AQA GCSE Chemistry 8462: The rate and extent of chemical change – Revision Notes

These notes condense section 4.6 The rate and extent of chemical change of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 onwards and exams in 2018 onwards (version 1.1). They cover spec points 4.6.1.1 to 4.6.2.7 and required practical activity 5. The topic is examined on Paper 2 at Foundation and Higher Tier; Higher tier only content is labelled. For full explanations and worked examples, use the rate and extent of chemical change study guide.

Test yourself afterwards with the practice questions. The course hub is AQA GCSE Chemistry, the printable checklist lists every spec point, and the free 10-minute diagnostics show where to focus. Activation energy is also in the Energy changes revision notes.

Key definitions

Term	Definition to learn
Mean rate of reaction	Quantity of reactant used (or product formed) ÷ time taken
Activation energy	The minimum amount of energy that particles must have to react
Catalyst	Changes the rate of a reaction but is not used up
Enzyme	A catalyst in biological systems
Reversible reaction	The products can react to produce the original reactants
Equilibrium	In a closed system, the forward and reverse reactions occur at exactly the same rate
Le Chatelier’s Principle (Higher tier only)	If a condition is changed at equilibrium, the system responds to counteract the change

4.6.1.1 Rate formulas and units

Quantity measured	Formula	Unit
Mass (g)	mass ÷ time	g/s
Gas volume (cm³)	volume ÷ time	cm³/s
Moles (Higher tier only)	(mass ÷ Mr) ÷ time	mol/s
Rate at one moment	slope of the tangent	as the graph axes
Colour or cloudiness method	1 ÷ time	s⁻¹ (relative rate)

Small worked reminders

• 84 cm³ of gas in 60 s → 84 ÷ 60 = 1.4 cm³/s.
• 0.36 g of magnesium used in 120 s → 0.36 ÷ 120 = 0.003 g/s.
• (Higher tier only) 0.44 g of CO₂ (Mr 44) in 50 s → 0.44 ÷ 44 = 0.010 mol → 0.010 ÷ 50 = 2.0 × 10⁻⁴ mol/s.

Tangent gradient – method in steps (Higher tier only)

1. Place a ruler so it touches the curve at the chosen time only.
2. Draw the tangent long – across most of the grid.
3. Read two points far apart on the tangent.
4. Gradient = (y₂ − y₁) ÷ (x₂ − x₁). Add the unit.

Worked reminder. Tangent through (10 s, 20 cm³) and (30 s, 50 cm³): (50 − 20) ÷ (30 − 10) = 1.5 cm³/s.

Reading rate graphs

• Steep at the start → fastest rate (highest concentration).
• Curve flattens → rate falls as reactants are used up.
• Horizontal → reaction has stopped (one reactant used up).
• A faster reaction with the same amounts reaches the same final amount, sooner.

Comparing two runs on one graph

Questions often show two curves for the same reaction under different conditions.

• The steeper curve at the start shows the faster reaction (for example higher temperature, powder instead of lumps, or a catalyst).
• If the same amounts of reactants are used, both curves level off at the same final volume or mass.
• If the final amount is half as big, half the amount of the limiting reactant was used – not a slower rate.
• A curve that levels off earlier means the reaction finished sooner.

4.6.1.2 Factors – recall list

Rate increases with: higher concentration, higher pressure (gases), larger surface area (smaller pieces), higher temperature, and adding a catalyst.

Required practical 5 – method in steps

Hypothesis first: for example, “the higher the concentration of acid, the faster the rate”.

Gas method

1. Measure a fixed volume of acid of known concentration into a conical flask.
2. Connect a gas syringe. Add a fixed length of magnesium ribbon and start the timer.
3. Record the gas volume every 10 s. Repeat for other concentrations.

Turbidity method

1. Put a flask on a paper cross. Add sodium thiosulfate solution.
2. Add hydrochloric acid and start the timer.
3. Stop the timer when the cross disappears. Calculate 1/time. Repeat for other concentrations.

Control: temperature, volumes, same cross and same observer.

4.6.1.3 Collision theory – method in steps

To explain any rate change:

1. Say what happens to the particles (more in the same volume, closer, more exposed, faster).
2. Say more frequent collisions (collisions per second).
3. For temperature, add more energetic collisions, so more collisions have at least the activation energy.
4. Conclude: so the rate increases.

Surface area to volume ratio. A 3 cm cube: surface area 54 cm², volume 27 cm³, ratio 2 : 1. Smaller pieces have a larger ratio, so the rate is faster.

Proportionality. Double the concentration → about double the collision frequency → about double the rate. Example: times of 50 s and 25 s give 1/t of 0.02 and 0.04 s⁻¹ – the rate doubled.

4.6.1.4 Catalysts

• Provide a different pathway with a lower activation energy.
• Not used up; mass unchanged at the end; not in the chemical equation.
• Catalysed reaction profile: same reactant and product levels, lower hump.

4.6.2.1–4.6.2.3 Reversible reactions and equilibrium

• ⇌ means reversible. Changing conditions can change the direction.
• Ammonium chloride ⇌ ammonia + hydrogen chloride (heat forward, cool back).
• Exothermic one way = endothermic the other way, and the same amount of energy is transferred.
• Hydrated copper sulfate (blue) ⇌ anhydrous copper sulfate (white) + water. Forward endothermic; adding water to the white solid is exothermic.
• Equilibrium needs a closed system; forward rate = reverse rate. Reactions do not stop.

4.6.2.4–4.6.2.7 Changing conditions (Higher tier only)

Change	Equilibrium response
Increase concentration of a reactant	More products form
Decrease concentration of a product	More reactants react
Increase temperature	Shifts in the endothermic direction
Decrease temperature	Shifts in the exothermic direction
Increase pressure (gases)	Shifts to the side with fewer molecules
Decrease pressure (gases)	Shifts to the side with more molecules

Worked reminder. N₂O₄(g) ⇌ 2NO₂(g), forward endothermic. Increasing the temperature increases the relative amount of NO₂. Increasing the pressure shifts left (1 molecule vs 2).

Must-know distinctions

• Mean rate vs rate at a time: whole interval vs tangent at one point.
• More frequent vs more energetic collisions: concentration, pressure and surface area give only the first; temperature gives both.
• Rate vs yield: how fast vs how much product at equilibrium.
• Reversible vs at equilibrium: a reaction can be reversible without being at equilibrium; equilibrium needs a closed system and equal rates.

Quick self-test

1. 45 cm³ of gas is made in 90 s. Calculate the mean rate.
2. Give the unit of rate when mass is measured in grams.
3. What does the slope of a tangent to a product–time curve measure?
4. Name the five factors that affect rate.
5. Explain, using collision theory, why powder reacts faster than lumps.
6. Give the two reasons a higher temperature increases rate.
7. How does a catalyst increase the rate?
8. Why is a catalyst not written in the chemical equation?
9. What is true about the forward and reverse rates at equilibrium?
10. (Higher tier only) 2NO₂(g) ⇌ N₂O₄(g). Which way does the equilibrium shift if the pressure is increased?
11. (Higher tier only) A forward reaction is exothermic. What happens to the relative amount of products if the temperature is raised?
12. (Higher tier only) 1.1 g of CO₂ (Mr 44) is made in 25 s. Calculate the rate in mol/s.

Answers

1. 45 ÷ 90 = 0.5 cm³/s.
2. g/s.
3. The rate of reaction at that time.
4. Concentration, pressure, surface area, temperature, catalyst.
5. Larger surface area, so more particles exposed and more frequent collisions.
6. More frequent collisions; more energetic collisions (more with at least the activation energy).
7. It provides a different pathway with a lower activation energy.
8. It is not used up in the reaction.
9. They are exactly equal.
10. Right, towards N₂O₄ (1 molecule instead of 2).
11. It decreases.
12. 1.1 ÷ 44 = 0.025 mol; 0.025 ÷ 25 = 0.001 mol/s (1.0 × 10⁻³ mol/s).

Where marks are usually lost

• Rate answers with no unit, or cm³ and g mixed up.
• Tangents that cross the curve, or gradients from two points too close together.
• Writing “more collisions” instead of “more frequent collisions”.
• Missing the energy point for temperature: collisions are more energetic, so more exceed the activation energy.
• Saying catalysts “give particles energy” or “are not involved in the reaction”.
• Stating that reactions stop at equilibrium.
• Forgetting that equilibrium needs a closed system.
• Counting atoms, not molecules, when predicting pressure effects.
• Giving the direction of shift but not the effect on the amount of the named product.

Official syllabus

AQA GCSE Chemistry (8462) specification, for teaching from September 2016 onwards, exams in 2018 onwards, version 1.1 (4 October 2019), published by AQA – section 4.6 The rate and extent of chemical change.