# Does hypochlorous acid oxidize hydrogen peroxide?

The following reaction is from https://ncert.nic.in/textbook/pdf/kech202.pdf, p. 11:

$$\ce{H2O2 + HOCl -> H3O+ + Cl- + O2}$$

I have to find if $$\ce{H2O2}$$ is acting like a reducing agent or an oxidising agent in this reaction.

If we see, the $$\ce{Cl}$$ atom in $$\ce{HOCl}$$ got reduced from $$+1$$ to $$-1$$ oxidation state. That tells me that $$\ce{H2O2}$$ is acting as a reducing agent. But if I see the $$\ce{O}$$ atom, it is oxidised from $$-2$$ to $$0$$, now this tells me that $$\ce{H2O2}$$ is acting as a oxidising agent.

But how is this possible? I know I'm doing something fundamentally wrong.

• Evaluate formal oxidation numbers of chlorine and oxygen, the answer to the question should be clear from the result. Mar 16, 2022 at 11:57

\begin{align} \ce{H2\overset{-1}{O}_2 &<=> \overset{0}{O}_2 + 2 H+ + 2 e-} &\quad E^\circ_1 &= \pu{-0.695 V} \tag{1}\\ \ce{H\overset{+1}{Cl}O + H+ + 2 e- &<=> \overset{-1}{Cl}^- + H2O} &\quad E^\circ_2 &= \pu{+1.482 V} \tag{2}\\ \hline \ce{H2\overset{-1}{O}_2 + H\overset{+1}{Cl}O &-> H3O+ + \overset{-1}{Cl}^- + \overset{0}{O}_2} &\quad E^\circ &= +\pu{0.787 V} \tag{3} \end{align}
Since resulting $$E^\circ = \pu{0.787 V} > 0,$$ free Gibbs energy $$Δ_\mathrm{r}G^\circ = -nFE^\circ < 0,$$ and the redox reaction where hydrogen peroxide is oxidized by chlorate(I) can be considered a thermodynamically favorable process.