University Chemistry II
Spring 2006
Instructor:
Office:
Phone:
Email address:
Office hours:
Dr. Sarah A. Green
Chem Sci. 607
487-2048
sgreen@mtu.edu
Wednesday 1:00–3:00 pm
Class time: MWF 11:05-11:55
Place: DOW 641
Lab Supervisor: Lorri Reill y, Chemical Sci. 508B, lareill y@mtu.edu ; 7-2044
Learning Center Coordinator: Lois Blau, Chem Sci. 206A lablau@mtu.edu; 7-2297
Textbook: Chemistry: The Central Science, 10th edition, by Brown, LeMay, and Bursten.
Week
1
2
3
4
5
6
EXAM 1
7
8
BREAK
9
10
EXAM II
11
12
13
EXAM III
14
Dates
Jan 9-13
Jan 16-20
Jan 23-27
Jan 30-Feb 3
Feb 6-8
Feb 13-15
Feb 15, 6:00 pm
Feb 20-24
Feb 27-March 3
March 6-10
March 13-17
March 20-22
March 22, 6:00 pm
March 27-31
April 3-7
April 10-12
April 12, 6:00 pm
April 17-21
Chapter
13
14
15
16
17
17, Review
13-16
18
19
20
Review
17-20
21
22
23
21-23
25
Topic
Solutions
Chemical Kinetics
Chemical Equilibrium
Acid-Base Equili bria
No class Friday: Winter Carnival
No class Friday, Feb 17
Environmental Chemistry
Thermodynamics
Electrochemistry
No class Friday, March 24
Nuclear Chemistry
Nonmetals
Metals
No class Friday, April 14
Organic/Biochem
FINAL EXAM: Wednesday, April 26 from 10:15 am - 12:15 pm
This schedule is subject to modification. Any changes will be announced in class and posted to
the class via WebCT.
Solutions
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 13
Properties of Solutions
Adapted by SA Green from:
John D. Bookstaver
St. Charles Community College
St. Peters, MO
2006, Prentice Hall, Inc.
Solutions
Solutions
• Solutions are homogeneous mixtures of two
or more pure substances.
• In a solution, the solute is dispersed uniformly
throughout the solvent.
Solutions
Solutions
How does a solid dissolve
into a liquid?
What „drives‟ the
dissolution process?
What are the energetics of
dissolution?
Solutions
How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy ( H) changes with each interaction broken or
formed.
Ionic solid dissolving in water
Solutions
How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy ( H) changes with each interaction broken or
formed.
Solutions
How Does a Solution Form
The ions are solvated
(surrounded by
solvent).
If the solvent is water,
the ions are
hydrated.
The intermolecular
force here is iondipole.
Solutions
Energy Changes in Solution
To determine the enthalpy
change, we divide the
process into 3 steps.
1. Separation of solute
particles.
2. Separation of solvent
particles to make
„holes‟.
3. Formation of new
interactions between
solute and solvent.
Solutions
Enthalpy Changes in Solution
The enthalpy
change of the
overall process
depends on H for
each of these steps.
Start
End
Solutions
Start
End
Enthalpy changes during dissolution
Hsoln = H1 + H2 + H3
The enthalpy of
solution, Hsoln, can
be either positive or
negative.
Hsoln (MgSO4)= -91.2 kJ/mol --> exothermic
Hsoln (NH4NO3)= 26.4 kJ/mol --> endothermic
Solutions
Why do endothermic processes
sometimes occur spontaneously?
Some processes,
like the dissolution
of NH4NO3 in water,
are spontaneous at
room temperature
even though heat is
absorbed, not
released.
Solutions
Enthalpy Is Only Part of the Picture
Entropy is a measure of:
• Dispersal of energy in
the system.
• Number of microstates
(arrangements) in the
system.
b. has greater entropy,
is the favored state
(more on this in chap 19)
Solutions
Entropy changes during dissolution
Each step also involves a
change in entropy.
1. Separation of solute
particles.
2. Separation of solvent
particles to make
„holes‟.
3. Formation of new
interactions between
solute and solvent.
Solutions
SAMPLE EXERCISE 13.1 Assessing Entropy Change
In the process illustrated below, water vapor reacts with excess solid sodium
sulfate to form the hydrated form of the salt. The chemical reaction is
Does the entropy of the system increase or decrease?
Solutions
Dissolution vs reaction
Ni(s) + HCl(aq)
NiCl2(aq) + H2(g)
dry
NiCl2(s)
• Dissolution is a physical change—you can get back the
original solute by evaporating the solvent.
• If you can‟t, the substance didn‟t dissolve, it reacted.
Solutions
Degree of saturation
• Saturated solution
 Solvent holds as much
solute as is possible at
that temperature.
 Undissolved solid
remains in flask.
 Dissolved solute is in
dynamic equilibrium
with solid solute
particles.
Solutions
Degree of saturation
• Unsaturated Solution
 Less than the
maximum amount of
solute for that
temperature is
dissolved in the
solvent.
 No solid remains in
flask.
Solutions
Degree of saturation
• Supersaturated
 Solvent holds more solute than is normally
possible at that temperature.
 These solutions are unstable; crystallization can
often be stimulated by adding a “seed crystal” or
scratching the side of the flask.
Solutions
Degree of saturation
Unsaturated, Saturated or Supersaturated?
How much solute can be dissolved in a solution?
More on this in Chap 17
(solubility products, p 739)
Solutions
Factors Affecting Solubility
• Chemists use the axiom
“like dissolves like”:
 Polar substances tend to
dissolve in polar solvents.
 Nonpolar substances tend
to dissolve in nonpolar
solvents.
Solutions
Factors Affecting Solubility
Example: ethanol in water
The stronger the
intermolecular
attractions between
solute and solvent,
the more likely the
solute will dissolve.
Ethanol = CH3CH2OH
Intermolecular forces = H-bonds; dipole-dipole; dispersion
Ions in water also have ion-dipole forces.
Solutions
Factors Affecting Solubility
Glucose (which has
hydrogen bonding)
is very soluble in
water.
Cyclohexane (which
only has dispersion
forces) is not watersoluble.
Solutions
Factors Affecting Solubility
• Vitamin A is soluble in nonpolar compounds
(like fats).
• Vitamin C is soluble in water.
Solutions
Which
vitamin is
water-soluble
and which is
fat-soluble?
Solutions
Gases in Solution
• In general, the
solubility of gases in
water increases with
increasing mass.
Why?
• Larger molecules
have stronger
dispersion forces.
Solutions
Gases in Solution
QuickTime™ and a
TIFF (LZW) decompressor
are neede d to see this picture.
Solutions
Gases in Solution
Increasing
pressure
above
solution
forces
more gas
to dissolve.
• The solubility of
liquids and solids
does not change
appreciably with
pressure.
• But, the solubility of
a gas in a liquid is
directly proportional
to its pressure.
Solutions
Henry‟s Law
Sg = kPg
where
• Sg is the solubility of
the gas;
• k is the Henry‟s law
constant for that gas in
that solvent;
• Pg is the partial
pressure of the gas
above the liquid.
Solutions
Henry‟s Law
k for N2 at 25°
=6.8 x 10-4 mol/L atm
Sg = kPg
Solutions
Temperature
Generally, the
solubility of solid
solutes in liquid
solvents increases
with increasing
temperature.
Solutions
Temperature
• The opposite is true of
gases. Higher
temperature drives
gases out of solution.
 Carbonated soft drinks
are more “bubbly” if
stored in the
refrigerator.
 Warm lakes have less
O2 dissolved in them
than cool lakes.
Solutions
Chap 13:
Ways of Expressing
Concentrations of
Solutions
Solutions
Mass Percentage
mass of A in solution
Mass % of A =
total mass of solution
100
Solutions
Parts per Million and
Parts per Billion
Parts per Million (ppm)
mass of A in solution
ppm =
total mass of solution
106
Parts per Billion (ppb)
mass of A in solution
ppb =
total mass of solution
109
Solutions
Mole Fraction (X)
moles of A
XA =
total moles in solution
• In some applications, one needs the
mole fraction of solvent, not solute—
make sure you find the quantity you
need!
Solutions
Molarity (M)
M=
mol of solute
L of solution
• You will recall this concentration
measure from Chapter 4.
• Because volume is temperature
dependent, molarity can change with
temperature.
Solutions
Molality (m)
m=
mol of solute
kg of solvent
Because neither moles nor mass
change with temperature, molality
(unlike molarity) is not temperature
dependent.
Solutions
Solutions
SAMPLE EXERCISE 13.4 Calculation of Mass-Related Concentrations
(a) A solution is made by dissolving 13.5 g of glucose (C 6H12O6) in 0.100 kg of water. What is the mass
percentage of solute in this solution? (b) A 2.5-g sample of groundwater was found to contain 5.4 g of Zn2+
What is the concentration of Zn2+ in parts per million?
PRACTICE EXERCISE
(a) Calculate the mass percentage of NaCl in a solution containing 1.50 g of NaCl in 50.0 g of water. (b) A
commercial bleaching solution contains 3.62 mass % sodium hypochlorite, NaOCl. What is the mass of NaOCl
in a bottle containing 2500 g of bleaching solution?
PRACTICE EXERCISE
A commercial bleach solution contains 3.62 mass % NaOCl in water. Calculate (a) the molality and (b) the mole
fraction of NaOCl in the solution.
Solutions
Colligative Properties
• Colligative properties depend only on
the number of solute particles present,
not on the identity of the solute
particles.
• Among colligative properties are
Vapor pressure lowering
Boiling point elevation
Melting point depression
Osmotic pressure
Solutions
Vapor Pressure
As solute molecules are
added to a solution,
the solvent become
less volatile
(=decreased vapor
pressure).
Solute-solvent
interactions contribute
to this effect.
Solutions
Vapor Pressure
Therefore, the vapor
pressure of a solution
is lower than that of
the pure solvent.
Solutions
Raoult‟s Law
PA = XA P
A
where
• XA is the mole fraction of compound A
• P A is the normal vapor pressure of A at
that temperature
NOTE: This is one of those times when you
want to make sure you have the vapor
pressure of the solvent.
Solutions
SAMPLE EXERCISE 13.8 Calculation of Vapor-Pressure Lowering
Glycerin (C3H8O3) is a nonvolatile nonelectrolyte with a density of 1.26 g/mL at 25°C. Calculate the vapor
pressure at 25°C of a solution made by adding 50.0 mL of glycerin to 500.0 mL of water. The vapor pressure of
pure water at 25°C is 23.8 torr (Appendix B).
PRACTICE EXERCISE
The vapor pressure of pure water at 110°C is 1070 torr. A solution of ethylene glycol and water has a vapor
pressure of 1.00 atm at 110°C. Assuming that Raoult’s law is obeyed, what is the mole fraction of ethylene
glycol in the solution?
Solutions
Boiling Point Elevation and
Freezing Point Depression
Solute-solvent
interactions also
cause solutions to
have higher boiling
points and lower
freezing points than
the pure solvent.
Solutions
Boiling Point Elevation
The change in boiling
point is proportional to
the molality of the
solution:
Tb = Kb  m
Tb is added to the normal
boiling point of the solvent.
where Kb is the molal
boiling point elevation
constant, a property of
the solvent.
Solutions
Freezing Point Depression
• The change in freezing
point can be found
similarly:
Tf = Kf  m
• Here Kf is the molal
freezing point
depression constant of
the solvent.
Tf is subtracted from the normal
freezing point of the solvent.
Solutions
Boiling Point Elevation and
Freezing Point Depression
In both equations,
T does not depend
on what the solute
is, but only on how
many particles are
dissolved.
Tb = Kb  m
Tf = Kf  m
Solutions
Colligative Properties of
Electrolytes
Because these properties depend on the number of
particles dissolved, solutions of electrolytes (which
dissociate in solution) show greater changes than those
of nonelectrolytes.
e.g. NaCl dissociates to form 2 ion particles; its limiting
van‟t Hoff factor is 2.
Solutions
Colligative Properties of
Electrolytes
However, a 1 M solution of NaCl does not show
twice the change in freezing point that a 1 M
solution of methanol does.
It doesn‟t act like there are really 2 particles.
Solutions
van‟t Hoff Factor
One mole of NaCl in
water does not
really give rise to
two moles of ions.
Solutions
van‟t Hoff Factor
Some Na+ and Cl−
reassociate as
hydrated ion pairs,
so the true
concentration of
particles is
somewhat less than
two times the
concentration of
NaCl.
Solutions
The van‟t Hoff Factor
• Reassociation is
more likely at higher
concentration.
• Therefore, the
number of particles
present is
concentration
dependent.
Solutions
The van‟t Hoff Factor
We modify the
previous equations
by multiplying by the
van‟t Hoff factor, i
Tf = Kf  m  i
i = 1 for non-elecrtolytes
Solutions
Osmosis
• Semipermeable membranes allow some
particles to pass through while blocking
others.
• In biological systems, most
semipermeable membranes (such as
cell walls) allow water to pass through,
but block solutes.
Solutions
Osmosis
In osmosis, there is
net movement of
solvent from the area
of higher solvent
concentration (lower
solute concentration)
to the are of lower
solvent
concentration (higher
solute concentration).
Water tries to equalize the concentration on
both sides until pressure is too high.
Solutions
Osmotic Pressure
• The pressure required to stop osmosis,
known as osmotic pressure, , is
=(
n
)
RT = MRT
V
where M is the molarity of the solution
If the osmotic pressure is the same on both sides
of a membrane (i.e., the concentrations are the
same), the solutions are isotonic.
Solutions
Osmosis in Blood Cells
• If the solute
concentration outside
the cell is greater than
that inside the cell, the
solution is hypertonic.
• Water will flow out of
the cell, and crenation
results.
Solutions
Osmosis in Cells
• If the solute
concentration outside
the cell is less than
that inside the cell, the
solution is hypotonic.
• Water will flow into the
cell, and hemolysis
results.
Solutions
Solutions
Molar Mass from
Colligative Properties
We can use the
effects of a colligative
property such as
osmotic pressure to
determine the molar
mass of a compound.
Solutions
Colloids:
Suspensions of particles larger than
individual ions or molecules, but too small to
be settled out by gravity.
Solutions
Tyndall Effect
• Colloidal suspensions
can scatter rays of light.
• This phenomenon is
known as the Tyndall
effect.
Solutions
Colloids in Biological Systems
Some molecules have
a polar, hydrophilic
(water-loving) end and
a nonpolar,
hydrophobic (waterhating) end.
Solutions
Colloids in Biological Systems
Sodium stearate
is one example
of such a
molecule.
Solutions
Colloids in Biological Systems
These molecules
can aid in the
emulsification of fats
and oils in aqueous
solutions.
Solutions
END Chap 13
Solutions