1. Intermolecular Forces
Solids & Liquids

2. Warm-Up If 8 g of hydrogen and 32 g of methane (CH4) are
injected into an empty, rigid 100 L container at a
temperature of 100oC, determine the following:
The partial pressure of each gas.
The total pressure in the container.
The mole fraction of each gas.
How many times faster hydrogen would effuse than methane.

3. Liquids and Solids Intermolecular Forces (IMF)
Force of attraction between molecules in liquids and solids
Intramolecular forces (chemical bonds) are about
10x stronger than intermolecular

6. London-Dispersion Forces:He

7. Dipole-Dipole Forces:
HCl

8. Hydrogen Bonds:

9. London Dispersion Only
Which IM Forces?
YES NO
Nonpolar
Is it polar?
Dipole-Dipole
London Dispersion Only
H-F, H-N or H-O?
Hydrogen Bonds
Weakest
Low BP, MP
Stronger
Higher BP, MP

10. Examples:
CH4 HF
NH3

11. Warm-up Which type(s) of IM forces are in the following liquids: HCl
CO2
CF4
H2O

12. IMF BP MP
IMF BP MP

13. London Dispersion Forces
Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

14. London Dispersion Forces
London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

15. London Dispersion Forces
These forces are present in all molecules, whether they are polar or nonpolar.
The tendency of an electron cloud to distort in this way is called polarizability.

16. Factors Affecting London Forces
The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).
This is due to the increased surface area in n-pentane.

17. Factors Affecting London Forces
The strength of dispersion forces tends to increase with increased molecular weight.
Larger atoms have larger electron clouds, which are easier to polarize.

18. 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 be the dominating force.
If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

20. How Do We Explain This?
The nonpolar series (SnH4 to CH4) follow the expected trend.
The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

21. Hydrogen Bonding
The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
We call these interactions hydrogen bonds.

22. Hydrogen Bonding
Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.
Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

23. Heating & Cooling Curves Phase Diagrams

24. Warm-up: Identify the Intermolecular Forces found in the
following substances:
HCl CO2 NH3 SO3
2. Cl2 is a gas at room temperature, whereas Br2 is a solid.
Explain:
H2S is a gas at room temperature. Water is much lighter
but is a liquid at room temperature. Explain:

25. Phase Changes

26. Energy Changes Associated with Changes of State
Heat of Fusion: Energy required to change a solid at its melting point to a liquid.

27. Energy Changes Associated with Changes of State
Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas.

28. Heating/Cooling Curves (Water)
T (oC) 0o
During a phase change Temperature does not change!!!!
Heat Added (kJ)

29. Energy needed to BOIL: q = nDHVAP
Energy needed to MELT: q = nDHFUS
Energy needed to HEAT: q = mcpDT
During a phase change Temperature does not change!!!!

30. Energy Changes Associated with Changes of State
The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.
The temperature of the substance does not rise during the phase change.

31. Warm-up:

33. 2. How much energy does it take to melt
1. How much energy does it take to boil
100 g of water? (DHVAP = 41 kJ/mole)
2. How much energy does it take to melt
100 g of water? (DHFUS = 6 kJ/mole)

34. Heating/Cooling Curve
T (oC)
20o
Heat Added (kJ)

35. When water boils are chemical bonds being broken?
Explain:

38. Phase Diagrams The phase of a substance depends on T and P
Critical Point
Impossible to
Liquefy
Triple Point

39. Normal Pressure

40. What is the phase when T = -90 oC and P = 6 atm?
What is T at the triple point?
What is the critical temperature?
What happens if CO2 at -90 oC and 1 atm is heated at
constant P?

41. Vapor Pressure & Boiling & Freezing Points

42. What is the phase when T = 20 oC and P = 50 atm?
What is P at the triple point?
What is the critical temperature?
4. Can liquid CO2 exist above 32oC?
5. What happens if CO2 at -90 oC and 8 atm is heated at
constant P?

43. Vapor Pressure(VP,PVAP): Pressure exerted by a gas
above the liquid
PVAP depends only on TEMPERATURE!!!!

44. T VP
Weak IMF
Strong IMF

45. When does a substance boil?

48. Intermolecular Forces
IMF BP MP
HVAP
HFUS VP

52. Bonding in Solids

53. Types of Bonding in Crystalline Solids

54. Bonding in Crystalline Solids
Particles arranged in a orderly repeating pattern
Molecular: LD, DD, H-bonds- MP varies Ionic: Ions, very strong forces, Lattice Energy (charge, size of ion)-High MP Covalent Network: held together by very strong COVALENT bonds, very high melting points Diamond [C] and Quartz [SiO2] Metallic: Mostly high MP, “Delocalized Electrons)

55. Ionic Solids: ELat MP
MP = 2852oC
MP = 801oC

56. Amorphous Solids

57. DIAMOND (C) GRAPHITE (C)
MP = 3550oC !!!

58. QUARTZ (SiO2)
MP = 1670oC !!!

59. Metallic Solids Al 660oC Hg -38oC W 3440oC
Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.
In metals, valence electrons are delocalized throughout the solid. Al
660oC Hg
-38oC W
3440oC

60. What IM Forces are in these Solids?
CO C (diamond)
H2O 8. NaCl
MgO Au
SiO2 Ar

61. Viscosity: measure of resistance to flow

62. Surface tension: cohesiveness of molecules

63. Capillary Action: movement of liquids up surfaces
Ex) Paper towel in water
water up roots

64. Give examples of :
Viscosity in a molecule
Capillary action in a molecule
Surface Tension in a molecule

65. “Like-Dissolves-Like”
Polar dissolves in polar
Nonpolar dissolves in nonpolar
Oil = nonpolar
Water = polar

66. Why did the polar bear disappear in water?

72. Which IMF are in the following:
Test Review:
Which IMF are in the following:
NaCl has a M.P of 801oC and MgO has a M.P. of 2852oC. Why is
MgO so much higher.
Which would you expect to have a higher BP, F2 or Cl2?
What is the vapor pressure of water when it boils?
Vapor pressure depends only on ______________
Ar (l)
CO2 (s)
H2O (l)
SiO2 (s)
Ag (s)
Br2 (l)
C (diamond)
MgO (s)

73. Heating/Cooling Curve
T (oC)
20o
Heat Added (kJ)

75. Natural Logarithm
Calculating Vapor Pressure using Natural Logarithm and the Clausius-Clapeyron Equation (Appendix 1.2)

77. Wet Dry Ice Lab

78. What is DRY ICE?

79. In order to observe liquid CO2 what is the lowest
pressure you would need?

80. Look for the liquid CO2!!!!!!!