1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Intermolecular Forces Forces between (rather than within) molecules.  dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “  ” ends of the dipoles are close to each other. Ô hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

2. Figure 10.2: (a) The electrostatic interaction of two polar molecules. (b) The interaction of many dipoles in a condensed state.

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Figure 10.3: (a) The polar water molecule. (b) Hydrogen bonding among water molecules.

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Figure 10.4: The boiling points of the covalent hydrides of the elements in Groups 4A, 5A, 6A, and 7A. General trend to increase down the group Small size Large electronegativity

5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 London Dispersion Forces 4 relatively weak forces that exist among noble gas atoms and nonpolar molecules. (Ar, C 8 H 18 ) 4 caused by instantaneous dipole, in which electron distribution becomes asymmetrical. 4 the ease with which electron “cloud” of an atom can be distorted is called polarizability.

6. Figure 10.5: (a) An instantaneous polarization can occur on atom A, creating an instantaneous dipole. This dipole creates an induced dipole on neighboring atom B. (b) Nonpolar molecules such as H2 also can develop instantaneous and induced dipoles.

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7

8. 8 Some Properties of a Liquid Surface Tension: The resistance to an increase in its surface area (polar molecules). Capillary Action: Spontaneous rising of a liquid in a narrow tube. Adhesion- stick to walls Cohesion- stick to each other Viscosity: Resistance to flow (molecules with large intermolecular forces).

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Figure 10.6: A molecule in the interior of a liquid is attracted by the molecules surrounding it, whereas a molecule at the surface of a liquid is attracted only by molecules below it and on each side.

10. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Figure 10.7: Nonpolar liquid mercury forms a convex meniscus in a glass tube, whereas polar water forms a concave meniscus. Which has stronger cohesive forces?water

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Figure 10.8: Two crystalline solids: pyrite (left), amethyst (right).

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Types of Solids Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS 2 )]. Amorphous solids: considerable disorder in their structures (glass).

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

14. TiO 2 Crystal lattice of alternating positive and negative ions. CsClNaCl

15. Coordination number of an ion is the number of ions of opposite charge That surround the ion in a crystal Sodium Chloride: Each Na + is surrounded by 6 Cl -, therefore the Coordination # = 6 Each Cl - is surrounded by 6 Na + Cubic centered

17. Metallic Crystals  Metals are made up of closely packed cations rather than neutral atoms.  Chemical bonding results from the attraction between metal ions and the surrounding sea of electrons.  Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move about the metal.

18. Because of this arrangement  Metals are good conductors of heat and electricity  Metals are malleable  Metals are ductile  Metals have high tensile strength  Metals have luster

19. Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

20. Every atom has 8 neighbors Every atom has 12 neighbors Every atom has 12 neighbors

21. Metal Alloys  Substitutional Alloy: some metal atoms replaced by others of similar size.

22. Metal Alloys  Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.

23. Figure 10.9: Three cubic unit cells and the corresponding lattices.

24. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Figure 10.12: Examples of three types of crystalline solids. (a) An atomic solid. (b) An ionic solid. (c) A molecular solid.

25. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 Representation of Components in a Crystalline Solid Unit Cell: The smallest repeating unit of the lattice. 4 simple cubic 4 body-centered cubic 4 face-centered cubic

26. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Types of Crystalline Solids Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl). Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice). Atomic Solid: have atoms at the lattice points. metallic, network, and group 8A solids

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27

28. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Figure 10.13: The closest packing arrangement of uniform spheres, (a) aba packing (b) abc packing.

29. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Figure 10.14: When spheres are closest packed so that the spheres in the third layer are directly over those in the first layer (aba), the unit cell is the hexagonal prism illustrated here in red.

30. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Figure 10.15: When spheres are packed in the abc arrangement, the unit cell is face-centered cubic.

31. Figure 10.16: The indicated sphere has 12 nearest neighbors.

32. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Figure 10.17: The net number of spheres in a face-centered cubic unit cell.

33. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Crystalline silver contains cubic closest packed silver atoms.

34. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 Figure 10.18: The electron sea model for metals postulates a regular array of cations in a "sea" of valence electrons. (a) Representation of an alkali metal (Group 1A) with one valence electron. (b) Representation of an alkaline earth metal (Group 2A) with two valence electrons.

35. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors. 4 hexagonal closest packed (“aba”) 4 cubic closest packed (“abc”)

36. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 Bonding Models for Metals Electron Sea Model: A regular array of metals in a “sea” of electrons. Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.

37. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 Figure 10.20: (left) A representation of the energy levels (bands) in a magnesium crystal. (right) Crystals of magnesium grown from a vapor.

38. Figure 10.21: Two types of alloys.

39. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Metal Alloys 1. Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Zn (1/3 of Cu replaced by Zn) sterling silver= Ag/Cu pewter=Sn/Cu.Bi.At Substances that have a mixture of elements and metallic properties.

40. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 Metal Alloys (continued) 2.Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon (change from indirectional to directional bonds) 3.Both types: Alloy steels contain a mix of substitutional (carbon) and interstitial (Cr, Mo) alloys.

41. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Network Solids Composed of strong directional covalent bonds that are best viewed as a “giant molecule”. 4 brittle 4 do not conduct heat or electricity 4 carbon, silicon-based graphite, diamond, ceramics, glass

42. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 Diamond (sp 3 ) tetrahedral arrangement hard, colorless, an insulator Graphite (sp 2 ) layers of six member rings slippery, black, conductor (delocalization ofthe π bonds)

43. Figure 10.22: The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

44. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 Figure 10.23: Partial representation of the molecular orbital energies in (a) diamond and (b) a typical metal.

45. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 Figure 10.24: The p orbitals (a) perpendicular to the plane of the carbon ring system in graphite can combine to form (b) an extensive π-bonding network.

46. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 Figure 10.25: Graphite consists of layers of carbon atoms.

47. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Silicon Silicon-oxygen compound SiO 4 silica (tetrahedral) Silicates- found in rocks, soil, and clay. Silica is heated and cooled rapidly- glass

48. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 Computer-generated model of silica.

49. Figure 10.26: (top) The structure of quartz (empirical formula SiO 2 ). Quartz contains chains of SiO 4 tetrahedra (bottom) that share oxygen atoms.

50. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50 Figure 10.27: Examples of silicate anions, all of which are based on SiO 4 4- tetrahedra.

51. Figure 10.28: Two- dimensional representations of (a) a quartz crystal and (b) a quartz glass.

52. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 52 ceramics From clay (contain silicates, like glass) and firing at high temp. non metallic, strong, brittle, and resistant to heat and chemicals Glass- homogeneous Ceramics- heterogeneous ( crystal of silicates suspended in a glassy cement)

53. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 Semiconductors 4 Conductivity is enhanced by doping with group 3a or group 5a elements. 4 N-type (As; has 1 more e - ) 4 P-type (B; has 1 less e - ) A substance in which some electrons can cross the band gap. Conductivity increases w/increase in temp.

54. Figure 10.29: (a) A silicon crystal doped with arsenic, which has one more valence electron than silicon. (b) A silicon crystal doped with boron, which has one less electron than silicon.

55. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 55 Figure 10.30: Energy-level diagrams for (a) an n-type semiconductor and (b) a p- type semiconductor.

56. Figure 10.31: The p-n junction involves the contact of a p-type and an n-type semiconductor.

57. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57 Figure 10.32: A schematic of two circuits connected by a transistor. The signal in circuit 1 is amplified in circuit 2.

58. Figure 10.33: The steps for forming a transistor in a crystal of initially pure silicon.

59. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 59 Molecular Solids Ice- lattice is occupied by H 2 O molecules Dry ice- solid CO 2 Sulfur- contain S 8 Phosphorus- P 4 (white phosphorus) CO 2, I 2, S 8, P 4 (London dispersion) (gas)

60. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 60 A “steaming” piece of dry ice.

61. Figure 10.34: (a) Sulfur crystals (yellow) contain S 8 molecules. (b) White phosphorus (containing P 4 molecules) is so reactive with the oxygen in air that it must be stored under water.

62. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 62

63. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 63 Ionic Solids Stable, high melting point substances, held together by strong electrostatic forces between oppositely charged ions. Closest packing model, where the smaller cation fits in the holes.

64. Figure 10.35: The holes that exist among closest packed uniform spheres. (a) The trigonal hole formed by three spheres in a given plane. (b) The tetrahedral hole formed when a sphere occupies a dimple formed by three spheres in an adjacent layer. (c) The octahedral hole formed by six spheres in two adjacent layers.

65. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 65 Trigonal hole never reached by ionic solids.

66. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 66 Figure 10.36: (a) The location (X) of a tetrahedral hole in the face-centered cubic unit cell. (b) One of the tetrahedral holes. (c) The unit cell for ZnS where the S 2- ions (yellow) are closest packed with the Zn 2+ ions (red) in alternating tetrahedral holes.

67. Figure 10.37: (a) The locations (gray X) of the octahedral holes in the face-centered cubic unit cell. (b) Representation of the unit cell for solid NaCl.

68. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 68

69. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 69 Vapor Pressure a nd Changes of State Molecules of a liquid escape the liquid’s surface and form a gas- Vaporization or Evaporation. Endothermic (overcome intermolecular forces)

70. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 70 Heat of vaporization, or enthalpy of vaporization (ΔH vap ) Energy required to vaporize 1mol of liquid at 1atm. H 2 O has a large heat of vaporization(40.7kJ)

71. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 71 Figure 10.38: Behavior of a liquid in a closed container. Initial transfer of liquid to the vapor phase. Initially the # of vapor molecules increase and so does the rate of return. (condensation). Eventually equilibrium occurs

72. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 72 Figure 10.39: The rates of condensation and evaporation over time for a liquid sealed in a closed container.

73. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 73 Vapor Pressure... is the pressure of the vapor present at equilibrium.... is determined principally by the size of the intermolecular forces in the liquid.... increases significantly with temperature. Volatile liquids have high vapor pressures.

74. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 74 Figure 10.40: (a) The vapor pressure of a liquid can be measured easily using a simple barometer of the type shown here. (b) The three liquids, water, ethanol (C 2 H 5 OH), and diethyl ether [(C 2 H 5 ) 2 O], have quite different vapor pressures.

75. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 75 In general substances with large molar masses have low vapor pressures (Large dispersion forces) ( the more e - the more polarizable) H 2 O smaller molar mass but has H-bonding (stronger!!!)

76. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 76 Figure 10.41: Diagram showing the reason vapor pressure depends on temperature. Vapor Pressure increases w/increased Temp.

77. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 77 Figure 10.42: (a) The vapor pressure of water, ethanol, and diethyl ether as a function of temperature. (b) Plots of In(P vap ) versus 1/T (Kelvin temperature) for water, ethanol, and diethyl ether.

78. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 78

79. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 79 Figure 10.43: Iodine being heated, causing it to sublime onto an evaporating dish cooled by ice. From solid to a gas

80. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 80 Figure 10.44: The heating curve (not drawn to scale) for a given quantity of water where energy is added at a constant rate.

81. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 81 Melting Point Vibration of water molecules increases as T increases. Molecules break loose from lattice points and solid changes to liquid. (Temperature is constant as melting occurs.) vapor pressure of solid = vapor pressure of liquid Plateau at 0°C.

82. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 82 Heat of fusion, enthalpy of fusion ΔH fus Enthalpy change that occurs at MP

83. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 83

84. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 84 Boiling Point At 100°C the liquid water reaches BP. Constant temperature when added energy is used to vaporize the liquid. vapor pressure of liquid = pressure of surrounding atmosphere ***All intermolecular changes (physical) no intermolecular bonds broken.

85. Figure 10.46: Solid and liquid phases in equilibrium with the vapor pressure.

86. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 86 Normal melting point- T at which the solid and liquid states have the same vapor pressure under conditions were the total pressure is 1atm. Normal boiling point- T at which the vapor pressure of the liquid is exactly 1atm

87. Figure 10.47: Water in a closed system with a pressure of 1 atm exerted on the piston. No bubbles can form within the liquid as long as the vapor pressure is less than 1 atm.

88. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 88 Super cooling Cooled below O°C at 1atm and remain in the liquid state. As it is cooled it may not reach the degree of organization necessary to form ice at O°C.

89. Figure 10.48: The supercooling of water. The extent of supercooling is given by S.

90. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 90 Boiling chip releasing air bubbles acts as a nucleating agent for the bubbles that form when water boils.

91. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 91 Figure 10.49: The phase diagram for water.

92. Figure 10.47: Water in a closed system with a pressure of 1 atm exerted on the piston. No bubbles can form within the liquid as long as the vapor pressure is less than 1 atm.

93. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 93 Pressure at 1atm Filled w/ice (no air space) T=-20°C Piston= 1atm pressure Since T<0, VP of ice is <1atm, no vapor present As heated ice is only component until T=0 VP(solid)=VP(liquid) melting Still no vapor ; P <1atm

94. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 94 Solid completely to liquid, temp changes Now only water (still no vapor) T=100C ; VP= 1atm; boiling occurs; change to vapor Liquid completely to vapor, temp rises.

95. Figure 10.50: Diagrams of various heating experiments on samples of water in a closed system.

96. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 96 Experiment 2 P=2torr T= -20 Just ice T increases to -10, sublimation occurs VP of ice = external pressure No liquid because the VP of liquid is always greater than 2torr

97. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 97 Experiment 3 Pressure 4.58 torr T= -20, only ice New phase when T=0.01°C; Triple Point solid and liquid water have identical VP of 4.58. All three phases exist.

98. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 98 Experiment 4 P=225atm T=300°C ; liquid As T increases the liquid changes to vapor, goes through intermediate “fluid” region.

99. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 99 Phase Diagram critical temperature: temperature above which the vapor can not be liquefied. critical pressure: pressure required to liquefy AT the critical temperature. critical point: critical temperature and pressure (for water, T c = 374°C and 218 atm).

100. Figure 10.51: The phase diagram for water. At point X on the phase diagram, water is a solid.

101. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 101

102. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 102 Figure 10.52: The phase diagram for carbon dioxide.