Redox

Redox reactions, also known as reduction–oxidation reactions, refer to a broad class of chemical processes in which electrons are transferred between reacting species. These changes in electron distribution are accompanied by corresponding changes in oxidation states. Although commonly associated with reactions involving oxygen, the term now encompasses all processes that involve electron loss or gain, regardless of whether oxygen is present. Redox principles underpin many chemical, biological, industrial, and environmental systems.

Fundamental Concepts

A redox reaction comprises two simultaneous half-reactions: oxidation, in which a species loses electrons, and reduction, in which a species gains electrons. These processes cannot occur independently, as the electrons lost by one species must be accepted by another. The substance that donates electrons is called the reducing agent or reductant, while the species that accepts electrons is known as the oxidising agent or oxidant.
Oxidation involves either an increase in oxidation state or direct electron loss, exemplified by sodium donating its outer electron to fluorine in the formation of sodium fluoride. In this reaction, sodium is oxidised and fluorine is reduced. Both steps proceed simultaneously, producing an overall reaction that reflects complete electron transfer. The reductant is consequently oxidised, while the oxidant is reduced.
Two major classes of redox reactions are recognised:

• Electron-transfer reactions, in which electrons move directly from a donor to an acceptor.
• Atom-transfer reactions, in which a full atom—often containing or associated with electrons—is transferred. Rusting illustrates atom transfer, where iron atoms transfer electrons to oxygen as iron converts to iron oxide.

Redox reactions can release or absorb considerable energy. Certain combinations of strong oxidants and reducing agents lead to violent reactions, such as the vigorous ignition observed when powdered potassium permanganate contacts mild reducing agents.

Terminology and Historical Development

The term redox is a portmanteau of reduction and oxidation, introduced in 1928. Historically, oxidation referred to reactions with oxygen, such as the formation of metal oxides, while reduction described processes that removed oxygen from a compound, typically by heating ores to obtain metals. Antoine Lavoisier’s experiments clarified that the change in mass reflected oxygen loss rather than the creation of metal, leading to a modern understanding of reduction as electron gain.
Current definitions emphasise electron transfer. Reducing agents are electron donors; oxidising agents are electron acceptors. A redox pair consists of a reduced species and its corresponding oxidised form. Understanding these pairs is essential for analysing both chemical and electrochemical systems.
In living organisms, the concept of reducing equivalents is widely used. These refer to species—such as electrons or hydrogen atoms—that can transfer the equivalent of one electron during biochemical reactions. Compounds such as NADH and FADH₂ carry reducing equivalents critical to cellular respiration.

Oxidants and Reductants

Oxidants are substances capable of causing oxidation by removing electrons from another species. They are typically highly electronegative or exist in high oxidation states. Common examples include oxygen, fluorine, chlorine, nitrates, and permanganates. In industrial and laboratory contexts, oxidisers may initiate reactions or cause combustion, with nitric acid serving as a strong oxidising agent.
Reductants act oppositely, supplying electrons to another species. Metal elements of low electronegativity—lithium, sodium, magnesium, iron, zinc, and aluminium—readily donate electrons and thus function as effective reducing agents. Hydride-donating reagents such as sodium borohydride and lithium aluminium hydride are widely used in organic synthesis, particularly in the reduction of carbonyl compounds to alcohols. Hydrogen gas also serves as a reductant in catalytic hydrogenation processes.

Mechanisms and Reaction Rates

Redox reactions display varied mechanisms and proceed at different rates. Electron-transfer reactions are often rapid, occurring almost instantaneously upon mixing reactants. In contrast, atom-transfer reactions may involve multiple intermediate steps and proceed more slowly.
Two primary pathways govern electron-transfer mechanisms:

• Inner-sphere electron transfer, where a ligand bridges the two species and facilitates electron movement.
• Outer-sphere electron transfer, where electron transfer occurs without direct bonding between reactants.

Reaction energetics may be evaluated using bond energies, ionisation energies, and solvation characteristics. These assessments help predict whether a reaction will proceed and determine its thermodynamic feasibility.

Standard Electrode Potentials

In electrochemistry, redox reactions are analysed through standard electrode potentials, or reduction potentials, denoted E°. These potentials measure the tendency of a species to gain electrons under standard conditions. The standard hydrogen electrode is assigned a potential of zero volts, serving as the reference point.
Oxidising agents with strong tendencies to gain electrons exhibit positive reduction potentials; for example, fluorine has a high potential of +2.866 V. Weaker oxidising agents exhibit lower or negative potentials. Conversely, oxidation potentials measure the tendency of reducing agents to lose electrons, although they do not represent actual electrode potentials.
The overall cell voltage in an electrochemical system is obtained by combining the potentials of the reduction and oxidation half-reactions. These values determine the direction and magnitude of electron flow in galvanic and electrolytic cells.

Examples of Redox Reactions

A wide range of redox reactions underpin industrial processes, environmental phenomena, and everyday chemical transformations.
Hydrogen–fluorine reaction: Hydrogen gas reacts spontaneously with fluorine to produce hydrofluoric acid, releasing large quantities of energy. Hydrogen is oxidised to protons, and fluorine is reduced to fluoride ions. Combining these half-reactions yields the overall product HF.
Metal displacement reactions: In displacement processes, a metal atom replaces another metal in solution. A classic example occurs when zinc metal is placed in copper(II) sulfate solution. Zinc donates electrons, becoming Zn²⁺, while copper ions accept electrons and are deposited as metallic copper. This reaction underlies galvanic cell operation, where zinc serves as the anode and copper as the cathode.
Biological and environmental redox processes: Denitrification, the reduction of nitrate to nitrogen gas in the presence of acids, is a key step in the nitrogen cycle. Combustion reactions, including the burning of hydrocarbons in engines, produce water and carbon dioxide through extensive oxidation processes. Incomplete oxidation may yield carbon monoxide and other partially oxidised intermediates.