Chemicals/Reactions/Oxidation numbers

Oxidation numbers are the numbers of electrons that are released or gained from the outer shell electronic configuration of an atom. They refer to how an atom will bond with another. Due to this, they help in balancing oxidation-reduction reaction equations. There are several rules for how these are assigned.

Rules

1. The oxidation number (o.n.) of an uncombined element is 0.
2. The total o.n. of all of the atoms in a neutral compound is 0.
3. The total o.n. of all of the atoms in a polyatomic ion is equal to its charge.
4. In compounds, group one alkali metals have an o.n. of +1 and group two alkali earth metals have an o.n. of +2.
5. In compounds, fluorine atom has an o.n. of -1, hydrogen atom has an o.n. of +1, and oxygen atom has an o.n. of -2.
6. In two element compounds with metals, group 17 elements have an o.n. of -1, group 16 elements have an o.n. of -2, and group 15 elements have an o.n. of -3.
7. The o.n. of hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals
8. The o.n. of fluorine is -1 in all compounds. The other halogens have an o.n. of -1 in most binary compounds When combined with oxygen, as in oxyanions, however, they have a positive oxidation state.

Practice

What are the oxidation numbers of the following atoms:

1. Oxygen in O₂?

Oxygen is an uncombined element, so by rule #1, the oxidation number is 0.

2. Carbon in CO₂?

Because CO₂ is a neutral compound, the overall oxidation number must be 0. According to rule #5, oxygen has an oxidation number of -2. Therefore, carbon must have an oxidation number of +4.

3. Oxygen in H₂O?

Oxygen in a combined element is -2, by rule #5.

4. Sulfur in PbSO₄?

Oxygen has an oxidation number of -2. Because we know that lead must have an oxidation number of +2, being an ion, the oxidation number of sulfur must then be +6. (-2 * 4 oxygen = -8 + +2 lead = -6 * 1 sulfur)

5. Oxygen in H₂O₂?

In this rare case, where the overall o. n. must still equal 0, hydrogen is still given an oxidation number of +1, forcing oxygen to have a charge of -1, in order for it to balance to 0.

Reference

1. Petrucci, Harwood, and Herring. General Chemistry: Principles and Modern Applications. Eighth edition.