Catalysts

Collision Theory Recap

• Reactions happen when particles:

  1. Collide with each other.

  2. Collide with sufficient energy (≥ activation energy).

• Rate of reaction depends on the frequency of successful collisions (collisions per second that have enough energy).

Activation Energy

• Definition: The minimum amount of energy that particles need to react.

• Shown as the "energy barrier" on an energy profile diagram.

What is a Catalyst?

• Definition:
  A catalyst increases the rate of a chemical reaction but is not used up during the reaction.

How Catalysts Work

• Provide an alternative reaction pathway with a lower activation energy.

• This means:

  1. More particles have enough energy to collide successfully.

  2. More frequent successful collisions → faster reaction.

Key Points About Catalysts

• Not included in chemical equations (they’re not reactants).

• Catalysts can be reused.

• Different reactions require different catalysts.

• Enzymes are biological catalysts (special case in living organisms).

Energy Profile with Catalyst

• Without catalyst → higher activation energy "hump".

• With catalyst → lower activation energy hump.

• Reactants and products’ energy stays the same → only the pathway changes.

✅ Conclusion:
Catalysts speed up reactions by lowering the activation energy, making more collisions successful. They are not used up and can be reused.