Le Châtelier's principle: how equilibria respond to stress
Learn Le Châtelier's principle: how equilibrium systems respond to changes in concentration, pressure, and temperature — and why catalysts are the exception.
Quick answer
Le Châtelier's principle states that when a stress (concentration, pressure, or temperature change) disturbs an equilibrium, the system shifts in the direction that relieves the stress — mathematically, the direction that restores Q to K.
Example: For N₂ + 3H₂ ⇌ 2NH₃ (ΔH = −92.4 kJ/mol): raising the pressure shifts right (4 gas moles → 2), while raising the temperature shifts left (heat favors the endothermic reverse direction).
What you will learn
- State Le Châtelier's principle in terms of Q returning to K.
- Predict the shift caused by adding or removing a reactant or product.
- Predict the shift caused by pressure/volume changes using gas mole counts.
- Predict the shift caused by temperature changes for exothermic versus endothermic reactions.
The principle behind the shifts
Le Châtelier's principle: when a change (a 'stress') is applied to a system at equilibrium, the system shifts to partially counteract the change. Behind the words is simple math: the stress changes Q instantly while K stays fixed, and the reaction runs in whichever direction brings Q back to K.
The three stresses that matter are concentration, pressure (or volume), and temperature. Catalysts are not a stress — they change neither Q nor K, so they never shift equilibrium.
Concentration changes
Adding a reactant (or removing a product) drops Q below K, so the system shifts right, consuming the added material to make more product. Removing a reactant (or adding a product) raises Q above K, shifting the system left.
Example: in the equilibrium Fe³⁺ + SCN⁻ ⇌ FeSCN²⁺ (pale yellow + colorless ⇌ blood red), adding Fe³⁺ deepens the red color as more complex forms; adding a reagent that removes SCN⁻ fades it. The color change makes the invisible shift visible — a classic demonstration.
Pressure and volume changes
Changing pressure (by changing volume) matters only when the two sides have different numbers of gas moles. Higher pressure favors the side with fewer gas moles; lower pressure favors the side with more.
For N₂(g) + 3H₂(g) ⇌ 2NH₃(g): 4 gas moles become 2, so compressing the mixture shifts the equilibrium right, toward ammonia. For H₂(g) + I₂(g) ⇌ 2HI(g): 2 gas moles on each side, so pressure changes do nothing at all.
Adding an inert gas at constant volume changes nothing (concentrations unchanged); at constant pressure it dilutes everything, which can shift the equilibrium.
Temperature: the stress that changes K
Temperature is unique: it changes K itself. Treat heat as a reactant or product:
- Exothermic (ΔH < 0, heat is a 'product'): raising T shifts left; lowering T shifts right.
- Endothermic (ΔH > 0, heat is a 'reactant'): raising T shifts right; lowering T shifts left.
Ammonia synthesis (ΔH = −92.4 kJ/mol) is exothermic, so low temperatures favor ammonia yield — but also slow the reaction to a crawl. The industrial compromise (~450 °C, 200 atm, iron catalyst) balances thermodynamics against kinetics, a trade-off Le Châtelier alone cannot resolve.
Variables and units
| Symbol | Name | Unit | Meaning |
|---|---|---|---|
| Enthalpy change | kJ/mol | Enthalpy change of the reaction; negative for exothermic, positive for endothermic |
Common mistakes
Systems do not want anything. The shift is the automatic consequence of Q being pushed away from K: the reaction proceeds in whichever direction restores Q = K. The principle is a summary of the math, not a motive.
Usually yes, but only if the added reactant actually changes Q. Adding more of a pure solid or liquid does nothing (they are not in the Q expression), and adding reactant to a system already at Q > K can shift it left instead.
It favors the side with fewer moles of gas. Solids and liquids do not count. For a reaction like C(s) + CO₂(g) ⇌ 2CO(g), higher pressure shifts left (1 gas mole versus 2) even though the left has more total 'molecules.'
Real-world applications
Haber process: high pressure and low temperature favor ammonia yield — Le Châtelier in industry.
Carbonated drinks: bottled under CO₂ pressure to keep the gas dissolved.
Contact process: sulfuric acid manufacture optimizes temperature against equilibrium.
Physiology: the bicarbonate buffer system shifts to stabilize blood pH.
Frequently asked questions
Does adding reactant change the equilibrium constant?
No — Q changes instantly (concentrations or pressures changed) while K stays fixed at constant temperature. The system then shifts to bring Q back to K. Temperature is the special case: it changes K itself, which is why heating shifts equilibrium differently.
Does a catalyst shift the equilibrium position?
No. A catalyst lowers the barrier for both directions equally, so equilibrium is reached faster but the final concentrations are identical. Catalysts change the journey, never the destination.
Do solids and liquids count when predicting pressure effects?
No. Count only gas moles: for N₂ + 3H₂ ⇌ 2NH₃, 4 gas moles become 2, so higher pressure favors ammonia. Solids and liquids are omitted from the count because their concentrations do not change with pressure.
Why is the Haber process run hot if heat hurts the ammonia yield?
The two effects fight: higher temperature speeds the reaction (kinetics) but lowers the ammonia yield (thermodynamics, since the reaction is exothermic). Industry compromises around 450 °C with an iron catalyst — fast enough, with acceptable yield.