Chemical Equilibrium

Full Notes

Reversible reactions can go forward (reactants → products) and backward (products → reactants).

A dynamic equilibrium is reached in a closed system when:

• The rate of the forward reaction = the rate of the reverse reaction.

At this point, the concentrations of reactants and products remain constant (but are not necessarily equal).

It is dynamic because both reactions continue, but there is no overall change in amounts.

Example: The Haber Process (Ammonia Production)

Forward reaction: N₂ + H₂ → NH₃
Reverse reaction: NH₃ decomposes into N₂ and H₂.

Equilibrium will be reached when forward and reverse reaction rates are equal, so concentrations remain constant.

Le Chatelier’s Principle

Le Chatelier’s Principle states: “If a system at equilibrium is subjected to a change, the position of equilibrium will shift to oppose the change”.

This is a very important principle for dealing with equilibrium systems and we can use it to influence a position of equilibrium.

Changing Concentration:

• Increasing reactant concentration shifts equilibrium right (products increase).
• Increasing product concentration shifts equilibrium left (reactants increase).

Example: Haber Process
Adding more N₂ shifts equilibrium right, producing more NH₃.

Changing Pressure (for Gaseous Equilibria):

• Increasing pressure shifts equilibrium to the side with fewer gas molecules.
• Decreasing pressure shifts equilibrium to the side with more gas molecules.
• No effect if there are equal numbers of gas molecules on both sides.

Example: Haber Process

4 moles (N₂ + 3H₂) ⇌ 2 moles (NH₃)

Higher pressure shifts equilibrium right, increasing NH₃ yield.

Changing Temperature:

• Increasing temperature favours the endothermic direction (+ΔH).
• Decreasing temperature favours the exothermic direction (−ΔH).

Example: Haber Process — forward reaction is exothermic (−ΔH).

Increasing temperature shifts equilibrium left, reducing NH₃ yield.
Decreasing temperature shifts equilibrium right, increasing NH₃ yield.

Catalysts and Equilibrium

A catalyst speeds up both forward and reverse reactions by providing an alternative route with lower activation energy.

• Catalysts do not change the position of equilibrium.
• They allow equilibrium to be reached faster.
• Important in industrial processes for efficiency.

Investigating Equilibrium Shifts

Equilibrium shifts can be observed experimentally using colour changes.

Example: Chromate–Dichromate Equilibrium

CrO₄²⁻ (yellow) + 2H⁺ ⇌ Cr₂O₇²⁻ (orange)

• Add acid (more H⁺): equilibrium shifts right → orange solution.
• Add alkali (removes H⁺): equilibrium shifts left → yellow solution.

Equilibrium in Industry

In industry, conditions must balance rate of production and yield of product. High yield may not be practical if rate is too slow, so compromises are made.

Example: Haber Process (N₂ + 3H₂ ⇌ 2NH₃)

• Forward reaction is exothermic.
• High pressure favours NH₃ but is costly and dangerous.
• High temperature increases rate but favours reverse reaction.
• Compromise: 450°C, 200 atm, iron catalyst.
• This ensures reasonable yield at a practical rate.

Summary

• Dynamic equilibrium occurs when forward and reverse rates are equal, and concentrations remain constant.
• Le Chatelier’s Principle predicts shifts in equilibrium position when conditions change.
• Equilibrium position is affected by concentration, pressure, and temperature.
• Catalysts do not shift equilibrium but help it reach faster.
• Industrial conditions are compromises between yield, rate, safety, and cost.