Le Chatelier’s principle states that if a system at equilibrium is subjected to a change in temperature, pressure or the concentration of a reactant or product, then the system readjusts itself to counteract the effect of the change and a new equilibrium is established.
If the concentration of a species is increased, then the equilibrium moves in the direction that reduces the concentration of that species. If the concentration of a species is decreased, then the equilibrium moves in the direction that increases the concentration of that species.
Example:
For the equilibrium:
\[\ce{H2(g) + I2(g) <=> 2HI(g)}\]
Changing pressure affects equilibria involving gases. Increasing the pressure shifts the equilibrium toward the side with fewer moles of gas. Decreasing pressure shifts it toward the side with more moles of gas.
Example 1:
\[\ce{H2(g) + I2(g) <=> 2HI(g)}\]
Changing the pressure will have no effect on the position of this equilibrium, as there is the same number of moles on both sides
Note: If the number of gas moles on both sides is equal, pressure change has no effect on equilibrium.
Example 2:
\[\ce{PCl5(g) <=> PCl3(g) + Cl2(g)}\]
If the pressure is increased, the system will move to the left, where there are fewer moles. If the pressure is decreased, the system will move to the right, where there are more moles.
For endothermic reactions (\( \Delta H > 0 \)), increasing temperature shifts equilibrium to the right (favoring products). For exothermic reactions (\( \Delta H < 0 \)), increasing temperature shifts it to the left (favoring reactants).
Example:
\[\ce{2NO2(g) <=> N2O4(g)} \quad \Delta H = -57.2\, \text{kJ/mol}\]
A catalyst speeds up the rate of both the forward and reverse reactions equally. It does not affect the position of equilibrium, only the rate at which equilibrium is achieved.
Written by Fillios Memtsoudis