Redox (source code)
= Redox
{wiki}
= Oxidation-reduction reaction
{synonym}
A <redox> <chemical reaction> couples <oxidation> and <reduction>. In the <electrode> examples here, <oxidation> releases <electrons> and <reduction> accepts them. <Chemistry 2e by Paul Flowers and others>[§4.2, “Oxidation-Reduction Reactions”: paired electron loss and gain]{ref}
<Redox> can be recognized by changes in <oxidation state>: an increase records <oxidation>, a decrease <reduction>. These are formal <electron> assignments, distinct from measured <partial atomic charges>. <Half-reactions> expose the matching loss and gain, making <electric charge conservation> visible.<Chemistry 2e by Paul Flowers and others>[§4.2, “Oxidation-Reduction Reactions”: oxidation-number definition and complementary processes]{ref}
The <electron> accounting need not describe free <electrons> travelling between spatially separated <chemical species>. For example, in $2\mathrm H_2+\mathrm O_2\longrightarrow2\mathrm H_2\mathrm O$, <hydrogen> changes from <oxidation state> $0$ to $+1$ and <oxygen> from $0$ to $-2$. Four <hydrogen> <chemical atoms> lose four formally assigned <electrons> in total; two <oxygen> <chemical atoms> gain four. The <water> product has <covalent bonds>, and these assignments do not turn it into separated <ions>. See <ionic approximation of bonds> and <electron transfer and ionic character>.
By contrast, <neutralization> can rearrange <chemical bonds> without changing <oxidation states>:
$$
\mathrm H^+(aq)+\mathrm{OH}^-(aq)\longrightarrow\mathrm H_2\mathrm O(l).
$$
<Hydrogen> remains $+1$ and <oxygen> remains $-2$. The aqueous <proton> notation abbreviates its interaction with <water>. This <acid-base chemistry> example conserves <electric charge> but has no matching increase and decrease of <oxidation state>; movement of <ions> or a change of <chemical bonds> alone does not establish <redox>.
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