1. CHAPTER 13 Intermolecular Forces, Liquids, and Solids
Chemistry and Chemical Reactivity 6th Edition
John C. Kotz Paul M. Treichel Gabriela C. Weaver
CHAPTER 13
Intermolecular Forces, Liquids, and Solids
Lectures written by John Kotz
© 2006 Brooks/Cole Thomson

2. WHY? Why is water usually a liquid and not a gas?
Why does liquid water boil at such a high temperature for such a small molecule?
Why does ice float on water?
Why do snowflakes have 6 sides?
Why is I2 a solid whereas Cl2 is a gas?
Why are NaCl crystals little cubes?

3. Liquids, Solids & Intermolecular Forces Chap. 13

4. Inter- molecular Forces
Have studied INTRAmolecular forces—the forces holding atoms together to form molecules.
Now turn to forces between molecules — INTERmolecular forces.
Forces between molecules, between ions, or between molecules and ions.

5. Ion-Ion Forces for comparison of magnitude
Na+—Cl- in salt
These are the strongest forces.
Lead to solids with high melting temperatures.
NaCl, mp = 800 oC
MgO, mp = 2800 oC

6. Covalent Bonding Forces for comparison of magnitude
C=C, 610 kJ/mol
C–C, 346 kJ/mol
C–H, 413 kJ/mol
CN, 887 kJ/mol

7. Attraction Between Ions and Permanent Dipoles
Water is highly polar and can interact with positive ions to give hydrated ions in water.

8. Attraction Between Ions and Permanent Dipoles
Water is highly polar and can interact with positive ions to give hydrated ions in water.

9. Attraction Between Ions and Permanent Dipoles
Many metal ions are hydrated. This is the reason metal salts dissolve in water.

10. Attraction Between Ions and Permanent Dipoles
Attraction between ions and dipole depends on ion charge and ion-dipole distance.
Measured by ∆H for Mn+ + H2O --> [M(H2O)x]n+
-1922 kJ/mol
-405 kJ/mol
-263 kJ/mol

11. Dipole-Dipole Forces
Such forces bind molecules having permanent dipoles to one another.

12. Dipole-Dipole Forces
Influence of dipole-dipole forces is seen in the boiling points of simple molecules.
Compd Mol. Wt. Boil Point
N oC
CO oC
Br oC
ICl oC

13. H-bonding is strongest when X and Y are N, O, or F
Hydrogen Bonding
A special form of dipole-dipole attraction, which enhances dipole-dipole attractions.
H-bonding is strongest when X and Y are N, O, or F

14. H-Bonding Between Methanol and Water-
H-bond + -

15. H-Bonding Between Two Methanol Molecules- + -
H-bond

16. H-Bonding Between Ammonia and Water- + -
H-bond
This H-bond leads to the formation of NH4+ and OH-

17. Hydrogen Bonding in H2O
H-bonding is especially strong in water because
the O—H bond is very polar
there are 2 lone pairs on the O atom
Accounts for many of water’s unique properties.

18. Hydrogen Bonding in H2O Ice has open lattice-like structure.
Ice density is < liquid.
And so solid floats on water.
Snow flake:

19. Hydrogen Bonding in H2O Ice has open lattice-like structure.
Ice density is < liquid and so solid floats on water.
One of the VERY few substances where solid is LESS DENSE than the liquid.

20. A consequence of hydrogen bonding

21. Hydrogen Bonding in H2O
H bonds ---> abnormally high specific heat capacity of water (4.184 J/g•K)
This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

22. Hydrogen Bonding
H bonds leads to abnormally high boiling point of water.
See Screen 13.7

23. Boiling Points of Simple Hydrogen-Containing Compounds
Active Figure 13.8

24. Methane Hydrate

25. Methane Clathrate

26. Hydrogen Bonding in Biology
H-bonding is especially strong in biological systems — such as DNA.
DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O.
Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases.
—adenine with thymine
—guanine with cytosine

27. Double helix of DNA
Portion of a DNA chain

28. Base-Pairing through H-Bonds

29. Base-Pairing through H-Bonds

30. Double Helix of DNA

31. Discovering the Double Helix
Rosalind Franklin,
Maurice Wilkins,
James Watson and Francis Crick, 1953

32. Hydrogen Bonding in Biology
Hydrogen bonding and base pairing in DNA.
See Screen 13.6

33. FORCES INVOLVING INDUCED DIPOLES
How can non-polar molecules such as O2 and I2 dissolve in water?
The water dipole INDUCES a dipole in the O2 electric cloud.
Dipole-induced dipole

34. FORCES INVOLVING INDUCED DIPOLES
Solubility increases with mass the gas

35. FORCES INVOLVING INDUCED DIPOLES
Process of inducing a dipole is polarization
Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.

36. IM FORCES — INDUCED DIPOLES
Consider I2 dissolving in ethanol, CH3CH2OH. O H - d +
I-I R
The alcohol temporarily creates or INDUCES a dipole in I2.

37. FORCES INVOLVING INDUCED DIPOLES
Formation of a dipole in two nonpolar I2 molecules.
Induced dipole-induced dipole

38. FORCES INVOLVING INDUCED DIPOLES
The induced forces between I2 molecules are very weak, so solid I2 sublimes (goes from a solid to gaseous molecules).

39. FORCES INVOLVING INDUCED DIPOLES
The magnitude of the induced dipole depends on the tendency to be distorted.
Higher molec. weight ---> larger induced dipoles.
Molecule Boiling Point (oC)
CH4 (methane)
C2H6 (ethane)
C3H8 (propane)
C4H10 (butane)

40. Boiling Points of Hydrocarbons
CH4
C2H6
C3H8
C4H10
Note linear relation between bp and molar mass.

41. Summary of Intermolecular Forces
Ion-dipole forces
Dipole-dipole forces
Special dipole-dipole force: hydrogen bonds
Forces involving nonpolar molecules: induced forces

42. Intermolecular Forces Summary

43. Intermolecular Forces
Figure 13.13

44. Liquids Section 13.5 In a liquid • molecules are in constant motion
• there are appreciable intermolec. forces
• molecules close together
• Liquids are almost incompressible
• Liquids do not fill the container
Liquids Section 13.5

45. Liquids evaporation---> <---condensation LIQUID Add energy VAPOR
The two key properties we need to describe are EVAPORATION and its opposite—CONDENSATION
evaporation--->
LIQUID
Add energy
VAPOR
break IM bonds
make IM bonds
Remove energy
<---condensation

46. Liquids—Evaporation
To evaporate, molecules must have sufficient energy to break IM forces.
Breaking IM forces requires energy. The process of evaporation is endothermic.

47. Liquids— Distribution of Energies
lower T
higher T
Number of molecules
Molecular energy
Minimum energy req’d to break IM forces and evaporate
Distribution of molecular energies in a liquid.
KE is propor-tional to T.
See Figure 13.14

48. Distribution of Energy in a Liquid
Figure 13.14

49. Liquids
At higher T a much larger number of molecules has high enough energy to break IM forces and move from liquid to vapor state.
High E molecules carry away E. You cool down when sweating or after swimming.

50. When molecules of liquid are in the vapor state, they exert a VAPOR PRESSURE
EQUILIBRIUM VAPOR PRESSURE is the pressure exerted by a vapor over a liquid in a closed container when the rate of evaporation = the rate of condensation.
Liquids

51. Measuring Equilibrium Vapor Pressure
Liquid in flask evaporates and exerts pressure on manometer.
Active Fig

52. Vapor Pressure CD, Screen 13.9

53. Equilibrium Vapor Pressure Active Figure 13.18

54. Liquids Equilibrium Vapor Pressure
FIGURE 13.18: VP as a function of T.
1. The curves show all conditions of P and T where LIQ and VAP are in EQUILIBRIUM
2. The VP rises with T.
3. When VP = external P, the liquid boils.
This means that BP’s of liquids change with altitude.

55. Boiling Liquids
Liquid boils when its vapor pressure equals atmospheric pressure.

56. Boiling Point at Lower Pressure
When pressure is lowered, the vapor pressure can equal the external pressure at a lower temperature.

57. Consequences of Vapor Pressure Changes
When can cools, vp of water drops. Pressure in the can is less than that of atmosphere, so can is crushed.

58. Liquids Figure 13.18: VP versus T
4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING POINT
5. VP of a given molecule at a given T depends on IM forces. Here the VP’s are in the order C 2 H 5
water
alcohol
ether
increasing strength of IM interactions
extensive
H-bonds
dipole-
dipole O

59. Liquids
HEAT OF VAPORIZATION is the heat req’d (at constant P) to vaporize the liquid.
LIQ heat ---> VAP
Compd. ∆Hvap (kJ/mol) IM Force
H2O (100 oC) H-bonds
SO (-47 oC) dipole
Xe (-107 oC) induced dipole

60. Equilibrium Vapor Pressure & the Clausius-Clapeyron Equation
Clausius-Clapeyron equation — used to find ∆H˚vap.
The logarithm of the vapor pressure P is proportional to ∆Hvaporiation and to 1/T.
ln P = –(∆H˚vap/RT) + C

61. Liquids
Molecules at surface behave differently than those in the interior.
Molecules at surface experience net INWARD force of attraction.
This leads to SURFACE TENSION — the energy req’d to break the surface.

62. Surface Tension
SURFACE TENSION also leads to spherical liquid droplets.

63. Liquids concave meniscus ADHESIVE FORCES between water and glass H 2
Intermolec. forces also lead to CAPILLARY action and to the existence of a concave meniscus for a water column.
concave
meniscus
ADHESIVE FORCES
between water
and glass H 2
O in
glass
tube
COHESIVE FORCES
between water
molecules

64. Capillary Action
Movement of water up a piece of paper depends on H-bonds between H2O and the OH groups of the cellulose in the paper.

65. Metallic and Ionic Solids Sections 13.6-8

66. Types of Solids Table 13.6 TYPE EXAMPLE FORCE
Ionic NaCl, CaF2, ZnS Ion-ion
Metallic Na, Fe Metallic
Molecular Ice, I2 Dipole Ind. dipole
Network Diamond Extended Graphite covalent

67. Network Solids
Diamond
Graphite

68. Network Solids
A comparison of diamond (pure carbon) with silicon.

69. Properties of Solids
1. Molecules, atoms or ions locked into a CRYSTAL LATTICE
2. Particles are CLOSE together
3. STRONG IM forces
4. Highly ordered, rigid, incompressible
ZnS, zinc sulfide

70. Crystal Lattices
Regular 3-D arrangements of equivalent LATTICE POINTS in space.
Lattice points define UNIT CELLS
smallest repeating internal unit that has the symmetry characteristic of the solid.

71. Cubic Unit Cells
There are 7 basic crystal systems, but we are only concerned with CUBIC.
All sides
equal length
All angles
are 90 degrees

72. Cubic Unit Cells of Metals Figure 13.27
Simple cubic (SC)
Body-centered cubic (BCC)
Face-centered cubic (FCC)

73. Cubic Unit Cells of Metals Figure 13.27

74. Units Cells for Metals
Figure 13.28

75. Simple Cubic Unit Cell Figure 13.26
E atom is at a corner of a unit cell and is shared among 8 unit cells.
Each edge is shared with 4 cells
Each face is part of two cells.

76. Atom Packing in Unit Cells
Assume atoms are hard spheres and that crystals are built by PACKING of these spheres as efficiently as possible.

77. Atom Packing in Unit Cells
Assume atoms are hard spheres and that crystals are built by PACKING of these spheres as efficiently as possible.

78. Atom Packing in Unit Cells

79. Crystal Lattices— Packing of Atoms or Ions
FCC is more efficient than either BC or SC.
Leads to layers of atoms.

80. Packing of Atoms or Ions
Crystal Lattices—
Packing of Atoms or Ions
Packing of C60 molecules. They are arranged at the lattice points of a FCC lattice.

81. Number of Atoms per Unit Cell
Unit Cell Type Net Number Atoms SC
BCC
FCC 1 2 4

82. Atom Sharing at Cube Faces and Corners
Atom shared in corner --> 1/8 inside each unit cell
Atom shared in face
--> 1/2 inside each unit cell

83. Simple Ionic Compounds
Lattices of many simple ionic solids are built by taking a SC or FCC lattice of ions of one type and placing ions of opposite charge in the holes in the lattice.
EXAMPLE: CsCl has a SC lattice of Cs+ ions with Cl- in the center.

84. Simple Ionic Compounds
CsCl unit cell has a SC lattice of Cl- ions with Cs+ in the center.
1 unit cell has 1 Cs+ ion plus
(8 corners)(1/8 Cl- per corner) = 1 net Cl- ion.

85. Two Views of CsCl Unit Cell
Lattice can be SC lattice of Cl- with Cs+ in hole
OR SC lattice of Cs+ with Cl- in hole
Either arrangement leads to formula of 1 Cs+ and 1 Cl- per unit cell

86. Simple Ionic Compounds
Salts with formula MX can have SC structure — but not salts with formula MX2 or M2X

87. NaCl Construction Na+ in octahedral holes
FCC lattice of Cl- with Na+ in holes

88. Octahedral Holes - FCC Lattice

89. The Sodium Chloride Lattice
Na+ ions are in OCTAHEDRAL holes in a face-centered cubic lattice of Cl- ions.

90. The Sodium Chloride Lattice
Many common salts have FCC arrangements of anions with cations in OCTAHEDRAL HOLES — e.g., salts such as CA = NaCl
• FCC lattice of anions ----> 4 A-/unit cell
• C+ in octahedral holes ---> 1 C+ at center
+ [12 edges • 1/4 C+ per edge]
= 4 C+ per unit cell

91. Comparing NaCl and CsCl
Even though their formulas have one cation and one anion, the lattices of CsCl and NaCl are different.
The different lattices arise from the fact that a Cs+ ion is much larger than a Na+ ion.

92. Common Ionic Solids Titanium dioxide, TiO2
There are 2 net Ti4+ ions and 4 net O2- ions per unit cell.

93. Common Ionic Solids Zinc sulfide, ZnS
The S2- ions are in TETRAHEDRAL holes in the Zn2+ FCC lattice.
This gives 4 net Zn2+ ions and 4 net S2- ions.

94. Common Ionic Solids Fluorite or CaF2 FCC lattice of Ca2+ ions
This gives 4 net Ca2+ ions.
F- ions in all 8 tetrahedral holes.
This gives 8 net F- ions.

95. Barium titanate, a perovskite
BaTiO3
Ba2+
Barium titanate, a perovskite
Ti4+

96. Common Ionic Solids
Magnesium silicate, MgSiO3

97. Biotite Layers of linked octahedra of MgO6 and FeO6.
Layers of linked SiO4 tetrahedra.
K ions between layers

98. Phase Diagrams

99. TRANSITIONS BETWEEN PHASES Section 13.10
Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line.
(At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example.)

100. Phase Diagram for Water
Liquid phase
Solid phase
Gas phase

101. Phase Equilibria — Water
Solid-liquid
Gas-Liquid
Gas-Solid

102. Triple Point — Water
At the TRIPLE POINT all three phases are in equilibrium.

103. Phases Diagrams— Important Points for Water
T(˚C) P(mmHg)
Normal boil point
Normal freeze point
Triple point

104. Critical T and P
As P and T increase, you finally reach the CRITICAL T and P
Above critical T no liquid exists no matter how high the pressure.

105. Critical T and P
COMPD Tc(oC) Pc(atm)
H2O CO CH
Freon (CCl2F2)
Notice that Tc and Pc depend on intermolecular forces.

106. Solid-Liquid Equilibria
In any system, if you increase P the DENSITY will go up.
Therefore — as P goes up, equilibrium favors phase with the larger density (or SMALLER volume/gram).
Liquid H2O Solid H2O
Density 1 g/cm g/cm3
cm3/gram

107. Solid-Liquid Equilibria
Raising the pressure at constant T causes water to melt.
The NEGATIVE SLOPE of the S/L line is unique to H2O. Almost everything else has positive slope.

108. Solid-Liquid Equilibria
The behavior of water under pressure is an example of
LE CHATELIER’S PRINCIPLE
At Solid/Liquid equilibrium, raising P squeezes the solid.
It responds by going to phase with greater density, i.e., the liquid phase.
Solid-Liquid Equilibria

109. Solid-Vapor Equilibria
At P < 4.58 mmHg and T < ˚C
solid H2O can go directly to vapor. This process is called SUBLIMATION
This is how a frost-free refrigerator works.

110. CO2 Phase Diagram

111. CO2 Phases Separate phases Increasing pressure More pressure
Supercritcal CO2
Figure 13.41