Collision theory states that a chemical reaction can only occur when reactant particles collide with sufficient energy and the correct orientation. The minimum energy required for a collision to result in a reaction is the activation energy. Temperature, concentration, surface area, and catalysts all change the rate of reaction by affecting how many collisions meet this energy threshold.
What is collision theory?
Collision theory provides a particle-level explanation of why chemical reactions occur and how fast they proceed. Two key conditions must both be satisfied for a collision to be successful:
- Sufficient energy — the colliding particles must together possess at least the activation energy (Ea). If they collide with less energy, they simply bounce apart unchanged.
- Correct orientation — the reacting parts of the molecules must be aligned suitably during the collision. For simple ions in solution this is less critical, but for more complex molecules orientation matters greatly.
The vast majority of collisions at room temperature are unsuccessful because most particles have less energy than the activation energy.
What is activation energy?
Activation energy (Ea) is the minimum amount of energy that colliding particles must have for the collision to result in a chemical reaction. It represents the energy barrier that must be overcome to break the bonds in the reactants and begin forming the bonds of the products.
On an energy profile (reaction coordinate diagram), the activation energy is the difference in energy between the reactants and the peak (the transition state or activated complex). The transition state is an unstable, high-energy arrangement of atoms that exists momentarily at the top of the energy barrier.
| Feature of energy profile | What it represents |
|---|---|
| Left side (reactant energy level) | Energy stored in bonds of reactants |
| Peak (transition state) | Activation energy — energy needed to initiate the reaction |
| Right side (product energy level) | Energy stored in bonds of products |
| Peak − reactant level | Activation energy Ea |
| Product level − reactant level | Overall energy change (ΔH); negative = exothermic |
How does the Maxwell–Boltzmann distribution explain reaction rates?
At any given temperature, particles in a gas or solution do not all have the same kinetic energy — they have a distribution of energies. The Maxwell–Boltzmann distribution is a curve showing the number of particles at each energy level:
- The curve starts at zero (no particles have zero energy).
- It rises to a peak (the most probable energy), then tails off gradually to the right.
- A small fraction of particles have very high energies — these are in the long tail to the right.
- Only particles with energy ≥ Ea (in the tail to the right of the activation energy line) can react.
Effect of increasing temperature: When temperature rises, particles move faster and the distribution shifts to the right:
- The peak shifts to a higher energy.
- The tail becomes longer.
- A much larger fraction of particles now have energy ≥ Ea.
Because the number of successful collisions increases dramatically with temperature, reaction rate rises steeply. A 10 °C increase roughly doubles the rate for many reactions.
How does each factor affect rate — a collision theory explanation?
| Factor | Change | Collision theory explanation |
|---|---|---|
| Temperature | Increase | Particles have more kinetic energy → more frequent collisions AND greater fraction with energy ≥ Ea → more successful collisions |
| Concentration (solutions) | Increase | More particles per unit volume → more frequent collisions per unit time → more successful collisions |
| Pressure (gases) | Increase | Same effect as concentration — particles are closer together, more frequent collisions |
| Surface area (solids) | Increase | More surface particles exposed → more frequent collisions with reactant solution or gas |
| Catalyst | Add | Provides alternative reaction pathway with lower Ea → larger fraction of particles exceed the new, lower threshold → more successful collisions without changing temperature |
How do catalysts lower activation energy?
A catalyst speeds up a reaction by providing an alternative reaction pathway that has a lower activation energy than the uncatalysed route. On an energy profile, the catalysed pathway shows a lower peak.
On the Maxwell–Boltzmann distribution, lowering Ea shifts the threshold line to the left, so a much larger area under the curve lies to the right of the threshold — meaning many more particles have enough energy to react. The result is a significant increase in rate without any change in temperature.
Key points about catalysts:
- They are not consumed — they can be recovered unchanged at the end.
- They do not change the overall energy change (ΔH) of the reaction — the reactant and product energy levels are the same; only the peak (transition state) is lower.
- Biological catalysts are called enzymes; they are proteins that catalyse reactions in living cells and have very high specificity.
What is the difference between homogeneous and heterogeneous catalysts?
- Homogeneous catalyst — in the same phase (state) as the reactants. For example, the acid catalyst in esterification is in solution, as are the reactants.
- Heterogeneous catalyst — in a different phase from the reactants. For example, iron catalyst (solid) in the Haber process for ammonia synthesis, where the reactants (nitrogen and hydrogen) are gases. Reactant molecules adsorb onto the catalyst surface, react, and products desorb.
Heterogeneous catalysts provide an enormous surface area for reaction and can be reused — the catalyst in a catalytic converter in a car exhaust is a platinum and palladium alloy on a ceramic support.
Frequently asked questions
What is activation energy in GCSE chemistry?
Activation energy (Ea) is the minimum energy that colliding particles must possess for a collision to result in a chemical reaction. Even if particles collide, no reaction occurs unless their combined kinetic energy meets or exceeds this threshold. On an energy profile diagram, it is shown as the height of the peak (transition state) above the reactant energy level. Reactions with a high activation energy are slow at room temperature because few particles have sufficient energy.
How does increasing temperature affect the Maxwell–Boltzmann distribution?
Increasing temperature shifts the Maxwell–Boltzmann energy distribution curve to the right — the peak moves to a higher energy and the curve broadens. The key effect is that the area under the curve to the right of the activation energy line increases substantially: a much larger fraction of particles now have energy ≥ Ea and can react successfully. This explains why even a modest temperature rise produces a large increase in reaction rate.
How does a catalyst appear on an energy profile diagram?
On an energy profile diagram, adding a catalyst lowers the peak (the transition state) compared with the uncatalysed route, because the catalyst provides an alternative pathway with a lower activation energy. The reactant and product energy levels remain the same — the catalyst does not change the overall energy released or absorbed (ΔH). The catalysed profile shows a lower, smaller peak between the reactant and product levels, while the uncatalysed profile shows a higher peak.
Why do catalysts not change the amount of product formed?
A catalyst speeds up the reaction by lowering the activation energy, increasing the proportion of successful collisions per unit time. However, it does not alter the thermodynamics of the reaction — the energy difference between reactants and products is unchanged. This means the equilibrium position is unchanged: the same amounts of products and reactants are present at equilibrium. Catalysts simply allow equilibrium to be reached faster, not by shifting it in favour of products.
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