oxidations

Oxidations are processes in which the oxidation state of an element increases, typically through loss of electrons. In classical terms, oxidation was associated with reaction with oxygen, but in modern chemistry oxidation is defined by electron transfer and is always paired with reduction. In redox reactions, the species that donates electrons is the reducing agent and the species that accepts electrons is the oxidizing agent.

Common examples include the corrosion of iron by oxygen to form iron oxides, where iron is oxidized from 0 to a positive oxidation state, and combustion, in which fuels are oxidized by oxygen to produce carbon dioxide and water. Oxidation states provide a bookkeeping system for tracking electron transfer during these processes.

Oxidations occur in many contexts. Chemical oxidations use oxidants such as oxygen, hydrogen peroxide, permanganate, or dichromate to increase the oxidation state of target molecules. Biological oxidations occur in metabolism, such as cellular respiration, where substrates are oxidized and oxygen serves as the final electron acceptor. Electrooxidation and photooxidation describe oxidation driven by electrical energy or light, respectively. Industrially important oxidations include the conversion of ethylene to ethylene oxide and the oxidation of alcohols to aldehydes, ketones, or acids, often implemented with catalysts.

Key concepts include oxidation states, oxidizing and reducing agents, and the balance of electron transfer in redox couples. Oxidations can be exothermic and form reactive intermediates, with implications for corrosion, energy production, synthesis, and environmental chemistry.