Explain how a catalyst changes the rate of a reaction, with an energy diagram. State the collision theory and the two conditions for an effective collision.

1. 1 · Define a catalyst. A catalyst speeds up a reaction without permanent chemical change; e.g. MnO₂ in 2KClO₃ → 2KCl + 3O₂.
2. 2 · Explain the mechanism. It forms an intermediate complex, giving a path with lower Ea; the catalyst is regenerated.
3. 3 · Draw the energy profile. Two humps, catalysed one lower; mark Ea and Ea′; ΔH the same for both.
4. 4 · What it does not change. Not ΔG, not the equilibrium constant; it only helps equilibrium come sooner.
5. 5 · State collision theory. Reaction needs collisions; rate = P Z_AB e^(−Ea/RT).
6. 6 · Give the two conditions. Energy ≥ threshold energy and proper orientation (example: OH⁻ attacking CH₃Br from the side opposite to Br).
7. 7 · Link to daily life. Enzymes in saliva and the gut are catalysts that digest food at 37 °C.

Scoring tip: Keep the reactant and product levels the same for both humps; a changed ΔH loses the diagram mark.