# Key Ideas

• Oxidants: substances that cause oxidation and are reduced themselves
• Reductants: substances that cause reduction and are oxidised themselves
• Reduction: the gain of electrons
• Oxidation: the loss of electrons

# Oxidation Numbers

• The oxidation number of a free element is $$0$$
• The oxidation number of a simple ion is equal to the charge on the ion
• In compounds, some elements have oxidation numbers that are regarded as fixed, except in a few exceptions
• Main group metals have an oxidation number equal to the charge on their ions
• Hydrogen has an oxidation number of $$+1$$ when it forms compounds with non-metals
• In metal hydrides, the oxidation number of hydrogen is $$-1$$ e.g. $$NaH$$, $$CaH_2$$
• Oxygen usually has an oxidation number of $$-2$$
• In compounds with fluorine, oxygen will have a positive oxidation number
• In peroxides, oxygen has an oxidation number of $$-1$$ e.g. $$H_2O_2$$, $$BaO_2$$
• The sum of the oxidation numbers in a neutral compound is $$0$$
• The sum of the oxidation numbers in a polyatomic ion is equal to the charge on the ion
• The most electronegative element is assigned the negative oxidation number

# Writing Redox Equations

 1 Balance key elements in the half equations 2 Balance oxygen atoms by adding water molecules 3 Balance the hydrogen atoms by adding $$H^+$$ ions 4 Balance the overall charge on each side with electrons 5 Add states to complete the half equation 6 Make the number of electrons in the two half equations equal, then add them together