Unit 4 Practice Problems Key
Intermolecular Forces
1.
2.
3.
4.
5.
As volume decreases (pressure increase) and/or
temperature decreases real gases act less ideally. CO2
is a nonpolar molecule with only weak LDF and
therefore, would not experience as strong as IMF’s as
SO2, which is a polar molecule with dipole-dipole
interactions. In the decreasing volume of non-ideal
conditions, the SO2 molecules would attract and their
collisions with the container would be fewer and with
𝑃𝑉
less force causing the
ratio to decrease.
𝑅𝑇
High pressure is generated by crowded molecules where
the volume of empty space (Videal) is significantly less
than 100 % of the total volume. The gas molecules
themselves take up an increasing proportion of the
𝑃𝑉
volume. The
ratio would increase.
𝑅𝑇
Close to boiling point (very cold temperatures << 0ºC),
molecules slow down and clump and collide less often,
which generates less pressure  Preal < Pideal
Dipole-Dipole
d
Pair
H2O & H2S
Dispersion
b
H-bond
a
ion-dipole
c
Justification
H-bonding > dipole-dipole forces
Ne & Kr
6.
7.
6.
Both LDF, Greater atomic mass
LDF for Cl2 weaker and dipole-dipole
Cl2 & SO2
for SO2
Propanol molecules H-bond, which is a stronger
attraction then ethyl methyl ether's dipole forces.
H2= none, H2S= dipole-dipole, CHF3= dipole-dipole,
NH3= H-bonding.
15. At constant pressure and increasing temperature; solid
changing to liquid (melting) and liquid to gas
(vaporization).
16. solid
17. 0.01ºC.
18. 0ºC
19. Above this temperature and pressure the gas and liquid
phase are indistinguishable, called supercritical fluid.
20.
a)
b)
21.
Can liquid CO2 exist at room pressure? no
What happens to CO2(s) at -78.5oC?
b)
c)
8.
a)
b)
c)
d)
e)
9.
10.
11.
12.
55 ºC
90ºC
the phase changes
5.11 atm
What is the critical temperature for CO2?
31.1 oC
22. Water boils when vapor pressure equals atmospheric
pressure. As air is pumped out of the bell jar, its
pressure decreases until it reaches the vapor pressure.
23. Water boils at a lower temperature at high altitudes,
which means that water based foods cook slower
because they cook at the boiling point.
24. 𝑞 = 𝑚𝐶𝑝 ∆𝑇
10𝑔𝐻2 𝑂 ×
10𝑔𝐻2 𝑂 ×
18.02 𝑔
26.
27.
28.
29.
30.
×
6.01𝑘𝐽
1𝑚𝑜𝑙
= 3.34𝑘𝐽
(10.0𝑔)(4.18 𝐽/𝑔𝐶)(100𝐶) = 4180𝐽
1𝑚𝑜𝑙
18.02 𝑔
𝑞 = 𝑚𝐶𝑝 ∆𝑇
25. .
a)
b)
c)
d)
e)
f)
(10.0𝑔)(2.09)(5𝐶) = 104.5𝐽
1𝑚𝑜𝑙
𝑞 = 𝑚𝐶𝑝 ∆𝑇
𝑞 = 𝑚𝐶𝑝 ∆𝑇 = (2.0 𝑔)(. 565 𝐽 ∙ 𝑔/𝐶)(25𝐶) = 𝟐𝟖. 𝟐𝟓 𝑱
𝑞 = 𝑚𝐶𝑝 ∆𝑇 = (2.0 𝑔)(. 565 𝐽 ∙ 𝑔/𝐶)(−20𝐶) = − 𝟐𝟐. 𝟔 𝑱
KE remains constant
KE increasing
Segment D-E is longer than segment B-C
a) sublimation b) vaporization c) condensation
d) solidification e) deposition f) melting (enthalpy of
fusion)
13. gas
14. At constant temperature the pressure is increasing; gas
changing to a solid (deposition), then solid changing to a
liquid (melting).
solid
What is the triple point pressure for CO2?
7.
Q-R is liquid only, R-S is phase change, S-T is gas
only.
As heat is absorbed the KE of the molecules in
increasing.
The molecules are undergoing a phase change,
vaporization ∆𝐻𝑣𝑎𝑝 = 40.7 𝑘𝐽/𝑚𝑜𝑙
sublimates
Which is the most dense phase for CO2?
NaCl is ionic bonding, therefore; it would have the
highest heat of vaporization. H2O and HF use Hbonding, HF is more polar because F is more
electronegative than O. HF would have the higher heat
of vaporization.
a)
The phase diagram for water shows that as the
pressure increases the MP decreases, the slope of
the line separating the solid and liquid phase is
negative. Reason: water is unique in that the solid
state is less dense than the liquid state.
The solid/liquid line would have a positive slope.
×
40.4𝑘𝐽
1𝑚𝑜𝑙
= 22.59𝑘𝐽
(10.0𝑔)(2.01 𝐽/𝑔𝐶)(5𝐶) = 100.5 𝐽
𝟑𝟎. 𝟑𝟐𝒌𝑱 𝒕𝒐𝒕𝒂𝒍
26.7 kPa
6.65 kPa
75ºC
26.7 kPa
118ºC
temperature at which the vapor pressure is 1 atm or
101.2kPa
g) approx. 72ºC
h) approx. 116ºC
i) 70ºC (remember boiling occurs when Pvap=Patm)
j) approx. 20kPa
water
pentane
First pentane would boil away, then tetrachloromethane
and finally water.
a) approx. 40kJ/mol, based on extrapolation of graph.
b) All show dispersion forces
c) As number of electrons increase “polarizability” of
atom also increases, therefore the dispersion forces
for I2 are stronger, more energy has to be absorbed
to vaporize.
a) H2O=H-bonding, all others dipole-dipole
b) H2O has high Hvap, then drops down and climbs again
for H2Te.
c) The increase for the large molecules is a result of
increasing LDF for the larger, more polarizable
molecules.
32.
31.
a)
b)
c)
All have LDF present. ethanethiol is LDF only,
ethane is LDF only, ethanol has H-bonding.
False statement. Boiling is a phase change and
doesn’t involve breaking the covalent intramolecular
bonding. Only the IMF’s are broken during phase
changes.
The water molecules (H-bonding) are able to attract
to the ethanol because it too has H-bonding,
therefore “like dissolves like” is appropriate. The
other nonpolar molecules (LDF) will dissolve in
another nonpolar molecule.
Molecular Crystal
London
forces
Dipoledipole
attractions
NH3
Kr
Hydrogen
Bonds
Metallic
Bonds
Ionic
Crystal
Ionic
Bonds
X
X
X
KMnO4
X
NaCl
X
SO2
X
CO2
X
C3H8
X
CH4
X
CH3Cl
X
HF
C6H6
Network
Solid
Covalent
Bonds
X
HCl
F2
Metal
X
X
NO
X
H2SO4
X
WC
X
Si
X
SiO2
X
C(graphite)
X
N2
X
CH3OH
X
Ag
X
(C2H5)2NH
X
NaOH
X
Al
X
PCl3
X
33.
Structural Unit
Bond name
Bond strength
Melting point
Solubility
Conductivity
Malleability
Metallic
ion
metallic
variable
variable
low
high
high
Covalent Network
atom
covalent
strong
high
low
low
low
Molecular
molecule
molecular
weak
low
variable
low
variable
Ionic
ion
ionic
strong
high
high
low
low
34.
35.
a)
SO2 forms dispersion and dipole forces between
molecules whereas, SiO2, a covalent network solid,
forms covalent bonds throughout. The stronger
covalent bonds, which break during melting of SiO2,
require more energy (higher temperature) to break.
b)
Liquid Cl2 is held together by dispersion forces,
which although weak increase in strength as the
number of electrons increases. Liquid HCl is held
together by dipole forces in addition to dispersion
forces, but the addition of dipole forces between HCl
molecules must not make up for the fewer electrons
around the HCl.
c)
The stronger ionic bond in NaCl is due to the
smaller Na+ ion compared to K+, which allows the
Cl- ion to get closer and strengthens the attraction
between ions, making the NaCl bond stronger and
the melting point higher than KCl.
d)
Si is a covalent network solid with strong covalent
bonds between atoms. Cl2 has discrete molecules
with weak London dispersion forces between
molecules. Therefore, melting Si requires a higher
temperature than Cl2.