Oxidation Numbers

Oxidation number (ON), also called oxidation state (OS), is the hypothetical total number of electrons that an atom either gains or loses in order to form a fully ionic bond with another atom. Conceptually, ON may be positive (loss of electrons), negative (gain of electrons), or zero (equal share of electrons, e.g., in a covalent bond between two atoms of the same element). The ON of an atom does not represent the “real” formal charge on that atom, it is not the same ionic charge, and ON values in a given compound vary depending on the choice of electronegativity scale used in their calculation. Nevertheless, ON is important in understanding the nomenclature conventions of inorganic compounds, and a useful tool for keeping track of electron transfers, which is essential in the study of redox reactions.

Oxidation numbers are typically represented by integers, although the average oxidation number of an element may be a fraction if atoms of the same element in a compound have different oxidation numbers. For instance, the average oxidation number of C in propane (CH3CH2CH3) is -8/3, because the first and third carbon atoms in propane each has an oxidation number of -3, while that of the central one is -2.

Rules for calculations of oxidation numbers

  1. Atoms in free elements (or neutral compounds that contains atoms of only one element) have an oxidation number of zero (e.g., Na, H2, O2, O3, P4, S8)
  2. All oxidation numbers of atoms in a neutral compound must add up to zero.
  3. The oxidation number of a monatomic ion is equal to the ion charge (e.g., the oxidation number of iron in the Fe3+ ion is +3, and the oxidation number of iron in the Fe2+ ion is +2)
  4. In a polyatomic ion the sum of the oxidation numbers of all atoms is equal to the total charge on the ion.
  5. Hydrogen has oxidation number of +1 when bonding with nonmetals (e.g., H2O); and of -1 if bonding with metals (e.g., CaH2)
  6. Atoms in Group 1 (or 1A) in the Periodic Table always has an oxidation number of +1 when bonding with atoms of other elements (e.g., Li3N and Na2S).
  7. Atoms in Group 2 (or 2A) have an oxidation number of +2 when forming compounds with atoms of other elements (e.g., Mg in Mg3N2)
  8. Oxygen has an oxidation number of -2 in almost all known compounds of oxygen. However, other oxidation numbers of oxygen can also be found in uncommon compounds, for instance, −1 in peroxides (H2O2), −1⁄3 in ozonides (O3–), 0 in hypofluorous acid (HOF), +1 in dioxygen difluoride (O2F2), and +2 in oxygen difluoride (OF2).
  9. Atoms in Group 17 (7A or Halogens) usually have an oxidation number of -1 (e.g., Chlorine in HCl or Bromine in ZnBr2), except when bonding with oxygen, for example, chlorine has an oxidation number of +1 in compound Cl2O.
  10. The oxidation number of Florine (F) is always -1

For cases that cannot be found in the above rules, knowledge of electronegativity can become handy. Electronegativity is a useful tool in determining ON for atoms in complex compounds, especially organic compounds which are formed mostly by covalent bonds between atoms. Using this approach, covalent bonds are treated as if they were fully ionic, that is, the shared electrons in the covalent bond are transferred to the more electronegative atom, as such the less electronegative atom would completely lose electrons (increase in positive ON). In other words, “the more electronegative atom gets ALL of the shared electrons”. For bonds between identical atoms (atoms of the same element) there are no electron transfers. For example, electronegativity of oxygen is less than that of fluorine (F), therefore in compound OF2, oxygen loses electrons in the O–F bonds to F atoms. As a result, oxygen has positive oxidation number (+2), not the usual -2.