# Why do single elements in a molecule have an oxidation state of zero?

For example, take $\ce{O2}$.

The oxidation state of oxygen is -2, yet once its in a molecule its oxidation state becomes zero?

How is this so?

• Very strictly speaking, I don't think this must necessarily be true. It is possible to imagine allotropes of a pure element which would consist of ions and thus would require non-zero oxidation numbers. For example, the hypothetical octanitrogen $\ce{N8}$ allotrope could be composed of the azide anion $\ce{N3^{-}}$ and the pentazenium cation $\ce{N5^{+}}$. Granted, I do not believe this to be the case for any element, at least not in ambient conditions. – Nicolau Saker Neto Jul 16 '17 at 12:13

So when you say oxygen's oxidation state is -2, this is only with respect to other elements of lower electronegativity. Now in the case of oxygen, the only other element with higher electonegativity is fluorine. For example, the molecule Oxygen Difluoride ($\ce{OF2}$) has oxygen in the +2 oxidation state because fluorine is more electronegative! For any other element in combination with oxygen, however, we assume the electrons spend their time with the oxygen filling its outer valence shell, thus the -2 oxidation state.