Le Chatelier’s Principle

Henri Le Chatelier's principle states that when a system at equilibrium is subjected to stress (changes in concentration, temperature, or pressure), it adjusts to counteract the stress. Favouring the forward reaction increases product concentrations, while favouring the reverse reaction increases reactant concentrations.

1.0Le Chatelier’s Principle

Le Chatelier's principle states that if a change is applied to a system at dynamic equilibrium, the system adjusts to counteract that change and reestablish equilibrium.
This principle is a powerful tool for predicting how temperature, pressure, or concentration changes affect the equilibrium position.

Position of Equilibrium

The position of equilibrium describes the relative amounts of reactants and products in a mixture at equilibrium:

• Shifting to the left: Indicates an increase in the concentration of reactants.
• Shifting to the right: Indicates an increase in the concentration of products.

2.0Effects of Concentration Change

When the concentration of a substance in the reaction changes, the equilibrium shifts to counteract this change:

• Increase in reactant concentration:
• • The forward reaction rate increases.
  • More products are formed, shifting the equilibrium to the right.
• Decrease in reactant concentration:
• • The backward reaction rate increases.
  • More reactants are formed, shifting the equilibrium to the left.
• Increase in product concentration:
• • The backward reaction rate increases.
  • The equilibrium shifts to the left to reduce the product concentration.
• Decrease in product concentration:
• • The forward reaction rate increases.
  • The equilibrium shifts to the right to produce more products.

1. Effects of Concentration on Equilibrium

2. Effect of Concentration on the Equilibrium Constant (K)

The equilibrium constant, KKK, remains unchanged by changes in the concentration of reactants or products as long as other conditions (e.g., temperature) stay constant.

Example: Decomposition of Hydrogen Iodide

The reaction:

$2HI\leftrightharpoons H_2+I_2$

The equilibrium expression:

$K=[H2][I2][HI]2=6.25\times 10-3K=\frac{[H_2][I_2]}{[HI{]}^2}=6.25\times 10^{-3}K=[HI]2[H2][I2]=6.25\times 10-3$

• Adding more HIHIHI:
• • The ratio $[H2][I2][HI]2\frac{[H_2][I_2]}{[HI{]}^2}[HI]2[H2][I2]$​ initially decreases.
  • To restore equilibrium, $[H2][H_2][H2]$ and $[I2][I_2][I2]$ increase, while [HI][HI][HI] decreases.

Equilibrium is reestablished when the ratio returns to $6.25\times 10^{-3}$.

3.0Effects of Pressure Change

Pressure changes affect only reactions involving gases.

• Understanding pressure in gases:
• • In a fixed volume, the pressure increases as the number of gas molecules increases.
  • Changes in pressure cause the equilibrium to shift in a way that counteracts the change.
• Increase in pressure:
• • The equilibrium shifts towards the side with fewer gas molecules to reduce pressure.
• Decrease in pressure:
• • The equilibrium shifts towards the side with more gas molecules to increase pressure.

1. Effects of Pressure on Equilibrium

2. Effects of Pressure on the Equilibrium Constant

The equilibrium constant K remains unchanged when pressure changes, provided all other factors (e.g., temperature) stay constant.

Note: If the number of gas molecules is the same on both sides of the equation, changes in pressure do not affect the equilibrium position.

4.0Effects of Temperature  Change

Temperature changes affect equilibrium by causing the system to shift in a direction that absorbs or releases energy to counteract the change.

Effects of Temperature on Equilibrium Position

• Increase in temperature:
• • The equilibrium shifts in the endothermic direction to absorb energy and counteract the change.
• Decrease in temperature:
• • The equilibrium shifts in the exothermic direction to release energy and counteract the change.

Effects of Temperature on Equilibrium

Effects of Temperature on the Equilibrium Constant (K)

Unlike pressure or concentration changes, temperature directly affects the equilibrium constant K.

1. For an Endothermic Reaction

Example:

$2HI(g)\rightleftharpoons H_2(g)+I_2(g)$

$K=\frac{[H_2][I_2]}{[HI{]}^2}K=[HI]2[H2][I2]$

• Increase in temperature:
• • The equilibrium shifts right (in the endothermic direction), increasing the concentrations of [H2]$[H2][H_2][H2]and[I2][I_2][I2]$ while [HI][HI][HI] decreases.
  • Result: K increases.

2. For an Exothermic Reaction

Example:
2SO₂(g)+O₂(g)⇌2SO₃(g)

$K=\frac{[S{O}_3{]}^2}{[S{O}_2{]}^2[O_2]}$

• Increase in temperature:
• • The equilibrium shifts left (endothermic direction), decreasing [SO3][SO_3][SO3​] while [SO2][SO_2][SO2​] and [O2][O_2][O2​] increase.
  • Result: K decreases.

5.0Effect of Inert Gas Addition

The equilibrium position remains unaffected when an inert gas, such as argon, is added to a system at constant volume.

• Adding an inert gas does not alter the partial pressures or molar concentrations of the substances involved in the reaction.
• The reaction quotient is only affected if the added gas participates as a reactant or product in the reaction.

6.0Effects of a Catalyst

A catalyst is a substance that speeds up the rate of a chemical reaction by equally increasing the rates of both the forward and reverse reactions.

• Catalysts do not alter the equilibrium position or the equilibrium constant's value (K).
• They simply help the system reach equilibrium more quickly.

Industrial Processes and Catalysts

In industrial processes, Le Chatelier’s principle predicts conditions that shift the equilibrium toward the desired products, maximising yield. However, reaction kinetics must also be considered, as the reaction rate needs to be practical for efficient production.

Example:
For a reversible reaction where the forward reaction is exothermic:

• Le Chatelier’s principle states lower temperatures favour the forward reaction, producing a higher equilibrium yield.
• However, lower temperatures slow the reaction rate.

Important: A moderate temperature is chosen to balance yield and reaction speed, resulting in a slightly lower yield but faster production.

7.0Heterogeneous Equilibria

Le Chatelier’s principle applies to heterogeneous equilibria as well.

Example:
In a fizzy drink bottle, equilibrium exists between dissolved $C{O}_2$ and gaseous $C{O}_2$:

${\text{CO}}_2(g)\rightleftharpoons {\text{CO}}_2(aq)$

• When the bottle is opened, gaseous $C{O}_2$​ escapes, causing the equilibrium to shift left to replace the lost gas.
• This shift produces visible bubbles of $C{O}_2$.

8.0Solved Examples 

Example 1 

Q. Use the reaction below:

$C{O}_2(g)+H_{2}O(l)\leftrightharpoons H_{2}C{O}_3(aq)$

Explain what happens to the position of equilibrium when the system is diluted by adding more water.

Explanation

When more water is added to the system, the concentration of $H_{2}O$ effectively increases. The position of equilibrium shifts to the right, favouring the formation of $H_{2}C{O}_3(aq)$, to reduce the effect of the change according to Le Chatelier's principle.

Example 2 

Q. Consider the equilibrium reaction:

$COB{r}_2(g)\rightleftharpoons CO(g)+B{r}_2(g)$

What happens when inert argon gas, Ar(g), is added?

Explanation

 When an inert gas like $Ar(g)$ is introduced to a system at equilibrium, it does not alter the concentrations of the reactants or products, as it does not participate in the reaction. Consequently, the reaction remains at equilibrium.

Example 3

Q. Consider the endothermic reaction:

$N{H}_{4}HS(s)+heat\rightleftharpoons N{H}_3(g)+H_{2}S(g)$

What happens when the temperature is increased?

Explanation

 For an endothermic reaction, heat can be considered a reactant:

$N{H}_{4}HS(s)+heat\rightleftharpoons N{H}_3(g)+H_{2}S(g)$

When the temperature increases, the system shifts to favour the forward reaction to absorb the added heat and reestablish equilibrium