Chemistry, The Central Science, 11th edition
Theodore L. Brown, H. Eugene LeMay, Jr.,
and Bruce E. Bursten
Chapter 11
Intermolecular Forces,
Liquids, and Solids
Intermolecular
Forces
© 2009, Prentice-Hall, Inc.
Overview of Chapter 11 –
Intermolecular Forces
Review/On your own:
• compare liquids and solids with gases in the context of the KMT
•
describe the three intermolecular forces and apply this knowledge to predictions
about substances’ properties, such as melting and boiling points, solubilities,
surface tension, etc.
•
explain what vapor pressure is, and how it differs between contained and
uncontained containers
•
interpret heating curves, vapor pressure graphs, and phase diagrams
•
calculate enthalpy changes associated with phase changes
•
relate vapor pressure to volatility of a liquid and the intermolecular forces
•
interpret vapor pressure graphs
Intermolecular
Forces
© 2009, Prentice-Hall, Inc.
Overview of Chapter 11 –
Intermolecular Forces
More depth:
• interpret phase diagrams
• describe the three intermolecular forces
and apply this knowledge to predictions
about substances’ properties, such as
melting and boiling points, solubilities,
surface tension, etc.
• NOTHING ON SOLIDS (11.7 and 11.8)
Intermolecular
Forces
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States of Matter
Shape
Volume
Density
Compressible?
Mass?
Solid
Liquid
Gas
Intermolecular
Forces
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States of Matter
The fundamental difference between states of
matter is the distance between particles.
Intermolecular
Forces
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States of Matter
Because in the solid and liquid states
particles are closer together, we refer to them
as condensed phases.
Intermolecular
Forces
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The States of Matter
• The state a substance is
in at a particular
temperature and
pressure depends on
two antagonistic entities:
– the kinetic energy of the
particles;
– the strength of the
attractions between the
particles.
Intermolecular
Forces
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Intermolecular Forces
The attractions between molecules are not
nearly as strong as the intramolecular
attractions that hold compounds together.
Intermolecular
Forces
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Intermolecular Forces
They are, however, strong enough to control
physical properties such as boiling and
melting points, vapor pressures, and
viscosities.
Intermolecular
Forces
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Intermolecular Forces
These intermolecular forces as a group are referred to as van
der Waals forces.
Note: Sometimes scientists use the term van der Waals force to
refer to dispersion forces only. Be sure to describe the
interaction as thoroughly as possible to clearly distinguish
amongst them.
Intermolecular
Forces
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van der Waals Forces
• London dispersion forces (weakest,
present in all molecules)
• Dipole-dipole interactions
• Hydrogen bonding
Intermolecular
Forces
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Ion-Dipole Interactions
• Ion-dipole interactions (a fourth type of force),
are important in solutions of ions.
• The strength of these forces are what make it
possible for ionic substances to dissolve in
polar solvents.
Intermolecular
Forces
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Dipole-Dipole Interactions
• 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.
Intermolecular
Forces
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Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling
point.
Note: There is a stronger relationship between the
strength of IMFs and boiling points, more so than with
Intermolecular
melting points.
Forces
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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
Intermolecular
same side of the atom.
Forces
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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.
Intermolecular
Forces
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London Dispersion Forces
Another helium nearby, then, would have a
dipole induced in it, as the electrons on the
left side of helium atom 2 repel the electrons
in the cloud on helium atom 1.
Intermolecular
Forces
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London Dispersion Forces
London dispersion forces, or dispersion
forces, are attractions between an
instantaneous dipole and an induced dipole.
Intermolecular
Forces
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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.
Intermolecular
Forces
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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.
• Note the differences in boiling
points
Intermolecular
Forces
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Factors Affecting London Forces
• The strength of dispersion forces tends to
increase with increased molecular weight.
• BECAUSE larger atoms have larger electron
clouds which are easier to polarize.
Intermolecular
Forces
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Sample Exercise 11.1 (p. 443)
The dipole moments of acetonitrile, CH3CN, and methyl iodide,
CH3I, are 3.9 D and 1.62 D, respectively.
a)
Which of these substances will have the greater
dipole-dipole attractions among its molecules?
b)
Which of these substances will have the greater
London dispersion attractions?
c)
The boiling points of CH3CN and CH3I are 354.8 K and
315.6 K, respectively. Which substance has the greatest
overall attractive forces?
Intermolecular
Forces
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Practice Exercise 11.1
Of Br2, Ne, HCl, HBr, and N2, which is likely to
have
a) the largest intermolecular dispersion
forces?
b) the largest dipole-dipole attractive
forces?
Intermolecular
Forces
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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.
• If one molecule is much larger than
another, dispersion forces will likely
determine its physical properties.
Intermolecular
Forces
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How Do We Explain This?
• The nonpolar series
(SnH4 to CH4) follow
the expected trend.
• The polar series
follows the trend
from H2Te through
H2S, but water is
quite an anomaly.
Intermolecular
Forces
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Hydrogen Bonding
• The dipole-dipole interactions
experienced when H is bonded to
N, O, or F are unusually strong.
• We call these interactions
hydrogen bonds.
Intermolecular
Forces
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Hydrogen Bonding
• Hydrogen bonding
arises in part 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, small and can approach an electronegative atom
very closely, making it much easier to interact with it. Intermolecular
Forces
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Sample Exercise 11.2 (p. 444)
In which of the following substances is
hydrogen bonding likely to play an
important role in determining physical
properties: methane (CH4), hydrazine
(H2NNH2), methyl fluoride (CH3F), or
hydrogen sulfide (H2S)?
Intermolecular
Forces
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Practice Exercise 11.2
In which of the following substances is
significant hydrogen bonding possible:
methylene chloride (CH2Cl2), phosphine
(PH3), hydrogen peroxide (HOOH), or
acetone (CH3COCH3)?
Intermolecular
Forces
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Summarizing Intermolecular Forces
Intermolecular
Forces
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Sample Exercise 11.3 (p. 447)
List the substances BaCl2, H2, CO, HF,
and Ne in order of increasing boiling
points.
Intermolecular
Forces
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Practice Exercise 11.3
a) Identify the intermolecular forces
present in the following substances,
and
b) select the substance with the highest
boiling point:
CH3CH3, CH3OH, CH3CH2OH.
Intermolecular
Forces
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Intermolecular Forces Affect
Many Physical Properties
The strength of the
attractions between
particles can greatly
affect the properties
of a substance or
solution.
Intermolecular
Forces
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Viscosity
• Resistance of a liquid
to flow is called
viscosity.
• It is related to the ease
with which molecules
can move past each
other.
• Viscosity increases
with stronger
intermolecular forces
and decreases with
higher temperature.
Intermolecular
Forces
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Surface Tension
Surface tension
results from the net
inward force
experienced by the
molecules on the
surface of a liquid.
Intermolecular
Forces
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Phase Changes
Intermolecular
Forces
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Energy Changes Associated
with Changes of State
The heat of fusion is the energy required to
change a solid at its melting point to a liquid.
Intermolecular
Forces
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Energy Changes Associated
with Changes of State
The heat of vaporization is defined as the energy required to
change a liquid at its boiling point to a gas.
Note how much more energy is needed to vaporize a substance
than to melt it.
Intermolecular
Forces
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Energy Changes Associated
with Changes of State
• The heat added to the
system at the melting
and boiling points goes
into pulling the
molecules farther apart
from each other.
• The temperature of the
substance does not rise
during a phase change.
Intermolecular
Forces
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Sample Exercise 11.4 (p. 451)
Calculate the enthalpy change upon converting 1.00
mol of ice at -25oC to water vapor (steam) at 125oC
under a constant pressure of 1 atm. The specific
heats of ice, water, and steam are 2.03 J/g-K, 4.18
J/g-K, and 1.84 J/g-K, respectively. For H2O, DHfus =
6.01 kJ/mol, and DHvap = 40.67 kJ/mol.
Note: multi-step problem, tedious, but not hard
(56.0 kJ)
Intermolecular
Forces
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•
•
•
•
•
0.91 kJ
6.01 kJ  melting
7.52 kJ
40.7 kJ  vaporization
0.83 kJ
Intermolecular
Forces
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Practice Exercise 11.4
What is the enthalpy change during the
process in which 100.0 g of water at
50.0oC is cooled to ice at -30.0oC?
(Use the specific heats and enthalpies
for phase changes given in Sample
Exercise 11.4.)
(-60.4 kJ)
Intermolecular
Forces
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Vapor Pressure
• At any temperature some molecules in a
liquid have enough energy to escape.
• As the temperature rises, the fraction of
molecules that have enough energy to
escape increases.
Intermolecular
Forces
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Vapor Pressure
As more molecules
escape the liquid,
the pressure they
exert increases.
Intermolecular
Forces
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Vapor Pressure
The liquid and vapor
reach a state of
dynamic equilibrium:
liquid molecules
evaporate and vapor
molecules condense
at the same rate.
Intermolecular
Forces
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Vapor Pressure
• 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.
Intermolecular
Forces
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Sample Exercise 11.5 (p. 455)
Use the preceding figure to estimate the
boiling point of diethyl ether under an
external pressure of 0.80 atm.
Intermolecular
Forces
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Practice Exercise 11.5
At what external pressure will ethanol
have a boiling point of 60oC?
Intermolecular
Forces
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Phase Diagrams
Phase diagrams display the state of a
substance at various pressures and
temperatures and the places where equilibria
exist between phases.
Intermolecular
Forces
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Phase Diagrams
• The circled line is the liquid-vapor interface.
• It starts at the triple point (T), the point at
which all three states are in equilibrium.
Intermolecular
Forces
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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.
Intermolecular
Forces
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Phase Diagrams
Each point along this line is the boiling point
of the substance at that pressure.
Intermolecular
Forces
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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.
Intermolecular
Forces
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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.
Intermolecular
Forces
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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.
Intermolecular
Forces
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Phase Diagram of Water
• The slope of the solidliquid line is negative.
– This means that as the
pressure is increased at a
temperature just below the
melting point, water goes
from a solid to a liquid.
– Why?
Intermolecular
Forces
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• Read from passage from “Dirt”.
Intermolecular
Forces
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Phase Diagram of Carbon Dioxide
Carbon dioxide
cannot exist in the
liquid state at
pressures below
5.11 atm; CO2
sublimes at normal
pressures.
Intermolecular
Forces
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Phase Diagram of Carbon Dioxide
The low critical
temperature and
critical pressure for
CO2 make
supercritical CO2 a
good solvent for
extracting nonpolar
substances (like
caffeine)
Intermolecular
Forces
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• Complete the rest of the Chapter 11
packet on your own. Note that the
section on packing of crystalline
structures is not required.
(most of Sec. 11.7)
Intermolecular
Forces
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Sample Exercise 11.6 (p. 458)
Referring to the figure at
right, describe any
changes in the phases
present when H2O is
a) kept at 0oC while the
pressure is increased from
that at point 1 to that at
point 5 (vertical line);
b) kept at 1.00 atm while the
temperature is increased
from that at point 6 to that
at point 9 (horizontal line).
Intermolecular
Forces
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Practice Exercise 11.6
Using the phase diagram of CO2, describe what
happens when the following changes are made
in a CO2 sample initially at 1 atm and -60oC:
a) pressure increases at constant
temperature of -60oC from 1 atm to 60 atm,
b) temperature increases from -60oC to -20oC
at constant 60 atm pressure.
Intermolecular
Forces
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Solids
• We can think of
solids as falling into
two groups:
– crystalline, in which
particles are in highly
ordered arrangement.
Intermolecular
Forces
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Solids
• We can think of
solids as falling into
two groups:
– amorphous, in which
there is no particular
order in the
arrangement of
particles.
Intermolecular
Forces
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Attractions in Ionic Crystals
In ionic crystals, ions
pack themselves so as
to maximize the
attractions and
minimize repulsions
between the ions.
Intermolecular
Forces
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Crystalline Solids
Because of the
ordered in a crystal,
we can focus on the
repeating pattern of
arrangement called
the unit cell.
Intermolecular
Forces
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Crystalline Solids
There are several types of basic
arrangements in crystals, like the ones
depicted above.
Intermolecular
Forces
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Crystalline Solids
We can determine
the empirical
formula of an ionic
solid by determining
how many ions of
each element fall
within the unit cell.
Intermolecular
Forces
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Ionic Solids
• 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
Intermolecular
Forces
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Types of Bonding in
Crystalline Solids
Intermolecular
Forces
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Covalent-Network and
Molecular Solids
• Diamonds are an example of a covalentnetwork solid, in which atoms are covalently
bonded to each other.
– They tend to be hard and have high melting
points.
Intermolecular
Forces
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Covalent-Network and
Molecular 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.
Intermolecular
Forces
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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.
Intermolecular
Forces
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Sample Integrative Exercise
(p. 469)
The substance CS2 has a melting point of -110.8oC and a boiling
point of 46.3oC. Its density at 20oC is 1.26 g/cm3. It is highly
flammable.
a) What is the name of this compound?
b) List the intermolecular forces that CS2 molecules would have
with each other.
c) Predict what type of crystalline solid CS2(s) would form.
Intermolecular
Forces
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Sample Integrative Exercise
(p. 469)
d) Write a balanced equation for the combustion of this
compound in air. (You will have to decide on the
most likely oxidation products.)
e) The critical temperature and pressure for CS2 are
552 K and 78 atm, respectively. Compare these
values with those for CO2 and discuss the possible
origins of the differences.
f) Would you expect the density of CS2 at 40oC to be
greater or less than at 20oC? What accounts for the
difference?
Intermolecular
Forces
© 2009, Prentice-Hall, Inc.