An elementary reaction between $A$ and $B$ is a second order reaction. Which of the following rate equations must be correct?

Why can’t molecularity of any reaction be equal to zero ?

For the reaction : $2A + B \to  A_2B$ ; the rate $= K[A]\, [B]^2$ with $K = 2.0\times10^{-6}\, lit^2\, mol^{-2}\, sec^{-1}$. Initial concentration of $A$ and $B$ are $0.2\, mol/lit$ and $0.4\, mol/lit$ respectively. Calculate the rate of reaction after $[A]$ is reduced to $0.12\, mol/litre$.

${A_2} + {B_2} \to 2AB;R.O.R = k{[{A_2}]^a}{[{B_2}]^b}$

Order of reaction with respect to $A_2$ and $B_2$ are respectively 

For a reaction $A \to$ Products, a plot of $log\,t_{1/2}$ versus $log\,a_0$ is shown in the figure. If the initial concentration of $A$ is represented by $a_0,$ the order of the reaction is

Reaction : $2Br^{-} + H_2O_2 + 2H^{+} \to  Br_2 + 2H_2O$

take place in two steps :

$(a)$ $Br^{-} + H^{+} + H_2O_2 \xrightarrow{{slow}} HOBr + H_2O$

$(b)$ $HOBr + Br^{-} + H^{+} \xrightarrow{{fast}} H_2O + Br_2$

The order of the reaction is