Reductionoxidation

Reduction-oxidation, or redox, refers to chemical processes in which electrons are transferred between species, resulting in changes to oxidation states. In a redox reaction, one substance loses electrons (oxidation) while another gains electrons (reduction). The substance that donates electrons is the reducing agent and is oxidized; the substance that accepts electrons is the oxidizing agent and is reduced. Because these changes occur in paired steps, oxidation state numbers rise for the oxidized species and fall for the reduced species.

Redox reactions are often analyzed using half-reactions, separately listing oxidation and reduction steps. When combined, the electrons lost in the oxidation half-reaction cancel with those gained in the reduction half-reaction, yielding a balanced overall equation. Redox chemistry can be represented by standard reduction potentials, which indicate the tendency of a species to be reduced and help predict reaction direction and energy yield. Such information underpins the operation of electrochemical cells, including galvanic (voltaic) cells that generate electricity and electrolytic cells that drive non-spontaneous processes.

Common examples include the rusting of iron, where iron is oxidized by oxygen in the presence of water, and cellular respiration, where glucose is oxidized to carbon dioxide with the reduction of oxygen. Redox processes are central to numerous industrial methods, environmental chemistry, and biological systems, where controlled electron transfer governs energy production, metabolism, corrosion, and pollutant transformations.