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Collision theory provides the scientific explanation for why the rate of a chemical reaction changes when conditions are altered. It is a core part of the AQA GCSE Chemistry specification and underpins everything you have learned about factors affecting rate. This lesson explains the theory in full, introduces the concept of activation energy, and shows you how to use collision theory to explain experimental observations.
Collision theory states that for a chemical reaction to occur, the reacting particles must:
If either condition is not met, the collision is unsuccessful and no reaction takes place. The particles simply bounce apart.
| Condition | What It Means |
|---|---|
| Collision occurs | The reacting particles must physically hit each other |
| Sufficient energy | The kinetic energy of the colliding particles must be equal to or greater than the activation energy |
| Correct orientation | The particles must be facing the right way so that bonds can break and reform |
Exam Tip: In AQA GCSE Chemistry, you must use the term "successful collision" to describe a collision that leads to a reaction. An unsuccessful collision is one where the particles bounce apart without reacting because they did not have enough energy or the correct orientation.
The activation energy (Ea) is the minimum amount of energy that colliding particles must have in order for a reaction to take place. It is the energy needed to break the bonds in the reactants so that new bonds can form to make the products.
Every chemical reaction has its own activation energy. Some reactions have a low activation energy (they happen easily, even at room temperature) and some have a high activation energy (they need a lot of energy input to get started, such as heating with a flame).
| Reaction | Activation Energy | Ease of Reaction |
|---|---|---|
| Rusting of iron | Low | Happens at room temperature (but slowly) |
| Neutralisation (acid + alkali) | Very low | Happens immediately on mixing |
| Burning magnesium | High | Needs a Bunsen flame to start |
| Decomposition of limestone | Very high | Needs sustained high temperature |
graph TD
A[Reactant Particles Collide] --> B{Energy >= Activation Energy?}
B -->|Yes| C[Bonds break in reactants]
C --> D[New bonds form to make products]
D --> E[Successful collision - Reaction occurs]
B -->|No| F[Particles bounce apart]
F --> G[Unsuccessful collision - No reaction]
Exam Tip: When drawing or interpreting energy profile diagrams, the activation energy is shown as the difference between the energy of the reactants and the peak (top) of the energy curve. Make sure you can identify it on a diagram and label it correctly.
An energy profile diagram (also called a reaction profile) shows the energy changes during a reaction. It plots energy on the y-axis against the progress of the reaction on the x-axis.
For an exothermic reaction, the products have less energy than the reactants. Energy is released to the surroundings.
| Feature | Description |
|---|---|
| Reactants | Start at a certain energy level |
| Activation energy (Ea) | The rise from reactants to the peak |
| Peak | The maximum energy point — the transition state |
| Products | Lower energy than reactants |
| Overall energy change | Negative (energy released) |
For an endothermic reaction, the products have more energy than the reactants. Energy is absorbed from the surroundings.
| Feature | Description |
|---|---|
| Reactants | Start at a certain energy level |
| Activation energy (Ea) | The rise from reactants to the peak |
| Peak | The maximum energy point |
| Products | Higher energy than reactants |
| Overall energy change | Positive (energy absorbed) |
When concentration increases, there are more reactant particles in the same volume. This means particles are closer together and there are more collisions per second (higher frequency of collisions). With more collisions per second, there are more successful collisions per second, so the rate increases.
When temperature increases, particles have more kinetic energy and move faster. This has two effects:
The second effect is more significant. Even a small rise in temperature greatly increases the proportion of particles that can overcome the activation energy barrier.
graph TD
A[Temperature Increases] --> B[Particles gain kinetic energy]
B --> C[Particles move faster]
C --> D[More frequent collisions]
C --> E[More particles have energy >= Ea]
D --> F[More successful collisions per second]
E --> F
F --> G[Rate of reaction increases]
When a solid reactant is broken into smaller pieces, the surface area exposed to the other reactant increases. More particles on the surface are available to collide with particles in the solution or gas. This means more collisions per second, leading to more successful collisions per second and a higher rate.
When the pressure of a gas is increased, the same number of particles occupy a smaller volume. The particles are closer together, so they collide more frequently. More frequent collisions means more successful collisions per second and a higher rate.
A catalyst provides an alternative reaction pathway with a lower activation energy. This means a greater proportion of collisions have energy equal to or above the new, lower activation energy. More collisions are therefore successful, and the rate increases.
At any given temperature, the particles in a substance have a range of energies. Some particles move slowly (low energy) and some move quickly (high energy). Most particles have energies somewhere in the middle.
This distribution is important because:
| Temperature | Average Particle Energy | Proportion Exceeding Ea | Rate |
|---|---|---|---|
| Low | Low | Small | Slow |
| High | High | Large | Fast |
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