Activation energy (AQA GCSE Chemistry): Revision Notes
Activation energy
What is activation energy?
Activation energy is the minimum amount of energy that particles need to start a chemical reaction. Think of it like a hill that particles must climb over before they can react with each other.
The symbol for activation energy is Ea.
Without enough activation energy, even if particles bump into each other, no reaction will happen. The particles need this energy "boost" to break their existing bonds and form new ones.
The hill analogy is particularly helpful for visualising activation energy. Just like you need energy to climb up a hill before you can roll down the other side, chemical particles need activation energy to "get over the hump" before the reaction can proceed.
Energy diagrams
Scientists use energy diagrams (also called reaction profiles) to show how energy changes during a reaction. These diagrams have two main parts:
- Energy on the y-axis (vertical)
- Progress of reaction on the x-axis (horizontal)
Exothermic reactions
In exothermic reactions, the products have less energy than the reactants. This means energy is given out to the surroundings.
On the energy diagram:
- Reactants start at a higher energy level
- There's still an energy barrier (activation energy) to overcome
- Products end up at a lower energy level than the reactants
- The difference shows energy has been released
Endothermic reactions
In endothermic reactions, the products have more energy than the reactants. This means energy is taken in from the surroundings.
On the energy diagram:
- Reactants start at a lower energy level
- There's an energy barrier (activation energy) to overcome
- Products end up at a higher energy level than the reactants
- The difference shows energy has been absorbed
How reactions happen
For any chemical reaction to occur, particles must meet specific requirements. Understanding these conditions helps explain why some reactions happen easily while others need extra help.
The Three Essential Requirements for Chemical Reactions:
- Particles must collide - the reactant particles need to bump into each other
- They must collide with enough energy - they need at least the activation energy
- They must collide in the right way - the particles need to hit each other at the correct angle
If particles don't have enough energy when they collide, they just bounce off each other. No reaction takes place.
Catalysts and activation energy
Catalysts are special substances that speed up reactions without being used up themselves. They work by providing a different pathway for the reaction that has a lower activation energy.
This means:
- More particle collisions will have enough energy to react
- The reaction happens faster
- The catalyst isn't changed by the reaction
- The overall energy change stays the same
On an energy diagram, a catalyst creates a lower "hill" for the particles to climb over.
Worked Example: Understanding Catalyst Effect
Imagine a reaction normally needs 100 kJ/mol of activation energy:
- Without catalyst: Only high-energy collisions (≥100 kJ/mol) can react
- With catalyst: The same reaction now only needs 60 kJ/mol
- Result: Many more collisions now have sufficient energy to react, so the reaction rate increases significantly
Key Points to Remember:
- Activation energy is the minimum energy needed for particles to react
- All reactions need activation energy, even exothermic ones that give out energy overall
- Energy diagrams show the energy barrier that must be overcome
- Catalysts lower the activation energy, making reactions faster
- Particles must collide with enough energy for successful reactions