If $a$ is the initial concentration and $n$ is the order of the reaction and the half-life period is $T$,then:

In the reaction $X \to Y$,if the concentration of reactant $X$ is increased by $1.5$ times,the rate of reaction increases by $1.837$ times. Determine the order of the reaction with respect to $X$.

The rate law for a reaction between reactants $A$,$B$,and $C$ is $r = K[A][B][C]^2$. If the concentration of $A$ is halved,then the rate of reaction:

For the reaction: $2 A + B \rightarrow A_2 B$,the rate $= k[A][B]^2$ with $k = 2.0 \times 10^{-6} \ mol^{-2} \ L^2 \ s^{-1}$. Calculate the initial rate of the reaction when $[A] = 0.1 \ mol \ L^{-1}$ and $[B] = 0.2 \ mol \ L^{-1}$. Calculate the rate of reaction after $[A]$ is reduced to $0.06 \ mol \ L^{-1}$.

The rate law for the decomposition of hydrogen iodide is $-\frac{d[HI]}{dt}=k[HI]^2$. The units of rate constant $k$ are

For the non-stoichiometric 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 $(B)$	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: