Reaction Rates
Investigate the factors that control how fast chemical reactions proceed — from temperature and concentration to surface area and catalysts.
Collision Theory
For a chemical reaction to occur, reacting particles (atoms, molecules or ions) must collide with each other. But not every collision leads to a reaction — the collision must have:
Sufficient energy
The collision must provide at least the activation energy (Ea) — the minimum energy needed to break existing bonds and start the reaction.
Correct orientation
Particles must collide in the right orientation so that reactive parts of each molecule meet and bonds can form.
Rate of reaction is a measure of how quickly reactants are used up or products are formed. It can be measured by tracking changes in mass, volume of gas produced, colour change, or pH over time.
Rate = change in quantity ÷ time taken
Factors Affecting Reaction Rate
1. Temperature
Increasing temperature gives particles more kinetic energy. This means they move faster, collide more frequently, and a greater proportion of collisions have sufficient energy to exceed the activation energy.
Australian example: Food spoils much faster left on the bench in a hot Australian summer than it does in the fridge, because higher temperature speeds up bacterial chemical reactions.
2. Concentration
A higher concentration of reactants means more particles are present in the same volume. This increases the frequency of collisions, so the reaction proceeds faster.
Example: Concentrated hydrochloric acid reacts much more vigorously with marble chips than dilute hydrochloric acid.
3. Surface Area
For solid reactants, only particles on the surface can collide with other reactants. Breaking a solid into smaller pieces increases the total surface area exposed, so more collisions can occur per second.
Smaller pieces = greater surface area = faster reaction
4. Catalysts
A catalyst speeds up a chemical reaction by providing an alternative reaction pathway with a lower activation energy. Crucially, the catalyst is not consumed — it is regenerated at the end of the reaction.
Example: Manganese dioxide (MnO2) catalyses the decomposition of hydrogen peroxide into water and oxygen gas. Enzymes in your body act as biological catalysts (e.g., amylase breaks down starch in your saliva).
Note: A catalyst lowers Ea but does not change the overall energy released or absorbed by the reaction.
Measuring Reaction Rate
The rate of a reaction can be measured by monitoring a change in a measurable quantity over time. Common methods include:
Volume of gas
Collect gas produced in a measuring cylinder. The faster the gas accumulates, the faster the reaction. Used for reactions producing CO2 or O2.
Change in mass
Place the reaction vessel on a balance. If a gas is produced and escapes, the mass decreases. Rate = mass lost per second.
Colour change (colorimetry)
Measure how quickly a coloured solution becomes clear (or vice versa). The “disappearing cross” experiment uses this method with sodium thiosulfate and HCl.
Key Vocabulary
| Term | Definition |
|---|---|
| Activation energy | The minimum amount of energy that colliding particles must possess for a reaction to occur. |
| Catalyst | A substance that increases the rate of a reaction by providing a lower-energy alternative pathway, without being consumed. |
| Concentration | The amount of a substance dissolved in a given volume of solution, usually expressed in mol/L. |
| Collision theory | The theory that chemical reactions occur when reactant particles collide with sufficient energy and correct orientation. |
Worked Examples
Calculating average rate of reaction from collected gas data.
Given: 48 mL of CO2 gas was collected in 4 minutes when marble chips reacted with HCl.
Step 1: Convert time to seconds: 4 min × 60 = 240 s.
Step 2: Apply the rate formula: Rate = 48 mL ÷ 240 s.
Answer: Rate = 0.2 mL/s of CO2 produced.
Explaining why powdered iron rusts faster than iron filings.
Concept: Surface area and reaction rate.
Explanation: Powdered iron has a much greater surface area than iron filings. More iron particles are exposed to oxygen and water vapour in the air, so more collisions occur per second between iron atoms and O2/H2O molecules.
Answer: Powdered iron rusts faster because the greater surface area increases the collision frequency between iron and the reactants (oxygen and water), leading to a faster rate of oxidation.
Identifying the independent variable in a reaction rate experiment.
Scenario: A student investigates how temperature affects the time for a sodium thiosulfate and HCl reaction. She uses 25 mL of sodium thiosulfate and 5 mL of HCl at temperatures of 20°C, 30°C, 40°C, 50°C and 60°C.
Independent variable: Temperature (°C) — what is being changed deliberately.
Dependent variable: Time for the cross to disappear (s) — what is being measured.
Controlled variables: Volume and concentration of both solutions, size of cross, same observer — kept the same to ensure a fair test.
Knowledge Check
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Key Concepts Summary
- •Reactions occur when particles collide with sufficient energy (above activation energy) and correct orientation.
- •Temperature: higher temperature → more energetic particles → more successful collisions → faster rate.
- •Concentration: more particles per volume → more frequent collisions → faster rate.
- •Surface area: smaller pieces expose more surface → more collisions → faster rate.
- •Catalysts lower activation energy and are not consumed, speeding up reactions without being used up.