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12-1
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Intermolecular Forces:
Liquids, Solids, and Phase Changes
12-2
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Intermolecular Forces:
Liquids, Solids, and Phase Changes
12.1 An Overview of Physical States and Phase Changes
12.2 Quantitative Aspects of Phase Changes
12.3 Types of Intermolecular Forces
12.4 Properties of the Liquid State
12.5 The Uniqueness of Water
12.6 The Solid State: Structure, Properties, and Bonding
12.7 Advanced Materials
12-3
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ATTRACTIVE FORCES
electrostatic in nature
Intramolecular forces bonding forces
These forces exist within each molecule.
They influence the chemical properties of the substance.
Intermolecular forces nonbonding forces
These forces exist between molecules.
They influence the physical properties of the substance.
12-4
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Phase Changes
solid liquid gas
melting
freezing
vaporizing
condensing
sublimination
endothermic
exothermic
12-5
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Table 12.1
A Macroscopic Comparison of Gases, Liquids, and Solids
State Shape and Volume Compressibility Ability to Flow
Gas Conforms to shape and volume
of container
high high
Liquid Conforms to shape of container;
volume limited by surface
very low moderate
Solid Maintains its own shape and
volume
almost none almost none
12-6
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Figure 12.1
Heats of vaporization and fusion for several common substances.
12-7
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Figure 12.2 Phase changes and their enthalpy changes.
12-8
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Figure 12.3
A cooling curve for the conversion of gaseous water to ice.
q = (amount)(molar heat capacity)(T) within a phase
q = (amount)(enthalpy of phase change) during a phase change
12-9
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Within a phase, a change in heat is accompanied by a change in
temperature which is associated with a change in average Ek as
the most probable speed of the molecules changes.
Quantitative Aspects of Phase Changes
During a phase change, a change in heat occurs at a constant
temperature, which is associated with a change in Ep, as the
average distance between molecules changes.
q = (amount)(molar heat capacity)(T)
q = (amount)(enthalpy of phase change)
12-10
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Figure 12.4 Liquid-gas equilibrium.
12-11
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Figure 12.5The effect of temperature on the distribution of
molecular speed in a liquid.
12-12
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Figure 12.6 Figure 12.7
Vapor pressure as a function of
temperature and intermolecular forces.
A linear plot of vapor
pressure- temperature
relationship.
12-13
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The Clausius-Clapeyron Equation
ln P =
-Hvap
R
1
T
C
ln P2P1
= -Hvap
R
1
T2
1
T1
12-14
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SAMPLE PROBLEM 12.1 Using the Clausius-Clapeyron Equation
SOLUTION:
PROBLEM: The vapor pressure of ethanol is 115 torr at 34.90C. If Hvap of
ethanol is 40.5 kJ/mol, calculate the temperature (in 0C) when
the vapor pressure is 760 torr.
PLAN: We are given 4 of the 5 variables in the Clausius-Clapeyron
equation. Substitute and solve for T2.
ln P2P1
= -Hvap
R
1
T2
1
T1
34.90C = 308.0K
ln760 torr
115 torr=
-40.5 x103 J/mol
8.314 J/mol*K
1
T2
1
308K-
T2 = 350K = 770C
12-15
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Figure 12.8 Iodine subliming.
iodine solid
iodine vapor
iodine solid
test tube with ice
12-16
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Figure 12.9 Phase diagrams for CO2 and H2O.
CO2 H2O
12-17
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Types of
Intermolecular forces
12-18
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The nature of the phases and their changes are due primarily to forces among molecules.
Bonding and nonbonding (intermolecular)
The two type of forces differ and Coulomb’s law explains why:
Bonding – strong – larger charges closer together
Intermolecular – weak – smaller charges farther apart.
12-19
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bond length
covalent radius
van der Waal’s distance
van der Waal’s radius
Figure 12.10 Covalent and van der Waals radii.
How far apart are the charges between molecules
that give rise to intermolecular forces?
12-20
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Figure 12.11
Periodic trends in covalent and van der Waals radii (in pm).
The van der Waals radius
of an atom is always larger
than its covalent radius,
But van der Waals radii
decrease across a period
and increase down a group,
just as covalent radii do.
12-21
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12-22
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12-23
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Ion – dipole forcesWhen an ion and a nearby polar molecule (dipole)
attract each other, an ion –dipole force results.
Example: when an ionic compound dissolves in water.
Ep α - | z | µ
r2
The – sign means that the potential energy of the molecule is lowered by its interaction with the polar
solvent.
1/r2 means that the interaction depends more strongly in distance than between two ions.
12-24
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Dipole – dipole forces
When polar molecules lie near each other, as in
liquids and solids, their partial charges act as
tiny electric fields that orient them and give
rise to dipole – dipole forces
12-25
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Figure 12.12 Polar molecules and dipole-dipole forces.
solid
liquid
12-26
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THE HYDROGEN BOND
a dipole-dipole intermolecular force
The elements which are so electronegative are N, O, and F.
A hydrogen bond may occur when an H atom in a molecule,
bound to small highly electronegative atom with lone pairs of
electrons, is attracted to the lone pairs in another molecule.
..
F..
..
..H O..
N.. FH
..
..
..
O.. ..
..
NH
hydrogen bond
donor
hydrogen bond
acceptor
hydrogen bond
acceptor
hydrogen bond
donor
hydrogen bond
donor
hydrogen bond
acceptor
12-27
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Figure 12.13 Dipole moment and boiling point.
For molecular compounds of approximately the same size and molar mass, the
greater the dipole-dipole forces between the molecules are, and so the more
energy it takes to separate them.
12-28
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SAMPLE PROBLEM 12.2 Drawing Hydrogen Bonds Between Molecules
of a Substance
SOLUTION:
PROBLEM: Which of the following substances exhibits H bonding? For
those that do, draw two molecules of the substance with the H
bonds between them.
C2H6(a) CH3OH(b) CH3C NH2
O
(c)
PLAN: Find molecules in which H is bonded to N, O or F. Draw H
bonds in the format -B: H-A-.
(a) C2H6 has no H bonding sites.
(c)(b)C O H
H
H
H
COH
H
H
H
CH3C N
O
H
H
CH3CN
O
H
H
CH3C
N
O
H
H
CH3C
N
O
H
H
12-29
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Figure 12.14 Hydrogen bonding and boiling point.
12-30
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Polarizability and Charged-Induced Dipole Forces
distortion of an electron cloud
•Polarizability increases down a group
size increases and the larger electron clouds are further
from the nucleus
•Polarizability decreases left to right across a period
increasing Zeff shrinks atomic size and holds the electrons
more tightly
•Cations are less polarizable than their parent atom
because they are smaller.
•Anions are more polarizable than their parent atom
because they are larger.
12-31
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Figure 12.15 Dispersion forces among nonpolar molecules.
separated
Cl2molecules
instantaneous
dipoles
12-32
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London or dispersion forces
• Dispersion forces are caused by momentary oscillation of electron charge in atoms and, therefore, are present between all particles (atoms, ions and molecules)
• Instantaneous dipole – induced dipole forces.
• Except in cases involving small polar molecules or those with strong hydrogen bonding, the dispersion force is the dominant force
• Example: HCl - 85% dispersion and 15% dipole-dipole.
12-33
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Figure 12.16
Molar mass and boiling point.
Higher molar mass, higher
polarizability
Intermolecular forces between
the molecules.
12-34
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Figure 12.17 Molecular shape and boiling point.
more points for
dispersion
forces to act
fewer points for
dispersion
forces to act
12-35
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SAMPLE PROBLEM 12.3 Predicting the Type and Relative Strength of
Intermolecular Forces
PROBLEM: For each pair of substances, identify the dominant
intermolecular forces in each substance, and select the
substance with the higher boiling point.
(a) MgCl2 or PCl3
(b) CH3NH2 or CH3F
(c) CH3OH or CH3CH2OH
(d) Hexane (CH3CH2CH2CH2CH2CH3)
or 2,2-dimethylbutaneCH3CCH2CH3
CH3
CH3PLAN: Use the formula, structure and Table 2.2 (button).
•Bonding forces are stronger than nonbonding(intermolecular) forces.
•Hydrogen bonding is a strong type of dipole-dipole force.
•Dispersion forces are decisive when the difference is molar mass or
molecular shape.
12-36
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SOLUTION:
SAMPLE PROBLEM 12.3 Predicting the Type and Relative Strength of
Intermolecular Forcescontinued
(a) Mg2+ and Cl- are held together by ionic bonds while PCl3 is covalently
bonded and the molecules are held together by dipole-dipole interactions. Ionic
bonds are stronger than dipole interactions and so MgCl2 has the higher boiling
point.
(b) CH3NH2 and CH3F are both covalent compounds and have bonds which are
polar. The dipole in CH3NH2 can H bond while that in CH3F cannot. Therefore
CH3NH2 has the stronger interactions and the higher boiling point.
(c) Both CH3OH and CH3CH2OH can H bond but CH3CH2OH has more CH for
more dispersion force interaction. Therefore CH3CH2OH has the higher boiling
point.
(d) Hexane and 2,2-dimethylbutane are both nonpolar with only dispersion
forces to hold the molecules together. Hexane has the larger surface area,
thereby the greater dispersion forces and the higher boiling point.
12-37
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Figure 12.18
Summary diagram for analyzing the intermolecular forces in a sample.
INTERACTING PARTICLES
(atoms, molecules, ions)
ions only
IONIC BONDING
(Section 9.2)
ion + polar molecule
ION-DIPOLE FORCES
ions present ions not present
polar molecules only
DIPOLE-DIPOLE
FORCES
HYDROGEN
BONDING
polar + nonpolar
molecules
DIPOLE-
INDUCED DIPOLE
FORCES
nonpolar
molecules only
DISPERSION
FORCES only
DISPERSION FORCES ALSO PRESENT
H bonded toN, O, or F
12-38
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OTHER PROPERTIES OF
LIQUIDS
12-39
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Figure 12.19 The molecular basis of surface tension.
hydrogen bonding
occurs in three
dimensions
hydrogen bonding
occurs across the surface
and below the surfacethe net vector
for attractive
forces is downward
12-40
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Surface tension
• Is the energy required to increase the surface area by a unit amount – the stronger the forces are between particles in a liquid, the greater the surface tension.
• Tendency to minimize surface area (spherical drops)
• The relatively high surface tension of water accounts for the ease with which it can be nebulized, or placed into aerosol form. Low surface tension liquids tend to evaporate quickly and are difficult to keep in an aerosol form. All liquids display surface tension to some degree.
12-41
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Table 12.3 Surface Tension and Forces Between Particles
Substance Formula
Surface Tension
(J/m2) at 200C Major Force(s)
diethyl ether
ethanol
butanol
water
mercury
dipole-dipole; dispersion
H bonding
H bonding; dispersion
H bonding
metallic bonding
1.7x10-2
2.3x10-2
2.5x10-2
7.3x10-2
48x10-2
CH3CH2OCH2CH3
CH3CH2OH
CH3CH2CH2CH2OH
H2O
Hg
12-42
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Figure 12.20 Shape of water or mercury meniscus in glass.
adhesive forces
stronger
cohesive forces
H2O
capillarity
Hg
concaveconvex
12-43
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Capillarity
• The rising of a liquid through a narrow space
against the pull of gravity is called capillary
action, or capillarity.
• It is the result from a competition between
intermolecular forces within the liquid
(cohesive forces) and those between the liquid
and the tube walls (adhesive forces)
12-44
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Viscosity
• Its resistance to flow – intermolecular
attractions impede this movement.
• Gases and liquids flow, but liquid viscosities
are much higher because of intermolecular
forces.
• Viscosity decreases with heating
• Molecular shape – small, spherical molecules
make little contact and pour easily.
12-45
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Table 12.4 Viscosity of Water at Several Temperatures
Temperature(0C)
Viscosity
(N*s/m2)*
20
40
60
80
1.00x10-3
0.65x10-3
0.47x10-3
0.35x10-3
*The units of viscosity are newton-seconds per square meter.
viscosity - resistance to flow
12-46
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Figure 12.21 The H-bonding ability of the water molecule.
hydrogen bond donor
hydrogen bond acceptor
12-47
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The Unique Nature of Water
•great solvent properties due to polarity and
hydrogen bonding ability
•exceptional high specific heat capacity
•high surface tension and capillarity
•density differences of liquid and solid states
12-48
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Figure 12.22 The hexagonal structure of ice.
12-49
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Figure 12.23 The expansion and contraction of water.
12-50
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Figure 12.24The macroscopic properties of water and their atomic
and molecular “roots”.
12-51
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Solids:
- crystalline
- amorphous
12-52
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Figure 12.25 The striking beauty of crystalline solids.
celestite pyrite amethyst halite
12-53
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Structures of solids
• Crystalline solids have a regularly repeating
pattern of atoms.
• We will consider only cubic unit cells, with
equal edges meeting at right angles (90°) when
we calculate in detail the packing efficiency.
• There are 7 crystal systems and 14 types of
unit cells that occur in nature.
12-54
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Cubic system
• Majority of metallic elements
• Some covalent compounds
• Several ionic compounds
There are 3 types of cubic unit cells within the cubic system:
1. Simple cubic unit cell
2. Body-centered cubic unit cell
3. Face-centered cubic unit cell
12-55
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portion of a 3-D lattice
Figure 12.26 The crystal lattice and the unit cell.lattice point
unit
cell
portion of a 2-D lattice
unit
cell
Lattice - Consists of all points with identical surroundings
Unit cell – The smallest portion of the crystal that, if repeated in all
three directions, gives the crystal.
12-56
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Figure 12.27 (1 of 3) The three cubic unit cells.
Simple Cubic
coordination number = 6
Atoms/unit cell = 1/8 * 8 = 1
1/8 atom at
8 corners
Coordination number is the number of nearest neighbors surrounding it.
12-57
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Figure 12.27 (2 of 3) The three cubic unit cells.
Body-centered
Cubic
coordination number = 8
1/8 atom at
8 corners
1 atom at
center
Atoms/unit cell = (1/8*8) + 1 = 2
12-58
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Figure 12.27 (3 of 3) The three cubic unit cells.
Face-centered
Cubic
coordination number = 12 Atoms/unit cell = (1/8*8)+(1/2*6) = 4
1/8 atom at
8 corners
1/2 atom at
6 faces
12-59
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Figure 12.28 Packing of spheres.
simple cubic
(52% packing efficiency)
body-centered cubic
(68% packing efficiency)
12-60
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hexagonal
unit cell
Figure 12.28 (continued)
closest packing of first
and second layers
layer a
layer a – shifting
every other row a
layer b
layer c
hexagonal
closest
packingcubic closest
packing
abab… (74%)abcabc… (74%)
expanded
side views
face-centered
unit cell
covers
the white
space.
Hexagonal( Mg, Zn) and FCC (Ni, Cu) is the most efficient packing for identical spheres
12-61
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Packing efficiency %
= (volume occupied by spheres) / (total volume) x 100
This calculation is carried out on the contents of the
unit cell. Therefore, our first step is to count the
number of atoms contained within the unit cell. To
do this you should keep in mind that atoms located
at the corners, edges and faces of the unit cell are
not fully contained in a single unit cell, so that they
only partially occupy the unit cell.
12-62
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• Atom Location Fraction Inside Unit Cell*
Corner 1/8
Edge 1/4
Face 1/2
Anywhere else 1
*For unit cells with non-orthogonal axes it is not
strictly true that an atom on a corner (edge) is
exactly 1/8 (1/4) inside the unit cell, but it is still
true that 8 (4) such atoms add up to one atom in the
unit cell
12-63
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• Now we can determine the number of atoms
inside the unit cell for each of the packing
arrangements we have discussed.
Face Centered Cubic (ccp) (8 ×1/8) + (6 ×1/2) = 4
Body Centered Cubic (bcc) (8 ×1/8) + 1 = 2
Simple Cubic (sc) (8 ×1/8) = 1
Hexagonal closed packed (hcp) = 6
12-64
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• The volume of occupied space contained
within the unit cell can now be easily
calculated by multiplying the number of atoms
per unit cell by the volume of a sphere.
V = 4/3 π r3 x ?
12-65
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• The next step is to determine the size
of the unit cell as a function of the
atomic radius, r:
SC, BCC, FCC
12-66
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a = b = c = 2r
a = 2r
V = (2r)3
V = 8r3
Packing efficiency % = 4/3 π r3 x 1 X 100 = 52 %
8r3
1/8 atom at
8 corners
Atoms/unit cell = 1/8 * 8 = 1
The volume of a sphere is
V = 4/3 π r3
Simple cubic
12-67
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The volume of a sphere is
V = 4/3 π r3
Atoms/unit cell = (1/8*8) + 1 = 2
Packing efficiency % = 4/3 π r3 x 2 X 100 = 68 %
(2.3r)3
Body-centered cubic
a2 + a2 = ?
a2 + ? = (4r)2
a2 + 2a2 = (4r)2
3a2 = 16r2
a = √(16/3)r
12-68
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a2 + a2 = (4r)2
2a2 = 16r2
a = 81/2 r
V = (81/2 r)3
V = 83/2r3
The volume of a sphere is
V = 4/3 π r3
Packing efficiency % = 4/3 π r3 x 4 X 100 = 74 %
83/2r3
Atoms/unit cell = (1/8*8)+(1/2*6) = 4
Face-centered cubic
12-69
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Density
• d = m/V
• m = M/NAx #of spheres per unit
• V = unit cell (sc, bcc, fcc)
12-70
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SAMPLE PROBLEM 12.4 Determining Atomic Radius from Crystal Structure
PROBLEM: Barium is the largest nonradioactive alkaline earth metal. It has
a body-centered cubic unit cell and a density of 3.62 g/cm3.
What is the atomic radius of barium?
(Volume of a sphere: V = 4/3pr3)
PLAN: We can use the density and molar mass to find the volume of 1 mol of
Ba. Since 68%(for a body-centered cubic) of the unit cell contains
atomic material, dividing by Avogadro’s number will give us the volume
of one atom of Ba. Using the volume of a sphere, the radius can be
calculated.
density of Ba (g/cm3)
volume of 1 mol Ba metal volume of 1 Ba atom
radius of a Ba atom
multiply by packing efficiency
reciprocal divided by M V = 4/3pr3
volume of 1 mol Ba atoms
divide by Avogadro’s number
12-71
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SAMPLE PROBLEM 12.4 Determining Atomic Radius from Crystal Structure
SOLUTION:
continued
Volume of Ba metal =137.3 g Ba
mol Ba= 37.9 cm3/mol Ba
37.9 cm3/mol Ba x 0.68 = 26 cm3/mol Ba atoms
mol Ba atoms
6.022x1023 atoms= 4.3x10-23 cm3/atom
r3 = 3V/4p = 2.2 x 10-8cm
1 cm3
3.62 gx
26 cm3
mol Ba atomsx
r =3V
4p3
3(4.3x10-23cm3 )
4 x 3.143
12-72
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Cu d = 8.93 g/cm3 atomic radius 128pm
M = 63.55 g/mol
Is the metal close-packed or body centered cubic?
Strategy :
• Calculate the density of the metal assuming first that is ccp and then that it is bcc.
• The structure with the density closer to the experimental value is more likely to be the actual structure.
• The mass of a unit cell is the sum of the masses of the atoms that it contains.
• The mass of each atom is given by M/NA
• Length of FCC (a = 81/2r)
• Length of BCC (a = √(16/3) r )
12-73
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Cu d = 8.93 g/cm3
Is the metal close-packed or body centered cubic?
• FCC d = m/a3 m = 4 x M/NA
(81/2r)3
• BCC d = m/a3 m = 2 x M/NA
(√(16/3) r )3
Are these metals BCC or FCC ?
Fe d = 7.87 g/cm3 atomic radius 124 pm
Ag d = 10.5 g/cm3 atomic radius 144 pm
12-74
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Table 12.5 Characteristics of the Major Types of Crystalline Solids
Particles
Interparticle
Forces
Physical
Behavior Examples (mp,0C)
Atomic
Molecular
Ionic
Metallic
Network
Group 8A(18)
[Ne-249 to Rn-71]
Molecules
Positive &
negative ions
Atoms
Atoms
Soft, very low mp, poor
thermal & electrical
conductors
DispersionAtoms
Dispersion,
dipole-dipole,
H bonds
Fairly soft, low to moderate
mp, poor thermal &
electrical conductors
Nonpolar - O2[-219],
C4H10[-138], Cl2
[-101], C6H14[-95]
Polar - SO2[-73],
CHCl3[-64], HNO3[-
42], H2O[0.0]
Covalent bond
Metallic bond
Ion-ion
attraction
Very hard, very high mp,
usually poor thermal and
electrical conductors
Soft to hard, low to very
high mp, excellent thermal
and electrical conductors,
malleable and ductile
Hard & brittle, high mp,
good thermal & electrical
conductors when molten
NaCl [801]
CaF2 [1423]
MgO [2852]
Na [97.8]
Zn [420]
Fe [1535]
12-75
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Figure 12.29 Figure 12.30
Cubic closest packing for
frozen argon.
Atomic solids
Individual atoms held together by
dispersion forces. Noble gases
Cubic closest packing of
frozen methane.
Molecular solids
The lattice points are occupied
by individual molecules
12-76
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The sodium chloride structure.
expanded view
space-filling
Ionic solids
• The unit cell contains particles with whole charge and has the same
cation:anion ratio as the empirical formula.
• The ionic bonding is stronger than intermolecular forces in atomic or molecular
solids.
• Many ionic solids use the cubic closest packing.
• Coordination number is 6 - # of ions of opposite charge surrounding a specific
ion.
1:1 ratio(8 x 1/8) + (6 x 1/2) = 4 Cl-
(12 x 1/4) + 1 (center) = 4 Na+
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Figure 12.32 The zinc blende structure.
zinc sulfide1:1 ratio
(8 x 1/8) + (6 x 1/2) = 4 S2-
4 Zn2+ tetrahedrally surrounded by four
ions of opposite charge.
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Figure 12.33 The fluorite (CaF2) structure.
1:2 ratio(8 x 1/8) + (6 x 1/2) = 4 Ca2+
F- occupy all 8 available holes.
Antifluorite is often see in a 2:1 ratio.
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Coordination number in an Ionic solid
• (6,6) – coordination -NaCl, KBr, RbI, MgO
In this notation the first number is the cation and the
second is the anion.
0.4 to 0.7 a structure similar to NaCl is expected
> 0.7 more anions can fit around the cation where
the two ions are almost the same size.
CsCl (8,8) coordination
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Figure 12.34 Crystal structures of metals.
cubic closest packing hexagonal closest packing
Metallic solids – Metallic bonding forces hold individual atoms together.
Wide range of melting point and hardnesses related to packing
efficiency and the number of valence e- available for bonding.
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Network covalent solids – where strong covalent bonds link the atoms together
throughout a network.
High melting point and boiling point.
Conductivity and hardness depends on the bonding
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Figure 12.35 Crystalline and amorphous silicon dioxide.
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Figure 12.36 The band of molecular orbitals in lithium metal.
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MOLECULAR ORBITAL BAND THEORY
The electrons in a substance must move within the conduction band in order to conduct electricity.
• The conduction band is a partially filled band or a band of vacant energy levels just higher in energy than a filled band.
• Metals conduct electricity because the highest energy electrons can easily move in a conduction band.
• For a metal, the electrical conductivity decreases as the temperature increases. The metal ions move more, which slows down the flow of electrons when an electric field is applied.
• A typical metalloid is a semiconductor where there is a gap between the highest energy electrons and the conduction band and they do not conduct electricity at low temperatures. A small increase in temperature will excite some of the highest energy electrons into the empty conduction band. Therefore, electrical conductivity increases with increasing temperature for a metalloid.
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Figure 12.37
Electrical conductivity in a conductor, semiconductor, and insulator.
conductor
semiconductor
insulator
•Metallic luster
•Malleability
•Thermal conductivity
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Figure 12.38
The levitating power of a superconducting oxide.
rare earth magnet
superconducting ceramic disk
liquid nitrogen
Superconductivity –
to conduct with no energy loss
requires extreme cooling to minimize atom movement.
It has been observed by cooling metals to near absolute zero.
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Figure 12.39
Crystal structures and band
representations of doped
semiconductors.
Doping – adding small amounts of other elements to increase or decreae the
number of valence electrons in the bands.
extra negative charges (electrons) are present the empty orbitals act as positive holes
INCREASING CONDUCTIVITY
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Forward bias
Reverse bias
p-n junctionFigure 12.40 The p-n junction.
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heat in furnace with O2
treat with photoresist apply template •expose to light
and solvent
•remove template
•etch SiO2 with HF
• remove photoresist
•treat with Ga vapor
•remove SiO2
Figure 12.41 Steps in manufacturing a p-n junction.
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Figure 12.42
Structures of two typical molecules that form liquid crystals.
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Figure 12.43 The three common types of liquid crystal phases.
nematic smecticcholesteric
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Figure 12.44 Liquid crystals in biological systems.
nematic arrays of
tobacco mosaic virus particles
actin and myosin protein
filaments in voluntary muscle cells
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Figure 12.45
Schematic of a liquid
crystal display (LCD).
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Table 12.7 Some Uses of Modern Ceramics and Ceramic Mixtures
Ceramic Applications
SiC, Si3N4, TiB2, Al2O3 Whiskers(fibers) to strength Al and other ceramics
Si3N4 Car engine parts; turbine rotors for “turbo” cars;
electronic sensor units
Si3N4, BN, Al2O3 Supports or layering materials(as insulators) in
electronic microchips
SiC, Si3N4, TiB2, ZrO2,
Al2O3, BN
ZrO2, Al2O3
Cutting tools, edge sharpeners(as coatings and
whole devices), scissors, surgical tools, industrial
“diamond”
BN, SiC Armor-plating reinforcement fibers(as in Kevlar
composites)
Surgical implants(hip and knee joints)
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Figure 12.46 Unit cells of some modern ceramic materials.
SiC
silicon
carbide
BN
cubic boron
nitride
(borazon)
YBa2Cu3O7
high
temperature
superconductor
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Table 12.8 Molar Masses of Some Common Polymers
Name Mpolymer (g/mol) n Uses
Acrylates 2 x105 2 x103 Rugs, carpets
Polyamide(nylons) 1.5 x104 1.2 x102 Tires, fishing line
Polycarbonate 1 x105 4 x102 Compact discs
Polyethylene 3 x105 1 x104 Grocery bags
Polyethylene (ultra-
high molecular weight)
5 x106 2 x105 Hip joints
Poly(ethylene
terephthalate)
2 x104 1 x102 Soda bottles
Polystyrene 3 x105 3 x103 Packing; coffee cups
Poly(vinyl chloride) 1 x105 1.5 x103 Plumbing
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Figure 12.47 The random coil shape of a polymer chain.
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Figure 12.48 The semicrystallinity of a polymer chain.
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Figure 12.49 The viscosity of a polymer in solution.
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Table 12.9 Some Common Elastomers
Name Tg (0C)*
*Glass transition temperature
Uses
Poly(dimethyl siloxane) -123
-106
-65
-43
Polybutadiene
Polyisoprene
Polychloroprene (neoprene)
Breast implants
Rubber bands
Surgical gloves
Footwear, medical tubing
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Figure 12.50 Manipulating atoms.
tip of an atomic force microscope (AFM)
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Figure 12.50 Manipulating atoms.
nanotube gear
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Figure B12.1
Diffraction of x-rays by crystal planes.
Tools of the Laboratory
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Figure B12.2
Formation of an x-ray diffraction pattern of the
protein hemoglobin.
Tools of the Laboratory
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Tools of the Laboratory
Figure B12.3 Scanning tunneling micrographs.
gallium arsenide semiconductor metallic gold
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Other properties of liquids
• Surface tension- energy required to disrupt a
drop and spread it out as a film- tendency to
minimize surface area (spherical drops)
• Capillary action- creeping up a surface
cohesive vs. adhesive
• Viscosity: resistance to flow
depends on intermolecular forces
less inter.forces high viscosity.
