When a redox operation proceeds, one species is oxidised and the other is reduced. It is useful to think of these two processes separately. For example, in the dissolution of zinc metal in acid:

Oxidation:
Reduction:

The zinc metal is oxidized, and the protons are reduced to hydrogen gas. The splitting up of the overall process into separate parts for the oxidant and reductant is conceptual, as the individual reactions may not occur on their own. When the process is split up like this, each of the reactions for the oxidant and the reductant is known as a half-reaction. The reactant and product of each of these half-reactions are known as a redox couple, ie. H⁺/H₂ and Zn²⁺/Zn.

If we write the reaction for the oxidation of zinc in reverse, then it too is a reduction, and the overall process is the difference between two reduction processes.

Overall process is - Reduction 1
- Reduction 2

Standard reduction potentials

The overall reaction in a redox process has a characteristic Gibbs free reaction energy. If we split up the redox process into a reduction part and an oxidation part, then we can assign the overall Gibbs free reaction energy into contributions from the reduction process and the oxidation process. The overall Gibbs reaction energy is therefore the difference between the reaction energies for the two reduction half-reactions.

Just as splitting up the overall redox process into two separate reduction half-reactions is conceptual, this splitting up of the reaction energy is also conceptual, and the ability to do this relies on knowing the reaction energy of one reduction half-reaction so that those of all the others can be calculated from it, as the reactions always occur in pairs.

We define the free energy for the H⁺/H₂ half-reaction as zero, and then all other reduction free energies may be calculated. In the reaction above, the overall Gibbs reaction energy may be measured, and as the H⁺/H₂ reduction couple does not contribute, the entire reaction energy is due to the Zn²⁺/Zn couple.

	ΔG = + 147 kJ
	ΔG = 0
	ΔG = + 147 kJ

Thus, the reaction energy for the Zn2+/Zn reduction couple is known, and may be converted into a standard reduction potential, E, using the relationship

ΔG = -nFE

where n is the number of electrons transferred, 2 in this case, and F is  Faraday's constant.

So we can say the the standard reduction potential of the H⁺/H₂ couple is zero, and that for the Zn²⁺/Zn couple is E = -0.76 V.

The potential for the overall reaction is therefore the difference between the standard reduction potentials for the individual half-reactions.

	E = 0 V
	E = -0.76 V
	E = E(H+/H2) - E(Zn2+/Zn) = (0) - (-0.76) = +0.76 V

The reaction as written has a positive electrode potential, and this corresponds to a negative free energy of reaction, and so the reaction is spontaneous as written.

This means that the reduction of H⁺ ions by zinc metal is favourable under standard conditions.

This further means that any reduction couple with a negative standard reduction potential, as is that for Zn²⁺/Zn, also has the thermodynamic tendency to reduce H⁺ ions.

When the standard reduction potential has been calculated, we notice that those couples with negative potentials are able to reduce H⁺, and those with positive reduction potentials are able to oxidize H₂. We can rank the couples in the order of their strength as reducing and oxidizing agents. This list is known as the electrochemical series: Ox/Red couple with strongly positive E^*, where Ox is strongly oxidizing, through to Ox/Red couple with strongly negative E^*, where Red is strongly reducing.

Strongly Oxidizing		E* = +2.87V
Strongly Reducing		E* = -3.04V