What you'll learn
This topic examines how quickly chemical reactions occur and the factors that control their speed. Understanding reaction rates is essential for CIE IGCSE Chemistry papers, where you'll analyse experimental data, interpret graphs, and explain observations using collision theory. Questions typically appear in both Paper 2 (core) and Paper 4 (extended), often combined with practical skills.
Key terms and definitions
Rate of reaction — the speed at which reactants are converted into products, measured as the change in concentration of a reactant or product per unit time.
Collision theory — the principle that particles must collide with sufficient energy and correct orientation for a reaction to occur.
Activation energy — the minimum energy that colliding particles must possess for a successful collision that leads to reaction.
Catalyst — a substance that increases the rate of a chemical reaction without being chemically changed or used up in the process.
Successful collision — a collision between reactant particles that results in a chemical reaction, requiring both sufficient energy (equal to or greater than activation energy) and appropriate orientation.
Surface area — the total area of a solid substance exposed to other reactants, which increases when the solid is divided into smaller pieces.
Concentration — the amount of solute dissolved in a given volume of solution, measured in mol/dm³ or g/dm³.
Precipitate — an insoluble solid formed when two solutions react, often used in rate experiments where the formation can be timed.
Core concepts
Measuring rates of reaction
The rate of reaction can be measured by monitoring changes in reactants or products over time. CIE IGCSE Chemistry examinations focus on three principal methods:
Measuring gas volume produced: When a reaction produces gas, collect it using a gas syringe or inverted measuring cylinder over water. Record the volume at regular time intervals. The steeper the gradient on a volume-time graph, the faster the reaction rate.
Measuring mass loss: When gas escapes from a reaction vessel on a balance, the mass decreases. Record mass at regular intervals. This method works well for reactions producing carbon dioxide, such as marble chips with hydrochloric acid.
Measuring time for a precipitate to form: For reactions producing cloudiness (like sodium thiosulfate with hydrochloric acid), measure the time taken for a cross marked under the flask to disappear from view. A shorter time indicates a faster reaction rate.
Measuring colour change: Some reactions involve colour changes that can be timed, though this method appears less frequently in CIE papers.
Graphs plotting these measurements against time show characteristic curves. Initially steep, the gradient decreases as reactants are used up, eventually becoming horizontal when the reaction completes.
Collision theory and successful collisions
For a chemical reaction to occur, reactant particles must collide. However, not all collisions result in reaction. Collision theory explains the conditions necessary:
- Particles must collide with energy equal to or greater than the activation energy
- Particles must collide with the correct orientation for bonds to break and form
- Only collisions meeting both criteria are successful collisions that produce products
Increasing the rate of reaction means increasing the frequency of successful collisions per unit time. Every factor that affects reaction rate does so by changing either:
- The total number of collisions per second
- The proportion of collisions with sufficient energy
Factors affecting reaction rate: concentration and pressure
Concentration (for solutions) and pressure (for gases) affect the number of particles in a given volume.
When concentration increases:
- More reactant particles occupy the same volume
- Particles are closer together
- Collisions occur more frequently
- More successful collisions per second
- Rate of reaction increases
For gaseous reactants, increasing pressure has the same effect as increasing concentration — more molecules in the same space leads to more frequent collisions.
CIE IGCSE exam questions often present data tables showing reaction time decreasing as concentration increases. Remember: a shorter reaction time means a faster rate. If asked to calculate rate, use rate = 1/time.
Factors affecting reaction rate: temperature
Temperature increases have a dramatic effect on reaction rate. Even a 10°C rise typically doubles or triples the rate.
When temperature increases:
- Particles possess more kinetic energy
- Particles move faster
- Collisions become more frequent
- A greater proportion of particles now possess energy equal to or exceeding activation energy
- More collisions are successful
- Rate of reaction increases significantly
The primary reason for increased rate is not simply more collisions, but that more collisions now have sufficient energy to overcome the activation energy barrier. This distinction is crucial for extended-tier responses.
Factors affecting reaction rate: surface area
Surface area only applies to solid reactants. Breaking a solid into smaller pieces increases the total surface area exposed to other reactants.
When surface area increases (by using smaller pieces, powders rather than lumps):
- More particles on the solid's surface are exposed and available for collision
- Collisions occur more frequently
- More successful collisions per second
- Rate of reaction increases
Classic CIE practical investigations compare marble chips of different sizes reacting with hydrochloric acid. Large chips react slowly; powder reacts rapidly and may even appear violent.
Catalysts and how they work
A catalyst speeds up a chemical reaction without being permanently changed. The same mass of catalyst remains at the end as at the start, though it may appear physically different.
Catalysts work by providing an alternative reaction pathway with lower activation energy:
- The activation energy for the catalysed route is lower than for the uncatalysed route
- More collisions now possess sufficient energy to react
- The proportion of successful collisions increases
- Rate of reaction increases
- The catalyst itself is not consumed, so can be recovered and reused
Examples frequently appearing in CIE papers:
- Manganese(IV) oxide (MnO₂) catalyses the decomposition of hydrogen peroxide to water and oxygen
- Iron catalyses the Haber process (ammonia production)
- Vanadium(V) oxide catalyses the Contact process (sulfuric acid manufacture)
- Enzymes act as biological catalysts
Catalysts are economically valuable in industry because they increase production rates without being consumed, though they may need replacing due to physical degradation or poisoning.
Energy distribution curves and activation energy
Energy distribution curves (Maxwell-Boltzmann distributions) show the range of energies possessed by particles in a sample. These appear in extended-tier CIE papers.
The curve shows:
- Most particles have intermediate energies (peak of curve)
- Few particles have very low or very high energies
- The area under the curve represents the total number of particles
- The activation energy is marked as a vertical line
- Particles with energy ≥ activation energy (area to the right of this line) can react upon collision
When temperature increases, the curve shifts and flattens — significantly more particles now exceed the activation energy threshold. When a catalyst is used, the activation energy line shifts left, meaning more particles now have sufficient energy without changing temperature.
Worked examples
Example 1: Interpreting rate data
A student investigates how concentration of hydrochloric acid affects its reaction with magnesium ribbon. She measures the time taken to collect 50 cm³ of hydrogen gas.
| Concentration (mol/dm³) | Time (s) | Rate (1/time) |
|---|---|---|
| 0.5 | 120 | |
| 1.0 | 60 | |
| 1.5 | 40 | |
| 2.0 | 30 |
(a) Complete the table by calculating the rate for each concentration. [2 marks]
Solution:
- 0.5 mol/dm³: rate = 1/120 = 0.0083 s⁻¹
- 1.0 mol/dm³: rate = 1/60 = 0.017 s⁻¹
- 1.5 mol/dm³: rate = 1/40 = 0.025 s⁻¹
- 2.0 mol/dm³: rate = 1/30 = 0.033 s⁻¹
[1 mark for method, 1 mark for all values correct to 2-3 significant figures]
(b) Explain, using collision theory, why increasing concentration increases the rate of reaction. [3 marks]
Solution:
- Higher concentration means more acid particles in the same volume [1]
- Particles collide more frequently [1]
- More successful collisions per second, so faster reaction [1]
Example 2: Catalyst investigation
Hydrogen peroxide decomposes slowly to form water and oxygen. A student adds manganese(IV) oxide to hydrogen peroxide solution.
(a) State the role of manganese(IV) oxide in this reaction. [1 mark]
Solution: Catalyst [1]
(b) Explain how manganese(IV) oxide increases the rate of this decomposition. [2 marks]
Solution:
- It provides an alternative pathway [1]
- With lower activation energy [1]
Alternatively: More particles now have sufficient energy to react / greater proportion of successful collisions [2]
(c) A student uses 0.5 g of manganese(IV) oxide. State the mass of catalyst remaining after the reaction finishes. Explain your answer. [2 marks]
Solution:
- 0.5 g [1]
- Catalyst is not used up / not chemically changed in the reaction [1]
Example 3: Practical design
Describe how you would investigate the effect of temperature on the rate of reaction between sodium thiosulfate solution and dilute hydrochloric acid. Include how you would make the investigation a fair test. [6 marks]
Solution:
- Measure fixed volumes of sodium thiosulfate and hydrochloric acid [1]
- Heat sodium thiosulfate to required temperature using water bath [1]
- Place flask on paper marked with a cross [1]
- Add hydrochloric acid and start timing immediately [1]
- Record time taken for cross to disappear [1]
- Keep volumes/concentrations constant; change only temperature [1]
Alternative acceptable methods include measuring mass loss or gas production with different reactions.
Common mistakes and how to avoid them
Mistake: Stating that "particles have more energy to collide" when explaining temperature effects. Correction: Particles already collide at all temperatures. Increased temperature means particles move faster and collide with more energy, so more collisions exceed the activation energy.
Mistake: Confusing rate with time — writing that rate increases when time increases. Correction: Rate and time are inversely related. When time decreases, rate increases. Rate = 1/time.
Mistake: Claiming catalysts are used up or that they "speed up particles." Correction: Catalysts are not consumed (though they may become physically degraded). They provide an alternative pathway with lower activation energy; they do not affect particle speed or energy.
Mistake: Explaining surface area by saying "more surface area means more particles." Correction: The number of particles remains constant. Increasing surface area means more particles are exposed on the surface and available for collision.
Mistake: Writing that concentration or pressure changes the activation energy. Correction: Only catalysts and temperature changes affect whether collisions have sufficient energy. Concentration and pressure affect collision frequency only.
Mistake: Drawing rate-of-reaction graphs with straight lines or unusual shapes. Correction: Volume-time and mass-time graphs start steep (fast rate), gradually level off (slowing rate), and become horizontal (reaction complete). The curve is smooth, not angular.
Exam technique for Rates of Reaction
Command word 'Explain': These questions require reasoning, not just statements. Use collision theory language: reference particle collisions, energy, frequency, and successful collisions. For example, don't just write "higher temperature increases rate" — explain that particles move faster, collide more frequently, and more collisions have energy ≥ activation energy. Expect 2-4 marks.
Practical questions: CIE frequently asks you to describe methods for investigating rates. Always specify measurements (volumes, masses, times), controlled variables, and safety. State how you would make the test fair. Questions typically worth 4-6 marks — check the mark allocation to gauge required detail.
Graph interpretation: You may need to calculate rate from gradient (change in y / change in x) or compare steepness of different curves. Steeper gradient = faster rate. State which line represents which condition clearly.
Using 'collision theory' in answers: Extended-tier questions often require collision theory explanations. Structure your answer: state what changes (e.g., particle proximity, particle energy), explain the effect on collisions (frequency or energy), conclude with successful collisions and rate. This three-step structure typically matches the mark scheme.
Quick revision summary
Rate of reaction measures how quickly reactants form products. Measured by volume of gas produced, mass loss, or precipitation time. Collision theory: particles must collide with sufficient energy (≥ activation energy) and correct orientation. Increasing concentration, pressure, temperature, or surface area increases collision frequency. Temperature also increases the proportion of particles exceeding activation energy. Catalysts provide alternative pathways with lower activation energy, increasing successful collision proportion without being consumed. Rate = 1/time for simple calculations.