7.5 Liquid/Gas Systems and Henry’s Law
7.26.Convert the units of the Henry’s law constant for CO 2 ,
in Table 7.1, to units of mmHg, atm, and bar. In which case(s)
does the numerical value of the constant change?
7.27.What is the difference between hydrogen chloride and
hydrochloric acid? Do you expect that either of them acts as
an ideal substance?
7.28.The Henry’s law constant for methyl chloride, CH 3 Cl, in
aqueous solution is 2.40 106 Pa. What pressure of methyl
chloride is necessary to establish a mole fraction of 0.0010 in
an aqueous solution?
7.29.The mole fraction of CCl 2 F 2 , a compound once used as
a refrigerant, in an aqueous solution was found to be 4.17
10 ^5 at normal pressure. What is the approximate molarity of
this solution and what is the Henry’s law constant for this gas?
Use a density of 1.00 g/cm^3 for water.
7.30.At 25°C, the mole fraction of air in water is about
1.388 10 ^5. (a)What is the molarity of this solution?
(b)What is the Henry’s law constant for air? (c)Would you
expect the solubility of air to increase or decrease with an in-
crease in temperature? Compare your numerical answers to
the constants for nitrogen and oxygen in Table 7.1.
7.31.At 25°C, the mole fraction of nitrogen, N 2 (g), in water
is 1.274 10 ^5. (a)Compare this with the number in the
previous problem and comment. (b)Calculate the solubility
of oxygen, O 2 (g), in water given the fact that air is approxi-
mately 80% nitrogen and 20% oxygen. (c)Calculate the
Henry’s law constant for oxygen. Compare your answer to the
number in Table 7.1.
7.32.Does a higher Henry’s law constant mean that a gas is
more soluble in a liquid, or less soluble? Be able to defend
your answer.
7.6 & 7.7 Liquid/Solid and
Solid/Solid Solutions
7.33.What is the approximate molarity of a saturated solu-
tion of phenol, C 6 H 5 OH, for which 87.0 g can be dissolved in
100 mL of water? The density of phenol is 1.06 g/cm^3 ; assume
ideal behavior with respect to the total volume of the solution.
7.34.Calculate the solubility of phenol, C 6 H 5 OH, in water at
25°C if (^) fusHfor phenol is 11.29 kJ/mol and its melting point
is 40.9°C. Compare the calculated solubility with the numbers
from the previous exercise. Can you explain any deviations?
7.35. (a)Convert the calculated mole fraction of naphtha-
lene dissolved in toluene from Example 7.10 into molarity, as-
suming that the volumes are strictly additive. The density of
toluene is 0.866 g/mL and the density of naphthalene is 1.025
g/mL. Assume the volumes are additive.
(b)Estimate the solubility, in g/100 mL and molarity, of naph-
thalene in n-decane, C 10 H 22 , which has a density of 0.730
g/mL.
7.36.Will equation 7.39 work for the solubility of gases in liq-
uids? Why or why not?
7.37.Consider the following solutions:
Sodium chloride (s) in water
Sucrose (s) in water
C 20 H 42 (s) in cyclohexane
Water in carbon tetrachloride
For which solution(s) do you think that a calculated solubility
will be close to the experimental solubility? Explain your rea-
soning.
7.38.Determine how ideal the following solutions are by cal-
culating the mole fraction of solute in each solution, and com-
paring that to the expected mole fractions. All data are for
25.0°C.
(a)14.09 weight percent of I 2 in C 6 H 6 , MP of I 2 is 112.9°C
(sublimes), and (^) fusH15.27 kJ/mol
(b)2.72 weight percent of I 2 in C 6 H 12 , MP of I 2 is 112.9°C
(sublimes), and (^) fusH15.27 kJ/mol
(c)20.57 weight percent of para-dichlorobenzene, C 6 H 4 Cl 2 ,
in hexane, MP of C 6 H 4 Cl 2 is 52.7°C, and (^) fusH17.15 kJ/mol
7.39.Iron metal has a (^) fusHvalue of 14.9 kJ/mol and is solu-
ble in mercury to the level of xFe8.0 10 ^3 at 25.0°C.
Estimate the melting point of iron. Compare the estimate to
the literature value of 1530°C.
7.40.How many degrees of freedom are required to specify
the eutectic for a two-component system?
7.41.Do communities that use salt in the winter use enough
to form the low-melting eutectic between NaCl and H 2 O, or
are they taking advantage of the freezing-point depression
phenomenon in general? How can you tell?
7.42.Starting from xNa0.50 in Figure 7.23 in the liquid re-
gion, describe what happens as the temperature is decreased
until the entire solution is solid.
7.43.Construct a qualitative phase diagram for the Sn/Sb
system, which has binary eutectics at 92% and 95% Sn that
melt at 199°C and 240°C, respectively. The melting points of
tin and antimony are 231.9°C and 630.5°C.
7.44.Explain why zone refining, used to make ultra-pure sil-
icon, would not be a practical method of making ultra-pure
carbon.
7.45.Estimate the solubility of Na in Hg at 0°C. The heat of
fusion of sodium is 2.60 kJ/mol and its melting point is 97.8°C.
7.46.Show how the formula of the stoichiometric compound
in Figure 7.23 was determined.
7.8 Colligative Properties
7.47.Explain how the unit molarity automatically includes
the concept of partial molar volumes.
7.48.Why do you think people who live at high altitudes are
advised to add salt to water when boiling food like pasta?
What mole fraction of NaCl is needed to raise the boiling point
of H 2 O by 3°C? Does the amount of salt added to water (typ-
ically about one teaspoon to four quarts of water) substantially
change the boiling point? Kb(H 2 O) 0.51°C/molal.
204 Exercises for Chapter 7