science

1. 1
Chapter 12
Intermolecular
Forces and the
Physical Properties
of Liquids and
Solids
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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