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 $kJ \ mol^{-1}$) is
$R=8.3 \ J \ K^{-1} \ mol^{-1}, \log 2=0.3$ (in $.33$)

For a reaction,given below is the graph of $\ln k$ vs $\frac{1}{T}$. The activation energy for the reaction is equal to $...... \ cal \ mol^{-1}$. (Nearest integer). (Given : $R = 2 \ cal \ K^{-1} \ mol^{-1}$)

$A$ graph plotted between $\log \,K$ vs $\frac{1}{T}$ for calculating activation energy is shown by:

The relation between the rate constant and temperature according to the Arrhenius equation is:

Given below are two statements: one is labelled as Assertion $A$ and the other is labelled as Reason $R$.
Assertion $A$: $A$ reaction can have zero activation energy.
Reason $R$: The minimum extra amount of energy absorbed by reactant molecules so that their energy becomes equal to threshold value,is called activation energy.
In the light of the above statements,choose the correct answer from the options given below:

Which statement is incorrect for collision theory?