# Collision Theory

> CIE A-Level Chemistry · Reaction kinetics (AS)
> Source: https://www.owlsprep.com/study/cie-9701-u8-collision-theory/

This sub-topic explains how particle collisions govern the rate of chemical reactions, introducing key requirements for successful reactions and linking particle properties to observable reaction rate changes.

**Prerequisites:** [Basic rate of reaction definitions and measurement](https://www.owlsprep.com/study/cie-9701-u8-rates-of-reaction-intro/)

## Learning objectives

- Explain the two requirements for a successful reaction collision
- Define activation energy and its role in determining reaction rate
- Predict how changing reaction conditions affects rate via collision theory
- Link collision theory to observed experimental rate trends

## Core Principles of Collision Theory

**Collision Theory** — A model that explains reaction rates in terms of collisions between reactant particles, which must meet two key criteria to produce products.

*Example:* Collision theory correctly predicts that higher concentrations increase reaction rate by increasing collisions per second.

For a chemical reaction to occur, reactant particles must first collide with each other. Not all collisions result in a reaction; only a small fraction of all collisions are successful, meaning they form products.

- Collisions must have **sufficient energy** equal to or greater than the activation energy ($E_a$) of the reaction
- Colliding particles must have the **correct orientation** relative to each other to allow breaking of old bonds and formation of new bonds

**Worked example:** A student observes that magnesium reacts faster with 1 mol/dm³ HCl than 0.1 mol/dm³ HCl. Explain this observation using collision theory.

1. Step 1: Relate concentration to number of particles per unit volume
2. Higher concentration of HCl means more H⁺ ions per unit volume in the mixture.
3. Step 2: Relate number of particles to collision frequency
4. More particles mean more frequent collisions between H⁺ ions and magnesium atoms per second.
5. Step 3: Relate collision frequency to successful collisions and rate
6. A higher frequency of collisions gives a higher frequency of successful collisions, so the reaction rate is faster.

> **Exam tip:** Always state both requirements for successful collisions when asked for an explanation.

## Activation Energy and Collision Energy

**Activation Energy** — The minimum total kinetic energy that colliding reactant particles must have for a collision to be successful and lead to a reaction.

*Notation:* $E_a$

*Example:* The activation energy for the combustion of methane is very high, so methane does not spontaneously react with air at room temperature.

Activation energy accounts for the energy required to break the existing bonds in reactant molecules before new bonds can form to make products. Only collisions where the combined kinetic energy of the particles is equal to or higher than $E_a$ can overcome this energy barrier and produce a reaction.

> **info**
>
> At room temperature, only around 1 in $10^{14}$ collisions between gas molecules have enough energy to be successful. This explains why most reactions do not happen instantly even at high concentrations.

**Worked example:** Explain why increasing the temperature of a reaction increases the rate according to collision theory.

1. Step 1: Increasing temperature increases the average kinetic energy of all reactant particles.
2. Step 2: A much larger proportion of all particles now have kinetic energy greater than or equal to the activation energy $E_a$.
3. Step 3: This leads to a large increase in the frequency of successful collisions per second, even though total collision frequency only increases slightly.
4. Step 4: The net effect is a large increase in overall reaction rate.

## Effect of Conditions on Collision Rate

Collision theory can be used to predict how changing common reaction conditions (concentration, pressure, surface area, temperature) changes reaction rate. Each condition affects either the total frequency of collisions or the proportion of collisions that have enough energy to be successful.

| Condition | Change | Effect on collisions | Effect on rate |
| --- | --- | --- | --- |
| Concentration | Increase | More particles per volume → higher collision frequency | Increases |
| Gas Pressure | Increase | Same particles in smaller volume → higher collision frequency | Increases |
| Solid Surface Area | Increase | More exposed particles → more collisions per second | Increases |
| Temperature | Increase | Higher particle energy → more collisions with ≥$E_a$ | Greatly increases |

**Worked example:** Powdered calcium carbonate reacts faster with HCl than large pieces of the same total mass. Explain this using collision theory.

1. Step 1: Large pieces of calcium carbonate only have their surface particles exposed to H⁺ ions from the acid; internal particles cannot collide with H⁺.
2. Step 2: When ground into powder, total exposed surface area increases, so many more calcium carbonate particles are available to collide with H⁺ ions.
3. Step 3: This increases the frequency of successful collisions per second, increasing the overall reaction rate.

## Common pitfalls

- **Wrong:** Claiming increasing temperature only increases collision frequency to increase rate
  - Why it fails: The main effect of temperature is increasing the proportion of collisions with energy ≥ $E_a$. The increase in total collision frequency is very small by comparison.
  - Correct: When explaining rate increase with temperature, emphasise the higher proportion of successful collisions with sufficient energy, not just higher collision frequency.
- **Wrong:** Claiming all collisions with enough energy are successful
  - Why it fails: Orientation of colliding particles is also required for successful bond breaking and formation. Even with enough energy, wrong alignment gives no reaction.
  - Correct: Always mention both sufficient energy (≥ $E_a$) and correct orientation as requirements for a successful collision.
- **Wrong:** Claiming increasing pressure increases the energy of individual gas particles, increasing rate
  - Why it fails: Increasing pressure only compresses the gas, increasing collision frequency. It does not change the average energy of particles unless temperature also increases.
  - Correct: Explain pressure effects on rate in terms of higher collision frequency from more particles per unit volume.
- **Wrong:** Claiming collision theory only applies to gas phase reactions
  - Why it fails: Collision theory can explain reaction rates for reactions in solution and heterogeneous reactions involving solids just as well.
  - Correct: Apply collision theory to all reaction types, adjusting for particle availability (e.g. surface area for solids).

## Cheatsheet

| Change in Condition | Collision Frequency | Proportion of Successful Collisions | Net Effect on Rate |
| --- | --- | --- | --- |
| Increase concentration | Increases | Unchanged | Increases |
| Increase gas pressure | Increases | Unchanged | Increases |
| Increase solid surface area | Increases | Unchanged | Increases |
| Increase temperature | Slightly increases | Greatly increases | Greatly increases |
| Add catalyst | Unchanged | Increases (lower $E_a$) | Increases |

## What's next

Collision theory is the foundational model for all rate behaviour in AS level reaction kinetics, and it forms the basis for understanding more advanced topics like the Maxwell-Boltzmann distribution and catalysis. You will use the core ideas of successful collisions, activation energy, and collision frequency to explain all rate changes you encounter, so mastering this topic is critical for answering exam questions on kinetics. Understanding how different conditions affect reaction rate via collisions also prepares you for calculating rate constants and investigating the effect of temperature on rate via the Arrhenius equation at A2 level.

- [Catalysis](https://www.owlsprep.com/study/cie-9701-u8-catalysis/)
- [Periodicity](https://www.owlsprep.com/study/cie-9701-u9-overview/)
- [Physical properties periodic trends](https://www.owlsprep.com/study/cie-9701-u9-physical-properties-periodic-trends/)

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