A catalyst increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy, making it easier for reactants to convert into products. Here’s a detailed explanation:
1. Activation Energy Concept
- Every chemical reaction requires a minimum energy called activation energy (Ea) to start.
- Higher Ea → slower reaction; lower Ea → faster reaction.
- A catalyst does not change the reactants or products—it just lowers the energy barrier.
2. How Catalysts Work
- Adsorption of Reactants (for heterogeneous catalysts)
- Reactant molecules stick to the catalyst surface.
- This brings them closer together and weakens bonds that need to break.
- Formation of Intermediate
- The catalyst may form a temporary complex with reactants.
- This stabilizes the transition state, making it easier for the reaction to proceed.
- Reaction Along a Lower-Energy Pathway
- Because the activation energy is reduced, more molecules have enough energy to react.
- The reaction happens faster even at the same temperature.
- Desorption of Products
- Products leave the catalyst surface.
- The catalyst remains unchanged and can participate in another reaction cycle.
3. Key Points
- Catalysts do not change the overall energy change (ΔG or ΔH) of the reaction.
- They increase the reaction rate by lowering the activation energy, not by providing energy themselves.
- Can be homogeneous (same phase as reactants) or heterogeneous (different phase, usually solid).
4. Simple Analogy
Think of a catalyst as a slide in a playground:
- Without a slide, you must climb over a wall (high energy) to reach the bottom.
- With a slide, you take an easier path (lower energy) and reach the bottom faster.