The Kinetic Molecular Theory of
Liquids & Solids – Chapter 9
 Distance between gas molecules are so great at
ordinary temperatures and pressures (25 *C and
1atm) that there is no real interaction between
gas molecules.
 Liquids – the molecules are so close together
that there is little empty space. Allowing for a
definite volume but taking the shape of it’s
container.
 Solids – molecules are held rigidly in a position
with virtually no freedom of motion. So that they
have a definite volume and shape.
Intermolecular Forces
 These are attractive forces between molecules:
Dispersion (London) Forces
Dipole-Dipole Forces
 Hydrogen Bonding
 These forces are partly responsible for the nonideal gas law behavior discussed earlier. And
these forces are why liquids and solids do not
present “ideal behavior”.
 Keep in mind that intramolecular forces are
forces within a molecule, and intermolecular
forces are between molecules.
Intermolecular Forces…………(cont.)
 Boiling points and melting points often reflect the
strength of these intermolecular forces.
 Molecular substances tend to have the following
characteristics:
 Non-conductors of electricity when pure, for example pure
water and ethyl alcohol are 2 examples of molecular
substances that, when pure, will not conduct electricity.
 Generally, are insoluble in water but soluble in non-polar
solvents.
 They tend to have low melting and boiling points. The
stronger the intermolecular force the higher the boiling or
melting point.
1. Dispersion (London) Forces:
The attractive forces that arise as a result
of temporary dipoles induced in atoms or
molecules.
This happens when something (cation or polar
compound) “throws off” the distribution of electrons in
an atom or molecule.
These forces can only occur in non-polar molecules.
As molar mass (MM) increases the dispersion forces
become stronger, consequently the boiling point of
non-polar molecules tend to increase with MM.
Reference: Table 9.2 – page 234.
2. Dipole-Dipole Forces:
These are attractive forces between polar
molecules.
The larger the dipole moment then the greater the
dipole force.
Polar molecules tend to have slightly higher boiling
and melting points than non-polar substances of
similar MM.

Reference: Table 9.3 – page 235.
3. Hydrogen Bonding:
This a special, strong type of dipole-dipole
interaction between hydrogen and polar
bond in elements nitrogen, oxygen, and
fluorine. Only between N—H, O—H, and
F—H bonds have hydrogen bonding.
These types of bonds tend to not follow the
relationship between MM and boiling points.
Reference: Table 9.4 – page 237.
 *It is important to remember that the 3 intermolecular forces
mentioned above are relatively weak in comparison to the
covalent bond within a molecule.
Polarizability
 The ease with which the electron cloud of a
particle can be distorted is called its
polarizability. Smaller atoms (or ions) are less
polarizable than larger ones because their
electrons are closer to the nucleus and therefore
are held more tightly. Thus, we observe several
trends:
Polarizability increases down a group because atomic
size increases, so the larger electron clouds are farther
from the nucleus and, thus, easier to distort.
Polarizability decreases from left to right across a period
because the increasing Zeff shrinks atomic size and
holds the electrons more tightly.
Cations are less polarizable than their parent atoms
because they are smaller; anions are more polarizable
because they are larger.
Example
What types of intermolecular forces exist
between the following pairs:
a. Hydrogen bromide and hydrogen sulfide
b. Diatomic chlorine and carbon
tetrabromide
c. Diatomic iodine and the nitrate ion
d. ammonia and benzene
Example 9.5 – page 237
Would you expect to find hydrogen bonds
in:
Acetic acid
Diethyl ether
Hydrazine
Example 9.6 – page 238
What types of intermolecular forces are
present in:
nitrogen
chloroform
carbon dioxide
ammonia
Properties of Liquids
 Intermolecular forces give rise to a number of
structural features and properties of liquids.
Notable, two features: surface tension and
viscosity.
 Surface Tension: the amount of energy required to stretch
or increase the surface of a liquid be a unit area (for
example 1 cm2).
• Liquids that have strong intermolecular forces have high
surface tensions. So because of hydrogen bonding, water has
a greater surface tension than most other liquids.
 Viscosity: a measure of a fluids resistance to flow
(unit = N * s/m2).
• Liquids that have strong intermolecular forces have higher
viscosities than those with weaker intermolecular forces.
• Examples: Water = 1.01 x 10-3
Acetone = 4 x 10-3
Blood = 3.16 x 10-4
Glycerol= 1.49
Network Covalent Solids
Most covalent solids are molecular
however some form a network of repeating
patterns, these are called network
covalent solids.
Plastics, and allotropes of carbon
(diamond and graphite) do this.
They tend to have high melting points
compared to molecular covalent.
Ionic Solids
 Consist of anions and cations that are held in a regular
repeating arrangement by strong ionic bonds, or
electrostatic interactions.
 Ionic solids tend to have high boiling and melting points.
 Ionic solids do not conduct electricity, because the ions
are fixed.
 Many are soluble in water. Would solutions of ionic solids
be good or poor conductors of electricity.
 The strength of the ionic bond depends on 2 things:
 The charge of the ion, CaO has a stronger ionic bond than NaCl
 The size of the ions: d = rcation + ranion
Metallic Solids
 There is no real “bond” between metals.
Basically the electrons of metals arrange around
each other’s nuclei. This forms a “sea” of
electrons, so that the electrons move about in
the crystal.

They are good conductors of electricity, have
high thermal conductivity, tend to have luster
(reflect light), are ductile and malleable, and are
generally insoluble in water.
Example 9.7- page 243
Reference: Table 9.5- page 244
For each species in a column A, choose
the description that best applies in column
B.
Crystal Structures
Ionic solids tend to crystallize in definite
geometric forms. These geometric forms
are made up of unit cells (the smallest
structural unit in the 3-D repeating
pattern).
There are 14 possible crystalline
structures but we will focus on the main 3.
3 Main Crystal Structures
Simple Cubic (SC)
Face Centered Cubic (FCC)
Body Centered Cubic (BCC)
How they differ…
 Number of atoms per unit cell
SC
FCC
BCC
 The relation between side of cell, s (in nm), and the
atomic radius, r. This relation is expressed in the
following equation and offers and experimental way to
determine the atomic radius of a metal:
SC: 2r = s
BCC: 4r = s1/3
FCC: 4r = s21/2
 The percentage of empty space – the greater the
amount of empty space the more unstable the structure.
 Reference table 9.6 – page 245
Example 9.8 – page 246
Silver crystallizes with a face centered
cubic unit cell 0.407 nm on an edge.
Calculate the atomic radius of silver.
A Vapor
A gas is a substance that is normally in the
gaseous state at ordinary temperatures
and pressures.
A vapor is the gaseous form of a
substance that is a liquid or a solid at
normal temperature and pressure –
generally 25*C and 1 atm.
Vapor Pressure
 The pressure exerted by the vapor over the liquid
remains constant (in a sealed container the rate of
condensation becomes equal to the rate of evaporation).
 Diagram:
 Thus there is a state of dynamic equilibrium between the
liquid and the vapor.
 So as long as the pressure exerted by the liquid is less
than that of the atmosphere than the liquid will not boil.
 Diagram:
 Once the vapor pressure exerted by the liquid is equal to
that of the atmosphere the liquid boils (this is considered
the normal boiling point)
 Diagram:
Vapor Pressure………………..(cont.)
 Vapor pressure always increases with
temperature. The higher the average kinetic
energy then the more collisions and the stronger
the force exerted by the liquid (the vapor
pressure).
What pressure would be required to boil water at a
temperature of 70*C if water exerts a pressure of
24mmHg?
 The stronger the intermolecular forces in the
liquid the lower the vapor pressure of that liquid.
Examples: Water – 24 mmHg
Ether – 537
mmHg
Why?
The Clausius-Clapeyron Equation:
Refer to page 227
R = 8.31 J/mol K
T = temperature in K
P = pressure in atm or mmHg
Example 9.2 – page 228
Benzene has a vapor pressure of 183
mmHg at 40 *C. Taking its heat of
vaporization to be 30.8 kJ/mol. Calculate
its vapor pressure at 25 *C.
Important Terms
 Critical temperature – the temperature above which the
liquid phase of a pure substance cannot exist. Basically,
above this temperature the substance is a gas and
below it the substance is a liquid.
 Reference: Table 9.1 – page 230
 Critical pressure – conversely, this is the pressure that
must be applied at the critical temperature to cause
condensation.
 Supercritical Fluid – just means that the substance is
above the critical temperature (refers to a gas).
 Melting point – pressure has little effect on melting point.
Most solids are denser than their liquid form. There is a
very important exception to the aforementioned, guess
what it is?
Heating Curve
Triple Point Diagrams