The Chemistry of Solvents and Solutions

Solutions

  • Definition: A solution is a homogeneous mixture of substances.

  • Components:

    • Solvent: The component present in the greatest amount.

    • Solute: All other components, which may include more than one substance.

  • Interactions: The ability of a substance to dissolve in a particular solvent is determined by the interactions between solvent molecules and solute molecules, which relate to solubility and intermolecular forces.

Types of Solutions

  • States: Solutions can exist in all three physical states: solid, liquid, and gas.

  • Examples of Solutions:

    • Gas in Gas: Air

    • Gas in Liquid: Carbonated drinks

    • Gas in Solid: Hydrogen gas in palladium (H2(g) in Pd(s))

    • Liquid in Liquid: Motor oil, vinegar

    • Solid in Liquid: Ocean water, sugar-water

    • Solid in Solid: Bronze, pewter, 14K gold

Solubility and Intermolecular Forces

  • Like Dissolves Like:

    • Similar Polarity: Polar solutes dissolve in polar solvents, and non-polar solutes dissolve in non-polar solvents.

    • Dissimilar Polarity: Polar and non-polar substances are generally insoluble in one another.

Solute-Solvent Interactions

  • Substances dissolve when the attraction between solvent and solute molecules is stronger than the attractions between the solvent-solvent and solute-solute molecules.

  • The energy changes involved can be considered in terms of breaking existing interactions and forming new ones, shifting the balance towards solubility.

Miscibility

  • Miscible Liquids: Liquids that can dissolve in all proportions (e.g., ethanol and water).

  • Immiscible Liquids: Liquids that do not mix and form distinct phases (e.g., gasoline and water).

Effect of Chain Length in Alcohols on Solubility

  • As the size of alcohol molecules increases, their solubility in water decreases because:

    • The solute-solute attraction (between larger alcohol molecules) increases.

    • The solute-solvent attraction remains approximately constant.

  • Solubility Data for Alcohols:

    • Methanol (CH3OH): Miscible

    • Ethanol (C2H5OH): Miscible

    • 1-Propanol (C3H7OH): Miscible

    • 1-Butanol (C4H9OH): 7.9g/100g H2O @20°C

    • 1-Pentanol (C5H11OH): 2.7g/100g H2O @20°C

    • 1-Hexanol (C6H13OH): 0.6g/100g H2O @20°C

Enthalpy, Entropy, and Dissolving Solutes

  • Enthalpy of Solution ( \Delta H_{soln}): Can be endothermic or exothermic depending on the interactions between solute and solvent.

  • Endothermic Process: More energy required to separate solute, leading to an increase in temperature.

  • Exothermic Process: Energy released when solute and solvent interact, leading to a decrease in temperature.

  • Entropy (S): When solute and solvent molecules mix, they spread out, resulting in increased entropy, which drives the solution formation.

  • Example: NH4NO3 + H2O leads to a positive enthalpy change but greater than zero entropy change, suggesting it will dissolve.

Types of Solutions Based on Solute Concentration

  • Unsaturated Solutions: Can still dissolve more solute.

  • Saturated Solutions: No more solute can dissolve; undissolved solids may be present.

  • Supersaturated Solutions: Contain more solute than what can normally be dissolved; often created under high-temperature conditions and then gradually cooled down.

Solubility & Equilibrium Dynamics

  • At saturation, the solute concentration remains constant while the solute is in dynamic equilibrium with the solvent. Solute molecules constantly enter and exit from the solution.

  • Supersaturation can occur when a solution is cooled slowly, allowing for temporary solute retention in solution before crystallization occurs.

Dissolving Ionic Solids in Liquids

  • When an ionic compound dissolves in water:

    • Ions overcome lattice forces and become hydrated (surrounded by water molecules).

Colligative Properties

  • Colligative properties are changes that depend on the number of solute particles but not their identity:

    • Vapor Pressure Lowering: The vapor pressure of a solvent is lowered upon addition of a non-volatile solute.

    • Boiling Point Elevation: Non-volatile solutes raise the boiling point of the solvent.

    • Freezing Point Depression: Non-volatile solutes lower the freezing point.

  • Example: Addition of solute urea lowers the vapor pressure of water, which is reflected in Raoult's law.

Osmotic Pressure

  • Definition: The pressure required to stop osmosis, the net movement of solvent through a semipermeable membrane from a dilute to a more concentrated solution.

  • Relationships similar to the ideal gas law can be applied for calculations of osmotic pressure.