Collision Theory and Rate of Reaction

Collision theory explains why reactions happen at different speeds. It says particles must collide to react, and only some collisions are successful.

What makes a collision successful?

• Enough energy: Particles must collide with energy at least equal to the activation energy, written as $E_a$.
• Correct orientation: Particles must hit in a way that lets old bonds break and new ones form.

Think of pushing open a heavy door. You need to push hard enough (energy) and in the right place (orientation).

Key idea (in symbols)

$$\text{Successful collision} \iff E_{\text{collision}} \geq E_a \text{ and correct orientation}$$

How factors change rate using collision theory

• Concentration (solutions): More particles in the same volume means more frequent collisions, so a faster rate.
• Pressure (gases): Higher pressure squeezes gas particles closer, increasing collision frequency and rate.
• Surface area (solids): Smaller pieces expose more surface for collisions. Powdered solids react faster than lumps.
• Temperature: Particles move faster and have more kinetic energy. Collisions happen more often, and a larger fraction have energy $\geq E_a$, so rate increases strongly.
• Catalyst: Provides an alternative pathway with lower activation energy $E_a$. More collisions are successful. The catalyst is unchanged at the end.

Real-world connections

• Finely ground sugar dissolves faster than a sugar cube (greater surface area).
• Food spoils slower in the fridge (lower temperature reduces successful collisions).

Practical link

• Measure gas made using a gas syringe or measure mass lost on a balance to compare rates.

Common misconceptions

• “More collisions always mean reaction.” Only collisions with energy $\geq E_a$ and correct orientation react.
• “Catalysts make particles faster.” Catalysts lower $E_a$; they do not increase particle kinetic energy.