Transcript Chapter Twelve

Chapter Twelve
Physical Properties
Of Solutions
Outline of the Chapter
• The fundamental properties of Solutions
• Working with solutions
– Concentration Units
• Saturated Solutions and Equilibrium
• Colligative Properties
– Freezing point depression
– Boiling point elevation
Some Types Of Solutions
Solution Concentration
• Molarity is the moles of solute per liter of solution.
moles of solute (mol)
Molarity (M) = _____________________________
volume of solution (L)
• Molarity is the most frequently used concentration in
chemical situations. Needed in many calculations to
determine how many moles are present in a certain
volume of solution:
moles of solute = volume of solution x molarity
More Concentration Units
amount of solute (g)
Percent by mass = ____________________________ x 100%
amount of solution (g)
volume of solute (L)
Percent by volume = ____________________________ x 100%
volume of solution (L)
amount of solute (g)
Mass/volume percent = ___________________________ x 100%
volume of solution (L)
Solutions By “Parts”
• Parts per million is the number of particles
of solute per one-million particles of solution.
Fluoridated drinking water contains ~ 1 ppm
of fluoride ion.
1 ppm = 1 mg/L (10-3)
• Parts per billion is the number of particles of
solute per one-billion particles of solution.
1 ppb = 1 µg/L (10-6)
• Parts per trillion is the number of particles of
solute per one-trillion particles of solution.
1 ppt = 1 ng/L (10-9)
Molality
• Molarity varies with
moles of solute (mol)
temperature due to the
_____________________________
M
=
expansion or contraction in
volume of solution (L)
the volume of the solution.
• To be independent of
temperature, a
concentration unit must be
based on mass only, not
volume.
• Molality (m) is the number
of moles of solute per one
moles of solute
kilogram of solvent (not
m
solution!). Molality does
mass of solvent (kg)
not vary with temperature.
An Example
How many milliliters of water (d = 0.998 g/mL)
are required to dissolve 25.0 g of urea and
thereby produce a 1.65 m solution of urea,
CO(NH2)2?
Mole Fraction And Mole Percent
• The mole fraction (xi) of a solution component i is the
fraction of all the molecules in the solution that are
molecules of i.
amount of component i (mol)
xi =
xi < 1;
total amount of solution components (mol)
= ni / ntotal
x1 + x2 + x3 + … = 1
• The mole percent of a solution component is its mole
fraction multiplied by 100%.
An Example
Example 12.6 (p.516)
An aqueous solution of ethylene glycol used as an
automobile engine coolant is 40.0% HOCH2CH2OH by
mass and has a density of 1.05 g/mL. What are the
(a) molarity, (b) molality, and mole fraction of
HOCH2CH2OH.
Formation Of Solution
Consider two important factors in solution formation:
• Enthalpy of solution
• Intermolecular forces in mixtures
Enthalpy Of Solution
• Solution formation can be considered to take
place in three steps:
– Move the molecules of solvent apart to make
room for the solute molecules. H1 > 0
– Separate the molecules of solute to the
distances found between them in the solution.
H2 > 0
– Allow the separated solute and solvent
molecules to mix randomly. H3 < 0
Hsoln = H1 + H2 + H3
Illustrating
enthalpy
terms when
solute
dissolves in
solvent
Intermolecular Forces In Solution Formation
• If all intermolecular forces are of comparable strength, this
type of solution is called an ideal solution and Hsoln = 0.
• If the intermolecular forces between solute and solvent
molecules are stronger than other intermolecular forces,
H3 > H1 + H2,
Hsoln < 0 - exothermic.
• If the intermolecular forces between solute and solvent
molecules are weaker than other intermolecular forces,
H3 < H1 + H2, Hsoln > 0 – endothermic.
• If the intermolecular forces between solute and solvent are
much weaker than other intermolecular forces, the solute
does not dissolve in the solvent. The compound is
described as relatively insoluble in the solvent.
Intermolecular Forces In Solution
Aqueous Solutions Of Ionic Compounds
• Two forces
- inter-ionic attractions that hold ions together
in crystal
- ion-dipole forces, the attraction of water
dipoles for cations and anions, that cause
an ionic solid to dissolve in water
• The extent to which an ionic solid dissolves in
water is determined largely by the
competition between inter-ionic attractions
and the ion-dipole attractions.
Ion-Dipole Forces in Dissolution
An Example
Example 12.8
Predict whether each of the following is likely to be a
solution or a heterogeneous mixture.
a. ethanol, CH3OH, and water, HOH
b. Pentane, CH3(CH2)3CH3, and octane, CH3(CH2)6CH3
c. Sodium chloride, NaCl, and carbon tetrachloride,
CCl4
d. 1-decanol, CH3(CH2)8CH2OH, and water, HOH
Formation of a Saturated Solution
Dynamic equilibrium: rate of crystallization = rate of dissolving
Some Solubility Terms
• Liquids that mix in all proportions are called
miscible.
• When there is a dynamic equilibrium between
an undissolved solute and a solution, the
solution is saturated.
• The concentration of the solute in a saturated
solution is the solubility of the solute.
• A solution which contains less solute than can
be held at equilibrium is unsaturated.
Solubility As A Function of Temperature
• About 95% of all ionic compounds have
aqueous solubilities that increase significantly
with increasing temperature.
• Most of the remainder have solubilities that
change little with temperature.
• A very few have solubilities that decrease with
increasing temperature.
• A supersaturated solution is created when a
warm, saturated solution is allowed to cool
without the precipitation of the excess solute.
Some
Solubility
Curves
A Supersaturated Solution
Selective Crystallization
When KNO3(s) is
crystallized from an
aqueous solution of
KNO3 containing
CuSO4 as an impurity,
CuSO4 remains in the
solution.
The Solubilities Of Gases
The solubilities of gases depend on
temperature and pressure.
• Most gases become less soluble in liquids as the
temperature increases.
• At a constant temperature, the solubility (S) of a
gas is directly proportional to the pressure of the
gas (Pgas) in equilibrium with the solution.
S = k Pgas (k > 1)
The value of k depends on the particular gas and
the solvent.
• The effect of pressure on the solubility of a gas is
known as Henry’s Law.
Effect of Temperature on Solubility of Gases
Effect of Pressure on Solubility of Gases
As a gas is
compressed into a
smaller volume,
increasing the number
of molecules per unit
volume, the number of
dissolved molecules
per unit volume – the
concentration of the
solution – also
increases.
Aqueous Solubility As A
Function Of Gas Pressure
Vapor Pressures of Solutions
• Raoult’s law states that the vapor pressure of the
solvent above a solution (Psolv) is the product of the
vapor pressure of the pure solvent (Posolv) and the
mole fraction of the solvent in the solution (xsolv):
Psolv = xsolv . Posolv
• The vapor in equilibrium with an ideal solution of
two volatile components has a higher mole fraction
of the more volatile component than is found in the
liquid.
Fractional Distillation
An Interesting Phenomenon
Two different aqueous solutions in a closed container. After
a time, the levels are unequal. Explain why this
spontaneous change occurred.
Colligative Properties
Colligative Properties – physical properties of
solutions that depend on the number of solute
particles present but not on the identity of the
solute.
• Freezing point depression
• Boiling point elevation
Freezing Point Depression
and Boiling Point Elevation
• Consider solutions with
- a volatile solvent
- a solute: nonvolatile, nonelectrolytic, and soluble
in liquid solvent but not frozen solvent
• The vapor pressure of the solution is that of the
solvent in the solution, and at all temperatures this
vapor pressure is lower than that of the pure
solvent.
• The presence of the solute lowers (depresses) the
freezing point of the solvent (Tf) and increases
(elevates) the boiling point of the solvent (Tb).
Vapor Pressure Lowering
By A Nonvolatile Solute
In a solution, the
vapor pressure of
the solvent is
lowered
and
the fusion curve is
displaced to lower
temperatures
(red curves)
Equations for FPD, BPE
• Tf = Tf, solution– Tf, solvent = – Kf x m
• Tf = Tb, solution– Tb, solvent = Kb x m
Tf – fusion temperature (melting)
Tb – boiling temperature
m – molality of solute
Kf, Kb – constants
FPD And BPE Constants
An Example
What mass of sucrose, C12H22O11, should be added
to 75.0 g H2O to raise the boiling point to 100.35
0C?
Summary
• Molarity is expressed as moles per liter and
molality is expressed as moles of solute per
kilogram of solvent.
• Units used for very dilute concentrations of
solutes include parts per million (ppm), parts per
billion (ppb), and parts per trillion (ppt).
• The type and magnitude of intermolecular
forces are important in solution composition.
• The solubility of a solute is its concentration in a
saturated solution.
Summary (continued)
• Solubility of solids vs. gasses
• Solubility of ionic solids increases with
increasing temperature
• Solubility of gasses decreases with increasing
temperature
• Solubility of gases in solutions:
•decreases with an increase in temperature
•increases linearly with pressure – Henry’s Law
•Solubility (S)  Pgas
S = k  P (as long as T is constant)
Summary (continued)
• Vapor pressure of solutions vs. pure solvents
• Raoult’s Law:
•Vapor pressure of solvent above solution is
equal to the product of the mole fraction of
solvent and the vapor pressure of pure
solvent:
Psolution = xsolvent  Psolvent
Summary (continued)
• Colligative properties depend on the
particular solvent and the number of solute
particles present.
• The presence of solutes lowers the vapor
pressure of the solution and causes both a
freezing point depression and a boiling point
elevation of the solvent.