Le Châtelier's principle is related ultimately to the rates of the forward and reverse steps in a reaction. Explain (a) why an increase in reactant concentration shifts the equilibrium position to the right but does not change K; (b) why a decrease in V shifts the equilibrium position toward fewer moles of gas but does not change K; (c) why a rise in T shifts the equilibrium position of an exothermic reaction toward reactants and also changes K; and (d) why a rise in temperature of an endothermic reaction from ${}^{{\mathbf{T}}_{\mathbf{1}}}$to ${}^{{\mathbf{T}}_{\mathbf{2}}}$results in ${}^{{\mathbf{K}}_{\mathbf{2}}}$ being larger than${}^{{\mathbf{K}}_{\mathbf{1}}}$ .

Le-Chatelier's principle states that when the equilibrium is disturbed by changing the determining factors, the system will adjust itself to counteract the effect of the change.

This principle is used to carry out reactions with maximum yield in industries.

The favored direction of a reversible reaction depends on the reaction condition, and using Le-Chatelier's principle; the reaction conditions can be set such that the forward reaction is favored.

The equilibrium of a reaction is determined by factors like temperature, pressure, concentration, volume, etc. Any change in these variables causes a change in equilibrium and the value of the reaction quotient changes.

Then, the system adjusts itself such that the value of the reaction quotient becomes equal to the equilibrium constant again. This is done by favoring the forward or backward reaction (increasing its rate). Therefore, the equilibrium shifts to either the reactant or the product side.

(a)

When the reactant concentration of a reaction in equilibrium is increased, the value of the reaction quotient decreases, which is given by:

$\mathrm{Q}=\frac{\left[\mathrm{Products}\right]}{\left[\mathrm{Reactans}\right]}$

The system then adjusts itself such that the value of the reaction quotient increases to the value of the equilibrium constant. This is done by favoring the forward reaction and thus increasing the product concentration.

The equilibrium constant of a reaction is a constant at constant temperature and does not change with initial reactant concentration.

(b)

The pressure of a system is inversely proportional to the volume of a gas. Therefore, a decrease in volume causes an increase in pressure on the system. The pressure of a system increases with an increase in the number of moles of gas.

Therefore, an increase in pressure of a system is counteracted by decreasing the number of moles. The reaction shifts towards the side of the reaction with a fewer number of moles.

The equilibrium constant of the reaction is only dependent on the temperature and, therefore, does not change with the change in volume of the gas.

(c)

A reaction in which heat is liberated from the system is an exothermic reaction. Temperature is a measure of heat, and an increase in temperature involves an increase in heat content. Therefore, when the temperature of a system is increased, the reaction will favor the endothermic reaction, which will be the backward reaction.

This causes an increase in the reactant concentration and a decrease in product concentration. Therefore, the equilibrium constant at the new equilibrium position will have a different value.

When the equilibrium is increased, the system cannot re-establish the previous equilibrium; therefore, the K value changes.

(d)

Reactions in which the system absorbs heat are called endothermic reactions. When the temperature of a system increases, the heat content in the system increases. Therefore, more heat will be available for the reactants to absorb, and the rate of the forward endothermic reaction will increase.

This causes an increase in the concentration of the product and a decrease in the concentration of reactants. The equilibrium constant, which gives the ratio of product concentration to reactant concentration, increases.

Therefore, at a higher temperature, the equilibrium constant of an endothermic reaction will have a greater value.