Le Châtelier's principle
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In chemistry, Le Châtelier's principle can be used to predict the effect of a change in conditions on a chemical equilibrium. The principle is named after Henri Louis Le Châtelier.
Le Châtelier's principle can be summarized thus:
If a chemical system at equilibrium experiences a change in concentration, temperature, or total pressure; the equilibrium will shift in order to minimize that change.
The principle is used by chemists in order 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. [1]
In economics, a similar concept also named after Le Châtelier was introduced by American economist Paul Samuelson in 1947. It deals with constraints on maximizing behavior, explaining that short-run demands have lower elasticity than those in the long run since a longer time frame allows new factors and prices to change.
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[edit] Examples
[edit] Concentration
Changing the concentration of an ingredient will shift the equilibrium to the side that would reduce that change in concentration.
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 collisions of that reactant would increase also, allowing for an increase in forward reaction, and generation of the product.
[edit] Temperature
Let us take for example the reaction of nitrogen gas with hydrogen gas. This is a reversible reaction, in which the two gases react to form ammonia:
- N2 + 3 H2 → 2 NH3 ΔH = -92kJ
This is an exothermic reaction when producing ammonia. If we were to lower the temperature, the equilibrium would shift in such a way as to produce heat. Since this reaction is exothermic to the right, it would favour the production of more ammonia. This reaction is used in the Haber process, which is a good example of the way chemists utilize Le Châtelier's principle.
[edit] Total Pressure
Once again, let us refer to the reaction of nitrogen gas with hydrogen gas to form ammonia:
- N2 + 3 H2 → 2 NH3 ΔH = -92kJ
Note the number of moles of gas on the left hand side, and the number of moles of gas on the right hand side. We know that gases at the same temperature and pressure will occupy the same volume. We can use this fact to predict the change in equilibrium that will occur if we were to change the total pressure.
Suppose we increase total pressure on the system, now by Le Châtelier's principle the equilibrium would move to decrease the pressure. Noting that 4 moles of gas occupy more volume than 2 moles of gas, we can deduce that the reaction will move towards the products if we were to increase the pressure.
[edit] Inert Gases
An inert gas (or noble gas) such as helium is one which does not react with other elements or compounds. To add an inert gas into a closed system at equilibrium may or may not result in a shift. For example, consider adding helium to the same reaction:
- N2 + He → N2 + He
The main effect of adding an inert gas to a closed reaction is that it will increase the pressure of the reaction. It could also possibly act as a catalyst.
[edit] Volume Held Constant
If volume is held constant, the concentrations of the above gases do not change. The partial pressures of the gases in the reaction do not change, even though we have increased the total pressure by adding helium. Therefore, there is no change in equilibrium.
[edit] Volume Allowed to Increase
If the volume is allowed to increase, the concentrations, as well as the partial pressures, all decrease. Because there are more stoichiometric moles on the lefthand side of the equation, the decrease in concentration affects the lefthand side more than the righthand side. Therefore, the reaction would shift to the left until the system is at equilibrium again.
