OxfordAQA A-Level Chemistry: Kinetics (9620) – Revision Notes

For full explanations and worked examples, read the kinetics study guide first. These notes condense section 3.1.6 Kinetics of the OxfordAQA International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards. They cover 3.1.6.1 to 3.1.6.5. All of it is International AS content, so none of it is International A-level only. The specification assesses it in Unit 2 (Organic 1 and Physical 1), and the International A2 Unit 5 covers all content.

Test yourself with the kinetics practice questions. Course hub: OxfordAQA A-Level Chemistry. Tick off each point on the printable checklist, and use the free 10-minute diagnostics to find weak topics. Rate equations and Required practical 8 belong to 3.1.11 (International A-level only), not to this section, and 3.1.6 lists no required practical of its own.

Definitions to learn word for word

Term	What to write
Activation energy, Ea (3.1.6.1)	The minimum energy that particles must have when they collide for a reaction to occur
Rate of reaction (3.1.6.3)	The change in concentration (or amount) of a reactant or product per unit time
Catalyst (3.1.6.5)	A substance that increases the rate of a chemical reaction without being changed in chemical composition or amount
How a catalyst works (3.1.6.5)	It provides an alternative reaction route of lower activation energy
Successful collision	A collision between particles with energy at or above Ea, which leads to a reaction

Rate: meaning and units

What you measure	Mean rate	Unit
Concentration of a reactant or product	change in concentration ÷ time	mol dm⁻³ s⁻¹
Volume of gas formed	change in volume ÷ time	cm³ s⁻¹
Amount of a substance	change in amount ÷ time	mol s⁻¹
Mass lost as a gas escapes	change in mass ÷ time	g s⁻¹

• Rate is quoted as a positive number whether a reactant falls or a product rises.
• If the equation ratio is 1 : 1, a reactant is used up at the same rate as the product forms.
• Rate is highest at the start and falls as reactants are used up.

Worked reminder. The concentration of a product rises from 0 to 0.036 mol dm⁻³ in 120 s. Mean rate = 0.036 ÷ 120 = 3.0 × 10⁻⁴ mol dm⁻³ s⁻¹.

3.1.6.1 Collision theory

Method: a three-link chain for any rate question

• Link A: do particles now meet more or less often?
• Link B: is a bigger or smaller fraction of those meetings carrying at least Ea?
• Link C: finish with the effect on successful collisions per second, then on rate.

Why most collisions fail: in a typical mixture only a tiny minority of particles carry Ea or more. A collision between two lower-energy particles cannot break the necessary bonds, so the pair rebounds with no change.

3.1.6.2 Maxwell–Boltzmann distribution: drawing rules

Feature	What to show	Why
Start	Line leaves (0, 0)	Zero-energy molecules do not exist
Peak	One maximum	It marks the most probable energy
Shape	Lopsided, long right-hand side	The tail pulls the mean energy above the peak energy
Tail	Flattens towards the energy axis, never reaching it	No upper energy limit
Area under line	Fixed for a given sample	It represents every molecule present
Area beyond Ea	Small sliver in the tail	Molecules able to react on collision

Method: comparing a cool and a warm sample on shared axes

• Cool curve: taller, narrower, peak nearer the origin.
• Warm curve: shorter, broader, peak displaced to higher energy.
• One crossing point; to the right of it the warm curve is on top, so place Ea there.
• Enclosed areas match, because no molecules are gained or lost.

Reading areas from a grid. Proportion with E ≥ Ea = (squares beyond Ea ÷ total squares) × 100. Compare two curves or two Ea lines by dividing one count by the other.

3.1.6.3 Temperature

• Warming raises the average kinetic energy of the molecules.
• Big effect: the distribution moves to higher energy, so the fraction at or above Ea multiplies.
• Small effect: faster molecules meet slightly more often.
• Why “small rise, big effect”: Ea sits in the thin tail, and a modest sideways shift of the curve swells that thin sliver many times over.
• For a reaction with Ea near 50 kJ mol⁻¹, a 10 K rise near room temperature roughly doubles the rate, while collision frequency rises by only about 2 %.

3.1.6.4 Concentration and pressure

• Concentrated solution: each cm³ holds more reactant particles, so they meet more often; with the same fraction succeeding, the number of successes per second, and so the rate, increases.
• Compressed gas (same temperature): the molecules are packed into less space, so they meet more often and the rate increases.
• From pV = nRT: n/V = p/(RT), so at fixed T, doubling p doubles the amount of gas per unit volume.
• Neither change moves the Maxwell–Boltzmann curve or changes the proportion of collisions with E ≥ Ea.

3.1.6.5 Catalysts on a distribution curve

Method: catalyst sketch for a gas reaction at a fixed temperature

• A single distribution line only, since temperature (and so the spread of energies) is unchanged.
• Two energy markers in the tail: the original Ea, and Ea(cat) at a lower energy.
• Shade the band between them: molecules that only the new pathway lets react.
• Words to include: different pathway, smaller energy barrier, bigger fraction able to react, more successful collisions each second.

• The catalyst is left with the same mass and chemical composition at the end.
• On an enthalpy profile, the catalysed hump is lower; reactant and product levels, and so ΔH, do not change.
• A catalyst does not change the position of equilibrium (3.1.7); see equilibria revision notes.

Must-know distinctions

• Hotter: meetings slightly more frequent; fraction with E ≥ Ea much bigger; curve flattens and moves right; barrier unchanged.
• More concentrated, or gas compressed: meetings more frequent; fraction unchanged; curve unchanged; barrier unchanged.
• Catalyst added: meeting frequency unchanged; fraction bigger because the barrier is lower on a new pathway; curve unchanged, only the Ea marker moves.
• Peak vs average: the peak gives the commonest energy; the long tail lifts the average above it.
• Frequency vs proportion: concentration and pressure change how often particles collide; temperature and catalysts mainly change what fraction of collisions succeed.
• Curve vs Ea line: temperature moves the curve; a catalyst moves the Ea line.

Quick self-test

1. Define activation energy.
2. Why do most collisions in a reaction mixture not lead to a reaction?
3. On a Maxwell–Boltzmann sketch, what does the area beyond the Ea line stand for?
4. Why is the mean energy greater than the most probable energy?
5. A reactant concentration falls from 0.64 mol dm⁻³ to 0.40 mol dm⁻³ in 80 s. Calculate the mean rate.
6. A gas syringe collects 66 cm³ of gas in the first 40 s. Calculate the mean rate.
7. A curve has a total area of 2400 squares, with 30 squares beyond Ea. What percentage of molecules have E ≥ Ea?
8. Beyond Ea there are 5 squares; beyond Ea(cat) there are 40. By what factor has the catalyst increased the proportion of molecules that can react?
9. A catalyst lowers Ea from 120 kJ mol⁻¹ to 75 kJ mol⁻¹. By how much has the barrier fallen?
10. Does doubling the pressure of a gas change the proportion of collisions with E ≥ Ea?
11. When a reaction mixture is warmed, which effect contributes most to the rise in rate?
12. Name two things about a catalyst that are the same at the end of a reaction as at the start.

Answers

1. The minimum energy that particles must have when they collide for a reaction to occur.
2. Most particles have less energy than the activation energy.
3. The molecules with enough energy to react if they collide.
4. The curve is not symmetrical; the long high-energy tail pulls the mean to the right of the peak.
5. (0.64 − 0.40) ÷ 80 = 0.24 ÷ 80 = 3.0 × 10⁻³ mol dm⁻³ s⁻¹.
6. 66 ÷ 40 = 1.65 cm³ s⁻¹.
7. 30 ÷ 2400 × 100 = 1.25 %.
8. 40 ÷ 5 = 8 times.
9. 120 − 75 = 45 kJ mol⁻¹.
10. No. It increases collision frequency only.
11. The larger proportion of molecules with energy at or above Ea.
12. Its chemical composition and its amount (mass).

Where marks are usually lost

• An Ea definition without “minimum”.
• A warm curve drawn taller than the cool one, or with more area beneath it.
• Lines that begin part-way up the y-axis or land on the energy axis.
• Temperature explained by collision frequency alone.
• Concentration or pressure credited with raising particle energy.
• A catalyst said to “supply energy” or to shift the distribution.
• “Lower activation energy” with no mention of a different pathway.
• Rate units missing the “per second” part.
• Calling the peak the mean energy.
• Stopping before the final link: successful collisions per second.

Next steps

Return to the study guide for any point you missed, then work through the practice questions. Related International AS pages: amount of substance for pV = nRT, and energetics for enthalpy profiles.

Official syllabus

OxfordAQA, International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards: section 3.1.6 Kinetics (International AS), points 3.1.6.1 to 3.1.6.5.