Catalysis

A catalyst works by offering the reaction a detour. Instead of forcing molecules over the same tall energy barrier, it lets them react through a different mechanism, often by binding reactants temporarily, that has a lower peak energy. Lower barrier means a larger fraction of collisions have enough energy to react, so the rate constant k jumps up, by the Arrhenius relationship k = Ae^−Ea/RT. The catalyst itself is consumed in an early step and regenerated in a later one, so it never appears in the overall balanced equation and never gets used up.

On an energy profile, draw the same starting reactants and same final products at the same heights as the uncatalyzed curve. The catalyzed curve dips through a lower, often multi-humped, hill in between. Because start and end points are unchanged, the overall energy released or absorbed by the reaction (delta H) is identical for both paths — only the hump in the middle shrinks. Catalysts are classed by where they act: homogeneous catalysts dissolve in the same phase as the reactants (an acid catalyst in solution), heterogeneous catalysts work at a solid surface where gases or solutions adsorb onto it (platinum in a catalytic converter), and enzymes are biological catalysts with an active site shaped for one specific substrate.

Because the catalyst speeds up the forward path, it necessarily speeds up the reverse path by the same new mechanism. Both directions reach their faster pace together, so equilibrium is reached sooner but the equilibrium constant and the final product amount stay exactly what they would have been anyway.