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:   Whitten, Kenneth, Davis, Raymond, Peck, and Stanley,

Thompson-Brooks/Cole/Cengage Learning with On-line access, 10th Ed., 2014


General Chemistry I: Course Manual

Alfare, Carlo MCCC, 12th Edition (2018 - 2019)


General Chemistry I: Laboratory Manual

Alfare, Carlo MCCC, 14th Edition (2018 - 2019)


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  AND  MAT 146 or higher


Carlo Alfare                                           Michael Dorneman

Professor of Chemistry                                     Professor of Chemistry

Course Coordinator

MS 123 609-570-3381                                MS 120      609-570-3369




E-mails accepted only from MCCC assigned student E-mail address





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




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 Max. 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%)

+ 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.







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. All but the last column will be found in Whitten, Davis, Peck, and Stanley


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.



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 ), 100, 101(ans. 2.350 )

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 41, 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 5-61-11, 13-21, 24, 26-28, 30, 32,          Unit 6: 14 - 17
36, 38, 41, 42

10a 6-3, 6-427-32, 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 81-4, 7-10, 13-16, 20-22          Unit 7: 4 - 5

10d 824, 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, 6-92-12, 14, 18, 20         Unit 8: 1

11b 5-9 67 - 71

11c 101-4, 6-13, 17-20, 23-27         Unit 8: 2

11d 1031, 32, 34-36, 40, 41, 43-47,         Unit 8: 3 - 6
50, 57, 61, 63

12a 5-746, 48, 49         Unit 8: 7 9


12b 6-549, 50, 52

12c 111-4, 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, 11-552a,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 1446, 48, 50, 60, 68, 81                Unit 9: 8 - 12




Textbook OWLv2 On-Line Mastery and Homework Assignments:

This should be your major source of extra help in this course


Tutoring sites on my home page:


Science Learning Center (MS 211) for help with labs and homework




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.

         For those who study chemistry, the benefits to be expected are an increased capacity for

         enjoyment , a livelier interest in the world in which we live, and a more intelligent attitude

         toward the great questions of the day.

        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:


You should be able to:


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

Unit II: Stoichiometry:


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. (Maxwell-Boltzman Distribution).


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 Dalton's Law of partial pressures, and use it in calculations, including when water vapor is one of the gases.


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. Normal boiling point

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, OH-, and that it will turn a blue litmus to pink.


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




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




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 solutions by working alone.


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.


 9.  Students must show work, thought process and/or justification  
      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:


                     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.