1. 1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 13 Intermolecular Forces, Liquids, and Solids © 2006 Brooks/Cole Thomson Lectures written by John Kotz

2. 2 © 2006 Brooks/Cole - Thomson WHY? Why is water usually a liquid and not a gas? Why does liquid water boil at such a high temperature for such a small molecule? Why does ice float on water? Why do snowflakes have 6 sides? Why is I 2 a solid whereas Cl 2 is a gas? Why are NaCl crystals little cubes?

3. 3 © 2006 Brooks/Cole - Thomson Liquids, Solids & Intermolecular Forces Chap. 13

4. 4 © 2006 Brooks/Cole - Thomson Inter- molecular Forces Have studied INTRA molecular forces—the forces holding atoms together to form molecules. Now turn to forces between molecules — INTER molecular forces. Forces between molecules, between ions, or between molecules and ions.

5. 5 © 2006 Brooks/Cole - Thomson Ion-Ion Forces for comparison of magnitude Na + —Cl - in salt These are the strongest forces. Lead to solids with high melting temperatures. NaCl, mp = 800 o C MgO, mp = 2800 o C

6. 6 © 2006 Brooks/Cole - Thomson Covalent Bonding Forces for comparison of magnitude C–H, 413 kJ/mol C=C, 610 kJ/mol C–C, 346 kJ/mol CN, 887 kJ/mol

7. 7 © 2006 Brooks/Cole - Thomson Attraction Between Ions and Permanent Dipoles Water is highly polar and can interact with positive ions to give hydrated ions in water.

8. 8 © 2006 Brooks/Cole - Thomson Attraction Between Ions and Permanent Dipoles Water is highly polar and can interact with positive ions to give hydrated ions in water.

9. 9 © 2006 Brooks/Cole - Thomson Attraction Between Ions and Permanent Dipoles Many metal ions are hydrated. This is the reason metal salts dissolve in water.

10. 10 © 2006 Brooks/Cole - Thomson Attraction between ions and dipole depends on ion charge and ion-dipole distance. Measured by ∆H for M n+ + H 2 O --> [M(H 2 O) x ] n+ -1922 kJ/mol -405 kJ/mol -263 kJ/mol Attraction Between Ions and Permanent Dipoles

11. 11 © 2006 Brooks/Cole - Thomson Dipole-Dipole Forces Such forces bind molecules having permanent dipoles to one another.

12. 12 © 2006 Brooks/Cole - Thomson Dipole-Dipole Forces Influence of dipole-dipole forces is seen in the boiling points of simple molecules. CompdMol. Wt.Boil Point N 2 28-196 o C CO28-192 o C Br 2 16059 o C ICl16297 o C

13. 13 © 2006 Brooks/Cole - Thomson Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole-dipole attractions. H-bonding is strongest when X and Y are N, O, or F

14. 14 © 2006 Brooks/Cole - Thomson H-Bonding Between Methanol and Water H-bondH-bond ---- ++++ ----

15. 15 © 2006 Brooks/Cole - Thomson H-Bonding Between Two Methanol Molecules H-bondH-bond ---- ++++ ----

16. 16 © 2006 Brooks/Cole - Thomson H-Bonding Between Ammonia and Water H-bondH-bond ---- ++++ ---- This H-bond leads to the formation of NH 4 + and OH -

17. 17 © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O H-bonding is especially strong in water because the O—H bond is very polarthe O—H bond is very polar there are 2 lone pairs on the O atomthere are 2 lone pairs on the O atom Accounts for many of water’s unique properties.

18. 18 © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O Ice has open lattice-like structure. Ice density is < liquid. And so solid floats on water. Snow flake: www.snowcrystals.com

19. 19 © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O Ice has open lattice-like structure. Ice density is < liquid and so solid floats on water. One of the VERY few substances where solid is LESS DENSE than the liquid.

20. 20 © 2006 Brooks/Cole - Thomson A consequence of hydrogen bonding

21. 21 © 2006 Brooks/Cole - Thomson Hydrogen Bonding in H 2 O H bonds ---> abnormally high specific heat capacity of water (4.184 J/gK) This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

22. 22 © 2006 Brooks/Cole - Thomson Hydrogen Bonding H bonds leads to abnormally high boiling point of water. See Screen 13.7

23. 23 © 2006 Brooks/Cole - Thomson Boiling Points of Simple Hydrogen- Containing Compounds Active Figure 13.8

24. 24 © 2006 Brooks/Cole - Thomson Methane Hydrate

25. 25 © 2006 Brooks/Cole - Thomson Methane Clathrate

26. 26 © 2006 Brooks/Cole - Thomson Hydrogen Bonding in Biology H-bonding is especially strong in biological systems — such as DNA. DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O. Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases. —adenine with thymine —adenine with thymine —guanine with cytosine —guanine with cytosine

27. 27 © 2006 Brooks/Cole - Thomson Portion of a DNA chain Double helix of DNA

28. 28 © 2006 Brooks/Cole - Thomson Base-Pairing through H-Bonds

29. 29 © 2006 Brooks/Cole - Thomson Base-Pairing through H-Bonds

30. 30 © 2006 Brooks/Cole - Thomson Double Helix of DNA

31. 31 © 2006 Brooks/Cole - Thomson Discovering the Double Helix James Watson and Francis Crick, 1953 Rosalind Franklin, 1920- 1958 Maurice Wilkins, 1916 - 2004

32. 32 © 2006 Brooks/Cole - Thomson Hydrogen Bonding in Biology Hydrogen bonding and base pairing in DNA. See Screen 13.6

33. 33 © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES How can non-polar molecules such as O 2 and I 2 dissolve in water? The water dipole INDUCES a dipole in the O 2 electric cloud. Dipole-induced dipole

34. 34 © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES Solubility increases with mass the gas

35. 35 © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES Process of inducing a dipole is polarizationProcess of inducing a dipole is polarization Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.

36. 36 © 2006 Brooks/Cole - Thomson IM FORCES — INDUCED DIPOLES Consider I 2 dissolving in ethanol, CH 3 CH 2 OH. O H -  +  I-I R -  +  O H +  -  R The alcohol temporarily creates or INDUCES a dipole in I 2.

37. 37 © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES Formation of a dipole in two nonpolar I 2 molecules. Induced dipole- induced dipole

38. 38 © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES The induced forces between I 2 molecules are very weak, so solid I 2 sublimes (goes from a solid to gaseous molecules).

39. 39 © 2006 Brooks/Cole - Thomson FORCES INVOLVING INDUCED DIPOLES The magnitude of the induced dipole depends on the tendency to be distorted. Higher molec. weight ---> larger induced dipoles. MoleculeBoiling Point ( o C) MoleculeBoiling Point ( o C) CH 4 (methane) - 161.5 CH 4 (methane) - 161.5 C 2 H 6 (ethane)- 88.6 C 2 H 6 (ethane)- 88.6 C 3 H 8 (propane) - 42.1 C 3 H 8 (propane) - 42.1 C 4 H 10 (butane) - 0.5 C 4 H 10 (butane) - 0.5

40. 40 © 2006 Brooks/Cole - Thomson Boiling Points of Hydrocarbons Note linear relation between bp and molar mass. CH 4 C2H6C2H6C2H6C2H6 C3H8C3H8C3H8C3H8 C 4 H 10

41. 41 © 2006 Brooks/Cole - Thomson Summary of Intermolecular Forces Ion-dipole forcesIon-dipole forces Dipole-dipole forcesDipole-dipole forces –Special dipole-dipole force: hydrogen bonds Forces involving nonpolar molecules: induced forcesForces involving nonpolar molecules: induced forces

42. 42 © 2006 Brooks/Cole - Thomson Intermolecular Forces Summary

43. 43 © 2006 Brooks/Cole - Thomson Intermolecular Forces Figure 13.13

44. 44 © 2006 Brooks/Cole - Thomson Liquids Section 13.5 In a liquid molecules are in constant motionmolecules are in constant motion there are appreciable intermolec. forcesthere are appreciable intermolec. forces molecules close togethermolecules close together Liquids are almost incompressibleLiquids are almost incompressible Liquids do not fill the containerLiquids do not fill the container

45. 45 © 2006 Brooks/Cole - Thomson Liquids The two key properties we need to describe are EVAPORATION and its opposite— CONDENSATION break IM bonds make IM bonds Add energy Remove energy LIQUID VAPOR<---condensation evaporation--->

46. 46 © 2006 Brooks/Cole - Thomson Liquids—Evaporation To evaporate, molecules must have sufficient energy to break IM forces. Breaking IM forces requires energy. The process of evaporation is endothermic.

47. 47 © 2006 Brooks/Cole - Thomson Liquids— Distribution of Energies Distribution of molecular energies in a liquid. KE is propor- tional to T. Distribution of molecular energies in a liquid. KE is propor- tional to T. 0 Number of molecules Molecular energy higher Tlower T See Figure 13.14 Minimum energy req’d to break IM forces and evaporate

48. 48 © 2006 Brooks/Cole - Thomson Distribution of Energy in a Liquid Figure 13.14

49. 49 © 2006 Brooks/Cole - Thomson Liquids At higher T a much larger number of molecules has high enough energy to break IM forces and move from liquid to vapor state. High E molecules carry away E. You cool down when sweating or after swimming.

50. 50 © 2006 Brooks/Cole - Thomson Liquids When molecules of liquid are in the vapor state, they exert a VAPOR PRESSURE EQUILIBRIUM VAPOR PRESSURE is the pressure exerted by a vapor over a liquid in a closed container when the rate of evaporation = the rate of condensation.

51. 51 © 2006 Brooks/Cole - Thomson Measuring Equilibrium Vapor Pressure Liquid in flask evaporates and exerts pressure on manometer. Active Fig. 13.17

52. 52 © 2006 Brooks/Cole - Thomson Vapor Pressure CD, Screen 13.9

53. 53 © 2006 Brooks/Cole - Thomson Equilibrium Vapor Pressure Active Figure 13.18

54. 54 © 2006 Brooks/Cole - Thomson Liquids Equilibrium Vapor Pressure FIGURE 13.18: VP as a function of T. 1. The curves show all conditions of P and T where LIQ and VAP are in EQUILIBRIUM 2. The VP rises with T. 3. When VP = external P, the liquid boils. This means that BP’s of liquids change with altitude. This means that BP’s of liquids change with altitude.

55. 55 © 2006 Brooks/Cole - Thomson Boiling Liquids Liquid boils when its vapor pressure equals atmospheric pressure.

56. 56 © 2006 Brooks/Cole - Thomson Boiling Point at Lower Pressure When pressure is lowered, the vapor pressure can equal the external pressure at a lower temperature.

57. 57 © 2006 Brooks/Cole - Thomson Consequences of Vapor Pressure Changes When can cools, vp of water drops. Pressure in the can is less than that of atmosphere, so can is crushed.

58. 58 © 2006 Brooks/Cole - Thomson 4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING POINT 5. VP of a given molecule at a given T depends on IM forces. Here the VP’s are in the order C 2 H 5 H 5 C 2 H H 5 C 2 H H wateralcoholether increasing strength of IM interactions extensive H-bonds dipole- dipole O O O Liquids Figure 13.18: VP versus T

59. 59 © 2006 Brooks/Cole - Thomson Liquids HEAT OF VAPORIZATION is the heat req’d (at constant P) to vaporize the liquid. LIQ + heat ---> VAP Compd.∆H vap (kJ/mol) IM Force H 2 O40.7 (100 o C)H-bonds SO 2 26.8 (-47 o C)dipole Xe12.6 (-107 o C)induced dipole

60. 60 © 2006 Brooks/Cole - Thomson Equilibrium Vapor Pressure & the Clausius-Clapeyron Equation Clausius-Clapeyron equation — used to find ∆H˚ vap.Clausius-Clapeyron equation — used to find ∆H˚ vap. The logarithm of the vapor pressure P is proportional to ∆H vaporiation and to 1/T.The logarithm of the vapor pressure P is proportional to ∆H vaporiation and to 1/T. ln P = –(∆H˚ vap /RT) + Cln P = –(∆H˚ vap /RT) + C

61. 61 © 2006 Brooks/Cole - Thomson Liquids Molecules at surface behave differently than those in the interior. Molecules at surface experience net INWARD force of attraction. This leads to SURFACE TENSION — the energy req’d to break the surface.

62. 62 © 2006 Brooks/Cole - Thomson Surface Tension SURFACE TENSION also leads to spherical liquid droplets.

63. 63 © 2006 Brooks/Cole - Thomson Liquids Intermolec. forces also lead to CAPILLARY action and to the existence of a concave meniscus for a water column. concave meniscus H 2 O in glass tube ADHESIVE FORCES between water and glass COHESIVE FORCES between water molecules

64. 64 © 2006 Brooks/Cole - Thomson Capillary Action Movement of water up a piece of paper depends on H-bonds between H 2 O and the OH groups of the cellulose in the paper.

65. 65 © 2006 Brooks/Cole - Thomson Metallic and Ionic Solids Sections 13.6-8

66. 66 © 2006 Brooks/Cole - Thomson Types of Solids Table 13.6 TYPEEXAMPLEFORCE Ionic NaCl, CaF 2, ZnSIon-ion MetallicNa, FeMetallic MolecularIce, I 2 Dipole Ind. dipole NetworkDiamondExtended Graphitecovalent

67. 67 © 2006 Brooks/Cole - Thomson Network Solids Diamond Graphite

68. 68 © 2006 Brooks/Cole - Thomson Network Solids A comparison of diamond (pure carbon) with silicon.

69. 69 © 2006 Brooks/Cole - Thomson Properties of Solids 1. Molecules, atoms or ions locked into a CRYSTAL LATTICE 2. Particles are CLOSE together 3. STRONG IM forces 4. Highly ordered, rigid, incompressible ZnS, zinc sulfide

70. 70 © 2006 Brooks/Cole - Thomson Crystal Lattices Regular 3-D arrangements of equivalent LATTICE POINTS in space.Regular 3-D arrangements of equivalent LATTICE POINTS in space. Lattice points define UNIT CELLSLattice points define UNIT CELLS –smallest repeating internal unit that has the symmetry characteristic of the solid.

71. 71 © 2006 Brooks/Cole - Thomson Cubic Unit Cells All angles are 90 degrees All sides equal length There are 7 basic crystal systems, but we are only concerned with CUBIC.

72. 72 © 2006 Brooks/Cole - Thomson Cubic Unit Cells of Metals Figure 13.27 Simple cubic (SC) Body- centered cubic (BCC) Face- centered cubic (FCC)

73. 73 © 2006 Brooks/Cole - Thomson Cubic Unit Cells of Metals Figure 13.27

74. 74 © 2006 Brooks/Cole - Thomson Units Cells for Metals Figure 13.28

75. 75 © 2006 Brooks/Cole - Thomson E atom is at a corner of a unit cell and is shared among 8 unit cells. Each edge is shared with 4 cells Each face is part of two cells. Simple Cubic Unit Cell Figure 13.26

76. 76 © 2006 Brooks/Cole - Thomson Atom Packing in Unit Cells Assume atoms are hard spheres and that crystals are built by PACKING of these spheres as efficiently as possible.

77. 77 © 2006 Brooks/Cole - Thomson Atom Packing in Unit Cells Assume atoms are hard spheres and that crystals are built by PACKING of these spheres as efficiently as possible.

78. 78 © 2006 Brooks/Cole - Thomson Atom Packing in Unit Cells

79. 79 © 2006 Brooks/Cole - Thomson Crystal Lattices— Packing of Atoms or Ions FCC is more efficient than either BC or SC.FCC is more efficient than either BC or SC. Leads to layers of atoms.Leads to layers of atoms.

80. 80 © 2006 Brooks/Cole - Thomson Packing of C 60 molecules. They are arranged at the lattice points of a FCC lattice. Crystal Lattices— Packing of Atoms or Ions

81. 81 © 2006 Brooks/Cole - Thomson Number of Atoms per Unit Cell Unit Cell Type Net Number Atoms Unit Cell Type Net Number Atoms SC SCBCCFCC 1 2 4

82. 82 © 2006 Brooks/Cole - Thomson Atom Sharing at Cube Faces and Corners Atom shared in corner --> 1/8 inside each unit cell Atom shared in face --> 1/2 inside each unit cell

83. 83 © 2006 Brooks/Cole - Thomson Simple Ionic Compounds Lattices of many simple ionic solids are built by taking a SC or FCC lattice of ions of one type and placing ions of opposite charge in the holes in the lattice. EXAMPLE: CsCl has a SC lattice of Cs + ions with Cl - in the center.

84. 84 © 2006 Brooks/Cole - Thomson Simple Ionic Compounds CsCl unit cell has a SC lattice of Cl - ions with Cs + in the center. 1 unit cell has 1 Cs + ion plus (8 corners)(1/8 Cl - per corner) = 1 net Cl - ion.

85. 85 © 2006 Brooks/Cole - Thomson Two Views of CsCl Unit Cell Lattice can be SC lattice of Cl - with Cs + in hole OR SC lattice of Cs + with Cl - in hole Either arrangement leads to formula of 1 Cs + and 1 Cl - per unit cell

86. 86 © 2006 Brooks/Cole - Thomson Simple Ionic Compounds Salts with formula MX can have SC structure — but not salts with formula MX 2 or M 2 X

87. 87 © 2006 Brooks/Cole - Thomson NaCl Construction FCC lattice of Cl - with Na + in holes Na + in octahedral holes

88. 88 © 2006 Brooks/Cole - Thomson Octahedral Holes - FCC Lattice

89. 89 © 2006 Brooks/Cole - Thomson The Sodium Chloride Lattice Na + ions are in OCTAHEDRAL holes in a face-centered cubic lattice of Cl - ions.

90. 90 © 2006 Brooks/Cole - Thomson Many common salts have FCC arrangements of anions with cations in OCTAHEDRAL HOLES — e.g., salts such as CA = NaCl FCC lattice of anions ----> 4 A - /unit cellFCC lattice of anions ----> 4 A - /unit cell C + in octahedral holes ---> 1 C + at centerC + in octahedral holes ---> 1 C + at center + [12 edges 1/4 C + per edge] = 4 C + per unit cell The Sodium Chloride Lattice

91. 91 © 2006 Brooks/Cole - Thomson Comparing NaCl and CsCl Even though their formulas have one cation and one anion, the lattices of CsCl and NaCl are different. The different lattices arise from the fact that a Cs + ion is much larger than a Na + ion.

92. 92 © 2006 Brooks/Cole - Thomson Common Ionic Solids Titanium dioxide, TiO 2 There are 2 net Ti 4+ ions and 4 net O 2- ions per unit cell.

93. 93 © 2006 Brooks/Cole - Thomson Common Ionic Solids Zinc sulfide, ZnSZinc sulfide, ZnS The S 2- ions are in TETRAHEDRAL holes in the Zn 2+ FCC lattice.The S 2- ions are in TETRAHEDRAL holes in the Zn 2+ FCC lattice. This gives 4 net Zn 2+ ions and 4 net S 2- ions.This gives 4 net Zn 2+ ions and 4 net S 2- ions.

94. 94 © 2006 Brooks/Cole - Thomson Common Ionic Solids Fluorite or CaF 2 FCC lattice of Ca 2+ ions This gives 4 net Ca 2+ ions. F - ions in all 8 tetrahedral holes. This gives 8 net F - ions.

95. 95 © 2006 Brooks/Cole - Thomson Barium titanate, a perovskite Ba 2+ Ti 4+ BaTiO 3

96. 96 © 2006 Brooks/Cole - Thomson Common Ionic Solids Magnesium silicate, MgSiO 3

97. 97 © 2006 Brooks/Cole - Thomson BiotiteBiotite Layers of linked octahedra of MgO 6 and FeO 6. Layers of linked SiO 4 tetrahedra. K ions between layers

98. 98 © 2006 Brooks/Cole - Thomson Phase Diagrams

99. 99 © 2006 Brooks/Cole - Thomson TRANSITIONS BETWEEN PHASES Section 13.10 Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line. (At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example.)

100. 100 © 2006 Brooks/Cole - Thomson Phase Diagram for Water Solid phase Liquid phase Gas phase

101. 101 © 2006 Brooks/Cole - Thomson Phase Equilibria — Water Solid-liquid Gas-Liquid Gas-Solid

102. 102 © 2006 Brooks/Cole - Thomson Triple Point — Water At the TRIPLE POINT all three phases are in equilibrium.

103. 103 © 2006 Brooks/Cole - Thomson Phases Diagrams— Important Points for Water T(˚C)P(mmHg) Normal boil point 100760 Normal freeze point0760 Triple point 0.00984.58

104. 104 © 2006 Brooks/Cole - Thomson Critical T and P Above critical T no liquid exists no matter how high the pressure. As P and T increase, you finally reach the CRITICAL T and P

105. 105 © 2006 Brooks/Cole - Thomson Critical T and P COMPDT c ( o C)P c (atm) H 2 O374218 CO 2 3173 CH 4 -8246 Freon-1211241 (CCl 2 F 2 ) Notice that T c and P c depend on intermolecular forces.

106. 106 © 2006 Brooks/Cole - Thomson Solid-Liquid Equilibria In any system, if you increase P the DENSITY will go up. Therefore — as P goes up, equilibrium favors phase with the larger density (or SMALLER volume/gram). Liquid H 2 OSolid H 2 O Liquid H 2 OSolid H 2 O Density1 g/cm 3 0.917 g/cm 3 cm 3 /gram11.09

107. 107 © 2006 Brooks/Cole - Thomson Solid-Liquid Equilibria Raising the pressure at constant T causes water to melt. The NEGATIVE SLOPE of the S/L line is unique to H 2 O. Almost everything else has positive slope.

108. 108 © 2006 Brooks/Cole - Thomson Solid-Liquid Equilibria The behavior of water under pressure is an example of LE CHATELIER’S PRINCIPLE At Solid/Liquid equilibrium, raising P squeezes the solid. It responds by going to phase with greater density, i.e., the liquid phase.

109. 109 © 2006 Brooks/Cole - Thomson Solid-Vapor Equilibria At P < 4.58 mmHg and T < 0.0098 ˚C solid H 2 O can go directly to vapor. This process is called SUBLIMATION This is how a frost-free refrigerator works.

110. 110 © 2006 Brooks/Cole - Thomson CO 2 Phase Diagram

111. 111 © 2006 Brooks/Cole - Thomson CO 2 Phases Separate phases Increasing pressure More pressure Supercritcal CO 2 Figure 13.41