Which one of the following is true for an exothermic reaction $A \rightleftharpoons B$,if $E_f$ and $E_b$ are the activation energies of forward and backward reactions respectively?

Consider a complex reaction taking place in three steps with rate constants $k_1$,$k_2$,and $k_3$ respectively. The overall rate constant $k$ is given by the expression $k = \sqrt{\frac{k_1 k_3}{k_2}}$. If the activation energies of the three steps are $60$,$30$,and $10 \ kJ \ mol^{-1}$ respectively,then the overall energy of activation in $kJ \ mol^{-1}$ is $..........$ $(Nearest \ integer)$

For a reaction at a given temperature $T$,the rate constant $K$ is given by the Arrhenius equation. What is the value of $\lim_{T \to \infty} \log K$?

What is the value of the slope in the graph of $\log_{10} K$ against $\frac{1}{T}$?

The rate of reaction becomes double when the temperature of the reaction is increased from $27\,^{\circ}C$ to $57\,^{\circ}C$. The activation energy of the reaction will be ....... $k\,cal$.

For a reversible reaction $A \rightleftharpoons B$,which one of the following statements is wrong from the given energy profile diagram?