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An oxidising agent (also oxidant, oxidizer or oxidiser) is a substance that oxidizes (removes electrons from) another reactant in a redox chemical reaction. The oxidising agent is reduced by taking electrons onto itself and the reactant is oxidised by having its electrons taken away. Oxygen is the prime (and eponymous) example among the varied types of oxidising agents.
- The oxidizing agent is reduced.
- The reducing agent is oxidized.
- All atoms in a molecule can be assigned an oxidation number. This number changes when an oxidant acts on a substrate.
- Redox reactions occur when oxidation states of the reactants change.
Example of oxidation 
The formation of iron(III) oxide;
- 4Fe + 3O2 → 2Fe2O3
In the above equation, the iron (Fe) has an oxidation number of 0 before and 3+ after the reaction. For oxygen (O) the oxidation number began as 0 and decreased to 2−. These changes can be viewed as two "half-reactions" that occur concurrently:
- Oxidation half reaction: Fe0 → Fe3+ + 3e−
- Reduction half reaction: O2 + 4e− → 2 O2−
Iron (Fe) has become oxidised because its oxidation number increased and was the reducing agent because it gave electrons to the oxygen (O). Oxygen (O) has been reduced because the oxidation number has decreased and was the oxidising agent because it took electrons from iron (Fe).
Electron acceptor 
In one definition, an oxidising agent accepts - or gains - electrons. In this context, the reducing agent is called an electron donor. A classic oxidising agent is the ferrocenium ion [Fe(C5H5)2]+, which accepts an electron to form Fe(C5H5)2. Of great interest to chemists are the details of the electron transfer event, which can be described as inner sphere or outer sphere.
In more colloquial usage, an oxidising agent transfers oxygen atoms to the substrate. In this context, the oxidising agent can be called an oxygenation reagent or oxygen-atom transfer agent. Examples include [MnO4]− (permanganate), [CrO4]2− (chromate), OsO4 (osmium tetroxide), and especially [ClO4]− (perchlorate). Notice that these species are all oxides, and are in fact polyoxides. In some cases, these oxides can also serve as electron acceptors, as illustrated by the conversion of [MnO4]− to [MnO4]2−, manganate.
Dangerous materials definition 
The dangerous materials definition of an oxidising agent is a substance that is not necessarily combustible, but may, generally by yielding oxygen, cause or contribute to the combustion of other material. By this definition some materials that are classified as oxidising agents by analytical chemists are not classified as oxidising agents in a dangerous materials sense. An example is potassium dichromate, which does not pass the dangerous goods test of an oxidising agent.
Common oxidising agents 
- Oxygen (O2)
- Ozone (O3)
- Hydrogen peroxide (H2O2) and other inorganic peroxides
- Fluorine (F2), chlorine (Cl2), and other halogens
- Nitric acid (HNO3) and nitrate compounds
- Sulfuric acid (H2SO4)
- Peroxydisulfuric acid (H2S2O8)
- Peroxymonosulfuric acid (H2SO5)
- Chlorite, chlorate, perchlorate, and other analogous halogen compounds
- Hypochlorite and other hypohalite compounds, including household bleach (NaClO)
- Hexavalent chromium compounds such as chromic and dichromic acids and chromium trioxide, pyridinium chlorochromate (PCC), and chromate/dichromate compounds
- Permanganate compounds such as potassium permanganate
- Sodium perborate
- Nitrous oxide (N2O)
- Silver oxide (Ag2O)
- Osmium tetroxide (OsO4)
- Tollens' reagent
- 2,2'-Dipyridyldisulfide (DPS)
Common oxidising agents and their products 
|O2 oxygen||Various, including the oxides H2O and CO2|
|O3 ozone||Various, including ketones, aldehydes, and H2O; see ozonolysis|
|I2 iodine||I−, I3−|
|OCl− hypochlorite||Cl−, H2O|
|ClO3− chlorate||Cl−, H2O|
|HNO3 nitric acid||NO nitric oxide
NO2 nitrogen dioxide
CrO3 chromium trioxide
|Mn2+ (acidic) or MnO2 (basic)|
|H2O2, other peroxides||Various, including oxides and H2O|
See also 
- Australian Dangerous Goods Code, 6th Edition