Chemistry: The Molecular Nature of Matter, 6E
Jespersen/Brady/Hyslop
Intermolecular Forces
 Important differences between gases, solids,
and liquids:
 Gases

Expand to fill their container
 Liquids

Retain volume, but not shape
 Solids

Retain volume and shape
2
Intermolecular Forces
 Physical state of molecule depends on
 Average kinetic energy of particles

Recall KE  Tave
 Intermolecular Forces

Energy of Inter-particle attraction
 Physical properties of gases, liquids and solids
determined by
 How tightly molecules are packed together
 Strength of attractions between
molecules
3
Intermolecular Attractions
 Converting gas  liquid or solid
 Molecules must get closer together

Cool or compress
 Converting liquid or solid  gas
 Requires molecules to move farther apart

Heat or reduce pressure
 As T decreases, kinetic energy of molecules
decreases
 At certain T, molecules don’t have enough
energy to break away from one another’s
attraction
4
Inter vs. Intra-Molecular Forces
 Intramolecular forces
 Covalent bonds within molecule
 Strong
 Hbond (HCl) = 431 kJ/mol
 Intermolecular forces
 Attraction forces between molecules
 Weak
 Hvaporization (HCl) = 16 kJ/mol
Covalent Bond (strong)
Cl
H
Intermolecular attraction (weak)
Cl
H
5
Electronegativity Review
Electronegativity: Measure of attractive force that
one atom in a covalent bond has for electrons of the
bond
6
Bond Dipoles
 Two atoms with different electronegativity values
share electrons unequally
 Electron density is uneven
 Higher charge concentration around more
electronegative atom
 Bond dipoles
 Indicated with delta (δ) notation
 Indicates partial charge has arisen
H
F




7
Net Dipoles
 Symmetrical molecules
 Even if they have polar bonds
 Are non-polar because bond dipoles cancel
 Asymmetrical molecules
 Are polar because bond dipoles do not cancel
 These molecules have permanent, net dipoles
 Molecular dipoles
 Cause molecules to interact
 Decreased distance between molecules increases amount
of interaction
8
COVALENT
BOND
CHCl3
TiO2
F2
CaBr2
POLAR
COVALENT
BOND
IONIC
BOND
Group
Problem
Identify the overall dipole moment for CHCl3
Jesperson, Brady, Hyslop. Chemistry: The
Molecular Nature of Matter, 6E
10
Group
Problem
Identify the overall dipole moment for these molecules:
Jesperson, Brady, Hyslop. Chemistry: The
Molecular Nature of Matter, 6E
11
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
12
Intermolecular Forces

When substance melts or boils




Intermolecular forces are broken
Not covalent bonds
Responsible for non-ideal behavior of gases
Responsible for existence of condensed states of
matter
Responsible for bulk properties of matter



Boiling points and melting points
Reflect strength of intermolecular forces
13
Three Important Types of
Intermolecular Forces
1. London dispersion forces
2. Dipole-dipole forces
 Hydrogen bonds
3. Ion-dipole forces
 Ion-induced dipole forces
14
London Forces
 When atoms near one another,
their valence electrons interact
 Repulsion causes electron clouds
in each to distort and polarize
 Instantaneous dipoles result from
this distortion
 Effect enhanced with increased
volume of electron cloud size
 Effect diminished by increased
distance between particles and
compact arrangement of atoms
15
London Forces

Affect All molecules, both polar and 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
 Distance between particles
 Polarizability of electron cloud
 Points of attraction


Number of atoms
Molecular shape (compact or elongated)
16
Polarizability
 Ease with which the electron
cloud can be distorted
 Larger molecules often more
polarizable
 Larger number of less tightly held
electrons
 Magnitude of resulting partial
charge is larger
 Larger electron cloud
17
Group
Problem
Arrange the following atoms in order of
increasing polarizability: Ar, He,
Kr, Ne and Xe.
18
Table 12.1 Boiling Points of Halogens and
Noble Gases
Larger molecules have stronger London forces
and thus higher boiling points.
19
Number of Atoms in Molecule
 London forces depend on number atoms in molecule
 Boiling point of hydrocarbons demonstrates this trend
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
20
Group
Problem
Which of the following molecules will
have the highest boiling point?
More sites (marked with *) along its chain where attraction to other molecules can
occur
21
Molecular Shape
 Increased surface area available for contact = increased
London forces
 London dispersion forces between spherical molecules
are lower than chain-like molecules
 More compact molecules
Hydrogen atoms not as free to interact with
hydrogen atoms on other molecules
 Less compact molecules
 Hydrogen atoms have more chance to interact
with hydrogen atoms on other molecules

22
Physical Origin of Shape Effect
 Small area for interaction  Larger area for
interaction
More compact – lower BP
Less compact – higher BP
23
Dipole-Dipole Attractions
 Occur only between polar
+

+


+

+
+

+

molecules
 Possess dipole moments
 Molecules need to be
close together
 Polar molecules tend to
align their partial charges
 Positive to negative
 As dipole moment
increases, intermolecular
force increases
24
Dipole-Dipole Attractions
 Tumbling molecules
 Mixture of attractive and
repulsive dipole-dipole
forces
 Attractions (- -) are
maintained longer than
repulsions(- -)
 Get net attraction
 ~1–4% of covalent bond
25
Dipole-Dipole Attractions
 Interactions between net dipoles in polar
molecules
 About 1–4% as strong as a covalent bond
 Decrease as molecular distance increases
 Dipole-dipole forces increase with increasing
polarity
26
Hydrogen Bonds
 Special type of dipole-dipole Interaction
 Very strong dipole-dipole attraction
 ~10% of a covalent bond
 Occurs between H and highly electronegative atom (O,
N, or F)
 H—F, H—O, and H—N bonds very polar
 Electrons are drawn away from H, so high partial charges
 H only has one electron, so +H presents almost bare proton
 –X almost full –1 charge
 Element’s small size, means high charge density
 Positive end of one can get very close to negative end of another
27
Examples of Hydrogen Bonding
H
H
O
H
H
N
H
H
H
O
H
H
H
H
H
O
H
H
F
N
H
H
O
H
H
H
F
H
N
N
H
H
H
O
H
H
H
N
H
28
Hydrogen Bonding in Water
 Responsible for expansion of water as it freezes
 Hydrogen bonding produces strong attractions in
liquid
 Hydrogen bonding (dotted lines) between
water molecules in ice form tetrahedral configuration
29
Hydrogen Bonding in Water
1.97 Å
0.957 Å
Your Turn
List all intermolecular forces for CH3CH2OH.
A. Hydrogen-bonds
B. Hydrogen-bonds, dipole-dipole attractions,
London dispersion forces
C. Dipole-dipole attractions
D. London dispersion forces
E. London dispersion forces, dipole-dipole
attractions
31
Your Turn
In the liquid state, which species has the strongest
intermolecular forces, CH4, Cl2, O2 or HF?
A. CH4
B. Cl2
C. O2
D. HF
32
Ion-Dipole Attractions
 Attractions between ion and charged end of polar
molecules
 Attractions can be quite strong as ions have full charges
(a) Negative ends of water dipoles surround cation
(b) Positive ends of water dipoles surround anion 33
Ex. Ion-Dipole Attractions: AlCl3·6H2O
 Attractions between ion and polar molecules
 Positive charge of Al3+ ion
attracts partial negative
charges – on O of water
molecules
 Ion-dipole attractions hold
water molecules to metal
ion in hydrate
 Water molecules are found
at vertices of octahedron
around aluminum ion
34
Ion-Induced Dipole Attractions
 Attractions between ion and dipole it induces on
neighboring molecules
 Depends on
Ion charge and
 Polarizability of its neighbor

 Attractions can be quite strong as ion charge is
constant, unlike instantaneous dipoles of ordinary
London forces
 E.g., I– and Benzene
35
Group
Problem
List the intermolecular forces and rank in
order of strength.
36
Summary of Intermolecular Attractions
Dipole-dipole
 Occur between neutral molecules with permanent dipoles
 About 1–4% of covalent bond
 Mid range in terms of intermolecular forces
Hydrogen bonding



Special type of dipole-dipole interaction
Occur when molecules contain N—H,
H—F and O—H bonds
About 10% of a covalent bond
37
Summary of Intermolecular Attractions
London dispersion
 Present in all substances
 Weakest intermolecular force
 Weak, but can add up to large net attractions
Ion-dipole
 Occur when ions interact with polar molecules
 Strongest intermolecular attraction
Ion-induced dipole
 Occur when ion induces dipole on neighboring particle
 Depend on ion charge and polarizability of its neighbor
38
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
39
Physical Properties that Depend on How
Tightly Molecules Pack
 Compressibility
 Measure of ability of substance to be forced into
smaller volume
 Determined by strength of intermolecular forces
 Gases highly compressible


Molecules far apart
Weak intermolecular forces
 Solids and liquids nearly incompressible
 Molecules very close together
 Stronger intermolecular forces
40
Intermolecular Forces Determine
Strength of Many Physical Properties
 Retention of volume and shape
 Gases, expand to fill their containers
 Weakest intermolecular attractions
 Molecules farthest apart
 Liquids retain volume, but not shape
 Attractions intermediate
 Solids retain both volume and shape
 Strongest intermolecular attractions
 Molecules closest
41
Intermolecular Forces and
Temperature
 Decrease with increasing temperature
 Increasing kinetic energy overcomes attractive
forces
 If allowed to expand, increasing temperature
increases distance between gas particles and
decreases attractive forces
42
Diffusion
 Movement that spreads
one gas though another
gas to occupy space
uniformly
 Spontaneous
intermingling of
molecules of one gas
with molecules of
another gas
 Occurs more rapidly in gases than in liquids
 Hardly at all in solids
43
Diffusion
 In Gases
 Molecules travel long
distances between
collisions
 Diffusion rapid
 In Liquids
 Molecules closer
 Encounter more
collisions
 Takes a long time to
move from place to
place
 In Solids
 Diffusion close to zero at
room temperature
 Will increase at high
44
temperature
Surface Tension
Why does H2O bead up on a freshly waxed car
instead of forming a layer?
 Inside body of liquid
 Intermolecular forces are
the same in all directions
 Molecules at surface
 Potential energy increases
when removing neighbors
 Molecules move together
to reduce surface area and
potential energy
45
Surface Tension
 Causes a liquid to take
the shape (a sphere) that
minimizes its surface
area
 Molecules at surface
have higher potential
energy than those in bulk
of liquid and move to
reduce the potential
energy
 Wax = nonpolar
 H2O = polar
 Water beads in order
to reduce potential
energy by reducing
surface area
46
Surface Tension
 Liquids containing molecules
with strong intermolecular forces
have high surface tension
 Allows us to fill glass above
rim



Gives surface rounded
appearance
Surface acts as “skin” that lets
water pile up
Surface resists expansion and
pushes back
 Surface tension
increases as
intermolecular
forces increase
 Surface tension
decreases as
temperature
increases
47
Wetting
 Ability of liquid to spread
across surface to form
thin film
 Greater similarity in
attractive forces between
liquid and surface, yields
greater wetting effect
 Occurs only if
intermolecular attractive
force between surface
and liquid about as
strong as within liquid
itself
48
Wetting
Ex. H2O wets clean glass surface as it forms
H–bonds to SiO2 surface
 Does not wet greasy glass, because grease is
nonpolar and water is very polar


Only London forces
Forms beads instead
Surfactants
 Added to detergents to lower surface tension of H2O
 Now water can spread out on greasy glass
49
Surfactants (Detergents)
 Substances that have both polar and non-polar
characteristics
 Long chain hydrocarbons with polar tail
O
O Na+
O
O
S
O
O Na+
 Nonpolar end dissolves in nonpolar grease
 Polar end dissolves in polar H2O
 Thus increasing solubility of grease in water
50
Viscosity
 Resistance to flow
 Measure of fluid’s
resistance to flow or
changing form
 Related to
intermolecular
attractive forces
www.chemistryexplained.com
 Also called internal friction
 Depends on intermolecular attractions
51
Viscosity
 Viscosity decreases when temperature increases
 Most people associate liquids with viscosity
 Syrup more viscous than water
 Gases have viscosity
 Respond almost instantly to form-changing forces
 Solids, such as rocks and glass have viscosity
 Normally respond very slowly to forces acting to
change their shape
52
Effect of Intermolecular Forces on Viscosity
Acetone
 Polar molecule
Ethylene glycol
 Polar molecule
 Dipole-dipole and
 Hydrogen-bonding
 London forces
 Dipole-dipole and
Which is more viscous?
 London forces
53
Your Turn
For each pair given, which is has more viscosity?
Pair 1. CH3CH2CH2CH2OH,
Pair 2. C6H14,
Pair 3. NH3(l ),
A. CH3CH2CH2CH2OH
B. CH3CH2CH2CH2OH
C. CH3CH2CH2CHO
D. CH3CH2CH2CHO
E. CH3CH2CH2CH2OH
CH3CH2CH2CHO
C12H26
PH3(l )
C6H14
C12H26
C6H14
C12H26
C12H26
NH3(l )
NH3(l )
PH3(l )
NH3(l )
PH3(l )
54
Solubility
 “Like dissolves like”
 To dissolve polar substance, use polar solvent
 To dissolve nonpolar substance, use nonpolar solvent
 Compare relative polarity
 Similar polarity means greater ability to dissolve in each
other
 Differing polarity means that they don’t dissolve, they are
insoluble
 Surfactants
 Both polar and non-polar characteristics
 Used to increase solubility
55
Your Turn
Which of the following are not expected to be soluble in
water?
A. HF
B. CH4
C. CH3OH
D. All are soluble
56
Phase Changes
 Changes of physical state
 Deal with motion of molecules
 As temperature changes
 Matter will undergo phase changes
 Liquid  Gas
 Evaporation, vaporization
 As heat is added, H2O, forms steam or water
vapor
 Requires energy or source of heat
57
Phase Changes
 Solid  Gas
 Sublimation
 Ice cubes in freezer, leave in long enough disappear
 Endothermic
 Gas  Liquid
 Condensation
 Dew is H2O vapor condensing onto cooler ground
 Exothermic
 Often limits lower night time temperature
58
Phase Changes = changes of physical state with
temperature ( α to KE)
fusion
SOLID
evaporation
LIQUID
freezing
GAS
condensation
deposition
sublimation
endothermic
exothermic
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
59
Phase Changes of Water
Fusion/melting
ICE
evaporation
WATER
freezing
VAPOR
Condensation/
forming dew
deposition
sublimation
endothermic
exothermic
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
60
Rate of Evaporation
 Depends on
 Temperature
 Surface area
 Strength of
intermolecular
attractions
 Molecules that escape
from liquid have larger
than minimum escape
KE
 When they leave
 Average KE of
remaining molecules
is less and so T
lower
61
Effect of Temperature on Evaporation Rate
 For given liquid
 Rate of evaporation per
unit surface area
increases as T
increases
 Why?
 At higher T, total
fraction of molecules
with KE large enough to
escape is larger
 Result: rate of
evaporation is larger
62
Kinetic Energy Distribution in Two Different Liquids
A
B
 Smaller intermolecular
 Larger intermolecular
forces
 Lower KE required to
escape liquid
 A evaporates faster
forces
 Higher KE required to
escape liquid
 B evaporates slower
63
Changes Of State Involve Equilibria
 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
64
Before System Reaches Equilibrium
 Liquid is placed in empty,
closed, container
 Begins to evaporate
 Once in gas phase
 Molecules can condense
by
 Striking surface of liquid and
giving up some kinetic
energy
65
System At Equilibrium
 Rate of evaporation =
rate of condensation
 Occurs in closed
systems where
molecules cannot
escape
66
Similar Equilibria Reached in Melting
Melting Point (mp)
 Solid begins to change
into liquid as heat added
 Dynamic equilibria exists
between solid and liquid
states
 Melting (red arrows) and
freezing (black arrows)
occur at same rate
 As long as no heat added or
removed from equilibrium
mixture
67
Equilibria Reached in Sublimation
At equilibrium
 Molecules sublime from
solid at same rate as
molecules condense
from vapor
68
Phase Changes
Energy of System
Gas
Vaporization
Condensation
Sublimation
Deposition
Liquid
Melting
or Fusion
Freezing
Solid
 Exothermic, releases heat
 Endothermic, absorbs heat
69
Energy Changes Accompanying Phase
Changes
 All phase changes are possible under the right
conditions
 Following sequence is endothermic
heat solid  melt  heat liquid  boil  heat gas
 Following sequence is exothermic
cool gas  condense  cool liquid  freeze  cool
solid
70
Enthalpy Of Phase Changes
Endothermic Phase Changes
1.
2.
Must add heat
Energy entering system (+)
Sublimation: Hsub > 0
Vaporization: Hvap > 0
Melting or Fusion: Hfus > 0
Exothermic Phase Changes
1.
2.
Must give off heat
Energy leaving system (–)
Deposition: H < 0 = –Hsub
Condensation: H < 0 = –Hvap
Freezing: H < 0 = –Hfus
71
Phase Changes
 As T changes, matter undergoes phase changes
 Phase Change
 Transformation from one phase to another
 Liquid-Vapor Equilibrium
 Molecules in liquid




Not in rigid lattice
In constant motion
Denser than gas, so more collisions
Some have enough kinetic energy to escape,
some don’t
72
 At any given T,
 Average kinetic energy
of molecules is constant
 But particles have a
distribution of kinetic
energies
 Certain number of
molecules have enough
KE to escape surface
Fraction of molecules
Liquid-Vapor Equilibrium
Kinetic Energy
 As T increases, average KE increases and number
molecules with enough KE to escape increases
73
Vapor Pressure
 Pressure molecules exert when they evaporate or
escape into gas (vapor) phase
 Pressure of gas when liquid or solid is at
equilibrium with its gas phase
 Increasing temperature increases vapor pressure
because vaporization is endothermic
 liquid + heat of vaporization ↔ gas
Equilibrium Vapor Pressure
 VP once dynamic equilibrium reached
 Usually referred to as simply vapor pressure
74
Measuring Vapor Pressure
To measure pressures inside vessels, a manometer is
used.
75
Vapor Pressure Diagram
 Variation of vapor
pressure with T
 Ether
 Volatile
 High vapor pressure
near RT
 Propylene glycol
 Non-volatile
 Low vapor pressure
near RT
RT = 25 C
76
Effect of Volume on VP
A. Initial V
 Liquid – vapor
equilibrium exists
B. Increase V
 Pressure decreases
 Rate of
condensation
decreases
C. More liquid
evaporates
 New equilibrium
established
77
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
 Clausius-Clapeyron equation
 Measure pressure at various temperatures, then
plot
 H

ln P  

vap
R
1
 C
T

 Two point form of Clausius-Clapeyron equation
 Measure pressure at two temperatures and solve
equation
P1 Hvap
ln

P2
R
1
1
  
T 2 T1 
79
Learning Check
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
 P2
401 mm Hg
2
P2 =
= 6.6 ´10 mm Hg
0.6109
80
Your Turn
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
81
Your Turn - Solution
82
Do Solids Have Vapor Pressures?
 Yes
 At given temperature
 Some solid particles have enough KE to escape into
vapor phase
 When vapor particles collide with surface
 They can be captured
 Equilibrium vapor pressure of solid
 Pressure of vapor in equilibrium with solid
83
Boiling Point (bp)
 T at which vapor pressure of liquid = atmospheric
pressure.
 Bp increases as strength of intermolecular forces
increase
 Normal Boiling Point
 T at which vapor pressure of liquid = 1 atm
84
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
85
Your Turn
Which of the following will affect the boiling point of
a substance?
A. Polarizability
B. Intermolecular attractions
C. The external pressure on the material
D. All of these
E. None of these
86
Heating Curve
 Heat added at constant rate
 Horizontal lines
 Phase changes
 Melting point
 Boiling point
 Diagonal
lines
 Heating of
solid, liquid
or gas1
 Superheating
 Temperature of liquid rises slightly above boiling point
87
Cooling Curve
 Heat removed at constant rate
 Horizontal lines
 Phase changes
 Melting point
 Boiling point
 Diagonal lines
 Cooling of solid,
liquid or gas
 Supercooling
 Temperature of liquid dips below its freezing point
88
Your Turn
How much heat, in J, is required to convert 10.00 g
of ice at -10.00 °C to water at 50.00 °C?
Specific heat (J/g K): ice, 2.108, water, 4.184
Enthalpy of fusion = 6.010 kJ/mol
A. 5483 J
B. 5643 J
C. 2304 J
D. 2364 J
E. 62,400 J
89
Energies of Phase Changes
 Expressed per mole
 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
 All of these quantities tend to increase with
increasing intermolecular forces
90
Le Chatelier’s Principle
 Equilibria are often disturbed or upset
 When dynamic equilibrium of system is upset by
a disturbance
 System responds in direction that tends to
counteract disturbance and, if possible, restore
equilibrium
 Position of equilibrium
 Used to refer to relative amounts of substance
on each side of double (equilibrium) arrows
91
Liquid Vapor Equilibrium
Liquid + Heat  Vapor
 Increasing T
 Increases amount of vapor
 Decreases amount of liquid
 Shifts equilibrium to the right
 More vapor is produced at expense of liquid
 Temperature-pressure relationships can be
represented using a phase diagram
92
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
93
Phase Diagram
E
 X axis – temperature
 Y axis – pressure
 As P increases
(T constant), solid most
likely
 More compact
 As T increases
(P constant), gas most
likely
F
 Higher energy
 Each point = T and P
B=
E=
F=
0.01 °C, 4.58 torr
100 °C, 760 torr
–10 °C, 2.15 torr
94
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
95
Case Study: An Ice Necklace
 A cube of ice may the string into
the ice cube suspended on a string
simply by pressing be. As the
string is pressed onto the surface, it
becomes embedded into the ice.
 Why does this happen?
96
Phase Diagram – CO2
 Now line between
solid and liquid
slants to right
 More typical
 Where is triple
point?
 Where is critical
point?
97
Supercritical Fluid
 Substance with temperature above its critical
temperature (TC) and density near its liquid
density
 Have unique properties that make them
excellent solvents
 Values of TC tend to increase with increased
intermolecular attractions between particles
98
Your Turn
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
99
Types of Solids
 Crystalline Solids
 Solids with highly regular arrangements of
components
 Amorphous Solids
 Solids with considerable disorder in their
structures
100
Crystalline Solids
 Unit Cell
 Smallest segment
that repeats
regularly
 Smallest repeating
unit of lattice
 Two-dimensional
unit cells
101
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

102
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
center
 Edge length
a=
4r
3
 Face-centered cubic (FCC)
 Has one atom centered in each face, and
one at each corner
 Edge length
a = 4r / 2
103
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
104
Two Ways to Put on Third Layer
Cubic lattice: 3-dimensional arrays
1. Directly above
spheres in first
layer
2. Above holes in first
layer
 Remaining holes not
covered by second layer
105
3-D Simple Cubic Lattice
Unit Cell
Portion of lattice—
open view
Space filling
model
106
Other Cubic Lattices
Face Centered
Cubic
Body Centered
Cubic
107
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
108
Common Ionic Solids
Rock salt or NaCl
 Face centered cubic lattice of Cl– ions (green)
 Na+ ions (blue) in all octahedral holes
109
Other Common Ionic Solids
Cesium
Chloride, CsCl
Zinc Sulfide,
ZnS
Calcium
Fluoride, CaF2
110
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
111
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
112
Example: NaCl
Face
Edge
Corner
Site
Center
# of
6  12   3
+
Na
#
of Cl–
Body
1
0
Face
0
Edge
12  1 4   3
Corner
0
8  1 8  1
Total
4
4
0
113
Learning Check:
Determine the number of each type of ion in
the unit cell.
114
Some Factors Affecting Crystalline
Structure
 Size of atoms or ions involved
 Stoichiometry of salt
 Materials involved
 Some substances do not form crystalline solids
115
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
116
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
117
X-ray Diffraction
Experimental Setup
Diffraction Pattern
118
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
119
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
120
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.
121
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.
122
Ex. Using Diffraction Data (cont.)
123
Learning Check
Silver packs together in a faced center cubic fashion.
The interplanar distance, d, corresponds to the length of
a side of the unit cell, and is 407 pm. What is the radius
of a silver atom?
a
124
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
125
Sample Homework 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.
126
Your Turn
Yitterbium crystallizes with a face centered cubic
lattice. The atomic radius of yitterbium is 175 pm.
Determine the unit cell length.
A. 495 pm
B. 700 pm
C. 350 pm
D. 990 pm
E. 247 pm
127
Your Turn - Solution
128
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
129
Ex. Covalent (Network) Solid
 Diamond (all C)
 Shown
 SiO2 silicon oxide
 Alternating Si and O
 Basis of glass and quartz
 Silicon carbide (SiC)
130
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
131
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
132
Learning Check:
Classify the following in terms of most likely type of solid.
Substance
ionic molecular covalent metallic
X: Pulverizes when struck;
non-conductive of heat
and electricity
Y: White crystalline solid
that conducts electrical
current when molten or
dissolved
Z: Shiny, conductive,
malleable with high
melting temperature
133
Your Turn
Molecular crystals can contain all of the listed
attraction forces except:
A. Dipole-dipole attractions
B. Electrostatic forces
C. London forces
D. Hydrogen bonding
134