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Activation energy is the minimum amount of energy that particles must have to react when they collide, measured in kJ/mol. Free notes and 8 questions answered.
Activation Energy and Collision Theory is part of Exothermic and endothermic reactions in GCSE AQA Higher Chemistry.
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Activation energy is the minimum amount of energy that particles must have to react when they collide, measured in kJ/mol. Free notes and 8 questions answered.
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Use this lesson to stay inside Exothermic and endothermic reactions while connecting the detail back to the wider Topic 5: Energy changes topic.
Read the core explanations from Activation Energy and Collision Theory before testing yourself from memory.
Think about a match. Wood and oxygen are all around us, but they don't burst into flames until you strike the match. Why not?
The answer lies in energy barriers. Even when particles want to react, they need enough energy to get the reaction started. This is where activation energy comes in - it's like the energy hill that particles must climb before they can react.
Picture particles as bouncing balls. They're constantly moving and colliding, but most collisions don't cause reactions. Only the high-energy collisions - those with energy equal to or greater than the activation energy - can break bonds and form new ones.
Reaction profiles are energy maps that show the journey from reactants to products. They reveal three crucial pieces of information: the activation energy, whether the reaction is exothermic or endothermic, and the overall energy change.
The peak of the curve is the highest energy point where bonds are breaking and forming. The height of this peak above the reactants gives us the activation energy.
Worked example: Reading a Reaction Profile
From the reaction profile, identify: (a) activation energy, (b) overall energy change, (c) whether the reaction is exothermic or endothermic.
Answer: Read from left to right: reactants → peak (activation energy) → products (energy change)
For a reaction to occur, two conditions must be met simultaneously: particles must collide AND they must have sufficient energy. This explains why increasing temperature speeds up reactions - more particles have enough energy to overcome the activation energy barrier.
Think of it like a high jump competition. The bar (activation energy) is set at a certain height. Only athletes (particles) who can jump high enough (have sufficient energy) will clear the bar (react successfully).
Not all high-energy collisions lead to reactions. Particles must also collide with the correct orientation - the reactive parts must hit each other. This is why even with sufficient energy, some collisions still fail to produce products.
Exam questions often ask you to calculate activation energy from reaction profiles. The key is identifying the correct energy levels and subtracting carefully.
Worked example: Calculating Activation Energy from a Reaction Profile
A reaction profile shows reactants at 100 kJ/mol, a peak at 180 kJ/mol, and products at 60 kJ/mol. Calculate the activation energy. [2 marks]
Answer: 80 kJ/mol
Activation energy questions appear regularly in GCSE exams. Here's how to maximise your marks and avoid common pitfalls.
Mark-winning phrases: Use 'sufficient energy' rather than 'enough energy', 'overcome the activation energy barrier' rather than 'get over the hump', and always specify 'successful collisions' when explaining rate changes.
Exam precision note: In Activation Energy and Collision Theory - Reaction profiles - Exothermic and endothermic reactions, always identify the key chemical idea first, then link it to evidence, data, or an equation. A high-scoring GCSE answer names the particles involved, states what changes, and justifies the conclusion using correct scientific language. Where calculations are used, show each step with units and keep significant figures sensible. Finally, check that terminology matches AQA wording so your answer is unambiguous and mark-scheme friendly.
Successful collisions need both enough energy and the right orientation. Increasing temperature does not lower the activation energy itself, but it increases the proportion of particles with enough energy to overcome that barrier, so more collisions become successful each second.
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Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
Activation energy is the minimum amount of energy that particles must have to react when they collide, measured in kJ/mol. On a reaction profile it is the height of the peak measured from the reactants level.
Activation energy is measured in kilojoules per mole (kJ/mol). It is the minimum amount of energy particles must have to react when they collide, shown on a reaction profile as the height of the peak over the reactants level.
Activation energy = peak energy - reactant energy. For a reaction profile with reactants at 100 kJ/mol and a peak at 180 kJ/mol, the activation energy is 180 - 100 = 80 kJ/mol.
A successful collision is a collision between particles that has enough energy, equal to or greater than the activation energy, and the correct orientation to cause a reaction. Even with sufficient energy, collisions with the wrong orientation fail to produce products.
Increasing temperature does not lower the activation energy itself. It increases the proportion of particles with enough energy to overcome the activation energy barrier, so more collisions become successful each second.
Reactions do not all happen instantly because particles need enough energy to overcome an energy barrier, the activation energy, before they can react. Wood and oxygen do not burst into flames until a match is struck to give the energy to get the reaction started.
Correct orientation means the reactive parts of the particles must hit each other when they collide. Even with sufficient energy, a collision with the wrong orientation still fails to produce products.
Not all collisions lead to a reaction because a successful collision needs both sufficient energy, equal to or greater than the activation energy, and the correct orientation. Most collisions between particles don't have enough energy to break bonds.