1. Intermolecular
Forces
Chapter 11
Intermolecular Forces,
Liquids, and Solids

2. Intermolecular
Forces
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States of Matter
The fundamental difference between states of
matter is the distance between particles.

3. Intermolecular
Forces
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States of Matter
Because in the solid and liquid states
particles are closer together, we refer to them
as condensed phases.

4. Intermolecular
Forces
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The States of Matter
• The state a substance is
in at a particular
temperature and
pressure depends on
two antagonistic entities:
– the kinetic energy of the
particles;
– the strength of the
attractions between the
particles.

5. Intermolecular
Forces
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Intermolecular Forces
The attractions between molecules are not
nearly as strong as the intramolecular
attractions that hold compounds together.

6. Intermolecular
Forces
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Intermolecular Forces
They are, however, strong enough to control
physical properties such as boiling and
melting points, vapor pressures, and
viscosities.

7. Intermolecular
Forces
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Intermolecular Forces
These intermolecular forces as a group are
referred to as van der Waals forces.

8. Intermolecular
Forces
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van der Waals Forces
• Dipole-Dipole interactions
• Hydrogen bonding
• London dispersion forces

9. Intermolecular
Forces
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Ion-Dipole Interactions
• Ion-dipole interactions (a fourth type of force),
are important in solutions of ions.
• The strength of these forces are what make it
possible for ionic substances to dissolve in
polar solvents.

10. Intermolecular
Forces
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Dipole-Dipole Interactions
• Molecules that have
permanent dipoles are
attracted to each other.
– The positive end of one is
attracted to the negative
end of the other and vice-
versa.
– These forces are only
important when the
molecules are close to
each other.

11. Intermolecular
Forces
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Dipole-Dipole Interactions
The more polar the molecule, the higher
is its boiling point.

12. Intermolecular
Forces
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London Dispersion Forces
While the electrons in the 1s orbital of helium
would repel each other (and, therefore, tend
to stay far away from each other), it does
happen that they occasionally wind up on the
same side of the atom.

13. Intermolecular
Forces
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London Dispersion Forces
At that instant, then, the helium atom is polar,
with an excess of electrons on the left side
and a shortage on the right side.

14. Intermolecular
Forces
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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.

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

16. Intermolecular
Forces
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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.

17. Intermolecular
Forces
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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.

18. Intermolecular
Forces
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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.

19. Intermolecular
Forces
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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.

20. Intermolecular
Forces
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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. Intermolecular
Forces
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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. Intermolecular
Forces
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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. Intermolecular
Forces
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Summarizing Intermolecular Forces

24. Intermolecular
Forces
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Intermolecular Forces Affect
Many Physical Properties
The strength of the
attractions between
particles can greatly
affect the properties
of a substance or
solution.

25. Intermolecular
Forces
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Viscosity
• Resistance of a liquid
to flow is called
viscosity.
• It is related to the ease
with which molecules
can move past each
other.
• Viscosity increases
with stronger
intermolecular forces
and decreases with
higher temperature.

26. Intermolecular
Forces
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Surface Tension
Surface tension
results from the net
inward force
experienced by the
molecules on the
surface of a liquid.

27. Intermolecular
Forces
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Phase Changes

28. Intermolecular
Forces
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Energy Changes Associated
with Changes of State
The heat of fusion is the energy required to
change a solid at its melting point to a liquid.

29. Intermolecular
Forces
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Energy Changes Associated
with Changes of State
The heat of vaporization is defined as the
energy required to change a liquid at its
boiling point to a gas.

30. Intermolecular
Forces
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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 a phase change.

31. Intermolecular
Forces
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Vapor Pressure
• At any temperature some molecules in a
liquid have enough energy to escape.
• As the temperature rises, the fraction of
molecules that have enough energy to
escape increases.

32. Intermolecular
Forces
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Vapor Pressure
As more molecules
escape the liquid,
the pressure they
exert increases.

33. Intermolecular
Forces
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Vapor Pressure
The liquid and vapor
reach a state of
dynamic equilibrium:
liquid molecules
evaporate and vapor
molecules condense
at the same rate.

34. Intermolecular
Forces
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Vapor Pressure
• The boiling point of a
liquid is the
temperature at which
it’s vapor pressure
equals atmospheric
pressure.
• The normal boiling
point is the
temperature at which
its vapor pressure is
760 torr.

35. Intermolecular
Forces
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Phase Diagrams
Phase diagrams display the state of a
substance at various pressures and
temperatures and the places where equilibria
exist between phases.

36. Intermolecular
Forces
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Phase Diagrams
• The circled line is the liquid-vapor interface.
• It starts at the triple point (T), the point at
which all three states are in equilibrium.

37. Intermolecular
Forces
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Phase Diagrams
It ends at the critical point (C); above this
critical temperature and critical pressure the
liquid and vapor are indistinguishable from
each other.

38. Intermolecular
Forces
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Phase Diagrams
Each point along this line is the boiling point
of the substance at that pressure.

39. Intermolecular
Forces
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Phase Diagrams
• The circled line in the diagram below is the
interface between liquid and solid.
• The melting point at each pressure can be
found along this line.

40. Intermolecular
Forces
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Phase Diagrams
• Below the triple point the substance cannot
exist in the liquid state.
• Along the circled line the solid and gas
phases are in equilibrium; the sublimation
point at each pressure is along this line.

41. Intermolecular
Forces
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Phase Diagram of Water
• Note the high critical
temperature and critical
pressure.
– These are due to the
strong van der Waals
forces between water
molecules.

42. Intermolecular
Forces
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Phase Diagram of Water
• The slope of the solid-
liquid line is negative.
– This means that as the
pressure is increased at a
temperature just below the
melting point, water goes
from a solid to a liquid.

43. Intermolecular
Forces
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Phase Diagram of Carbon Dioxide
Carbon dioxide
cannot exist in the
liquid state at
pressures below
5.11 atm; CO2
sublimes at normal
pressures.

44. Intermolecular
Forces
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Phase Diagram of Carbon Dioxide
The low critical
temperature and
critical pressure for
CO2 make
supercritical CO2 a
good solvent for
extracting nonpolar
substances (like
caffeine)

45. Intermolecular
Forces
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Solids
• We can think of
solids as falling into
two groups:
– crystalline, in which
particles are in highly
ordered arrangement.

46. Intermolecular
Forces
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Solids
• We can think of
solids as falling into
two groups:
– amorphous, in which
there is no particular
order in the
arrangement of
particles.

47. Intermolecular
Forces
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Attractions in Ionic Crystals
In ionic crystals, ions
pack themselves so as
to maximize the
attractions and
minimize repulsions
between the ions.

48. Intermolecular
Forces
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Crystalline Solids
Because of the
ordered in a crystal,
we can focus on the
repeating pattern of
arrangement called
the unit cell.

49. Intermolecular
Forces
49
A crystalline solid possesses rigid and long-range order. In a
crystalline solid, atoms, molecules or ions occupy specific
(predictable) positions.
An amorphous solid does not possess a well-defined
arrangement and long-range molecular order.
A unit cell is the basic repeating structural unit of a crystalline
solid.
lattice
point
Unit Cell Unit cells in 3 dimensions
At lattice points:
• Atoms
• Molecules
• Ions

50. Intermolecular
Forces
50
Seven Basic Unit Cells

51. Intermolecular
Forces
51
Three Types of Cubic Unit Cells

52. Intermolecular
Forces
52
Arrangement of Identical Spheres in a Simple Cubic Cell

53. Intermolecular
Forces
53
Arrangement of Identical Spheres in a Body-Centered
Cubic Cell

54. Intermolecular
Forces
54
Shared by 8
unit cells
Shared by 4
unit cells
A Corner Atom, a Edge-Centered Atom and a
Face-Centered Atom
Shared by 2
unit cells

55. Intermolecular
Forces
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Crystalline Solids
There are several types of basic
arrangements in crystals, like the ones
depicted above.

56. Intermolecular
Forces
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Crystalline Solids
We can determine
the empirical
formula of an ionic
solid by determining
how many ions of
each element fall
within the unit cell.

57. Intermolecular
Forces
57
Number of Atoms Per Unit Cell
1 atom/unit cell
(8 x 1/8 = 1)
2 atoms/unit cell
(8 x 1/8 + 1 = 2)
4 atoms/unit cell
(8 x 1/8 + 6 x 1/2 = 4)

58. Intermolecular
Forces
58
Relation Between Edge Length and Atomic Radius

59. Intermolecular
Forces
59
Closet Packing: Hexagonal and Cubic
hexagonal cubic

60. Intermolecular
Forces
60
Exploded Views

61. Intermolecular
Forces
61
When silver crystallizes, it forms face-centered cubic cells. The
unit cell edge length is 409 pm. Calculate the density of silver.
d =
m
V
V = a3
= (409 pm)3
= 6.83 x 10-23
cm3
4 atoms/unit cell in a face-centered cubic cell
m = 4 Ag atoms
107.9 g
mole Ag
x
1 mole Ag
6.022 x 1023
atoms
x = 7.17 x 10-22
g
d =
m
V
7.17 x 10-22
g
6.83 x 10-23
cm3
= = 10.5 g/cm3

62. Intermolecular
Forces
62
An Arrangement for Obtaining the X-ray Diffraction Pattern
of a Crystal.

63. Intermolecular
Forces
63
Extra distance = BC + CD = 2d sinθ = nλ (Bragg Equation)
Reflection of X rays from Two Layers of Atoms.

64. Intermolecular
Forces
64
X rays of wavelength 0.154 nm are diffracted from a crystal at an
angle of 14.17o
. Assuming that n = 1, what is the distance (in pm)
between layers in the crystal?
nλ = 2d sin θ n = 1 θ = 14.17o
λ = 0.154 nm = 154 pm
d =
nλ
2sinθ
=
1 x 154 pm
2 x sin14.17
= 314.0 pm

65. Intermolecular
Forces
65
Types of Crystals
Ionic Crystals
• Lattice points occupied by cations and anions
• Held together by electrostatic attraction
• Hard, brittle, high melting point
• Poor conductor of heat and electricity
CsCl ZnS CaF2

66. Intermolecular
Forces
66
Types of Crystals
Covalent Crystals
• Lattice points occupied by atoms
• Held together by covalent bonds
• Hard, high melting point
• Poor conductor of heat and electricity
diamond graphite
carbon
atoms

67. Intermolecular
Forces
67
Types of Crystals
Molecular Crystals
• Lattice points occupied by molecules
• Held together by intermolecular forces
• Soft, low melting point
• Poor conductor of heat and electricity
water benzene

68. Intermolecular
Forces
68
Types of Crystals
Metallic Crystals
• Lattice points occupied by metal atoms
• Held together by metallic bonds
• Soft to hard, low to high melting point
• Good conductors of heat and electricity
Cross Section of a Metallic Crystal
nucleus &
inner shell e-
mobile “sea”
of e-

69. Intermolecular
Forces
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Metallic Solids
• 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.

70. Intermolecular
Forces
70
Crystal Structures of Metals

71. Intermolecular
Forces
71
Types of Crystals

72. Intermolecular
Forces
72
An amorphous solid does not possess a well-defined
arrangement and long-range molecular order.
A glass is an optically transparent fusion product of inorganic
materials that has cooled to a rigid state without crystallizing
Crystalline
quartz (SiO2)
Non-crystalline
quartz glass

73. Intermolecular
Forces
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Covalent-Network and
Molecular Solids
• Diamonds are an example of a covalent-
network solid, in which atoms are covalently
bonded to each other.
– They tend to be hard and have high melting
points.

74. Intermolecular
Forces
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Covalent-Network and
Molecular Solids
• Graphite is an example of a molecular solid,
in which atoms are held together with van der
Waals forces.
– They tend to be softer and have lower melting
points.