Kinetics — collision theory, the Maxwell–Boltzmann distribution and catalysts.

HookBritain's most stolen car part is a kinetics lesson

ModelCollision theory — almost every collision fails

Collision theory says a reaction can only happen when reactant particles collide, and only when the collision carries at least a minimum energy called the activation energy, \(E_a\) — the energy needed to break the necessary bonds and begin forming products. Collisions below that threshold are elastic non-events: the molecules bounce apart unchanged.

The numbers show how brutal the filter is. A molecule in a gas at room conditions undergoes around a billion (10⁹) collisions every second — yet a sealed mixture of hydrogen and oxygen can sit unreacted for years. For a typical activation energy of about 50 kJ mol⁻¹ at 298 K, only around two collisions in every billion carry enough energy to react. Reaction rates are set not by how often molecules meet but by how often they meet hard enough.

That framing turns rate into a counting problem: anything that increases the number of successful collisions per second increases the rate. There are exactly two levers — make collisions more frequent (concentration, pressure, surface area) or make a larger fraction of them energetic enough to clear \(E_a\) (temperature, or lower the bar itself with a catalyst). Every explain-question in 3.1.5 is answered by naming which lever is being pulled.

ModelThe Maxwell–Boltzmann distribution — the whole section in one curve

Molecules in a gas do not share one speed; they carry a spread of kinetic energies, redistributed at every collision. The Maxwell–Boltzmann distribution plots number of molecules against energy, and its shape is examined line by line. The curve starts at the origin — no molecules have zero energy. It rises to a single peak at the most probable energy, the energy carried by more molecules than any other. The mean energy sits slightly to the right of the peak, dragged over by the long high-energy tail. And that tail approaches the x-axis but never touches it — there is no upper limit to the energy a lucky molecule can accumulate. The total area under the curve equals the total number of molecules.

Now mark the activation energy, far out along the axis. The shaded area beyond \(E_a\) is the population of molecules that can react on their next collision — for our 50 kJ mol⁻¹ reaction at room temperature, a sliver of about two-billionths of the total area. Rate questions are area questions: whatever grows that shaded region speeds the reaction. Temperature reshapes the curve to push molecules into the tail; a catalyst leaves the curve alone and moves the \(E_a\) line instead. Learn to read the diagram both ways and the rest of the section is bookkeeping.

DataTemperature — why +10 K roughly doubles the rate

Heating a gas from 300 K to 310 K makes molecules move faster, so they do collide more often — but only about 2% more often, because collision frequency scales with the square root of temperature. If frequency were the whole story, a 10 K rise would barely register. Measured rates instead roughly double. The missing factor is the shape-change of the Maxwell–Boltzmann curve: at the higher temperature the peak shifts right and drops lower, the curve flattens and broadens, and the area stays fixed (same number of molecules) — but the high-energy tail fattens enormously. Because the tail is exponentially thin, a small shift of the whole distribution multiplies the population beyond \(E_a\).

So the creditworthy explanation is always: at higher temperature a much greater proportion of molecules have energy greater than or equal to the activation energy, so there are many more successful collisions per second — with the small rise in collision frequency mentioned as the minor effect, if at all. Answers that lead with 'particles collide more often' invert the physics and cap the marks.

MechanismConcentration and pressure — more collisions, same energy

Raising the concentration of a solution crowds more reactant particles into each cubic decimetre, so collisions between them come more frequently and the rate rises. For gases, raising the pressure does the identical job — compressing the same molecules into a smaller volume is raising their concentration. Nothing about the energy of any individual collision changes: the Maxwell–Boltzmann distribution depends only on temperature, so the curve does not move. The shaded fraction beyond \(E_a\) is unchanged; there are simply more collisions per second for that fraction to act on.

This is also why almost every reaction decelerates as it proceeds: reactants are consumed, their concentrations fall, collisions become rarer, and the rate decays — the reason a concentration–time graph starts steep and flattens towards the end. When an exam question asks you to compare two moments during the same run, concentration is the lever that moved.

Keep the two levers separate in written answers. Concentration and pressure change collision frequency only. Temperature changes frequency slightly and the energy distribution massively. One clean sentence on which lever operates — and which does not — is what distinguishes a level-3 explanation from a muddle.

MechanismCatalysts — a different road over the mountain

A catalyst increases the rate of a reaction without being used up, by providing an alternative reaction pathway with a lower activation energy. That exact clause is the one AQA credits — the catalyst does not 'lower the activation energy' of the existing route, and it certainly does not give molecules more energy. The original path still exists; the catalyst opens a cheaper one, often by binding a reactant to a surface and weakening its bonds. On the Maxwell–Boltzmann diagram the curve is untouched, but the \(E_a\) marker slides left — and because the tail is exponentially thin, even a modest slide multiplies the qualifying population many times over.

Your converter runs reactions like \(2\text{CO} + 2\text{NO} \rightarrow 2\text{CO}_2 + \text{N}_2\) on a platinum–rhodium surface: gases adsorb onto the metal, bonds loosen, products form and desorb, and the metal emerges unchanged — which is why three grams of coating lasts the life of the car, and why it is worth stealing intact. Industry is built on the same trick: iron in ammonia synthesis, vanadium(V) oxide in sulfuric acid manufacture, nickel in the hydrogenation of vegetable oils. Around 90% of manufactured chemicals pass over a catalyst somewhere in production, because a lower-energy pathway means lower temperatures, smaller fuel bills and less CO₂ per tonne of product — the economic and environmental argument examiners expect you to make in a sentence.

CaseRequired practical 3 — the disappearing cross, done properly

RP3 investigates how temperature changes rate, using the reaction \(\text{Na}_2\text{S}_2\text{O}_3\text{(aq)} + 2\text{HCl(aq)} \rightarrow 2\text{NaCl(aq)} + \text{S(s)} + \text{SO}_2\text{(g)} + \text{H}_2\text{O(l)}\). The solid sulfur turns the mixture cloudy, so you stand the flask on a paper cross, mix, and time how long the cross takes to vanish. Rate is proportional to \(1/t\): a fixed amount of sulfur is needed to hide the cross, so the reciprocal of the time to make it is a fair rate proxy — but only while everything except temperature is held fixed.

That caveat is where the AO3 marks live. Independent variable: temperature, set by warming the thiosulfate in a water bath before mixing. Dependent variable: time for the cross to disappear. Control variables: concentrations and volumes of both solutions, the same flask (liquid depth changes the opacity needed), the same cross, and the same observer. The endpoint is the method's soft spot — 'disappeared' is a human judgement, so one person should judge every run, and a colorimeter or light sensor with a data logger is the standard improvement because it replaces the eye with a threshold. Temperature drifts during the run too: the fix is recording the temperature at mixing and at the endpoint and using the mean. Safety: the reaction releases toxic sulfur dioxide, so keep concentrations low, ventilate, and avoid running it hot enough to gas the room.

VocabularyKey terms the mark scheme pays for

TrapsMisconceptions that cost marks

ExamWhat examiners want