oxidationreduction

Oxidationreduction, commonly abbreviated as redox, encompasses chemical processes in which electrons are transferred between species. In a redox reaction, one substance is oxidized (loses electrons) and another is reduced (gains electrons). The modern framework uses oxidation states—formal charges that track electron transfer—rather than the older oxygen-centered definitions. Redox concepts underpin energy flow in chemistry, biology, and environmental processes, and they explain phenomena from metal corrosion to metabolism-driven energy production. Understanding redox requires recognizing that oxidation and reduction always occur simultaneously as electrons move through the system.

A redox reaction can be split into half-reactions: an oxidation half that releases electrons and a reduction half that consumes them. The species that accepts electrons is the oxidizing agent; the one that donates electrons is the reducing agent. Common illustrations include the corrosion of iron by oxygen, where iron is oxidized, and the burning of hydrogen in oxygen, where hydrogen is oxidized and oxygen is reduced. In biology, cellular respiration transfers electrons from organic molecules to oxygen, harvested as chemical energy.

Redox chemistry is central to electrochemistry, energy storage, and environmental chemistry. Redox potentials quantify the tendency of a species to gain or lose electrons, with standard electrode potentials providing reference values. Applications include batteries, electroplating, and wastewater treatment, as well as metabolic pathways. Equations are balanced by electron-counting methods; in solution, half-reactions are balanced within acidic or basic media, often guided by the Nernst equation.