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Analysis of Homogeneous Water Oxidation Catalysis with Collector− Generator Cells Benjamin D. Sherman,† Matthew V. Sheridan,† Kyung-Ryang Wee, Na Song, Christopher J. Dares, Zhen Fang, Yusuke Tamaki, Animesh Nayak, and Thomas J. Meyer* Department of Chemistry, University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599, United States ABSTRACT: A collector−generator (C−G) technique has been applied to determine the Faradaic efficiencies for electrocatalytic O2 production by the homogeneous water oxidation catalysts Ru(bda)(isoq)2 (1; bda = 2,2′-bipyridine and isoq = isoquinoline) and [Ru(tpy)(bpz)(OH2)]2+ (2; tpy = 2,2′:6′,2″-terpyridine and bpz = 2,2′-bipyrazine). This technique uses a custom-fabricated cell consisting of two fluorine-doped tin oxide (FTO) working electrodes separated by 1 mm with the conductive sides facing each other. With a catalyst in solution, water oxidation occurs at one FTO electrode under a sufficient bias to drive O2 formation by the catalyst; the O2 formed then diffuses to the second FTO electrode poised at a potential sufficiently negative to drive O2 reduction. A comparison of the current versus time response at each electrode enables determination of the Faradaic efficiency for O2 production with high concentrations of supporting electrolyte important for avoiding capacitance effects between the electrodes. The C−G technique was applied to electrocatalytic water oxidation by 1 in the presence of the electron-transfer mediator Ru(bpy)32+ in both unbuffered aqueous solutions and with the added buffer bases HCO3−, HPO42−, imidazole, 1-methylimidazole, and 4methoxypyridine. HCO3− and HPO42− facilitate water oxidation by atom-proton transfer (APT), which gave Faradaic yields of 100%. With imidazole as the buffer base, coordination to the catalyst inhibited water oxidation. 1-Methylimidazole and 4methoxypyridine gave O2 yields of 55% and 76%, respectively, with the lower Faradaic efficiencies possibly due to competitive C−H oxidation of the bases. O2 evolution by catalyst 2 was evaluated at pH 12 with 0.1 M PO43− and at pH 7 in a 0.1 M H2PO4−/HPO42− buffer. At pH 12, at an applied potential of 0.8 V vs SCE, water oxidation by the RuIV(O)2+ form of the catalyst gave O2 in 73% yield. In a pH 7 solution, water oxidation at 1.4 V vs SCE, which is dominated by RuV(O)3+, gave O2 with an efficiency of 100%. The lower efficiency for RuIV(O)2+ at pH 12 may be due to competitive oxidation of a polypyridyl ligand.
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INTRODUCTION In electrocatalytic or photocatalytic water oxidation, a significant challenge arises in acquiring accurate and rapid determination of the amount of O2 produced. An increasingly popular approach is the use of the collector−generator (C−G) method in which O2 generated at an anode or photoanode (eq 1) is detected in situ at a second electrode at a defined distance from the source.1−7 The benefits of this approach include the following: (1) catalytically produced O2 is detected in solution, at close proximity to the point of generation, which enables a straightforward analysis of the Faradaic efficiency; (2) relative confinement of product O2 in a small volume results in a welldefined electrochemical response; (3) in contrast to a Clarktype electrode or analysis by gas chromatography, this approach does not require tightly sealed reaction cells to provide accurate measurements. 2H 2O → O2 + 4H+ + 4e−
Figure 1. Schematic diagram of the C−G cell for O2 production and analysis.
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Special Issue: Small Molecule Activation: From Biological Principles to Energy Applications Part 3
The C−G cells described here operate with the conductive sides of the two working electrodes facing one another, separated by 1 mm (Figure 1). The “generator” electrode © XXXX American Chemical Society
Received: September 24, 2015
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DOI: 10.1021/acs.inorgchem.5b02182 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry operates under a sufficiently positive bias to drive water oxidation (eq 1). The O2 produced diffuses to the “collector” electrode, poised at a sufficiently negative bias to monitor the current resulting from the reduction of O2 to water (O2 + 4H+ + 4e− → 2H2O). When performed under an inert gas atmosphere, these measurements allow for the accurate quantification of the current passed at both the collector and generator electrodes. With the use of a bipotentiostat, the independent measurement of the current and, by integration over time, the charge passed at each electrode allows for the straightforward, direct measurement of the Faradaic efficiency for O2 production. Here we extend the use of the C−G technique to homogeneous water oxidation catalysis both to demonstrate the generality of the approach and to provide further mechanistic insight into two molecular catalysts. The catalysts, shown in Figure 2, are Ru(bda)(isoq)2 (1; bda = 2,2′-
[Ru V (O)]+ + HOH‐‐‐OP(O)2 OH− → [Ru III(OOH)] + HOP(O)2 OH−
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kET
Ru(bpy)33 + + [Ru IV (O)] ⎯→ ⎯ Ru(bpy)32 + + [Ru V (O)]+ (7)
Water oxidation catalysis by 2 has also been reinvestigated by use of the C−G technique.12,13 In addition to water oxidation catalysis by RuV(O)3+, 2 provides a rare example of a ruthenium polypyridyl complex with reactivity toward water oxidation by RuIV(O)2+. Mechanistically, water oxidation by this catalyst occurs through RuIV(O)2+ by an initial proton-coupled electron-transfer (PCET) oxidation of RuII(OH2)2+ (eq 8). After oxidation to RuIV(O)2+, O---O bond formation occurs (eq 9), followed by oxidation of the peroxide intermediate and O2 release (eq 10). For RuIV(O)2+ as the oxidant, an important role for APT has also been found with bases such as HPO42−. The results with 1 and 2 show that the C−G method can be used reliably to determine the Faradaic efficiency for O2 generation for electrochemically driven homogeneous catalysts. −2e−, − 2H+
[Ru II−OH 2]2 + ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ [Ru IV (O)]2 +
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[Ru IV (O)]2 + + OH− → [Ru II(OOH)]+
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−4e, + H 2O, − 3H+
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Figure 2. Water oxidation catalysts 1 and 2.
−2e−, − 2H+
−e −
[Ru IV (O)] ⎯⎯⎯→ [Ru V (O)]+ k O−O
[Ru V (O)]+ + H 2O ⎯⎯⎯⎯→ [Ru III(OOH)] + H+
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EXPERIMENTAL SECTION
General Procedures. Complex 1 and the triflate salt of 2 were prepared according to literature procedures.9,12 Ru(bpy)32+ was purchased from Sigma-Aldrich as the chloride salt and used as received. Fluorine-doped tin oxide (FTO)-coated glass slides were purchased from Hartford Glass (15 Ω/cm2). All other reagents were purchased from Sigma-Aldrich and used without further purification. Distilled water was additionally purified using a Milli-Q Ultrapure water filtration system. All electrochemical experiments were performed using a CH Instruments 760E bipotentiostat, a platinum wire counter electrode, and a saturated calomel reference electrode (SCE). C−G Cell. The C−G cell was constructed from two FTO electrodes (1 cm × 4 cm) first using conductive silver epoxy (Chemtronics CW2400) to attach wire leads. Thinly cut pieces of microscope glass (1 mm thick, Fisher) were attached to the outer edges of the collector FTO using nonconductive epoxy (Hysol E00CL). Finally, the generator FTO was bonded to the collector FTO using nonconductive epoxy, with the conductive sides facing each other and care taken to ensure that the wire leads did not touch. The lateral edges of the cell were completely sealed with epoxy, leaving a void space between the two electrode faces open at the top and bottom. Upon immersion in solution, the interior space fills with electrolyte by capillary action. In a typical experiment, the catalyst (concentrations indicated in the text) along with (where indicated) 3 equiv of Ru(bpy)32+ was dissolved in 0.5 M NaClO4 solutions containing 0.1 M buffer, with the pH indicated in the text. The solutions were degassed with N2, and the C−G cell was introduced into the test solution. For the C−G experiments, the ionic strength should be greater than 0.5 M to avoid capacitive interference between the two working electrodes.
bipyridine-6,6′-dicarboxylate and isoq = isoquinoline) and [Ru(tpy)(bpz)(OH2)]2+ (2; tpy = 2,2′:6′,2″-terpyridine and bpz = 2,2′-bipyrazine). Catalyst 1 is a member of the family of Ru(bda)(L)2 catalysts initially described by Sun and coworkers.8,9 Complexes in this series undergo rapid water oxidation catalysis in acidic solutions with cerium(IV) as a chemical oxidant. The proposed mechanism for 1 involves rate-limiting bimolecular O---O bond formation following oxidation of the complex to RuV(O)+ (eqs 2−7).10 At carbon electrodes, facile electrocatalytic water oxidation by single-site catalysis has been observed. At oxide electrodes, oxidation of RuIV(O) to RuV(O)+ occurs slowly and electron-transfer mediation by the Ru(bpy)33+/2+ couple (eq 7) activates the complex toward water oxidation.10,11 In the O---O bond-forming step for these catalysts, an important role for atom-proton transfer (APT) has been identified, with O-atom transfer occurring in concert with proton transfer to an acceptor base. An example is shown in eq 6 with HPO42− as the acceptor base. In the current study, the buffer base dependence for water oxidation by 1 is extended to a series of bases (pKa in parentheses): H2PO4−/HPO42− (7.2), CO2/bicarbonate (6.4), imidazole (7.0), 1-methylimidazole (7.4), and 4-methoxypyridine (6.6). [Ru−OH 2]2 + ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ [Ru IV (O)]
[Ru II(OOH)]+ ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ [Ru IV (O)]2 + + O2
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RESULTS AND DISCUSSION The C−G method has previously been applied to electrochemical and photoelectrochemical water oxidation in dyesensitized photoelectrosynthesis cells, with the catalyst bound to the electrode surface.1−7 Application of the technique to
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−4e−, + H 2O, − 3H+
[Ru III(OOH)] ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ [Ru V (O)]+ + O2 (rapid) (5) B
DOI: 10.1021/acs.inorgchem.5b02182 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry homogeneous water oxidation requires that additional experimental factors be taken into account. For example, in this study with 1 and Ru(bpy)32+ added as an electron-transfer mediator, the charge passed at the collector electrode from rereduction of the oxidized forms of the mediator and catalyst needs to be accounted for upon determination of the total charge resulting from O2 reduction. Rereduction of the oxidized mediator and catalyst can be distinguished from O2 reduction by parallel charge measurements at two different collector potentials. Under one bias regime, the potential is sufficiently negative to reduce both the oxidized catalyst/mediator and evolved O2 (QCH). At the other, the collector is poised at a more positive potential before the onset of O2 reduction to measure the current passed due solely to the rereduction of the oxidized mediator and catalyst (QCL). In our experiments, QCL measurements were made at −0.35 V vs SCE, a sufficiently negative potential for complete reduction of the oxidized mediator and catalyst, and QCH was measured at −0.85 V vs SCE, sufficiently negative for O2 reduction. Figure 3 shows current−potential traces at an FTO electrode in air-saturated and N2-degassed solutions. With O2 present in
Figure 4. CVs of 0.15 mM 1 with (black) and without (gray) 0.45 mM Ru(bpy)32+ at a 1 cm2 FTO electrode in pH 6.5, 0.5 M NaClO4 unbuffered solution, and ν = 0.05 V/s.
solution with 0.5 M NaClO4 as the supporting electrolyte at pH 6.5 with 0.15 mM 1 with (black) and without (gray) 0.45 mM Ru(bpy)32+ at an FTO electrode. A comparison of the CVs illustrates the dramatic effect of electron-transfer mediation at the oxide electrode. In a C−G experiment, a generator potential of 1.2 V vs SCE was used to drive water oxidation by the RuV(O)+ form of the catalyst with the assistance of the mediator couple, Ru(bpy) 33+/2+ with E1/2 = 1.02 V vs SCE.11 Figure 5 shows current−time traces at a C−G electrode in an unbuffered pH 6.5 solution containing 0.15 mM 1, 0.45 mM
Figure 3. CVs at an FTO electrode in the presence (black) and absence (gray) of atmospheric O2 in pH 6.5 (unbuffered), 0.5 M NaClO4 solution, and ν = 0.05 V/s.
solution, O2 reduction is observed at −0.85 V vs SCE but not at −0.35 V vs SCE. Without O2 in solution, minimal current passes at −0.85 V vs SCE, demonstrating the electrochemical sensitivity of the FTO electrode to O2. This behavior provides the basis for C−G analysis of O2 with current/charge measurements at −0.35 V vs SCE attributed to oxidized mediator, catalyst, or non-O2 oxidation products and the additional current observed at −0.85 V vs SCE due to O2 produced at the generator. Applying eq 11 allows for determination of the Faradaic efficiency for homogeneous O2 production (ηO2). In this analysis, ηO2 is given by the charge attributed to O2 reduction (QCH − QCL) divided by the total charge passed at the generator electrode (QGH), corrected for the collection efficiency (70%) of the C−G cell as determined in an earlier study.1 ηO = [(Q CH − Q CL)/Q GH]/0.70 2
Figure 5. Current−time traces at the generator (top, solid traces) and collector (bottom, dashed traces) with the generator at 1.2 V vs SCE from 0 to 600 s and 0.2 V vs SCE from 600 to 1200 s (solid black, GH) and the collector at −0.85 V vs SCE (dashed black, CH) and with the generator at 1.2 V vs SCE from 0 to 600 s and 0.2 V vs SCE from 600 to 1200 s (solid gray, GL) and the collector at −0.35 V vs SCE (dashed gray, CL). The experiments shown were carried out with 0.15 mM 1 and 0.45 mM Ru(bpy)32+ in 0.5 M NaClO4 at pH 6.5.
Ru(bpy)32+, and 0.5 M NaClO4 with the potential of the generator electrode held at 1.2 V vs SCE and the collector electrode at −0.85 V vs SCE and −0.35 V vs SCE in two separate experiments. In these experiments, the potential at the generator electrode was held at 1.2 V vs SCE from 0 to 600 s and then switched to 0.2 V vs SCE from 600 to 1200 s, a potential cathodic of O2 generation. The 600 s time interval at a lower generator potential in the second half of the trace allows for the complete capture of oxidized products formed within the electrolyte volume between the collector and generator
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In a first case study, water oxidation catalysis by 1 was examined with and without the electron-transfer mediator Ru(bpy)32+. Figure 4 shows cyclic voltammograms (CVs) in an unbuffered C
DOI: 10.1021/acs.inorgchem.5b02182 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry electrodes. The current at the collector electrode at −0.85 V vs SCE (black dashed trace) rises initially as O2 produced at the generator traverses across the cell and then decreases after t = 600 s as O2 production stops at the generator and O2 is gradually depleted from the volume in the C−G cell. With the collector at a more positive potential (gray dashed trace), a negligible increase in the cathodic current is observed with a negligible response from oxidized 1 or Ru(bpy)33+. Figure 6 shows integration of the current−time traces in Figure 5. At the lower collector potential of −0.35 V vs SCE,
Figure 7. Current−time traces for 0.15 mM 1 at 1 cm2 FTO electrodes in a pH 6.5, 0.5 M NaClO4 solution without added Ru(bpy)32+ with (top, solid black) the generator electrode at 1.2 V vs SCE from 0 to 600 s and 0.2 V vs SCE from 600 to 1200 s and (bottom, dashed black) the collector at −0.85 V vs SCE.
imidazole as the added base, the collector electrode does not detect O2 as an electrolysis product. With 1-methylimidazole, O2 is observed, however, at a lower Faradaic efficiency of 55% compared with the inorganic bases. CVs with imidazole and 1-methylimidazole are shown in Figure 9. The CV with added imidazole shows the appearance waves at E1/2 = 0.5 and 0.75 V vs SCE consistent with the coordination of imidazole and deactivation of the catalyst.14−18 With 1-methylimidazole (or 4-methoxypyridine), evidence for catalysis is observed, but the less than quantitative production of O2 points to a competitive process or processes. Although not investigated in detail, there is an extensive oxidation chemistry of activated C−H bonds by high-oxidation-state ruthenium oxo catalysts, which, presumably, is in competition with water oxidation.19−21 Water oxidation by ruthenium(II) polypyridyl complexes is dominated by the 3e−/2H+ oxidation of Ru−OH2 to RuV(O) followed by water attack and O---O bond formation. There are limited examples of oxidation by ruthenium(IV) including a report by Thummel and co-workers with the equatorial 4-tertbutyl-2,6-bis(10,80-naphthyrid-20-yl)pyridyl ligand complex shown in Figure 1022 and by RuIV(tpy)(bpz)(O)2+.13 As shown in Figure 10, for the latter, the mechanism involves PCET oxidation of 2 to RuIV(O)2+ followed by water oxidation. CVs in pH 12, 0.1 M phosphate buffer solution, at an FTO electrode are shown in Figure 11 with and without 0.97 mM 2. No electron-transfer mediator is needed to activate the cationic ruthenium(II) precursor complex at the oxide electrode. From the results of previous electrochemical studies, 2 undergoes a 2e−/2H+ oxidation to RuIV(O)2+ at E1/2 = 0.46 V at pH 11 followed by catalytic water oxidation at a potential of ∼0.6 V below the background.13 The results of C−G experiments for 2 under similar conditions are shown in Figure 12 with 0.8 V vs SCE applied at the generator electrode. On the basis of the average of two experiments, the Faradaic efficiency under these conditions was 73%. The C−G experiment was repeated with 2 as the oxidant but at pH 7 at a generator applied potential of 1.4 V vs SCE in a 0.1 M H2PO4−/HPO42− buffer. Under these conditions, water oxidation by RuIV(O)2+ is slow and RuV(O)3+ is accessible and dominates reactivity. With RuV(O)3+ in a H2PO4−/HPO42− buffer, water oxidation occurs with a Faradaic efficiency of 100%.
Figure 6. Charge−time traces for the experiments in Figure 5 with the generator at 1.2 V vs SCE from 0 to 600 s and 0.2 V vs SCE from 600 to 1200 s (solid gray, QGL) and the collector electrode at −0.35 V vs SCE (dashed gray, QCL) and with the generator at 1.2 V vs SCE from 0 to 600 s and 0.2 V vs SCE from 600 to 1200 s (solid black, QGH) and the collector at −0.85 V vs SCE (dashed black, QCH). The experiments were carried out with 0.15 mM 1 and 0.45 mM Ru(bpy)32+ in 0.5 M NaClO4 at pH 6.5.
the charge collected at the collector electrode accounted for 1% of that collected at the generator. By contrast, at −0.85 V vs SCE, a substantial charge is observed at the collector. The application of eq 11 to these results gave a Faradaic efficiency of 100% for O2 production for the unbuffered pH 6.5 solution. Without the electron-transfer mediator Ru(bpy)32+ and consistent with the CVs in Figure 4, an almost negligible amount of catalytic current was obtained in the C−G experiment, as shown in Figure 7. The catalytic current was significantly lower than that in experiments in the unbuffered solution with Ru(bpy)32+ present (Figure 5). Steady-state currents at 1.2 V vs SCE were 2 μA without a mediator and 20−60 μA with a mediator. Using the same set of C−G experiments, the influence of different buffer bases was investigated with 1 as the oxidant, with the results summarized in Table 1. In all cases, increased catalytic currents were observed relative to the buffer-free solution, consistent with rate enhancements through baseassisted APT pathways. Mechanistic details were reported previously for 1 with acetate and HPO42− as the proton acceptor bases.10,11 The results of C−G experiments in a CO2/ HCO3− buffer are shown in Figure 8. From Table 1, Faradaic efficiencies of 100% are observed with the inorganic bases HCO3− and HPO42− consistent with quantitative water oxidation under these conditions. The data also highlight potential complications inherent in water oxidation catalysis that arise from competing processes. An example appears with the nitrogen heterocycle bases; for D
DOI: 10.1021/acs.inorgchem.5b02182 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry
Table 1. Summary of the Solution Conditions and Faradaic Efficiencies for 1 (0.15 mM) in the Presence of Ru(bpy)32+ (0.45 mM) with the Indicated Buffer buffer
concentration (M)
active base
pKa
concentration of base (M)
Faradaic efficiency (%)
None H2PO4− /HPO42− CO2/HCO3− imidazole 1-methylimidazole 4-methoxypyridine
0.1 0.5 0.1 0.1 0.1
HPO42− HCO3− imidazole 1-methylimidazole 4-methoxypyridine
7.2 6.4 7.0 7.4 6.6
0.04 0.40 0.05 0.03 0.07
100 100 100 0 55 76
Figure 8. As in Figure 5, current−time traces for 0.15 mM 1 and 0.45 mM Ru(bpy)32+ at 1 cm2 FTO electrodes in a pH 7, 0.5 M CO2/ bicarbonate buffer solution.
Figure 11. CVs of a 1 cm2 FTO electrode with (black) and without (gray) 0.97 mM 2 in pH 12, 0.1 M phosphate buffer, 0.5 M NaClO4, and ν = 0.05 V/s.
Figure 9. CVs of 0.11 mM 1 at a boron-doped diamond electrode in pH 7, 0.1 M imidazole (black) and 1-methylimidazole (gray), 0.5 M NaClO4 solution, and ν = 0.05 V/s.
Figure 12. Current−time traces at the generator (top, solid traces) and collector (bottom, dashed traces) with the generator potential at 0.8 V vs SCE from 0 to 900 s and at 0.2 V vs SCE from 900 to 1800 s (solid black, GH) and collector at −0.85 V vs SCE (dashed black, CH) and with the generator at 0.8 V vs SCE from 0 to 900 s and at 0.2 V vs SCE from 900 to 1800 s (solid gray, GL) and collector at −0.35 V vs SCE (dashed gray, CL). The solution was 0.97 mM 2 at pH 12 in 0.1 M PO43− with a 0.5 M NaClO4 supporting electrolyte.
The comparison between the two results highlights a second complication in water oxidation catalysis. The “missing” oxidative equivalents at the collector presumably arise from competitive ligand oxidation at pH 12 for the RuIV(O)2+ form of 2. A related, base-induced ligand oxidation chemistry is wellknown for ruthenium(III) polypyridyl complexes.23−25
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Figure 10. Structure of 3 and mechanism of water oxidation through RuIV(O)2+ by 2.13,22
CONCLUSIONS Our results describe an in situ procedure for O2 analysis in a small, confined volume in a C−G electrochemical cell. E
DOI: 10.1021/acs.inorgchem.5b02182 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry
(15) Oyama, N.; Anson, F. C. J. Electroanal. Chem. Interfacial Electrochem. 1978, 88, 289−297. (16) Mahanti, B.; De Sankar, G. Transition Met. Chem. 1992, 17, 23− 28. (17) Davies, N. R.; Mullins, T. L. Aust. J. Chem. 1967, 20, 657−668. (18) Davies, N. R.; Mullins, T. L. Aust. J. Chem. 1968, 21, 915−925. (19) Meyer, T. J.; Huynh, M. H. V. Inorg. Chem. 2003, 42, 8140− 8160. (20) Gunay, A.; Theopold, K. H. Chem. Rev. 2010, 110, 1060−1081. (21) Kojima, T.; Nakayama, K.; Ikemura, K.; Ogura, T.; Fukuzumi, S. J. Am. Chem. Soc. 2011, 133, 11692−11700. (22) Lewandowska-Andralojc, A.; Polyansky, D. E.; Zong, R.; Thummel, R. P.; Fujita, E. Phys. Chem. Chem. Phys. 2013, 15, 14058−14068. (23) Hyde, J. T.; Hanson, K.; Vannucci, A. K.; Lapides, A. M.; Alibabaei, L.; Norris, M. R.; Meyer, T. J.; Harrison, D. P. ACS Appl. Mater. Interfaces 2015, 7, 9554−9562. (24) Roecker, L.; Kutner, W.; Gilbert, J. A.; Simmons, M.; Murray, R. W.; Meyer, T. J. Inorg. Chem. 1985, 24, 3784−3791. (25) Ghosh, P. K.; Brunschwig, B. S.; Chou, M.; Creutz, C.; Sutin, N. J. Am. Chem. Soc. 1984, 106, 4772−4783.
Application of the cell to homogeneous water oxidation catalysis by both mediated and unmediated catalysts at oxide electrodes with and without inorganic buffers results in the quantitative appearance of O2 at the collector cathode, validating the technique. The C−G procedure also provides a basis for identifying limitations in homogeneous water oxidation catalysis arising from coordinative deactivation with added organic buffer bases and competitive oxidation of the added base. Under basic conditions at pH 12, water oxidation by the RuIV(O)2+ form of 2 is also less than quantitative, presumably because of competitive oxidation of a polypyridyl ligand.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Author Contributions †
The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. These authors contributed equally. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This research was supported solely by the University of North Carolina Center for Solar Fuels, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, under Award DESC0001011.
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REFERENCES
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