Collision Theory and Reaction Rates

Why are some reactions fast (like burning) and others slow (like rusting)? Collision theory explains this using moving particles.

Key idea of collision theory

• Particles must collide to react. No collision → no reaction.
• Not every collision works. A collision is successful only if particles have enough energy and the right orientation.

Activation energy is the minimum energy needed to start a reaction. It is written as $E_a$. A useful summary is:

$$\text{Rate of reaction} \propto \text{number of successful collisions per second}$$

Only collisions with energy $E \ge E_a$ and correct orientation are successful.

Energy hurdle picture

Imagine pushing a ball over a hump. If you do not push hard enough, the ball rolls back. Particles also must “get over” an energy hump ($E_a$) to form products.

How different factors change rate (explained by collisions)

• Concentration (solutions): More particles in the same volume → more frequent collisions → higher chance of successful collisions → faster rate.
• Pressure (gases): Particles are squeezed closer → collide more often → faster rate.
• Surface area (solids): Smaller pieces have more exposed surface → more places for particles to hit → more frequent collisions → faster rate.
• Temperature: Particles move faster and have more kinetic energy. This gives (1) more collisions per second and (2) a bigger fraction of particles with energy $\ge E_a$. Rate increases strongly.
• Catalyst: Provides a different pathway with lower $E_a$, so more collisions are successful. The catalyst is unchanged at the end.

Real-world links

• Food in a fridge lasts longer because lower temperature reduces successful collisions.
• Powdered medicines act quickly because of greater surface area.