Electrode Potential and Standard Electrode Potential

Every metal electrode dipped into a solution of its own ions develops a characteristic electrical potential, arising from the competing tendencies of metal atoms to lose electrons and enter solution as ions, and of metal ions in solution to gain electrons and deposit back onto the metal. Different metals have different, measurable tendencies to do this — some metals (like zinc) lose electrons relatively easily, while others (like copper or silver) hold onto their electrons more strongly and instead preferentially gain electrons from solution.

This tendency cannot be measured for a single electrode in isolation — voltage, by its nature, is always a difference between two points, so only the potential difference between two electrodes joined into a complete cell can actually be read off a voltmeter. To build a usable, universal scale of individual electrode potentials anyway, chemists adopt a single reference electrode (the standard hydrogen electrode, discussed next) and, by convention, assign it a potential of exactly $0.00\ \text{V}$. Every other electrode's potential is then defined as the EMF of the cell formed by pairing that electrode with the standard hydrogen electrode, under carefully specified standard conditions.

The standard electrode potential, $E^{\circ}$, of a half-cell is this potential, measured with all solute concentrations at $1\ \text{M}$, any gas at $1\ \text{atm}$ pressure, and the temperature at $298\ \text{K}$ ($25^{\circ}\text{C}$). By strict convention, every $E^{\circ}$ value is tabulated as a reduction potential — that is, for the half-reaction written in the direction of gaining electrons, e.g. $\text{Cu}^{2+}(aq) + 2e^{-} \to \text{Cu}(s)$, $E^{\circ}=+0.34\ \text{V}$ — never as an oxidation potential, so that every value in a standard table can be compared on the same footing without having to remember which direction each individual value was originally written in. …