Half-life periods for a reaction at initial concentrations of $0.1 \ M$ and $0.01 \ M$ are $5$ and $50$ minutes,respectively. The order of reaction is

For a hypothetical reaction $A + B + C \rightarrow \text{Product}$,the rate is given by $r = -\frac{d[A]}{dt} = K[A]^{1/2}[B]^{1/3}[C]^{1/4}$. The order of the reaction is:

For a reaction $A \rightarrow 2 B + C$ the half-lives are $100 \ s$ and $50 \ s$ when the concentration of reactant $A$ is $0.5 \ mol \ L^{-1}$ and $1.0 \ mol \ L^{-1}$ respectively. The order of the reaction is (Nearest Integer).

For the reaction $2A + B \to C$,the values of initial rate at different reactant concentrations are given in the table below: The rate law for the reaction is

$[A] \ (mol \ L^{-1})$	$[B] \ (mol \ L^{-1})$	Initial Rate $(mol \ L^{-1} \ s^{-1})$
$0.05$	$0.05$	$0.045$
$0.10$	$0.05$	$0.090$
$0.20$	$0.10$	$0.72$

In the reaction $A + B \to \text{Products}$,if $B$ is taken in excess,then it is an example of

The reaction $2N_2O_5 \rightleftharpoons 4NO_2 + O_2$ follows first order kinetics. Hence,the molecularity of the reaction is