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Syllabus Chemistry 102 Spring 2011 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: firstname.lastname@example.org
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Chemistry 102 Spring 2011
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
(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: Mary Hesse
Kotz and Treichel 7th ed. TEXTBOOKS Averill and Eldridge
Online Dictionary of Chemistry
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.
The way the real world works
Individual Mastery compared
to a large population
What you are used to and I will
1) Raw scores are determined. Sum of points assigned to correct responses
2) Individual scores are normalized. A context-free evaluation of relative performance
3) Normalized scores are transformed. An absolute score is assigned to a defined scale
4) Letter grades are assigned>89.501 A
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.
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.
E = n =
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.
Na+—Cl- in salt
These are the strongest forces.
Lead to solids with high melting temperatures.
NaCl, mp = 800 oC
MgO, mp = 2800 oC
Many metal ions are hydrated. This is the reason metal salts dissolve in water.
Attraction between ions and dipole depends on ion charge and ion-dipole distance.
Measured by ∆H for Mn+ + H2O --> [M(H2O)x]n+
Such forces bind molecules having permanent dipoles to one another.
• 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 nonpolarsolvents
– The difference between hydrophilic and hydrophobic substances has substantial consequences in biological systems.
– Vitamins can be classified as either fat soluble or water soluble.
– 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.
A special form of dipole-dipole attraction, which enhances dipole-dipole attractions.
H-bonding is strongest when X and Y are N, O, or F
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.
Active Figure 13.8
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
How can non-polar molecules such as O2 and I2 dissolve in water?
The water dipole INDUCES a dipole in the O2 electric cloud.
The alcohol temporarily creates or INDUCES a dipole in I2.
dFORCES INVOLVING INDUCED DIPOLES
Consider I2 dissolving in ethanol, CH3CH2OH.
Solubility increases with mass the gas
Formation of a dipole in two nonpolar I2 molecules.
The induced forces between I2 molecules are very weak, so solid I2 sublimes (goes from a solid to gaseous molecules).
Induced dipole-induced dipole
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
– 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.
the KE distribution of the molecules in the liquid.
the average KE of the particles in a liquid
and the range of KE for the molecules.
can escape from the liquid to enter
the vapor phase
where it is physically possible for the
molecule to leave the liquid surface.
– Does not depend on the amount of liquid
– Depends strongly on the temperature and intermolecular forces
increasing strength of IM interactionsLiquids 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
HEAT OF VAPORIZATION is the heat req’d (at constant P) to vaporize the liquid.
LIQ + heat VAP
Compd. ∆Hvap (kJ/mol) IM Force
H2O 40.7 (100 oC) H-bonds
SO2 26.8 (-47 oC) dipole
Xe 12.6 (-107 oC) induced dipole
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.)
At the TRIPLE POINTall three phases are in equilibrium.
Normal boil point 100 760
Normal freeze point 0 760
Triple point 0.0098 4.58
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.
• 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.
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.
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).
1. melting solid → liquid
2. freezing liquid → solid
3. vaporization liquid → gas
4. condensation gas → liquid
5. sublimation solid → gas
6. deposition gas → solid
Heating Curve – A plot of temperature
versus time where heat is added
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
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
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
– 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
• 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.
• 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.
1. Boiling-point elevation
2. Freezing-Point depression
3. Vapor pressure
4. Osmotic pressure
i = apparent number of particles in solution
number of moles of solute dissolved
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 diagramBoiling Point Elevation Freezing Point Depression
• 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.
• 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
– 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 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.