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Intermolecular Forces

Intermolecular Forces. Kinetic Molecular Theory. Describes the behavior of subatomic particles Liquids, solids, and gases are composed of small particles that have mass. Particles are in constant, random, rapid motion. Particles have collisions.

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Intermolecular Forces

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  1. Intermolecular Forces

  2. Kinetic Molecular Theory • Describes the behavior of subatomic particles • Liquids, solids, and gases are composed of small particles that have mass. • Particles are in constant, random, rapid motion. • Particles have collisions. • Particles have an avg. KE directly related to temperature. • The state of a substance at room temperature depends on the strength of the attractions between its particles.

  3. Definition of IMF • Attractive forces between molecules. • Much weaker than chemical bonds within molecules.

  4. Definition of IMF • Intramolecular forces: Covalent Bonding • Much stronger than chemical bonds between atoms. • Examples : nonpolar, polar sharing

  5. Definition of IMF • Intermolecular Forces • Attractive forces between molecules. • Much weaker than chemical bonds within molecules.

  6. Intermolecular Forces • Attractive forces between molecules or particles (ions, metal atoms, etc…) • Examples: • dispersion, (London /Vander Waals); • dipole-dipole, • dipole-ion, • hydrogen “bonding”, • metallic bonding, • ion-ion

  7. Intermolecular Forces • Relative Strength: Weakest Strongest • Examples: • dispersion, (London /Vander Waals); • dipole-dipole, • dipole-ion, • hydrogen “bonding”, • metallic bonding, • ion-ion

  8. Types of IMF

  9. Types of IMF • London Dispersion Forces View animation online.

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

  11. Types of IMF • Hydrogen Bonding

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

  13. Liquids & Solids Physical Properties

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

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

  16. water mercury Liquid Properties • Capillary Action • attractive force between the surface of a liquid and the surface of a solid

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

  18. Types of Solids Ionic (NaCl) Metallic

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

  20. Liquids & Solids Changes of State

  21. Phase Changes

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

  23. # of Particles temp volatility IMF volatility Kinetic Energy Phase Changes Boltzmann Distribution p. 477

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

  25. temp v.p. IMF v.p. Phase Changes p.478 • Vapor Pressure • pressure of vapor above a liquid at equilibrium v.p. • depends on temp & IMF • directly related to volatility temp

  26. Patm b.p. IMF b.p. 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

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

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

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

  30. 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 • “Volcano” clip - water has a very high heat capacity

  31. 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.

  32. Heating Curves • Heat of Vaporization (Hvap) • energy required to boil 1 gram of a substance at its b.p. • usually larger than Hfus…why? • EX: sweating, steam burns, the drinking bird

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

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