A team of Japanese and South Korean researchers has pioneered a way to use seawater to obtain hydrogen peroxide (H2O2) instead of using pure water as a solar fuel. Their paper, “Seawater usable for production and consumption of hydrogen peroxide as a solar fuel,” was published in the May 4 edition of Nature Communications. “It is highly desired to utilize the most earth-abundant seawater instead of precious pure water for the practical use of H2O2 as a solar fuel,” the researchers said in the paper.
Because it is pollution-free, renewable, and ubiquitous, solar energy is a prized energy resource. And while converting solar energy directly to electricity is desirable, it is not always as convenient as using that energy to create a photoelectrochemical cell, such as a battery. Hydrogen gas (H2) would seem to be the obvious choice to hold that solar energy, but it is difficult to store at room temperature as a gas and requires a large amount of energy to compress or change its state. Using H2O2 is a viable alternative to gaseous H2 since it is easier to store and transport, has a slightly higher energy density than even compressed hydrogen, and still results in clean emissions—just dioxygen (O2) and water (H2O).
“In contrast to gaseous H2, H2O2 can be produced as an aqueous solution from water and O2 in the air by the combination of the photocatalytic two-electron reduction of O2 and the catalytic four-electron oxidation of water,” according to the paper. The overall photovoltaic catalytic reaction (below) produces hydrogen peroxide.
Until now this process has required pure water, but generally sodium-chloride (NaCl)-containing seawater is easier to obtain. Seawater is not only much more abundant than pure water but also can produce a higher yield of H2O2. The research team showed that the rate of H2O2 produced with seawater was much higher than with pure water: 24 times the amount when compared to a 2013 experiment, according to the paper.
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| Blue circles plot the rate of H2O2 production in seawater. The blue circles show the rate of H2O2 production in lab-mixed salt water, while the red squares indicate the rate of production in pure H2O. All of the mixtures had their pHs adjusted to 1.3 for this experiment, making them highly acidic and encouraging the redox reactions that charge a fuel cell. Image Credit: Nature Communications |
Now some basics about fuel cells. Photolectrochemical cells generate electrical energy through reduction-oxidation, or redox, reactions, which involve transferring electrons from one molecule to another to rearrange chemical bonds into a higher-energy state. Each half of the reaction occurs within isolated half-cells. Unlike voltaic (traditional battery) cells in which the redox reaction occurs in the spontaneous (or predisposed) direction, electrolysis pushes the chemical reaction in the opposite direction. Going against the flow, it requires an external power source, which is provided by sunlight.
A cell is composed of two electrodes, an anode and a cathode—one each within each half-cell—separated by a membrane that’s specially designed to let ions through while preventing the contents of the half-cells from spilling into one another, which would stop the reaction. During electrolysis, oxidation occurs at the anode, composed of tungsten trioxide (WO3), while reduction happens at the cathode, a cobalt-chlorin complex (CoII(Ch)). Oxidation removes the electrons from the water, leaving positively charged ions in the solution. Those electrons make their way across the connecting wire, and are then taken up from the anode by O2 as it combines with H2O during reduction to form H2O2.
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| A generic photoelectrochemical cell diagram, showing production of H2. Image Credit: University of Pittsburgh |
—Clinton Parks