Solutions
Chapter 13 and 14
Honors Chemistry
Solution
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Definition: a homogeneous mixture of 2 or
more substances in a single physical state
Parts: solute and solvent (usually water)
Types of solutions
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Physical states: solid (alloys), liquid, gas
Miscible vs. Immiscible
Miscible - liquids that dissolve freely in one
another in any proportion
 Immiscible - liquid solutes and solvents that are
not soluble
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Dilute vs. Concentrated
Electrolyte vs. Nonelectrolyte
Saturated, Unsaturated and Supersaturated
Supersaturated Solution demo
Electrolyte vs. Nonelectrolyte
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Saturated – soln
containing the max
amt of solute
Unsaturated – soln
containing less
solute than a sat
soln under the
existing conditions
Supersaturated –
contains more
dissolved solute
than a saturated
solution under the
same conditions
Solubility Curves
supersaturated solution
(stirred)
Solubility
(physical change)

Definition: mass of
solute needed to make
a saturated solution at
a given temperature

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solution equilibrium in
a closed system
dissolution ↔
crystallization
Unit = g solute/100 g
H2O
At 20oC, a saturated
solution contains how
many grams of NaNO3
in 100 g of water?
90 g
What kind of solution is
formed when 90 g
NaNO3 is dissolved in
100 g water at 30oC?
unsaturated
What kind of solution
is formed when 120 g
NaNO3 is dissolved in
100 g water at 40oC?
supersaturated
180
Saturated sol’n
170
160
150
140
Supersaturated
solution
130
120
Solubility ( g/100 g water )
What is the solubility
at 70oC?
135 g/100 g water
Solubility Graph for NaNO3
110
100
90
80
70
Unsaturated solution
60
50
40
30
20
10
0
0
10
20
30
40
50
60
70
Temperature (deg C)
80
90
100
110
Solubility of solids in liquids

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For most solids, increasing
temperature, increases solubility.
In general, “like dissolves like”.
Depends on
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Type of bonding
Polarity of molecule
Intermolecular forces between solute and
solvent
Solubility of Gases
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Gases are less
soluble at high
temperatures than
at low temperatures
Increasing
temperature,
decreases
solubility.
Increasing
pressure, increases
solubility.

The quantity of gas that dissolves in a
certain volume of liquid is directly
proportional to the pressure of the gas
(above the solution).

Effervescence – rapid escape of gas
dissolved in liquid
Factors Affecting Solubility

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Increase surface area of solute
(crushing)
Stir/shake
Increase temperature
Dissolution Process

electrolyte
Ionic Compounds
NaCl(s)  Na+1(aq) + Cl-1(aq)
nonelectrolyte

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For dissolution to occur, must overcome solute attractions
and solvent attractions.
Dissociation Reaction: the separation of IONS when an
ionic compound dissolves (ions already present)
Try calcium chloride
Dissolving
NaCl in water
hexahydrated for Na+1;
Solvation: process of solvent molecules
most cations have 4-9 H2O molecules
surrounding solute
6 is most common
Hydration: solvation with water
Dissolution Process
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Molecular Compounds

Nonpolar molecular solids do not dissolve in
polar solvents
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Polar molecule
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naphthalene
C12H22O11(s)  C12H22O11(aq)
Molecular solvation
Nonelectrolyte
Polar molecule
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HCl(g)  H+1(aq) + Cl-1(aq) or
HCl(g) + H2O  H3O+1(aq) + Cl-1(aq)
Ionization: ions formed from solute molecules by
action of solvent (no ions initially present)
Nonelectrolyte (HCl)  electrolyte (ions)
Energy Changes

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Heat of solution = Hsoln
Endothermic
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Solute particles separating in solid
Solvent particles moving apart to allow solute to enter
liquid
Energy absorbed
Exothermic

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Solute particles separating in solid
Solvent particles attracted to solvating solute particles
Energy released
Hsoln = heat of solvation –
crystal lattice energy
Solution Reactions
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Single replacement
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Activity Series
Double replacement
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Solubility chart (soluble vs. insoluble)
Neutralization (water)
Gases (CO2 and H2S)
Concentration

Percent concentration by mass (mass %)
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Molarity (M)
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Moles of solute/mass of solvent = mol/kg
ppm and ppb

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Moles of solute/Liters of solution = mol/L
Molality (m)
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(solute/solution) x 100% = % Concentration
Used for very dilute solutions
Dilution – a process in which more solvent is
added to a solution

How is this solution different?

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How is it the same?

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Volume, color, molarity
Same mass of solute, same moles of solute
In Dilution ONLY: M1V1 = M2V2
Colligative Properties

Definition: physical properties of
solutions that differ from properties
of its solvent.


Property depends upon the number of
solute particles in solution.
Types:
1.
2.
3.
Vapor Pressure
Boiling Point ELEVATION
Freezing Point DEPRESSION
Vapor Pressure
A measure of the tendency of molecules to escape from a liquid
For nonvolatile liquids or solid solutes
A nonvolatile solute will typically increase the
boiling point and decrease the freezing point.
Adding a nonvolatile solute lowers the
concentration of water molecules at the surface
of the liquid.
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This lowers the tendency of the water molecules to leave
the solution and enter the gas phase.
Therefore the vapor pressure of the solution is LESS
than pure water.
H2O
H2O
H2O
Sugar
H 2O
Same Temperature
Vapor Pressure (kPa)
100
H2O
80
60
40
20
Temperature (ºC)
100
solution
Boiling Point Elevation
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tb = boiling point elevation
tb = iKbm
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i = molality conversion factor; for electrolytes
adjust for # of ions actually present in solution
(dissociation process)
Kb = molal bp elevation constant
Kb = 0.512°C·kg H2O
moles of solute (ions or molecules)
m = molality = moles solute
kg of solvent
bp of solution = bp of solvent + tb
Boiling Point Elevation and
Freezing Point Problems
At what temperature will a solution begin
to boil if it is composed of 1.50 g
potassium nitrate in 35.0 g of water?
1.

Solute:
At what temperature will a solution begin
to freeze when 18.0 g ammonium
phosphate is dissolved in 200.0 g water?
2.

Solute:
Freezing Point Depression
when a solution freezes, the solvent solidifies as a pure
substance; deviates for more concentrated solutions

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tf = freezing point depression
tf = iKfm
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i = molality conversion factor; for electrolytes
adjust for # of ions actually present in solution
(dissociation process)
Kf = molal freezing point depression constant
Kf = 1.858°C·kg H2O
moles of solute (ions or molecules)
m = molality = moles solute
kg of solvent
fp of solution = fp of solvent - tf
Osmotic Pressure

The external pressure needed to stop
osmosis

Osmosis is the movement of solvent through a
semi permeable membrane from lower
concentration to higher concentration
the pressure required to allow for no
transport of solvent across the membrane is
called the OSMOTIC pressure