Question

In a reaction carried out at 400 K, 0.01% of the total collisions are effective. What is the energy of activation of the reaction?

• (A) 13.3 kJ/mol
• (B) 23.5 kJ/mol
• (C) 3.2 kJ/mol
• (D) 30.6 kJ/mol

Answer 1

The Correct Option is D

Explanation:
In the Arrhenius equation,$$k=A{e}^{\frac{-k_n}{nT}}$$'A' is the pre-exponential factor which represents the number of collisions that are favourably oriented.$e^{\frac{-E_a}{RT}}$ represents the fraction of molecules that possess energy equal to or greater than the threshold energy.$${\mathrm{e}}^{-\frac{E_2}{RT}}=\frac{0.01}{100}\Rightarrow -\frac{E_3}{RT}=2.303\log {10}^{-4}\Rightarrow E_3=2.303\times 4\times 8.314\times 400=30.6\mathrm{kJ}$$$/\mathrm{mol}$