Chapter 12, Worksheet 3 (ws12.3)
Factors that Affect Solubility
Review
1. A solid was added to water and the mixture was stirred until no more solid would dissolve and
the excess solid remained at the bottom of the container. At this point, the solution is
______________. We say that the dissolving reaction has “reached equilibrium”. What does
this mean?
2. Starting with the solid compound, describe how to make 250.0 mL of 0.200 M NaCl. You have
an analytical balance, a 250 mL volumetric flask, and some water.
3. What is the molality of NaCl in the solution made in question 2? Assume that the density of the
solution is 1.00 g/mL.
4. Why don’t water and oil mix?
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5. Consider the trend shown in the table below. Which of the alcohols would be most soluble in a
non-polar solvent?
6. Recall the solubility rules for ionic compounds dissolved in water. Many of the compounds
listed as “insoluble” in the table below are actually slightly soluble in water (the solubility is less
than 0.01 M) What type of ionic compounds would you expect to be only slightly soluble in
water? Consider solute-solute and solute-solvent interactions.
7. Why are network solids such as diamond (carbon) and SiO2 (sand, glass) totally insoluble in
water?
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Effect of temperature on the solubility of solids and gases in water
The solubility of most solids increases as the temperature increases.
The solubility of all gases decreases as the temperature increases.
To understand the effect of temperature on the solubility, we must consider the effect of temperature on
the entropy of a system.
1. Let’s first review how the second law of thermodynamics allows us to explain why ammonium
nitrate is very soluble in water. Ammonium nitrate dissolving in water is an endothermic
process. This means that the solute-solvent interactions are weaker than the sum of the solutesolute and solvent-solvent interactions. Therefore, you might think that very little ammonium
nitrate would dissolve in water. In reality, it is very soluble (s = 19 M!) so there must be some
“driving force” for the process that we have not yet considered. What is it?
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2. Remember this catch phrase: Increasing the temperature favors the more disordered state.
(We will talk a lot about this idea as the course progresses.)
a. Which is more disordered, a pure crystalline solid or the same solid dissolved in water?
b. Which is more disordered, a pure gas or the same gas dissolved in water?
c. Now explain the effect of temperature on the solubility of solids and gases.
d. Why do you think some solids buck the trend? (Look at cerium sulfate in the graph above.)
e. How do you think temperature affects the solubility of a liquid in water? Explain.
f. Why does an open soda taste better if stored in the frig rather than on the kitchen table?
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Henry’s law: Effect of partial pressure of a gas on its solubility in water
Henry’s Law: The solubility of a gas in liquid is directly proportional to the partial pressure of the gas
above the solution.
c = kP
c is the molar solubility of the gas
P is the partial pressure of the gas over the solution (in atm)
k is the Henry’s law constant. The value of k depends only on temperature and the identity of the gas.
1. If, at equilibrium, the partial pressure of an undissolved gas triples, what has happened to its
solubility?
2. Explain why the solubility of a gas increases as the partial pressure of the gas increases.
(Remember that in a saturated solution, the rate of dissolving is equal to the rate of leaving the
solution. What happens to these rates as P increases?)
3. Why does soda taste better if you store it on the kitchen table with the lid on than with the lid
off?
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4. The partial pressure of nitrogen in the atmosphere at sea level is 0.78 atm. At 25oC, the
solubility of nitrogen at sea level is 5.3 x 10-4 M.
a. What is the value of Henry’s law constant for nitrogen at 25oC. Include the correct unit.
b. What is the solubility of nitrogen in water when the partial pressure of nitrogen is 5.0 atm (in
a pressurized tank)?
c. What must happen to the value of k as the temperature increases?
5. Why do divers experience decompression sickness (“the bends”) if they surface too quickly?
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