# Le Chatelier's principle

> CIE A-Level Chemistry · 9701 AS Level
> Source: https://www.owlsprep.com/study/cie-9701-u7-le-chatelier-s-principle/

This module explains Le Chatelier's principle for predicting how dynamic chemical equilibria respond to external changes in concentration, pressure, and temperature. You will apply the principle to common exam problems and industrial process scenarios.

**Prerequisites:** [Dynamic equilibrium](https://www.owlsprep.com/study/cie-9701-u7-dynamic-equilibrium/)

## Learning objectives

- State Le Chatelier's principle correctly and apply it to external changes
- Predict equilibrium shifts for changes in concentration, pressure and temperature
- Explain the effect of catalysts on equilibrium position
- Apply the principle to industrial equilibrium processes

## Core Definition of Le Chatelier's Principle

Le Chatelier's principle is a simple predictive rule that describes how a system at dynamic equilibrium responds to external changes to conditions. It does not explain the underlying kinetics or thermodynamics of the shift, only what the final change to equilibrium position will be.

**Le Chatelier's Principle** — When a system at dynamic equilibrium is subjected to an external change in conditions, the position of equilibrium shifts in the direction that counteracts the applied change.

*Example:* An increase in temperature adds heat to the system, so equilibrium shifts in the endothermic direction to absorb the extra heat.

**Worked example:** State what Le Chatelier's principle predicts when the total pressure of a gaseous equilibrium system is decreased.

1. 1. Identify the external change: the total pressure of the system has decreased.
2. 2. The system will shift to counteract the decrease, so it will act to increase the total pressure.
3. 3. Higher pressure comes from a greater number of gaseous moles in the system.
4. 4. Conclusion: equilibrium shifts to the side of the reaction with more moles of gas.

> **Exam tip:** Remember the principle states the shift counteracts, not completely cancels, the external change.

## Effect of Concentration Changes

Changing the concentration of any reactant or product in an equilibrium system will shift the position of equilibrium, but it does not change the value of the equilibrium constant $K_c$, as long as temperature remains constant.

If you increase the concentration of a reactant, equilibrium shifts right to use up the extra reactant and produce more product. If you increase the concentration of a product, equilibrium shifts left to use up the extra product.

**Worked example:** For the equilibrium $\text{Fe}^{3+}(aq) + \text{SCN}^-(aq) \rightleftharpoons [\text{Fe(SCN)}]^{2+}(aq)$, what happens to the position of equilibrium when $\text{Fe}^{3+}$ ions are removed from solution? The product $[\text{Fe(SCN)}]^{2+}$ is dark red.

1. 1. The external change is a decrease in concentration of $\text{Fe}^{3+}$, which is a reactant.
2. 2. By Le Chatelier's principle, the system shifts to counteract the decrease in $\text{Fe}^{3+}$.
3. 3. To counteract the decrease, more $\text{Fe}^{3+}$ must be produced by breaking down product.
4. 4. Conclusion: equilibrium shifts left. The concentration of dark red product falls, so the solution becomes paler.

## Effect of Pressure and Temperature Changes

Pressure changes only affect the position of equilibrium if there is a different total number of moles of gas on the reactant and product sides. If the number of moles of gas is the same on both sides, changing pressure has no effect on equilibrium position.

All changes to temperature change the value of $K_c$. If the forward reaction is endothermic, an increase in temperature shifts equilibrium right (to absorb extra heat) and $K_c$ increases. If the forward reaction is exothermic, an increase in temperature shifts equilibrium left and $K_c$ decreases.

> **note**
>
> Only temperature changes the value of $K_c$. Concentration and pressure changes at constant temperature never change $K_c$.

**Worked example:** For the Haber process equilibrium: $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \Delta H = -92\ \text{kJ mol}^{-1}$. Predict the effect of (i) increasing total pressure, (ii) increasing temperature on equilibrium position.

1. Part (i): Increasing pressure
2. 1. Count total moles of gas: 4 mol on the left, 2 mol on the right.
3. 2. The change is increased pressure, so the system shifts to counteract the increase by decreasing pressure.
4. 3. Fewer moles of gas give lower total pressure, so equilibrium shifts right (towards ammonia).
5. Part (ii): Increasing temperature
6. 1. $\Delta H$ for forward reaction is negative, so forward is exothermic, reverse is endothermic.
7. 2. The change is increased temperature, so the system shifts to absorb the extra heat.
8. 3. Heat is absorbed by the endothermic reverse reaction, so equilibrium shifts left.

## Effect of Catalysts and Industrial Applications

Catalysts increase the rate of both the forward and reverse reactions equally. They reduce the time taken to reach equilibrium, but do not change the position of equilibrium or the value of $K_c$. They have no effect on equilibrium yield.

In industrial chemistry, Le Chatelier's principle is used to select conditions that balance high yield with economic factors like reaction rate and equipment cost. For example, the Haber process uses a compromise temperature to get both an acceptable yield and acceptable rate.

**Worked example:** Methanol is produced via this equilibrium: $\text{CO}(g) + 2\text{H}_2(g) \rightleftharpoons \text{CH}_3\text{OH}(g) \Delta H = -91\ \text{kJ mol}^{-1}$. What pressure and temperature conditions give the highest equilibrium yield?

1. 1. Count gaseous moles: 3 mol on reactant side, 1 mol on product side.
2. 2. Increasing pressure shifts equilibrium to the side with fewer moles, so high pressure gives the highest yield.
3. 3. The forward reaction is exothermic, so decreasing temperature shifts equilibrium right to release extra heat.
4. 4. Conclusion: High pressure and low temperature give the highest equilibrium yield. In practice, low temperature is too slow, so a moderate temperature with a catalyst is used.

## Common pitfalls

- **Wrong:** Claiming a catalyst shifts equilibrium right to increase yield
  - Why it fails: Catalysts speed up forward and reverse reactions equally, they only reduce time to reach equilibrium
  - Correct: State that catalysts have no effect on equilibrium position or equilibrium yield
- **Wrong:** Predicting a pressure shift when moles of gas are equal on both sides
  - Why it fails: Equal moles of gas mean pressure change affects both sides equally, no net counteraction possible
  - Correct: State that changing pressure has no effect on equilibrium position when total gaseous moles are equal
- **Wrong:** Saying increasing temperature shifts equilibrium to the exothermic direction
  - Why it fails: The shift must counteract the change: increasing temperature adds heat, so the system absorbs heat
  - Correct: Always shift to the endothermic direction when temperature increases, and exothermic when temperature decreases
- **Wrong:** Claiming changing concentration changes $K_c$ at constant temperature
  - Why it fails: $K_c$ is only dependent on temperature, even after a shift concentrations adjust to keep $K_c$ constant
  - Correct: State that $K_c$ remains constant at constant temperature, regardless of concentration or pressure changes

## Cheatsheet

| Change | Equilibrium Shift | Effect on $K_c$ |
| --- | --- | --- |
| Increase [reactant] | Right | No change |
| Increase [product] | Left | No change |
| Increase total pressure ($\Delta n_{gas} \neq 0$) | To side with fewer gas moles | No change |
| Increase temperature | Endothermic direction | Changes: increases if forward endo |
| Add catalyst | No shift | No change |

## What's next

Le Chatelier's principle is the core foundation for all equilibrium topics in A-Level Chemistry, from industrial processes to acid-base and solubility equilibria. Mastering how to predict shifts correctly is essential for both multiple choice and structured exam questions, and you will build on this knowledge to calculate equilibrium constants and solve quantitative equilibrium problems next.

- [Equilibrium constants](https://www.owlsprep.com/study/cie-9701-u7-equilibrium-constants/)
- [Reaction kinetics (AS)](https://www.owlsprep.com/study/cie-9701-u8-overview/)
- [Rates of Reaction](https://www.owlsprep.com/study/cie-9701-u8-rates-of-reaction/)

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