For the non - stoichimetre reaction $2A + B \rightarrow C + D,$ the following kinetic data were obtained in three separate experiments, all at $298\, K.$

Initial Concentration $(A)$	Initial Concentration $(A)$	Initial rate of formation of $C$ $(mol\,L^{-1}\,s^{-1})$
$0.1\,M$	$0.1\,M$	$1.2\times 10^{-3}$
$0.1\,M$	$0.2\,M$	$1.2\times 10^{-3}$
$0.2\,M$	$0.1\,M$	$2.4 \times 10^{-3}$

The rate law for the formation of $C$ is :

The rate of reaction $A + 2B \to 3C$ becomes $72\, times$ when concentration of $A$ is tripled and concentration of $B$ is doubled then the order of reaction with respect to $A$ and $B$ respectively is

Diazonium salt decomposes as ${C_6}{H_5}N_2^ + C{l^ - } \to {C_6}{H_5}Cl + {N_2}$ At ${0\,^o}C$, the evolution of ${N_2}$ becomes two times faster when the initial concentration of the salt is doubled. Therefore, it is

The rate constant for the reaction, $2N_2O_5 \to 4NO_2 + O_2$ is $3.0\times 10^{- 4}\,s^{-1}$ . If start made with $1.0\,mol\,L^{-1}$ of $N_2O_5$, calculate the rate of formation of $NO_2$ at the moment of the reaction when concentration of $O_2$ is $0.1\, mol\, L^{-1}$.

The given data are for the reaction :

$2NO(g) + Cl_2(g) \to 2NOCl(g)$ at $298\, K$

$[Cl_2]$                $[NO]$                       Rate $(mol\, L^{-1} \sec^{-1})$

$I$    $0.05\, M$       $0.05\,M$                           $1 \times 10^{-3}$

$II$   $0.15\, M$      $0.05\,M$                           $3 \times 10^{-3}$

$III$  $0.05\, M$      $0.15\,M$                           $9 \times 10^{-3}$

The rate law for the reaction is

Consider the kinetic data given in the following table for the reaction $A + B + C \rightarrow$ Product.

The rate of the reaction for $[ A ]=0.15 mol dm ^{-3},[ B ]=0.25 mol dm ^{-3}$ and $[ C ]=0.15 mol dm ^{-3}$ is found to be $Y \times 10^{-5} mol dm d ^{-3} s ^{-1}$. The value of $Y$ i. . . . . . .