Chapter 5: Chemical Energetics

Chapter 5: Chemical Energetics

Energy Changes in Reactions

- Chemical reactions involve energy changes, either exothermic (release heat) or endothermic (absorb heat).

- Exothermic: Products have less energy than reactants (ΔH is negative).

- Endothermic: Products have more energy than reactants (ΔH is positive).

Enthalpy Change (ΔH)

- ΔH: Heat change at constant pressure.

- Measured in kilojoules per mole (kJ/mol).

- Standard conditions: 298 K, 1 atm, 1 mol quantities.

Types of Enthalpy Changes

- ΔHᶠ: Enthalpy of formation, of 1 mole of compound from elements.

- ΔHᶜ: Enthalpy of combustion, burning 1 mole of substance in oxygen.

- ΔHⁿᵉᵘ: Enthalpy of neutralisation, acid + base → water.

- ΔHᵣ: Enthalpy of reaction, general reaction enthalpy.

Measuring Enthalpy Changes

- Use calorimetry: measure temperature change in water or solution.

- Equation:  

  q = mcΔT

  - q: heat energy (J)  

  - m: mass of water/solution (g)  

  - c: specific heat capacity (4.18 J/g·K)  

  - ΔT: temperature change (K)

Hess’s Law

- Total enthalpy change is the same, regardless of reaction path.

- Use known ΔH values to calculate unknown ones by constructing energy cycles.

Bond Enthalpies

- Bond enthalpy: Energy needed to break one mole of a bond in the gaseous state.

- Average values used for estimates.

- ΔH ≈ Σ(bonds broken) – Σ(bonds formed)

Energy Profile Diagrams

- Show energy changes during a reaction.

- Activation energy (Ea): Minimum energy needed to start a reaction.

- Catalysts lower Ea, making reactions faster.