Electrolytes and Colligative Properties

Electrolytes

  • Definition - General: Compounds that ionize in any liquid solution or melt.

    • Specific: Compounds that dissociate into ions in aqueous solutions (dissolved in water).

Types of Electrolytes

  • Ionic Electrolytes:

    • Examples: NaCl, MgCO3, Al2(SO4)3

    • Dissociate into ions when dissolved in water:

      • NaCl → Na⁺ + Cl⁻

      • Al2(SO4)3 → 2Al³⁺ + 3(SO4)²⁻

  • Molecular Electrolytes:

    • Examples: HCl, HBr, NH3, CO2

    • Dissociation can occur in two ways:

      • Direct ionization: HCl → H⁺ + Cl⁻

      • Reaction with water: NH3 + H2O → NH4⁺ + OH⁻

Classification of Electrolytes

  • Strong Electrolytes: Fully dissociate into ions.

    • Examples: Inorganic salts, strong acids/bases

  • Weak Electrolytes: Only partially dissociate into ions; many solute molecules remain undissociated.

    • Examples: CH3COOH (acetic acid), NH4OH (ammonium hydroxide)

  • Non-electrolytes: Do not produce ions in solution; consist of whole molecules only.

    • Examples: Sugars (like glucose)

Conductivity of Electrolyte Solutions

  • Electrolyte solution contains ions: The movement of ions allows the solution to conduct electricity.

  • Stronger electrolytes produce a higher current due to complete dissociation of ions in solution.

Colligative Properties

  • Dependent on the number of solute particles in a solution, not the identity of the solute.

  • Key examples:

    1. Vapour Pressure Lowering

    2. Boiling Point Elevation

    3. Freezing Point Depression

    4. Osmosis

Vapour Pressure

  • Defined as the pressure exerted by a vapor in equilibrium with its liquid or solid phase.

  • High for volatile substances, low for non-volatile substances.

  • Effect of solute: Adding a non-volatile solute lowers the vapour pressure of a solvent due to fewer solvent molecules escaping into the vapor phase.

    • Example: If a solution contains 0.1 moles of a non-volatile solute (like glucose) and 0.9 moles of water, the mole fraction of water (Xsolvent) would be 0.9. If P°solvent is the vapour pressure of pure water at a specific temperature, say 23 mmHg, then according to Raoult's law:

      • Psolution = Xsolvent * P°_solvent = 0.9 * 23 mmHg = 20.7 mmHg.

  • This illustrates that the addition of the solute lowers the vapour pressure compared to pure water (23 mmHg).

Raoult's Law

  • Applies to solutions containing a non-volatile solute:

    • Psolution = Xsolvent * P⁰_solvent

Boiling Point Elevation

  • Boiling point of a solution is higher than that of the pure solvent.

  • Given by the formula:

    • ΔTb = i * Kb * m

    • Example: Dissolving 1 mole of NaCl in 1 kg of water results in:

      • ΔT_b = 2 * 0.513 °C/m * 1 = 1.026 °C.

  • Thus, the boiling point of water elevates from 100 °C to approximately 101.03 °C.

Freezing Point Depression

  • The freezing point of a solution is lower than that of the pure solvent.

  • Given by the formula:

    • ΔTf = i * Kf * m

    • Example: Adding 1 mole of glucose (C6H12O6) to 1 kg of water results in:

      • ΔT_f = 1 * 1.86 °C/m * 1 = 1.86 °C.

  • The freezing point of water drops from 0 °C to approximately -1.86 °C.

Important Constants for Water

  • K_b = 0.513 °C/m

  • K_f = 1.86 °C/m

Summary of Key Points

  • Electrolytes are essential for understanding solution chemistry, particularly in aqueous environments.

  • Strong electrolytes fully dissociate into ions, while weak electrolytes do not.

  • Colligative properties help predict changes in physical properties of solutions based on solute concentration rather than identity, critical for many scientific and industrial applications.