Ch 11: Intermolecular Forces and
Types of Solids
Brown, LeMay
AP Chemistry

11.1: Intermolecular Forces (IMF)
 IMF < intramolecular forces (covalent, metallic, ionic bonds)
 IMF strength: solids > liquids > gases
 Boiling points and melting points are good indicators of relative IMF strength.
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11.2: Types of IMF
1. Electrostatic forces: act over larger distances in accordance with Coulomb’s law
3
F

Q

Q
 d 2 a. Ion-ion forces: strongest; found in ionic crystals
(i.e. lattice energy)

b. Ion-dipole: between an ion and a
neutral, polar molecule/has separated
(a partial charges)
 Increase with increasing polarity of molecule and increasing ion charge.
4
F

Q

Q
 d 2
Ex: Compare IMF in Cl (aq) and S 2(aq).
d  d  d  d  d 
Cl d  d  d  d 
< d  d  d 
S 2d  d  d  d  d  d 

c. Dipole-dipole: weakest electrostatic force; exist between neutral polar molecules
 Increase with increasing polarity (dipole moment) of molecule
Ex: What IMF exist in NaCl (aq)?
5

6 d. Hydrogen bonds (or H-bonds):
 H is unique among the elements because it has a single e that is also a valence e .
– When this e is “hogged” by a highly EN atom
(a very polar covalent bond), the H nucleus is partially exposed and becomes attracted to an e -rich atom nearby.

 H-bonds form with H-X•••X', where X and X'
7 have high EN and X' possesses a lone pair of e-
 X = F, O, N (since two molecules:
elements) on
F -H
O -H
N -H
:
:
:
F
O
N

 * There is no strict cutoff for the ability to form H-bonds (S forms a biologically important hydrogen bond in proteins).
 * Hold DNA strands together in double-helix
8
Nucleotide pairs form Hbonds
DNA double helix

 H-bonds explain why ice is less dense than water.
9

Ex: Boiling points of nonmetal hydrides
Conclusions:
 Polar molecules have higher BP than nonpolar molecules
∴ Polar molecules have stronger IMF
 BP increases with increasing MW
∴ Heavier molecules have stronger IMF
 NH
3
, H
2
O, and HF have unusually high BP.
∴ H-bonds are stronger than dipole-dipole IMF
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2. Inductive forces:
 Arise from distortion of the e cloud particle or molecule nearby.
by the electrical field produced by another
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 London dispersion: between polar or nonpolar molecules or atoms
– * Proposed by Fritz London in 1930
– Must exist because nonpolar molecules form liquids
Fritz London
(1900-1954)

How they form:
1. Motion of e- creates an instantaneous dipole moment, making it “temporarily polar”.
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2. Instantaneous dipole moment induces a dipole in an adjacent atom
• * Persist for about 10 -14 or 10 -15 second
Ex: two He atoms

* Geckos!
 Geckos’ feet make use of
London dispersion forces to climb almost anything.
 A gecko can hang on a glass surface using only one toe.
 Researchers at Stanford
University recently developed a gecko-like robot which uses synthetic setae to climb walls http://en.wikipedia.org/wiki/Van_der_Waals%27_force
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London dispersion forces increase with:
 Increasing MW, # of e , and # of atoms (increasing # of e orbitals to be distorted)
Boiling points:
Effect of MW: pentane 36ºC hexane 69ºC heptane 98ºC
Effect of # atoms:
Ne –246°C
CH
4
–162°C
??? effect:
H
2
D
2
O
O
100°C
101.4°C
 “Longer” shapes (more likely to interact with other molecules)
C
5
H
12 isomers: 2,2-dimethylpropane 10°C pentane 36°C
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Summary of IMF
Van der Waals forces

Ex: Identify all IMF present in a pure sample of each substance, then explain the boiling points.
BP(⁰C
)
IMF
HF 20
London, dipoledipole, H-bonds
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Explanation
Lowest MW/weakest London, but most polar/strongest dipole-dipole and has H-bonds
HCl -85
London, dipoledipole
Low MW/weak London, moderate polarity/dipole-dipole and no H-bonds
HBr -67
London, dipoledipole
Medium MW/medium London, moderate polarity/dipole-dipole and no H-bonds
HI -35
London, dipoledipole
Highest MW/strongest
London, but least polar bond/weakest dipole-dipole and no H-bonds

11.3: Properties resulting from IMF
1. Viscosity: resistance of a liquid to flow
2. Surface tension: energy required to increase the surface area of a liquid
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3. Cohesion: attraction of molecules for other molecules of the same compound
4. Adhesion: attraction of molecules for a surface

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5. Meniscus: curved upper surface of a liquid in a container; a relative measure of adhesive and cohesive forces
Ex:
Hg
(cohesion rules)
H O

11.4: Phase Changes
Processes:
 Endothermic: melting, vaporization, sublimation
 Exothermic: condensation, freezing, deposition
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I
2
(s) and (g) Microchip

Water: Enthalpy diagram or heating curve
Q
 mc

T Q
 mH
21
Q
 m ( 1 .
87 J/g

C )

T
Q
 m ( 2 260 J/g)
Q
 m ( 4.18
J/g

C )

T
Q
 m ( 334 J/g)
Q
 m ( 2 .
06 J/g

C )

T

11.5: Vapor pressure
Pressure cooker ≈ 2 atm
Normal BP = 1 atm
10,000’ elev ≈ 0.7 atm
29,029’ elev (Mt. Everest)
≈ 0.3 atm
 A liquid will boil when the vapor pressure equals the atmospheric pressure, at any T above the triple point.
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11.6: Phase diagrams: CO
2
 Lines: 2 phases exist in equilibrium
 Triple point: all 3 phases exist together in equilibrium (X on graph)
 Critical point, or critical temperature & pressure: highest T and P at which a liquid can exist (Z on graph) Temp (ºC)
 For most substances, inc P will cause a gas to condense
(or deposit), a liquid to freeze, and a solid to become more dense (to a limit.)
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Phase diagrams: H
2
O
• For H
2
O, inc
P will cause ice to melt.
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*
25

*
26

11.7-8: Structures of solids
 Amorphous: without orderly structure
Ex: rubber, glass
 Crystalline: repeating structure; have many different stacking patterns based on chemical formula, atomic or ionic sizes, and bonding
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Types of crystalline solids (Table 11.6)
Type
Atomic
Particles
Atoms
Forces
London dispersion
Notable properties
 Poor conductors
 Very low
MP
Examples
Ar (s),
Kr (s)

Molecular
Molecules
(polar or non-polar)
London dispersion, dipoledipole, Hbonds
 Poor conductors
 Low to moderate
MP
CO
2
(s),
C
12
H
22
O
11
,
H
2
O (s)
Carbon dioxide, dry ice
(g at room T)
Ice
(liq at room T)
Sucrose
(liq at room T)

Ionic
Anions and cations
Ion-ion (ionic bonding)
 Hard & brittle
 High MP
 Poor conductors
 Some solubility in
H
2
O
NaCl,
Ca(NO
3
)
2

Covalent
(a.k.a. covalent network)
Atoms bonded in a covalent network
Covalent bonds
 Very hard
 Very high
MP
 Generally insoluble
 Variable conductivity
C (diamond
& graphite)
SiO
2
(quartz)
Ge, Si, SiC,
BN
Graphite Diamond SiO
2

Metallic
Metal cations in a diffuse, delocalized e cloud
Metallic bonds
 Excellent conductors
 Malleable
 Ductile
 High but wide range of MP
Cu, Al, Fe