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Chapter 13

Chapter 13. The Properties of Mixtures: Solutions and Colloids. Terms Associated with Solutions. solute the material dissolved; the lesser amount solvent the material the solute is dissolved in the larger amount; usually water. Terms Associated with Solutions. miscible.

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Chapter 13

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  1. Chapter 13 The Properties of Mixtures: Solutions and Colloids

  2. Terms Associated with Solutions solute the material dissolved; the lesser amount solvent the material the solute is dissolved in the larger amount; usually water

  3. Terms Associated with Solutions • miscible

  4. Solubility (S) Solubility: # moles solute dissolved / L solution or… # grams solute dissolved / 100 ml water “Like Dissolves Like”

  5. Factors Affecting Solubility Given the structures of Vitamin A and Vitamin C, which vitamin must be consumed regularly? Which vitamin will be stored in the fatty tissues of the body? Vitamin C

  6. Figure 13.3 Like dissolves like: solubility of methanol in water

  7. Demos • Salting Out • Water, Iodine, Hexane, Ethanol

  8. Figure 13.1 Ion-Dipole Forces Explains solubility of ionic compounds in water Example: Any ion in water

  9. Figure 13.1 Hydrogen Bonding Chemistry is FON! Example: Organic compounds (sugars/alcohols/amino acids) in cellular fluid

  10. Dipole-Dipole Forces Solubility of polar molecules in each other Figure 13.1 Example: Polar organic molecules in polar non-aqueous solvents

  11. Ion-Induced Dipole An ion’s charge distorts the electron cloud of a nearby polar molecule Example: Fe2+ binding to hemoglobin Solubility of salts in ethanol

  12. Dipole-Induced Dipole Forces Polar molecule distorts electron cloud in nonpolar substance Example: Solubility of gases in water

  13. Dispersion Forces Contribute to the solubility of ALL solutes in ALL solvents The principal type of IMF in solutions of nonpolar substances Example: Petroleum and gasoline

  14. LIKE DISSOLVES LIKE Substances with similar types of intermolecular forces dissolve in each other. When a solute dissolves in a solvent, solute-solute interactions and solvent-solvent interactions are being replaced with solute-solvent interactions. The forces must be comparable in strength in order to have a solution occur.

  15. Liquid Solutions and the role of Molecular Polarity • Liquid – Liquid solutions When a solute dissolves in a solvent, solute-solute interactions and solvent-solvent interactions are being replaced with solute-solvent interactions. The forces must be comparable in strength in order to have a solution occur.

  16. PROBLEM: Predict which solvent will dissolve more of the given solute: (a) Methanol - NaCl is ionic and will form ion-dipoles with the -OH groups of both methanol and propanol. However, propanol is subject to the dispersion forces to a greater extent. (b) Water - Hexane has no dipoles to interact with the -OH groups in ethylene glycol. Water can H bond to the ethylene glycol. (c) Ethanol - Diethyl ether can interact through a dipole and dispersion forces. Ethanol can provide both while water would like to H bond. SAMPLE PROBLEM 13.1 Predicting Relative Solubilities of Substances (a) Sodium chloride in methanol (CH3OH) or in propanol (CH3CH2CH2OH) (b) Ethylene glycol (HOCH2CH2OH) in hexane (CH3CH2CH2CH2CH2CH3) or in water. (c) Diethyl ether (CH3CH2OCH2CH3) in water or in ethanol (CH3CH2OH) SOLUTION:

  17. Gas – Liquid Solutions Why aren’t gases very soluble in water? Table 13.3 Correlation Between Boiling Point and Solubility in Water Gas Solubility (M)* bp (K) He 4.2 x 10-4 4.2 Ne 6.6 x 10-4 27.1 N2 10.4 x 10-4 77.4 CO 15.6 x 10-4 81.6 O2 21.8 x 10-4 90.2 NO 32.7 x 10-4 121.4 * At 273K and 1 atm

  18. Gas Solutions and Solid Solutions • Gas – Gas Solutions • Infinitely soluble in each other • Gas – Solid Solutions • Gas molecules occupy spaces between the solid particles • Solid – Solid solutions • Alloys

  19. The arrangement of atoms in two types of alloys Figure 13.4

  20. Memory Metal Products

  21. solute (aggregated) + heat solute (separated) DHsolute > 0 solvent (aggregated) + heat solvent (separated) DHsolvent > 0 solute (separated) + solvent (separated) solution + heat DHmix < 0 Heats of solution and solution cycles 1. Solute particles separate from each other - endothermic 2. Solvent particles separate from each other - endothermic 3. Solute and solvent particles mix - exothermic DHsoln = DHsolute + DHsolvent + DHmix

  22. Solution cycles and the enthalpy components of the heat of solution Figure 13.5

  23. H2O M+ (g) [or X-(g)] M+(aq) [or X-(aq)] DHhydr of the ion < 0 M+ (g) + X-(g) MX(s) DHlattice is always (-) Heats of Hydration solvation of ions by water - always exothermic DHhydr is related to the charge density of the ion, that is, coulombic charge and size matter. Lattice energy is the DH involved in the formation of an ionic solid from its gaseous ions.

  24. Heats of Hydration: Ionic Solids in Water Solvation/Hydration: The process of surrounding a solute particle with solvent particles DHsoln = DHsolute + DHsolvent + DHmix DHsoln = DHlattice + DHhydr of the ions

  25. Figure 13.2 Hydration shells around an aqueous ion

  26. Table 13.4 Trends in Ionic Heats of Hydration Ion Ionic Radius (pm) DHhydr (kJ/mol) Group 1A(1) Li+ 76 -510 Na+ 102 -410 K+ 138 -336 Rb+ 152 -315 Cs+ 167 -282 Group 2A(2) Mg2+ 72 -1903 Ca2+ 100 -1591 Sr2+ 118 -1424 Ba2+ Group 7A(17) F- 133 -431 Cl- 181 -313 Br- 196 -284 I- 220 -247

  27. Figure 13.6 Dissolving ionic compounds in water NaCl NH4NO3 NaOH

  28. The Solution Process and the Change in Entropy • The natural tendency of a system to disperse matter or energy • Which has highest entropy: solid, liquid, or gas? • Which has highest entropy: pure solute, pure solvent, solution?

  29. Entropy, S and Enthalpy (H)

  30. Enthalpy diagrams for dissolving NaCl and octane in hexane Figure 13.7 Entropy drives the reaction (DH negligible) Entropy not big enough to overcome DH

  31. Solubility as an Equilibrium Process • Saturated • Unsaturated • Supersaturated • Equilibrium

  32. solute (undissolved) solute (dissolved) Figure 13.8 Equilibrium in a saturated solution

  33. Figure 13.9 Sodium acetate crystallizing from a supersaturated solution

  34. Figure 13.10 The relation between solubility and temperature for several ionic compounds

  35. Fighting thermal pollution Figure 13.11

  36. Factors Affecting Solubility equilibrium Increase pressure New equilibrium William Henry Henry’s Law: As Pressure increases, concentration of dissolved gases ____

  37. Henry’s Law Sgas = kH X Pgas The solubility of a gas (Sgas) is directly proportional to the partial pressure of the gas (Pgas) above the solution.

  38. Examples of Henry’s Law

  39. PROBLEM: The partial pressure of carbon dioxide gas inside a bottle of cola is 4 atm at 250C. What is the solubility of CO2? The Henry’s law constant for CO2 dissolved in water is 3.3x10-2 mol/L*atm at 250C. S = (3.3x10-2 mol/L*atm)(4 atm) = CO2 SAMPLE PROBLEM 13.2 Using Henry’s Law to Calculate Gas Solubility SOLUTION: 0.1 mol/L

  40. amount (mol) of solute Molarity (M) volume (L) of solution amount (mol) of solute Molality (m) mass (kg) of solvent mass of solute Parts by mass mass of solution volume of solute Parts by volume volume of solution amount (mol) of solute Mole fraction  amount (mol) of solute + amount (mol) of solvent Table 13.5 Concentration Definitions Concentration Term Ratio

  41. PROBLEM: Hydrogen peroxide is a powerful oxidizing agent used in concentrated solution in rocket fuels and in dilute solution a a hair bleach. An aqueous solution H2O2 is 30.0% by mass and has a density of 1.11 g/mL. Calculate its 0.882 mol H2O2 SAMPLE PROBLEM 13.5 Converting Concentration Units (a) Molality (b) Mole fraction of H2O2 (c) Molarity PLAN: (a) To find the mass of solvent we assume the % is per 100 g of solution. Take the difference in the mass of the solute and solution for the mass of peroxide. (b) Convert g of solute and solvent to moles before finding c. (c) Use the density to find the volume of the solution. 70.0 g H2O SOLUTION: (a) g of H2O = 100. g solution - 30.0 g H2O2 = 30.0 g H2O2 mol H2O2 34.02 g H2O2 molality = = 12.6 m H2O2 kg H2O 70.0 g H2O 103 g

  42. mol H2O 18.02 g H2O mL 1.11 g L 103 mL SAMPLE PROBLEM 13.5 Converting Concentration Units continued (b) 70.0 g H2O = 3.88 mol H2O 0.882 mol H2O2 = 0.185 c of H2O2 0.882 mol H2O2 + 3.88 mol H2O (c) = 90.1 mL solution 100.0 g solution 0.882 mol H2O2 = 9.79 M H2O2 90.1 mL solution

  43. A little harder… • The electrolyte in car batteries is a 3.75 M sulfuric acid solution that has a density of 1.230 g/ml. Calculate the molality of the sulfuric acid.

  44. Colligative Properties of Sol’ns • The presence of solute particles changes the physical properties of the solution • Depends on NUMBER of solute particles • Colligative (“collective”) Properties: • Vapor pressure lowering, BP elevation, FP depression, osmotic pressure

  45. Vapor Pressure of Solutions Which has higher vapor pressure: water or aqueous solution? Why?

  46. Collig Properties of Nonvolatile Nonelectrolyte Sol’ns • i.e. sugar Effect on Vapor Pressure: Less surface area— less evaporation

  47. Figure 13.15 In terms of entropy? l -> g increases entropy Mixture already has higher entropy Than the pure solvent, so Less evaporation is necessary

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