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Chapter 11
Intermolecular Forces and
Liquids and Solids
Scientists are interested in how matter behaves under unusual
circumstances. For example, before the space station could
be built, fundamental research into materials properties had to
be undertaken.
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In this chapter intermolecular forces will be
used to explain (among other things):
• How can carbon exist as the hardest (diamond)
and one of the softest substances known
(graphite) ?
• Why do boiling points decrease with altitude?
• Why is ice less dense than water?
• What are the structures of crystals?
• Why does dry ice (CO2) sublimate?
Solids and Liquids
Intermolecular Forces
Ion-dipole, dipole-dipole and H-bonding,
dipole-induced dipole, induced dipoleinduced dipole
Liquids
Phase Diagrams
Vapor pressure and
temperature, Critical
T & P, Surface
tension and
viscosity
Show relation of solid,
liquid, and gas phases
with change in T and
P
Solids
Unit cells, metal
structures, formulas
and structures of ionic
compounds,
Molecular, network,
and amorphous solids
Properties of Solids
Lattice energy, heat of
fusion, melting point
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Dispersion forces
Van der Waal’s forces
Summary of Intermolecular Forces
1. Ion-dipole Forces
••
O
δ+ H
Hδ
+
-
+
-
-
+
-
+
+
-
+
-
H
••
O
••
H
δ
δ••
+
+
δ
-
δ
+
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2. Dipole-Dipole Interactions
-
The positive and negative ends of polar molecules interact
with each other, resulting in a net force of attraction.
-
+
+
+
-
-
+
+
+
-
-
The strongest type of dipole-dipole interaction
involving a hydrogen on one molecule (attached to a
F, O, or N) and either F, O, or N on another
molecule.
About 10 % as strong as an ordinary covalent
bond so approximately 15-40 kJ/mol.
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δ+
+
H δ
O
δ-
H δ+
δ+
H
δ+
H
O
δ-
H δ+
+
H δ
O
δ-
O
δδ+ H
H
+
H δ
O
δ-
δ+ H
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The Double Helix of DNA is held together
by hydrogen bonding
adenine
thymine
white = hydrogen
blue = nitrogen
black = carbon
red = oxygen
The polymer Nylon is also held together by hydrogen bonding
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Why is the hydrogen bond considered a
“special” dipole-dipole interaction?
Decreasing molar mass
Decreasing boiling point
DISPERSION FORCES
3. Dipole-Induced Dipole Interactions
Polarizability - a measure of the extent to which the
electron cloud of an atom or molecule can be
distorted by an external electric charge.
In general, larger atoms or molecules are more
easily polarizable than smaller ones (more shells,
etc), and so experience larger dipole-induced dipole
interactions.
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4. Induced dipole - Induced dipole
Interactions (London Forces)
Even nonpolar molecules
and uncombined atoms
have attractive forces
between them, otherwise
they would never
condense or solidify.
Which straight-chain alkane (CnH2n+2) molecule has
greater attractive intermolecular forces ?
H H
H
H-C-C-C-H
H H
H H H
H
H - C - C - C - C - C - C -H
H H
H
H H
H H H
H
H H
H
H H
H H H
H
H-C-C-C-H
H H
H
H - C - C - C - C - C - C -H
H H
H H H
H
The molecule with the longer chain because there are
more points of “attachment” via London Forces.
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Straight-Chain
Alkane
Effect of Dispersion Forces on Melting Points
of Nonpolar Compounds
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Properties of Liquids
Surface tension is the amount of energy required to stretch
or increase the surface of a liquid by a unit area.
Strong
intermolecular
forces
High
surface
tension
Photo credit: Microsoft Encarta
Properties of Liquids
Cohesion is the intermolecular attraction between like molecules
Adhesion is an attraction between unlike molecules
Adhesion
Cohesion
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Properties of Liquids
Viscosity is a measure of a fluid’s resistance to flow.
Strong
intermolecular
forces
High
viscosity
Crystals and Solids
Crystal structures covered in the
laboratory. Read Sections 11.4 – 11.7,
p. 446-462
CsCl
ZnS
CaF2
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Condensation
Evaporation
T2 > T1
The equilibrium vapor pressure is the vapor pressure
measured when a dynamic equilibrium exists between
condensation and evaporation
H2O (l)
H2O (g)
Dynamic Equilibrium
Rate of
Rate of
= evaporation
condensation
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Enthalpy of Vaporization, ∆Hvap
(Molar Heat of Vaporization)
Liquid
vaporization
heat energy absorbed
by liquid
Vapor
∆Hvap = amount of energy required to evaporate 1
mole of a liquid under constant pressure
Molar Heat of Vaporization and Boiling Points
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Before
Evaporation
At
Equilibrium
The tendency for a liquid to
evaporate increases as 1. the temperature rises
2. the surface area
increases
3. the intermolecular
forces decrease
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Vapor Pressure and Temperature
Diethyl ether - dipole-dipole
interactions
Water - hydrogen
bonding (stronger)
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Boiling Point
“If you have a beaker of water
open to the atmosphere, the
mass of the atmosphere is
pressing down on the surface. As
heat is added, more and more
water evaporates, pushing the
molecules of the atmosphere
aside. If enough heat is added, a
temperature is eventually
reached at which the vapor
pressure of the liquid equals the
atmospheric pressure, and the
liquid boils.”
Normal boiling point - the temperature at
which the vapor pressure of a liquid is equal to the
external atmospheric pressure of 1 atm.
Increasing the external atmospheric pressure
increases the boiling point
Decreasing the external atmospheric pressure
decreases the boiling point
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Location
Elevation (ft)
San Francisco
Salt Lake City
Denver
Mt. Everest
sea level
4400
5280
29,028
Boiling Point H2O (oC)
100.0
95.6
95.0
76.5
The Clausius-Clapeyron Equation
Relationship between vapor pressure, T, and ∆Hvap
ln P = −
∆H vap
RT
R = 8.314 J/mol·K
+C
C = constant that depends on
the compound’s volatility
We can use this equation to measure ∆Hvap for a given
compound if we know P and T at two different points:
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Problem 11.88 – Estimate the molar heat of vaporization of a
liquid whose vapor pressure doubles when the temperature is
raised from 85oC to 95oC.
The melting point of a solid
or the freezing point of a
liquid is the temperature at
which the solid and liquid
phases coexist in equilibrium
Freezing
H2O (l)
Melting
H2O (s)
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Liquid – Solid Equilibrium
solid → liquid
→
The melting point of the solid, or the freezing point of the
liquid, is the temperature at which the two phases are in
equilibrium.
rate melting → rate freezing
(dynamic equilibrium)
→
∆Hfusion = molar heat of fusion, energy in kJ
required to melt one mole of solid
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H2O (g)
Molar heat of sublimation
(∆Hsub) is the energy required
to sublime 1 mole of a solid.
Deposition
Sublimation
H2O (s)
∆Hsub = ∆Hfus + ∆Hvap
( Hess’s Law)
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Heating – Cooling Curves
Temperature
Vapor
Boiling point
Melting point
Solid and
liquid in
equilibrium
A
C
D
Liquid and vapor
in equilibrium
Liquid
B
Solid
Time
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Phase Diagrams
States of matter as a function of temperature and pressure
Pressure →
liquid
solid
gas
Temperature →
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Phase Diagram for Water
Phase Diagram for Water
1. Curve AD is the equilibrium vapor pressure
curve where the liquid and gas are in equilibium.
2. Line AC is the solid-liquid equilibrium curve. Note
that the slope is negative for water - most
substances have a positive slope here. A negative
slope means that the substance becomes a liquid
when the pressure is increased.
3. Point A is the Triple Point - all three phases are
in equilibrium at once.
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On the left is water; water is unusual because the solid is less dense
than the liquid. For most substances, the solid is more dense than the
liquid like for benzene (on the right).
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Water has a Maximum Density at 4 oC
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Phase Diagram for CO2
Critical
point
1. Positive slope for the solid-liquid interface (normal)
2. Sublimation occurs at room temperature under 1
atm of pressure.
3. There is a “critical point” at 73 atm and 31 oC
(varies from substance to substance).
4. Above the critical point the substance is known as
a “supercritical fluid” (not a liquid; not a gas).
Increasing the pressure normally would convert a
gas to a liquid but doesn’t happen.
5. Supercritical fluids have enhanced solvation
abilities, i.e. the fluid will dissolve greater amounts
of solute than normal.
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