Topic 9 · page 4 of 15

Collisions and activation energy

Particles must collide to react, but most collisions do nothing. Only a collision with at least the activation energy is successful, and that splits "rate" into how often collisions happen and how many are successful.

11 min read9.19.2

Petrol vapour and air can sit together in a car's fuel tank for months. Burning petrol is strongly exothermic, so plenty of energy is available. Yet nothing happens. A tiny spark, and the whole mixture reacts in a fraction of a second. The chemicals are the same and the energy released is the same. Something else decides how fast a reaction goes, and this page names it.

Reactions need collisions 9.1

For two particles to react they must first collide. In a gas or a solution, particles are moving all the time and colliding very often, billions of times each second. If every collision caused a reaction, everything would react at once. It does not, because most collisions are unsuccessful: the particles simply bounce apart unchanged.

Two collisions between particle A and particle B. Left, collision energy below the activation energy: short arrows show the particles approaching, then they bounce apart unchanged. Right, collision energy of at least the activation energy: long arrows show the particles approaching, then they join to form a product.
A collision is successful only if the particles have enough energy.

Activation energy 9.2

To react, existing bonds must be stretched or broken as new bonds begin to form. That takes energy, and it has to come from the kinetic energy of the colliding particles.

On an enthalpy profile, the activation energy is the height from the reactants up to the top of the hump. The reaction can only get over the hump if the collision supplies that much energy.

Reaction profile for an exothermic reaction without a catalyst. Enthalpy is on the vertical axis and progress of reaction on the horizontal axis. The curve starts at the reactants on a high level, rises over a single hump, and falls to the products on a lower level. A vertical arrow from the reactants level to the top of the hump is labelled activation energy. A downward arrow between the reactants and products levels is labelled delta H, negative.
The activation energy is the barrier between the reactants and the top of the hump. ΔH is the difference between reactants and products.

The size of EaE_\text{a} and the sign of ΔH are separate things. An exothermic reaction can have a large activation energy. That is why petrol and air do not react at room temperature: very few collisions have enough energy to get over the barrier. The spark gives a few molecules enough energy to react. Those reactions release energy, which gives more molecules enough energy, and the reaction spreads.

Rate in two parts 9.1

Putting the two ideas together, the rate depends on:

  1. How often particles collide (the collision frequency).
  2. What fraction of those collisions have energy of at least EaE_\text{a}.

The successful collisions per second are the product of the two. Every factor in the specification changes one or both:

Factor Changes How
Concentration, and pressure of a gas collision frequency more particles in the same volume
Surface area of a solid collision frequency more particles exposed to the other reactant
Temperature mainly the fraction with E≥EaE \geq E_\text{a} the energy distribution changes
Catalyst the value of EaE_\text{a} an alternative route with a lower barrier

The next pages take these one by one.

Check yourself

  1. Define activation energy.
  2. Explain why most collisions between reactant particles do not result in a reaction.
  3. Petrol vapour and air are mixed at room temperature and do not react, although burning is exothermic. Explain why, and why a spark starts the reaction.
  4. A student writes "collisions cause reactions". Rewrite this so that it is correct.
  5. State where the activation energy is measured on an enthalpy profile.
  6. State the two things that the rate of a reaction depends on, according to collision theory.

Answers cover these until you have tried the questions

  1. The minimum energy that colliding particles must have for a reaction to occur.
  2. Most colliding particles have less energy than the activation energy, so the bonds cannot be broken and the particles bounce apart unchanged. (Some collisions also have the wrong orientation.)
  3. The reaction is exothermic, but it has a high activation energy, and at room temperature very few collisions have energy of at least EaE_\text{a}, so the reaction is extremely slow. The spark gives a few molecules enough energy to react. Their reaction releases energy, which gives other molecules enough energy, so the reaction spreads rapidly.
  4. "Only collisions between particles with energy greater than or equal to the activation energy cause a reaction."
  5. On the vertical axis, from the level of the reactants up to the top of the hump.
  6. The frequency of collisions (how often particles collide) and the fraction of collisions that have energy of at least EaE_\text{a}.