The oxidation of bromide ions is thought to occur by the following mechanism:

$H^{+}+H_2{O}_{2}harr{H}_{2}O+-OH(rapid$ equilibrium$)$

$H_2{O}^{+}-OH+B{r}^{-}ra\frac{r}{r}HOBr+H_{2}O(slow)$

HOBr$+H^{+}+B{r}^{-}B{r}_2+H_{2}O(fast)$

Which of the following statements is correct?

Rate $=k[H_2{O}_2][H^{+}][B{r}^{-}]$

The overall reaction equation is $2H^{+}+H_2{O}_2+2B{r}^{-}B{r}_2+H_{2}O$

$H_{2}O$ is a reaction intermediate.

$H^{+}$is a catalyst.