Collision Theory

Summary: Collision theory explains reaction rates by stating that particles must collide with energy equal to or greater than the activation energy, and with the correct orientation, for a reaction to occur. Tags: igcse chemistry Created: 2026-07-18

Collision theory is the model chemists use to explain why chemical reactions happen and why they proceed at different rates. It states that for a reaction to occur, reactant particles must physically collide with one another. However, simply colliding is not enough: the particles must collide with a minimum amount of energy, called the activation energy (Ea), and they must be oriented correctly so that the right atoms or bonds meet. Collisions that satisfy all of these conditions are called successful (or effective) collisions, and only these lead to products. In reality, the vast majority of collisions in a reaction mixture are unsuccessful because the particles either lack sufficient energy or approach each other at the wrong angle. The rate of a reaction therefore depends on how frequently particles collide and on what fraction of those collisions are successful, which is why changing temperature, concentration, pressure, surface area, or adding a Catalyst alters the rate.

The Three Requirements for a Reaction

According to collision theory, a reaction between two particles can only happen if all three of the following conditions are met:

1. The particles must collide. Reactant particles that never come into contact cannot react. Particles in liquids and gases are in constant random motion, so collisions happen continuously.
2. The collision must have enough energy. The combined kinetic energy of the colliding particles must be equal to or greater than the activation energy (Ea). If the collision energy is below Ea, the particles simply bounce apart unchanged.
3. The particles must collide with the correct orientation. The reactive parts of the molecules must meet. If two molecules collide “back to front”, no bonds can break or form even if the energy is sufficient.

A collision that meets all three requirements is called a successful collision (or effective collision). A crucial point for IGCSE:

Most collisions are NOT successful. In a typical gas or solution, particles collide billions of times per second, yet only a tiny fraction of these collisions actually lead to a reaction. If every collision were successful, virtually all reactions would be explosively fast.

This gives us a simple relationship for reaction rate:

$\text{Rate}\propto \text{collision frequency}\times \text{fraction of collisions with }E\ge E_a$

Anything that increases either factor — more collisions per second, or a greater proportion of energetic collisions — increases the rate of reaction.

Activation Energy and the Energy Barrier

Activation Energy (Ea) is defined as the minimum energy that colliding particles must have for a reaction to occur. It can be pictured as an energy “barrier” or “hill” between reactants and products on a reaction profile (energy level diagram):

• Colliding particles must have at least Ea to break the existing bonds in the reactants and start forming new bonds.
• Particles with less energy than Ea collide and bounce off each other unchanged — the collision is unsuccessful.
• Once the barrier is crossed, new bonds form and energy is released as products form.

The energies of particles in a substance are not all the same — they follow a spread known as the Maxwell–Boltzmann distribution. At any given temperature, some particles move slowly (low energy), most have moderate energy, and a small number move very fast (high energy). Only the particles in the “tail” of the distribution — those with energy at or above Ea — are capable of reacting when they collide. This idea is the key to understanding both the effect of temperature and the effect of a catalyst:

• Raising the temperature shifts the distribution towards higher energies, so a much larger fraction of particles exceeds Ea.
• Adding a catalyst does not change the particles’ energies at all; instead it lowers Ea, so a larger fraction of the existing particles already qualifies as having enough energy.

How Collision Theory Explains Rate Factors

Every factor covered in Rates of Reaction can be explained using collision theory, by asking two questions: does it change the frequency of collisions? and does it change the fraction of collisions with enough energy?

Temperature

Increasing temperature increases the rate of reaction for two reasons at once, which is why temperature has such a dramatic effect:

1. Particles move faster, so they collide more frequently.
2. A much greater proportion of particles have energy ≥ Ea, so a higher fraction of collisions are successful.

The second effect is by far the more important one. A modest rise in temperature causes a large increase in the number of particles in the high-energy tail of the Maxwell–Boltzmann distribution. This is why, as a rough rule of thumb, a 10 °C rise in temperature approximately doubles the rate of many reactions — far more than the small increase in collision frequency alone could account for.

Concentration and Pressure

• Increasing the concentration of a solution means there are more reactant particles per unit volume. The particles are more crowded, so collisions happen more frequently, and the rate increases.
• Increasing the pressure of a gas has exactly the same effect: compressing a gas squeezes the same number of particles into a smaller volume, so there are more particles per unit volume and more frequent collisions.

Note that changing concentration or pressure does not change the energy of the particles — the fraction of successful collisions stays the same. Only the frequency of collisions increases.

Surface Area

For reactions involving a solid, only the particles on the surface of the solid can be collided with. Breaking a solid into smaller pieces (e.g. grinding lumps into powder) exposes more particles at the surface for the same total mass:

• More exposed particles → more frequent collisions between the solid and the surrounding liquid or gas particles → faster rate.

This is covered in detail in Surface Area and Rate. As with concentration, only the collision frequency changes — the energy of the collisions is unaffected.

Catalyst

A Catalyst speeds up a reaction without being used up in the process. Collision theory explains this differently from the other factors:

• A catalyst provides an alternative reaction pathway with a lower activation energy.
• Because Ea is lower, many more of the colliding particles now have energy ≥ Ea.
• Therefore a greater fraction of collisions are successful, and the rate increases.

On a Maxwell–Boltzmann distribution diagram, the curve itself does not move — instead, the Ea line moves to the left (to lower energy), so the shaded area representing “particles able to react” becomes much larger. On a reaction profile diagram, the catalysed route shows a lower “hump” between reactants and products, while the energies of the reactants and products (and therefore the overall energy change, ΔH) are unchanged.

Sources

• Cambridge IGCSE Chemistry (0620/0971) syllabus — Topic: Chemical reactions, rate of reaction.
• Edexcel International GCSE Chemistry (4CH1) specification — Topic: Rates of reaction.
• CGP IGCSE Chemistry Revision Guide — Rates of reaction and collision theory.
• BBC Bitesize GCSE Chemistry — “Collision theory and rates of reaction”.

Common Misconceptions

• “All collisions cause a reaction.” False — only collisions with energy ≥ Ea and the correct orientation are successful. The overwhelming majority of collisions are unsuccessful.
• “Increasing temperature only makes particles collide more often.” Incomplete — the bigger effect is that a far greater proportion of particles have energy above Ea. Both effects matter, but the energy effect dominates, which is why temperature changes rate so dramatically.
• “Increasing concentration gives particles more energy.” False — concentration only increases the frequency of collisions. The energy of individual particles depends on temperature, not concentration.
• “A catalyst gives particles more energy.” False — a catalyst lowers the activation energy by providing an alternative pathway. The particles’ energies are unchanged; more of them simply now exceed the (lower) barrier.
• “A catalyst is used up in the reaction.” False — a catalyst is chemically unchanged and can be recovered at the end; a small amount can catalyse a large quantity of reactants.
• “Catalysts change the amount of product or the overall energy change.” False — a catalyst changes only the rate; the yield and ΔH are unaffected.
• “Activation energy changes with temperature.” False — Ea is fixed for a given reaction pathway. Raising the temperature changes how many particles exceed Ea, not the value of Ea itself.

Related pages: Rates of Reaction | Activation Energy | Catalyst | Surface Area and Rate