Le Chatelier’s Principle: Effect of Temperature, Pressure, Concentration

Le Chatelier’s Principle: Let’s consider a reversible reaction of melting of the ice. Ice, when kept open, changes into the water due to the absorption of heat energy. It can be reversed by decreasing temperature. In this article, you will explore the effect of concentration, temperature, and pressure on equilibrium by Le Chatelier’s principle and its applications.

Equilibrium

Equilibrium represents the state of a process in which the properties like temperature, pressure, and the concentration of the system do not show any change with time. Attaining an equilibrium involves two opposing processes. Processes involving only physical change are called physical equilibrium, and the processes involving a chemical change are called chemical equilibrium.

What is Le Chatelier’s Principle?

The state of a system can be changed by changing the conditions like temperature, pressure, and concentration of different species present in the system was studied by Le Chatelier and Braun in \(1884.\)

According to Le Chatelier’s principle process, if a system under equilibrium is subjected to a change in temperature, pressure, or concentration, then the equilibrium shifts itself in such a way to undo or neutralise the effect of change.

Effect of Change of Concentration

The increase in the concentration of any of the reactants shifts the equilibrium towards the forward direction. The increase in the concentration of any of the products shifts the equilibrium towards the backward direction.

Consider the general reaction, \({\text{A + B}} \rightleftharpoons {\text{C + D}}\)

If more reactant \({\text{A}}\) or \({\text{B}}\) is added to the reaction in equilibrium, then according to Le Chatelier’s principle, the effect will decrease the concentration of \({\text{A}}\) or \({\text{B}}\) Therefore, the equilibrium shifts in the forward direction. Similarly, the impact of adding more of products \({\text{C}}\) or \({\text{D}}\) to the reaction in equilibrium will shift the equilibrium in the backward direction; hence the equilibrium is re-established after some time.

Example: The reaction between yellow-coloured ferric nitrate and colourless potassium thiocyanate forms a red-coloured iron thiocyanate complex.

\({\text{F}}{{\text{e}}^{{\text{3 + }}}}\left( {{\text{aq}}} \right){\text{ + SC}}{{\text{N}}^{\text{-}}}\left( {{\text{aq}}} \right) \rightleftharpoons {\left[ {{\text{Fe}}\left( {{\text{SCN}}} \right)} \right]^{{\text{2 + }}}}\left( {{\text{aq}}} \right)\)

By adding a small amount of ferric chloride to the equilibrium mixture, the intensity of the red colour will increase. This indicates that ferric ion has combined with thiocyanate ions to give iron thiocyanate complex, and equilibrium has shifted to the right. On adding iron thiocyanate, the intensity of the colour decreases indicating the equilibrium is shifted to the left.

Effect of Change of Temperature

The equilibrium of exothermic and endothermic reactions can be altered by changing the temperature. Exothermic reactions are favoured by low temperature, whereas endothermic reactions are favoured by high temperature. On the altering temperature, if the forward reaction is exothermic, the backward reaction will be endothermic and vice versa.

In other words, the increase in temperature shifts the equilibrium in the direction of the endothermic reaction and the decrease in temperature shifts the equilibrium in the direction of an exothermic reaction.

Example 1: The formation of ammonia is an exothermic process, while the reverse is an endothermic process.
\({{\text{N}}_{\text{2}}}\left( {\text{g}} \right){\text{ + 3}}{{\text{H}}_{\text{2}}}\left( {\text{g}} \right) \rightleftharpoons {\text{2N}}{{\text{H}}_{\text{3}}}\left( {\text{g}} \right);\Delta {\text{H=-92}}{\text{.5}}\,{\text{kJmo}}{{\text{l}}^{{\text{-1}}}}\)
Suppose the temperature of the reaction is increased, the equilibrium shifts to the left to neutralise or undo the effect of an increase in temperature. The backward reaction is accomplished by decreasing temperature or cooling. Similarly, if the temperature is increased, then the equilibrium will shift to the right because heat is evolved in the forward reaction.

Example 2: The formation of nitric oxide by the reaction of nitrogen and oxygen is an endothermic reaction.
\({{\text{N}}_{\text{2}}}\left( {\text{g}} \right){\text{ + }}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right) \rightleftharpoons {\text{2NO}}\left( {\text{g}} \right);\Delta {\text{H= + 180}}{\text{.0}}\,{\text{kJmo}}{{\text{l}}^{{\text{-1}}}}\)
The increase in temperature will shift the equilibrium to the right. In contrast, the decrease in temperature will shift the equilibrium to the left to neutralise or undo the effect of change in the temperature.

Effect of Change of Pressure

Pressure has hardly any effect on the reactions carried in solid and liquid. However, it does influence the equilibrium of the reactions that are carried in the gases. The effect of pressure depends upon the number of moles of the reactants and the product involved in a particular reaction.

Increasing pressure shifts the equilibrium in the direction of a decrease in the gaseous moles. In contrast, a decrease in pressure shifts the equilibrium in increasing gaseous moles. However, pressure does not affect an equilibrium reaction that proceeds with no change in the total number of moles.

Example 1: Dissociation of \({{\text{N}}_{\text{2}}}{{\text{O}}_{\text{4}}}\) into \({\text{N}} {{\text{O}}_2}\)
\({{\text{N}}_{\text{2}}}{{\text{O}}_{\text{4}}} \rightleftharpoons {\text{2N}}{{\text{O}}_{\text{2}}}\)

At constant temperature, the volume will decrease on increasing the pressure, i.e., the number of moles per unit volume will increase. According to Le Chatelier’s principle, the equilibrium point will shift in that direction, accompanied by a decrease in the number of moles per unit volume.

Since backward reaction takes place with a decrease in the number of moles, so an increase in pressure favours the combination of \({\text{N}} {{\text{O}}_2}\) molecules to produce \({{\text{N}}_{\text{2}}}{{\text{O}}_{\text{4}}}\) i.e., suppresses the dissociation of \({{\text{N}}_{\text{2}}}{{\text{O}}_{\text{4}}}\) into \({\text{N}} {{\text{O}}_2}\) .

Example 2:\( {\text{PC}}{{\text{l}}_ {\text{3}}}\left({\text{g}} \right) {\text{ + C}} { {\text{l}}_ {\text{2}}}\left({\text{g}} \right) \rightleftharpoons {\text{PC}} {{\text{l}}_{\text{5}}}\left({\text{g}} \right)\)
In the above reaction, the number of moles of the product is less than the reactants. Therefore, the increase in pressure will shift the equilibrium to the right and decrease pressure when shifting the equilibrium to the left.

Applications of Le Chatelier’s Principle

Le Chatelier’s principle has great practical significance in the equilibrium state in chemical and physical systems.

Application of Le Chatelier’s principle in chemical equilibria

Le Chatelier’s principle is instrumental in predicting the condition of temperature, pressure, and concentration to get high yields in certain industrial reactions. A few examples are,

1. Formation of ammonia in Haber’s process: The chemical equilibrium taking place is
   \({{\text{N}}_ {\text{2}}}\left({\text{g}} \right){\text{ + 3}} { {\text{H}}_ {\text{2}}}\left({\text{g}} \right) \rightleftharpoons {\text{2N}} {{\text{H}}_{\text{3}}}\left({\text{g}} \right){\text{;}}\Delta {\text{H=-92}}{\text{.5}}\, {\text{kJmo}} {{\text{l}}^{{\text{-1}}}}\)

This reaction is favoured at optimum temperature \(750 {\text{K}}\) and pressure \(200\) atmospheres. By increasing the concentration of reactants and continuous removal of products formation of ammonia is increased. Finely divided iron is used as a catalyst to achieve equilibrium rapidly, and promoter molybdenum is used to increase the catalyst’s efficiency.

2. Formation of nitric oxide: The formation of nitric oxide by a direct combination of nitrogen and oxygen is favoured at \(2773 {\text{K,}}\) and a high concentration of \({{\text{N}}_{\text{2}}}\) and \({{\text{O}}_{\text{2}}}\). Pressure does not affect this reaction since the number of moles of reactants and products are the same.

\({{\text{N}}_{\text{2}}}\left({\text{g}} \right){\text{ + }}{{\text{O}}_{\text{2}}}\left({\text{g}} \right) \rightleftharpoons {\text{2NO}}\left({\text{g}} \right);\Delta {\text{H= + 180}} {\text{.7}}\, {\text{kJmo}} { {\text{l}}^{{\text{-1}}}}\)

This process is used in the Binkland-Eyde process for the manufacture of nitric acid.

3. Formation of sulphur trioxide in contact process of the manufacture of sulphuric acid:  The equilibrium reaction involving the oxidation of sulphur dioxide to sulphur trioxide is represented as:

\({\text{2S}}{{\text{O}}_{\text{2}}}\left({\text{g}} \right){\text{ + }}{{\text{O}}_{\text{2}}}\left({\text{g}} \right) \rightleftharpoons {\text{2S}} { {\text{O}}_ {\text{3}}}\left({\text{g}} \right);\Delta {\text{H= + 193}} {\text{.2}}\, {\text{kJmo}} {{\text{l}}^ {{\text{-1}}}}\)

The reaction is favoured at low temperature. Usually, an optimum temperature of \(673 – 723{\text{K}}\) and pressure of \(1.5-1.7\) atm. The finely divided platinum or vanadium pentoxide is used as a catalyst to attain the equilibrium state rapidly. The concentration of \({\text{S}}{{\text{O}}_{\text{2}}}\) and \({{\text{O}}_{\text{2}}}\) increases the rate of reaction.

4. Bosch process for the manufacture of hydrogen: Hydrogen gas is produced from the water gas and steam according to the following equilibrium reaction.

\({{\text{H}}_{\text{2}}}\left({\text{g}} \right){\text{ + CO}}\left({\text{g}} \right){\text{ + }}{{\text{H}}_{\text{2}}}{\text{O}}\left({\text{g}} \right) \rightleftharpoons {\text{2}}{{\text{H}}_{\text{2}}}\left( {\text{g}} \right){\text{ + C}}{{\text{O}}_{\text{2}}}\left({\text{g}} \right),\Delta {\text{H= + 42}}{\text{.0}}\,{\text{kJmo}}{{\text{l}}^{{\text{-1}}}}\)

The forward reaction being endothermic is favoured by high temperature, and usually, a temperature of \(673 – 723{\text{K}}\) is used. Pressure does not affect this reaction since the number of reactants and products are the same. A high concentration of reactants increases the forward reaction.

Application of Le Chatelier’s Principle in Physical Equilibria

1. Melting of ice: At \(0^\circ {\text{C}}\), both ice and water are present and are in a state of equilibrium.

\({\rm{Ice}}\left( {\rm{s}} \right){\rm{water}}\left( {\rm{l}} \right);\Delta {\rm{H =   +  6}}{\rm{.0}}{\mkern 1mu} \,{\rm{kJ}}\,{\rm{mo}}{{\rm{l}}^{{\rm{ – 1}}}}\)

The change of ice into water is an endothermic process. The forward reaction increases the temperature and pressure.

2. Evaporation of water: The evaporation of water is an endothermic process.

\({{\text{H}}_{\text{2}}}{\text{O}}\left({\text{s}} \right) \rightleftharpoons {{\text{H}}_{\text{2}}}{\text{O}}\left({\text{g}} \right);\Delta {\text{H= + 40}}{\text{.6}}\,{\text{kJmo}}{{\text{l}}^{{\text{-1}}}}\)

On increasing the temperature, the equilibrium shifts to the right, i.e., forward reaction will increase due to the endothermic nature of the reaction. Therefore, according to Le Chatelier’s principle, the increase in temperature increases the evaporation rate. The increase in pressure will favour the condensation of water because water molecules in the liquid state have less pressure than the molecules in the vapour state.

3. The dissolution of solids in water: The dissolution of solid in water is affected by the temperature.
If the dissolution process is endothermic, then the forward reaction can be increased by increasing the temperature. On increasing the temperature, most of the solids dissolve in water due to the absorption of heat energy which overcomes the attractive forces that bind the oppositely charged ions in salt.

Example: \({\text{NaCl}}\left({\text{s}} \right){\text{ + }}\left({{\text{aq}}} \right) \rightleftharpoons {\text{NaCl}}\left({{\text{aq}}} \right);\Delta {\text{H= + 5}}{\text{.0}}\,{\text{kJmo}}{{\text{l}}^{{\text{-1}}}}\)

If the dissolution process is exothermic, then a decrease in the temperature will favour the dissolution of solid and favours forward reaction.
\({\text{NaOH}}\left({\text{s}} \right){\text{ + }}\left({{\text{aq}}} \right) \rightleftharpoons {\text{NaOH}}\left({{\text{aq}}} \right);\Delta {\text{H=-ve}}\)

4. Solubility of gas in liquid: When gas is dissolved in a liquid, a decrease in volume and dissolution process is exothermic. The dissolution increases with an increase in pressure and a decrease in temperature.

\({\text{Gas + Liquid}} \rightleftharpoons {\text{Dissolved}}\,{\text{gas}};\,\Delta {\text{H=-ve}}\)

Summary

In this article, we understood the concept of equilibrium, i.e., Le Chatelier’s principle, the effect of concentration, temperature and pressure on the equilibrium based on Le Chatelier’s principle. We also studied the examples of Le Chatelier’s Principle and its applications on physical and chemical equilibria.

FAQs on Le Chatelier’s Principle

Q.1. How is Le Chatelier’s principle used in real life?
Ans: Le Chatelier’s principle is instrumental in predicting the condition of temperature, pressure, and concentration to get high yields in specific industrial reactions like the preparation of ammonia, sulphuric acid, hydrogen, etc.

Q.2. What is Le Chatelier’s principle? Explain with examples.
Ans: According to Le Chatelier’s principle, if a system under equilibrium is subjected to a change in temperature, pressure or concentration, then the equilibrium shifts itself in such a way to undo or neutralise the effect of change.
Example: The reaction between yellow-coloured ferric nitrate and colourless potassium thiocyanate forms a red-coloured iron thiocyanate complex.
\({\text{Ice}}\left({\text{s}} \right) \rightleftharpoons {\text{Water}}\left({\text{l}} \right);\Delta {\text{H= + 6}}{\text{.0}}\,{\text{kJmo}}{{\text{l}}^{{\text{-1}}}}\)

Adding a small amount of ferric chloride to the equilibrium mixture will increase the red colour’s intensity. This indicates that equilibrium has shifted to the right. On adding iron thiocyanate, the intensity of the colour decreases indicating the equilibrium is shifted to the left.

Q.3. What are the applications of Le Chatelier’s principle?
Ans: Le Chatelier’s principle is useful in predicting the condition of temperature, pressure, and concentration to get high yields in certain industrial reactions like preparation of ammonia, sulphuric acid, hydrogen, etc. It is also used to predict the favourable conditions for melting of ice, evaporation of water, the solubility of a gas in the liquid, the dissolution of solids in water, etc.

Q.4. Does Le Chatelier’s principle apply to gases?
Ans: Yes, Le Chatelier’s principle is applicable to gases. Since, a gaseous system under equilibrium when subjected to a change in temperature, pressure or concentration, then the equilibrium shifts itself in such a way so as to undo or neutralise the effect of change.

Q.5. How does temperature affect Le Chatelier’s principle?
Ans: According to Le Chatelier’s principle, the increase in temperature shifts the equilibrium in the direction of the endothermic reaction. The decrease shifts the equilibrium in the direction of an exothermic reaction.

Study Rate Law Expression: Law of Mass Action

We hope this article on ‘Le Chatelier’s principle’ has helped you. If you have any queries, drop a comment below and we will get back to you.