Nernst equation: a standalone construct

Walther Nernst introduced the relation in 1889 in Die elektromotorische Wirksamkeit der Ionen (Zeitschrift für physikalische Chemie), deriving how the electrical potential of an electrode depends on the concentrations, or more precisely the activities, of the chemical species taking part in the reaction at its surface. The equation states that this potential equals a fixed reference value plus a term proportional to the absolute temperature and to the logarithm of the ratio of the reacting species, so that as reactant and product amounts shift, the measured voltage shifts in a predictable way. In its modern form, standardised in graduate electrochemistry texts such as Bard and Faulkner (2001), it links the standard electrode potential, the gas constant, the temperature, the number of electrons transferred, and the Faraday constant, and it reduces to the thermodynamic condition for equilibrium when no net current flows. The same relation is used across chemistry, physics, and physiology, where it gives the equilibrium potential of an ion across a cell membrane, yet it names one construct rather than several, since these are one equation applied in different settings. What makes it a landmark is that it ties a bench measurement, a voltage, directly to the underlying thermodynamics and composition of a system, which is why it grounds both analytical electrochemistry and the quantitative study of membrane potentials.[1][2]

Unit

A physical relationship, expressed as an equation, giving electrode or ion equilibrium potential in volts as a function of temperature and species concentration

Recovery means

Not applicable

Measured by

Applied by measuring a cell or membrane voltage with a potentiometer or high-impedance voltmeter and comparing it to the value computed from known concentrations, temperature, and the standard potential