Principle of the smallest compulsion

The principle of Le Chatelier , also known as the principle of least constraint , was formulated by Henry Le Chatelier and Ferdinand Braun between 1884 and 1888:

"If you exert a compulsion on a chemical system in equilibrium , it reacts in such a way that the effect of the compulsion is minimal."

or more precisely:

"If one exerts a compulsion on a system that is in chemical equilibrium by changing the external conditions, a new equilibrium is established as a result of this disturbance of the equilibrium, evading the compulsion."

The principle is therefore very general, so that it does not allow any quantitative statements. Nevertheless, it is often used, as a qualitative prediction is sufficient for the first steps in many areas. It is also very easy to use.

Examples:

The compulsion to raise or lower the temperature is evaded with heat consumption or generation.
If the pressure is increased for a mixture of liquid and gas (compulsion), part of the gas changes into the liquid phase (smaller particle distances → less in the gas phase)

In this sense, "constraints" are changes in temperature , pressure or substance concentration :

If the temperature is increased, the heat-producing reaction is suppressed and vice versa.
If the pressure is increased, the system evades in such a way that the volume-reducing reaction is promoted, and vice versa.
If you change the concentration, e.g. B. by removing a product from the approach, the equilibrium system reacts in that this product is subsequently produced.

The correctness of this concept can be confirmed empirically, i.e. in experiments, as well as by calculating the temperature, pressure and concentration dependence of the free enthalpy of reaction .

Temperature change

Heat supply and heat withdrawal cause a shift in equilibrium, i. H. the setting of a new equilibrium with changed concentrations. Heat extraction favors the heat-supplying ( exothermic ) reaction, heat supply the heat-consuming ( endothermic ) reaction. As a result, the temperature change in the system is smaller than without a shift in equilibrium.

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:

Disorder	Type of response	Increase in
Temperature increase	exothermic	Educts
Temperature increase	endothermic	Products
Temperature decrease	exothermic	Products
Temperature decrease	endothermic	Educts

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

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

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

If the temperature is now increased while the volume is constant, the reaction will proceed in the opposite direction, i.e. in the endothermic direction, which shifts the equilibrium to the left and the gas mixture becomes darker. Lowering the temperature causes the exothermic reaction, which shifts the equilibrium to the right and the gas mixture brightens.

Change in volume or pressure

The chemical equilibrium of reactions in which no gases are involved is hardly influenced by an externally induced change in volume. If, however, gaseous substances are involved, the equilibrium is only influenced if the number of particles in the gas phase changes due to the shift in equilibrium.

The reaction constant changes with changes in pressure according to the following formula:

{\ displaystyle \ left ({\ frac {\ partial \ ln K} {\ partial p}} \ right) _ {T} = - {\ frac {\ Delta _ {R} {\ overline {V}} _ { \ text {fl}} ^ {0}} {RT}}}

, where is the molar reaction volume .

{\ displaystyle {\ overline {V}} _ {\ text {fl}} ^ {0}}

A change in pressure only affects equilibrium in a closed system . Depending on the reaction conditions, a change in pressure or a change in volume can be determined: The system reduces the pressure generated by a volume reduction by running down in favor of the side that has the lower number of particles and thus requires the smaller volume. As a result, the pressure increase is less pronounced than if the gases were not able to react. Correspondingly, an increase in volume shifts the equilibrium in the direction of larger particle numbers.

The position of equilibrium can be influenced from outside by increasing the pressure:

with a constant reaction volume by adding more starting materials

when the reaction volume changes due to compression .

If the reaction takes place in an open system , the gas produced during the reaction can continuously escape. As a result, new gas is constantly produced, which in turn escapes. This disturbance of the equilibrium means that it cannot adjust itself: 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 :

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

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

The same principle can also be applied to the nitrogen dioxide / nitrous tetroxide equilibrium.

Change of substance amount

By adding or removing a reactant, the equilibrium is disturbed; the reaction consequently runs increasingly in one direction until equilibrium is reached again. If you change the concentration of one of the substances involved in the equilibrium, the concentrations of all other partners also change. If an equilibrium reaction is to take place completely in favor of a product, it is sufficient to multiply one of the starting materials from the reaction mixture or to remove one of the products from the reaction mixture. The reverse reaction is prevented until the original equilibrium is restored.

Since the equilibrium only depends on the temperature and, if necessary, on the pressure , the reaction takes place after the change in concentration in such a way that the original equilibrium is restored. For an equilibrium reaction:

{\ displaystyle {\ ce {a A + b B <=> c C + d D}}}

With

{\ displaystyle K_ {C} = {\ frac {k _ {\ mathrm {hin}}} {k _ {\ mathrm {r {\ ddot {u}} ck}}}} = {\ frac {c ^ {c} ({\ text {C}}) \ cdot c ^ {d} ({\ text {D}})} {c ^ {a} ({\ text {A}}) \ cdot c ^ {b} ({ \ text {B}})}}}

the following cases can be distinguished:

modification	causes
Add A or B	Increase in products
Addition of C or D	Increase in starting materials
Withdrawal from A or B	Acceptance of the products
Withdrawal of C or D	Decrease in the starting materials

A change in the reaction conditions of temperature and pressure leads to a shift in the equilibrium and thus to a change in the equilibrium concentrations.

The influence of changes in the amount of substance on the position of an equilibrium can be illustrated by the esterification of carboxylic acids or the hydrolysis of carboxylic acid esters .

{\ displaystyle {\ ce {R-COOH + R'-OH <=> R-COO-R '+ H2O}}}

When a carboxylic acid is dissolved in an alcohol , the system is initially out of equilibrium. If the equilibrium has been established (e.g. after adding a catalyst or after - very long - waiting), the amount of alcohol has hardly changed due to the large excess; the ester and a corresponding amount of water have formed, and there is a very small amount of carboxylic acid left. If, for example, the ester is removed from the equilibrium by distillation, the ester is reproduced due to the law of mass action . The addition of sulfuric acid as a catalyst can also influence the equilibrium by binding the water produced. The ester yield based on the carboxylic acid can therefore be optimized very well in this way.

On the other hand, adding water to the reaction mixture shifts the equilibrium towards the starting materials.

This means that by adding a component in excess (alcohol or water) you can control which product (ester or acid) predominates in equilibrium.

Further examples:

Combination of changes in temperature, pressure and amount of substance

A very good example of the influence of external conditions on a production process is the Haber-Bosch process mentioned above. As described, the increase in pressure increases the product yield. However, since the reaction requires a high activation energy and the reaction is exothermic, the high temperature shifts the equilibrium to the side of the starting materials, i.e. to the higher volumes. When optimizing processes, an optimal combination of pressure and temperature has to be found. In the Haber-Bosch process, this reaction is therefore carried out at a pressure of approx. 300 bar and a temperature of 550 ° C. In addition, the yield is increased by removing the ammonia, i.e. reducing the amount of substance in the reaction mixture.

Influence of light

Photochromism of the Fulgids

Light-induced isomerization of spiropyran ( 1 ) to merocyanine ( 2 )

A photostationary equilibrium is a state that is created by exposure to light in the visible or ultraviolet range. Examples are photochromic equilibria with certain dyes. The state of equilibrium and thus the color changes depending on the light irradiation. The reverse reaction can take place either by emitting light or heat. If you irradiate z. B. colorless spiropyran with UV light , an intense blue color occurs with the formation of merocyanine , an isomeric compound of spiropyran. In this case one speaks of an isomer equilibrium. The establishment of equilibrium is usually also dependent on the wavelength of the light. The orange triphenyl-fulgide turns bluish when irradiated with short-wave light; when irradiated with red light, the original color is restored.

Web links

More information and examples (Prof. Blum's educational server for chemistry)
Video: Principle of LE CHATELIER BRAUN - How can equilibria be shifted? . Jakob Günter Lauth (SciFox) 2013, made available by the Technical Information Library (TIB), doi : 10.5446 / 15668 .

Individual evidence

^ Le Châtelier, H.-L .: [A general statement of the laws of chemical equilibrium] . In: Comptes rendus hebd. Seanc. Acad. Sci. Paris . tape 99 , 1884, pp. 786-789 .
↑ Braun, F .: Studies on the solubility of solid bodies and the volume and energy changes accompanying the process of dissolution . In: Z. Phys. Chem. 1U, no. 1 , 1887, p. 259-272 , doi : 10.1515 / zpch-1887-0131 .
↑ Andreas Heintz: Thermodynamics of Mixtures: mixed phases, interfaces, reactions, electrochemistry, external force fields . Springer-Verlag, 2017, ISBN 978-3-662-49924-5 ( limited preview in Google book search).
^ Wiberg, Egon., Wiberg, Nils ,: Textbook of Inorganic Chemistry . 102nd, heavily reworked and verb. Ed. De Gruyter, Berlin 2007, ISBN 978-3-11-017770-1 .
↑ Simone Krees: Chemical equilibrium and photostationary equilibrium. Praxis der Naturwissenschaften Chemie, 2/61, Aulis Verlag, March 2012, pp. 18–24.

[1] Le Châtelier, H.-L .: [A general statement of the laws of chemical equilibrium] . In: Comptes rendus hebd. Seanc. Acad. Sci. Paris . tape 99 , 1884, pp. 786-789 .

[2] Braun, F .: Studies on the solubility of solid bodies and the volume and energy changes accompanying the process of dissolution . In: Z. Phys. Chem. 1U, no. 1 , 1887, p. 259-272 , doi : 10.1515 / zpch-1887-0131 .

[3] Andreas Heintz: Thermodynamics of Mixtures: mixed phases, interfaces, reactions, electrochemistry, external force fields . Springer-Verlag, 2017, ISBN 978-3-662-49924-5 ( limited preview in Google book search).

[4] Wiberg, Egon., Wiberg, Nils ,: Textbook of Inorganic Chemistry . 102nd, heavily reworked and verb. Ed. De Gruyter, Berlin 2007, ISBN 978-3-11-017770-1 .

[5] Simone Krees: Chemical equilibrium and photostationary equilibrium. Praxis der Naturwissenschaften Chemie, 2/61, Aulis Verlag, March 2012, pp. 18–24.