Standard electrode potential

The IUPAC "Gold Book" defines it as; "the value of the standard emf (electromotive force) of a cell in which molecular hydrogen under standard pressure is oxidized to solvated protons at the left-hand electrode".

[1] The basis for an electrochemical cell, such as the galvanic cell, is always a redox reaction which can be broken down into two half-reactions: oxidation at anode (loss of electron) and reduction at cathode (gain of electron).

The galvanic cell potential results from the voltage difference of a pair of electrodes.

The electrode potentials are independent of the number of electrons transferred —they are expressed in volts, which measure energy per electron transferred—and so the two electrode potentials can be simply combined to give the overall cell potential even if different numbers of electrons are involved in the two electrode reactions.

For practical measurements, the electrode in question is connected to the positive terminal of the electrometer, while the standard hydrogen electrode is connected to the negative terminal.

However, if the solicitations exerted on the system are sufficiently small and applied slowly, one can consider an electrode to be reversible.

Such a system is far from equilibrium and continuously submitted to important and constant changes in a short period of time The larger the value of the standard reduction potential, the easier it is for the element to be reduced (gain electrons); in other words, they are better oxidizing agents.

For example, F2 has a standard reduction potential of +2.87 V and Li+ has −3.05 V: The highly positive standard reduction potential of F2 means it is reduced easily and is therefore a good oxidizing agent.

In contrast, the greatly negative standard reduction potential of Li+ indicates that it is not easily reduced.

Instead, Li(s) would rather undergo oxidation (hence it is a good reducing agent).