Rate of Reaction » Collision Theory and Activation Energy

Collision theory

You already know that some factors make reactions faster. Collision theory explains why. It is the idea behind every exam answer about rate.

According to collision theory, a chemical reaction can only happen when two things are true:

• the reacting particles collide with each other, and
• they collide with enough energy.

Many collisions do not lead to a reaction. The particles just bounce off each other, like two snooker balls. Only the collisions with enough energy are successful.

Key terms:

• Collision theory: the idea that reactions happen only when reacting particles collide with sufficient energy.
• Frequency of collisions: how many collisions happen in each second.

Activation energy

Remember, the minimum energy particles must have to react is called the activation energy. A collision with less energy than this does not cause a reaction. A collision with at least this much energy can.

Think of a football match. Passing the ball gently along the ground does not score a goal. Only a hard enough shot gets it into the net. In the same way, only a hard enough collision breaks the bonds in the reactants.

So there are two ways to make a reaction faster:

The rate of reaction goes up when either of these happens. Look for them in every answer.

Concentration

In a solution, the reacting particles are spread out in the water. In a more concentrated solution there are more reacting particles in the same volume. They are packed closer together.

So the particles collide more often. The frequency of collisions increases, so the rate of reaction increases.

Notice that the particles are not moving any faster. Each collision has the same energy as before. There are just more of them every second.

Simple proportionality: if you double the concentration, there are twice as many particles in the same space, so there are about twice as many collisions each second. The rate roughly doubles.

Pressure of gases

Gas particles are far apart. If you raise the pressure, the same number of gas particles is squeezed into a smaller volume. They are closer together.

So the particles collide more often. The frequency of collisions increases, so the rate of reaction increases. This is the same idea as concentration, because squeezing a gas makes it more concentrated.

Pressure only matters for reactions where a gas is a reactant. It makes no real difference to reactions with only solids and liquids.

Surface area of solids

In a reaction between a solid and a liquid or gas, the reaction happens at the surface of the solid. Particles inside a lump cannot be hit.

If you break a solid into smaller pieces, more of its particles are on the outside. There is more surface for the other reactant to hit. So the frequency of collisions increases and the rate of reaction increases.

Remember the surface area to volume ratio. Small pieces have a bigger surface area compared with their volume. Here are two cubes of the same solid:

• A cube with sides of 2 cm has a surface area of 24 cm² and a volume of 8 cm³. The ratio is 3 : 1.
• A cube with sides of 1 cm has a surface area of 6 cm² and a volume of 1 cm³. The ratio is 6 : 1.

The smaller cube has the bigger ratio. Cutting the 2 cm cube into eight 1 cm cubes gives the same mass, but twice the surface area to volume ratio. So smaller pieces react faster than one big lump.

Temperature

Temperature is special because it does two things.

Both effects make the rate of reaction increase.

Summary

• Higher concentration: more particles in the same volume, so more frequent collisions.
• Higher pressure (gases): particles closer together, so more frequent collisions.
• Larger surface area: more particles exposed, so more frequent collisions.
• Higher temperature: more frequent collisions and more energetic collisions.

Exam-style question

A student reacts magnesium ribbon with dilute hydrochloric acid. Explain, using collision theory, why the reaction is faster when the acid is warmed from 20 °C to 40 °C. [4 marks]