Chapter 13 Reading Guide
Section 13.1
Define the following terms:
solvation:
hydration:
entropy:
What are the two main factors that determine the ability of substances to form solutions?
How does the strength of IMFs between solute and solvent determine the extent to which a solution will
form? Use this to explain why non-polar substances do not dissolve in water, yet many ionic compounds do
dissolve in water.
How do water molecules align near ions? Be specific!
Explain, in terms of energy changes, the three components of the solution process. Look carefully at Figures
13.3 and 13.4 to help your understanding of energy changes in the solution process.
Why are exothermic processes more likely to occur spontaneously than endothermic processes? How is it
possible that a endothermic solution process may still occur spontaneously?
Section 13.2
Define the following terms:
crystallization:
saturated:
solubility:
unsaturated:
supersaturated:
What does it mean to say that a solution that is saturated is in a “state of dynamic equilibrium”?
How is a supersaturated solution formed? Why is it unstable?
Section 13.3
Define the following terms:
miscible:
immiscible:
State the effect of solute-solvent interactions on solubility.
Explain the increase of solubility of NON-polar gases as shown in Table 13.2. WHY does the solubility
increase?
Explain the decrease in solubility of the alcohols as seen in Table 13.3.
Why are the trends in table 13.2 and 13.3 so different? (In other words, why does the solubility of non-polar
gases increase as MM increases, but the solubility of alcohols decreases as MM increases??)
What effect does increased pressure have on the solubility of liquids and solids? On gases?
Write the mathematical expression of Henry’s Law. Identify each variable.
For most solids, how does an increase in temperature affect solubility?
For gases, how does an increase in temperature affect solubility?
Section 13.4
Define each of the following with its mathematical formula:
mass percent:
ppm:
mole fraction:
molarity:
molality:
mass percent:
ppm:
mole fraction:
molarity:
molality:
Why is molality use in place of molarity if they are such similar measurement of concentration?(Under what
circumstances is molality preferable?)
Why is density of a solution a useful means of converting concentration units?
Section13.5
Define the following terms:
colligative properties:
ideal solution:
semi-permeable:
osmosis:
osmotic pressure:
isotonic:
hypotonic:
hypertonic:
What effect does the presence of a non-volatile solute have on each of the following properties of a solvent?
vapor pressure:
boiling point:
freezing point:
Write down the mathematical expression of Raoult’s Law and define each variable.
Look closely at Figure 13.22. Note how the effect of a solute on the vapor pressure of a solvent is responsible
for the effect on the boiling point of the solvent. (Recall that a liquid will boil at a temperature in which vapor
pressure is equal to atmospheric pressure)
How is the effect of an ionic solute different than the effect of a non-electrolyte on the colligative properties
of a solvent. WHY? (What does “colligative” mean?)
Look at Figure 13.23 (labeled osmosis). Why does the solvent move from a “more concentrated” solution to a
“less concentrated” solution?
Read carefully through sample exercises13.12 to note how the concept of colligative properties can be used to
determine the molar mass of a solute. This is a VERY important calculation! Try the practice exercise (the
answer is given)