrogen-evolution catalysts in a photocatalytic system.34-37 These metals in bulk form have been reported to need larger overpotential for proton reduction compared with that with Pt,38 however, reduction of their size into nanoscale may improve their intrinsic catalysis by quantum size effects, which are typically observed in the catalysis of Au.39 Although Pt nanoparticles (PtNPs) have been the most frequently employed as an H2-evolution catalyst, precious and expensive Pt should be replaced with earth-abundant and less expensive metals for practical applications. Recently, photocatalytic H2 evolution has been successfully achieved by employing 2-phenyl-4-(1-naphthyl)quinolinium ion (QuPh+–NA), dihydronicotinamide adenine dinucleotide (NADH) and ruthenium nanoparticles (RuNPs), iron nanoparticles (FeNPs) or nickel nanoparticles (NiNPs) as an organic photocatalyst, a sacrificial electron donor and an H2-evolution catalyst, respectively.35,37 Scheme 1 shows the chemical structure

Scheme 1. Structure of QuPh+–NA and the overall catalytic cycle for the photocatalytic H2 evolution with QuPh+–NA and metal nanoparticles (MNPs, M = Pt, Ru, Fe or Ni). 4 ACS Paragon Plus Environment

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of QuPh+–NA ion and the overall catalytic cycle for the photocatalytic H2 evolution. An important feature of this photocatalytic system is the absence of an electron mediator such as methylviologen dication, because QuPh+–NA affords the long-lived electron-transfer (ET) state40 upon photoexcitation, which can reduce the H2-evolution catalyst without an electron mediator.35-37 RuNPs and FeNPs were prepared by thermal decomposition of Ru3(CO)12 and Fe(CO)5 in the presence of tri-n-octylamine for RuNPs and oleylamine for FeNPs as capping agents in an organic solvent.35 The RuNPs have a spherical shape with the diameter about 4.1 ± 0.6 nm and the FeNPs contained cubic and spherical particles with the diameter of 12 ± 5 nm as evidenced by transmission electron microscope (TEM) observations.35

An

organic

agent

capping

nanoparticles

was

exchanged

to

polyvinylpyrrolidone (PVP) before the photocatalytic H2-evolution measurements in order to increase the dispersibility in an aqueous solution.35 Figure 1 compares the time profiles of the photocatalytic H2 evolution catalyzed by

Figure 1. Time profiles of H2 evolution under photoirradiation (300 W Xe lamp, λ > 340 nm) of a deaerated mixed solution (2.0 mL) of a phthalate buffer (pH 4.5) and MeCN [1:1 (v/v)] containing QuPh+–NA (0.22 mM) and NADH (1.0 mM) with metal nanoparticles [12.5 mg L-1; RuNPs blue circle; PtNPs, red square and FeNPs, green diamond] at 298 K.35



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PtNPs, RuNPs and FeNPs. The photoirradiation (λ > 340 nm) of a mixed solution (2.0 mL) of a phthalate buffer (pH 4.5) and acetonitrile (MeCN) [1:1 (v/v)] containing NADH (1.0 mM), QuPh+–NA (0.22 mM) and MNPs (12.5 mg L-1; M = Pt, Ru or Fe) resulted in H2 evolution.35 When PtNPs were used as an H2-evolution catalyst (Figure 1, square), highly efficient H2 evolution to yield the stoichiometric amount of H2 based on the amount of NADH in solution (2.0 µmol) was observed within 4 min, in which the H2-evolution rate normalized by the weight concentration of PtNPs (VH2) was 2.2 µmol h-1 mg-1 L.35 Virtually the same VH2 was observed when the PtNPs were replaced by the RuNPs (Figure 1, circle),35 whereas the amount of H2 evolved with FeNPs was less than the stoichiometric amount (Figure 1, green diamonds) with VH2 lower than 0.02 µmol h-1 mg-1 L.35 The slow H2-evolution with FeNPs resulted from slow proton-coupled electron transfer from the photogenerated radical species to FeNPs due to the smaller binding energy of Fe–H (190 kJ mol–1) as compared with that of Pt–H (350 kJ mol–1).34 However, it is certainly required to improve the catalytic reactivity of FeNPs and to elucidate the catalytic mechanism. The main drawback of those metal nanoparticles is lack of stability, partly because of agglomerates formation of the metal nanoparticles capped with organic molecules during the photocatalytic H2 evolution. The RuNPs have been reported to be less stable than PtNPs in the photocatalytic H2 evolution.36 The total amount of evolved H2 normalized by the catalyst weight was 2.0 mol g-Pt-1 in the photocatalytic H2 evolution system using PtNPs (1.0 mg L-1), while that was only 0.5 mol g-Ru-1 when PtNPs was replaced with RuNPs.36 The durability of the RuNPs can be improved by supporting RuNPs on various metal oxides as the suppression of agglomeration of RuNPs.36 Figure 2 compares the time courses of H2 evolution in the photocatalytic H2 evolution using 3 wt% RuNPs supported on various metal oxides. When RuNPs supported on MgO (black diamonds), TiO2 (purple reversed triangle) and CeO2 (green triangles) were


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employed, non-stoichiometric amount of H2 was evolved, indicating that these metal oxides deactivated RuNPs.36 On the other hand, the stoichiometric amount of H2 (2.0 µmol) was evolved by employing Ru/SiO2 and Ru/Al2O3-SiO2 as H2-evolution catalsyts.36 Moreover, VH2 for Ru/SiO2 was 5.3 µmol h-1 mg-Ru-1 L, which is similar to VH2 (5.2 µmol h-1 mg-Ru-1 L) for the H2 evolution system using the same amount of RuNPs.36

Figure 2. Time courses of H2 evolution performed by photoirradiation (300 W Xe lamp, λ > 340 nm) of a deaerated mixed solution (2.0 mL) of a phthalate buffer (pH 4.5) and MeCN [1:1 (v/v)] containing QuPh+–NA (0.88 mM), NADH (1.0 mM) and RuNPs supported on metal oxides (3.0 wt% Ru; 100 mg L-1; Ru/SiO2, red; Ru/Al2O3-SiO2, blue; Ru/CeO2, green; Ru/TiO2, purple and Ru/MgO, black). Adapted with permission from reference 36. The catalytic activity of RuNPs is highly sensitive to the morphology of SiO2 supports.32 SiO2 supports with various morphology can be synthesized by choosing an appropriate directing reagent and reaction conditions.36 On these SiO2 supports, Ru was loaded by the chemical vapor deposition (CVD) method.36 The loading amount of Ru was 1% for each Ru/SiO2 determined by X-ray fluorescence measurements.36 The structures of the Ru/SiO2 catalysts were observed by transmission electron microscopy (TEM) as indicated in Figures 3a-c; Ru/SiO2 with undefined shape (u-SiO2), Ru/mesoporous SiO2 with hexagonally packed array (m-SiO2) and Ru/SiO2 with spherical shape (s-SiO2).36 By



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employing these Ru/SiO2 catalysts as H2-evolution catalysts in the photocatalytic H2-evoluiton system under the same reaction conditions described above, the fastest H2 evolution was observed for Ru/u-SiO2 (Figure 3d, red circles) with VH2 9.1 µmol h-1 mg-Ru-1 L.36 Slower H2 evolution rates were observed for Ru/s-SiO2 (Figure 3d, square) and Ru/m-SiO2 (Figure 3d, triangle) with VH2 of 6.3 and 4.7 µmol h-1 mg-Ru-1 L, respectively.36

Figure 3. TEM images of silica supporting RuNPs in the shape of (a) undefined, (b) hexagonally packed mesoporous and (c) nonporous sphere. (d) Time courses of H2 evolution performed by photoirradiation (300 W Xe lamp, λ > 340 nm) of a deaerated mixed solution (2.0 mL) of a phthalate buffer (pH 4.5) and acetonitrile [1:1 (v/v)] containing QuPh+–NA (0.88 mM), NADH (1.0 mM) and RuNPs supported on SiO2 prepared by the chemical vapor deposition method (1 wt%; 100 mg L-1; Ru/SiO2 in the shape of undefined shape, red circles; mesoporous SiO2, green triangles and spherical SiO2, blue squares). Reprinted with permission from reference 36.


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VH2 for Ru/m-SiO2 was only half VH2 for Ru/u-SiO2.36 This suggested that electron injection from QuPh•–NA to RuNPs supported inside mesopores hardly occurs due to a steric reason.36 In the case of Ru/s-SiO2, the fused particles observed by TEM (Figure 3c) suggested that some RuNPs located between particles, resulting in the reduction of effective surface area for H2 evolution.36 The RuNPs between two large SiO2 spheres can also hardly interact with QuPh•–NA.36 Thus, morphology of SiO2 supports significantly affects the catalytic activity of Ru/SiO2.36 The robustness of Ru/u-SiO2 was compared with that of PtNPs capped with PVP in the photocatalytic H2 evolution, which was performed by photoirradiation (λ > 340 nm) of a deaerated mixed solution (2.0 mL) of a phthalate buffer (pH 4.5) and acetonitrile [1:1 (v/v)]

Figure 4. Amount of H2 evolved in the repetitive photocatalytic H2 evolution using Ru/u-SiO2 (red, left) and PtNPs (blue, right). The photocatalytic H2-evolution was performed by photoirradiation (300 W Xe lamp, λ > 340 nm) of a deaerated mixed solution (2.0 mL) of a phthalate buffer (pH 4.5) and acetonitrile [1:1 (v/v)] containing QuPh+–NA (0.88 mM), NADH (1.0 mM) and Ru/u-SiO2 or PtNPs [1.0 mg-M L-1 (M = Pt or Ru)]. A mixed solution containing NADH was added to the reaction solution after each run.32 The initial NADH concentration at each run was fixed to 1.0 mM. Reprinted with permission from reference 36. 9 ACS Paragon Plus Environment


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containing QuPh+–NA (0.88 mM), NADH (1.0 mM) and Ru/u-SiO2 or PtNPs [1.0 mg-M L-1 (M = Pt or Ru)].36 A mixed solution containing NADH was added to the reaction solution to make initial concentraion of NADH in the reaction solution 1.0 mM after each run. When these catalysts were repeatedly used, H2 evolution was observed up to 4th cycle for Ru/u-SiO2 and 5th cycle for PtNPs as shown in Figure 4.36 PtNPs exhibited higher stability than RuNPs, however, aggregation of PtNPs after the 5th cycle was confirmed by dynamic light scattering (DLS) measurements.36 The total amount of evolved H2 normalized by the Pt weight was 2.0 mol g-Pt-1, which is only slightly higher than the amount for Ru/u-SiO2 (1.7 mol g-Ru-1).36 When RuNPs was used as the H2 evolution catalyst instead of Ru/u-SiO2 under the same reaction conditions, the total amount of evolved H2 normalized by the Ru weight was only 0.5 mol g-Ru-1. Thus, the durability of Ru/u-SiO2 becomes closer to that of PtNPs by supporting on u-SiO2 support.36 Although Ru is not an earth-abundant metal, the main component of Ru/SiO2 is SiO2 which is a most earth-abundant material. Another candidate to be employed as an H2-evolution catalyst is nickel metal nanoparticles (NiNPs). Nickel is much earth-abundant metal compared with precious Pt and Ru. Various NiNPs were prepared by thermal decomposition of nickel acetylacetonate complexes in the presence of a capping reagent.37 The crystal structure of NiNPs can be controlled to hexagonal closed packing (hcp) or face-centered cubic (fcc) structure by choosing appropriate preparation conditions.37 A capping reagent of as-prepared NiNPs was replaced with PVP before the catalysis examination to increase the dispersibility in water.37 The photocatalytic H2 evolution was conducted under photoirradiation of a mixed solution (2.0 mL) of a deaerated buffer (pH 4.5) and MeCN [1:1 (v/v)] containing QuPh+–NA, NADH and NiNPs (12.5 mg L-1) with different crystal structures (hcp and fcc) and sizes (6.6 – 210 nm).37 With hcp-NiNPs, a nearly stoichiometric amount of H2 based on the amount of NADH in the solution (2.0 µmol) was evolved by employing smaller NiNPs


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with the size of 6.6 nm or 11 nm as H2-evolution catalysts (closed and open circles in Figure 5a). The amount of evolved H2 was decreased to ~80% of the stoichiometric amount when larger hcp-NiNPs with the size of 40 nm (closed square in Figure 5a) was employed as the H2 evolution catalyst. Only 50% of the stoichiometric amount of H2 was evolved with the largest NiNPs with the size of 210 nm in 20 min.37 Photocatalytic H2 evolution with fcc-NiNPs as an H2-evolution catalyst was also investigated by photoirradiation of a

(a)

(b)

(c)

VH2, µmol h-1 mg-1 L

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Figure 5 (a, b) Time courses of H2 evolution under photoirradiation (300 W Xe lamp, λ > 340 nm) of mixed solutions (2.0 mL) of deaerated phthalate buffer (pH 4.5) and MeCN [1:1 (v/v)] containing NADH (1.0 mM), QuPh+–NA (0.44 mM), and (a) hcp-NiNPs and (b) fcc-NiNPs with different sizes (12.5 mg L–1) at 298 K. (c) Plots of H2-evolution rates with NiNPs (hcp, square ; fcc, circle) vs the size of NiNPs. Reproduced from reference 37 with permission from The Royal Society of Chemistry.



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mixed solution (2.0 mL) of a phthalate buffer (pH 4.5) and MeCN [1:1 (v/v)] containing NADH (1.0 mM) and QuPh+–NA (0.44 mM).37 Figure 5b shows the time courses of H2 evolution with fcc-NiNPs (12.5 mg L-1) at 298 K.37 The fastest H2 evolution was observed with fcc-NiNPs (16 nm), whereas the rate became the slowest with fcc-NiNPs (80 nm).37 Figure 5c plots VH2 for hcp-NiNPs or fcc-NiNPs against the size of NiNPs, indicating that catalytic activity of NiNPs depends on both the crystal structure and the particles size. The smallest hcp-NiNPs with the size of 6.6 nm exhibited the highest VH2 of 1.1 µmol h-1 mg-1 L, which is 40% of that obtained with PtNPs under the same reaction conditions.37 NADH has been proved to act as a two-electron reductant suitable for mechanistic study,41 however, its high cost and low stability under ambient conditions are obstacles for practical applications. In photosynthesis, carbon dioxide and water are converted to carbohydrates by utilizing solar energy.42 Some carbohydrates are converted to oxalate, which is used as a reductant of O2 to produce H2O2 at the active center of oxalate oxidase. Oxalate and its conjugate acid can act as a two-electron donor to produce two equiv of CO2 and H2 (eq 1). Because CO2 evolved with H2 in eq 1 was that fixed by photosynthesis, (COOH)2 → 2 CO2 + H2

(1)

oxalate is a carbon-neutral electron donor. In addition, one-electron oxidized species of oxalate irreversibly undergoes decarboxylation, which suppresses the back electron transfer and facilitates the photocatalytic reaction. Thus, oxalate was used instead of NADH in the photocatalytic H2-evolution.42 Figure 6 shows the time course of H2 evolution under photoirradiation of a deaerated mixed solution (2.0 mL) of a phosphate buffer (pH 6.0) and MeCN [1:1 (v/v)] containing QuPh+–NA (0.22 mM), oxalate (3.0 mM) and Pt, Ru or NiNPs (12.5 mg L-1).42 Among these H2-evolution catalysts, PtNPs showed the highest activity with the H2-evolution rate of 0.18 µmol h-1 g-Pt-1 L, which is determined from the initial (1 h) slope of Figure 6, and


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Figure 6. Time course of H2 evolution under photoirradiation (300 W Xe lamp, λ > 340 nm) of a deaerated mixed solution (2.0 mL) of a phosphate buffer (pH 6.0) and acetonitrile [1:1 (v/v)] containing QuPh+–NA (0.22 mM), oxalate (3.0 mM) and Pt, Ru or Ni nanoparticles (12.5 mg L-1). Reproduced from reference 42 with permission from the PCCP Owner Societies. the amount of evolved H2 reached 4.8 µmol, which is 80% of the stoichiometric amount based on the amount of oxalate in the solution (6.0 µmol), in which oxalate acts as a two-electron donor (eq 1).42 When RuNPs were used as an H2-evolution catalyst, the H2-evolution rate was 0.11 µmol h-1 mg-Ru-1 L.42 NiNPs were less active than PtNPs and RuNPs.35 However, judging from the amount of Ni reserve, which is four-orders of magnitude more abundant than Pt,42 the catalytic reactivity of NiNPs with the initial H2-evolution rate of 0.056 µmol h-1 mg-Ni-1 L, which corresponds to the value of 32 % of PtNPs, is quite remarkable.42 The photocatalytic H2-evolution systems described above uses organic solvent to dissolve and isolate QuPh+-NA ion, because back electron transfer of photogenerated QuPh•–NA•+ occurs via intermolecular reactions.43 For the QuPh•–NA•+ ion, the intramolecular back electron transfer is too slow to compete with the intermolecular back electron transfer.43 The lifetime of the electron-transfer state in bulk solid state becomes shorter than that in solution due to strong intermolecular interaction in the solid state,



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indicating that a poor solvent of QuPh+-NA ion, such as pure water, is hardly used for the photocatalytic H2-evolution.44,45 However, supporting the QuPh+-NA on nanosized mesoporous silica-alumina (AlMCM-41) by cation exchange enables the photocatalytic H2 evolution in water owing to the suppression of intermolecular interaction between QuPh•–NA•+ ions on the surfaces of AlMCM-41 and high dispersibility of the AlMCM-41 in water.45 A problem arisen in this reaction system is incorporation of H2-evolution catalysts. Metal nanoparticles used in the reported photocatalytic H2-evolution systems cannot be employed due to their larger size than the window size of the mesoporous silica-alumina. A possible solution for this issue is in situ production of H2-evolution catalysts by the reduction of metal ions precursors in close proximity of the donor-acceptor linked molecules (Scheme 2). With this preparation method, not only Pt but also Cu which is more earth abundant can act as an H2-evolution catalyst (unpublished results).

Scheme 2. Overall photocatalytic cycle of H2-evolution system using Pt or Cu catalyst by in situ reduction of metal ion precursor.45 Figure 7a shows the TEM image of Na+-exchanged spherical AlMCM-41 (sAlMCM-41). The sizes of the particles are in the range between 200 nm to 700 nm.45 The pore size determined by the N2-isotherm measurements was around 1.8 nm.44,45 QuPh+-NA was adsorbed on the sAlMCM-41 by ion exchange with Na+ in a mixed solution of water and MeCN.45 The adsorption of QuPh+-NA on sAlMCM-41 was confirmed by the UV-vis absorption spectra as indicated in Figure 7b.45 The characteristic absorption peaks observed 14 ACS Paragon Plus Environment

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(a)

(b)

1 µm

Figure 7. (a) Transmission electron microscope (TEM) image of Na+-exchanged sAlMCM-41. (b) UV-vis absorption spectra of QuPh+–NA in MeCN (black dashed line) and DRS spectra of QuPh+–NA@sAlMCM-41(red solid line). Reproduced from reference 45 with permission from The Royal Society of Chemistry.

in MeCN solution was also observed in diffused reflectance spectra (DRS) of QuPh+-NA/sAlMCM-41.45 H2 evolution was examined under photoirradiation (λ > 340 nm) of a phthalate buffer dispersion (2.0 mL) containing QuPh+–NA@sAlMCM-41 (5.0 mg, QuPh+-NA: 0.22 mM), sodium oxalate (50 mM) and K2PtCl6 (0.05 mM) as a photocatalyst, a sacrificial electron donor and the precursor of an H2-evolution catalyst, respectively (unpublished results). Under photoirradiation, continuous H2 evolution was observed as shown in Figure 8 (red circle). The total amount of evolved H2 reached 8.0 µmol in 12 h with an H2-evolution rate of 0.85 µmol h-1 for initial 6 h. An apparent VH2 normalized by the weight of Pt in the reaction solution was 0.068 µmol h-1 mg-Pt-1 L. Similarly, when Cu(NO3)2 was used instead of K2PtCl6 in the photocatalytic H2-evolution system, significant H2 evolution was observed although

the

H2-evolution

rate

(0.11

µmol

h-1)

with

the

Cu-deposited

QuPh+–NA/sAlMCM-41 was rather modest as compared with the rate (0.85 µmol h-1) of the Pt-deposited QuPh+-NA/sAlMCM-41 (Figure 8 (square)). With this preparation method, Cu can also act as an H2-eveolution catalyst.



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Time, h

Figure 8. Time courses of H2 evolution under photoirradiation (300 W Xe lamp, λ > 340 nm) of a deaerated dispersion (2.0 mL) of a phthalate buffer (pH 4.5) containing QuPh+–NA@sAlMCM-41 (QuPh+–NA: 0.22 mM) and oxalate (50 mM) with K2PtCl6 (red circle) or Cu(NO3)2 (blue square) (unpublished results).

Photocatalytic Water Oxidation with Earth-Abundant Metals. As compared with catalytic water reduction described above, water oxidation is a more challenging process. The process involves two water molecules and four-electron transfer coupled with four-proton transfer (eq 2). In this process, O-O bond formation is thought to be the 2H2O = O2 + 4e– + 4H+

Eo = + 1.223 V vs NHE

(2)

rate-determining step, however, the reaction mechanism how O-O bond forms between two water moelcules has yet to be fully understood. More efficient and robust water oxidation catalysts (WOCs), which can facilitate the O-O bond formation, are certainly required to develop an artificial photosynthesis system. There have been extensive studies on water oxidation using homogeneous and heterogeneous WOCs.46-50 Homogeneous WOCs have advantage to elucidate the detailed catalytic mechanisms including detection of active intermediates for water oxidation. Recently, there has been a critical concern if the homogeneous WOCs maintain their structures during the water


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oxidation, especially when homogeneous metal complexes have organic ligands. Generally, oxidants used in water oxidation possess the ability to oxidize organic ligands of metal complexes, resulting in formation of metal oxide or metal hydroxide particles.51,52 A series of mononuclear iridium complexes, [IrIII(Cp*)(4,4’-R2-2,2’-bpy)(H2O)]SO4 (R = OH, OMe, Me and COOH, bpy = 2,2’-bipyridine, Cp* = η5-pentamethylcyclopentadienyl), have been reported to convert to Ir(OH)3 nanoparticles, which act as the true catalysts for the water oxidation by cerium(IV) ammonium nitrate (CAN).52 Even metal complexes with inorganic ligands such as polyoxometalate, which has high durability against oxidative stress, would elute off metal ions to form metal oxides under several limited conditions.53 Thus, careful examination would be required to analyze kinetics and to detect intermediates for molecular metal complexes in water oxidation. As compared with homogeneous WOCs, heterogeneous WOCs have advantage for practical applications because of high catalytic activity and easy separation of catalysts by filtration. Many metal oxides such as iridium, ruthenium, rhodium, cobalt and manganese oxides were reported to exhibit catalytic activity for the photocatalytic water oxidation using [Ru(bpy)3]2+ as photosensitizer and Na2S2O8 as an electron acceptor.48 Among the metal oxides, iridium oxides have exhibited the highest catalytic activity, however, iridium is a precious and expensive metal.48 Thus, efficient WOCs based on earth-abundant metals, especially the first-row transition metals such as cobalt, nickel, manganese and iron, are highly desired for practical applications. Among them cobalt-based WOCs have attracted much attention over the years due to their high catalytic activity for water oxidation.54 Various methods, such as structural controls of size, shape or phase, have been studied to improve the catalytic activity of cobalt oxide for water oxidation.54,55 For example, nano-sized Co3O4 loaded on mesoporous silica exhibited higher activity than micron-sized Co3O4 particles.54 A cobalt cluster with cubic Co4O4 core modified by organic ligands can



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act as a catalyst for the photocatalytic water oxidation.55 Recently, cobalt-based nanoparticles produced from cobalt complexes with organic ligands have been reported to show high catalytic activity for water oxidation.56 The organic ligands of the complex decomposed during the reaction play a unique role to improve the catalytic activity by preventing the generated nanoparticles from aggregation.56 The cobalt-based nanoparticles were prepared by photoirradiation (λ > 420 nm) of a buffer solution containing [Ru(bpy)3]2+ (0.50 mM), Na2S2O8 (10 mM) and a cobalt complex ([CoII(Me6tren)(OH2)]2+ (1): Me6tren = tris(N,N’-dimethylaminoethyl)amine; 2.5 mM) as shown in Figure 9a.56 No O2 was evolved with 1 (2.5 mM) in the buffer solution (1st run),

Figure 9. (a) A cobalt complex (1) used as a precatalyst in photocatalytic water oxidation. (b) Time courses of O2 evolution under photoirradiation (300 W Xe lamp, λ > 420 nm) of a buffer solution (2.0 mL, 100 mM borate, pH 9.0) containing [Ru(bpy)3]2+ (0.50 mM) and Na2S2O8 (10 mM) (a) with 1 (2.5 mM) (1st run), (b) with nanoparticles (~0.12 mg) derived from 1 in a fresh solution (2nd run) and (c) with the nanoparticles collected after 2nd run with another fresh solution (3rd run).56

although CO2 evolution was observed in the reaction (Figure 9b). This indicates that the oxidation of the organic ligands of 1 proceeds prior to water oxidation.56 Particles generated at the 1st run were separated from the reaction solution by centrifugation.56 The particles 18 ACS Paragon Plus Environment

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collected from the solution exhibited catalytic activity in the photocatalytic water oxidation (2nd and 3rd runs).56 The catalytic cycle of photocatalytic water oxidation is shown in Scheme 3. Photoinduced electron transfer from the excited state of [Ru(bpy)3]2+ ([Ru(bpy)3]2+*:

*

denotes the excited state) to S2O82– affords [Ru(bpy)3]3+, SO42– and SO4•–. The produced SO4•– can oxidize another [Ru(bpy)3]2+ to produce two equiv of [Ru(bpy)3]3+, which can oxidize water in the presence of a WOC to evolve O2.56 When the collected particles (~0.12 mg) were employed as a WOC, efficient O2 evolution was observed (2nd run in Figure 9b).56 The robustness of the collected particles was confirmed by using them at further reaction cycles.56 A significant amount of O2 evolution was still obtained in the 3rd run with the nanoparticles collected from the reaction solution of 2nd run.56 These results clearly demonstrate that the particles prepared from 1 exhibit high catalytic activity for the photocatalytic water oxidation.56

Scheme 3. Photocatalytic cycle of water oxidation with S2O82– and [Ru(bpy)3]2+ by using water oxidation catalysts (WOCs).

The particles derived from 1 were characterized by TEM, powder X-ray diffraction, X-ray photoelectron spectroscopy and thermogravimetry/differential thermal analysis (TG/DTA) measurements, which reveal that the particles were composed of Co(OH)x nanoparticles and carbonaceous residues.56 TG/DTA measurements suggested that the content of carbonaceous residue in the particles derived from 1 was 14 wt%.56 The



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remaining carbonaceous residues in the Co(OH)x nanoparticles can prevent particle aggregation during the photocatalytic water oxidation as demonstrated by the comparison with particles prepared from Co(NO3)2.56 The catalytic activity of Co(OH)x nanoparticles in the photocatalytic water oxidation was compared to that of particles prepared from Co(NO3)2 without organic materials (Figure 10a).56 The Co(OH)x nanoparticles with carbonaceous residues exhibited higher catalytic activity than that of particles prepared from Co(NO3)2 for the photocatalytic water oxidation. The O2 yield of the reaction with Co(NO3)2 in the 2nd run was only 10%, which is significantly lower than the O2 yield of 22% in the 1st run, whereas the O2 yield of the reaction with Co(OH)x in the 2nd run was 54%, which is similar to that of 1st run (61 %).56 The O2 yields (2[O2]/[Na2S2O8]) were determined based on the amount of Na2S2O8 in the reaction solutions.56 The decrease of O2

Figure 10. (a) Time courses of O2 evolution under photoirradiation (Xe lamp, λ > 420 nm) of a buffer solution (pH 10, 2.0 mL) containing [Ru(bpy)3]2+ (0.50 mM) and Na2S2O8 (10 mM) with a precatalyst (i and ii) 1 (50 µM) and (iii and iv) Co(NO3)2 (50 µM). 2nd runs were performed by adding Na2S2O8 (5.0 µmol) to the solutions after 1st run. (b) Particles size and their distribution determined by DLS measurements of particles derived from 1 and (c) Co(NO3)2. Particles formed by photoirradiation (300 W Xe lamp, λ > 420 nm) of a buffer solution (2.0 mL, 100 mM borate, pH 9.0) containing 1 or Co(NO3)2 (50 µM), [Ru(bpy)3]2+ (0.50 mM) and Na2S2O8 (10 mM) for 3 min (black solid line), 10 min (blue broken line) and 30 min (red dotted line).56


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yield with Co(NO3)2 may be ascribed to particle aggregation observed by DLS measurements as shown in Figures 10b and 10c, where the size of particles derived from Co(NO3)2 increased during the photoirradiation, whereas no increase in the particle size was observed with Co(OH)x.56 It seems beneficial to keep the particles size smaller for maintaining the high surface area. These results indicate that Co(OH)x nanoparticles derived from 1 exhibit high activity and durability due to the presence of carbonaceous residues.56 Another method to improve catalytic activity was reported by doping foreign metal ions to cobalt oxides as observed in the manganese-calcium cluster in OEC.57 The catalytic activity of La3+ doped cobalt oxide was examined in the photocatalytic water oxidation by Na2S2O8 with [Ru(bpy)3]2+ as well as the thermal water oxidation with [Ru(bpy)3]3+.57 The concentration of [Ru(bpy)3]3+ decreased even in the absence of a catalyst, however, the

Figure 11. (a) Time courses of absorbance decay at 670 nm due to [Ru(bpy)3]3+ (pH 7.0, 0.20 mM) in the absence and presence of water oxidation catalysts (63 mg L-1; no catalyst, black; La0.7Sr0.3CoO3, blue; LaCoO3, red) at 298 K. (b) Time courses of O2 evolution under visible light irradiation (300 W Xe lamp, λ > 420 nm) of a phosphate buffer solution (50 mM, 2 mL, pH 7.0) containing [Ru(bpy)3]2+ (0.25 mM), Na2S2O8 (5.0 mM) with a cobalt-containing catalyst (0.25 g L-1, Co3O4, LaCoO3, La0.7Sr0.3CoO3 and CoWO4). Reproduced from reference 57 with permission from the PCCP Owner Societies. 21 ACS Paragon Plus Environment


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decay rate was much faster in the presence of LaCoO3 and La0.7Sr0.3CoO3 as shown in Figure 11a.57 The faster decay observed for LaCoO3 and La0.7Sr0.3CoO3 compared with no catalyst conditions indicates that LaCoO3 and La0.7Sr0.3CoO3 can quench [Ru(bpy)3]3+. O2 evolution observed in Figure 11b demonstrated LaCoO3 and La0.7Sr0.3CoO3 can act as catalysts for the photocatalytic water oxidation. The O2 yields in the photocatalytic water oxidation with LaCoO3, CoWO4, Co3O4 and La0.7Sr0.3CoO3 at pH 7.0 were 74%, 19%, 59% and 47%, respectively (Figure 11b).57 The O2 yield with LaCoO3 was improved to 78% at pH 8.0.57 Thus, LaCoO3 exhibited the highest catalytic activity compared with other cobalt-containing catalysts. The partial replacement of La3+ with Sr2+ allowing the formation of CoIV species in cobalt-containing perovskite led to decrease in catalytic activity.57 On the other hand, the high catalytic activity of LaCoO3 was maintained by the replacement of La3+ with Nd3+ or Y3+, NdCoO3 and YCoO3.57 The trivalent metal ions enhanced the oxidation activity of CoIV species, which should be formed during the water oxidation, by stabilizing the CoIII state.57 Among the first-row transition metals, iron is the most earth-abundant, thus, often used as a catalyst in various oxidation reactions. However, the catalytic activity of iron oxides for the photocatalytic water oxidation was reported to be lower than that of cobalt oxides.48 As described above, doping La3+ ion to cobalt oxides has improved the catalytic activity of cobalt oxides.57 Similarly, the catalytic activity of iron oxides was improved by doping Ni2+ ion to iron oxides (NiFe2O4).58 The catalytic activity of NiFe2O4 was compared with those of Fe3O4, Fe2O3, Co3O4 and NiO in the photocatalytic water oxidation with by Na2S2O8 with [Ru(bpy)3]2+ as shown in Figure 12a, where the O2 yield is the highest for NiFe2O4.58 The catalytic activity of NiFe2O4 was even slightly higher than that of Co3O4. No significant change in the total amount of O2 evolution was observed at 2nd and 3rd run from the reaction solutions.58 This result demonstrates that NiFe2O4 acts as an efficient and


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robust catalyst for the photocatalytic water oxidation.58 The apparent TOF value of NiFe2O4 (0.11 µmol s-1 m–2), which is defined as the TOF value divided by the surface area of a catalyst, was higher than those of Fe3O4 (0.042 µmol s-1 m-2), Fe2O3 (0.065 µmol s-1 m-2) and NiO (0.020 µmol s-1 m-2).58 These results support that the high activity of NiFe2O4 results from the composite effect of nickel and iron oxides. The water oxidation catalysis of NiFe2O4 was also scrutinized under electrocatalytic conditions.58 Figure 12b shows cyclic voltammetry of water using a working electrode modified with a metal oxide catalyst in a buffer solution of pH 8.0. The anodic currents with NiFe2O4 started growing around 0.8 V (vs saturated calomel electrode: SCE) and reached more than 650 µA at 1.5 V, which is larger than those with iron-based oxides.58 The overpotential of NiFe2O4 (η = 0.43 V) for

Figure 12. (a) Time courses of O2 evolution under photoirradiation (300 W Xe lamp, λ > 420 nm) of a phosphate buffer solution (50 mM, 2.0 mL, pH 8.0) containing Na2S2O8 (5.0 mM) and [Ru(bpy)3]SO4 (0.25 mM) with NiFe2O4, Fe3O4, Fe2O3, Co3O4 or NiO (0.50 g L-1) at room temperature in three repetitive examinations. (b) Cyclic voltammograms (CV) in a phosphate buffer solution (50 mM, 2.0 mL, pH 8.0) with a carbon-paste working electrode (gray, surface area = 0.071 cm-2) containing 5% of NiFe2O4, Fe3O4 or Fe2O3, a Pt wire (counter electrode) and a standard calomel electrode (scan rate of 100 mV s-1). Inset shows the initial range of the electrocatalytic current.58 Reprinted from reference 58.



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the electrochemical water oxidation is comparable to the reported overpotential of catalysts such as cobalt phosphate,59 CoOx54 and nickel borate.60 Around 0.8 V, a small redox couple assignable to Ni2+/Ni3+ appeared as reported for the nickel oxide electrode formed on nickel metal in alkaline solution.61 Recently, the oxidized nickel species under anodic potential are assigned to Ni4+ in nickel borate by X-ray absorption near edge structure spectra.60 These results indicate that Ni2+ ion incorporated into iron oxide enhances the water oxidation ability of iron oxides for the electrocatalytic water oxidation, which may be able to exhibit high activity for the photocatalytic water oxidation. The improvement in the catalytic activity of manganese oxides was also observed by doping Ni2+ ion to manganese oxides (NiMnO3) (unpublished results). The use of effective doping of Ni2+ was inspired by manganese-oxo-calcium clusters (Mn4CaO5) in OEC of photosystem II. This has important implications for the exploitation of efficient WOCs with bimetallic metal oxides of earth-abundant metals.

Conclusions and Perspectives. Organic electron-donor acceptor dyads, which afford the long-lived CS states, have been combined with metal nanoparticles acting as water reduction catalysts for the photocatalytic H2 evolution with earth-abundant natural electron donors such as oxalic acid. The efficient catalytic reactivity of hydrogenases in the reversible proton reduction has inspired us to replace precious and scarce platinum by earth-abundant metal catalysts such as Ni and Cu for the photocatalytic H2 evolution. The catalysis improvement in robustness and activity was obtained by supporting Ru metal on earth-abundant metal oxide. Such improvement by the combination of metal and metal oxide, both of which are earth abundant, will be obtained in future. Additionally, shape control of nanoparticles would be a promising method to improve the catalysis, which have only been proved for PtNPs.34 The essential role of Ca2+ in the OEC of photosystem II has also inspired us to develop efficient water oxidation catalysts by using earth-abundant metal oxides mixed with Lewis acid metal ions, which enhance the water oxidation reactivity. Such bioinspired approaches provide promising strategies to develop more efficient


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catalysts based on earth-abundant metals for water reduction and oxidation, both of which should be combined with appropriate organic photocatalysts in future to produce H2 from water. At this stage, direction control of electron flow is very important issue, because both half reactions use irreversibly decomposable an electron donor or acceptor, which are beneficial to drive the photocatalytic reactions to only one-way. Rational alignment of light-harvesting, water reduction and water oxidation units like natural systems is needed to realize the overall photocatalytic water splitting.

Biographies Shunichi Fukuzumi earned Ph.D. degree in applied chemistry at Tokyo Institute of Technology in 1978. He has been a Full Professor of Osaka University since 1994. He is a Specially Distinguished Professor and the leader of a Global COE program, Global Education and Research Centre for Bio-Environmental Chemistry at Osaka University. He is also the director of an ALCA (Advanced Low Carbon Technology Research and Development)

project

of

Japan

Science

Technology

Agency

(JST);

http://www-etchem.mls.eng.osaka-u.ac.jp/ Dachao Hong received his BS degree in applied chemistry at Osaka University in 2010. He is a Ph.D. course student and a JSPS research fellow at Osaka University from 2012. Yusuke Yamada earned his Ph.D. degree in macromolecular science from Osaka University in 1998. He was working as a researcher and senior researcher at Osaka National Research Institute (currently AIST) from 1998 to 2009. From 2007 to 2008, he was a visiting scholar at UC Berkeley. Currently he is an associate professor of Osaka University.

ACKNOWLEDGMENT The authors are most grateful to the collaborators and coworkers who performed the researches reported in this Perspective and support by Grant-in-Aids (Nos. 24350069 and 25600025) from the Ministry of Education, Culture, Sports, Science 25 ACS Paragon Plus Environment


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and Technology, Japan and KOSEF/MEST through WCU project (R31-2008-000-10010-0) and GRL (2010-00353) Program, Korea.

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Quotes to highlight in paper;


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1. Yamada, Y.; Shikano, S.; Fukuzumi, S. Robustness of Ru/SiO2 as a Hydrogen-Evolution Catalyst in a Photocatalytic System Using an Organic Photocatalyst. J. Phys. Chem. C 2013, 117, 13143-13152. 2. Yamada, Y.; Miyahigashi, T.; Kotani, H.; Ohkubo, K.; Fukuzumi, S. Photocatalytic Hydrogen Evolution with Ni Nanoparticles by Using 2-Phenyl-4-(1-Naphthyl)Quinolinium Ion as a Photocatalyst. Energy Environ. Sci. 2012, 5, 6111-6118. 3. Yamada, Y.; Yano, K.; Hong, D.; Fukuzumi, S. LaCoO3 Acting as an Efficient and Robust Catalyst for Photocatalytic Water Oxidation with Persulfate. Phys. Chem. Chem. Phys. 2012, 14, 5753-5760. 4. Hong, D.; Yamada, Y.; Nagatomi, T.; Takai, Y.; Fukuzumi, S. Catalysis of Nickel Ferrite for Photocatalytic Water Oxidation Using [Ru(bpy)3]2+ and S2O82–. J. Am. Chem. Soc. 2012, 134, 19572-19575.


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