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

$$\ce{4 H3PO3 -&gt; 3 H3PO4 + PH3}$$

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

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

$$\ce{2 KClO3 -&gt; 2 KCl + 3 O2}$$

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:

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

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

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