Exothermic and Endothermic Reactions

5.1 Exothermic and Endothermic Reactions

Definition and Energy Transfer

• Exothermic reaction:
  • Transfers thermal energy to the surroundings.
  • Surroundings get hotter → temperature increases.
  • Examples: combustion of fuels, neutralisation between acids and alkalis, respiration.
  • Energy change sign: ΔH < 0.
• Endothermic reaction:
  • Takes in thermal energy from the surroundings.
  • Surroundings get cooler → temperature decreases.
  • Examples: thermal decomposition, photosynthesis, dissolving some salts like NH₄NO₃ in water.
  • Energy change sign: ΔH > 0.

Enthalpy Change (ΔH)

• Definition: Amount of heat energy transferred during a chemical reaction under constant pressure.
• Units: kJ/mol.
• Sign conventions:
  • ΔH negative → exothermic (energy released).
  • ΔH positive → endothermic (energy absorbed).

Activation Energy (Eₐ)

• Definition: Minimum energy that colliding particles must have for a reaction to occur.
• Determines reaction rate — higher Eₐ means slower reaction unless temperature or catalyst changes.
• Catalysts lower the Eₐ, speeding up the reaction without being used up.

Reaction Pathway Diagrams

Exothermic Reaction

• Reactants start at a higher energy level than products.
• Energy is released → ΔH is negative.
• Peak represents activation energy barrier (Eₐ).

Features to label:

1. Reactants – higher energy.
2. Products – lower energy.
3. ΔH – vertical drop from reactants to products.
4. Eₐ – from reactants to the peak of the curve.

Endothermic Reaction

• Reactants start at a lower energy level than products.
• Energy is absorbed → ΔH is positive.
• Peak represents activation energy barrier (Eₐ).

Features to label:

1. Reactants – lower energy.
2. Products – higher energy.
3. ΔH – vertical rise from reactants to products.
4. Eₐ – from reactants to the peak of the curve.

Bond Breaking and Bond Making

• Bond breaking: Endothermic → energy is absorbed to overcome attractive forces between atoms.
• Bond making: Exothermic → energy is released when new bonds are formed.
• Overall enthalpy change:
  ΔH = Total energy absorbed (bond breaking) − Total energy released (bond making).

Bond Energy Calculations Example

Data given:

• H–H: 436 kJ/mol
• Cl–Cl: 243 kJ/mol
• H–Cl: 431 kJ/mol

Reaction: H₂ + Cl₂ → 2HCl

Step 1 – Energy absorbed to break bonds:
H–H = 436 kJ
Cl–Cl = 243 kJ
Total absorbed = 436 + 243 = 679 kJ

Step 2 – Energy released to make bonds:
2 × H–Cl = 2 × 431 = 862 kJ

Step 3 – ΔH calculation:
ΔH = 679 − 862 = −183 kJ/mol → Exothermic.

Examples in Real Life

• Exothermic:
  • Combustion of fuels (e.g., CH₄ + 2O₂ → CO₂ + 2H₂O).
  • Neutralisation (e.g., HCl + NaOH → NaCl + H₂O).
• Endothermic:
  • Thermal decomposition (e.g., CaCO₃ → CaO + CO₂).
  • Photosynthesis (6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂).