Assign oxidation numbers to individual elements in all of the compounds, both on the reactant and product sides of the equation. Oxidation numbers represent only one atom at a time and can be positive, negative or zero. If the compound is neutral, the sum of the oxidation numbers must be zero. An element in its natural form has an oxidation number of zero, such as when carbon is present with its four electrons. If the atom is an ion in the equation, its charge is the oxidation number, such as Al^3- having an oxidation number of negative three (-3). The periodic table indicates the oxidation numbers for other elements: Group 1A metals are +1, Group 2A metals are +2, hydrogen is +1 when bonded to a nonmetal and -1 when bonded to a metal, oxygen is -2 unless with a peroxide and fluorine, which is always -1.
Add the oxidation numbers of the elements within each compound, identifying the overall charge of the compound relative to the elements in it. For example, in the equation that demonstrates the chemical reaction causing iron (symbolized by Fe on the periodic table) to rust, Fe is oxidized -- the iron combines with oxygen, reducing Fe's number of electrons. Fe has an oxidation number of zero on the reactant side, but the number changes to 3+ on the product side in the compound FeO2, iron oxide. The product side also shows water as a byproduct.
Compare similar compounds on both sides of the equation. Sometimes, the elements have switched around and new compounds were created. You are still able to tell if the various elements were oxidized or reduced based on their individual oxidation numbers changing. If the equation is more simple and the compound is only losing electrons, meaning it was oxidized, the change in overall oxidation number from the reactant to the product side of the equation reflects such a change. A compound becoming more positive from the reactant side of the equation to the product side confirms that the compound has been oxidized.