Intermolecular Attractions & the
Properties of Liquids & Solids
CHAPTER 12
Chemistry: The Molecular Nature of Matter, 6th edition
By Jesperson, Brady, & Hyslop
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
 Understand, describe, and rank in order of strength the types of
intermolecular forces.
 Difference between bonds and intermolecular forces
 Changes of state: heat of vaporization, fusion, & sublimation
 Clausius-Clapyron equation
 Heating and cooling curves: ΔH, phase transition temperatures
 Phase diagrams
 Solids: Unit cell, stoichiometry, packing patterns, XRD, common
types and their properties
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
2
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
Lecture Road Map:
① Properties of gas, liquids, solids
② Intermolecular forces
③ Changes of state
④ Dynamic Equilibrium
⑤ Structure & Characterization of a solid
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
3
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
Properties of
gases, liquids, &
solids
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
4
Intermolecular Forces
Important differences between gases,
solids, and liquids:
o Gases
oExpand to fill their container
o Liquids
oRetain volume, but not shape
o Solids
o Retain volume and shape
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
5
Intermolecular Forces
o Physical state of molecule depends on
o Average kinetic energy of particles
oRecall KE  Tave
o Intermolecular Forces
oEnergy of Inter-particle attraction
o Physical properties of gases, liquids and
solids determined by
o How tightly molecules are packed together
o Strength of attractions between
molecules
6
Intermolecular Attractions
o Converting gas  liquid or solid
o Molecules must get closer together
o Cool or compress
o Converting liquid or solid  gas
o Requires molecules to move farther
apart
o Heat or reduce pressure
o As T decreases, kinetic energy of
molecules decreases
o At certain T, molecules don’t have
enough energy to break away from one
another’s attraction
7
Inter vs. Intra-Molecular Forces
o Intramolecular forces
o Covalent bonds within molecule
o Strong
o Hbond (HCl) = 431 kJ/mol
o Intermolecular forces
o Attraction forces between molecules
o Weak
o Hvaporization (HCl) = 16 kJ/mol
Covalent Bond (strong)
Cl
8
H
Intermolecular attraction (weak)
Cl
H
Electronegativity Review
Electronegativity: Measure of attractive
force that one atom in a covalent bond has
for electrons of the bond
9
Bond Dipoles
o Two atoms with different electronegativity
values share electrons unequally
o Electron density is uneven
o Higher charge concentration around more
electronegative atom
o Bond dipoles
o Indicated with delta (δ) notation
o Indicates partial charge has arisen
H
F




10
o Net Dipoles
o Symmetrical molecules
o Even if they have polar bonds
o Are non-polar because bond dipoles cancel
o Asymmetrical molecules
o Are polar because bond dipoles do not cancel
o These molecules have permanent, net dipoles
o Molecular dipoles
o Cause molecules to interact
o Decreased distance between molecules increases
amount of interaction
o 11
COVALENT
BOND
✔
TiO2
CaBr2
IONIC
BOND
✔
CHCl3
F2
POLAR
COVALENT
BOND
✔
✔
Group
Problem
Identify the overall dipole moment for CHCl3
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
13
Group
Problem
Identify the overall dipole moment for these molecules:
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
14
Solubility
LIKE DISSOLVES LIKE
polar molecules dissolve in polar solvents
nonpolar molecules dissolve in nonpolar solvents
Polar Solvents
Water: H2O
Methanol: CH3OH
Ethanol: CH3CH2OH
Acetone: (CH3)2CO
Acetic Acid: CH3CO2H
Ammonia: NH3
Acetonitrile: CH3CN
Nonpolar Solvents
Pentane: C5H12
Hexane: C6H14
Cyclohexane: C6H12
Benzene: C6H6
Toluene: CH3C6H5
Chloroform: CHCl3
Diethylether: (CH3CH2)2O
15
Group
Problem
Which molecule will dissolve in
water?
Vitamin A
Vitamin B12
16
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
Intermolecular
Forces
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
17
Intermolecular Forces
The forces of attraction or repulsion between
neighboring particles (atoms or molecules).
+ / - charges attract one another
- / - or + / + forces repel each other
KE
r
18
r
Intermolecular Forces
o When substance melts or boils
o Intermolecular forces are broken, not covalent
bonds
o Responsible for non-ideal behavior of gases
o Responsible for existence of condensed
states of matter
o Responsible for bulk properties of matter
o Boiling points and melting points reflect strength
of intermolecular forces
19
Types of Intermolecular Forces
①
②
③
④
London dispersion forces
Dipole-dipole forces
Hydrogen bonds
Ion-dipole forces
o Ion-induced dipole forces
20
London-Dispersion Forces
o When atoms near one another,
their valence electrons interact
o Repulsion causes electron clouds
in each to distort and polarize
o Instantaneous dipoles result from
this distortion
o Effect enhanced with increased
volume of electron cloud size
o Effect diminished by increased
distance between particles and
compact arrangement of atoms
21
London Dispersion Forces
Affects ALL molecules, both polar & nonpolar
Boiling Point (BP) is an indication of relative intermolecular
force strength.
Ease with which dipole moments can be induced and thus
London Forces depend on
① Distance between particles
② Polarizability of electron cloud
③ Points of attraction
o Number atoms
o Molecular shape (compact or elongated)
22
Polarizability = Ease with which the electron
cloud can be distorted
Larger molecules often more polarizable
o Larger number of less tightly held
electrons
o Magnitude of resulting partial
charge is larger
o Larger electron cloud
23
Group
Problem
Which is more polarizable?
F2 or I2?
24
Table 12.1 Boiling Points of Halogens and
Noble Gases
Larger molecules have stronger London forces
and thus higher boiling points.
Number of Atoms in Molecule
o London dispersion forces increase with the number
atoms in molecule because more points of
attraction
Formula BP at 1 atm, C Formula BP at 1 atm, C
CH4
–161.5
C5H12
36.1
C2H6
–88.6
C6H14
68.7
C3H8
–42.1
:
:
C4H10
–0.5
C22H46
327
26
Group
Problem
Which of the following molecules will have the
highest boiling point?
Hexane, C6H14
Propane, C3H8
BP 68.7 °C
BP –42.1 °C
27
Molecular Shape
o Increased surface area available for contact =
increased points of contact = increase in
London Dispersion forces.
o More compact molecules:
Less surface area to interact with other
molecules
o Less compact molecules:
More surface area to interact with other
molecules
28
• Small area for
interaction
• Larger area for
interaction
More
compact – lower BP
29
Less compact – higher BP
Group
Problem
Which of the following molecules experience
the strongest Dispersion forces?
30
Types of Intermolecular Forces
①
②
③
④
London dispersion forces
Dipole-dipole forces
Hydrogen bonds
Ion-dipole forces
o Ion-induced dipole forces
31
Dipole-Dipole Attractions
o Occurs only between
polar molecules
o Proportional to distance
between molecules
o Polar molecules tend to
align their partial
charges: + / o As dipole moment
increases, intermolecular
force increases
32
+

+


+

+
+

+

Dipole-Dipole Attractions
Tumbling molecules
o Mixture of attractive and
repulsive dipole-dipole
forces
o Attractions (- -) are
maintained longer than
repulsions(- -)
o Get net attraction
o ~1–4% of covalent bond
33
Group
Problem
In the liquid state, which species has the
strongest intermolecular forces, CH4, Cl2, O2
or HF?
HF
The polar molecule
34
Types of Intermolecular Forces
①
②
③
④
London dispersion forces
Dipole-dipole forces
Hydrogen bonds
Ion-dipole forces
o Ion-induced dipole forces
35
Hydrogen Bonds
o Very strong dipole-dipole attraction: ~10% of a covalent
bond
o Occurs between H and highly electronegative atom (O, N, or
F): H—F, H—O, and H—N bonds very polar
o Electrons are drawn away from H giving atoms high
partial charges
o H only has one electron, so +H presents almost bare
proton
o –X almost full –1 charge
o Element’s small size, means high charge density
36
Examples of Hydrogen Bonding
H
O
H
H
H
O
H
H
N
H
H
H
H
H
H
H
F
N
O
H
H
H
O
H
H
H
F
H
N
N
H
H
H
O
H
H
H
N
H
37
Hydrogen Bonding in Water
Hydrogen Bonds are strong!
o Responsible for the high boiling point of water
o Responsible for expansion of water as it freezes
o Hydrogen bonding (dotted lines) between
water molecules in ice form tetrahedral configuration
38
Hydrogen Bonding in Water
1.97 Å
0.957 Å
Group
Problem
List all intermolecular forces for CH3CH2OH.
Hydrogen-bonds, dipole-dipole
attractions, London dispersion forces
40
Types of Intermolecular Forces
①
②
③
④
London dispersion forces
Dipole-dipole forces
Hydrogen bonds
Ion-dipole forces
o Ion-induced dipole forces
41
Ion-Dipole Attractions
o Attractions between ion and charged end of
polar molecules
o Ions have full charges, increasing the attraction
(a) Negative ends of water dipoles surround cation
(b) Positive ends of water dipoles surround anion
42
AlCl3·6H2O
Attractions between ion and polar molecules
o Positive charge of Al3+ ion
attracts partial negative
charges – on O of water
molecules
o Ion-dipole attractions hold
water molecules to metal ion in
hydrate
o Water molecules are found
at vertices of octahedron
around aluminum ion
43
Ion-Induced Dipole Attractions
o Attractions between ion and dipole it induces on
neighboring molecules
o Depends on
oIon charge and
oPolarizability of its neighbor
o Attractions can be quite strong as ion charge is
constant, unlike instantaneous dipoles of Londondispersion forces
44
Group
Problem
How many water molecules would be attracted to
this molecule by Ion-Dipole interactions?
45
Group
Problem
List the intermolecular forces and rank in order of
strength for the liquids of each molecule.
46
Group
Problem
o Ion-Dipole
o Hydrogen Bonding
o Dipole-Dipole
o London Forces
Strongest
Weakest
• Larger, longer, and therefore heavier molecules
often have stronger intermolecular forces
• Smaller, more compact, lighter molecules have
generally weaker intermolecular forces
Intermolecular Forces and
Temperature
Decrease with increasing temperature
o Increasing kinetic energy overcomes attractive
forces
o If allowed to expand, increasing temperature
increases distance between gas particles and
decreases attractive forces
48
Group
Problem
GROUP PROBLEM SET 12.1
49
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
Diffusion
Compressibility
Boiling
Point
Surface Tension
More properties
of gases,
liquids, & solids
Wetting
Melting
Point
Viscosity
Retention of Volume
& shape
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
50
Melting & Boiling Point
Often can predict physical properties by comparing
strengths of intermolecular attractions:
Boiling Point increases when
intermolecular forces increase
Melting Point increases when
intermolecular forces increase
51
Compressibility
Measure of the ability of a substance to be
forced into smaller volume
o Determined by strength of intermolecular forces
o Gases highly compressible
o Molecules far apart
o Weak intermolecular forces
o Solids and liquids nearly incompressible
o Molecules very close together
o Stronger intermolecular forces
52
Retention of volume and shape
o Solids retain both volume and shape
o Strongest intermolecular attractions
o Molecules closest
o Liquids retain volume, but not shape
o Attractions intermediate
o Gases, expand to fill their containers
o Weakest intermolecular attractions
o Molecules farthest apart
53
Diffusion
In Gases
o Molecules travel long
distances between collisions
o Diffusion rapid
In Liquids
o Molecules closer
o Encounter more collisions
o Takes a long time to move
from place to place
In Solids
o Diffusion close to zero at room
temperature
o Will increase at high
temperature
54
Surface Tension
Why does H2O bead up on a freshly waxed car
instead of forming a layer?
Inside body of liquid
o Intermolecular forces are
the same in all directions
Molecules at surface
o Potential energy
increases when removing
neighbors
o Molecules move together
to reduce surface area
and potential energy
 sphere
55
Surface Tension
Liquids containing molecules
with strong intermolecular
forces have high surface
tension
Allows us to fill glass above rim o Surface tension
o Gives surface rounded
appearance
o Surface resists expansion and
pushes back
56
increases as
intermolecular
forces increase
o Surface tension
decreases as
temperature
increases
Wetting
o Ability of liquid to spread
across surface to form
thin film
o Greater similarity in
attractive forces
between liquid and
surface, yields greater
wetting effect
o Occurs only if
intermolecular attractive
force between surface
and liquid about as
strong as within liquid
itself
57
Wetting: Surfactants (Detergents)
o Detergents added to water to lower surface tension so water can
spread on greasy glass
o Substances that have both polar and non-polar characteristics
o Long chain hydrocarbons with polar tail
O
O Na+
O
O
S
O
O Na+
o Nonpolar end dissolves in nonpolar grease
o Polar end dissolves in polar H2O
o Thus increasing solubility of grease in water
58
Viscosity
o Resistance to flow
o Measure of fluid’s
resistance to flow or
changing form
o Decreases as Temp
increases
o Not just a property of
liquids:
o Gas: respond to instantly
to form changing force
o Amorphous solids, like
glass
59
Group
Problem
Viscosity
Acetone
Polar molecule
o Dipole-dipole and
o London forces
Which is more viscous?
60
Ethylene glycol
Polar molecule
o Hydrogen-bonding
o Dipole-dipole and
o London forces
Group
Problem
For each pair given, which is has more viscosity?
CH3CH2CH2CH2OH,
C6H14,
NH3(l ),
61
C12H26
PH3(l )
CH3CH2CH2CHO
Group
Problem
GROUP PROBLEM SET 12.2
62
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
Phase Diagrams
Heating/Cooling
Curves
ΔH
Changes of
State
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
63
Phase Changes = changes of physical state with
temperature ( α to KE)
fusion
SOLID
evaporation
LIQUID
freezing
GAS
condensation
deposition
sublimation
endothermic
exothermic
64
System absorbs energy from surrounds in the form of heat
o Requires the addition of heat
System releases energy into surrounds in the form of heat or light
o Requires heat to be decreased
Phase Changes of Water
melting
ICE
evaporation
WATER
freezing
forming dew
deposition
sublimation
65
VAPOR
Phase Changes
Energy of System
Gas
Vaporization
Condensation
Sublimation
Liquid
Melting
or Fusion
Freezing
Solid
 Exothermic, releases heat
 Endothermic, absorbs heat
66
Deposition
Heating Curve
If heat added at constant rate
Horizontal lines
o Phase changes
o Melting point
o Boiling point
Diagonal lines
o Heating of solid,
liquid or gas
Cooling Curve
Heat removed at constant rate
Horizontal lines
o Phase changes
o Melting point
o Boiling point
Diagonal lines
o Cooling of
solid, liquid or
gas
Supercooling
o Temperature of liquid dips below its freezing point
o68
Boiling Point (bp)
Bp increases as strength of intermolecular
forces increase
Normal Boiling Point
• T at which vapor pressure of liquid = 1 atm
69
Rate of Evaporation
o Depends on
o Temperature
o Surface area
o Strength of
intermolecular
attractions
o Molecules that escape
from liquid have larger
than minimum escape
KE
o When they leave the
average KE of remaining
molecules is less and so
T lower
70
Effect of Temperature on Evaporation Rate
For given liquid the rate of
evaporation per unit surface
area increases as T increases
Why?
o At higher T, total fraction
of molecules with KE
large enough to escape
is larger
o Result: rate of
evaporation is larger
71
Kinetic Energy Distribution
in 2 different liquids
A
o Smaller intermolecular
forces
o Lower KE required to
escape liquid
o A evaporates faster
72
B
o Larger intermolecular
forces
o Higher KE required to
escape liquid
o B evaporates slower
Group
Problem
What is an example of gas A and gas B?
73
Effects of Hydrogen Bonding
• Boiling points of
hydrogen compounds
of elements of Groups
4A, 5A, 6A, and 7A.
• Boiling points of
molecules with
hydrogen bonding are
much higher than
expected
74
Energies of Phase Changes
Hfus
Hvap
fusion
evaporation
SOLID
LIQUID
freezing
GAS
condensation
deposition
sublimation
Hsub
Molar heat of fusion (Hfus)
Heat absorbed by one mole of solid when it melts to give liquid at constantT
and P
Molar heat of vaporization (Hvap )
Heat absorbed when one mole of liquid is changed to one mole of vapor at
constant T and P
Molar heat of sublimation (Hsub )
Heat absorbed by one mole of solid when it sublimes to give one mole of
vapor at constant T and P
Measuring Hvap
o Clausius-Clapeyron equation
o Measure pressure at various temperatures, then
plot
 Hvap
ln P  
 R
1
 C
T

o Two point form of Clausius-Clapeyron equation
o Measure pressure at two temperatures and solve
equation
P1 Hvap
ln

P2
R
1
1
  
T 2 T1 
76
Vapor Pressure Diagram
o Variation of vapor
pressure with T
o Ether
o Volatile
o High vapor pressure
near RT
o Propylene glycol
o Non-volatile
o Low vapor pressure
near RT
RT = 25 C
77
 Hvap
ln P  
 R
Temp (K) Vapor P
280
32.4
300
92.5
320
225
330
334
340
483
1/T
0.003571429
0.003333333
0.003125
0.003030303
0.002941176
1
 C
T

lnP
3.478158423
4.527208645
5.416100402
5.811140993
6.180016654
Slope = Hvap/R = -4288.1K
Hvap = 8.3145 Jmol/K x 4288.1K
Hvap = 35.65 x 103 J/mol
Hvap = 35.65 kJ/mol
The vapor pressure of diethyl ether is 401 mm Hg at
18 °C, and its molar heat of vaporization is 26 kJ/mol.
Calculate its vapor pressure at 32 °C.
P1 Hvap
ln

P2
R
1
1  T1 = 273.15 + 18 = 291.15 K
  
T 2 T1  T2 = 273.15 + 32 = 305.15 K
ö
2.6 ´ 104 J/mol æ
1
1
çç
÷÷ = -0.4928
ln =
P2 8.314 J/(K × mol) è 305.15 K 291.15 K ø
P1
P1
 e 0.4928  0.6109
P2
P1
0.6109
401 mm Hg
2
P2 =
= 6.6 ´10 mm Hg
0.6109
 P2
79
Group
Problem
Determine the enthalpy of vaporization, in
kJ/mol, for benzene, using the following vapor
pressure data.
T = 60.6 °C; P = 400 torr
T = 80.1 °C; P = 760 torr
A. 32.2 kJ/mol
B. 14.0 kJ/mol
C. –32.4 kJ/mol
D. 0.32 kJ/mol
E. –14.0 kJ/mol
80
Group
Problem
DHvap æ 1 1 ö
ç - ÷
ln =
P2
R çèT2 T1 ÷ø
P1
400 mm Hg
ln
=
760 mm Hg
DHvap
81
DHvap
æ
ö
1
1
çç
÷÷
è 353.1 K 333.6 K ø
J
8.314
K mol
= 32,235 J/mol or 32.2 kJ/mol
Phase Diagrams
• Show the effects of both pressure and temperature
on phase changes
• Boundaries between phases indicate equilibrium
• Triple point:
– The temperature and pressure at which s, l, and g are all
at equilibrium
• Critical point:
– The temperature and pressure at which a gas can no
longer be condensed
– TC = temperature at critical point
– PC = pressure at critical point
82
Phase Diagram
F
X axis – temperature
Y axis – pressure
o As P increases
(T constant), solid most
E
likely more compact
o As T increases
(P constant), gas most likely
higher energy
o Each point = T and P
o B = 0.01 °C, 4.58 torr
o E = 100 °C, 760 torr
o F = –10 °C, 2.15 torr
Phase Diagram of Water
AB = vapor pressure
curve for ice
BD = vapor pressure
curve for liquid water
BC = melting point line
B = triple point: T and P
where all three phases are
in equilibrium
D = critical point
T and P above which liquid
does not exist
84
Phase Diagram – CO2
o Now line
between solid
and liquid slants
to right
o More typical
o Where is triple
point?
o Where is critical
point?
85
Supercritical Fluid
o Substance with temperature above its critical
temperature (TC) and density near its liquid
density
o Have unique properties that make them
excellent solvents
o Values of TC tend to increase with increased
intermolecular attractions between particles
86
Group
Problem
At 89 °C and 760 mmHg, what
physical state is present?
A.Solid
B.Liquid
C.Gas
D.Supercritical fluid
E.Not enough information
is given
87
Group
Problem
GROUP PROBLEM SET 12.3
88
The Before & After of Phase Changes
fusion
SOLID
evaporation
LIQUID
freezing
GAS
condensation
deposition
sublimation
endothermic
exothermic
89
System absorbs energy from surrounds in the form of heat
o Requires the addition of heat
System releases energy into surrounds in the form of heat or light
o Requires heat to be decreased
The molar heat of a phase
change (H) describes the heat
needed for a phase change to
go to completion.
The specific heat of a phase
change (q) describes the heat
needed for an amount of a
substance to completely
undergo a phase change.
q = n x H
90
Enthalpy Of Phase Changes
Endothermic Phase Changes
1. Must add heat
2. Energy entering system (+)
Sublimation: Hsub > 0
Vaporization: Hvap > 0
Melting or Fusion: Hfus > 0
Exothermic Phase Changes
1. Must give off heat
2. Energy leaving system (–)
91
Deposition: H < 0 = –Hsub
Condensation: H < 0 = –Hvap
Freezing: H < 0 = –Hfus
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
Dynamic
Equilibria
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
92
Equilibria Exist During a Phase Change
• Fraction of molecules in condensed state is higher when
intermolecular attractions are higher
• Intermolecular attractions must be overcome to separate
the particles, while separated particles are simultaneously
attracted to one another
condensed
phase
separated
phase
93
Le Chatelier’s Principle
o Equilibria are often disturbed or upset
o When dynamic equilibrium of system is
upset by a disturbance
o System responds in direction that tends to
counteract disturbance and, if possible, restore
equilibrium
o Position of equilibrium
o Used to refer to relative amounts of substance on
each side of double (equilibrium) arrows
94
Liquid Vapor Equilibrium
Liquid + Heat  Vapor
• Increasing T
– Increases amount of vapor
– Decreases amount of liquid
• Equilibrium has shifted
– Shifted to the right
– More vapor is produced at expense of liquid
• Temperature-pressure relationships can be
represented using a phase diagram
95
Equilibrium & Phase Diagrams
T1 = 78°C
P1 = 330 atm
To increase
T2 = 100°C
The system must
respond by
increasing
P2 = 760 to restore
equilibrium:
o T is higher
o Volume of liquid
is lower
o P of vapor higher
96
Le Chatelier’s Principle
Liquid + Heat  Vapor
Initial
V1
T1
P1
Change
Volume lost
in
evaporation
Increase
Temperature
Pressure
increases
Final
V2
T2
P2
Evaporation Rate
98
Before System Reaches Equilibrium
o Liquid is placed in empty, closed,
container
o Begins to evaporate
o Once in gas phase
o Molecules can condense by
o Striking surface of liquid and
giving up some kinetic energy
99
System At Equilibrium
o Rate of evaporation = rate of
condensation
o Occurs in closed systems
where molecules cannot
escape
100
Enthalpy Of Phase Changes
Endothermic:
Liquid+ heat of vaporization ↔ Gas
Liquid + Hvap ↔ Gas
Solid + heat of fusion ↔ Liquid
Solid + Hfus ↔ Liquid
Solid + heat of sublimation ↔ Gas
Solid + Hsub ↔ Liquid
Exothermic:
Liquid ↔ Gas - Hvap
Solid ↔ Liquid - Hfus
Solid ↔ Liquid - Hsub
101
liquid + heat of vaporization ↔ gas
Equilibrium Vapor Pressure
o Pressure of gas when liquid or solid is at
equilibrium with its gas phase
o Usually referred to as simply vapor pressure
o Increasing temperature increases vapor pressure
because vaporization is endothermic
102
Vapor Pressure Diagram
• Variation of vapor
pressure with T
• Ether
– Volatile
– High vapor pressure
near RT
• Propylene glycol
RT = 25 C
– Non-volatile
– Low vapor pressure
near RT
103
Effect of Volume on Vapor Pressure
Initial
(equilibrium
exists)
Volume of
Container
Volume of
liquid
Change
Volume
manually
increased
Rate
condensation
decreases
Pressure
decreases
System changes
to establish new
equilibrium
Volume
of container
greater
Volume of
liquid
decreases
P2
(P2 = P1)
P1
104
Similar Equilibria Reached in Melting
Melting Point (mp)
o Solid begins to change into
liquid as heat added
Dynamic equilibria exists
between solid and liquid states
o Melting (red arrows) and
freezing (black arrows)
occur at same rate
o As long as no heat added or
removed from equilibrium
mixture
105
Equilibria Reached in Sublimation
At equilibrium
molecules sublime from
solid at same rate as
molecules condense from
vapor
106
Do Solids Have Vapor Pressures?
o At given temperature some solid particles
have enough KE to escape into vapor
phase
o When vapor particles collide with surface
they can be captured
o Yes equilibrium vapor pressure of solid
exists
107
CHAPTER 12 Intermolecular Attractions &
the Properties of Liquids & Solids
Solid
Structures
Jesperson, Brady, Hyslop. Chemistry:
The Molecular Nature of Matter, 6E
108
Types of Solids
• Crystalline Solids
– Solids with highly regular arrangements of
components
• Amorphous Solids
– Solids with considerable disorder in their
structures
109
Crystalline Solids
• Unit Cell
– Smallest
segment that
repeats
regularly
– Smallest
repeating unit of
lattice
– Twodimensional unit
cells
110
Crystal Structures Have Regular Patterns
• Lattice
– Many repeats of unit cell
– Regular, highly
symmetrical system
– Three (3) dimensional
system of points
designating positions of
components
• Atoms
• Ions
• Molecules
111
Three Types Of 3-D Unit Cells
• Simple cubic
– Has one host atom at each corner
– Edge length a = 2r
– Where r is radius of atom or ion
• Body-centered cubic (BCC)
– Has one atom at each corner and one in
4r
center
a=
– Edge length
3
• Face-centered cubic (FCC)
– Has one atom centered in each face, and
one at each corner
– Edge length
a = 4r / 2
112
Close Packing of Spheres
1st
layer
2nd layer
 Most efficient arrangement of spheres in two
dimensions
 Each sphere has 6 nearest neighbors
 Second layer with atoms in holes on the first
layer
113
Two Ways to Put on Third Layer
Cubic lattice: 3-dimensional arrays
1. Directly above
spheres in first
layer
114
2. Above holes in first
layer
 Remaining holes not
covered by second layer
3-D Simple Cubic Lattice
Unit
Cell
Portion of lattice—
open view
115
Space filling
model
Other Cubic Lattices
Face Centered
Cubic
116
Body Centered
Cubic
Ionic Solids
Lattices of alternating charges
• Want cations next to anions
– Maximizes electrostatic attractive forces
– Minimizes electrostatic repulsions
• Based on one of three basic lattices:
– Simple cubic
– Face centered cubic
– Body centered cubic
117
Common Ionic Solids
Rock salt or NaCl
– Face centered cubic lattice of Cl– ions (green)
– Na+ ions (blue) in all octahedral holes
118
Other Common Ionic Solids
Cesium
Chloride, CsCl
119
Zinc Sulfide,
ZnS
Calcium
Fluoride, CaF2
Spaces In Ionic Solids Are Filled With
Counter Ions
• In NaCl
– Cl– ions form facecentered cubic unit
cell
– Smaller Na+ ions fill
spaces between Cl–
ions
• Count atoms in unit
cell
– Have 6 of each or
1:1 Na+:Cl– ratio
120
Counting Atoms per Unit Cell
• Four types of sites in unit cell
– Central or body position – atom is completely contained
in one unit cell
– Face site – atom on face shared by two unit cells
– Edge site – atom on edge shared by four unit cells
– Corner site – atom on corner shared by eight unit cells
Site
Body
Face
Edge
Corner
Counts as Shared by X unit cells
1
1/2
1/4
1/8
1
2
4
8
121
Example: NaCl
Face
Center
Edge
Corner
Site
# of Na+
# of Cl–
Body
1
0
Face
0
Edge
12  1 4   3
6  12   3
Corner
0
8  1 8  1
Total
4
4
0
122
Determine the number of each type of ion in
the unit cell.
1:1
CsCl
123
4:4
ZnS
4:8
CaF2
Some Factors Affecting Crystalline
Structure
• Size of atoms or ions involved
• Stoichiometry of salt
• Materials involved
– Some substances do not form crystalline solids
124
Amorphous Solids (Glass)
• Have little order, thus referred to as “super cooled
liquids”
• Edges are not clean, but ragged due to the lack of
order
125
X-Ray Crystallography
• X rays are passed through
crystalline solid
• Some x rays are absorbed,
most re-emitted in all
directions
• Some emissions by atoms
are in phase, others out of
phase
• Emission is recorded on film
126
X-ray Diffraction
Experimental Setup
127
Diffraction Pattern
Interpreting Diffraction Data
• As x rays hit atoms
in lattice they are
deflected
• Angles of
deflections related
to lattice spacing
• So we can
estimate atomic
and ionic radii from
distance data
128
Interpreting Diffraction Data
Bragg Equation
• nλ=2d sinθ
– n = integer (1, 2, …)
–  = wavelength of
X rays
– d = interplane
spacing in crystal
–  = angle of incidence
and angle of
reflectance of
X rays to various
crystal planes
129
Example: Diffraction Data
The diffraction pattern of copper metal was
measured with X-ray radiation of wavelength of
131.5 pm. The first order (n = 1) Bragg
diffraction peak was found at an angle θ of
50.5°. Calculate the spacing between the
diffracting planes in the copper metal.
n = 2d sin 
1(131.5 pm) = 2 × d × sin(50.5)
d = 283 pm
130
Example: Using Diffraction Data
X-ray diffraction measurements reveal that
copper crystallizes with a face-centered cubic
lattice in which the unit cell length is 362 pm.
What is the radius of a copper atom expressed
in picometers?
This is basically a geometry problem.
131
Ex. Using Diffraction Data (cont.)
Pythagorean theorem: a2 + b2 = c2
Where a = b = 362 pm sides and c = diagonal
2a2 = c2
and
c  2a 2  2a
diagonal = 2 ´ (362 pm) = 512 pm
diagonal = 4  rCu = 512 pm
rCu = 128 pm
132
Ionic Crystals (e.g. NaCl, NaNO3)
•
•
•
•
•
•
•
Have cations and anions at lattice sites
Are relatively hard
Have high melting points
Are brittle
Have strong attractive forces between ions
Do not conduct electricity in their solid states
Conduct electricity well when molten
133
Group
Problem
Potassium chloride crystallizes with the rock
salt structure. When bathed in X rays, the
layers of atoms corresponding to the surfaces
of the unit cell produce a diffracted beam of
X rays (λ=154 pm) at an angle of 6.97°. From
this, calculate the density of potassium
chloride in g/cm3.
134
Covalent Crystals
• Lattice positions occupied by atoms that are
covalently bonded to other atoms at
neighboring lattice sites
• Also called network solids
– Interlocking network of covalent bonds extending
all directions
• Covalent crystals tend to
– Be very hard
– Have very high melting points
– Have strong attractions between covalently
bonded atoms
135
Ex. Covalent (Network) Solid
• Diamond (all C)
– Shown
• SiO2 silicon oxide
– Alternating Si and O
– Basis of glass and quartz
• Silicon carbide (SiC)
136
Metallic Crystals
• Simplest models
– Lattice positions of metallic
crystal occupied by positive
ions
– Cations surrounded by
“cloud” of electrons
• Formed by valence electrons
• Extends throughout entire solid
137
Metallic Crystals
• Conduct heat and electricity
– By their movement, electrons transmit kinetic
energy rapidly through solid
• Have the luster characteristically associated
with metals
– When light shines on metal
– Loosely held electrons vibrate easily
– Re-emit light with essentially same frequency
and intensity
138
Group
Problem
GROUP PROBLEM SET 12.3
139