In chemistry, Le Chatelier's principle, also called the Chatelier's principle, can be used to predict the effect of a change in conditions on a chemical equilibrium. The principle is named after Henry Louis Le Chatelier and sometimes Karl Ferdinand Braun who discovered it independently. It can be summarized as:
If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or partial pressure, then the equilibrium shifts to counteract the imposed change and a new equilibrium is established.
This principle has a variety of names, depending upon the discipline using it. See, for example, Lenz's law and homeostasis. It is common to take Le Chatelier's principle to be a more general observation, 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 Chatelier's principle, thereby explaining the diverse phenomena of receptor activation and desensitization.[1] 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 Chatelier principle can occur; these can be resolved by the theory of Response reactions.
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Le Chatelier's principle qualitatively describes systems of non-instantaneous change; the duration of adjustment depends on the strength of the negative feedback to the initial shock. Where a shock initially induces positive feedback (such as thermal runaway), the new equilibrium can be far from the old one, and can take a long time to reach. In some dynamic systems, the end-state cannot be determined from the 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 must have a finite half-life.[2] The principle has analogs throughout Systems theory.
Changing the concentration of an ingredient 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.
Suppose we were to increase the concentration of CO in the system. Using Le Chatelier'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.
The effect of changing the temperature in the equilibrium can be made clear by incorporating heat as either a reactant or a product. When the reaction is exothermic (ΔH is negative, puts energy out), we include heat as a product, and, when the reaction is endothermic (ΔH is positive, takes energy in), we include it as a reactant. Hence, we can determine whether increasing or decreasing the temperature would favour the forward or reverse reaction by applying the same principle as with concentration changes.
For example, the reaction of nitrogen gas with hydrogen gas. This is a reversible reaction, in which the two gases react to form ammonia:
If you put heat as a product:
This is an exothermic reaction (hence the minus sign) when producing ammonia. If we were to lower the temperature, the equilibrium would shift to produce more heat. Since making ammonia is exothermic, this would favour the production of more ammonia. In practice, in the Haber process, the temperature is set at a compromise value, so ammonia is made quickly, even though less would be present at equilibrium.
In exothermic reactions, increase in temperature decreases the equilibrium constant, K, whereas, in endothermic reactions, increase in temperature increases the K value.
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Changes in pressure are attributable to changes in volume. The equilibrium concentrations of the products and reactants do not directly depend on the pressure subjected to the system. However, a change in pressure due to a change in volume of the system will shift the equilibrium.
Considering the reaction of nitrogen gas with hydrogen gas to form ammonia:
Note the number of moles of gas on the left-hand side and the number of moles of gas on the right-hand side. When the volume of the system is changed, the partial pressures of the gases change. If we were to decrease pressure by increasing volume, the equilibrium of the above reaction will shift to the left, because the reactant side has greater number of moles than does the product side. The system tries to counteract the decrease in partial pressure of gas molecules by shifting to the side that exerts greater pressure. Similarly, if we were to increase pressure by decreasing volume, the equilibrium shifts to the right, counteracting the pressure increase by shifting to the side with fewer moles of gas that exert less pressure. If the volume is increased because there are more moles of gas on the reactant side, this change is more significant in the denominator of the equilibrium constant expression, causing a shift in equilibrium.
Thus, an increase in system pressure due to decreasing volume causes the reaction to shift to the side with the fewer moles of gas.[3] A decrease in pressure due to increasing volume causes the reaction to shift to the side with more moles of gas. There is no effect on a reaction where the number of moles of gas is the same on each side of the chemical equation.
An inert gas (or noble gas) such as helium is one that does not react with other elements or compounds. Adding an inert gas into a gas-phase equilibrium at constant volume does not result in a shift.[4] This is because the addition of a non-reactive gas does not change the partial pressures of the other gases in the container. While it is true that the total pressure of the system increases, the total pressure does not have any effect on the equilibrium constant; rather, it is a change in partial pressures that will cause a shift in the equilibrium. 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. There is a short form to remember this: LBMF (little boy married fiona); L stands for less pressure, B - backward reaction, M - more pressure, and F - forward reaction.
A catalyst speeds up the rate of a reaction by providing additional mechanism(s). Adding a catalyst allows for alternative pathways to be made, where the particles can be adsorbed onto the catalyst temporarily before being re-bonded into a new arrangement. The intended effect in adding a catalyst is to lower the activation energy, which frequently increases the rate of reaction. However, the activation energy is lowered by the same amount for the forward and reverse reactions. There is the same increase in reaction rates for both reactions. As a result, a catalyst does not affect the position of the equilibrium. It only effects the time or energy that is required to achieve equilibrium.
In economics, a similar concept also named after Le Chatelier was introduced by US 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. Thus, factor-demand and commodity-supply elasticities are hypothesized to be lower in the short run than in the long run because of the fixed-cost constraint in the short run.[5]