Chemistry, The Central Science, 11th edition
Theodore L. Brown, H. Eugene LeMay, Jr.,
and Bruce E. Bursten
Chapter 11
Intermolecular Forces,
Liquids, and Solids
John D. Bookstaver
St. Charles Community College
Cottleville, MO
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.
Intermolecular
Forces
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van der Waals Forces
• Dipole-dipole interactions
• Hydrogen bonding
• London dispersion forces
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.
Intermolecular
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.
Intermolecular
Forces
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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.
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 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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Sample Exercise 11.1 Comparing Intermolecular Forces
The dipole moments of acetonitrile, CH3CN, and methyl iodide, CH3I, are 3.9 D and 1.62 D, respectively.
(a) Which of these substances has greater dipole–dipole attractions among its molecules? (b) Which of these
substances has 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 greater overall attractive forces?
Solution
(a) Dipole–dipole attractions increase in magnitude as the dipole moment of the molecule increases. Thus,
CH3CN molecules attract each other by stronger dipole–dipole forces than CH3I molecules do. (b) When
molecules differ in their molecular weights, the more massive molecule generally has the stronger
dispersion attractions. In this case CH3I (142.0 g/mol) is much more massive than CH3CN (41.0 g/mol), so
the dispersion forces will be stronger for CH3I. (c) Because CH3CN has the higher boiling point, we can
conclude that more energy is required to overcome attractive interactions between CH 3CN molecules. Thus,
the total intermolecular attractions are stronger for CH3CN, suggesting that the energies resulting from
dipole–dipole forces are decisive when comparing these two substances. Nevertheless, the attractive
interactions due to dispersion forces play an important role in determining the properties of CH 3I.
Practice Exercise
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?
Answer: (a) Br2 (largest molecular weight), (b) HCl (largest polarity)
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.
• ..\Media
Gallery\Animations\Hydro
genBonding\HydrogenBon
ding.html
Also, when hydrogen is bonded to one of those
very electronegative elements, the hydrogen
nucleus is exposed.
Intermolecular
Forces
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Sample Exercise 11.2 Identifying Substances that Can Form Hydrogen Bonds
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)?
Solution
Analyze: We are given the chemical formulas of four substances and asked to predict whether they can
participate in hydrogen bonding. All of these compounds contain H, but hydrogen bonding usually occurs
only when the hydrogen is covalently bonded to N, O, or F.
Plan: We can analyze each formula to see if it contains N, O, or F directly bonded to H. There also needs to
be a nonbonding pair of electrons on an electronegative atom (usually N, O, or F) in a nearby molecule,
which can be revealed by drawing the Lewis structure for the molecule.
Solve: The criteria listed above eliminate CH4 and H2S, which do not contain H bonded to N, O, or F. They
also eliminate CH3F, whose Lewis structure shows a central C atom surrounded by three H atoms and an F
atom. (Carbon always forms four bonds, whereas hydrogen and fluorine form one each.) Because the
molecule contains C—F a bond and not a H—F bond, it does not form hydrogen bonds. In H2NNH2,
however, we find bonds. If we draw the Lewis structure for the molecule, we see that there is a nonbonding
pair of electrons on each N atom. Therefore, hydrogen bonds can exist between the molecules as depicted
below.
Intermolecular
Forces
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Sample Exercise 11.2 Identifying Substances that Can Form Hydrogen Bonds
Solution (continued)
Check: While we can generally identify substances that participate in hydrogen bonding based on their
containing N, O, or F covalently bonded to H, drawing the Lewis structure for the interaction, as shown
above, provides a way to check the prediction.
Practice Exercise
In which of the following substances is significant hydrogen bonding possible: methylene chloride (CH 2Cl2),
phosphine (PH3), hydrogen peroxide (HOOH), or acetone (CH3COCH3)?
Answer: HOOH
Intermolecular
Forces
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Summarizing Intermolecular Forces
Intermolecular
Forces
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Sample Exercise 11.3 Predicting the Types and Relative Strengths of
Intermolecular Attractions
List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling points.
Solution
Analyze: We need to relate the properties of the listed substances to boiling point.
Plan: The boiling point depends in part on the attractive forces in the liquid. We need to order these
according to the relative strengths of the different kinds of intermolecular attractions.
Solve: The attractive forces are stronger for ionic substances than for molecular ones, so BaCl 2 should have
the highest boiling point. The intermolecular forces of the remaining substances depend on molecular
weight, polarity, and hydrogen bonding. The molecular weights are H 2 (2), CO (28), HF (20), and Ne (20).
The boiling point of H2 should be the lowest because it is nonpolar and has the lowest molecular weight.
The molecular weights of CO, HF, and Ne are roughly the same. Because HF can hydrogen bond, however,
it should have the highest boiling point of the three. Next is CO, which is slightly polar and has the highest
molecular weight. Finally, Ne, which is nonpolar, should have the lowest boiling point of these three. The
predicted order of boiling points is therefore
H2 < Ne < CO < HF < BaCl2
Check: The actual normal boiling points are H2 (20 K), Ne (27 K), CO (83 K), HF (293 K), and BaCl2
(1813 K)—in agreement with our predictions.
Practice Exercise
(a) Identify the intermolecular attractions present in the following substances, and (b) select the substance
with the highest boiling point: CH3CH3, CH3OH, and CH3CH2OH.
Answer: (a) CH3CH3 has only dispersion forces, whereas the other two substances have both dispersion forces
and hydrogen bonds; (b) 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.
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 Calculating ΔH for Temperature and Phase Changes
Calculate the enthalpy change upon converting 1.00 mol of ice at –25 °C to water vapor (steam) at 125 °C
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, ΔHfus = 6.01 kJ/mol and ΔHvap = 40.67 kJ/mol.
Solution
Analyze: Our goal is to calculate the total heat required to convert 1 mol of ice at -25 °C to steam at 125
°C.
Plan: We can calculate the enthalpy change for each segment and then sum them to get the total enthalpy
change (Hess’s law, Section 5.6).
Solve: For segment AB in Figure 11.19, we are adding
enough heat to ice to increase its temperature
by 25 °C. A temperature change of is 25 °C
the same as a temperature change of 25 K, so
we can use the specific heat of ice to calculate
the enthalpy change during this process:
For segment BC in Figure 11.19, in which we convert
ice to water at 0 °C, we can use the molar enthalpy of
fusion directly:
The enthalpy changes for segments CD, DE, and EF
can be calculated in similar fashion:
Intermolecular
Forces
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Sample Exercise 11.4 Calculating ΔH for Temperature and Phase Changes
Solution (continued)
The total enthalpy change is the sum of the changes of the individual steps:
Check: The components of the total energy change are reasonable in comparison with the lengths of the
horizontal segments of the lines in Figure 11.19. Notice that the largest component is the heat of vaporization.
Practice Exercise
What is the enthalpy change during the process in which 100.0 g of water at 50.0 °C is cooled to ice at
–30.0 °C ? (Use the specific heats and enthalpies for phase changes given in Sample Exercise 11.4.)
Answer: –20.9 kJ – 33.4 kJ – 6.09 kJ = –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
it’s 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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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.
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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Sample Exercise 11.6 Interpreting a Phase diagram
Referring to Figure 11.28, describe any changes in the phases present when H 2O is (a) kept at 0 °C 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).
Solution
Analyze: We are asked to use the phase diagram provided to deduce what phase changes might occur when
specific pressure and temperature changes are brought about.
Plan: Trace the path indicated on the phase diagram, and note what phases and phase changes occur.
Solve:
(a) At point 1, H2O exists totally as a vapor. At point 2 a solid–vapor equilibrium exists. Above that
pressure, at point 3, all the H2O is converted to a solid. At point 4 some of the solid melts and equilibrium
between solid and liquid is achieved. At still higher pressures all the H 2O melts, so only the liquid phase is
present at point 5.
(b) At point 6 the H2O exists entirely as a solid. When the temperature reaches point 4, the solid begins to
melt and equilibrium exists between the solid and liquid phases. At an even higher temperature, point 7, the
solid has been converted entirely to a liquid. A liquid–vapor equilibrium exists at point 8. Upon further
heating to point 9, the H2O is converted entirely to the vapor phase.
Check: The indicated phases and phase changes are consistent with our knowledge of the properties of
water.
Practice Exercise
(Using Figure 11.27(b), describe what happens when the following changes are made in a CO 2 sample: (a)
Pressure increases from 1 atm to 60 atm at a constant temperature of –60 °C. (b) Temperature increases from
–60 °C to –20 °C at a constant pressure of 60 atm.
Answer: (a) CO2(g) → CO2(s); (b) CO2(s) → CO2(l)
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 Putting Concepts Together
The substance CS2 has a melting point of –110.8 °C and a boiling point of 46.3 °C. Its density at 20 °C 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.
(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 (Table 11.5), and discuss the possible origins of the
differences. (f) Would you expect the density of CS2 at 40 °C to be greater or less than at 20 °C? What
accounts for the difference?
Solution
(a) The compound is named carbon disulfide, in analogy with the naming of other binary molecular
compounds such as carbon dioxide . (Section 2.8)
(b) Only London dispersion forces affect CS2; it does not have a dipole moment, based upon its molecular
shape, and obviously cannot undergo hydrogen bonding.
(c) Because CS2(s) consists of individual CS2 molecules, it will be a molecular solid.
(d) The most likely products of the combustion will be CO2 and SO2. (Sections 3.2 and 7.8) Under some
conditions SO3 might be formed, but this would be the less likely outcome. Thus, we have the following
equation for combustion:
CS2(l) + 3 O2(g) → CO2(g) + 2 SO2(g)
(e) The critical temperature and pressure of CS2 (552 K and 78 atm) are both higher than those given for
CO2 in Table 11.5 (304 K and 73 atm). The difference in critical temperatures is especially notable. The
higher values for CS2 arise from the greater London dispersion attractions between the CS 2 molecules
compared with CO2. These greater attractions are due to the larger size of the sulfur compared to oxygen
and therefore its greater polarizability.
Intermolecular
Forces
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Sample Integrative Exercise Putting Concepts Together
Solution (continued)
(f) The density would be lower at the higher temperature. Density decreases with increasing temperature
because the molecules possess higher kinetic energies. Their more energetic movements result in larger
average distances between molecules, which translate into lower densities.
Intermolecular
Forces
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