Le Châtelier's Principle

Poke a reaction at equilibrium and it pushes back. A single rule predicts exactly how it responds to every kind of stress.

High schoolIntro Gen ChemUni Year 1
⏱️ About 18 min

A system sitting at equilibrium is like a balance beam that's settled level. Push down on one side — add a chemical, squeeze the container, heat it up — and it tilts, then settles at a new level that partly undoes your push. Le Châtelier's principle is the rule that tells you, every time, which way it tilts.

💡
The big idea: If you disturb a system at equilibrium, it shifts in the direction that partially counteracts the disturbance. Add something and it consumes some of it; remove something and it replaces some; heat it and it absorbs some heat. The system always pushes back — but only partway.
🎯 By the end, you'll be able to
  • State Le Châtelier's principle and apply it to a disturbed equilibrium
  • Predict the shift from a concentration, pressure/volume, or temperature change
  • Explain why only a temperature change alters the value of K
  • Explain why a catalyst speeds equilibrium up but does NOT shift it
📎 Helpful to know first

The principle: shift to counteract the stress

Take a reaction that has reached equilibrium and then stress it — change a concentration, the pressure or volume, or the temperature. Le Châtelier's principle says the equilibrium will shift in whatever direction partially opposes the change you made.

The word partially matters: the system relieves some of the stress but never fully cancels it. It settles at a new equilibrium that is a compromise between the disturbance and the reaction's natural balance.

🔑 Le Châtelier's principle
When a system at equilibrium is disturbed, it responds by shifting in the direction that partially counteracts the disturbance and restores a (new) equilibrium.

1. Changing a concentration

Add a reactant (or product) and the system shifts to consume some of it. Remove a substance and the system shifts to replace some of it.

For N₂ + 3H₂ ⇌ 2NH₃: pump in extra H₂ and the equilibrium shifts right, using up some of that H₂ to make more NH₃. Pull NH₃ out as it forms and the system shifts right again, replacing it. This is exactly why industrial plants continuously remove product — to keep dragging the reaction forward.

✨ This is just Q vs K again
Concentration shifts are the reaction quotient in action. Add reactant and you shrink Q below K (Q < K), so the reaction runs forward. Add product and you push Q above K (Q > K), so it runs reverse. Le Châtelier and the Q-versus-K rule give the same answer — because they're the same physics.

2. Changing pressure or volume (gases)

For reactions involving gases, squeezing the container (raising the pressure by lowering the volume) shifts the equilibrium toward the side with fewer moles of gas — that's the side that relieves the pressure. Expanding the volume (lowering pressure) shifts toward the side with more moles of gas.

For N₂ + 3H₂ ⇌ 2NH₃: the left has 1 + 3 = 4 gas molecules, the right has 2. Increase the pressure and the equilibrium shifts right (toward the 2-molecule side). If both sides have equal moles of gas, a pressure change causes no shift.

\[ \ce{N2(g) + 3H2(g) <=> 2NH3(g)} \qquad 4\ \text{mol gas} \;\longrightarrow\; 2\ \text{mol gas} \]
Raising the pressure favours the side with fewer gas molecules — here, the products.
✨ Adding an inert gas at constant volume does nothing
Pumping in an unreactive gas (like argon) at fixed volume raises the total pressure but leaves each reacting gas's own concentration unchanged — so Q is unchanged and there is no shift. What matters is the partial pressures of the species in the reaction, not the total.

3. Changing the temperature

Temperature is the special one — it's the only change that actually alters the value of K. The trick is to treat heat as a reactant or a product:

  • Exothermic reaction (releases heat): heat is a product. Raising the temperature is like adding product, so the equilibrium shifts reverse (left) and K decreases.
  • Endothermic reaction (absorbs heat): heat is a reactant. Raising the temperature is like adding reactant, so it shifts forward (right) and K increases.

Lowering the temperature does the opposite in each case.

⚠️ A catalyst does NOT shift the equilibrium
A catalyst speeds up the forward and reverse reactions equally. It helps the system reach equilibrium faster, but it does not change where that equilibrium sits and does not change K. Catalysts affect the rate, never the position.
📝 Worked example: The Haber process, N₂(g) + 3H₂(g) ⇌ 2NH₃(g), is exothermic. Predict the effect of (a) adding N₂, (b) increasing the pressure, and (c) raising the temperature.
  1. (a) Adding N₂ is adding a reactant, so the system shifts right to consume some of it — more NH₃ forms.
  2. (b) The left side has 4 mol of gas, the right has 2. Higher pressure favours fewer gas molecules, so it shifts right — again more NH₃.
  3. (c) Exothermic means heat is a product. Raising the temperature adds 'product' heat, shifting the equilibrium left and lowering K — less NH₃ at equilibrium.
✓ (a) right, (b) right, (c) left. Real plants use high pressure but only a moderate temperature — a compromise between yield and speed.
✏️ Practice: For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), increasing the pressure shifts the equilibrium toward the side with fewer gas molecules. How many gas molecules are on the product (right-hand) side?
molecules
Solution
  1. Count the gas molecules on each side using the coefficients.
  2. Reactants: 1 N₂ + 3 H₂ = 4 molecules. Products: 2 NH₃ = 2 molecules.
  3. Because the right side has fewer (2 < 4), raising the pressure shifts the equilibrium toward the products.

Check your understanding

1. You add more reactant to a system at equilibrium. Which way does it shift?
Adding reactant is a stress the system counteracts by consuming some of it — it shifts toward the products (forward).
2. Which of these changes actually changes the value of K?
Only temperature changes K. Concentration and pressure/volume changes shift the position but leave K fixed; a catalyst changes neither.
3. What does adding a catalyst do to a reaction at equilibrium?
A catalyst speeds the forward and reverse reactions equally, so equilibrium is reached faster but its position and K are unchanged.
4. An exothermic reaction is at equilibrium. You raise the temperature. What happens?
For an exothermic reaction, heat is a product. Adding heat shifts the equilibrium reverse (toward reactants) and lowers K.
🔎 Related lessons
Dynamic Equilibrium & K The Reaction Quotient Q ICE-Table Calculations pH, pOH & Strong Acids/Bases
✅ Key takeaways
  • Le Châtelier's principle: a disturbed equilibrium shifts to partially counteract the disturbance.
  • Add a substance → shift away from it; remove one → shift toward it.
  • Raise the pressure → shift toward the side with fewer moles of gas (equal moles → no shift).
  • Temperature is the only change that alters K — treat heat as a reactant (endothermic) or product (exothermic).
  • A catalyst reaches equilibrium faster but does NOT shift its position or change K.
➡️ You can now predict the direction of any shift. The next step is putting numbers on it — calculating the actual equilibrium concentrations from the starting amounts and K. The ICE table is the bookkeeping tool that makes those calculations routine.
Want to test yourself on this? Try the Chemistry practice test →