Le Chatelier’s Principle states that if a system at dynamic equilibrium is subjected to a change in conditions, the position of equilibrium shifts to oppose the imposed change. An increase in temperature shifts the equilibrium in the endothermic direction, absorbing the added heat, while a decrease in temperature shifts it in the exothermic direction, releasing heat. An increase in pressure shifts the equilibrium toward the side with fewer moles of gas, reducing the pressure, whereas a decrease in pressure favours the side with more moles of gas. Changing the concentration of a reactant or product causes the equilibrium to shift in the direction that consumes the added substance or replenishes a removed one. A catalyst has no effect on the position of equilibrium because it speeds up the forward and reverse reactions equally, lowering the activation energy for both pathways by the same amount. This principle underpins the design of major industrial processes such as the Haber process (ammonia production) and the Contact process (sulfuric acid manufacture), where engineers must balance rate and yield when selecting operating conditions, and it also explains laboratory phenomena such as the reversible colour change observed when hydrated copper(II) sulfate is heated and rehydrated.
Statement of the Principle
When a system at dynamic equilibrium experiences a change in concentration, temperature, or pressure, the equilibrium position adjusts so as to counteract the change. This is often summarised by the phrase “the equilibrium shifts to oppose the change.” It is important to understand that the principle describes the direction of shift, not the rate at which it occurs, and that the system must already be at equilibrium for the principle to apply.
When describing the response, use precise language:
- The equilibrium shifts to the right (forward reaction favoured, more products formed, yield increases).
- The equilibrium shifts to the left (reverse reaction favoured, more reactants formed, yield decreases).
The word “oppose” is the term required by IGCSE mark schemes — avoid phrases like “the equilibrium helps” or “the system wants to.”
Effect of Temperature Change
Temperature changes affect the position of equilibrium depending on whether the forward reaction is exothermic or endothermic.
- Increasing temperature shifts the equilibrium in the endothermic direction (the direction that absorbs heat).
- Decreasing temperature shifts the equilibrium in the exothermic direction (the direction that releases heat).
Worked Example: The Haber Process
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹
The forward reaction is exothermic (ΔH is negative). Therefore:
- Lowering the temperature favours the forward reaction, increasing the yield of ammonia.
- Raising the temperature favours the reverse reaction, decreasing the yield of ammonia.
Rate vs yield note: Although a low temperature gives a higher yield of ammonia in the Haber process, the rate of reaction becomes unacceptably slow. In industrial practice, a compromise temperature of around 450 °C is used to achieve a reasonable rate while still obtaining an economically viable yield. An iron catalyst is also used to increase the rate further, though it does not affect the equilibrium position.
Effect of Pressure Change (Gaseous Systems Only)
Pressure changes only affect equilibria involving gases. If the number of gas molecules is the same on both sides of the equation, pressure has no effect on the position of equilibrium. Solids and liquids are virtually incompressible and are ignored when counting moles for pressure predictions.
- Increasing pressure shifts the equilibrium toward the side with fewer moles of gas.
- Decreasing pressure shifts the equilibrium toward the side with more moles of gas.
Counting Moles of Gas
Count only the stoichiometric coefficients of gaseous species:
| Reaction | Moles of gas (reactants) | Moles of gas (products) | Favoured by high pressure? |
|---|---|---|---|
| N₂ + 3H₂ ⇌ 2NH₃ | 1 + 3 = 4 | 2 | Yes — forward reaction |
| 2SO₂ + O₂ ⇌ 2SO₃ | 2 + 1 = 3 | 2 | Yes — forward reaction |
| H₂ + I₂ ⇌ 2HI | 1 + 1 = 2 | 2 | No shift — equal moles |
| CaCO₃(s) ⇌ CaO(s) + CO₂(g) | 0 (solids ignored) | 1 | No — reverse reaction favoured |
Contact Process Example
2SO₂(g) + O₂(g) ⇌ 2SO₃(g) (3 moles of gas on left, 2 moles of gas on right)
A high pressure shifts the equilibrium to the right, increasing the yield of SO₃. In practice, the reaction is carried out at only 1–2 atm because the equilibrium already lies far to the right under these conditions, and the use of a vanadium(V) oxide catalyst ensures a fast enough rate without the expense of high-pressure equipment.
No-Effect Scenario
For the reaction H₂(g) + I₂(g) ⇌ 2HI(g), there are 2 moles of gas on each side. Changing the pressure has no effect on the equilibrium position because both sides contain the same number of gaseous molecules. Both the forward and reverse reactions are affected equally by compression.
Effect of Concentration Change
Changing the concentration of one species in an equilibrium mixture causes a shift to oppose that change.
- Adding a reactant shifts the equilibrium to the right (toward products), using up the added reactant.
- Adding a product shifts the equilibrium to the left (toward reactants), consuming some of the added product.
- Removing a product (e.g., by allowing a gas to escape or by precipitation) shifts the equilibrium to the right, as the system replaces what was removed.
- Removing a reactant shifts the equilibrium to the left.
This effect is exploited in industrial processes: in the Haber process, ammonia is continuously liquefied and removed from the reaction vessel, pulling the equilibrium to the right and increasing overall yield. Similarly, in esterification reactions, removing water as it forms drives the equilibrium toward more ester production.
Effect of a Catalyst
This is one of the most tested IGCSE misconceptions.
A catalyst has NO effect on the position of equilibrium.
A catalyst provides an alternative reaction pathway with a lower activation energy. Crucially, it lowers the activation energy by the same amount for both the forward and reverse reactions:
- The rate of the forward reaction increases.
- The rate of the reverse reaction increases by the same factor.
- The equilibrium is reached more quickly, but the position of equilibrium is unchanged.
- The equilibrium mixture has the same composition as it would have without a catalyst.
This is a common examination point: adding a catalyst does not increase the yield of products; it only reduces the time taken to reach equilibrium.
Why? The equilibrium position is determined by thermodynamic factors (the relative free energies of reactants and products, summarised by the equilibrium constant Kc). A catalyst is a kinetic factor — it changes how fast equilibrium is reached, not where equilibrium lies. Because the activation energy barrier for both forward and reverse reactions is lowered by the same amount, the ratio of rate constants (k_forward / k_reverse) remains unchanged, and so the equilibrium composition remains unchanged.
If a question asks “What effect does adding a catalyst have on the yield of ammonia in the Haber process?”, the correct answer is “No effect” or “None.” Adding “the catalyst speeds up the rate but does not change the equilibrium position” earns full marks.
Summary Table: All Changes and Their Effects
| Change imposed | System response | Direction of equilibrium shift | Effect on yield of forward product |
|---|---|---|---|
| Increase temperature | Absorbs heat | Endothermic direction | Increases if forward is endothermic; decreases if forward is exothermic |
| Decrease temperature | Releases heat | Exothermic direction | Increases if forward is exothermic; decreases if forward is endothermic |
| Increase pressure | Reduces pressure | Side with fewer moles of gas | Increases if product side has fewer moles |
| Decrease pressure | Increases pressure | Side with more moles of gas | Increases if product side has more moles |
| Add reactant | Consumes added reactant | Right (toward products) | Increases |
| Remove reactant | Replaces removed reactant | Left (toward reactants) | Decreases |
| Add product | Consumes added product | Left (toward reactants) | Decreases |
| Remove product | Replaces removed product | Right (toward products) | Increases |
| Add catalyst | No change in composition | No shift | No change |
Industrial Applications
Haber Process (N₂ + 3H₂ ⇌ 2NH₃, ΔH = −92 kJ mol⁻¹)
| Condition | Value used | Reason |
|---|---|---|
| Temperature | ~450 °C | Compromise: low temperature favours yield (exothermic forward reaction) but would make the rate too slow; 450 °C gives an acceptable rate |
| Pressure | ~200 atm | High pressure favours the forward reaction (4 moles → 2 moles of gas), shifting equilibrium to the right; limited by equipment cost and safety |
| Catalyst | Finely divided iron (Fe) | Speeds up attainment of equilibrium; does not affect yield |
| Product removal | NH₃ liquefied and removed | Removing product shifts equilibrium right, improving overall yield |
| Recycling | Unreacted N₂ and H₂ recycled | Re-adding reactants shifts equilibrium right on each pass |
Contact Process (2SO₂ + O₂ ⇌ 2SO₃, ΔH = −197 kJ mol⁻¹)
| Condition | Value used | Reason |
|---|---|---|
| Temperature | ~450 °C | Compromise: low temperature increases yield but reduces rate; 450 °C is a practical balance. Reaction already has a high equilibrium yield |
| Pressure | 1–2 atm | Forward reaction is favoured by high pressure (3 moles → 2 moles), but the equilibrium already lies well to the right at low pressure, so high pressure is not economically justified |
| Catalyst | Vanadium(V) oxide (V₂O₅) | Speeds up attainment of equilibrium; essential for a viable rate at the chosen temperature |
The Rate vs Yield Distinction (Key IGCSE Concept)
A common IGCSE examination theme is the distinction between how a change affects the rate of reaction and how it affects the yield (the position of equilibrium).
- Rate of reaction — how fast equilibrium is reached (kinetics).
- Position of equilibrium / yield — the proportion of products present at equilibrium (thermodynamics).
Conditions that give a fast rate may give a poor yield, and vice versa:
| Factor | Effect on rate | Effect on yield / equilibrium position |
|---|---|---|
| Temperature increase | Always increases rate (more particles have energy ≥ Eₐ) | Shifts equilibrium in the endothermic direction |
| Pressure increase (gases) | Increases rate (more frequent collisions) | Shifts equilibrium toward fewer moles of gas |
| Concentration increase | Increases rate (more frequent collisions) | Shifts equilibrium away from the added substance |
| Catalyst addition | Increases rate | No effect on position of equilibrium or yield |
| Surface area increase | Increases rate | No effect on equilibrium position |
The distinction explains why industrial conditions are almost always a compromise: the conditions that give the best yield are rarely those that give the fastest rate, and vice versa. The cost of equipment, energy, and raw materials must all be weighed against the value of the product.
Laboratory Example: Hydrated Copper(II) Sulfate
The interconversion of hydrated and anhydrous copper(II) sulfate provides a clear visual demonstration of Le Chatelier’s Principle in a reversible reaction:
CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)
Hydrated copper(II) sulfate crystals are blue. Anhydrous copper(II) sulfate is a white powder.
- Heating (increase temperature): The forward reaction is endothermic (dehydration requires energy input). Equilibrium shifts right — blue hydrated crystals turn to white anhydrous powder.
- Adding water (increase concentration of a product on the right): Equilibrium shifts left — white anhydrous powder rehydrates and turns blue again. Heat is released (the reverse reaction is exothermic).
This example illustrates both the temperature and concentration predictions of Le Chatelier’s Principle in a simple, visually striking system. It is also the basis of the chemical test for water using anhydrous copper(II) sulfate: the colour change from white to blue indicates the presence of water.
Sources
- Cambridge IGCSE Chemistry 0620 Syllabus (2023–2025), Topic 7: Chemical Reactions — Reversible Reactions and Equilibrium
- Harwood, R. & Lodge, I. (2014) Cambridge IGCSE Chemistry Coursebook, 3rd ed. Cambridge University Press
- OpenStax (2016) Chemistry, Chapter 13: Fundamental Equilibrium Concepts, Section 13.3 — Le Chatelier’s Principle
- BBC Bitesize (2024) “Reversible reactions and equilibria — Le Chatelier’s principle”
- Equilibrium
- Reversible Reactions
- Haber Process and Ammonia
- Contact Process
Common Misconceptions
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“The equilibrium shifts to help or assist the change” — The equilibrium does not assist the applied change; it shifts to oppose or counteract it. The system partially undoes what was done to it, but it cannot fully reverse the change. For example, if temperature is increased, the equilibrium shifts in the endothermic direction, absorbing some heat but not restoring the original temperature. IGCSE mark schemes penalise any wording that implies the system “helps” or “wants” to undergo the change — use the word “oppose.”
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“A catalyst increases the yield of products” — A catalyst lowers the activation energy equally for both forward and reverse reactions. It allows equilibrium to be reached more quickly but has no effect whatsoever on the position of equilibrium or the percentage yield. This is one of the most frequently tested points in IGCSE examinations. If asked about the effect of a catalyst on yield, the correct answer is “No effect.”
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“Higher temperature always favours the forward reaction” — Temperature shifts are determined by enthalpy, not by which direction is “forward.” If the forward reaction is exothermic, increasing temperature favours the reverse (endothermic) reaction, reducing the yield of products. The direction depends entirely on whether the forward reaction is exothermic or endothermic.
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“Pressure changes affect all equilibria” — Pressure changes only affect the position of equilibrium when the reaction involves gases and the number of gas molecules differs between reactants and products. Equilibria in solution, or gaseous equilibria with equal moles on each side (e.g., H₂ + I₂ ⇌ 2HI), are unaffected by pressure changes.
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“Rate and yield are the same thing — a fast reaction means more product” — Many students conflate rate and yield. A fast reaction does not mean a high yield, and conditions that maximise rate (e.g., high temperature) may simultaneously reduce yield if the forward reaction is exothermic. Lowering temperature may increase equilibrium yield but always decreases the rate. These two concepts must be distinguished clearly in examination answers.