Reversible Reactions and Equilibrium

Reversible Reactions

• Definition:
  • A reversible reaction is one in which the products can react together to form the original reactants.
  • Represented by the double arrow symbol: ⇌
  • Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
• Explanation:
  • The forward reaction converts reactants into products.
  • The backward (reverse) reaction converts products back into reactants.
  • Both reactions can occur simultaneously in opposite directions.

Examples of Reversible Reactions in Chemistry

1. Thermal decomposition and rehydration of hydrated compounds
   • Hydrated copper(II) sulfate:
     • Heating hydrated copper(II) sulfate crystals (CuSO₄·5H₂O) drives off water, forming white anhydrous copper(II) sulfate (CuSO₄).
       CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(g)
     • Adding water to the white powder reforms the blue crystals.
   • Hydrated cobalt(II) chloride:
     • Heating pink hydrated cobalt(II) chloride (CoCl₂·6H₂O) forms blue anhydrous cobalt(II) chloride.
       CoCl₂·6H₂O(s) ⇌ CoCl₂(s) + 6H₂O(g)
     • Adding water to the blue powder reforms the pink hydrated salt.
2. Addition of water to anhydrous compounds
   • The reverse of heating hydrated compounds.
   • Adding water changes colour and physical state as the hydration occurs.

Equilibrium in a Closed System

• Definition of a closed system:
  • A system where no substances can enter or leave, but energy can be exchanged with surroundings.
• Conditions at equilibrium:
  • The rate of the forward reaction = the rate of the reverse reaction.
  • The concentrations of reactants and products remain constant (but not necessarily equal).
  • Microscopic changes still occur — particles are continually reacting, but overall quantities remain unchanged.
• Dynamic nature:
  • Even at equilibrium, bonds are breaking and forming; the system is not static.

Factors Affecting the Position of Equilibrium (Le Chatelier’s Principle)

• Principle: If a system at equilibrium is subjected to a change, the system shifts to oppose the change.

1. Effect of Temperature

• Increasing temperature favours the endothermic direction of the reaction.
• Decreasing temperature favours the exothermic direction.
• Example: In the Haber process, the forward reaction is exothermic, so higher temperature shifts equilibrium to the left (less ammonia), but increases rate.

2. Effect of Pressure (for gases only)

• Increasing pressure shifts equilibrium towards the side with fewer moles of gas.
• Decreasing pressure shifts equilibrium towards the side with more moles of gas.
• Example: In the Haber process, 4 moles (N₂ + 3H₂) → 2 moles (NH₃). Higher pressure favours ammonia production.

3. Effect of Concentration

• Increasing the concentration of reactants shifts equilibrium to the right (more products formed).
• Increasing concentration of products shifts equilibrium to the left (more reactants formed).

4. Effect of a Catalyst

• A catalyst speeds up both forward and reverse reactions equally.
• It reduces the time taken to reach equilibrium but does not change the position of equilibrium.

The Haber Process (Ammonia Production)

• Equation: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
• Source of reactants:
  • Nitrogen: from air (78% nitrogen).
  • Hydrogen: from methane (CH₄) via steam reforming:
    CH₄ + H₂O → CO + 3H₂
• Typical conditions:
  • Temperature: 450°C
  • Pressure: 20 000 kPa / 200 atm
  • Catalyst: Finely divided iron
• Reason for conditions:
  • High pressure favours ammonia formation (fewer moles of gas), but higher cost and safety risks limit maximum pressure.
  • Moderate temperature balances yield and rate — low temperature gives high yield but slow rate; high temperature gives fast rate but low yield.
  • Iron catalyst speeds up reaction without changing equilibrium position.

The Contact Process (Sulfuric Acid Production)

• Step 2 equation: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
• Sources of reactants:
  • Sulfur dioxide: burning sulfur (S + O₂ → SO₂) or roasting sulfide ores (2ZnS + 3O₂ → 2ZnO + 2SO₂).
  • Oxygen: from air.
• Typical conditions:
  • Temperature: 450°C
  • Pressure: 200 kPa / 2 atm
  • Catalyst: Vanadium(V) oxide (V₂O₅)
• Reason for conditions:
  • The forward reaction is exothermic — lower temperature favours SO₃ yield, but too low slows the rate.
  • Slightly above atmospheric pressure is used for economic and safety reasons; higher pressures increase yield only slightly but cost more.
  • Catalyst ensures fast rate at moderate temperature.

Industrial Considerations for Choosing Conditions

• Rate vs yield trade-off:
  • Too low temperature = high yield but too slow rate.
  • Too high temperature = fast rate but low yield.
• Economic factors:
  • Higher pressures require expensive, stronger equipment.
  • Operating costs and safety must be balanced with output.
• Catalyst use:
  • Reduces energy costs by allowing operation at lower temperature.
• Continuous removal of products:
  • Increases yield by shifting equilibrium to favour more product formation.