1. Chapter 8 Liquids and Solutions As already mentioned in chapter 2, a lot of chemistry is done in solution, especially aqueous solution. In this chapter we address issues that arise when dealing with solutions.

2. The Structure of Gases, Liquids and Solids Figure 8.1

3. The Structure of Gases, Liquids and Solids Table 8.1

4. The Structure of Gases, Liquids and Solids ● Intramolecular bond ● Intermolecular force Figure 8.2

5. Intermolecular Forces ● Absent in kinetic molecular theory. ● In their absence, all matter is in the gas phase. ● Relative strength of intermolecular forces established using boiling points. Low bp means weak intermolecular forces. High bp means strong intermolecular forces.

6. Intermolecular Forces Five important types dipole-dipole dipole-induced dipole induced dipole-induced dipole van der Waals (aka London, dispersion) Hydrogen bonding

7. Intermolecular Forces ● dipole-dipole Figure 8.3

8. Intermolecular Forces ● dipole-induced dipole Figure 8.4

9. Intermolecular Forces ● induced dipole-induced dipole Figure 8.5

10. Intermolecular Forces ● van der Waals Weak Present in all systems. Proportional to the number of electrons in the molecules. Table 8.2

11. Intermolecular Forces ● van der Waals MW, Figure 8.8 Shape, Figure 8.7 – n-pentane (bp 36.1 ⁰ C) vs neopentane (bp 9.5 ⁰ C) Figure 8.8 Figure 8.7

12. Intermolecular Forces ● Hydrogen bonding Misleading name Possible in molecules with H - X bond where X is F, O, or N. – Highlights importance of Lewis structure.

13. Intermolecular Forces...importance of Lewis structure Two isomers, only one participates in hydrogen bonding.

14. Intermolecular Forces ● Hydrogen bonding Profound consequences Figure 8.9

15. Intermolecular Forces ● Hydrogen bonding...results in liquid water on Earth!!

16. Relative Strengths of Intermolecular Forces Table 8.3

17. Relative Strengths of Intermolecular Forces Table 8.5

18. The Kinetic Theory of Liquids ● Average KE  T (section 6.2). ● Range of KE. ● Intermolecular forces present. That's why it's a liquid.

19. The Kinetic Theory of Liquids ● Enthalpy of vaporization, ΔH° vap. ● Enthalpy of fusion, ΔH° fus. ● ΔH° vap >> ΔH° fus. Why?

20. The Vapor Pressure of a Liquid ● Introduced with Dalton's law, section 6.14. ● Properly called equilibrium vapor pressure of a liquid. ● Increases with temperature. ● Reason liquids in open containers (non equilibrium situation) evaporate.

21. The Vapor Pressure of a Liquid Figure 8.11

22. The Vapor Pressure of a Liquid Figure 8.12

23. The Vapor Pressure of a Liquid Figure 8.13

24. Melting Point and Freezing Point ● Should be the same. Some liquids supercool. Solids don't superheat. ● Melting points used to characterize compounds. Purity Identification, especially in organic chemistry

25. Melting Point and Freezing Point ● During melting, heat added to the system does not raise the temperature. ● Where does it go?

26. Melting Point and Freezing Point ● During melting, heat added to the system does not raise the temperature. ● Where does it go? Into ΔH° fus

27. Melting Point and Freezing Point Figure 8.15

28. Boiling Point ● Indication of strength of intermolecular forces. ● Vapor pressure of liquid = external pressure. Therefore, bp varies with external pressure. ● When external pressure is 1 bar, the boiling point is called the normal boiling point.

29. Boiling Point Figure 8.17

30. Phase Diagrams ● Plot of equilibrium phase as a function of P and T. ● Axes often not linear. ● Determined experimentally.

31. Phase Diagrams Figure 8.18

32. Hydrogen Bonding and the Anomalous Properties of Water ● Water is a strange substance. Density decreases upon freezing. Boiling point is high. Specific heat is high. ● Many of its strange properties are the result of the hydrogen bonding present in water.

33. Hydrogen Bonding and the Anomalous Properties of Water ● HF has a larger ΔEN, but fewer H per X. ● NH 3 has more H per X, but a smaller ΔEN. ● H 2 O has just the right balance of H per X and ΔEN to make it such an unusual molecule.

34. Solutions: Like Dissolves Like ● Move from pure liquids to solutions. ● Emphasis on solubility: Important property in chemistry and biochemistry. ● Characterize solvents as Polar.Nonpolar. ● This terminology was first used in section 4.17.

35. Solutions: Like Dissolves Like ● Polarity of solvent will determine what kind of solutes dissolve in it. ● Hence the title of the section.

36. Solutions: Like Dissolves Like ● Iodine molecules (I 2 ) are bound to each other through van der Waals interactions. Intermolecular force ● KMnO 4 is made up of K + and MnO 4 - ions which are bound to each other through ionic bonding.

37. Solutions: Like Dissolves Like Table 8.6

38. Solutions: Like Dissolves Like Figure 8.24

39. Hydrophilic and Hydrophobic Molecules ● Hydrophilic Example: molecules which hydrogen bond. Soluble ionic compounds. ● Hydrophobic Example: hydrocarbons, C x H y.

40. Hydrophilic and Hydrophobic Molecules ● Portions of a single molecule can be hydrophilic and hydrophobic: OH part of an alcohol is hydrophilic. The alkyl part (C x H y ) is hydrophobic. Table 8.7

41. Hydrophilic and Hydrophobic Molecules Table 8.8

42. Soaps, Detergents, and Dry-Cleaning Agents ● Involve two fundamental principles Solubility Intermolecular interactions

43. Soaps, Detergents, and Dry-Cleaning Agents ● “Dirt” is not soluble in water. ● It is soluble in hydrocarbons, but no one wants to wash their clothes with lighter fluid or gasoline. ● Trick the “dirt” into dissolving in a hydrocarbon which has been slipped into a water medium.

44. Soaps, Detergents, and Dry-Cleaning Agents ● …a hydrocarbon which has been slipped into a water medium. Figure 8.28 Figure 8.31

45. Soaps, Detergents, and Dry-Cleaning Agents ● Major problem with soap: hard water

46. Soaps, Detergents, and Dry-Cleaning Agents ● Water softening ● Synthetic soaps Figure 8.32

47. Why Do Some Solids Dissolve in Water? ● Both ionic and covalent solids will dissolve in water. ● But not all ionic and covalent solids!

48. Why Do Some Solids Dissolve in Water? ● Energy required to break up solid. ● Energy produced by interaction of solid components with solvent. The relative magnitude of these two energy terms determines solubility.

49. Solubility Equilibria ● Already seen an equilibrium, section 8.5: liquid liquid ⇄ vapor. ● Now we have pure solid pure solid ⇄ solute in solution. ● Reversible and dynamic in both cases.

50. Solubility Equilibria ● Precipitation reaction Soluble species form an insoluble product. ● Saturated Solution rate of precipitation = rate of dissolution ● Solubility Maximum amount of solute which can be dissolved at a given temperature.

51. Solubility Equilibria ● Electrolytes Strong electrolytes – All the solutes break up into ions. Weak electrolytes – Some of the solutes break up into ions. ● Nonelectrolytes – None of the solutes break up into ions.

52. Solubility Rules Table 8.9

53. Solubility Rules ● Solubility is a subjective term. Figure 8.38

54. Net Ionic Equations ● Condensed BaCl 2 (aq) +Na 2 SO 4 (aq) → BaSO 4 (s)↓ +2NaCl(aq) BaCl 2 (aq) +Na 2 SO 4 (aq) → BaSO 4 (s)↓ +2NaCl(aq) ● Ionic Ba +2 (aq) + 2Cl - (aq) + 2Na + (aq) + SO 4 -2 (aq) → Ba +2 (aq) + 2Cl - (aq) + 2Na + (aq) + SO 4 -2 (aq) → BaSO 4 (s)↓ + 2Na + (aq) + 2Cl - (aq) BaSO 4 (s)↓ + 2Na + (aq) + 2Cl - (aq) ● Net Ionic Ba +2 (aq) + SO 4 -2 (aq) → BaSO 4 (s)↓ Ba +2 (aq) + SO 4 -2 (aq) → BaSO 4 (s)↓

55. Net Ionic Equations ● Each of the previous three types has its virtues and limitations. ● For example, the net ionic lacks information about the spectator ions: Ba +2 (aq) + SO 4 -2 (aq) → BaSO 4 (s)↓ Ba +2 (aq) + SO 4 -2 (aq) → BaSO 4 (s)↓