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Chemical equilibrium is the steady state reached by a reversible reaction in a closed system, when the forward and reverse processes proceed at identical rates. Crucially, equilibrium is dynamic — both reactions continue at the molecular level, but their effects cancel macroscopically, leaving the concentrations of reactants and products constant (though not generally equal). This lesson introduces Le Chatelier's principle, first articulated in 1884 by the French chemist whose name it bears, as a qualitative rule for predicting how an equilibrium responds to disturbance. We will analyse the effects of changing concentration, total pressure, temperature, and adding a catalyst on the position of equilibrium, and apply these ideas to three industrially decisive systems: the Haber synthesis of ammonia, the Contact process for sulfur trioxide, and the esterification equilibrium that underpins flavour and fragrance chemistry. The quantitative companion to this lesson — equilibrium constants Kc and Kp — follows immediately, but Le Chatelier's principle is the conceptual scaffold on which all of that quantitative work is hung.
Spec mapping (AQA 7405): This lesson maps to §3.1.6 (chemical equilibria, Le Chatelier's principle and Kc — qualitative treatment). The quantitative Kc expression and ICE-table calculations are developed in lesson 5 of this course; the partial-pressure formulation Kp for gaseous equilibria in lesson 6; the equilibrium criterion ΔG = 0 (and Kc/Kp from thermodynamics) in lesson 5 of the energetics & thermodynamics course. Organic equilibria — esterification, acid hydrolysis of esters — anchor at §3.3.8; industrial equilibria (Haber, Contact) recur across §3.1.6 and §3.3. Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: Stating Le Chatelier's principle and the conditions for dynamic equilibrium are AO1 recall items that appear almost every series. Predicting the direction of shift for changes in concentration, temperature, and pressure — and the corresponding effect on yield — is AO2 and constitutes the bulk of mark-scheme credit. AO3 questions ask students to rationalise industrial choices: why 400-450°C and not 200°C for the Haber process? Why 1-2 atm for the Contact process when the Haber runs at 200 atm? Why a catalyst when it cannot shift the position? These yield-vs-rate trade-offs are the highest-tariff item in the section and discriminate Grade A* from Grade B.
When a reversible reaction is carried out in a closed system — one that exchanges energy but not matter with its surroundings — the concentrations of reactants and products eventually settle into time-independent values. At this point the system is at chemical equilibrium. The defining microscopic feature is that both the forward and reverse reactions are still occurring; equilibrium is not a state of inactivity. Individual reactant molecules continue to convert to products, and product molecules continue to revert to reactants, but the two rates are now equal, so there is no net change in concentration. This is what is meant by a dynamic equilibrium.
Three criteria must be met:
Key Definition: A dynamic equilibrium occurs in a closed system when the rate of the forward reaction equals the rate of the backward reaction, so the concentrations of reactants and products remain constant on the macroscopic scale, even though both reactions continue at the molecular level.
A useful evidential argument for the dynamic (rather than static) nature of equilibrium comes from isotopic-labelling experiments. If a sample of water containing some D₂O is added to a saturated salt solution containing dissolved D-free HCl, the chloride and hydrogen scramble across H/D positions despite there being no net change in any measurable concentration. The continuous interconversion would be impossible if the equilibrium were static. Similar isotope-scrambling experiments at the C=O of esters, or across the proton positions of NH₃/ND₃ mixtures, all confirm the dynamic picture.
In 1884 Henry Louis Le Chatelier proposed an empirical rule, since vindicated by thermodynamics, for predicting the qualitative response of an equilibrium to disturbance. The principle is sometimes called Le Chatelier's-Braun principle to credit Karl Ferdinand Braun's independent 1887 formulation, but the simpler name dominates in school textbooks.
Key Definition: Le Chatelier's principle states that if a system at chemical equilibrium is subjected to a change in conditions (concentration, temperature, or pressure), the position of equilibrium shifts to partially oppose the change.
Two cautions are essential. First, the principle predicts only the direction of shift, not its magnitude. Calculating how far the equilibrium moves requires the equilibrium constant and an ICE table (next lesson). Second, the principle predicts partial opposition, never complete cancellation. The system shifts enough to relieve some of the imposed stress but never restores the original concentrations or temperature exactly — there must be a net difference, because the equilibrium constant typically also changes (for temperature changes) or stays fixed (for concentration/pressure changes).
The principle should not be applied to:
We will systematically analyse each variable using the exothermic, gas-phase Haber equilibrium as a running example:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹
Adding more of any reactant shifts the equilibrium right (towards products); adding more of any product shifts it left (towards reactants). The system responds to the imposed excess by consuming some of the added species, partially restoring the previous concentration. Removing a species (e.g. by condensing out a volatile product, or by precipitation) has the opposite effect.
For the Haber equilibrium:
Crucially, a concentration change does not alter the equilibrium constant Kc. When a system at equilibrium is disturbed by a concentration change, the reaction quotient Q is briefly different from Kc, and the system shifts to re-establish Q = Kc with new concentrations.
Increasing the total pressure on a gas-phase equilibrium shifts the position towards the side with fewer moles of gas. Decreasing the pressure shifts it towards the side with more moles of gas. If the number of moles of gas is the same on both sides, pressure has no effect on the position.
For the Haber equilibrium:
For the dissociation of dinitrogen tetroxide N₂O₄(g) ⇌ 2NO₂(g):
For H₂(g) + I₂(g) ⇌ 2HI(g):
Note: pressure changes also do not alter Kc (or Kp) — only temperature does. The position shifts because the concentrations (and partial pressures) of each species change differently upon compression, transiently making Q ≠ K.
Temperature is the only variable that changes the value of the equilibrium constant. Increasing temperature shifts the equilibrium in the endothermic direction (i.e. towards the side that absorbs heat); decreasing temperature shifts it in the exothermic direction.
For an exothermic forward reaction (ΔH < 0), such as the Haber synthesis:
For an endothermic forward reaction (ΔH > 0), such as the dissociation of phosphorus pentachloride PCl₅(g) ⇌ PCl₃(g) + Cl₂(g) with ΔH = +124 kJ mol⁻¹:
The mnemonic that catches the sign: "heat is on the endothermic side." Treat heat as if it were a chemical species that appears on the reactant side of an endothermic reaction (and on the product side of an exothermic one). Adding heat then shifts the equilibrium "away" from heat, by Le Chatelier — i.e. into the endothermic direction. The quantitative version of this is the van't Hoff equation, dlnK/dT = ΔH°/RT², which appears in Going Further below.
A catalyst does not change the position of equilibrium and does not change the value of K. It lowers the activation energy of the forward and reverse reactions by the same amount, so it speeds up both directions equally. The equilibrium is therefore reached sooner, but the equilibrium composition is identical to that without a catalyst.
This is occasionally questioned ("surely a tiny effect?") but the equality is exact. By thermodynamics, K depends only on standard-state Gibbs free energy differences between reactants and products — not on the activation energy of the path connecting them. A catalyst, which alters only the activation energy, cannot change K. The "tiny effect" intuition arises from confusing reaction rate with reaction extent.
In industry, catalysts are essential not for yield but for economic time-to-equilibrium. Without a catalyst at 450°C, the Haber process would take days per equilibration cycle; with an iron catalyst it takes minutes.
| Disturbance | Effect on Position | Effect on K |
|---|---|---|
| Add reactant | Shift right | No change |
| Add product | Shift left | No change |
| Remove product | Shift right | No change |
| Increase pressure | Shift to fewer-moles side | No change |
| Increase temperature | Shift in endothermic direction | Changes (K depends on T) |
| Add catalyst | No shift | No change |
Common Misconception: Students sometimes claim that increasing pressure "always favours the forward reaction" or "always increases yield". Neither is true. Pressure shifts only towards the fewer-moles-of-gas side; whether that is forward or backward depends on the specific stoichiometry. And the meaning of "yield" is also specific — increasing yield of product P does not necessarily mean the position has shifted in the conventional "forward" direction.
The three classic industrial equilibria — Haber, Contact, and esterification — illustrate every facet of Le Chatelier's principle and the yield-vs-rate trade-off that dominates industrial chemistry.
The synthesis of ammonia from atmospheric nitrogen and hydrogen, devised by Fritz Haber and scaled up by Carl Bosch:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹
Operating conditions: approximately 400-450°C, 150-250 atm (typically 200 atm), iron catalyst (with K₂O and Al₂O₃ promoters).
The choice of conditions is a textbook example of compromise:
Pressure: 150-250 atm. Equilibrium position favours NH₃ (fewer moles of gas on the right), so high pressure is desirable. Going higher than ~250 atm increases yield further but the marginal yield gain shrinks while capital costs (thicker steel vessels, more robust seals) and safety risks (compressor failure, runaway exotherm) rise steeply. 200 atm is the economic optimum.
Temperature: 400-450°C. The forward reaction is exothermic, so low temperature favours the equilibrium yield of NH₃. However, at low temperatures the reaction rate is too slow to be commercially viable. 400-450°C is a deliberate compromise: yield is reduced (~15-20% per pass) but the rate is acceptable, and unreacted N₂ and H₂ are recycled, so the overall conversion approaches 97-99% after many passes. A purely Le Chatelier-driven student would say "use low temperature for maximum yield"; the industrially-aware A student recognises that the rate at low T makes the plant economically unworkable.*
Catalyst: iron with K₂O/Al₂O₃ promoters. Iron supplies the catalytic surface; potassium oxide and aluminium oxide promote activity and prevent sintering. The catalyst does not shift the equilibrium but allows it to be reached in seconds rather than days. Without the catalyst, the Haber process at 450°C would not be commercially viable.
Continuous removal of NH₃. The product stream is cooled, condensing liquid ammonia (b.p. −33°C) which is drawn off. By Le Chatelier, removing NH₃ continually shifts the equilibrium right and drives further conversion.
The Haber process produces around 175 million tonnes of ammonia annually, most of which is converted to fertiliser. The process is estimated to underpin the food supply of nearly half the world's population.
The catalytic oxidation of sulfur dioxide to sulfur trioxide, en route to sulfuric acid:
2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = −197 kJ mol⁻¹
Operating conditions: approximately 450°C, 1-2 atm, vanadium(V) oxide (V₂O₅) catalyst.
Pressure: 1-2 atm. Equilibrium position favours SO₃ (fewer moles of gas on the right: 3 → 2). However, at moderate temperatures the yield is already >99% even at ambient pressure, so the modest yield gain from increasing pressure does not justify the cost of compression and pressure-rated reactors. Contrast with Haber, where the equilibrium is much less favourable and high pressure is essential.
Temperature: 450°C. Forward reaction is exothermic, so low T favours yield. As with Haber, the catalyst is not effective at low temperatures (sulfur compounds poison vanadium oxide below ~400°C, and surface kinetics are too slow), so 450°C is the rate/yield compromise.
Catalyst: V₂O₅. Vanadium(V) oxide on a silica or kieselguhr support. The catalysis proceeds via a redox cycle: V₂O₅ + SO₂ → V₂O₄ + SO₃, then V₂O₄ + ½O₂ → V₂O₅, regenerating the catalyst. The catalyst surface accepts and releases oxygen in a continuous cycle.
The SO₃ produced is dissolved in concentrated H₂SO₄ to form oleum (H₂S₂O₇), which is then diluted to give 98% sulfuric acid. SO₃ is not bubbled directly into water because the reaction is so exothermic that it produces a corrosive acid mist that cannot be condensed efficiently. Sulfuric acid is the most-produced industrial chemical in the world by mass (over 250 million tonnes per year).
The acid-catalysed condensation of a carboxylic acid with an alcohol to form an ester:
CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O (Kc ≈ 4 at room temperature)
Unlike the Haber and Contact equilibria, esterification:
To drive the equilibrium forward, industrial and laboratory chemists use Le Chatelier-driven strategies:
Industrial production of ethyl acetate (a major solvent for paints, lacquers, and pharmaceuticals) relies on continuous water removal. Industrial production of methyl methacrylate (the monomer for PMMA / acrylic) also runs through esterification equilibria driven by water removal and reactant excess.
Example 1 — Contact equilibrium under pressure. For 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), ΔH = −197 kJ mol⁻¹, predict the effect of increasing pressure at constant temperature.
Left side: 2 + 1 = 3 moles of gas. Right side: 2 moles of gas. Increasing pressure shifts the equilibrium towards the side with fewer moles — to the right, increasing the yield of SO₃. Kp is unaffected (only T changes Kp).
Example 2 — PCl₅ dissociation under heating. For PCl₅(g) ⇌ PCl₃(g) + Cl₂(g), ΔH = +124 kJ mol⁻¹, predict the effect of increasing temperature.
The forward reaction is endothermic. Increasing T shifts the equilibrium in the endothermic direction — to the right, increasing the yields of PCl₃ and Cl₂. The value of Kc increases. Observation: the pale yellow-green colour of Cl₂ intensifies as the sealed sample is heated.
Example 3 — N₂O₄ compression at constant T. For N₂O₄(g) ⇌ 2NO₂(g), ΔH = +57 kJ mol⁻¹, predict the effect of compressing the sample.
Left side: 1 mole. Right side: 2 moles. Compression shifts the equilibrium to the side with fewer moles — to the left (towards colourless N₂O₄). Observation: the brown colour due to NO₂ fades as the gas is compressed. Kc is unchanged because temperature is constant.
Example 4 — Haber catalyst rationale. Explain why an iron catalyst is used in the Haber process despite having no effect on equilibrium position.
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