The movements of electrons between cellular reductants and oxidants represent another form of energy transfer in cells.
A reductant (or reducing agent) is a substance that loses or donates electrons to another substance; the latter substance is the oxidant (or oxidizing agent).
Conversely, an oxidant is a substance that accepts electrons from another substance, the latter being the reductant.
Reactions that involve the movement of electrons between reductant and oxidant are called redox reactions.
Different chemical substances have different potentials for donating or accepting electrons. The tendency of hydrogen to dissociate
H2<===> 2H+ +2 e–
thereby releasing electrons, is used as a standard against which the tendencies of other substances to release or accept electrons is measured. The electron donor (e.g., H2 in the above reaction) and the electron acceptor (e.g., 2 H + in the above reaction) are called a redox couple or half cell. The tendency of any chemical substance to lose or gain electrons is called the redox potential and is measured in volts (V). Measurements are made using an electrode that has been standardized against the H2-2 H + couple whose redox potential is set at 0.0 under standard conditions (pH 0.0, 1M [H+], 25°C, and 1 atmosphere pressure).
This potential is noted by the symbol E0. For biochemical reactions, which normally occur at pH 7.0, the redox potential of the H2-2 H + couple is – 0.421 V; standard redox potentials at pH 7.0 are noted by the symbol The E’0 values of a number of biologically important redox couples are given in Table 9-4.
Any substance with a more positive E’0 value than another has the potential for oxidizing that substance (i.e., removing electrons from the substance with the more negative E’0 value). The greater the difference in redox potentials, the greater the energy changes involved.
The change in standard free energy, ∆G0‘, is related to E’0 as follows:
∆G0‘ = nf∆E’0,
Where n is the number of electrons exchanged per molecule, is the Faraday (96,406 J/V), and E’0 is the difference in redox potential between the more positive and the more negative members of the redox couple. For example, the oxidized form of cytochrome c can oxidize the reduced form of cytochrome b by removal of two electrons. The difference between the redox potentials of the two is
+ 0.254-(+ 0.030)= + 0.224 V
Therefore,
∆G0‘ = – 2 (96,406 J/VX0.224 V)
= – 43.19 kJ (per mole of each cytochrome)
In cells, the energy changes during redox reactions may be coupled to the synthesis of ATP.