In all examples of disproportionation I have seen, there is one particular atom whose oxidation state is both increasing and decreasing to give rise to different products

For example, on heating phosphorous acid

$$\ce{4 H3PO3 -> 3 H3PO4 + PH3}$$

the oxidation state of phosphorous changes from +3 to +5 and -3.

However, in the case of heating $\ce{KClO3}$

$$\ce{2 KClO3 -> 2 KCl + 3 O2}$$

the oxidation state of oxygen changes from -2 to 0, while that of chlorine changes from +5 to -1. So, will such a reaction be called a disproportionation since its the same molecule giving rise to two products, or will it be simply a redox reaction?


Thermal decomposition of potassium chlorate is not disproportionation, just a redox reaction. Disproportionation refers to the same element acting both as oxidizing agent and a reducing agent, resulting in compounds which contain the same element in different oxidation states.

On the other hand, if you consider preliminary decomposition stage involving perchlorate formation, then it is a disproportionation reaction:

$$\ce{4 K\overset{+5}{Cl}O3 ->[\pu{400 °C}] 3 K\overset{+7}{Cl}O4 + K\overset{-1}{Cl}}$$

Also, synthesis of $\ce{KClO3}$ is a disproportionation reaction:

$$\ce{6 KOH + 3 \overset{0}{Cl}_2 -> 3 H2O + 5 K\overset{-1}{Cl} + K\overset{+5}{Cl}O3}$$

| improve this answer | |

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.