# Oxidation Numbers

Oxidation numbers are used by chemists to track how many electrons an atom has.

They don’t always correspond to the actual numbers of electrons an atom has.

## Rules

1. Atoms in their elemental state have an oxidation number of `0`
2. Atoms in monatomic (consisting of one atom) ions have an oxidation number equal to their charge
3. In compounds, Fluorine is assigned a `-1` oxidation number;
• Oxygen is usually assigned a `-2 `oxidation number (except in peroxide, where it is `-1`)
• Hydrogen is usually assigned a `+1` oxidation number, except when existing as a hydride ion H
4. In compounds, all other atoms are assigned an oxidation number so that the sum of the oxidation numbers on all atoms in the species equals the charge of the species

Side Note: A peroxide is a compound that has an oxygen-oxygen single bond. They are uncommon due to O-O single bonds being weak.

### Example:

Question: What is the oxidation state of the Hydrogen atom in 1) H2, and 2) H20?

1) 0 – Because it is in its elemental state (rule #1).

2) +1 – The total is 0 yes, but the question asked for the oxidation state of the Hydrogen (+1), as per rules 3, and 4.

## OIL RIG

Oxidation is loss of electrons.

Reduction is gain of electrons.

### Example

2H2 (g) + O2 (g) → 2H20 (l)

Oxidation numbers:

2H2 : `0`
O2 : `0`

2H2 : `+1` and 0: `-2`

The Hydrogen loses electrons, and therefore the oxidation number increases.

Oxygen gained 2 electrons which is evident from the `-2`

Negative oxidation number = gain in electrons.

Chemical reactions that involve electron transfer are called oxidation-reduction, or redox reactions.