Contributed by:
The highlights are:
1. Solutions
2. Solvation and Hydration
3. Energy changes and Entropy
4. Saturated Solutions and Solubility
5. Hydrocarbon molecules
6. Henry's law
7. Colligative properties
1.
Chapter 13 – Properties
of Solutions
Jennie L. Borders
2.
Section 13.1 – The Solution Process
• A solution is a homogeneous mixture.
• The solvent is present in the largest
amount.
• The solutes are the other components.
• Aqueous solutions are have water as
the solvent.
• Solutions can be solids, liquids, or
gases.
3.
Solutions
• The ability of substances to form
solutions depends on two factors:
1. The intermolecular forces involved.
2. The tendency of the substance to spread
into larger volumes.
4.
Intermolecular Forces
• Ion-dipole forces are present when ionic
substance dissolve in water.
• Dispersion forces are present when
nonpolar substances form solutions.
• Solutions form when the attractive
forces between the solute and solvent
particles are comparable to the forces
that exist between 2 solute particles or
2 solvent particles.
5.
Solvation and Hydration
• Solvation is the interaction between the
solute and solvent particles.
• Hydration is the interaction between the
solute and solvent particles when the
solvent is water.
6.
Energy Changes
Hsoln = H1 + H2 + H3
Hsoln is the energy change when the
solution forms.
H1 is the energy required to separate the
solute particles (endothermic).
H2 is the energy required to separate the
solvent particles (endothermic).
H3 is the energy released when the solute
and solvent particles attract (exothermic).
8.
Will a solution form?
• Processes that are exothermic tend to
be spontaneous, but endothermic
processes still occur.
Ex: ammonium nitrate spontaneously
dissolves in water even though Hsoln =
26.4 kJ/mol
9.
Entropy
• Spontaneous processes involve two factors:
1. energy (exothermic)
2. tendency to spread out
• Entropy is the degree of randomness or
disorder of a system.
• Processes in which the entropy
(randomness) of the components increases
tend to be spontaneous.
10.
Solutions
• The formation of solutions involves
enthalpy (energy change) and entropy
(randomness).
11.
Sample Exercise 13.1
• The picture below shows the following
reaction:
Na2SO4(s) + 10H2O(g) Na2SO4.10H2O(s)
Essentially all of the water vapor in the
closed container is consumed in this
reaction.
12.
Sample Exercise 13.1 con’t
• If we consider our system to consist
initially of Na2SO4(s) and 10H2O(g)
a. does the system become more or
less ordered in the process?
b. does the entropy of the system
increase or decrease?
13.
Practice Exercise
• Does the entropy of the system
increase or decrease when the
stopcock is opened to allow mixing of
the two gases in the apparatus?
14.
Section 13.2 – Saturated
Solutions and Solubility
• Dissolving and crystallization are
opposite processes.
Solute + solvent solution
15.
Solubility
• The maximum amount of solute that will
dissolve in a given amount of solvent at
a specified temperature is the solubility.
16.
Unsaturated
• An unsaturated solution contains less
than the maximum amount of solute at a
given temperature.
• If more solute is added, then it
will dissolve.
17.
Saturated
• A saturated solution contains the
maximum amount of solute at a given
temperature.
• If more solute is added, then
it will not dissolve and it will
settle at the bottom.
18.
Supersaturated
• A supersaturated solution contains
more than the maximum amount of
solute at a given temperature.
• A supersaturated solution is made by
heating the solvent to a higher
temperature so that more solute can be
dissolved.
• The solution is then carefully cooled so
that crystallization does not occur.
19.
Supersaturated
• If more solute is added, then rapid
crystallization occurs.
• Supersaturated solutions are
extremely unstable.
20.
Section 13.3 – Factors Affecting
Solubility
• The stronger the attractions between
the solute and solvent molecules, the
greater the solubility.
• Polar solvents tend to dissolve polar or
ionic solutes.
• Nonpolar solvents tend to dissolve
nonpolar solutes.
• “Like dissolves like.”
21.
Miscible vs. Immiscible
• Liquids that are miscible will dissolve
one another. Ex: water and rubbing
alcohol
• Liquids that are immiscible will not
dissolve one another. Ex: water and oil
22.
Hydrocarbon Molecules
• C-H bonds are considered nonpolar due
to the small difference in
electronegativity.
• Hydrocarbons tend to be insoluble in
polar solvents.
• One way to increase the solubility is to
increase the number of polar groups.
Ex: an alcohol group –OH.
23.
Alcohols
• As the carbon chain of an alcohol
increases, the solubility decreases due
to the increased London dispersion
forces. The molecule starts to act more
nonpolar.
24.
Network Solids
• Network solids are not soluble in polar
or nonpolar substances due to the
strong bonding forces within the solid.
25.
Sample Exercise 13.2
• Predict whether each of the following
substances is more likely to dissolve in
the nonpolar solvent carbon
tetrachloride or water: C7H16, Na2SO4,
HCl, and I2.
26.
Practice Exercise
• Arrange the following substances in
order of increasing solubility in water:
27.
Pressure
• The solubilities of liquids and solids are
not affected by pressure, but the
solubility of a gas is affected.
• The solubility of a gas increases when
pressure is increased.
28.
Henry’s Law
• Henry’s Law shows the relationship
between pressure and solubility of a gas:
Sg = kPg
Sg = solubility of a gas (M = molarity)
k = Henry’s law constant (different for each
solute-solvent pair and changes with
temperature) (mol/L.atm)
Pg = partial pressure of the gas (atm)
29.
Sample Exercise 13.3
• Calculate the concentration of CO2 in a
soft drink that is bottled with a partial
pressure of CO2 of 4.0 atm over the
liquid at 25oC. The Henry’s law constant
for CO2 in water at this temperature is
3.1 x 10-2 mol/L.atm.
30.
Practice Exercise
• Calculate the concentration of CO2 in a
soft drink after the bottle is opened and
equilibrates at 25oC under a CO2 partial
pressure of 3.0 x 10-4 atm.
31.
Temperature
• The solubility of most solids increases
as temperature increases.
• The solubility of most gases decreases
with increasing temperature.
32.
Section 13.4 – Ways of
Expressing Concentration
• We do not need to cover parts per
million (ppm), parts per billion (ppb),
mass percent, molality, or normality.
33.
Mole Fraction and Molarity
• Mole Fraction = mole of component
total moles
• Molarity = moles of solute
liters of solution
34.
Molarity
• The molarity of a solution does change
with temperature because the
contraction or expansion of the solution
changes the volume.
35.
Sample Exercise 13.6
• An aqueous solution of hydrochloric
acid contains 36% HCl by mass.
Calculate the mole fraction of HCl in the
solution.
36.
Practice Exercise
• A commercial bleach solution contains
3.62 mass % NaOCl in water. Calculate
the mole fraction of NaOCl in the
solution.
37.
Section 13.5 – Colligative
Properties
• Colligative properties are properties of a
solvent that change when solute
particles are added.
• Colligative properties are only based on
the number of solute particles, not the
identity of the solute.
38.
Vapor Pressure
• The vapor pressure of a substance
lowers when solute particles are added
because the solvent particles are
attracted to the solute particles.
39.
Boiling Point
• The boiling point of a solvent elevates
when solute particles are added
because the solvent particles are
attracted to the solute particles.
40.
Freezing Point
• The freezing point of a solvent
depresses (lowers) when solute
particles are added because the solvent
particles are attracted to the solute
particles.
41.
Osmosis
• Materials that only allow some
molecules to pass through are called
semipermeable.
• For example, a membrane may allow
small water molecules to pass through
but not larger solute particles.
• Osmosis is the net movement of solvent
toward the solution with the
higher solute concentration.
42.
Colligative Properties
• Colligative properties are considered
prior knowledge and should not be
directly assessed on the AP exam, so I
have just provided a brief overview to
make sure that you are familiar with
them.
43.
Section 13.6 - Colloids
• Solutions contain very small solute
particles.
• Colloids contain intermediate (medium-
sized) solute particles. Colloid particles
are between 5 to 1000 nm.
• Suspension contain large solute
particles that settle over time.
44.
Colloids
• Colloids particles are small enough to
seem uniform, but they are large
enough to scatter light.
• The scattering of light by colloid
particles is called the Tyndall effect.
45.
Hydrophilic and Hydrophobic
• Hydrophilic colloids are “water loving.”
• Hydrophobic colloids are “water
fearing.”
46.
Adsorption vs. Absorption
• Adsorption means to adhere to a
surface.
• Absorption means to pass into the
interior.
47.
Biology … INFERIOR!
• In the body, hydrophobic colloids can
mix with water by adsorbing ions onto
the surface.
48.
Removal of Colloid Particles
• Colloid particles are too small to be
removed by filtration.
• Coagulation is a process of enlarging
the colloid particles so that they can be
removed.
• Heating a colloid or adding electrolytes
can bring about coagulation.
50.
Sample Integrative Exercise
• A 0.100L solution is made by dissolving
0.441g of CaCl2(s) in water. The enthalpy
of solution for CaCl2 is H = -81.3
kJ/mol. If the final temperature of the
solution is 27.0oC, what was the initial
temperature? (Assume that the density
of the solution is 1.00 g/mL, that its
specific heat is 4.18 J/g.K, and that the
solution loses no heat to its
surroundings.)
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