1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Chemistry FIFTH EDITION by Steven S. Zumdahl University of Illinois

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Chemistry FIFTH EDITION Chapter 10 Liquids and Solids

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Figure 10.1 Schematic Representations of the Three States of Matter

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Bonds and intermolecular forces have one very fundamental thing in common, both are electrostatic forces of attractions. The primary difference between bonds and intermolecular forces is the locations of the attraction and the magnitudes of the attraction.

5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Intermolecular Forces Definition: Intermolecular Forces are electrostatic forces of attraction that exist between an area of negative charge on one molecule and an area of positive charge on a second molecule. Chemists tend to consider three fundamental types of bonding. Ionic bonding Covalent bonding Metallic bonding

6. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 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)

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 : Dipole-Dipole Forces: Only polar covalent molecules have the ability to form dipole- dipole attractions between molecules. Polar covalent molecules have positive ends and negative ends which attract each other.

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Figure 10.2 Dipole-Dipole Attractions

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9

10. 10

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12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Hydrogen Bonding : These occur between polar covalent molecules that possess a hydrogen bonded to an extremely electronegative element, specifically - N, O, and F.

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Figure 10.3 A Water Molecule

14. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Figure 10.4 The Boiling Points of the Covalent Hydrides of the Elements in Groups 4A, 5A, 6A, and 7A

15. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 NONPOLAR STRUCTURES Nonpolar systems lack partial charges. Yet, they are also held together by the electrostatic forces.. London: cause by induced dipoles London: cause by induced dipoles.

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 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.  the ease with which electron “ cloud ” of an atom can be distorted is called polarizability.

17. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 London Dispersion Force

18. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 London:

19. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Ap CHEM free response From the guru himself: “ …greater surface area/electrons and greater polarizability (and with it increased LDF's) will score the points. ”

20. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Figure 10.5 London Dispersion Forces

21. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 When chemical bonds break, it ’ s a chemical reaction 1 molecule  2 molecules When intermolecular forces break, it ’ s not a reaction 2 together  2 separated Intermolecular and Intramolecular forces  

22. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Approximate range of strengths of attractive forces ionic and covalent bonding ion-dipole forces dipole-dipole forces hydrogen bonding dispersion forces

23. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 IMF Behavior controlling magnitude of properties of a substance Surface tension Vapor pressure Viscosity Capillary action Boiling pt Melting pt Critical temp

24. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Defintions needed to understand behavior Surface tension- the strength of the molecules held together to form a surface on the liquid Vapor pressure- pressure of vapor escaping from the liquid- Viscosity Capillary action Boiling pt Melting pt Critical temp

25. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 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. Viscosity: Resistance to flow (molecules with large intermolecular forces).

26. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Figure 10.6 Molecules in a Liquid

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 COMMENTS All bonds are stronger than all intermolecular forces. Systems that use bonding only will have the strongest structures. Systems that have intermolecular forces will have weaker structures, where the strength will depend upon the type of intermolecular force being used.

29. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 SOLID PARTICLESBinding forces Expt Identification Ionic +/- ionsElectrostatic attraction conductivity of fused salt metallic + ionsElectrostatic attraction between atoms and ions conductivity of the solid Covalent network atomscovalenthigh melting pt., extreme hardness, etc. molecular moleculesVan Der Waals low melting pt., non-conductivity of fused salt, etc.

30. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 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).

31. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 Infinite array of positive and negative ions held in place by ionic bonding Ionic bonding (strong) Properties –high melting point, boiling point –solid is electrical insulator, liquid is conductor –hard –brittle ex: NaCl Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Na + Cl – Ionic Solid

32. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32

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34. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 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.

35. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 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

36. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 Figure 10.9 Three Cubic Unit Cells and the Corresponding Lattices

37. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 Bragg Equation Used for analysis of crystal structures. n = 2d sin  d = distance between atoms n = an integer = wavelength of the x-rays

38. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Figure 10.10 Interference of Light Rays

39. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Figure 10.11 Diagram to Support the Bragg Equation

40. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 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).

41. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 10.12 Examples of Three Types of Crystalline Solids

42. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.  hexagonal closest packed ( “ aba ” )  cubic closest packed ( “ abc ” )

43. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 Figure 10.13 The Closest Packing Arrangement of Uniform Spheres

44. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 Figure 10.14 Hexagonal Closest Packing

45. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 Figure 10.15 Cubic Closest Packing

46. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 Figure 10.16 The Indicated Sphere Has 12 Nearest Neighbors

47. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Figure 10.17 The Net Number of Spheres in a Face-Centered Cubic Unit Cell

48. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 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.

49. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Figure 10.18 The Electron Sea Model for Metals Postulates a Regular Array of Cations in a “ Sea ” of Valence Electrons

50. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50 Figure 10.19 The Molecular Orbital Energy Levels Produced When Various Numbers of Atomic Orbitals Interact

51. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 51 Figure 10.20 The Band Model for Magnesium

52. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 52 Metal Alloys 1. Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Zn Substances that have a mixture of elements and metallic properties.

53. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 Metal Alloys (continued) 2.Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon 3.Both types: Alloy steels contain a mix of substitutional (carbon) and interstitial (Cr, Mo) alloys.

54. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 54 Figure 10.21 Two Types of Alloys

55. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 55 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

56. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 56 Figure 10.22 The Structures of Diamond and Graphite

57. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57 Semiconductors 4 Conductivity is enhanced by doping with group 3a or group 5a elements. A substance in which some electrons can cross the band gap.

58. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 58 Figure 10.23 Partial Representation of the Molecular Orbital Energies in A) Diamond and B) a Typical Metal

59. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 59 Figure 10.24 The p Orbitals

60. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 60 Figure 10.26 The Structure of Quartz

61. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 61 Figure 10.27 Silicate Anions

62. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 62 Figure 10.28 Two Dimensional Representations of (a) a Quartz Crystal and (b) a Quartz Glass

63. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 63 Figure 10.29 Silicon Crystal Doped with (a) Arsenic and (b) Boron

64. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 64 Figure 10.31 The p-n Junction

65. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 65 Figure 10.30 Energy Level Diagrams for (a) an n-Type Semiconductor and (b) a p-Type Semiconductor

66. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 66 Figure 10.50 A Schematic of Two Circuits Connected by a Transistor

67. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 67 Figure 10.51 The Steps for Forming a Transistor in a Crystal of Initially Pure Silicon

68. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 68 Figure 10.33 The Holes that Exist Among Closest Packed Uniform Spheres

69. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 69 Figure 10.34 The Position of Tetrahedral Holes in a Face-Centered Cubic Unit Cell

70. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 70 Figure 10.35 Cubic Closest Packing in NaCl

71. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 71 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.

72. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 72 Figure 10.36 Behavior of a Liquid in a Closed Container

73. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 73 Figure 10.37 The Rates of Condensation and Evaporation

74. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 74 Figure 10.38 Vapor Pressure

75. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 75 Figure 10.39 The Number of Molecules in a Liquid With a Given Energy Versus Kinetic Energy at Two Temperatures

76. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 76 Figure 10.40 The Vapor Pressure of Water

77. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 77 Melting Point 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

78. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 78 Boiling Point Constant temperature when added energy is used to vaporize the liquid. vapor pressure of liquid = pressure of surrounding atmosphere

79. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 79 Figure 10.42 Heating Curve for Water

80. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 80 Figure 10.43 The Vapor Pressures of Solid and Liquid Water

81. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 81 Figure 10.44 An Apparatus that Allows Solid and Liquid Water to Interact Only Through the Vapor State

82. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 82 Figure 10.45 Water in a Closed System

83. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 83 Figure 10.46 The Supercooling of Water

84. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 84 Phase Diagram Represents phases as a function of temperature and pressure. 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).

85. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 85 Figure 10.47 The Phase Diagram for Water

86. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 86 Figure 10.48 Diagrams of Various Heating Experiments

87. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 87 Figure 10.49 The Phase Diagram for Water

88. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 88 Figure 10.52 The Phase Diagram for Carbon Dioxide