7-_ch_12_-_physical_properties_of_solutions_0.ppsx

CHEM 101
(3+1+0)
Dr. Mohamed El-Newehy
http://fac.ksu.edu.sa/melnewehy
General Chemistry

Chapter 7
Physical Properties of
Solutions

Types of Solutions
A solution is a homogenous mixture of 2 or more substances.
The solute is(are) the substance(s) present in the smaller
amount(s).
The solvent is the substance present in the larger amount.
Types of solutions, depending on the original states
(solid, liquid, or gas) of the solution components.

Types of Solutions
An unsaturated solution contains less solute than the solvent
has the capacity to dissolve at a specific temperature.
A supersaturated solution contains more solute than is
present in a saturated solution at a specific temperature.
Sodium acetate crystals rapidly form when a seed crystal
is added to a supersaturated solution of sodium acetate.
A saturated solution contains the maximum amount of a solute
that will dissolve in a given solvent at a specific temperature.
Crystallization is the process in which dissolved solute comes
out of solution and forms crystals

A Molecular View of the Solution Process
Three types of interactions in the solution process:
o solvent-solvent interaction
o solute-solute interaction
o solvent-solute interaction
DHsoln = DH1 + DH2 + DH3
Step 1 is the separation
of solvent molecules.
(endothermic)
Step 2 is the separation of
solute molecules.
(endothermic)
Step 3 the solvent and
solute molecules mix.
(exothermic or endothermic)
The heat of solution DHsoln is given by

o If the solute-solvent attraction is stronger than the solvent-
solvent attraction and solute-solute attraction,
The solution process is favorable, or exothermic (ΔHsoln = 0).
o If the solute-solvent interaction is weaker than the solvent-
solvent and solute-solute interactions,
The solution process is endothermic (ΔHsoln = 0).

“like dissolves like”
Two substances with similar intermolecular forces are likely
to be soluble in each other.
o non-polar molecules are soluble in non-polar solvents
CCl4 in C6H6
o polar molecules are soluble in polar solvents
C2H5OH in H2O
o ionic compounds are more soluble in polar solvents
NaCl in H2O or NH3 (l)
Two liquids are said to be miscible if they are completely soluble in
each other in all proportions.

Concentration Units
Types of Concentration Units
Percent by Mass
The percent by mass (also called percent by weight or weight
percent) is the ratio of the mass of a solute to the mass of the
solution, multiplied by 100 percent:
% by mass = x 100%
mass of solute
mass of solute + mass of solvent
= x 100%
mass of solute
mass of solution
The concentration of a solution is the amount of solute present
in a given quantity of solvent or solution.

Concentration Units
Mole Fraction (X)
The mole fraction of a component of a solution, say,
component A, is written XA and is defined as
XA =
moles of A
sum of moles of all components

Concentration Units
Molality (m)
Molality is the number of moles of solute dissolved in 1 kg (1000 g) of solvent
Molarity (M)
Molarity was defined as the number of moles of solute in 1 L of solution
The units of molarity are mol/L.
M =
moles of solute
liters of solution
m =
moles of solute
mass of solvent (kg)

Concentration Units
Comparison of Concentration Units
For example;
o The mole fraction is not used to express the concentrations
of solutions for titrations and gravimetric analyses.
o It is appropriate for calculating partial pressures of gases
and for dealing with vapor pressures of solutions.
o Percent by mass is similar to molality in that it is
independent of temperature.
o It is defined in terms of ratio of mass of solute to mass of
solution, We do not need to know the molar mass of the
solute in order to calculate the percent by mass.

Concentration Units
o The advantage of molarity is that it is generally easier to measure the
volume of a solution.
For this reason, molarity is often preferred over molality.
o Molality is independent of temperature, because the concentration is
expressed in number of moles of solute and mass of solvent.
The volume of a solution typically increases with increasing temperature, so that
a solution that is 1.0 M at 25°C may become 0.97 M at 45°C because of the
increase in volume on warming.
o This concentration dependence on temperature can significantly affect
the accuracy of an experiment.
Therefore, it is sometimes preferable to use molality instead of molarity.

Concentration Units
What is the molality of a 5.86 M ethanol (C2H5OH)
solution whose density is 0.927 g/mL?
m =
moles of solute
M =
moles of solute
liters of solution
Assume 1 L of solution:
5.86 moles ethanol = 270 g ethanol
927 g of solution (1000 mL x 0.927 g/mL)
mass of solvent = mass of solution – mass of solute
= 927 g – 270 g = 657 g = 0.657 kg
m =
moles of solute
=
5.86 moles C2H5OH
0.657 kg solvent
= 8.92 m

Concentration Units
To convert one concentration unit of a solution to another;
Example;
Express the concentration of a 0.396 m glucose (C6H12O6) solution
in molarity.
0.396 m glucose (C6H12O6) i.e. there is 0.396 mole of glucose
in 1000 g of the solvent.
To calculate molarity; we need to determine the volume of this
solution.
First, we calculate the mass of the solution from the molar mass of
glucose:

Concentration Units
Determine the density of the solution, which is found to be
1.16 g/mL.
We can now calculate the volume of the solution in liters
The molarity of the solution is given by

The Effect of Temperature on Solubility
Solid solubility and temperature
o Solubility is the maximum amount of a solute that will dissolve in a
given quantity of solvent at a specific temperature.
o Temperature affects the solubility of most substances.
solubility increases with
increasing temperature
solubility decreases with
increasing temperature

Fractional crystallization is the separation of a mixture of substances into
pure components on the basis of their differing solubilities.
Suppose you have 90 g KNO3
contaminated with 10 g NaCl.
Fractional crystallization:
1. Dissolve sample in 100 mL of
water at 600C
2. Cool solution to 00C
3. All NaCl will stay in solution
(s = 34.2g/100g)
4. 78 g of PURE KNO3 will
precipitate (s = 12 g/100g).
90 g – 12 g = 78 g
Solid solubility and temperature

Gas Solubility and Temperature
O2 gas solubility and temperature
solubility usually decreases
with increasing temperature
The solubility of gases in water usually decreases with increasing
temperature.

The Effect of Pressure on the Solubility of Gases
The solubility of a gas in a liquid is proportional to the pressure of the
gas over the solution (Henry’s law).
c = kH P
c is the concentration (M) of the dissolved gas
P is the pressure of the gas over the solution
kH is a constant for each gas (mol/L•atm) that
depends only on temperature
low P
low c
high P
high c
A practical demonstration of Henry’s law is the effervescence of a soft drink
when the cap of the bottle is removed (air and CO2 saturated with water vapor).

o Most gases obey Henry’s law, but there are some important
exceptions.
o For example, if the dissolved gas reacts with water, higher
solubilities can result.
Carbon dioxide also reacts with water, as follows:
The solubility of ammonia is much higher than
expected because of the reaction
Another interesting example is the dissolution of
molecular oxygen in blood.

Colligative Properties of Nonelectrolyte Solutions
Colligative properties (or collective properties) are properties
that depend only on the number of solute particles in solution
and not on the nature of the solute particles.
o The colligative properties are
 vapor-pressure lowering,
 boiling-point elevation,
 freezing-point depression,
 osmotic pressure.
we are talking about relatively dilute solutions, that is, solutions whose
concentrations are ≤ 0.2 M.

Vapor-Pressure Lowering
Raoult’s law
If the solution contains only one solute:
X1 = 1 – X2
P 1
0
- P1 = DP = X2 P 1
0
P 1
0 = vapor pressure of pure solvent
X1 = mole fraction of the solvent
X2 = mole fraction of the solute
P1 = X1 P 1
0
Raoult’s† law, which states that the vapor pressure of a solvent over a
solution, P1, is given by the vapor pressure of the pure solvent, P°1, times
the mole fraction of the solvent in the solution, X1:
P1 = (1-X2)P1
0
P1 = P1-X2P1
0
0
We see that the decrease in vapor pressure, DP, is directly proportional to
the solute concentration (measured in mole fraction).

Boiling-Point Elevation
o At any temperature the vapor
pressure of the solution is lower
than that of the pure solvent
regardless of temperature.
o The boiling point of the solution is
higher than that of water.
o The boiling point of a solution is the temperature at which its vapor
pressure equals the external atmospheric pressure.
o Because the presence of a nonvolatile solute lowers the vapor
pressure of a solution, it must also affect the boiling point of the
solution.

Boiling-Point Elevation
o The boiling point elevation (DTb) is defined as the boiling point of
the solution (Tb) minus the boiling point of the pure solvent (T°b):
DTb = Tb – T b
0
Tb > T b
0
DTb > 0
T b is the boiling point of the pure solvent
0
T b is the boiling point of the solution
DTb = Kb m
m is the molality of the solution
Kb is the molal boiling-point elevation
constant (0C/m) for a given solvent

Freezing-Point Depression
o Ice on frozen roads and sidewalks melts when sprinkled with salts
such as NaCl or CaCl2.
This method of thawing succeeds because it depresses the
freezing point of water.
o Lowering the vapor pressure of the
solution shifts the solid-liquid curve to
the left

Freezing-Point Depression
The freezing point depression (DTf) is defined as the freezing point
of the pure solvent (T f°) minus the freezing point of the solution (Tf):
DTf = T f – Tf
0
T f > Tf
0
DTf > 0
T f is the freezing point of the pure solvent
0
T f is the freezing point of the solution
DTf = Kf m
m is the molality of the solution
Kf is the molal freezing-point depression
constant (0C/m) for a given solvent

What is the freezing point of a solution containing 478 g
of ethylene glycol (antifreeze) in 3202 g of water? The
molar mass of ethylene glycol is 62.01 g/mol.
DTf = Kf m
m =
moles of solute
= 2.41 m
=
3.202 kg solvent
478 g x
1 mol
62.01 g
Kf water = 1.86 0C/m
DTf = Kf m = 1.86 0C/m x 2.41 m = 4.48 0C
DTf = T f – Tf
0
Tf = T f – DTf
0
= 0.00 0C – 4.48 0C = -4.48 0C

Osmotic Pressure (p)
Osmosis is the selective passage of solvent molecules through
a porous membrane from a dilute solution to a more
concentrated one.
A semipermeable membrane allows the passage of solvent
molecules but blocks the passage of solute molecules.
Osmotic pressure (p) is the pressure required to stop osmosis.

High
P
Low
P
p = MRT
M is the molarity of the solution
R is the gas constant (0.0821 L . atm/K . mol)
T is the temperature (in K)

Summary - Colligative Properties of
Nonelectrolyte Solutions
Colligative properties are properties that depend only on the
number of solute particles in solution and not on the nature of
the solute particles.
Vapor-Pressure Lowering P1 = X1 P 1
0
Boiling-Point Elevation DTb = Kb m
Freezing-Point Depression DTf = Kf m
Osmotic Pressure (p) p = MRT

Using Colligative Properties to Determine Molar Mass
From the experimentally determined freezing-point depression
or osmotic pressure,
We can calculate the molality or molarity of the solution.
Knowing the mass of the solute, we can readily determine its
molar mass,

Using Colligative Properties to Determine Molar Mass

7-_ch_12_-_physical_properties_of_solutions_0.ppsx

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