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