Intermolecular Forces
and
Liquids and Solids
Chapter 12
1. See Table 12.3 in Chang. Note that glycerol has the
highest viscosity of the liquids shown. Look up the
structure and provide a reason why glycerol is more
viscous than water.
•
Viscosity increases with the
strength of intermolecular forces.
•
Glycerol and water are both polar
compounds: dispersion forces and
dipole-dipole forces.
•
Both are capable of hydrogen
bonding, a special kind of dipoledipole force.
•
Difference in viscosity must be
explained in terms of the number of
hydrogen bonds that can be
formed.
12.13 Arrange the following compounds in order of
increasing boiling point: RbF, CO2, CH3OH,
CH3Br. Explain your arrangement.
Are any of the compounds ionic? If so, they will exhibit
the stronger intermolecular forces. RuF
Are any of the compounds nonpolar? If so, they exhibit
the weakest intermolecular forces. CO
2
Are any compounds polar? If so, they will exhibit
intermediate molecular forces. CH3OH, CH3Br
Are any of the polar compounds capable of hydrogen
bonding? If so, they will exhibit the stronger intermediate
forces. CH OH
3
weaker
stronger
12.19 These nonpolar molecules have the same number
and type of atoms. Which one would you expect to have a
higher boiling point? (Hint: Molecules that can be stacked
together more easily have greater intermolecular
attraction.)
linear
branched
The molecule with the larger surface area has stronger
van der Waals forces; for two molecules with the same molecular
formula but different connectivities (structural isomers), the less
compact molecule has the larger surface area.
A linear isomer is less compact than a branched isomer.
12.15 Which member of each of these pairs of
substances would you expect to have the
highest boiling point: (a) O2 or N2, (b) SO2 or
CO2, (c) HF or HI?
12.17 Explain in terms of intermolecular forces
why (a) NH3 has a higher boiling point than CH4
and (b) KCl has a higher melting point than I2.
Solids
Crystalline—high order
(regular repeating pattern)
Amorphous—low order
Attractions in Ionic Crystals
In ionic crystals, ions pack themselves so as to maximize the
attractions and minimize repulsions between the ions.
Crystalline Solids
basic repeating structural unit of a crystalline solid— unit cell
spheres— lattice points
At lattice points:
•
Atoms
•
Molecules
•
Ions
12.4
12.4
Shared by 8
unit cells
Shared by 2
unit cells
12.4
12.4
1 atom/unit cell
2 atoms/unit cell
4 atoms/unit cell
(8 x 1/8 = 1)
(8 x 1/8 + 1 = 2)
(8 x 1/8 + 6 x 1/2 = 4)
12.4
We can determine the empirical formula of an ionic solid by
determining how many ions of each element fall within the unit cell.
What are the empirical formulas for these
ionic solids?
(a) Green: chlorine; Gray: cesium
(b) Yellow: sulfur; Gray: zinc
(c) Green: calcium; Gray: fluorine
(a)
(b)
CsCl
ZnS
(c)
CaF2
Know diagonal of cube (c) in terms of r
Know diagonal of one face (b) in terms of a
If we can relate (c) to a then, we know (c) in
terms of r
12.4
When silver crystallizes, it forms face-centered cubic
cells. The unit cell edge length is 409 pm. Calculate
the density of silver.
d=
m
V
V = a3 = (409 pm)3 = 6.83 x 10-23 cm3
4 atoms/unit cell in a face-centered cubic cell
m = 4 Ag atoms x
1 mole Ag
107.9 g = 7.17 x 10-22 g
x
23
6.022 x 10 atoms mole Ag
7.17 x 10-22 g
m
3
=
=
10.5
g/cm
d=
V
6.83 x 10-23 cm3
12.4