Strength of oxoacids of Chlorine
When you are considering redox reactions, you should get familiar with electrochemical series (Ref.1). Oxidizing and reducing characters of certain reagent also depends on the conditions  used as well. For example, according to Ref.1, $\ce{HClO}$ in acidic medium:
$$\ce{HClO + H+ + e- &lt;=&gt;  1/2Cl2 + H2O} ~~~~~~~~ \mathrm{E}^\circ = \pu{+1.594 V} $$
On the other hand, $\ce{ClO-}$ in basic or neutral medium:
$$\ce{ClO- + H2O + 2e- &lt;=&gt;  Cl- + 2OH-} ~~~~~~~~ \mathrm{E}^\circ = \pu{+0.890 V} $$

Consider the following half reaction:
$$\ce{Br2 + 2e- &lt;=&gt;  Br- } ~~~~~~~~ \mathrm{E}^\circ = \pu{+1.098 V} $$

Now, you can say that $\ce{HClO}$ in acidic medium spontaneously oxidizes $\ce{Br-}$ to $\ce{Br2}$, but $\ce{ClO-}$ in basic or neutral medium cannot spontaneously oxidize $\ce{Br-}$ to $\ce{Br2}$.

In similar way, you may also say that $\ce{HClO}$ in acidic medium is a better oxidizing agent than that of $\ce{ClO4-}$ in acidic medium, because: 
$$\ce{ClO4- + 8H+ + 7e- &lt;=&gt;  1/2Cl2 + 4H2O} ~~~~~~~~ \mathrm{E}^\circ = \pu{+1.392 V} $$



Reference:


Steven G. Bratsch, “Standard Electrode Potentials and Temperature Coefficients in Water at 298.15 K,” Journal of Physical and Chemical Reference Data 1989, 18(1), 1–21 (https://doi.org/10.1063/1.555839).