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I. Intermolecular Forces

Ch. 12 – States of Matter. I. Intermolecular Forces. A. Definition of IMF. Attractive forces between molecules. Much weaker than chemical bonds within molecules. a.k.a. van der Waals forces. B. Types of IMF. B. Types of IMF. London Dispersion Forces. View animation online.

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I. Intermolecular Forces

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  1. Ch. 12 – States of Matter I. Intermolecular Forces

  2. A. Definition of IMF • Attractive forces between molecules. • Much weaker than chemical bonds within molecules. • a.k.a. van der Waals forces

  3. B. Types of IMF C. Johannesson

  4. B. Types of IMF • London Dispersion Forces View animation online.

  5. + - B. Types of IMF • Dipole-Dipole Forces View animation online.

  6. B. Types of IMF • Hydrogen Bonding

  7. C. Determining IMF • NCl3 • polar = dispersion, dipole-dipole • CH4 • nonpolar = dispersion • HF • H-F bond = dispersion, dipole-dipole, hydrogen bonding

  8. Ch. 12 - Liquids & Solids II. Physical Properties

  9. LIQUIDS Stronger than in gases Y high N slower than in gases SOLIDS Very strong N high N extremely slow A. Liquids vs. Solids IMF Strength Fluid Density Compressible Diffusion

  10. B. Liquid Properties • Surface Tension • attractive force between particles in a liquid that minimizes surface area

  11. B. Liquid Properties • Capillary Action • attractive force between the surface of a liquid and the surface of a solid

  12. B. Liquid Properties • Viscosity • Measure of the resistance of a liquid to flow

  13. water mercury B. Liquid Properties • Cohesion • Force of attraction between identical molecules • Adhesion • Force of attraction between different molecules

  14. decreasing m.p. C. Types of Solids • Crystalline - repeating geometric pattern • covalent network • metallic • ionic • covalent molecular • Amorphous - no geometric pattern

  15. C. Types of Solids Ionic (NaCl) Metallic

  16. C. Types of Solids Covalent Molecular (H2O) Covalent Network (SiO2 - quartz) Amorphous (SiO2 - glass)

  17. water mercury C. Types of Solids • Allotrope • In a covalent network, different forms in the same state Carbon allotropes

  18. water mercury C. Types of Solids • Amorphous Solid • No geometric pattern

  19. Ch. 12 - Liquids & Solids III. Changes of State

  20. A. Phase Changes

  21. A. Phase Changes • Evaporation • molecules at the surface gain enough energy to overcome IMF • Volatility • measure of evaporation rate • depends on temp & IMF

  22. # of Particles temp volatility IMF volatility Kinetic Energy A. Phase Changes Boltzmann Distribution

  23. A. Phase Changes • Equilibrium • trapped molecules reach a balance between evaporation & condensation

  24. temp v.p. IMF v.p. A. Phase Changes • Vapor Pressure • pressure of vapor above a liquid at equilibrium v.p. • depends on temp & IMF • directly related to volatility temp

  25. Patm b.p. IMF b.p. A. Phase Changes • Boiling Point • temp at which v.p. of liquid equals external pressure • depends on Patm & IMF • Normal B.P. - b.p. at 1 atm

  26. IMF m.p. A. Phase Changes • Melting Point • equal to freezing point • Which has a higher m.p.? • polar or nonpolar? • covalent or ionic? polar ionic

  27. A. Phase Changes • Sublimation • solid  gas • v.p. of solid equals external pressure • EX: dry ice, mothballs, solid air fresheners

  28. Gas - KE  Boiling - PE  Liquid - KE  Melting - PE  Solid - KE  B. Heating Curves

  29. B. Heating Curves • Temperature Change • change in KE (molecular motion) • depends on heat capacity • Heat Capacity • energy required to raise the temp of 1 gram of a substance by 1°C

  30. B. Heating Curves • Phase Change • change in PE (molecular arrangement) • temp remains constant • Heat of Fusion (Hfus) • energy required to melt 1 gram of a substance at its m.p.

  31. B. Heating Curves • Heat of Vaporization (Hvap) • energy required to boil 1 gram of a substance at its b.p. • EX: sweating, steam burns, the drinking bird

  32. C. Phase Diagrams • Show the phases of a substance at different temps and pressures.

  33. The following slides… same information, different explanation and examples

  34. Phase Changes

  35. Why do liquids and solids form at all? • KMT postulates • A gas is a collection of small particles traveling in straight-line motion and obeying Newton's Laws. • The molecules in a gas occupy no volume. • Collisions between molecules are perfectly elastic • no energy is gained or lost during the collision • There are no attractive or repulsive forces between the molecules. • Kinetic energy is proportional to temperature X

  36. What is a phase? • Region of matter that is: • chemically uniform • physically distinct • mechanically separable. • Often synonymous with (same meaning as) “state of matter”

  37. Changing phases Distinguish liquid vs. solid?

  38. Properties of … • Gas phase • Like/unlike soccer players on field • Liquid phase • Like/unlike crowd at a rally • Like/unlike gases • Solid phase • Like/unlike movie theatre • Compare intermolecular interactions • gases vs. liquids vs. solids

  39. Phases of matter Add more energy

  40. A - melting B - freezing C – boiling/ evap D - condensation E - sublimation F - solidification Phase change vocabulary

  41. Phase changes and IMF’s • As InterMolecular Forces increase, melting and boiling temperatures _________? (increase or decrease)? Network covalent bonding Larger sphere, higher melting point Metallic bonding

  42. Strength of interactions • Which simulation has stronger intermolecular interactions? A or B • How do you know? Same temperature A B

  43. Ranking of intermolecular interactions • Water • Wood • Iron • Air • Gold • Mercury • Carbon dioxide • Oxygen • Gasoline • Lead vs. Why?

  44. Why is water special? • Periodic trends • Boiling and melting points of hydrides

  45. KMT, energy and phase changes

  46. Total energy

  47. Water phases present?

  48. Heating curves Why are b and d flat?

  49. Boiling vs. melting Which takes more energy? (same mass)

  50. Consider liquid  gasEvaporation removes energy

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