# Rates of Reaction

> Edexcel International GCSE Chemistry · Edexcel IGCSE Chem 4CH1 (2017)
> Source: https://www.owlsprep.com/study/edexcel-igcse-chemistry-s3-rates-of-reaction/

This guide covers all Edexcel IGCSE 4CH1 content for rates of reaction, including core rate factors, collision theory explanations, required practicals, and higher-tier reaction profile diagrams. You will learn to answer both practical and theoretical exam questions on the topic.

**Prerequisites:** [Basic particle theory](https://www.owlsprep.com/study/edexcel-igcse-chemistry-s1-particle-model/); [Enthalpy changes (for higher tier reaction profiles)](https://www.owlsprep.com/study/edexcel-igcse-chemistry-s3-enthalpy-changes/)

## Learning objectives

- Describe experiments to investigate effects of surface area, concentration, temperature and catalysts on reaction rate
- Explain rate factor effects using collision theory (collision frequency and energy)
- Define catalysts and explain their mode of action via lower activation energy pathways
- Draw and label higher-tier reaction profile diagrams showing ΔH and activation energy
- Carry out and analyse data from the two required rate of reaction practicals

## Measuring Reaction Rate

**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.

**Worked example:** 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
2. $$Rate = 45/cm^3 \div 30//s$$
3. Final answer: 1.5 cm³/s

*Calculator:* allowed

## Core Factors Affecting Reaction Rate

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)

**Worked example:** A student uses large marble chips with dilute HCl, then repeats with the same mass of powdered marble. Describe the difference in observed rate.

1. Observation: Powdered marble reacts much faster, so bubbles of CO₂ are produced more rapidly, and the reaction finishes sooner.
2. Explanation: Powdered marble has a larger surface area exposed to the HCl particles, so more frequent collisions occur between reactant particles.

## Collision Theory Explanations

**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.

> **info**
>
> Remember two key collision properties: frequency (how often collisions happen) and energy (proportion of collisions that have enough energy to react).

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.

**Worked example:** Explain why increasing the pressure of a gaseous reaction increases the rate, using collision theory.

1. Increased pressure pushes gas particles closer together, so the concentration of gas particles increases.
2. This increases the frequency of collisions between reactant gas particles, leading to more successful collisions per second, so rate increases.

## Catalysts & Reaction Profiles (Core + Higher Tier)

**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.

*Notation:* Common catalysts include MnO₂ for H₂O₂ decomposition.

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.

> **Higher Tier (Chemistry Only 3.14C)**
>
> You must be able to draw and label reaction profile diagrams showing activation energy (Eₐ) for both catalysed and uncatalysed reactions, plus ΔH (enthalpy change) for the reaction.

**Worked example:** 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).
2. 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.
3. Label the vertical difference between reactant and product energy levels as ΔH (negative for exothermic reactions).

## Required Practicals for Rates of Reaction

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.

**Worked example:** A student tests the effect of MnO₂, CuO and ZnO on H₂O₂ decomposition. State the independent, dependent and control variables for this experiment.

1. Independent variable: Type of solid catalyst used.
2. Dependent variable: Volume of oxygen produced per unit time (rate of reaction).
3. Control variables: Concentration and volume of hydrogen peroxide, mass of catalyst used, temperature of the reaction mixture.

## Common pitfalls

- **Wrong:** Stating catalysts are used up during reactions
  - Why it fails: Catalysts only provide an alternative lower Eₐ pathway, they do not react permanently with reactants
  - Correct: State catalysts remain chemically unchanged and can be reused at the end of the reaction
- **Wrong:** Explaining temperature effect only by increased collision frequency
  - Why it fails: Higher temperature increases both collision frequency and proportion of collisions with energy ≥ Eₐ, the energy effect is larger
  - Correct: Mention both frequency and energy factors when explaining temperature effect on rate
- **Wrong:** Confusing surface area with total mass of solid
  - Why it fails: Two samples of marble with the same mass can have different surface areas if one is powdered and one is lumps
  - Correct: State that larger surface area comes from smaller pieces of the same mass of solid
- **Wrong:** Drawing catalysed Eₐ higher than uncatalysed Eₐ on reaction profiles
  - Why it fails: Catalysts lower activation energy, so the catalysed peak is lower
  - Correct: Always draw the catalysed reaction peak below the uncatalysed peak on reaction profiles
- **Wrong:** Claiming catalysts change the enthalpy change (ΔH) of a reaction
  - Why it fails: Catalysts only change activation energy, the energy of reactants and products stays the same, so ΔH is unchanged
  - Correct: Note that ΔH is identical for catalysed and uncatalysed versions of the same reaction

## 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 |

## 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.

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