Chapter 12: Solutions
I.
Types of Mixtures
A. Solutions
1. Solution: a homogeneous mixture of two or more substances in a
single phase.
2. Soluble: capable of being dissolved.
3. A solution has two components:
a. Solute: substance dissolved in a solution.
b. Solvent: dissolving medium in a solution.
4. Solutions exist as gases, liquids, and solids. ALL are mixtures of
one or two of these.
B. Suspensions
1. Suspension: particles of a solute are so large that they settle out
unless the mixture is constantly stirred or agitated.
2. For example, muddy water.
C. Colloids
1. Colloid: Particles that are intermediate in size between those in
solutions and suspensions.
2. For example, muddy water after big particles are removed, still
cloudy.
3. Colloids exhibit the Tyndall effect, a scattering of light.
3. Several classes. See Tables 2 and 3 on page 404.
D. Solutes: Electrolytes vs. Nonelectrolytes
1. Electrolyte: A substance that dissolves in water to give a solution
that conducts electric current.
2. Nonelectrolyte is the opposite.
E.
Ionization
1. Ionization is process in which ions are formed from solute
molecules by the action of the solvent.
2. Note ionization is different from dissociation. Ionization is the
creation of ions where there were none. When dissociation occurs
the ions were already present.
3. The Hydronium Ion
a. H+ ions attract molecules so strongly that they don’t normally
exist alone.
b.
When HCl is mixed with water, it dissociates. The resulting H+
attaches to the water molecule, making H3O+. This is the
hydronium ion.
C. Strong and Weak Electrolytes
1. Any compound of which all or almost all of the dissolved compound
exists as ions in an aqueous solution is called a strong
electrolyte.
a. The distinguishing factor of strong electrolytes is that, to
whatever extent they dissolve in water, they yield only ions.
b. These ions conduct electricity.
c. For example, NaCl.
2. A weak electrolyte is a compound of which a relatively small
amount of the dissolved compound exists as ions in an aqueous
solution.
a. Many times the dissociation occurs, but the reverse reaction
also occurs, making less ions present, therefore not as much
conductivity.
b. For example, HF, and acetic acid.
3. Strong and weak electrolytes differ in the degree of ionization or
dissociation. Concentrated and dilute solutions differ in the amount
of solute dissolved in a given amount of solvent.
a. HCl: Strong electrolyte, no matter the conc. (0.00001 M)
b. Acetic Acid: weak electrolyte, no matter the conc. (10M).
c. Refer to Figure 5 on page 442.
II. The Solution Process
A. Factors Affecting the Rate of Dissolution
1. Increasing the surface area of solute.
2. Agitation (stirring).
3. Heating solvent.
B. Solubility
1. For every combination of solute and solvent at a given temperature,
there is a limit to the amount of solute that can be dissolved. once
beyond this point no more dissolves under those conditions.
2. A solution that contains the maximum amount of solute is said to be
saturated.
3. A solution that contains less than the maximum amount of solute is
said to be unsaturated.
4. Supersaturated: a solution that contains more dissolved solute
than a saturated solution contains under the same conditions.
a.
5.
Made by dissolving a large amount of solute in hot water, and
then allowing the solution to cool slowly and undisturbed.
b. Seeding or disturbing the solution will result in rapid formation
of crystals of the excess solute.
The solubility of a substance is the amount of that substance
required to form a saturated solution with a specific amount of
solvent at a specific temperature.
C. Solute-Solvent Interactions
1. Always remember “Like dissolves like”. Polar solvents will dissolve
polar solutes, but usually not nonpolar solutes.
2. When ionic crystals dissolve in water, they break up into their
component ions. This is called hydration.
3. Some ionic substances trap water when they crystallize. They are
called hydrates.
4. Ionic compounds are generally insoluble in nonpolar solvents.
5. Liquid solutes and solvents
a. Liquids that dissolve freely in one another in any proportion are
said to be completely miscible. For example, benzene and
CCl4.
b. Liquid solutes and solvents that are not soluble in each other
are immiscible. For example, water and oil.
6. Increases in pressure increase gas solubility in liquids, but have no
significant effect on other solubilities.
a. Henry’s Law: the solubility of a gas in a liquid is directly
proportional to the partial pressure of that gas on the surface of
the liquid.
b. Effervescence: rapid escape of gas from a liquid in which it is
dissolved.
7. Effects of Temperature on Solubility
a. Increase of T decreases gas solubility.
b. Generally increase of T increases solid in liquid solubility. The
degree varies. See Figure 15 on page 414.
D. Enthalpies of Solution
1. The formation of a solution is accompanied by an energy change.
Refer to Figure 16 on page 415.
2. A solute particle that is surrounded by solvent molecules is said to
be solvated.
3. Enthalpy of Solution: net amount of energy absorbed as heat by
the solution when a specific amount of solute dissolves in a solvent.
III. Concentration of Solutions
A. Concentration
1. The concentration of a solution is a measure of the amount of
solute in a given amount of solvent or solution.
2. We will discuss 2 ways to express concentration: Molarity and
Molality.
B. Molarity
1. Molarity is the number of moles of solute in one liter of solution.
2. Molarity (M) = amount of solute (mol)/volume of solution (L).
3. A 1 M solution of NaOH would contain 1 mol of NaOH (40.0g) in 1 L
of solution.
4. Note that a 1 M solution is not made by adding 1 mol of solute to 1
L of solvent. The total volume when all is done is 1 L.
5. Practice, practice, practice.
C. Molality
1. Molality is the concentration of a solution expressed in moles of
solute per kg of solvent.
2. Molality (m) = moles solute / mass solvent.
3. A 1 m solution of NaOH would contain 1 mol of NaOH (40.0g) and
1 kg of water.
4. Molality is used when studying properties of solutions related to
temperature changes because it does not change with changes in
temperature.
5. Practice, Practice, Practice.
IV. Colligative Properties
A. Colligative properties are properties that depend on the
concentration of solute particles but not on their identity.
B. Freezing Point Depression
1. The freezing point depression (tf) is the difference between the
freezing points of the pure solvent and a solution of a nonelectrolyte
in that solvent, it is directly proportional to the molality.
2. For any concentration of a nonelectolyte solute in water, the
decrease in FP can be determined by using the value of -1.86oC/m.
3. This value is the molal FP constant (Kf) for water.
4. tf = Kfm. To get the FP of the impure solution, determine the FP
of the pure solution and insert the FP depression.
5. Practice, Practice, Practice.
C. Boiling Point Elevation
1.
2.
3.
4.
The boiling point elevation (tb) is the difference between the
boiling points of the pure solvent and a solution of a nonelectrolyte
in that solvent, it is directly proportional to the molality.
The molal BP constant (Kb) for water is 0.51oC/m
tb = Kbm. To get the BP of the impure solution, determine the BP
of the pure solution and insert the BP elevation.
Practice, Practice, Practice.
D. Solutions with electrolyte solutes
1. Freezing point depression and boiling point elevation values
depend on the concentration of solute particles not on their
identity.
2. When electrolytes dissolve in water, they break apart into their
component ions. All of these ions affect the freezing/boiling point.
3. When doing problems involving electrolytes, multiply or divide the
concentration by the number of ions present. Which you do will
depend on the problem.