Chapter 13
Ions in Aqueous Solutions and
Colligative Properties
13-1 Compounds in Aqueous
Solutions
Dissociation
The separation of ions that occurs when
an ionic compound dissolves
• One formula unit of NaCl produces two ions:
NaCl (s) → Na+ (aq) + Cl-(aq)
• One mole of NaCl produces two moles of ions
• One formula unit of CaCl2 produces three ions:
CaCl2 (s) → Ca+2 (aq) + 2Cl- (aq)
• One mole of CaCl2 produces three moles of ions
Dissociation Equations
NaCl(s)  Na+(aq) + Cl-(aq)
AgNO3(s)  Ag+(aq) + NO3-(aq)
MgCl2(s) 
Mg2+(aq) + 2 Cl-(aq)
Na2SO4(s) 
AlCl3(s) 
2 Na+(aq) + SO42-(aq)
Al3+(aq) + 3 Cl-(aq)
Precipitation Reactions
Solubility Rules
• No compound is completely insoluble
• Compounds of very low solubility can be
considered insoluble
• Dissociation equations cannot be written
for insoluble compounds
More Solubility Guidelines
• Oxalates are generally insoluble except for
those of alkali metals. (alkaline earth
oxalates are insoluble.)
• Chromates of silver, lead & mercury are
insoluble. Those of alkali metals, calcium,
magnesium & ammonium are soluble.
Double replacement forming a precipitate…
Double replacement (ionic) equation
Pb(NO3)2 (aq) + 2KI (aq)  PbI2 (s) + 2KNO3 (aq)
Spectator ions are those ions that do not take
Complete ionicpart
equation
shows
in a chemical
rxn compounds
and are found inas
solution
aqueous ions both before and after the rxn- unchanged!
Pb2+ (aq) + 2 NO3- (aq) + 2 K+ (aq) +2 I- (aq) 
PbI2 (s) + 2 K+ (aq) + 2 NO3- (aq)
Net ionic equation eliminates the spectator ions
Pb2+ (aq) + 2 I- (aq)  PbI2 (s)
Includes only those compounds and ions that
undergo a chemical change
Ionization
• Ions are formed from solute
compounds by the action of the
solvent
• Polar water molecules are attracted
to polar or charged solute particles
• Electronegative oxygen of water is
attracted to electropositive portion of
a solute compound
• Electropositive hydrogen of water is
attracted to the electronegative
portion of a solute compound
The Hydronium Ion
• H3O+ is called the "hydronium" ion
H2O (l) + HCl (g) → H3O+ (aq) + Cl- (aq)
Strong Electrolytes
• Any compound of which all or almost all of
the dissolved compound exists as ions in
an aqueous solution
• All soluble ionic compounds are strong
electrolytes
• Hydrogen halides :
HCl, HBr, HI
Weak Electrolytes
• A compound of which a relatively small
amount of the dissolved compound exists
as ions in an aqueous solution
• Examples:
HF, organic acids
Chapter 13
Ions in Aqueous Solutions and
Colligative Properties
13-2 Colligative Properties
Colligative Properties
• These are properties that are dependent
only on the number & concentration of
solute particles.
Vapor Pressure Lowering
• Effect of Solutes on Vapor-Pressure
Any nonvolatile solute will lower the vapor
pressure of a solution, having two
noticeable effects:
• Raising the boiling point of the solution
• Lowering the freezing point of the solution
Solute particles take up space @ the liquid-air surface.
So vapor pressure is lowered, boiling occurs at a
higher temp b/c more energy is required for the vapor
pressure to equal the atmospheric pressure.
Freezing Point Depression
• Solutions have lower freezing points than pure
solvents.
• If the solution is aqueous, its freezing point will
always be lower than 0C.
• How much lower?
• Depends on the # & concentration of solute
particles.
• Ex: Salt on ice in the winter
Colligative Properties
• Colligative properties depend on:
number & concentration of solute particles
• Since ionic substances dissolve into
multiple particles, their colligative effects
are greater than those of covalent
substances.
Freezing Point Depression
• Ionic solutes depress the freezing point more than
covalent solutes. Look at their solubility rxns, note
the number of particles formed:
• Ionic: NaCl (s) + H2O (l)  Na+ (aq) + Cl- (aq)
2 particles formed
• Covalent: C6H12O6 (s) + H2O (l)  C6H12O6 (aq)
1 particle formed
Molal Freezing-Point Constant for Water
• The freezing-point depression of the solvent in a 1molal solution of a nonvolatile, nonelectrolyte solute
Kf = -1.86 °C/m
• Freezing-Point Depression
• The difference between the freezing points of the pure
solvent and a solution of a nonelectrolyte in that
solvent
Δtf = Kf m
Where: m = molality
Δtf = change in freezing point
Boiling Point Elevation
• Solutions have higher boiling points than pure
solvents.
• If the solution is aqueous, its boiling point will
always be higher than 100C.
• Boiling: Temperature at which vapor pressure
equals atmospheric pressure.
Molal Boiling-Point Constant for Water
• The boiling point elevation of the solvent in a 1-molal
solution of a nonvolatile, nonelectrolyte solute
Kb = 0.51 °C/m
• Boiling-Point Elevation
• The difference between the boiling points of the pure
solvent and a solution of a nonelectrolyte in that
solvent
Δtb = Kb m
Where: m = molality
Δtb = change in boiling point
The van’t Hoff Factor, i
Electrolytes may have two, three or more
times the effect on boiling point and freezing
point, depending on its dissociation.
T = i  K  m
Freezing Point Depression and Boiling
Point Elevation Constants
Dissociation Equations
NaCl(s)  Na+(aq) + Cl-(aq)
i=2
i=2
AgNO3(s)  Ag+(aq) + NO3-(aq)
MgCl2(s)  Mg2+(aq) + 2 Cl-(aq)
Na2SO4(s) 
i=3
i=3
2 Na+(aq) + SO42-(aq)
AlCl3(s)  Al3+(aq) + 3 Cl-(aq)
i=4
Ideal vs. Real van’t Hoff Factor
•
•
•
•
•
Attractive forces between ions cause clustering
Clustering is greatest in concentrated solutions
Ideal and real results are closest in very dilute solutions
The Debye-Huckel Theory
Clustering hinders the movements of ions, so fewer ions
appear to be present
Ideal vs. Real van’t Hoff Factor
The ideal van’t Hoff Factor is only achieved in
VERY DILUTE solution.
Preventing
icing of roads
using CaCl2
Osmotic Pressure
• Semipermeable membranes
• Membranes that allow the movement of some
particles while blocking the movement of others
• Osmosis
• The movement of solvent through a semipermeable
membrane from the side of lower solute concentration
to the side of higher solute concentration
• Osmosis occurs when two solutions of different
concentration are separated by a semipermeable
membrane
Osmotic Pressure
• The external pressure that must be applied to
stop osmosis
• Osmotic pressure increases with the
concentration of solute particles
• Osmotic pressure is not dependent on the
TYPE of solute particles