State and explain Le-Chatelier’s principle.

Le-Chatelier’s principle. This principle may be stated as if a stress (such as a change in concentration, temperature or pressure) is applied to a system in equilibrium, the equilibrium shifts in a way to undo or nullify the effect of the imposed stress.

(i) Effect of change of concentration on equilibrium. If the concentration of any one or all the reactants is increased, the equilibrium shifts towards right hand side to form more products whereas increase in the concentration of any one or all the products shifts the equilibrium towards left hand side to form more reactants in order to nullify the effect of increase in the concentration of reactants or products respectively. For example, consider the reaction,


Increase in concentration of reactants (N₂, H₂) will shift the equilibrium in the forward direction in order to decrease their concentration. The addition of extra NH₃ from outside the equilibrium mixture will shift the equilibrium in the backward direction.

(ii) Effect of temperature on equilibrium. According to Le-Chatelier’s principle of increasing the temperature, the equilibrium shifts towards that direction where absorption of heat (endothermic change) takes place in order to nullify the effect of the rise in temperature. On the other hand, on decreasing the temperature, the equilibrium shifts towards that direction where the evolution of heat (exothermic change) takes place in order to nullify the effect of a decrease in temperature. For example.
   

On decreasing the temperature, the equilibrium shifts in the forward direction i.e. towards the exothermic reaction (evolution of heat). Thus, a decrease in temperature favours the formation of sulphur trioxide.


On increasing the temperature, the equilibrium shifts in the forward direction i.e. towards the endothermic reaction (absorption of heat). Thus, an increase in the temperature favours the formation of nitric oxide.

(iii) Effect of change in pressure on equilibrium. On increasing the pressure, the number of moles per unit volume increases and thus according to Le-Chatelier’s principle the equilibrium shifts towards the side where the number of moles per unit volume decreases in order to nullify the effect of an increase in pressure. For example.

On increasing the pressure, the number of moles per unit volume increases and thus according to Le-Chaterlier's principle the equilibrium shifts towards the right-hand side (i.e. towards the formation of ammonia) where the number of moles per unit volume decreases.
(iv) Effect of the catalyst. A catalyst has no effect on equilibrium point. This is because it increases the rate of the forward as well as backward reaction to the same extent. Thus, a catalyst does not affect the position of equilibrium, but simply helps to achieve the equilibrium in a shorter time i.e. quickly.

The given dissociation equilibrium is

   

           

(i) Adding inert gas at constant volume. If an inert gas like helium or nitrogen etc. is added to a system at equilibrium at constant volume, then the total pressure increases. However, there will be no change in the position of equilibrium, even though the pressure has changed. This is because the concentrations of reactants and products (number of moles/volume) will not change. Hence the values of concentrations will continue to satisfy the equilibrium law. Hence the state of equilibrium will remain unaffected.

(ii) Adding inert gas at constant pressure. When an inert gas is added to a system at equilibrium keeping the pressure constant, the volume of the system increases. This results in a decrease in a number of moles of reactants per unit volume. According to Le-Chatelier’s principle, the equilibrium shifts in a direction in which there is an increase in the number of moles of gases. Hence the equilibrium shifts towards the forward direction where the number of moles of gases increases. In other words, more PCl₅ dissociates to give PCl₃ and Cl₂. Hence dissociation of PCl₅ increases with the addition of an inert gas.

(i) Effect of concentration. Increase in concentration of reactants (N₂, H₂) will shift the equilibrium in the forward direction to form more ammonia in order to decrease their concentrations. The addition of extra NH₃ from outside to the equilibrium mixture will shift the equilibrium in the backward direction. Thus, the addition of N₂ and H₂ favours the formation of ammonia.

(ii) Effect of temperature. The forward reaction is exothermic in nature while the backward reaction is endothermic in nature. According to Le-Chatelier’s principle, on decreasing the temperature, the equilibrium shifts towards that direction where the evolution of heat takes place in order to nullify the effect of decreasing temperature. Thus, a decrease in temperature favours the formation of ammonia.
On the other hand, on increasing the temperature the equilibrium shifts towards the backward direction where absorption of heat takes place in order to nullify the effect of the rise in temperature. Thus, lower temperature favours the formation of ammonia.

(iii) Effect of pressure. On increasing the pressure, the number of moles per unit volume increases and thus according to Le-Chatelier’s principle, the equilibrium shifts towards that side where the number of moles per unit volume decreases in order to nullify the effect of an increase in pressure. On the other hand, on decreasing the pressure, the equilibrium shifts towards the backwards direction i.e. ammonia decomposes to give N₂ and H₂. Thus, an increase in pressure favours the formation of ammonia.
Hence favourable conditions in the formation of ammonia are:

(i) The addition of N₂ and H₂ 
(ii) Lower temperature
(iii) Higher pressure.