Editing Chemical reaction (section)

===Oxidation and reduction===
[[File:redox reaction.png|thumb|right|250px|Illustration of a redox reaction]]
[[File:Common-salt.jpg|thumb|right|250px|[[Sodium chloride]] is formed through the redox reaction of sodium metal and chlorine gas]]

[[Redox]] reactions can be understood in terms of the transfer of electrons from one involved species ([[reducing agent]]) to another ([[oxidizing agent]]). In this process, the former species is ''oxidized'' and the latter is ''reduced''. Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation is better defined as an increase in [[oxidation state]] of atoms and reduction as a decrease in oxidation state. In practice, the transfer of electrons will always change the oxidation state, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving [[covalent]] bonds).<ref>{{cite encyclopedia | author = Glusker, Jenny P. |author-link=Jenny Glusker| contribution = Structural Aspects of Metal Liganding to Functional Groups in Proteins | editor = Christian B. Anfinsen | url = https://books.google.com/books?id=HvARsi6S-b0C&pg=PA7 | title = Advances in Protein Chemistry | volume = 42 | publisher = [[Academic Press]] | location = San Diego | year = 1991 | isbn = 978-0-12-034242-6 | page = 7}}</ref><ref>{{ cite encyclopedia | author = Guo, Liang-Hong | author2 = Allen, H. | author3 = Hill, O. | contribution = Direct Electrochemistry of Proteins and Enzymes | editor = A.G. Sykes | url = https://books.google.com/books?id=qnRkjATn0YUC&pg=PA359 | title = Advances in Inorganic Chemistry | volume = 36 | publisher = [[Academic Press]] | location = San Diego | year = 1991 | isbn = 978-0-12-023636-7 | page = 359}}</ref>

In the following redox reaction, hazardous [[sodium]] metal reacts with toxic [[chlorine]] gas to form the ionic compound [[sodium chloride]], or common table salt:
<chem display="block">2Na(s) + Cl2(g)->2NaCl(s)</chem>

In the reaction, sodium metal goes from an oxidation state of 0 (a pure element) to +1: in other words, the sodium lost one electron and is said to have been oxidized. On the other hand, the chlorine gas goes from an oxidation of 0 (also a pure element) to −1: the chlorine gains one electron and is said to have been reduced. Because the chlorine is the one reduced, it is considered the electron acceptor, or in other words, induces oxidation in the sodium – thus the chlorine gas is considered the oxidizing agent. Conversely, the sodium is oxidized or is the electron donor, and thus induces a reduction in the other species and is considered the ''reducing agent''.

Which of the involved reactants would be a reducing or oxidizing agent can be predicted from the [[electronegativity]] of their elements. Elements with low electronegativities, such as most metals, easily donate electrons and oxidize – they are reducing agents. On the contrary, many oxides or ions with high oxidation numbers of their non-oxygen atoms, such as {{chem|link=hydrogen peroxide|H|2|O|2}}, {{chem|link=permanganate|MnO|4|-}}, {{chem|link=chromium trioxide|CrO|3}}, {{chem|link=dichromate|Cr|2|O|7|2-}}, or {{chem|link=Osmium(VIII) oxide|OsO|4}}, can gain one or two extra electrons and are strong oxidizing agents.

For some [[main-group element]]s the number of electrons donated or accepted in a redox reaction can be predicted from the [[electron configuration]] of the reactant element. Elements try to reach the low-energy [[noble gas]] configuration, and therefore alkali metals and halogens will donate and accept one electron, respectively. Noble gases themselves are chemically inactive.<ref>[[#Wiberg|Wiberg]], pp. 289–290</ref>

The overall redox reaction [[Electrochemistry#Balancing redox reactions|can be balanced]] by combining the oxidation and reduction half-reactions multiplied by coefficients such that the number of electrons lost in the oxidation equals the number of electrons gained in the reduction.

An important class of redox reactions are the electrolytic [[Electrochemistry|electrochemical]] reactions, where electrons from the power supply at the negative electrode are used as the reducing agent and electron withdrawal at the positive electrode as the oxidizing agent. These reactions are particularly important for the production of chemical elements, such as [[chlorine]]<ref>[[#Wiberg|Wiberg]], p. 409</ref> or [[aluminium]]. The reverse process, in which electrons are released in redox reactions and chemical energy is converted to electrical energy, is possible and used in [[Electric battery|batteries]].