1.8.2 Le Chatelier's Principle

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Le Chatelier's principle helps predict how a change in conditions affects the position of equilibrium in a reversible reaction.

It states that "if a system at equilibrium experiences a change in concentration, pressure, or temperature, the equilibrium will shift in a direction that opposes this change."

Effects of Changing Conditions

Change in Pressure

Increasing pressure shifts the equilibrium toward the side with fewer gas molecules to reduce the pressure.
Decreasing pressure shifts the equilibrium toward the side with more gas molecules to increase the pressure.

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Example: For the reaction:

2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)

Increasing pressure shifts the equilibrium to the right (fewer moles of gas on the product side), increasing the yield of $SO_3$

Decreasing pressure shifts the equilibrium to the left (more moles of gas on the reactant side), increasing the yield of $SO_2$ and $O_2$

Change in Temperature

If the forward reaction is exothermic, an increase in temperature shifts the equilibrium toward the reverse (endothermic) reaction, reducing the yield of products.
A decrease in temperature favours the exothermic reaction, shifting the equilibrium toward the products.

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Example: For the reaction

2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)

where

\Delta H = :highlight[-197 \, kJ/mol]

Increasing temperature shifts the equilibrium to the left (endothermic direction).

Decreasing temperature shifts the equilibrium to the right (exothermic direction), increasing the yield of $SO_3$

Change in Concentration

Increasing the concentration of reactants shifts the equilibrium to the right, increasing product formation.
Increasing the concentration of products shifts the equilibrium to the left, increasing reactant formation.

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Example: For the reaction

2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)

Adding more $SO_2$ shifts the equilibrium to the right, producing more $SO_3$

Adding more $SO_3$ shifts the equilibrium to the left, increasing the concentrations of $SO_2$ and $O_2$

Use of a Catalyst

A catalyst increases the rate of both the forward and reverse reactions equally. This means that equilibrium is reached more quickly, but it does not affect the position of equilibrium or the concentrations of reactants and products.

Industrial Applications: Compromise Conditions

In industrial processes, achieving both high yield and fast reaction rates is often necessary. Compromise conditions are used to balance the demands of maximizing yield and maintaining a practical reaction rate.

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Example: Ethanol Production (Hydration of Ethene) The production of ethanol from ethene is a reversible reaction:

C_2H_4(g) + H_2O(g) \rightleftharpoons C_2H_5OH(g)

\quad \Delta H = :highlight[-46 \, kJ/mol]

Catalyst: Phosphoric acid speeds up the reaction but does not affect the equilibrium position.

Pressure: High pressure shifts the equilibrium to the right (fewer gas molecules), increasing ethanol yield. However, very high pressures increase costs and require stronger equipment.

Temperature: A lower temperature favours the exothermic forward reaction, increasing yield. But, too low a temperature slows the reaction, so a compromise temperature of 300°C is used to balance yield and rate.

Excess Reactants: Using excess ethene or steam shifts the equilibrium to the right, increasing ethanol yield.

Practical Example for Students

You can observe the effect of concentration and temperature changes on equilibrium by performing a test-tube experiment with the following system:

[Cu(H_2O)_6]^{2+}(aq) + 4Cl^-(aq) \rightleftharpoons [CuCl_4]^{2-}(aq) + 6H_2O(l)

Adding concentrated HCl shifts the equilibrium to the right, changing the solution colour from blue to green.

Le Chatelier's Principle Simplified Revision Notes for A-Level AQA Chemistry

1.8.2 Le Chatelier's Principle

Effects of Changing Conditions

Change in Pressure

Change in Temperature

Change in Concentration

Use of a Catalyst

Industrial Applications: Compromise Conditions

Practical Example for Students

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