Collision theory and reaction rates

IB Chemistry SL· Reactivity 2.2: Collision theory· 35 min read

1. Key Principles of Collision Theory★★☆☆☆⏱ 10 min

📘 Definition

Collision Theory

A qualitative model that explains reaction rates based on collisions between reactant particles. Only collisions that meet specific criteria lead to product formation.

For a chemical reaction to occur, reactant particles must first collide with one another. However, not all collisions produce new products. Only a small fraction of all collisions are classified as successful, meaning they result in chemical change.

Colliding particles must have total kinetic energy equal to or greater than the activation energy ( $E_{a}$ ) of the reaction.
Colliding particles must have the correct spatial orientation to allow breaking of old bonds and formation of new bonds.

📐 Worked Example

Explain why two hydrogen iodide (HI) molecules may collide but not react to form $H_{2}$ and $I_{2}$ .

1
Recall the two requirements for a successful collision.
2
One possible reason: the collision has less kinetic energy than the activation energy for the reaction. Even with correct orientation, insufficient energy cannot break the H-I bonds to form products.
3
A second possible reason: the collision has enough energy but incorrect orientation. If the hydrogen end of one HI does not collide with the hydrogen end of the other HI, bonds cannot rearrange to form $H_{2}$ .

Exam tip:

Always mention both energy and orientation when explaining unsuccessful collisions to earn full marks.

2. Activation Energy★★☆☆☆⏱ 10 min

📘 Definition

Activation Energy

$E_{a}$

The minimum total kinetic energy that colliding reactant particles must have to undergo a successful reaction. It acts as an energy barrier that only some collisions can overcome.

Example:

Activation energy for decomposition of $H_{2} O_{2}$ is ~75 kJ mol⁻¹ at 298 K.

When colliding particles meet both requirements for a successful reaction, they form an unstable, high-energy intermediate state called the activated complex (or transition state) before rearranging into final products. The activation energy is the energy difference between the reactants and this activated complex.

📐 Worked Example

Reaction A has an activation energy of 40 kJ mol⁻¹, and Reaction B has an activation energy of 120 kJ mol⁻¹, at the same temperature and equal concentration. Which reaction is faster? Explain using collision theory.

1
Recall that lower activation energy means a larger fraction of collisions meet the energy requirement for success.
2
Reaction A has a lower activation energy, so more collisions per unit time have energy ≥ $E_{a}$ .
3
A higher frequency of successful collisions means a faster overall reaction rate, so Reaction A is faster.

3. Collision Theory and Rate Factors★★★☆☆⏱ 15 min

Collision theory can be used to explain how every common factor affecting reaction rate changes the frequency of successful collisions, which directly determines overall reaction rate.

Concentration/Pressure: Higher concentration (solutions) or pressure (gases) increases particles per unit volume, increasing total collision frequency and therefore successful collision frequency.
Temperature: Higher temperature increases average kinetic energy, greatly increasing the fraction of particles with energy ≥ $E_{a}$ , leading to a large increase in successful collisions.
Surface Area: Larger surface area of solid reactants exposes more particles to collision, increasing total and successful collision frequency.

📐 Worked Example

Use collision theory to explain why increasing temperature increases reaction rate much more than increasing concentration.

1
Increasing concentration only increases total collision frequency, so successful collisions increase proportionally to the change in concentration.
2
Increasing temperature causes two changes: a small increase in total collision frequency, and a much larger increase in the fraction of collisions that meet or exceed activation energy.
3
Since normally only a small fraction of collisions are successful, a large increase in this fraction leads to a much larger overall increase in reaction rate than a comparable increase in concentration.

Exam tip:

You must explicitly link any factor change to the frequency of successful collisions (not just total collisions) to earn full marks.

4. Maxwell-Boltzmann Distribution and Collision Theory★★★☆☆⏱ 10 min

The Maxwell-Boltzmann distribution describes the spread of kinetic energies across particles in a sample at a given temperature. The area under the curve to the right of the activation energy line equals the fraction of particles that have enough energy to react.

📐 Worked Example

Explain how the fraction of successful collisions changes when temperature increases, using a Maxwell-Boltzmann curve.

1
Draw the skewed Maxwell-Boltzmann curve, with kinetic energy on the x-axis. Add a vertical line for activation energy $E_{a}$ to the right of the curve's peak.
2
The area to the right of $E_{a}$ is the fraction of particles with enough energy for successful collisions.
3
When temperature increases, the curve shifts right and flattens. The area to the right of $E_{a}$ increases significantly, so the fraction of successful collisions increases, leading to a faster reaction rate.

5. Common Pitfalls

Wrong move:

Forgetting to mention orientation as a requirement for successful collisions.

Why:

Exam questions expect you to state both conditions (energy and orientation) for full marks.

Correct move:

Always confirm that both energy ≥ activation energy and correct orientation are required for a successful collision.

Wrong move:

Referring to an increase in total collisions instead of successful collisions when explaining rate changes.

Why:

More total collisions do not guarantee more successful collisions, so this answer is incomplete.

Correct move:

Explicitly link the change in conditions to an increase or decrease in the frequency of successful collisions.

Wrong move:

Claiming that increasing temperature decreases the activation energy of a reaction.

Why:

Activation energy is a fixed property of a given reaction and does not change with temperature.

Correct move:

State that increasing temperature increases the fraction of particles with energy equal to or greater than $E_{a}$ , which increases reaction rate.

Wrong move:

Confusing activation energy with the enthalpy change of a reaction.

Why:

Enthalpy change is the total energy difference between reactants and products, not the energy barrier.

Correct move:

Activation energy is the energy difference between reactants and the transition state (energy peak), not reactants and products.

6. Quick Reference Cheatsheet

Concept	Key Fact
Successful collision	Energy ≥ $E_{a}$ ; 2. Correct orientation
Activation energy	Fixed for a reaction; minimum energy for reaction
Higher concentration	More collisions per unit time → more successful collisions
Higher temperature	More particles ≥ $E_{a}$ → large rate increase
Lower activation energy	Higher fraction of successful collisions → faster rate

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2022 · Paper 1
Conditions for successful collisions
2023 · Paper 2
Explain temperature effect via collision theory
2021 · Paper 1
Activation energy definition

Going deeper

syllabusIB Chemistry SL Syllabus: Reactivity 2

What's Next

Collision theory is the foundational qualitative model for all further study of chemical kinetics in IB Chemistry SL. It provides the framework for understanding how reaction conditions change rate, and underlies quantitative concepts like the rate law and Arrhenius equation that you will explore next. Mastery of collision theory is required for exam questions on enzyme activity, catalysis, industrial reaction design, and energy profiles, which are common core topics in SL paper 1 and paper 2. Building on this model will let you connect qualitative rate changes to quantitative calculations and predict the behavior of multi-step reactions.