Download Course Syllabus Course Number and Name: CHM:160 Chemistry I

Course Syllabus
Course Number and Name: CHM:160 Chemistry I
Semester: 2
Classroom & Class Time: Room 604 WDHS
Instructor:
Phone:
Email Address:
Office/ Hours:
Gerald J. Ross
563-876-3442 ext 334
[email protected]
Room 604 8:00 AM – 1:45 PM
Course Description:
This course deals with the structure of the atom, elements and the periodic table, chemical formulas and
nomenclature, chemical equations, stoichiometry, bonding, thermochemistry, gases, liquids and solids, and solution
chemistry.
Course Objectives:
Unit One
At the end of this unit, the student will be able to:
1. Write units and abbreviations used to measure length, volume, and mass.
2. Write numerical values of metric prefixes.
3. Write numerical conversion factors for units that describe the same quantity.
4. Write numerical relationships between two metric units.
5. Use conversion factors to change from one unit of measure to another.
6. Calculate the density and/or specific gravity of a substance.
7. Compute relationships between Celsius, Fahrenheit, and Kelvin scales.
Unit Two
At the end of this unit, the student will be able to:
1. Write and explain the three main postulates of modern atomic theory.
2. Write and explain three basic laws of chemistry.
3. Describe properties of the three subatomic particles for a given element.
4. Identify atomic number, mass, and number of subatomic particles for a given element.
5. Identify molecules and ions.
6. Explain the concepts of atomic mass, isotopic abundance, and significance of Avogadro's number.
7. Define and explain the concepts of mole, molar mass, formula mass, and molarity.
Unit Three
At the end of this unit, the student will be able to:
1. Define and give examples of various forms of chemical formulas.
2. Calculate percent composition from the formula and the formula from the percent composition.
3. Calculate molecular formulas for ionic compounds.
4. Recognize and predict formulas for ionic compounds.
5. Assign correct names to simple ionic and binary (nonmetal) molecular inorganic compounds.
6. Write and balance inorganic chemical equations.
7. Calculate mass relations of reactants and products in a chemical equation.
8. Determine the limiting reactant and theoretical yield when mass quantities of reactants are given.
Unit Four
At the end of this unit, the student will be able to:
1. State useful units of volume, amount, temperature and pressure as applied to the behavior of gases.
2. State and use the Ideal Gas Law for the determination of R, application to initial and final state problems, the
calculation of P, V, n, or T, and the determination of molar mass or density of a gas.
3. Use the Ideal Gas Law or the Law of Combining Volumes to determine the volumes of gases involved in
reactions.
4. Calculate the pressures of a gas mixture or its component gases by the application of Dalton's Law.
5. Explain and use the kinetic theory of gases to determine the rates of gaseous effusion and diffusion and
molecular velocities and energies.
6. Describe real gas departures from the ideal in terms of particle volume and attractive forces and use Vander
Waals equation as a better approximation to real gas behavior and high pressures.
Unit Five
At the end of this unit, the student will be able to:
1. 10.5.1 State and describe the basic postulates of the Quantum Theory and its relationship to observed atomic
spectra.
2. Perform calculations relating wave length, frequency and amplitude.
3. Describe the Bohr model of the hydrogen atom.
4. Describe the Quantum Mechanical atom. Discuss the wave nature of the electron, DeBroglie waves, wave
functions, and electron clouds.
5. Use Quantum number, energy level, and orbital concepts to describe the electron configuration of the elements
and relate these to the Pauli Exclusion Principle.
6. Assign quantum numbers to each electron in an atom and draw orbital diagrams utilizing Hund's Rule.
7. Derive orbital diagrams and quantum numbers for each electron in a monatomic ion.
Unit Six
At the end of this unit, the student will be able to:
1. Trace the development of the Periodic Table from its inception to modern times, stating the principle difference
between the Mendeleev arrangement and the modern arrangement.
2. Show the direct relationship between the electron arrangement within an atom and its place in the Periodic
Table and, therefore, its general chemical behavior.
3. Describe the trends of atomic radius, ionic radius, ionization energy and electron affinity of atoms with respect
to their position in the Periodic Table.
4. Define the properties that distinguish metals from nonmetals from metalloids with regard to their position in the
Periodic Table, physical properties and metallic bonding.
5. Describe the physical and chemical properties of the main group metals and their preparation as pure metals.
6. State some (commercially) important compounds of the main group metals, their preparation and uses.
Unit Seven
At the end of this unit, the student will be able to:
1. Define and explain the basic concepts of thermochemistry: state properties, sign convention for heat flow, heat
capacity and specific heat, and heat flow in reactions.
2. Describe the techniques and principles of calorimetry, including the similarities and differences between the
coffee-cup and bomb calorimeter.
3. Define enthalpy and relate this concept to the Law of Conservation of Energy.
4. Apply the three basic laws of thermochemistry to set up thermochemistry equations, taking into consideration
where appropriate, the heat of fusion and heat of vaporization of the material.
5. Define the concept of heat of formation for chemical reactions and for ions in solution; utilize the heat of
formation to calculate the standard enthalpy change of a reaction.
6. Explain the First Law of Thermochemistry and its relationship to enthalpy and energy in chemical reactions.
7. Compare energy sources in the U.S. with respect to advantages and disadvantages.
Unit Eight
At the end of this unit, the student will be able to:
1. Give a general description of the covalent bond.
2. Discuss Lewis Structures and the Octet Rule, including writing Lewis Structures, and resonance forms; and
exceptions to the Octet Rule, including expanded octets.
3. Explain covalent bond properties in terms of polar and nonpolar bonds, bond distances, bond energies, and
estimated change in enthalpy from bond energies.
Unit Nine
At the end of this unit, the student will be able to:
1. Predict the major features of molecular geometry according to VSEPR theory, including molecules involving
atoms with unshared electron pairs and multiple bonds.
2. Use the molecular geometry and the electronegativity of atoms forming a molecule to predict the polarity of the
molecules.
3. Describe hybridization of atomic orbitals and use it to describe single, double, and triple covalent bonds.
4. Explain the nature of sigma and pi bonds.
5. Apply molecular orbital theory to predict properties of diatomic molecules of second period elements and to
explain "resonance."
Unit Ten
At the end of this unit, the student will be able to:
1. Discuss aspects of liquid-vapor equilibrium, including vapor pressure, temperature vs. vapor pressure, boiling
point, critical temperature and pressure.
2. Interpret phase diagrams and explain the phenomena of sublimation and fusion.
3. State the general class properties of molecular substances and explain the role that dispersive forces, dipole
forces and hydrogen bonds play in their physical properties.
4.
State the general class properties of network covalent substances. Apply this to the specific examples of
elemental carbon and silicon dioxide.
5. State the general class properties of ionic substances and use Coulomb's Law to predict the relative strength of
ionic bonds as a function of interionic distance.
6. State the general class properties of metals.
7. Describe the common crystal structures and use this information, with xray diffraction data for crystallized
elements to calculate the radius of the atom.
Unit Eleven
At the end of this unit, the student will be able to:
1. Correctly define and use the terminology specific to the chemistry of solutions.
2. Calculate the concentration of solutions in terms of molarity, mole fraction, and molality and be able to convert
from one form to another as required.
3. Recognize the factors which control solubility of a solute in a solvent and be able to calculate temperature and
pressure effects given sufficient data.
4. Recognize and define the colligative properties of solutions and be able to calculate the effect of solute,
electrolyte or non-electrolyte, on vapor pressure, osmotic pressure, boiling point elevation, and freezing point
depression (or molar mass of solute, given this information) from sufficient data.
Unit Twelve
At the end of this unit, the student will be able to:
1. State the principles of solubility of ionic solids, predict the formation of precipitates when ionic solutions are
mixed, and write chemical equations for precipitation reactions.
2. Define "acid" and "base" and describe, using chemical equations, the interactions: strong acid--strong base
strong acid--weak base weak acid--strong base.
3. State the principles of oxidation-reduction reactions and the role of oxidation numbers in redox reactions, and
be able to balance redox equations.
4. Perform correct stoichiometric calculations for solutions which react when mixed.
5. Calculate the equivalence point or solution concentration by the techniques of volumetric analysis.
Unit Thirteen
At the end of this unit, the student will be able to:
1. Discuss the N2O4-NO2 gaseous equilibrium system and the calculation of the equilibrium constant for the
system.
2. Explain the Coefficient Rule, the Rule of Multiple Equilibria, and the principles of heterogeneous equilibria as
they are reflected in the calculation of Kc for any gaseous equilibrium, and be able to calculate a Kc for a given
equilibrium.
3. Describe how Kc may be used to determine the direction of a given gas phase reaction and the equilibrium
concentrations of reactants and products.
4. Explain and calculate effect of a change in partial pressure, volume, or temperature, or removal of chemical
specie upon a gaseous equilibrium (Le Chatelier’s Principle).
Unit Fourteen
At the end of this unit, the student will be able to:
1. Contrast and compare theoretical and operational definitions of the Arrhenius, Bronstead-Lowry and Lewis
theories of acids and bases.
2. Explain the ionization constant for water (Kw) and the relationship between H30+ and OH- concentration in
aqueous solutions.
3. Understand the difference between strong and weak acids and bases and use Ka and Kb in calculations
involving strong and weak acids and bases.
4. Identify buffers and the conditions necessary for the existence of a buffer and understand how buffers resist
change in pH at any point in a titration.
5. Use the solubility product constants to determine to what extent a sparingly soluble salt will dissolve in
water.
Required Reading Materials:
Textbook chapters
Assignments:
Questions and problems will be assigned for each chapter.
Methods of Assessment:
Tests
Assignments
80%
20%
Grading Scale:
Grades will be assigned for work completed using the following grading scale:
A
93%
B80%
A90%
C+
77%
B+
87%
C
73%
B
83%
C70%
D
D
DF
67%
63%
60%
<60%
Methods of delivery:
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Readings, lectures, discussions, small group activities, audiovisual materials, power point presentations and
computer software may be used to present the subject matter.
Written assignments, role plays, simulations and /or other projects may be necessary to complete course
requirements.
Student-instructor conferences may be scheduled as needed.
Tests and other evaluation devices will be used to monitor student progress.
Students who are having difficulty with the subject matter are encouraged to contact the teacher for individual
help.
Disclaimer:
(Example: Faculty choosing to list activities to course calendar, such as chapters covered, field trips, test, quizzes,
etc., need a disclaimer on their course syllabus. This disclaimer should state that because of illness/death in family,
weather conditions, power outages, etc., the course calendar is a guide for activities and subject to change.)
Accommodation Policy: (High Schools Contracted Classes should use HS Accommodation Policy)
The Americans with Disabilities Act (ADA) provides protection from illegal discrimination for qualified students
with disabilities. Northeast Iowa Community College is committed to the equal provision of education for all
students. Any student who needs instructional accommodation because of physical or learning disability is
encouraged to contact the Coordinator of Developmental Education: Connie Swift, Peosta Campus, at 563-5565110 or 1-800-728-7367, ext. 280, or Janet Reis, Calmar Campus, at 563-562-3263 or 1-800-728-2256, ext. 258.