# Collision theory

> IB Chemistry HL · R2: How much / how fast / how far?
> Source: https://www.owlsprep.com/study/ib-chemistry-hl-u5-collision-theory/

This subtopic explains the requirements for chemical reactions to occur between reactant particles, and how changing reaction conditions alters reaction rate according to collision theory. It forms the foundation for all kinetics study in IB Chemistry HL.

**Prerequisites:** [Introduction to reaction kinetics and rate definitions](https://www.owlsprep.com/study/ib-chemistry-hl-u5-introduction-to-kinetics/)

## Learning objectives

- State the core postulates of collision theory
- Explain the effect of changing reaction conditions on rate via collision theory
- Distinguish between successful and unsuccessful collisions
- Explain how catalysts increase reaction rate using collision theory

## Core Postulates of Collision Theory

Collision theory is a simple kinetic model developed to explain reaction rates, originally for gas-phase reactions but applicable to almost all reaction systems. It makes testable predictions about how reaction conditions change rate, which match experimental observations.

**Collision Theory** — A model that links reaction rate to the frequency and properties of collisions between reactant particles.

*Example:* It correctly predicts that increasing temperature always increases reaction rate for non-catalyzed reactions.

1. For a reaction to occur, reactant particles must first collide with one another.
2. Only collisions with energy equal to or greater than the reaction's activation energy ($E_a$) can produce a reaction.
3. Colliding particles must also approach each other with the correct orientation to break existing bonds and form new product bonds.

**Worked example:** State which collision between two HI molecules will result in a successful reaction, given the activation energy for $2HI \rightarrow H_2 + I_2$ is 235 kJ mol⁻¹: (a) 250 kJ mol⁻¹, correct orientation, (b) 200 kJ mol⁻¹, correct orientation, (c) 250 kJ mol⁻¹, incorrect orientation

1. Recall the three requirements for successful collision: collision occurs, energy ≥ $E_a$, correct orientation.
2. Evaluate option (a): Energy (250 kJ mol⁻¹) > 235 kJ mol⁻¹, orientation is correct. This collision is successful.
3. Evaluate option (b): Energy (200 kJ mol⁻¹) < activation energy. Even with correct orientation, no reaction occurs.
4. Evaluate option (c): Energy meets the requirement, but orientation is incorrect. Bonds cannot break to form product, so collision is unsuccessful.
5. Conclusion: Only collision (a) is successful.

> **Exam tip:** Always mention both energy and orientation when explaining successful collisions — exam markers require both points for full marks.

## Effect of Reaction Conditions on Rate

Collision theory can explain how changing common reaction conditions alters reaction rate, which is a very common exam question.

**Successful Collision** — A collision that meets all three requirements to form new product molecules.

**Worked example:** Use collision theory to explain why increasing concentration of a reactant increases reaction rate at constant temperature.

1. Increasing concentration means more reactant particles per unit volume of the reaction mixture.
2. More particles per unit volume leads to a higher frequency of collisions between reactant particles per second.
3. At constant temperature, the proportion of collisions with energy ≥ $E_a$ stays the same, so higher collision frequency gives more successful collisions per second.
4. More successful collisions per unit time equals a higher reaction rate.

- **Temperature**: The dominant effect of increasing temperature is that a much greater proportion of particles have energy ≥ $E_a$, leading to more successful collisions per second. There is also a small increase in collision frequency.
- **Surface area (solid reactants)**: Breaking a solid into smaller pieces increases the surface area of reactant exposed to other reactants, increasing collision frequency and rate.
- **Pressure (gases)**: Increasing pressure compresses gas into a smaller volume, which increases concentration, collision frequency and rate.

> **Exam tip:** When explaining temperature effects, emphasize the change in proportion of particles with energy ≥ $E_a$, not just more collisions. The proportion change is the dominant effect.

## Catalysts and Activation Energy

Catalysts increase reaction rate without being consumed in the reaction, and their action is easily explained by collision theory principles.

**Activation Energy** — The minimum energy that colliding reactant particles must have to form products in a successful collision.

*Notation:* $E_a$

**Worked example:** Use collision theory to explain why a catalyst increases reaction rate at the same temperature.

1. A catalyst does not change the energy of reactant particles, instead it provides an alternative reaction mechanism (pathway) for the reaction.
2. This alternative pathway has a lower activation energy than the uncatalyzed pathway:
3. $$E_{a(\text{catalyzed})} < E_{a(\text{uncatalyzed})}$$
4. At a given temperature, a much larger proportion of reactant particles have energy equal to or greater than the lower activation energy.
5. This leads to a higher frequency of successful collisions per second, so reaction rate increases, without changing temperature or total collision frequency.

> **note**
>
> Catalysts do not change the enthalpy change of the reaction, or the position of equilibrium for reversible reactions. They only increase the rate at which equilibrium is reached.

## Common pitfalls

- **Wrong:** Claiming that increasing temperature increases activation energy.
  - Why it fails: Activation energy is a fixed property of the reaction pathway, it does not change with temperature.
  - Correct: State that increasing temperature increases the proportion of particles with energy ≥ $E_a$.
- **Wrong:** Forgetting to mention orientation when explaining successful collisions.
  - Why it fails: Most IB exam mark schemes require both energy and orientation conditions for full marks.
  - Correct: Always explicitly state that successful collisions require both sufficient energy (≥ $E_a$) and correct orientation.
- **Wrong:** Claiming that catalysts work by increasing the frequency of collisions.
  - Why it fails: Catalysts do not change the total number of collisions between reactant particles, only the pathway.
  - Correct: Explain that catalysts provide an alternative lower activation energy pathway, increasing the proportion of successful collisions.
- **Wrong:** Stating that increasing concentration increases the proportion of particles with energy ≥ $E_a$.
  - Why it fails: At constant temperature, the energy distribution of particles does not change when concentration increases.
  - Correct: Explain that higher concentration increases collision frequency, leading to more successful collisions per second.

## Cheatsheet

| Condition Change | Key Change per Collision Theory | Effect on Rate |
| --- | --- | --- |
| Increase concentration | Higher collision frequency | Increase |
| Increase temperature | Higher proportion of collisions ≥ $E_a$ | Increase |
| Increase surface area | Higher collision frequency | Increase |
| Increase pressure (gases) | Higher collision frequency | Increase |
| Add catalyst | Alternative lower $E_a$ pathway | Increase |

## What's next

Collision theory is the core foundation for all further study of reaction kinetics in IB Chemistry HL. You will next build on this model to explore the Arrhenius equation, which quantifies the relationship between temperature, activation energy and the rate constant. Understanding collision theory is also critical for interpreting reaction profile diagrams and studying the mechanisms of organic reactions, and it is assessed in almost every kinetics exam question.

- [Dynamic equilibrium](https://www.owlsprep.com/study/ib-chemistry-hl-u5-dynamic-equilibrium/)
- [The equilibrium constant](https://www.owlsprep.com/study/ib-chemistry-hl-u5-the-equilibrium-constant/)

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