AHL: Extended Le Chatelier's principle

IB Chemistry HL· Topic 17: Equilibrium (AHL)· 20 min read

1. Core Principles of Extended Le Chatelier's★★☆☆☆⏱ 5 min

Le Chatelier's principle is a qualitative tool to predict how an equilibrium system adjusts to external changes. The AHL extended version requires you to connect the direction of shift to changes in equilibrium constants, not just position of equilibrium.

📘 Definition

Extended Le Chatelier's Principle

When a system at dynamic equilibrium is subjected to a change in concentration, temperature, or total pressure, the system shifts the position of equilibrium to partially counteract the imposed change. Only temperature changes alter the value of the equilibrium constant.

Example:

For an exothermic reaction, increasing temperature shifts equilibrium left to absorb added heat.

✓ Quick check

Check your understanding of core rules

Which change will alter the value of $K_{c}$ for the reaction $N_{2} (g) + 3 H_{2} (g) ⇌ 2 N H_{3} (g)$ ?
- Increase in total pressure
- Addition of a catalyst
- Increase in temperature
- Addition of more $N_{2}$
Reveal answer
Increase in temperature —
Correct! Only temperature changes change the value of the equilibrium constant, all other changes leave K unchanged at constant temperature.

2. Effect of Individual Condition Changes★★☆☆☆⏱ 6 min

Each condition change has a predictable effect on equilibrium position, which you need to be able to explain clearly for exam marks.

Concentration: Increasing concentration of a reactant shifts equilibrium right to consume extra reactant; increasing product concentration shifts left.
Pressure (gases only): Increasing total pressure shifts equilibrium to the side with fewer moles of gas, to reduce overall pressure. No shift if moles of gas are equal on both sides.
Temperature: Increasing temperature shifts equilibrium in the endothermic direction, to absorb the added heat. Decreasing temperature shifts in the exothermic direction.
Catalyst: No shift in equilibrium position, no change to $K$ , only faster attainment of equilibrium.

📐 Worked Example

For the reaction $P C l_{5} (g) ⇌ P C l_{3} (g) + C l_{2} (g)$ ΔH = +92 kJ mol⁻¹, predict the effect of increasing temperature on the position of equilibrium and $K_{c}$ .

1
1. Identify the direction of the endothermic reaction:
2
$\Delta H = +92 \text{ kJ mol}^{-1}, \text{so the forward reaction is endothermic}.$
3
1. Increasing temperature adds heat, so equilibrium shifts to counteract this by absorbing heat in the endothermic direction.
4
1. The forward reaction is endothermic, so equilibrium shifts right.
5
1. Shifting right increases product concentration and decreases reactant concentration, so $K_{c}$ increases.

Exam tip:

Always link your temperature shift prediction to the sign of ΔH, exam markers require this explicit connection.

3. Multiple Concurrent Condition Changes★★★★☆⏱ 7 min

A common higher difficulty exam question asks you to predict the net shift when two conditions change at the same time. You need to analyze each change separately, then combine their effects correctly.

📐 Worked Example

For the reaction $2 S O_{2} (g) + O_{2} (g) ⇌ 2 S O_{3} (g)$ ΔH = -197 kJ mol⁻¹, predict the net effect on equilibrium position if we increase temperature and increase total pressure at the same time.

1
1. Analyze the effect of increased temperature first:
2
The forward reaction is exothermic, so increased temperature shifts equilibrium left (to the endothermic direction to absorb heat).
3
1. Analyze the effect of increased total pressure next:
4
$\text{Moles of gas reactants} = 2 + 1 = 3, \text{Moles of gas products} = 2. \text{Fewer moles on product side, so increased pressure shifts right.}$
5
1. Net effect: Unless the magnitude of each change is given, we cannot predict the final position, only the individual effects of each change. State this explicitly in an exam.

4. Industrial Applications★★★☆☆⏱ 2 min

Extended Le Chatelier's principle is commonly applied to industrial equilibrium processes like the Haber process (ammonia production) and Contact process (sulfuric acid production) to explain the choice of operating conditions that balance yield, rate, and cost.

📐 Worked Example

Explain why the Haber process $N_{2} (g) + 3 H_{2} (g) ⇌ 2 N H_{3} (g)$ ΔH = -92 kJ mol⁻¹ uses high pressure rather than low pressure.

1
1. Count moles of gas on each side: 1 + 3 = 4 moles of gaseous reactants, 2 moles of gaseous product.
2
1. Increasing total pressure shifts equilibrium to the side with fewer moles of gas, which is the product side, to counteract the pressure increase.
3
1. Shifting right increases the equilibrium yield of ammonia, so high pressure is used to maximize product yield.
4
1. A compromise pressure of ~200 atm is typically used, as extremely high pressure is too expensive for industrial plant infrastructure.

Exam tip:

Always mention compromise conditions in industrial application questions, this is a common hidden marking point.

5. Common Pitfalls

Wrong move:

Claiming that increasing pressure shifts equilibrium even when moles of gas are equal on both sides.

Why:

Pressure change only affects equilibrium position if there is a difference in total moles of gas. Equal moles mean no shift.

Correct move:

State that there is no change to the position of equilibrium when moles of gas are equal on both sides, even if pressure increases.

Wrong move:

Claiming that adding a catalyst shifts the position of equilibrium.

Why:

Catalysts speed up forward and reverse reactions equally, so they do not change equilibrium position or K.

Correct move:

State that a catalyst only reduces the time taken to reach equilibrium, with no effect on position or equilibrium yield.

Wrong move:

Claiming that changing concentration changes the value of $K_{c}$ .

Why:

K is only affected by temperature, it remains constant at constant temperature even when concentration changes.

Correct move:

State that K is constant at constant temperature, after a concentration change the system re-establishes equilibrium with the same K value.

Wrong move:

Predicting a net shift for two opposing changes without magnitude data.

Why:

Without knowing how large each change is, you cannot determine which shift dominates.

Correct move:

Describe the effect of each individual change, then state the net direction cannot be determined from the given information.

Wrong move:

Saying increasing temperature always increases K.

Why:

K only increases if the forward reaction is endothermic. If forward is exothermic, increasing temperature decreases K.

Correct move:

Link K change to the sign of ΔH for the forward reaction to get the direction of change right.

6. Quick Reference Cheatsheet

Change Type	Equilibrium Shift Direction	Change to $K$
Increase reactant concentration	Right	No change
Increase product concentration	Left	No change
Increase total pressure (gas)	Side with fewer moles of gas	No change
Increase T (forward endothermic)	Right	Increases
Increase T (forward exothermic)	Left	Decreases
Add catalyst	No shift	No change

7. Frequently Asked

Does a catalyst change the position of equilibrium?

No. A catalyst increases the rate of both forward and reverse reactions equally, so it only speeds up reaching equilibrium, and does not shift the position or change the value of $K$ .

Does changing pressure always shift equilibrium?

No. Changing pressure only shifts equilibrium if there is a difference in the total number of moles of gas on the reactant and product sides. If moles of gas are equal, pressure change has no effect on equilibrium position.

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.

2025 · 2
6 mark Haber process shift explanation
2024 · 1
MCQ on K change with temperature
2023 · 2
Pressure change equilibrium prediction

Going deeper

syllabusIB Chemistry HL Topic 17 SyllabusOfficial AHL equilibrium content outline
articleHaber Process Operating ConditionsCase study of Le Chatelier's application

What's Next

Extended Le Chatelier's principle is a foundational concept for all further equilibrium topics in IB Chemistry HL, including acid-base equilibria, solubility equilibria, and buffer calculations. Mastering the rules here will make predicting shifts in these more complex systems much easier, as the same core logic applies. Common exam questions combine this concept with equilibrium constant calculations, so understanding how K changes (or does not change) with different conditions is critical for solving calculation-based problems correctly.