Rate of reaction
Collision theory and the effects of concentration, temperature, surface area, pressure and catalysts.
Review these first
Learning objectives
What you need to be able to do
Teacher-mapped phrasing — check against the official Cambridge syllabus for exact wording.
- 6.2.1Describe how to measure rate of reaction by gas volume or mass loss.
- 6.2.2Explain the effects of concentration, pressure, surface area, temperature and catalysts using collision theory.
- 6.2.3Interpret rate graphs, including the meaning of the gradient and the point where the reaction finishes.Supplement
7 minute read
Collision theory and rates
For a reaction to happen, particles must collide, with enough energy (the activation energy) and in the right orientation.
Anything that increases either the frequency of collisions or the proportion of successful collisions will increase the rate.
The factors
Concentration ↑ → more particles in the same volume → more frequent collisions → faster.
Pressure ↑ (gases) → particles pushed closer together → more frequent collisions → faster.
Surface area ↑ (smaller pieces) → more particles exposed on the surface → more frequent collisions → faster.
Temperature ↑ → this one has two effects, and good answers give both:
- Particles move faster, so they collide more frequently.
- More particles have energy greater than or equal to the activation energy, so a greater proportion of collisions are successful.
The second effect is by far the larger one, and it is the one most students omit.
Catalyst → provides an alternative reaction pathway with a lower activation energy, so a greater proportion of collisions succeed. The catalyst itself is not used up.
Reading rate graphs
Plot volume of gas (or mass lost) against time:
- The steeper the curve, the faster the reaction.
- The curve is steepest at the start, when concentration is highest.
- It levels off when a reactant is used up. The height of the plateau shows how much product was made in total.
Two experiments with the same amount of reactant but different conditions will level off at the same height — only the steepness differs.
Think of it like this
Collisions are like a busy corridor. More people (concentration) or faster walking (temperature) both mean more bumps — but only bumps hard enough to knock a book out of someone's hand count as a reaction.
Worked examples
Method, step by step
Explain why powdered calcium carbonate reacts faster with acid than the same mass of large lumps.
- 1Powder has a much larger total surface area than lumps of the same mass.
- 2More calcium carbonate particles are exposed at the surface for acid particles to collide with.
- 3The frequency of successful collisions is therefore higher.
The greater surface area gives more frequent collisions, so the rate is higher — though the total volume of gas produced is the same.
Common misconceptions
- Saying a catalyst "lowers the energy of the reaction". It lowers the activation energy by providing a different pathway; ΔH is unchanged.
- Explaining a temperature increase only by "particles move faster". The larger effect is that more particles exceed the activation energy.
- Thinking a faster reaction produces more product. It produces the same amount, more quickly.
In the exam
- Rate explanations need two steps: what happens to the particles, then what happens to the collisions. "More collisions" alone is usually 1 mark of 2.
- Say "more frequent collisions", not "more collisions" — over the whole reaction the total number is the same.
- On a rate graph, comparing plateau heights is how you show that the amount of reactant was unchanged.