Redox

Redox, short for reduction-oxidation, describes chemical processes in which electrons are transferred between species. Oxidation is an increase in oxidation state and is typically accompanied by loss of electrons; reduction is a decrease in oxidation state and involves gain of electrons. Because electrons cannot appear or disappear, redox reactions couple an oxidation with a corresponding reduction. The species that donates electrons is the reducing agent; the species that accepts electrons is the oxidizing agent.

Chemists balance redox changes using half-reactions and track them with oxidation states. Redox couples have potentials that indicate the tendency to gain or lose electrons; standard electrode potentials are measured under standard conditions. The Nernst equation relates these potentials to concentrations, allowing calculation of cell potentials and reaction feasibility.

Common examples include rusting of iron, where iron is oxidized and oxygen is reduced, and combustion, where fuel molecules are oxidized by oxygen. In electrochemical cells, redox chemistry powers devices: the anode undergoes oxidation and the cathode reduction; electrons flow through an external circuit producing electrical work.

Redox is central to biology, with organisms carrying out redox reactions to extract energy from nutrients. In metabolism, NAD+/NADH and FAD/FADH2 act as carriers. In environmental and materials science, redox potential governs metal speciation, contaminant mobility, corrosion, and the operation of batteries and fuel cells. Analytical chemistry uses redox titrations and colorimetric redox indicators to monitor oxidation state changes.