CHEM 162 Chp 12 for class
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Transcript CHEM 162 Chp 12 for class
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