Le Chatelier's principle
This article or subsequent section is not sufficiently supported by evidence (e.g., anecdotal evidence). Information without sufficient evidence may be removed in the near future. Please help Wikipedia by researching the information and adding good supporting evidence.
This article deals with the principle of least constraint according to Le Chatelier. The principle according to Gauss can be found under the lemma Principle of least constraint.
Le Chatelier's principle, also called the principle of least constraint, was formulated by Henry Le Chatelier and Ferdinand Braun between 1884 and 1888:
"If you apply a constraint to a chemical system at equilibrium, it will react in such a way that the effect of the constraint becomes minimal."
or more specifically:
"If a constraint is exerted on a system which is in chemical equilibrium by a change in external conditions, a new equilibrium, evading the constraint, is established as a result of this disturbance of the equilibrium."
The principle is thus very general, so that it does not allow any quantitative statements. Nevertheless, it is often used because a qualitative prediction is sufficient for first steps in many areas. Furthermore, it is very easy to apply.
Examples:
- The constraint of increasing or decreasing temperature is avoided by heat consumption or generation.
- If the pressure is increased for a mixture of liquid and gas (forcing), part of the gas passes into the liquid phase (smaller particle distances → less in the gas phase)
"Constraints" in this sense are changes in temperature, pressure, or mass concentration:
- If you increase the temperature, the heat-producing reaction is pushed back and vice versa.
- If you increase the pressure, the system deviates in such a way that the volume-reducing reaction is promoted, and vice versa.
- If the concentration is changed, e.g. by removing a product from the preparation, the equilibrium system reacts by producing this product again.
The correctness of this concept can be confirmed both empirically, i.e. in experiments, and by calculations of temperature, pressure and concentration dependence of the free enthalpy of reaction.
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 |
| endothermic | Products | |
| Temperature reduction | exothermic | Products |
| endothermic | Educts |
The gas mixture from the equilibrium between the brown nitrogen dioxide and the colourless dinitrogen tetroxide can serve as an example:
The enthalpy of the outward reaction is Δ 
, i.e., it is an exothermic reaction because energy is released. The reverse reaction is endothermic: Δ 
.
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:
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.
