fossil-fuel economy, a number of technical aspects must be improved: for
instance, the cost of the fuel cells must be significantly lowered, means must
be established for the storage and delivery of hydrogen, and approaches
for the industrial-scale production of hydrogen must be established.
Although hydrogen is abundant, it is almost never found by itself, but is
present in chemical compounds such as hydrocarbons or water.
Both the formation of hydrogen during electrolysis and its consump-
tion in a fuel cell must be run as spontaneous reactions and therefore
will involve some free-energy loss. However, in current practice, the amount
of energy lost during these processes is considerably larger than simple
thermodynamics would require. This is due to the substantial activation
energy inherent in the redox half reactions performed at the surface of
bare metal electrodes. One of the greatest losses occurs during the four-
electron oxidation of water to generate molecular oxygen and protons
during electrolysis (the half reaction at the oxygen-producing electrode).
The overall reaction for electrolysis is:
2H 2 O →2H 2 +O 2 (12.37)
For the simplest case, this overall chemical reaction at pH 7 can be expressed
as two half reactions:
2H 2 O →4H++O 2 +4e− E 0 =−0.82 V (12.38)
4H++4e−→2H 2 E 0 =−0.42 V
The half reaction midpoint potentials listed are given under standard con-
ditions at pH 7. Considering only the free-energy difference between the
reactants and products, the magnitude of the thermodynamic potential
required for the overall reaction is about 1.24 V under standard conditions.
However, in practice, electrolysis only occurs at a useful rate between two
metal electrodes when much larger voltages, typically 2 V or more, are
applied (US Department of Energy 2003). Because the power required to
perform this process is simply given by the product of the voltage and the
current (the current determines the rate of hydrogen production), the higher
the voltage required to obtain a particular rate of hydrogen production, the
more power is dissipated per quantity of hydrogen formed. The voltage that
must be applied beyond that demanded by the free-energy difference
between reactant and products is the so-called overpotential and the greater
the overpotential used, the greater the loss of energy as heat.
The necessity for using a substantial overpotential during electrolysis
largely results from the activation energy associated with the splitting
of water on a metal surface, generating molecular oxygen and protons
(eqn 12.37). The overpotential that is required for useful levels of
hydrogen production from protons can be quite low using a platinum
electrode (in the order of 0.1 V), but for oxygen evolution the numbers