This document discusses intermolecular forces and the physical properties of liquids and solids. It covers topics such as the different types of intermolecular forces, how these forces determine properties of liquids like viscosity and vapor pressure, crystal structure of solids, and phase changes between the different states of matter. The key points are that intermolecular forces are weaker than ionic or covalent bonds but still significantly impact properties, liquids and solids exist as a result of these intermolecular interactions, and phase changes occur with the addition or removal of heat to overcome these attractive forces.

1. 1
Chapter 12
Intermolecular
Forces and the
Physical Properties
of Liquids and
Solids
Insert picture from
First page of chapter
Copyright McGraw-Hill 2009

2. Copyright McGraw-Hill 2009 2
12.1 Intermolecular Forces
• Intermolecular forces are the attractive forces
holding particles together in the condensed
(liquid and solid) phases of matter
• Result from coulombic attractions
– Dependent on the magnitude of the charge
– Dependent on distance between charges
• Weaker than forces of ionic bonding
• Involve partial charges

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The Three Phases of Matter
condensed phases

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Types of Intermolecular Forces
• van der Waals forces –between atoms
and molecules of pure substances
– Dipole-dipole interactions – attractive forces
between polar molecules
– Hydrogen bonding – attractive force in polar
molecules containing a H atom bonded to a
small, highly electronegative element (N, O
and F)
– (London) Dispersion forces – attractive
forces arising from instantaneous dipoles and
induced dipoles

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Arrangement of Polar Molecules in a Liquid
and a Solid.

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• Intermolecular forces determine certain
physical properties.

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Hydrogen Bonds Between HF Molecules

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Effect of Molar Mass and Hydrogen Bonding
on Boiling Points

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Instantaneous and Induced Dipoles
Magnitude depends on the ability to be polarized which
is greater for larger molecules.

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Polarization and Molar Mass

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What kind(s) of intermolecular forces exist in
CH2ClCH2COOH(l)
C C
H
Cl
H
C
H
H
O
O H
dispersion forces
dipole-dipole interactions
hydrogen bonding

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–Occur in mixtures of ionic and polar
species
–Coulombic attraction between ions and
polar molecules
–Dependent upon
• Size and charge of ion
• Dipole moment of the molecule
• Size of the molecule
–Can also be repulsive
•Ion-dipole Interactions

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12.2 Properties of Liquids
• Surface Tension – a quantitative measure of the
elastic force at the surface of a liquid
• Manifestations
– Formation of a mensicus
– Capillary action which
results from a combination of
• Cohesion (attractions between
like molecules, cohesive forces)
• Adhesion (attractions between
unlike molecules, adhesive forces)

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Effect of Surface Tension

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Intermolecular Forces: Surface
versus Interior of a Liquid

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Cohesion and Adhesion
Water Mercury
Adhesion > Cohesion Cohesion > Adhesion

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• Viscosity – a measure of a fluid’s
resistance to flow
–Units: N.s/m2
–The higher the viscosity the greater the
resistance to flow
–Varies inversely with temperature
–Stronger intermolecular forces produce
higher viscosities

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Glycerol – high viscocity
due to
• Three hydrogen bonding
sites
• Molecular shape

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• Vapor Pressure of a Liquid
–Depends on the magnitude of intermolecular
forces
–Temperature dependent
T1<T2

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• Equilibrium Vapor Pressure – a dynamic
state
− Evaporation (vaporization) – liquid
molecules escape into the gas phase
− Condensation – gas molecules return to
the liquid phase

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Comparison of Rates of Evaporation and
Condensation at Constant Temperature

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Effect of Temperature and Intermolecular Forces
on Vapor Pressure

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• Clausius-Clapeyron Equation – linear relation
between temperature and vapor pressure
− At two temperatures, T1 and T2:
where R = 8.314 J/K . mol

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An unknown compound exhibits a vapor pressure of
255 mmHg at 25.5oC and 434 mmHg at 48.8oC.
What is DHvap of this substance?

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K
298.65
273.15
C
25.5
1 

 o
T
K
321.95
273.15
C
48.8
2 

 o
T








D

K
298.65
1
K
321.95
1
mol
8.314J/K
mmHg
434
mmHg
255 vap
H
ln
 
3
3
3.3484x10
3.1061x10
mol
8.314J/K
0.53178 



D


vap
H









D

1
2
vap
2
1 1
1
T
T
R
H
P
P
ln

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vap
H
D





 1
4
K
2.423x10
mol)
.314J/K
0.53178)(8
(
 
1
4
K
2.423x10
mol
8.314J/K
0.53178 



D


vap
H
vap
4
J/mol
10
x
1.82 H
D

vap
kJ/mol
18.2 H
D


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12.3 Crystal Structure
• Crystalline solid – possesses rigid and
long-range order
• Lattice structure – arrangement of
particles in a crystalline solid
– Depends on nature of particles
– Depends on size of particles
• Stability depends on type of force between
particles (ionic or covalent bonds and/or
intermolecular forces)

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unit cell – basic repeating structural unit of
a crystalline solid
lattice point – each atom ion or molecule in
a unit cell
single unit cell 3-D array of unit cells

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Seven Types of Unit Cells

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Coordination number – number of atoms
(particles) surrounding an atom in a crystal
lattice
• Indicates how tightly atoms pack
• Larger coordination numbers indicate
tighter packing

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• Types of cubic unit cells
–simple or primitive (scc)
–body-centered (bcc)
–face-centered (fcc)

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Alternate Perspective of bcc Arrangement

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Sharing of Atoms by Adjacent Unit Cells
corner atom edge atom face atom

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• Number of atoms per unit cell
–scc: 1 atom
–bcc: 2 atoms
–fcc: 4 atoms
•Allocation of atoms among unit cells
− corner atoms – 1/8 atom within unit cell
− face atoms – 1/2 atom within unit cell
− body atoms – 1 atom within unit cell

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• Closest Packing – most efficient way to
arrange atoms in a crystal
–hexagonal closest packed (ABA)
–cubic closest packed (ABC)

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Closest Packing and Cubic Unit Cells
fcc
cubic
hexagonal

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Geometric Relationships

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When silver crystallizes, it forms face-
centered cubic cells. The unit cell edge
length is 4.087A . Calculate the density of
silver.
o

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g
10
7.167
amu
10
6.022
g
1
atom
amu
107.9
cell
unit
atoms
4 22
23







m
  3
23
3
8
3
cm
10
6.827
cm
10
4.087 





 a
V
cm
10
4.087
m
1
cm
00
1
A
10
1
m
1
A
4.087 8
10






 

a
3
3
23
22
g/cm
10.5
cell
/unit
cm
10
6.827
cell
g/unit
10
7.167




 

V
m
d
Mass of unit cell
Volume of unit cell
Density of unit cell

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• X-ray diffraction utilizes the scattering of
X-rays and the resulting scattering patterns
to deduce arrangement of particles

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− Bragg equation

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12.4 Types of Crystals
• Ionic Crystals
– Composed of anions and cations
– Held together by coulombic forces
– Anions generally are bigger than cations
– Size and relative number of each ion
determines the crystal structure

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Unit cell of NaCl as Defined by Cl− or Na+

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Examples of Ionic Crystal Lattices
CsCl ZnS CaF2

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How many of each ion are contained within
a unit cell of CaF2?

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Ca2+
8 corner ions x 1/8= 1 ion
6 face ions x ½ = 3 ions
4 ions of Ca2+
F−
8 body ions x 1 = 8 ions of F−
Ca2+
F−

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• Covalent crystals
– Held together by covalent bonds
diamond graphite

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• Molecular crystals
– Lattice points occupied by molecules
– Held together by intermolecular forces
water

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• Metallic crystals
– Lattice points occupied by atoms
– Generally bcc, fcc, hexagonal closest packed
– Very dense
– Bonding arises from delocalized electrons

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12.5 Amorphous Solids
• Lack regular arrangement of atoms
• Glass is a familiar and important
amorphous solid
– Transparent fusion of inorganic materials
– Chief component SiO2
– Behaves more as a liquid than a solid

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Comparison of crystalline quartz and amorphous
quartz glass
crystalline amorphous

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12.6 Phase Changes
• Phase – homogenous part of a system
that is separated from the rest of the
system by a well-defined boundary
• Phase change – transition from one
phase to another
–Caused by the removal or addition of
energy
–Energy involved is usually in the form of
heat

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The Six Possible Phase Changes

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• Liquid-Vapor Phase Transition
– Boiling point – the temperature at which the
vapor pressure of liquid equals atmospheric
pressure
– Molar heat of vaporization (DHvap) – the
amount of heat required to vaporize one mole
of a substance at its boiling point usually in
kJ/mol
– Dependent on the strength of intermolecular
forces
– Condensation – opposite of vaporization

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• Critical temperature (Tc) – the
temperature above which a gas cannot be
liquified by application of pressure
• Critical pressure (Pc) – the pressure that
must be applied to liquefy a gas at Tc.
• Supercritical fluid – the fluid that exists
above Tc and Pc.

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• Solid-Liquid Phase Transition
– Freezing – transformation of liquid to solid
– Melting (fusion) – opposite of freezing
– Melting point of solid (or freezing point of
liquid) – temperature at which the solid and
liquid phases coexist in equilibrium
• Dynamic equilibrium in which the forward
and reverse processes are occurring at the
same rate
– Molar heat of fusion (DHfus) – energy to melt
one mole of a solid usually in kJ/mol

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Typical Heating Curve

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• Solid-Vapor Phase Transition
–Sublimation – process by which
molecules go directly from the solid
phase to the vapor phase
–Deposition – reverse of sublimation
–Molar heat of sublimation (DHsub) –
energy required to sublime one mole of
solid usually in kJ/mol
iodine

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Heating Curve for Water

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Calculate the amount of energy (in kJ)
required to convert 125 g of ice at 10.0oC
to liquid water at the normal boiling point.
Assume that the specific heat of ice is
2.050 J/goC.

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Energy to warm ice from 10oC to 0oC
C
10.0
C)
10.0
(
C
0.0 o
o
o




DT
kJ
10
2.563
C
10.0
C
g
J
2.050
g
125 3
o
o






D
 T
ms
q
Energy to melt ice at 0oC
kJ
10
4.169
mol
kJ
6.01
mol
937
.
6 1
vap 



D
 H
n
q
kJ
2.563
J
10
1
kJ
J
10
2.563 3
3




mol
6.937
g
18.02
mol
g
125 


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Energy to warm water from 0.0oC to 100.0oC
C
100.0
C
0.0
C
100.0 o
o
o



DT
J
10
5.230
C
100.0
C
g
J
4.184
g
125 4
o
o






D
 T
ms
q
Total energy required
kJ
10
5.230
J
10
1
kJ
1
J
10
5.230 1
3
4





96.6kJ
)
kJ
10
5.230
(
)
kJ
10
4.169
(
kJ
2.563 1
1






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12.7 Phase Diagrams
• Phase diagram – summarizes the
conditions (temperature and pressure) at
which a substance exists as a solid, liquid
or gas
– Divided into three regions (solid, liquid, gas)
– Phase boundary line – line separating any
two regions
– Triple point – the point at which all three
phase boundary lines meet

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Phase Diagram of CO2

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Heating CO2 Starting at 100oc and 1 atm

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Phase Diagram of H2O

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What is a) the normal* melting point, b) the normal*
boiling point and c) the physical state of the substance
at 2.0 atm and 110o C?
*normal – measured at 1.00 atm

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normal melting point normal boiling point
~135oC ~200oC

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solid physical state

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Key Points
• Intermolecular forces
– Dipole-dipole interactions
– Hydrogen bonding
– (London) dispersion forces
• Properties of liquids
– Surface tension
– Viscosity
– Vapor pressure
• Clausius-Clapeyron equation

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• Crystal structure
– Unit cells
• Lattice point
– Packing spheres
• Coordination number
• Cubic unit cells
– Closest Packing
• Types of crystals
– Ionic
– Covalent
– Molecular
– Metallic

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• Amorphous solids
• Phase changes
– Liquid-vapor transitions
• Boiling point
• Heat of vaporization
• Critical temperature and pressure
– Solid-liquid transitions
• Melting point
• Heat of fusion
• Phase diagrams