For a reaction, the graph of $\ln k$ (on $y$-axis) and $\frac{1}{T}$ (on $x$-axis) is a straight line with a slope $-2 \times 10^4 \mathrm{~K}$. The activation energy of the reaction (in $\mathrm{kJ} \mathrm{mol}^{-1}$ ) is ( $R=8.3 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}$ )

The following graph is obtained for a first order reaction $(A \rightarrow P)$. The activation energy ( $E_a$ in $\left.\mathrm{kJ} \mathrm{mol}^{-1}\right)$ and heat of reaction $\left(|\Delta H|\right.$ in $\left.\mathrm{kJ} \mathrm{mol}^{-1}\right)$ for this reaction are respectively

$\left(x=\right.$ reaction coordinate; $y=E$ in $\left.\mathrm{kJ} \mathrm{mol}^{-1}\right)$

For a first order reaction, the ratio between the time taken to complete $\frac{3}{4}$ th of the reaction and time taken to complete half of the reaction is

The following equation is obtained for a first order reaction at 300 K

$ \log _{10} \frac{k}{A}=0.00174 $

What is the activation energy (in $\mathrm{J} \mathrm{mol}^{-1}$ ) of the reaction?

$ \left(R=8314 \mathrm{~J} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}\right) $

In a first order reaction, the concentration of the reactant is reduced to $1 / 8$ of the initial concentration in 75 minutes. The $t_{1 / 2}$ of the reaction (in minutes) is $(\log 2=0.30, \log 3=0.47, \log 4=0.60)$

At $T(\mathrm{~K})$ the following equation is obtained for a first order reaction $\log \frac{k}{A}=-\frac{x}{T}$. The activation energy for this reaction is equal to ( $R=$ gas constant)

Consider the reaction given below

$ A+2 B \longrightarrow 3 C+2 D $

If rate of disappearance of $B$ is $x \times 10^{-2} \mathrm{~mol} \mathrm{~L}^{-1} \mathrm{~s}^{-1}$, the ratio of rate of reaction and rate of appearance of $C$ is

Activation energy for the hydrolysis of sucrose by acid is $X \mathrm{~kJ} \mathrm{~mol}^{-1}$ whereas activation energy for the hydrolysis of sucrose by sucrase is $Y \mathrm{~kJ} \mathrm{~mol}^{-1} . X$ and $Y$ respectively are

$A \rightarrow$ products, is a first order reaction. The following data is obtained for this reaction at $T(\mathrm{~K})$. The value of $x: y$ is

$ \begin{array}{cc} \hline \text { Rate }\left(\mathrm{molL}^{-1} \mathrm{~min}^{-1}\right) & {[A]} \\ \hline 0.2 & 0.02 \mathrm{M} \\ \hline 0.4 & x \mathrm{M} \\ \hline 1.0 & y \mathrm{M} \\ \hline \end{array} $

Isomerisation of gaseous cyclobutene to butadiene is first order reaction. At $T(\mathrm{~K})$. The rate constant of reaction is $33 \times 10^{-4} \mathrm{~s}^{-1}$. What is the time required (in min ) to complete $90 \%$ of this reaction at the temperature? $(\log 2=03)$

For a first order reaction the concentration of reactant was reduced from $0.03 \mathrm{molL}^{-1}$ to $0.02 \mathrm{molL}^{-1}$ in 25 min . What is its rate (in $\mathrm{molL}^{-1} \mathrm{~s}^{-1}$ )?

The rate constant of a reaction at 500 K and 700 K are $0.02 \mathrm{~s}^{-1}$ and $0.2 \mathrm{~s}^{-1}$ respectively. The activation energy of the reaction (in $\left.\mathrm{kJ} \mathrm{mol}^{-1}\right)$ is $\left(R=8.3 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}\right)$

The time required for completion of $93.75 \%$ of a first order reaction is $x$ minutes. The half-life of it (in minutes) is

. The rate constant for a zero order reaction $A \longrightarrow$ products is $0.0030 \mathrm{~mol} \mathrm{~L}^{-1} \mathrm{~S}^{-1}$. How long it will take for the initial concentration of $A$ to fall from 0.10 M to 0.075 M ?

For a $A+B \rightarrow$ products, the rate of the reaction is given by rate $=k[A][B]^2$. The units of rate constant $(k)$ will be

For an elementary reaction, $X(g) \longrightarrow Y(g)+Z(g)$, the $t_{1 / 2}$ is $10 \mathrm{~min}$. In what period of time would the concentration of $X$ be reduced to $10 \%$ of its original concentration?

Which statement among the following is incorrect?

For zero order reaction, a plot of $t_{1 / 2}$ versus $[A]_0$ will be

If the rate constant for a first order reaction is $2.303 \times 10^{-3} \mathrm{~s}^{-1}$. Find the time required to reduce $4 \mathrm{~g}$ of the reactant to $0.2 \mathrm{~g}$.