IB Chemistry SL — Reactivity 2.2: The Rate of Chemical Change (Collision Theory & Maxwell–Boltzmann)

The rate of reaction is how fast reactants turn into products — the change in concentration per unit time. Collision theory explains what controls it: particles must collide, often enough, hard enough (energy ≥ activation energy) and the right way round.

👆 Read rate off the concentration–time graph · change each factor · reshape the Maxwell–Boltzmann curve

1. Rate of reaction & how to measure it R2.2.1

Rate is the change in concentration of a reactant or product per unit time (mol dm⁻³ s⁻¹). You can follow it by measuring anything that changes: gas volume produced, mass lost, colour (colorimetry), pH, or conductivity. On a concentration–time graph the gradient is the rate — steepest at the start (most reactant, most collisions) and levelling to zero as reactant runs out:

The reaction is fastest at t = 0 (steepest gradient) and stops when the limiting reactant is used up.

2. Collision theory R2.2.2

For particles to react they must collide, and the collision must be successful: the colliding particles need combined kinetic energy of at least the activation energy (Eₐ) — the minimum energy to start breaking bonds — and the correct orientation (geometry). Most collisions are too gentle or wrongly aligned and simply bounce off. Anything that increases the frequency of collisions or the proportion that are successful will increase the rate.

3. Factors that change the rate R2.2.3

Click each factor to see how it works through collision theory:

4. Activation energy & the Maxwell–Boltzmann distribution R2.2.4–5

In any sample, particles have a spread of kinetic energies — the Maxwell–Boltzmann distribution. The curve starts at the origin (no particle has zero energy), rises to a peak (the most probable energy), then tails off (a few very fast particles). The area under the whole curve is the total number of particles; the shaded area beyond Eₐ is the fraction with enough energy to react. Change the conditions:

A catalyst speeds a reaction up by providing an alternative pathway with a lower activation energy, so a larger fraction of particles can react — without being used up itself. Note it does not change the shape of the distribution; it moves the Eₐ line to the left.

Doing Higher Level? HL extends R2.2 with the rate equation (rate = k[A]ᵐ[B]ⁿ), order of reaction and the rate constant, reaction mechanisms and the rate-determining step, and the Arrhenius equation. → Open the HL lesson

Common mistakes examiners see

What makes a collision "successful"?✗ Any collision leads to reaction. ✓ The particles need energy ≥ Eₐ and the correct orientation. Most collisions fail on one or both.

How does temperature speed up a reaction?✗ Only by making particles collide more often. ✓ Mainly because a much larger fraction of particles now exceed Eₐ (the big effect); they also collide slightly more often.

Does a catalyst change the shape of the Maxwell–Boltzmann curve?✗ Yes, it makes the curve taller. ✓ No — the curve is unchanged. The catalyst lowers Eₐ (moves the Eₐ line left), so more particles fall beyond it.

Where does the Maxwell–Boltzmann curve start?✗ At a positive value on the y-axis. ✓ At the origin — no particle has zero kinetic energy — and it never touches the x-axis again (asymptotic tail).

Is a catalyst used up in the reaction?✗ Yes, gradually. ✓ No — it is regenerated, so it is not consumed overall and only a small amount is needed.

Check your understanding

Six questions — instant feedback, nothing saved.

Answered 0 / 6

Practise this topic — free games

▶ Open the Subjects Arcade ▶ All IB Chemistry topics

The Rate of Chemical Change