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Le Châtelier's principle

A qualitative principle in chemistry and physics describing how a system at equilibrium shifts to oppose applied changes in concentration, pressure, temperature, or other external constraints.

Author: Leandro Alegsa Creation date: November 13, 2022 Last updated: May 14, 2026

en.wikipedia.org · Public domain

Le Châtelier's principle is a qualitative rule used to predict the direction in which a system at dynamic equilibrium will respond when it is subjected to an external change or stress. The principle states that a disturbed equilibrium will shift in the direction that tends to reduce the effect of the disturbance. It applies to chemical equilibria, phase equilibria and many physical processes in which balance between opposing influences is established.

Common stresses and typical responses

Typical disturbances and the expected Le Châtelier responses include:

Concentration: adding a reactant or removing a product generally shifts the equilibrium toward products; removing a reactant or adding a product shifts it toward reactants.
Pressure and volume: for equilibria involving gases, increasing pressure (or decreasing volume) favors the side with fewer moles of gas; decreasing pressure favors the side with more gas moles.
Temperature: changing temperature shifts an equilibrium depending on the reaction enthalpy. Heating favors the endothermic direction, cooling favors the exothermic direction. Temperature changes also alter the equilibrium constant.

How the principle relates to thermodynamics

Le Châtelier's principle is qualitative and does not give the magnitude of the shift. A quantitative description uses the reaction quotient Q and the equilibrium constant K: the system shifts to make Q approach K. Thermodynamic relations, such as changes in Gibbs free energy, give the numerical direction and extent of change; Le Châtelier provides an intuitive guide to the direction only.

Applications and examples

In industry the principle guides conditions for optimum yield. For example, the Haber process (N2 + 3 H2 ⇌ 2 NH3) is operated at high pressure to favor ammonia formation and at temperatures chosen to balance yield and reaction rate. Continuous removal of a desired product is another practical way to drive a reaction forward.

Limitations and cautions

The principle does not apply to systems far from equilibrium, to kinetically controlled reactions where rates limit product formation, or to open systems exchanging matter without reaching a closed equilibrium. When multiple stresses act together, or when activities rather than concentrations are important, quantitative thermodynamic analysis is needed for accurate predictions.

Historical note

The rule is named after the French chemist Henri Louis Le Châtelier, who formulated the idea in the late 19th century. It remains a foundational heuristic in chemistry education and practice.

Temperature change

→ Main article: Van 't Hoff equation

Heat addition and heat removal cause a shift in equilibrium, i.e. the setting of a new equilibrium with changed concentrations. Heat extraction favours the heat-supplying (exothermic) reaction, heat supply the heat-consuming (endothermic) reaction. As a result, the temperature change of the system is less than without an equilibrium shift.

A change in temperature always leads to a change in the equilibrium concentrations. Which concentration increases or decreases depends on whether the formation of the products is exothermic or endothermic:

Malfunction	Nature of the reaction	Increase in
Temperature increase	exothermic	Educts
Temperature increase	endothermic	Products
Temperature reduction	exothermic	Products
Temperature reduction	endothermic	Educts

The gas mixture from the equilibrium between the brown nitrogen dioxide and the colourless dinitrogen tetroxide can serve as an example:

${\ce {2 NO2 -> N2O4}}$

The enthalpy of the outward reaction is Δ $\Delta {\overrightarrow {H}}=-58{\rm {\tfrac {kJ}{mol}}}$ , i.e., it is an exothermic reaction because energy is released. The reverse reaction is endothermic: Δ $\Delta {\overleftarrow {H}}=+58{\rm {\tfrac {kJ}{mol}}}$ .

If the temperature is increased while the volume remains constant, the reaction will take place in the opposite direction, i.e. in the endothermic direction, with the result that the equilibrium shifts to the left and the gas mixture becomes darker. Lowering the temperature causes the exothermic reaction, whereby the equilibrium shifts to the right and the gas mixture lightens.

Volume or pressure change

The chemical equilibrium of reactions in which no gases are involved is hardly influenced by an externally induced change in volume. If, on the other hand, gaseous substances are involved, the equilibrium is only influenced if the number of particles in the gas phase changes as a result of the equilibrium shift.

A change in pressure only affects equilibrium in a closed system. Depending on the reaction condition, one can see a pressure change or a volume change: The system reduces the pressure created by a reduction in volume by running off in favor of the side that has the smaller number of particles and thus requires the smaller volume. As a result, the pressure increase is less than if the gases were incapable of any reaction. Accordingly, an increase in volume shifts the equilibrium towards larger numbers of particles.

The position of the equilibrium can be influenced by an increase in pressure from outside:

at constant reaction volume by further addition of reactants
with variable reaction volume due to compression.

If the reaction takes place in an open system, the gas produced during the reaction can constantly escape. This constantly produces new gas, which in turn escapes. This disturbance of the equilibrium leads to the fact that it can not be set: the reaction proceeds completely to the product side.

A well-known reaction is the production of ammonia in the Haber-Bosch process from nitrogen and hydrogen:

${\ce {N2 + 3 H2 <=> 2 NH3}}$

Thus, 4 gas molecules on the educt side on the left, 2 gas molecules on the product side on the right are created. If the pressure is now increased, the system moves to the volume-reducing side - i.e. the side with fewer molecules. Thus, the formation of ammonia can be promoted by increasing the pressure.

The same principle can be applied to the nitrogen dioxide-nitrogen tetroxide equilibrium.