Chemical Energetics

This note explains enthalpy change (ΔH) in chemical reactions, covering exothermic and endothermic reactions, energy profile diagrams, bond breaking and making, and how to calculate overall enthalpy change using bond energies.

1. Definition of Enthalpy Change
  • 🔹 Enthalpy Change (ΔH):

    Heat energy change at constant pressure during a chemical reaction. Indicates whether energy is released or absorbed when reactants form products.

  • 🔹 Exothermic Reactions:
    • 🔹 Description:

      Releases heat to surroundings; temperature of surroundings increases; ΔH negative (ΔH < 0).

    • 🔹 Examples:

      Combustion of fuels, neutralization of acids and bases.

  • 🔹 Endothermic Reactions:
    • 🔹 Description:

      Absorbs heat from surroundings; temperature of surroundings decreases; ΔH positive (ΔH > 0).

    • 🔹 Examples:

      Thermal decomposition of calcium carbonate, photosynthesis.

2. Energy Profile Diagrams
  • 🔹 Purpose:

    Graphically represent energy changes during reactions.

  • 🔹 Components:
    • 🔹 Reactants:

      Starting energy level before reaction.

    • 🔹 Products:

      Energy level after reaction.

    • 🔹 Activation Energy (Ea):

      Minimum energy required for reaction to occur (energy needed to break bonds).

    • 🔹 Overall ΔH:

      Difference between energy of products and reactants.

  • 🔹 Interpretation:
    • 🔹 Exothermic:

      Products at lower energy than reactants; ΔH negative; diagram slopes downward.

    • 🔹 Endothermic:

      Products at higher energy than reactants; ΔH positive; diagram slopes upward.

3. Bond Breaking and Bond Making
  • 🔹 Bond Breaking:

    Requires energy input; endothermic; energy absorbed to overcome atomic forces; e.g., H–Cl bonds in HCl.

  • 🔹 Bond Making:

    Releases energy; exothermic; new bonds release potential energy as heat; e.g., formation of H₂O from H and O atoms.

4. Overall Enthalpy Change (Bond Energies)
  • 🔹 Formula:

    ΔH = Total energy absorbed in breaking bonds – Total energy released in forming bonds

  • 🔹 Exothermic:

    More energy released in bond formation than absorbed in bond breaking; ΔH negative.

  • 🔹 Endothermic:

    More energy absorbed in bond breaking than released in bond formation; ΔH positive.

  • 🔹 Example Calculation:
    • 🔹 Reaction:

      H₂ + Cl₂ → 2HCl

    • 🔹 Bonds Broken:

      H–H (436 kJ/mol) + Cl–Cl (243 kJ/mol) = 679 kJ/mol

    • 🔹 Bonds Formed:

      2 × H–Cl (2 × 431 kJ/mol) = 862 kJ/mol

    • 🔹 ΔH:

      679 – 862 = –183 kJ/mol → Exothermic

  • 🔹 Exothermic neutralization: HCl + NaOH → NaCl + H₂O; ΔH < 0
  • 🔹 Endothermic decomposition: CaCO₃ → CaO + CO₂; ΔH > 0
  • 🔹 Bond enthalpy: ΔH = Σ(Bonds broken) - Σ(Bonds formed)

  • ⚠️ Bond breaking releases energy instead of requiring it.
  • ⚠️ Activation energy is the same as overall ΔH.
  • ⚠️ Exothermic reactions always produce visible heat or flames.
  • ⚠️ Endothermic reactions cannot happen spontaneously.

  • 👉 Memorize ΔH sign conventions: negative for exothermic, positive for endothermic.
  • 👉 Practice drawing and labeling energy profile diagrams accurately.
  • 👉 Use bond energies carefully for ΔH calculations.
  • 👉 Explain energy changes clearly in terms of bond breaking and making.
  • 👉 Remember to show activation energy on diagrams even if not asked for ΔH.

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