Properties of Mixtures: Solutions and Colloids

Properties of Mixtures: Solutions and Colloids

Solution Process

  • Solution Formation Principles:

    • Depends on the strengths of intermolecular forces.
    • A solution forms if the energy change is favorable.

  • Three-Step Process (Hess's Law):

    1. Separation of Solute Particles:
    • Enthalpy change: ( \Delta H_{solute} ) (endothermic)
    1. Separation of Solvent Particles:
    • Enthalpy change: ( \Delta H_{solvent} ) (endothermic)
    1. Mixing Solute and Solvent:
    • Enthalpy change: ( \Delta H_{mix} ) (exothermic)

  • Overall Enthalpy Change:
    ( \Delta H{SOLN} = \Delta H{solute} + \Delta H{solvent} + \Delta H{mix} )

Cases of Solution Formation

  • Case #1: Ionic Solids in Water

    • Enthalpy considerations greatly influence solubility.

  • Case #2: Covalent Compounds Mixtures

    • Entropy plays a key role; solutions may form even if ( \Delta H_{SOLN} \approx 0 ).

  • Case #3: No Solution Formation

    • Ionic compounds do not dissolve in nonpolar solvents.
    • Insolubility arises when ( \Delta H{solute} ) is significantly larger than ( \Delta H{mix} ).
    • Rule of Thumb: LIKE dissolves LIKE.
Factors Affecting Solubility

  1. Structure Effects

    • Molecular structure impacts solubility.

  2. Pressure Effects (Henry’s Law)

    • ( S{gas} = kH P_{gas} )
    • Gases are more soluble in liquids at higher pressures.

  3. Temperature Effects

    • Increasing temperature generally increases the rate of dissolving:
      • Gases: More soluble at lower temperatures.
      • Solids: Solubility changes depending on ( \Delta H_{SOLN} ).
Non-Liquid Solutions
  • Gas-Gas Solutions: e.g., air (infinitely soluble in other gases).
  • Solid-Solid Solutions: e.g., alloys.
    • Types:
    • Substitutional alloy
    • Interstitial alloy
Concentration Measurements
  • Molarity and Molality Calculations:
    • Example: What is the molarity of a 28% by mass solution of ammonia with a density of 0.90 g•mL-1?
    • Example: What is the molality of a 15% by mass solution of sodium chloride (NaCl: 58.44 g•mol-1)?
Colligative Properties of Solutions

  • Definition:

    • Depends on the concentration of solute, not its identity.

  • Assumptions:

    • Solute particles are nonvolatile.
    • Ideal behavior in dilute solutions.

  • Raoult’s Law:
    ( P{solvent} = X{solvent} P_{solvent}^\circ )

  • Vapor Pressure Lowering:
    ( \Delta P = X{solute} P{solvent}^\circ )

    • Higher entropy in solutions leads to less solvent evaporation.
    • Results in lower freezing points and higher boiling points than pure solvents:
    • ( \Delta TB = TB - T_{B}^{\circ} )
    • ( \Delta TF = TF - T_{F}^{\circ} )

  • Freezing Point Depression and Boiling Point Elevation:
    ( \Delta TB = kB m )
    ( \Delta TF = -kF m )

    • Constants (( kB ), ( kF )) are solvent-specific.
Electrolytes as Solutes
  • van't Hoff Factor (i):
    • Adjusts colligative property values based on dissociation into ions.
    • General formulas:
    • ( \Delta P = i X{solute} P{solvent}^\circ )
    • ( \Delta TB = i kB m )
    • ( \Delta TF = -i kF m )
    • For NaCl, ideally ( i = 2 ).
Colloids and Suspensions

  • Suspensions:

    • Heterogeneous mixtures that settle out.
    • Particles visible to the naked eye.

  • Colloids:

    • Dispersed substance larger than simple molecules but small enough not to settle out (sizes 1–1000 nm).
    • Are heterogeneous and scatter light beams.

  • Solutions:

    • Homogeneous mixtures that do not separate and are visually uniform.