Chapter 17: Solutions

Chapter 17: Solutions

Recommended Textbook Problems

• Solution Composition: 13, 15, 17, 19, 21
• Thermodynamics of Solutions and Solubility: 25, 27, 33, 35, 37, 39
• Vapor Pressures of Solutions: 43, 45, 47, 49, 53
• Colligative Properties: 59, 61, 63, 65, 69, 71
• Properties of Electrolyte Solutions: 79, 81, 85, 87

Solution: Definition and Examples (17.1)

• Solution: Homogeneous mixture of two or more substances.
• Solvent: Substance in excess.
• Solute: Substance not in excess.
• States of matter in solutions:
  • Liquid solvent:
    • Liquid solute: Ethanol in water.
    • Solid solute: Seawater.
    • Gas solute: Carbonated $H_2O$.
  • Gas solvent:
    • Gas solute: Air ($N<em>2$, $O</em>2$, $H_2O$…).
  • Solid solvent:
    • Solid solute: Brass (Cu/Zn alloy), Pt/H2 catalyst

Electrolytic Solutions

• An electrolyte "ionizes" in $H_2O$ (dissociates into ions).
• Ionic compounds are electrolytes.
• Electrolytic cell:
  • Electrons are carriers.
  • Oxidation occurs at the (+) electrode (anode).
  • Reduction occurs at the (-) electrode (cathode).
• Strong Electrolyte:
  • Completely ionizes in solution.
  • Example: $NaCl<em>{(s)} \rightarrow Na^+</em>{(aq)} + Cl^-_{(aq)}$
• Weak Electrolyte:
  • Partially ionizes in solution.
  • Example: $CH<em>3COOH</em>{(aq)} \rightleftharpoons CH<em>3COO^-</em>{(aq)} + H^+_{(aq)}$
• Non-Electrolyte:
  • Does not ionize.
  • Example: Sucrose

Not Everything Forms A Solution (17.2)

• Colloid:
  • A homogeneous mixture formed by "ultramicroscopic" particles in a dispersing medium.
• Miscible:
  • If two substances form a solution, they are miscible
• Immiscible:
  • If two substance do not form a solution, they are immiscible
• To form a solution, solvent-solute IMFs must be stronger than solvent-solvent + solute-solute IMFs.
• "Like dissolves like".

Examples of Colloids (17.8)

Solvent-Solute Interactions

• "Like dissolves like"
• Solvent-solute IMFs must be stronger than Solvent-Solvent + Solute-Solute IMFs
• Matching IMFs close in strength.
• Example: Dissolution of $NaCl$ in water
  • $NaCl<em>{(s)} \rightarrow Na^+</em>{(aq)} + Cl^-_{(aq)}$ (Ion-Dipole interactions)
• $NaCl_{(s)}$ will not dissolve in oil.
  • $NaCl$ must "give up" +/- interaction to dissolve

Review: Ways to Define Solution Concentration (17.1)

• Molarity (M):
  • $M = \frac{mols \; A}{L \; sol'n}$
  • $[H<em>2O] = 55.5 \; M$ (pure $H</em>2O$)
  • Useful when volume changes with temperature.
• Molality (m):
  • $m = \frac{mols \; A}{kg \; solvent}$
• Mass %:
  • $mass \;$X_A):
    • XA = \frac{nA}{n_{total}}

  Solubility of Gases vs Solids (17.3)

  • Solids dissolved in liquids:
    • T \uparrow$, Solubility$\uparrow
    • Why? -T\Delta S term
    • \Delta G = \Delta H - T\Delta S
    • Related to IMF strength
    • \Delta S \approx 0 for solid & solution
    • \Delta G more negative, more spontaneous at a higher T
  • Solubility of gases predicted by Henry's Law:
    • C = kH \cdot PA
      • C = concentration of gas A
      • k_H = Henry's Law constant for gas A
      • P_A = Partial pressure of gas A
  • Gases dissolved in liquids:
    • T \uparrow$, Solubility$\downarrow
    • Reason: \Delta G \uparrow$as$T \uparrow
    • Q=K

  Henry's Law

  • The Henry's law constant for CO2$is$3.4 \times 10^{-2} \; M/atm$at 25 °C. Assuming ideal solution behavior, what pressure of carbon dioxide is needed to maintain a$CO2 concentration of 0.10 M in a can of lemon-lime soda?
    • C = kH \cdot P{CO_2}
    • P{CO2} = \frac{C}{k_H} = \frac{0.10 \; M}{3.4 \times 10^{-2} \; M/atm} = 2.9 \; atm$$