🔬

Heats of Reactions

Understanding Energy Changes in Chemical Processes

4 Comprehensive Periods of Study

🌡️

Endothermic and Exothermic Reactions

Periods 1-2: Understanding Energy Flow

🔥

Exothermic Reactions

Energy is RELEASED to surroundings

  • Temperature increases
  • Products have lower energy than reactants
  • ΔH is negative (ΔH < 0)
  • Energy flows OUT of the system

Common Examples:

CH₄ + 2O₂ → CO₂ + 2H₂O + ENERGY

Combustion of methane

2H₂ + O₂ → 2H₂O + 286 kJ/mol

Formation of water

C + O₂ → CO₂ + 393.5 kJ/mol

Carbon combustion

🧊

Endothermic Reactions

Energy is ABSORBED from surroundings

  • Temperature decreases
  • Products have higher energy than reactants
  • ΔH is positive (ΔH > 0)
  • Energy flows INTO the system

Common Examples:

CaCO₃ → CaO + CO₂ - 178 kJ/mol

Limestone decomposition

NH₄Cl(s) → NH₄⁺(aq) + Cl⁻(aq)

Dissolving ammonium chloride

H₂O(l) → H₂O(g) - 40.7 kJ/mol

Water evaporation

Key Differences Comparison

Aspect Exothermic Endothermic
Energy Change (ΔH) Negative (< 0) Positive (> 0)
Temperature Effect Increases Decreases
Energy Level Products < Reactants Products > Reactants
Spontaneity Usually spontaneous Usually non-spontaneous
🌡️

Temperature Change in Chemical Reactions

Period 3: Measuring and Understanding Heat Changes

🧪 Calorimetry

Science of measuring heat changes in chemical reactions

Heat = m × c × ΔT

m = mass of solution (g)

c = specific heat capacity (J/g°C)

ΔT = temperature change (°C)

🥤 Simple Calorimeter

  • Polystyrene cup
  • Thermometer
  • Stirrer
  • Insulation material
  • Lid to minimize heat loss

Used for solution-based reactions

💣 Bomb Calorimeter

  • Steel bomb chamber
  • Water bath
  • Stirrer
  • Thermometer
  • Ignition system

Used for combustion reactions

Calculation Examples

Example 1: Neutralization

HCl + NaOH → NaCl + H₂O

50 mL of 1M HCl + 50 mL of 1M NaOH

Initial temperature: 20°C

Final temperature: 27°C

Solution:

Mass = 100g (assuming density = 1g/mL)

ΔT = 27 - 20 = 7°C

Heat = 100g × 4.18 J/g°C × 7°C = 2,926 J

Molar enthalpy = -58.5 kJ/mol

Example 2: Dissolution

NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq)

5g of NH₄NO₃ in 100 mL water

Initial temperature: 25°C

Final temperature: 21°C

Solution:

Mass = 105g (salt + water)

ΔT = 21 - 25 = -4°C

Heat = 105g × 4.18 J/g°C × (-4°C) = -1,756 J

Molar enthalpy = +28.1 kJ/mol

📊

Energy Level Diagrams

Periods 4-5: Visualizing Energy Changes

Exothermic Energy Diagram

Reaction Progress Energy Reactants Products Ea ΔH < 0

• Products are at lower energy than reactants

• Energy is released (negative ΔH)

• Activation energy still required to start reaction

Endothermic Energy Diagram

Reaction Progress Energy Reactants Products Ea ΔH > 0

• Products are at higher energy than reactants

• Energy is absorbed (positive ΔH)

• Continuous energy input often required

Key Components of Energy Diagrams

Activation Energy (Ea)

  • • Minimum energy needed to start reaction
  • • Height of energy barrier
  • • Always positive value
  • • Can be lowered by catalysts

Enthalpy Change (ΔH)

  • • Energy difference: Products - Reactants
  • • Negative for exothermic
  • • Positive for endothermic
  • • Independent of pathway

Transition State

  • • Highest energy point
  • • Unstable intermediate
  • • Cannot be isolated
  • • Determines reaction rate
🎯

Summary: Periods 1-4

You've covered the fundamental concepts of heat in chemical reactions, from basic endothermic and exothermic processes through temperature measurement and energy visualization. These concepts form the foundation for understanding chemical energetics.