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Some reactions never quite finish. Mix nitrogen and hydrogen in a sealed container under the right conditions and ammonia forms — but some of it immediately breaks back down into nitrogen and hydrogen. After a while the amounts stop changing, even though molecules are still reacting in both directions.
That state is called dynamic equilibrium, and a rule stated by the French chemist Henri Le Chatelier in 1884 lets you predict how it responds when you disturb it.
What equilibrium means
A reversible reaction is written with a double arrow:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
At equilibrium:
- the forward and reverse reactions are happening at the same rate, and
- so the concentrations of reactants and products stay constant (not necessarily equal!).
It’s “dynamic” because reactions are still happening. It’s a balance of two rates, not a standstill.
The principle
If a system at equilibrium is disturbed, the equilibrium shifts in the direction that tends to counteract the disturbance.
The system doesn’t fully undo the change, but it partly compensates for it. Here’s what that means for each kind of change.
1. Changing concentration
- Add more reactant → the equilibrium shifts right (towards products) to use some of it up.
- Remove product → the equilibrium shifts right to replace it.
- Add more product → shifts left.
In the Haber process, ammonia is continuously cooled and liquefied out of the gas mixture. Removing the product keeps pulling the equilibrium towards making more.
A nice visible example is the iron(III) thiocyanate equilibrium:
Fe³⁺ (pale yellow) + SCN⁻ ⇌ FeSCN²⁺ (blood red)
Add more Fe³⁺ or SCN⁻ and the solution turns deeper red.
2. Changing pressure (gases only)
Pressure changes only matter when the two sides have different numbers of gas molecules.
- Increase pressure → shifts towards the side with fewer gas molecules (to reduce the pressure).
- Decrease pressure → shifts towards the side with more gas molecules.
For N₂ + 3H₂ ⇌ 2NH₃, there are 4 gas molecules on the left and 2 on the right. High pressure favours ammonia.
For H₂ + I₂ ⇌ 2HI, there are 2 molecules on each side, so pressure has no effect on the position of equilibrium.
Adding an unreactive gas like argon at constant volume raises the total pressure but doesn’t change the partial pressures of the reacting gases — so it has no effect either.
3. Changing temperature
Temperature is special: it’s the only change that alters the equilibrium constant itself. The key is whether the forward reaction releases or absorbs heat (see exothermic vs. endothermic reactions).
Treat heat as if it were a reactant or product:
- For an exothermic forward reaction, heat is a product. Raising the temperature shifts equilibrium left; lowering it shifts right.
- For an endothermic forward reaction, heat is a reactant. Raising the temperature shifts equilibrium right.
The ammonia synthesis is exothermic, so lower temperatures give a higher proportion of ammonia at equilibrium.
A visible demonstration: brown nitrogen dioxide and colourless dinitrogen tetroxide.
2NO₂ (brown) ⇌ N₂O₄ (colourless), exothermic forward
Put a sealed tube in ice and it fades; put it in hot water and it turns dark brown.
4. Adding a catalyst
A catalyst speeds up the forward and reverse reactions equally. It gets the system to equilibrium faster, but it doesn’t shift the equilibrium position or change the yield at equilibrium. More on this in reaction rates and catalysts.
The Haber process: a real compromise
Making ammonia from nitrogen and hydrogen feeds much of the world, through fertilizers. Le Chatelier’s principle says:
- High pressure favours ammonia (fewer gas molecules). ✓
- Low temperature favours ammonia (exothermic). ✓
But at low temperature the reaction is painfully slow, even with a catalyst, and very high pressure is expensive and dangerous. So industry compromises: around 400–450 °C and 150–300 atmospheres, with an iron catalyst, continuous removal of ammonia, and recycling of the unreacted gases. Each pass converts only about 15% of the gases to ammonia, but recycling brings the overall conversion to about 97%.
Quick answers
Does “equilibrium” mean equal amounts of reactants and products? No. It means the amounts are constant. Some equilibria lie far to the right (mostly products), others far to the left.
Does Le Chatelier’s principle change the equilibrium constant? Only a temperature change does. Concentration and pressure changes shift the position but leave K the same.
Why do catalysts not change the yield? Because they lower the energy barrier for both directions equally.
Related
Equilibrium ideas are also behind weak acids and buffers, where adding acid or base shifts an equilibrium that then resists the pH change.
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