Will heating of KClO3 be considered disproportionation?

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](https://goldbook.iupac.org/html/D/D01799.html) 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:


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


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