Dual Catalytic Cycle of H2 and H2O Oxidations by a Half-Sandwich

Article Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

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Dual Catalytic Cycle of H2 and H2O Oxidations by a Half-Sandwich Iridium Complex: A Theoretical Study Kei Ikeda,† Yuta Hori,‡ Muhammad Haris Mahyuddin,‡ Yoshihito Shiota,‡ Aleksandar Staykov,§ Takahiro Matsumoto,†,§,∥ Kazunari Yoshizawa,*,‡ and Seiji Ogo*,†,§,∥ †

Department of Chemistry and Biochemistry, Graduate School of Engineering, Kyushu University, Fukuoka 819-0395, Japan Institute for Materials Chemistry and Engineering, Kyushu University, Fukuoka 819-0395, Japan § International Institute for Carbon-Neutral Energy Research (WPI-I2CNER), Kyushu University, Fukuoka 819-0395, Japan ∥ Center for Small Molecule Energy, Kyushu University, Fukuoka 819-0395, Japan

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‡

S Supporting Information *

ABSTRACT: While hydrogenase and photosystem II enzymes are known to oxidize H2 and H2O, respectively, a recently reported iridium aqua complex [IrIII(η5-C5Me5){bpy(COOH)2}(H2O)]2+ is able to oxidize both of the molecules and generate energies as in the fuel and solar cells (Ogo et al.ChemCatChem 2017, 9, 4024−4028). To understand the mechanism behind such an interesting bifunctional catalyst, in the present study, we perform density functional theory (DFT) calculations on the dual catalytic cycle of H2 and H2O oxidations by the iridium aqua complex. In the H2 oxidation, we found that the H−H bond is easily cleaved in a heterolytic fashion, and the resultant iridium hydride complex is significantly stabilized by the presence of H2O molecules, due to dihydrogen bond. The rate-determining step of this reaction is found to be the H2O → H2 ligand substitution with an activation energy of 10.7 kcal/mol. In the H2O oxidation, an iridium oxo complex originating from an oxidation of the iridium aqua complex forms a hydroperoxide complex, where an O−O bond is formed with an activation energy of 21.0 kcal/mol. Such a relatively low activation barrier is possible only when at least two H2O molecules are present in the reaction, allowing the water nucleophilic attack (WNA) mechanism to take place. The present study suggests and discusses in detail six reaction steps required for the dual catalytic cycle to complete.

1. INTRODUCTION Renewable energies generated from small molecules such as H2 and H2O are important for addressing the environmental and energy-related issues.1−11 Hydrogenase (H2ase) and photosystem II (PSII) enzymes are known to catalyze the oxidations of H2 and H2O, respectively. The H2ase and its model complexes consist of Ni and/or Ru and/or Fe centers10,12−20 that heterolytically cleave the H−H bond of H2 and generate energy as in the fuel cell.14 In contrast, the PSII consists of Mn and Ca centers that oxidize H2O to O2 by absorbing energy as in the solar cell.21,22 Despite their high catalytic activity, H2ase and PSII have some issues related to enzyme isolation, industrial applications, and catalytic deactivation. Therefore, researchers have attempted to develop alternative, better catalysts.23−35 Iridium (Ir) complexes36−42 are promising candidates for substituting H2ase as the catalyst for H2 oxidation because Ir has a good affinity toward H2 that promotes the heterolytic H−H bond cleavage. Hou et al.39 theoretically demonstrated that the IrIII(η5-C5Me5)bpy(OH)2H2O complex,40 where bpy(OH)2 is 6,6′-dihydroxy-2,2′-bipyridine, interacts with an H2 molecule through an electron donation from the H2-σ © XXXX American Chemical Society

bonding orbital to the Ir-d unoccupied orbital as well as an electron back-donation from the occupied Ir-d orbital to the antibonding H2-σ* orbital. Moreover, they suggested that the OH− groups of the bpy(OH)2 ligand assist the heterolytic cleavage of H−H bond, which then leads to the formation of an iridium hydride complex [IrIII(η5-C5H5) bpy(OH)2H]− with an activation energy of only 13.4 kcal/mol. Different half-sandwich Ir complexes have also been reported to oxidize H2O and show a good performance in stabilizing the high-valent iridium−oxo species that reacts with H2O to form an O−O bond, which is the most important step in the H2O oxidation in water solution.43 Blakemore et al., for instance, reported the formation of an iridium oxo complex leading to an O−O bond formation from the reaction of two H2O molecules with a CpIr(O)(ppy)+ complex, where Cp is cyclopentadienyl and ppy is 2-phenylpyridine.43,44 However, the oxidant used for forming this complex may lead to a catalyst deactivation, due to the reaction between the oxidant and the Cp and Cp* ring.45,46 Recently, Joya et al. succeeded Received: February 1, 2019

A

DOI: 10.1021/acs.inorgchem.9b00307 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

Figure 1. Dual catalytic cycle of H2 and H2O oxidations by iridium aqua complex (AC) proposed by Ogo’s group.48

Figure 2. (a) X-ray crystal and (b) DFT-optimized structures of the Ir aqua complex (AC) in the closed-shell singlet ground state. d denotes the shortest bond distance between the Ir center and the C atoms of the 5-membered ring, highlighted by a red pentagon. Distances are in Å.

Figure 1 shows the dual catalytic cycle of H2 and H2O oxidations proposed by the Ogo’s group.48 In the H2 oxidation cycle, insertion of H2 to AC leads to the formation of an iridium hydride complex Ir III (η 5 -C 5 Me 5 ){bpy(COO)(COOH)}(H) (HC), in which a proton is subsequently released during an oxidation on the carbon electrode in the presence of an H2O molecule. In the H2O oxidation cycle, on the other hand, an iridium oxo complex [IrV(η5-C5Me5){bpy(COOH)2}(O)]+2 (OC) oxidizes two H2O molecules to form an O2 molecule and to regenerate AC. In this reaction, it is assumed that the OC formation and the AC regeneration are triggered by a light irradiation leading to the formation of electron holes on the WO3 electrode. In this study, we employ density functional theory (DFT) to calculate energy diagrams for the dual catalytic cycle of H2 and H2O oxidations by AC and to suggest detailed reaction

in synthesizing an oxidant-free iridium oxo complex that oxidizes H2O on an indium tin oxide (ITO) electrode at mild conditions.47 While the half-sandwich iridium complexes described above are able to catalyze either the H2 or H2O oxidation, a single compound of the recently reported iridium aqua complex [IrIII(η5-C5Me5){bpy(COOH)2}(H2O)]2+ (AC), where bpy(COOH)2 is 2,2′-bipyridine-4,4′-dicarboxylic acid, is able to catalyze both of the molecular oxidations.48 This catalyst, together with a carbon or tungsten oxide (WO3) electrode on the anode and a carbon-supported platinum (Pt/C) electrode on the cathode (the carbon and WO3 electrodes are used for the H2 and H2O oxidations, respectively), creates a system capable of switching functionalities between the fuel and solar cells. Moreover, this Ir complex has a distorted octahedral geometry with a robustness and a stability in a water solution. B


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Inorganic Chemistry

Figure 3. Currently proposed reaction pathways for the dual catalytic cycle of H2 and H2O oxidations by the iridium aqua complex (AC).

Figure 4. Computed energy diagram for the formation of H2C (A1) and HC (A2) in the closed-shell singlet state. Energies (in kcal/mol) include the zero-point energy correction (ZPC). Values in parentheses are the self-consistent field (SCF) energies.

In AC, the Ir−Cp*, Ir−O, Ir−N1, and Ir−N2 bond lengths are calculated to be 2.149, 2.266, 2.122, and 2.123 Å, respectively, which are slightly longer than the corresponding experimental values of 2.141, 2.161, 2.091, and 2.107 Å.48 However, it is widely accepted that these small errors are due to the tendency of DFT methods.38,46,49 Moreover, Kazaryan et al. demonstrated that the B3LYP-D3 method, compared to the CCSD(T) method as a benchmark, is accurate enough for calculating energy diagrams of H2O oxidation over a similar half-sandwich iridium complex.50 In Figure 3, we show a schematic representation of our computational suggestions for the dual catalytic cycle of the H2 and H2O oxidations, which are referred to as Cycle A and Cycle B, respectively. In Cycle A, we suggest the formation of a hydrogen complex (H2C) prior to the HC formation, resulting in two thermal reactions (black lines, A1 and A2) and one electrode reaction (red line, A3). In reaction A1, the H2O ligand of AC is substituted by H2 via the formation of an intermediate structure preceding the H2C formation. Sub-

mechanisms. We rationally show that AC is indeed able to catalyze both the H2 and H2O oxidations with low activation barriers. Also, we discuss the high importance of water molecules in both oxidation reactions. The present study provides a deep mechanistic understanding on the design of dual catalytic systems capable of oxidizing both H2 and H2O molecules.

2. RESULTS AND DISCUSSION Parts a and b of Figure 2 show the structures of AC obtained from the X-ray measurement48 and our DFT calculations, respectively. In the real experimental conditions,48 one H atom of the bpy(COOH)2 moiety is separated during the formation of HC, and the carbon and WO3 electrodes are used to immobilize the iridium complex. However, for simplicity, in this study, we kept the number of atoms in the bpy(COOH)2 moiety the same for all reaction steps and did not consider the effects of the electrodes. C


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Inorganic Chemistry

Figure 5. Optimized structures of (a) TS2, (b) H2C, (c) Int4, (d) TS3, (e) HC-{H+}wat, and (f) HC. In parts c−f, Cp* and bpy(COOH)2 ligands are omitted for clarity. Bond lengths are in Å. The natural populations (charges) of H1 and H2 atoms are also shown.

an activation energy of only 0.7 kcal/mol (see the selfconsistent-field (SCF) energies in the parentheses) or no barrier if we take the zero-point energy correction (ZPC) into consideration, both of which suggest that the reaction readily occurs. This is consistent with the fact that H2C is not observed in experiments.48 In the final step of the reaction, the {H+}wat cluster is removed from HC, which requires a desorption energy of 7.6 kcal/mol (see the ZPC energies). The overall reaction is exothermic by 2.2 kcal/mol where the rate-limiting step is the H2 insertion instead of the H−H bond cleavage, suggesting a plausible reaction pathway under mild conditions and showing that the Ir center has a good affinity for the H2 oxidation. Figure 5 shows the optimized structures of TS2, H2C, Int4, TS3, HC-{H+}wat, and HC. Let us first discuss the changes in the bonding nature of H2 after binding with the iridium center in H2C. From the natural population analysis,51 we found an atomic charge of 0.17 for each of the H1 and H2 atoms. As suggested by Hou et al.,39 this indicates that the H2-σ bonding and H2-σ* antibonding orbitals respectively make donation and back-donation interactions with the Ir-d unoccupied and occupied orbitals, resulting in the H1−H2 bond elongation from 0.744 Å in the gas phase52 to 0.839 Å in H2C. The Ir−H1 and Ir−H2 bond lengths of H2C are found to be the same, i.e., 1.792 Å, suggesting that both of the H atoms have an equivalent ligand field. When the O atom of { }wat cluster binds the H2 atom of H2C to form Int4 with a O···H2 distance of 1.868 Å, the H1− H2 bond is further elongated to 0.873 Å, due to a driving force originated from the O···H2 interaction. In Int4, the natural population of the H1 atom is decreased to 0.13, whereas that of the H2 atom is increased to 0.23, indicating that the H2O fragment of the { }wat cluster activates the H2 ligand heterolytically. The binding between { }wat and H2C leads to a proton transfer via TS3 and thus results in a further elongation of the H1−H2 bond to 0.947 Å, whereas the O··· H2 distance and the Ir−H1 bond length are decreased to 1.498 and 1.698 Å, respectively. As also reported by other

sequently, in reaction A2, the H2 ligand of H2C undergoes a heterolytic H−H bond cleavage to form HC, after which a two-electron oxidation of HC occurs to release a proton to the carbon electrode (A3). In Cycle B, on the other hand, we suggest the formation of a hydroperoxide complex (PC) subsequent to the OC formation, resulting in one thermal reaction (black line, B2) and two electrode reactions (red lines, B1 and B3). First, a two-electron oxidation of AC (B1) occurs to form OC. After that, an O−O bond formation (B2) proceeds by reacting two H2O molecules with OC to form PC, in which a subsequent two-electron oxidation by H2O (B3) takes place to form AC and O2. Since Ogo’s group suggested that the light irradiation plays a key role in the formation of electron holes on the WO3 electrode,48 we consider reactions B1 and B3 as oxidation processes catalyzed by a photogenerated positive charge. To keep our discussions clear and concise, we first discuss the formation of HC from AC (simplified A1 and complete A2 processes). Next, we discuss the O−O bond formation (B2 process) and then briefly describe the reaction pathways for the electrode oxidation reactions (A3, B1, and B3 processes). For details of the A1, A3, B1, and B3 processes, please see the provided Supporting Information. 2.1. Formation of H2C and HC from AC (A1 and A2). Figure 4 shows a computed energy diagram for the formation of H2C and HC in the closed-shell singlet state, where the energies are measured relatively from the dissociation limit (AC + H2). First, an exchange between the H2O ligand of AC and an incoming H2 molecule occurs. This reaction involves an H2O dissociation from the Ir center and an H2 insertion to a vacant site via transition state-2 (TS2) with an activation energy of 10.7 kcal/mol (see Figure S2 for details). After the ligand exchange forming H2C, a (H2O)3 water cluster denoted as { }wat is adsorbed on the vicinity of the H2 ligand of H2C to form Int4 with an adsorption energy of −10.2 kcal/mol. Subsequently, the H1−H2 bond of Int4 is cleaved in a heterolytic fashion via TS3, where the separated H2 atom approaches the { }wat to form a H+(H2O)3 cluster that is bound to HC (HC-{H+}wat in Figure 4). This reaction step requires D


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Inorganic Chemistry

Figure 6. Calculated energy diagrams for B2 process in the singlet and triplet states. All energies are in kcal/mol. The red line corresponds to the direct formation of PC. Values in parentheses are the SCF energies.

Figure 7. Optimized structures of (a) tOC, (b) tInt5, (c) sTS4, (d) sInt6, (e) sTS5, (f) sInt7, (g) PC, and (h) HPC, where the superscript t and s stand for the triplet and singlet states, respectively. The spin densities for Ir and O1 atoms of parts a and b are also shown. In parts b−h, Cp* and bpy(COOH)2 ligands are omitted for clarity. Bond lengths are in Å.

groups,39,41,42 such structural changes are quite reasonable for an H−H bond cleavage. In HC-{H+}wat, the separated H2 atom is now fully bound to the O atom of the { }wat with an O−H2 bond length of 1.000 Å. Interestingly, a dihydrogen bond53 between the H1 hydride and the H2 proton with an H1···H2 distance of 1.520 Å is formed while the Ir−H1 bond is further shortened to 1.611 Å. As the natural population for the H1 and H2 atoms in HC-{H+}wat is calculated to be 0.01 and 0.53, respectively, we assign the formal charge of these atoms to −1 and +1, respectively. This confirms that the H1− H2 bond is cleaved in a heterolytic fashion. HC is then formed by removing the {H+}wat moiety from the iridium center, which

reduces the total electronic charge from +2 to +1 and shortens the Ir−H1 bond to 1.598 Å. The energy required for such a removal (7.6 kcal/mol) corresponds to the binding energy of the dihydrogen bond formed in HC-{H+}wat, indicating that HC is stabilized by the surrounding H2O molecules. 2.2. O−O Bond Formation from OC and H2O (B2). We next investigate the H2O oxidation in B2 process, where an O− O bond is formed through an H2O binding to the O ligand of OC, which facilitates one of the π* bonds in the IrO species of OC to break. Such an H2O oxidation catalyzed by a highvalent metal oxo species is known as the water nucleophilic attack (WNA) mechanism.35,43,54 Blakemore et al. reported E


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Inorganic Chemistry Scheme 1. Suggested Reaction Pathways for the A3, B1, and B3 Processesa

a

Superscript Roman number following the Arabic number indicates the oxidation stage (see text). PCET stands for proton-coupled electron transfer.

abstraction of the proximal hydrogen by { }wat to form PC or by a desorption of the coordinated H2O molecule to form a hydrogen peroxide complex (HPC), the former of which is energetically more favorable with a deprotonation energy of 7.8 kcal/mol. Although the overall reaction of B2 is slightly endothermic by 3.4 kcal/mol, we expect that the H2O oxidation by OC forming PC with an effective activation barrier of only 21.0 kcal/mol can proceed facilely. Figure 7 shows optimized structures of tOC, tInt5, sTS4, s Int6, sTS5, sInt7, PC, and HPC, where the superscript t and s stand for the triplet and singlet states, respectively. As can be seen from Figure 7b, the two H2O molecules in tInt5 are adsorbed with one of the H atoms being closer to the oxo ligand (O1). Here, the Ir−O1 bond length and the O1···O2 distance are calculated to be 1.836 and 3.023 Å, respectively. However, upon the O1−O2 bond formation via TS4, the O1 atom approaches the O2 atom with a shortened O1···O2 distance of 1.762 Å, while the Ir−O1 bond is elongated to 1.948 Å. In Int6, a bond between the O1 and O2 atoms is formed with a bond length of 1.505 Å, which is by 0.298 Å longer than that for O2 molecule in the gas phase.52 Such a weak O−O bond explains the low activation energy found for the reverse reaction from Int6 to Int5. During the O−O bond formation from Int5 to Int6, no radical species are formed, and the formal charge of Ir atom is changed from +5 to +3 by a two-electron transfer from the O1 atom to the Ir center. This indicates that the O−O bond formation occurs through the WNA mechanism and confirms that the high-valent IrVO species reacts with the H2O molecule to form an O−O bond in the thermal reaction. After the O−O bond formation, the H3 atom migrates to the O3 atom to form PC or Int7. In either case, the O−O bond is strengthened, as indicated by its shortened bond length to 1.463 and 1.455 Å, respectively for the PC and Int7 formations. In contrast, the Ir−O1 bond length in PC and Int7 are shortened by 0.014 Å and elongated by 0.167 Å, respectively, indicating that the H3 atom is dissociated as a proton that enhances the Ir−O1 bond strength. 2.3. Electrode Reactions (A3, B1, and B3). Having discussed the H−H bond cleavage forming HC and the O−O bond formation leading to PC by reacting OC with an H2O molecule, we now discuss the three electrode reactions of A3,

that the activation barriers for the O−O bond formation on a CpIr(O)(ppy)+ complex in the presence of one and two H2O molecules are 32.2 and 24.0 kcal/mol (B3LYP), respectively,43 suggesting that the presence of two H2O molecules in the O− O bond formation is essential. Following this suggestion, we therefore calculate the reaction path for the B2 process using two H2O molecules. For comparison, we also calculate the reaction path using only one H2O molecule (see Figure S9 in the Supporting Information). However, in contrast to the water-assisted mechanism (i.e., the O−O bond formation proceeds in conjunction with the proton relay migrating a hydrogen atom from the distal oxygen to the proximal oxygen via an insertion of an H2O molecule) suggested by Blakemore et al.,43 here we propose two separate processes including the O−O bond formation and a subsequent removal of the distal hydrogen atom. As shown in Figure 6, the B2 process begins with an exothermic adsorption of two H2O molecules on OC to form Int5 in the triplet state with an adsorption energy of −9.5 kcal/ mol, showing that OC is stabilized by the H2O molecules. Subsequently, the oxo ligand of Int5 reacts with one of the adsorbed H2O molecules via TS4 to form a dioxygen complex (Int6) that involves an oxo−water ligand (−OOH2). The activation energy for this reaction step is calculated to be 21.0 kcal/mol with a spin inversion from the triplet state to the closed-shell singlet state being expected to occur before TS4. We consider this activation barrier low enough for the reaction to proceed under mild conditions, although the reverse reaction from Int6 to Int5 is more likely to occur with an activation barrier of only 1.5 kcal/mol, which indicates a weak O−O bond in Int6. Nonetheless, Int6 may readily be stabilized by an insertion of two additional H2O molecules, where the {H+}wat cluster is released without barrier to form PC (see red line in Figure 6). Alternatively, a proton relay may migrate one of the H atoms in Int6 to the neighboring H2O molecule, forming a H2O2 ligand and a rather separated H2O molecule (Int7) via TS5. However, the activation energy required for this process (4.1 kcal/mol measured from Int6 or 23.6 kcal/mol measured from Int5 as suggested by Blakemore et al.43) is higher than that for the Int6 → Int5 reverse reaction, suggesting that this alternative pathway is less likely to occur. The formation of Int7 may be followed either by an F


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Inorganic Chemistry

OH2]2+ + O2 + H+ + 2e−. As shown in Scheme 1 (B3), PC initially releases one proton to the water solution and one electron to the WO3 electrode to form a dioxygen complex (8II) through the PCET mechanism, where E1/2 is calculated to be 0.83 V. The O2 moiety in 8II has a O1−O2 bond length of 1.311 Å and a spin density of 1.02 (see Figure S12), indicating that it is a superoxo species (−OO•−1). Then, a ligand exchange between the O2 moiety and an H2O molecule occurs to form 3II, where the Ir center is reduced to IrII with a reaction free energy of 7.9 kcal/mol. The final step is the one-electron oxidation of 3II to AC as discussed in the A3 process.

B1, and B3 which correspond to the regeneration of AC from HC, the formation of OC from AC, and the regeneration of AC from PC, respectively. In these three reactions, the number of total electrons changes, which thus makes us difficult to construct energy diagrams. To overcome this difficulty, we calculate oxidation potentials which involve a proton release by an H2O molecule and an electron transfer to the cathode electrode. Scheme 1 shows the reaction scheme for A3, B1, and B3 processes, where ΔG° and E1/2 stand for free energy of the deprotonation reaction and oxidation half potential, respectively (the computational methods are described in section 4). In order to distinguish the number of electrons released, we define three oxidation stages, namely, stages I, II, and III, corresponding to zero, one, and two electron(s) released. The detailed reaction pathways as well as geometrical and electronic parameters related to these processes are shown in Figure S4, S5, S6, S7, S11, and S12 of the Supporting Information. Below, we briefly discuss the mechanisms of these reactions. Two-Electron Oxidation of HC (A3). In this process (see Scheme 1, A3), AC is regenerated by a release of the H ligand from the Ir center (deprotonation) and a subsequent twoelectron oxidation of HC on the carbon electrode according to the following reaction: [(IrIII)−H]+ + H2O → [(IrIII)−OH2]2+ + H+ + 2e−. Specifically, HC releases one proton to the water solution and one electron to the carbon electrode to form 2II through the proton-coupled electron transfer (PCET) mechanism55 with E1/2 of 0.45 V. Subsequently, an incoming H2O molecule interacts with 2II to form 3II, where the Ir···O distance is found to be 3.649 Å (see Figure S5), indicating that the H2O molecule is uncoordinated to Ir atom. This might be due to the large ligand field splitting effects from the squarepyramidal geometry of the Ir center.56 The low reaction free energy of this step (ΔG° = 6.9 kcal/mol) indicates that this reaction is likely to occur. The final step is a one-electron oxidation of 3II to AC, where E1/2 is calculated to be 0.24 V. Considering that the oxidation potential of H2O (O2 + 4H+ + 4e− → 2H2O) is 1.23 V,55,57 the A3 process is expected to proceed spontaneously. Two-Electron Oxidation of AC (B1). In this reaction, a twoelectron oxidation of AC to OC takes place on the WO3 electrode in the presence of light irradiation according to the following reaction: [(IrIII)−OH2]2+ → [(IrV)O]2+ + 2H+ + 2e−. As shown in Scheme 1 (B1), in the first deprotonation reaction (AC to 5I), a deprotonation free energy is found to be 14.7 kcal/mol. Subsequently, a one-electron oxidation of 5I proceeds to form a hydroxyl complex (5II) with a total spin density for Ir and O1 of 0.86 (see Figure S7) and a calculated E1/2 of 1.44 V, which is higher than the oxidation potential of H2O (1.23 V),55,57 suggesting that this reaction should occur with the assistance of light irradiation. Afterward, a second deprotonation occurs to form 6II, requiring a high deprotonation free energy of 22.0 kcal/mol. The computed total spin density for Ir and O1 in 6II is 0.99, showing that an iridium− oxyl radical (IrO•) is formed. Finally, 6II is oxidized by one electron to OC (t6III), where the ground state is changed from the double state to the triplet state. The oxidation half potential of this final step is calculated to be 1.19 V. Two-Electron Oxidation of PC (B3). In this reaction, AC is regenerated from PC that was resulted in the B2 process. Specifically, PC is oxidized on the WO3 electrode in the presence of light irradiation to form AC and O2 according to the following reaction: [(IrIII)−OOH]+ + H2O → [(IrIII)−

3. CONCLUSIONS Using DFT calculations, we have proposed and discussed detailed mechanisms for the dual catalytic cycle of H2 and H2O oxidations by the iridium aqua complex [IrIII(η5-C5Me5){bpy(COOH)2}(H2O)]2+ (AC) toward the design of devices capable of switching functionalities between the fuel and solar cells. We have suggested five intermediate structures (i.e., aqua complex AC, hydrogen complex H2C, hydride complex HC, oxo complex OC, and hydroperoxide complex PC) and six reaction steps (i.e., AC ← HC ← H2C ← AC → OC → PC → AC) to be involved in the dual catalytic cycle. In the H2 oxidation reaction, our DFT calculations show that the H2 molecule is bound to the Ir center as a ligand replacing the H2O ligand of AC to form H2C, after which the H−H bond is heterolytically cleaved by an incoming threewater-molecule cluster to form HC that involves a stable dihydrogen bond in water solution. While the H−H bond cleavage is surprisingly barrierless, the H2O → H2 ligand substitution is found to be the rate-determining step with an activation energy of 10.7 kcal/mol. In the H2O oxidation reaction, OC is initially formed through an electrode reaction, after which two incoming H2O molecules weaken the Ir− O(oxo) bond of OC, allowing one O atom of the water molecules to bind the oxo species. Such an O−O bond formation is referred to as the water nucleophilic attack (WNA) mechanism. Subsequently, an oxo−water ligand (−OOH2) is formed with a required activation energy of 21.0 kcal/mol. Owing to its instability, this ligand is spontaneously deprotonated to −OOH ligand (PC) with a stronger O−O bond. We found that H2O molecules play important roles on the heterolytic H−H bond cleavage and the nucleophilic O−O bond formation. The three electrode reactions, HC → AC, AC → OC, and PC → AC, are predicted to proceed spontaneously in the presence of H2O molecules and light irradiation. 4. COMPUTATIONAL DETAILS All reaction intermediates and transition states were calculated by using the B3LYP functional,58 as implemented in the Gaussian 16 program package.59 The Grimme’s D3 method was used to account for the dispersion correction.60 We used the TZVP basis set61 for the H, C, N, and O atoms and the SDD basis set62 for the Ir atom. By calculating the analytical harmonic vibration frequencies, we confirmed that the obtained ground states and transition states have, respectively, no and one imaginary frequency. The obtained energy values include the zero-point energy correction (ZPC) except those in the electrode reactions, where free energies under 333.25 K and 1 atm were considered.48 Considering that the high humidity in the real operation conditions of fuel cells results in G


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water vapor around the electrode and AC, implicit solvent effects of H2O with a dielectric constant of 78.39 were considered and predicted by using the polarized continuum model.63 Because hydronium ion (H3O+) as a proton source is highly acidic, we considered a H+(H2O)3 cluster as a more reasonable and widely accepted64−67 model of H+ in water solution (see Figure S1 in the Supporting Information). The attachment and detachment of (H2O)3 and H+(H2O)3 clusters considered in this study are required to rationalize the experimental findings and have been widely used by other researchers.67,68 To evaluate the electrode reactions, we calculated oxidation potentials which involves electron transfer from the Ir complex to out of the system and proton release. Oxidation half potentials (E1/2) were calculated using the following equation:

Muhammad Haris Mahyuddin: 0000-0002-8017-7847 Kazunari Yoshizawa: 0000-0002-6279-9722 Seiji Ogo: 0000-0003-2078-6349 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Ministry of Education, Culture, Sports, Science, and Technology (MEXT), Japan, though Grants-in-Aid [26000008 (Specially Promoted Research) and 17H06928 (Research Activity Start-up)], the World Premier International Research Center Initiative (WPI), Japan, and JST CREST Grant Numbers JPMJCR15P5 and JPMJCR18R2, Japan.

AH(nnHH − ne) → A + nee− + nH H+

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Here ne and nH stand for the number of electrons and protons, respectively. The oxidation potential is related to the standard oxidation free energy in the water solution by ° = G°(A−ne) − G°(A) −neFE° = ΔG EA

° is free energy change associated with an oxidation where ΔG EA in a solvent environment at standard condition and F is the Faraday constant. The standard oxidation half potential ° , is related to ΔG EA ° by (assuming nH = 0), i.e., E1/2 ° = ΔG1/2 ° = G°(AH(nnHH − ne)) − G°(A) − nHG°(H+) ΔG EA ° = neFE1/2

In our model, we described G°(H+) using the following equation: G°(H+) = G°({H+}wat ) − G°({}wat )

Here G°({H+}wat) and G°({ }wat) are the free energies of {H+}wat and { }wat, respectively. The dependence of the oxidation half potential on pH is obtained from the Nernst equation, where nonstandard E1/2 is calculated using the following equation: ij ΔG1/2 ° yz zz − RT ln(10)nH pH − E ° ,ref E1/2 = −jjj 1/2 j ne zz ne k {

Here E°,ref 1/2 is the reference potential (usually the normal hydrogen electrode (NHE) at 4.24 V).57,69 This approach is a widely accepted method.69,70

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.9b00307. Descriptions, energy diagrams, optimized structures, natural populations, and spin densities of the A3, B2, and B3 processes and Cartesian coordinates for all optimized structures (PDF)

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REFERENCES

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*(K.Y.) E-mail: [email protected]. *(S.O.) E-mail: [email protected]. H


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