AS Chemistry: Reaction Kinetics — Revision Notes

Condensed for the final weeks. For the full explanation, use the Reaction Kinetics study guide.

Collision theory

A reaction occurs only when particles collide with:

1. Energy greater than or equal to the activation energy, Eₐ, and
2. The correct orientation.

Most collisions are unsuccessful on both counts.

Worked example — rate from a graph. A student measures gas volume every 10 s. Between 20 s and 30 s, volume rises from 24 cm³ to 39 cm³.

rate = change in volume / change in time = (39-24)/(30-20) = 15/10 = 1.5 cm^3 s^-1

Reading a rate directly from experimental data (a gradient over an interval, or a tangent to a graph) requires no rate equation.

The four factors and their correct explanations

Factor	Rate	Why — be precise
Temperature ↑	Faster	Particles move faster → more frequent collisions and, far more importantly, a greater proportion exceed Eₐ
Concentration ↑	Faster	More particles per unit volume → more frequent collisions → more effective collisions per second, since the proportion effective is unchanged by concentration
Pressure ↑ (gases)	Faster	Same as concentration — particles closer together
Surface area ↑	Faster	More particles exposed → more frequent collisions
Catalyst	Faster	Alternative pathway of lower activation energy

Only temperature increases the energy of collisions. The others change frequency alone — and the temperature effect is dominated by the proportion exceeding Eₐ, not by the modest increase in collision frequency.

The Boltzmann distribution

Key features to draw and label:

• Starts at the origin — no particle has zero energy.
• Rises to a peak (the most probable energy), then falls asymptotically, never touching the x-axis — there is no maximum energy.
• Area under the curve = total number of particles, so it must stay constant when temperature changes.
• Eₐ marked on the x-axis; the area to the right of Eₐ is the number of particles able to react.

At higher temperature: the peak moves right and becomes lower and broader; the area to the right of Eₐ increases significantly.

With a catalyst: the curve is unchanged, but Eₐ moves left, so a larger area lies beyond it.

That distinction — temperature moves the curve, a catalyst moves the line — is one examiners often check carefully.

Catalysts

• Increase rate by providing an alternative route of lower Eₐ.
• Are not consumed — chemically unchanged at the end.
• Do not change ΔH, yield or the position of equilibrium.
• Homogeneous — same phase as reactants, e.g. aqueous iodide ions catalysing the decomposition of aqueous hydrogen peroxide. Heterogeneous — different phase, working by adsorption onto the surface, reaction, then desorption, e.g. the solid iron catalyst in the gas-phase Haber process.

A reaction-pathway diagram compares catalysed and uncatalysed routes: both start and end at the same enthalpy levels (ΔH is unchanged), but the catalysed pathway’s peak sits lower, showing the reduced activation energy directly.

Exam traps

• Saying “more collisions” rather than “more frequent collisions”.
• Attributing the temperature effect mainly to collision frequency — it is the proportion exceeding Eₐ.
• Drawing a Boltzmann curve that touches the x-axis or starts above the origin.
• Changing the area under the curve when redrawing at a higher temperature.
• Saying a catalyst “lowers the energy of the reactants” — it lowers the activation energy by an alternative pathway.
• Claiming a catalyst increases yield.
• Confusing homogeneous/heterogeneous with “pure vs mixed” — the classification is only about whether the catalyst shares the same phase as the reactants.
• Drawing a reaction-pathway diagram where the catalysed and uncatalysed routes start or end at different enthalpy levels.
• Forgetting concrete examples of homogeneous (I⁻(aq) with H₂O₂) and heterogeneous (iron in the Haber process) catalysis.

Self-test

1. State the two conditions for a successful collision.
2. Why does a small temperature rise produce a large rate increase?
3. Sketch (describe) how the Boltzmann distribution changes at higher temperature.
4. How does a catalyst appear on a Boltzmann diagram?
5. Distinguish homogeneous from heterogeneous catalysis.
6. Gas volume rises from 24 cm³ to 39 cm³ between 20 s and 30 s. Find the average rate over this interval.
7. On a reaction-pathway diagram, what stays the same and what changes between the catalysed and uncatalysed routes?
8. Explain, as a full reasoning chain, why increasing concentration speeds up a reaction.

Answers: 1. Energy at least equal to the activation energy, and the correct orientation. 2. It substantially increases the proportion of particles with energy exceeding Eₐ — the area beyond Eₐ grows sharply — which matters far more than the modest rise in collision frequency. 3. The peak moves right and becomes lower and broader; the total area stays the same; the area beyond Eₐ increases. 4. The curve is unchanged; the Eₐ line moves to the left, so a greater area lies beyond it. 5. Homogeneous catalysts are in the same phase as the reactants; heterogeneous catalysts are in a different phase and act by adsorption onto their surface. 6. rate = (39−24) ÷ (30−20) = 15 ÷ 10 = 1.5 cm³ s⁻¹. 7. The start and end enthalpy levels stay the same (ΔH unchanged); the peak height (activation energy) is lower for the catalysed route. 8. More particles per unit volume → more frequent collisions → more effective collisions per second, since the proportion of collisions that are effective is unchanged by concentration → faster rate.