This document discusses different states of matter and properties of liquids and solids. It defines key terms like phases, intermolecular forces, and boiling point. It describes different types of solids like ionic, molecular, metallic and network solids. It also discusses properties of liquids like surface tension, capillary action, and viscosity.

1. Liquids and Solids Chapter 10

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. Schematic representations of the three states of matter.

7. Intermolecular Forces Dipole-Dipole Forces Attractive forces between polar molecules Orientation of Polar Molecules in a Solid

9. (a) The polar water molecule. (b) Hydrogen bonding among water molecules.

11. Why is the hydrogen bond considered a “special” dipole-dipole interaction? Why this sudden increase? The boiling point represents the magnitude and type of bonding Decreasing molar mass Decreasing boiling point

12. Intermolecular Forces Ion-Dipole Forces Attractive forces between an ion and a polar molecule Ion-Dipole Interaction

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.

16. 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. S O O

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. 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.

20. More Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules Adhesion Cohesion

21. More Properties of Liquids Viscosity is a measure of a fluid’s resistance to flow. Strong intermolecular forces High viscosity

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

28. Shared by 8 unit cells Shared by 2 unit cells

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. Three cubic unit cells and the corresponding lattices.

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. 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 Simple Body-centered Face-centered

33. Analysis of crystal structure using X-Ray Diffraction

34. BC + CD = 2 d sin  = n  (Bragg Equation) Extra distance =

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  = 2 d sin  n = 1  = 14.17 0  = 0.154 nm = 154 pm d = = 77.0 pm 11.5 n  2sin  = 1 x 154 pm 2 x sin14.17

41. The closest packing arrangement of uniform spheres, (a) aba packing (b) abc packing.

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. When spheres are packed in the abc arrangement, the unit cell is face-centered cubic.

44. The indicated sphere has 12 nearest neighbors.

45. When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 409 pm. Calculate the density of silver. 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 = 7.17 x 10 -22 g = 10.5 g/cm 3 d = m V 107.9 g mole Ag x 1 mole Ag 6.022 x 10 23 atoms x d = m V 7.17 x 10 -22 g 6.83 x 10 -23 cm 3 =

48. Examples of three types of crystalline solids (a) An atomic solid. (b) An ionic solid (c) A molecular solid

49. Types of Crystals

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. The molecular orbital energy levels produced when various numbers of atomic orbitals interact.

54. (left) A representation of the energy levels (bands) in a magnesium crystal. (right) Crystals of magnesium grown from a vapor.

57. Two types of alloys: (a) substitutional (b) interstitial

59. The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

60. Partial representation of the molecular orbital energies in (a) diamond and (b) a typical metal.

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. Graphite consists of layers of carbon atoms.

63. Computer-generated model of silica.

64. (top) The structure of quartz (empirical formula SiO2). Quartz contains chains of SiO4 tetrahedral (bottom) that share oxygen atoms.

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. Energy-level diagrams for (a) an n-type semiconductor and (b) a p-type semiconductor.

68. The p-n junction involves the contact of a p-type and an n-type semiconductor.

70. A “steaming” piece of dry ice CO 2

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.

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. (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. (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.

78. Behavior of a liquid in a closed container.

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. Before Evaporation At Equilibrium

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. (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. Molar heat of vaporization (  H vap ) is the energy required to vaporize 1 mole of a liquid. P = (equilibrium) vapor pressure T = temperature (K) R = gas constant (8.314 J/K • mol) ln P = -  H vap RT + C Clausius-Clapeyron Equation

84. ln P T1 P T2 =  H vap R 1 1 T 2 T 1

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.

87. Melting 11.8 Freezing 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 H 2 O ( s ) H 2 O ( l )

88. Molar heat of fusion (  H fus ) is the energy required to melt 1 mole of a solid substance.

89. The supercooling of water. The extent of supercooling is given by S .

90. Sublimation Deposition  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. H 2 O ( s ) H 2 O ( g )

93. The phase diagram for water.

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. Diagrams of various heating experiments on samples of water in a closed system. Negative Slope

96. The phase diagram for carbon dioxide. Positive Slope