Reaction Kinetics: Collision Theory and Catalysis

This guide covers subtopics 8.1, Rate of reaction, 8.2, Effect of temperature on reaction rates and the concept of activation energy, and 8.3, Homogeneous and heterogeneous catalysts, from Topic 8, Reaction kinetics, of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. All three are AS Level content, covered together here because they build one continuous argument: what makes a collision effective, what activation energy has to do with it, and how a catalyst changes the picture.

Before studying this

At IGCSE or O Level, you met the qualitative factors that speed up a reaction — concentration, temperature, surface area, catalysts — explained loosely in terms of “more collisions.” Rates of Reaction and Reversible Reactions covers that treatment.

AS Level replaces “more collisions” with a precise model: not every collision reacts, only effective ones, and why a collision is effective or not is what this whole topic explains. Formal rate equations, orders of reaction and the rate constant k are A Level (A2) content, not AS — this resource stays entirely at the collision-theory/Boltzmann level the AS syllabus requires, without introducing rate equations.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 8

8.1 Rate of reaction — explaining and using the terms rate of reaction, frequency of collisions, effective and non-effective collisions; explaining qualitatively, in terms of frequency of effective collisions, the effect of concentration and pressure changes on rate; using experimental data to calculate the rate of a reaction.

8.2 Effect of temperature on reaction rates and the concept of activation energy — defining activation energy, Eₐ, as the minimum energy required for a collision to be effective; sketching and using the Boltzmann distribution to explain the significance of activation energy; explaining qualitatively, in terms of the Boltzmann distribution and frequency of effective collisions, the effect of temperature change on rate.

8.3 Homogeneous and heterogeneous catalysts — explaining and using the terms catalyst and catalysis: that a catalyst provides a different mechanism of lower activation energy; explaining this in terms of the Boltzmann distribution; constructing and interpreting a reaction-pathway diagram for a reaction with and without an effective catalyst.

Collisions: effective and non-effective

Particles must collide to react, but not every collision leads to a reaction. A collision is effective only if the particles collide with at least the minimum required energy (activation energy) and the correct orientation. Most collisions are non-effective — either too little energy, wrong orientation, or both.

Effect of concentration and pressure: increasing the concentration of a reactant (or the pressure of a gas reaction) increases the number of particles per unit volume, which increases the frequency of collisions — and since a roughly constant proportion of collisions are effective at a given temperature, more collisions overall means more effective collisions per second, and a faster rate.

Worked example. A student measures the volume of gas produced by a reaction every 10 seconds. Between 20 s and 30 s, the volume rises from 24 cm³ to 39 cm³. Calculate the average rate over this interval.

Rate = change in volume / change in time = (39 − 24) cm³ / (30 − 20) s = 15 / 10 = 1.5 cm³ s⁻¹

This is the kind of calculation outcome 8.1.3 expects — reading a rate directly from experimental data (a gradient over an interval, or from a tangent to a graph), without needing a rate equation.

Activation energy and the Boltzmann distribution

Activation energy, Eₐ, is the minimum energy a collision must have to be effective. The Boltzmann distribution shows how the kinetic energies of particles in a sample are spread out at a given temperature: starting at zero, rising to a peak, then falling away in a long tail towards higher energies. The curve never touches the energy axis — some particles always have very high energy, however few. Importantly:

• The area under the whole curve represents the total number of particles.
• The area under the curve beyond Eₐ represents the number of particles with enough energy to react in an effective collision.

Effect of temperature. Raising the temperature shifts the whole distribution to the right and flattens/broadens it (the peak moves right and lowers, since the same total number of particles now spreads over a wider energy range). The critical point examiners look for: the area beyond Eₐ increases by a much larger proportion than the small shift in the peak would suggest — a modest rise in temperature produces a disproportionately large increase in the fraction of particles with E ≥ Eₐ, which is why rate is so sensitive to temperature. A small increase in collision frequency also occurs (particles move faster on average), but it is the change in the proportion of effective collisions that dominates the rate increase.

Catalysts

A catalyst increases the rate of a reaction, without itself being used up, by providing an alternative reaction pathway (mechanism) with a lower activation energy. On a Boltzmann distribution, a lower Eₐ means a larger area lies beyond it — more particles already have enough energy to react via the catalysed pathway than could react via the uncatalysed one, at the same temperature.

A reaction-pathway diagram compares the two routes: both start and end at the same enthalpy levels (a catalyst does not change ΔH, or the position of equilibrium — only the route taken and how quickly it’s reached), but the catalysed pathway’s peak sits lower than the uncatalysed pathway’s peak, showing the reduced activation energy directly.

Homogeneous catalysis occurs when the catalyst is in the same phase as the reactants — for example, aqueous iodide ions, I⁻(aq), catalysing the decomposition of aqueous hydrogen peroxide, where catalyst and reactant are both in solution. Heterogeneous catalysis occurs when the catalyst is in a different phase from the reactants — the solid iron catalyst in the gas-phase Haber process, from Chemical Equilibria: Kc, Kp and Le Chatelier’s Principle, is a heterogeneous example you’ve already met.

Common mistakes

• Explaining a faster reaction at higher concentration purely by “more collisions,” without connecting it to effective collisions. The reasoning chain is: more particles per volume → more frequent collisions → more effective collisions per second (since the proportion effective is unchanged by concentration) → faster rate.
• Explaining a faster reaction at higher temperature purely by “particles move faster and collide more often.” That’s a real but minor effect — the dominant reason is the disproportionate increase in the fraction of particles exceeding Eₐ, shown by the Boltzmann distribution’s changing shape.
• Drawing a reaction-pathway diagram where the catalysed curve starts or ends at a different energy level. A catalyst changes only the height of the peak (activation energy) — never the enthalpy of reactants or products.
• Confusing homogeneous/heterogeneous with the states-of-matter idea of “pure vs mixed.” The classification is about whether the catalyst shares the same phase as the reactants, nothing else.

Quick revision checklist

• Effective vs non-effective collisions; why not every collision reacts
• Concentration/pressure and rate: more frequent collisions, same proportion effective
• Activation energy, Eₐ: minimum energy for an effective collision
• The Boltzmann distribution: shape, and area beyond Eₐ as “particles that can react”
• Why temperature has such a large effect on rate: the disproportionate change in area beyond Eₐ
• Catalysts: lower Eₐ via an alternative pathway; no change to ΔH or equilibrium position
• Homogeneous (same phase) vs heterogeneous (different phase) catalysis

Related resources

• Rates of Reaction and Reversible Reactions — the IGCSE/O Level qualitative treatment this resource builds on
• Chemical Equilibria: Kc, Kp and Le Chatelier’s Principle — why a catalyst affects rate but never equilibrium position
• Chemical Energetics: Hess’s Law and Enthalpy Cycles — the enthalpy levels a reaction-pathway diagram is built from
• Reaction Kinetics: Rate Equations and Catalysis — the A Level topic that adds rate equations and orders of reaction
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.