Electrolysis of Anhydrous Sodium Hypochlorite
Obviously, sodium does not have a choice other than getting reduced. Oxidising it to sodium(II) is not an easy feat considering that the electron would have to be extracted from a core orbital. That means, the hypochlorite must be oxidised.

Usually, an oxidation of hypochlorite will yield chlorate, $\ce{ClO3-}$. If we attempt to form a redox half-reaction out of that, we can only use $\ce{ClO-}$ as a charge balancing agent:

$$\begin{align}\ce{ClO- \phantom{\ce{+ 4 ClO-}} &amp;-&gt; ClO3- + 4e-}\tag{Ox1}\\
\ce{ClO- + 4 ClO- &amp;-&gt; ClO3- + 4 e-}\tag{Ox2}\end{align}$$

If we do this, we have four extraneous chlorine atoms and two extraneous oxygen atoms. Thus, it makes most sense to do the mass balance by adding $\ce{Cl2O}$:

$$\ce{ClO- + 4 ClO- -&gt; ClO3- + 4 e- + 2 Cl2O}\tag{Ox3}$$

However, it will be difficult confirming this. Sodium hypochlorite is not stable, especially not at elevated temperatures, especially not in pure form, especially not anhydrous sodium hypochlorite (sensing a pattern?). You would need a certain amount of recklessness to attempt this and sufficient luck to be able to report the result. Which is possibly why I cannot find any scientific publication on the topic; so take the equation with two grains (or more) of salt.