In electrochemistry, a half-cell is a structure that contains a conductive electrode and a surrounding conductive electrolyte separated by a naturally occurring Helmholtz double layer. Chemical reactions within this layer momentarily pump electric charges between the electrode and the electrolyte, resulting in a potential difference between the electrode and the electrolyte. The typical anode reaction involves a metal atom in the electrode being dissolved and transported as a positive ion across the double layer, causing the electrolyte to acquire a net positive charge while the electrode acquires a net negative charge. The growing potential difference creates an intense electric field within the double layer, and the potential rises in value until the field halts the net charge-pumping reactions. This self-limiting action occurs almost instantly in an isolated half-cell; in applications two dissimilar half-cells are appropriately connected to constitute a Galvanic cell.

A standard half-cell consists of a metal electrode in an aqueous solution where the concentration of the metal ions is 1 molar (1 mol/L) at 298 kelvins (25 °C).[1] In the case of the standard hydrogen electrode (SHE), a platinum electrode is used and is immersed in an acidic solution where the concentration of hydrogen ions is 1M, with hydrogen gas at 1atm being bubbled through solution.[2] The electrochemical series, which consists of standard electrode potentials and is closely related to the reactivity series, was generated by measuring the difference in potential between the metal half-cell in a circuit with a standard hydrogen half-cell, connected by a salt bridge.

The standard hydrogen half-cell:

2H+(aq) + 2e → H2(g)

The half-cells of a Daniell cell:

Original equation
Zn + Cu2+ → Zn2+ + Cu
Half-cell (anode) of Zn
Zn → Zn2+ + 2e
Half-cell (cathode) of Cu
Cu2+ + 2e → Cu

See also


  1. ^ "an introduction to redox equilibria and electrode potentials". Retrieved 2024-02-11.
  2. ^ "Untitled Document". Retrieved 2024-02-11.