1. Intermolecular Forces and Liquids and Solids Chapter 11

2. These intermolecular forces as a group are referred to as van der Waals forces Ion-dipole forces Dipole-dipole interactions London dispersion forces Hydrogen bonding

3. Which Have a Greater Effect? Dipole-Dipole Interactions or Dispersion Forces If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

4. Intermolecular Forces Hydrogen bond A special dipole-dipole interaction between the hydrogen atom in a polar N−H O−H F−H bond and an highly electronegative O, N, or F atom on another molecule Fig 11.8

5. Fig 11.9 Comparing densities of liquid and solid phases paraffin ice d ice < d water

6. Fig 11.10 Hydrogen bonding in ice

7. Fig 11.11 Expansion of water upon freezing

8. Fig 11.12 Flowchart for determining intermolecular forces

9. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution: Viscosity Surface tension Capillary action

10. Resistance of a liquid to flow Related to the ease with which molecules can move past each other Increases with stronger intermolecular forces and decreases with higher temperature Viscosity Fig 11.13 SAE 40 SAE 10

11. Strong intermolecular forces High viscosity Viscosity

12. Surface Tension Results from the net inward force experienced by the molecules on the surface of a liquid

13. Surface Tension Fig 11.16 Two meniscus shapes Adhesive forces Cohesive forces Adhesive forces Cohesive forces H2OH2O Hg > > adhesion > cohesioncohesion > adhesion

14. Fig 11.17 Phase changes and associated names Phase changes

15. Fig 11.18 Comparison of enthalpy changes for fusion and vaporization In all cases: ΔH vap > ΔH fus Energy Changes Accompanying Phase Changes

16. Fig 11.19 Heating curve for water Indicates changes when 1.00 mol H 2 O is heated from 25°C to 125°C at constant P.

17. Vapor pressure - measured when a dynamic equilibrium exists between evaporation and condensation Rate of condensation Rate of evaporation = Dynamic Equilibrium Fig 11.22 Equilibrium vapor pressure over a liquid

18. Fig 11.23 Effect of temperature on the distribution of kinetic energies in a liquid

19. Vapor Pressure and Boiling Point Boiling point ≡ temperature at which its vapor pressure equals atmospheric pressure Normal boiling point ≡ temperature at which its vapor pressure is 760 torr. Fig 11.24

20. Molar heat of vaporization (  H vap ) ≡ the energy required to vaporize 1.00 mole of a liquid ln P = −  H vap RT + C Clausius-Clapeyron Equation P = (equilibrium) vapor pressure T = temperature (K) R = gas constant (8.314 J/K mol) ln P 1/T m =  H vap /R