Chapter 14
Solutions
Overview
• Solution formation
– Types of solution
– Solubility and the solution process
– Effect of temperature and pressure on solubility
• Ways of expressing concentration
• Colligative properties
– Vapor pressure of a solution
– Boiling-point elevation and freezing point
depression.
– Osmosis
– Colligative properties of ionic solutions
Types of Solution
• Solution – homogeneous mixture of two or more
substances of ions or molecules. E.g. NaCl (aq)
• Colloid – appears to be a homogeneous mixture, but
particles are much bigger, but not filterable. E.g. Fog,
smoke, whipped cream, mayonnaise, etc.
• Suspension: larger particle sizes, filterable. E.g. mud,
freshly squeezed orange juice.
Solution Composition
Solvent = component which is the
component in greater amount.
Solute = component which is present in
the smaller amount.
Aqueous Solutions = aqueous solutions
are those in which water is the solvent.
Solution Composition
– Gaseous = gases are completely
miscible in each other.
– Liquid = gas, liquid or solid solute
dissolved in solvent.
– Solid = mixture of two solids that are
miscible in each other to form a single
phase.
Electrolytes
• Electrolytes are solutes that dissolve in water to form
ions and consequently are capable of conducting
electricity. Electrolytes consist of ionic compounds as
well as binary molecules of hydrogen (acids).
•
•
•
•
Examples:
NaCl(s) ----------> Na+(aq) + Cl-(aq)
HF(g) ------------> H+(aq) + F-(aq)
Ca(NO3)2(s) ----------> Ca2+(aq) + 2NO3-(aq)
Nonelectrolytes
• Nonelectrolytes are solutes that dissolve in
water to form molecules and consequently
are incapable of conducting electricity.
• CH3OH(l) ----------> CH3OH(aq)
Concentration of Solutes Qualitative
• Dilute - this means that the solution contains only a
small amount of solute
• Concentrated - this means that the solution contains
more solute
• As you may have guessed - these are not real good
ways of describing solution concentrations. But they do
occasionally serve useful purposes. For example, in the
case of acids and bases the terms dilute and
concentrated refer to special concentrations of the
solutions:
•
•
•
•
•
Solution
HCl
HNO3
H2SO4
NH3
Dilute
6.00 M
6.00 M
3.00 M
6.00 M
Concentrated
12.0 M
15.0 M
18.0 M
15.0 M
Solubility and the Solution Process
• The solid dissolves rapidly
at first but as the solution
approaches saturation the
net rate of dissolution
decreases since the
process is in dynamic
equilibrium.
• When the solution has
reached equilibrium the
amount of solute does not
change with time;
• At equilibrium: the rate of
dissolution = rate of
solution
Fig. 12.2 Solubility Equilibrium
• http://www.mhhe.com/physsci/chemistry/e
ssentialchemistry/flash/molvie1.swf
Concentration of Solutes Using
Equilibria
• Concentrations in terms of equilibria rely
heavily on the solubility of the solute in the
particular solvent. Solubility is defined as
the maximum amount of solute that will
dissolve in a particular solvent under the
specified conditions of temperature and
pressure. Solubility is generally given in
terms of grams of solute per 100 grams of
solvent.
• Saturated solution: a solution in which
a dynamic equilibria exists between the
undissolved solute and the solution.
The solution contains the maximum
amount of dissolved solute according to
solubility.
• Unsaturated solution: a solution that
contains less than the maximum amount
of dissolved solute.
• Supersaturated: a solution that contains
a greater concentration of solute than a
saturated solution. This type of solution is
unstable. Generally, if you add a tiny
crystal of the solute to the solution all the
excess solute will very rapidly come out of
solution.
Molarity, M
Molarity is defined as the moles of solute
per liter of solution. (Chapter 5 - see
pages 228 - 235 if you can't remember.)
Molality, m
Molality is defined as the moles of solute
per kilogram of solvent. This is new to this
chapter and will be very important when
determining certain properties of solutions.
mol solute
Molality (m) 
mass of solvent (kg)
Determine the molality of a solution
prepared by dissolving 1.44 g NaCl into
exactly 100.0 mL of water. Assume the
density of water is 1.00 g/mL.
Mole Fraction, X
You can have a mole fraction of the solute or a
mole fraction of the solvent. Mole fraction is
defined as the moles of solute divided by the
total number of moles or the moles of solvent
divided by the total number of moles. The mole
fraction is expressed as a decimal. An
explanation of this can also be found on page
656.
mol of A
X
mol A  mol B
Determine the mole fraction of a solution
made from dissolving 30.0 g H2O2 with
70.0 g H2O.
Weight percent/PPM/PPB
Weight percent is defined as the mass of component of
the solution divided by the total mass of the solution multiplied by 100. Parts per million/parts per billion are
determined by dividing the mass of the solute by the total
mass and then multiplying of 1 million or 1 billion.
wt% 
mass of solute
x10 2
mass of solution
ppm 
mass of solute
x10 6
mass of solution
mass of solute
ppb 
x10 9
mass of solution
Determine the mass % of a solution
prepared by dissolving 1.44 g of NaCl in
100.0 mL of water. Assume that the
density of water is 1.00 g/mL
Determine mass % of solution made
from dissolving 30.0 g H2O2 with 70.0
g H2O.
Determine molality of 30% H2O2(aq)
Concentrated ammonia is 14.8 M and
has a density of 0.900 g/mL. What is
the molality?
Factors Affecting Solubility
• Nature of the Solute and Solvent
• Temperature
• Pressure
Nature of the Solute and Solvent
• Gases = generally completely soluble in each other
because of entropy (tendency towards maximum
randomness).
• Molecules in gas phase are far apart from each other
and not interacting strongly with each other in solution.
Mixing of Gas Molecules
Electrolytes and Nonelectrolytes
• Most nonelectrolytes that are appreciably
soluble in water are hydrogen bonded
(CH3OH, H2O2, sugars)
• Other types of nonelectrolytes are
generally more soluble in nonpolar or
slightly polar solvents (C6H6, CCl4)
Molecular Solutions
• Molecular compounds with similar chemical
structures and polarities tend to be miscible.
• Homologous alcohol series have polar and
non-polar ends.
Ionic Solutions
• Solubility affected by:
– Energy of attraction (due Ion-dipole force)
affects the solubility. Also called hydration
energy,
– Lattice energy (energy holding the ions
together in the lattice. Related
• to the charge on ions; larger charge means
higher lattice energy.
• Inversely proportional to the size of the ion; large
ions mean smaller lattice energy.
• Solubility increases with increasing ion size,
due to decreasing lattice energy;
Mg(OH)2(least soluble), Ca(OH)2, Sr(OH)2,
Ba(OH)2(most soluble) (lattice energy
changes dominant).
• Energy of hydration increases with for smaller
ions than bigger ones; thus ion size.
MgSO4(most soluble),... BaSO4 (least
soluble.) Hydration energy dominant.
Heat of Solution and Solubility
• Hsoln is sometimes negative and
sometimes positive.
– Solvent – solvent interactions: energy required to
break weak bonds between solvent molecules.
– Solute – solute interactions: energy required to
break intermolecular bonds between the solute
molecules.
– Solute – solvent interactions: H is negative since
bonds are formed between them.
Energy Changes and Solution Formation
Effect of Temperature on Solubility
1. An increase in temperature favors the endothermic
process
2. Solid + water -----------> Solution; DH for these
solutions is usually endothermic so generally an
increase in temperature will increase solubility.
3. Gas + water ------------> Solution; H for these
solutions is usually exothermic (WHY?) so generally
an increase in temperature will decrease solubility.
Solubility: Temperature
Dependence
• All solubilities are temperature dependent; must report
temperatures with solubilities.
• Temperature related to sign of Hsoln;
Predict the temperature dependence of
the solubility of Li2SO4, Na2SO4 and K2SO4
if their Hsoln are 29.8 kJ/mol, 2.4 kJ/mol
and +23.8 kJ/mol, respectively.
Solubility: Pressure Dependence
• Pressure has little effect on the solubility of a
liquid or solid, but has dramatic effect on gas
solubility in a liquid.
• Henry’s law S = kHP. Allows us to predict the
solubility of a gas at any pressure.
At 25C P(O2 in air) = 0.21 atm. Its solubility in
water is 3.2x104M. Determine its solubility when
pressure of O2 = 1.00 atm.
Physical Behavior of Solutions:
Colligative Properties
• Compared with the pure solvent the
solution’s:
– Vapor pressure is lower
– Boiling point is elevated
– Freezing point is lower
– Osmosis occurs from solvent to solution
when separated by a membrane.
Vapor-Pressure Lowering of
Solutions: Raoult’s Law
• Raoult’s Law: Psoln = PsolvxXsolv
• Non–volatile solute: vapor pressure
decreases upon addition of solute.
• Linear for dilute solutions
• http://dwb4.unl.edu/ChemAnime/solutio
ns.htm
• Vapor pressure lowering : P = Po  P =
Po(1Xsolv)
Determine vapor pressure lowering when
5.00 g of sucrose added to 100.0 g of
H2O. MM(sucrose) = 342.3 g/mol. The
vapor pressure of water at 25°C is 23.8
mmHg.
Determine the mass of sucrose dissolved in
100.0 g of water if the vapor pressure was
20.0 mmHg.
BP Elevation of Solutions
The magnitude of the change in FP and BP is directly
proportional to the concentration of the solute
(molality) – expressed in terms of the total number of
particles in the solution.
• BP Elevation
The magnitude of the BP increase
is given by the equation:
Tb  Kb  m
where Kb has units of °Ckg/mol or °C/m
Determine boiling point elevation when 5.00 g of sucrose is
added to 100.0 g of H2O. MM(sucrose) = 342.3 g/mol.
Kb = 0.521 C/m.
FP Depression of Solutions
• FP Depression: linear variation with
composition and given by:
Tf  K f  m
where the units for this constant are the same
as for Kb
Determine freezing point depression when
5.00 g of sucrose is added to 100.0 g of H2O.
FM(sucrose) = 342.3 g/mol. Kf = 1.86°C/m.
Osmosis and Osmotic Pressure
Osmosis: the passage of solvent through a membrane
from the less concentrated side to the more
concentrated side.
Osmotic pressure: the amount of pressure necessary to
stop Osmosis.
• Small molecules such as water can move through
certain types of materials (membranes).
• The tendency for this to occur is related to the molarity of
the solution, is also a function of the temperature and is
measured with a device called a Thistle tube.
  MRT
where M = is molarity of solute particles
Determine osmotic pressure of a solution
containing 0.100 g of hemoglobin
(molecular mass = 6.41x104 amu) in
0.0100 L at 1.00C.
Reverse Osmosis
• Application of a pressure to the
solution (that is equal to or
greater than the Osmotic
pressure) and the solvent flows
from the more concentrated
side to the other one.
• This process is used to obtain
pure water from salt water.
Colligative Properties and Molar
Mass
• molar mass = (kb x grams of
solute)/(kilograms of solvent x Tb)
• van’t Hoff factor can be something other
than integer under certain circumstances,
but for completely ionic solutions is equal
to the number of ions/ionic compound to
be found in solution:
E.g. NaCl: i = 2; Na2SO4: i = 3;
Colligative Properties and Molar
Mass II
Osmotic pressure of a solution containing
50.0 mg of a compound in 10.0 mL of
water was 4.80 torr at 5.00C. Determine
FM of the compound.
Colligative Properties of Ionic Solutions
• Colligative properties of solutions depends upon the
total concentration of particles.
• Each equation describing colligative properties must
be modified to account for this with ionic solutions
since each ionic compound gives more than one
mole of ions for every mole of compound.
– BP elevation:
– Freezing point depression:
– Osmotic pressure:
Where i = van’t Hoff factor.