Upon completion of the course, the student should be able to:
In the lecture upon completion of the course, the student should be
1. Solve problems involving the concepts listed under Course Content
2. Solve problems using SI units and dimensional analysis
3. Write concise explanations describing various chemical phenomena
4. Write and interpret balanced chemical equations.
5. Describe and identify various types of colloids.
6. Write balanced chemical equations for oxidation-reduction reactions.
7. Write balanced chemical equations for precipitation reactions.
8. Describe different models of atomic structure.
9. Use standard chemical notation and nomenclature.
10. Predict the outcomes of combination, decomposition, single-
displacement, metathesis and combustion reactions.
11. Calculate the mass percentages of the elements from the formula of
12. Determine the empirical formula of a compound from elemental
13. Derive the molecular formula of a compound from the empirical formula.
14. Predict the amounts of reactants and products involved in a chemical
15. Solve limiting-reagent problems.
16. Calculate theoretical and actual percentage yields.
17. Perform calculations involving mass percentage, molarity, normality,
molality and mole fraction units.
18. Solve solution stoichiometry problems.
19. Derive predictions of total pressure, partial pressures, volume,
temperature, moles or mass utilizing the ideal gas laws.
20. Relate properties of gases to the kinetic-molecular theory.
21. Predict deviations from ideal behavior in real gases.
22. Calculate molecular weights of gases from Graham's law.
23. Apply the First Law of thermodynamics.
24. Describe colligative properties of solutions.
25. Write and interpret thermochemical equations.
26. Use Hess's Law to calculate enthalpies of reaction from standard heats
27. Calculate wavelength, frequency, speed and energy of electromagnetic
28. Describe the energy level diagram and spectral series for atomic
29. Calculate wavelength and momentum using the deBroglie relationship.
30. Use the Heisenberg principle to predict uncertainty in position or
31. Describe the significance of the four quantum numbers.
32. Use the Aufbau Principle to derive the ground-state electronic
configurations of the elements.
33. Apply Hund's rule and predict the number of unpaired electrons in an
34. Distinguish between diamagnetic and paramagnetic behavior.
35. Describe the relationship between electronic configuration and atomic
radius, ionization energy, electron affinity and electronegativity.
36. Define ionic bonding, and apply the Born-Haber cycle to predict the
stability of ionic crystalline solids.
37. Define covalent bonding utilizing Lewis dot structures.
38. Predict the existence of polar bonds and dipole moments in molecules.
39. Describe the delocalized bonding and resonance structures.
40. Explain the factors that affect solubility.
41. Calculate formal charges, bond orders, oxidation numbers and
42. Calculate enthalpies of reaction using bond dissociation energies.
43. Apply the valence-shell electron-pair repulsion model to predict
In the laboratory upon completion of the course, the student should be
1. Observe all of the fundamental safety procedures.
2. Properly dispose of waste chemicals.
3. Manipulate standard laboratory apparatus.
4. Perform gravimetric and titrimetric analyses.
5. Collect and analyze scientific data using graphical and statistical
6. Summarize lab results in both formal and informal report formats.
7. Use a Macintosh personal computer (or equivalent) to perform word
processing, spreadsheet computations, graphing and statistical
calculations for lab reports.
1. Keys to the Study of Chemistry
a. Fundamental definitions
b. Chemical arts and origins of modern chemistry
c. The scientific approach
d. Chemical problem solving
e. Measurement in Scientific Study
f. Significant Figures
2. The Components of Matter
a. Elements, Compounds and Mixtures
b. The atomic view of matter
c. The nuclear atom model
d. The atomic theory today
e. Elements and the periodic chart
f. Introduction to bonding
g. Compounds - formulas, names, and masses
a. The mole
b. Determining the formula of an unknown compound
c. Writing and balancing chemical equations
d. Calculating the amounts of reactant and product
e. Fundamentals of solution stoichiometry
4. The Major Classes of Chemical Reactions
a. Types of chemical reactions
b. The role of water as a solvent
c. Some important aqueous ionic reactions
d. Redox reactions
e. Reversible reactions, equilibrium
5. Gases and Kinetic-Molecular Theory
a. The physical states of matter
b. Measuring the pressure of a gas
c. The gas laws and their experimental foundations
d. Further applications of the ideal gas law
e. The ideal gas law and reaction stoichiometry
f. The kinetic-molecular theory
g. Real gases: deviations from ideal behavior
a. Forms of Energy and their inter conversion
d. Stoichiometry of thermochemical equations
e. Hess's law of heat summation
f. Standard heats of reactions
7. Quantum Theory and Atomic Structure
a. The nature of light
b. Atomic spectra and the Bohr model of the atom
c. The wave-particle duality of matter and energy
d. The quantum-mechanical model of the atom
8. Electron Configuration and Chemical Periodicity
a. Characteristics of many-electron atoms
b. The quantum-mechanical atom and the periodic table
c. Trends in some key periodic atomic properties
d. The connection between atomic structure and chemical reactivity
9. Models of Chemical Bonding
a. Atomic properties and chemical bonds
b. The ionic bonding model
c. The covalent bonding model
d. Between the extremes: electronegativity and bond polarity
e. Depicting molecules and ions with Lewis structures
f. Using Lewis structures and bond energies to calculate heats of
g. An introduction to metallic bonding
10. Molecular Shape and Theories of Covalent Bonding
a. VSEPR theory
b. Molecular shape and molecular polarity
c. Valence bond theory and orbital hybridization
d. Molecular orbital theory and electron delocalization
11. Intermolecular Forces
a. Physical states and phase changes
b. Types of intermolecular forces
c. Properties of the liquid state
d. Properties of the solid state
e. Quantitative aspects of changes in state
f. The uniqueness of water
12. The Properties of Mixtures
a. Types of solutions: intermolecular forces and the prediction of
b. Energy changes in the solution process
c. Solubility as an equilibrium process
d. Quantitative ways of expressing concentration
e. Colligative properties of solutions
f. The structure and properties of colloids
1. Laboratory safety, techniques and maintaining data notebooks
2. Writing formal and informal laboratory reports
3. Word processing, spreadsheets, graphing and curve-fitting software
4. Computer interfacing experiments
5. Calibration of volumetric glassware
6. Basic laboratory skills: weighing, pipetting, filtration, melting
7. Graphical determination of density
8. Empirical formula of a compound
9. Gravimetric analysis of iron, sulfate or chloride
10. Observing & classifying types of chemical reactions
12. Molar mass of a volatile compound by vapor density
13. Boyle•s and Charles• laws
14. Atomic spectroscopy
15. Acid-base titrations
16. Limiting reactants
17. Synthesis and analysis of an inorganic compound
18. Molecular geometry: VSEPR theory
19. Molar mass by freezing point depression
CHEMISTRY: PRINCIPLES & PRACTICE by Daniel Reger, Scott Goode and Edward
Mercer; Saunders College Publishing, 1997.
PRINCIPLES OF MODERN CHEMISTRY by David Oxtoby, H. P. Gillis & Norman
Nachtrieb, Saunders College Publishing, 1999.
CHEMICAL PRINCIPLES by Steven Zumdahl, Houghlin Mifflin Publishing, 1998.
CHEMISTRY IN THE LABORATORY by J. A. Beran, John Wiley Publishing, 1995.
EXPERIMENTS IN GENERAL CHEMISTRY by Frank Milio, Nordulf Debye & Clyde
Metz, Saunders College Publishing, 1991.
QUANTITATIVE CHEMICAL ANALYSIS by Daniel Harris, W. H. Freeman Publishing,
FUNDAMENTALS OF ANALYTICAL CHEMISTRY by Douglas Skoog, Donald West & James
Holler, Saunders College Publishing, 1996.
SPECIAL STUDENT MATERIALS:
Laboratory data notebook