CHAPTER TEN ANSWERS 6-40
6.
There is an electrostatic attraction between the permanent dipoles of polar molecules. The greater
the polarity, the greater the attraction among molecules.
7.
London dispersion (LD) < dipole-dipole < H-bonding < metallic bonding, covalent network, ionic.
Yes, there is considerable overlap. Consider some of the examples in Exercise 10.86. Benzene (onh
LD forces) has a higher boiling point than acetone (dipole-dipole forces). Also, there is even more
overlap between the stronger forces (metallic, covalent, and ionic).
8.
As the size of the molecule increases, the strength of the London dispersion forces also increases. As
the electron cloud gets larger, it is easier for the electrons to be drawn away from the nucleus (more
polarizable).
9.
As the strengths of interparticle forces increase, surface tension, viscosity, melting point and boiling
point increase, while the vapor pressure decreases.
10.
Dipole forces are generally weaker than hydrogen bonding. They are similar in that they arise from
an unequal sharing of electrons. We can look at hydrogen bonding as a particularity strong dipole
force.
11.
a. Polarizability of an atom refers to the ease of distorting the electron cloud. It can also refer to
, distorting the electron clouds in molecules or ions. Polarity refers to the presence of a perma
dipole moment in a molecule.
b. London dispersion forces are present in all substances. LD forces can be referred to as
accidental dipole - induced dipole forces. Dipole - dipole forces involve the attraction of
molecules with permanent dipoles for each other.
c. inter: between; intra: within; For example, in Br2 the covalent bond is an intramolecular fonx
holding the two Br-atoms together in the molecule. The much weaker London dispersion forces
are the intermolecular forces of attraction which hold different molecules of Br2 together in the
liquid phase.
12) Liquids and solids both have characteristic volume and are not very compressible. Liquids and gases
flow and assume the shape of their container.
13) Atoms have an approximately spherical shape (on the average). It is impossible to pack spheres
together without some empty space between the spheres.
14) Critical temperature:
The temperature above which a liquid cannot exist, i.e., the gas cannot be
liquified by increased pressure.
Critical pressure: The pressure that must be applied to a substance at its critical temperature to
produce a liquid.
The kinetic energy distribution changes as one raises the temperature (T4 > Tc > T3 > T2 > Tj). At the critical
temperature, Tc, all molecules have kinetic energies greater than the intermolecular forces, F, and a liquid
can't form. Note: The distributions above are not to scale.
15) As the intermolecular forces increase, the critical temperature increases.
16) Evaporation takes place when some molecules at the surface of a liquid have enough energy to break
the intermolecular forces holding them in the liquid phase. When a liquid evaporates, the molecules that
escape have high kinetic energies. The average kinetic energy of the remaining molecules (whimps) is
lower, thus, the temperature of the liquid is lower.
17) Types of substances by bonding:
a. Crystalline solid:
Amorphous solid:
b. Ionic solid:
Molecular solid:
c. Molecular solid:
Regular, repeating structure
Irregular arrangement of atoms or molecules
Made up of ions held together by ionic bonding.
Made up of discrete covalently bonded molecules held together in
the solid phase by weaker forces (LD, dipole or hydrogen bonds).
Discrete, individual molecules
Covalent network solid: No discrete molecules; A covalent network solid is one large
molecule. The interparticle forces are the covalent bonds between
atoms.
d. Metallic solid:
Covalent network solid:
18.
19.
20.
Completely delocalized electrons, conductor of electricity (ions in a
sea of electrons)
Localized electrons; Insulator or semiconductor
A crystalline solid because a regular, repeating arrangement is necessary to produce planes of atoms that
will diffract the x-rays in regular patterns. An amorphous solid does not have a regular repeating
arrangement and will produce a complicated diffraction pattern.
No, an example is common glass which is primarily amorphous Si02 (a covalent network solid) as
compared to ice (a crystalline solid held together by weaker H-bonds). The intermolecular forces in the
amorphous solid in this case are stronger than those in the crystalline solid. Whether or not a solid is
amorphous or crystalline depends on the long range order in the solid and not on the strengths of the
intermolecular forces.
Conductor: The energy difference between the filled and unfilled molecular orbitals is minimal. We call
this energy difference the band gap. Since the band gap is minimal, electrons can easily
move into the conduction bands (the unfilled molecular orbitals).
Insulator: Large band gap; Electrons do not move from the filled molecular orbitals to the
conduction bands since the energy difference is large.
Semiconductor: Small band gap; Since the energy difference between the filled and unfilled
molecular orbitals is smaller than in insulators, some electrons can jump into the
conduction bands. The band gap, however, is not as small as with conductors, so
semiconductors have intermediate conductivity.
21.
a. As the temperature is increased, more electrons in the filled molecular orbitals have sufficient kinetic
energy to jump into the conduction bands (the unfilled molecular orbitals).
b. A photon of light is absorbed by an electron which then has sufficient energy to jump into the
conduction bands.
c.
An impurity either adds electrons at an energy near that of the conduction bands (n-type) or
creates holes (empty energy levels) at energies in the previously filled molecular orbitals (ptype).
22.
In conductors the electrical conductivity is inversely proportional to the temperature. Increases in
temperature increases the motions of the atoms which gives rise to increased resistance (decreased
conductivity). In a semiconductor the electrical conductivity is directly proportional to the temperature.
An increase in temperature provides more electrons with enough kinetic energy to jump from the filled
molecular orbitals to the conduction bands, increasing conductivity.
23.
a. Condensation: vapor -> liquid
c. Sublimation:
b. Evaporation:
liquid -> vapor
solid -> vapor
d. A supercooled liquid is a liquid which is at a temperature below its freezing point.
24.
Equilibrium:
Dynamic:
25.
There is no change in composition; the vapor pressure is constant.
Two processes, vapor -> liquid and liquid ->• vapor, are both occurring but with
equal rates so the composition of the vapor is constant.
a. As the intermolecular forces increase, the rate of evaporation decreases. b.
As temperature increases, the rate of evaporation increases, c. As surface area
increases, the rate of evaporation increases.
26.
Sublimation will occur allowing water to escape as H2O(g).
27.
The phase change, H2O(g) -> H2O(1), releases heat that can cause additional damage. Also steam can be at
a temperature greater than 100°C.
28.
Fusion refers to a solid converting to a liquid and vaporization refers to a liquid converting to a gas. Only
a fraction of the hydrogen bonds are broken in going from the solid phase to the liquid phase. Most of the
hydrogen bonds are still present in the liquid phase and must be broken during the liquid to gas phase
transition. Thus, the enthalpy of vaporization is much larger than the enthalpy of fusion since more
intermolecular forces are broken during the vaporization process.
29.
a. ionic
b. dipole, LD (LD = London dispersion)
c. LD only
d. LD only; For all practical purposes, we consider a C - H bond to be nonpolar.
e. ionic
f. LD only
g. H-bonding, LD
a. ionic
30.
b. LD mostly; C - F bonds are polar, but polymers like tuflon are so large the LD forces are the
predominant intermolecular forces.
31.
c. LD
d. dipole, LD
f. dipole, LD
g. LD
e. H-bonding, LD
a. OCS; OCS is polar and has dipole-dipole forces in addition to London dispersion (LD) forces. All
polar molecules have dipole forces. CO2 is nonpolar and only has LD forces. In all of the
following (b-d), only one molecule is polar and, in turn, has dipole-dipole forces. To predict
polarity, draw the Lewis structure and deduce if the individual bond dipoles cancel.
b. PF3; PF3 is polar (PF5 is nonpolar).
c. SF2; SF2 is polar (SF6 is nonpolar).
d. SO2; SO2 is polar (SO3 is nonpolar).
32.
a. H2NCH2CH2NH2; More extensive hydrogen bonding is possible.
b. H2CO; H2CO is polar while CH3CH3 is nonpolar. H2C03 has dipole forces in addition to LD forces.
c. CH3OH; CH3OH can form relatively strong H-bonding interactions, unlike H2CO. d. HF; HF is
capable of forming H-bonding interactions, HBr is not.
33.
a. Neopentane is more compact than n-pentane. There is less surface area contact.between neopentane
molecules. This leads to weaker LD forces and a lower boiling point.
b. Ethanol is capable of H-bonding, dimethyl ether is not. c. HF is
capable of H-bonding, HC1 is not.
d. LiCl is ionic and HC1 is a molecular solid with only dipole forces and LD forces. Ionic forces are much
stronger than the forces for molecular solids.
e. n-pentane is a larger molecule so has stronger LD forces.
f. Dimethyl ether is polar so has dipole forces, in addition to LD forces, unlike n-propane which only has LD
forces.
34.
CH3C02H:
H-bonding + dipole forces + LD forces
CH2C1CO2H: H-bonding + larger electronegative atom replacing H (greater dipole) + LD forces
CH3CO2CH3:
dipole forces (no H-bonding) + LD forces
From the intermolecular forces listed above, we predict CH3CO2CH3 to have the weakest intermolecular forces and
CH2C1CO2H to have the strongest. The boiling points are consistent with this view.
35.
See Question 10.9 to review the dependence of some physical properties on the strength of the intermolecular
forces.
a. HC1; HC1 is polar while Ar and F2 are nonpolar. HC1 has dipole forces unlike Ar and F2. b. NaCl; Ionic
forces are much stronger than molecular forces.
c. I2; All are nonpolar so the largest molecule (I2) will have the strongest LD forces and the lowest vapor
pressure.
d. N2; Nonpolar and smallest, so has weakest intermolecular forces, e. CH4;
Smallest, nonpolar molecule so has weakest LD forces.
f. HF; HF can form relatively strong H-bonding interactions unlike the others.
g. CH3CH2CH2OH; H-bonding unlike the others so has strongest intermolecular forces.
36.
a. NO; NO is polar while N2 and O2 are nonpolar. NO has dipole forces unlike N2 and O2.
b. CH3CN; Polar, but no H-bonding so has weaker forces than H20 and CH3OH which can both Hbond.
c. H2; Nonpolar like CH4 but H2 is smaller than CH4, so H2 has weaker LD forces, d.
H2O; H2O can H-bond unlike the others so has strongest intermolecular forces.
e. HOCH2CH2OH (ethylene glycol); Greatest amount of H-bonding since two -OH groups are present.
f.
NH3; Need N - H, F - H, or 0 - H covalent bond for hydrogen bonding. NH3 can form H-bonds.
g. H20; H2O can H-bond unlike others.
h. CO2; Nonpolar substance, unlike others, so has weakest intermolecular (LD) forces.
37.
The attraction of H2O for glass is stronger than the H20 - H2O attraction. The miniscus is concave to
increase the area of contact between glass and H20. The Hg - Hg attraction is greater than the Hg - glass
attraction. The miniscus is convex to minimize the Hg - glass contact. Polyethylene is a nonpolar
substance. The H20 - H2O attraction is stronger than the H20 - polyethylene attraction. Thus, the
miniscus will have a convex shape.
38.
H2O will rise higher in a glass tube because of a greater attraction for glass than for polyethylene.
39.
The structure of H2O2 is H - O - O - H, which produces greater hydrogen bonding than water. Long
chains of hydrogen bonded H2O2 molecules then get tangled together.
40.
CO2 is a gas at room temperature. As mp and bp increase, the strength of the intermolecular forces also
increases. Therefore, the strength of forces is CO2 < CS2 < CSe2. From a structural standpoint this is
expected. All three are linear, nonpolar molecules. Thus, only London dispersion forces are present.
Since the molecules increase in size from CO2 < CS2 < CSe2, the strength of the intermolecular forces will
increase in the same order.