Collision Theory, Activation Energy and Catalysts

For a reaction to happen, reacting particles must first bump into each other. But not every collision causes a reaction. Collision theory explains which ones do.

Two conditions must both be met for a successful collision:

• The particles must collide with each other.
• They must collide with enough energy to react.

A collision with too little energy simply bounces the particles apart unchanged. So the rate of a reaction depends on how often particles collide and on what fraction of those collisions have enough energy.

Reactions occur only when particles collide with sufficient energy. A collision with too little energy does not cause a reaction.

The more frequent the successful collisions, the faster the reaction. Everything that increases rate does so by increasing the frequency of collisions, the energy of collisions, or both.

The "enough energy" in collision theory has a name: the activation energy.

Activation energy is the minimum energy that colliding particles must have for a reaction to occur.

Think of it as an energy barrier the particles have to get over. If a collision brings less than the activation energy, the reaction cannot start and the particles separate again. If it brings at least the activation energy, the collision is successful and products form.

This idea links directly to reaction rate. At any moment, only a fraction of particles are moving fast enough for a collision between them to exceed the activation energy. Anything that raises the proportion of collisions that clear this barrier will speed the reaction up. The reaction profile below shows the activation energy as the "hump" reactants must climb before turning into products.

Each factor from the previous lesson can now be explained. The key is to say whether it changes the frequency of collisions, the energy of collisions, or both.

Factor increased	Effect on collisions	Why the rate rises
Concentration (solution)	More frequent collisions	More particles in the same volume, so they collide more often
Pressure (gas)	More frequent collisions	Gas particles squeezed closer together, so they collide more often
Surface area of a solid	More frequent collisions	More of the solid is exposed, so more particles are available to be hit
Temperature	More frequent and more energetic collisions	Particles move faster, so they collide more often and more collisions exceed the activation energy

Concentration, pressure and surface area increase only the frequency of collisions. Temperature increases both the frequency and the energy of collisions, which is why it has such a strong effect.

For surface area, it is the surface area to volume ratio of the solid that matters: smaller pieces have a higher ratio, exposing proportionally more particles.

Collision theory lets you make rough predictions using proportionality. If doubling the concentration doubles the number of particles in a given volume, it roughly doubles how often they collide, so the rate roughly doubles.

For example, if a reaction produces gas at 2 cm³/s at a certain acid concentration, then doubling that concentration would be expected to give a rate of about 4 cm³/s, because collisions happen about twice as often.

If the concentration of a reactant doubles, the frequency of collisions roughly doubles, so the rate roughly doubles.

This is a simple proportional relationship, not an exact rule for every reaction, but it is the reasoning the exam expects when a graph or table shows how rate changes with concentration.

A catalyst increases the rate of a reaction without being used up. It does this by providing a different reaction pathway with a lower activation energy.

Lowering the activation energy means the barrier is easier to clear, so a greater proportion of collisions now have enough energy to be successful. More successful collisions per second means a faster reaction.

The energy profile below shows both pathways. The uncatalysed route (higher hump) needs more energy to start; the catalysed route (lower hump) needs less. The reactants and products sit at the same energy levels in both cases, so the catalyst does not change the overall energy released or absorbed, only the height of the barrier between them.

A catalyst provides a pathway with a lower activation energy. It does not lower the energy of the reactants or products, only the barrier between them.

Catalysts have a set of properties the exam tests directly.

• A catalyst is not used up in the reaction, so it can be reused. A small amount catalyses a large amount of reaction.
• A catalyst does not appear in the chemical equation, because it is not a reactant or a product.
• Different reactions need different catalysts. A catalyst that speeds up one reaction may do nothing for another.
• Enzymes are biological catalysts. They speed up reactions in living organisms, such as digestion.

You can identify a catalyst in an experiment because it increases the rate but its mass is unchanged at the end and it is absent from the balanced equation.

You do not need to recall the names of catalysts beyond those specified in the course. What you must explain is that a catalyst works by lowering the activation energy and is not used up.

1. Saying a catalyst lowers the energy of the reactants

A catalyst lowers the activation energy, which is the height of the barrier. It does not change the energy of the reactants or products, so the overall energy change stays the same.

2. Saying a catalyst is used up

A catalyst is chemically unchanged at the end of the reaction. It is not a reactant, does not appear in the equation, and can be used again.

3. Explaining temperature the same way as concentration

Concentration and pressure only increase how often particles collide. Temperature also gives each collision more energy, so more collisions exceed the activation energy. State both effects for temperature.

4. Forgetting that collisions need enough energy

More collisions alone do not guarantee more reaction. The collisions must have at least the activation energy. A full answer mentions both the frequency and the energy of collisions.

5. Confusing catalyst and fuel

A catalyst is not consumed and provides no energy. It simply lowers the activation energy so existing collisions succeed more often.