9.04 Chapter Summary

3. Reversible Reactions and Equilibrium

3.1 Reversible Reactions

• Definition: Reactions where products can revert to reactants.
• Symbol: ⇌
• Example: Haber Process – N₂ + 3H₂ ⇌ 2NH₃
• Energy Consideration: If the forward reaction is exothermic, the reverse is endothermic.

3.2 Hydrated and Anhydrous Salts

• Hydrated Salts:
  • Definition: Salts containing water of crystallization.
  • Example: Copper(II) sulfate pentahydrate (CuSO₄·5H₂O) is blue.
• Anhydrous Salts:
  • Definition: Salts without water of crystallization.
  • Example: Anhydrous copper(II) sulfate (CuSO₄) is white.
• Transformations:
  • Heating Hydrated Salt:
    • Reaction: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O (endothermic).
    • Observation: Blue crystals become white.
  • Adding Water to Anhydrous Salt:
    • Reaction: CuSO₄ + 5H₂O → CuSO₄·5H₂O (exothermic).
    • Observation: White crystals become blue.

3.3 Equilibrium

• Definition: A state where the forward and reverse reaction rates are equal, and concentrations of reactants and products remain constant.
• Characteristics:
  • Dynamic Nature: Continuous formation and reversion of products and reactants.
  • Closed System: No exchange of substances with the environment.
• Le Chatelier’s Principle:
  • Statement: If a change is made to the conditions of a system at equilibrium, the system adjusts to counteract the change.
  • Applications:
    • Concentration Changes:
      • Increase Reactant: Shifts equilibrium to the right (more products).
      • Increase Product: Shifts equilibrium to the left (more reactants).
    • Pressure Changes (for Gases):
      • Increase Pressure: Shifts equilibrium towards fewer gas molecules.
      • Decrease Pressure: Shifts equilibrium towards more gas molecules.
    • Temperature Changes:
      • Increase Temperature: Shifts equilibrium towards the endothermic direction.
      • Decrease Temperature: Shifts equilibrium towards the exothermic direction.
    • Catalysts:
      • Effect: Speed up both forward and reverse reactions equally, no shift in equilibrium position.

4. Oxidation and Reduction (Redox) Reactions

4.1 Oxidation Numbers

• Definition: Numbers assigned to atoms to indicate the number of electrons lost or gained.
• Rules for Assigning Oxidation Numbers:
  1. Elements: Oxidation number is 0 (e.g., O₂, H₂).
  2. Monatomic Ions: Equal to the ion’s charge (e.g., Na⁺ = +1).
  3. Oxygen: Usually -2; -1 in peroxides (e.g., H₂O₂).
  4. Hydrogen: +1 when bonded to non-metals; -1 when bonded to metals (e.g., NaH).
  5. Fluorine: Always -1.
  6. Sum in Compound: Total oxidation numbers equal the overall charge (0 for neutral compounds).
• Examples:
  • Fe₂O₃:
    • Oxygen: -2 each.
    • Iron (Fe): +3 each.
  • CuSO₄·5H₂O:
    • Sulfur (S): +6.
    • Oxygen (O): -2 each.

4.2 Redox Reactions

• Definition: Reactions involving the transfer of electrons from one species to another.
• Processes:
  • Oxidation: Loss of electrons (increase in oxidation number).
  • Reduction: Gain of electrons (decrease in oxidation number).
• Agents:
  • Oxidizing Agent: Gains electrons (is reduced).
  • Reducing Agent: Loses electrons (is oxidized).
• Example Reaction:
  • Reaction: Fe + Cu²⁺ → Fe²⁺ + Cu
    • Iron (Fe): Oxidized (Fe → Fe²⁺ + 2e⁻).
    • Copper Ion (Cu²⁺): Reduced (Cu²⁺ + 2e⁻ → Cu).
    • Oxidizing Agent: Cu²⁺.
    • Reducing Agent: Fe.

4.3 Identifying Redox Reactions

• By Oxidation Numbers: Look for changes in oxidation states of elements.
• By Electron Transfer: Identify species losing or gaining electrons.
• By Color Changes: Certain oxidizing/reducing agents cause observable color changes.
  • Example: Acidified KMnO₄ (purple) becomes colorless when reduced.

4.4 Redox in Electron Transfer

• Half-Equations:
  • Oxidation Half-Reaction: Shows loss of electrons.
    • Example: Ag → Ag⁺ + e⁻.
  • Reduction Half-Reaction: Shows gain of electrons.
    • Example: O₂ + 4e⁻ → 2O²⁻.
• Overall Redox Reaction Example:
  • Reaction: Fe + Cu²⁺ → Fe²⁺ + Cu.
  • Oxidation: Fe → Fe²⁺ + 2e⁻.
  • Reduction: Cu²⁺ + 2e⁻ → Cu.

4.5 Testing for Redox Reactions

• Using Potassium Manganate(VII) (KMnO₄):
  • Oxidizing Agent: KMnO₄ is reduced, changing color from purple to colorless.
• Using Potassium Iodide (KI):
  • Reducing Agent: KI is oxidized, producing iodine (I₂), which imparts a red-brown color.

5. Industrial Processes

5.1 Haber Process (Ammonia Synthesis)

• Reaction: N₂ + 3H₂ ⇌ 2NH₃ (exothermic).
• Conditions:
  • Temperature: 450°C (compromise between reaction rate and yield).
  • Pressure: 200 atm (favors formation of fewer gas molecules, i.e., NH₃).
  • Catalyst: Iron (speeds up both forward and reverse reactions without being consumed).
• Stages:
  1. Gas Preparation:
     • Nitrogen (N₂): Obtained from air.
     • Hydrogen (H₂): Obtained from methane (CH₄) via steam reforming.
  2. Compression: Gases are compressed to high pressure (200 atm).
  3. Reaction: Compressed gases pass over iron catalyst at 450°C to form ammonia.
  4. Cooling: Ammonia is condensed and separated; unreacted gases are recycled.
  5. Recycling: Unreacted N₂ and H₂ are reused in the process.

5.2 Contact Process (Sulfuric Acid Production)

• Key Reaction: 2SO₂ + O₂ ⇌ 2SO₃ (exothermic).
• Conditions:
  • Temperature: 450°C (balanced for optimal rate and yield).
  • Pressure: 2 atm (favors formation of SO₃ while avoiding high pressures).
  • Catalyst: Vanadium(V) oxide (V₂O₅) to speed up reaction.
• Stages:
  1. Sulfur Dioxide Production:
     • Sources: Burning sulfur or roasting sulfide ores.
     • Reaction: S + O₂ → SO₂.
     • Roasting Ores: Metal sulfide + O₂ → Metal oxide + SO₂.
  2. Oxidation of SO₂ to SO₃:
     • Reaction: 2SO₂ + O₂ ⇌ 2SO₃.
     • Catalyst: V₂O₅ used to increase reaction rate.
  3. Absorption of SO₃:
     • Reaction: SO₃ + H₂SO₄ → H₂S₂O₇ (oleum).
     • Formation of Oleum: Prevents formation of a fine mist by directly absorbing SO₃ into concentrated H₂SO₄.
  4. Production of Sulfuric Acid:
     • Reaction: H₂S₂O₇ + H₂O → 2H₂SO₄.
     • Final Product: Concentrated sulfuric acid (H₂SO₄).

6. Equilibrium Concepts and Le Chatelier’s Principle

6.1 Equilibrium in Reversible Reactions

• Dynamic Equilibrium: Both forward and reverse reactions continue at equal rates.
• Closed System Requirement: Prevents loss or gain of reactants/products.
• Graphical Representation: Rates of forward and reverse reactions balance, resulting in constant concentrations.

6.2 Le Chatelier’s Principle Applications

• Concentration Changes:
  • Increase Reactant Concentration: Shifts equilibrium to the right (more products).
  • Decrease Reactant Concentration: Shifts equilibrium to the left (more reactants).
  • Increase Product Concentration: Shifts equilibrium to the left (more reactants).
  • Decrease Product Concentration: Shifts equilibrium to the right (more products).
• Pressure Changes (for Gaseous Systems):
  • Increase Pressure: Shifts equilibrium toward the side with fewer gas molecules.
  • Decrease Pressure: Shifts equilibrium toward the side with more gas molecules.
• Temperature Changes:
  • Increase Temperature: Favors the endothermic direction.
  • Decrease Temperature: Favors the exothermic direction.
• Catalysts:
  • Effect: Speed up attainment of equilibrium without shifting its position.

6.3 Worked Examples

Example 1: Effect of Temperature

• Reaction: ICl + Cl₂ ⇌ ICl₃ (exothermic forward reaction).
• Question: What happens to the color when the temperature is increased?
• Answer:
  • Application: Increase temperature favors the endothermic reverse reaction.
  • Effect: Equilibrium shifts to the left, producing more ICl (dark brown), making the mixture darker.

Example 2: Effect of Pressure

• Reaction: 2NO₂ (brown) ⇌ N₂O₄ (colorless).
• Question: What happens to the color when pressure is increased?
• Answer:
  • Application: Increase in pressure favors the formation of fewer gas molecules (N₂O₄).
  • Effect: More N₂O₄ is formed, shifting equilibrium to the right, making the mixture more colorless.

Example 3: Effect of Concentration

• Reaction: ICl + Cl₂ ⇌ ICl₃.
• Questions:
  1. Increase in ICl₃ Concentration:
     • Effect: Shifts equilibrium to the left, producing more ICl and Cl₂.
     • Observation: Mixture becomes darker (more ICl).
  2. Removal of Cl₂:
     • Effect: Shifts equilibrium to the left to produce more Cl₂ and ICl.
     • Observation: Mixture becomes darker (more ICl).