Oxidant - Related Links
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Journal of Clinical & Experimental PharmacologyCardioprotective Role of Saxagliptin through Antioxidant Mechanism in Experimental Myocardial Infarction in STZ Induced Diabetic Rats
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Journal of Biosensors & BioelectronicsElectrochemical Evaluation of the Antioxidant Capacity of Phenolic Compounds in Virgin Olive Oil
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Journal of Microbial & Biochemical TechnologyTamarix nilotica (Ehrenb) Bunge: A Review of Phytochemistry and Pharmacology
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An oxidizing agent (also oxidant, oxidizer or oxidiser) is the element or compound in an oxidation-reduction (redox) reaction that accepts an electron from another species. Because the oxidizing agent is gaining electrons (and is thus often called an electron acceptor), it is said to have been reduced.
The oxidizing agent itself is reduced, as it is taking electrons onto itself, but the reactant is oxidized by having its electrons taken away by the oxidizing agent. Oxygen is the prime (and eponymous) example among the varied types of oxidizing agents, but oxidisers (e.g., chlorine trifluoride) do not necessarily donate or contain oxygen.
- The oxidizing agent takes electrons from another species, and thus itself is reduced.
- The reducing agent gives electrons to another species, and thus itself 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 (rust) through the oxidation of iron;
- 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 (balanced) "half-reactions" that occur concurrently:
- Oxidation half reaction: 4Fe → 4Fe3+ + 12e−
- Reduction half reaction: 3O2 + 12e− → 6O2−
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).
In one definition, an oxidising agent accepts - or gains - electrons. In this context, the oxidizing agent is called an electron acceptor and 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 oxidizing 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.
The U.S. Department of Transportation defines Oxidizing agent specifically. There are two definitions for oxidizing agents governed under DOT regulations. These two are Class 5; Division 5.1 and Class 5; Division 5.2. Division 5.1 "means a material that may, generally by yielding oxygen, cause or enhance the combustion of other materials." Division 5.1 of the DOT code applies to solid oxidizers "if, when tested in accordance with the UN Manual of Tests and Criteria (IBR, see § 171.7 of this subchapter), its mean burning time is less than or equal to the burning time of a 3:7 potassium bromate/cellulose mixture." 5.1 of the DOT code applies to liquid oxidizers "if, when tested in accordance with the UN Manual of Tests and Criteria, it spontaneously ignites or its mean time for a pressure rise from 690 kPa to 2070 kPa gauge is less than the time of a 1:1 nitric acid (65 percent)/cellulose mixture."
Common oxidizing 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)
- Potassium nitrate (KNO3), the oxidizer in black powder
- Tollens' reagent
- 2,2'-Dipyridyldisulfide (DPS)
Common oxidizing 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−|
|ClO− 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|
- Australian Dangerous Goods Code, 6th Edition
- 49 CFR 172.127 General Requirements for Shipments and Packagings; Subpart D
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