The big idea: The rate of reaction tells you how fast reactants are turned into products.
As a reaction proceeds, reactants are used up (their concentration falls) and products are made (their concentration rises). The rate is the change in concentration per unit time.
Reactions start fast and slow down as the reactants run out — until they stop.
Define every term: - Rate of reaction — the change in concentration of a reactant or product per unit time. - Concentration — the amount of substance dissolved per unit volume, in mol dm⁻³. - Instantaneous rate — the rate at one moment (the gradient of the curve at that point). - Average rate — the rate over an interval (total change ÷ total time).
If you plot the concentration of a product against time, the rate is the gradient (steepness) of the graph. The curve is steepest at the start and flattens to zero as the reaction finishes.
Animated graph
Watch the graph build step by step in study mode.
- change in concentration of a reactant (used up) or product (formed), in mol dm⁻³
- change in time over which it is measured, in s
- rate of reaction, in mol dm⁻³ s⁻¹ (the gradient of a concentration–time graph)
Units of rate: Rate = a concentration ÷ a time, so its units are mol dm⁻³ s⁻¹.
If you follow a gas by its volume instead, the rate comes out as cm³ s⁻¹; if you follow it by mass lost, as g s⁻¹. The units always follow the quantity you measured ÷ time.
Average vs instantaneous: - Average rate over an interval = (change in the quantity) ÷ (time taken) — the slope of the straight chord joining the two points. - Instantaneous rate at one moment = the slope of the tangent drawn at that point. The initial rate (the tangent at t = 0) is the steepest and the most useful for comparing experiments.
Worked example — average rate from a graph
In a reaction, the concentration of a product rises from 0 to 0.36 mol dm⁻³ in the first 30 s. Calculate the average rate over this interval.
Solution
- Formula first — average rate is the change in concentration over the time taken:
- Substitute the values:
- Work it out — keep the unit:
Final answer
Average rate = 0.012 mol dm⁻³ s⁻¹.
| What changes | How to follow it | Example reaction |
|---|---|---|
| A gas is produced | collect the gas and measure its volume vs time, or measure the mass lost vs time | carbonate + acid → CO2 |
| The mixture changes colour | measure light absorbed (colorimeter) vs time | reactions of coloured ions |
| A precipitate forms | time how long until a cross disappears (turbidity) | thiosulfate + acid → sulfur |
| pH changes | track the pH with a probe vs time | acid–base / hydrolysis |
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Particles must collide: Collision theory says that for particles to react they must collide with each other. But not every collision leads to a reaction — only effective (successful) collisions do.
A collision is effective only if both of these are true:
- the colliding particles have enough energy — at least the activation energy, Eₐ, and - they collide in the correct orientation (the right way round).
Activation energy: The activation energy (Eₐ) is the minimum energy that colliding particles must have for a reaction to occur. It is the energy needed to start breaking the bonds in the reactants. Collisions with less energy than Eₐ simply bounce apart unchanged.
Only molecules with energy to the RIGHT of Eₐ have enough energy to react on collision — most collisions (to the left) just bounce off.
Interactive diagram
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| Requirement | What it means | If it is not met |
|---|---|---|
| Sufficient energy | the combined collision energy must be ≥ the activation energy, Eₐ | the molecules just bounce apart unchanged |
| Correct orientation | the molecules must collide the right way round so the reacting parts meet | no reaction — they collide but miss the reactive sites |
Why orientation matters: Even an energetic collision fails if the molecules are pointing the wrong way. For two molecules to react, the reactive parts must line up and meet. A 'glancing' collision in the wrong orientation does nothing — the molecules just separate again.
How this is tested: R2.2 appears across all three papers.
- Paper 1A (MCQ): pick the correct units of rate, or read an average rate off a volume–time graph. - Paper 1B (data): draw a tangent at t = 0 to find the initial rate, or compute an average rate over an interval with units. - Paper 2: 'outline/explain why not every collision between reactant molecules leads to a reaction' — a short collision-theory answer.
The marks are lost by quoting rate without units, and by giving only one of the two collision requirements.
Score it cleanly: For rate: state the change ÷ time and always carry the unit. For collision theory: give both requirements — energy ≥ Eₐ AND correct orientation — to bank both marks.
IB-style question — effective collisions (a)
(a) Outline, using collision theory, why not every collision between reactant molecules results in a reaction. [2]
How to score the marks
- Mark 1 — energy. The colliding particles must have at least the activation energy (Eₐ); collisions with less energy are not successful and the particles simply bounce apart.
- Mark 2 — orientation. The particles must also collide in the correct orientation (the right way round) so that the reactive parts meet. A collision that meets only one of these fails.
Final answer
Only collisions with energy ≥ Eₐ AND the correct orientation are effective; collisions failing either requirement do not react.
IB-style question — average rate from data (b)
(b) A piece of marble reacts with acid. The mass of the flask falls from 84.50 g to 83.30 g in the first 40 s as carbon dioxide escapes. Determine the average rate of mass loss over this interval, including units. [2]
How to score the marks
- Mark 1 — find the change. Mass lost = 84.50 − 83.30 = 1.20 g of CO2 over the first 40 s.
- Mark 2 — divide by the time and give units. Average rate = mass lost ÷ time:
Final answer
Average rate = 0.030 g s⁻¹ (the units follow mass ÷ time).