Transcript Introductory Chemistry, 2nd Edition Nivaldo Tro

Introductory Chemistry, 3rd Edition
Nivaldo Tro
Chapter 12
Liquids, Solids, and
Intermolecular Forces
Roy Kennedy
Massachusetts Bay Community College
Wellesley Hills, MA
2009, Prentice Hall
Interactions Between Molecules
• Many of the phenomena we observe are
related to interactions between molecules
that do not involve a chemical reaction.
Your taste and smell organs work because
molecules in the thing you are sensing interact
with the receptor molecule sites in your tongue
and nose.
• In this chapter, we examine the physical
interactions between molecules and the
factors that effect and influence them.
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The Physical States of Matter
• Matter can be classified as solid, liquid, or
gas based on what properties it exhibits.
State
Shape
Volume
Compress
Flow
Solid
Fixed
Fixed
No
No
Liquid
Indef.
Fixed
No
Yes
Gas
Indef.
Indef.
Yes
Yes
•Fixed = Keeps shape when placed in a container.
•Indefinite = Takes the shape of the container.
Tro's Introductory Chemistry, Chapter
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Structure Determines Properties
• The atoms or molecules have different
structures in solids, liquids, and gases,
leading to different properties.
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Properties of the States of Matter:
Gases
• Low densities compared to solids and liquids.
• Fluid.
The material exhibits a smooth, continuous flow
as it moves.
• Take the shape of their container(s).
• Expand to fill their container(s).
• Can be compressed into a smaller volume.
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Properties of the States of Matter:
Liquids
• High densities compared to gases.
• Fluid.
 The material exhibits a smooth, continuous flow as it
moves.
• Take the shape of their container(s).
• Keep their volume, do not expand to fill their
container(s).
• Cannot be compressed into a smaller volume.
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Properties of the States of Matter:
Solids
• High densities compared to gases.
• Nonfluid.
 They move as an entire “block” rather than a smooth,
continuous flow.
• Keep their own shape, do not take the shape of their
container(s).
• Keep their own volume, do not expand to fill their
container(s).
• Cannot be compressed into a smaller volume.
Tro's Introductory Chemistry, Chapter
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The Structure of
Solids, Liquids, and Gases
Tro's Introductory Chemistry, Chapter
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Gases
• In the gas state, the particles have complete
freedom from each other.
• The particles are constantly flying around,
bumping into each other and their
container(s).
• In the gas state, there is a lot of empty space
between the particles.
On average.
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Gases, Continued
• Because there is a lot of empty space, the particles
can be squeezed closer together. Therefore, gases
are compressible.
• Because the particles are not held in close contact
and are moving freely, gases expand to fill and take
the shape of their container(s), and will flow.
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Liquids
• The particles in a liquid are closely
packed, but they have some ability to
move around.
• The close packing results in liquids
being incompressible.
• But the ability of the particles to move
allows liquids to take the shape of their
container and to flow. However, they
don’t have enough freedom to escape
and expand to fill the container(s).
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Solids
• The particles in a solid are packed close together
and are fixed in position.
 Though they are vibrating.
• The close packing of the particles results in solids
being incompressible.
• The inability of the particles to move around
results in solids retaining their shape and volume
when placed in a new container, and prevents the
particles from flowing.
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Solids, Continued
• Some solids have their particles
arranged in an orderly geometric
pattern. We call these crystalline
solids.
Salt and diamonds.
• Other solids have particles that do
not show a regular geometric
pattern over a long range. We call
these amorphous solids.
Plastic and glass.
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Why Is Sugar a Solid, But
Water Is a Liquid?
• The state a material exists in depends on the
attraction between molecules and their ability to
overcome the attraction.
• The attractive forces between ions or molecules
depends on their structure.
 The attractions are electrostatic.
 They depend on shape, polarity, etc.
• The ability of the molecules to overcome the
attraction depends on the amount of kinetic energy
they possess.
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Properties and Attractive Forces
Shape
Volume
Relative
strength of
attractive
forces
Phase
Density
Gas
Low
Indefinite Indefinite
Weakest
Liquid
High
Indefinite
Definite
Moderate
Solid
High
Definite
Definite
Strongest
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Phase Changes:
Melting
• Generally, we convert a material in the solid state
into a liquid by heating it.
• Adding heat energy increases the amount of
kinetic energy of the molecules in the solid.
• Eventually, they acquire enough energy to
partially overcome the attractive forces holding
them in place.
• This allows the molecules enough extra freedom
to move around a little and rotate.
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Phase Changes:
Boiling
• Generally, we convert a material in the liquid state
into a gas by heating it.
• Adding heat energy increases the amount of
kinetic energy of the molecules in the liquid.
• Eventually, they acquire enough energy to
completely overcome the attractive forces holding
them together.
• This allows the molecules complete freedom to
move around and rotate.
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Properties of Liquids:
Surface Tension
• Liquids tend to minimize their
surface—a phenomenon we call
surface tension.
• This tendency causes liquids to
have a surface that resists
penetration.
• The stronger the attractive force
between the molecules, the larger
the surface tension.
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Surface Tension
• Molecules in the interior of a liquid
experience attractions to surrounding
molecules in all directions.
• However, molecules on the surface
experience an imbalance in attractions,
effectively pulling them in.
• To minimize this imbalance and maximize
attraction, liquids try to minimize the number
of molecules on the exposed surface by
minimizing their surface area.
• Stronger attractive forces between the
molecules = larger surface tension.
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Properties of Liquids:
Viscosity
• Some liquids flow more easily than
others.
• The resistance of a liquid’s flow is
called viscosity.
• The stronger the attractive forces
between the molecules, the more
viscous the liquid is.
• Also, the less round the molecule’s
shape, the larger the liquid’s viscosity.
 Some liquids are more viscous because
their molecules are long and get tangled in
each other, causing them to resist flowing.
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Escaping from the Surface
• The process of molecules of a
liquid breaking free from the
surface is called evaporation.
Also known as vaporization.
• Evaporation is a physical change
in which a substance is
converted from its liquid form to
its gaseous form.
The gaseous form is called a
vapor.
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Evaporation
• Over time, liquids evaporate—the molecules of
the liquid mix with and dissolve in the air.
• The evaporation happens at the surface.
• Molecules on the surface experience a smaller
net attractive force than molecules in the interior.
• All the surface molecules do not escape at once,
only the ones with sufficient kinetic energy to
overcome the attractions will escape.
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Escaping the Surface
• The average kinetic energy is directly proportional to
the Kelvin temperature.
• Not all molecules in the sample have the same amount
of kinetic energy.
• Those molecules on the surface that have enough
kinetic energy will escape.
 Raising the temperature increases the number of molecules
with sufficient energy to escape.
23
Escaping the Surface, Continued
• Since the higher energy molecules from the
liquid are leaving, the total kinetic energy of
the liquid decreases, and the liquid cools.
• The remaining molecules redistribute their
energies, generating more high energy
molecules.
• The result is that the liquid continues to
evaporate .
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Factors Effecting the
Rate of Evaporation
• Liquids that evaporate quickly are called volatile
liquids, while those that do not are called nonvolatile.
• Increasing the surface area increases the rate of
evaporation.
 More surface molecules.
• Increasing the temperature increases the rate of
evaporation.
 Raises the average kinetic energy, resulting in more
molecules that can escape.
• Weaker attractive forces between the molecules =
faster rate of evaporation.
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Reconnecting with the Surface
• When a liquid evaporates in a closed
container, the vapor molecules are trapped.
• The vapor molecules may eventually bump
into and stick to the surface of the container
or get recaptured by the liquid. This process
is called condensation.
A physical change in which a gaseous form is
converted to a liquid form.
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Dynamic Equilibrium
• Evaporation and condensation are opposite
processes.
• Eventually, the rate of evaporation and rate of
condensation in the container will be the same.
• Opposite processes that occur at the same rate
in the same system are said to be in dynamic
equilibrium.
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Evaporation and Condensation
When water is just
added to the flask
and it is capped, all
the water molecules
are in the liquid.
Shortly, the water
starts to evaporate.
Initially the rate
of evaporation is
much faster than rate
of condensation
Eventually, the condensation
and evaporation reach the
same speed. The air in the
flask is now saturated with
water vapor.
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Vapor Pressure
• Once equilibrium is reached, from that time
forward, the amount of vapor in the
container will remain the same.
As long as you don’t change the conditions.
• The partial pressure exerted by the vapor is
called the vapor pressure.
• The vapor pressure of a liquid depends on
the temperature and strength of
intermolecular attractions.
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Boiling
• In an open container, as you heat a
liquid the average kinetic energy of
the molecules increases, giving more
molecules enough energy to escape
the surface.
 So the rate of evaporation increases.
• Eventually, the temperature is high
enough for molecules in the interior
of the liquid to escape. A
phenomenon we call boiling.
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Boiling Point
• The temperature at which the vapor pressure of
the liquid is the same as the atmospheric pressure
is called the boiling point.
 The normal boiling point is the temperature required
for the vapor pressure of the liquid to be equal to 1
atm.
• The boiling point depends on what the
atmospheric pressure is.
 The temperature of boiling water on the top of a
mountain will be cooler than boiling water at sea level.
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Temperature and Boiling
• As you heat a liquid, its
temperature increases until it
reaches its boiling point.
• Once the liquid starts to boil, the
temperature remains the same
until it all turns to a gas.
• All the energy from the heat
source is being used to overcome
all of the attractive forces in the
liquid.
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Energetics of Evaporation
• As it loses its high energy molecules through
evaporation, the liquid cools.
• Then the liquid absorbs heat from its surroundings to
raise its temperature back to the same as the
surroundings.
• Processes in which heat flows into a system from the
surroundings are said to be endothermic.
• As heat flows out of the surroundings, it causes the
surroundings to cool.
 As alcohol evaporates off your skin, it causes your skin to
cool.
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Energetics of Condensation
• As it gains the high energy molecules through
condensation, the liquid warms.
• Then the liquid releases heat to its surroundings to
reduce its temperature back to the same as the
surroundings.
• Processes in which heat flows out of a system into the
surroundings are said to be exothermic.
• As heat flows into the surroundings, it causes the
surroundings to warm.
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Heat of Vaporization
• The amount of heat needed to vaporize one mole of a
liquid is called the heat of vaporization.
 DHvap
 It requires 40.7 kJ of heat to vaporize one mole of water at
 100 °C.
 Always endothermic.
 Number is +.
 DHvap depends on the initial temperature.
• Since condensation is the opposite process to evaporation,
the same amount of energy is transferred but in the
opposite direction.
 DHcondensation = −DHvaporization
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Heats of Vaporization of Liquids
at Their Boiling Points and at 25 °C
DHvap at
Normal
Chemical boiling
formula point, °C
H2O
100
Liquid
Water
Isopropyl
C3H7OH
alcohol
Acetone C3H6O
Diethyl
C4H10O
ether
DHvap at
boiling
point,
(kJ/mol)
+40.7
25 °C,
(kJ/mol)
+44.0
82.3
+39.9
+45.4
56.1
+29.1
+31.0
34.5
+26.5
+27.1
Tro's Introductory Chemistry, Chapter
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Example 12.1—Calculate the Mass of Water that
Can Be Vaporized with 155 KJ of Heat at 100 °C.
Given:
Find:
155 kJ
g H2O
Solution Map:
kJ
1 mol
40.7 kJ
Relationships:
Solution:
Check:
mol H2O
18.02 g
1 mol
g H2O
1 mol H2O = 40.7 kJ, 1 mol = 18.02 g
1 mol H 2 O 18.02 g
155 kJ 

 68.6 g H 2 O
40.7 kJ
1 mol
Since the given amount of heat is almost 4x the
DHvap, the amount of water makes sense.
Example 12.1:
• Calculate the amount of water in grams that can be
vaporized at its boiling point with 155 kJ of heat.
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Example:
Calculate the amount of
water in grams that can be
vaporized at its boiling
point with 155 kJ of heat.
• Write down the given quantity and its units.
Given:
155 kJ
Tro's Introductory Chemistry, Chapter
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Example:
Calculate the amount of
water in grams that can be
vaporized at its boiling
point with 155 kJ of heat.
Information:
Given: 155 kJ
• Write down the quantity to find and/or its units.
Find: ? g H2O
Tro's Introductory Chemistry, Chapter
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Example:
Calculate the amount of
water in grams that can be
vaporized at its boiling
point with 155 kJ of heat.
Information:
Given: 155 kJ
Find: g H2O
• Collect needed conversion factors:
DHvap = 40.7 kJ/mol  40.7 kJ  1 mol H2O
18.02 g H2O = 1 mol H2O
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Example:
Calculate the amount of
water in grams that can be
vaporized at its boiling
point with 155 kJ of heat.
Information:
Given: 155 kJ
Find: g H2O
Conversion Factors:
40.7 kJ = 1 mol; 18.02 g = 1 mol
• Write a solution map for converting the units:
kJ
mol H2O
1 mol H 2O
40.7 kJ
g H2O
18.02 g H 2O
1 mol H 2O
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Example:
Calculate the amount of
water in grams that can be
vaporized at its boiling
point with 155 kJ of heat.
Information:
Given: 155 kJ
Find: g H2O
Conversion Factors:
40.7 kJ = 1 mol; 18.02 g = 1 mol
Solution Map: kJ → mol → g
• Apply the solution map:
1 mol H 2O 18.02 g H 2O
155 kJ 

40.7 kJ
1 mol H 2O
= 68.626 g H2O
• Significant figures and round:
= 68.6 g H2O
Tro's Introductory Chemistry, Chapter
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Example:
Calculate the amount of
water in grams that can be
vaporized at its boiling
point with 155 kJ of heat.
Information:
Given: 155 kJ
Find: g H2O
Conversion Factors:
40.7 kJ = 1 mol; 18.02 g = 1 mol
Solution Map: kJ → mol → g
• Check the solution:
155 kJ of heat can vaporize 68.6 g H2O.
The units of the answer, g, are correct.
The magnitude of the answer makes sense
since it is more than one mole.
Tro's Introductory Chemistry, Chapter
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Practice—How Much Heat Energy, in kJ, is Required to
Vaporize 87 g of Acetone, C3H6O, (MM 58.08) at 25 C?
(DHvap = 31.0 kJ/mol)
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Practice—How Much Heat Energy, in kJ, Is Required to Vaporize
87 g of Acetone, C3H6O, (MM 58.08) at 25 C?
(DHvap = 31.0 kJ/mol), Continued
Given:
Find:
Solution Map:
Relationships:
Solution:
Check:
87 g C3H6O
kJ
g C3H6O
1 mol
58.08 g
mol C3H6O
kJ
31.0 kJ
1 mol
1 mol C3H6O = 31.0 kJ at 25 C, 1 mol = 58.08 g
1 mol C3H 6O
31.0 kJ
87 g C3H 6O 

 46 kJ
58.08 g
1 mol C3H 6O
Since the given mass is than one mole, the answer
being greater than DHvap makes sense.
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Temperature and Melting
• As you heat a solid, its temperature
increases until it reaches its
melting point.
• Once the solid starts to melt, the
temperature remains the same until
it all turns to a liquid.
• All the energy from the heat source
is being used to overcome some of
the attractive forces in the solid
that hold them in place.
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Energetics of Melting and Freezing
• When a solid melts, it absorbs heat from its
surroundings, it is endothermic.
• As heat flows out of the surroundings, it causes the
surroundings to cool.
 As heat flows out of your drink into the ice cubes (causing
them to melt), the liquid gets cooler.
• When a liquid freezes, it releases heat into its
surroundings, it is exothermic.
• As heat flows into the surroundings, it causes the
surroundings to warm.
 Orange growers often spray their oranges with water when a
freeze is expected. Why?
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Heat of Fusion
• The amount of heat needed to melt one mole of a solid
is called the heat of fusion.
 DHfus
 Fusion is an old term for heating a substance until it melts, it
is not the same as nuclear fusion.
• Since freezing (crystallization) is the opposite process
of melting, the amount of energy transferred is the
same, but in the opposite direction.
 DHcrystal = -DHfus
• In general, DHvap > DHfus because vaporization requires
breaking all attractive forces.
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Heats of Fusion of Several Substances
Liquid
Water
Isopropyl alcohol
Acetone
Diethyl ether
Chemical
formula
H2O
C3H7OH
C3H6O
C4H10O
Melting
point, °C
0.00
-89.5
-94.8
-116.3
Tro's Introductory Chemistry, Chapter
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DHfusion,
(kJ/mol)
6.02
5.37
5.69
7.27
50
Example 12.2—Calculate the Mass of Ice that Can
Be Melted with 237 kJ of Heat.
Given:
Find:
237 kJ
g H2O
Solution Map:
kJ
1 mol
6.02 kJ
Relationships:
Solution:
Check:
mol H2O
18.02 g
1 mol
g H2O
1 mol H2O = 6.02 kJ, 1 mol = 18.02 g
1 mol H2O 18.02 g
237 kJ 

 709 g H2O
6.02 kJ
1 mol
Since the given amount of heat is almost 4x the
DHvap, the amount of water makes sense.
51
Practice—How Much Heat Energy, in kJ, is Required to
Melt 87 g of Acetone, C3H6O, (MM 58.08)?
(DHfus = 5.69 kJ/mol)
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Practice—How Much Heat Energy, in kJ, Is Required to
Melt 87 g of Acetone, C3H6O, (MM 58.08)?, Continued
Given:
Find:
Solution Map:
Relationships:
Solution:
Check:
87 g C3H6O
kJ
g C3H6O
1 mol
58.08 g
mol C3H6O
kJ
5.69 kJ
1 mol
1 mol C3H6O = 5.69 kJ at -94.8 C, 1 mol = 58.08 g
1 mol C3H 6O
5.69 kJ
87 g C3H 6O 

 8.5 kJ
58.08 g
1 mol C3H 6O
Since the given mass is more than one mole, the
answer being greater than DHvap makes sense.
53
Sublimation
• Sublimation is a physical
change in which the solid
form changes directly to the
gaseous form.
Without going through the
liquid form.
• Like melting, sublimation is
endothermic.
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Intermolecular
Attractive Forces
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Effect of the Strength of
Intermolecular Attractions on
Properties
• The stronger the intermolecular attractions
are, the more energy it takes to separate the
molecules.
• Substances with strong intermolecular
attractions have higher boiling points,
melting points, and heat of vaporization;
they also have lower vapor pressures.
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Practice—Pick the Substance in Each Pair
with the Stronger Intermolecular Attractions.
• sugar or water
water.
• water or acetone
acetone.
• ice or dry ice.
ice
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Why Are Molecules Attracted to
Each Other?
• Intermolecular attractions are a result of attractive
forces between opposite charges.
• + ion to – ion.
• + end of one polar molecule to − end of another
polar molecule.
 H-bonding is especially strong.
 Even nonpolar molecules will have temporary induced
dipoles.
• Larger charge = stronger attraction.
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Dispersion Forces
• Also known as London forces or instantaneous
dipoles.
• Caused by distortions in the electron cloud of
one molecule inducing distortion in the electron
cloud on another.
• Distortions in the electron cloud lead to a
temporary dipole.
• The temporary dipoles lead to attractions
between molecules—dispersion forces.
• All molecules have attractions caused by
dispersion forces.
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Instantaneous Dipoles
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Strength of the Dispersion Force
• Depends on how easily the electrons can
move, or be polarized.
• The more electrons and the farther they are
from the nuclei, the larger the dipole that
can be induced.
• Strength of the dispersion force gets
larger with larger molecules.
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Dispersion Force and Molar Mass
Noble Gas Molar Mass
(g/mol)
4.00
He
Boiling Point
(K)
4.2
Ne
20.18
27
Ar
39.95
87
Kr
83.80
120
Xe
131.29
165
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Relationship Between Dispersion Force and
Molecular Size
250
200
BP, Noble Gas
150
BP, Halogens
Boiling Point, °C
100
BP, XH4
50
0
1
2
3
4
5
6
-50
-100
-150
-200
-250
-300
Period
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Practice—The Following Are All Made of Non–
Polar Molecules. Pick the Substance in Each Pair
with the Highest Boiling Point.
• CH4 or C3H8.
• BF3 or BCl3.
• CO2 or CS2.
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Permanent Dipoles
• Because of the kinds of
atoms that are bonded
together and their relative
positions in the molecule,
some molecules have a
permanent dipole.
 Polar molecules.
• The size of the molecule’s
dipole is measured in
debyes, D.
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Dipole-to-Dipole Attraction
• Polar molecules have a
permanent dipole.
A + end and a – end.
• The + end of one molecule
will be attracted to the –
end of another.
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Polarity and
Dipole-to-Dipole Attraction
CH3CH2CH3
Molar Mass
(g/mol)
44
Boiling
Point, °C
-42
Dipole
size, D
0
CH3-O-CH3
46
-24
1.3
CH3 - CH=O
44
20.2
2.7
CH3-CN
41
81.6
3.9
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Attractive Forces
Dispersion forces—All molecules.
+
+
+
+
_
_
_
_
Dipole-to-dipole forces—Polar molecules.
+
-
+
-
+
Tro's Introductory Chemistry, Chapter
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- +
-
68
Intermolecular Attraction and
Properties
• All molecules are attracted by dispersion
forces.
• Polar molecules are also attracted by dipoledipole attractions.
• Therefore, the strength of attraction is
stronger between polar molecules than
between nonpolar molecules of the same size.
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Practice—Determine Which of the Following Has
Dipole–Dipole Attractive Forces.
(EN C= 2.5, F = 4, H = 2.1, S = 2.5)
• CS2
Nonpolar bonds = nonpolar molecule.
• CH2F2
• CF4

S



C
S
Polar bonds and asymmetrical = polar molecule.
H
C
F
F
H
Polar bonds and symmetrical shape = nonpolar
molecule.
F
F
F
C
F
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70
Attractive Forces and Properties
• Like dissolves like.
 Miscible = Liquids that do not separate, no matter what
the proportions.
• Polar molecules dissolve in polar solvents.
 Water, alcohol, CH2Cl2.
 Molecules with O or N higher solubility in H2O due to
H-bonding with H2O.
• Nonpolar molecules dissolve in nonpolar solvents.
 Ligroin (hexane), toluene, CCl4.
• If molecule has both polar and nonpolar parts, then
hydrophilic-hydrophobic competition.
Tro's Introductory Chemistry, Chapter
12
71
Immiscible Liquids
When liquid pentane, a
nonpolar substance, is
mixed with water, a polar
substance, the two liquids
separate because they are
more attracted to their
own kind of molecule
than to the other.
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72
Hydrogen Bonding
• HF, or molecules that have
OH or NH groups have
particularly strong
intermolecular attractions.
 Unusually high melting and
boiling points.
 Unusually high solubility in
water.
• This kind of attraction is
called a hydrogen bond.
Tro's Introductory Chemistry, Chapter
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73
Properties and H-Bonding
Name
Molar
Formula mass
(g/mol)
Ethane C2H6
30.0
Boiling
point,
°C
Structure
H
H
H
C
C
H
H
Melting
point, Solubility
in water
°C
H
-88
-172
Immiscible
H
64.7
-97.8
Miscible
H
Ethanol CH4O
32.0
H
C
O
H
Tro's Introductory Chemistry, Chapter
12
74
Intermolecular H-Bonding
Tro's Introductory Chemistry, Chapter
12
75
Hydrogen Bonding
• When a very electronegative atom is bonded
to hydrogen, it strongly pulls the bonding
electrons toward it.
• Since hydrogen has no other electrons, when
it loses the electrons, the nucleus becomes
deshielded.
Exposing the proton.
• The exposed proton acts as a very strong
center of positive charge, attracting all the
electron clouds from neighboring molecules.
Tro's Introductory Chemistry, Chapter
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76
H-Bonds vs. Chemical Bonds
• Hydrogen bonds are not chemical bonds.
• Hydrogen bonds are attractive forces
between molecules.
• Chemical bonds are attractive forces that
make molecules.
Tro's Introductory Chemistry, Chapter
12
77
Relationship Between H-Bonding and
Intermolecular Attraction
150
H2O
Boiling Point, °C
100
50
HF
H2Te
0
1
NH3 2
-50
3
H2S
-200
5
SnH4
-100
-150
H2Se 4
GeH4
SiH4
CH4
BP, HX
BP, H2X
BP, H3X
Period
Tro's Introductory Chemistry, Chapter
12
BP, XH4
78
Attractive Forces and Properties
CH3CH2OCH2CH3
Molar Boiling
Solubility
Mass
Point,
in water
(g/mol)
°C
(g/100 g H2O)
74
34.6
7.5
CH3CH2CH2CH2CH3
72
36
Insoluble
CH3CH2CH2CH2OH
74
117
9
Tro's Introductory Chemistry, Chapter
12
79
Example 12.5—Which of the Following Is a Liquid
at Room Temperature? (The Other Two Are Gases.)
• formaldehyde, CH2O.
O
 30.03 g/mol.
g/mol
 polar molecule  dipole–dipole
dipole–dipole attractions present
present.
 Polar
polar C=O bond &
andasymmetric
asymmetric.
O
C
H
H
• fluoromethane, CH3F.
F
 34.03 g/mol.
g/mol
 polar molecule  dipole–dipole
dipole–dipole attractions present
present.
 Polar
polar C−F bond &
andasymmetric
asymmetric.
• hydrogen peroxide, H2O2
F
 34.02 g/mol.
g/mol
 polar molecule  dipole–dipole
dipole–dipole attractions present
present.
 Polar
polar H−O bonds &
andasymmetric
asymmetric.
 H−O bonds  hydrogen-bonding
Hydrogen bonding present.
present
Tro's Introductory Chemistry, Chapter
12
C
H
H
H
O
O
H
H
80
Practice–Pick the Compound in Each Pair Expected
to Have the Higher Solubility in H2O.
• CH3CH2OCH2CH3 or CH3CH2CH2CH2CH3.
• CH3CH2NHCH3 or CH3CH2CH2CH3.
• CH3CH2OH or CH3CH2CH2CH2CH2OH.
Tro's Introductory Chemistry, Chapter
12
81
Practice–Pick the Compound in Each Pair Expected to
Have the Higher Solubility in H2O, Continued.
• CH3CH2OCH2CH3 or CH3CH2CH2CH2CH3 contains polar O.
• CH3CH2NHCH3 or CH3CH2CH2CH3 contains polar N.
• CH3CH2OH or CH3CH2CH2CH2CH2OH contains less nonpolar parts.
Tro's Introductory Chemistry, Chapter
12
82
Types of Intermolecular Forces
Type of
force
Relative
strength
Present
in
Example
Weak, but
All atoms
Dispersion increases
and
H2
force
with molar
molecules
mass
Dipole–
Dipole
force
Hydrogen
Bond
Moderate
Only
polar
HCl
molecules
Strong
Molecules
having H
HF
bonded to
F, O, or N
Crystalline Solids
Tro's Introductory Chemistry, Chapter
12
84
Types of Crystalline Solids
Tro's Introductory Chemistry, Chapter
12
85
Molecular Crystalline Solids
• Molecular solids are
solids whose composite
units are molecules.
• Solid held together by
intermolecular attractive
forces.
Dispersion, dipole-dipole,
or H-bonding.
• Generally low melting
points and DHfusion.
Tro's Introductory Chemistry, Chapter
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86
Ionic Crystalline Solids
• Ionic solids are solids whose
composite units are formula units.
• Solid held together by electrostatic
attractive forces between cations
and anions.
 Cations and anions arranged in a
geometric pattern called a crystal
lattice to maximize attractions.
• Generally higher melting points and
DHfusion than molecular solids.
 Because ionic bonds are stronger than
intermolecular forces.
Tro's Introductory Chemistry, Chapter
12
87
Atomic Crystalline Solids
• Atomic solids are solids whose
composite units are individual
atoms.
• Solids held together by either
covalent bonds, dispersion
forces, or metallic bonds.
• Melting points and DHfusion vary
depending on the attractive
forces between the atoms.
Tro's Introductory Chemistry, Chapter
12
88
Practice—Classify Each of the Following Crystalline
Solids as Molecular, Ionic, or Atomic.
• H2O(s)—molecular.
O(s)
• Si(s)—atomic.
Si(s)
• C12H22O11(s)—molecular.
(s)
• CaF2(s)—ionic.
(s)
• Sc(NO3)3(s)—ionic.
(s)
Tro's Introductory Chemistry, Chapter
12
89
Types of Atomic Solids
Tro's Introductory Chemistry, Chapter
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90
Types of Atomic Solids:
Covalent
• Covalent atomic solids have their atoms attached by
covalent bonds.
 Effectively, the entire solid is one giant molecule.
• Because covalent bonds are strong, these solids have
very high melting points and DHfusion.
• Because covalent bonds are directional, these substances
tend to be very hard.
• Elements found as covalent atomic solids are C, Si, and
B.
• Compounds that occur as covalent atomic solids include
SiO2 and SiC.
Tro's Introductory Chemistry, Chapter
91
12
Types of Atomic Solids:
Nonbonding
• Nonbonding atomic solids are held
together by dispersion forces.
• Because dispersion forces are
relatively weak, these solids have
very low melting points and
DHfusion.
• All the noble gases form
nonbonding atomic solids.
Tro's Introductory Chemistry, Chapter
12
92
Types of Atomic Solids:
Metallic
• Metallic solids are held together
by metallic bonds.
• Metal atoms release some of
their electrons to be shared by
all the other atoms in the
crystal.
• The metallic bond is the
attraction of the metal cations
for the mobile electrons.
 Often described as islands of
cations in a sea of electrons.
Tro's Introductory Chemistry, Chapter
12
93
Metallic Bonding
• The model of metallic bonding
can be used to explain the
properties of metals.
• The luster, malleability, ductility,
and electrical and thermal
conductivity are all related to the
mobility of the electrons in the
solid.
• The strength of the metallic bond
varies, depending on the charge
and size of the cations, so the
melting points and DHfusion of
metals vary as well.
94
Substances with Both Bonding and
Nonbonding Attractions
• Some substances have chains or layers of
bonded atoms that are then attracted by
dispersion forces.
Chain substances include grey selenium,
polymeric SO3, and asbestos.
Layer substances include graphite, black
phosphorus, and mica.
Tro's Introductory Chemistry, Chapter
12
95
Practice—Decide if Each of the Following
Atomic Solids Is Covalent, Metallic, or
Nonbonding.
• diamond covalent.
• neon nonbonding.
• iron metallic.
Tro's Introductory Chemistry, Chapter
12
96
Water:
A Unique and Important Substance
• Water is found in all three states
on Earth.
• As a liquid, it is the most
common solvent found in nature.
• Without water, life as we know it
could not exist.
The search for extraterrestrial life
starts with the search for water.
Tro's Introductory Chemistry, Chapter
12
Water
• Liquid at room temperature.
Most molecular substances that have a molar
mass (18.02 g/mol) similar to water’s are gaseous.
• Relatively high boiling point.
• Expands as it freezes.
Most substances contract as they freeze.
Causes ice to be less dense than liquid water.
Tro's Introductory Chemistry, Chapter
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98