What factors affect the rate of reaction in chemical reactions?

The rate of reaction in chemical reactions can be affected by several factors, including temperature, concentration of reactants, surface area of solid reactants, and the presence of catalysts. In the Cambridge IGCSE Coordinated Science curriculum, students learn that increasing temperature generally increases reaction rates because particles move faster and collide more frequently. Higher concentration means more particles are available to collide, while a larger surface area allows more collisions to occur. Catalysts speed up reactions without being consumed by providing an alternative pathway with a lower activation energy.

How is the rate of reaction measured in experiments?

The rate of reaction can be measured by observing the change in concentration of reactants or products over time. In the Cambridge IGCSE Coordinated Science syllabus, students might measure the rate by monitoring changes in mass, volume of gas produced, or color intensity. For example, if a gas is produced, the rate can be determined by measuring the volume of gas collected over a specific time period. Alternatively, if a reaction involves a color change, a colorimeter can be used to measure the change in color intensity over time.

Why does increasing temperature generally increase the rate of reaction?

Increasing temperature generally increases the rate of reaction because it provides reactant particles with more kinetic energy, causing them to move faster. This results in more frequent and energetic collisions between particles, which increases the likelihood of successful collisions that lead to a reaction. In the context of the Cambridge IGCSE Coordinated Science curriculum, students learn that a higher temperature means more particles have the energy needed to overcome the activation energy barrier, thus speeding up the reaction.

What role do catalysts play in the rate of reaction?

Catalysts play a crucial role in increasing the rate of reaction by providing an alternative reaction pathway with a lower activation energy. This means that more reactant particles have enough energy to react, leading to an increased rate of reaction. Importantly, catalysts are not consumed in the reaction and can be used repeatedly. In the Cambridge IGCSE Coordinated Science curriculum, students explore how catalysts are used in various industrial processes to enhance efficiency and reduce costs.

How does surface area affect the rate of reaction in solids?

Surface area affects the rate of reaction in solids because it determines how many particles are available to collide with reactant molecules. A larger surface area allows more particles to be exposed and available for collisions, thus increasing the reaction rate. For example, powdered solids react faster than larger chunks because the powder has a greater surface area. In the Cambridge IGCSE Coordinated Science curriculum, students learn that increasing the surface area of a solid reactant can significantly speed up the rate of reaction.

Why is "Saying higher temperature only causes more collisions (without mentioning higher energy)" a common mistake in Rate (Speed) of Reaction ?

Why it happens: Students know temperature increases collision frequency but forget the second, more significant effect on collision energy. How to avoid it: Higher temperature → (1) more frequent collisions AND (2) more particles have energy ≥ activation energy. Both effects must be stated for full marks. Source: 0654 Examiner Report 2023

Why is "Saying a catalyst is used up or consumed in the reaction" a common mistake in Rate (Speed) of Reaction ?

Why it happens: Students think catalysts participate and therefore must be depleted. How to avoid it: Catalysts lower activation energy but are regenerated at the end of the reaction; their mass and chemical composition are unchanged.

Why is "Confusing surface area with concentration" a common mistake in Rate (Speed) of Reaction ?

Why it happens: Students write 'more concentrated' when describing why a smaller particle size increases rate. How to avoid it: Surface area (solid reactant) and concentration (dissolved/gas reactant) are different factors. Use 'surface area' for solids; 'concentration' for solutions/gases.

Cambridge IGCSE Coordinated Science (0654)

Rate (Speed) of Reaction

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Short Notes - Rate (Speed) of Reaction

Detailed Study Notes

Detailed notes on Chemical Reactions for Cambridge IGCSE Coordinated Science, covering key concepts, explanations, examples, and exam-focused revision points.

Rates of Reaction — Cambridge IGCSE Coordinated Science 0654

The rate of reaction depends on how frequently and energetically reactant particles collide. Cambridge tests collision theory, the four factors (concentration, temperature, surface area, catalyst), and how to measure and interpret rate data.

At a glance

Rate of reaction = amount of reactant used (or product formed) per unit time.
Collision theory: reaction occurs when particles collide with SUFFICIENT ENERGY (≥ activation energy) and CORRECT ORIENTATION.
Increasing concentration/pressure: more particles per unit volume → more frequent collisions → faster rate.
Increasing temperature: particles have more kinetic energy → more frequent AND more energetic collisions → faster rate.
Increasing surface area: more particles exposed at surface → more frequent collisions → faster rate.
Catalyst: lowers activation energy → more particles have sufficient energy → faster rate (without being consumed).

What you’ll learn

Mapped to the Cambridge IGCSE 0654 syllabus (2025-2027).

C7.1 — Define rate of reaction and state how it is measured.
C7.2 — Explain the effect of concentration, temperature, surface area, and catalysts on rate using collision theory.
C7.3 — Describe and interpret graphs of rate of reaction experiments.
C7.4 — Describe how to investigate rate of reaction experimentally.

Measuring Rate of Reaction

Rate can be measured by following how quickly a reactant is used up or a product is formed over time.

Rate of reaction: the change in the amount of reactant (or product) per unit time.

Rate = change in quantity / time taken

Methods of measuring rate:

Method	Suitable for
Volume of gas collected (gas syringe or displacement of water)	Reactions producing gas (H₂, CO₂, O₂)
Mass loss (top-pan balance)	Reactions producing a gas that escapes (CO₂)
Change in colour/transmission (colorimeter)	Reactions with a colour change
Change in pH	Acid-consuming reactions
Turbidity (cloudiness)	Precipitation reactions (sulfur from thiosulfate + HCl)

Interpreting rate graphs:

Steep gradient = fast rate (early in reaction or under conditions favouring fast rate)
Gradient becomes zero = reaction complete (plateau = no more reactant or product change)
Total amount of product formed = same as long as same initial quantities used (rate affects speed, not yield, unless reactant is limiting)

Rate = change in amount / time. Measure by gas volume, mass loss, colour change.
Steep graph gradient = fast rate. Flat plateau = reaction complete.
Faster reaction = steeper initial gradient + reaches plateau sooner.

Factors Affecting Rate — Collision Theory Explanations

All factors are explained by collision theory: more frequent collisions with sufficient energy = faster rate.

Collision theory: For a reaction to occur, reactant particles must:

Collide
With sufficient energy (≥ activation energy, Ea)
With correct orientation (for complex molecules)

1. Concentration (or pressure for gases):

Higher concentration → more particles per unit volume → more frequent collisions
Rate increases with increasing concentration (at constant temperature)

2. Temperature:

Higher temperature → particles have more kinetic energy → move faster
Collisions are MORE FREQUENT (more movement)
AND more collisions have energy ≥ Ea (more particles above activation energy threshold)
Rule of thumb: 10°C rise roughly DOUBLES rate (for many reactions)

3. Surface area (particle size):

Smaller particles → larger surface area per unit mass → more particles exposed at surface → more frequent collisions with other reactants
Example: marble chips vs marble powder with HCl — powder reacts much faster

4. Catalyst:

Provides an alternative reaction pathway with lower activation energy
At the same temperature, MORE particles have sufficient energy to react
Rate increases without increasing temperature or concentration
Catalyst is NOT consumed — recovered unchanged at end of reaction
Different catalysts used: MnO₂ for H₂O₂ decomposition; Fe for Haber process; V₂O₅ for Contact process

Important: Rate only affects HOW QUICKLY the reaction reaches completion — NOT the final amount of product (yield) — unless a catalyst is so selective it changes the products.

A catalyst provides an alternative pathway with lower activation energy so more particles can react per second.

Higher concentration/pressure, temperature, surface area → more frequent collisions → faster rate.
Temperature also increases collision energy → more particles exceed Ea.
Catalyst: lower Ea → more particles react → faster rate. Not consumed.

Quick recap

Rate = amount / time. Measure by gas volume, mass loss, or colour change.
Collision theory: react when collide with sufficient energy and correct orientation.
Faster rate: higher concentration, higher temperature, larger surface area, catalyst present.
Catalyst: lowers Ea, unchanged at end. Does not change yield.

Memorise this

Verbatim phrases and definitions Cambridge mark schemes credit.

Activation energy (Ea) — minimum energy required for a successful collision.
Collision theory — reactions occur when particles collide with sufficient energy and correct orientation.
Catalyst — substance that increases reaction rate by providing a lower activation energy pathway; not consumed.
Rate of reaction — change in amount of reactant used or product formed per unit time.
Surface area — increased by using smaller particles; more reactant surface exposed to collisions.

How it’s examined

Paper 4: 'Explain, using collision theory, how increasing temperature increases the rate of reaction' (3 marks — more kinetic energy, more frequent collisions AND more collisions with energy ≥ Ea, faster rate). Graph interpretation: 'Compare the rates in experiments A and B from the graph' (steeper gradient = faster rate; same final volume = same amount of product; reach endpoint sooner). Practical planning: 'Describe how you would investigate the effect of concentration on rate' (must identify variables, method of measurement, how to make it fair test). Catalyst explanation questions score well if they include: lower Ea, more particles react.

Sources: Cambridge IGCSE Coordinated Sciences 0654 syllabus 2025-2027 (C7); 0654 Examiner Reports 2022-2024. Last reviewed 2026-05-14.

Take this whole topic with you

Step-by-step worked examples — Rate (Speed) of Reaction

Step-by-step solutions to past-paper-style questions on rate (speed) of reaction , written exactly the way a tutor would explain them at the board.

1Calculating rate of reaction
Core
Question
A student collects $60\,\text{cm}^3$ of gas in $30\,\text{s}$ . Calculate the mean rate of reaction.
Step-by-step solution
1. Step 1
  Apply the rate formula.
  $\text{rate} = \frac{\text{change in quantity}}{\text{time}} = \frac{60}{30} = 2\,\text{cm}^3/\text{s}$
Answer
Rate $= 2\,\text{cm}^3\text{/s}$
Examiner tip
Include units in the answer. Rate can also be expressed as change in mass per second (g/s) or change in concentration per second (mol/dm³/s).
2Effect of concentration — collision theory
Extended
Question
Explain, using collision theory, why increasing the concentration of a reactant in solution increases the rate of reaction.
Step-by-step solution
1. Step 1
  Increasing the concentration means more reactant particles are present per unit volume.
2. Step 2
  The greater number of particles per unit volume leads to more frequent collisions between reactant particles.
3. Step 3
  More collisions per second with energy ≥ activation energy → more successful collisions → higher rate of reaction.
Answer
Higher concentration → more particles per dm³ → more frequent collisions → more successful collisions per second → faster rate.
Examiner tip
Three marking points: more particles per volume / more frequent collisions / more successful collisions (or more per second with energy ≥ Ea).
3Surface area effect
Extended
Question
Powdered calcium carbonate reacts faster with hydrochloric acid than marble chips of the same mass. Explain why, using collision theory.
Step-by-step solution
1. Step 1
  Powdering the calcium carbonate breaks it into much smaller pieces, greatly increasing the total surface area exposed to the acid.
2. Step 2
  More CaCO₃ particles are in contact with (exposed to) the acid molecules at the same time.
3. Step 3
  More surface particles exposed → more frequent collisions between CaCO₃ and HCl → more successful collisions per second → faster rate.
Answer
Powder has much larger surface area than chips; more CaCO₃ particles are exposed to HCl; more frequent collisions; faster rate.
4Investigating temperature and rate (sodium thiosulfate)
Extended
Question
Describe an experiment to investigate the effect of temperature on the rate of reaction between sodium thiosulfate solution and hydrochloric acid.
Step-by-step solution
1. Step 1
  Place a conical flask on a piece of paper marked with a cross. Add a fixed volume (e.g. 50 cm³) of sodium thiosulfate solution.
2. Step 2
  Heat the thiosulfate solution to a chosen temperature (measure with thermometer). Add a fixed volume of HCl (e.g. 5 cm³) and start the stopwatch.
3. Step 3
  Look down through the solution from above. Stop the timer when the cross is no longer visible (sulfur precipitate has formed).
4. Step 4
  Repeat at different temperatures (e.g. 20, 30, 40, 50, 60 °C), keeping volume and concentration of both solutions constant.
5. Step 5
  Calculate rate = 1/time for each temperature. Plot rate vs temperature.
Answer
Measure time for cross to disappear at different temperatures (fixed volume and concentration); rate = 1/time; plot rate vs temperature; rate increases with temperature.
Examiner tip
Control variables: same volumes, same concentrations, same cross, same observer. Rate = 1/t is a common Cambridge mark-scheme answer.

Key Formulae — Rate (Speed) of Reaction

The formulae you need to memorise for rate (speed) of reaction on the Cambridge IGCSE 0654 paper, with every variable defined in plain English and a note on when to use it.

Mean rate of reaction
$\text{rate} = \frac{\text{change in quantity}}{\text{time taken}}$
$\text{change in quantity}$
e.g. volume of gas (cm³ or dm³), mass lost (g), or change in concentration
$\text{time taken}$
in seconds (s)
When to use
Calculating the mean rate from experimental data.
Example
60 cm³ gas in 30 s → rate = 2 cm³/s

Key Definitions and Keywords — Rate (Speed) of Reaction

Definitions to memorise and the exact keywords mark schemes credit for rate (speed) of reaction answers — sharpened from recent examiner reports for the 2026 0654 sitting.

Rate of reaction
Examiner keyword
The change in amount of a reactant or product per unit time; can be measured as change in volume of gas (cm³/s), change in mass (g/s), or change in concentration (mol/dm³/s).
Collision theory
Examiner keyword
A model that states that chemical reactions occur only when particles collide with sufficient energy (at least equal to the activation energy) and with the correct orientation.
Activation energy
Examiner keyword
The minimum energy that colliding particles must have for a reaction to occur. Only collisions with energy ≥ activation energy are successful.
Successful collision
A collision between particles that results in a chemical reaction; requires energy ≥ activation energy and correct orientation.
Surface area
Examiner keyword
The total area of solid material exposed to a reactant. Increasing surface area (e.g. by powdering) increases the rate of reaction by exposing more particles.

Common Mistakes and Misconceptions — Rate (Speed) of Reaction

The traps other students keep falling into on rate (speed) of reaction questions — taken from recent Cambridge IGCSE 0654 examiner reports and mark schemes — and how to avoid them.

✕Saying higher temperature only causes more collisions (without mentioning higher energy)
0654 Examiner Report 2023
Why it happens
Students know temperature increases collision frequency but forget the second, more significant effect on collision energy.
How to avoid it
Higher temperature → (1) more frequent collisions AND (2) more particles have energy ≥ activation energy. Both effects must be stated for full marks.
✕Saying a catalyst is used up or consumed in the reaction
Why it happens
Students think catalysts participate and therefore must be depleted.
How to avoid it
Catalysts lower activation energy but are regenerated at the end of the reaction; their mass and chemical composition are unchanged.
✕Confusing surface area with concentration
Why it happens
Students write 'more concentrated' when describing why a smaller particle size increases rate.
How to avoid it
Surface area (solid reactant) and concentration (dissolved/gas reactant) are different factors. Use 'surface area' for solids; 'concentration' for solutions/gases.