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Why does combining hydrogen and oxygen typically produce water rather than hydrogen peroxide?



Chemists Joel Rosenthal and Daniel G. Nocera of the Massachusetts Institute of Technology provide this answer.

When molecular hydrogen (H2) and oxygen (O2) are combined and allowed to react together, energy is released and the molecules of hydrogen and oxygen can combine to form either water or hydrogen peroxide. These two processes are represented by the two chemical equations shown at right. Chemists use redox half-reactions to describe thermodynamic processes like the ones embodied by such equations. For both of the reactions shown, the hydrogen molecules are oxidized and the oxygen atoms are reduced. Accordingly, each of the reactions below is described by a combination of two half-reactions--one corresponding to a chemical oxidation and another corresponding to a reduction.

The redox half-reaction for hydrogen oxidation is relatively simple and is shown on the left side of the scheme below. In this oxidation, a molecule of hydrogen gas is ionized to two electrons and two protons. Writing a half-reaction for oxygen reduction is more complicated, since oxygen can be reduced by either one, two or four electrons, as shown by the square redox scheme to the right, below. In most chemical reactions, molecular oxygen is reduced along the red and blue pathways highlighted in this redox scheme. The complete reduction of O2 by four electrons (4e- + 4H+, blue horizontal pathway) generates two equivalents of water whereas the corresponding two-electron reduction (2e- + 2H+, red diagonal pathway) yields hydrogen peroxide. Both the two- (¿G¿ = -0.695 V) and four-electron (¿G¿ = -1.229 V) reductions of O2 are energetically downhill, but more than half a volt of energy is squandered in the former reaction. Accordingly, biological processes coupled to O2 reduction, such as cellular respiration, are highly selective for the complete 4e- + 4H+ pathway in order to maximize the energy available for ATP synthesis. The selective reduction of oxygen to water in such biological systems is crucial, not only in order to maximize the energy produced for cellular metabolism but also because hydrogen peroxide is a powerful oxidant and cytotoxin, which harms living cells.

Given the energetics presented above, there is a strong thermochemical bias for the production of water over hydrogen peroxide when H2 and O2 are reacted together. For instance, when hydrogen gas is burned in the presence of oxygen, a large amount of energy is released and water is produced as the major product. In cases where the reaction is more controlled, however, such as the consumption of hydrogen and oxygen in a fuel cell, the mechanism and kinetics of the O2 reduction process can complicate issues greatly. For instance, the delivery of the protons and electrons derived from the ionization of hydrogen (see redox half-reaction above) to a molecule of oxygen has to be precisely controlled via a process know as proton-coupled electron transfer in order to ensure that the complete four-electron reduction of O2 dominates. Platinum metal is capable of serving as a catalyst that brandishes exquisite selectivity for the four-electron reduction of oxygen to water, and accordingly lies at the heart of fuel cell design and function. Given that platinum is rare and extremely expensive, current research is aimed at the development of structural and functional models for oxygen activation and reduction to water via proton-coupled electron transfer. Similar strategies are also being exploited to drive the energetically uphill reverse reaction, in which hydrogen is produced from water using solar energy. The success of both these areas of work may ultimately prove crucial to the development and sustainability of a global hydrogen economy.

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