Why does the presence of a catalyst increase the observed rate of a chemical reaction?

Introduction / Context:
This question targets the core kinetic reason a catalyst accelerates reactions: barrier lowering. It is foundational for understanding reactor performance and catalyst design.

Given Data / Assumptions:

• No external temperature change; only catalyst addition.
• Catalyst is regenerated and not consumed.

Concept / Approach:
Rate constant k depends exponentially on activation energy or free energy. A catalyst provides an alternative pathway with smaller ΔE or ΔG‡, thus increasing k at a given temperature.

Step-by-Step Solution:
1) Arrhenius form: k = A * exp(-Ea/(RT)); transition-state form: k = (kBT/h) * exp(-ΔG‡/(R*T)).2) Lower Ea or ΔG‡ ⇒ larger exponent ⇒ higher k ⇒ faster rate.3) Equilibrium (ΔG°) stays the same; catalyst does not alter final composition, only the speed toward it.

Verification / Alternative check:
Experiments show catalysts change rates without affecting equilibrium constants. This directly supports barrier lowering rather than thermodynamic driving force changes.

Why Other Options Are Wrong:
Option b contradicts observed acceleration; higher barrier would slow rates.
Option c is not the general mechanism; collision geometry matters but barrier lowering is key.
Option d changes equilibrium, which catalysts do not do.
Option e violates energy conservation.

Common Pitfalls:
Confusing kinetic effects (rate) with thermodynamic effects (equilibrium) and assuming catalysts are consumed.

Final Answer:
It lowers the activation barrier.