Chemistry · Ch 7 — Redox Reactions
Balancing of Redox Reactions
Balancing of Redox Reactions
Chemical equations for redox reactions cannot be balanced by simple inspection, because
both mass and charge must be conserved at the same time. Two systematic methods exist.
The first is based on the change in the oxidation numbers of the reducing agent and the
oxidising agent; the second splits the reaction into two half reactions — one for
oxidation, one for reduction — balances each separately, and then adds them. Both methods
are in use, and the choice between them rests with the person using them.
(a) Oxidation Number Method
The fundamental principle: total increase in oxidation number = total decrease in oxidation number. The electrons lost by the reducing agent must be exactly the
electrons gained by the oxidising agent.
In outline, the method runs through five steps:
- Write the correct formula for each reactant and product.
- Identify the atoms whose oxidation number changes in the reaction, by assigning oxidation numbers to all elements on both sides.
- Equalise the increase and the decrease in oxidation number by multiplying the species involved with suitable numbers. (If at this point you find two substances being reduced and nothing oxidised — or vice versa — something is wrong: either a formula is incorrect or the oxidation numbers have not been assigned properly.)
- Balance the ionic charges. If the reaction happens in water, add H⁺ ions (for an acidic solution) or OH⁻ ions (for a basic solution) to the appropriate side so that the total ionic charges of reactants and products are equal.
- Balance hydrogen by adding H₂O molecules to the reactants or products, then check the oxygen count. If the oxygen atoms also balance, the equation represents the fully balanced redox reaction.
Problems 7.8 and 7.9 below walk this method through two complete examples — dichromate
oxidising sulphite in acidic solution, and permanganate oxidising bromide in basic
solution.
(b) Half Reaction Method
In this method, the two half equations are balanced separately and then added together so
that the electrons cancel, giving the balanced overall equation.
General procedure for the half reaction method:
- Produce the unbalanced (skeletal) equation in ionic form.
- Separate it into an oxidation half reaction and a reduction half reaction.
- Balance the atoms other than O and H in each half reaction.
- Balance O by adding H₂O, and H by adding H⁺ (for acidic medium).
- Balance the charge by adding electrons.
- Multiply the half reactions by suitable numbers so the electrons cancel, and add.
- Verify that atoms and charges balance on both sides.
Basic medium: first balance the atoms as in acidic medium; then, for each H⁺ ion,
add an equal number of OH⁻ ions to both sides, and combine H⁺ + OH⁻ on the same side
into H₂O.
The chapter develops the method on the oxidation of Fe²⁺ ions by dichromate ions in
acidic medium, where Cr₂O₇²⁻ is reduced to Cr³⁺:
Step 1 — the unbalanced equation in ionic form:
Step 2 — separate into half reactions:
Step 3 — balance atoms other than O and H. The oxidation half is already balanced in
Fe. In the reduction half, multiply Cr³⁺ by 2 to balance chromium:
Step 4 — for a reaction in acidic medium, add H₂O to balance O and H⁺ to balance H:
Step 5 — add electrons to balance the charges. The oxidation half needs one electron
on the right; the reduction half has net +12 on the left against +6 on the right, so six
electrons go on the left:
To equalise the electrons, multiply the oxidation half reaction by 6: …