Collision Theory & Activation Energy
Section 2 told you what speeds reactions up. Collision theory is the single idea that explains why — and it is worth learning properly, because almost every “explain” question in this topic is marked against it.
Chemical reactions can occur only when reacting particles collide with each other, and with sufficient energy. The minimum amount of energy that particles must have to react is called the activation energy.
So a reaction’s rate is set by two things: how often particles collide (the frequency of collisions), and what fraction of those collisions carry enough energy to count (at least the activation energy). Speed either one up and the reaction gets faster.
Explaining each factor
| Factor increased | What changes for the particles | Effect |
|---|---|---|
| Concentration | More particles in the same volume | Collisions more frequent → rate increases |
| Pressure (gases) | Same particles squeezed into a smaller volume | Collisions more frequent → rate increases |
| Surface area (solids) | Smaller pieces → larger surface area : volume ratio, more particles exposed | Collisions more frequent → rate increases |
| Temperature | Particles move faster and carry more energy | Collisions more frequent and more energetic (a bigger fraction beat the activation energy) → rate increases a lot |
“More collisions” on its own is usually not enough for the mark. Left to run for long enough, even a slow reaction clocks up more collisions. What matters for rate is collisions per second — so write “collisions are more frequent” (or “more collisions per second”). The same discipline applies to energy: it’s not that particles “hit harder”, it’s that a greater proportion of collisions have energy ≥ the activation energy — these are often called successful collisions. Put the two together and you have the phrase that reliably earns the mark: rate depends on the frequency of successful collisions — how many collisions with energy ≥ the activation energy happen per second. A bare “number of successful collisions” still names no time period, so on its own it can’t describe a rate.
Concentration, pressure and surface area work through frequency only. Temperature is the only factor that works through both frequency and energy — and the energy effect dominates, which is why a small temperature rise can have a dramatic effect on rate. An “explain the effect of temperature” answer that only mentions faster-moving particles colliding more often is doing half the job: the second mark almost always needs “more particles collide with energy greater than or equal to the activation energy, so more collisions are successful”.
Surface area — smaller pieces, bigger ratio
Cutting a solid into smaller pieces doesn’t change how much of it there is — it changes how much of it is on the outside. You need to describe this in terms of the surface area to volume ratio: smaller pieces have a larger surface area to volume ratio, so more particles are exposed for collisions and the rate increases. The extreme case is a powder, which reacts far faster than the same mass in lumps.
Cut a 2 cm cube into eight 1 cm cubes and the volume is unchanged — but every cut exposes two new faces, doubling the surface area and the SA:V ratio. More exposed particles → more frequent collisions → faster reaction.
Proportionality — the simple maths of collisions
For factors that work through frequency alone, the maths is direct: double the concentration and there are twice as many particles in the same volume, so collisions are twice as frequent and the rate roughly doubles — exactly the straight-line-through-the-origin pattern from the required practical in section 2. Temperature does not follow this simple proportionality: because it also boosts the fraction of successful collisions, a mere 10 °C rise can roughly double the rate of many reactions.
rate ∝ concentration and rate ∝ 1time
The symbol ∝ means “is directly proportional to” — double one side and the other doubles too. It’s the compact way to write the straight-line-through-the-origin patterns from the required practical in section 2.
Watch all of this happen below. Choose a scenario — a solid in a solution, or two gases — then flip the factors that apply to it and watch how each one changes the rate of successful collisions.
🔬 Collision theory — explore the five factors
Pick a scenario, then explore the factors. Particles move and collide; a collision with at least the activation energy reacts — a ✦ spark and a purple product. Each factor only appears for the state of matter it actually applies to, so the science stays true. Watch the rate of successful collisions respond.
All five factors are at their baseline. Increase any one and watch the rate of successful collisions respond — and notice why it changes.
🧪 Exam-style questions
Which factor increases the rate of reaction by increasing both the frequency and the energy of collisions? Tick (✓) one box.
Powdered calcium carbonate reacts faster with acid than the same mass of marble chips because the powder has… Tick (✓) one box.
The concentration of an acid is doubled (everything else unchanged), and the initial rate of its reaction with magnesium roughly doubles. Why? Tick (✓) one box.
Increasing the pressure of a reacting gas increases the rate of reaction. Which statement explains why? Tick (✓) one box.
Explain, using collision theory, why increasing the temperature increases the rate of a reaction.
Show a model answer
- At a higher temperature the particles move faster… 1 mark
- …so collisions are more frequent (more collisions per second). 1 mark
- The particles also have more energy, so a greater proportion of collisions have energy ≥ the activation energy… 1 mark
- …so there are more successful collisions per second, and the rate increases. 1 mark
Examiner note — this is the full causal chain: faster particles → more frequent collisions; more energy → bigger fraction over Eₐ. Answers that say only “more collisions”, or that describe (“the reaction speeds up”) without explaining, stall at half marks.
Catalysts
The fifth way to speed up a reaction is the cleverest: add a substance that makes the reaction faster without being used up itself.
A catalyst changes the rate of a chemical reaction but is not used up during the reaction. Because it isn’t a reactant, it is not included in the chemical equation. Different reactions need different catalysts — for example, enzymes act as catalysts in biological systems.
How do they do it? A catalyst provides a different pathway for the reaction — one with a lower activation energy. The particles themselves are no faster and no more energetic than before; the bar has simply been lowered, so a greater proportion of the collisions that were already happening now have enough energy to react. More successful collisions per second — a faster rate. You can picture it on a reaction profile, where the catalysed pathway climbs a noticeably lower hump:
- It does not give the particles more energy — it lowers the activation energy by providing a different pathway. (Raising the temperature does the reverse: it leaves the barrier where it is but gives the particles more energy. A catalyst lowers the bar; heating lifts the particles over it.)
- It does not change the overall energy change of the reaction, and it does not make more product — only faster product.
- It is not used up — but avoid saying it “doesn’t take part”. It does take part (that’s how it provides the new pathway); it is regenerated by the end, which is why it never appears in the equation.
Spotting a catalyst in exam data
Questions rarely say “X is a catalyst” — they make you identify it from its fingerprints. Look for a substance that:
- speeds the reaction up (gas produced faster, time shorter, initial gradient steeper);
- has the same mass at the end as at the start (not used up) — often it can be filtered out and reused;
- does not appear in the chemical equation for the reaction.
A classic demonstration is the decomposition of hydrogen peroxide, which is very slow on its own but fizzes vigorously with a little manganese(IV) oxide powder:
2H2O2(aq) → 2H2O(l) + O2(g)
(MnO2 catalyst — written above the arrow, never in the equation.)
You don’t need to memorise lists of catalysts, but one pattern helps you identify them in data questions: transition metals and their compounds are very often the catalyst (manganese(IV) oxide above; iron, nickel, platinum, copper and their ions elsewhere) — acting as catalysts is one of the characteristic properties of the transition elements. So when a question shows several substances and asks which is most likely to be the catalyst, a transition-metal compound or ion is the safe bet over a main-group salt. (A classic class investigation compares different metal salts added to hydrogen peroxide.)
Catalysts matter to industry for a very practical reason: a faster reaction at a lower temperature means less energy bought from burning fuel — lower costs and less environmental impact. Enzymes do the same job in living things (and in industry too — they are the catalysts in fermentation).
🧪 Exam-style questions
A catalyst increases the rate of a reaction by… Tick (✓) one box.
0.5 g of manganese(IV) oxide is added to hydrogen peroxide solution. The mixture fizzes rapidly. When the reaction has finished, the manganese(IV) oxide is filtered off, dried and reweighed. What mass is found? Tick (✓) one box.
Which of these does a catalyst change? Tick (✓) one box.
A student adds a small amount of each of these substances to separate samples of hydrogen peroxide solution. Which is most likely to act as a catalyst for its decomposition? Tick (✓) one box.
On a reaction profile, what does adding a catalyst change, and what stays the same?
Show a model answer
- Changes: the height of the “hump” — the catalysed pathway has a lower activation energy. 1 mark
- Stays the same: the energy levels of the reactants and products — so the overall energy change is identical. 1 mark