Types of Reactions in Relation to Heat Changes

1. Exothermic Reactions
   • Definition: Reactions that release heat energy to the surroundings.
   • Characteristics:
     • Surroundings get hotter.
   • Examples:
     • Combustion
     • Neutralization reactions
     • Rusting
     • Dehydration of sucrose by sulfuric acid
     • Reaction between sodium hydroxide and water
2. Endothermic Reactions
   • Definition: Reactions that absorb heat energy from the surroundings.
   • Characteristics:
     • Surroundings feel colder.
   • Examples:
     • Photosynthesis
     • Dissolution of salts in water (e.g., ammonium nitrate)
     • Evaporation of water
     • Melting of ice cubes

Investigating Temperature Changes

Aim: To describe temperature changes involved in exothermic and endothermic reactions.

Materials:

• 2 test tubes, measuring cylinder, thermometer, spatula, tap water, ammonium chloride, sodium hydroxide.

Procedure:

1. Pour 5 cm³ of water into each test tube.
2. Measure initial temperatures and record.
3. Add ammonium chloride to one test tube; measure and record final temperature.
4. Repeat with sodium hydroxide in the other test tube.

Observations:

• Ammonium Chloride: Temperature decreases, test tube feels cold → Endothermic
• Sodium Hydroxide: Temperature increases, test tube feels hot → Exothermic

Interpretation:

• Decrease in temperature: heat absorbed (endothermic).
• Increase in temperature: heat released (exothermic).

Conclusion:

• Endothermic reactions: temperature decreases.
• Exothermic reactions: temperature increases.

Identifying Reactions from Thermo-Chemical Equations

• Enthalpy (H): Energy stored in bonds; measured in kilojoules (kJ).
• Enthalpy Change (∆H): Change in energy from reactants to products.
  • Exothermic: ∆H is negative.
  • Endothermic: ∆H is positive.

Examples of Thermo-Chemical Equations:

1. NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l), ∆H = -57 kJ/mol

• Solution: Exothermic (negative ∆H).

1. H₂O (l) → 2H₂ (g) + O₂ (g), ∆H = +575 kJ/mol

• Solution: Endothermic (positive ∆H).

Energy Level Diagrams

Energy Level Diagram for Exothermic Reactions

• Description: Shows a decrease in energy as the reaction proceeds.
• Diagram Representation:
  • Reactants at a higher energy level than products.
  • Arrow points downward, indicating heat energy is released.

Energy Level Diagram for Endothermic Reactions

• Description: Shows an increase in energy as the reaction proceeds.
• Diagram Representation:
  • Reactants at a lower energy level than products.
  • Arrow points upward, indicating heat energy is absorbed.

Worked Examples

Draw Energy Level Diagrams for the Following Reactions:

1. H₂ (g) + Cl₂ (g) → 2HCl (aq), ∆H = -92 kJ/mol
   • Diagram: Reactants at a higher level, products at a lower level, arrow pointing down.
2. NH₄NO₃ (s) + H₂O (l) → NH₄⁺ (aq) + NO₃⁻ (aq), ∆H = +28.1 kJ/mol
   • Diagram: Reactants at a lower level, products at a higher level, arrow pointing up.

Bond Energies

• Definition: Energy required to break one mole of chemical bonds in a compound. Measured in kJ/mol.

Examples of Bond Energies

• H–H: 436 kJ/mol
• O=O: 498 kJ/mol
• O–H: 464 kJ/mol
• C–H: 413 kJ/mol
• C–C: 346 kJ/mol
• Cl–Cl: 242 kJ/mol
• C=O: 805 kJ/mol

Energy Changes in Reactions

• Bond Breaking: Endothermic (energy required).
• Bond Formation: Exothermic (energy released).

Calculating Overall Energy Changes

Formula:

ΔH=Energy required to break bonds−Energy given out when bonds are made

Example Calculation: Methane combustion

• Reaction: CH₄ (g) + 2O₂ (g) → CO₂ (g) + 2H₂O (l)

Calculations:

1. Energy to Break Bonds:
   • 4 C–H: 4×413 kJ=1652 kJ
   • 2 O=O: 2×498 kJ=996 kJ
   • Total: 1652+996=2648 kJ
2. Energy Released in Forming Bonds:
   • 2 C=O in CO₂: 2×805 kJ=1610 kJ
   • 4 O–H in H₂O: 4×464 kJ=1856 kJ
   • Total: 1610+1856=3466 kJ
3. Overall Energy Change:

ΔH=2648−3466=−818 kJ

Conclusion: The reaction is exothermic (negative ∆H).