Chapter 14 – Solutions (Principles of Chemistry: A Molecular Approach)

Seawater, Dehydration & the Driving Force for Mixing

• Drinking seawater leads to dehydration and diarrhea.

  • Seawater contains higher salt concentration than human cells.

  • Cell membranes block ion (solute) movement ➔ water must leave the cells to equalize concentration.

  • Flow of water out of cells into the gut arises from nature’s tendency toward spontaneous, uniform mixing (entropy-driven).

Fundamental Vocabulary

• Solution = homogeneous mixture of ≥2 substances.

  • Solvent = majority component.

  • Solute = minority component.

• In aqueous solutions water is always the solvent.

• Solubility = maximum amount of solute that dissolves in given solvent quantity at specific T,P.

• Miscible / Immiscible (liquids) – mutually soluble / insoluble.

• Homogeneous vs. Heterogeneous mixtures

  • Air, seawater, brass = homogeneous (solutions).

  • Suspensions, colloids = heterogeneous (later).

Spontaneous Mixing & Entropy

• Ideal gases mix even without enthalpy change (negligible IMFs) because of entropy increase.

  • Entropy (S) = measure of energy dispersal; systems favor greater number of microstates.

• Conceptual Q14.1 answer: Ideal gases mix because mixing increases entropy.

Common Types of Solutions (Phase Combinations)

• Gas in gas (air), gas in liquid (club soda), liquid in liquid (vodka), solid in liquid (seawater), solid in solid (brass).

Intermolecular Forces (IMFs) & Solution Formation

• Types: dispersion, dipole–dipole, hydrogen bond, ion–dipole.

• Relative strength considerations:

  • \text{If }\,\text{Solvent–solute} > \text{Solvent–solvent & Solute–solute} \Rightarrow \text{solution forms readily.}

  • $=\Rightarrow$ still forms.

  • $<$ may or may not form; must weigh against entropy gain.

• Rule of thumb: Like dissolves like (polarity compatibility).

Energetics: Enthalpy of Solution (ΔH_soln)

• Three conceptual steps:

  1. Separate solute particles $\Delta H_{solute}>0$ (endothermic; = –lattice energy for ionic solids).

  2. Separate solvent particles $\Delta H_{solvent}>0$ (endothermic; overcome H-bonds, etc.).

  3. Mix solute + solvent $\Delta H_{mix}<0$ (exothermic; new attractions form).

• 
  

  • Exothermic if step-3 release exceeds steps-1&2 costs.

• Heat of Hydration (ΔHhydration) (ionic solutes in water):

  • Lattice energy always exothermic (negative), but sign is reversed in step-1, so cost.

  • Very negative ΔH_hydration from strong ion–dipole.

  • Example KF:

• Conceptual Q14.4 (CsF): ΔHsoln = –36.8 kJ/mol ⇒ |ΔHhydration| > |ΔH_solute|.

Solution Equilibrium & Saturation

• Dynamic equilibrium when dissolution rate = recrystallization rate.

• Definitions:

  • Saturated – at equilibrium; extra solute won’t dissolve.

  • Unsaturated – below equilibrium; more solute can dissolve.

  • Supersaturated – above equilibrium; unstable, precipitates when disturbed.

  • Making supersaturated solutions: dissolve at non-room conditions then cool slowly.

Temperature Dependence of Solubility

• Solids in water

  • Most endothermic dissolutions: solubility ↑ with T (see solubility curves: KNO₃ vs NaCl etc.).

  • Purification by recrystallization exploits this.

• Gases in water

  • ΔH_solution exothermic ⇒ solubility ↓ with ↑T (warm soda fizzes).

  • Conceptual Q14.5: cooling saturated mix ➔ gas bubbles out, salt precipitates.

Pressure Dependence for Gases – Henry’s Law

• $S<em>{gas}=k</em>H P<em>{gas}$ (S in molarity; $P</em>{gas}$ in atm).

  • Larger partial pressure ⇒ higher solubility.

  • Table constants (25 °C): $k_H$(NH₃)=5.8 M atm⁻¹ (largest) due to polarity.

• Conceptual Q14.6: NH₃ constant largest because NH₃ is polar while others are nonpolar.

Concentration Units & Conversions

• Molarity (M) = mol solute / L solution. (Temp-dependent).

• Molality (m) = mol solute / kg solvent. (Temp-independent).

• Mole fraction (χ) = mol component / total mol; mole % = χ×100.

• Mass or volume percentage

  • % w/w, % v/v, ppm (×10⁶), ppb (×10⁹).

• Unit conversions: treat definitions as conversion factors; convert numerator & denominator separately.

• Conceptual Q14.7: 25 g solute (25 g mol⁻¹) in 100 g solvent ⇒ 10 m.

Colligative Properties Overview

• Depend solely on quantity of solute particles, not their identity.

  1. Vapor pressure lowering.

  2. Boiling point elevation.

  3. Freezing point depression.

  4. Osmotic pressure.

• Electrolytes vs nonelectrolytes: account for dissociation via van’t Hoff factor (i).

Vapor Pressure Lowering & Raoult’s Law

• For nonvolatile solute:
  No ideal

• For volatile solutions, total:
  Ideal

• Deviations:

  • Negative (strong A–B IMFs): P_tot < Raoult prediction.

  • Positive (weak A–B): P_tot > prediction.

• Conceptual Q14.8: χsolute =0.200 ⇒ Psoln = 80 torr (but slide mis-key showed 20 torr; correct calc uses χ_solvent=0.8).

• Conceptual Q14.9: Measured 120 mmHg < ideal 150 mmHg ⇒ solute–solvent > like-like (negative deviation).

Boiling Point Elevation & Freezing Point Depression

• Equations (use molality of particles):
  $\Delta T<em>b = i K</em>b m$ (positive; raises Tb). $\Delta T</em>f = i K<em>f m$ (positive; lowers Tf). Remember sign.

  Conceptual Q14.10: 1 m in ethanol shows larger ΔT_b vs water.

Osmosis & Osmotic Pressure

• Osmosis: solvent flows from low [solute] → high [solute] through semipermeable membrane.

• Osmotic pressure formula:
  $\Pi = i M R T$
  (M = molarity of particles).

• Medical relevance:

  • Isosmotic: equal Π ⇒ cells unchanged.

  • Hyperosmotic: external Π > internal ⇒ cell shrivels.

  • Hyposmotic: external Π < internal ⇒ cell swells.

Van’t Hoff Factor (i)

• Theoretical: NaCl → 2, MgCl₂ → 3, FeCl₃ → 4.

• Measured lower due to ion pairing; table provided (e.g., MgSO₄ expected 2, measured 1.3).

• Conceptual Q14.11: Highest ΔT_b comes from solution with greatest i·m ⇒ 0.50 M MgCl₂ (i≈3).

Colloids, Suspensions & the Tyndall Effect

• Solutions: <1 nm particles; transparent; pass membranes.

• Colloids: 1–1000 nm; heterogeneous but stable; show Tyndall scattering & Brownian motion.

  • Hydrophilic colloids stabilized by solvation.

  • Hydrophobic colloids stabilized by surface charge repulsion (can coagulate if charge neutralized).

  • Classification matrix (aerosol, foam, emulsion, solid sol, etc.).

• Suspensions: >1000 nm; settle out.

Soap, Micelles & Emulsification

• Soap = sodium/potassium salt of fatty acid.

  • Structure: hydrophilic ionic head + hydrophobic hydrocarbon tail.

  • In water, form micelles: tails inward, heads outward.

  • Micelles solubilize grease/oil; charged heads repel ➔ stable dispersion.

• Micelle repulsions influenced by ionic strength; reduce repulsion ➔ coagulation.

Wrap-Up & Connections

• Solution behavior combines thermodynamics (ΔH, ΔS) with IMFs to dictate whether, and how much, solute dissolves.

• Colligative properties provide powerful experimental tools to infer molar mass & electrolyte behavior.

• Real-world links: seawater dehydration, antifreeze, IV fluids, carbonated beverages, environmental ppm/ppb limits, detergents.