Le Chatelier's Principle: Definition and Examples

Le Chatelier’s principle answers a practical question: if you change something in a chemical equilibrium, which way does the reaction go? According to Jim Clark’s page on Chemistry LibreTexts, the French chemist and engineer Henry-Louis Le Chatelier proposed the idea in 1884, and the page calls it one of the central concepts of chemical equilibrium.

Le Chatelier’s principle: definition and statement

A chemical equilibrium is dynamic. The forward and reverse reactions both keep running. Change the conditions and the position shifts to oppose the change, until a new equilibrium sets in.

Clark warns that this is a guide for predicting what happens and that it does not say why at the molecular level.

Effect of concentration

When the products increase, the reaction quotient Qc grows and rises above the equilibrium constant Kc. Adding a reactant, say A in the reaction A + 2B ⇌ C + D, sends the equilibrium to the right and uses up part of what was added.

Taking a product away as soon as it forms pushes the equilibrium further and further toward the products, until the reaction is effectively one-way. Chemguide points out that this is handy when one reactant is expensive: use a cheap one in excess and the costly one gets converted as fully as possible.

Practical example: in the iron thiocyanate demonstration (Fe³⁺ + SCN⁻ ⇌ FeSCN²⁺), adding ferric or thiocyanate ions moves the equilibrium toward the complex, and the color deepens where the addition fell. A substance that binds Fe³⁺, such as a phosphate, removes a reactant. The equilibrium goes left and the solution fades. This comes from the UC Davis demonstration page.

Effect of pressure

Pressure matters only when gases are involved. Chemguide explains that pressure depends on the number of gas molecules, so raising it sends the reaction to whichever side has fewer moles of gas, where the system can lower the pressure by making fewer molecules. Shrinking the volume pushes toward the side with fewer moles, and enlarging it pushes toward the side with more.

If both sides have the same moles of gas, pressure does not shift the equilibrium. The same goes for an inert gas added at constant volume: the total pressure goes up, yet the partial pressures of the reacting gases stay as they were and the constant is unchanged. The link between pressure, volume and moles of gas is the one covered in gas laws and kinetic theory.

Practical example: in ammonia synthesis, N₂ + 3H₂ ⇌ 2NH₃, there are 4 moles of gas on the left and 2 on the right. In a textbook example reported by Clark, going from a 1.5 L flask to a 5 L flask lowers the pressure, and the shift goes toward the reactants, where there are more moles of gas.

Effect of temperature

According to Clark, its main effect is to change the value of the equilibrium constant. Heating a system at equilibrium favors the endothermic reaction, which absorbs the extra heat, and cooling it favors the exothermic one.

Clark admits that treating heat as a reactant or a product is strictly incorrect, but he accepts it as a convenient shortcut for predicting the shift. Chemguide offers a teaching reaction whose forward step is exothermic (250 kJ per mole of A): heating from 300 to 500 °C sends the equilibrium left, toward the reaction that absorbs heat, and more A and B remain.

Disturbance	Where the equilibrium goes	Does K change?
More reactant	Toward the products	No
Product removed	Toward the products	No
Higher pressure (gases)	Toward the side with fewer moles of gas	No
Inert gas at constant volume	No shift	No
Heating	Toward the endothermic reaction	Yes
Cooling	Toward the exothermic reaction	Yes
Catalyst	No shift	No

Practical example: in the equilibrium N₂O₄ ⇌ 2NO₂, endothermic toward the right with ΔH° = +58 kJ/mol according to the RSC, heating moves the equilibrium toward NO₂ and the brown color deepens. Raising the pressure sends it toward N₂O₄, which has fewer moles of gas, and the color fades. The RSC reports that Kp rises with temperature and does not depend on pressure. In the table on George Bodner’s page, for 2NO₂ ⇌ N₂O₄, Kp is 58 at 0 °C and 0.066 at 100 °C (values estimated from standard thermodynamic data; the formation of N₂O₄ falls as it warms), so K changes by orders of magnitude with temperature.

Chemical equilibrium and Le Chatelier’s principle in industry: the Haber process

To make ammonia, the Haber process takes nitrogen from the air and hydrogen, produced mostly from methane. The reaction is reversible and exothermic, and nitrogen and hydrogen enter the reactor in a 1 to 3 ratio by volume, as the equation requires. Chemguide, in a page updated in April 2013, gives a reference pressure of 200 atmospheres. It is always high because the ammonia side has fewer moles of gas, and higher pressures call for sturdier vessels and cost too much.

For the best yield at equilibrium the temperature would have to be as low as possible, but Chemguide explains that the reaction is too slow when it is colder. Plants therefore run at roughly 400-450 °C and accept a smaller yield: each pass through the reactor converts about 15% of the nitrogen and hydrogen, with figures that vary from plant to plant. Ammonia liquefies easily under pressure if it is not too hot, and that is how it is separated from the unreacted gases, which stay gaseous and are recycled. It also amounts to removing the product continuously. With recycling, the overall conversion rises to about 98%, according to Chemguide.

A 2026 review by Arroyo-Caire and colleagues in Molecules describes ammonia production as still based mainly on the Haber-Bosch process at 400-500 °C and 150-250 bar. The authors trace the harsh conditions to kinetics, meaning speed, because the N≡N bond is very stable. According to the review, research aims at catalysts that work at 200-300 °C and 20-30 bar.

What Le Chatelier’s principle leaves out

A catalyst falls outside the principle. According to Clark it does not change the position of equilibrium, because it speeds up the forward and reverse reactions equally. For a very slow reaction, getting to equilibrium can take years, and a catalyst shortens that. In the Haber process, iron with a promoter makes the reaction fast enough and leaves the percentage of ammonia at equilibrium as it was.

Chemists also voice reservations. Eric Scerri, in a 2015 opinion piece on RSC Education, observes that textbooks do not give two identical statements of the principle, and that many replace “counter” with “accommodate” the change, to avoid wrong predictions. For changes in concentration the classic wording seems to work well, he writes, which adds to the confusion, but temperature needs another approach because K itself depends on temperature. Scerri proposes replacing the principle with a calculation of Kp, using mole fractions and total pressure.