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Question
Justify that the following reaction is redox reaction; identify the species oxidized/reduced, which acts as an oxidant and which acts as a reductant.
\[\ce{I2_{(aq)} + 2S2O^{2-}_{3(aq)}->S4O^{2-}_{6(aq)} + 2I^-_{ (aq)}}\]
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Solution
\[\ce{I2_{(aq)} + 2S2O^{2-}_{3(aq)}->S4O^{2-}_{6(aq)} + 2I^-_{ (aq)}}\]
- Write oxidation number of all the atoms of reactants and products.
- Identify the species that undergoes a change in oxidation number.
- The oxidation number of S increases from +2 to +2.5 and that of I decreases from 0 to –1. Because oxidation number of one species increases and that of the other decreases, the reaction is a redox reaction.
- The oxidation number of S increases by loss of electrons and therefore, S is a reducing agent and itself is oxidized. On the other hand, the oxidation number of I decreases by a gain of electrons, and therefore, I is an oxidizing agent and itself is reduced.
Result:
- The given reaction is a redox reaction.
- Oxidant/oxidizing agent (Reduced species): I2
- Reductant/reducing agent (Oxidized species): \[\ce{S2O^2-_3}\]
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