IB Chemistry SLTopic 5 — The Rate of ReactionPaper 1 & 2Core idea~11 min read
Collision Theory
Particles in a gas or solution collide millions of times a second. If every collision produced a reaction, everything would react instantly. The interesting question is not why reactions happen, but why almost every collision achieves nothing.
📚 What you need to know
Particles must collide before they can react.
A collision is successful only if the particles have energy ≥ the activation energyand the correct orientation.
Rate depends on the number of successful collisions per second, not the total number.
The vast majority of collisions are unsuccessful — usually because of insufficient energy.
Particles at a given temperature have a range of kinetic energies, not a single value.
Kinetic energy is KE = ½mv2, so at the same temperature lighter particles move faster.
Colliding is not enough
An AND, not an OR. Failing either condition means the particles simply separate again unchanged.
The first requirement is energy. Reacting means breaking bonds, and bonds do not break for free. The colliding particles have to bring at least a certain minimum amount of energy between them — the activation energy — or the collision merely bounces.
The second is orientation. Molecules are not featureless spheres; the reaction happens at a particular part of each one. Approach from the wrong angle, with all the energy in the world, and the right atoms never meet.
The oxygen being transferred has to leave one molecule and arrive at the carbon of the other. Only one approach makes that geometrically possible.
Orientation matters most for large, complex molecules, where the reactive part may be a small site on a big structure. This is exactly why enzymes are so fussy: the substrate has to fit the active site, and any other approach is wasted.
Frequency and proportion
Two separate quantities decide the rate, and keeping them apart is the whole skill in this topic.
Collision frequency — how many collisions occur per second. Raise it and there are simply more chances.
The proportion that succeed — what fraction of those collisions clear the activation energy with a workable orientation.
What sets the rate
rate ∝ number of successful collisions per second
= collision frequency × the fraction that succeed
Multiplying the two is what matters. A change that doubles the collision frequency doubles the rate; so does a change that doubles the proportion succeeding. As you will see on the next page, most factors move only one of these, and one moves both.
Perspective on how wasteful this is: in a typical gas-phase reaction at room temperature, something like one collision in a billion has enough energy to react. Reactions proceed at a sensible speed only because collisions are so extraordinarily frequent.
Energy is spread out
Particles do not all move at the same speed. Collisions constantly redistribute energy, so at any instant some particles are moving very slowly, a few very fast, and most somewhere in the middle. Temperature fixes the average kinetic energy, not each individual one.
That spread is why activation energy acts like a threshold rather than an on-off switch: at any moment, some fraction of the particles is above it. Raise the temperature and that fraction grows — the idea that the Maxwell–Boltzmann distribution puts on a graph later in this sub-topic.
One consequence of KE = ½mv2 is worth noting: at the same temperature, two particles of different mass have the same average kinetic energy, so the lighter one must be moving faster. This is why light gases diffuse and effuse more quickly than heavy ones.
WORKED EXAMPLE
Explain, in terms of collision theory, why the reaction between hydrogen and oxygen does not happen at room temperature even though it is strongly exothermic.
Step 1 — separate the two ideasHow much energy a reaction RELEASES says nothing about how much it needs to START.Step 2 — apply collision theoryThe activation energy is very high, because strong H–H and O=O bonds must be broken first.Step 3 — the conclusionalmost no collisions reach EₐAt room temperature virtually no particles have enough energy, so the mixture sits unreacted. A spark supplies enough to start it, and the energy released then keeps it going.
WORKED EXAMPLE
Classify each change as affecting collision frequency, the proportion of successful collisions, or both: (a) increasing concentration, (b) raising temperature, (c) adding a catalyst, (d) grinding a solid to a powder.
(a) concentrationfrequency onlyMore particles per unit volume means more collisions, but each collision is no more energetic than before.(b) temperaturebothFaster particles collide more often AND a greater proportion carry enough energy. The second effect is by far the larger.(c) catalystproportion onlyLowering Eₐ means more of the existing collisions qualify. Nothing collides any more often than before.(d) powdering a solidfrequency onlyMore surface exposed, so more particles are available to be hit — but the energy of each collision is unchanged.
💡 Exam tip
Say successful collisions, not just “collisions”. The word does real work in a mark scheme.
Give both conditions when defining a successful collision: sufficient energy and correct orientation.
Be precise about frequency versus proportion — examiners write questions specifically to separate them.
Never say a catalyst or a temperature rise “gives the particles more energy to break bonds” without linking it to Ea.
Remember that ΔH and Ea are independent: an exothermic reaction can still be immeasurably slow.
⚠️ Common mix-up
Thinking all collisions react. Almost none do.
Assuming exothermic means fast. The two have nothing to do with each other.
Saying particles “gain energy” from a catalyst. A catalyst lowers the bar; it does not raise the particles.
Treating orientation as the main reason collisions fail. Insufficient energy is the usual culprit.
Assuming all particles at one temperature have the same energy. They have a wide distribution.
Up next: Factors Affecting Reaction Rate — five practical changes, each explained by exactly the two quantities you have just separated.
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