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Illustration of a redox reaction

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Redox (shorthand for reduction/oxidation reaction) describes all chemical reactions in which atoms have their oxidation number (oxidation state) changed.

This can be either a simple redox process such as the oxidation of carbon to yield carbon dioxide, or the reduction of carbon by hydrogen to yield methane (CH4), or it can be a complex process such as the oxidation of sugar in the human body through a series of very complex electron transfer processes.



Non-redox reactions, which do not involve changes in formal charge, are known as metathesis reactions.

The two parts of a redox reaction
Rusting iron
A bonfire. Combustion consists of redox reactions involving free radicals.

Oxidizing and reducing agents

Substances that have the ability to oxidize other substances are said to be oxidative and are known as oxidizing agents, oxidants or oxidizers. Put another way, the oxidant removes electrons from another substance, and is thus reduced itself. And because it "accepts" electrons it is also called an electron acceptor.

Oxidants are usually chemical substances with elements in high oxidation numbers (e.g., H2O2, MnO4, CrO3, Cr2O72−, OsO4) or highly electronegative substances that can gain one or two extra electrons by oxidizing a substance (O, F, Cl, Br).

Substances that have the ability to reduce other substances are said to be reductive and are known as reducing agents, reductants, or reducers. Put in another way, the reductant transfers electrons to another substance, and is thus oxidized itself. And because it "donates" electrons it is also called an electron donor. Reductants in chemistry are very diverse. Metal reduction—electropositive elemental metals can be used (Li, Na, Mg, Fe, Zn, Al). These metals donate or give away electrons readily. Other kinds of reductants are hydride transfer reagents (NaBH4, LiAlH4), these reagents are widely used in organic chemistry[1][2], primarily in the reduction of carbonyl compounds to alcohols. Another useful method is reductions involving hydrogen gas (H2) with a palladium, platinum, or nickel catalyst. These catalytic reductions are primarily used in the reduction of carbon-carbon double or triple bonds.

The chemical way to look at redox processes is that the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized and the oxidant or oxidizing agent gains electrons and is reduced. The pair of an oxidising and reducing agent that are involved in a particular reaction is called a redox pair.

Examples of redox reactions

A good example is the reaction between hydrogen and fluorine:

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and the reduction reaction:

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Elements, even in molecular form, always have an oxidation number of zero. In the first half reaction, hydrogen is oxidized from an oxidation number of zero to an oxidation number of +1. In the second half reaction, fluorine is reduced from an oxidation number of zero to an oxidation number of −1.

When adding the reactions together the electrons cancel:

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And the ions combine to form hydrogen fluoride:

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Other examples

  • iron(II) oxidizes to iron(III):
Fe2+ → Fe3+ + e
H2O2 + 2 e → 2 OH

overall equation for the above:

2Fe2+ + H2O2 + 2H+ → 2Fe3+ + 2H2O
2NO3 + 10e + 12 H+ → N2 + 6H2O
  • iron oxidizes to iron(III) oxide and oxygen is reduced forming iron(III) oxide (commonly known as rusting, which is similar to tarnishing):
4Fe + 3O2 → 2 Fe2O3

Redox reactions in industry

The primary process of reducing ore to produce metals is discussed in the article on Smelting.

Oxidation is used in a wide variety of industries such as in the production of cleaning products.

Redox reactions are the foundation of electrochemical cells.

Redox reactions in biology

Many important biological processes involve redox reactions.

Cellular respiration, for instance, is the oxidation of glucose (C6H12O6) to CO2 and the reduction of oxygen to water. The summary equation for cell respiration is:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O
The process of cell respiration also depends heavily on the reduction of NAD+ to NADH and the reverse reaction (the oxidation of NADH to NAD+). Photosynthesis is essentially the reverse of the redox reaction in cell respiration:
6 CO2 + 6 H2O + light energy → C6H12O6 + 6 O2



Redox cycling


Examples of redox cycling-inducing molecules are the herbicide paraquat and other viologens and quinones such as menadione. [2]PDF (2.76 MiB)

Balancing redox reactions

Describing the overall electrochemical reaction for a redox process requires a balancing of the component half reactions for oxidation and reduction. For reactions in aqueous solution, this general involves adding H+ , OH- ion, H2O and electrons to compensate the oxidation changes.

Acid medium

In acid medium H+ ions and water are added to half reactions to balance the overall reaction. For example, when manganese (II) reacts with sodium bismuthate.

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The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (i.e. multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa). Addition gives:

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Reaction balanced:

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Similarly for a propane fuel cell under acidic conditions:

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Balancing the number of electrons involved gives:

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Equation balanced:

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Basic medium

In basic medium OH- ions and water are added to half reactions to balance the overall reaction. For example on reaction between potassium permanganate and sodium sulfite.

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Balancing the number of electrons in the two half-cell reactions gives:

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Equation balanced:

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See also


  1. Hudlický, Miloš (1996). Reductions in Organic Chemistry. Washington, D.C.: American Chemical Society, 429. ISBN 0-8412-3344-6. 
  2. Hudlický, Miloš (1990). Oxidations in Organic Chemistry. Washington, D.C.: American Chemical Society, 456. ISBN 0-8412-1780-7. 

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