5.2 Collision Theory and Activation Energy

Introduction

In a chemical reaction, bonds between atoms are broken and new bonds are formed. But how do these reactions actually happen? Why do some collisions between particles lead to a reaction while others do not?

Collision theory helps us understand how chemical reactions happen. It explains that for a reaction to occur, particles must collide with each other — but not every collision leads to a reaction. The collision must meet certain conditions.

Definition

Collision Theory: Collision theory is a model that explains how chemical reactions occur. It states that for a reaction to happen, the reacting particles must collide with each other, and the collision must have sufficient energy and correct orientation.

Activation Energy: The smallest amount of energy needed for a reaction to happen is called activation energy.

In simple words:

• Collision theory = particles must collide to react.
• Activation energy = the minimum energy needed for a collision to cause a reaction.

Main Idea

For a reaction to occur, the atoms or molecules must collide with each other. However, not all collisions lead to a reaction. The success of a collision depends on two things:

• Energy — the colliding particles must have enough energy to overcome the repulsion between their outer electrons.
• Orientation — the particles must be aligned in a way that allows the necessary atoms to come together and form new bonds.

Key point:

• Not every collision leads to a reaction.
• Only collisions with sufficient energy and correct orientation are effective collisions.

Core Concept

Key point: Only effective collisions — those with sufficient energy and correct orientation — lead to a reaction.

Detailed Explanation

What is Collision Theory?

Collision theory helps us understand how chemical reactions happen. For a reaction to occur, the atoms or molecules must collide with each other.

However, not all collisions lead to a reaction. The success of a collision depends on two things:

1. Energy:

• The colliding particles must have enough energy to overcome the repulsion between their outer electrons.
• If the energy of the colliding particles is less than the activation energy, the reaction will not occur.

2. Orientation:

• The particles must be aligned in a way that allows the necessary atoms to come together and form new bonds.
• If the orientation is wrong, the collision will not lead to a reaction.

What is Activation Energy?

The smallest amount of energy needed for a reaction to happen is called activation energy.

• If the energy of the colliding particles is less than this activation energy, the reaction will not occur.
• Reactions happen faster when more particles have enough energy to collide effectively.
• If the activation energy is high, fewer particles will have enough energy, and the reaction will be slower.
• If the activation energy is low, more particles can collide effectively, and the reaction will happen faster.

Simple explanation: Think of activation energy like a hill. The reacting molecules need to climb this hill (using energy) before they can roll down the other side and form products. If the molecules don't have enough energy to reach the top, they roll back and remain unchanged.

Example: Reaction Between Hydrogen and Chlorine

For example, when hydrogen (H₂) and chlorine (Cl₂) react to form hydrogen chloride (HCl), the molecules need enough energy to break their bonds and form new ones.

The reaction:

H₂ + Cl₂ → 2HCl

What happens:

• H₂ and Cl₂ molecules collide.
• If the collision has enough energy, the bonds in H₂ and Cl₂ break.
• A temporary, high-energy state called the activated complex (or transition state) is formed.
• This complex quickly breaks down to form the final products (HCl).

Activated Complex: When an effective collision occurs, the particles form a temporary, high-energy state called the activated complex (or transition state). This complex quickly breaks down to form the final products.

Effective Collision

An effective collision is a collision that:

• Has enough energy (≥ activation energy).
• Has the correct orientation.

Only effective collisions lead to a reaction.

Key point: Not all collisions are effective. Only those with sufficient energy and correct orientation result in a reaction.

Visualising Activation Energy

You can think of activation energy like a hill:

• The reacting molecules need to climb this hill (using energy) before they can roll down the other side and form products.
• If the molecules don't have enough energy to reach the top, they roll back and remain unchanged.
• If they have enough energy, they reach the top (activated complex) and then roll down to form products.

Summary Table

Intermediate Level

Conditions for an effective collision:

• Energy: The colliding particles must have energy equal to or greater than the activation energy.
• Orientation: The particles must be aligned correctly so that the necessary atoms can come together and form new bonds.

Activation energy:

• The minimum energy required for a reaction to occur.
• High activation energy → slower reaction.
• Low activation energy → faster reaction.

Activated complex:

• A temporary, high-energy state formed during an effective collision.
• It quickly breaks down to form the final products.

Key point:

• Not all collisions are effective.
• Only collisions with sufficient energy and correct orientation lead to a reaction.

Example: H₂ + Cl₂ → 2HCl

• The collision must have enough energy to break the H–H and Cl–Cl bonds.
• The orientation must allow the H and Cl atoms to come together.

Important Points

• - Collision theory = particles must collide to react.
• - Not all collisions lead to a reaction.
• - An effective collision must have:
• 1. Enough energy (≥ activation energy)
• 2. Correct orientation
• - Activation energy = minimum energy needed for a reaction.
• - If energy < activation energy → no reaction.
• - High activation energy → slower reaction.
• - Low activation energy → faster reaction.
• - Activated complex = temporary, high-energy state formed during an effective collision.
• - Reactions happen faster when more particles have enough energy to collide effectively.

Remember

• - Collision theory = particles must collide to react.
• - Effective collision = enough energy + correct orientation.
• - Activation energy = minimum energy needed.
• - Energy < activation energy → no reaction.
• - High activation energy → slower reaction.
• - Low activation energy → faster reaction.
• - Activated complex = temporary, high-energy state.
• - More effective collisions → faster reaction.

Common Mistakes

Quick Revision

Collision Theory and Activation Energy:

In a chemical reaction, bonds between atoms are broken and new bonds are formed. Collision theory helps us understand how these reactions happen. For a reaction to occur, the atoms or molecules must collide with each other. However, not all collisions lead to a reaction. The success of a collision depends on two things:

• Energy: The colliding particles must have enough energy to overcome the repulsion between their outer electrons.
• Orientation: The particles must be aligned in a way that allows the necessary atoms to come together and form new bonds.

The smallest amount of energy needed for a reaction to happen is called activation energy. If the energy of the colliding particles is less than this activation energy, the reaction will not occur.

Reactions happen faster when more particles have enough energy to collide effectively. If the activation energy is high, fewer particles will have enough energy, and the reaction will be slower. If the activation energy is low, more particles can collide effectively, and the reaction will happen faster.

For example, when hydrogen (H₂) and chlorine (Cl₂) react to form hydrogen chloride (HCl), the molecules need enough energy to break their bonds and form new ones. When an effective collision occurs, the particles form a temporary, high-energy state called the activated complex (or transition state). This complex quickly breaks down to form the final products.

You can think of activation energy like a hill. The reacting molecules need to climb this hill (using energy) before they can roll down the other side and form products. If the molecules don't have enough energy to reach the top, they roll back and remain unchanged.

60-Second Revision

• Collision theory = particles must collide to react.
• Effective collision = enough energy + correct orientation.
• Activation energy = minimum energy needed for a reaction.
• Energy < activation energy → no reaction.
• High activation energy → slower reaction.
• Low activation energy → faster reaction.
• Activated complex = temporary, high-energy state.
• More effective collisions → faster reaction.
• H₂ + Cl₂ → 2HCl is an example.
• ⚠️ Warning: not all collisions lead to a reaction — only effective collisions do.