Chemistry

Chemistry

Problem

A reaction has ΔG∘=−40 kJ mol−1\Delta G^\circ = -40\,\text{kJ}\,\text{mol}^{-1} and an uncatalyzed activation energy of 200 kJ mol−1200\,\text{kJ}\,\text{mol}^{-1}. Introducing a suitable catalyst lowers the activation energy to 60 kJ mol−160\,\text{kJ}\,\text{mol}^{-1} but does not change ΔG∘\Delta G^\circ. Which statement most accurately describes the effect of adding the catalyst at 298 K298\,\text{K}? a) The catalyst lowers ΔG∘\Delta G^\circ, thereby increasing the equilibrium constant for the reaction. b) The catalyst makes the reaction thermodynamically favorable when it was not previously. c) The catalyst increases the reaction rate at 298 K298\,\text{K} while leaving the reaction’s thermodynamic favorability unchanged. d) The catalyst has no effect because the reaction was already spontaneous (ΔG∘<0\Delta G^\circ<0).
Intuitive explanation

The key insight is that a catalyst speeds up a reaction by lowering the activation energy barrier, but it doesn't change the thermodynamic favorability (ΔG∘\Delta G^\circ) or the position of equilibrium. Since ΔG∘\Delta G^\circ remains negative, the reaction is still spontaneous and will reach the same equilibrium state—the catalyst just gets there much faster by providing an alternative reaction pathway with lower energy requirements.

Intuitive explanation

The key insight is that a catalyst speeds up a reaction by lowering the activation energy barrier, but it doesn't change the thermodynamic favorability (ΔG∘\Delta G^\circ) or the position of equilibrium. Since ΔG∘\Delta G^\circ remains negative, the reaction is still spontaneous and will reach the same equilibrium state—the catalyst just gets there much faster by providing an alternative reaction pathway with lower energy requirements.

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