Spencer L. Seager
Michael R. Slabaugh
www.cengage.com/chemistry/seager
Chapter 7:
Solutions and Colloids
Jennifer P. Harris
LEARNING OBJECTIVES/ASSESSMENT
When you have completed your study of this chapter, you should be able to:
1. Classify mixtures as solutions or nonsolutions based on their appearance. (Section 7.1; Exercise 7.4)
2. Demonstrate your understanding of terms related to the solubility of solutes in solution. (Section 7.2;
Exercises 7.6 and 7.12)
3. Predict in a general way the solubilities of solutes in solvents on the basis of molecular polarity.
(Section 7.3; Exercise 7.16)
4. Calculate solution concentrations in units of molarity, weight/weight percent, weight/volume
percent, and volume/volume percent. (Section 7.4; Exercises 7.22 b, 7.30 c, 7.34 a, and 7.38 c)
5. Describe how to prepare solutions of specific concentration using pure solutes and solvent, or
solutions of greater concentration than the one desired. (Section 7.5; Exercises 7.46 and 7.48 b)
6. Do stoichiometric calculations based on solution concentrations. (Section 7.6; Exercise 7.56)
7. Understand colligative solution properties of boiling point, freezing point, and
osmotic pressure and how to determine osmolarity; (Section 7.7; Exercises 7.64 a & c and 7.74)
8. Describe the characteristics of colloids. (Section 7.8; Exercise 7.82)
9. Describe the process of dialysis, and compare it to the process of osmosis. (Section 7.9; Exercise 7.84)
Mixtures
• Solutions (sometimes called true solutions) are homogeneous
mixtures of two or more substances in which the components are present
as atoms, molecules, or ions.
• Particles in these solutions are:
• too small to reflect light, are transparent (always clear but can be
colored).
• in constant motion and not settled by the influence of gravity.
• Heterogeneous mixtures
• Particles in these “solutions” are:
• Large aggregates, reflect light (and are turbid) and do settle out
due to gravity
• Colloidal solutions
• Particles in these solutions are:
• Large aggregates, reflect light (and are turbid) but do not settle
out due to gravity (homogeneous mixture)
The dissolving process
• For true solutions
SOLUTION TERMINOLOGY
• Solutions can be solids, liquids or gases but we will deal
primarily with liquid solutions.
• Solutions are composed of a solvent (main component) and one
or more solutes (substance that is dissolved in the solvent)
• Solubility – the degree to which a solute dissolves (insoluble,
slightly soluble, soluble, very soluble)
• Miscible/immiscible – terms that describe whether a liquid solute
dissolves or not
• Saturated, unsaturated, supersaturated
SOLUBILITY
• The solubility of a solute is the maximum amount of the solute
that can be dissolved in a specific amount of solvent under
specific conditions of temperature and pressure.
EXAMPLES OF SOLUTE SOLUBILITES
AT 0°C
EXAMPLES OF SOLUTE SOLUBILITES
AT 0°C (continued)
EFFECT OF TEMPERATURE ON
SOLUBILITY
THE SOLUTION PROCESS
• The solution process involves interactions between solvent
molecules (often water) and the particles of solute.
• An example of the solution process for an ionic solute in water:
THE SOLUTION PROCESS (continued)
• An example of the solution process for a polar solute in water:
THE SOLUTION PROCESS (continued)
• A solute will not dissolve in a solvent if:
• the forces between solute particles are too strong to be
overcome by interactions with solvent particles.
• the solvent particles are more strongly attracted to each
other than to solute particles.
• A good rule of thumb for solubility is “like dissolves like.”
• Polar solvents dissolve polar or ionic solutes.
• Nonpolar solvents dissolve nonpolar or nonionic solutes.
INCREASING THE RATE OF DISSOLVING
• Crush or grind the solute.
• Small particles provide more surface area for solvent attack
and dissolve more rapidly than larger particles.
• Heat the solvent.
• Solvent molecules move faster and have more frequent
collisions with solute at higher temperatures.
• Stir or agitate the solution.
• Stirring removes locally saturated solution
from the vicinity of the solute and allows
unsaturated solvent to take its place.
HEAT AND SOLUTION FORMATION
• Endothermic: Solute
• Exothermic:
+ solvent + heat→ solution
Solute + solvent → solution + heat
SOLUTION CONCENTRATIONS
• Solution concentrations express a quantitative relationship
about the amount of solute contained in a specific amount of
solution.
• Concentration units discussed include molarity and
percentage.
MOLARITY
• The molarity of a solution expresses the number of moles of
solute contained in one liter of solution.
• The mathematical calculation of the molarity of a solution
involves the use of the following equation:
moles of solute
M
liters of solution
• In this equation, the number of moles of solute in a sample of
solution is divided by the volume in liters of the same sample of
solution.
PERCENT CONCENTRATIONS
• Percent concentrations express the amount of solute
contained in 100 parts of solution. The parts of solution may
be expressed in different units.
part
%
 100
total
• Three variations exist for % concentrations. W/W, W/V or V/V
generally expressed in grams and mL; if not then at least the
same units.
• To make solutions of M, W/V% and V/V%, the amount of
solute is measured out and solvent added to the volume
needed.
• To make solutions of W/W% , the grams of solvent and solute
are added to determine the weight of the solution.
CONCENTRATON CALCULATIONS
• Example 1: A 250-mL sample of solution contains 0.134
moles of solute. Calculate the molarity of the solution.
• 9.45 g of methyl alcohol, CH3OH, was dissolved in enough
pure water to give 500 mL of solution. What was the molarity
of the solution?
• Calculate the %(w/w) of a solution prepared by dissolving
15.0 grams of table sugar in 100 mL of water. The density of
the water is 1.00 g/mL.
• Calculate the %(w/v) of a solution prepared by dissolving 8.95
grams of sodium chloride in enough water to give 50.0 mL of
solution.
• A solution is made by dissolving 250 mL of glycerin in enough
water to give 1.50 L of solution. Calculate the %(v/v) of the
resulting solution
SOLUTION PREPARATION
• Solutions of known concentration are usually prepared in one
of two ways.
• In one method, the necessary quantity of pure solute is
measured using a balance or volumetric equipment. The solute
is put into a container and solvent, usually water, is added until
the desired volume of solution is obtained.
SOLUTION PREPARATION EXAMPLE
• Calculation example: Describe how to prepare 500 mL of
0.250 M NaCl solution.
• Solution: The mass of NaCl needed must first be
determined. The volume and concentration of the desired
solution are known, so the equation for molarity is rearranged
to solve for the number of moles of solute needed. The result
is:
moles of solute = M x liters of solution
= 0.250 M x 0.500 L = 0.125 mole
• Thus, 0.125 moles of NaCl is needed. NaCl has a formula
weight of 58.4 u, so 0.125 moles has a mass of 0.125 x 58.4g
or 7.30 grams. The solution is prepared by weighing a
sample of NaCl with a mass of 7.30 grams. The sample is put
into a 500 mL volumetric flask and pure water is added up to
the mark on the flask.
SOLUTION PREPARATION (continued)
• In a second method, a quantity of solution with
a concentration greater than the desired
concentration is diluted with an appropriate
amount of solvent to give a solution with a
lower concentration. This type of problem is
made simpler by using the following equation:
(Cc)(Vc) = (Cd)(Vd)
• In this equation, Cc is the concentration of the
concentrated solution that is to be diluted, Vc
is the volume of concentrated solution that is
needed, Cd is the concentration of the dilute
solution, and Vd is the volume of dilute
solution.
SOLUTION PREPARATION EXAMPLE
• Calculation example: Describe how to prepare 250 mL of 0.500
M HCl solution from a 1.50 M HCl solution.
• Solution: According to the definitions given above, Cc = 1.50 M,
Cd = 0.500 M, and Vd = 250 mL. The equation given above can
be solved for Vc, the volume of concentrated solution needed:
Vc

C d Vd  0.500 M250 mL 


 83.3 mL
Cc 
1.50 M
• The solution is prepared by measuring 83.3 mL of 1.50 M HCl and
pouring it into a 250 mL volumetric flask. Pure water is then added
up to the mark on the flask to give 250 mL of 0.500 M solution.
SOLUTION STOICHIOMETRY
• As shown earlier, the number of moles of solute in a volume of
solution of known molarity can be obtained by multiplying
together the known molarity and the solution volume in liters.
• Molarity is a ratio of moles of solute to liters of solution. This
ratio can be written as two conversion factors:
moles solute
liters solution
or
liters solution
moles solute
• The conversion factor on the left is used to multiply by the
molarity. It is selected to cancel the units of liters of solution
and obtain the units of moles of solute.
• The conversion factor on the right is used to divide by the
molarity. It is selected to cancel the units of moles of solute
and obtain the units of liters of solution.
SOLUTION STOICHIOMETRY EXAMPLE
• Calculation example: Consider the balanced equation
HCl(aq) + NaOH(aq)
NaCl(aq) + H2O(l)
How many mL of 0.100 M HCl solution would exactly react
with 25.00 mL of 0.125 M NaOH solution?
SOLUTION PROPERTIES
• Absolutely pure water conducts electricity very poorly.
• Some solutes called electrolytes produce water solutions that
conduct electricity well.
• Some solutes called nonelectrolytes produce water solutions
that do not conduct electricity.
A solution of a
strong electrolyte
conducts electricity
well.
A solution of a weak
electrolyte conducts
electricity poorly.
A solution of a
nonelectrolyte does not
conduct electricity.
ELECTROLYTES
• STRONG ELECTROLYTES
• Strong electrolytes form solutions that
conduct electricity because they
dissociate completely into charged ions
when they dissolve.
• WEAK ELECTROLYTES
• Weak electrolytes form weakly
conducting solutions because they
dissociate into ions only slightly when
they dissolve.
• NONELECTROLYTES
• Nonelectrolytes form nonconducting
solutions because they do not
dissociate into ions at all when they
dissolve.
COLLIGATIVE PROPERTIES OF SOLUTIONS
• Colligative solution properties are properties that depend
only on the concentration of solute particles in the solution.
Three colligative properties are boiling point, freezing point,
and osmotic pressure.
• Experiments demonstrate that the vapor pressure of water
(solvent) above a solution is lower than the vapor pressure of
pure water.
SOLUTION BOILING POINT
• The boiling point of a solution is always higher than the
boiling point of the pure solvent of the solution.
• Since BP elevation is a colligative property, dependent only
on the concentration of particles, the following substances
would have an equal impact on BP elevation (or any other
colligative property)
• 0.1 M CaCl2
• 0.15 M NaCl
• 0.3 M sugar (non-electrolyte)
SOLUTION BOILING POINT (continued)
• For example, the dissociation of calcium chloride is represented
as:
CaCl2
Ca2+ + 2 Cl• Thus, when 1 mole of CaCl2 dissolves, 3 moles of particles (ions)
are put into the solution.
SOLUTION FREEZING POINT
• The freezing point of a solution is always lower than the
freezing point of the pure solvent of the solution.
OSMOTIC PRESSURE OF SOLUTIONS
• When solutions having different concentrations
of solute are separated by a semipermeable
membrane, solvent tends to flow through the
membrane from the less concentrated solution
into the more concentrated solution in a process
called osmosis.
• When the more concentrated solution involved
in osmosis is put under sufficient pressure, the
net osmotic flow of solvent into the solution can
be stopped.
• The pressure necessary to prevent the osmotic
flow of solvent into a solution is called the
osmotic pressure of the solution and can be
calculated by using the following equation, which
is similar to the ideal gas law given earlier:
π = nMRT
OSMOTIC PRESSURE OF SOLUTIONS
(continued)
• In this equation, π is the osmotic pressure, n is the number of
moles of solute particles put into solution when 1 mole of
solute dissolves, M is the molarity of the solution, R is the
universal gas constant written as 62.4 L torr/K mol, and T is the
solution temperature in Kelvin.
• The product of n and M is called the osmolarity of the solution.
COLLOIDS
• Colloids are homogeneous mixtures of two or more
components called the dispersing medium and the
dispersed phase. The dispersed phase substances in a
colloid are in the form of particles larger than those found in
solutions.
• DISPERSING MEDIUM OF A COLLOID
• The dispersing medium of a colloid is the substance
present in the largest amount. It is analogous to the
solvent of a solution.
• DISPERSED PHASE OF A COLLOID
• The dispersed phase of a colloid is the substance
present in a smaller amount than the dispersing
medium. It is analogous to the solute of a solution.
COLLOID PROPERTIES
• In colloids, the dispersed phase particles cannot be seen and
do not settle under the influence of gravity.
• Colloids appear to be cloudy because the larger particles in the
dispersed phase scatter light.
• Colloids demonstrate the Tyndall effect in which the path of the
light through a colloid is visible because the light is scattered.
TYPES OF COLLOIDS
STABILIZING COLLOIDS
• Substances known as emulsifying agents or stabilizing
agents are used to prevent some colloids from coalescing
(e.g. egg yolk in oil and water to form mayonnaise, soap/
detergent ions forming a charged layer around nonpolar oils
and greases).
STABILIZING COLLOIDS (continued)
DIALYSIS
• Dialysis can be used to separate small particles from
colloids (e.g. cleaning the blood of people suffering from
kidney malfunction).
DIALYSIS
• A dialyzing membrane is a semipermeable membrane with
larger pores than osmotic membranes that allow solvent
molecules, other small molecules, and hydrated ions to pass
through.
• Dialysis is a process in which solvent molecules, other small
molecules, and hydrated ions pass from a solution through a
membrane.