Le Chatelier's Principle

What is it?

Le Chatelier’s principle: if a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the position of equilibrium shifts in the direction that partly counteracts the change.

“Shifts to the right” means more products form; “shifts to the left” means more reactants form. The system never fully undoes the change; it only reduces its effect.

Key idea

Changes in concentration and pressure shift the equilibrium without changing K: they make $Q \ne K$ , and the reaction runs until $Q = K$ again. A change in temperature changes K itself. A catalyst changes neither.

Why does it matter?

Industry. Chemists choose temperatures and pressures that push equilibria toward the product, as in the Haber process for ammonia.
Everyday chemistry. A fizzy drink goes flat when opened because lowering the pressure shifts $\ce{CO2(aq) <=> CO2(g)}$ toward the gas.
Living systems. The body’s buffers respond to added acid or base following the same principle.

How does it work?

1. Changing a concentration

Add a reactant (or remove a product): $Q$ becomes smaller than $K$ , so the equilibrium shifts right to use some of it up.
Add a product (or remove a reactant): $Q > K$ , so it shifts left.

2. Changing the pressure (gases)

Increasing the pressure by reducing the volume shifts the equilibrium toward the side with fewer moles of gas, which lowers the pressure:

\ce{N2(g) + 3H2(g) <=> 2NH3(g)}

The left side has 1 mol + 3 mol = 4 mol of gas; the right side has 2 mol.

Higher pressure shifts this equilibrium right, toward ammonia. If both sides have the same number of moles of gas (as in $\ce{H2 + I2 <=> 2HI}$ ), pressure has no effect on the position. Adding an inert gas at constant volume does not change any concentration, so it has no effect either.

3. Changing the temperature

Treat heat as a reactant or product:

Exothermic forward reaction (ΔH negative): heat is a “product”. Raising the temperature shifts the equilibrium left and decreases K.
Endothermic forward reaction (ΔH positive): heat is a “reactant”. Raising the temperature shifts it right and increases K.

For $\ce{N2O4(g) <=> 2NO2(g)}$ , ΔH = +57 kJ: on heating, more brown $\ce{NO2}$ forms, so the colourless-to-brown mixture gets darker.

4. Adding a catalyst

A catalyst speeds up the forward and reverse reactions equally. Equilibrium is reached sooner, but its position and $K$ are unchanged.

Change	Effect on position	Effect on K
Add reactant / remove product	shifts right	none
Add product / remove reactant	shifts left	none
Increase pressure (smaller volume)	toward fewer moles of gas	none
Increase temperature	in the endothermic direction	changes
Add a catalyst	none (equilibrium reached faster)	none

Think of it like this

Think of a crowded room with two doors and people moving between it and the corridor at equal rates. Push more people into the room and, for a while, more people leave than enter, until the flows balance again. The room ends up a bit more crowded than before, but less crowded than right after the push: the change is partly counteracted.

More precisely

Le Chatelier’s principle predicts the direction of a shift; comparing $Q$ with $K$ explains why, and an ICE table gives the new concentrations. How strongly $K$ depends on temperature is described by the van ‘t Hoff equation, which links it to ΔH°.

Visualise it

Graph of concentration against time for H2 + I2 forming 2HI at equilibrium, Kc = 49.0. At first HI is constant at 0.156 M and H2 and I2 at 0.0222 M. Then 0.100 M H2 is added: H2 jumps up and then falls to 0.107 M, I2 falls to about 0.0067 M, and HI rises to 0.187 M, until a new equilibrium is reached. — Adding H₂ shifts the equilibrium to the right: some of the added H₂ is used up and more HI forms. Calculated from a simple rate model.

Worked example

Worked example: Predicting shifts in the Haber process

Question: For $\ce{N2(g) + 3H2(g) <=> 2NH3(g)}$ , ΔH = −92 kJ, predict the effect on the amount of ammonia of (a) adding $\ce{N2}$ (b) increasing the pressure (c) increasing the temperature (d) adding an iron catalyst.

Adding $\ce{N2}$ (a reactant): shifts right, so more $\ce{NH3}$ .
Higher pressure: 4 mol gas on the left, 2 mol on the right, so it shifts toward fewer moles: right, more $\ce{NH3}$ .
The forward reaction is exothermic, so heating shifts left: less $\ce{NH3}$ (and $K$ decreases).
A catalyst does not shift the equilibrium: no change in the amount, but equilibrium is reached faster.

Worked example: Checking a shift with Q

Question: An equilibrium mixture of $\ce{H2(g) + I2(g) <=> 2HI(g)}$ ( $K_\text{c} = 49.0$ ) contains 0.0222 M $\ce{H2}$ , 0.0222 M $\ce{I2}$ and 0.156 M HI. 0.100 M $\ce{H2}$ is added. Which way does the equilibrium shift?

Just after the addition, $[\ce{H2}] = 0.0222\ \text{M} + 0.100\ \text{M} = 0.122\ \text{M}$
$Q = \dfrac{(0.156\ \text{M})^2}{(0.122\ \text{M})(0.0222\ \text{M})} = 8.99$ (the units of M² cancel)
$Q = 8.99 < K_\text{c} = 49.0$ , so the reaction goes forward: $\ce{H2}$ and $\ce{I2}$ are used up and more HI forms, as Le Chatelier’s principle predicts.

Common mistake

Common mistake: Thinking concentration changes alter K

Adding a reactant shifts the position, but $K$ stays the same. Only temperature changes $K$ .

Common mistake: Thinking a catalyst increases the yield

A catalyst makes equilibrium arrive faster but does not change how much product is present at equilibrium.

Common mistake: Counting all moles, not moles of gas

For pressure changes, count only gases. In $\ce{CaCO3(s) <=> CaO(s) + CO2(g)}$ there is 0 mol of gas on the left and 1 mol on the right.

Notation note

“Shifts to the right / forward” = more products; “to the left / backward” = more reactants.
ΔH refers to the forward reaction as written; the reverse reaction has the opposite sign.

Remember this

A disturbed equilibrium shifts to partly counteract the change.
Add reactant or remove product: right. Higher pressure: fewer moles of gas.
Heating favours the endothermic direction and is the only change that alters $K$ .
A catalyst gives equilibrium faster but does not move it.

Practise this topic with flashcards and a quiz at chemistryclarity.com/chemistry/le-chateliers-principle/