Chapter 13 The Properties of Mixtures: Solutions and Colloids

Chapter 13 The Properties of Mixtures: Solutions and Colloids

Definitions

• Solution: A homogeneous mixture where the composition is uniform down to the molecular level.

  • A mixture is a physical intermingling of two or more substances with variable composition.

  • Solvent: The major component in a solution, usually present in the greatest amount (in terms of moles).

  • Solute: The minor component(s) present in a solution.

Types of Solutions

• There are six commonly encountered types of solutions based on phase combinations:

  • Gas in gas: Air (e.g., O₂ in N₂)

  • Liquid in gas: Soda water (e.g., CO₂ in water)

  • Liquid in liquid: Ethanol in water

  • Solid in liquid: NaCl in water

  • Solid in gas: H₂ in Pd or Pt

  • Solid in solid: Brass (alloys)

Intermolecular Forces

• Intermolecular forces determine the physical states of substances (e.g., Cl₂ is gas at 25°C, Br₂ is liquid, I₂ is solid).

• To form a solution, substances with distinct intermolecular forces must mix uniformly:

  • Particles of each component will interact, changing their surroundings to include particles of the other component, leading to energy changes.

Energy Changes in Solution Formation

• Formation involves a series of steps, each with associated energy changes:

  • Separation of solute particles: Endothermic (∆H_solute > 0)

  • Separation of solvent particles: Endothermic (∆H_solvent > 0)

  • Solute and solvent particles intermingle: Exothermic (∆H_mix < 0)

Favorability for Mixing

• Mixing leads to increased entropy, a thermodynamic quantity that favors spontaneous processes.

• If a dissolution process is exothermic, it usually occurs naturally because energy leaves the system (increased dispersal).

• Endothermic processes can occur if the increase in material dispersal compensates for energy loss.

Generalization for Mixing

• "Like dissolves like" guideline:

  • Polar solvents dissolve polar solutes.

  • Nonpolar solvents dissolve nonpolar solutes.

  • Polar and nonpolar substances generally do not mix, although exceptions exist.

Immiscibility and Solubility

• Miscible liquids: Two liquids that mix in any proportion (e.g., CCl₄ and C₆H₆).

• Immiscible liquids: Liquids that do not mix readily (e.g., CCl₄ and H₂O).

• Dissolution Process: Typically, solid and ionic solutes dissolve in liquid solvents through solvation (hydration in the case of water).

• Solubility limits:

  • Saturated solution: Contains maximum solute at a given temperature.

  • Unsaturated solution: Contains less than the maximum solute at a given temperature.

  • Supersaturated solution: Contains more solute than can be sustained at equilibrium; typically unstable.

Solubility Factors

• Temperature effects:

  • Solubility of solids generally increases with temperature, but this is not universal (e.g., NaCl is relatively unaffected).

  • Fractional crystallization can separate components based on varying temperature dependencies.

• Gases in liquid solubility decreases with increased temperature (e.g., carbonated beverage).

Pressure Effects on Solubility

• Pressure has negligible effects on solids and liquids.

• Gas solubility is directly proportional to the partial pressure of the gas above the solution: $c \propto P$

  • Henry's Law: $c = kP$

  • At a fixed temperature, $S<em>{gas} = k</em>H P<em>{gas}$ where $kH$ is the Henry's law constant.

Concentration Terms

• Solutions have variable composition as opposed to compounds.

• Concentration types include:

  • Mass percent:
    $ext{Mass percent} = \frac{m<em>{solute}}{m</em>{solution}} \times 100 = \frac{m<em>{solute}}{m</em>{solute} + m_{solvent}} \times 100$

  • Parts per million (ppm):
    $\text{ppm} = \frac{m<em>{solute}}{m</em>{solution}} \times 10^6$

  • Mole fraction (X): Useful in calculating partial pressures.

  • Molality (m): $\text{Molality} = \frac{n<em>{solute}}{m</em>{solvent} \text{ (in kg)}}$

Molarity

• $\text{Molarity (M)} = \frac{n<em>{solute}}{V</em>{solution} \text{ (in L)}}$

  • Laboratory commonly uses molarity due to volume measuring convenience but is temperature dependent.

Colligative Properties of Solutions

• Colligative properties depend on solute particle concentration rather than characteristics of the solute:

  • Vapor pressure lowering

  • Boiling point elevation

  • Freezing point depression

  • Osmotic pressure

Vapor Pressure Lowering

• Addition of solute to a solvent lowers vapor pressure:
  $P<em>{solvent} = X</em>{solvent} P^0_{solvent}$

Boiling Point Elevation and Freezing Point Depression

• Boiling point elevation: $\Delta T<em>b = K</em>b m$

  • $\Delta T<em>b = T</em>b - T^0_b$

• Freezing point depression:
  $\Delta T<em>f = -K</em>f m$

Osmotic Pressure

• Defined as the pressure needed to stop osmosis (flow of solvent across a semipermeable membrane):
  $\pi = MRT$

Colligative Properties of Electrolyte Solutions

• Electrolytes dissociate into ions influencing the colligative properties.

• Incorporate the van’t Hoff factor $i$:

  • $\Delta T<em>b = i K</em>b m$

  • $\Delta T<em>f = -i K</em>f m$

  • $\pi = i MRT$

Structure and Properties of Colloids

• Colloids are stable, heterogeneous mixtures where colloidal particles (10 Å to 10000 Å) remain dispersed due to energies involved (Brownian motion).

• Colloids scatter light, demonstrating Tyndall effect - a visible light beam differentiates colloids from true solutions.

Types of Colloids

• Typically categorized based on dispersion medium, with water being a prominent solvent:

  • Hydrophilic: Polar parti0cles mixing well in water.

  • Hydrophobic: Non-polar particles resisting water dispersion.

Colloid Stability and Flocculation

• Stabilizing hydrophobic suspensions involves adsorption of ions onto particle surfaces (e.g., surfactants).

• Coagulation/destabilization occurs through temperature changes or ion concentration changes leading to flocculation.