Collision Theory & Activation Energy (OCR GCSE Chemistry A, Combined (Gateway Science Suite)): Revision Notes
📚 Revision Notes
6.1.7 Collision Theory & Activation Energy
Collision Theory:
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Collision Theory helps us understand how and why chemical reactions happen. It states that for a reaction to occur, the particles (atoms, molecules, or ions) of the reactants must collide with each other.
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Just bumping into each other isn't enough. For the collision to lead to a reaction, the particles must hit with enough energy and in the right direction.
Successful Collisions:
- Not all collisions result in a reaction. A successful collision is one where particles collide with enough energy to break their bonds and form new ones, leading to the creation of products.
- The rate of reaction depends on how often these successful collisions occur. The more frequent the successful collisions, the faster the reaction.
Activation Energy:
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Activation Energy is the minimum amount of energy that particles need to collide successfully and start a reaction. Think of it as the energy "barrier" that must be overcome for a reaction to happen.
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If the particles don't have enough energy, they'll just bounce off each other without reacting.
Impact on Reaction Rate:
- Reactions with low activation energy happen more quickly because more particles have enough energy to collide successfully.
- Reactions with high activation energy are slower because fewer particles can overcome the energy barrier during collisions.