Solutions & Solubility Concepts

Terminology

• Solution: A homogeneous mixture of one or more solutes uniformly distributed in a solvent; examples include air, which consists of nitrogen, oxygen, and trace gases, showcasing that solutions aren't limited to liquids.
• Solute: The dissolved material in a solution, typically present in lesser amounts.
• Solvent: The dissolving medium in a solution, generally present in larger amounts.
• Molar concentration: Measured in moles per liter of solution.
• Molal concentration: Measured in moles per kilogram of solvent, not the total solution.
• Percent by weight to volume: Represents grams of solute in 100 mL of solution.
• Percent by weight to weight: Represents grams of solute in 100 grams of solution.
• Normality: Equivalent concentration per liter; usage has declined in chemistry and medicine.
• Parts per million (ppm): Concentration measurement indicating solute amount in one million parts of solution.
• Miscible liquids: Liquids that can mix in any proportion, e.g., water and ethyl alcohol.
• Immiscible liquids: Liquids that do not mix well, e.g., water and oil.
• Colloids: Mixtures where one phase is dispersed in another phase, not fully dissolved.

Solubility

• Solubility: Refers to the maximum amount of solute dissolvable in a given solvent at a specific temperature.
• Saturated solution: Contains the maximum solute possible under defined conditions.
• Supersaturated solution: Contains more solute than a saturated solution, unstable as excess solute can crystallize out.
• Polarity principle: Polar solutes are more soluble in polar solvents and nonpolar solutes in nonpolar solvents.

Energy Changes & the Solution Process

• Heat of solution: Energy change occurring when a solute dissolves; it’s the enthalpy change for one mole of solute.
• Types of enthalpy of solution:
  • Exothermic: Energy released during solvation exceeds energy needed for separation.
  • Endothermic: Energy required for separation surpasses energy released during solvation.
• Ion interaction: Dissolution involves separating ionic compounds and their attraction to solvent molecules influencing overall energy balance.

Factors Affecting Solubility

• Pressure impact: Increases in pressure enhance the solubility of gaseous solutes, described by Henry’s Law: (S = KH \times P_{gas}).
• Temperature effect: Increased temperature generally raises solubility for solids and liquids but decreases it for gases in liquid solvents.

Colligative Properties of Solutions

• Colligative properties: Physical property changes resulting from solute addition include vapor pressure, boiling point, freezing point, and osmotic pressure.
• Vapor pressure: Decreases when solute molecules occupy surfaces, quantified by Raoult's Law.
• Boiling point elevation: Higher energy is required to raise vapor pressure due to solute introduction.
• Freezing point depression: Solute presence disrupts solid crystallization, lowering the freezing point of the solution.
• Osmotic pressure: Directly proportional to the number of solute particles present in a solution.

Solubility Product

• Solubility product constant (Ksp): Describes the equilibrium concentration of ions in saturated solutions of sparingly soluble ionic compounds.
• Ksp formula: (Ksp = [M^+]^x [A^-]^y), where M is cation concentration, A is anion concentration, and x & y represent their stoichiometric coefficients.
• Example: For (PbCl_2), the Ksp expression reflects its ionic dissociation: (Ksp = [Pb^{2+}][Cl^-]^2).
• Ksp values: Lower values indicate lower solubility of the substance in a solvent.