1. 1 Intermolecular Forces and Liquids and Solids Chapter 11 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. CONTENTS 2 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

3. 1.The kinetic molecular theory of liquids and solids 3 Kinetic energy attractive forces between molecules gaseous state Kinetic energy attractive forces between molecules (atoms or ions) condensed(liquids and solids) state ＞ ＜

4. CONTENTS 4 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

5. 2.Chemical bonds and Intermolecular forces 5 Chemical bonds (intramolecular forces) Metallic bonds Ionic bonds Covalent bonds

6. 6 Comparing Intermolecular with intramolecular Forces Intermolecular forces are attractive forces between molecules. Intramolecular forces hold atoms(ions) 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

7. 7 Types of Intermolecular Forces Intermolecular Forces Van der Waals forces Ion-dipole forces Hydrogen bonding

8. 8 Types of van der Waals Forces van der Waals Forces Dipole-dipole forces Dipole-induced dipole forces Dispersion forces

9. 9 Types of van der Waals Forces(1) Dipole-Dipole Forces Attractive forces between polar molecules Orientation of Polar Molecules in a Solid

10. 10 Types of van der Waals Forces(2) Induced diple ion-induced dipole interaction (not van der Waals forces) dipole-induced dipole interaction (van der Waals forces)

11. 11 Types of van der Waals Forces(3) Dispersion Forces Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules Induced Dipoles Interacting With Each Other Instantaneous dipole(last for just a tiny fraction of a second)

12. 12 Types of van der Waals Forces(3) 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.

13. 13 Intermolecular Forces Ion-Dipole Forces Attractive forces between an ion and a polar molecule Ion-Dipole Interaction

14. 14 in solution Interaction Between Water and Cations

15. 15 Intermolecular Forces Hydrogen Bond The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H … B A H … A or A & B are N, O, or F

16. 16 Hydrogen Bond HCOOH and water

17. 17 Why is the hydrogen bond considered a “special” dipole-dipole interaction? Decreasing molar mass Decreasing boiling point

18. 18 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. Example and exercise

19. CONTENTS 19 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

20. 20 3.Properties of Liquids Intermolecular forces contribute Surface tension viscosity To liquids

21. 21 Properties of Liquids(1) Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. Strong intermolecular forces High surface tension

22. 22 Example of surface tension-capillary action Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules Adhesion Cohesion

23. 23 Properties of Liquids(2) Viscosity is a measure of a fluid’s resistance to flow. Strong intermolecular forces High viscosity

24. 24 Density of Water Maximum Density 4 0 C Ice is less dense than water Water is a Unique Substance3-D Structure of Water Structure and properties of water

25. CONTENTS 25 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

26. 4.Crystal structure 26 Types of solids Crystalline solids Amorphous solids

27. The general structure of crystalline The types of unit cells Packing spheres Closest packing 27 subtitles

28. 28 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 repeating structural unit of a crystalline solid. lattice point Unit Cell Unit cells in 3 dimensions At lattice points: Atoms Molecules Ions The general structure of crystalline

29. 29 Seven Basic Unit Cells ( 七大晶系 ) The types of unit cells 立方 ( 等軸 ) 四方 ( 正方 ) 正交 ( 斜方 ) 三方 單斜三斜 六方

30. 30 Three Types of Cubic Unit Cells Packing spheres

31. 31 Arrangement of Identical Spheres in a Simple Cubic Cell

32. 32 Arrangement of Identical Spheres in a Body-Centered Cubic Cell

33. 33 Shared by 8 unit cells Shared by 4 unit cells A Corner Atom, a Edge-Centered Atom and a Face-Centered Atom Shared by 2 unit cells Lattice point Edge Face

34. 34 Number of Atoms Per Unit Cell 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)

35. 35 Relation Between Edge Length and Atomic Radius

36. 36 Closest packing Hexagonal (hcp)Cubic (ccp)

37. 37 Exploded Views

38. 38 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 Example : application of information of cubic unit cell

39. CONTENTS 39 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

40. 40 An Arrangement for Obtaining the X-ray Diffraction Pattern of a Crystal.

41. 41 Extra distance = BC + CD = 2d sin  = n (Bragg Equation) Reflection of X rays from Two Layers of Atoms.

42. 42 X rays of wavelength 0.154 nm are diffracted from a crystal at an angle of 14.17 o. Assuming that n = 1, what is the distance (in pm) between layers in the crystal? n = 2d sin  n = 1  = 14.17 o = 0.154 nm = 154 pm d = n 2sin  = 1 x 154 pm 2 x sin14.17 = 314.0 pm

43. CONTENTS 43 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

44. 44 Chemical bonds or intermolecular forces Types of crystals Ionic bond covalent bond Intermolecular forces Metallic bond Ionic crystals Covalent crystals Molecular crystals Metallic crystals

45. 45 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

46. 46 Types of Crystals Covalent 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

47. 47 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 waterbenzene

48. 48 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 -

49. 49 Crystal Structures of Metals

50. 50 Types of Crystals and General Properties

51. CONTENTS 51 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

52. 52 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

53. 53 Chemistry In Action: High-Temperature Superconductors MgB 2 YBa 2 Cu 3 O X X=6 or 7 Superconducting around 95 K Coolant: liquid nitrogen Superconducting at about 40 K Coolant: liquid neon

54. 54 Chemistry In Action: And All for the Want of a Button white tin grey tin T < 13 0 C stableweak

55. CONTENTS 55 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

56. subtitles 56 Liquid-vapor equilibrium vapor pressure molar heat of vaporization and boiling point critical temperature and pressure Liquid-solid equilibrium molar heat of fusion Solid-vapor equilibrium molar heat of sublimation

57. 57 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

58. 58 Greatest Order Least Order Phase Changes : transformation from one phase to another, occur when energy is added or removed

59. 59 Evaporation : the process in which a liquid is transformed into a gas Liquid gas Evaporation condensation

60. 60 The factors influence the evaporation rate Factors Intermolecular forces (Attractive force evaporation ) The kinetic energy of molecules (depend on temperature)

61. 61 T 2 > T 1 Effect of Temperature on Kinetic Energy T evaporation

62. 62 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

63. 63 Before Evaporation At Equilibrium Measurement of Vapor Pressure

64. 64 Molar heat of vaporization (  H vap ) is the energy required to vaporize 1 mole of a liquid at its boiling point. ln P = -  H vap RT + C Clausius-Clapeyron Equation P = (equilibrium) vapor pressure T = temperature (K) R = gas constant (8.314 J/K mol) Vapor Pressure Versus Temperature T vapor pressure Steeper line represents larger  H vap

65. 65 Alternate Forms of the Clausius-Clapeyron Equation At two temperatures or

66. 66 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.

67. 67 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.

68. 68 The Critical Phenomenon of SF 6 T < T c T > T c T ~ T c T < T c

69. 69 The relationship between intermolecular forces and physical properties of liquids In general Intermolecular forces Vapor pressure Boiling temperature Critical temperature  H vap

70. 70 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 Solid-Liquid Equilibrium

71. 71 Molar heat of fusion (  H fus ) is the energy required to melt 1 mole of a solid substance at its freezing point.

72. 72 Heating Curve

73. 73 H 2 O (s) H 2 O (g) Molar heat of sublimation (  H sub ) is the energy required to sublime 1 mole of a solid.  H sub =  H fus +  H vap ( Hess’s Law) Solid-Gas Equilibrium

74. CONTENTS 74 1.The kinetic molecular theory of liquids and solids 2.Chemical bonds and Intermolecular forces 3.Properties of liquids 4.Crystal structure 5.X-Ray diffraction by crystals 6.Types of crystals 7.Amorphous solids 8.Phase changes 9.Phase diagrams

75. 75 A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas. Phase Diagram of Water

76. 76 Phase Diagram of Carbon Dioxide At 1 atm CO 2 (s) CO 2 (g)

77. 77 Effect of Increase in Pressure on the Melting Point of Ice and the Boiling Point of Water

78. 78 Chemistry In Action: Ice Skating

79. 79 Chemistry In Action: Liquid Crystals