What you'll learn
Rates of reaction tell us how fast chemical reactions happen and how we can control them. In this guide you will learn how to measure reaction rate, the factors that affect it (concentration, pressure, surface area, temperature and catalysts), how collision theory explains each factor, how to interpret rate graphs and calculate rates from them, and how catalysts work. These ideas are tested in calculations, graphs and required practicals.
Key terms and definitions
Rate of reaction — how quickly reactants are used up or products are formed.
Collision theory — the idea that reactions happen when particles collide with enough energy and the correct orientation.
Activation energy — the minimum energy particles must have to react when they collide.
Catalyst — a substance that speeds up a reaction without being used up, by lowering the activation energy.
Surface area — the total exposed area of a solid reactant.
Concentration — the amount of a substance dissolved in a given volume.
Core concepts
Measuring rate
The rate of reaction is the change in amount of reactant or product over time. It can be measured by:
- Measuring the volume of gas produced over time (gas syringe).
- Measuring the loss in mass as gas escapes (balance).
- Measuring the time for a solution to turn cloudy (e.g. the sodium thiosulfate "disappearing cross" experiment).
Rate = amount of reactant used or product formed ÷ time. Common units are cm³/s or g/s.
Collision theory
Reactions occur when particles collide with enough energy (at least the activation energy) and the correct orientation. Anything that increases the frequency of successful collisions, or the energy of collisions, increases the rate.
Factors affecting rate
- Concentration / pressure: higher concentration (or, for gases, higher pressure) means more particles in a given volume, so more frequent collisions and a faster rate.
- Surface area: breaking a solid into smaller pieces (or powder) increases the surface area, so more particles are exposed and collisions are more frequent, increasing the rate.
- Temperature: higher temperature gives particles more energy, so they move faster (more frequent collisions) and a greater proportion have at least the activation energy. Temperature has a large effect.
- Catalyst: a catalyst provides an alternative pathway with a lower activation energy, so more collisions are successful, speeding up the reaction without being used up.
Interpreting rate graphs
On a graph of product formed against time, a steeper line means a faster rate. The line levels off when the reaction finishes (a reactant is used up). The reaction is fastest at the start (most reactant present) and slows as reactants are used up. To find the rate at a particular time, draw a tangent to the curve and calculate its gradient (change in y ÷ change in x). The overall rate can be found from total product ÷ total time.
Catalysts in detail
A catalyst speeds up a reaction by lowering the activation energy, but is not used up and does not appear in the overall equation. Enzymes are biological catalysts. Catalysts are valuable in industry because they allow reactions to run faster (or at lower temperatures), saving energy and money. Different reactions need different catalysts.
Worked examples
Example 1: Effect of surface area
Explain why powdered calcium carbonate reacts faster with acid than large lumps.
Powder has a larger surface area, so more particles are exposed to the acid. This means more frequent collisions per second, so the rate is faster.
Example 2: Calculating rate from a graph
40 cm³ of gas is produced in 20 seconds. What is the mean rate?
Rate = volume ÷ time = 40 ÷ 20 = 2 cm³/s.
Example 3: Effect of temperature
Why does increasing temperature speed up a reaction?
Particles gain more energy and move faster, so they collide more often, and a greater proportion of collisions have at least the activation energy — so more collisions are successful.
Common mistakes and how to avoid them
Saying higher temperature only makes particles collide more often. It also means more collisions exceed the activation energy — mention both.
Confusing concentration and surface area. Concentration is about particles per volume in solution; surface area is about exposed solid.
Forgetting catalysts lower activation energy. That is how they speed up the reaction.
Misreading rate graphs. Steeper = faster; levelling off = finished. Use a tangent for the rate at a point.
Saying a catalyst is used up. It is not consumed and is unchanged at the end.
Exam technique for Rate of Reaction
Use collision theory to explain every factor — frequency and/or energy of successful collisions.
Describe measurement methods (gas volume, mass loss, cloudiness) appropriately.
Read and calculate from rate graphs, using gradients and tangents.
Explain catalysts as lowering activation energy without being used up.
State both effects of temperature — more frequent and more energetic collisions.
Quick revision summary
The rate of reaction is how fast reactants are used or products formed, measured by gas volume, mass loss, or time to turn cloudy. Collision theory says reactions need particles to collide with at least the activation energy in the correct orientation, so rate rises when successful collisions become more frequent or more energetic. Higher concentration/pressure packs more particles into a volume → more frequent collisions. Larger surface area (smaller pieces/powder) exposes more particles → more frequent collisions. Higher temperature gives particles more energy → faster movement (more frequent collisions) and more collisions exceeding the activation energy. A catalyst lowers the activation energy via an alternative pathway, speeding the reaction without being used up (enzymes are biological catalysts). On rate graphs, steeper = faster, the curve levels off when finished, and the rate is fastest at the start; find the rate at a point from the gradient of a tangent. Always explain factors through collision theory, give both temperature effects, and read graphs carefully.