# Redox Reactions¶

Oxidation Numbers Rules for calculating oxidation numbers (in this order)

1. Free or uncombined elements have an oxidation number of 0 $$\hskip{1cm}\ce{Na, Be, K, Pb, H2, O2, P4}$$
2. In monatomic ions, the oxidation number = the ion’s charge

$$\hskip{1cm}\ce{Li+} = +1 \hskip{1cm} \ce{Fe^{3+}} = +3 \hskip{1cm} \ce{O^{2−}} = −2$$

1. The oxidation number of oxygen is usually $$-2$$ except in $$\ce{H2O2}$$ and $$\ce{O2^{2−}}$$ where it′ s $$-1$$
2. The oxidation number of hydrogen is $$+1$$ except when it’s bonded to metals in binary compounds - hydrides. In these cases, its oxidation number is $$-1$$ $$\ce{LiH, CaH2, LiAlH4 = −1}$$
3. Group I metals are $$+1$$, group II metals are $$+2$$ and fluorine is always $$-1$$
4. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on that molecule or ion.

# Using Oxidation Numbers¶

If an oxidation number increases over a reaction, it has undergone oxidation If an oxidation number decreases over a reaction, it has undergone reduction

E.g. Identify which element is undergoing reduction and which is undergoing oxidation in the following equations.