Chapter 8: Effects of Intermolecular Forces
8.1 Effects of Intermolecular Forces
8.2 Types of Intermolecular Forces
8.3 Liquids
8.4 Forces in Solids
8.5 Order in Solids
8.6 Phase Changes
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8.1 Effects of Intermolecular Forces
Learning objective:
Understand the effects of intermolecular forces on
condensation, vapourization and melting and boiling
points
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
8.1 Effects of Intermolecular Forces
 Intermolecular forces: forces of attraction between
molecules which result in liquids and solids.
 At STP, only 11 elements are gases.
 So, intermolecular forces are important in all elements,
except for those which are gases.
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Phases of Elements at Room Temperature
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The Balance of Kinetic Energy and
Intermolecular Potential Energy
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Melting and Boiling Points
 Both are indicators of the strengths of intermolecular
forces:
Normal melting point (Tm): the temperature at which a solid
and liquid coexist at equilibrium under a pressure of 1 atm.
 Normal boiling point (Tb): the temperature at which a liquid
and vapour coexist at equilibrium under a pressure of 1 atm.

 Vapourization: l → g
 Condensation: g → l
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8.2 Types of Intermolecular Forces
Learning objectives:
Predict the relative magnitudes of intermolecular
forces and their effects on physical properties of
substances
Explain the origins of each type of intermolecular
force
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8.2 Types of Intermolecular Forces
1.
Dispersion Forces – attraction between the negatively charged
electron cloud of one molecule and the positively charged
nuclei of a neighbor molecule.
2.
Dipolar forces – attraction between negatively charged end of
one molecule with the positively charged end of another
molecule.
3.
Induced Dipoles – attraction between an ion or a permanent
dipole and a molecule which has had a dipole induced in it by
the ion or permanent dipole.
4.
Hydrogen bonding forces –attraction between lone pair
electrons on an O, N or F atom with a hydrogen atom.
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1. Dispersion Forces
 Dispersion forces are found in all molecular substances
 Such forces are electrostatic in nature and arise from
attractions involving induced dipoles.
 Dispersion forces help explain why such things as
nonpolar compounds dissolve in water or ethanol.
 The magnitude of dispersion forces depends on how
easy it is to polarize the electron cloud of a molecule.
 A larger molecule has a larger polarizability.
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Polarization
• The process of inducing a dipole is called polarization.
• The higher the molar mass, the higher the polarizability of the
molecule.
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Larger Molecules are More Polarizable
The larger the
molecule, the larger
the electron cloud.
The larger the
electron cloud, the
more polarizable the
molecule.
The more polarizable,
the stronger the
dispersion forces.
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Higher Polarization Causes Higher
Melting and Boiling Points
 The larger the atoms or molecules, the stronger
the dispersion forces, the higher the boiling point
because, larger molecules have higher
polarizability and therefore they have stronger
dispersion forces.
 The stronger forces require more energy to break
resulting in a higher boiling point.
 Similar arguments for melting points.
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Boiling Point vs Molecular Size
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Example 8 - 1
Neon and xenon are gases at room temperature, but
both become liquids if the temperature is low enough.
Draw a molecular picture showing the relative sizes and
polarizabilities of atoms of neon and xenon, and use
the picture to determine which substance has the
lower boiling point.
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2. Dipolar Forces
 Occur when one polar molecule encounters
another polar molecule.
 The positive ends will be attracted to the negative
ends.
 Dipolar forces are characteristically stronger than
dispersion forces.
 Dipolar forces increase with an increase in the
polarity of the molecule.
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2-methylpropane
Acetone is a polar
molecule, 2methylpropane is not.
Therefore, acetone will
experience dipolar forces,
but 2-methylpropane will
not.
Which of the two has a
higher boiling point?
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acetone
Example 8 - 2
The line structures of butane, methyl ethyl ether and
acetone follow. Explain the trend in boiling points:
butane (0°C), methyl ethyl ether (8°C) and acetone
(56°C).
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3. Hydrogen Bonding Forces
1.
One molecule has a hydrogen atom attached by a
covalent bond to an atom of oxygen, nitrogen, or
fluorine.
2.
The other molecule has an oxygen, nitrogen, or
fluorine atom.
Remember: they are not really bonds…
just attractive forces!
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Hydrogen Bonding
Hydrogen bonding can occur with any hydrogen that is bonded
to either oxygen, nitrogen or fluorine.
(N-H, O-H and H-F bonds are very polar)
+
H
Hydrogen bond
+
O H
-
O H
H +
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+
Hydrogen Bonding Examples
Intramolecular
H-bond
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Example 8 - 3
In which of following systems will hydrogen bonding play
an important role: CH3F, (CH3)2CO (acetone), CH3OH,
and NH3 dissolved in (CH3)2CO?
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Binary Hydrogen Compounds
Exceptions due to polarity!
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8.3 Liquids
Learning objective:
Explain trends in surface tension, capillary action,
viscosity and vapour pressure in terms of
intermolecular forces
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8.3 Liquids
Liquid Properties
1.
2.
3.
4.
Surface Tension
Capillary Action
Viscosity
Vapour Pressure
All are functions of intermolecular forces!
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1. Surface Tension
 The resistance of a liquid to an increase
in its surface area (thus the units J/m2).
 The surface molecules of a liquid have
a net inward force of attraction,
forming a “skin”.
 The toughness of the skin is called
surface tension.
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2. Capillary Action
Capillary Action


The upward movement of water inside a capillary
against the force of gravity.
Due to large forces between the glass (polar) and the
water (also polar) than among water molecules.
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3. Viscosity
Viscosity - resistance to flow



Dependent on intermolecular forces (e.g. hydrocarbons vs.
water)
Dependent on length of the carbon chain
Dependent on temperature (e.g. viscosity breakdown in
engine oils)
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Vapour Pressure
Vapour pressure (pvap): the pressure at which dynamic
equilibrium is achieved in a closed container.
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8.4 Forces in solids
Learning objective:
Explain the properties of solids in terms of the dominant
intermolecular forces present
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8.4 Forces in solids
 In liquids and gases, molecules are free to move
continually and randomly.
 In solids, molecules, atoms, and ions cannot move
freely, but they can vibrate and occasionally rotate.
 There are four major solid types: (1) molecular solids,
(2) network solids, (3) metallic solids, and (4) ionic
solids.
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Forces in Solids
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1. Molecular Solids
 Aggregates of molecules
bound together by
intermolecular forces.
 Molecules of the
molecular solids retain
their individual
properties.
 Gases under normal
conditions, but form
solids at low
temperatures.
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Crystals of benzoic acid
contain pairs of molecules
held together head to
head by hydrogen bonds.
These pairs then stack in
planes which are held
together by dispersion forces.
The effect of extensive Hbonding is revealed by the high
melting point of glucose (155°C).
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2. Network Solids
 Unlike molecular solids, network
solids have high melting points.
 They are held together by
covalent bonds which are much
harder to break than
intermolecular forces.
 Bonding patterns determine the
properties of the solid.
 Usually durable compounds
(Rock of Gibraltar, rubies,
sapphires, etc…)
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Network Solids of Carbon
 Diamond – three dimensional array of s bonds with
each tetrahedral (sp3) carbon linked to others through
covalent bonds, this array makes diamonds very strong
and abrasive.
 Graphite – has trigonal planar (sp2) carbons in a 2D
array of s bonds, each 2D layer is attracted to its
neighbours by dispersion forces.
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Network Solids of Carbon
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Example 8 - 4
Whereas SiO2 melts at
1710°C, other nonmetal
oxides melt at much
lower temperatures.
For example, P4O6 melts
at 25°C. Referring to
the accompanying
bonding pictures,
describe the forces that
hold these solids
together.
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3. Metallic Solids
 Bonding of electrons is delocalized.
 Therefore, the strength of the bonding is variable.
 Metals are ductile (able to be drawn into wires) and
malleable (able to be hammered into sheets)
 Metals have a range of melting points
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When a metal changes shape, its atoms shift position. However,
because the valence electrons are fully delocalized, the energy of
these electrons is unaffected.
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4. Ionic Solids
 Contain cations and anions strongly
attracted to each other through interionic
forces – in between ions.
 Many ionic solids contain metal cations
and polyatomic ions
 Superconductors: ionic solids composed
of oxides of rare earth metals: YBa2Cu3O7-x

They carry immense electrical current without
losses due to resistance.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
8.5 Order in Solids
Learning objective:
Understand amorphous and crystalline solids at the
molecular level
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8.5 Order in Solids
 Amorphous – No ordered structure to the
particles of the solid
No well defined faces, angles or shapes
 Often are mixtures of molecules which do not stack
together well, or large flexible molecules
 Examples include glass and rubber

 Crystalline - The atoms, molecules or ions pack
together in an ordered arrangement
Such solids typically have flat surfaces, with unique
angles between faces and a unique 3-dimensional shape
 Examples of crystalline solids include diamonds, and
quartz crystals

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The Crystal Lattice and the Unit Cell
Unit cell – the smallest
unit from which the
entire pattern can
be assembled
One unit cell
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Basic Definitions
 Lattice points – the corners of the unit cell
 Crystal lattice – a group of identical unit cells
 Cubic Unit Cell – a unit cell which has edges of equal
length (l = w = h) and angles of 90°
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Three Basic Cubic Crystals
1.
2.
3.
Simple cubic (sc) – layers of atoms stacked one
directly above another, so that all atoms lie along
straight lines at right angles.
Body-centered cubic (bcc) – simple cube with one
entire atom in the center of the cube (in the body).
Face-centered cubic (fcc) – simple cube with atoms in
the center of each face of the cube.
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Simple Cubic
# 1/8 of an atom in each
corner
# atoms = 8 atoms (1/8) =
1 atom
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Example 8 - 5
Calculate the mass of a single unit cell of polonium (Po)
metal, which crystallizes in a simple cubic structure.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Body Centred Cubic
1/8 of an atom at each
corner,
and 1 whole atom in the
centre
# atoms = 8 atoms (1/8) + 1 =
2 atoms
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Face Centred Cubic
1/8 of an atom at each corner
and 1/2 of an atom on each face
# atoms = 8 atoms (1/8) +
6 atoms (1/2) = 4 atoms
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 8 - 6
Calculate the density of silver metal (in g/cm3), if the
edge length of its face-centred cubic unit cell is 407 pm.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 8 - 7
Calculate the radius of a palladium (Pd) atom, given that
the density of Pd metal is 12.02 g/cm3 and that it has a
face-centred cubic unit cell.
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Close-Packed Crystals
 Close-packing maximizes intermolecular attractions.
 All empty space around the atoms or molecules is
minimized.
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Hexagonal Close-Packed (hcp)
and Cubic Close-Packed (ccp)
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Hexagonal Close-Packed (hcp)
and Cubic Close-Packed (ccp)
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Ionic Solids
 Ions of opposite charges alternate with one another to
maximize attractions among ions.
 Cations and anions are of different size (cations are
usually smaller).
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Crystalline Defects
 Defects can alter the
properties of the solid
material
 Examine what happens
when carbon is added to
iron to make steel.
 Iron is relatively soft, but
adding carbon atoms
reduces its ability to
become deformed by filling
in empty holes in the lattice.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
8.6 Phase Changes
Learning objective:
Explain enthalpies of phase changes in terms of
intermolecular forces
Interpret a pressure-temperature phase diagram of a
pure substance
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8.6 Phase Changes
 There are three states
 Solid
 Liquid
 Gas
 A transformation from one state to another is called a
phase change
 Each phase change is associated with a change in
energy of the system
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Temperature and Phase Changes
As a phase change occurs, temperature remains constant
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Heats of Phase Changes
 Molar heat of vaporization, DHvap: the heat needed to
vapourize one mole of a substance at its normal boiling
point.
 Molar heat of fusion, DHfus: the heat needed to melt
one mole of a substance at its normal melting point.
 Molar heat of sublimation, DHsub: the heat needed to
sublime one mole of a substance from the solid phase
to the gas phase (skips the liquid phase).
DHsub  DHvap + DHfus
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Example 8 - 8
A swimmer emerging from a pool is covered in a film
containing 75 g of water. How much heat must be
supplied to evaporate the water?
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Phase Diagrams
 Illustrate the relationship between phases of matter
and the pressure and temperature
 The lines identify the conditions under which two
phases exist in equilibrium
 Triple point – point at which all three phases coexist
 Critical point – point at which one cannot distinguish
between a gas and a liquid.
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A General Phase Diagram
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Example 8 - 9
Ammonia is a gas at normal temperature and pressure.
Its normal boiling point is 239.8 K, and it freezes at
195.5 K. The triple point for NH3 is p = 0.0610 bar and T
= 195.4 K. Use this information to construct an
approximate phase diagram for NH3.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 8 - 10
A chemist wants to perform a synthesis in a vessel at p =
0.50 atm using liquid NH3 as the solvent. What
temperature range would be suitable? When the
synthesis is complete, the chemist wants to boil off the
solvent without raising the T above 220 K. Is this
possible?
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Variations on Phase
Diagrams
 Some substances
have more than one
solid phase, while
others display liquid
crystal phases.
 The resulting phase
diagram must
account for all the
possible phases.
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Silica
Chapter 8 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 8 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 8 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 8 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 8 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 8 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.