aq)
It is important to remember, however, that oxidation and reduction reactions always
occur in pairs.* This relationship is formalized by the convention of calling the
species being oxidized a reducing agent, because it provides the electrons for the reduction
half-reaction. Conversely, the species being reduced is called an oxidizing
agent. Thus, in reaction 6.22, Fe3+ is the oxidizing agent and H2C2O4 is the reducing
agent.
The products of a redox reaction also have redox properties. For example, the
Fe2+ in reaction 6.22 can be oxidized to Fe3+, while CO2 can be reduced to H2C2O4.
Borrowing some terminology from acid–base chemistry, we call Fe2+ the conjugate
reducing agent of the oxidizing agent Fe3+ and CO2 the conjugate oxidizing agent of
the reducing agent H2C2O4.
Unlike the reactions that we have already considered, the equilibrium position
of a redox reaction is rarely expressed by an equilibrium constant. Since redox reactions
involve the transfer of electrons from a reducing agent to an oxidizing agent,
it is convenient to consider the thermodynamics of the reaction in terms of the
electron.
The free energy, ÆG, associated with moving a charge, Q, under a potential, E,
is given by