Consider the following reaction $A + B \rightarrow C$.
The time taken for $A$ to become $1/4$ of its initial concentration is twice the time taken to become $1/2$ of the same. Also,when the change of concentration of $B$ is plotted against time,the resulting graph gives a straight line with a negative slope and a positive intercept on the concentration axis. The overall order of the reaction is . . . . .

$A$ reaction $2NO + 2H_2 \longrightarrow N_2 + 2H_2O$ has the following mechanism:
Step-$I$: $2NO \longrightarrow N_2O_2$
Step-$II$: $N_2O_2 + H_2 \longrightarrow N_2O + H_2O$
Step-$III$: $N_2O + H_2 \longrightarrow N_2 + H_2O$
Which of the following substances is a reaction intermediate?

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)

Give two examples of zero-order reactions and two examples of first-order reactions.

The experimental data for the reaction $2A + B_2 \longrightarrow 2AB$ is given below:

Exp.	$[A] \ (mol \ L^{-1})$	$[B_2] \ (mol \ L^{-1})$	Rate $(mol \ L^{-1} \ S^{-1})$
$1$	$0.50$	$0.50$	$1.6 \times 10^{-4}$
$2$	$0.50$	$1.00$	$3.2 \times 10^{-4}$
$3$	$1.00$	$1.00$	$3.2 \times 10^{-4}$

Determine the rate law for the reaction.

For the reaction $A + B \longrightarrow \text{product}$,the rate law equation is $\text{rate} = k[A]^2[B]$. If the rate of reaction is $0.22 \ mol \ L^{-1} \ s^{-1}$,calculate the rate constant $k$. Given: $[A] = 1 \ mol \ L^{-1}, [B] = 0.25 \ mol \ L^{-1}$.