Le Chatelier’s - Principle and Its Examples for JEE

Equilibrium is a condition at which all opposing forces or processes balance each other. Both physical and chemical systems can reach equilibrium. A phase transformation system acquires physical equilibrium. Chemical equilibria are attained by chemical reactions. Chemical equilibria play a significant role in most biological, environmental, and industrial chemical processes.

Based on the experimental conditions, the given reaction may be fast or slow. A reversible reaction reaches dynamic equilibrium when there is no further change in the concentration of reactants and products and the concentration of reactants and products becomes constant. In a reversible reaction, dynamic equilibrium is reached when the rates of both forward and backward reactions become equal. To understand the implications of chemical equilibrium, the law of equilibrium and Le Chatelier’s Principle are very important.

Le Chatelier’s Principle 

When a system's equilibrium state is disturbed, a net reaction occurs in some direction causing the system to return to its original equilibrium state. Le Chatelier’s principle gives insight into how the system responds when equilibrium is altered. It states that if there is a change in any of the factors that govern the equilibrium conditions of a system, then the system will undergo a change to reduce the effect of the change. Le Chatelier’s law helps to give a qualitative prediction about the effect of a change in equilibrium conditions. It can be applied to both physical and chemical systems. 

Law of Chemical Equilibrium

The Law of chemical equilibrium states that at a particular temperature, the equilibrium constant Kc is equal to the product of concentrations of products raised to their  respective stoichiometric coefficients divided by the product of concentrations of reactants raised to their respective stoichiometric coefficients in the balanced chemical equation.

For a general reaction:

\[aA\text{ + bB }\rightleftharpoons \text{ cC + dD}\]

\[{{\text{K}}_{c}}\text{ = }{{\dfrac{{{\left[ C \right]}^{c}}\left[ D \right]}{{{\left[ A \right]}^{a}}{{\left[ B \right]}^{b}}}}^{d}}\]

Here, $[A]$ and $[B]$ are concentrations of reactants at equilibrium, while $[C]$ and $[D]$ are the equilibrium concentration of products. The equilibrium constant $K_c$ can be used to predict the direction of the reaction. It is only applicable at the equilibrium concentrations. Another constant called the reaction quotient Q_c can also be defined similarly to the equilibrium constant except that Q_c can be obtained anytime in the reaction.

If Q_c > K_c, when decreasing the value of Q_c, the reaction proceeds in the reverse direction.

If Q_c < K_c, when increasing the value of Q_c, the reaction proceeds in forward direction.

If Q_c = K_c, the reaction will be in equilibrium.

Discussion on Le Chatelier’s Principle 

Effect of Concentration Change

If the equilibrium is disturbed by the addition/removal of any reactant/ products,  according to Le Chatelier’s principle, the reaction will proceed to minimise the effect of concentration changes by the following methods:

• The concentration stress of an added reactant/product is nullified by changing the course of net reaction in the direction that consumes the added substance.

•  The concentration stress of a reactant/product which is removed is nullified by changing the direction of net reaction in such a way that it replenishes the removed substance. 

Consider the reaction: 

\[{{N}_{2}}\left( g \right)+3\text{ }{{H}_{2}}\left( g \right)\rightleftharpoons 2\text{ }N{{H}_{3}}\left( g \right)\]

• Addition of a Reactant: When we add more nitrogen to the system, the concentration stress is relieved by shifting the reaction towards more nitrogen being consumed, i.e., the reaction proceeds in a forward direction. Equilibrium shifts toward a forward direction. The addition of a reactant at equilibrium results in a lower Q_c compared to K_c. This favours a forward reaction.

• Removal of Product: To reduce the effect, the removal of product equilibrium shifts to the forward direction.

• Removal of Reactant/ Addition of Product: When a reactant is removed or the product formed is added in the reaction mixture, the reaction shifts to a backward direction, which favours the replenishment of reactant or consumption of the excess product. 

For heterogeneous reactions like dissociation of calcium carbonate:

$\mathrm{CaCO}_{3}(\mathrm{~s}) \rightleftharpoons \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{~g})$

The equilibrium doesn’t depend on the amount of solids as the concentration of pure solids doesn't change.

Effect of Pressure Change

A pressure change can affect the yield of reactants and products in the gaseous state when different moles of reactants and products are present in the gaseous state.

Consider the reaction:

\[{{N}_{2}}\left( g \right)+3\text{ }{{H}_{2}}\left( g \right)\rightleftharpoons 2\text{ }N{{H}_{3}}\left( g \right)\]

Here, $4$ moles reactants in the gaseous state yield $2$ moles of product in the gaseous state.

• Increase in Pressure: At constant temperature and volume, the amount of a gas is directly proportional to the pressure of the gas. If we increase the system pressure in the above written gas phase reaction, the equilibrium will shift to where there is less pressure. Here, the reaction shifts towards a forward direction as it produces less moles of gas molecules and hence decreases the pressure.  

• Decrease in Pressure: It will have a reverse effect and the system will shift towards increased pressure /a greater number of moles of gaseous reactant produced, which means backward reaction will increase.

Effect of Inert Gas

• If inert gas is added at a constant volume, then there will be no effect on the equilibrium state. This is because, at constant volume, if inert gas is added, then total pressure increases but the number of moles of reactants and products per unit volume won’t change.

• If inert gas is added at constant pressure, then the number of moles of reactants and products per unit volume will decrease, which decreases the individual partial pressure of each gas, and reaction will shift in direction where the number of moles increases.

For example: in reaction $\mathrm{N}_{2}+3 \mathrm{H}_{2} \rightleftharpoons 2 \mathrm{NH}_{3}$, the addition of inert gas at constant pressure will shift the reaction to the left (More number of moles) and will result in a decrease in product formation.

Effect of Temperature Change 

When a temperature change occurs during a reaction, the value of the equilibrium constant K_c is changed. The change depends on the sign of enthalpy change $\Delta H$ for the reaction.

• Exothermic Reaction (negative $\Delta H$): For this reaction, the equilibrium constant decreases as the temperature increases. 

• Endothermic Reaction (positive $\Delta H$): Here, the equilibrium constant increases as the temperature increases. 

Ammonia synthesis is an exothermic reaction. According to Le Chatelier’s principle, an increase in temperature shifts the equilibrium to a backward reaction and decreases the equilibrium concentration of ammonia in ammonia synthesis. So, we can say that low temperature gives a high yield of ammonia. Practically, very low temperatures slow down the reaction. Thus, a catalyst is used to get a high yield.

Effect of a Catalyst 

A catalyst increases the rate of the chemical reaction without taking part in the reaction. It is possible because the catalyst shifts the reaction to a low energy pathway. It does not affect equilibrium. The catalyst decreases the activation energy for the forward and reverse reactions by exactly the same amount. 

Effect of Catalyst

Le Chatelier’s Principle Examples 

These are some examples of the application of Le Chatelier’s principle:

• When we put clothes for drying, an equilibrium is established between the gaseous molecules of water evaporated and water in the liquid form present in clothes. When air breezes, this equilibrium is disturbed as gaseous molecules decrease. Hence, in order to increase gas molecules, the reaction shifts in the direction where gaseous molecules increase. Thus, clothes dry quickly.

• When blood carrying oxyhemoglobin reaches a tissue where oxygen partial pressure is less, oxyhemoglobin breaks down and gives oxygen to the tissue. 

• In Haber's process for the manufacture of ammonia in industries, Le Chatelier's principle is often used in order to increase the yield of ammonia by removing it. This leads to an increase in the yield of ammonia as the reaction moves forward.

$\mathrm{N}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NH}_{3}(\mathrm{~g})$

Conclusion

The equilibrium constant has a constant value at a fixed temperature and helps to predict the direction of a reaction. Le Chatelier’s principle states that the change in any condition of a reaction system such as temperature, pressure, concentration, etc. will cause the equilibrium to change the direction of reaction so that it will reduce the effect of the change. It is very helpful to determine the direction of equilibrium and to control the yield of products by controlling these factors.