Collisions and reactions
Collision theory explains rate of reaction in terms of colliding particles. For a reaction to occur, reactant particles must collide with one another. However, not every collision leads to a reaction.
A collision that actually produces products is called an effective collision. Two conditions must be met for a collision to be effective:
- The colliding particles must have energy equal to or greater than the activation energy (Ea) — the minimum energy needed to start a reaction.
- The particles must collide in the correct orientation (the right way round).
Key idea
Rate of reaction ∝ frequency of effective collisions. Anything that increases the number of effective collisions per second increases the rate.
Linking factors to collisions
Each factor that affects rate can be explained by collision theory:
- Higher concentration or pressure — more particles per unit volume, so collisions are more frequent, so more effective collisions per second.
- Larger surface area — more particles are exposed, so collisions are more frequent.
- Higher temperature — particles move faster (more frequent collisions) AND a larger fraction have energy ≥ Ea, so a much bigger share of collisions are effective. This is why temperature has a strong effect.
- Catalyst — provides a path with a lower activation energy, so a larger fraction of collisions have enough energy to be effective. The frequency of collisions is not changed.
Worked example
Two gas molecules collide but bounce apart unchanged. Explain why. Either their combined energy was below the activation energy, or they collided in the wrong orientation. Only collisions that satisfy BOTH conditions — enough energy and correct orientation — are effective and lead to products.
Remember
- Not all collisions are effective.
- Effective collision = energy ≥ Ea AND correct orientation.
- A catalyst lowers Ea; it does not lower the temperature.