1. Intermolecular Forces, Liquids & Solids Chapter 11

2. Overview  Liquids & Solids  Intermolecular Forces  Liquids  Phase Changes  Vapor Pressure  Phase Diagrams  Solids -- Structure  Bonding Types in Solids

3. Liquids & Solids  Solids particles close together locked into relative positions (crystalline) strong interactions (interparticle forces)  Liquids particles farther apart mobile relative to each other weaker interactions between particles Remember

4. GasLiquid + E - E condensation vaporization  H condensation = -  H vaporization = +

5. Liquid Solid + E - E freezing melting  H freezing = -  H melting = +

6. Intermolecular Forces  Strength of IM Forces determine boiling points and melting points  Ion-Dipole Forces occur between ions and dipoles between charged particles and neutral, polar covalent particles  Dipole-Dipole Forces occur between two dipoles between two, neutral, polar covalent particles

7. Ion - Dipole Force (40 - 600 kJ/mol) + O H  - ++ ++

8. Dipole - Dipole Force (5 - 25 kJ/mol)   “Hydrogen Bonding” Force (4 - 25 kJ/mol)

9.  “Hydrogen Bonding” Forces special case of Dipole-Dipole forces occur between two, neutral, polar covalent particles which – have a H atom (which is bound to an O, F or N atom) for the (+) dipole – have an O, F, or N atom for the (-) dipole extra strength due to –small size and large EN of O, F or N –and small size of H

10.  London Dispersion Forces occur between all particles even neutral, non- polar covalent particles occur between an instantaneous dipole and an induced dipole force is weak but strengthens with increasing polarizability of the particles polarizability of the particles increases with increasing size or mass

11. Dipole - Induced Dipole ( 2 - 10 kJ/mol)   polar molecule non-polar molecule

12.   Induced Dipole - Induced Dipole Force (0.05 - 40 kJ/mol) London Dispersion

13. Strength of Forces Ion - Dipole “Hydrogen Bonding” Dipole - Dipole Dipole - Induced Dipole London Dispersion Increasing

14. He 4.6 Ne 27.3 Ar 87.5 Kr 120.9 Xe 166.1 Boiling Points of Noble Gases Temp K MM

15. F 2 85.1 Cl 2 238.6 Br 2 332.0 I 2 457.6 Boiling Points of Halogen Diatomics Temp K MM

16. H 2 O 373 H 2 S 212 H 2 Se 231 H 2 Te 271 Boiling Points of Group 6 Dihydrides Temp K MM GeH 4 184 CH 4 109 SiH 4 161 SnH 4

17. What Does it Mean?  “hydration” is very important in the solvation process & compound formation  stronger the interaction, the more energy is released, more exothermic  the strength of the interaction determines the state of the substance  unusual properties of water are due to “hydrogen bonding”

18. What kind of forces are between:  O 2 molecules  H 2 O and NH 3 molecules  Ne atoms  HF and NH 3 molecules  CH 4 and Br 2 molecules London Dispersion Hydrogen Bonding London Dispersion Hydrogen Bonding London Dispersion

19. Properties of Liquids  Viscosity resistance to flow ease with which liquid particles move relative to one another related to attractive forces between particles and structural properties of the particles themselves decreases with increasing energy (temperature) of particles

20.  Surface Tension energy required to increase the surface area of a liquid by a unit amount a sphere produces the minimum surface area competition between cohesive forces vs adhesive forces –cohesive forces tend to minimize surface area –adhesive forces tend to maximize surface area high surface tension reflects strong cohesive forces capillary action -- low surface tension, strong adhesive forces

21. Phase Changes  Energy Changes melting, solid to liquid –endothermic --  H fusion vaporization, liquid to gas –endothermic --  H vaporization condensation, gas to liquid –exothermic --  H condensation = -  H vaporization freezing, liquid to solid –exothermic --  H freezing = -  H fusion

22. Phase Changes, cont’d sublimation, solid to gas –endothermic --  H sublimation deposition, gas to solid –exothermic --  H deposition = -  H sublimation

23. Properties of Liquids liquid solid freezing -- exothermic melting -- endothermic liquidgas vaporization -- endothermic condensation -- exothermic

24. Temp Time solid liquid gas solid + liquid in equilibrium liquid + gas in equilibrium constant temp

25.  Critical Temperature & Pressure critical temp. –highest temperature at which a substance can exist as a liquid critical pressure –pressure required to cause liquefaction at the critical temperature a gas cannot be liquefied above the critical temperature critical point –corresponds to T c and P c –the point above which a supercritical gas exists –the substance cannot be liquefied by increasing the pressure

26. Pressure Liquid H 2 O Vapor T c = 647.6 K Critical Point 217.7 atm = P c Temperature

27. Vapor Pressure  Molecular Description liquids have a distribution of energies for the liquid molecules at any temperature, some molecules have sufficient energy for vaporization the higher the temperature, the greater number of molecules with energy of vaporization at constant temperature, average energy of molecules is constant but in dynamic equilibrium vapor pressure is the pressure exerted by vaporized molecules when liquid and vapor states are in dynamic equilibrium

28. Evaporation Equilibrium Vapor Pressure liquid gas Eq. Vapor Pressure -- partial pressure of gas over a liquid at equilibrium

29.  Volatility, Vapor Pressure & Temp. substances with high vapor pressure are volatile in a open container, dynamic equilibrium cannot be established -- complete evaporation higher the temperature, greater vapor pressure, greater volatility  Vapor Pressure & Boiling Point boiling point -- temperature at which the vapor pressure = atmospheric pressure –boiling point increases with increasing external pressure, vice versa normal boiling point -- temp. at which the vapor pressure = 1 atm

30. Normal Boiling Point: temperature at which the vapor pressure is equal to 1.00 atmosphere Boiling Point: temperature at which the vapor pressure is equal to the external or atmospheric pressure

31. Which has highest boiling point:  H 2 O or H 2 S  BrCl or Cl 2  BrCl or HCl  CH 4 or C 2 H 6 H2OH2O BrCl HCl C2H6C2H6

32.  molar enthalpy of vaporization(kJ/mol) x amount of substance(moles)   H vap water = +40.7 kJ/mol @ 100 º  Clausius Clapeyron Equation: ln P = -  H vap / RT + C Calculate the Energy of Vaporization:

33. Phase Diagrams normal freezing pt. solid liquid normal boiling pt. vapor Triple Point Pressure Temperature H2OH2O solid vapor solidliquid vapor

34. Structures of Solids  Unit Cells characteristic 3-dimensional repeating unit of a crystalline substance primitive cubic (sc) 1 atom/u.c. body-centered cubic (bcc) 2 atoms/u.c. face-centered cubic (fcc) 4 atoms/u.c.

35. primitive cubic body-centered cubic face-centered cubic

36. Close Packing of Spheres  Layered as ABABABAB or  ABCABC

37. Physical Properties of Solids: solidliquid melting -- endothermic freezing -- exothermic solidgas sublimation -- endothermic deposition -- exothermic

38. Melting Point: Temperature at which the crystal lattice breaks down  H fusion  H freezing meltingfreezing

39. Which one of each pair will have the higher melting point?  NaF or NaI  HCl or HI  O 2 or Br 2  NaCl or BrCl  H 2 S or H 2 O NaF HCl Br 2 NaCl H2OH2O

40. Ioniccations & anions Metallic metal atoms Molecularmolecules Networkatoms Amorphousirregular networks Types of Solids: TypesUnits

41. Network Solid Diamond

42. Network Solid Graphite

43. NaCl Na Cl