Transcript Document 7145972

Chapter 12
Intermolecular Forces: Liquids, Solids and
Phase Changes
12-1
Intermolecular Forces:
Liquids, Solids and Phase Changes
12.1 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
12-2
Types of Molecular Forces
Intramolecular: bonding forces within a molecule; influence
chemical properties
Intermolecular: forces between molecules; influence
physical properties
Three states of water: water vapor, liquid water, ice
Phase change: liquid water
ice
Liquid water and ice are examples of condensed phases.
12-3
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
Importance of interplay between potential and kinetic energies
12-4
Types of Phase Changes
condensation/vaporization
freezing/melting (or fusion)
Enthalpy changes accompany phase changes!
condensation and freezing: exothermic processes
vaporization and melting: endothermic processes
DHofus = heat of fusion (+)
DHovap = heat of vaporization (+)
DHosubl = heat of sublimation = DHofus + DHovap (Hess’s Law)
solid
gas
gas
12-5
solid (called deposition)
Heats of vaporization and fusion for some common substances
Figure 12.1
12-6
It takes more energy to vaporize than to melt!
Phase changes and their associated enthalpy changes
Figure 12.2
12-7
Quantitative Aspects of Phase Changes
Within a phase, a change in heat is accompanied by a change in
temperature which is associated with a change in the average Ek
as the most probable speed of the molecules changes.
q = (amount in moles)(molar heat capacity)(DT)
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 in moles)(enthalpy of phase change)
12-8
A cooling curve for the conversion of gaseous water to ice
Figure 12.3
12-9
Five stages, two phase changes
Phase changes are reversible and reach equilibrium.
A. Liquid-Gas Equilibria
Pressure at equilibrium =
vapor pressure
Figure 12.4
12-10
At the BP, the rate of evaporation
equals the rate of condensation!
The effect of temperature on the distribution of
molecular speed in a liquid
Figure 12.5
12-11
higher temperature = higher vapor pressure
Vapor pressure as a function of
temperature and intermolecular
forces
Weaker intermolecular forces
translate into higher vapor
pressure at a given temperature
Figure 12.6
12-12
A linear plot of the vapor
pressure-temperature
relationship
Plotting ln P against
1/T yields a straight
line with slope equal to
-DHvap/R (the ClausiusClapeyron equation)
Figure 12.7
12-13
Full form of the Clausius-Clapeyron equation:
ln P =
-DHvap 
1 
   C
R
T 
Two-point version of the equation:
P2
-DHvap  1
1 
ln
=
  
R
P1
T2 T1
12-14
Utility of the Clausius-Clapeyron Equation
It provides a means to determine experimentally the heat of vaporization,
which is the energy required to vaporize 1 mole of molecules in the liquid state.
or
If DHvap is known, and vapor pressure at one T is known, then
vapor pressure at a new T can be calculated.
12-15
SAMPLE PROBLEM 12.1
PROBLEM:
PLAN:
Using the Clausius-Clapeyron equation
The vapor pressure of ethanol is 115 torr at 34.9 oC. If DHvap of
ethanol is 40.5 kJ/mol, calculate the temperature (in oC) when
the vapor pressure is 760 torr.
We are given four of the five variables in the Clausius-Clapeyron
equation. Substitute and solve for T2.
SOLUTION:
P2
-DHvap  1
1 
ln
=
  
P1
R T2 T1
760 torr
ln
115 torr
=
-40.5 x103 J/mol
8.314 J/mol.K
T2 = 350 K = 77 oC
12-16
T1 = 34.9 oC = 308.0 K
1
1
308 K
T2
Vapor Pressure and Boiling Point
If we assume an open container, then the boiling point (BP) is the temperature
at which the vapor pressure equals the external pressure (usually
atmospheric pressure, 760 mmHg).
Thus, the BP depends on
the applied pressure
(see Figure 12.6)
Water boils at 100 oC at sea level,
but at 72 oC on the peak of Mt.
Everest!
12-17
B. Liquid-Solid Equilibria
Characterized by a melting point (temperature at which the
rate of melting equals the rate of freezing)
The MP is not significantly affected by pressure (two condensed
phases are involved).
12-18
Iodine subliming
I2 vapor in contact
with a cold finger at
atmospheric pressure
C. Solid-Gas Equilibria
Why??
External pressure and
intermolecular forces maintain the
liquid phase after melting. These
are too weak in some cases.
Figure 12.8
12-19
QuickTime™ and a
Photo - JPEG decompressor
are needed to see this picture.
Bringing It All Together: A Phase Diagram
A graph that describes phase changes of a substance under
various combinations of temperature and pressure
Key Characteristics
Regions (bounded areas) (one phase)
Interfaces (lines) between different regions (equilibria between two phases)
Isolated Points (critical point, triple point) (unique T/P combinations)
12-20
Phase diagrams for CO2 and H2O
H2O
CO2
Figure 12.9
12-21
(ice is less dense than
liquid water)
Video: Phase Diagrams
12-22
Defining/Quantifying Intermolecular Forces
Intramolecular (bonding) forces: strong,
involve larger charges closer together
Intermolecular forces: weak, involve
smaller charges farther apart
12-23
Covalent and van der Waals radii
solid Cl2
Figure 12.10
12-24
Periodic trends
in covalent and
van der Waals
radii (in pm)
blue: covalent radius
black: van der Waals radius
Figure 12.11
12-25
Types of intermolecular (van der Waals) forces
ion-dipole
hydrogen bonding
dipole-dipole
ion-induced dipole
dipole-induced dipole
dispersion (London)
12-26
decreasing
strength
12-27
Importance of polarizability!
12-28
Orientation of polar molecules caused by
dipole-dipole forces
More orderly in the solid phase than in the liquid phase
Figure 12.12
12-29
Dipole moment and boiling point
Higher dipole moment translates into higher BP.
Figure 12.13
12-30
Hydrogen bonding
Involves molecules that have an H atom bound at a small,
highly electronegative atom with lone electron pairs
N-H O-H H-F
General Model
_ :
B
H_A
electronegative atom bearing
hydrogen (donor)
H-bond
electronegative atom with
lone electron pair (acceptor)
12-31
SAMPLE PROBLEM 12.2
PROBLEM:
Which of the following substances exhibits hydrogen bonding?
For those that do, draw two molecules of the substance with
the H-bonds between them.
O
(a)
PLAN:
SOLUTION:
(b)
Drawing hydrogen bonds between molecules
of a substance
C2H6
(c) CH3C NH2
(b) CH3OH
Find molecules in which hydrogen is bonded to N, O or F.
Draw H-bonds in the format, B: ----- HA.
(a) C2H6 has no H-bonding sites (a non-polar molecule).
H
H C O H
H
H
H O C H
H
H
(c)
H
O
O
H N CH3C
CH3C N H
H
O
CH3C
CH3C
N H
H
Note: more than one H-bond per molecule is possible!
12-32
H N
O
Hydrogen bonding and boiling point
H2O, HF and NH3
exhibit aberrant
behavior due to their
ability to form H-bonds.
binary hydrides of Groups 4-7
Figure 12.14
12-33
Covalent and
hydrogen bonding
in the helical
structure of DNA
A single H-bond is relatively weak,
but the existence of many such bonds
in a molecule can influence molecular
structure significantly.
Figure 12.15
12-34
QuickTime™ and a
Photo - JPEG decompressor
are needed to see this picture.
strength in
numbers!
Polarizability
The ease with which a particle’s electron cloud can be distorted
Pertinent to charge-induced dipole forces
(ion-induced dipole and dipole-induced dipole)
Increases down a group of atoms or ions (size)
Decreases from left to right in a period (effective nuclear charge)
Cations are less polarizable than their parent atoms; anions are
more polarizable than their parent atoms.
12-35
Dispersion
forces among
non-polar
molecules
Caused by momentary
oscillations of electron
charge
More electrons, larger
molecule, greater mass,
greater dispersion forces
Figure 12.16
12-36
separated Cl2
molecules
instantaneous
dipoles
Molar mass and
boiling point
Figure 12.17
12-37
Molecular shape and boiling point
Figure 12.18
12-38
SAMPLE PROBLEM 12.3
PROBLEM:
Predicting the Type and Relative Strength of
Intermolecular Forces
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-dimethylbutane
CH3
CH3CCH2CH3
CH3
PLAN:
• Bonding forces are stronger than nonbonding (intermolecular) forces.
• Hydrogen bonding is a strong dipole-dipole force.
• Dispersion forces are decisive when the difference is molar mass or
molecular shape.
12-39
SAMPLE PROBLEM 12.3
(continued)
SOLUTION:
(a) Mg2+ and Cl- are held together by ionic bonds (a salt) while PCl3 is
covalently bonded and the molecules are held together by dipole-dipole
interactions. Ionic attractions are much stronger than dipole interactions and so
MgCl2 has the higher boiling point.
(b) CH3NH2 and CH3F are both covalent compounds and have polar bonds.
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
bonds for greater dispersion force interactions. Therefore, CH3CH2OH has the
higher boiling point.
(d) Hexane and 2,2-dimethylbutane are both non-polar with only dispersion
forces to hold the molecules together. Hexane has a larger surface area, and
therefore the greater dispersion forces and higher boiling point.
12-40
Properties of the Liquid State



12-41
surface tension
capillarity
viscosity
The molecular basis
of surface tension
The energy required to increase
surface area by a unit amount
Surface molecules experience
a net attraction downward.
Stronger intermolecular
forces translate into
greater surface tension.
Figure 12.19
12-42
Table 12.3
Surface Tension and Forces Between Particles
substance
formula
surface tension
(J/m2) at 20 oC
diethyl ether
CH3CH2OCH2CH3
1.7 x 10-2
dipole-dipole; dispersion
ethanol
CH3CH2OH
2.3 x 10-2
H-bonding
butanol
CH3CH2CH2CH2OH
2.5 x 10-2
H-bonding; dispersion
H2O
7.3 x 10-2
H-bonding
Hg
48 x 10-2
metallic bonding
water
mercury
12-43
major force(s)
Capillarity: the rising of a liquid through a narrow
space against the pull of gravity
12-44
Shape of a water or mercury meniscus in glass
stronger
adhesive
forces
water-glass forces >
water-water forces
Figure 12.20
12-45
stronger
cohesive forces
Hg-Hg forces >
Hg-glass forces
Viscosity: a liquid’s resistance to flow


12-46
Affected by temperature (viscosity decreases at higher T)
Affected by molecular shape (longer molecules exhibit higher viscosity)
Table 12.4
Viscosity of Water at Several Temperatures
temperature (oC)
viscosity
(N.s/m2)*
20
1.00 x 10-3
40
0.65 x 10-3
60
0.47 x 10-3
80
0.35 x 10-3
*The units of viscosity are newton-seconds per square meter.
12-47
Water







12-48
H-bonding ability
solvent power
high specific heat capacity
high heat of vaporization
high surface tension
high capillarity
density of liquid water vs ice
The H-bonding ability of the
water molecule
Four H-bonds per
molecule in the solid
state; fewer in the liquid state
Figure 12.21
12-49
acceptor
donor
The hexagonal structure of ice
Figure 12.22
12-50
End of Assigned Material
12-51
The macroscopic properties of water and their atomic
and molecular “roots”
12-52
Figure 12.24
The crystal lattice and the unit cell
Figure 12.26
12-53
The three cubic unit cells
Figure 12.27
simple cubic
1/8 atom at
8 corners
Atoms/unit cell = 1/8 x 8 = 1
coordination number = 6
12-54
The three cubic unit cells
body-centered
cubic
1/8 atom at
8 corners
1 atom at
center
coordination number = 8
Atoms/unit cell = (1/8 x 8) + 1 = 2
Figure 12.27
12-55
The three cubic unit cells
face-centered
cubic
1/8 atom at
8 corners
1/2 atom at
6 faces
coordination number = 12
12-56
Atoms/unit cell = (1/8 x 8) + (1/2 x 6) = 4
Figure 12.27
Packing of spheres
simple cubic
52% packing efficiency
body-centered cubic
Figure 12.28
12-57
68% packing efficiency
Figure 12.26
layer a
layer b
hexagonal
closest
packing
layer a
cubic closest
packing
layer c
closest packing of first
and second layers
abab… (74%)
hexagonal
unit cell
12-58
expanded
side views
abcabc… (74%)
face-centered
unit cell
SAMPLE PROBLEM 12.4
PROBLEM:
PLAN:
Determining atomic radius from crystal structure
Barium is the largest non-radioactive 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)
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 the volume of one
atom of Ba. Using the volume of a sphere, the radius can be
calculated.
density of Ba (g/cm3)
radius of a Ba atom
V = 4/3pr3
reciprocal divided by M
volume of 1 mol Ba metal
volume of 1 Ba atom
multiply by packing efficiency
volume of 1 mol Ba atoms
12-59
divide by Avogadro’s number
SAMPLE PROBLEM 12.4
(continued)
SOLUTION:
volume of Ba metal =
1 cm3
3.62 g
x
137.3 g Ba
mol Ba
= 37.9 cm3/mol Ba
37.9 cm3/mol Ba x 0.68 = 26 cm3/mol Ba atoms
26 cm3
mol Ba atoms
x
r3 = 3V/4p
12-60
mol Ba atoms
6.022 x
1023
atoms
= 4.3 x 10-23 cm3/atom
-23
3
3V
3(4.3x10
cm
)
r=3
3
= 2.2 x 10-8 cm
4p
4 x 3.14
Cubic closest packing for
frozen argon
Figure 12.29
12-61
Cubic closest packing of
frozen methane
Figure 12.30
12-62
The sodium chloride structure
Figure 12.31
12-63
The zinc blende structure
Figure 12.32
12-64
The fluorite (CaF2) structure
Figure 12.33
12-65
Crystal structures of metals
cubic closest packing
hexagonal closest packing
Figure 12.34
12-66
12-67
12-68
The random coil shape of a polymer chain
12-69
Figure 12.47
The semi-crystallinity of a polymer chain
Figure 12.48
12-70
The viscosity of a polymer in solution
Figure 12.49
12-71
Manipulating atoms
Figure 12.50
tip of an atomic force microscope (AFM)
12-72
Manipulating atoms
Figure 12.50
nanotube gear
12-73
Tools of the Laboratory
Figure B12.1
Diffraction of x-rays by crystal planes
12-74
Tools of the Laboratory
Formation of an x-ray diffraction pattern
of the protein hemoglobin
12-75
Figure B12.2
Figure B12.3
Tools of the Laboratory
Scanning tunneling micrographs
gallium arsenide semiconductor
12-76
metallic gold