Statement I: A catalyst increases the rate of a reaction.
Statement II: A catalyst provides an alternate pathway (reaction mechanism) for the reaction, which has a lower activation energy.
Answer & explanation
Correct answer: option 1
The correct answer is Option (1) → Both Statement I and Statement II are correct and Statement II is the correct explanation of Statement I.
The activation energy (Ea) is the energy barrier the reactants must cross to form products. The smaller the barrier, the larger the fraction of molecules that can cross it.
Statement I: A catalyst increases the rate of a reaction. — True. A catalyst increases the rate of a reaction without itself undergoing any permanent chemical change.
Statement II: A catalyst provides an alternate pathway (reaction mechanism) for the reaction, which has a lower activation energy. — True. The catalyst opens a new path with a lower potential energy barrier between the same reactants and products.
Why Statement II explains Statement I?
In the Arrhenius equation, k = Ae−Ea/RT, the factor e−Ea/RT is the fraction of molecules with energy equal to or greater than Ea.
A lower Ea makes this fraction larger, so k rises and the reaction goes faster. The rate increases because the catalyst lowers the activation energy.
NCERT (2026-27 edition):
3.4.1 Effect of Catalyst – “A catalyst is a substance which increases the rate of a reaction without itself undergoing any permanent chemical change. … It is believed that the catalyst provides an alternate pathway or reaction mechanism by reducing the activation energy between reactants and products and hence lowering the potential energy barrier as shown in Fig. 3.11. It is clear from Arrhenius equation (3.18) that lower the value of activation energy faster will be the rate of a reaction.”
\( \displaystyle k=A\,e^{-E_a/RT} \qquad (3.18) \)
Fig. 3.11: Effect of catalyst on activation energy