Le Chatelier's principle
The principle is named after French chemist Henry Louis Le Chatelier who enunciated the principle in 1884 by extending the reasoning from the Van 't Hoff relation of how temperature variations changes the equilibrium to the variations of pressure and what's now called chemical potential,[3][4] and sometimes also credited to Karl Ferdinand Braun, who discovered it independently in 1887.
Le Chatelier–Braun principle analyzes the qualitative behaviour of a thermodynamic system when a particular one of its externally controlled state variables, say
The principle can be stated in two ways, formally different, but substantially equivalent, and, in a sense, mutually 'reciprocal'.
[8] The duration of adjustment depends on the strength of the negative feedback to the initial shock.
The principle is typically used to describe closed negative-feedback systems, but applies, in general, to thermodynamically closed and isolated systems in nature, since the second law of thermodynamics ensures that the disequilibrium caused by an instantaneous shock is eventually followed by a new equilibrium.
Shear pins and other such sacrificial devices are design elements that protect systems against stress applied in undesired manners to relieve it so as to prevent more extensive damage to the entire system, a practical engineering application of Le Chatelier's principle.
The chemical system will attempt to partly oppose the change affected to the original state of equilibrium.
In turn, the rate of reaction, extent, and yield of products will be altered corresponding to the impact on the system.
This can be illustrated by the equilibrium of carbon monoxide and hydrogen gas, reacting to form methanol.
Using Le Chatelier's principle, we can predict that the concentration of methanol will increase, decreasing the total change in CO.
Even if the desired product is not thermodynamically favored, the end-product can be obtained if it is continuously removed from the solution.
The effect of a change in concentration is often exploited synthetically for condensation reactions (i.e., reactions that extrude water) that are equilibrium processes (e.g., formation of an ester from carboxylic acid and alcohol or an imine from an amine and aldehyde).
When the reaction is exothermic (ΔH is negative and energy is released), heat is included as a product, and when the reaction is endothermic (ΔH is positive and energy is consumed), heat is included as a reactant.
Hence, whether increasing or decreasing the temperature would favor the forward or the reverse reaction can be determined by applying the same principle as with concentration changes.
The equilibrium concentrations of the products and reactants do not directly depend on the total pressure of the system.
With a pressure increase due to a decrease in volume, the side of the equilibrium with fewer moles is more favorable[10] and with a pressure decrease due to an increase in volume, the side with more moles is more favorable.
There is no effect on a reaction where the number of moles of gas is the same on each side of the chemical equation.
Adding an inert gas into a gas-phase equilibrium at constant volume does not result in a shift.
If, however, the volume is allowed to increase in the process, the partial pressures of all gases would be decreased resulting in a shift towards the side with the greater number of moles of gas.
For example, consider the Haber process for the synthesis of ammonia (NH3): In the above reaction, iron (Fe) and molybdenum (Mo) will function as catalysts if present.
The latter are stable against perturbations that satisfy certain criteria; this is essential to the definition of thermodynamic equilibrium.
[11][12][13] In theory and, nearly, in some practical scenarios, a body can be in a stationary state with zero macroscopic flows and rates of chemical reaction (for example, when no suitable catalyst is present), yet not in thermodynamic equilibrium, because it is metastable or unstable; then Le Chatelier's principle does not necessarily apply.
The Gibbs approach is reliable within its proper scope, thermodynamic equilibrium, though of course it does not cover non-equilibrium scenarios.
Thermodynamic non-equilibrium scenarios can contradict an over-general statement of Le Chatelier's Principle.
"or, "roughly stated":[16] Any change in status quo prompts an opposing reaction in the responding system.The concept of systemic maintenance of a stable steady state despite perturbations has a variety of names, and has been studied in a variety of contexts, chiefly in the natural sciences.
In chemistry, the principle is used to manipulate the outcomes of reversible reactions, often to increase their yield.
[17] In biology, the concept of homeostasis is different from Le Chatelier's principle, in that homoeostasis is generally maintained by processes of active character, as distinct from the passive or dissipative character of the processes described by Le Chatelier's principle in thermodynamics.
In economics, a similar concept also named after Le Chatelier was introduced by American economist Paul Samuelson in 1947.
There the generalized Le Chatelier principle is for a maximum condition of economic equilibrium: Where all unknowns of a function are independently variable, auxiliary constraints—"just-binding" in leaving initial equilibrium unchanged—reduce the response to a parameter change.
[18] Since the change of the value of an objective function in a neighbourhood of the maximum position is described by the envelope theorem, Le Chatelier's principle can be shown to be a corollary thereof.
When heat is removed and the temperature decreases, the reaction shifts to the right and the flask turns colorless due to an increase in N 2 O 4 . This demonstrates Le Chatelier's principle: the equilibrium shifts in the direction that releases energy.
When heat is added and the temperature increases, the reaction shifts to the left and the flask turns reddish brown due to an increase in NO 2 : again, according to Le Chatelier's principle.