The Kinetic Molecular Theory of Liquids & Solids
 Distance between gas molecules are so great at
ordinary temperatures and pressures (25 oC 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.
Gases, Liquids, & Solids
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.
Dipole-Dipole
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.
Solids vs Liquids
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.
 *It is important to remember that the 3 intermolecular forces
mentioned above are relatively weak in comparison to the
covalent bond within a molecule.
Hydrogen Bonding
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 and the greater the
dispersion force.
 Polarizability decreases from left to right across a period
because the increasing Zeff(effective nuclear charge) shrinks atomic
size and holds the electrons more tightly and the weaker the
dispersion force.
 Cations are less polarizable than their parent atoms because
they are smaller; anions are more polarizable because they
are larger.
 This concept also applies to nonpolar molecules such as H2,
Example 1
Which of the following have the highest
boiling point:
a. Argon
b. Neon
c. Xenon
d. Krypton
Example 2
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 3
Would you expect to find hydrogen bonds
in:
Acetic acid
Diethyl ether
Hydrazine
Example 4
What types of intermolecular forces are
present in:
nitrogen
chloroform
carbon dioxide
ammonia
water vapor and hydrogen gas
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.
Crystal Structures
Ionic and atomic 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.
Example 5
Classify each of the following substances
as atomic solid, molecular solid, ionic
solid, or network covalent solid according
to the type of solid it forms:
Gold
Carbon dioxide
Lithium Fluoride
Krypton
Quartz (empirical formula SiO2)
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 25oC 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:
 T1 and P1 are a corresponding temperature (in Kelvin) and vapor pressure (in
atm of mmHg)
 T2 and P2 are the corresponding temperature and pressure at another point
 ΔHvap is the molar enthalpy of vaporization (energy associated with
vaporization)
 R is the gas constant (8.31 J mol−1K−1)
 This can be used to predict the temperature at a certain pressure, given the
temperature at another pressure, or vice versa. Alternatively, if the
corresponding temperature and pressure is known at two points, the enthalpy
of vaporization can be determined.
Example 6
Benzene has a vapor pressure of 183
mmHg at 40 oC. Taking its heat of
vaporization to be 30.8 kJ/mol. Calculate
its vapor pressure at 25 oC.
Important Phase Change Terms
 Critical temperature (also called the critical point) – 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.
 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).
 Normal Boiling Point – (of a liquid) is the temperature at
which the vapor pressure of the liquid is exactly 1 atm.
 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?
Triple Point Diagrams
Heating Curve
MC #’s 1-5
 Use the following answers for questions 1 - 5.
(A) A network solid with covalent bonding
(B) A molecular solid with zero dipole moment
(C) A molecular solid with hydrogen bonding
(D) An ionic solid
(E) A metallic solid
1. Solid ethyl alcohol, C2H5OH
2. Silicon dioxide, SiO2
3. Silver
4. Diamond
5. Carbon tetrachloride
MC #’s 6-9
 Use these answers for questions 6-9.
(A) hydrogen bonding
(B) hybridization
(C) ionic bonding
(D) resonance
(E) van der Waals forces (London dispersion forces)
6. Is used to explain why iodine molecules are held
together in the solid state
7. Is used to explain why the boiling point of HF is greater
than the boiling point of HBr
8. Is used to explain the fact that the four bonds in methane
are equivalent
9. Is used to explain the fact that the carbon-to-carbon
bonds in benzene, C6H6, are identical
MC #10
 CH3CH2OH boils at 78 °C and CH3OCH3 boils
at - 24 °C, although both compounds have the
same composition. This difference in boiling
points may be attributed to a difference in
(A) molecular mass
(B) density
(C) specific heat
(D) hydrogen bonding
(E) heat of combustion
MC #11
 The melting point of MgO is higher than that of
NaF. Explanations for this observation include
which of the following?
I. Mg2+ is more positively charged than Na+
II. O2¯ is more negatively charged than F¯
III. The O2¯ ion is smaller than the F¯ ion
(A) II only
(B) I and II only
(C) I and III only
(D) II and III only
(E) I, II, and III
MC #’s 12 - 14
 Questions 12-14 refer to the phase diagram below of a pure
substance.
(A) Sublimation
(B) Condensation
(C) Solvation
(D) Fusion
(E) Freezing
12. If the temperature increases from 10° C to 60° C at a constant
pressure of 0.4 atmosphere, which of the processes occurs?
13. If the temperature decreases from 110° C to 40° C at a constant
pressure of 1.1 atmospheres, which of the processes occurs?
14. If the pressure increases from 0.5 to 1.5 atmospheres at a constant
temperature of 50° C, which of the processes occurs?
FRQ #1
 The phase diagram for a pure
substance is shown above. Use this
diagram and your knowledge about
changes of phase to answer the
following questions.
 (a) What does point V represent?
What characteristics are specific to
the system only at point V?.
 (b) What does each point on the
curve between V and W represent?
 (c) Describe the changes that the
system undergoes as the
temperature slowly increases from X
to Y to Z at 1.0 atmosphere.
 (d) In a solid-liquid mixture of this
substance, will the solid float or
sink? Explain.
FRQ #2
 For each of the following, use appropriate chemical
principles to explain the observation.
(a) At room temperature, NH3 is a gas and H2O is a
liquid, even though NH3 has a molar mass of 17 grams
and H2O has a molar mass of 18 grams.
(b) C (graphite) is used as a lubricant, whereas C
(diamond) is used as an abrasive.
(c) Pouring vinegar onto the white residue inside a
kettle used for boiling water results in a fizzing/bubbling
phenomenon.