In a disproportionation reaction, the same element is both oxidized and reduced. To check if an element is undergoing a disproportionation reaction, the oxidation states of the element in the reaction equation can be checked. Keep in mind that:
Oxidation is the increase in oxidation state by the loss of electrons
Reduction is the decrease in oxidation state by the gain of electrons
Thus, the oxidation state of the element in the reactant side is compared with the oxidation states of the element in the product side to see if there is both an increase and decrease in the elements’ oxidation state. Let’s take a look at an example:
Is copper undergoing a disproportionation reaction in the following reaction?
Cu2O(aq) + H2SO4(aq) --> Cu (s) + CuSO4 (aq) + H2O(l)
We can draw a quick table to compare the oxidation states of copper in the reactant and products side.
| ---- | Reactant | Product | -- |
Compound
| Cu2O
| Cu
| CuSO4
|
Oxidation State
| +1
| 0
| +2
|
Here we can see that from Cu2O -> Cu, the oxidation state decreased from +1 -> 0
From Cu2O -> CuSO4, the oxidation state increased from +1 -> +2
Since both oxidation and reduction of copper is occurring, we can say that Cu is undergoing a disproportionation reaction.
