This lesson explores energy changes in chemical reactions, enthalpy concepts, Hess’s Law, and the fundamentals of thermodynamics, including entropy and Gibbs free energy.

1. Core Concepts (Short Notes)

7.1 Energy in Chemical Reactions

Chemical reactions involve breaking and forming bonds, which leads to energy changes:

• Exothermic reactions: Release heat (ΔH is negative).

• Endothermic reactions: Absorb heat (ΔH is positive).

Examples:

• Exothermic: Combustion, neutralization, respiration.

• Endothermic: Photosynthesis, thermal decomposition.

• 
  

7.2 Enthalpy (ΔH)

Enthalpy is the heat content of a system at constant pressure.

• ΔH = H(products) – H(reactants)

• Units: kJ mol⁻¹

  
  

7.3 Standard Enthalpy Change

Measured under standard conditions:

• 25°C (298 K)

• 1 atm pressure

• 1 mol dm⁻³ concentration

Includes:

• Standard Enthalpy of Formation (ΔHf°): Formation of 1 mole of compound from elements.

• Standard Enthalpy of Combustion (ΔHc°): Complete combustion of 1 mole of compound.

• Neutralization Enthalpy: Formation of 1 mole of water from acid + base.

  
  

7.4 Bond Enthalpies

Energy required to break one mole of bonds in gaseous molecules.

• Breaking bonds → endothermic.

• Forming bonds → exothermic.

Approximate calculation: ΔH = Σ(bonds broken) – Σ(bonds formed)

7.5 Hess’s Law

The total enthalpy change of a reaction is the same, regardless of the pathway.

• Useful when direct measurement of ΔH is impossible.

• Construct enthalpy cycles using known ΔH values.

  
  

7.6 Entropy (ΔS)

A measure of disorder or randomness.

• Gases have highest entropy.

• Solids have lowest entropy.

• Reactions increasing number of gas molecules → higher entropy.

  
  

7.7 Gibbs Free Energy (ΔG)

Determines whether a reaction is spontaneous. Formula: ΔG = ΔH – TΔS

Interpretation:

• ΔG < 0: spontaneous reaction.

• ΔG > 0: non-spontaneous.

• ΔG = 0: equilibrium.

2. Key Formulas to Memorize

Enthalpy Change

ΔH = H(products) – H(reactants)

Bond Enthalpy Equation

ΔH = Σ(bonds broken) – Σ(bonds formed)

Gibbs Free Energy

ΔG = ΔH – TΔS

Entropy Trends

ΔS > 0 when:

• Solids → liquids → gases

• Fewer → more moles of gas

3. Tips & Tricks for Exams

• Always identify whether ΔH is positive or negative.

• Bond enthalpy questions require careful counting of each bond.

• For Hess’s Law, reverse equations by changing the sign of ΔH.

• If a reaction absorbs heat, temperature increase makes it less favorable (Le Chatelier’s principle).

• Units matter: convert kJ to J in Gibbs energy calculations when needed.

• Look for entropy changes: more gas = higher entropy.

• A reaction may be spontaneous even if ΔH is positive, provided TΔS is large.

4. Important Points to Remember

• Enthalpy, entropy, and free energy are central to predicting reaction behavior.

• Hess’s Law is a powerful tool for indirect enthalpy calculations.

• Endothermic ≠ non-spontaneous; exothermic ≠ always spontaneous.

• Entropy naturally increases in spontaneous processes.

ΔG combines enthalpy and entropy to determine feasibility.