Chapter 13 Properties of Solutions

1. CHEMISTRYThe Central Science 9th Edition Chapter 13Properties of Solutions

2. Text, P. 417, review (Chapter 11)

3. 13.1: The Solution Process • Solutions • homogeneous mixtures • Solution formation is affected by • strength and type of intermolecular forces • forces are between and among the solute and solvent particles

4. Text, P. 486

5. Hydration of solute • Attractive forces between solute & solvent particles are comparable in magnitude with those between the solute or solvent particles themselves • Note attraction of charges • What has to happen to: • Water’s H-bonds? • NaCl? • What intermolecular • force is at work in • solvation? Text, P. 486

6. Energy Changes and Solution Formation • There are three energy steps in forming a solution: • the enthalpy change in the solution process is • Hsoln = H1 + H2 + H3 • Hsoln can either be + or - depending on the intermolecular forces Text, P. 487

7. Text, P. 488 MgSO4 Hot Pack NH4NO3 Cold Pack

8. Breaking attractive intermolecular forces is always endothermic • Formingattractive intermolecular forces is always exothermic • To determine whether Hsoln is positive or negative, consider the strengths of all solute-solute and solute-solvent interactions: • H1 and H2 are both positive • H3 is always negative

9. Rule: Polar solvents dissolve polar solutes • Non-polar solvents dissolve non-polar solutes • (like dissolves like) • WHY? • If Hsoln is too endothermic a solution will not form • NaCl in gasoline: weak ion-dipole forces (gasoline is non-polar) • The ion-dipole forces do not compensate for the separation of ions

10. Solution Formation, Spontaneity, and Disorder • A spontaneous process occurs without outside intervention • When energy of the system decreases, the process is spontaneous • Some spontaneous processes do not involve the system moving to a lower energy state (e.g. an endothermic reaction) • If the process leads to a greater state of disorder, then the process is spontaneous • Entropy

11. Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids • Therefore, they spontaneously mix even though Hsoln is very close to zero Text, P. 489

12. Solution Formation and Chemical Reactions • Example: • Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g) • When all the water is removed from the NiCl2 solution, no Ni is found only NiCl2·6H2O(achemical reaction that results in the formation of a solution) • Water molecules fit into the crystal lattice in places not specifically occupied by a cation or an anion • Hydrates • Water of hydration • Think about it: What happens when NaCl is dissolved in water and then heated to dryness?

13. NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq) • When the water is removed from the solution, NaCl is found • NaCl dissolution is a physical process

14. Sample problem # 3

15. 13.2: Saturated Solutions and Solubility • Dissolve: solute + solvent  solution • Crystallization: solution  solute + solvent • Saturation: crystallization and dissolution are in equilibrium • Solubility: amount of solute required to form a saturated solution • Supersaturated: a solution formed when more solute is dissolved than in a saturated solution

16. 13.3: Factors Affecting Solubility • 1. Solute-Solvent Interaction • “Like dissolves like” • Miscible liquids: mix in any proportions • Immiscible liquids: do not mix

17. Generalizations: • Intermolecular forces are important: • Water and ethanol are miscible • broken hydrogen bonds in both pure liquids are • re-established in the mixture • The number of carbon atoms in a chain affects solubility: the more C atoms in the chain, the less soluble the substance is in water

18. Generalizations, continued: • The number of -OH groups within a molecule increases solubility in water • The more polar bonds in the molecule, the better it dissolves in a polar solvent (like dissolves like) • Network solids do not dissolve • the strong IMFs in the solid are not re-established in any solution

19. Text, P. 493

20. Read “Chemistry & Life”, P. 494 Fat soluble vitamin Water soluble vitamin

21. 2. Pressure Effects • Solubility of a gas in a liquid is a function of the pressure of the gas

22. High pressure means • More molecules of gas are close to the solvent • Greater solution/gas interactions • Greater solubility • If Sg is the solubility of a gas • k is a constant • Pg is the partial pressure of a gas • then Henry’s Law gives: • Carbonated Beverages!

23. 3. Temperature Effects • As temperature increases • Solubility of solids generally increases • Solubility of gases decreases • Thermal pollution Text, P. 497

24. Figure 13.17, P. 497

25. Sample problem # 17

26. 13.4: Ways of Expressing Concentration • All methods involve quantifying amount of solute per amount of solvent (or solution) • Amounts or measures are masses, moles or liters • Qualitatively solutions are dilute or concentrated

27. Definitions: • 1.

28. 2. • 3. • Recall mass can be converted to moles using the molar mass

29. 4. • Converting between molarity (M) and molality (m) requires density • Molality doesn’t vary with temperature • Mass is constant • Molarity changes with temperature • Expansion/contraction of solution changes volume

30. Text, P. 501

31. Sample Problems #31, 33, 37, 39, 41

32. 13.5: Colligative Properties • Colligative properties depend on quantity of solute particles, not their identity • Electrolytes vs. nonelectrolytes • 0.15m NaCl  0.15m in Na+ & 0.15m in Cl-  0.30m in particles • 0.050m CaCl2  0.050m in Ca+2 & 0.1m in Cl-  0.15m in particles • 0.10m HCl  0.10m in H+ & 0.10m in Cl-  0.20m in particles • 0.050m HC2H3O2  between 0.050m & 0.10m in particles • 0.10m C12H22O11 0.10m in particles • Compare physical properties of the solution with those of the pure solvent

33. 1. Lowering Vapor Pressure • Non-volatile solutes reduce the ability of the surface solvent molecules to escape the liquid • Vapor pressure is lowered • Raoult’s Law: • PA is the vapor pressure with solute • PA is the vapor pressure without solute • A is the mole fraction of solvent in solution A Increase X of solute, decrease vapor pressure above the solution

34. Ideal solution: one that obeys Raoult’s law • Raoult’s law breaks down (Real solutions) • Real solutions approximate ideal behavior when • solute concentration is low • solute and solvent have similar IMFs • Assume ideal solutions for problem solving • 2. Boiling-Point Elevation • The triple point - critical point curve is lowered

35. At 1 atm (normal BP of pure liquid) there is a lower vapor pressure of the solution • A higher temperature is required to reach a vapor pressure of 1 atm for the solution (Tb) • Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m:

36. Text, P. 505

37. 3. Freezing Point Depression • The solution freezes at a lower temperature (Tf) than the pure solvent • lower vapor pressure for the solution • Decrease in FP (Tf) is directly proportional to molality (Kfis the molal freezing-point-depression constant):

38. Text, P. 505 Applications: Antifreeze!

40. Examples: # 45, 47, 49, 51 & 53 • A neat link

41. 4. Osmosis • Semipermeable membrane: permits passage of some components of a solution • Example: cell membranes and cellophane • Osmosis: the movement of a solvent from low solute concentration to high solute concentration • There is movement in both directions across a semipermeable membrane • “Where ions go, water will flow” ~ Mrs. Moss

42. Eventually the pressure difference between the arms stops osmosis Text, P. 507

43. Osmotic pressure, , is the pressure required to stop osmosis: • It is colligative because it depends on the concentration of the solute in the solvent

44. Isotonic solutions: two solutions with the same  separated by a semipermeable membrane • Hypertonic solution: a solution that is more concentrated than a comparable solution • Hypotonic solution: a solution of lower  than a hypertonic solution • Osmosis is spontaneous • Read text, P. 508 – 509 for practical examples

45. Examples: #57, 59 & 61

46. There are differences between expected and observed changes due to colligative properties of strong electrolytes • Electrostatic attractions between ions • “ion pair” formation temporarily reduces the number of particles in solution • van’t Hoff factor (i): measure of the extent of ion dissociation

47. Ratio of the actual value of a colligative property to the calculated value (assuming it to be a nonelectrolyte) • Ideal value for a salt is the # of ions per formula unit • Factors that affect i: • Dilution • Magnitude of charge on ions • lower charges, less deviation

48. Sample Problem, # 63, 82

49. 11.6: Colloids • Read Text, Section 13.6, P. 511 – 515 • Terms/Processes: • Tyndall effect • Hydrophilic • Hydrophobic • Adsorption • Coagulation

50. 11.6: Colloids • Read Text, Section 13.6, P. 511 – 515 • Suspensions in which the suspended particles are larger than molecules • too small to drop out of the suspension due to gravity • Tyndall effect: ability of a colloid to scatter light • The beam of light can be seen through the colloid