Disproportionation of hydrogen peroxide Hydrogen peroxide is decomposed as follows:
$$\ce{2H2O2 -&gt; 2H2O +O2}$$
This is a disproportionation redox reaction of $\ce{H2O2}$ involving the 2 half reactions
$$\ce{H2O2 -&gt; O2 + 2H^+ + 2e-}$$
$$\ce{H2O2 + 2e- + 2H^+ -&gt; 2H2O}$$
But I noticed semantically that it can also be the sum of the two following half reactions:
$$\ce{2H2O2 -&gt; 2O2 + 4H^+ + 4e-}$$
$$\ce{O2 + 4H^+ + 4e- -&gt; 2H2O}$$
Both reactions are feasible according to the following:

$$\begin{array}{ccc}
\text{Oxidised species} &amp; \text{Reduced species} &amp; E^\circ (\mathrm{V}) \\
\hline
\ce{H2O2} &amp; \ce{H2O} &amp; 1.763 \\
\ce{O2} &amp; \ce{H2O} &amp; 1.23 \\
\ce{O2} &amp; \ce{H2O2} &amp; 0.695 \\
\end{array}$$

Can it happen in the second way?
If yes, can you explain the mechanism in simple terms?
In principle you could find a lot of equations, all leading to the same result. But these are only (half-)equations. Most of the times the reaction mechanism is much more complicated than just these reaction equations.

The decomposition of hydrogenperoxide is still an active field of research in various media with various catalysts.

Unfortunately with the massive overflow of the above, I was only able to find a quite old PhD thesis by Trice Walter Haas (Iowa State University, 1960)

Here is a small line-up of reaction equations that will play a role in the decomposition, at least to that part that I can imagine it. And I am just assuming that there are radicals involved, because there usually are.
\begin{align}
\ce{H2O2 &amp;&lt;=&gt; H+ + HO2- }\\
\ce{HO2- + H2O2 &amp;&lt;=&gt; HO2 + HO + OH-}\\
\ce{HO2 + OH &amp;&lt;=&gt; H2O + O2}\\
\end{align}