(Redirected from
Oxidation)
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"Reduced" redirects here. For other senses of that word, see reduction.
Redox reactions include all chemical processes in which atoms have their oxidation number (oxidation state) changed.
This can be a simple redox process, such as the oxidation of carbon to yield carbon dioxide, it could be the reduction of carbon by hydrogen to yield methane, or it could be the oxidation of sugar in the human body, through a series of very complex electron transfer processes.
The term redox comes from the two concepts of reduction and oxidation. It can be explained in simple terms:
- Oxidation describes the loss of an electron by a molecule, atom or ion
- Reduction describes the uptake of an electron by a molecule, atom or ion
This can be remembered easily with the mnemonic OIL RIG Oxidation Is Losing (electrons) and Reduction Is Gaining (electrons)
However, these descriptions (though sufficient for many purposes) are not truly correct. Oxidation and reduction properly refer to a change in oxidation number—the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation number, and reduction as a decrease in oxidation number. In practice, the transfer of electrons will always cause a change in oxidation number, but there are many reactions which are classed as "redox", though no electrons are transferred (such as those involving covalent bonds).
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 in another way, the oxidant removes electrons from the other substance, and is thus reduced itself. 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 (O2, O3, F2, Cl2, Br2).
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 the substance. 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, 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.
Oxidation in industry
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.
Former meaning (oxygen/hydrogen)
Formerly, oxidation simply meant the addition of oxygen or the removing of hydrogen (hence the name oxidation), and reduction was removal of oxygen or the addition of hydrogen. Currently, however, the terms are normally used in a more general sense, describing electron movement.
That said, the definitions of oxidation and reduction mentioned here are still pervasive in Organic Chemistry.
Examples of redox reactions
A good example is the reaction between hydrogen and fluorine:
- H2 + F2 → 2HF
We can write this overall reaction as two half-reactions: an oxidation reaction:
- H2 → 2H+ + 2e-
and a reduction reaction:
- F2 + 2e- → 2F-
Analysing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above).
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:
- H2 → 2H+ +
2e-
- +
2e- + F2 → 2F-
- ---------------------
- H2 + F2 → 2H+ + 2F-
And the ions combine to form hydrogen fluoride:
2H+ + 2F- → 2HF
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 or tarnishing):
- 4Fe + 3O2 → 2 Fe2O3.
Redox reactions in biology
Much biological energy is stored and released by means of redox reactions. Photosynthesis involves the reduction of carbon dioxide into sugars and the oxidation of water into molecular oxygen. The reverse reaction, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce nicotinamide adenine dinucleotide (NAD+), which then contributes to the creation of a proton gradient, which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, mitochondria perform similar functions. See Membrane potential article.
The term redox state is often used to describe the balance of NAD+/NADH and NADP+/NADPH in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., lactate and pyruvate, beta-hydroxybutyrate and acetoacetate) whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis.
See also
