Oxidizer/reducer in decomposition of potassium chlorate
If I remember my introductory lab course correctly, there are two reactions going on, one after the other. First, $\ce{ClO3-}$ disproportions according to the following equation:

$$\ce{4 ClO3- -&gt; Cl- + 3 ClO4-}$$

Which has these two half-equations:

$$\ce{ClO3- + 6 e- + 6 H+ -&gt; Cl- + 3 H2O}\\
3(\ce{ClO3- + H2O -&gt; ClO4- + 2 e- + 2 H+})$$

Where the formally required water is merely spectator.

Only upon further heating does the perchlorate anion decompose according to:

$$\ce{ClO4- -&gt; Cl- + 2 O2}$$

Which again can be broken down to:

$$\ce{ClO4- + 8 e- + 8 H+ -&gt; Cl- + 4 H2O}\\
2(\ce{2H2O -&gt; O2 + 4 e- + 4 H+})$$

And again formally requires water as a spectator.

Note that it is not at all unusual for one entity to be both; the simplest example is probably the disproportion of hydrogen peroxide according to:

$$\ce{2 H2O2 -&gt; O2 + 2 H2O}$$

Where one molecule is reduced by another of the same type.