According to collision theory, for a chemical reaction to occur:
Particles must collide, and those collisions must be effective.
- Not every collision leads to a reaction.
- In fact, most do not.
- Three key conditions must be satisfied.
This seems obvious, but it’s fundamental:
No collision → no reaction.
- Reactant particles are always moving and colliding.
- Only when two (or more) reactant particles meet can a reaction even be possible.
- Not all collisions have enough energy to break bonds.
- To react, colliding particles must have energy ≥ activation energy, $E_a$.
- Activation energy ($E_a$) is the minimum energy needed to start breaking old bonds so that new ones can form.
- If the particles collide with too little energy, they simply bounce off unchanged.
- If you push the ball gently, it rolls partway up the hill and comes back → no reaction.
- If you push hard enough to get it over the top, it rolls down the other side → reaction occurs.
- Same idea with particles: only collisions with enough energy to get over the “hill” ($E_a$) lead to products.
Key idea: Even exothermic reactions that release lots of energy still need some activation energy to get started.
Even if particles collide with enough energy, they must also hit in the right way.
Particles must have the correct orientation during collision.
- Many molecules have specific shapes and bonding sites.
- If they hit “sideways” or at the wrong angle, the bonds that need to break or form may not be properly aligned.
- Two puzzle pieces only fit together if you turn them the right way.
- If they’re misaligned, they bump into each other but don’t connect—just like an ineffective collision.
For a reaction to occur:
- Particles must collide.
- They must collide with enough energy ($\geq E_a$).
- They must collide with the correct orientation.
Only collisions that satisfy all three conditions are called successful (effective) collisions.
- The rate of a chemical reaction depends on:
- How often particles collide (collision frequency).
- How many of those collisions have enough energy and the correct orientation.
- Three main factors influence these:
- Temperature
- Concentration
- Surface area (for solids)
- When you increase the temperature:
- Particles move faster → more collisions per second.
- Particles have more kinetic energy → a greater proportion of them have energy $\geq E_a$.
- So higher temperature → more collisions and more effective collisions → reaction rate increases.
If you heat a reaction mixture from 20 °C to 40 °C:
- The particles move faster.
- A larger fraction of them has enough energy to overcome the activation energy.
- The reaction may happen much faster (sometimes even twice as fast or more).
- The relationship between temperature and reaction rate is not linear.
- A small rise in temperature can cause a large increase in reaction rate because many more particles suddenly have energy $\geq E_a$.
- Concentration (solutions):
- Increasing the concentration of a reactant means more particles per unit volume.
- Particles are closer together, so collisions occur more frequently.
- More collisions per second → more effective collisions → faster reaction.
- Pressure (for gases) behaves similarly:
- Increasing pressure squeezes gas particles closer together.
- Collision frequency increases, and so the reaction speeds up.
If one reactant is a solid:
- Only particles at the surface can collide with particles in a gas or solution.
- Breaking a solid into smaller pieces (or a powder) increases its surface area.
More surface area → more exposed particles → more collisions per second → faster reaction.
- A large lump of calcium carbonate reacts slowly with hydrochloric acid.
- The same mass of calcium carbonate in powder form reacts much faster because many more particles are exposed at the surface.
Grinding solids into powders or using small chips is a common way to increase surface area and speed up reactions.
Collision theory explains reaction rates by combining:
- Collision frequency (how often particles collide), and
- Collision effectiveness (how many have enough energy and the right orientation).
- High $E_a$ → fewer particles have enough energy to react.
- Fewer effective collisions → slower reaction.
- Low $E_a$ → more particles have enough energy.
- More effective collisions → faster reaction.
- Explosive reactions (like combustion of gases such as hydrogen or petrol vapour) usually have low activation energies and are highly exothermic, so once started they proceed very rapidly.
- Even if $E_a$ is not very high, reactions can still be slow if:
- Concentrations are low → few collisions.
- One reactant is in a large solid lump → small surface area → few collisions at a time.
- By increasing concentration, pressure (for gases), or surface area (for solids), we can increase collision frequency and speed the reaction up.
A reaction is fast when:
- It has low activation energy, so many collisions have enough energy.
- Reactant particles collide frequently (high concentration, high pressure, large surface area).
- Temperature is relatively high, so many particles have energy ≥ Eₐ.
A reaction is slow when:
- It has high activation energy.
- Collisions are infrequent (low concentration, low surface area, lower temperature).
- Or both.
- Students often think higher temperature only means “more collisions”.
- In reality, higher temperature also means more energetic collisions, so a greater fraction of collisions are successful.
- According to collision theory, what three conditions must be met for a collision to lead to a reaction?
- How does increasing temperature affect:
- The frequency of collisions?
- The energy of collisions?
- Explain why a powdered solid reacts faster than a single large lump of the same substance.
- Using collision theory, explain why the combustion of a fuel gas can be very fast, while the rusting of iron is very slow, even though both are reactions with oxygen.