Transcript Slide 1

Syllabus
Chemistry 102 Spring 2010
Sec. 501, 503 (MWF 8-8:50, 12:40-1:30)
RM 100 HELD
Professor: Dr. Earle G. Stone
Office: Room 123E Heldenfels (HELD)
Telephone: 845-3010 (no voice mail) or leave message at 845-2356
email: [email protected]
(put CHEM 101-Sec. # + subject in subject line of your email)
Office Hours: HELD 408: Tue. And Thurs. 8:00-10:50 AM
I.A. Esther Ocola
S.I. Leader: Analise Castellano
All
College
BIMS
Science
GEST
Ag BICH, NUSC, GENE
Engineering
Education
Geosciences
Liberal Arts
Agriculture other
Architecture
Business
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149
118
57
35
55
18
20
10
25
0
5
100%
30%
24%
12%
7%
11%
4%
4%
2%
5%
0%
1%
501
College
BIMS
Science
Ag BICH, NUSC, GENE
GEST
Engineering
Education
Geosciences
Liberal Arts
Agriculture other
Architecture
Business
whoop
RIP
2012
2011
2010
2009
2008
260
193
33
9
0
495
82
56
28
13
28
9
14
9
9
0
2
82
33%
22%
11%
5%
11%
4%
6%
4%
4%
0%
1%
33%
73% pre-something
27% widely diverging
academic foci
503
242
67
62
22
31
27
9
6
5
10
0
3
100%
28%
26%
13%
9%
11%
4%
2%
2%
4%
0%
1%
BIMS
Science
GEST
Ag BICH, NUSC, GENE
Engineering
Education
Geosciences
Liberal Arts
Agriculture other
Architecture
Business
BIMS
Kotz and Treichel 7th ed.
TEXTBOOKS
Averill and Eldridge
Hardbound ~$200
Solution Manual ~$40
Online Tutor ~$45
Total ~$285
Ebook $45 per semester
Includes
Text
Solution manual
Online tutorial
Yvette Freeman
Publisher's Representative
Pearson Education
[email protected]
Helpful
Online Dictionary of Chemistry
Useful
As A Second Language
General Chemistry I and
Organic Chemistry I
(There are O-chem II and Physics
books in this series if you find these
useful and will have to take those classes.
Chang’s Essentials
http://slc.tamu.edu/
Tutoring
Supplemental Instruction
Courses
Texas Success Initiative
About Us
Contact Us
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Neeley Hall and the Northside Post Office
(979) 845-2724
The Student Learning Center has won the 2008 National College Learning Center Association
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The award recognizes the center's commitment to supporting and strengthening the
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The Student Learning Center provides Supplemental Instruction and tutoring free of charge to all
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students how to improve their study skills and prepare for the job market. The SLC manages
Developmental programs for students who have not yet passed the assessment tests required by the state.
Study Tips
•General
•Time Management
•Reading Textbooks
•Setting Goals
•Preparing for Exams
•Success Tips from Fellow Aggies
Tutoring
During the Fall 2008 semester, drop-in tutoring will be offered Sunday nights 5-8pm and Monday through Thursday
nights from 5-10pm. Tutoring will begin on Monday, September 1st. Tutor Zones are currently planned for Studio 12 of The
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are also available to help out with many other courses. If you need help in a particular course and would like to check to see if
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Week
Date
Averill & Eldridge
6th Edition
7th Edition
15-Mar
9
17-Mar
19-Mar
Spring Break
22-Mar 16
10
24-Mar
26-Mar
Acid/Base Equilibrium
17
Chapter 16–3, 9, 11, 13, 33, 37,
47, 55, 111, 113
Acid/Base Equilibrium
Chapter 17 –7, 11, 15, 23, 27,
3561, 69, 93, 107, 109
17
Chapter 17 –7, 11, 15, 23, 27,
35, 61, 69, 93, 103, 111, 109
29-Mar
11
12
13
31-Mar
2-Apr
5-Apr
7-Apr
9-Apr
12-Apr
14-Apr
16-Apr
14
19-Apr
21-Apr
15
23-Apr
26-Apr
28-Apr
16
17
Reading Day
17 Solubility Equilibrium
Chapter 17–1, 5, 7, 15, 23, 27,
37, 53
Exam # 3
19 Electrochemistry
Chapter 19–2, 11, 13, 43, 47,
53, 55, 99
30-Apr
3-May
Exam #4
4-May
Reading Day
7-May
Section 501 Final 10 AM
Reading Day
10-May Section 503 Final 10:30 AM
18
Solubility Equilibrium
18
Chapter 18 –1, 3, 9, 15, 33, 35, 37,
43, 53, 69, 75, 85, 103
Chapter 18 –1, 3, 9, 15, 33, 35,
37, 43, 53, 71, 77, 81, 109
20
Electrochemistry
20
Chapter 20 –1, 3, 5, 13, 25, 31, 45
Chapter 20 –1, 3, 5, 13, 25, 31,
43
Klein
How grades are determined
The way the real world works
Individual Mastery compared
to a large population
What you are used to and I will
report
A++
A+
A
B
C
D
F
F-
1)
2)
3)
4)
Raw scores are determined.
Individual scores are normalized.
Normalized scores are transformed.
Letter grades are assigned
>3
>2
>1
>0
<0
<-1
<-2
<-3
Normal
0.1%
2.1%
13.6%
34.1%
34.1%
13.6%
2.1%
0.1%
Exam 1
0.0%
1.4%
18.5%
32.7%
31.3%
14.7%
1.4%
0.0%
Exam 2
0.0%
0.4%
19.9%
33.7%
28.7%
14.9%
1.0%
0.0%
Exam 3
0.0%
0.6%
13.5%
25.7%
34.1%
20.9%
1.8%
0.0%
Exam 4
0.0%
0.0%
0.0%
0.0%
0.0%
0.0%
0.0%
0.0%
Final
0.0%
0.0%
0.0%
0.0%
0.0%
0.0%
0.0%
0.0%
Sum of points assigned to correct responses
A context-free evaluation of relative performance
An absolute score is assigned to a defined scale
>89.501 A
>79.501 B
>69.501 C
>59.501 D
<59.501 F
Grading:
Your grade will be based on
•Four one-hour examinations (200 raw points – T score 100)
•A final examination (400 raw points – 2 x T score 100)
Major Examination Schedule Spring 2010:
Major Exam No.1
Mon. Feb. 8
Major Exam No.2
Major Exam No.3
Major Exam No.4
Mon. Mar. 1
Fri. Apr. 16
Fri. Apr. 30
Final Exams
Section 501
Fri. May 7, 10:00 to 12:00
Section 503
Mon. May 10, 10:30 to 12:30
The mere formulation of a problem is far more
often essential than its solution, which may be
merely a matter of mathematical or experimental
skill. To raise new questions, new possibilities, to
regard old problems from a new angle requires
creative imagination and marks real advances in
science.
~Albert Einstein
Problem - A situation that presents difficulty, uncertainty, or perplexity:
Question - A request for data: inquiry, interrogation, query.
Answer - A spoken or written reply, as to a question.
Solution - Something worked out to explain, resolve, or provide a method
for dealing with and settling a problem.
1. Numbers – Significant Figures, Rounding Rules, Accuracy, Precision, Statistical
Treatment of the Data
2. Units – 5 of the 7
1. Time – seconds
2. Length – Meters
Density?
3. Mass – grams
Molecular Weight (Mass)
4. Amount – Moles
Mole Ratio, Molarity, molality
5. Temperature – Kelvins
3. Vocabulary – Approximately 100 new terms or words and applying new or
more rigid definitions to words you may already own.
4. Principles (Theories and Laws) – Stoichiometry, Quantum Theory, Bonding,
Chemical Periodicity, Solutions, Thermodynamics, Intermolecular Forces, Gas
Laws, Collogative Properties, Kinetics, Equilibrium, Electrochemistry
cp = q/mDT
DG = DH – TDS
PV = nRT
DT = Kmi
rate = k[A]m[B]n
∆E = q + w
Eocell = Ecathode = Eanode
[C]c[D]d
%yield = actual/theoretical * 100% K = [A]a[D]b
c (ms-1)
E = n =
l (m)
Last Semester you studied
INTRAmolecular forces—the forces
holding atoms together to form
molecules.
Now turn to forces between molecules —
INTERmolecular forces.
Forces between molecules, between ions, or
between molecules and ions.
Intermolecular
Forces
Ion-Ion Forces
for comparison of magnitude
Na+—Cl- in salt
These are the strongest
forces.
Lead to solids with high
melting temperatures.
NaCl, mp = 800 oC
MgO, mp = 2800 oC
Ion – Permanent Dipole
Attractions
••
•• water
dipole
-d
O
H
H +d
Water is highly polar and can
interact with positive ions to
give hydrated ions in water.
Attraction Between Ions
and Permanent Dipoles
Many metal ions are hydrated.
This is the reason metal salts
dissolve in water.
Attraction Between Ions
and Permanent Dipoles
Attraction between ions and dipole depends on ion charge
and ion-dipole distance.
Measured by ∆H for Mn+ + H2O --> [M(H2O)x]n+
d- H
O
H
d+
•••
Mg2+
-1922 kJ/mol
d- H
O
H
d+
d- H
O
H
d+
•••
Na +
-405 kJ/mol
•••
Cs+
-263 kJ/mol
Dipole-Dipole Forces
Such forces bind molecules having permanent dipoles
to one another.
Influence of dipole-dipole forces is seen in
the boiling points of simple molecules.
Compd
Mol. Wt.
Boil Point
N2
28
-196 oC
CO
28
-192 oC
Br2
160
59 oC
ICl
162
97 oC
Interactions in Liquid Solutions
• Hydrophilic and hydrophobic solutes
– A solute can be classified as hydrophilic, meaning that there is an electrostatic
attraction to water, or hydrophobic, meaning that it repels water.
1. Hydrophilic substance is polar and contains O–H or N–H groups that can form hydrogen bonds to
water; tend to be very soluble in water and other strongly polar solvents
2. Hydrophobic substance may be polar but usually contains C–H bonds that do not interact favorably
with water; essentially insoluble in water and soluble in nonpolar solvents
– The difference between hydrophilic and hydrophobic substances has substantial
consequences in biological systems.
– Vitamins can be classified as either fat soluble or water soluble.
1.
2.
Fat-soluble vitamins (Vitamin A) are nonpolar, hydrophobic molecules and tend to be absorbed into
fatty tissues and stored there.
Water-soluble vitamins (Vitamin C) are polar, hydrophilic molecules that circulate in the blood and
intracellular fluids and are excreted from the body and must be replenished in the daily diet.
– Ionic substances are most stable in polar solvents.
– Water is the most common solvent for ionic compounds because of its high polarity.
– A more useful measure of the ability of a solvent to dissolve ionic compounds is its dielectric
constant (), which is the ability of a bulk substance to decrease the electrostatic forces
between two charged particles.
Hydrogen Bonding
A special form of dipole-dipole attraction,
which enhances dipole-dipole attractions.
H-bonding is especially
strong in water because
• the O—H bond is very
polar
• there are 2 lone pairs on
the O atom
Accounts for many of
water’s unique
properties.
H-bonding is strongest when X and Y are N, O, or F
Hydrogen Bonding in H2O
Ice has open lattice-like structure.
Ice density is < liquid and so solid floats on water.
One of the VERY few
substances where solid is
LESS DENSE than the
liquid.
H bonds  abnormally high specific heat capacity of water (4.184 J/g•K)
This is the reason water is used to put out fires, it is the reason lakes/oceans
control climate, and is the reason thunderstorms release huge energy.
Boiling Points of Simple
Hydrogen-Containing
Compounds
Active Figure 13.8
Double helix of
DNA
Portion of a
DNA chain
H-bonding is especially strong in biological systems — such as DNA.
DNA — helical chains of phosphate groups and sugar molecules. Chains are helical
because of tetrahedral geometry of P, C, and O.
Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases.
—adenine with thymine
—guanine with cytosine
Base-Pairing through H-Bonds
Base-Pairing through H-Bonds
FORCES INVOLVING
INDUCED DIPOLES
How can non-polar molecules such as O2 and I2 dissolve in
water?
The water dipole INDUCES a dipole in the O2
electric cloud.
Dipole-induced dipole
FORCES INVOLVING INDUCED DIPOLES
Consider I2 dissolving
in ethanol, CH3CH2OH.
I-I
d
R
Solubility increases with mass the gas
d
-
I-I
O
H
d
+
The alcohol
temporarily
creates or
INDUCES a
dipole in I2.
+
d
d
O
R
• Process of inducing a dipole is polarization
• Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.
• Non-polar gases are most soluble in non-polar solvents.
H
+
d
FORCES INVOLVING
INDUCED DIPOLES
Formation of a dipole in two nonpolar I2 molecules.
Induced dipole-induced dipole
The induced forces between I2 molecules are very weak, so
solid I2 sublimes (goes from a solid to gaseous molecules).
FORCES INVOLVING
INDUCED DIPOLES
The magnitude of the induced dipole depends on the
tendency to be distorted.
Higher molec. weight → larger induced dipoles.
Molecule
Boiling Point (oC)
CH4 (methane) - 161.5
C2H6 (ethane)
- 88.6
C3H8 (propane) - 42.1
C4H10 (butane)
- 0.5
Intermolecular Forces Summary
– London dispersion forces, dipole-dipole interactions, and hydrogen bonds that hold
molecules to other molecules are weak.
– Energy is required to disrupt these interactions, and unless some of that energy is
recovered in the formation of new, favorable solute-solvent interactions, the increase in
entropy on solution formation is not enough for a solution to form.
Vapor Pressure
• When the liquid is heated, its molecules obtain sufficient kinetic energy to
overcome the forces holding them in the liquid and they escape into the
gaseous phase.
• The result of this phenomenon is that the molecules from the liquid phase
generate a population of molecules in the vapor phase above the liquid that
produces a pressure called the vapor pressure of the liquid.
• Plotting the fraction of molecules with a given KE against their KE gives
the KE distribution of the molecules in the liquid.
• Increasing the temperature increases both
the average KE of the particles in a liquid
and the range of KE for the molecules.
• Molecules with KE greater than Eo
can escape from the liquid to enter
the vapor phase
• Molecules must also be at the surface
where it is physically possible for the
molecule to leave the liquid surface.
Equilibrium Vapor Phase
• Volatile liquids have high vapor pressures and tend to evaporate readily;
nonvolatile liquids have low vapor pressures and evaporate more slowly.
• Substances with vapor pressures higher than that of water are volatile,
whereas those with vapor pressures lower than water are nonvolatile.
• Equilibrium vapor pressure of a substance at a particular temperature is a
characteristic of the material.
– Does not depend on the amount of liquid
– Depends strongly on the temperature and intermolecular forces
Liquids Equilibrium Vapor Pressure
VP as a function of T.
1. The curves show all conditions of P and T where LIQ and VAP are in
EQUILIBRIUM
2. The VP rises with T.
3. When VP = external P, the liquid boils.
This means that BP’s of liquids change with altitude.
4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING POINT
5. VP of a given molecule at a given T depends on IM forces. Here the VP’s are in
the order
ether
O
H5C2
alcohol
O
C2H5
dipole- dipole
H5C2
H-bonds
water
O
H
H
H
extensive H-bonds
increasing strength of IM interactions
Liquids
HEAT OF VAPORIZATION is the heat req’d (at
constant P) to vaporize the liquid.
LIQ + heat VAP
Compd.
∆Hvap (kJ/mol)
H2O
40.7 (100 oC)
SO2
26.8 (-47 oC)
Xe
12.6 (-107 oC)
IM Force
H-bonds
dipole
induced dipole
Phase Diagrams
Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases
on either side of the line. (At equilibrium particles move from liquid to gas as fast as they
move from gas to liquid, for example.)
Phase Equilibria — Water
Solid-liquid
Gas-Liquid
Gas-Solid
Triple Point
— Water
At the TRIPLE POINT all three
phases are in equilibrium.
Phases Diagrams—
Important Points for Water
Normal boil point
Normal freeze point
Triple point
T(˚C)
100
0
0.0098
P(mmHg)
760
760
4.58
Critical T and P
As P and T increase, you finally reach the CRITICAL T and P
Above critical T no liquid
exists no matter how high
the pressure. The gas and
liquid phases are intermixed
and indistinguishable.
At P < 4.58 mmHg and T < 0.0098 ˚C solid H2O can go directly to
vapor. This process is called SUBLIMATION
This is how a frost-free refrigerator works.
Terminology and Supercritical Fluids
• Critical temperature (Tc)
•
•
•
•
•
– Temperature above which the gas can no longer be liquefied, regardless
of pressure
– The highest temperature at which a substance can exist as a liquid
– Above the critical temperature, the molecules have too much kinetic
energy for the intermolecular attractive forces to hold them together in a
separate liquid phase
Critical pressure (Pc)
– Minimum pressure needed to liquefy a substance at the critical
temperature
Critical point — combination of critical temperature and critical pressure
As the temperature of a liquid increases, its density decreases.
As the pressure of a gas increases, its density increases.
At the critical point, the liquid and gas phases have exactly the same density,
and only a single phase exists, called a supercritical fluid, which exhibits many
of the properties of a gas but has a density typical of a liquid.
Solid-Liquid Equilibria
P
Solid
H 2O
Liquid
H 2O
Normal
freezing
point
760
mmHg
Raising the pressure at constant T
causes water to melt.
The NEGATIVE SLOPE of the S/L
line is unique to H2O. Almost
everything else has positive slope.
0 °C
T
The behavior of water under pressure is an example of LE CHATELIER’S
PRINCIPLE At Solid/Liquid equilibrium, raising P squeezes the solid.
It responds by going to phase with greater density, i.e., the liquid phase.
In any system, if you increase P the DENSITY will go up.
Therefore — as P goes up, equilibrium favors phase with the larger density
(or SMALLER volume/gram).
Changes of State
• Changes of state are examples of phase changes, or phase
transitions
• Any of the three forms of matter (gas, liquid, solid) is converted
to either of the other two
• Six most common phase changes:
1. melting
solid → liquid
2. freezing
liquid → solid
3. vaporization
liquid → gas
4. condensation
gas → liquid
5. sublimation
solid → gas
6. deposition
gas → solid
Temperature
Curves
Heating Curve – A plot of temperature
versus time where heat is added
1. Sample is ice at -23ºC
2. As heat is added, the temperature of the ice increases linearly with time
3. Slope of the line depends on both the mass of the ice and the specific heat of ice, the number of joules required to raise
the temperature of 1 g of ice by 1ºC
4. As the temperature of the ice increases, the water molecules in the ice crystal absorb more and more energy and at the
melting point, they have enough kinetic energy to overcome attractive forces and to move
5. As more heat is added, the temperature of the system does not increase further, but remains constant at 0ºC until all the
ice has melted
6. Once all the ice has been converted to liquid water, the temperature of the water begins to increase – temperature
increases more slowly than before because the heat capacity of water is greater than that of ice
7. At 100ºC, water begins to boil; temperature remains constant until all the water has been converted to steam
8. Temperature again begins to rise, but at a faster rate because the heat capacity of steam is less than that of ice or water
Temperature Curves
• Cooling curves – A plot of temperature vs. time when heat is removed
1. As heat is removed from steam at 200ºC, the temperature falls until it reaches 100ºC,
where the steam begins to condense to liquid water
2. No further temperature change occurs until all the steam is converted to the liquid;
then the temperature again decreases as the water is cooled
3. Temperature drops below the freezing point for some time — region corresponds to
an unstable form of the liquid, a supercooled liquid that will convert to a solid
4. As water freezes, temperature increases due to heat evolved and then holds constant
at the melting point
Units of Concentration
• There are several different ways to quantitatively describe the concentration of a
solution, which is the amount of solute in a given quantity of solution.
1. Molarity – Useful way to describe solution concentrations for reactions that are carried out in solution or
for titrations. Volume of a solution depends on its density, which is a function of temperature
Molarity = moles of solute = mol/L = mmol/mL
liter of solution
2. Molality – Concentration of a solution can also be described by its molality (m), the number of moles of
solute per kilogram of solvent. Depends on the masses of the solute and solvent, which are independent of
temperature.
Molality = moles of solute
kilogram solvent
3. Mole fraction – Used to describe gas concentrations and determine the vapor pressures of mixtures of
similar liquids. Depends on only the masses of the solute and solvent and is temperature independent.
Mole fractions sum to one for a given mixture.
Mole fraction () = moles of component .
total moles in the solution
4. Mass percentage (%) – The ratio of the mass of the solute to the total mass of the solution. Result can be
expressed as mass percentage, parts per million (ppm), or parts per billion (ppb). ppm and ppb are used
for highly dilute solutions, and correspond to milligrams (10-3) and micrograms (10-6) of solute per
kilogram of solution, respectively
mass percentage = mass of solute  100%
mass of solution
parts per million (ppm)= mass of solute  106
mass of solution
parts per billion (ppb)= mass of solute  109
mass of solution
Units of Concentration
– Mass percentage and parts per million or billion can express
the concentrations of substances even if their molecular mass is
unknown because these are simply different ways of expressing
the ratios of the mass of a solute to the mass of the solution
Molality and Mole Fraction
• Calculate the molarity and the molality of an aqueous solution that is 10.0% glucose,
C6H12O6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has
several medical uses. 1 mol C6H12O6 = 180 g
Molality and Mole Fraction
• Calculate the molality of a solution that contains 7.25 g of benzoic acid
C6H5COOH, in 2.00 x 102 mL of benzene, C6H6. The density of benzene is
0.879 g/mL. 1 mol C6H5COOH = 122 g
Molality and Mole Fraction
• What are the mole fractions of glucose and water in a 10.0%
glucose solution?
Vapor Pressure of Solutions and
Raoult’s Law
• The addition of a nonvolatile solute, one whose vapor pressure is
too low to measure readily, to a volatile solvent decreases the
vapor pressure of the solvent.
•
The relationship between solution composition and vapor is
known as Raoult’s law,
PA = AP0A,
where PA is the vapor pressure of component A of the solution
(the solvent), A is the mole fraction of A in solution, and P0A
is the vapor pressure of pure A.
• This equation is used to calculate the actual vapor pressure
above a solution of a nonvolatile solute.
•
•
•
•
•
•
•
Vapor Pressure of Solutions and
Raoult’s Law
Plots of the vapor pressures of both components versus the mole fractions are straight lines
that pass through the origin.
A plot of the total vapor pressure of the solution versus the mole fraction is a straight line
that represents the sum of the vapor pressures of the pure components.
The vapor pressure of the solution is always greater than the vapor pressure of either
component.
Solutions that obey Raoult’s law are called ideal solutions, in which the intermolecular
forces in the two pure liquids are almost identical in both kind and magnitude and the
change in enthalpy on solution formation is essentially zero (DHsoln  0).
Real solutions exhibit positive or negative deviations from Raoult’s law because the
intermolecular interactions between the two components A and B differ.
Negative deviation
– If the A–B interactions are stronger than the A–A and B–B interactions, each component of the
solution exhibits a lower vapor pressure than expected for an ideal solution, as does the solution as
a whole.
– The favorable A–B interactions stabilize the solution compared with the vapor.
Positive deviation
– If the A–B interactions are weaker than the A–A and B–B interactions yet the entropy increase is
enough to allow the solution to form, both A and B have an increased tendency to escape from
the solution into the vapor phase.
– The result is a higher vapor pressure than expected for an ideal solution.
Colligative Properties of Solutions
• The colligative properties of a solution depend primarily on the number of
solute particles present rather than the kind of particles.
• Not all solutions with the same molarity contain the same concentration of
solute particles.
• The sum of the concentrations of the various dissolved solute particles
dictates the physical properties of a solution.
• Colligative properties do not depend on the kinds of particles dissolved.
• Colligative properties are a physical property of solutions.
• There are four colligative properties:
1. Boiling-point elevation
2. Freezing-Point depression
3. Vapor pressure
4. Osmotic pressure
• Vapor pressure lowering is the key to all four of the colligative properties.
Lowering of Vapor Pressure and
Raoult’s Law
• Addition of a nonvolatile solute to a solution lowers the vapor pressure of the
solution.
– The effect is simply due to fewer solvent molecules at the solution’s surface.
– The solute molecules occupy some of the spaces that would normally be
occupied by solvent.
• Raoult’s Law models this effect in ideal solutions.
• This graph shows how the solution’s
vapor pressure is changed by the mole
fraction of the solute, which is
Raoult’s law.
Effect of Temperature on the
Solubility
• Solubility of a substance generally increases with increasing temperature.
• No relationship between the structure of a substance and temperature
dependence.
• Solubility may increase or decrease with temperature; the magnitude varies widely
among compounds.
• The variation of solubility with temperature is used to separate the components
of a mixture by fractional crystallization, a technique for purifying compounds.
• For fractional crystallization to be effective the compound must be more soluble
at high T than at low T: lowering the temperature causes it to crystallize out of
solution.
• Solubility of gases in liquids decreases with increasing temperature
• Attractive intermolecular interactions in the gas phase are essentially zero for
most substances.
• When a gas dissolves, its molecules interact with solvent molecules and heat is
released when these new attractive interactions form, therefore, dissolving most
gases in liquids is an exothermic process (DHsoln < 0)
• Adding heat to the solution provides thermal energy that overcomes the attractive
forces between the gas and the solvent molecules, thereby decreasing the
solubility of the gas
Fractional Distillation
• Distillation is a technique used to
separate solutions that have two or
more volatile components with
differing boiling points.
• A simple distillation has a single
distilling column.
– Simple distillations give reasonable
separations.
• A fractional distillation gives increased
separations because of the increased
surface area.
– Commonly, glass beads or steel wool are
inserted into the distilling column.
Pressure and Solubility of Gases:
Henry’s Law
• External pressure has little effect on the solubility of liquids and solids. The solubility
of gases increases as the partial pressure of the gas above a solution increases.
•
•
•
•
The concentration of molecules in the gas phase increases with increasing pressure, as does the
concentration of dissolved gas molecules in the solution at equilibrium at higher pressures.
Henry’s Law: C = kP, where C is the concentration of dissolved gas at equilibrium; P is the
partial pressure; and k is the Henry’s law constant, which is determined experimentally for each
combination of gas, solvent, and temperature: units are M/atm.
Concentration of a gas in water at a given pressure depends strongly on its physical properties.
Gases that react chemically with water do not obey Henry’s law; all of these gases are much more
soluble than predicted by Henry’s law
Boiling-Point Elevation
•
•
•
•
The normal boiling point of a substance is the temperature at which the vapor
pressure equals 1 atm.
Because the vapor pressure of the solution at a given temperature is lower than the
vapor pressure of the pure solvent, achieving a vapor pressure of 1 atm for the
solution requires a higher temperature than the normal boiling point of the solvent.
The boiling point of a solution is always higher than that of the pure solvent.
The magnitude of the increase in the boiling point is related to the magnitude of the
decrease in the vapor pressure; the decrease in the vapor pressure is proportional to
the concentration of the solute in solution.
Boiling Point Elevation
• Addition of a nonvolatile solute to a solution raises the boiling point of the
solution above that of the pure solvent.
– This effect is because the solution’s vapor pressure is lowered as described by
Raoult’s law.
– The solution’s temperature must be raised to make the solution’s vapor
pressure equal to the atmospheric pressure.
• The amount that the temperature is elevated is determined by the number of
moles of solute dissolved in the solution.
• Boiling point elevation relationship is:
Freezing Point Depression
• Dissolving a nonvolatile solute in water not only raises the boiling point of the water
but also lowers its freezing point.
• Freezing-point depression depends on the total number of dissolved nonvolatile
solute particles.
• The molar mass of an unknown compound can be determined by measuring the
freezing point of a solution that contains a known mass of solute, which is accurate
for dilute solutions.
• Changes in the boiling point are smaller than changes in the freezing point for a
given solvent, so boiling point elevations are difficult to measure and are not used to
determine molar mass.
• The freezing-point depression (DTf) is defined as the difference between the freezing
point of the pure solvent and the freezing point of the solution:
• Relationship for freezing point depression is:
Freezing Point Depression
Colligative Properties of Electrolyte Solutions
• The relationship between the actual number of moles of solute added to form a
solution and the apparent number as determined by colligative properties is called the
van’t Hoff factor (i) and is defined as
•
•
•
•
i = apparent number of particles in solution
number of moles of solute dissolved
The van’t Hoff factors for ionic compounds are
lower than expected, their solutions apparently
contain fewer particles than predicted by the
number of ions per formula unit.
Deviation from the expected value increases as the
concentration of the solute increases because ionic
compounds do not totally dissociate in aqueous
solution.
Some of the ions exist as ion pairs, for a brief time
are associated with each.
The actual number of solvated ions present in a
solution can be determined by measuring a
colligative property at several solute concentrations.
Boiling Point Elevation
Freezing Point Depression
• Notice the similarity of the two relationships for freezing point
depression and boiling point elevation.
• Fundamentally, freezing point depression and boiling point
elevation are the same phenomenon.
– The only differences are the mechanism, and the size of the
effect which is reflected in the sizes of the constants, Kf & Kb.
• The effects are easily seen in a phase diagram
Boiling Point Elevation
• What is the normal boiling point of a 2.50 m glucose,
C6H12O6, solution?
Freezing Point Depression
• Calculate the freezing point of a solution that contains 8.50 g of benzoic acid
(C6H5COOH, MW = 122) in 75.0 g of benzene, C6H6.
Determination of Molecular Weight
by Freezing Point Depression
•
The size of the freezing point depression depends on two
things:
1. The size of the Kf for a given solvent, which are well
known.
2. And the molal concentration of the solution which
depends on the number of moles of solute and the kg of
solvent.
•
If Kf and kg of solvent are known, as is often the case in
an experiment, then we can determine # of moles of solute
and use it to determine the molecular weight.
Determination of Molecular Weight
by Freezing Point Depression
• A 37.0 g sample of a new covalent compound, a nonelectrolyte, was dissolved
in 2.00 x 102 g of water. The resulting solution froze at -5.58oC. What is the
molecular weight of the compound?
Osmotic Pressure
• When a solution and a pure solvent are separated by a semipermeable membrane,
a barrier that allows solvent molecules but not solute molecules to pass through, the
flow of solvent in opposing directions is unequal and produces an osmotic
pressure, which is the difference in pressure between the two sides of the
membrane.
•
•
Because of the large magnitude of
osmotic pressures, osmosis is important
in biochemistry, biology, and medicine;
every barrier that separates an organism
or cell from its environment acts like a
semipermeable membrane, permitting
the flow of water but not solutes.
Examples of semipermeable membranes include:
1.
2.
3.
4.
5.
6.
cellophane and saran wrap
skin
cell membranes
Dialysis uses a semipermeable membrane.
Preserving fruits and meats employing cell membranes and using salt prevents bacterial growth.
Reverse Osmosis can be used to produce pure water from seawater.
Osmotic Pressure
• Osmosis is the net flow of solvent through a membrane due to
different solute concentrations; the direction of net solvent flow is
always from higher solvent concentration to lower solvent concentration
or from the side with the lower concentration of solute to the side with
the higher solute concentration.
• Osmosis is a rate controlled phenomenon. The flow from high to low
solvent is faster than low to high solvent concentration until equilibrium
is reached.
• Osmotic pressure () of a solution depends on the concentration of
dissolved solute particles:
Osmotic Pressure
•
Osmotic pressures can be very large.
– For example, a 1 M sugar solution has an osmotic pressure of
22.4 atm or 330 p.s.i.
•
Since this is a large effect, the osmotic pressure measurements
can be used to determine the molar masses of very large
molecules such as:
1. Polymers
2. Biomolecules like
• proteins
• ribonucleotides
Osmotic Pressure
• A 1.00 g sample of a biological material was dissolved in enough water
to give 1.00 x 102 mL of solution. The osmotic pressure of the solution
was 2.80 torr at 25oC. Calculate the molarity and approximate
molecular weight of the material.
Solutions of Solids
• Solutions are not limited to gases and liquids; solid solutions also exist.
• Amalgams, which are usually solids, are solutions of metals in liquid mercury.
• Network solids are insoluble in all solvents with which they do not react
chemically; covalent bonds that hold the network together are too strong to
be broken and are much stronger than any combination of intermolecular
interactions that might occur in solution.
• Most metals are insoluble in all solvents but do react with solutions such as
aqueous acid or base to produce a solution; in these cases the metal
undergoes a chemical transformation that cannot be reversed by removing
the solvent.
Aggregate Particles in Aqueous Solution
• A mixture of gases is the only combination of substances that cannot
produce a suspension or colloid because their particles are small and
form true solutions
A colloid can be classified as:
1. a sol or gel, a dispersion of solid particles in a liquid or solid in which all the
solvent has been absorbed by the solid particles, thus preventing the mixture
from flowing readily;
2. An aerosol, a dispersion of solid or liquid particles in a gas;
3. An emulsion, a dispersion of one liquid phase in another liquid with which it is
immiscible.
Interactions in Liquid Solutions
– All common organic liquids, whether polar or not, are miscible; the
strengths of the intermolecular attractions are comparable, the enthalpy
of solution is small, and the increase in entropy drives the formation of a
solution.
– If predominant intermolecular interactions in two liquids are very
different from one another, they may be immiscible, and when shaken with
water, they form separate phases or layers separated by an interface.
– Only the three lightest alcohols are completely miscible with water; as
the molecular mass of the alcohol increases, so does the proportion of
hydrocarbon in the molecule, which leads to fewer favorable electrostatic
interactions with water
•
Emulsions
Emulsions are colloids formed by the dispersion of a hydrophobic
liquid in water, thereby bringing two mutually insoluble liquids in close
contact.
• Various agents have been developed to stabilize emulsions, the most
successful being molecules that combine a relatively long hydrophobic
“tail” with a hydrophilic “head.” Examples of emulsifying agents
include soaps and detergents.
• Micelles
Micelles
– In the absence of a dispersed hydrophobic liquid phase, solutions of
detergents in water form organized spherical or cylindrical aggregates called
micelles, which minimize contact between the hydrophobic tails and water.
– In a micelle, only the hydrophilic heads are in direct contact with water, and
the hydrophobic tails are in the interior of the aggregate.
• Phospholipids – a large class of biological molecules that consist of detergent-like
molecules that contain a hydrophilic head and two hydrophobic tails; additional tail
results in a cylindrical shape that prevents phospholipids from forming a spherical
micelle
• Cells
– are collections of molecules that are surrounded by a phospholipid bilayer called
a cell membrane and are able to reproduce themselves