Le Chatelier’s Principle is a fundamental concept in chemistry that describes how a chemical system at equilibrium responds to changes in temperature, pressure, or concentration of reactants or products. This principle is named after the French chemist Henry Louis Le Chatelier, who formulated it in the late 19th century. Equilibrium refers to the state of balance which tells that there is equal weight on both sides of the scale. Chemical equilibrium is attained when the rate of forward reaction is equal to the rate of backward reaction.

In this article, we will have a thorough knowledge of equilibrium, Le Chatelier’s principle, and the effect of various factors on chemical equilibrium.

What is Le Chatelier’s Principle of Equilibrium?

Le Chatelier’s Principle was proposed by Le Chatelier, a renowned French chemist to describe the effect of change in pressure, temperature, and concentration on any reversible system. In other words, a change in any of the factors that determine the equilibrium conditions of a system will cause the system to change in such a manner to reduce the effect of change.

Le Chatelier’s Principle definition

According to Le Chatelier’s principle, If a system at equilibrium is subjected to a change of any one of the factors such as concentration, pressure, or temperature, the system adjusts itself in such a way as to oppose the effect of that change.

Le Chatelier’s Example

Consider the example of synthesis of ammonia which is given by the following equation:

N₂(g) + 3H₂(g) ⇄ 2NH₃(g) + 22400cal

Then, as per Le Chatelier’s Principle, following parameters affect the reaction:

• Effect of Concentration: When concentration of N₂ or H₂ increases, rate of forward reaction increases thereby increasing the concentration of NH₃.
• Effect of Pressure: When NH3 is formed, there is a decrease in volume. When pressure is increased, the system will move in direction of lower volume so as to follow the Le Chatelier’s principle. Therefore, high pressure helps in formation of NH₃.
• Effect of Temperature: It is an exothermic reaction. Hence, the concentration of NH₃ will increase at low temperature.

Le Chatelier’s Principle

Le Chatelier’s Principle considers following factors affecting equilibrium:

• Effect of Concentration Change
• Effect of Pressure Change
• Effect of Temperature Change
• Effect of Volume Change
• Effect of Catalyst
• Effect of Addition of Inert Gas

Let’s discuss them in detail below:

Effect of Concentration Change on Equilibrium

In chemical equilibrium, increasing the concentrations of the reactants results in shifting the reaction in favor of products so as to oppose the effect of increase of concentration of reactants. Hence, the rate of forward reaction increases. Similarly, increasing the concentrations of the products results in shifting the reaction in favor of reactants so as to oppose the effect of increase of concentration of reactants. Hence, the rate of backward reaction increases.

Effect of Change in Concentration of Reactant

• If the concentration of reactant increases, the equilibrium will shift in the direction so as to favor the formation of product.

2SO₂(g) + O₂(g) ⇄ 2SO₃(g)

• If the concentration of reactant decreases, the equilibrium will shift in the direction to favor the increase in concentration of reactant. Hence, the reaction will be in backward direction.

For example, If the concentration of SO₂ or O₂ is increased then the reaction will shift in the direction so as to oppose the increase in concentration of reactants. Hence, the reaction will shift in forward direction and vice versa will be in the case of decrease in concentration of reactant.

Effect of Change in Concentration of Product

• If the concentration of product increases, the equilibrium will shift in the direction so as to favor the formation of reactant. Hence, the direction of reaction will shift to backward direction

PCl₃(g) + Cl₂(g) ⇄ PCl₅(g)

• If the concentration of product decreases, the equilibrium will shift to favour increase in product. Hence, the direction of reaction will be in forward direction

For example, If the concentration of PCl₅ is decreased, then the rate of reaction is also increased which leads to decrease in concentration of chlorine and increase in concentration of phosphorus pentachloride. Hence, equilibrium will shift to right so rate of forward reaction increases and vice-versa will be in the case of increase of concentration of product.

Effect of Change in Pressure on Equilibrium

In chemical equilibrium, an increase in pressure applied to a system at equilibrium, favors the reaction in the direction that produces lower number of moles of gases and a decrease in pressure favors the opposite reaction. Since, pressure and number of moles are related. The relation between equilibrium constants k_p and k_c is given by the following formula:

K_p = K_c (RT)^Δn = K_c (p/v)^Δn

where,

• kp and kc are equilibrium constant in partial pressure and number of moles
• R is gas constant
• T is temperature
• n is number of moles

High pressure is favorable for the reaction in which there is increase in volume. Change in pressure only affect the gaseous reaction.

• Increase in pressure shifts the equilibrium in the direction of lesser number of gaseous moles.
• Decrease in pressure shifts the equilibrium in the direction of greater number of gaseous moles.

Note: If there is no change in number of gas molecules in a reaction, then the pressure change doesn’t affect the equilibrium.

• n₁ = Number of gaseous moles of reactants
• n₂ = Number of gaseous moles of products

△n_g = n₂ – n₁

• △n_g = 0; No effect of change in pressure
• △n_g = +ve; Increase in pressure will decrease product formation and vice versa.
• △n_g = -ve; Increase in pressure will increase product formation.

Let us understand it with examples

When △n_g is -ve

Increase in pressure will increase product formation and shift the equilibrium towards products and hence concentration of NH₃ and SO₃ will increase.

• High pressure favors the shift of equilibrium towards products.

N₂(g) + 3H₂(g) ⇄ 2NH₃(g); n₁ = 4, n₂ = 2; △n_g = -2

2SO₂(g) + O₂(g) ⇄ 2SO₃(g); n₁ = 3, n₂ = 2; △n_g = -1

When △n_g is +ve

Decrease in pressure will increase reactants and shift the equilibrium towards reactants and hence concentration of PCl5 and N2O4 will increase.

• Low pressure favors the shift of equilibrium towards reactants.

PCl₅(g) ⇄ PCl₃(g) + Cl₂(g); n₁ = 1, n₂ = 2; △n_g = 1

N₂O₄(g) ⇄ 2NO₂(g); n₁ = 1, n₂ = 2; △n_g = 1

When △n_g is Zero

There will be no effect of pressure if number of moles of gases is the same on both sides of reaction. Here, the concentration of both reactants and products remain same. Hence, △n_g = 0

H₂(g) + I₂(g) ⇄ 2HI(g) ; n₁ = 2, n₂ = 2; △n_g = 0

N₂(g) + O₂(g) ⇄ 2NO(g); n₁ = 2, n₂ = 2; △n_g = 0

If the species are not in their gaseous state hence there will be no effect of pressure change.

2NaNO₃(s) ⇄ 2NaNO₃(s) + O₂(g); n₁ = 2, n₂ = 3; △n_g = 1

NH₄HS(s) ⇄ NH₃(g) + H₂S(g) ; n₁ = 1, n₂ = 2; △n_g = 1

Effect of Volume Change on Product Formation

According to Le Chatelier’s Principle

• When there is a decrease in volume, the equilibrium will shift in the direction with fewer number of gaseous moles.
• When there is a increase in volume, the equilibrium will shift in the direction with more number of gaseous moles.

Some examples of volume change are:

• Equilibrium will shift in forward direction as products have lesser number of moles.

2H₂(g) + CO(g) ⇄ CH₃OH(g)

• Equilibrium will shift in backward direction as reactants have lesser number of moles.

4NH₃(g) + 5O₂(g) ⇄ 4NO(g) + 6H₂O(g)

Effect of Change in Temperature on Equilibrium

According to Le Chatelier’s Principle, the position of equilibrium moves in such a way as to tend to undo the change that we have made. If we increase the temperature, the position of equilibrium will move in such a way as to reduce the temperature again and it will be done by favoring the reaction which absorbs heat. A rise in temperature favors the endothermic reaction as the heat speeds up the rate of reaction by providing sufficient energy to reactant molecules to complete reaction faster.

Le Chatelier’s Principle for temperature change is governed by the following equation:

log[k₁/k₂] = -△H/R[{1/T₂}-{1/T₁}]

where,

• k is equilibrium Constant
• △H is change in enthalpy
• T is Temperature
• R is gas constant

According to this relation, increase in temperature will decrease k₂ and vice versa.

As per Le Chatelier’s Principle, increase in temperature shifts the equilibrium in the direction of endothermic reaction and decrease in temperature shifts the equilibrium in the direction of exothermic reaction.

• △H > 0 (Heat is absorbed): Endothermic reaction in which increase in temperature will shift the reaction towards right.
• △H < 0 (Heat is evolved): Exothermic reaction in which increase in temperature will shift the reaction towards left.

Let’s see some examples of this

High temperature favors forward reaction owing to its endothermic nature. In this case, the equilibrium will shift in backward direction and there will be decrease in product

N₂(g) + O₂(g) ⇄ 2NO (g); △H = +180 KJ (Heat is absorbed)

PCl₅(g) ⇄ PCl₃(g) + Cl₂(g); △H = +63 KJ (Heat is absorbed)

Low temperature favors forward reaction owing to its exothermic nature. In this case, the equilibrium will shift in forward direction and there will be increase in product formation.

H₂(g) + I₂(g) ⇄ 2HI(g) ; △H = -10.4 KJ (Heat is evolved)

N₂(g) + 3H₂(g) ⇄ 2NH₃(g) ; △H = -93.6 KJ (Heat is evolved)

2SO₂(g) + O₂(g) ⇄ 2SO₃(g) ; △H = -189 KJ (Heat is evolved)

Effect of Catalyst on Equilibrium

The use of appropriate catalyst will increase the rate of reaction thereby decreasing the time required to reach the equilibrium position. A catalyst can’t change the equilibrium constant Hence, the relative ratio of concentration of reactants and products remain unaffected.

For reversible reaction, catalyst helps in reaching the equilibrium position by increasing the rate of forward and backward reaction.

Eg: Haber’s Process

N₂(g) + 3H₂(g) ⇄ 2NH₃(g)

In Haber’s process the yield of ammonia is low at high temperature and reaction is slow at low temperature. So, in order to increase the speed of reaction at low temperature mixture of iron and molybdenum is added as a catalyst.

Effect of Addition of an Inert gas on Equilibrium

All the gases which do not participate in reaction are considered as inert gas. The effect of addition of inert gas on equilibrium at different conditions in discussed below:

Addition of an Inert Gas at Constant Volume

When an inert gas is added to the equilibrium at a constant volume, then the total pressure of the system will increase. But the concentrations of reactants and products will remain unaffected. Therefore, there will be no effect on the equilibrium.

Addition of an Inert Gas at Constant Pressure

When an inert gas is added to the equilibrium at a constant pressure, the volume increases. The number of moles per unit volume of various reactants and products decreases. Thus to overcome this the equilibrium will shift to direction with increased number of moles. Addition of inert gas at constant pressure will shift the equilibrium to backward direction.

Application of Le Chatelier’s Principle

Le Chatelier’s principle is a fundamental concept in chemistry that describes how a system at equilibrium responds to changes in temperature, pressure, or concentration of reactants or products. This principle is applied in various chemical processes and industries to optimize reactions, maximize yields, and control equilibrium conditions. Some key applications of Le Chatelier’s principle include:

• Chemical Equilibrium: Le Chatelier’s principle helps chemists predict and understand the effect of changes in concentration, temperature, or pressure on the equilibrium position of a chemical reaction.
• Industrial Chemical Processes: Le Chatelier’s principle is extensively used in industrial chemical processes such as the Haber process for ammonia synthesis, the contact process for sulfuric acid production, and the synthesis of methanol. By manipulating reaction conditions, such as temperature and pressure, production efficiency and yield can be maximized
• Catalysis: Understanding Le Chatelier’s principle is crucial in catalysis, where catalysts are used to increase the rate of chemical reactions. By adjusting reaction conditions to favor the desired equilibrium, catalytic processes can be optimized for improved efficiency and selectivity.
• Environmental Chemistry: Le Chatelier’s principle is relevant in environmental chemistry, particularly in the study of air pollution control and atmospheric chemistry. For example, in the atmosphere, changes in temperature and pressure can affect the equilibrium between various gases, leading to the formation or removal of pollutants.

Related Reads

• Equilibrium in Chemical Processes
• Homogeneous and Heterogeneous Equilibrium

Le Chateliers Principle FAQs

What is Chemical Equilibrium?

Chemical equilibrium is a state where rate of forward reaction is equal to the rate of backward reaction.

State Le Chatelier’s Principle.

According to Le Chatelier’s principle, If a system at equilibrium is subjected to a change of any one of the factors such as concentration, pressure or temperature, the system adjusts itself in such a way as to oppose the effect of that change.

What is Effect of Addition of Inert Gas in a Reaction with Equal moles of Reactants and Products?

There will be no effect of addition of inert gas in a reaction with equal number of moles of reactants and products.

Where the Equilibrium will shift if the Concentration of Products is Increased?

Increase in concentration of products shifts the reaction in backward direction.

Decreasing the Temperature will shift the Equilibrium in Which Direction?

Decrease in temperature i.e. low temperature shifts the equilibrium in direction of reactant.