Collision theory Notes

What Conditions Must Be Met for Particles to React Successfully?

According to collision theory, for a chemical reaction to occur:

Particles must collide, and those collisions must be effective.

1. Not every collision leads to a reaction.
2. In fact, most do not.
3. Three key conditions must be satisfied.

Condition 1: Particles Must Collide

This seems obvious, but it’s fundamental:

No collision → no reaction.

1. Reactant particles are always moving and colliding.
2. Only when two (or more) reactant particles meet can a reaction even be possible.

Condition 2: Collisions Must Have Sufficient Energy

1. Not all collisions have enough energy to break bonds.
2. To react, colliding particles must have energy ≥ activation energy, $E_a$.
3. Activation energy ($E_a$) is the minimum energy needed to start breaking old bonds so that new ones can form.
4. If the particles collide with too little energy, they simply bounce off unchanged.

Condition 3: Correct Orientation

Even if particles collide with enough energy, they must also hit in the right way.

Particles must have the correct orientation during collision.

1. Many molecules have specific shapes and bonding sites.
2. If they hit “sideways” or at the wrong angle, the bonds that need to break or form may not be properly aligned.

How Do Temperature, Concentration, and Surface Area Affect Collisions?

1. The rate of a chemical reaction depends on:
   1. How often particles collide (collision frequency).
   2. How many of those collisions have enough energy and the correct orientation.
2. Three main factors influence these:
   1. Temperature
   2. Concentration
   3. Surface area (for solids)

Temperature

1. When you increase the temperature:
   1. Particles move faster → more collisions per second.
   2. Particles have more kinetic energy → a greater proportion of them have energy $\geq E_a$.
2. So higher temperature → more collisions and more effective collisions → reaction rate increases.

Concentration (for Solutions) / Pressure (for Gases)

1. Concentration (solutions):
   1. Increasing the concentration of a reactant means more particles per unit volume.
   2. Particles are closer together, so collisions occur more frequently.
   3. More collisions per second → more effective collisions → faster reaction.
2. Pressure (for gases) behaves similarly:
   1. Increasing pressure squeezes gas particles closer together.
   2. Collision frequency increases, and so the reaction speeds up.

Surface Area (for Solids)

If one reactant is a solid:

1. Only particles at the surface can collide with particles in a gas or solution.
2. 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.

Why Are Some Reactions Fast and Others Slow? (Collision Theory View)

Collision theory explains reaction rates by combining:

1. Collision frequency (how often particles collide), and
2. Collision effectiveness (how many have enough energy and the right orientation).

Factor 1: Activation Energy

1. High $E_a$ → fewer particles have enough energy to react.
   1. Fewer effective collisions → slower reaction.
2. Low $E_a$ → more particles have enough energy.
   1. More effective collisions → faster reaction.
3. 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.

Factor 2: Frequency of Collisions

1. Even if $E_a$ is not very high, reactions can still be slow if:
   1. Concentrations are low → few collisions.
   2. One reactant is in a large solid lump → small surface area → few collisions at a time.
2. By increasing concentration, pressure (for gases), or surface area (for solids), we can increase collision frequency and speed the reaction up.