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activation energy catalysis
A LevelChemistry~4 min read
Overview
This lesson explores activation energy, the minimum energy required for a reaction to occur, and how catalysts provide an alternative reaction pathway with a lower activation energy, thereby increasing reaction rates without being consumed.
Understanding Activation Energy (Ea)
Activation energy (Ea) is a fundamental concept in chemical kinetics, representing the **minimum kinetic energy** that colliding reactant particles must possess for a reaction to occur. Not all collisions between reactant particles lead to a reaction; only those with sufficient energy to overcome th...
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Key Concepts
- Activation Energy (Ea): The minimum amount of energy that reactant particles must possess for a successful collision to occur, leading to a chemical reaction.
- Transition State (Activated Complex): A high-energy, unstable intermediate formed during a reaction, existing at the peak of the activation energy barrier.
- Catalyst: A substance that increases the rate of a chemical reaction without being chemically changed or consumed in the overall process.
- Homogeneous Catalysis: Catalysis where the catalyst and reactants are in the same physical state (e.g., all gases or all liquids).
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Exam Tips
- →Always draw and label reaction profile diagrams accurately, showing both catalysed and uncatalysed pathways, Ea, and ΔH. Ensure the Ea for the catalysed path is lower, but ΔH is the same.
- →Clearly explain *how* a catalyst works: by providing an alternative reaction pathway with a lower activation energy. Do not just state that it 'lowers activation energy'.
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