Properties of Solutions Notes Packet
The Solution Process
Solutions, are homogeneous mixtures of two or more substances. The substance
present in greatest amount is called the solvent, and the other components present are
called solutes. Although solutions consisting of a liquid solvent and a solid solute
(dissolved in the solvent) are most familiar, solutions can involve many different
combinations of the three states of matter.
The intermolecular forces act between solvent and solute particles in solutions.
When sodium chloride dissolves in water, the sodium and chloride ions and the polar
water molecules are powerfully attracted to one another by ion-dipole interactions. Polar
water molecules surround the ions in solution in a process called solvation. (When water
is the solvent, the process is called specifically hydration.)
Whether a substance is soluble in a particular solvent and whether the dissolution
process will be exothermic or endothermic depends on the relative magnitudes of the
energy changes for the three steps. In general, a substance will be soluble unless the
energy expended to break apart the solvent and solute particles is significantly greater
than the energy given off when solute-solvent interactions are established.
Saturated Solutions and Solubility
As a solid dissolves in water, the number of aqueous solute particles increases. As
it dissolves, though, the reverse process (recrystallization of the solute) also occurs. As
the concentration of dissolved solute increases, so does the rate of recrystallization. After
a period of time the rate of recrystallization eventually equals the rate of dissolution, and
a dynamic equilibrium is achieved. No further net increase in the amount of solute in
solution occurs. When this dynamic equilibrium has been established, we call the solution
saturated. (By definition, a saturated solution must be in contact with undissolved solid.)
Furthermore, the concentration of solute in the saturated solution is the solubility of that
solute in that solvent and at that temperature.
Unsaturated solutions contain less dissolved solute than is needed to form a
saturated solution. Under certain conditions it is possible to prepare a supersaturated
solution. A supersaturated solution contains a greater amount of solute than that needed
to form a saturated solution. One way to prepare a supersaturated solution is to warm the
solution and saturate it at a higher temperature and then to cool it gradually. A
supersaturated solution is unstable and can be made to crystallize by the addition of a
single crystal of the solid.
Factors Affecting Solubility
The solubilities of some gases in water are given in Table 13.2. Note that the
solubility increases with increasing molar mass. The primary interactions between gas
molecules and water molecules in solutions are dispersion forces. Recall that dispersion
forces increase with increasing molar mass. For two gases of roughly equal molar mass
(N2 and CO), the more polar molecule (CO) will be more soluble in water because of a
dipole-dipole component in addition to the solute-solvent attractions.
Liquids that mix with each other in all proportions are termed miscible, and those
that do not mix with one another are immiscible. Polar and ionic substances tend to be
soluble in polar solvents. Nonpolar substances tend to be soluble in nonpolar solvents.
This observation is summarized in the expression, "Like dissolves like."
The solubilities of liquids and solids in water are not appreciably affected by increased
pressure. The solubilities of gases are significantly affected by pressure. The solubility of
a gas is directly proportional to the partial pressure of the gas over the solution. This
relationship is known as Henry's law.
In familiar examples, such as stirring sugar into tea before the ice is added versus
after the ice is added, we recognize that solubility of a solid in water typically increases
with increasing temperature. However, there are some solids that actually become less
soluble at higher temperatures. Sodium hydroxide is an example. What governs the way
increased temperature impacts solubility is whether the dissolution process for the
substance is endothermic or exothermic. When the enthalpy of dissolution is
endothermic, a temperature increase will increase solubility. When the enthalpy of
dissolution is exothermic, a temperature increase will decrease solubility.
The solubilities of all gases decrease at higher temperatures because the
enthalpies of dissolution for all gases are exothermic. If you have ever left a glass of tap
water on the kitchen counter for a period of time, you may have noticed the formation of
bubbles on the inside of the glass. These bubbles are dissolved gases coming out of the
water as the temperature increases.
Ways of Expressing Concentration
Mass percentage is
simply the ratio of solute
mass to total mass, times
100
For very dilute solutions
the concentration might
be expressed in parts per
million (ppm). A 1 ppm
aqueous solution
contains approximately 1
milligram of solute per
liter of solution
For even more dilute
solutions the term parts
per billion (ppb) may be
used. A 1 ppb aqueous
solution contains 1
microgram of solute per
liter of solution
Mole fraction is the ratio
of moles of a solution
component to total moles
of all components in the
solution. Often
symbolized X, mole
fraction has no units
Molarity, is the ratio
of moles of solute to
liters of solution.
Molarity (M) can vary
slightly with temperature
because the density of a
solution can vary slightly
Molality, symbolized
m, is the ratio of moles
of solute to kilograms of
solvent. (For dilute
aqueous solutions,
molarity and molality are
roughly equal.) Molality
does not vary with
temperature, because
mass is not temperaturedependent
Colligative Properties
Certain properties of solutions depend only on the number of dissolved particles
and do not depend on the identity of those particles. Such properties are called colligative
properties.
The vapor pressure of a pure solvent is reduced by the presence of a nonvolatile
solute. A volatile substance is one that escapes readily into the vapor phase. A nonvolatile
substance is one that exerts essentially no vapor pressure. The amount by which the vapor
pressure is lowered depends on the concentration of the solution. Quantitatively, this is
expressed by Raoult's law, which says that the vapor pressure of a solution, PA, is
proportional to the mole fraction of the solvent, XA. PA° is the vapor pressure of the pure
solvent. Just as an ideal gas was one that obeyed the ideal gas law, an ideal solution is
one that obeys Raoult's law.
A solution of a nonvolatile solute will boil at a higher temperature than will pure
solvent. (Salting the water used for cooking increases the temperature at which it boils.)
This can best be understood by recalling that the normal boiling point of a liquid is the
temperature at which its vapor pressure is equal to 1 atm pressure. You just learned that
the vapor pressure is depressed by the presence of a solute. Therefore, it should make
sense that in order for the vapor pressure over a solution to equal atmospheric pressure,
the temperature will have to be higher. A similar argument, although it is not quite as
intuitive, can be made to explain the depression of the freezing point. (Salting the roads
during the winter lowers the freezing point of water and turns icy roads into wet roads.)
When a semipermeable membrane (one through which solvent can move but
solute particles cannot) separates a solution and pure solvent, solvent will flow through
the membrane into the solution. This phenomenon is known as osmosis. The pressure
required to just stop the flow of osmosis is called the osmotic pressure. Osmotic pressure,
denoted , depends on molarity.
Colloids
Colloids, or colloidal dispersions, are suspensions of particles so small that they
do not settle out of the medium in which they are dispersed. Colloid particles range in
size from 10 to 2000 Å. Colloids dispersed in water may be hydrophilic or hydrophobic.
Hydrophilic (water-loving) colloids include suspensions of biologically important
molecules such as enzymes and antibodies in the human body. These giant molecules
have exterior hydrophilic groups that help keep them suspended in water. Such molecules
also have hydrophobic (water-fearing) components, but they fold in such a way as to
prevent the hydrophobic groups from coming into contact with the water.
Hydrophobic colloids can be stabilized in water by adsorption of ions onto their
surfaces or by addition of molecules such as sodium stearate. Sodium stearate has a
hydrophobic end that interacts with the hydrophobic colloidal particle and a hydrophilic
end that can interact with the water. The net effect is a stabilization of the hydrophobic
colloidal suspension.
Distinguishing some colloids from solutions can be difficult–even with a
microscope. One way to tell a colloid from a solution is by using the Tyndall effect. As
small as they are, colloid particles scatter light. (If you've ever driven at night in fog,
you've seen the Tyndall effect.) Dissolved solutes are much smaller– too small to scatter
light.