Procedure for balancing redox equations

Although there is no single ‘correct’ method of balancing a redox reaction, the following systematic procedure (appropriate for reactions taking place in aqueous solution) is recommended.

1. Begin the equation by writing the formula of one of the reactants on the left hand side, and on the right the formula of the product to which it is converted.

2. Balance the equation with respect to the principal atom.

3. Balance the oxygen atoms (if any) by adding the appropriate number of H2O molecules to the oxygen-deficient side of the equation.

4. Balance the hydrogen atoms by adding H+ ions to the appropriate side of the equation.

5. Balance ion charges by adding electrons to the appropriate side of the equation. Call this equation (A).

6. Repeat steps 1-5 for the other reaction, to obtain another equation.

7. Multiply equations (A) and (B) by suitable factors such that the number of electrons on the left of one of the equations is equal to the number on the right of the other.

8. Add the equations together to obtain the required overall equation for the reaction.

This approach is appropriate for reactions in acidic or neutral solutions. For reactions taking place in alkaline solution, it is unrealistic to write an equation involving H+ ions.

To balance an equation for an alkaline solution reaction we proceed by applying the series of steps 1-8, then add one further step:

9. Note the number of H+ ions which appear in your equation, add the same number of OH- ions to each side of the equation, then write one H2O in place of each H+, OH- pair.

You might sometimes have to face problems in which the reactive species are not specified, but rather the names of compounds are given. For example, if one of the compounds is iron(III) chloride you will have to decide if the reactant species is Fe3+ or Cl-. This can be done rigorously using information about redox potentials, but without having recourse to this, some rough working rules may be useful.

1. If the compound is a salt of a Main Group 1 or Main Group 2 metal, the reactant will be the associated anion, because the metal ion cannot change its oxidation state to anything else which would be stable in aqueous solution.

2. If it is a salt of any other non-transition metal the reactant is likely to be the anion, unless it is a reaction in which the metal ion is reduced to the free metal.

3. In a transition metal compound it is likely that the ion which contains the transition metal will be the reactant. In some instances this will be a cation (e.g., Mn2+) and in others an anion (e.g., MnO4-).

4. If there is still doubt, this can often be resolved by considering whether the other reactant is likely to be an oxidizing agent or a reducing agent, e.g., iron (III) chloride will react with a reducing agent:

Fe3+ + e- Æ Fe2+

and with an oxidizing agent:

2Cl- Æ Cl2 + 2e-