IB Chemistry Notes
Chapter 11: Properties of Solutions
Solutions
 ___________________ – a homogeneous mixture of pure substances (occur in all phases, but we
will focus on aqueous solutions)
 The
is the medium in which the
are dissolved.
o (The solvent is usually the most abundant substance.)
Concentration of Solution - refers to the amount of solute dissolved in a solution.
MOLARITY (M) =
MOLALITY (m) =
MASS PERCENT (%) =
MOLE FRACTION () =
NORMALITY (N) =
Energy of Making Solutions
 Heat of solution ( Hsoln ) is the
 Most easily understood if broken into steps.
o Break apart solvent
 Have to overcome attractive forces.
o Break apart solute
 Have to overcome attractive forces.
o Mixing solvent and solute
 H3 depends on what you are mixing.
 Molecules can attract each other –
 Molecules can’t attract o This explains the rule
for making a solution.
Size of H3 determines whether a solution will form
 Types of Solvent and solutes
 If Hsoln is small and positive, a solution will still form because of entropy.
 There are many more ways for them to become mixed than there is for them to stay separate.
Solution Formation – Factors Favoring Spontaneity
 Processes in which the energy content of the system decreases (exothermic) tend to occur
spontaneously.
 Processes in which the disorder (entropy) of the system increases tend to occur
spontaneously.
Structure and Solubility
 Water soluble molecules must have dipole moments -
IB Chemistry Notes

Chapter 11: Properties of Solutions
To be soluble in non polar solvents the molecules must be
.
Pressure
 Changing the pressure doesn’t affect the amount of solid or liquid that dissolves
o They are incompressible.
 Pressure does affect solubility of gases.
Dissolving Gases
 Pressure affects the amount of
that can dissolve in a liquid.
 The dissolved gas is at equilibrium with the gas above the liquid.
 If you increase the pressure the gas molecules dissolve faster.
o The equilibrium is disturbed.
o The system reaches a new equilibrium with more gas dissolved.
 Henry’s Law:
Temperature Effects
 Increased temperature increases the rate at which a solid dissolves.
 We can’t predict whether it will increase the amount of solid that dissolves.
 We must read it from a graph of experimental data.
 Gases are predictable
 As temperature increases, solubility decreases.
 Gas molecules can move fast enough to escape.
 Thermal pollution.
Vapor Pressure of Solutions
 A nonvolatile solvent lowers the vapor pressure of the solution.
 The molecules of the solvent must overcome the force of both the other solvent molecules
and the solute molecules.
Raoult’s Law:

Applies only to an ideal solution where the solute doesn’t contribute to the vapor pressure.
To determine whether a sol’n is IDEAL…
 Liquid-liquid solutions where both are volatile.
 Modify Raoult’s Law to:
Ptotal =
•Ptotal = vapor pressure of mixture
• PA0= vapor pressure of pure A


If this equation works then the solution is ideal.
Solvent and solute are alike.
IB Chemistry Notes
Chapter 11: Properties of Solutions
Colligative Properties of Solutions = physical properties of solutions that depend on the # of particles
dissolved, not the kind of particle.
 Lowering vapor pressure
 Raising boiling point
 Lowering freezing point
 Generating an osmotic pressure
Boiling Point Elevation: a solution that contains a nonvolatile solute has a higher boiling pt than the pure
solvent; the boiling pt elevation is proportional to the # of moles of solute dissolved in a given mass of
solvent.
where: Tb = elevation of boiling pt
m = molality of solute
kb = the molal boiling pt elevation constant for a particular solvent
kb for water = 0.52 °C/m
Freezing/Melting Point Depression: the freezing point of a solution is always lower than that of the pure
solvent.
where: Tf = lowering of freezing point
m = molality of solute
kf = the freezing pt depression constant
kf for water = 1.86 °C/m
Ex: An antifreeze solution is prepared containing 50.0 cm3 of ethylene glycol, C2H6O2, (d = 1.12 g/cm3), in
50.0 g water. Calculate the freezing point of this 50-50 mixture. Would this antifreeze protect a car in
Chicago on a day when the temperature gets as low as –10° F?
(-10 °F = -23.3° C)
Electrolytes and Colligative Properties
 Colligative properties depend on the # of particles present in solution.
 Because ionic solutes dissociate into ions, they have a greater effect on freezing pt and boiling pt
than molecular solids of the same molal conc.
o For example, the freezing pt of water is lowered by 1.86°C with the addition of any
molecular solute at a concentration of 1 m, such as C6H12O6, or any other covalent
compound
o However, a 1 m NaCl solution contains 2 molal conc. of IONS. Thus, the freezing pt
depression for NaCl is 3.72°C…double that of a molecular solute.
IB Chemistry Notes
Chapter 11: Properties of Solutions
The relationships are given by the following equations:
Tf/b = f.p. depression/elevation of b.p.
m = molality of solute
kf/b = b.p. elevation/f.p depression constant
i = # particles formed from the dissociation of each formula unit of the solute
(van’t Hoff factor)
Ex: What is the freezing pt of:
a) a 1.15 m sodium chloride solution?
Ex: What is the freezing pt of:
b) a 1.15 m calcium chloride solution?
Ex: What is the freezing pt of:
c) a 1.15 m calcium phosphate solution?
Osmotic Pressure:
Experiments show that dependence of the osmotic pressure on solution concentration is expressed by the
eqn:
Where,
 = osmotic pressure (atm)
M = molarity (mol/L)
R = gas law constant = 0.08206
T = temp (K)