Chapter 11
Intermolecular Forces,
Liquids, and Solids
CONTENTS:
1. State of Matter
2. Intermolecular Forces
3. Some Properties of liquids
4. Phase Changes
5. Vapor Pressure
6. Phase Diagrams
7. Structures of Solids
8. Bonding in Solids
1
States of Matter
The fundamental difference between states of matter is the
distance between particles.

Liquid molecules are held closer together than gas
molecules, but not so rigidly that the molecules can slide
past each other.

Solid molecules are packed closely together. Molecules
are so rigidly packed that they cannot easily slide past
each other.
States of Matter
 Converting a gas into a liquid or solid requires the
molecules to get closer to each other: cool or
compress.
 Converting a solid into a liquid or gas requires the
molecules to move further apart from each other:
heat or decrease the pressure
2
States of Matter
Because in the solid and liquid states particles are
closer together, we refer to them as condensed
phases.
The States of Matter
 The state of a substance, at
a particular temperature and
pressure, depends on :


the kinetic energy of the
particles;
the strength of the
attractions between the
particles.
3
Intermolecular Forces
 Intermolecular forces are the forces that exist
between molecules; they are the attractive forces that
hold together the particles of a liquid or a solid. They
are typically much weaker than covalent or ionic
bonds (intramolecular forces).
Examples
So it takes much more energy to break the covalent bond between H and
Cl than it does to vaporize HCl
4
Intermolecular Forces
These intermolecular forces as a group are referred
to as van der Waals forces.
Effect of Intermolecular Forces on Physical
Properties of Liquids and Solids
 When a substance undergoes a physical change,
the intermolecular forces are broken and not the
intramolecular forces.
 When a substance condenses the intermolecular
forces are formed.
 The strength of the intermolecular forces in a
compound are reflected in its melting point and
boiling point. The stronger the intermolecular
forces, the more energy is required to overcome
those forces
5
van der Waals Forces
 Dipole-dipole interactions
 London dispersion forces
6
ClF3
(polar molecule)
7
What is Ion-Dipole Force (IDF)?
 An IDF exists between an ion and the
partial charge on the end of a polar
molecule.
 IDF (and Ionic Bonding) are the strongest
of all intermolecular forces
Ion-Dipole Forces (IDF)
The positive ends will orient themselves toward a negative ion, and the negative
ends will orient themselves toward a positive ion. As with all electrostatic
attraction, the magnitude of the attraction increases with the size of the charge
(both on the ion and on the dipole).
8
Dipole-Dipole Interactions (DDF)
 Molecules that have
permanent dipoles are
attracted to each other.


The positive end of one is
attracted to the negative
end of the other and viceversa.
These forces are only
important when the
molecules are close to each
other.
Interactions Between two Polar
Molecules
There is a mix of attractive and repulsive forces as the molecules tumble. The
overall effect is a net attraction
The greater the polarity of the molecules, the stronger the attractions between
them.
9
Dipole-Dipole Interactions
For molecules of approximately the same size, boiling
point increases with increasing dipole moment.
The more polar the molecule, the higher is its boiling
point.
London Dispersion Forces
Many nonpolar substances exist as condensed
phases.
The intermolecular forces holding nonpolar
atoms or molecules together are known as
London dispersion forces (LDF).
10
London Dispersion Forces
While the electrons in the 1s orbital of helium would
repel each other (and, therefore, tend to stay far
away from each other), it does happen that they
occasionally wind up on the same side of the atom.
London Dispersion Forces
At that instant, then, the helium atom is polar, with an
excess of electrons on the left side and a shortage on
the right side.
11
London Dispersion Forces
Another helium nearby, then, would have a dipole
induced in it, as the electrons on the right side of
helium atom 1 repel the electrons in the cloud on
helium atom 2.
London Dispersion Forces
London dispersion forces, or dispersion forces, are
attractions between an instantaneous dipole and an
induced dipole.
12
London Dispersion Forces
 These forces are present in all molecules, whether
they are polar or nonpolar.
 The tendency of an electron cloud to distort in this
way is called polarizability.
Factors Affecting London Forces
 The strength of dispersion forces tends to increase
with increased molecular weight.
 Larger atoms have larger electron clouds which are
easier to polarize.
13
Factors Affecting London Forces
 The shape of the molecule affects the
strength of dispersion forces: long,
skinny molecules (like n-pentane tend
to have stronger dispersion forces
than short, fat ones (like neopentane).
This is due to the increased surface
area in n-pentane.
Example
 Long thin molecules can develop bigger
temporary dipoles.
Butane:
b.p.: -0.5°C
H3C
2-methylpropane:
b.p.: -11.7°C
H2
C
C
H2
CH3
CH3
H3 C
CH
CH3
14
Which Have a Greater Effect?
Dipole-Dipole Interactions or Dispersion Forces ?
 If two molecules are of comparable size and
shape, dipole-dipole interactions will likely be
the dominating force.
Example
15
Which Have a Greater Effect?
Dipole-Dipole Interactions or Dispersion Forces ?
 If one molecule is much larger than another,
dispersion forces will likely determine its
physical properties.
Example
 Which of the following would have a higher
boiling point CCl4 or CHCl3 ?
16
CCl4 is less polar and bigger than CHCl3 but the
LDF for CCl4 is greatly larger than the combined
LDF and DDF for CHCl3
CHCl3 61.2°C
CCl4 76.8°C
Hydrogen bonding (HB)
 The dipole-dipole interactions experienced
when H is bonded to N, O, or F are unusually
strong.
 We call these interactions hydrogen bonds.
 The HB occurs when H is covalently
bonded to F, O, or N.
17
Example
Example:
5,6, and 7 Groups’ Boiling Points
18
Hydrogen Bonding
 Hydrogen bonding
arises from the high
electronegativity of
nitrogen, oxygen, and
fluorine.
Also, when hydrogen is bonded to one of those
very electronegative elements, the hydrogen
nucleus is exposed.
Hydrogen Bonding
Dark blue represents oxygen and red hydrogen. Magenta bonds
are bonds within each water molecule, light blue bonds are
hydrogen bonds between molecules.
19
Density of water
Temp
(oC)
Density
(kg/m3)
+100
+80
+60
+40
+30
+25
+22
+20
+15
+10
+4
0
−10
−20
−30
958.4
971.8
983.2
992.2
995.6502
997.0479
997.7735
998.2071
999.1026
999.7026
999.9720
999.8395
998.117
993.547
983.854
Strength of HB
 HB - relatively strong form of
intermolecular attractions
 HB are the strongest forces after IDF
20
Summarizing Intermolecular
Forces
Intermolecular Forces Affect Many
Physical Properties
The strength of the attractions between
particles can greatly affect the properties of
a substance or solution.
Viscosity, surface tension, meniscus, capillary action…
21
Viscosity
 Resistance of a liquid to
flow is called viscosity.
 It is related to the ease with
which molecules can move
past each other.
Example 1
 Pushing a spoon with a small force moves it
easily through a bowl of water, but the same
force moves mashed potatoes very slowly
 Different materials have different viscosities
22
Viscosity Units
The international unit is Kg·m−1·s−1 .
Other unit is the Poise.
1 Poise = 100 centipoise (cP) = 1 g·cm−1·s−1
Viscosity tends to fall as temperature increases
Water viscosity goes from 1.79 cP to 0.28 cP in the
temperature range from 0 °C to 100 °C
23
Example 2
 The stronger the intermolecular forces,
the higher the viscosity.
Surface Tension
The molecules at the surface of this sample of
liquid water are not surrounded by other
water molecules. The bulk molecules are
surrounded by other molecules and are
equally attracted to their neighbors.
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Surface Tension
 The unbalanced attraction of molecules at the
surface of a liquid tends to pull the molecules
back into the bulk liquid .
 Surface tension results from the net inward
intermolecular attractive forces experienced
by the molecules on the surface of a liquid
causing it to behaves like a skin.
Surface Tension
 This forces liquid surfaces to contract to the
minimal area.
 Surface tension is responsible for the shape
of liquid droplets.
25
Units of Measurements
 Surface tension is the amount of energy
required to increase the surface area of a
liquid
 Surface tension is typically measured in
dynes/cm.
 Other unit include J/m2
Talk ?
Example
 Water at 20°C has a surface tension of
72.9 dynes/cm (7.29 J/ m2) compared to
22.3 dynes/cm for ethyl alcohol and 465
dynes/cm for mercury.
The surface tension decreases significantly as
temperature increases
26
Cohesive and Adhesive Forces

Cohesive forces results from the
intermolecular attractions between the
molecules. Cohesive Forces bind the
molecules to each other.

Forces of attraction between a liquid and
a solid surface (container) are called
adhesive forces
Cohesive and Adhesive Forces
 The difference in strength between cohesive forces
and adhesive forces determine the behavior of a
liquid in contact with a solid surface:
 Water does not wet waxed surfaces because the
cohesive forces within the drops are stronger than
the adhesive forces between the drops and the wax.
 Water wets glass and spreads out on it because the
adhesive forces between the liquid and the glass are
stronger than the cohesive forces within the water.
27
Formation of a Meniscus
 When liquid water is confined in a tube, its
surface (meniscus) has a concave shape
because water wets the surface.
Formation of a Meniscus
In liquid Hg the cohesive forces within the drops
are stronger than the adhesive forces
between the drops and glass. When liquid
mercury is confined in a tube, its surface
(meniscus) has a convex shape because
Mercury does not wet glass
28
Capillary Action
 Even though we usually think of water as
running downhill, it can indeed flow upwards
using a process called capillary action.
 Capillary action is the tendency of a liquid to
rise in narrow tubes
 The liquid climbs up the inside of the tube
(as a result of adhesive forces between the
liquid and the inner walls of the tube).
The smaller the diameter of the tube, the
higher the liquid rises.
Example
 Plants contain many vein like tubes that carry water from the
plant's roots upwards to the plant's highest leaves via capillary
action.
 Capillary action is responsible for moving groundwater from wet
areas of the soil to dry areas.
 The small pores of a sponge act as small capillaries, causing it
to absorb a comparatively large amount of fluid.
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Phase Changes
Energy Changes Accompanying
Phase Changes







DHsub > 0 (endothermic).
Vaporization: DHvap > 0 (endothermic).
Melting or Fusion: DHfus > 0 (endothermic).
Deposition: DHdep < 0 (exothermic).
Condensation : DHcon < 0 (exothermic).
Freezing: DHfre < 0 (exothermic).
Sublimation:
Generally enthalpy of fusion is less than enthalpy of
vaporization: it takes more energy to completely
separate molecules, than partially separate them
30
Heating Curves
 Plot of temperature changes versus heat added is a
heating curve.
The temperature of the substance does not rise during
a phase change. These points are used to calculate
DHfus and DHvap
Supercooling

Supercooling: When a liquid is cooled
below its melting point and it still
remains a liquid.

..\..\..\..\..\..\Desktop\Supercooling.mp4

Also, superheating…
Talk ?
31
Example
 Freon-11, which has the chemical formula
CCl3F, has a normal boiling point of 23.8°C.
The specific heats of CCl3F(l) and CCl3F(g)
are 0.87 J/g-K and 0.59 J/g-K, respectively.
The heat of vaporization is 24.75 kJ/mol.
 Calculate the heat required to convert 10.0 g
of Freon-11 from a liquid at -50.0°C to a gas
at 50.0°C.
Volatility and Vapor Pressure
We can commonly observe that any liquid left in an open container will,
under most conditions, eventually evaporate, even if the temperature of the
liquid is well below the normal boiling point
Some of the molecules on the surface of a liquid have enough kinetic energy
to escape the attraction of the liquid molecules on the surface into the gas
phase. As the number of molecules in the gas phase increases, some of the
gas phase molecules strike the surface and return to the liquid.
32
Volatility and Vapor Pressure
Volatility is a measure of the tendency of a liquid to vaporize
Volatility 
gaseous fraction y

liquid fraction
x
A substance with a higher vapor pressure will vaporize more
readily than a substance with a lower vapor pressure.
Volatility and vapor pressure are directly proportional to one
another
Dynamic Equilibrium
As more molecules escape the liquid, the pressure they exert increases.
The liquid and vapor reach a state of dynamic equilibrium: liquid molecules
evaporate and vapor molecules condense at the same rate.
Vapor pressure is the pressure exerted when the liquid and vapor are in
dynamic equilibrium.
33
Vapor Pressure as Function of
Temperature
The tendency of a liquid to evaporate is referred as volatility: a more
volatile liquid evaporates more readily.
The higher the temperature, the higher the average kinetic energy,
the faster the liquid evaporates
Vapor pressure increases as temperature increases.
Vapor Pressure and Boiling Point



The boiling point of a liquid is the temperature
at which its vapor pressure equals atmospheric
pressure.
The normal boiling point is the temperature at
which its vapor pressure is 760 torr.
Pressure cookers operate at high pressure. At
high pressure the boiling point of water is
higher than 100 °C . Therefore, there is a higher
temperature over which the food would be
cooked faster.
34
Phase Diagrams
Plot of Pressure versus Temperature
Phase diagrams display the state of a substance at
various pressures and temperatures and the places
where equilibria exist between phases.
Phase Diagrams
 The circled line is the liquid-vapor interface (curve).
 It starts at the triple point (T), the point at which all
three states are in equilibrium.
35
Phase Diagrams
It ends at the critical point (C); above this critical
temperature and critical pressure the liquid and vapor
are indistinguishable from each other.
Phase Diagrams
Each point along this line is the boiling point of the
substance at that pressure.
36
Phase Diagrams
 The circled line in the diagram below is the interface
between liquid and solid.
 The melting point at each pressure can be found
along this line.
Phase Diagrams
 Below the triple point the substance cannot exist in
the liquid state.
 Along the circled line the solid and gas phases are in
equilibrium; the sublimation point at each pressure is
along this line.
37
Phase Diagrams
 Below the triple point the substance cannot
exist in the liquid state.
 Along the circled line the solid and gas
phases are in equilibrium; the sublimation
point at each pressure is along this line.
Phase Diagram of Water
 Note the high critical
temperature and critical
pressure.These are due to
the strong van der Waals
forces between water
molecules.
38
Phase Diagram of Water
 The slope of the solid-liquid
line is negative (slops to the
left). This means that as the
pressure is increased at a
temperature just below the
melting point, water goes
from a solid to a liquid.
Freeze drying
Technology developed by NASA for the long distance Apollo missions
39
Phase Diagrams
 Below the triple point the substance cannot
exist in the liquid state.
 Along the circled line the solid and gas
phases are in equilibrium; the sublimation
point at each pressure is along this line.
Freeze drying (coffee)
40
Supercritical fluids
 At ordinary pressures, a substance above its
critical temperature behaves like an ordinary
gas. However, as the pressure increases its
character changes. Like a gas , the
substances still expands but its density
approaches that of a liquid.
 A substance at temperature and pressure
higher than its critical temperature and
pressure is considered as supercritical fluid
Phase Diagram of Carbon Dioxide
Carbon dioxide cannot
exist in the liquid state
at pressures below 5.11
atm; CO2 sublimes at
normal pressures.
41
SOLIDS
Solids
 We can think of solids as
falling into two groups:

crystalline, in which
particles are in highly
ordered arrangement
Examples: quartz ,
sugar.
Crystalline solids tend to
melt at specific
temperatures. Therefore,
have a narrow range of
intermolecular forces
42
Solids

amorphous, in which there
is no particular order in the
arrangement of particles.
Tend to melt over a range
of temperatures.
Therefore, amorphous solids
have variable intermolecular
forces
Examples: polymer, glass
Crystalline Solids
Because of the order in
a crystal, we can focus
on the repeating
pattern of arrangement
called the unit cell.
43
Unit Cells in Crystalline Solid



The smallest repeating unit in a crystal is
called unit cell.
A crystal structure is composed of
periodically repeating units in three
dimensions
A crystal's structure play a role in
determining many of its properties
Example of Unit Cell
CsCl
unit cell
44
Crystal system
 There are seven unique crystal systems. The
simplest and most symmetric is the cubic
system that has the symmetry of a cube.
 The other six systems, are hexagonal,
tetragonal, rhombohedral (also known as
trigonal), orthorhombic, monoclinic and
triclinic
45
7 crystal systems
cubic
orthorhombic
tetragonal
rhombohedral
hexagonal
monoclinic
triclinic
7 systems = 14 Bravais lattices
46
47
Three types of cubic unit cells
Three types of cubic unit cells
48
Cubic Unit Cells

1. Primitive cubic, atoms are at the corners of a simple
cube. Each atom is shared by 8 unit cells.
 2. Body-centered cubic (bcc), atoms are at the corners of a
cube plus one at the center of the body of the cube. Corner
atoms are shared by 8 unit cells and the central one is
completely enclosed in one unit cell.
 3. In face-centered cubic (fcc), atoms are at the corners of a
cube and another one at the center of each face of the
cube. Corner atoms are shared by 8 unit cells, face atoms
are shared by 2 unit cells.
Empirical formula
We can determine the
empirical formula of an
ionic solid by
determining how many
ions of each element
fall within the unit cell.
49
Empirical formula
 What are the empirical formulas for these
compounds?



(a) Green: chlorine; Gray: cesium
(b) Yellow: sulfur; Gray: zinc
(c) Gray: calcium; Blue: fluorine
(a)
(b)
CsCl
ZnS
(c)
CaF2
X-Rays
 X-rays are electromagnetic radiations of
wavelength about 1 Å- 2 Å (about the same
size as an atom).
 X-ray crystallography is used to determine
solid structures
50
X-Rays
 Each crystalline solid has its unique
characteristic X-ray pattern which may be
used as a "fingerprint" for its identification.
 We can determine the size and the shape of
the unit cell for any compound most easily
using the diffraction of X-rays.
Bonding in Solids
Four types of solids:
1. Molecular Solids
2. Covalent Network Solids
3. Ionic Solids
4. Metallic Solids
51
Molecular Solids
 Molecular (formed from molecules) - usually soft





with low melting points and poor conductivity.
Intermolecular forces: dipole-dipole, London
dispersion and H-bonds.
Weak intermolecular forces give rise to low
melting points:
Naphthalene C10H8 Tf= 80°C
Benzoic Acid C6H5CO2H Tf= 122°C
Glucose C6H12O6 Tf= 155°C.
Covalent-Network Solid

Atoms or molecules linked together by covalent
bonds.
 In general very hard with high melting points
and poor to good conductivity.
Examples:
Carbon(diamond or graphite); quartz (SiO2);
silicon carbide (SiC); boron nitride (BN).
52
Covalent-Network Solids
 Diamond is an example of a covalent-network solid,
in which atoms are covalently bonded to each other.

They tend to be hard and have high melting points.
Covalent-Network Solids
 Graphite is an example of a molecular solid, in which
atoms are held together with van der Waals forces.

They tend to be softer and have lower melting points.
53
Why Phosphorus melts at lower
temperature than Silicon?
Phosphorus melts at = 317 K (Molecular
solid)
Silicon melts at 1683 K (Covalent Network
Solid)
Ionic Solids
 (formed from ions) – brittle to hard, high
melting points and poor conductivity .
 Ions held together by electrostatic forces
of attraction:
 F = k Q1Q2
d2
54
simple classifications for ionic
lattice types:
 NaCl Structure
 Face-centered cubic lattice.
 Cation to anion ratio is 1:1.
Examples: LiF, KCl, AgCl and CaO.
Example
The geometric arrangement of ions in crystals
of LiF is the same as that in NaCl. The unit
cell of LiF is 4.02 Å on an edge. Calculate the
density of LiF.
The volume of a cube of length a on an edge is a3, so the volume of the unit cell is (4.02 Å)3. We can
now calculate the density, converting to the common units of g/cm3.
55
CsCl Structure
Cs+ ions occupy the center and Cl- ions
occupy the corners of the cubic unit
cell. CsCl has a different structure than
NaCl (Cs+ is larger than Na+). Cation to
anion ratio is 1:1.
Zinc Blende Structure
 Typical example ZnS.
 S2- ions adopt a fcc arrangement.
 around the Zn2+ ions that occupy the
center of a tetrahedron .
 Example: CuCl.
56
Fluorite Structure
 Fluorite Structure CaF2
 Ca2+ ions in a fcc arrangement.
 F- occupies the center of a Octahedron
 Examples: BaCl2, PbF2
Structure of CsCl, Zns, and CaF2
57
Metallic Solids
 Metals are not covalently
bonded, but the attractions
between atoms are too strong
to be van der Waals forces.
 In metals valence electrons
are delocalized throughout
the solid.
the metal nuclei float in a sea
of electrons.
58