Consider the following gas-phase reaction.
$2HI_{(g)} \longrightarrow H_{2(g)} + I_{2(g)}$
and the following experimental data obtained at $555 \ K$. What is the order of the reaction with respect to $HI_{(g)}$?

$[HI] \ (M)$	Rate $(M \ s^{-1})$
$0.0500$	$8.80 \times 10^{-10}$
$0.1000$	$3.52 \times 10^{-9}$
$0.1500$	$7.92 \times 10^{-9}$

Consider the following single step reaction in gas phase at constant temperature.
$2 \ A_{(g)} + B_{(g)} \rightarrow C_{(g)}$
The initial rate of the reaction is recorded as $r_1$ when the reaction starts with $1.5 \ atm$ pressure of $A$ and $0.7 \ atm$ pressure of $B$. After some time,the rate $r_2$ is recorded when the pressure of $C$ becomes $0.5 \ atm$. The ratio $r_1 : r_2$ is $\qquad$ $\times 10^{-1}$.
(Nearest integer)

For a reaction,$A + 2B \longrightarrow C$,the rate is given by $+\frac{d[C]}{dt} = k[A][B]$. Hence,the order of the reaction is:

If $a$ is the initial concentration of the reactant,the half-life period of the reaction of $n^{th}$ order is inversely proportional to

The rate law for the reaction between substances $A$ and $B$ is given by $\text{Rate} = k[A]^n[B]^m$. If the concentration of $A$ is doubled and the concentration of $B$ is halved,what is the ratio of the new rate to the initial rate?

For the reaction, $A + B \rightarrow P$, the rate law is $\text{rate} = k[A][B]^2$. The rate of reaction is $0.25 \text{ Ms}^{-1}$ when $[A] = 1 \text{ M}$ and $[B] = 0.2 \text{ M}$ at $25^\circ \text{C}$. Calculate the rate constant $k$ of the reaction at the same temperature.