1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Liquids and Solids Chapter 10

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary. 2 Phases Solid phase - ice Liquid phase - water

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Schematic representations of the three states of matter.

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4

5. 5 Intermolecular Forces Intermolecular forces are attractive forces between molecules. Intramolecular forces hold atoms together in a molecule. Intermolecular vs Intramolecular 41 kJ to vaporize 1 mole of water (inter) 930 kJ to break all O-H bonds in 1 mole of water (intra) Generally, intermolecular forces are much weaker than intramolecular forces. “Measure” of intermolecular force boiling point melting point  H vap  H fus  H sub

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 bonded to a highly electronegative atom (F, O, N).

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Intermolecular Forces Dipole-Dipole Forces Attractive forces between polar molecules Orientation of Polar Molecules in a Solid

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 (a)The electrostatic interaction of two polar molecules. (b) The interaction of many dipoles in a condensed state.

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 (a) The polar water molecule. (b) Hydrogen bonding among water molecules.

10. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Why is the hydrogen bond considered a “special” dipole-dipole interaction? Decreasing molar mass Decreasing boiling point Why this sudden increase? The boiling point represents the magnitude and type of bonding

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Intermolecular Forces Ion-Dipole Forces Attractive forces between an ion and a polar molecule Ion-Dipole Interaction

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 London Dispersion Forces 4 relatively weak forces that exist among noble gas atoms and nonpolar molecules (Ar, sugar, 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.

14. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 (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.

15. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Intermolecular Forces Dispersion Forces Continued Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted. Polarizability increases with: greater number of electrons more diffuse electron cloud Dispersion forces usually increase with molar mass.

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 S O O What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH 4 CH 4 is nonpolar: dispersion forces. SO 2 SO 2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO 2 molecules.

17. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 Properties of Liquids Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. Or The resistance to an increase in its surface area (polar molecules). Strong intermolecular forces High surface tension

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

19. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 More Properties of Liquids Capillary Action: Spontaneous rising of a liquid in a narrow tube. Nonpolar liquid mercury forms a convex meniscus in a glass tube, whereas polar water forms a concave meniscus.

20. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 More Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules Adhesion Cohesion

21. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 More Properties of Liquids Viscosity is a measure of a fluid’s resistance to flow. Strong intermolecular forces High viscosity

22. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 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).

23. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing Crystalline quartz (SiO 2 ) Non-crystalline quartz glass

24. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions. An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A unit cell is the basic (smallest) repeating structural unit of a crystalline solid. Unit Cell lattice point At lattice points: Atoms Molecules Ions Unit cells in 3 dimensions

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

28. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Shared by 8 unit cells Shared by 2 unit cells

29. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 1 atom/unit cell (8 x 1/8 = 1) 2 atoms/unit cell (8 x 1/8 + 1 = 2) 4 atoms/unit cell (8 x 1/8 + 6 x 1/2 = 4)

30. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Three cubic unit cells and the corresponding lattices.

31. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 RELATIONSHIP BETWEEN ATOMIC RADIUS AND EDGE LENGTH IN THREE DIFFERENT UNIT CELLS l = 2r c = 4r b 2 = l 2 + l 2 c 2 = l 2 + b 2 c 2 = 3s 2 c = l √3 = 4r l = 4r/√3 b = 4r b 2 = l 2 + l 2 (4r) 2 = 2 l 2 16r 2 = 2 l 2 l = r√8

32. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 RELATIONSHIP BETWEEN ATOMIC RADIUS AND EDGE LENGTH IN THREE DIFFERENT UNIT CELLS s = 2r c = 4r b 2 = s 2 + s 2 c 2 = s 2 + b 2 c 2 = 3s 2 c = s√3 = 4r s = 4r/√3 b = 4r b 2 = s 2 + s 2 (4r) 2 = 2s 2 16r 2 = 2s 2 s = r√8 SimpleBody-centered Face-centered

33. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Analysis of crystal structure using X-Ray Diffraction

34. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 Extra distance = BC + CD = 2d sin  = n (Bragg Equation)

35. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 Bragg Equation Used for analysis of crystal structures. n = 2d sin  d = distance between atoms n = an integer = wavelength of the x-rays

36. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 X rays of wavelength 0.154 nm are diffracted from a crystal at an angle of 14.17 0. Assuming that n = 1, what is the distance (in pm) between layers in the crystal? n = 2d sin  n = 1  = 14.17 0 = 0.154 nm = 154 pm d = n 2sin  = 1 x 154 pm 2 x sin14.17 = 77.0 pm 11.5

37. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 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).

38. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Types of Crystals Ionic Crystals Lattice points occupied by cations and anions Held together by electrostatic attraction Hard, brittle, high melting point Poor conductor of heat and electricity CsClZnSCaF 2

39. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Types of Crystals Atomic Crystals Lattice points occupied by atoms Held together by covalent bonds Hard, high melting point Poor conductor of heat and electricity diamond graphite carbon atoms

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

41. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 The closest packing arrangement of uniform spheres, (a) aba packing (b) abc packing.

42. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 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.

43. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 When spheres are packed in the abc arrangement, the unit cell is face-centered cubic.

44. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 The indicated sphere has 12 nearest neighbors.

45. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 409 pm. Calculate the density of silver. d = m V V = a 3 = (409 pm) 3 = 6.83 x 10 -23 cm 3 4 atoms/unit cell in a face-centered cubic cell m = 4 Ag atoms 107.9 g mole Ag x 1 mole Ag 6.022 x 10 23 atoms x = 7.17 x 10 -22 g d = m V 7.17 x 10 -22 g 6.83 x 10 -23 cm 3 = = 10.5 g/cm 3

46. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 Types of Crystals Molecular Crystals Lattice points occupied by molecules Held together by intermolecular forces Soft, low melting point Poor conductor of heat and electricity

47. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Types of Crystals Metallic Crystals Lattice points occupied by metal atoms Held together by metallic bonds Soft to hard, low to high melting point Good conductors of heat and electricity Cross Section of a Metallic Crystal nucleus & inner shell e - mobile “sea” of e -

48. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 Examples of three types of crystalline solids (a) An atomic solid. (b) An ionic solid (c) A molecular solid

49. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Types of Crystals

50. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50

51. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 51 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.

52. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 52 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.

53. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 The molecular orbital energy levels produced when various numbers of atomic orbitals interact.

54. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 54 (left) A representation of the energy levels (bands) in a magnesium crystal. (right) Crystals of magnesium grown from a vapor.

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

56. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 56 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.

57. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57 Two types of alloys: (a) substitutional (b) interstitial

58. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 58 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

59. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 59 The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

60. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 60 Partial representation of the molecular orbital energies in (a) diamond and (b) a typical metal.

61. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 61 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.

62. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 62 Graphite consists of layers of carbon atoms.

63. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 63 Computer-generated model of silica.

64. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 64 (top) The structure of quartz (empirical formula SiO2). Quartz contains chains of SiO4 tetrahedral (bottom) that share oxygen atoms.

65. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 65 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.

66. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 66 (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.

67. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 67 Energy-level diagrams for (a) an n-type semiconductor and (b) a p-type semiconductor.

68. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 68 The p-n junction involves the contact of a p-type and an n-type semiconductor.

69. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 69 Molecular Solids Discrete molecular units at each lattice position e.g. ice, solid CO 2, S 8, P 4. Covalent bond within molecule while weak dipole-dipole or London dispersion forces (LD) e.g. Ice (Hydrogen bond), CO 2, P 4, S 8 (LD) very reactive. As the size of the molecule increases, the LD become large causing many substances to be solids at 25 0 C.

70. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 70 A “steaming” piece of dry ice CO 2

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

72. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 72 Ionic Solids Ionic solids are stable, high melting point and held ions together by strong electrostatic forces between opposite ions. Structure of most binary ionic solids (NaCl) can be explained by the closest packing of spheres –Anion (large ion) packed in one of the closest packing (hcp or ccp) –Cation (small ion) fits in holes among the anions.

73. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 73 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.

74. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 74 (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.

75. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 75 (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.

76. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 76

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

78. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 78 Behavior of a liquid in a closed container.

79. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 79 The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation H 2 O (l) H 2 O (g) Rate of condensation Rate of evaporation = Dynamic Equilibrium

80. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 80 Before Evaporation At Equilibrium

81. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 81 (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.

82. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 82 (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.

83. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 83 Molar heat of vaporization (  H vap ) is the energy required to vaporize 1 mole of a liquid. ln P = -  H vap RT + C Clausius-Clapeyron Equation P = (equilibrium) vapor pressure T = temperature (K) R = gas constant (8.314 J/K mol)

84. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 84 ln P T1 P T2 =  H vap R 1 1 T2T2 T1T1

85. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 85 The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure. The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm.

86. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 86 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

87. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 87 Melting 11.8 Freezing H 2 O (s) H 2 O (l) The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium Sublimation

88. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 88 Molar heat of fusion (  H fus ) is the energy required to melt 1 mole of a solid substance.

89. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 89 The supercooling of water. The extent of supercooling is given by S.

90. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 90 Sublimation Deposition H 2 O (s) H 2 O (g)  H sub =  H fus +  H vap ( Hess’s Law) Molar heat of sublimation (  H sub ) is the energy required to sublime 1 mole of a solid.

91. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 91

92. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 92 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).

93. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 93 The phase diagram for water.

94. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 94 The critical temperature (T c ) is the temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure. The critical pressure (P c ) is the minimum pressure that must be applied to bring about liquefaction at the critical temperature.

95. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 95 Diagrams of various heating experiments on samples of water in a closed system. Negative Slope

96. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 96 The phase diagram for carbon dioxide. Positive Slope