Rates of Reaction
Edexcel International GCSE Chemistry· 3.9–3.16 (sub-topic 3(b))· 25 min read
1. Measuring Reaction Rate★★☆☆☆⏱ 5 min
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Rate of reaction
The measure of how fast reactants are converted into products in a chemical reaction, calculated as change in quantity of reactant or product per unit time.
Example:
Rate of CO₂ production from marble + HCl = change in gas volume ÷ time taken
Reaction rate can be measured using three common methods: measuring mass lost over time (for reactions producing gas), measuring gas volume collected over time, or timing a colour change/precipitate formation. On a reaction graph, the steeper the gradient at any point, the faster the rate at that time.
A student collects 45 cm³ of CO₂ from a marble chip and HCl reaction in 30 seconds. Calculate the average rate of reaction over this period.
- 1
Use the average rate formula: rate = change in quantity ÷ time
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Final answer: 1.5 cm³/s
2. Core Factors Affecting Reaction Rate★★★☆☆⏱ 7 min
Five key factors change reaction rate, and you are expected to describe each effect and link it to practical observations:
Increased surface area of solid reactants: faster rate
Increased concentration of dissolved reactants: faster rate
Increased gas pressure: faster rate
Increased temperature: faster rate
Presence of a catalyst: faster rate (catalyst remains unchanged)
A student uses large marble chips with dilute HCl, then repeats with the same mass of powdered marble. Describe the difference in observed rate.
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Observation: Powdered marble reacts much faster, so bubbles of CO₂ are produced more rapidly, and the reaction finishes sooner.
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Explanation: Powdered marble has a larger surface area exposed to the HCl particles, so more frequent collisions occur between reactant particles.
3. Collision Theory Explanations★★★☆☆⏱ 6 min
Collision Theory
A model explaining reaction rate: reactions occur when reactant particles collide with enough energy (equal to or greater than activation energy) and the correct orientation.
Increased surface area, concentration, or pressure only increase collision frequency, so more collisions happen per second, increasing the rate. Increased temperature increases both collision frequency (particles move faster) and the proportion of collisions with energy ≥ activation energy, so it has a larger effect on rate than the other factors.
Explain why increasing the pressure of a gaseous reaction increases the rate, using collision theory.
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Increased pressure pushes gas particles closer together, so the concentration of gas particles increases.
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This increases the frequency of collisions between reactant gas particles, leading to more successful collisions per second, so rate increases.
4. Catalysts & Reaction Profiles (Core + Higher Tier)★★★★☆Core / C only⏱ 7 min
Catalyst
A substance that increases the rate of a chemical reaction, but remains chemically unchanged and in the same mass at the end of the reaction.
Catalysts work by providing an alternative reaction pathway with a lower activation energy. This means a higher proportion of collisions have enough energy to be successful, increasing rate without the catalyst being used up.
Draw a reaction profile for an exothermic reaction, labelling Eₐ for uncatalysed and catalysed reactions, and ΔH.
- 1
Draw a horizontal line for reactant energy level, then a peak for the activation energy, then a lower horizontal line for product energy level (exothermic, so products lower than reactants).
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Label the vertical difference between reactants and the peak as uncatalysed Eₐ. Draw a lower peak for the catalysed pathway, label its vertical difference from reactants as catalysed Eₐ.
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Label the vertical difference between reactant and product energy levels as ΔH (negative for exothermic reactions).
5. Required Practicals for Rates of Reaction★★★☆☆⏱ 5 min
Two required practicals are specified for this topic, and questions on their method, variables and results appear frequently in exams.
Practical 1: Effect of surface area of marble chips and HCl concentration on reaction rate: measure CO₂ volume or mass lost over time, keep other variables constant (temperature, mass of marble, volume of HCl).
Practical 2: Effect of different solid catalysts on H₂O₂ decomposition: measure volume of O₂ produced over time, test with glowing splint, keep H₂O₂ concentration, temperature and mass of catalyst constant.
A student tests the effect of MnO₂, CuO and ZnO on H₂O₂ decomposition. State the independent, dependent and control variables for this experiment.
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Independent variable: Type of solid catalyst used.
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Dependent variable: Volume of oxygen produced per unit time (rate of reaction).
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Control variables: Concentration and volume of hydrogen peroxide, mass of catalyst used, temperature of the reaction mixture.
6. Common Pitfalls
Wrong move:
Stating catalysts are used up during reactions
Why:
Catalysts only provide an alternative lower Eₐ pathway, they do not react permanently with reactants
Correct move:
State catalysts remain chemically unchanged and can be reused at the end of the reaction
Wrong move:
Explaining temperature effect only by increased collision frequency
Why:
Higher temperature increases both collision frequency and proportion of collisions with energy ≥ Eₐ, the energy effect is larger
Correct move:
Mention both frequency and energy factors when explaining temperature effect on rate
Wrong move:
Confusing surface area with total mass of solid
Why:
Two samples of marble with the same mass can have different surface areas if one is powdered and one is lumps
Correct move:
State that larger surface area comes from smaller pieces of the same mass of solid
Wrong move:
Drawing catalysed Eₐ higher than uncatalysed Eₐ on reaction profiles
Why:
Catalysts lower activation energy, so the catalysed peak is lower
Correct move:
Always draw the catalysed reaction peak below the uncatalysed peak on reaction profiles
Wrong move:
Claiming catalysts change the enthalpy change (ΔH) of a reaction
Why:
Catalysts only change activation energy, the energy of reactants and products stays the same, so ΔH is unchanged
Correct move:
Note that ΔH is identical for catalysed and uncatalysed versions of the same reaction
7. Quick Reference Cheatsheet
Factor | Effect on Rate | Collision Theory Explanation |
|---|---|---|
Increased solid surface area | Increase | Higher frequency of collisions (more exposed particles) |
Increased solution concentration / gas pressure | Increase | Higher frequency of collisions (particles closer together) |
Increased temperature | Large increase | Higher collision frequency + more collisions have energy ≥ Eₐ |
Addition of catalyst | Increase | Alternative pathway with lower activation energy, more successful collisions |
8. Frequently Asked
Do catalysts get used up during reactions?
No. Catalysts provide an alternative lower activation energy pathway but are chemically unchanged at the end of the reaction, so they can be reused.
Which rate factor increases both collision frequency and energy?
Increased temperature: particles move faster (more frequent collisions) and a higher proportion have energy ≥ activation energy, leading to a large rate increase.
Going deeper
What's Next
Now that you have mastered rates of reaction for Edexcel IGCSE Chemistry, you can move on to closely related physical chemistry topics that often appear alongside rate questions in exams. The next core topic is reversible reactions and equilibrium, where you will learn how rate of forward and reverse reactions equalise in closed systems, and how rate factors affect equilibrium position. You should also revise enthalpy changes to reinforce your understanding of reaction profile diagrams for higher tier exams, and practical design skills to answer required practical 6-mark questions. Past paper practice on rate calculation and collision theory explanation questions will help you secure full marks on this high-weight topic.