What you'll learn
This topic examines the factors that control how quickly chemical reactions proceed and the scientific principles behind them. Understanding collision theory and how to manipulate reaction conditions forms a substantial portion of the CIE IGCSE Chemistry practical and theory papers, appearing in both multiple-choice and structured questions worth significant marks.
Key terms and definitions
Rate of reaction — the speed at which reactants are converted into products, measured by the change in concentration of reactants or products 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 collision to result in a chemical reaction.
Catalyst — a substance that increases the rate of a chemical reaction without being chemically changed or used up at the end of the reaction.
Surface area — the total exposed area of a solid reactant available for collisions with other particles.
Concentration — the amount of solute dissolved in a given volume of solution, typically measured in g/dm³ or mol/dm³.
Successful collision — a collision between reactant particles that has sufficient energy (equal to or greater than activation energy) and correct orientation to break existing bonds and form products.
Core concepts
Collision theory and activation energy
For a chemical reaction to occur, reactant particles must collide with each other. However, not all collisions lead to a reaction. Collision theory states that two conditions must be met:
- Particles must collide with energy equal to or greater than the activation energy
- Particles must collide with the correct orientation so that bonds can break and reform
The rate of reaction increases when the frequency of successful collisions increases. Any factor that increases either the number of collisions or the proportion of collisions with sufficient energy will increase the reaction rate.
On an energy profile diagram, activation energy is shown as the energy barrier between reactants and products. Even in exothermic reactions where products have lower energy than reactants, particles must first overcome this activation energy barrier.
Effect of temperature on reaction rate
Increasing temperature increases the rate of reaction significantly. This occurs for two reasons:
Increased collision frequency: Particles move faster at higher temperatures, resulting in more frequent collisions between reactant particles per unit time.
Increased collision energy: More importantly, a greater proportion of particles possess energy equal to or exceeding the activation energy. This means more collisions are successful, dramatically increasing the reaction rate.
A small temperature increase (typically 10°C) can approximately double the rate of many reactions. This explains why:
- Food spoils faster in warm conditions
- Reactions in the laboratory are often heated to increase their rate
- Refrigeration slows down bacterial growth and decay reactions
In CIE IGCSE Chemistry examinations, questions often require students to explain temperature effects using both collision frequency and energy distribution, not just one factor.
Effect of concentration on reaction rate
For reactions involving solutions, increasing the concentration of reactants increases the rate of reaction.
Why this occurs: When concentration increases, there are more reactant particles in the same volume. This leads to more frequent collisions between reactant particles per unit time, resulting in more successful collisions per second.
For gases, increasing the pressure has the same effect as increasing concentration. Higher pressure forces the same number of gas particles into a smaller volume, increasing the number of particles per unit volume and therefore the collision frequency.
Common examples tested in examinations:
- Hydrochloric acid reacting with calcium carbonate (marble chips)
- Sodium thiosulfate reacting with hydrochloric acid
- Magnesium ribbon reacting with sulfuric acid
The relationship is directly proportional — doubling the concentration typically doubles the rate, though this depends on the specific reaction mechanism.
Effect of surface area on reaction rate
For reactions involving solids, increasing the surface area increases the rate of reaction. This only applies to heterogeneous reactions where a solid reacts with a liquid or gas.
Why this occurs: Breaking a solid into smaller pieces increases the total exposed surface area without changing the total mass. More surface area means more reactant particles are exposed and available for collisions with particles of the other reactant.
Practical examples:
- Powdered calcium carbonate reacts faster than marble chips with the same mass of acid
- Wood shavings burn faster than a log of the same mass
- Finely divided iron powder reacts with sulfur faster than iron filings
The surface area effect is commonly tested using marble chips of different sizes reacting with hydrochloric acid. Large chips react slowly, while powder reacts vigorously. The total volume of gas produced remains the same (determined by the limiting reactant), but the rate differs significantly.
Effect of catalysts on reaction rate
A catalyst is a substance that increases the rate of a chemical reaction without being permanently changed or used up. Catalysts are crucial in industrial processes and biological systems.
How catalysts work: Catalysts provide an alternative reaction pathway with lower activation energy. This means more colliding particles have sufficient energy to react, increasing the proportion of successful collisions without raising the temperature.
Key characteristics of catalysts:
- Only small amounts are needed
- They are chemically unchanged at the end of the reaction
- They do not affect the position of equilibrium or the yield of products
- Different reactions require different specific catalysts
- They can be reused
Industrial catalysts commonly examined:
- Iron in the Haber process (nitrogen + hydrogen → ammonia)
- Vanadium(V) oxide in the Contact process (sulfur dioxide → sulfur trioxide)
- Nickel in the hydrogenation of alkenes (adding hydrogen to unsaturated fats)
Biological catalysts (enzymes):
- Protein molecules that catalyse reactions in living organisms
- Highly specific to particular reactions
- Effective at body temperature (around 37°C)
- Denatured (permanently damaged) by high temperatures or extreme pH
In examinations, students must explain that catalysts lower activation energy but are not used up, and should be able to identify catalysts from reaction data where initial and final masses remain constant.
Measuring rate of reaction
CIE IGCSE Chemistry examinations frequently test practical methods for measuring reaction rate:
1. Measuring gas volume produced
Use a gas syringe or inverted measuring cylinder filled with water to collect gas produced. Record volume at regular time intervals.
Example: calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide
The rate is fastest at the start (steepest gradient on a graph of volume against time) and gradually decreases as reactants are used up.
2. Measuring mass loss
Place the reaction mixture on a balance. As gas escapes, the mass decreases. Record mass at regular intervals.
This method is suitable when the gas can safely be released to the atmosphere (e.g., carbon dioxide but not toxic gases).
3. Measuring time for a fixed change
The classic "disappearing cross" experiment: sodium thiosulfate reacts with hydrochloric acid to produce a precipitate of sulfur. Time how long it takes for a cross drawn under the flask to become invisible.
Shorter time = faster rate
4. Measuring colour change
Some reactions produce coloured products or use up coloured reactants. The time taken for a specific colour change can indicate the rate.
Worked examples
Example 1: Temperature and rate
Question: A student investigated the reaction between magnesium ribbon and dilute hydrochloric acid at 20°C and 40°C. At 20°C, 50 cm³ of gas was produced in 40 seconds. At 40°C, the same volume of gas was produced in 15 seconds.
(a) Calculate the rate of reaction at each temperature. [2 marks] (b) Explain, in terms of particles, why the rate increases at higher temperature. [3 marks]
Answer:
(a) Rate at 20°C = 50 ÷ 40 = 1.25 cm³/s ✓ Rate at 40°C = 50 ÷ 15 = 3.33 cm³/s ✓
(b) At higher temperature, particles have more kinetic energy ✓ Particles collide more frequently ✓ A greater proportion of collisions have energy equal to or greater than the activation energy / more successful collisions ✓
Example 2: Surface area
Question: Marble chips (calcium carbonate) react with hydrochloric acid to produce carbon dioxide gas. A student used 5.0 g of large marble chips, then repeated the experiment with 5.0 g of powdered marble.
(a) State which form of marble would react faster. [1 mark] (b) Explain your answer using collision theory. [2 marks] (c) State whether the total volume of gas produced would be different. Explain your answer. [2 marks]
Answer:
(a) Powdered marble ✓
(b) Powder has greater surface area ✓ More reactant particles are exposed / available for collision (with acid particles), so collision frequency is higher ✓
(c) No (difference in total volume) ✓ Both contain the same mass/moles of calcium carbonate, so same amount of carbon dioxide is produced ✓
Example 3: Catalysts
Question: Hydrogen peroxide decomposes slowly to produce water and oxygen. Manganese(IV) oxide can be added to speed up this reaction.
2H₂O₂(aq) → 2H₂O(l) + O₂(g)
(a) What name is given to a substance like manganese(IV) oxide? [1 mark] (b) Explain how this substance increases the rate of decomposition. [2 marks] (c) A student adds 0.5 g of manganese(IV) oxide at the start. Predict the mass of manganese(IV) oxide present after the reaction is complete. [1 mark]
Answer:
(a) Catalyst ✓
(b) It provides an alternative pathway/route (for the reaction) ✓ With lower activation energy / so more collisions are successful ✓
(c) 0.5 g ✓ (Catalyst is not used up/consumed in the reaction)
Common mistakes and how to avoid them
Mistake: Stating that catalysts "speed up particles" or "give particles more energy." Correction: Catalysts lower the activation energy by providing an alternative pathway. They do not increase particle energy or speed—that is the effect of temperature increase.
Mistake: Confusing concentration with temperature effects, or giving vague answers like "particles move faster" for concentration increases. Correction: Concentration increases collision frequency only (more particles per unit volume). Temperature increases both collision frequency AND the proportion with sufficient energy—always mention both when explaining temperature effects.
Mistake: Stating that larger surface area means "more particles" or changes the amount of product. Correction: Surface area affects the rate by exposing more particles at the surface, but the total number of particles (and therefore total product) remains unchanged. Mass of reactant and total volume of gas produced are the same for different surface areas.
Mistake: Writing that increasing concentration "makes collisions more powerful" or "increases energy of collisions." Correction: Concentration only increases the frequency of collisions, not their energy. The energy of individual collisions depends on temperature.
Mistake: In graph questions, confusing steep gradient (high rate) with high final value. Correction: The gradient of a volume-time or mass-time graph shows the rate. A steeper gradient means a faster rate. The final value (plateau) depends on the amount of limiting reactant, not the rate.
Mistake: Stating that catalysts are "used up slowly" or "last a long time." Correction: Catalysts are not chemically changed at all—they are fully recoverable in their original mass and chemical form after the reaction, allowing them to be reused indefinitely.
Exam technique for "Rate of reaction: factors affecting rate"
Command words and mark allocation: "Explain" questions typically award 2-3 marks and require reference to collision theory (collision frequency and/or collision energy). Simply stating "more collisions" without mentioning frequency or success rate often loses marks. "Describe" questions about graphs require reference to both gradient (rate) and final value (total amount).
Graph interpretation skills: Questions frequently show volume-time or mass-time graphs with different curves. The steepest initial gradient indicates the fastest rate. Where curves level off at the same height, the same amount of product was formed (same limiting reactant). Where they level off at different heights, different amounts of reactant were used.
Practical questions: When asked to describe a method for investigating factors affecting rate, include specific details: how you would change the independent variable (e.g., "use 1.0 mol/dm³, 1.5 mol/dm³, and 2.0 mol/dm³ acid"), what you would measure (e.g., "volume of gas every 30 seconds"), and which variables must be controlled (e.g., "same mass and surface area of marble chips, same volume of acid, same temperature").
Catalyst questions: Always state both key points: catalysts lower activation energy AND are not used up. Questions worth 2 marks typically award one mark for each point.
Quick revision summary
Rate of reaction depends on successful collision frequency. Temperature increases both collision frequency and energy—most significant effect. Concentration (solutions) and pressure (gases) increase particle density, raising collision frequency. Surface area of solids exposes more particles for collision without changing total amount. Catalysts lower activation energy by providing an alternative pathway and are not consumed. Measure rate using gas volume, mass loss, or time for observable change. Graphs show rate by gradient; steeper = faster. Always link explanations to collision theory for full marks.