1. John E. McMurry Robert C. Fay Lecture Notes Alan D. Earhart Southeast Community College Lincoln, NE General Chemistry: Atoms First Chapter 10 Liquids, Solids, and Phase Changes Copyright © 2010 Pearson Prentice Hall, Inc.

2. Chapter 10/2

3. Chapter 10/3 Polar Covalent Bonds and Dipole Moments

4. Chapter 10/4 Polar Covalent Bonds and Dipole Moments C-Cl bond has a bond dipole because of a difference in electronegativities.

5. Chapter 10/5 Polar Covalent Bonds and Dipole Moments

6. Chapter 10/6 Polar Covalent Bonds and Dipole Moments The individual bond polarities do not cancel. Therefore, the molecule has a dipole moment. In other words, the molecule is polar.

7. Chapter 10/7 Polar Covalent Bonds and Dipole Moments The individual bond polarities cancel. Therefore, the molecule does not have a dipole moment. In other words, the molecule is nonpolar.

8. Chapter 10/8

9. Chapter 10/9 Intermolecular Forces Intermolecular Forces: Attractions between “molecules” that hold them together. These forces are electrical in origin and result from the mutual attraction of unlike charges or the mutual repulsion of like charges. Types of Intermolecular Forces: Ion-Dipole Forces Van der Waals Forces dipole-dipole forces London dispersion forces hydrogen bonds

10. Chapter 10/10 Ion-Dipole Forces: The result of electrical interactions between an ion and the partial charges on a polar molecule. Intermolecular Forces

11. Chapter 10/11 Dipole-Dipole Forces: The result of electrical interactions between dipoles on neighboring molecules. Intermolecular Forces

12. Chapter 10/12 Dipole-Dipole Forces Intermolecular Forces As the dipole forces increase the intermolecular forces increase. As the intermolecular forces increase, the boiling points increase.

13. London Dispersion Forces: The result of motion of electrons which gives the molecule a short-lived dipole moment which induces temporary dipoles in neighboring molecules. Intermolecular Forces

14. Chapter 10/14 Intermolecular Forces London Dispersion Forces As the dispersion forces increase the intermolecular forces increase. As the intermolecular forces increase, the boiling points increase.

15. Intermolecular Forces London Dispersion Forces

16. Chapter 10/16 Intermolecular Forces Hydrogen Bond: An attractive force between a hydrogen atom bonded to a very electronegative atom (O, N, or F) and an unshared electron pair on another electronegative atom.

17. Chapter 10/17 Intermolecular Forces Hydrogen Bond

19. Intermolecular Forces Hydrogen Bond

20. Chapter 10/20 Intermolecular Forces

21. Some Properties of Liquids

23. Chapter 10/23 Some Properties of Liquids

24. Chapter 10/24 Phase Changes Phase Change (State Change): A change in physical form but not the chemical identity of a substance. liquid to solid gas to liquid gas to solid Freezing: Condensation: Deposition: solid to liquid liquid to gas solid to gas Fusion (melting): Vaporization: Sublimation:

25. Chapter 10/25

26. Chapter 10/26 Phase Changes Enthalpy (Heat) of Fusion (∆H fusion ): The amount of energy required to convert a solid into a liquid. Enthalpy (Heat) of Vaporization (∆H vap ): The amount of energy required to convert a liquid into a gas.

27. Phase Changes Heating Curve for Water

28. Chapter 10/28 Phase Changes

29. Evaporation, Vapor Pressure, and Boiling Point Vapor Pressure: The partial pressure of a gas in equilibrium with liquid at a constant temperature.

30. Chapter 10/30 Evaporation, Vapor Pressure, and Boiling Point Clausius-Clapeyron Equation + C 1 T - ∆H vap R ln P vap = mxb+y=

32. Chapter 10/32 Kinds of Solids Amorphous Solids: Particles are randomly arranged and have no ordered long-range structure. Example - rubber. Crystalline Solids: Particles have an ordered arrangement extending over a long range. ionic solids molecular solids covalent network solids metallic solids

33. Chapter 10/33

34. Kinds of Solids Ionic Solids: Particles are ions ordered in a regular three- dimensional arrangement and held together by ionic bonds. Example - sodium chloride.

35. Kinds of Solids Molecular Solids: Particles are molecules held together by intermolecular forces. Example - ice.

36. Kinds of Solids Covalent Network Solids: Particles are atoms linked together by covalent bonds into a giant three-dimensional array. Example: quartz

37. Chapter 10/37 Kinds of Solids Metallic Solids: Particles are metal atoms whose crystals have metallic properties such as electrical conductivity. Example - iron. Summary

39. Probing the Structure of Solids: X-Ray Crystallography Diffraction: Occurs when electromagnetic radiation is scattered by an object containing regularly spaced lines (such as a diffraction grating) or points (such as the atoms in a crystal). Interference: Occurs when two waves pass through the same region of space.

40. Probing the Structure of Solids: X-Ray Crystallography Bragg Equation: d = 2 sin  n

41. Unit Cells and the Packing of Spheres in Crystalline Solids Body-Centered Cubic Packing Simple Cubic Packing

42. Unit Cells and the Packing of Spheres in Crystalline Solids Cubic Closest- Packing Hexagonal Closest- Packing

43. Chapter 10/43 Unit Cells and the Packing of Spheres in Crystalline Solids Unit Cell: A small repeating unit that makes up a crystal.

44. Unit Cells and the Packing of Spheres in Crystalline Solids Body-Centered Cubic Primitive Cubic

45. Unit Cells and the Packing of Spheres in Crystalline Solids Face-Centered Cubic

46. Chapter 10/46 Unit Cells and the Packing of Spheres in Crystalline Solids

47. Chapter 10/47 Structures of Some Ionic Solids

49. Chapter 10/49 Structures of Some Covalent Network Solids Carbon Allotropes Allotropes: Different structural forms of an element.

50. Structures of Some Covalent Network Solids Carbon Allotropes

51. Structures of Some Covalent Network Solids Carbon Allotropes

52. Chapter 10/52 Structures of Some Covalent Network Solids Silica (SiO 2 ) quartz sand quartz glass

53. Chapter 10/53

54. Chapter 10/54 Phase Diagrams Normal: Occurs at 1 atm. Critical Point: A combination of temperature and pressure beyond which a gas cannot be liquefied. Critical Temperature: The temperature beyond which a gas cannot be liquefied regardless of the pressure. Critical Pressure: The pressure beyond which a liquid cannot be vaporized regardless of the temperature. Supercritical Fluid: A state of matter beyond the critical point that is neither liquid nor gas. Triple Point: A point at which three phases coexist in equilibrium.

55. Phase Diagrams Carbon Dioxide

56. Phase Diagrams Water