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Chapter 15 Solutions
I. What are Solutions
A. Solution - homogeneous mixture made up of individual molecules, atoms or ions.
B. Solute - the substance being ________________________
C. Solvent - the substance ________________________
D. Soluble - substance that _________________________________ in a solvent
E. Insoluble - substance that _____________________________________ in a solvent
F. Immiscible – liquids that are __________________________________ in each other
G. Miscible – two liquids that ___________________________________ in each other
H. Types of solutions
1. Aqueous solution ____________________ is the solvent
2. Tincture - ____________ is the solvent
3. Alloys - solid solution of two or more
metals
- _____________ - an alloy in which one
metal is Hg
II. Electrolytes vs. Nonelectroyles
A. Electrolytes: any substance that,
when it is dissolved in a solution, will
conduct an electric current by means of movement of ions
NaCl(aq) _____
6M HCl(aq)______
1M HCl(aq)___________
Glacial Acetic Acid(aq) _____6M HC2H3O2(aq) ______1M HC2H3O2(aq) _____
NaOH(aq) ____
tap H2O _____
distilled H2O _____
 Strong Electrolytes - substance that dissolves almost completely in water
to produce many ions to conduct electricity
1. NaCl(s)
2. HCl(g)
3. NaOH(s)
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 Weak (or non) electrolyte - poor conductor of electricity
1. HC2H3O2 (aq)
2. H2 O (l)
III. Heterogeneous Mixtures
1. Suspension (NOT A SOLUTION) – a mixture containing particles that remain
while thoroughly mixed, but then settles out when left undisturbed.
 Can you think of examples of a suspension?
2. Colloids (NOT A SOLUTION) –Colloids or colloidal dispersions are
suspensions in which the suspended particles are larger than molecules but
too small to separate out of the suspension due to gravity.
Particle size: 10 to 2000 Å.
 Can you think of examples colloids?
3. Tyndall Effect Dispersed colloid particles that are large enough to scatter
light. The path of a beam of light projected through a colloidal suspension can
be seen through the suspension.
 Watch and observe the demo?
Hydrophilic and Hydrophobic Colloids
Water-loving colloids: hydrophilic.
Water-hating colloids: hydrophobic.
• In the human body, large biological
molecules such as proteins are kept in
suspension by association with surrounding
water molecules.
• These macromolecules fold up so that hydrophobic groups are away from the
water (inside the folded molecule).
• Hydrophilic groups are on the surface of these molecules and interact with solvent
(water) molecules.
•Typical hydrophilic groups are polar (containing C–O, O–H, N–H bonds) or charged.
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• Consider a small drop of oil in water. They will not mix unless soaps act on the oil
and water.
• Soaps are molecules with long hydrophobic tails and hydrophilic heads that
remove dirt by stabilizing the colloid in water.
• Most dirt stains on people and clothing are oil-based.
• Biological application of this principle:
• The gallbladder excretes a fluid called bile.
• Bile contains substances (bile salts) that form an emulsion with fats in our small
intestine.
• Emulsifying agents help form an emulsion.
• Emulsification of dietary fats and fat-soluble vitamins is important in their
absorption and digestion by the body.
IV. The Solutions Process
 Can you draw how you think a salt crystal dissolves? What do the ions of NaCl and
molecules of water look like? How do they interact?
Consider NaCl (solute) dissolving in water (solvent). Water molecules
orient themselves on the NaCl crystals.
H-bonds between the water molecules have to be broken.
NaCl dissociates into Na+ and Cl–.
Ion-dipole forces form between the Na+ and the negative end of the water
dipole.
Similar ion-dipole interactions form between the Cl– and the positive end of
the water dipole.
.

The process by which the charged particles in an ionic solid
separate from one another is called _____________________________
Dissociation Equation:
NaCl (s)
 Solvation – the process of surrounding solute particles with solvent particles to
form a solution. If water is the solvent, the interaction is called hydration.
Factors that affect rate of solvation – must increase collisions between solute and
solvent
1. ____________________________________
2. ____________________________________
3. ____________________________________
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Solubility
 measure the maximum amount of solute that can be dissolved in a given amount of
solvent at a specific temperature and pressure.
 usually expressed a quantity of solute per definite amount of solvent at a
particular temperature (ex. 36g NaCl/100g H2O at 20 C)
Unsaturated
Saturated
 the amount of solute
 solution holds the maximum
dissolved is less than the
maximum that could be
dissolved
amount of solute per
amount of the solution
under the given conditions
Supersaturated
 solutions contain more solute

than the usual maximum amount
and are unstable
They cannot permanently hold
the excess solute in solution and
may release it suddenly
Solution Equilibrium is the physical state in which the opposing processes of dissolving and
crystallization of a solute occur at equal rates.
 Graphs and solubility tables are useful planning tools for making solutions.
 Can you draw the curve of a graph as a salt increases in solubility as temperature
increases?
Factors affecting solubility
The tendency of a substance to dissolve in another depends on:
1. Nature of the solute and solvent
example: at room temp, 1g of PbCl2/100g H2O and 200g
ZnCl2/100g H2O
 Which solute is more soluble and why, can you explain?
Intermolecular forces are an important factor in determining
solubility of a solute in a solvent.
The stronger the attraction between solute and solvent molecules, the
greater the solubility.
For example, polar liquids tend to dissolve in polar solvents.
Favorable dipole-dipole interactions exist (solute-solute, solvent-solvent, and solutesolvent).
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Water and ethanol are miscible because the broken hydrogen bonds in both pure liquids
are re-established in the mixture.
However, not all alcohols are miscible with water.
Why? The number of carbon atoms in a chain affects solubility.
The greater the number of carbons in the chain, the more the molecule behaves like a
hydrocarbon.
Thus, the more C atoms in the alcohol, the lower its solubility in water.
Increasing the number of –OH groups
within a molecule increases its
solubility in water.
The greater the number of –OH
groups along the chain, the more
solute-water H-bonding is possible.
Generalization: “Like dissolves like”. Substances with similar intermolecular
attractive forces tend to be soluble in one another.
• The more polar bonds in the molecule, the better it dissolves in a polar solvent.
•The less polar the molecule the less likely it is to dissolve in a polar solvent and the
more likely it is to dissolve in a nonpolar solvent.
• Network solids do not dissolve because the strong intermolecular forces in the
solid are not reestablished in any solution.
2. Pressure
The solubility of a gas in a liquid is a function of the pressure of the
gas over the solution.
a. solid in liquid?
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b. gas in liquid?
 What happen to when the pressure goes up?
c. Henry’s Law  At a given temperature the solubility (S) of a gas in a
liquid is directly proportional to the pressure (P) of
the gas above the liquid.
S=
P=
3. Temperature ___ temp, ____ kinetic energy, ____ randomness
a. solid in liquid?
what if cools back down?
b. gas in liquid?
 What happen to a gas as the liquid warms up?
 Heat of Solution: overall energy change that occurs during the solution process
Energy Changes and Solution Formation
There are three steps involving energy in the formation
of a solution:
•Separation of solute molecules (ΔH1),
•Separation of solvent molecules (ΔH2), and
•Formation of solute-solvent interactions (ΔH3).
•
•
•
•
•
•We define the enthalpy change in the solution process
as:
ΔHsoln = ΔH1 + ΔH2 +ΔH3
To determine whether ΔHsoln is positive or negative, we
consider the strengths of all solute-solute, solventsolvent and solute-solvent interactions:
Breaking attractive intermolecular forces is always
endothermic.
ΔH1 and ΔH2 are both positive.
Forming attractive intermolecular forces is always exothermic.
ΔH3 is always negative.
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It is possible to have either ΔH3 > (ΔH1
+ ΔH2) or ΔH3 < (ΔH1 + ΔH2).
Examples:
MgSO4 added to water has ΔHsoln = –
91.2 kJ/mol.
NH4NO3 added to water has ΔHsoln = +
26.4 kJ/mol.
MgSO4 is often used in instant heat
packs and NH4NO3 is often used in
instant cold packs.
How can we predict if a solution will form?
•In general, solutions form if the ΔHsoln is negative.
•If ΔHsoln is too endothermic a solution will not form.
•“Rule of thumb”: Polar solvents dissolve polar solutes.
•Nonpolar solvents dissolve nonpolar solutes.
•Consider the process of mixing NaCl in gasoline.
•Only weak interactions are possible because gasoline is nonpolar.
•These interactions do not compensate for the separation of ions from one another.
•Result: NaCl doesn't dissolve to any great extent in gasoline.
•Consider the process of mixing water in octane (C8H18).
•Water has strong H-bonds.
•The energy required to break these H-bonds is not compensated for by
interactions between water and octane.
•Result: Water and octane do not mix.
Factors affecting rate of solution
1. size of particle
a. dissolving only occurs at surface
b. smaller pieces, ___ surface area, ___ rate of solution
2. stirring
a. brings fresh parts of solvent in contact with solute, so ___ rate of sol
3. amount of solute already dissolved
a. the more solute dissolved, ______ rate of solution
4. temperature
a. solids: _____ temp, _____ rate of solution
b. gases: _____ temp, _____ rate of solution
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Ways of Expressing Concentration
• All methods involve quantifying the amount of solute per amount of solvent (or solution).
• Concentration may be expressed qualitatively or quantitatively.
• The terms dilute and concentrated are qualitative ways to describe concentration.
• A dilute solution has a relatively small concentration of solute.
• A concentrated solution has a relatively high concentration of solute.
• Quantitative expressions of concentration require specific information regarding such
quantities as masses, moles, or liters of the solute, solvent, or solution.
Solution
Concentrations
what does
concentration
mean?
Mass percentage, ppm, and ppb
• Mass percentage is one of the simplest ways to express concentration.
Mass % of component
mass of component in soln
100
total mass of solution
•By definition:
• Similarly, parts per million (ppm) can be expressed as the number of mg of solute per
kilogram of solution.
Parts per million (ppm) of component
mass of component in soln
106
total mass of solution
• If the density of the solution is 1g/ml, then 1 ppm = 1 mg solute per liter of solution.
• We can extend this again!
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• Parts per billion (ppb) can be expressed as the number of µg of solute per kilogram of solution.
Parts per billion (ppb) of component
mass of component in soln
109
total mass of solution
• If the density of the solution is 1g/ml, then 1 ppb = 1 µg solute per liter of solution.
Mole Fraction, Molarity, and Molality
• Common expressions of concentration are based on the number of moles of one or more
components.
• Recall that mass can be converted to moles using the molar mass.
• Recall:
Mole fraction of component , X
moles of component
total moles of all components
•
Note: Mole fraction has no units.
•
Note: Mole fractions range from 0 to 1.
• Recall:
M (molarity)
=
# of moles of solute
1 Liter of solution
example: What is the molarity of a solution which contains 4.9 g of H2SO4 in 250 ml of
solution?
example #2: How many moles of KNO3 are contained in 300 ml of a 0.50 M solution?
• Note: Molarity will change with a change in temperature (as the solution volume increases or
decreases).
• We can define molality (m), yet another concentration unit:
Molality, m
moles of solute
kilograms of solvent
• Molality does not vary with temperature.
• Note: converting between molarity (M) and molality
(m) requires density.
Colligative Properties:
Those people who live in the colder climates of the
world add antifreeze to the water in their automobiles
radiators. From the name of the product, it is obvious
that its purpose is to prevent the water from freezing.
How does the antifreeze prevent freezing? It is not that
the antifreeze itself has such a low freezing point. The
best protection is afforded by a mixture that is roughly
50% antifreeze and 50% water. Next we will investigate
the behavior of solutions in such situations as freezing
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and boiling. We will look at those properties of solvents that are changed by solutes.
Properties which are determined by the number of particles in solution rather than by the type of
particle in solution, are colligative properties.
Freezing point depression and boiling point elevation
A. Vapor Pressure Lowering:
1. VP: the pressure exerted in a closed container by liquid particles that have escaped the liquid’s
surface and have entered the gaseous state.
2. Vapor pressure for water shown in first beaker, a solute lowers the vapor pressure in the
second beaker.
3. There are fewer solvent particles on the surface of the second beaker to enter the gaseous
state and the vapor pressure is lowered.
4. Remember Boiling Point is defined as:
B. Boiling point elevation and freezing point depressions depend on:
1. molal concentration - number of moles of solute in a kilogram of solvent
2. nature of the solvent
 Can you think of examples of a BP elevation or FP depression?
Example: #2 What would lower the freezing point better on the roads, sugar
(C12H22O11), table salt (NaCl), another salt CaCl2, and why?
C. Molal freezing and boiling point constants
1. A 1 molal solution will lower the freezing point of water 1.86 C. Water
could freeze at -1.86 C instead of 0 C. This value is called the molal freezing
point constant for water (1.86 C/molal).
Tf = change in freezing point
Tf = kfm
m
kf = constant for water
= molal concentration of
solution
2. A 1
boiling
Tb = Kbm
molal solution will raise the
point of water 0.52 C.
Examples: If 85.0 grams of
dissolve in 392 grams of water,
be the boiling point and
points of the resulting solution?
molecular formula of sugar is
C12H22O11. (answer = 100.33
sugar are
what will
freezing
The
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C, -1.18 C)
Lowering the Vapor Pressure
Nonvolatile solutes (with no measurable vapor pressure) reduce the ability of the surface solvent
molecules to escape the liquid.
• Therefore, vapor pressure is lowered.
• The amount of vapor pressure lowering depends on the amount of solute.
•
Osmosis
•Semipermeable membranes permit passage of some components of a solution.
•Often they permit passage of water but not larger molecules or ions.
•Examples of semipermeable membranes are cell membranes and cellophane.
• Osmosis is the net movement of a solvent from an area of low solute
concentration to an area of high solute concentration.
•Consider a U-shaped tube with a two liquids separated by a semipermeable membrane.
One arm of the tube contains pure solvent.
The other arm contains a solution.
•There is movement of solvent in both directions across a
semipermeable membrane.
•The rate of movement of solvent from the pure solvent to the solution
is faster than the rate of movement in the opposite direction.
•As solvent moves across the membrane, the fluid levels in the arms
become uneven.
•The vapor pressure of solvent is higher in the arm with pure solvent.
•Eventually the pressure difference due to the difference in height of liquid
in the arms stops osmosis.
• Osmotic pressure, , is the pressure required to prevent
osmosis.
• Everyday examples of osmosis include:
•If a cucumber is placed in NaCl solution, it will lose water to shrivel up
and become a pickle.
•A limp carrot placed in water becomes firm because water enters via
osmosis.
•Eating large quantities of salty food causes retention of water and
swelling of tissues (edema).
•Water moves into plants, to a great extent, through osmosis.
•Salt may be added to meat (or sugar added to fruit) as a
preservative.
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•Salt prevents bacterial infection: A bacterium placed on the salt will lose water through osmosis and
die.
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