Le Chatelier’s principle predicts the effect of a change on system at dynamic equilibrium. Changing the conditions of a system at thermodynamic equilibrium (concentration, temperature, pressure, volume, etc.) causes the system to react in a way that counteracts the change and establishes a new equilibrium. While originally described for chemical reactions, Le Chatelier’s principle also applies to homeostasis in biology, economics, pharmacology, and other disciplines. Other names for Le Chatelier’s principle are Chatelier’s principle or the Equilibrium Law.
The Basics of Le Chatelier’s Principle
- The principle is credited to French chemist Henry Louis Le Chatelier and sometimes also to German scientist Karl Ferdinand Braun, who discovered it independently.
- Le Chatelier’s principle helps you predict the direction of the response to a change in equilibrium.
- The principle does not explain the reason why the equilibrium shifts, only the direction of the shift.
- Concentration: Increasing the concentration of reactants shifts the equilibrium to produce more products. Increasing the concentration of products shift the equilibrium to make more reactants.
- Temperature: The direction of equilibrium shift resulting from a temperature change depends on which reaction is exothermic and which is endothermic. Increasing temperature favors the endothermic reaction, while decreasing temperature favors the exothermic reaction.
- Pressure/Volume: Increasing the pressure or volume of a gas shifts the reaction toward the side with fewer molecules. Decreasing the pressure or volume of a gas shifts the reaction toward the side with more molecules.
Remember, Le Chatelier’s principle states that the equilibrium shifts toward the side of a reversible reaction that opposes the change. The equilibrium constant for the reaction does not change.
As an example, consider the equilibrium reaction where carbon dioxide and hydrogen gas react and form methanol:
CO + 2 H2 ⇌ CH3OH
If you increase the concentration of CO (a reactant), the equilibrium shifts to produce more methanol (a product), thus reducing the amount of carbon monoxide. Collision theory explains the process. When there is more CO, the frequency of successful collisions between reactant molecules increases, generating more product. Increasing the concentration of hydrogen has the same effect.
Decreasing the concentration of carbon monoxide or hydrogen has the opposite effect. The equilibrium shifts to compensate for the reduced reactants, favoring the decomposition of methanol into its reactants.
Increasing the amount of methanol favors the formation of reactants. Decreasing the concentration of methanol increases its formation. So, removing a product from a system aids in its production.
Le Chatelier’s principle predicts the equilibrium shift when you increase or decrease the pressure of a reaction involving gases. Note that the equilibrium constant for the reaction does not change. Increasing the pressure shifts the reaction in a way that reduces pressure. Decreasing pressure shifts the reaction in a way that increases pressure. The side of the reaction with more molecules exerts a greater pressure than the side of the reaction with fewer molecules. The reason is that the more molecules there are hitting the walls of a container, the higher the pressure.
For example, consider the general reaction:
A (g) + 2 B (g) ⇌ C (g) + D (g)
There are three moles of gas (1 A and 2 B) on the left side of the reaction arrow (reactants) and two moles of gas (1 C and 1 D) on the product side of the reaction arrow. So, if you increase the pressure of the reaction, the equilibrium shifts toward the right (fewer moles, lower pressure). If you increase the pressure of the reaction, the equilibrium shifts toward the left (more moles, higher pressure).
Adding an inert gas, such as helium or argon, at constant volume does not cause a shift in equilibrium. Even though the pressure increases, the unreactive gas does not participate in the reaction. So, Le Chatelier’s principle applies when the partial pressure of a reactant or product gas changes. If you add an inert gas and allow the gas volume to change, then adding this gas decreases the partial pressure of all gases. In this case, the equilibrium shifts toward the side of the reaction with the greater number of moles.
Unlike changing concentration or pressure, changing the temperature of a reaction shifts the magnitude of the equilibrium constant. The direction of the equilibrium shift depends on the enthalpy change of the reaction. In a reversible reaction, one direction is an exothermic reaction (evolves heat and has a negative ΔH) and the other direction is an endothermic reaction (absorbs heat and has a positive ΔH). Adding heat to a reaction (increasing the temperature) favors the endothermic reaction. Removing heat (lowering the temperature) favors the exothermic reaction.
For example, consider the general reaction:
A + 2 B ⇌ C + D; ΔH = -250 kJ/mol
The forward reaction (forming C and D) is exothermic, with a negative ΔH value. So, you know the reverse reaction (forming A and B) is endothermic. If you increase the temperature of the reaction, the equilibrium shifts to favor the endothermic reaction (C + D form A + B). If you decrease the temperature of the reaction, the equilibrium shifts to favor the exothermic reaction (A + 2 B forms C + D).
Le Chatelier’s Principle and Catalysts
Le Chatelier’s principle does not apply to catalysts. Adding a catalyst does not shift the equilibrium of a chemical reaction because it increases the rates of the forward and reverse reactions equally.
Le Chatelier’s Principle Example Problem
For example, predict the effect when changes occur in the reaction where gaseous SO3 decomposes into SO2 and O2:
2 SO3 (g) ⇌ 2 SO2 (g) + O2 (g); ΔH = 197.78 kJ/mol
(a) What happens if you increase the temperature of the reaction?
The shift of equilibrium favors the forward reaction because the decomposition reaction is endothermic.
(b) What happens if you increase the pressure on the reaction?
Increasing the pressure favors the side of the reaction with fewer moles of gas because it reduces the pressure, so the equilibrium shifts to the left (the reactant, SO3).
(c) What happens if you add more O2 to the reaction at equilibrium?
Adding more oxygen shifts the equilibrium toward forming the reactant (SO3).
(d) What happens if you remove SO2 from the reaction at equilibrium?
Removing SO2 shifts the equilibrium toward forming the products (SO2 and O2).
- Atkins, P.W. (1993). The Elements of Physical Chemistry (3rd ed.). Oxford University Press.
- Callen, H.B. (1985). Thermodynamics and an Introduction to Thermostatistics (2nd ed.) New York: Wiley. ISBN 0-471-86256-8.
- Le Chatelier, H.; Boudouard, O. (1898), “Limits of Flammability of Gaseous Mixtures.” Bulletin de la Société Chimique de France (Paris). 19: 483–488.
- Münster, A. (1970). Classical Thermodynamics (translated by E.S. Halberstadt). Wiley–Interscience. London. ISBN 0-471-62430-6.
- Samuelson, Paul A (1983). Foundations of Economic Analysis. Harvard University Press. ISBN 0-674-31301-1.