1. Identify the reactants and products in the reaction.
2. Determine the charges of the individual atoms in the reactants and products.
- Neutral atoms have an oxidation number of 0.
- Monoatomic ions have an oxidation number equal to their charge.
- For polyatomic ions, the sum of the oxidation numbers of the individual atoms must equal the overall charge of the ion.
3. Use the law of conservation of mass to balance the equation.
- The total number of atoms of each element must be the same on both sides of the equation.
4. Assign oxidation numbers to the elements in the reaction.
- Start by assigning oxidation numbers to the elements that have a known oxidation number, such as hydrogen (+1), oxygen (-2), and alkali metals (+1).
- Use the charges of the ions to help you assign oxidation numbers to the other elements.
- Check that the sum of the oxidation numbers of the atoms in each compound is equal to the overall charge of the compound.
5. Calculate the change in oxidation number for each element.
- The change in oxidation number is the difference between the oxidation number of the element in the product and the oxidation number of the element in the reactant.
- A positive change in oxidation number indicates that the element has been oxidized, while a negative change in oxidation number indicates that the element has been reduced.
Here is an example of how to find the oxidation numbers in a reaction:
- Reactants: Fe(s) + 2HCl(aq) --> FeCl2(aq) + H2(g)
- Products: Fe(II) has an oxidation number of +2 in FeCl2.
- Cl(-1) has an oxidation number of -1 in HCl and FeCl2.
- H(+1) has an oxidation number of +1 in HCl and H2.
- Since the reaction is balanced, the total number of atoms of each element is the same on both sides of the equation.
The change in oxidation number for each element is:
- Fe: 0 → +2 (2e- lost)
- H: +1 → 0 (2e- gained)
- Cl: 0 → -1 (1e- gained)
- Therefore, Fe is oxidized, H is reduced, and Cl is reduced.
