COURSE
OUTLINE
CHE
101________ Chemistry______
General Chemistry I
Course
Number Science Division Course
Title
__4__ ______3________
_________3_________
Credits Class Hours/Week Laboratory
Hours/Week
Required
Materials:
.
General
Chemistry:
Thompson-Brooks/Cole/Cengage Learning with On-line access,
General
Chemistry I: Course Manual
General
Chemistry I: Laboratory Manual
Any basic
calculator (scientific notation, log., and trig. functions recommended) may be
used (Programmable calculators, cell phones or computers will not be allowed for use
on tests or in laboratories, even if memories are cleared.)
Goggles
must be worn in the laboratory. They
will be supplied or can be purchased in
the
Bookstore.
14
Weeks____ ____Week
15___
Length
of Semester Final
Examination
Catalog
Description:
Basic concepts introduced. Theoretical principles of modern chemistry
integrated with descriptive and practical aspects. Topics discussed include: stoichiometry, atomic theory and the
structure of matter, Periodic Table, chemical bonding, kinetic-molecular theory
and the states of matter; gas laws, solutions, oxidation-reduction, acid-base
systems, and thermochemistry. Laboratory
work illustrates selected topics covered in lecture.
Prerequisite: High School Chemistry or CHE 100
Corequisite: ENG 101 or higher
Professor
of Chemistry
Course
Coordinator
MS
123 609-570-3381
HOME PAGE: http://www.mccc.edu/~alfarec/ http://www.mccc.edu/~dornemam/
E-mails accepted only from MCCC assigned student E-mail address alfarec@mccc.edu
TABLE OF CONTENTS
Course Outline
General Information 1
Grading and Testing 3
Topical Outline 4
Homework Schedule 6
Supplementary Materials 8
Performance Objectives 9
Specific Course Objectives 9
Lecture Notes
Unit I Introduction 31
Unit II Stoichiometry 47
Unit III Gases
75
Unit IV Crystal Structure 97
Unit V Liquids and Changes of State 125
Unit VI Atomic Structure and Periodic Properties 139
Unit VII Chemical Bonding 173
Unit VIII Chemical Reactions in Aqueous Solutions 211
Unit IX Properties of Solutions 257
Appendices:
Appendix A Names and Formulas of Inorganic Compounds 291
Appendix B Practice Exams 295
Appendix C Solutions to Practice Exams 323
Appendix D Worksheets for Self-Study Packages 335
Appendix E Grade Record Keeping Chart 361
Grading
Procedure:
Grading
will be based on the point system as indicated below.
Activity % of Total Basis for Points
Examination I 15.7 Percent 100
Examination II 15.7 Percent 100
Examination III 15.7 Percent 100
Quizzes 15.7 Total (best of 10) 100
Laboratory 21.6 1/2 of Total 140
Final Examination 15.7 Percent 100
TOTAL 640
Minimum Course Grade Assignment:*
A 595 Points (93%) B
-
510 Points (80%)
A - 575
Points (90%) C + 490 Points (77%)
B + 555 Points (87%) C 435 Points (68%)
B 530 Points (83%) D 370 Points (58%)
*Acceptable laboratory and recitation
participation and performance along with a passing grade on the final
examination are required to pass the course.
See the Course Objectives for more details.
Quizzes and
Examinations:
Weekly
quizzes will be given in recitation. No
make-up will be given. There will be
three hourly examinations given during the semester at approximately the
intervals shown in the Topical Outline.
Specific dates and locations for these examinations will be announced at
least one week in advance. It is your
responsibility to be present at all the examinations and the final exam. An unexcused absence will constitute a zero
score on any exam or quiz. An absence
will be considered unexcused if notification of the course coordinator is not
made before hand, where possible, or within 48 hours of the absence.
See
Appendix E for Grade Record Keeping Chart.
TOPICAL OUTLINE
Homework Schedule
Homework
assignments are on a weekly basis to help you learn the course material
according to the performance objectives and to help you to test your mastery of
the material considered. They are not to
be considered "exclusive" but representative of the material. If you feel the need to do additional reading
or problems, you can ask your instructor for guidance.
The only way to
learn the material is to do it yourself.
Chemistry is a problem solving course.
You can only master it with practice.
Read
Week Chap. Questions and Problems Course Manual
1a 1 3-5,
8, 9, 14-18, 21-23, 29-32, 35, Unit
1: 1 - 16
36, 45(ans. 2.0
x 104 g), 48, 53
1b 2 1,
2, 11, 18, 26, 40 Unit
2: 1
2a 2 28,
36, 44 (DO MANUAL FIRST) Unit 2: 2 - 10
2b 2 54,
58, 62, 76, 78, 112(optional) Unit
2: 11 - 13
3a 3 DO
THE MANUAL FIRST Unit
2: 14 - 16
1, 2, 5, 8, 12,
16, 24
3b 3 57(ans.
0.255 mole, 75.0 g), Unit
2: 15
62, 69(ans. 42
g), 70, 72 (Optional)
NOTE:
The programs described in Appendix D can be very helpful. If you notice ANY
difficulties in the early part of the course, I recommend immediate recourse to
these programs.
4a 12 1-4,
6, 12, 20, 22, 31, 32, 36, 60 Unit
3: 1 – 5
4b 12 14,
26, 34, 66, 67(ans. 32.0 mL), 111
5a 12 37,
38, 44, 85, 87, 88, 91, 93, 96 Unit
3: 6 – 8
5b 12 40,
45(ans. 5.91 g/L), 46, 56, 98, 108 Unit
3: 9, 10
6a 13 75,
76, 84-88, 89(ans. 3.47 x 10-22 g Unit
4: 1 - 21
& 3.79
x 10-23 cm3), 90- 92, 96, 98,
99(ans. 3.53 Å), 100, 101(ans. 2.350 Å)
6b 13 2,
4, 6, 8, 12, 15, 16 Unit
5: 1 - 5
7 13 21-24,
27, 28, 30, 31, 34, 39, Unit
5: 6 - 8
42, 45,
46, 65-70, 72, 120, 121
8 4 1,
2, 4, 7, 9, 15, 16, 18, 20, 22, Unit 6:
1 – 13
24, 39,
40, 42, 49-52, 53a, 54a,
65-67,
69, 70, 79-92, 94, 105,
107-109,
114-117, 126-128
9 5-1 to 5-6 1-11, 13-21, 24, 26-28, 30, 32, Unit 6:
14 - 17
36,
38, 41, 42
10a 6-3, 6-4 27-32, 34, 36, 38, 40, 45, 46 Unit 7: 1 - 2
10b 7 1-14,
16, 18-21, 23-25, 28, 30, 32, Unit 7: 3
34,
36, 52, 57, 59, 62a, 67-69, 71, 74
10c 8 1-4,
7-10, 13-16, 20-22 Unit
7: 4 - 5
10d 8 24,
26, 30, 32, 34, 38, 46, 48-50, Unit
7: 6 - 7
(include
hybridization, orbital &
molecular geometry in all questions)
11a 6-1, 6-2, 6-9 2-12, 14, 18, 20 Unit 8: 1
11b 5-9 67
- 71
11c 10 1-4,
6-13, 17-20, 23-27 Unit
8: 2
11d 10 31,
32, 34-36, 40, 41, 43-47, Unit
8: 3 - 6
50,
57, 61, 63
12a 5-7 46, 48, 49 Unit
8: 7 – 9
12b 6-5 49,
50, 52
12c 11 1-4,
6, 40, 41(also calc.N, ans. Unit 8:
10 – 13
4.56 N)
44, 47(ans. 0.07365 M)
12d 11-6 63(ans.
16.4 mL), 64 Unit
8: 14 - 17
(Eq.
Wt. Method for both)
13 11-4, 11-5 52a,b, 53b,c, 55a, 61d Unit
8: 18 - 24
14a 14 1,
2, 4, 5, 7, 8, 10, 14, 22, Unit
9: 1 - 7
23, 32,
36, 38, 40, 43. 44, 91
14b 14 46,
48, 50, 60, 68, 81 Unit 9: 8 - 12
SUPPLEMENTARY
MATERIALS
Textbook OWLv2 On-Line Mastery and Homework Assignments:
Tutoring
sites on my home page:
www.mccc.edu/~alfare
OPTIONAL SUPPLEMENTARY MATERIALS
AVAILABLE IN THE BOOKSTORE:
General
Chemistry as a Second Language, David R. Klein.
Chemistry
Survival Skills by
Margaret Brault and Margaret MacDevitt.
This book will help the chemistry student be more successful in the
course.
Student
Study Guide by Raymond Davis. Chapter summaries, study goals, 80 drill and
concept questions per chapter with answers.
Student
Solutions Manual by
Yi-Noo Tang and Wendy Keeney-Kennicutt.
Answers and solutions to all even-numbered end-of-chapter exercises.
Schaum's
Outline of College Chemistry
by Jerome Rosenberg. Theory and problems
with complete solutions.
Goggles
and gloves for the laboratory.
Mercer's Academic Integrity Policy:
Academic integrity is violated when a student:
A.
Uses or obtains unauthorized assistance in any academic work.
B.
Gives fraudulent assistance to another student.
C.
Knowingly represents the work of others as his/her own, or represents previously completed academic work as current.
D.
Fabricates data in support of an academic assignment
E.
Inappropriately or unethically uses technological means to gain academic advantage
Violators will be penalized in accordance with college policy.
General Chemistry I is intended to provide you with an
initial exposure to a broad realm of fundamental concepts in chemistry. It will assist you in attaining a basic
understanding of these concepts, and it will help you to develop essential
skills in these areas. The lectures,
recitation discussions, laboratory sessions, homework assignments, quizzes, and
examinations provide an integrated selection of activities which can lead you
to success, provided that you are conscientious. In order to re-ceive credit for the course,
you must at least meet the minimum requirements de-scribed below.
Additional effort and achievement will be especially rewarding however.
Participation in Chemistry laboratory courses is permitted provided the student has completed
the required prerequisites, is a minimum of 16 years of age, or by permission of the instructor
and the Dean of the division.
It is the college policy that a student taking the class as an Audit must declare this at the
time of registration, and may not attend the laboratory, may not take exams, and may not
have quizzes graded.
If you need an accommodation, you must bring the form at least 2 weeks before it will be used.
Students behavior deemed unsafe by the laboratory instructor will be grounds for removing a student from the course with a grade of WI or F.
Performance
Objectives:
1. You must satisfactorily complete on an individual basis the assigned laboratory experiments. (Missing 3 or more will constitute an F or W for the course)
2. You must participate in weekly recitations (missing 3 or more may constitute an F or W for
the course)
3. You must complete the weekly quizzes and hour tests, as assigned, showing work, thought process and/or justification for answers when necessary.
4. You must complete all of the assigned homework
5. You must achieve a passing grade on a comprehensive final examination.
6. You must complete a minimum of 6 hours of work (not counting lab write-ups) on chemistry
at home each week.
7. You must demonstrate your level of performance (see page 3 for "grading") by mastering a
large part of the material covered by lectures, videos, homework, laboratory work and the textbooks as detailed in the specific course objectives that follow.
8. Since you will need a C or better in CHE 101 to take CHE 102, it is expected that
you have a working knowledge of CHE 101 and can expect test questions on this material or
that depend on knowing it in CHE 102.
CHE 101: Specific Course Objectives BACK TO TOP
| Unit I: Introduction | Unit VI: Atomic Structure and Periodic Properties |
| Unit II: Stoichiometry | Unit VII: Chemical Bonding |
| Unit III: Gases | Unit VIII: Chemical Reactions in Aqueous Solution |
| Unit IV: Crystal Structure | Unit IX: Properties of Solutions |
| UnitV: Liquids and Changes of State | Course Laboratory Objectives |
Specific Course
Objectives:
Unit I: Introduction:
1. Convert
any number to its equivalent in scientific notation, or any number in
scientific notation to its equivalent in decimal notation.
2. Perform
mathematical operations (addition, subtraction, multiplication, division,
square root) using numbers in scientific notation.
3. Solve
simple algebraic equations for one unknown.
4. Perform
dimensional analyses to verify the proper units in a mathematical operation.
5. Identify
the number of significant figures for any number and for the result of any
mathematical operation.
6. Distinguish
between precision and accuracy.
7. Distinguish
between fundamental quantities and derived quantities.
8. Learn
the fundamental units of mass, distance, time, temperature, and charge; and the
derived quantities of area, volume, and density.
9. Learn
the metric prefixes from micro to mega.
10. Learn,
use and convert between the metric prefixes:
milli, centi, and kilo.
11. Convert
between the English and Metric systems for mass, length, volume, etc., given
the conversion factors.
12. Determine
the density, mass, or volume of any substance given two of the three
quantities.
13. Given
the density of two immiscible (do not mix) liquids ( or a solid and a liquid),
explain which would be on top (or bottom) and why.
14. State
the relationship between solid and liquid volume in the metric system.
15. Estimate
(in metric units) the approximate mass and approximate size of common household
items.
16. Distinguish
between chemical and physical properties and changes.
17. Differentiate
between extensive and intensive properties, including examples.
18. Define and give examples for the
following items:
a. Atom f. Homogeneous k. Isotope
b. Element g. Heterogeneous l. Atomic
number
c. Molecule h. Symbol m. Mass number
d. Compound i. Ion n. Atomic weight
e. Mixture j. Polyatomic
Ion
19. Write
or interpret isotopes in nuclear notation and the notation: C-12
20. Explain
the Law of Definite Proportions and the Law of Multiple Proportions.
21. Explain
formulas and equations, using examples, and be able to balance simple chemical
equations.
22. Distinguish
between kinetic and potential energy; and between exothermic and endothermic
processes, using examples.
23. Differentiate
between heat and temperature.
24. Define
the calorie and specific heat capacity and relate them.
25. Demonstrate
a knowledge of the temperature scales (Fahrenheit, Celsius, Kelvin) and be able
to convert from one scale to the others.
26. Graph
experimental data; determine the slope, intercept and equation of a straight
line graph.
27. Take
a measurement to the accuracy of the instrument with balances, meter sticks,
graduate cylinders, and quantitative glassware.
28. Name from the formula (or give the formula
from the name) the simple monatomic ions and the following common polyatomic
ions; as well as their partially hydrogenated forms:
a. NH4+ ammonium i. C2H3O2- acetate
b. OH- hydroxide j. MnO4- permanganate
c. NO3- nitrate k. CO32- carbonate
d. NO2- nitrite l. HCO3- bicarbonate
e. ClO4- perchlorate m. C2O42- oxalate
f. ClO3- chlorate n. SO42- sulfate
g. ClO2- chlorite o. SO32- sulfite
h. ClO- hypochlorite p. PO43- phosphate
1. Determine
the atomic number, atomic weight, or formula weight (using the Periodic Table)
of any substance, given its symbol or formula.
2. Explain
what the atomic weight of an element represents and why it is not an integer.
3. Distinguish
between molecular weight and formula weight.
4. Describe
the relationship between the mass of a substance and the number of moles that
mass represents.
5. Find
the mass, atomic or formula weight, or number of moles of any substance, given
two of the three quantities.
6. Describe
the relationship between the number of moles of a substance and Avogadro's
Number.
7. Calculate
the mass, moles, and number of atoms or molecules in a substance given any one
of them, a Periodic Table, and Avogadro's Number.
8. Find
the percent composition by weight of all elements of any pure compound, given
its formula and the Periodic Table.
9. Find
the mass (or moles) of an element in a compound given the mass (or moles) of
the compound, and vice versa. The
formula and Periodic Table are also given.
10. Find
the empirical formula of a compound, given the percentage composition of all
but one of the elements in the compound and the Periodic Table.
11. Find
the empirical formula of a compound given the relative number of grams or moles
of each element (or a way to get them) and a Periodic Table.
12. Given
the grams or moles of a compound, find the grams or moles of each element in
it, and vice versa.
13. Derive
the molecular formula of a compound, given the molecular weight and empirical formula.
14. Balance
simple chemical equations by inserting the proper coefficient before each
symbol or formula.
15. Explain
the meaning of a balanced chemical equation, especially with regard to the
coefficients of the substances involved.
16. Define
and determine the limiting reagent in an equation, given the equation and
quantities of reactants.
17. Determine
the number of moles, the number of particles, and the mass or volume of all
substances involved in a reaction, given the equation and quantities of one or
more reactants or products.
18. Define
solute, solvent, and solution, giving examples.
19. Define,
calculate, and use in calculations:
Weight(mass) percent, and molarity.
20. Define,
explain, and calculate with the dilution formula.
21. Given
two of the following: molarity, weight
percent, density, calculate the 3rd.
Unit
III: Gases:
1. Discuss
the three states of matter in terms of the motion and closeness of their
molecules and what determines the shape of the state.
2. Outline
the basic tenets of the Kinetic-Molecular Theory and relate the concepts in
this theory to observable properties of a gas such as temperature or pressure.
3. Relate
(and sketch) the kinetic energy (and velocity) of gases vs. temperature and
explain how they are affected by temperature.
(
4. Give
the formula for kinetic energy and use it to relate the speeds of two gases at
the same and different temperatures.
5. Define
pressure and explain its units and how it is measured.
6. Measure
the atmospheric pressure as well as the pressure of a given gas in an enclosed
container in units of millimeter mercury height (or torr), given the
appropriate Barometer and a demonstration.
7. Define
one atmosphere in terms of mm Hg or Torr.
8. Describe
qualitatively the effect on a sample of a gas if any one of the three
variables (pressure, volume, temperature) is increased or decreased, using the
Kinetic-Molecular Theory.
9. State
the ideal gas laws (Boyle's Law, Charles' Law, Gay-Lussac's Law, General law)
and use these relationships to find the resulting pressure, volume, or
temperature of a sample of a gas if the changes it undergoes in two of these
variables are adequately specified.
10. Be
able to plot and interpret Boyle's and Charles' Laws.
11. State
the generalized gas law relationship, PV = nRT, and use this equation to find
any one of the variables if the remaining quantities are specified or to
determine any of the laws in the previous objective.
12. Define
and explain S.T.P. conditions.
13. Derive
an absolute scale of temperature from the behavior of ideal gases at constant
pressure.
14. State
15. State
Avogadro's Principle and describe what is meant by the "molar volume"
of an ideal gas.
16. Calculate
the value of R in liter-atm/mole oK, given the molar volume of an
ideal gas at S.T.P.
17. Derive
the relationships between the molecular weight and the density of a gas and
measurable quantities, and use this equation in calculations.
18. State
Graham's Law and show how it can be obtained from the fact that absolute
temperature is proportional to the average kinetic energy of the molecules,
given by 1/2mv2.
19. Apply
Graham's Law to the separation of gases, the prediction of relative rates of
diffusion, or effusion, or finding a molecular (or atomic) weight.
20. Define
an ideal gas in terms of its lack of molecular volume and its lack of
attractions among its molecules.
21. Explain
why real gases do not obey the ideal gas laws perfectly, and explain the
meaning of the terms in the Van der Waals equation of state of a real gas.
22. Name
the temperature and pressure conditions (i.e., high, low) under which a real
gas behaves more ideally or non-ideally, and why this is so.
Unit
IV: Crystal Structure:
1. Define
and relate the wave properties (and their symbols); wavelength ( l ), velocity (n), frequency ( u ), and amplitude (A), and
illustrate them on a picture of a wave.
2. Be
able to use the relationships between wavelength, velocity and frequency of a
wave in calculations.
3. Define
and illustrate diffraction and interference, and state Bragg's Law for
constructive interference for both crystals and a diffraction grating.
4. Describe
the electromagnetic spectrum in terms of wavelength, frequency, energy, and
types of waves (especially the color of the visible portion.)
5. Describe
the external features of a crystal.
6. Explain
the use of x-rays in determining crystal structure.
7. Distinguish
between an amorphous solid and a crystalline solid.
8. State
and explain Bragg's Law as it applies to crystal lattices and use it in
calculations.
9. Define
and give examples of unit cell, lattice, and space lattice.
10. Describe
the geometries of the seven basic shapes of unit cells: cubic, tetragonal, orthorhombic, monoclinic,
triclinic, hexagonal, and rhombohedral.
11. Describe
the four types of unit cells: simple,
body-centered, face-centered, and end-centered, and state which ones exist for
cubic unit cells.
12. Explain
why there are only 14 crystal systems.
13. Determine
what fraction of a sphere will be in a unit cell for spheres at the following
lattice points: corner, edge,
face-center, body-center.
14. Determine
the number of metal atoms contained in each of the types of unit cells.
15. Explain
that face centered cubic (cubic closes packing) is the most efficient and most
common type of packing; especially for metals such as Cu, Al, Ag, Au, etc.
16. Be
able to apply the Pythagorean Theorem to unit cells.
17. Calculate
the radius of a metal atom, given the unit cell length (or vice versa), for a
simple cubic, face-centered cubic, and body-centered cubic unit cell.
18. Calculate
the face-diagonal of a cubic unit cell given the edge length or sphere size, or
vice versa.
19. Calculate
the % void (unoccupied) space in or density of a simple, face-centered, and
body-centered cubic unit cell with spheres of diameter or radius of one
angstrom.
20. Describe
the location, shape, number, and relative size of tetrahedral and octahedral
sites within the face-centered cubic unit cell, and the simple cubic site, and
describe how this leads to ionic crystals.
21. Describe
the packing, given the number of formula units per unit cell, and give the
general formula and an example for the four types of cubic closest packing of
ionic salts: rock salt, zinc blende,
fluorite, and antifluorite.
22. Find
the ionic radii (and the density) for a salt in the rock salt structure, given
the length of the unit cell, and vice versa.
23. Describe
the packing of the cesium chloride structure for ionic salts, given the number
of formula units for unit cell, and be able to do calculations relating the
unit cell length to the ionic radii.
Unit V: Liquids and Changes of State:
1.
Define and give examples for the following terms:
a. Compressibility j. Freezing
b. Diffusion k. Melting
c. Surface tension l. Fusion
d. Evaporation m. Crystallization
e. Condensation n. Sublimation
f. Vapor pressure o. Boiling point
g. Critical temperature p.
h. Critical pressure q. Melting
point
i. Boiling r. Triple point
2. Define
the four states of matter.
3. Explain
on a molecular level why diffusion in a liquid is slower than in a gas.
4. Explain
on a molecular level the cause of surface tension and which states of matter
exhibit it.
5. State
the differences among the three states of matter for the physical properties of
density, compressibility, and ability to flow.
6. Explain
on a molecular level why evaporation is a cooling process.
7. Define
the heat of vaporization and the heat of fusion, and use them in calculations.
8. Describe
the attractive forces in a liquid and relate them to such properties as heat of
fusion, heat of vaporization, vapor pressure, boiling point, melting point, and
surface tension.
9. Describe
vapor pressure in terms of the dynamic equilibrium that exists between the
liquid and gaseous states.
10. Define
the principle of Le Chatelier and apply it to a liquid-gas, liquid-solid, and
solid-gas equilibria in terms of volume or temperature stresses.
11. Describe
the effect of size on the polarizability of a series of like-compounds such as
the hydrides of Groups VA, VIA, and VIIIA.
12. Describe
the effects of the polarizability of a molecule on the physical properties
(e.g. boiling point, vapor pressure, heat of fusion, etc.) of that molecule.
13. Describe
the course and nature of hydrogen bonds and its effect on a series of boiling
points (and other physical properties) of like compounds such as the hydrides
of Groups VA, VIA, and VIIIA.
14. Sketch
and interpret all portions of heating and cooling curves.
15. Sketch
or interpret a "phase diagram" in terms of triple points, critical
points, normal boiling and melting points, and the lines that reflect the
equilibrium between the states or phases of a substance. All phase changes should also be included,
along with the Kinetic-Molecular Theory, in your description.
16. Explain
what would happen to a substance as its temperature or pressure were changed,
given its phase diagram.
Unit
VI: Atomic Structure and Periodic
Properties:
1.
Define and give examples of the following terms:
a. Electron f. Coulomb k.
Atomic number
b. Neutron g. Radioactivity l. Mass number
c. Proton h. Alpha particle m. Isotope
d. Cathode rays i.
Beta particle n.
X-rays
e. Nucleus j. Gamma ray
2. Determine
how many protons, neutrons, and electrons a given atom (or ion) has, given its
symbol, atomic number, and mass number (and vice versa).
3. Describe
the relative mass, charge, and location of the three basic particles in the
atom (a, b, g)
4. State
what isotope serves as the current standard for the atomic mass scale.
5. Explain
why atomic masses of some elements (i.e., Cu, Cl) are so far from whole
numbers.
6. Summarize
the contributions made to our early understanding of atomic structure by:
a. Joseph
J. Thompson e. Antoine Becquerel
b. Robert A. Millikan f. Ernest
Rutherford
c. Eugene Goldstein g. Henry
Mosley
d. James Chadwick h. Niels
Bohr
7. Describe
the three important kinds of radiation emitted by radioactive substances.
8. Summarize
Rutherford's gold-foil experiment and the conclusions that can be made from its
outcome.
9. Describe
the evidence for the existence of electronic energy levels or orbits in atoms.
10. Compare
and contrast line and continuous spectra, and list sources of each.
11. Explain
how atomic spectra are obtained, what they look like, and what they mean.
12. Outline
the Bohr concept of atomic structure.
13. Explain
the meaning of the Rydberg equation and perform a Rydberg calculation to
determine the quantum level change and corresponding energy change for a line
in the emission spectrum of an atom such as hydrogen from a measured or given
wavelength emitted. From this one should
be able to construct an energy level diagram for the atom.
14. State
the relationships between energy, wavelength, and frequency and use them in
calculations.
15. Explain
and state the differences among the Balmer, Lyman, and Paschen series in the
hydrogen spectrum and relate them to the Rydberg Equation.
16. State
Planck's quantum theory relating energy with wavelength (or frequency) of
radiation of a given frequency or wavelength.
17. State
the Heisenberg Uncertainty Principle and interpret how it describes the
limitation on our simultaneous knowledge of the momentum and position of a
moving electron.
18. State
Louis De Broglie's contribution relating the wavelength and momentum of a
particle.
19. State
Erwin Schrodinger's contribution to the quantum picture of the atom.
20. Explain
the wave function ( y ) in the solution to the
Schrodinger Wave Equation describes the electron in an atom, and how its square
( y2) relates to the probability of
finding the electron at any point around the nucleus.
21. Explain
the relationship between y2, orbitals, and electron clouds, and
how they reflect our current picture of the atom.
22. Give
the symbols for, and name the four quantum numbers which describe an electron;
state what values they may assume, and their relationship to each other. This should include the relationship between
n and K,L,M,..; and s,p,d,f,..
23. Arrange
the orbitals described by these quantum numbers in order of increasing energy.
24. Explain
what the four quantum numbers tell us in terms of energy, location, and shape of the electron in the atom,
and how many electrons can fit in each orbital.
25. Explain how many electrons fit into each
orbital and "construct" successively larger atoms or elements by
"filling" these orbitals with the appropriate number of electrons.
26. Define and give examples of the terms shell
and subshell.
27. State the similarities and differences
between the Bohr picture and the modern quantum picture of the atom.
28. State the Pauli Exclusion Principle and
Hund's Rule and apply them to the structure of the atom.
29. Draw an energy level diagram for the lowest
energy state of an element. This diagram
should be labeled as to the n and l values of each level, and with arrows up or
down for the spin of the electron in each occupied orbital.
30. Sketch the spatial arrangements of s,p, and
d orbitals.
31. Explain why there is only one s
orbital in each shell, why there are three and only three p orbitals in
each case, etc.
32. Write the electronic configuration in
spectroscopic notation (e.g. 1s2,2s22p6 etc.) for an atom of any element,
or its ion; or from the atomic number of the element.
33. Explain the few exceptions to the predicted
order of filling of subshells in terms of the stability of a half-filled
subshell.
34. Explain how the format of the Periodic Table
results from the energy levels of the orbitals.
35. Define valence electrons, and write the
valence electronic configuration for an element from its position in the
Periodic Table or locate its position from its valence configuration.
36. Write the symbols from the names of the
chemical elements and vice versa, for the first 20 elements.
37. State the Periodic Law and describe the
Periodic Table as an arrangement of the elements in the order of their atomic
numbers so that elements of similar electronic structure and similar chemical
and physical properties are in the same column.
38. Define period, group, family, A and B
family, and give examples from the Periodic Table.
39. Identify the following groups and series in
the Periodic Table and correlate their identities with their valence electronic
configurations:
a. Alkali
metals (ns1)
b. Alkaline
earth metals (ns2)
c. Halogens
(ns2np5)
d. Noble
gases (ns2np6)
e. Representative
elements (s and p groups)
f. Transition
metals (nd)
g. Lantanides
(rare earth elements) (4f)
h. Actinides (5f)
I. Inner
transition elements (nf)
40. Define the following terms, giving examples,
and describe the trends in any row and in any column of the
Periodic Table for each one:
a. Atomic
radius f. First ionization potential
b. Ionic
radius g. Electron affinity
c. Electronegativity h. Metallic nature
d. Density i. Acidic/basic strength
e. Melting
point j. Oxidizing/reducing power
41. Distinguish among the terms metal, nonmetal,
and semimetal (or metalloid) and determine which elements in the Periodic Table
fall into which category.
42. Define the term isoelectronic and pick out
from a series of atoms and ions those which are isoelectronic, and be able to
list them according to decreasing radius (or increasing ionization potential).
43. Describe a metallic lattice and metallic
conduction in terms of a "sea" of electrons, and relate this to
electrical and thermal conductivity.
44. Define the terms malleability and ductility.
45. Describe the similarities as well as the
range of chemical and physical properties of the metals (eg. ease of oxidation,
reactivity, melting points, etc.)
46. Describe the range of some of the properties
of the nonmetals.
Unit VII: Chemical Bonding:
1. Define,
describe and distinguish among ionic, covalent, polar covalent, and metallic
bonding.
2. Determine
the ion an atom will form from its position in the Periodic Table or its
valence electronic configuration, and determine the electron configuration of
an ion.
3. Name
from the formula (or give the formula from the name) the simple monatomic ions
and the following common polyatomic ions; as well as their partially
hydrogenated forms:
a. NH4+ ammonium i. C2H3O2- acetate
b. OH- hydroxide j. MnO4- permanganate
c. NO3- nitrate k. CO32- carbonate
d. NO2- nitrite l. HCO3- bicarbonate
e. ClO4- perchlorate m. C2O42- oxalate
f. ClO3- chlorate n. SO42- sulfate
g. ClO2- chlorite o. SO32- sulfite
h. ClO- hypochlorite p. PO43- phosphate
4. Determine
the formula for an ionic compound from the position of its elements in the
Periodic Table or the charges on the ions that form the compound. You should also be able to name these
compounds.
5. Describe
the Born-Haber Cycle including the terms lattice energy and heats of formation
of an ionic substance, and explain how it is used to describe this information.
6. Explain
the octet rule for ionic and covalent substances, and list which elements
usually obey it as well as those that violate it.
7. Define
Lewis Structure, and be able to write Lewis Strictire hem for elements,
monatomic ions, polyatomic ions, ionic compounds, and covalent compounds.
8. Apply
the concept of electronegativity to predict which compounds are predominantly
ionic or covalent, or more ionic or covalent.
9. Define
and give examples of the following terms:
a. Bond length d. Dipole g. Hydrogen bonding
b. Bond energy e.
Dipole moment h. Coordinate
c. Bond order f.
Polar
covalent bond
10. Define
ionic potential and apply it to determine the relative ionic-covalent character
of compounds.
11. Apply
the concept of electronegativity to predict if a bond will have a dipole or be polar,
then determine if the whole molecule will have a dipole moment, based on its
structure.
12. Use
the relative size or charge of a cation or anion to predict the relative
ionic-covalent character of a compound.
13. Relate
the relative ionic-covalent character of a compound to its properties: solubility, acid-base character, color,
melting point, and cation hydrolysis.
14. Summarize
the two most commonly used elementary descriptions of covalent bonding:
a. Valence bond (VB) theory b. Molecular
orbital (MO) theory
15. Define
sigma (s) and pi (p) bonds and sketch their formation
from atomic orbitals of appropriate symmetry.
16. Describe
and use the devices of promotion and hybridization to explain covalent bonds and
the geometries of molecules.
17. Describe
the hybridization, orbital geometry, and molecular geometry from the formula or
structure of a compound or ion.
18. Explain
how the valence bonds are formed in a molecule or ion, and list the atomic
orbitals from which they are formed.
19. Explain
the nature of multiple bonds and how they are formed in molecules, using
examples (or given an example).
20. Define
resonance and give valence bond structures for molecules or ions which exhibit
it. (e.g., NO2-, NO3-, SO2, SO3, CO3=).
21. Explain
the electron pair repulsion theory and use it to determine molecular geometry.
22. Describe
the elemental form and structure of the nonmetals, and relate this to some of
their chemical and physical properties.
23. Explain
why the nonmetal and semimetal elements of the Second Period of the Periodic
Table form stable p-p pi bonds while those of the Third Period cannot; and use
this to explain the elemental structure of these elements.
24. Define
the term "allotropic", and give examples.
25. Describe
and name the allotropic forms of oxygen and carbon, describing some of their
properties.
Unit
VIII: Chemical Reaction in Aqueous
Solutions:
A. Solution Concentrations:
1.
Define and give examples of the following terms:
a. Solution f. Saturated
b. Solvent g. Unsaturated
c. Solute h. Supersaturated
d. Concentrated I. Solubility
e. Dilute j. Equilibrium
2. Define,
give symbols for, calculate the values of, and describe the preparation of
solutions in the following concentration units:
(given appropriate data)
a. Weight percent (wt%) d. Molarity (M)
b. Parts per million (ppm) e.
Molality (m)
c. Mole fraction (X) f.
Normality (N)
3. Given
any two of the following three items:
Molarity, percent by weight, density; calculate the third one.
4. Convert
from one unit of concentration to another, given the information required.
5. Prepare
a solution of given concentration and accuracy, given the appropriate equipment
and chemicals.
6. Calculate
and be able to dilute one solution to obtain another, given all but one of the
concentrations and volumes, and the proper equipment.
B. Acids,
Bases, and Salts:
1. Define
and give examples of the following terms:
a. Acid k. Precipitation
b. Base l. Neutralization
c. Salt m. Dissociation
d. Anion n. pH
e. Cation o. Acid anhydride
f. Ionization p. Basic anhydride
g. Indicators q. Conjugate acid
h. Electrolyte r. Conjugate base
i. Nonelectrolyte s. Oxoacid
j. Hydration t. Binary acid
2. Explain, compare, and give examples for
"strong" and "weak" for the following: acids, bases, electrolytes.
3. Explain
the "dynamic chemical equilibrium" that reflects the ionization of
weak electrolytes, and compare this to a strong electrolyte.
4. Write
the balanced molecular and ionic equation for a neutralization reaction, given
the acid and base involved, or the salt produced.
5. Define
and give examples of monoprotic and polyprotic acids and write the equations
for the stepwise ionization of polyprotic acids.
6. Define
and give examples of acid salts.
7. Name
the formulas (and vice versa) of the common acids:
HCl Hydrochloric acid HMnO4 Permanganic Acid
HNO3 Nitric
acid HOAc Acetic Acid
HNO2 Nitrous
acid H2SO4 Sulfuric
acid
HClO Hypochlorous acid H2SO3 Sulfurous
acid
HClO2 Clorous
acid H2CO3 Carbonic
acid
HClO3 Chloric
acid H3PO4 Phosphoric
acid
HClO4 Perchloric
acid H2C2O4 Oxalic
acid
8. Define
an acid solution as one which contains an excess amount of hydrogen ions, H+
(sometimes called
hydronium ions) over hydroxyl ions,
9. Define
a basic solution as one which has an excess amount of hydroxyl ions, OH-, over hydrogen ions, H+, and that it will turn a pink litmus
paper blue.
10. Describe the pH scale in terms of the
relative acidic or basic strength of a solution.
11. Describe the use of indicators in
determining the pH of a solution.
12. Explain the pH of a solution of a salt made
from:
a. A strong acid
and a strong base
b. A strong acid
and a weak base
c. A weak acid
and a strong base
13. Write a precipitation equation given the
salts involved and their solubilities.
14. Give the Bronsted-Lowry definition of acids
and bases and use it to identify the acid, base, conjugate acid, and conjugate
base in a reaction.
15. Write a chemical equation to produce the
conjugate acid or conjugate base of a molecule.
16. Write the chemical equation for the
autoionization of water.
17. Give the Lewis definition of acids and
bases.
18. Describe and explain the relative strength
of the oxoacids in terms of their structure:
eg. O's without H's, central atoms and other
electron withdrawing groups.
19. Describe the relative strength of the binary
acids of a family or period in terms of the size and electronegativity of the
atoms.
20. List the formulas for the strong and weak
acids and bases in water.
C. Oxidation
- Reduction:
1. Define
and give examples of each of the following:
a. Oxidation e. Oxidation number
b. Reduction f. Oxidizing agent
c. Oxidation
state g. Reducing agent
d. Half-reaction
2. Assign
the following oxidation numbers:
a. 0 Pure elements
b. +1 Alkali metals
c. +2 Alkaline earth metals
d. -1 Halogens as halides (with metals)
e. +1 Hydrogen with nonmetals
f. -1 Hydrogen as hydride (with metals)
g. -2 Oxygen as oxide
h. -1 Oxygen as peroxide
3. Using
the above assignments and charges on polyatomic ions, assign oxidation numbers
to all other elements in a compound or ion.
4. Determine
the change in oxidation number of an element in a reaction and use this to
identify if it is oxidized or reduced and to isolate the oxidation
half-reaction and the reduction half-reaction in terms of the covalent
molecules or ions involved.
5. Balance
oxidation half-reactions and reduction half-reactions with respect to mass and
charge by the ion-electron method.
6. Balance
oxidation-reduction reactions by finding balanced ionic equations and balanced
molecular equations for both acidic and basic systems.
7. Describe
the reaction of metals with acids and give examples with chemical equations.
8. Describe
the relative ease of oxidation of the metals.
9. Describe
the activity series for metals and use it to explain single displacement
reactions.
10. Describe the reaction of metals with oxygen
and write chemical equations as examples.
11. Predict the oxidation states of a metal
based on its valence electronic configuration, and whether it is a Representative
or Transition Metal.
12. State the trend in oxidation states for the
Representative Metals that have two oxidation states; and use this to predict
which of two metals will be the stronger oxidizing or reducing agent.
D. Quantitative
Aspects of Reactions in Solutions:
1. Define
and give examples of the following terms:
a. Titration d. Equivalents
b. End
point e. Equivalent weight
c. Equivalence
point f. Normality
2. Relate
grams, equivalents, and equivalent weight; and calculate any one of them, given
the other two.
3. Determine
the number of equivalents that equals one mole of an acid, a base, an oxidized
species, and a reduced species; from its reaction (or other appropriate
information); and use it to connect between each of the following pairs:
moles equivalents
N
M 
eq. wt. 
M.W.
or F.W. 
4. Identify
the end point of a titration as when the indicator changes color and the
equivalence point as when:
# Eq. ACID = # Eq. BASE or
# Eq. OXIDIZED = # Eq. REDUCED
5. Use
the above equations to determine the number of equivalents, equivalent weight,
mass, or percentage purity of a solid titrated in an acid-base or
oxidation-reduction reaction, given the necessary information.
6. Use
the above equations to determine the number of equivalents, equivalent weight,
mass, percentage purity, volume, or normality of a solution titrated in an
acid-base or oxidation-reduction reaction, given the necessary information.
7. Properly
use burettes and perform a redox or acid-base titration in the laboratory with
a reasonable degree of accuracy.
8. Explain
the use of permanganate as an oxidizing agent, and list its advantages in a
redox titration.
.
Unit IX:
Properties of Solutions:
1. List
the three kinds of mixtures:
suspensions, colloids, and solutions; and describe them in terms of
particle size, filtration, and settling; and give examples of each.
2. Define
and give examples of each of the following:
a. Polar e. Tyndall effect
b. Nonpolar f. Emulsifying effect
c. Solvated g. Alloys
d. Hydrated
3. Name
the three types of solutions due to physical state differences and to give at
least one example for each of the three.
4. Describe
how Polar/Non-polar solute and solvent interactions can account for observed
solubilities and explain the term "like dissolves like".
5. Define
and give examples of miscible and immiscible liquid systems.
6. Define
the heat of solution (DHSOLN) and differentiate betweenan
endothermic solution process and an exothermic
solution process.
7. Describe
in detail, including molecular and energy considerations as well as molecular
attractions, the process of forming a solution with:
a. Two liquids
b. A solid and a liquid
c. A gas in a liquid
8. Apply
Le Chatelier's Principle to solubility equilibria for temperature and pressure
effects on the solubility of solution, given the state of the components and
the heat of solution.
9. Define
an ideal solution as one in which the heat of solution is zero.
10. State Henry's Law (Cg = kg Pg) and use it in calculations.
11. Define what is meant by colligative
properties of solutions and list four.
12. State Raoult's Law (PA = XA PAo), use it in calculations, and
explain why it is so on a molecular level.
13. Apply Raoult's Law to a two component system
in which one or both components are volatile, plot the vapor pressure of the
components against their mole fractions, find the total pressure on the graph,
and determine pressures of all components for any given amounts of the two
components. Or, given the pressures,
find the mole fractions.
14. Explain what is meant by positive and
negative deviations from Raoult's law in terms of:
a. The above
graph.
b. The heat of
solution.
c. The relative
attractions of the components before and after solution.
d. What happens
to the temperature as the solution is prepared.
15. Describe the freezing point depression or
boiling point elevation of a solution in terms of the effect a solute has on
the phase diagram of the solvent.
16. State the laws governing boiling point
elevation and freezing point depression of ideal solutions:
Dtb
= Kb m and
Dtf
= Kf m and to
a. Define the
terms in these equations
b. Calculate any
one term from the other two
c. Calculate
the molecular weight of a solute
d. Calculate the
freezing point or boiling point of a solution
17. Explain the effects of electrolytes and
solutes on colligative properties, including calculations.
18. Define and give examples for:
a. Osmosis e. Crenation
b. Osmotic
pressure f. Isotonic
c. Semi-permeable
membrane g. Hypotonic
d. Hemolysis h. Hypertonic
19. State and explain the van't Hoff equation
for osmotic pressure:
PV = nRT,
and use it in calculations, including the obtaining of molecular weights.
20. Sketch an experimental apparatus whereby
osmotic pressure may be measured and discuss three possible modes of action of
a semi-permeable membrane on a molecular level.
Course Laboratory
Objectives:
1. Expand your understanding of the Course
Objectives.
2. Learn to manipulate chemicals and glassware
by working alone.
3. Learn to collect and analyze data from an
experiment by working alone.
4. Learn how to use laboratory balances.
5. Learn how to do quantitative analysis such as
titrations, pipetting and preparation of
6. learn how to collect and treat data on the
computer.
7. Utilize critical thinking and quantitative
reasoning skills in observing, organizing and analyzing data, synthesizing
information, interpreting results, and communicating the results of the
analyses and laboratory investigations orally and in writing.
8. Perform chemical experimentation in a safe
and scientific manner, using proper scientific and laboratory safety
procedures.
for answers when necessary on laboratory reports. They should
also be clear and legible.
Specific Course Objectives
You should be able to:
Unit I: Chemical Thermodynamics:
1. Define the following terms, using examples where appropriate:
a. State function j. Temperature s. DE
b. Internal energy k. Standard State t. DH
c. Enthalpy l. System u. DHo
d. Entropy m. Surroundings v. DHfo
e. Free Energy (Gibbs) n. Isothermal w. DS
f. Endothermic o. Adiabatic x. DSo
g. Exothermic p. Heat capacity y. DG
h. Heat q. Specific Heat z. DGo
i. Work r. Reversible process aa. DGfo
2. Distinguish between heat and temperature and describe how each is measured (in cal and joules).
3. Define the thermodynamic standard state of 298 oK and 1 atm pressure.
4. Distinguish those properties of a system which are state functions (P, V, T, E, H, S, G) from those which are not (q, w); and those which are thermodynamic functions (E, H, S, G) from those which are not (P, V, T).
5. Use the first law of thermodynamics to calculate any of the quantities involved given the other two.
6. State the second and third law of thermodynamics and explain what they mean.
7. Define heat of formation and explain how they are obtained.
8. Define Hess's Law and discuss its implications.
9. Calculate DH for a reaction given the appropriate data, such as DHf data.
10. Distinguish between a chemical change and a physical change, especially in terms of thermodynamic state functions.
11. Relate the heat change to constant pressure (qp) and at constant volume (qv) to DH and DE.
12. Determine the enthalpy change for a substance which undergoes a temperature change and/or a change in state, given the appropriate heat capacities and DHFUS, DHVAP, or DHSUB values.
13. Determine the entropy change (DS) for a reaction or phase change, given the appropriate data such as So, DSFUS, DSVAP, and DSSUB.
14. Relate the concept of entropy to a physical system or event which involves an entropy change only, describing the relationship between entropy and disorder.
15. Describe the two fundamental laws of nature as:
a. A system tends to attain a state of minimum energy
b. A system tends towards a state of maximum disorder
16. Write a mathematical expression for the Gibbs Free Energy change, DG, at a constant temperature and use the relationship to find an unknown, given the values of the remaining quantities.
17. State the relationship between the sign of DG and the spontaneity of a reaction, and state the equilibrium condition.
18. Given values of DH and DS for a system, and assuming only small changes in these values with temperature, indicate the effect of a temperature change on the reaction.
19. Given a table of DGof values, calculate DGo for a reaction and predict the direction of spontaneous change.
20. Determine the temperature at which a particular reaction just becomes spontaneous, given DH and DS for the process.
21. List the essential parts of a calorimeter and describe how it functions, for both constant volume and constant pressure.
22. Calculate the heat equivalent of the calorimeter, given the observed temperature rise, the mass of water in the calorimeter, the total heat energy given to the water and the calorimeter, and other pertinent information.
23. Calculate the heat of a reaction, given the heat equivalent of the calorimeter, the mass of reactants, the heat capacity of the products, and the temperature rise.
Unit II: Chemical Kinetics:
1. Define the following terms (using examples where appropriate) especially in terms of a "reaction profile curve":
a. DH reaction e. Exothermic
b. Activation energy, Ea f. Endothermic
c. Reaction coordinate g. Activated complex
d. Transition state h. Reaction coordinates
2. Define the rate of a chemical reaction.
3. Given any two of the following: Eaf, Ear, DH; calculate the third, and locate them on a "reaction profile" or Arrhenius diagram.
4. Account for the rate or reaction in gas phase reactions in terms of collision of molecules.
5. Explain what is meant by effective collisions and why so few collisions result in product molecules being formed.
6. Name all six factors: Nature of reactants, state of subdivision, temperature, catalysis, concentration, and pressure (gas reactions) upon which the rate of reaction depends, and explain where they appear in the rate law.
7. Explain how each of the above factors affects the rate of reaction including your discussion of the "reaction profile curve", and of the "molecules eye-view" (collision theory).
8. Differentiate among a homogeneous catalyst, a heterogeneous catalyst, and an inhibitor.
9. With respect to rate laws, define the terms (using examples where appropriate):
a. Rate e. Reaction mechanism
b. Rate constant f. Elementary process or step
c. Order of the reaction g. Rate determining step
d. Molecularity
10. For a given reaction such as: A + 3 B ssssd 2 C
a. Describe the rate of reaction in terms of the disappearance of A or of B or the formation of C.
b. Quantitatively correlate the rate of disappearance of A to that of B as well as to the rate of formation of C.
c. Write a general rate law for the reaction.
11. Given the measured initial rates of a reaction:
k
a A + b B sssssd products
where the initial concentration of each reactant is varied over a sufficient number of trials, determine the rate law:
Rate = k [A]x [B]y including x, y, and k values and
the order of the reaction.
12. Once a rate law is known, determine an initial rate given any set of initial concentrations.
13. Write the rate law for an elementary process.
14. Describe reaction mechanisms as a sum of elementary processes.
15. Given a number of steps in a reaction mechanism, and the rate constant of each step, identify the rate-determining step with the slowest step in the entire mechanism and determine the rate law for the overall reaction.
16. Given the mechanism and rate law, determine which step is the rate determining step.
17. Explain the sequence in a chain reaction.
Unit III: A. Chemical Equilibrium:
1. Define and explain the law of mass action, equilibrium, and equilibrium constant.
2. Write the mass action expression for any reaction given a balanced chemical equation.
3. For a reversible chemical reaction: a A + b B qwe c C + d D
derive the equilibrium constant expression:
[C]c [D]d
Kc = ¾¾¾¾¾¾
[A]a [B]b
by utilizing the dynamic equilibrium concept, (i.e., rate forward = rate reverse at equilibrium).
4. Given the concentration of all the products and the reactants involved in a reversible reaction, determine the numerical value of the equilibrium constant for that reaction.
5. Know that the equilibrium constant is a constant at a given temperature.
6. State LeChatelier's principle in your own words and apply it to a given system at equilibrium under the change of one of the following factors: temperature, concentration, pressure or volume (gas reactions only), addition of inert gases, addition of a catalyst; and to predict the direction of shift in the equilibrium position as well as the change (or lack of change) in the value of Kc due to each of the above factors.
7. Differentiate between a homogeneous equilibrium and a heterogeneous equilibrium.
8. In a heterogeneous equilibrium between gaseous and liquid solutions, represent the concentration of each gaseous species by its partial pressure raised to the appropriate power and the concentration of each species in liquid solution by its concentration in moles per liter.
9. In a heterogeneous equilibrium between gaseous and liquids or solids, note that the concentrations of the liquids or solids is constant and write the appropriate law of mass action.
10. Where it applies, define the equilibrium constant in terms of partial pressures only (Kp) and note that, while Kp is still a constant, it has a different value than Kc.
11. Calculate Kc from Kp and vice versa.
12. Differentiate between concentrations and activities.
13. Given the equilibrium constant, Kc, numerically, and the initial concentrations and/or partial pressures of all reactants and products involved, calculate the concentrations and/or partial pressures of all species involved at equilibrium.
14. Given initial concentrations and one equilibrium concentration, calculate Kc and the other equilibrium concentrations, and vice versa.
15. Convert natural logarithm into logarithm to the base 10:
DGo = - 2.303 RT log K
16. State the relationship between the Gibbs free energy and the mass action expression, Q:
DG = DGo + RT In Q [C]c [D]d
Q = ¾¾¾¾¾¾
[A]a [B]b
for a reaction such as: a A + b B qwe c C + d D
and derive the relationship between the standard free energy change of the reaction and the equilibrium constant: DGo = -RT ln K
by using the thermodynamic criteria for an equilibrium:
DG = O and Q = K
17. Given the value of DGo for a given reaction at a given temperature you shall be able to calculate K for the reaction and vice versa.
18. Understand the meaning, use, and conditions of the relationship:
K2 DH (T2 - T1)
Log ¾¾ = ¾¾¾¾¾¾¾¾
K1 2.303 R T1T2
19. Assuming that DHo and DSo are independent of temperature and given the equilibrium constant at a temperature T1, and the above equation, calculate the equilibrium constant at a new temperature T2.
Unit III: B. Spectrophotometry:
1. Define the terms percent transmission (%T) and absorbance (A) and calculate one from the other.
2. Standardize and take data (%T or A) from a Spectronic 20 spectrophotometer.
3. Plot and explain the absorbance versus wavelength graph for a substance and determine the maximum absorbance.
4. State the Beer-Lambert law: A = abc and explain all letters in it.
5. Plot and explain a Beer-Lambert law graph and use it to find the concentration or absorbance of a substance, given one of them.
6. Calculate the absorbtivity constant, the absorbance, or the concentration of the solution from the Beer-Lambert Law given any two of them and the path length of the solution (b).
7. Apply the above to state and explain the equilibrium which forms Fe(SCN)2+, and how to find K for the equilibrium.
Unit IV: Electrochemistry:
1. Balance oxidation-reduction (Redox) reactions and name the oxidizing agent and reducing agent.
2. Given a balanced half reaction, find the equivalent weight of either the oxidizing agent or the reducing agent.
3. Given the number of equivalents of an oxidizing agent used in a redox titration at the end point, find the number of equivalents of the reducing agent, and vice versa.
4. Given the volume and the normality of the oxidizing agent, calculate the normality of the reducing agent, if you were given its volume at the end point in a titration, and vice versa.
5. Relate the weight, equivalent weight, and normality in a redox titration and use them in calculations.
6. Define the following terms, using examples where appropriate:
a. Oxidation h. Faraday
b. Reduction i. Coulomb
c. Electrolytic cell j. Electrode
d. Electrolysis k. Cell potential
e. Voltaic or galvanic cell l. Standard potential
f. Cathode m. Electromotive force or emf
g. Anode n. Reduction potential
7. Differentiate between a strong electrolyte and a weak electrolyte by their abilities to conduct a direct current of electricity.
8. In cells, distinguish between electrolytic and voltaic (or galvanic cells), metallic and electrolytic conduction, cathode and anode.
9. Predict the electrode reactions (at the anode and cathode) that occur during the electrolysis of molten sodium chloride and other salts, and aqueous sodium chloride (dilute and concentrated); and be able to draw a diagram of the cells.
10. Define Faraday's law of electrolysis.
11. Quantitatively associate the number of Faradays (Coulombs, or Amps) of electricity passing through the cell with the number of equivalents of the element being reduced at the cathode (mostly metallic elements and also the hydrogen gas from an acid solution), or oxidized at the anode, and with calculations in either direction.
12. Given the half reactions, construct and describe a diagram of a galvanic or voltaic cell.
13. Give the diagram, reaction, and purpose of the standard hydrogen electrode.
14. Given a standard hydrogen electrode (or any other half reaction) and an accurate differential voltmeter, measure the standard reduction potential of a cell formed with another half reaction, and calculate the reduction potential of that half reaction.
15. Diagram a complete Voltaic cell consisting of two half cells, label the anode, the cathode, the salt bridge, the direction of electron flow, the direction of the cation flow or the anion flow across the salt bridge, and explain what occurs at the anode and cathode.
16. Given the standard reduction potentials of two standard electrodes, couple them to obtain a positive standard cell potential, and determine the direction in which the reaction will be spontaneous.
17. Write the short-hand notation for an electrochemical cell according to the convention.
Example: Zn(s) ½ Zn2+(1M) ½½ Cu2+(1M) ½ Cu(s) Eo = 1.10 v
and to construct a diagram of the cell from the notation.
18. Define and use the electromotive series to determine cell potentials, reactions and spontaneity.
19. Predict the effect of concentration changes on the potential of a cell.
20. Give the relationship between Gibbs free energy and the cell potential:
DG = - nFE or DGo = - nFEo
and use it to calculate the Gibbs free energy or the cell potential.
21. Using the previous relationships and the thermodynamic relationship:
DG = DGo + RT 1n Q derive the Nernst equation:
E = Eo - 
where Q = 
for a reaction
22. Use the Nernst equation to:
a. Show that E = Eo when concentrations (or pressures) are the standard state values of 1M (or 1 atm).
b. Show that Eo = 
log Kc at 298 oK & at equilibrium.
c. Calculate equilibrium constants, solubility products, pH, free energy changes, and all potentials.
Unit V: Acids and Bases:
1. Give (with examples) the following definitions of acids and bases:
a. Arrhenius b. Bronsted-Lowry c. Lewis
and recognize acids and bases by applying the definitions.
2. Define the following terms, using examples where appropriate:
a. Acid anhydride f. Leveling effect
b. Basic anhydride g. Leveling solvent
c. Amphoteric h. Differentiating solvent
d. Conjugate acid (or base) i. Hydrolysis
e. Conjugate acid-base pair j. Solvolysis
3. Identify which elements tend to form acidic or basic anhydrides and illustrate their formation.
4. Illustrate the Bronsted-Lowry acid-base theory and identify conjugate acid-base pairs.
5. Illustrate, with reactions, the amphoteric nature of some substances.
6. Determine the effect of size and electronegativity on the strengths of binary acids of a Family and of a Period, and list the three binary acids that are strong in water.
7. Interpret the strengths of acids and bases by employing the leveling effect of the solvent, the electron withdrawing effect of electronegative atoms such as oxygen, chlorine and fluorine, the polarization of water by metal ions, and the oxidation number of the element bonded to an OH group.
8. Relate the strengths of acids or bases to their percent ionization in solution.
9. Differentiate between the acid/base strength of the series of oxo-acids of metals and nonmetals.
10. Distinguish among monoprotic, diprotic and triprotic acids and their stages of ionization.
11. List five examples of strong and weak acids and three of strong and weak bases.
12. Define and give examples of neutralization and the products of this reaction.
13. Relate grams, equivalents, and equivalent weight for an acid-base reaction; and calculate any one of these, given the other two.
14. Relate equivalents, equivalent weight, weight, normality, and volume in an acid base titration, and use it in calculations.
Unit VI: Ionic Equilibria:
1. Define the following terms, using examples where appropriate:
a. Ionization constant i. Common ion effect
b. pH and pOH j. Complex ions
c. Weak electrolyte k. Instability constant
d. Dissociation l. Formation constant
e. Polyprotic acid m. Hydrolysis
f. Indicator n. Solubility product constant
g. Buffer o. Equivalence point
h. Hydrolysis constant p. Endpoint
2. Describe the ionization of water and its ionization constant.
3. Calculate the hydrogen ion concentration and the hydroxide ion concentration of pure water.
4. Given one of the following: [H+], [OH-], pH, or pOH of a solution, calculate the other three.
5. Given the concentration of a strong acid or strong base, calculate the pH and pOH of the solution.
6. Calculate the pH of weak acids or bases, given their equilibrium concentrations, and vice versa.
7. Given the ionization constant and the initial (total) concentration of a monoprotic weak acid, or a monohydroxy weak base, calculate the hydrogen ion concentration and the hydroxide ion concentration and the concentration of all other species of the solution at equilibrium (you should make the appropriate assumptions).
8. From the information given in (7), calculate the percent ionization in a monoprotic weak acid solution or in a monohydroxy weak base solution.
9. Calculate the equilibrium concentrations of all species present when a polyprotic weak acid dissociates.
10. Given the initial (or total) concentration of a monoprotic weak acid or a monohydroxy weak base and given the pH or the solution at equilibrium, calculate the ionization constant of the weak acid or the weak base, and vice versa.
11. Explain the nature, preparation and use of buffer solutions.
12. Given the initial concentrations of a weak acid and its salt or that of a weak base and its salt (buffers) and the ionization constant(s), deduce the equili-brium conditions with proper assumptions and calculate the resulting pH of the mixture at equilibrium, and after dilution or small additions of acids or bases.
13. Use the common ion effect to calculate equilibrium concentrations or the equilibrium constant for weak acids or weak bases, given initial concentrations of all species.
14. Illustrate the three kinds of hydrolysis:
a. Salt of a strong base and a weak acid
b. Salt of a weak base and a strong acid
c. Salt of a weak base and a weak acid
and apply the ideas of hydrolysis to calculate the concentration of all ions and the pH at equilibrium, given initial conditions, or predict the acidic, basic, or neutral nature of salts in water.
15. Standardize a pH meter and correctly measure the pH of a solution with a pH meter.
16. Given the ionization constant of a weak acid or a weak base, and given the weak acid or the weak base and a salt (strong electrolyte) of the acid or of the base and necessary apparatus, prepare a required volume of a buffer solution of a desired pH.
17. Given the initial concentration of the titrants and the ionization constants where applicable, predict the end point and the shape of a titration curve of pH against volume of acid or base added to a base or to an acid, respectively, for the following cases:
a. Strong acid titrated with a strong base (or the reverse)
b. Weak acid titrated with a strong base
c. Strong acid titrated with a weak base.
18. Calculate the pH at any point of the addition in (17).
19. Perform any of the titrations in (17) in the laboratory, properly using burettes.
20. Explain the nature of the curve in the titrations of (17) in terms of the vertical rise and the two plateaus and why they are so.
21. Explain an alternate means of preparing the solution that exists at the end point in (17).
22. Select the right indicator according to the range of pH in which the end point of an acid-base titration lies.
23. Calculate the solubility (or concentration of the ions) given the solubility product constant, and vice versa.
24. Use the common ion effect to calculate equilibrium concentrations or the equilibrium constant for slightly soluble salts, given initial concentrations of all species.
25. Given the concentration of a solution of a cation and the concentration of a separate solution of an anion, of a slightly soluble salt, and given its Ksp, mathematically determine if a precipitate will form if given volumes of the two solutions are mixed.
26. Given the information in (25), for the case where a precipitate forms, calculate the number of moles (and grams) of the solid formed, the percent precipitation, and the final concentration of each of the ions remaining in solution.
27. Predict, mathematically, which ion will precipitate when a precipitating agent is added to a solution of two or more ions.
28. Determine the molar solubility of salts in solvents that form complex ions with the solute added.
29. Write instability constant and formation constant expressions from the chemical equation.
30. Relate the instability constant to the formation constant for complex ion formation.
Unit VII: Chemistry of the Representative Elements I: The Metals:
1. Distinguish among metals, nonmetals, and metalloids (semi-metals) with respect to chemical properties, physical properties, and positions in the Periodic Table.
2. Write the outer shell electron configuration of any of the representative elements.
3. From the electron configuration of any element, determine which family or group it belongs to and vice versa.
4. Describe the reactions of the representative metals, their oxides, and their hydroxides with water, acids, or bases.
5. Describe the trends in metallic behavior , electronegativity, ionization energy, electron affinity, and atomic radii throughout the periodic table.
6. Deduce, using simple thermodynamics, what type of chemical reaction can be used to produce free metals from their compounds.
7. Illustrate some similarities in chemical behavior of the Group IA, IIA and IIIA elements, especially diagonal relationships, and the relative reactivities within each group.
8. Interpret diagonal relationships in terms of ionic potential.
9. Predict and explain the values of the stable oxidation states for the representative metals, and which will be more stable.
10. Interpret the trends in oxidation states exhibited by the atoms within a group in terms of the relative stabilities of high and low oxidation states.
11. Describe trends in any row or column of the periodic table with respect to:
a. Atomic radius f. Metallic properties
b. Ionic radius properties g. Oxidizing/reducing properties
c. Ionization potential h. Ionic potential
d. Electron affinity i. Polarization of ions
e. Electronegativity j. Hydrolysis
12. Use ionic potential to compare the relative degree of ionic-covalent bonding and physical properties of compounds composed of the representative elements.
13. Discuss the Solvay Process
Unit VIII: Chemistry of the Representative Elements II: The Metalloids and Nonmetals:
1. Define the following terms, including examples where appropriate:
a. Allotropism e. Disproportionation
b. Catenation f. Polymers
c. Three center bonds g. Oxoanion
d. Amorphous h. Hydride
2. Compare metalloids and nonmetals in terms of the oxidation states displayed and the processes employed in their production.
3. Contrast the methods of preparation of the metalloids with those for the production of the nonmetals.
4. Describe the molecular structure, bonding, geometry and name of the allotropic forms of the pure metalloids and nonmetals.
5. Predict the important oxidation states of the nonmetals and metalloids.
6. Determine the oxidation state of the nonmetals and metalloids in ions and in neutral compounds.
7. Illustrate examples of catenation among nonmetals and metalloids by drawing structural formulas, and which element does it best.
8. Describe the two general methods for the preparation of nonmetals and metalloid hydrides.
9. Relate the ease of preparation and stability of nonmetal and metalloid hydrides to their standard enthalpies and free energies of formation.
10. Compare the relative acidic strength of the hydrides for the elements in both vertical columns and horizontal rows.
11. Write equations for the hydrolysis of nonmetal anions such as sulfide, nitride, phosphide and carbide.
12. Draw the structure of diborane and describe the bonding in this substance and why BH3 is not the simplest stable boron hydride.
13. Write the equation for the reaction of nonmetal oxides with water.
14. Describe three methods of preparation of nonmetal oxides, writing chemical equations for each.
15. Give the formulas of the important nonmetal oxides.
16. Give the structures, hybridization, and resonance forms of NO, NO2, CO, CO2, SO2, SO3 and show the valence bonds ( d and p ) that form.
17. Write equations for the preparation of the compounds in (16).
18. Compare the structures of nonmetal oxides on the bases of bonding preferences exhibited by the non-metals.
19. Compare the structures of P4, P4O6 and P4O10, and give reaction for the preparation of the oxides from phosphorous.
20. Discuss the molecular structure of quartz.
21. Give examples of 3 methods of preparing oxoacids and their anions.
22. Given the structure or the name or the formula for the following oxoacids (and the oxoanions), given one of them:
a. HClO f. H2SO4 k. H3PO3
b. HClO2 g. H2S2O3 l. H3PO4
c. HClO3 h. HNO2 m. H2CO3
d. HClO4 i. HNO3 n. H3BO3
e. H2SO3 j. H3PO2 o. H2C2O4
and extend these structures to other members of the same families where appropriate.
23. List the oxidation state for the central element in the oxoacids and oxoanions in (22).
24. Give the bonding, geometries, resonance forms, and hybridization (where appropriate) for the oxoacids and oxoanions in (22).
25. Give the names and structures for the salts which form from the oxoacids in (22).
26. Give the equation for the formation of the oxoacids from the anhydrides, where appropriate (from CO2, N2O3, N2O5, P4O6, P4O10, SO2, SO3).
27. Compare the acidic strengths and oxidizing abilities of the oxoacids of the nonmetals.
28. Predict formulas for the halogen compounds of the nonmetals. Give geometries for these compounds based on the electron repulsion theory and list hybridizations where appropriate.
29. Use electronic structure and relative size of the atoms to determine possibility for existing and relative stability of the nonmetal halogen compounds.
30. Compare and explain the relative reactivities among the halogens and among the noble gases.
31. Explain the valence bond formation in the N2 molecule.
32. Explain why nitrogen is relatively unreactive, when compared to other nonmetals.
33. Define and explain nitrogen fixation and the nitrogen cycle in nature.
34. Discuss the preparation of and bonding in noble gas compounds. Include geometries and hybridization.
35. Indicate the composition of the two most abundant components of the atmosphere.
36. Indicate the six major pollutants of the air.
37. Indicate five sources for these pollutants.
Unit IX: The Transition Elements:
1. Describe similarities and differences between A and B groups of the periodic table.
2. Distinguish between representative, transition, and inner transition elements.
3. Explain why Group IIB is sometimes considered a representative group.
4. Compare the properties among the transition elements horizontally as well as vertically including atomic radii, ionic radii, important oxidation states, ionization energy, hardness, melting points, and density.
5. List at least five characteristics that the transition elements have in common with each other.
6. Write the electronic configurations of the first row transition elements, noting the anomalies and the reason for them.
7. Define "lanthanide contraction" and predict its effect on the properties of the transition elements in period 6.
8. Predict the possible oxidation states of the transition metals and give the more important oxidation states of the first row transition metals.
9. Indicate the relative importance of higher and lower oxidation states as one moves horizontally or vertically among the transition metals.
10. Use the relative stabilities of oxidation states to determine which of 2 compounds will be more easily oxidized (or reduced) or which will be the better oxidizing agent (or reducing agent).
11. Give formulas and names to the more important oxides and hydroxides (and their anions) of the first row transition metals and compare their relative oxidizing abilities.
12. Compare the relative acidity of the oxides and of the hydroxides of each transition element that has more than one important oxide of hydroxide.
13. Explain the use of silver in the black and white photographic process.
14. Explain the physiological action of mercury.
15. Discuss the coinage metals and why they are called that.
16. Discuss the two oxidation states of mercury and the unique structure and bonding they produce.
17. List the iron, palladium, and platinum triads, and why Group VIII is structured that way.
18. List the platinum metals and some of their important properties.
19. Define or describe the following terms relating to metallurgy, giving examples where appropriate:
a. Ore i. Blast furnace
b. Amalgam j. Cast iron
c. Flotation process k. Pig iron
d. Gangue l. Steel
e. Slag m. Calcination
f. Flux n. Bessemer converter
g. Roasting o. Open hearth furnace
h. Smelting p. Mond process
20. Identify and describe the three main steps involved in extracting a metal from its ore, and give examples of each.
21. Define an alloy.
22. Name at least two important alloys and describe their composition and applications.
23. List at least two properties in which an alloy differs from its components.
24. Define and compare the terms paramagnetism, ferromagnetism and diamagnetism, and give examples of elements exhibiting each type.
25. Define the term domain and relate it to the degree of magnetism which a substance exhibits.
26. Define the following terms relating to coordination chemistry, giving examples where appropriate:
a. Complex compound l. Enantiomers
b. Coordinate covalent bond m. Racemic
c. Ligand n. Inner orbital complex
d. Coordination sphere o. Outer orbital complex
e. Chelating group p. High spin complex
f. Monodentate ligand q. Low spin complex
g. Polydentate ligand r. Degenerate
h. Coordination number s. Crystal Field Theory
i. Stereoisomerism t. Crystal field splitting
j. Geometrical isomerism u. Valence Bond Theory
k. Optical isomerism v. Donor atom
27. Given the formula or structure of a transition metal complex, identify the ligands, chelating groups, coordination sphere, coordination number, and donor atom.
28. Name transition metal complexes (using the rules of nomenclature) given the formula, and vice versa.
29. Identify and draw isomers of some transition metal complexes, identifying cis, trans, and optical isomers, or given the structure, identify which isomer is present.
30. State what nonsuperimposable mirror images means and how this relates to coordination compounds and optically active coordination compounds.
31. Define polarized light and explain what happens to it when it is passed through a solution of each of a pair (or a mixture) of optical isomers.
32. Use the valence bond theory to explain the nature of the bonding in coordination complexes.
33. For transition metal complexes, use the valence bond theory to explain:
a. The nature of the coordinate covalent bond.
b. Their electron configuration (before and after complexing).
c. Their structure, geometry, and hybridization.
d. Whether an inner or outer orbital complex will form (be more stable).
e. The number of unpaired electrons that results.
f. Their magnetic properties.
g. Their color.
h. The faults in the theory.
34. For transition metal complexes, use the crystal field theory (ligand field theory) to explain:
a. The nature of the bond formed and compare it to the coordinate
covalent bond
b. The effect of the ligands on the energy levels of the central
metal ion.
c. The splitting and labeling of the energy levels above.
d. The crystal field splitting, D, and its relationship to the
spectrochemical series of ligands.
e. Their geometry or structure.
f. The number of unpaired electrons that results, including
whether low spin or high spin complexes will result.
g. Their magnetic properties.
h. Their color.
i. Their relative stabilities.
Unit X: A. Nuclear Chemistry:
1. Define or describe the following terms, using examples where appropriate:
a. Nuclide m. Accelerator
b. Isotope n. Natural radioactivity
c. Alpha (a) particles o. Nuclear transformation
d. Beta (b) particles p. Band of stability
e. Gamma (g) rays q. Electron capture
f. Parent/daughter isotopes r. Magic numbers
g. Radioactive (decay) series s. Nuclear fission
h. Half-life t. Nuclear fusion
i. Atomic mass u. Critical mass
j. Mass number v. Plasma
k. Transmutation w. Chain reaction
l. Nuclear force y. Induced fission
2. Describe the composition of the nucleus.
3. Explain the factors influencing the change of an unstable nucleus to a more stable nucleus.
4. Describe natural radioactivity and the types of decay it produces.
5. Give the symbols and properties for the three basic emissions in natural radioactivity (a,b,g).
6. Complete and balance nuclear equations given all but one reactant or one product (using nuclear notation).
7. Explain the kinetics of radioactive decay using equations.
8. Apply the kinetics of radioactive decay to calculate half-lives, amount of sample left, original amount of sample, or elapsed time, given three of them.
9. Discuss the application of radioactive decay dealing with archaeological (carbon) dating.
10. Give reactions that tend to bring unstable nuclei into the band of stability.
11. Explain how "magic numbers" predict the stability of super-heavy elements.
12. Explain the principle of operation of: the Geiger Muller counter, the cyclotron, a nuclear reactor.
13. State and explain Einstein's equation relating energy to mass
(E = MC2) and relate it to nuclear fission and fusion.
14. Explain the relationship of nuclear fission to the atomic bomb and to nuclear energy.
15. Explain the relationship of nuclear fusion to the hydrogen bomb, to nuclear energy, and to the sun's energy.
16. Give examples of chemical applications of nuclear reactions.
17. Discuss all sources of energy in terms of safety, pollution and availability of fuel.
18. Discuss the basic principles in the operation of an atomic and hydrogen bomb.
Unit X: B. Qualitative Analysis:
1. Know the cations in the groups we studied in the laboratory.
2. Know the reagents and ions responsible for the precipitation of each of the five groups.
3. Explain and interpret a flow chart.
4. Given an appropriate flow chart and some qualitative test results, determine which cations (or anions) might be present or absent from an unknown solution containing one or more ions from that flow chart. Be able to also do this in the laboratory.
5. Given a flow chart, explain how one would determine whether a particular ion on it were present or absent from an unknown solution containing one or more ions from that flow chart.
6. Explain and describe the chemical properties used in qualitative analysis as they relate to the appropriate objectives in the units on: "Electrochemistry", "Acids and Bases", "Ionic Equilibria", "Chemistry of the Representative Elements I and II", and "The oxidation-reduction, weak electrolytes, precipitation, solubilization, neutralization, amphoterism, hydrolysis, indicators, and coordination chemistry.
Unit XI: Organic Chemistry:
1. Define the following terms, citing examples where appropriate:
a. Organic chemistry l. Homologous series
b. Aliphatic hydrocarbon m. Derivative
c. Aromatic hydrocarbon m. Structural isomer
d. Alkane o. Geometrical isomer
e. Alkene p. Optical isomer
f. Alkyne q. Asymmetric carbon atom
g. Alkyl group r. Functional group
h. Saturated s. Resonance hybrid
i. Unsaturated t. Resonance stabilization energy
j. Cyclic compound u. Open structure
k. Olefin v. Condensed structure
2. Describe what is meant by sp3, sp2, and sp hybridization of a carbon atom, and illustrate and name the geometry that results from these hybridizations.
3. Draw a representation of a sigma (s) and a pi (p) bond between carbon atoms.
4. Draw a valence bond representation of ethane, ethylene, and acetylene, labeling the bonds sigma or pi as appropriate.
5. Give four reasons for the fact that there are so many compounds of carbon.
6. Write the general molecular formula for any alkane or alkene.
7. Draw and name the structural isomers of the first ten alkanes, first ten alkenes, and the first ten alkynes.
8. Derive the correct names of any given compounds from the structures for the alkanes, alkenes, or alkynes, including cyclic hydrocarbons and substituent groups up to four carbons (and vice versa).
9. Given the formula for hydrocarbon, write the structures for the different isomers.
10. Distinguish among the boat and chair forms of cyclohexane.
11. State the common names of some of the simpler organic compounds.
12. Give uses for the first 10 alkanes.
13. Name and give the structure of benzene and its substituted derivatives, distinguishing between ortho, meta and para, where appropriate.
14. Sketch the resonance hybrids for benzene and explain how they relate to the real structure.
15. Identify and name the following compounds (and their functional groups) given the structure, and vice versa:
a. Alkanes e. Ethers j. Carboxylic acids
b. Alkenes f. Aldehydes k. Esters
c. Alkynes f. Ketones l. Mercaptans
d. Halides h. Amines m. Dissulfides
e. Alcohols i. Amides n. Amino acids
16. Describe the synthesis of some simple alkanes, alkenes, alkyl halides, alcohols, esters, carboxylic acids, aldehydes and ketones.
17. Define, recognize and give examples of the following organic reactions:
a. Substitution d. Oxidation
b. Addition e. Reduction
c. Elimination f. Esterification
18. Chemically and physically distinguish between:
a. Primary, secondary and tertiary alcohols
b. Aldehydes and ketones
c. Alkanes and alkenes
Course Laboratory Objectives:
1. Expand your understanding of the Course Objectives.
2. Demonstrate the ability to correctly and effectively manipulate chemicals and glassware by working alone.
3. Demonstrate the ability to correctly and effectively collect and analyze data from an experiment by working alone.
4. Demonstrate the ability to correctly and effectively use laboratory balances.
5. Demonstrate the ability to correctly and accurately do quantitative analysis such as titrations, pipetting and preparation of solutions by working alone.
6. Demonstrate the ability to correctly and effectively collect and treat data on the computer.
7. Demonstrate the ability to correctly and effectively use instruments like Spectrophotometers, voltmeters and pH meters.
8. Utilize critical thinking and quantitative reasoning skills in observing, organizing and analyzing data, synthesizing information, interpreting results, and communicating the results of the analyses and laboratory investigations.
9. Perform chemical experimentation in a safe and scientific manner, using proper scientific and laboratory safety procedures.
10. Students must show work, thought process and/or justification for answers when
necessary on laboratory reports. They should also be clear and legible.