Collision Theory

Collision theory is the model that chemists use to explain why reactions happen and why changing conditions can make them faster or slower. This theory underpins every rate-of-reaction topic in the Edexcel GCSE Chemistry (1CH0) specification. In this lesson you will learn the conditions needed for a successful collision, what activation energy means, and how collision theory explains the effects of different factors on the rate of reaction.

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The Basic Idea

For a chemical reaction to take place, the reacting particles must collide with each other. However, not every collision leads to a reaction. Two conditions must be met for a collision to be successful (i.e. for it to result in a reaction):

1. The particles must collide with sufficient energy — this minimum energy is called the activation energy.
2. The particles must collide with the correct orientation — the reactive parts of the molecules must be facing each other.

If either condition is not met, the particles simply bounce apart and no reaction occurs.

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Activation Energy

The activation energy (symbol Eₐ) is the minimum amount of energy that colliding particles must have in order for a reaction to occur.

• Every reaction has its own specific activation energy.
• If particles collide with energy less than the activation energy, they bounce off each other unchanged — this is an unsuccessful collision.
• If particles collide with energy equal to or greater than the activation energy, bonds can break and new bonds form — this is a successful collision and a reaction takes place.

Exam tip: The activation energy is not the total energy of the reaction. It is the minimum energy barrier that must be overcome for the reaction to begin. Make sure you use the term "activation energy" in your answers when explaining rate changes — this is a key term that examiners look for.

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Successful vs Unsuccessful Collisions

The following diagram shows the difference between successful and unsuccessful collisions:

flowchart TD
    A[Two particles approach each other] --> B{Do they collide?}
    B -->|No| C[No reaction — particles pass by]
    B -->|Yes| D{Is the energy ≥ activation energy?}
    D -->|No| E[Unsuccessful collision — particles bounce apart unchanged]
    D -->|Yes| F{Is the orientation correct?}
    F -->|No| G[Unsuccessful collision — particles bounce apart unchanged]
    F -->|Yes| H[Successful collision — reaction occurs, products formed]
    style H fill:#66cc66,stroke:#228B22
    style E fill:#ff9999,stroke:#cc0000
    style G fill:#ff9999,stroke:#cc0000
    style C fill:#cccccc,stroke:#999999

Key Points from the Diagram

• Particles must first collide — if they miss each other, nothing happens.
• Even if they collide, they need enough energy (≥ Eₐ) to react.
• Even with enough energy, the orientation must be correct.
• Only when all three conditions are met does a successful collision occur.

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Why Not All Collisions Are Successful

In a container of reacting particles, there are billions of collisions happening every second. However, only a small fraction of these collisions result in a reaction because:

• Most particles do not have enough kinetic energy to overcome the activation energy barrier.
• Even when particles do have sufficient energy, they may collide at the wrong angle (incorrect orientation).

This is why reactions do not happen instantaneously, even when all the reactants are present.

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How Collision Theory Explains Factors Affecting Rate

Collision theory provides a framework for understanding why changing certain conditions affects the rate of reaction. The rate of reaction depends on two things:

1. The frequency of collisions — how often particles collide per second
2. The proportion of successful collisions — what fraction of those collisions have enough energy and the correct orientation

Any change that increases the frequency of collisions or increases the proportion of successful collisions will increase the rate of reaction.

Factor	How It Increases the Rate (Collision Theory Explanation)
Increasing temperature	Particles move faster → more frequent collisions AND a greater proportion have energy ≥ Eₐ → more successful collisions
Increasing concentration	More particles in a given volume → more frequent collisions → more successful collisions per second
Increasing pressure (gases)	Particles are closer together → more frequent collisions → more successful collisions per second
Increasing surface area	More of the solid reactant is exposed → more frequent collisions between the solid and the other reactant
Using a catalyst	Provides an alternative pathway with a lower activation energy → a greater proportion of collisions are successful

Exam tip: When answering questions about factors affecting rate, always link your answer back to collision theory. Simply stating "the reaction goes faster" is not enough. You need to explain why it goes faster using the ideas of collision frequency and activation energy.

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Energy Distribution of Particles

At any given temperature, the particles in a substance have a range of kinetic energies. Some particles are moving slowly (low energy), most have a moderate energy, and a few are moving very fast (high energy).

This distribution of energies is important because it determines how many particles have enough energy to overcome the activation energy barrier at any given moment.

When the temperature increases:

• The average kinetic energy of the particles increases.
• The distribution shifts so that a greater proportion of particles have energy ≥ Eₐ.
• This means more collisions are successful, increasing the rate.

Exam tip: At Higher tier, you may be asked to explain the effect of temperature on the energy distribution of particles. Remember two effects: (1) more frequent collisions and (2) a greater proportion of particles exceed the activation energy. The second effect is actually the more important one.

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Applying Collision Theory: A Worked Explanation

Question: Explain, using collision theory, why increasing the concentration of hydrochloric acid increases the rate of its reaction with magnesium ribbon.

Model answer:

1. Increasing the concentration of HCl means there are more HCl particles (H⁺ and Cl⁻ ions) in the same volume of solution.
2. The magnesium atoms on the surface of the ribbon are therefore surrounded by more acid particles.
3. This leads to more frequent collisions between H⁺ ions and the magnesium surface per second.
4. A greater number of these collisions will have energy equal to or greater than the activation energy.
5. Therefore there are more successful collisions per second, and the rate of reaction increases.

Exam tip: When writing an extended answer about collision theory, structure your response in a logical chain: change → effect on particles → effect on collision frequency → effect on successful collisions → conclusion about rate. This step-by-step approach ensures you pick up all the available marks.

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Reaction Profiles and Activation Energy