Oxidation is a chemical process in which a substance loses electrons, increases in oxidation state or, in older chemical terminology, combines with oxygen. The most general modern definition is based on electron loss or an increase in oxidation number. The word originated from the historical importance of reactions involving oxygen, particularly combustion and the formation of metal oxides. Modern chemistry uses the electron and oxidation state definitions because they apply to reactions that contain no oxygen at all.
For example, when a sodium atom loses one electron to form a sodium ion, sodium is oxidised because its oxidation state increases from zero to plus one. Oxidation is normally coupled with reduction because electrons removed from one species must be accepted by another species in a redox process. Oxidation therefore describes one side of an electron transfer process rather than an isolated event in which electrons simply disappear.
Oxidation states provide a systematic way to track electron redistribution. An oxidation state is a formal bookkeeping value assigned according to established rules rather than a direct measurement of an atom's actual local charge in every chemical environment. When the oxidation state of an element increases during a reaction, that element has been oxidised. When it decreases, the element has been reduced.
For example, iron can be oxidised from an oxidation state of plus two to plus three by losing one electron in an appropriate chemical process. Carbon can undergo oxidation when its formal oxidation state increases, even if no electron is transferred as a freely moving particle between isolated atoms. Oxidation state changes are particularly useful for covalent compounds where electron density is shared rather than transferred completely. They provide a consistent framework for analysing complex redox reactions.
Oxidation can occur in many different chemical contexts. Metals can be oxidised when they react with oxygen, halogens, acids or other oxidising agents. Iron rusting involves oxidation of iron coupled with reduction of oxygen and associated electrochemical processes. Combustion commonly involves oxidation of a fuel, although combustion and oxidation are not identical concepts because oxidation can occur without fire or even without oxygen.
Organic compounds can undergo oxidation when carbon becomes bonded to more electronegative atoms, loses hydrogen in suitable transformations or experiences an increase in formal oxidation state. Alcohols can be oxidised to aldehydes, ketones or carboxylic acids depending on their structure and reaction conditions. Biological metabolism also contains many oxidation reactions, allowing cells to extract energy from nutrients through controlled electron transfer pathways.
An oxidising agent causes another substance to undergo oxidation by accepting electrons or otherwise participating in a process that increases its oxidation state. The oxidising agent itself is reduced. This relationship is essential because oxidation cannot occur in isolation in an ordinary redox reaction. If one species loses electrons, another species must gain them. Common oxidising agents include oxygen, halogens, permanganate under appropriate conditions, dichromate in suitable acidic systems and many other chemical species.
The strength and behaviour of an oxidising agent depend on the reaction environment, including solvent, pH, concentration, temperature and electrode potential where electrochemical analysis applies. Calling a substance an oxidising agent therefore describes its role in a particular chemical process rather than an absolute property independent of conditions.
Oxidation is fundamental to electrochemistry. At an anode, oxidation occurs by definition in both galvanic and electrolytic electrochemical cells. In a galvanic cell, a spontaneous redox reaction produces electrical energy, with oxidation at one electrode supplying electrons to the external circuit. In electrolysis, an external electrical source drives a nonspontaneous process, but oxidation still occurs at the anode.
The mnemonic oxidation at the anode is useful because electrode names are defined by reaction type rather than by whether the electrode is positive or negative. In a galvanic cell the anode is negative, while in an electrolytic cell the anode is positive. This distinction prevents a common misconception that oxidation always occurs at a negatively charged electrode. The defining feature is electron loss at the anode.
Biological systems use oxidation as part of controlled energy metabolism. During cellular respiration, carbon compounds are progressively oxidised while electron carriers such as NAD plus and FAD accept electrons in specific biochemical steps. The electrons can ultimately contribute to the formation of water through the electron transport chain, while the released energy is coupled to ATP production. Photosynthetic organisms also carry out oxidation reactions, including the oxidation of water during the light dependent reactions of photosynthesis.
Enzymes control these processes so that electron transfer occurs through multiple stages rather than through uncontrolled combustion. Biological oxidation is therefore not synonymous with burning. It is a broad category of controlled electron redistribution that can occur at ordinary biological temperatures through specialised catalysts and molecular pathways.
Oxidation has practical importance in corrosion, combustion, metallurgy, environmental chemistry and industrial synthesis. Corrosion occurs when materials, particularly metals, undergo electrochemical oxidation in contact with their environment. Preventing corrosion can involve coatings, alloy design, cathodic protection, corrosion inhibitors and control of environmental conditions. Oxidation reactions are also used deliberately to manufacture chemicals, modify organic molecules and treat pollutants. In analytical chemistry, oxidation states help identify reaction stoichiometry and determine electron balances. In atmospheric chemistry, oxidation processes transform volatile organic compounds and pollutants into other species, sometimes producing secondary pollutants. Oxidation is therefore neither inherently beneficial nor harmful. Its consequences depend on the chemical system, products formed, reaction rate and environmental context.
Oxidation is best distinguished from several related but narrower concepts. Oxidation is not simply reaction with oxygen, although oxygen reactions provided the historical name. It is not synonymous with combustion, because many oxidations occur slowly or without visible heat or flame. It is also not identical to corrosion, which is a particular set of usually unwanted degradation processes involving materials and their environment.
The most general modern definition is an increase in oxidation state or, where electron transfer can be described explicitly, loss of electrons. Oxidation is almost always paired with reduction in a redox reaction. Understanding this relationship allows apparently different phenomena such as rusting, battery operation, metabolic electron transfer and combustion to be described using a common chemical framework.
