Factors which affect the rates of chemical reactions — AQA GCSE Combined Science

The rate of a chemical reaction is determined by collision theory, requiring particles to collide with sufficient activation energy. Increasing temperature, concentration, pressure, and surface area to volume ratio, or adding a catalyst, accelerates the rate of reaction by increasing collision frequency or energy.

Rate of reaction measures how quickly reactants are transformed into products over time, governed directly by collision theory. According to this theory, chemical reactions occur only when reacting particles collide with sufficient energy (the activation energy) and the correct orientation.

Understanding the factors affecting reaction rates—including concentration, pressure, surface area, temperature, and catalysts—is central to chemical engineering and industrial processes. Mastering these concepts connects fundamental particle models to quantitative analysis, graphs, and required practical experiments across AQA GCSE Combined Science.

• →Collision Theory: For a reaction to happen, particles must collide with energy greater than or equal to the activation energy.
• →Temperature Effect: Higher temperatures increase particle kinetic energy, leading to both more frequent collisions and a significantly higher proportion of collisions exceeding the activation energy.
• →Concentration and Pressure: Increasing the concentration of solutions or the pressure of gases packs more particles into a given volume, resulting in more frequent collisions.
• →Surface Area to Volume Ratio: Breaking solids into smaller pieces exposes more surface area, allowing more reactant particles to collide simultaneously and increasing collision frequency.
• →Catalysts: Substances that speed up reactions without being consumed by providing an alternative pathway with lower activation energy.

• Concentration: a higher concentration of reactants in solution means more particles in a given volume, so more frequent collisions and a faster rate.
• Pressure: increasing the pressure of reacting gases brings particles closer together, increasing collision frequency and rate.
• Surface area: a solid with a larger surface area, such as a powder, allows more particles to collide at once, increasing the rate.
• Temperature: raising the temperature increases particle speed, so collisions are more frequent and more energetic, increasing the rate.
• Catalysts: a catalyst lowers the activation energy by providing an alternative reaction pathway, increasing the rate without being used up.
• Rate can be followed by measuring the volume of gas produced, the loss in mass, or the time taken for a colour change or precipitate to appear.
• State that increasing temperature increases rate because particles gain kinetic energy, move faster and collide more frequently and with greater energy.
• State that increasing concentration or pressure increases rate because particles are closer together, so collision frequency increases.
• State that increasing surface area of a solid increases rate because more reactant particles are exposed and available to collide.
• State that a catalyst increases rate by providing an alternative reaction pathway of lower activation energy, and is not used up.
• Link each factor to collision frequency or collision energy, not merely to 'more reactions happening'.
• Use correct units where relevant, such as mol/dm³ for concentration and °C or K for temperature.
• Describe measuring gas volume at regular time intervals using a gas syringe or inverted measuring cylinder, and plotting volume against time.
• Describe measuring the time for a cross to disappear as sulfur precipitate forms, and using 1 ÷ time as a measure of rate.
• Identify the independent variable as concentration of one reactant and the dependent variable as gas volume or time.
• State control variables such as temperature, volume of solution, mass or size of solid, and same cross or light intensity.
• Explain that a steeper initial gradient on a volume–time graph indicates a faster rate.
• Explain that a shorter time for the cross to disappear indicates a faster rate.
• State a clear hypothesis that links the chosen factor to the rate of reaction and predicts the direction of change.
• Identify the independent variable, dependent variable and at least two control variables in the investigation.
• Justify the hypothesis using collision theory, referring to frequency or energy of collisions between reacting particles.
• Describe a suitable method for measuring rate, such as volume of gas collected at timed intervals or loss in mass over time.
• Explain how the results will be used to judge whether the hypothesis is supported, for example by comparing gradients or times to completion.

• 💡Link each factor to collision frequency or collision energy, then to the rate, so the explanation shows the mechanism rather than just naming the factor.
• 💡When describing an experiment, name the independent variable, the dependent variable and at least two control variables.
• 💡Use comparative language such as 'more frequent collisions' or 'greater proportion of successful collisions' to make the reasoning explicit.
• 💡For each factor, write a cause-and-effect chain: change → particle behaviour → collision frequency or energy → rate.
• 💡Use comparative language such as 'faster', 'more frequent' and 'greater energy' rather than vague words like 'better'.
• 💡When asked to explain, refer to particles and collisions; when asked to state, a short accurate sentence is enough.
• 💡State clearly what is measured and how it is used to compare rates, for example initial gradient or 1 ÷ time.
• 💡Name the control variables explicitly; examiners look for temperature, volume and solid mass or particle size.
• 💡For the turbidity method, explain that the cross disappears when enough sulfur precipitate has formed to block the view.
• 💡Write the hypothesis as a single sentence containing the independent variable, the dependent variable and the predicted direction of change.
• 💡When describing the method, state exactly what is measured and how often, for example gas volume every 10 seconds.
• 💡Use collision theory vocabulary such as frequency of collisions and activation energy to justify your prediction.
• 💡Always state 'per unit time' or 'frequency of collisions' rather than merely saying 'more collisions' to ensure full marks on AQA mark schemes.
• 💡When interpreting rate graphs, remember that the gradient represents the rate: a steeper line means a faster reaction, and a horizontal line means the reaction has stopped.

• Saying a catalyst increases the rate by increasing collision frequency: it lowers activation energy so more collisions are successful, while collision frequency is largely unchanged.
• Confusing surface area with concentration: surface area applies to solid reactants, while concentration applies to reactants in solution.
• Claiming temperature increases rate only because particles collide more often: the main effect is that more particles have energy greater than or equal to the activation energy.
• Saying a catalyst increases the rate by increasing collision frequency; correction: a catalyst lowers activation energy by providing an alternative pathway, so more collisions are successful.
• Confusing surface area with concentration; correction: surface area applies to solids, while concentration applies to solutions and pressure to gases.
• Stating that increasing temperature makes particles bigger or increases their volume; correction: it increases their kinetic energy and speed, so collisions are more frequent and energetic.
• Using the total volume of gas collected as the rate; correction: rate is related to the initial gradient or to 1 ÷ time, not the final total volume.
• Changing temperature or volume of acid while changing concentration; correction: keep all variables except concentration constant.
• Reading the cross through the side of the flask or using different lighting; correction: view from directly above and keep lighting and cross identical.
• Writing a hypothesis as a question or a vague statement such as 'I think it will be faster'; correct this by making a precise prediction that names the factor and the expected effect on rate.
• Changing several variables at once; correct this by keeping all variables except the independent variable constant so the test is fair.
• Failing to link the prediction to collision theory; correct this by explaining that the factor changes either the frequency of collisions or the proportion of particles with energy greater than or equal to the activation energy.
• Misconception: 'Catalysts take part in the reaction and get used up.' Correction: Catalysts can be chemically involved in intermediate steps, but they remain chemically unchanged at the end and can be reused.
• Misconception: 'Increasing temperature only makes particles hit each other more often.' Correction: The dominant reason temperature increases rate is that far more collisions have energy exceeding the activation energy, not just higher collision frequency.
• Misconception: 'A reaction stops because the products block the reactants.' Correction: A reaction stops when one of the reactants is completely used up (the limiting reactant).

1. 1Day 1-2: Master collision theory fundamentals and memorise definitions for activation energy, catalyst, and rate of reaction.
2. 2Day 3-4: Draw and annotate reaction profile diagrams showing catalyzed vs uncatalyzed reactions and temperature effects.
3. 3Day 5-6: Practice plotting rate curves and calculating tangents to determine initial rates and instantaneous rates.
4. 4Day 7: Complete exam-style 4-mark and 6-mark questions on required practicals (disappearing cross or gas syringe methods).

• 📋Graph interpretation questions: Calculating the rate at a specific time using a tangent (gradient = dy/dx).
• 📋Extended response (6-mark) practical questions: Describing investigations into how surface area, concentration, or temperature affects reaction rate.
• 📋Comparative explanation questions: Explaining why changing a specific variable affects the rate using collision theory keywords.