Page 112 - Geochemical Remote Sensing of The Sub-Surface

Spontaneous potentials and electrochemical cells 89

Fig. 3-2. Differences in oxidation potential around the electrodes in an electrolytic cell (requires
external power) and a voltaic cell (spontaneous) (from Hamilton, 1998).

potential of half reactions measured against that of the H+-H2 half cell. Any reducing
agent in the centre column is capable of spontaneously reducing any oxidising agent in
the left-hand column that is lower in the table. Standard voltages are usually calculated
for 25~ and assume molar concentrations of all species in solution and partial pressures
of 1 atmosphere of all gases involved in the reaction. Concentration changes of up to
several orders of magnitude and temperature changes up to several 10s of degrees
Celsius have relatively small effects on the voltages in a given reaction. The effect of
concentration can be calculated using the Nernst equation. For the reaction:

aA + b B ~ cC + dD

the Nernst equation takes the following form (at 25~

E = E ~ {0.0591/n} 9 log,0{([C] c , [D]d)/([A]a,[B]b)}

where E = reaction voltage; E ~ standard electrode potential of the reaction; and n =
number of electrons involved in the redox reaction.
The Nernst equation demonstrates that a change in concentration of a species
involved in the reaction does not change the final voltage if the concentrations of all
species change by a similar factor. For example, a hundred-fold dilution of the species
involved in the Cu-Zn reaction to 0.01 molar still produces 1.10 volts because the ratio
of [ZnZ+]/[Cu 2+] is the same. It takes very large changes in the ratio to change voltage
significantly and therefore it is often acceptable to use standard voltages to qualitatively
predict the spontaneity of a reaction. However, a problem arises with the application of