According to reaction rate theory, a catalyst increases the reaction rate by lowering the activation energy required for the reaction to occur. Here’s a explanation:
How Catalysts Work:
- Activation energy barrier:
- For a reaction to happen, molecules must have enough energy to reach the transition state.
- This energy threshold is called activation energy.
- Catalyst lowers activation energy:
- A catalyst provides an alternative reaction pathway with a lower energy barrier.
- More molecules now have enough energy to react → reaction speeds up.
- Molecules collide more effectively:
- Catalysts can also orient molecules correctly or bring them closer together, increasing the chances of a successful collision.
- Catalyst is not consumed:
- It participates in the reaction temporarily but is regenerated at the end, ready to catalyze again.
Simple idea:
- A catalyst is like a shortcut for a reaction: it makes it easier for molecules to react without changing the final products.
- Because of this, more molecules can react per unit time → faster reaction.