$R \rightarrow P$ is a first order reaction. For this reaction a graph of $\ln [R]$ (on $y$-axis) and time (on x -axis) gave a straight line with negative slope. The intercept on $y$-axis is equal to ( $k=$ rate constant)

The half-life of a zero order reaction $A \rightarrow$ products, is 0.5 hour. The initial concentration of $A$ is $4 \mathrm{molL}^{-1}$.

How much time (in hr ) does it take for its concentration to come from $2.0 \mathrm{~mol} \mathrm{~L}^{-1}$ to $1.0 \mathrm{~mol} \mathrm{~L}^{-1}$ ?

For a first order decomposition of a certain reaction, rate constant is given by the equation. $\log k\left(s^{-1}\right)=7.14-\frac{1 \times 10^4 \mathrm{~K}}{T}$. The activation energy of the reaction ( in $\mathrm{kJ} \mathrm{mol}^{-1}$ ) is

$ \left(R=8.3 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}\right) $

Consider a general first order reaction,

$ A(g) \longrightarrow B(g)+C(g) $

If the initial pressure is 200 mm and after 20 minutes it is 250 mm , then the half-life period of the reaction (in minutes) is ( $\log 2=0.30, \log 3=0.48, \log 4=0.60$ )

For the reaction $R \rightarrow P$, half life is independent of initial concentration of the reactant, $R$. Which one of the following graphs is not correct for the reaction?

For the gaseous reaction, $\mathrm{N}_{2} \mathrm{O}_{5} \longrightarrow 2 \mathrm{NO}_{2}+\frac{1}{2} \mathrm{O}_{2}$

the rate can be expressed as

$ \begin{array}{l} -\frac{d\left[\mathrm{~N}_{2} \mathrm{O}_{5}\right]}{d t}=K_{1}\left[\mathrm{~N}_{2} \mathrm{O}_{5}\right] \\\\ +\frac{d\left[\mathrm{NO}_{2}\right]}{d t}=K_{2}\left[\mathrm{~N}_{2} \mathrm{O}_{5}\right] \\\\ +\frac{d\left[\mathrm{O}_{2}\right]}{d t}=K_{3}\left[\mathrm{~N}_{2} \mathrm{O}_{5}\right] \end{array} $

The correct relation between $K_{1}, K_{2}$ and $K_{3}$

A possible mechanism for the gaseous reaction $2 \mathrm{H}_2+2 \mathrm{NO} \longrightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{N}_2$ is

Step 1:2 $\mathrm{NO} \rightleftharpoons \mathrm{N}_2 \mathrm{O}_2$

Step 2 : $\mathrm{N}_2 \mathrm{O}_2+\mathrm{H}_2 \longrightarrow \mathrm{~N}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O}$ (slow)

Step 3: $\mathrm{N}_2 \mathrm{O}+\mathrm{H}_2 \longrightarrow \mathrm{~N}_2+\mathrm{H}_2 \mathrm{O}$

The rate law for this reaction is

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

For a zero order reaction $A \rightarrow$ product, a plot of $[A]$ (on $y$-axis) and time (on $x$-axis) gave a straight line with slope equal to $-3 \times 10^{-3} \mathrm{M} \mathrm{min}^{-1}$ and intercept equal to $2 \times 10^{-2} \mathrm{M}$ (on y -axis). What is the rate constant (in M $\mathrm{min}^{-1}$ ) of this reaction?

The rate of a first order reaction doubles when the temperature changes from 300 K to 310 K . The activation energy of the reaction (in $\mathrm{kJ} \mathrm{mol}^{-1}$ ) is ( $R=8.3 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}, \log 2=0.3$ )

The graph obtained between $\ln k$ ( $k=$ rate constant) on $y$-axis and $1 / T$ on $x$-axis is a straight line. The slope of it is $-4 \times 10^4 \mathrm{~K}$. The activation energy of the reaction (in $\left.\mathrm{kJ} \mathrm{mol}^{-1}\right)$ is $\left(R=831 \mathrm{~J} \mathrm{~K}^{-1} \mathrm{~mol}^{-1}\right)$

For a reaction, the threshold energy is $75 \mathrm{~kJ} /$ mole. If the internal energy of the reactants is 20 $\mathrm{kJ} /$ mole, the activation energy (in $\mathrm{kJ} /$ mole) is

145. Rate constants in the following reaction are Reaction 1 :

$ A \xrightarrow{\text { Catalyst } 1} P_1, k_1=1 \mathrm{~s}^{-1} $

Reaction 2 :

$ A \xrightarrow{\text { Catalyst } 2} P_2, k_2=0.1 \mathrm{~L} \mathrm{~mol}^{-1} \mathrm{~s}^{-1} $

Reaction 3 :

$ A \xrightarrow{\text { Catalyst } 3} P_3, k_3=0.01 \mathrm{~L} \mathrm{~mol}^{-1} \mathrm{~s}^{-1} $

The correct relations between the rate of the reactions at 1 M of $A$ are

Which of the following is a zero order reaction?

The reaction, $2 A \rightarrow 2 B+C$ has a rate constant of $1.2 \times 10^{-2} \mathrm{~s}^{-1}$. Which of the following is correct?

The rate constant of a reaction is increased by 4 times after addition of catalyst to the reaction mixture at the same temperature of $27^{\circ} \mathrm{C}$. The change in the activation energy of this reaction is (Take $\ln (1 / 4)=-1386, R=8.314$ )

Half-life periods for a reaction at initial concentrations of 0.1 M and 0.01 M are 5 and 50 minutes, respectively. The order of reaction is

If the definition of the temperature coefficient of the reaction holds good for a reaction between $27^{\circ} \mathrm{C}$ and $37^{\circ} \mathrm{C}$, the activation energy for the reaction in $\mathrm{kJ} \mathrm{mol}^{-1}$ is

What is the unit for the zero order rate constant?

The specific rate constant of decomposition of a compound is given by $\ln k=5.0-\frac{12000}{T}$. The activation energy of decomposition for this compound at 300 K is

For a zero order reaction, the plot of concentration of reactant vs time is (Hint: Consider the intercept on the concentration axis)