A-Level Chemistry: Thermodynamics - Enthalpy Changes

Introduction

• Focus on enthalpy changes related to thermodynamics.
• Topics include: Enthalpy of solution, lattice enthalpy, and Born-Haber cycles.
• Entropy and Gibbs free energy covered in a separate video.

Enthalpy Changes Overview

• Definition: Heat energy transferred at constant pressure during a chemical reaction.
• Units: Usually expressed in kJ/mol.
• Standard Conditions:
  • Temperature: 298 K
  • Pressure: 100 kPa (just under 1 atm)
  • Solution concentration: 1 mol/dm³
• Signs: Indicated by ΔH.
  • Standard reactions denoted with a circle: ΔH°.
  • Exothermic: Negative ΔH, temperature increases, chemical energy to thermal energy.
  • Endothermic: Positive ΔH, temperature decreases, thermal energy to chemical energy.

Lattice Enthalpy

• Relevance: Key for understanding ionic compounds.
• Structure: Ionic compounds form a 3D lattice (e.g., alternating positive and negative ions).
• Influences:
  • Boiling point
  • Solubility
  • Reactivity
• Types:
  • Lattice Formation Enthalpy: Exothermic, one mole of ionic solid is formed from gaseous ions.
  • Lattice Dissociation Enthalpy: Endothermic, one mole of ionic solid dissociates into gaseous ions.

Comparing Lattice Enthalpies

• Factors:
  • Size of ions
  • Charge magnitude
  • Charge Density: Concentration of charge on an ion.
• Examples:
  • NaCl vs. KCl: NaCl has a more exothermic lattice enthalpy due to smaller ion size and higher charge density.
  • MgCl₂ vs. NaCl: MgCl₂ more exothermic due to higher charge and smaller ionic radius.
  • MgO vs. MgCl₂: MgO has a more exothermic lattice enthalpy due to smaller, more charged O²⁻ ion.

Solubility and Lattice Enthalpy

• Dissolution: Involves breaking lattice (endothermic) and hydrating ions (exothermic).
• Hydration Enthalpy: Formation of attractions between water and ions, always exothermic.
• Factors Affecting Hydration:
  • Charge density of ions.
  • Number of water molecules around the ion.
• Enthalpy of Solution: Depends on balance between lattice dissociation and hydration enthalpies.
  • Exothermic if hydration enthalpies > lattice dissociation.
  • Endothermic if hydration enthalpies < lattice dissociation.

Hess's Law and Enthalpy Cycles

• Hess's Law: Enthalpy change is independent of the reaction path.
• H Cycles: Used to calculate enthalpy changes of solution.
  • Example: NaCl's enthalpy of solution.

Born-Haber Cycles

• Purpose: Calculate unknown enthalpy values.
• Structure:
  • Rectangular, with energy as Y-axis.
  • Arrows indicate exothermic (down) or endothermic (up) processes.
• Application: Calculate lattice enthalpy or other unknowns using multiple steps.

Calculating Lattice Enthalpy

• Theoretical Calculations: Use electric field strength equations.
• Experimental Calculations: Indirect methods through Born-Haber cycles.
• Experimental vs. Theoretical Values:
  • Variations due to ionic model assumptions (perfect spheres) and polarization.

Polarization and Covalent Character

• Polarization: Distortion of electron cloud, mainly affects negative ions.
• Covalent Character:
  • More polarization leads to more covalent character.
  • Example: AlCl₃ (high polarization) vs. KCl (low polarization).

Final Notes

• Importance of Language: Use "more exothermic" or "more negative" for clarity.
• Study Aid: Understanding the influence of lattice enthalpy on other properties is crucial.