Le Chatelelier's Principle
Le Chatelier's Principle can be stated as follows: If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance.
If a chemical system is at equilibrium and we increase the concentration of a substance (either a reactant or a product), the system reacts to consume some of the substance. Conversely, if we decrease the concentration of a substance, the system reacts to produce some of the substance.
As an example, consider an equilibrium mixture of N2, H2, and NH3:
Adding H2 would cause the system to shift so as to reduce the newly increased concentration of H2. This change can occur only by consuming H2 and simultaneously consuming N2 to form more NH3.
Adding NH3 to the system at equilibrium would cause the concentrations to shift in the direction that reduces the newly increased NH3 concentration; some of the added ammonia would decompose to form N2 and H2. In the Haber reaction, therefore, removing NH3 from an equilibrium mixture of N2, H2, and NH3 causes the reaction to shift from left to right to form more NH3. If the NH3 can be removed continuously, the yield of NH3 can be increased dramatically. In the industrial production of ammonia, the NH3 is continuously removed by selectively liquefying it; the boiling point of NH3 (-33oC) is much higher than that of N2 (-196 oC) and H2 (-253 oC). The liquid NH3 is removed, and the N2 and H2 are recycled to form more NH3. By continuously removing the product, the reaction is driven essentially to completion.
Effects of Volume and Pressure Changes
At constant temperature, reducing the volume of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas. Conversely, increasing the volume causes a shift in the direction that produces more gas molecules.
Consider the equilibrium reaction:
What happens if the total pressure of an equilibrium mixture is increased by decreasing the volume?
According to Le Chatelier's Principle, we expect the equilibrium to shift to the side that reduces the total number of moles of gas, which is the reactant side in this case. (Notice the coefficients in the chemical equation; 1 mol of N2O4 appears on the reactant side and 2 mol NO2 appears on the product side.) We therefore expect the equilibrium to shift to the left, so that NO2 is converted into N2O4 as equilibrium is reestablished.
Effect of Temperature Changes
Changes in concentrations or partial pressures cause shifts in equilibrium without changing the value of the equilibrium constant. In contrast, almost every equilibrium constant changes in value as the temperature changes.
We can deduce the rules for the temperature dependence of the equilibrium constant by applying Le Chatelier's Principle. A simple way to do this is to treat heat as if it were a chemical reagent. In an endothermic (heat-absorbing) reaction we can consider heat as a reactant, whereas in an exothermic (heat-releasing) reaction we can consider heat as a product.
When the temperature of a system at equilibrium is increased, the system reacts as if we added a reactant to an endothermic reaction or a product to an exothermic reaction. The equilibrium shifts in the direction that consumes the excess reactant (or product), namely heat.
Source: Chemistry: The Central Science, by Theodore L. Brown, et. al., (c) 2013 Pearson Higher Education AU.