 Redox

A redox chemical reaction couples oxidation and reduction. In the electrode examples here, oxidation releases Electrons and reduction accepts them. ^[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.^[ref]

The Electron accounting need not describe free Electrons travelling between spatially separated chemical species. For example, in $2{\mathrm{H}}_2+{\mathrm{O}}_2⟶2{\mathrm{H}}_2\mathrm{O}$, 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: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.

Redox, short for reduction-oxidation, refers to a class of chemical reactions in which the oxidation states of atoms are changed. These reactions involve the transfer of electrons between species, leading to the oxidation of one substance and the reduction of another. - **Oxidation** is the process where an atom or molecule loses electrons, resulting in an increase in oxidation state. - **Reduction** is the process where an atom or molecule gains electrons, resulting in a decrease in oxidation state.