For a hypothetical reaction $A \to B$,the activation energies for forward and backward reactions are $19 \ kJ/mol$ and $9 \ kJ/mol$ respectively. The heat of reaction is $.... \ kJ$.

The rate constant of a first order reaction was doubled when the temperature was increased from $300 \ K$ to $310 \ K$. What is its approximate activation energy (in $kJ \cdot mol^{-1}$)? ($R = 8.3 \ J \cdot mol^{-1} \cdot K^{-1}$; $\log 2 = 0.3$)

Define the following terms:
$1.$ Effective collision
$2.$ Probability or steric factor $(P)$

For the following two reactions,which statement is true?

The Arrhenius equation is represented as $k = A e^{-E_a/RT}$. The activation energy $E_a$ of the reaction can be calculated by plotting:

Which is a wrong statement?