Intermolecular Forces

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.