What Is Le Chatelier's Principle? Predicting the Shift

You have an equilibrium. You poke it. Which way does it move? There's one rule that answers this for every disturbance you'll be asked about, and it fits in a sentence.

The short answer: Le Chatelier's principle states that if a system at equilibrium is disturbed, the position of equilibrium shifts in the direction that partially opposes the change. Add something and the system consumes some of it; heat it and it absorbs some of that heat.

What "opposes the change" actually means

The system doesn't undo the change — it can't. It partially offsets it. Add 1 mol of a reactant and the shift will use up some of that mol, so the new equilibrium has more of it than before, just less than 1 mol more.

Think of it as pushing back, not resetting.

The three disturbances

1. Concentration

You do this	Equilibrium shifts	Because
Add a reactant	Right (towards products)	To use some of it up
Remove a reactant	Left	To replace some of it
Add a product	Left	To use some of it up
Remove a product	Right	To replace some of it

Continuously removing a product is the standard industrial trick for dragging a reaction to a high yield — the equilibrium never stops shifting right because the product never accumulates.

2. Pressure (gases only)

Count the moles of gas on each side. Raising the pressure shifts the equilibrium towards the side with fewer moles of gas, because that reduces the pressure again.

For N₂ + 3H₂ ⇌ 2NH₃, that's 4 moles of gas on the left and 2 on the right. High pressure favours ammonia.

If both sides have the same number of gas moles — as in H₂ + I₂ ⇌ 2HI, 2 and 2 — pressure has no effect on the position at all. Solids and liquids don't count in the tally.

3. Temperature

Look at which direction is exothermic and which endothermic. Raising the temperature shifts the equilibrium in the endothermic direction, because that absorbs the added heat.

For an exothermic forward reaction (ΔH negative), heating pushes it backwards, lowering the yield. Cooling pushes it forwards.

Temperature is special in one important way: it is the only one of the three that changes Kc. Concentration and pressure changes move the position of equilibrium while leaving Kc alone; a temperature change alters Kc itself.

And what does nothing

• A catalyst. It speeds up forward and reverse reactions equally, so equilibrium arrives sooner at exactly the same position. This is examined relentlessly.
• An inert gas added at constant volume. Total pressure rises, but the partial pressures of the reacting gases are unchanged, so nothing shifts.

The Haber process, worked through

N₂ + 3H₂ ⇌ 2NH₃ ΔH ≈ −92 kJ mol⁻¹

Apply the principle to each variable and see where the theory and the factory disagree.

• Pressure: 4 mol gas → 2 mol gas, so high pressure favours ammonia. Industry uses around 200 atm. Higher would be better still, but the cost and safety demands of the vessels put a ceiling on it.
• Temperature: the forward reaction is exothermic, so low temperature favours ammonia. But low temperature also makes the reaction unworkably slow. The compromise is roughly 400–450 °C — a deliberately reduced yield, obtained fast enough to be worth having.
• Catalyst: iron. It doesn't improve the yield at all; it just gets the mixture to equilibrium quickly enough for a continuous process.
• Removing product: the ammonia is condensed out as a liquid and drawn off, and the unreacted N₂ and H₂ are recycled. Removing product keeps the equilibrium shifting right.

That temperature choice is the classic exam question, and the answer is always the same shape: Le Chatelier gives the ideal condition, kinetics vetoes it, and the plant runs a compromise. The Contact process, 2SO₂ + O₂ ⇌ 2SO₃ (ΔH ≈ −196 kJ mol⁻¹, V₂O₅ catalyst, about 450 °C), is built on exactly the same trade-off.

Worked examples

Predict each before reading on.

1. N₂O₄ ⇌ 2NO₂ (ΔH positive, forward endothermic). The sealed tube is warmed. What do you see?
Shift right, towards the endothermic direction — the mixture goes darker brown.

2. Same tube, now compressed into half the volume. Which way?
Towards fewer moles of gas: 2 → 1, so left. Watch carefully, though: it darkens instantly as everything is concentrated, then fades over a few seconds as NO₂ pairs up — but it settles darker than before you compressed it. The shift only partly offsets the squeeze, which is the principle in one observation.

3. CH₃COOH ⇌ CH₃COO⁻ + H⁺. You add sodium ethanoate, which supplies CH₃COO⁻. What happens to [H⁺]?
Adding a product shifts it left, so [H⁺] falls and pH rises. This is the common-ion effect — and it's how a buffer gets set up in the first place.

4. [Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O, forward endothermic, pink ⇌ blue. Add concentrated HCl.
More Cl⁻ shifts it right — pink to blue. Warming it does the same thing.

5. A catalyst is added to a low-yield equilibrium to improve the yield.
It won't. Same position, reached faster.

Common mistakes to avoid

• Saying the system "cancels" the change. It only partially opposes it. After adding a reactant there is still more of that reactant present than before.
• Using pressure arguments where there's no gas. Only count moles of gas. Solids and liquids are ignored, and if the gas moles are equal on both sides, pressure does nothing.
• Letting a catalyst shift the equilibrium. It changes the rate, never the position. Nor does it change Kc.

FAQ

What is Le Chatelier's principle in simple terms?
If you disturb a system at equilibrium, it shifts in whichever direction partially opposes the disturbance.

How does temperature affect equilibrium?
Increasing temperature shifts the equilibrium in the endothermic direction; decreasing it shifts it in the exothermic direction. Temperature is also the only factor that changes the equilibrium constant.

Why doesn't a catalyst shift the equilibrium?
Because it speeds up the forward and reverse reactions by the same factor. Equilibrium is reached sooner, at exactly the same position.

Why is the Haber process run hot if low temperature gives more ammonia?
Because at low temperature the rate is far too slow to be useful. Around 450 °C is a compromise between a good yield and a workable speed.

The takeaway

One sentence covers the whole topic: an equilibrium shifts to partially oppose whatever you do to it. Add or remove a substance, squeeze the gases, or change the temperature, and you can predict the direction every time — while remembering that a catalyst changes only how quickly you get there.

Background → [What Is Chemical Equilibrium?] (sibling) and [Reversible vs Irreversible Reactions] (sibling). See also → What Is a Catalyst? How Reactions Get a Shortcut for why it can't shift anything, and [What Is a Buffer?] (sibling) for Le Chatelier doing useful work.

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