Balancing Equations by the Method of Oxidation-Reduction
Posted by Jim Clark on 22nd May and posted in Tutorial
| Preliminaries: You might wish to review oxidation numbers before continuing | |
| An oxidation-reduction reaction is a reaction in which some atoms oxidation numbers change. | |
| Reaction (1) is an oxidation-reduction reaction because H changes from 0 to +1; O changes from 0 to -2 | 0 0 +1 -2 (1) H2 + O2 ® H2O |
| Reaction (2) is not an oxidation-reduction reaction because the oxidation numbers do not change | +1-2+1 +1-1 +1-1 +1-2 (2) NaOH + HCl ® NaCl + H2O |
| Some definitions | |
| The atom oxidized: the atom whose oxidation number increases | In reaction (1), H |
| The atom reduced: the atom whose oxidation number decreases | In reaction (1), O |
| The half reactions show the gain or loss of electrons by one atom (oxidized or reduced): | |
| The oxidation half reaction: | H0 - 1 e- ® H+1 |
| The reduction half reaction: | O0 + 2 e- ® O-2 |
| The oxidizing agent is the reactant containing the atom reduced | In reaction (1), O2 |
| The reducing agent is the reactant containing the atom oxidized | In reaction (1), H2 |
| Balancing an oxidation-reduction reaction | |
| Find the number of electrons gained and lost by the agents. | Each H2 loses 2 e-, since each H loses 1 e-; Each O2 gains 4 e-, since each O gains 2. |
| Select coefficients in front of them so that the numbers of electrons gained and lost are equal. |
-2e- +4e- |
| Now complete the balanced equation in the usual way, by placing a 2 before H2O | 2 H2 + O2 ® 2 H2O |
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