Collision Theory and Activation Energy — AQA A-level Chemistry

SPEC 3.1.5.1 • KINETICS

Collision Theory & Activation Energy

Understanding why reactions happen and what determines reaction rate

What is Reaction Rate?

Before we explore why reactions happen, we need to understand what we mean by how fast they happen. Some reactions are almost instantaneous (like neutralisation), while others take years (like rusting). Reaction rate tells us how quickly reactants are being converted into products.

Definition

Reaction rate is defined as the change in concentration (or amount) of a reactant or product per unit time.

A simple formula for finding the rate of a chemical reaction is:

Formula

Rate of reaction = \frac{amount of reactant used up or product formed}{time}

Quantity MeasuredUnitsRate Units

Concentrationmol dm⁻³mol dm⁻³ s⁻¹

Massgg s⁻¹ or g min⁻¹

Volume of gascm³cm³ s⁻¹ or cm³ min⁻¹

Key Concept

The goal of studying kinetics is to understand what factors affect reaction rate and why. Collision theory provides the explanation for how these factors work at the particle level.

Why Do Reactions Happen?

For a chemical reaction to occur, the reacting particles must first come into contact with each other. This seems obvious, but it raises an important question: if particles are constantly colliding in a mixture, why don't all collisions lead to a reaction?

Consider a mixture of hydrogen and oxygen gases at room temperature. The molecules are moving rapidly and colliding billions of times per second, yet the mixture remains stable indefinitely. Clearly, not every collision results in a reaction. This observation led scientists to develop collision theory.

Definition

Collision theory states that for a reaction to occur, particles must collide with sufficient energy and the correct orientation.

The key insight is that collisions are necessary but not sufficient for a reaction. Most collisions are unsuccessful because they fail to meet one or both of two critical requirements. Understanding these requirements helps us explain why some reactions are fast while others are slow, and how we can control reaction rates.

Key Term

The collision frequency is the number of collisions per unit time. When more collisions per unit time take place, the number of particles with energy greater than the Eₐ increases. This causes an increase in the rate of reaction.

Key Concept

A successful collision (one that leads to a reaction) must satisfy TWO conditions simultaneously. If either condition is not met, the particles simply bounce apart unchanged.

Two Conditions for Reaction

Every successful collision must meet both of these requirements:

Sufficient Energy

The colliding particles must have kinetic energy equal to or greater than the activation energy (Eₐ). This energy is needed to break existing bonds and begin forming new ones.

Correct Orientation

The particles must collide in a specific geometric arrangement so that the reactive parts of the molecules come into contact. Random orientations usually fail.

Think of it like a key fitting into a lock. The key must be pushed with enough force (sufficient energy) AND it must be inserted at the correct angle (correct orientation). If you push a key against the side of a lock, no amount of force will open it.

Exam Tip

When explaining why increasing temperature increases rate, you must mention BOTH conditions: more particles have energy ≥ Eₐ, AND there are more collisions per second. Don't forget the energy condition - this is where most marks come from.

Common Mistake

Students often write "particles collide more frequently" as a complete answer. This is insufficient - you must explain that more collisions have energy ≥ Eₐ, making more collisions successful.

Effective CollisionIneffective Collision

OrientationCorrect orientationWrong orientation

EnergyEnough energy (≥ Eₐ)Not enough energy

Chemical ReactionYes - products formNo - particles bounce off

Activation Energy

Definition

Activation energy (Eₐ) is the minimum energy that colliding particles must possess for a reaction to occur.

Why is activation energy needed? Even in exothermic reactions that release energy overall, bonds must be broken before new bonds can form. Breaking bonds requires energy input. The activation energy represents this initial "investment" of energy needed to start the bond-breaking process.

Imagine pushing a boulder over a hill. Even though the boulder will roll downhill and end up at a lower position (releasing energy), you must first push it up to the top of the hill. The height of the hill represents the activation energy - the energy barrier that must be overcome.

SymbolQuantityUnitTypical Values

EₐActivation energykJ mol⁻¹40 - 200 kJ mol⁻¹

ΔHEnthalpy changekJ mol⁻¹Variable (+ or −)

Key Concept

Activation energy is ALWAYS positive - it's an energy barrier that must be overcome. Even highly exothermic reactions have an activation energy. The sign of ΔH tells you whether energy is released or absorbed overall, but Eₐ is the hurdle that must be crossed first.

Examiner Tip

The activation energy is the energy needed to 'activate' the reactant particles in order for them to collide effectively and cause a chemical reaction. Remember: Eₐ is measured from reactants to the peak, NOT from reactants to products.

Enthalpy Profile: Exothermic Reaction

Enthalpy profile diagrams show how the enthalpy of the reacting system changes during a reaction. They allow us to visualise both the activation energy and the overall enthalpy change.

Key Features - Exothermic

Products are LOWER than reactants (energy released). ΔH is negative. The activation energy Eₐ is measured from reactants UP to the peak (transition state). The overall energy released equals the difference in height between reactants and products.

In exothermic reactions, the reactants are higher in energy than the products. This means the energy needed for the reactants to go over the energy barrier (the activation energy) is relatively small compared to endothermic reactions. Reactions with low activation energies tend to occur more readily at room temperature because many collisions will have sufficient energy.

Worked Example01

The diagram below shows the enthalpy profile for a reaction. The enthalpy of the reactants is +50 kJ mol⁻¹, the peak is at +120 kJ mol⁻¹, and the products are at −30 kJ mol⁻¹. Determine (a) the activation energy and (b) the enthalpy change for this reaction.

1

Activation energy is measured from reactants to the peak (transition state):

2

Eₐ = Peak − Reactants = 120 − 50 = 70 kJ mol⁻¹

3

Enthalpy change is measured from reactants to products:

4

ΔH = Products − Reactants = (−30) − (+50) = −80 kJ mol⁻¹

Answer: (a) Eₐ = 70 kJ mol⁻¹, (b) ΔH = −80 kJ mol⁻¹

Note: The negative ΔH confirms this is an exothermic reaction.

Enthalpy Profile: Endothermic Reaction

In an endothermic reaction, the products have MORE energy than the reactants. Energy is absorbed from the surroundings, so the products sit higher on the enthalpy diagram.

Key Features - Endothermic

Products are HIGHER than reactants (energy absorbed). ΔH is positive. The activation energy is still measured from reactants to peak, and is typically larger than for exothermic reactions (though this varies).

In endothermic reactions, the reactants are lower in energy than the products. The energy needed for the reactants to go over the energy barrier is therefore relatively large. This is why endothermic reactions often require heating to proceed - more particles need to gain sufficient kinetic energy to overcome the larger activation energy barrier.

Important

Even though particles collide with each other in the same orientation, if they don't possess a minimum energy that corresponds to the Eₐ of that reaction, the reaction will not take place. Both conditions must be satisfied simultaneously.

Worked Example02

A reaction has activation energy Eₐ = 85 kJ mol⁻¹ and enthalpy change ΔH = +45 kJ mol⁻¹. If the enthalpy of the reactants is arbitrarily set to 0 kJ mol⁻¹, determine the enthalpy of (a) the transition state and (b) the products.

1

The transition state is Eₐ above the reactants:

2

Transition state = Reactants + Eₐ = 0 + 85 = +85 kJ mol⁻¹

3

Products are ΔH above reactants (positive ΔH means endothermic):

4

Products = Reactants + ΔH = 0 + 45 = +45 kJ mol⁻¹

Answer: (a) Transition state = +85 kJ mol⁻¹, (b) Products = +45 kJ mol⁻¹

The Transition State

The peak of the enthalpy profile represents the transition state (sometimes called the activated complex). This is a high-energy, unstable arrangement where old bonds are partially broken and new bonds are partially formed.

Definition

The transition state is the highest energy arrangement of atoms during a reaction. It represents the point of maximum instability where the system can either proceed to products or revert to reactants.

The transition state exists only momentarily - it cannot be isolated or observed directly. It represents the "point of no return" for the reaction. Once particles reach this state, they are equally likely to form products or reform reactants.

Important Distinction

The transition state is NOT an intermediate. An intermediate is a species that exists (briefly) between two transition states and has a defined structure. The transition state is a single point of maximum energy with partially formed bonds.

The Orientation Factor

Even if particles collide with enough energy, the reaction will only occur if they collide in the correct orientation. The reactive parts of the molecules must come into contact for bonds to break and form.

Definition

An ineffective collision is when particles collide in the wrong orientation or when they don't have enough energy and bounce off each other without causing a chemical reaction.

Consider the reaction H₂ + Cl₂ → 2HCl. For this reaction to occur, the molecules must collide so that H atoms can bond with Cl atoms. The diagram below shows both unsuccessful and successful collisions:

(A) shows an ineffective collision due to the particles not having enough energy, whereas (B) shows an effective collision where the particles have the correct orientation and enough energy for a chemical reaction to take place.

Exam Tip

For complex molecules with many atoms, only a small proportion of collisions have the correct orientation. This is why reactions between complex molecules are often slower than those between simple molecules, even at the same temperature.

Worked Example03

Explain why increasing the concentration of reactants increases the rate of reaction, with reference to collision theory.

1

Increasing concentration means there are more particles per unit volume.

2

With more particles in the same space, collisions between reactant particles occur more frequently.

3

The proportion of collisions with energy ≥ Eₐ and correct orientation remains the same.

4

However, since total collisions increase, the number of successful collisions per unit time also increases.

5

Therefore, the rate of reaction increases.

Key point: Concentration affects collision frequency, not the energy or orientation of individual collisions.

Worked Example04

Two reactions have the same enthalpy change, ΔH = −50 kJ mol⁻¹. Reaction A has Eₐ = 25 kJ mol⁻¹ and Reaction B has Eₐ = 120 kJ mol⁻¹. At room temperature, which reaction will be faster? Explain your answer using collision theory.

1

Reaction A will be faster.

2

For a reaction to occur, colliding particles must have kinetic energy ≥ Eₐ.

3

At room temperature, particles have a range of kinetic energies (Maxwell-Boltzmann distribution).

4

With Eₐ = 25 kJ mol⁻¹ (Reaction A), a larger proportion of particles have sufficient energy to react.

5

With Eₐ = 120 kJ mol⁻¹ (Reaction B), only a small fraction of particles have enough energy.

6

More successful collisions per unit time in A → faster rate.

Key point: ΔH tells us about energy released/absorbed overall, but Eₐ determines how fast the reaction occurs. Low Eₐ = fast reaction.

Why Activation Energy Matters

The size of the activation energy has a profound effect on reaction rate. A reaction with a low Eₐ proceeds quickly because many particles have enough energy to react. A reaction with a high Eₐ proceeds slowly because only a small fraction of particles can overcome the energy barrier.

Low EₐHigh Eₐ

Many particles have E ≥ EₐFew particles have E ≥ Eₐ

High proportion of successful collisionsLow proportion of successful collisions

Fast reaction rateSlow reaction rate

Example: Neutralisation reactionsExample: Rusting of iron

Connection to Next Lesson

In the next lesson, we'll explore the Maxwell-Boltzmann distribution, which shows exactly how molecular energies are distributed in a sample. This will help explain quantitatively how temperature changes affect the proportion of molecules with E ≥ Eₐ.

Common Mistake

Don't confuse activation energy with enthalpy change. A reaction can be highly exothermic (large negative ΔH) but still have a high Eₐ and proceed slowly. The two quantities are independent.

Exam Tip

When asked "why does increasing temperature increase rate?", always mention: (1) particles move faster so collide more frequently, AND (2) a greater proportion of collisions have energy ≥ Eₐ. The second point is more important and carries more marks.

Practice

Question 1: An enthalpy profile diagram shows the reactants at an enthalpy of +30 kJ mol⁻¹, a peak (transition state) at +145 kJ mol⁻¹, and products at −40 kJ mol⁻¹. Calculate: (a) the activation energy, (b) the enthalpy change ΔH, and (c) state whether the reaction is exothermic or endothermic.

a)

Eₐ = Peak − Reactants = 145 − 30 = 115 kJ mol⁻¹

b)

ΔH = Products − Reactants = (−40) − 30 = −70 kJ mol⁻¹

c)

ΔH is negative, so the reaction is exothermic (products are at lower enthalpy than reactants).

Practice

Question 2: State the two conditions that must be met for a collision between reactant particles to result in a reaction. Explain why both conditions are necessary.

1

Sufficient energy: Particles must have kinetic energy ≥ activation energy (Eₐ). This energy is needed to break existing bonds and initiate the reaction.

2

Correct orientation: Particles must collide so that reactive parts of molecules come into contact. Without correct geometry, bonds cannot break and form appropriately.

Both are necessary because even high-energy collisions fail if molecules are misaligned, and correct orientation is useless if there isn't enough energy to break bonds.

Practice

Question 3: From an enthalpy profile, the peak is at +180 kJ mol⁻¹, reactants at +60 kJ mol⁻¹, and products at +95 kJ mol⁻¹. Determine Eₐ and ΔH, and state whether the reaction is exothermic or endothermic.

1

Eₐ = Peak − Reactants = 180 − 60 = 120 kJ mol⁻¹

2

ΔH = Products − Reactants = 95 − 60 = +35 kJ mol⁻¹

3

Since ΔH is positive (products higher than reactants), this is an endothermic reaction.

Summary

Key Concepts

TermDefinition

Collision theoryParticles must collide with sufficient energy and correct orientation for reaction

Activation energy (Eₐ)Minimum energy required for a reaction to occur (always positive)

Transition stateHighest energy point; unstable arrangement with partial bonds

Successful collisionCollision meeting both energy and orientation requirements

Key Exam Skills

Calculate Eₐ and ΔH from enthalpy profile diagrams (Eₐ = peak − reactants, ΔH = products − reactants)

Sketch enthalpy profiles for exothermic and endothermic reactions

Explain rate changes using collision theory (mention BOTH energy AND frequency)

Describe the role of orientation using molecular diagrams