Le chateliers principle
In chemistry, Le Châtelier's principle (pronounced /lə ˈʃɑːtlieɪ/), also called Chatelier's principle or "The Equilibrium Law", can be used to predict the effect of a change in conditions on a chemical equilibrium. The principle is named after Henry Louis Le Châtelier and sometimes Karl Ferdinand Braun who discovered it independently. It can be stated as:
When any system at equilibrium is subjected to change in concentration, temperature, volume, or pressure, then the system readjusts itself to (partially) counteract the effect of the applied change and a new equilibrium is established.
or whenever a system in equilibrium is disturbed the system will adjust itself in such a way that the effect of the change will be nullified. (in short)
This principle has a variety of names, depending upon the discipline using it (see homeostasis, a term commonly used in biology). It is common to take Le Châtelier's principle to be a more general observation,[1] roughly stated:
Any change in status quo prompts an opposing reaction in the responding system.
In chemistry, the principle is used to manipulate the outcomes of reversible reactions, often to increase the yield of reactions. In pharmacology, the binding of ligands to the receptor may shift the equilibrium according to Le Châtelier's principle, thereby explaining the diverse phenomena of receptor activation and desensitization.[2] In economics, the principle has been generalized to help explain the price equilibrium of efficient economic systems. In simultaneous equilibrium systems, phenomena that are in apparent contradiction to Le Châtelier's principle can occur; these can be resolved by the theory of response reactions.
Effect of change in concentration[edit]
Changing the concentration of a chemical will shift the equilibrium to the side that would reduce that change in concentration. The chemical system will attempt to partially 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.
CO + 2 H2 ⇌ CH3OH
Suppose we were to increase the concentration of CO in the system. Using Le Châtelier's principle, we can predict that the amount of methanol will increase, decreasing the total change in CO. If we are to add a species to the overall reaction, the reaction will favor the side opposing the addition of the species. Likewise, the subtraction of a species would cause the reaction to "fill the gap" and favor the side where the species was reduced. This observation is supported by the collision theory. As the concentration of CO is increased, the frequency of successful collisions of that reactant would increase also, allowing for an increase in forward reaction, and generation of the product. Even if a desired product is not thermodynamically favored, the end-product can be obtained if it is continuously removed from the solution.
Effect of change in temperature[edit]
The effect of changing the temperature in the equilibrium can be made clear by a) incorporating heat as either a reactant or a product, and b) assuming that an increase in temperature increases the heat content of a system. When the reaction is exothermic (ΔH is negative, puts energy out), heat is included as a product, and, when the reaction is endothermic (ΔH is positive, takes energy in), 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.
Take, for example, the reversible reaction of nitrogen gas with hydrogen gas to form ammonia:
N2(g) + 3 H2(g) ⇌ 2 NH3(g) ΔH = -92 kJ mol−1
Because this reaction is exothermic, it produces heat:
N2(g) + 3 H2(g) ⇌ 2 NH3(g) + heat
If the temperature was increased, the heat content of the system would increase, so the system would consume some of that heat by shifting the equilibrium to the left, thereby producing less ammonia. More ammonia would be produced if the reaction was run at a lower temperature, but a lower temperature also lowers the rate of the process, so, in practice (the Haber process) the temperature is set at a compromise value that allows ammonia to be made at a reasonable rate with an equilibrium concentration that is not too unfavorable.
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