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Chapter #16 – Liquids and Solids
16.1) Intermolecular Forces
16.2) The Liquid State
16.3) An Introduction to Structures and Types of Solids
16.4) Structure and Bonding of Metals
16.5) Carbon and Silicon: Network Atomic Solids
16.6) Molecular Solids
16.7) Ionic Solids
16.8) Structures of Actual Ionic Solids
16.9) Lattice Defects
16.10) Vapor Pressure and Changes of State
16.11) Phase Diagrams
16.12) Nanotechnology
Macroscopic Properties of Gases,
Liquids, and Solids “Condensed
States of Matter”
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Solids, liquids and gases
• Intermolecular forces occur between molecules, not
within molecules (as in bonding).
• When a molecule changes phases it remains intact –
the molecule does not change, just the forces among
the molecules around it.
What determines the temperature at
which liquids boil and solids melt?
Intermolecular Forces
Intramolecular forces: chemical bonds holding atoms together
in a molecule
ionic bonds, e.g. Na+Cl-
~400 - 4000 kJ/mol
covalent bonds, e.g. H2, CH4, etc.
~150 - 1200 kJ/mol
(the O-H bond energy in H2O is 500 kJ/mol)
metallic bonds, e.g. Au, Fe, etc.
~75 - 1000 kJ/mol
Inter molecular forces: forces between molecules
Much weaker: ~0.05 - 40 kJ/mol
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Intermolecular forces determine:
Boiling Point/Melting Point
Vapor Pressure
Viscosity (measure of a liquid’s resistance to flow)
Surface Tension (resistance of a liquid to an increase in its
surface area)
Figure 16.6
Figure 16.44
Stronger forces
require more energy to disrupt
intermolecular interactions
Types of Intermolecular Forces
Dipole-Dipole forces
• Based on charge-charge
interactions
• Fall off rapidly over distance
(relatively unimportant in the
gas phase)
• Occur in molecules with
polar bonds and dipole
moments (H-N, H-O, H-F)
• Typically 5-25 kJ/mol
Figure 16.2
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Hydrogen Bonding
A special type of dipole-dipole force.
Occurs when a Hydrogen atom is bound to the highly
electronegative atoms O, N, or F, which have lone pairs (let’s call
these atoms X or Z).
These H-X bonds are very polar (think HF). The partially positive
H atom from one unit is attracted to the partially negative X
atom from another unit (or to another partially negative atom,
Z). This is the hydrogen bond.
To occur, the atom sequence must be - Z :…H- X -, where both Z
and X are one of the following three highly EN atoms: O, N, or F.
Even though the strength of one H bond may be small, the
combined strength of many H bonds can be large.
Water has particularly strong hydrogen
bonding intermolecular forces
Isolated water
molecule (gas phase)
(+0.46)
(-0.92)
(+0.46)
Hydrogen bonding:
Typically 10 - 40 kJ/mol
Figure 16.3
H-bonded water molecules
(liquid or solid phase)
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Boiling Points are High in H-bonded liquids
Exceptions
Trend
Figure 16.4
O, N, F 
small sizes and large electronegativities: strong H-bonds
Drawing Hydrogen bonds between Molecules
Problem: In which of the following molecules does hydrogen bonding
occur? Draw the hydrogen bonds where appropriate in the molecules.
a) C3H8
b) C2H5OH
c) Glycine: H2NCH2COOH
Plan: We examine each structure to see if F, N, or O is present, and if
hydrogen can be bonded to them.
Solution:
a) for C3H8, only C-C and C-H bonds exist. No hydrogen bonds can be
formed!
H
H
H
H
C
C
C
H
H
H
H
b) for C2H5OH, the H covalently bound to the oxygen can interact
with the pair of electrons on the oxygen of another molecule to
form a strong hydrogen bond. It can also hydrogen bond with
water molecules…
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Identifying Hydrogen Bonds Between Molecules - II
b) Ethanol:
H H
H C C O H
H H
H O
H H
H H H
O
C C H
H C C O
H
H H
H H
H
H
c) Glycine:
N
H
H
H
N
H
H
O
C
C O
H
H
Adenine (A)
N
H
H
O
C
C
H
O
C
C
O H
H
O
H
H O
H
Thymine (T)
T
Guanine (G)
H
Cytosine (C)
C
A
G
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London Dispersion Forces
• In molecules without a
permanent dipole moment, it is
possible to induce an
“instantaneous” dipole moment.
• The interaction is weak and
short-lived, but can be very
significant for large molecules
(often more than dipole-dipole).
Dispersion Forces
Figure 16.5
Magnitude depends on polarizability and shape of the atom
or molecule: Heavier atoms/molecules with more electrons
are generally more polarizable. Molecules with high surface
areas interact more with their neighbors.
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Bromine
Rank by energy
• Br2: m.p. -7 ˚C, b.p. 59 ˚C.
• Molecular substances have relatively low melting and
boiling points because the molecules are held
together by fairly weak forces.
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Note: There can be more than one type of intermolecular
interaction between molecules or ions
Tetramethylammonium ion
Nitrate ion
ion-ion
dispersion
Ammonium ion
Nitrate ion
ion-ion
hydrogen-bonding
dispersion
Predicting the Types and Relative Strengths of
Intermolecular Forces
Problem: Select the substance with the higher boiling point in each pair:
a) CH3Cl or CH3OH
b) CH3CH2OH or C2H4(OH)2
c) n-pentane (C5H12) or neopentane (C5H12)
Plan: Examine the formulas and structures to determine the types of
forces involved:
Are ions present? Are the molecules polar or nonpolar? Is F, O, or N
bound to H? Do the molecules have different masses or shapes?
Remember:
a) Bonding (intramolecular) forces are stronger than intermolecular
forces.
b) Hydrogen bonding is a strong type of dipole-dipole force.
c) Dispersion forces are decisive when the major difference is molar
mass or molecular shape.
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Solutions:
a) CH3Cl (methyl chloride, MM=50.48g/mol) and CH3OH
(methanol, MM=32.04g/mol) are both polar molecules. CH3OH
has an O-H bond, so it can form hydrogen bonds, which are
stronger than the dipole-dipole forces that exist between the
molecules in CH3Cl. CH3OH will have the greater boiling point.
Check: CH3Cl (b.p.= -24.220C) and CH3OH (b.p.= 64.650C)
b) CH3CH2OH (ethyl alcohol, MM=46.07g/mol) & HOCH2CH2OH
(ethylene glycol, MM=64.07g/mol) both contain an O-H group,
but ethylene glycol is a di-hydroxy alcohol so it can have twice
the hydrogen bonding- it will have the higher boiling point.
Check: CH3CH2OH (b.p.= 78.5 0C) and HOCH2CH2OH (b.p.= 290 0C)
c) n-pentane (C5H12) and neopentane (C5H12) are both non-polar
molecules with the same molecular mass. The tetrahedral
neopentane makes less intermolecular contact than the straightchain n-pentane, so
n-pentane should have greater dispersion forces and a slightly
higher boiling point.
Check: n-pentane (b.p.=36.1 0C) and neopentane (b.p.=9.5 0C)
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Liquids
• Many liquids have the
tendency to bead on a
surface.
• This is the result of
surface tension – the
liquid’s resistance to an
increase in surface area.
• Large intermolecular
forces lead to high
surface tension.
Surface Tensions and Intermolecular Forces
Figure 16.6
Surface tension: the resistance of a liquid to an increase in
surface area; “surface” can be vapor/solid ; molecule at the
surface only attracted by molecules next to and below it
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Adhesive
Cohesive
Capillary Action
• Polar liquids also exhibit
capillary action: the
spontaneous rising of liquid
up a narrow tube.
• This is the result of cohesive
(liquid/liquid) and adhesive
(liquid/container) forces at
work.
– The liquid “wets” the
surface  increases surface
area  balance between
cohesive and adhesive…
liquid “pulls itself” up the
tube.
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Surfactants: Polar + Non-Polar Hybrids
Non-polar “tail”
Polar “head”
-CH2-CH2-CH2-
-CO2- as in soap
-SO3- as in industrial detergents
Air
Air-liquid
interface
Air-water
interface
Weak non-covalent
forces
Polar liquid
Surfactant: surface acting agent; lowers the surface tension
of a liquid (easier spreading)
Nitrogen
• Boiling point: −196 °C (77 K; −321 °F)
• Melting point: -210 °C
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Changes of State
• When molecules escape from the liquid state to form a
gas, the process is called vaporization (evaporation).
– An endothermic process: heat is needed to overcome
the intermolecular forces of the liquid.
DHvap = energy needed to vaporize 1 mol of liquid
at 1 atm pressure.
• The energy needed to melt a solid is called the enthalpy
or heat of fusion DHfus
• Not all substances need to pass from solid to liquid to
gas. Some substances can go from solid directly to gas.
This is called sublimation.
Vapor
Pressure
Recall:
PV = nRT
See Figures 16.44
and 16.45
At equilibrium, Evaporation rate = Condensation rate
The pressure of the vapor at equilibrium is called the
vapor pressure (Pvap) of the liquid (or solid) – think of
this as the number of gas molecules pushing down on
the liquid or solid
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Measuring Pvap
Figure 16.46
• If Pvap is large
substance called “volatile"
• As intermolecular forces increase, Pvap decreases
Vapor Pressures and Intermolecular Forces
Figure 16.44
More volatile substances:
have higher Pvap at lower T
(more molecules in the vapor phase)
Figure 16.48a
TREND: higher Pvap means weaker IMFs
(molecules can more easily escape to the vapor phase)
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∆H°vap = standard enthalpy of vaporization
(or standard heat of vaporization)
Example: H2O (l)
vaporization
condensation
H2O (g)
∆H°vap = 40.7 kJ/mol
This is why water is such a good coolant!
Note: “Standard” pressure is 1 atm and the quantity of
material is 1 mole
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Figure 16.48b
From a Pvap vs T plot (Fig 16.48a),
how can we tell which substance
has a smaller DHvap? Let’s plot the
data another way…
Clausius-Clapeyron Equation:
ln P 
DHvap 1 
  C
R T 
Similar to equation for the temperature dependence of
the equilibrium constant (Zumdahl section 10.11)
Does the equation make sense?
ln P 
DHvap 1 
  C
R T 
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Does the equation make sense?
ln P 
DHvap 1 
  C
R T 
Clausius-Clapeyron Equation:
ln P 
DHvap 1 
  C
R T 
For 2 temperatures, T1 and T2, and vapor pressures, P1 and P2:
 P  DH vap  1 1 
  
ln  2  
P
R
 1
 T1 T2 
From these equations, knowledge of any 2 pressures at 2
temperatures allows calculation of ∆Hvap.
Problem: Calculate the heat of vaporization of diethyl ether
from the following vapor pressures: 400 mm Hg at 18 oC &
760 mm Hg at 35oC.
Plan: We are given P1, P2, T1, and T2 substitute into the
C.C. equation and calculate the value of Hvap!
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Changes of State
Melting point:
• Energy added disrupts the solid structure, increasing the
potential energy of the molecules
• Solid and liquid states have the same Pvap
• “normal”: total Pvap= 1 atm
Boiling point:
• Pvap of a liquid becomes equal to the pressure of its
environment
• “normal”: temperature at which Pvap = 1 atm
 For water: 100°C
Changes of State
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Changes of State – Heating Curve
q = CgDT = J/g
q = ClDT = J/g
q = DHvap
= kJ/mol
q = DHfus
= kJ/mol
q = CsDT = J/g
Figure 16.50
and Ch 9
The Cooling Curve is Sensitive to Pressure
Boiling occurs when the vapor pressure of a liquid
becomes equal to the pressure of its environment.
In the mountains, atmospheric pressure is reduced, so
water boils at a lower temperature.
(Less heat is required to achieve Pvap = Patm)
Wouldn’t it be great if we had a way to figure out in which
physical state(s) a substance will exist at a given
temperature and pressure?
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Phase diagrams are important "road maps" for
understanding the effects of temperature and
pressure on a substance
Phase Diagram Landmarks
• Phase boundary lines: two phases are in equilibrium
• Triple point: point at which all three phases exists
simultaneously
• Critical temperature: the temperature above which a
vapor cannot be liquefied, no matter what pressure is
applied
• Critical pressure: the pressure required to produce
liquefaction at the critical temperature
• Critical point: beyond this point the transition from
one state to another involves neither true liquid nor
vapor (defined by critical temperature and critical
pressure.
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Phase Changes for Water and their Enthalpies
Gas
H0sub
-
Condensation
H0vap
= -40.7 kJ/mol
Liquid
Freezing
- H0fus = -6.02 kJ/mol
H0sub
Deposition
Sublimation
Vaporization
H0vap= 40.7 kJ/mol
-
Melting
H0fus = 6.02 kJ/mol
Solid
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Why is water so awesome?
• (Near) universal solvent
– The high polarity (and, therefore, hydrogen
bonding power) of water means it can dissolve so
many compounds – ionic compounds, polar, nonionic compounds and even non-polar gases.
Why is water so awesome?
• Thermal properties
– Water has a high heat capacity (higher than
almost any liquid). High intermolecular forces
mean water can absorb a lot of heat before it
boils. This allows the earth to remain at a steady
temperature.
– Water has a high heat of vaporization which gives
it enormous cooling power. Example:
Adding 4 kJ of heat to 1000g of water causes the
temperature to rise 1 °C, but only 2g of that water
has to evaporate to keep the remaining 998g at a
constant temperature.
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Why is water so awesome?
• Surface properties
– High surface tension and high “capillarity” (both
results of hydrogen bonding) are critical to plant
life (land and aquatic).
Why is water so awesome?
• Density of solid and liquid water
– Because solid water (ice) is less dense than liquid
water (as a result of hydrogen bonding), ice floats
on water. This protects aquatic life, erodes rocks,
but sometimes freezes your pipes…
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