# How to Assign Oxidation Numbers The oxidation number of an element or neutral compound is zero. Otherwise, the total charge is ionic charge.

The oxidation number is the positive or negative number of an atom that indicates the electrical charge the atom has if its compound consists of ions. In other words, the oxidation number gives the degree of oxidation (loss of electrons) or reduction (gain of electrons) of the atom in a compound. Because they track the number of electrons lost or gained, oxidation numbers are a sort of shorthand for balancing charge in chemical formulas.

This is a list of rules for assigning oxidation numbers, with examples showing the numbers for elements, compounds, and ions.

### Rules for Assigning Oxidation Numbers

Various texts contain different numbers of rules and may change their order. Here is a list of oxidation number rules:

1. Write the cation first in a chemical formula, followed by the anion. The cation is the more electropositive atom or ion, while the anion is the more electronegative atom or ion. Some atoms may be either the cation or anion, depending on the other elements in the compound. For example, in HCl, the H is H+, but in NaH, the H is H.
2. Write the oxidation number with the sign of the charge followed by its value. For example, write +1 and -3 rather than 1+ and 3-. The latter form is used to indicate oxidation state.
3. The oxidation number of a free element or neutral molecule is 0. For example, the oxidation number of C, Ne, O3, N2, and Cl2 is 0.
4. The sum of all the oxidation numbers of the atoms in a neutral compound is 0. For example, in NaCl, the oxidation number of Na is +1, while the oxidation of Cl is -1. Added together, +1 + (-1) = 0.
5. The oxidation number of a monatomic ion is the charge of the ion. For example, the oxidation number of Na+ is +1, the oxidation number of Cl is -1, and the oxidation number of N3- is -3.
6. The sum of the oxidation numbers of a polyatomic ion is the charge of the ion. For example, the sum of the oxidation numbers for SO42- is -2.
7. The oxidation number of a group 1 (alkali metal) element in a compound is +1.
8. The oxidation number of a group 2 (alkaline earth) element in a compound is +2.
9. The oxidation number of a group 7 (halogen) element in a compound is -1. The exception is when the halogen combines with an element with higher electronegativity (e.g., oxidation number of Cl is +1 in HOCl).
10. The oxidation number of hydrogen in a compound is usually +1. The exception is when hydrogen bonds with metals forming the hydride anion (e.g., LiH, CaH2), giving hydrogen an oxidation number of -1.
11. The oxidation number of oxygen in a compound is usually -2. Exceptions include OF2 and BaO2.

### Examples of Assigning Oxidation Numbers

Example 1: Find the oxidation number of iron in Fe2O3.

The compound has no electrical charge, so the oxidation numbers of iron and oxygen balance each other out. From the rules, you know the oxidation number of oxygen is usually -2. So, find the iron charge that balances the oxygen charge. Remember, the total charge of each atom is its subscript multiplied by its oxidation number.

O is -2

There are 3 O atoms in the compound so the total charge is 3 x -2 = -6

The net charge is zero (neutral), so:

2 Fe + 3(-2) = 0
2 Fe = 6
Fe = 3

Example 2: Find the oxidation number for Cl in NaClO3.

Usually, a halogen like Cl has an oxidation number of -1. But, if you assume Na (an alkali metal) has an oxidation number of +1 and O has an oxidation number of -2, the charges don’t balance out to give a neutral compound. It turns out all of the halogens, except for fluorine, have more than one oxidation number.

Na = +1
O = -2
1 + Cl + 3(-2) = 0
1 + Cl -6 = 0
Cl -5 = 0
Cl = -5

### References

• IUPAC (1997) “Oxidation Number”. Compendium of Chemical Terminology (the “Gold Book”) (2nd ed.). Blackwell Scientific Publications. doi:10.1351/goldbook
• Karen, P.; McArdle, P.; Takats, J. (2016). “Comprehensive definition of oxidation state (IUPAC Recommendations 2016)”. Pure Appl. Chem88 (8): 831–839. doi:10.1515/pac-2015-1204
• Whitten, K. W.; Galley, K. D.; Davis, R. E. (1992). General Chemistry (4th ed.). Saunders.