1 Nature Chemistry 2013 Vol: 5(5):403-409. DOI: 10.1038/nchem.1621

Decoupling hydrogen and oxygen evolution during electrolytic water splitting using an electron-coupled-proton buffer

Hydrogen is essential to several key industrial processes and could play a major role as an energy carrier in a future ‘hydrogen economy’. Although the majority of the world's hydrogen supply currently comes from the reformation of fossil fuels, its generation from water using renewables-generated power could provide a hydrogen source without increasing atmospheric CO2 levels. Conventional water electrolysis produces H2 and O2 simultaneously, such that these gases must be generated in separate spaces to prevent their mixing. Herein, using the polyoxometalate H3PMo12O40, we introduce the concept of the electron-coupled-proton buffer (ECPB), whereby O2 and H2 can be produced at separate times during water electrolysis. This could have advantages in preventing gas mixing in the headspaces of high-pressure electrolysis cells, with implications for safety and electrolyser degradation. Furthermore, we demonstrate that temporally separated O2 and H2 production allows greater flexibility regarding the membranes and electrodes that can be used in water-splitting cells.

Mentions
Figures
Figure 1: Schematic of the ECPB-based approach to water splitting. a, Water is oxidized to give oxygen gas, protons and electrons. The electrons pass through the external circuit and the protons diffuse through the semipermeable membrane that separates the compartments. The ECPB in the other compartment is then reduced by the electrons and simultaneously accepts charge-balancing protons, turning from yellow to dark blue as it does so (V = the application of an external bias). b, Reoxidation of the ECPB releases protons, which can migrate through the membrane to the other electrode where they combine with the electrons removed from the ECPB to generate hydrogen gas. The ECPB returns to its original yellow colour as this happens. c, The ECPB in a cellulose-membrane electrolysis cell immediately after the start of reduction, showing the ECPB solution as bright yellow/green. d, After a few seconds, dark-blue reduced and protonated ECPB is visible near the cathode in the unstirred solution, indicated by the red arrow. e, After two minutes the entire ECPB solution turns dark blue (see Supplementary Fig. S16a–d for photographs of other cells). Figure 2: Cyclic voltammograms showing the reversible redox waves of phosphomolybdic acid. Cyclic voltammograms of 0.5 M phosphomolybdic acid (solid black line) and 1 M H3PO4 (dashed red line) are shown. A three-electrode, single-compartment set-up was used, with a 2 mm diameter platinum disc working electrode, platinum mesh counter electrode and an Ag/AgCl reference electrode at a scan rate of 100 mV s−1. The green arrows highlight the oxidation and reduction peaks associated with the first reduced state of phosphomolybdic acid. The inset shows a graph of the peak current versus square root of the scan rate for the reoxidation event associated with the first reversible two-electron wave (centred at +0.65 V) over the scan-rate range of 50–400 mV s−1. The linear fit is provided as a guide to the eye. The error associated with the measurement of the currents was ±0.1 mA on each reading, which corresponds to the size of the data markers. All experiments were performed without degassing and under air. Figure 3: Current–voltage curves obtained when stirring with and without an ECPB. a, Comparison of the reduction of a 0.5 M solution of 50:50 (H3O+)[H2PMo12O40]− and (H3O+)[H4PMo12O40]− (red) and a 1 M solution of H3PO4 (black). b, Comparison of the oxidation of a 0.5 M solution of 50:50 (H3O+)[H2PMo12O40]− and (H3O+)[H4PMo12O40]− (red) and a 1 M solution of H3PO4 (black). A three-electrode, single compartment set-up was used, with a 2 mm diameter platinum disc working electrode, platinum mesh counter electrode and an Ag/AgCl reference electrode, with iR compensation. Horizontal error bars correspond to the error associated with the iR compensation of the potentiostat (±3%) and the vertical error bars are based on the standard deviation of the currents from the mean current at that voltage. Figure 4: Comparison of current densities for H2 and O2 evolution with and without the use of 0.5 M 50:50 (H3O+)[H2PMo12O40]−:(H3O+) [H4PMo12O40]−. a, Potential–current curves for the HER on platinum and carbon electrodes in a two-electrode configuration. For platinum, both electrodes were discs of area 0.031 cm2 and for glassy carbon both electrodes were discs of area 0.071 cm2. Red = platinum with ECPB, green = glassy carbon with ECPB, black = platinum without ECPB, blue = glassy carbon without ECPB. (The corresponding anodic half-reaction of this last process is electrode degradation and not oxidation of water45.) The potential values reported are corrected for solution ohmic losses. b, Potential–current curves for the OER on platinum electrodes in a two-electrode configuration. Both electrodes were platinum discs of area 0.031 cm2. Red = platinum with ECPB, black = platinum without ECPB. Horizontal error bars correspond to the error associated with the iR compensation of the potentiostat (±3%) and the vertical error bars are based on the standard deviation of the currents from the mean current at that voltage. Figure 5: Current densities and gas chromatography for H2 evolution using an ECPB. a, Potential–current curves for platinum and carbon electrodes in a two-electrode configuration. For platinum, the electrode performing the HER had an area of 0.031 cm2, and for glassy carbon the area was 0.071 cm2. The counter electrodes were platinum mesh for the platinum experiments, and carbon felt for the carbon experiments. The counter electrode was placed in either 0.5 M ECPB/ECPB* or 1 M H3PO4. Red = platinum with ECPB, green = glassy carbon with ECPB, black = platinum without ECPB, blue = glassy carbon without ECPB. Error bars were determined as for . b, Representative trace of cumulative H2 build-up in the headspace during HER on platinum electrodes. ‘H2 calculated’ (red) was determined from the charge passed and ‘H2 measured’ (blue) was determined by GC (Supplementary Section SI-6). The time lag between hydrogen production and detection by the GC resulted from the slow effusion of H2 through the narrow tubing in the GC system. The error bars are within the data markers and correspond to ±2%.
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    • . . . This was borne out by gas chromatography headspace analysis (GCHA), which showed that H2 equivalent to 100% (±2%) of the charge passed (that is, full Faradaic efficiency, see Fig. 5b) was obtained when ECPB* was used in the counter-electrode compartment at an iR-corrected voltage22 of −1.4 V (uncorrected voltage = −5 V, two-electrode set-up; see Supplementary Section SI-6 and Fig . . .
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    • . . . The reduced and protonated ECPB need not necessarily be produced by direct water oxidation and routes for reduced ECPB production from sustainable fuel sources (such as biomass) could also be used43, 44 . . .
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    • . . . The reduced and protonated ECPB need not necessarily be produced by direct water oxidation and routes for reduced ECPB production from sustainable fuel sources (such as biomass) could also be used43, 44 . . .
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    • . . . Red = platinum with ECPB, green = glassy carbon with ECPB, black = platinum without ECPB, blue = glassy carbon without ECPB. (The corresponding anodic half-reaction of this last process is electrode degradation and not oxidation of water45.) The potential values reported are corrected for solution ohmic losses. b, Potential–current curves for the OER on platinum electrodes in a two-electrode configuration . . .
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