Le Chatelier's Principle is an observation about chemical equilibria of reactions. It states that changes in the temperature, pressure (or volume), or concentration of a system will result in predictable and opposing changes in the system in order to achieve a new equilibrium state. Le Chatelier's Principle can be used in practice to understand reaction conditions that will favor increased product formation. This idea was discovered and formulated independently by Henri Louis Le Chatelier (pictured below) and Karl Ferdinand Braun.
Changes in Concentration
According to Le Chatelier's Principle, the addition of product or the removal of reactants to a system will shift the equilibrium towards the reactants, while the addition of more reactants or the removal of product will provide the opposite shift towards product, or the right-hand side of the reaction equation. 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, thereby 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.
Changes in Pressure
A change in pressure (or volume) will result in an attempt to restore equilibrium by creating more or less moles of gas. For example, if the pressure in a system increases (or the volume decreases), the equilibrium will shift to favor the side of the reaction that involves fewer moles of gas. Similarly, if the volume of a system increases (or the pressure decreases), the production of additional moles of gas will be favored.
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 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.
Changes in Temperature
Finally, increases in temperature will favor the reaction direction that consumes heat, while decreases in reaction temperature will favor the direction that produces heat. In other words, the addition of heat to an exothermic reaction will shift the equilibrium towards the reactants, while the same change will shift an endothermic reaction towards product formation. This can be viewed in the endothermic reaction of
Production of NO2 consumes heat. When heat is added and the termperature increases, it will shift the equilibrium to the right as more NO2 is produced. The color of the gas changes, as shown in .