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CHAPTER 12

CHAPTER 12. Liquids and Solids. Intermolecular Forces. Forces of attraction between different molecules rather than bonding forces within the same molecule. Dipole-dipole attraction Hydrogen bonds Dispersion forces. Dipole-Dipole Attraction.

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CHAPTER 12

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  1. CHAPTER 12 Liquids and Solids

  2. Intermolecular Forces Forces of attraction between different molecules rather than bonding forces within the same molecule. • Dipole-dipole attraction • Hydrogen bonds • Dispersion forces

  3. Dipole-Dipole Attraction • dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “-” ends of the dipoles are close to each other.

  4. Dipole Forces

  5. Hydrogen Bonding • hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

  6. Hydrogen Bonding in Water

  7. Hydrogen Bonding in DNA

  8. London Dispersion Forces • Dispersion forces: relatively weakforces caused by instantaneous dipole, in which electron distribution becomes asymmetrical. • The ONLY forces of attractionthat exist among noble gas atoms and nonpolar molecules. (Ar, C8H18)

  9. London Dispersion Forces

  10. Boiling point as a measure of intermolecular attractive forces

  11. Some Properties of a Liquid • Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals).

  12. Some Properties of a Liquid • Capillary Action: Spontaneous rising of a liquid in a narrow tube.

  13. Some Properties of a Liquid • Viscosity: Resistance to flow (molecules with large intermolecular forces).

  14. Types of Solids • Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)].

  15. Types of Solids • Amorphous solids: considerable disorder in their structures (glass).

  16. 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.

  17. Types of Crystalline Solids Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).

  18. Types of Crystalline Solids Molecular Solid: discrete covalently bondedmolecules at each of its lattice points (sucrose, ice).

  19. Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

  20. Closest Packing Holes

  21. Metal Alloys • Substitutional Alloy: some metal atoms replaced by others of similar size. • brass = Cu/Zn

  22. Metal Alloys(continued) • Interstitial Alloy:Interstices (holes) in closest packed metal structure are occupied by small atoms. • steel = iron + carbon

  23. Network Solids • Composed of strong directional covalent bonds that are best viewed as a “giant molecule”. • brittle • do not conduct heat or electricity • carbon, silicon-based • graphite, diamond, ceramics, glass

  24. Sulfur – S8

  25. Phosphorus – P4

  26. Diamond

  27. Graphite

  28. Zirconia

  29. Equilibrium Vapor Pressure • The pressure of the vapor present at equilibrium. • Determined principally by the size of the intermolecular forces in the liquid. • Increases significantly with temperature. • Volatile liquidshave high vapor pressures.

  30. LeChatelier’s Principle When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress.

  31. Translation: When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away. When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

  32. LeChatelier’s Example #1 A closed container of ice and water at equilibrium. The temperature is raised. Ice + Energy  Water The equilibrium of the system shifts to the _______ to use up the added energy. right

  33. LeChatelier’s Example #2 A closed container of N2O4 and NO2 at equilibrium. NO2 is added to the container. N2O4 + Energy  2 NO2 The equilibrium of the system shifts to the _______ to use up the added NO2. left

  34. LeChatelier’s Example #3 A closed container of water and its vapor at equilibrium. Vapor is removed from the system. water + Energy  vapor The equilibrium of the system shifts to the _______ to make more vapor. right

  35. constant Temperature remains __________ during a phase change. Water phase changes

  36. Kinetic Energy

  37. 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, Tc = 374°C and 218 atm).

  38. Water Water

  39. Phase changes by Name

  40. Carbon dioxide Carbon dioxide

  41. Carbon Carbon

  42. Sulfur

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