AQA A-Level Chemistry: Kinetics (7405) – Revision Notes

These notes condense section 3.1.5 Kinetics of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards. They cover 3.1.5.1 to 3.1.5.5 and Required practical 3. Everything here is AS content, so none of it is A-level only; it is examined on Paper 2 at AS and A-level, and A-level Paper 3 can test any content. For full explanations and worked examples, read the kinetics study guide. To test yourself, use the kinetics practice questions.

Course hub: AQA A-Level Chemistry. Tick off each point on the printable checklist. Rate equations and orders belong to 3.1.9 (A-level only), not to this topic.

Definitions to learn word for word

Term	Definition
Activation energy, Ea (3.1.5.1)	The minimum energy that colliding particles must have for a reaction to occur
Rate of reaction (3.1.5.3)	The change in amount or concentration of a reactant or product per unit time
Catalyst (3.1.5.5)	A substance that increases the rate of a chemical reaction without being changed in chemical composition or amount
How a catalyst works (3.1.5.5)	It provides an alternative reaction route of lower activation energy

Rate calculations

Quantity measured	Mean rate	Typical unit
Concentration	change in concentration ÷ time	mol dm⁻³ s⁻¹
Gas volume	change in volume ÷ time	cm³ s⁻¹
Mass lost	change in mass ÷ time	g s⁻¹
Amount	change in moles ÷ time	mol s⁻¹
Time to a fixed point (clock or cross)	rate ∝ 1/t	s⁻¹

• Rate at one instant = gradient of the tangent to the curve at that time.
• Initial rate = gradient of the tangent at t = 0.
• The rate falls as a reaction goes on, because reactant concentration falls.

3.1.5.1 Collision theory in one box

Method: explaining any rate change

1. Say whether the frequency of collisions changes.
2. Say whether the proportion of particles with E ≥ Ea changes.
3. Conclude: more (or fewer) successful collisions per unit time, so rate rises (or falls).

Why most collisions fail: most particles have energy less than Ea, so no bonds break when they collide.

3.1.5.2 Maxwell–Boltzmann distribution: drawing rules

• Axes: energy (x) and number of molecules with that energy (y).
• Starts at the origin.
• One peak = most probable energy. The mean energy is to the right of the peak.
• Long tail; the curve never meets the energy axis.
• Area under the curve = total number of molecules.
• Area to the right of Ea = number of molecules that can react on collision.

Method: two temperatures on one set of axes

1. Draw the lower-temperature curve first.
2. Draw the higher-temperature curve with its peak lower and further right.
3. Make the two curves cross once, with the hot curve above the cold one at high energy.
4. Both start at the origin; neither touches the x-axis; areas look equal.
5. Mark Ea in the tail, to the right of both peaks.

3.1.5.3 Temperature

• Higher temperature → molecules have more kinetic energy.
• Curve shifts right → many more molecules have E ≥ Ea (main reason).
• Molecules also move faster → slightly more frequent collisions (minor reason).
• Small temperature rise → large rise in rate, because Ea is far out in the tail and the tail area grows by a large factor.

Required practical 3: rate and temperature

• The specification links it to AT a (measure mass, time, volume, temperature), AT b (water bath or electric heater) and AT k (safe handling).
• One suggested reaction (initial rate method): sodium thiosulfate + hydrochloric acid.

Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + SO₂(g) + S(s) + H₂O(l)

• Time how long until a cross under the flask is hidden by the sulfur precipitate.
• Same amount of sulfur each run, so rate ∝ 1/t.
• Control: volumes and concentrations of both solutions, total volume, the same cross and flask, the same observer.
• Heat both solutions separately to the temperature before mixing; record the temperature of the mixture.
• Hazard: sulfur dioxide is toxic and irritant, so use good ventilation and small volumes.
• Plot 1/t against temperature: the curve gets steeper as temperature rises.

3.1.5.4 Concentration and pressure

• Higher concentration → more particles per unit volume → more frequent collisions → more successful collisions per second → higher rate.
• Higher gas pressure → more molecules per unit volume → more frequent collisions → higher rate.
• The proportion of particles with E ≥ Ea does not change, and the Maxwell–Boltzmann curve does not move.
• Continuous monitoring example from the specification: CaCO₃ + hydrochloric acid, following CO₂ volume (gas syringe) or mass lost (balance).

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

3.1.5.5 Catalysts and the distribution curve

Method: catalyst on a Maxwell–Boltzmann diagram

1. Draw one curve (temperature unchanged, so it does not move).
2. Mark Ea (uncatalysed) and Ea(cat) to its left.
3. Shade or label the extra area between them.
4. State: alternative route, lower Ea, greater proportion of molecules with E ≥ Ea(cat), more successful collisions per second.

The catalyst is not used up: the same mass and composition remain at the end.

Must-know distinctions

Change	Collision frequency	Proportion with E ≥ Ea	Curve moves?	Ea changes?
Raise temperature	Up a little	Up a lot	Yes, right and lower	No
Raise concentration	Up	No change	No	No
Raise gas pressure	Up	No change	No	No
Add a catalyst	No change	Up (Ea marker moves left)	No	Lower route provided

• Most probable energy vs mean energy: the peak is the most probable; the mean is to its right.
• Initial rate method vs continuous monitoring: an initial rate method times one early fixed point (for example, a cross disappearing); continuous monitoring records the change throughout one run.
• Mean rate vs instantaneous rate: mean rate is a change divided by an interval; instantaneous rate is a tangent gradient.

Quick self-test

1. Define activation energy.
2. Give the main reason most collisions do not lead to a reaction.
3. On a Maxwell–Boltzmann curve, what does the total area represent?
4. Why must the higher-temperature curve have a lower peak?
5. A concentration falls from 0.80 mol dm⁻³ to 0.56 mol dm⁻³ in 60 s. Calculate the mean rate.
6. In Required practical 3, the cross disappears after 25 s. Calculate 1/t.
7. A gas syringe collects 90 cm³ of gas in the first 45 s. Calculate the mean rate.
8. A flask loses 0.33 g of CO₂ (Mr = 44.0) in 60 s. Calculate the mean rate in mol s⁻¹.
9. Does increasing the pressure of a gas change the proportion of molecules with E ≥ Ea?
10. State two ways a catalyst does not change during a reaction.
11. How does a catalyst change a Maxwell–Boltzmann diagram for a gas reaction at fixed temperature?

Answers

1. The minimum energy that colliding particles must have for a reaction to occur.
2. Most particles have energy less than the activation energy.
3. The total number of molecules in the sample.
4. The number of molecules is unchanged, so the area must stay the same; spreading the curve to higher energy lowers the peak.
5. (0.80 − 0.56) ÷ 60 = 0.24 ÷ 60 = 0.0040 mol dm⁻³ s⁻¹ (4.0 × 10⁻³).
6. 1 ÷ 25 = 0.040 s⁻¹.
7. 90 ÷ 45 = 2.0 cm³ s⁻¹.
8. 0.33 ÷ 44.0 = 0.0075 mol; 0.0075 ÷ 60 = 1.25 × 10⁻⁴ mol s⁻¹.
9. No. It raises collision frequency only.
10. Its chemical composition and its amount.
11. The curve stays the same; the activation energy line moves to the left (lower Ea by an alternative route), so the area to its right is larger.

Where marks are usually lost

• Leaving out “minimum” from the activation energy definition, or writing “energy to start a reaction”.
• Drawing a Maxwell–Boltzmann curve that touches or crosses the energy axis at high energy.
• Starting the curve above zero on the y-axis instead of at the origin.
• Drawing the hotter curve taller, which implies more molecules.
• Giving “particles collide more often” as the only reason for the temperature effect.
• Writing that higher concentration gives particles more energy, or a higher proportion with E ≥ Ea.
• Saying a catalyst “gives particles more energy” or moves the distribution curve.
• Missing “alternative route” when explaining how a catalyst works.
• Using t instead of 1/t as the measure of rate in a clock or disappearing-cross experiment.
• Leaving units off a rate, or writing mol dm⁻³ instead of mol dm⁻³ s⁻¹.

Next steps

Go back to the study guide for any point you missed, then try the practice questions. The free 10-minute diagnostics help you find weak topics across the course.

Official syllabus

AQA, AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards: section 3.1.5 Kinetics (3.1.5.1 to 3.1.5.5), with Required practical 3.