Reaction Rates & Collision Theory

Molecules only react when they collide with enough energy. Turn up the temperature, crowd them together or add a catalyst, and count the successful collisions.

Collision theory

For A + B → C to happen, an A and a B must collide, and hit hard enough to get over the activation energy E_a. Only the fraction of collisions above that barrier succeeds; for a gas this fraction is about e^−E_a/RT. That is why a modest rise in temperature can speed a reaction up a lot: it adds energy to exactly the fast molecules in the tail of the distribution.

What changes the rate?

Concentration: more molecules, more collisions per second. Temperature: harder collisions, so a bigger share gets over the barrier (and slightly more collisions). Catalyst: a different pathway with a lower E_a, so more collisions succeed. The catalyst is not used up and does not change ΔH. The Arrhenius equation k = A e^−E_a/RT sums all this up.

Temperature 300 K

Activation energy 10 kJ/mol

Molecules of A and of B 40 each

Catalyst off

Collisions A–B0

Successful0

Success rate—

Molecules above Eₐ—

Blue (A) + orange (B) → green (C). A flash means a collision had enough energy. "Molecules above Eₐ" is e^−Eₐ/RT, the shaded tail of the graph. The success rate of collisions comes out a little higher, because fast molecules also collide more often.

A low-barrier model reaction so you can watch it in seconds; real reactions often have E_a of 50–100 kJ/mol. The box is kept at constant temperature.