# Is hydrogen peroxide a better oxidizing agent than chlorine gas? [closed]

If we compare the standard reduction potentials of $$\ce{H_2O_2}$$ with that of $$\ce{Cl_2}$$:

$$\ce{H_2O_2 + 2H^+ + 2e^- <=> 2H_2O : E^o = 1.78V}$$

$$\ce{Cl_2 + 2e^- <=> 2Cl^- : E^o = 1.36V}$$

(Source : Wikipedia)

From the above data, as value of $$\ce{E^o}$$ of $$\ce{H_2O_2}$$ is greater, its oxidizing power is greater than that of $$\ce{Cl_2}$$.

But if $$\ce{H_2O_2}$$ is a better oxidizing agent, why does it act as a reducing agent when reacted with $$\ce{Cl_2}$$ (in the following reaction)?

$$\ce{H_2O_2 + Cl_2 -> 2HCl + O_2}$$

The above reaction is feasible, but my doubt is why does it preferentially occur even though $$\ce{H_2O_2}$$ is the better oxidizing agent.

• @IvanNeretin So then, how can we determine that $\ce{H_2O_2}$ is going to act as a reducing agent here, instead of an oxidizing agent? – Pal Jun 19 at 7:48