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Understanding Intermolecular Forces in Liquids and Solids

Explore the forces of attraction in different molecules, such as dipole-dipole, hydrogen bonding, and dispersion forces. Learn about their impact on properties like boiling point and phase changes. Understand the structures of crystalline and amorphous solids, metal alloys, and network solids. Discover LeChatelier's principle and its application to equilibria. Dive into phase diagrams and critical points for substances like water and carbon dioxide.

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Understanding Intermolecular Forces in Liquids and Solids

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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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