Question

(i) Define activation energy.
(ii) Explain the effect of catalyst on the rate of chemical reaction.

• A.
• B.
• C.
• D.

Hint:

Think of activation energy as a "hill" that reactants must climb. A catalyst doesn't push the reactants harder; it simply lowers the height of the hill!

Answer 1

The Correct Option is A

Concept: According to the Collision Theory, for a chemical reaction to occur, reactant molecules must collide with a certain minimum energy called Threshold Energy. Activation Energy ($E_a$): The difference between the threshold energy and the average kinetic energy of the reactant molecules. Catalysis: A process where a substance (catalyst) alters the reaction speed without being consumed.
Step 1: Understanding Activation Energy.\ Reactants do not automatically turn into products upon contact. They must overcome an energy barrier. Mathematically: $$ \textActivation Energy (E_a) = \textThreshold Energy - \textAverage energy of reactants $$ Step 2: Effect of Catalyst.\ A catalyst increases the rate of a chemical reaction by participating in the reaction mechanism to provide a new "shortcut." This new pathway has a lower activation energy ($E_a'$) than the uncatalyzed pathway.
Because the energy barrier is lower, a larger fraction of reactant molecules possess enough energy to cross the barrier at a given temperature, thereby increasing the reaction rate. It is important to note that a catalyst does \textitnot change the enthalpy ($\Delta H$) or the equilibrium constant of the reaction.