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?

The rate constant of a first order reaction is $3.46 \times 10^{-2} \ s^{-1}$ at $298 \ K$. What is the rate constant of the reaction at $350 \ K$ if its activation energy is $50.1 \ kJ \ mol^{-1}$ (in $s^{-1}$)? $(R = 8.314 \ J \ K^{-1} \ mol^{-1})$
$(\log 2 = 0.3010)$

According to the collision theory,which of the following statements is $NOT$ correct?

Explain: How is the value of activation energy determined based on the Arrhenius equation?

The energy profile diagram for a multi-step reaction,$A$ $\xrightarrow{1} B$ $\xrightarrow{2} C$ $\xrightarrow{3} D,$ is given below. The rate-determining step of the reaction is:

In the Arrhenius equation,the rate of reaction is given by $k = A{e^{ - {E_a}/RT}}$. What does $E_a$ represent?