Intermolecular Forces: Liquids, Solids and Phase Changes

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Chapter 12
Intermolecular Forces Liquids, Solids and Phase
Changes
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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
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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.
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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
Importance of interplay between potential and
kinetic energies
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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
(Hesss Law) solid gas
gas solid (called deposition)
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Heats of vaporization and fusion for some common
substances
It takes more energy to vaporize than to melt!
Figure 12.1
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Phase changes and their associated enthalpy
changes
Figure 12.2
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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)
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A cooling curve for the conversion of gaseous
water to ice
Five stages, two phase changes
Figure 12.3
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Phase changes are reversible and reach
equilibrium.
A. Liquid-Gas Equilibria
Pressure at equilibrium vapor pressure
At the BP, the rate of evaporation equals the
rate of condensation!
Figure 12.4
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The effect of temperature on the distribution of
molecular speed in a liquid
higher temperature higher vapor pressure
Figure 12.5
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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
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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
Clausius- Clapeyron equation)
Figure 12.7
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Full form of the Clausius-Clapeyron equation
Two-point version of the equation
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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.
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SAMPLE PROBLEM 12.1
Using the Clausius-Clapeyron equation
SOLUTION
T1 34.9 oC 308.0 K

T2 350 K 77 oC
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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!
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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).
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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
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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)
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Phase diagrams for CO2 and H2O
CO2
H2O
(ice is less dense than liquid water)
Figure 12.9
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Video Phase Diagrams
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Defining/Quantifying Intermolecular Forces
Intramolecular (bonding) forces strong,
involve larger charges closer together
Intermolecular forces weak, involve smaller
charges farther apart
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Covalent and van der Waals radii
solid Cl2
Figure 12.10
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Periodic trends in covalent and van der Waals
radii (in pm)
blue covalent radius black van der Waals
radius
Figure 12.11
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Types of intermolecular (van der Waals) forces
ion-dipole hydrogen bonding dipole-dipole ion-indu
ced dipole dipole-induced dipole dispersion
(London)
decreasing strength
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Importance of polarizability!
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Orientation of polar molecules caused by
dipole-dipole forces
More orderly in the solid phase than in the
liquid phase
Figure 12.12
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Dipole moment and boiling point
Higher dipole moment translates into higher BP.
Figure 12.13
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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)
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SAMPLE PROBLEM 12.2
Drawing hydrogen bonds between molecules of a
substance
Find molecules in which hydrogen is bonded to N,
O or F. Draw H-bonds in the format, B ----- HA.
PLAN
SOLUTION
(a) C2H6 has no H-bonding sites (a non-polar
molecule).
(c)
Note more than one H-bond per molecule is
possible!
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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
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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.
strength in numbers!
Figure 12.15
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Polarizability
The ease with which a particles 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.
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Dispersion forces among non-polar molecules
separated Cl2 molecules
instantaneous dipoles
Caused by momentary oscillations of
electron charge
More electrons, larger molecule, greater mass,
greater dispersion forces
Figure 12.16
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Molar mass and boiling point
Figure 12.17
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Molecular shape and boiling point
Figure 12.18
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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
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.

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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.
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Properties of the Liquid State

• surface tension
• capillarity
• viscosity

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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
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Table 12.3
Surface Tension and Forces Between Particles
surface tension (J/m2) at 20 oC
substance
formula
major force(s)
diethyl ether
dipole-dipole dispersion
CH3CH2OCH2CH3
1.7 x 10-2
ethanol
H-bonding
2.3 x 10-2
CH3CH2OH
butanol
H-bonding dispersion
2.5 x 10-2
CH3CH2CH2CH2OH
water
H-bonding
7.3 x 10-2
H2O
mercury
metallic bonding
48 x 10-2
Hg
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Capillarity the rising of a liquid through a
narrow space against the pull of gravity
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Shape of a water or mercury meniscus in glass
water-glass forces gt water-water forces
Hg-Hg forces gt Hg-glass forces
Figure 12.20
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Viscosity a liquids resistance to flow

• Affected by temperature (viscosity decreases at
  higher T)
• Affected by molecular shape (longer molecules
  exhibit higher viscosity)

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Table 12.4 Viscosity of Water at Several
Temperatures
viscosity (N.s/m2)
temperature (oC)
20
1.00 x 10-3
40
0.65 x 10-3
0.47 x 10-3
60
80
0.35 x 10-3
The units of viscosity are newton-seconds per
square meter.
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Water

• H-bonding ability
• solvent power
• high specific heat capacity
• high heat of vaporization
• high surface tension
• high capillarity
• density of liquid water vs ice

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The H-bonding ability of the water molecule
Four H-bonds per molecule in the solid state
fewer in the liquid state
acceptor
donor
Figure 12.21
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The hexagonal structure of ice
Figure 12.22
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End of Assigned Material
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The macroscopic properties of water and their
atomic and molecular roots
Figure 12.24
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The crystal lattice and the unit cell
Figure 12.26
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The three cubic unit cells
Figure 12.27
simple cubic
Atoms/unit cell 1/8 x 8 1
coordination number 6
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The three cubic unit cells
body-centered cubic
Atoms/unit cell (1/8 x 8) 1 2
Figure 12.27
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The three cubic unit cells
face-centered cubic
Atoms/unit cell (1/8 x 8) (1/2 x 6) 4
Figure 12.27
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Packing of spheres
simple cubic 52 packing efficiency
body-centered cubic 68 packing efficiency
Figure 12.28
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Figure 12.26
closest packing of first and second layers
abab (74)
abcabc (74)
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SAMPLE PROBLEM 12.4
Determining atomic radius from crystal structure
PLAN
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 Avogadros number will give the
volume of one atom of Ba. Using the volume of a
sphere, the radius can be calculated.
radius of a Ba atom
density of Ba (g/cm3)
volume of 1 mol Ba metal
volume of 1 Ba atom
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SAMPLE PROBLEM 12.4
(continued)
SOLUTION
volume of Ba metal
37.9 cm3/mol Ba
37.9 cm3/mol Ba
x 0.68
26 cm3/mol Ba atoms
4.3 x 10-23 cm3/atom
r3 3V/4p
2.2 x 10-8 cm
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Cubic closest packing for frozen argon
Figure 12.29
Figure 12.30
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The sodium chloride structure
Figure 12.31
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The zinc blende structure
Figure 12.32
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The fluorite (CaF2) structure
Figure 12.33
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Crystal structures of metals
cubic closest packing
Figure 12.34
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The random coil shape of a polymer chain
Figure 12.47
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The semi-crystallinity of a polymer chain
Figure 12.48
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The viscosity of a polymer in solution
Figure 12.49
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Manipulating atoms
Figure 12.50
tip of an atomic force microscope (AFM)
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Manipulating atoms
Figure 12.50
nanotube gear
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Figure B12.1
Tools of the Laboratory
Diffraction of x-rays by crystal planes
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Figure B12.2
Tools of the Laboratory
Formation of an x-ray diffraction pattern of the
protein hemoglobin
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Tools of the Laboratory
Figure B12.3
Scanning tunneling micrographs
gallium arsenide semiconductor
metallic gold