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Date Adopted: 1966 Date Reviewed: 1987
Date Revised: 1988, 1998, 2001, 2008
Representing Alabama’s Public Two-Year College System
Alabama
Department of
Postsecondary Education
Jefferson State Community College
CHM 111
College Chemistry I
I. CHM 111, College Chemistry I, 4 Semester Hours
Core Area III, ASCI TSCI (Lec 3 hrs, Lab 2 hrs) (***State guide has 3HR Labs)
II. Course Description
This is the first course in a two-semester sequence designed for the science or engineering major who is expected to have a strong background in mathematics. Topics in this course include measurement, nomenclature, stoichiometry, atomic structure, equations and reactions, basic concepts of thermochemistry, chemical and physical properties, bonding, molecular structure, gas laws, kinetic-molecular theory, liquids and solids, solutions, and colloids. Lab is required.
III. Prerequisite: MTH 112 (Precalculus Algebra) or equivalent math placement score.
IV. Textbook
Chemistry and Chemical Reactivity, Kotz and Treichel, 7th^ Ed. Modular Lab Notebook, Chemical Education Resources
V. Course Competencies
In the classroom the student will:
A. Understand the basic mathematical principles involved in chemical calculations and will have a thorough understanding of the metric system of measurement. B. Understand the classification of matter into various groups based upon similarity of chemical and physical properties. C. Comprehend and interpret chemical symbols, formulas, names, chemical equations, and calculations that apply, stressing stoichiometry. D. Understand the electronic arrangement of the atom in terms of the quantum theory, and will be able to use the periodic table to link electronic configuration to the properties of the element. E. Understand chemical periodicity. F. Understand the nature of covalent bonding, ionic bonding, and the main concepts of the three covalent bonding theories (VB, VSEPR, MO). G. Comprehend and apply the principles of gas behavior in ideal as well as real gas systems. H. Understand the characteristics of the solid and liquid states of matter and phase diagrams. I. Understand terms used in solution chemistry. J. Understand the nature of aqueous solution systems and apply the principles of solubility, colligative properties and concentration in problem solving. K. Understand and apply chemical principles of acids and bases.
In the laboratory the student will:
A. Develop an understanding of basic laboratory techniques and procedures. B. Understand basic laboratory safety and will follow all laboratory rules during experimental work. C. Acquire understanding of the physical and chemical properties of commonly used elements, compounds and mixtures. D. Be able to make precise measurements and evaluate experimental data through selected qualitative laboratory experiments. E. Be able to make careful observations, report and interpret experimental data through selected quantitative laboratory experiments. F. Be able to perform simple calculations from experimental data through selected quantitative laboratory experiments.
VI. Course Outline of Topics
Lecture Topics Stated in Performance Terms
The student will be required to demonstrate that he has attained each general course competency by performing the objectives listed under each competency.
and precision, heat and temperature.
C. Comprehend and interpret chemical symbols, formulae, names, chemical equations, and calculations that apply, stressing stoichiometry.
- Memorize the names and symbols of the elements given in a reference table.
- Interpret a chemical formula in terms of the type and number of atoms present.
- Given an ionic or molecular formula, determine the formula weight.
- Relate the numbers of particles (atoms, molecules, or ions) and the mass in grams of a sample of matter.
- Given the formula of a substance, relate the number of moles and the mass in grams of the sample.
- Write and interpret the formulas for some common substances.
- Given the formula of a compound, calculate the percentages by mass of the elements.
- Determine the empirical formula of compound, given the mass percentages of the elements or the analytical data from which these can be calculated.
- Determine the molecular formula of a compound, given the simplest formula and at least an approximated molecular mass.
- Use the Periodic Table to obtain the charges of ions formed by the main- group elements.
- Write the formula for an ionic compound given either the formulas of the ions or the name of the compound.
- Given the formula for a compound, give its name.
- Describe some experimental methods of determining percent composition.
- Write and balance chemical equations and interpret the various symbols used in chemical equations to represent the condition of the reaction system.
- Relate the number of moles of any two substances taking part in a reaction.
- Relate the masses of any two substances taking part in a reaction.
- Given or having calculated two of the three quantities, concentration, number of moles of solute, volume of solution, determine the other quantity.
- Given the balanced equation for a reaction involving species in solution, relate the volumes or concentrations of two reactant species.
- Describe water solution reactions involving precipitation, acid-base, and oxidation-reduction.
- Given the number of moles or masses of all reactants, determine which is the limiting reagent and calculate the theoretical yield of any product.
- Calculate the percent yield, given the actual or theoretical yields.
D. Understand the electronic arrangement of the atom in terms of quantum theory, and will be able to use the periodic table to link electronic configuration to the properties of the elements.
- State and describe the postulates of quantum theory and compare quantum and classical theories.
- Relate the wavelength and frequency of a spectral line to the energy of photons and to the change in energy of an atom.
- Discuss the contributions to the atomic theory made by Dalton, Thomson, Rutherford, Bohr, Chadwick, deBroglie and Schrodinger.
- State and apply the Aufbau principle.
- Determine the number of electrons that may be accommodated by any given principal energy level or sublevel.
- Given the atomic number of an element, write the electron configuration.
- Given the electron configuration, state and apply Hund's rule and draw orbital diagram of the atom.
- Describe the four quantum numbers, and the rules for assigning them.
- Apply the rules and assign them to each of the various electrons in an atom.
- State and apply Pauli's Exclusion principle.
- Relate electronic configurations to the periodic table and to periodicity.
- Using the periodic table, predict the relative values of ionization energy,
- Predict molecular polarity from Lewis structures.
- Predict molecular geometry from orbital hybridization.
- Predict the kind and number of sigma and pi bonds in a molecular species.
- Write molecular orbital diagrams for simple diatomic species.
G. Comprehend and apply the principles of gas behavior in ideal as well as real gas
systems.
- Describe and apply Boyle's, Charles', Gay-Lussac's and Avogadro's law.
- Apply the ideal gas law to predict the effect of a change in conditions upon a variable such as volume.
- Apply gas laws to calculate the density of a gas at a given temperature and pressure.
- Use the ideal gas law to calculate the molecular mass of a gas, knowing the mass of a given volume or the density at a known pressure and temperature.
- Relate volumes of gases involved in chemical reactions from information obtained from chemical equations.
- Apply Dalton's law of partial pressures of gases in mixtures.
- List the assumptions of kinetic-molecular theory and describe gas behavior in terms of the theory.
- Describe and apply Graham's law to relate molecular masses, rates of effu- sion, and times of effusion of gases.
- Describe how real gases deviate from the assumptions of the ideal gas law, and indicate the conditions where these deviations are most significant for most gases.
H. Understand the characteristics of the solid and liquid states of matter and phase
diagrams.
- Determine vapor pressure of liquids in given temperatures and pressures.
- Predict and describe the various intermolecular forces present in molecular substances.
- Classify a given substance as ionic, nonpolar, polar, macromolecular, or metallic.
- List the general physical properties associated with each of the five categories of substances listed above.
- Write equations for the thermal decomposition of carbonates, hydroxides, and hydrates.
- Determine enthalpy change associated with a given phase change.
- Interpret phase diagrams and apply them to predict phase changes associated with changes in temperature and pressure.
I. Understand terms used in solution chemistry.
- Describe and distinguish among solvent, solute, solution, dispersion and colloid.
- Give examples of various kinds of solutions involving different combinations of solids, liquids, and gases as dispersing medium and dispersed substances
- Describe the relative effects on solubility of the following kinds of interactions: solute-solute; solvent-solvent; solvent-solute.
- Describe and illustrate the mechanism of dissolution of ionic solids and polar covalent substances in water.
- State the effects of exo- or endothermicity and of an increase in disorder on the spontaneity of the dissolution process.
- Distinguish among unsaturated, saturated, and supersaturated solutions.
- Distinguish between exothermic and endothermic dissolution process.
J. Understand the nature of aqueous solution system and apply the principles of
solubility, colligative properties and concentration in problem solving.
- Given the formula for a substance, predict whether it will be an electrolyte or a non-elecrolyte in aqueous solution.
- Predict the relative solubilities of different solutes in water.
- Describe the ionization of a poly-protic acid in aqueous solutions.
- Predict the relative strengths of acids and bases from a given set of molecular structures.
- Use titration data for an acid-base reaction to determine: the concentration of an acid or a base in aqueous solutions, and molecular mass of an acid or a base.
- Select an acid-base indicator appropriate for a given acid-base titration.
- Given the following [H+], [OH-], pH, or pOH calculate any of the others.
Laboratory Topics
A. The student will develop an understanding of basic laboratory techniques and
procedures.
- Properly operate the Bunsen burner.
- Operate a single pan balance.
B. The student will understand basic laboratory safety and will follow all laboratory
rules during experimental work.
- Follow basic laboratory safety rules as set forth by the department and the instructor.
- Locate laboratory safety and first aid equipment.
C. The student will acquire understanding of the physical and chemical properties of
commonly used elements, compounds, and mixtures.
- Distinguish between physical and chemical properties of substances.
- Determine physical properties such as density, volume, mass, etc.
- Make specific and accurate observations of materials and reactions as to color, odor, energy changes, gas evolution, precipitation, etc.
- Identify evidence of chemical changes.
D. The student will be able to make precise measurements and evaluate experimental
data through selected quantitative laboratory experiments.
- Use a meter stick to measure length of any object in cm, mm, and meters.
- Read centigrade thermometers and convert to Kelvin and Fahrenheit.
- Read the volume contained in any graduated cylinder to within 0.5 ml.
- Use a laboratory balance to determine the mass of any object to within 0.01 g.
E. The student will be able to make careful observations, report and interpret
experimental results through selected qualitative laboratory experiments.
- Interpret evidence of solubility and miscibility.
- Collect a precipitate by filtration.
- Predict the formation of precipitates based on principles of solubility.
- Make accurate observations of state, color, and odor of elements, compounds, and mixtures.
- Distinguish between elements, compounds, and mixtures.
- Record evidence of chemical change occurring in a reaction.
- Determine the relative activities of two metals in a single replacement reaction.
- Arrange a group of metals from most active to least based upon observations of a series of single replacement reactions.
F. The student will be able to perform simple calculations from experimental
data through selected quantitative laboratory experiments.
- Calculate the densities of selected solids and water.
- Calculate the percent of water in selected unknown hydrated salts.
- Calculate the empirical formula for strontium iodide salt or other compound.
- Determine the concentration of a basic solution when titrated with acid of a known concentration.
- Determine the molar mass of an impurity due to freezing point depression.
- Apply the rules of significant figures, rounding off, exponential notation, and instrument precision to the numerical results of measurements and calculations.
- Measure volumes of liquids correctly with burettes, graduated cylinders and pipettes.
- Apply titration techniques to the standardization of solutions.
- Use pH paper, universal indicator, and special indicators in determinations of pH of solutions.
- Use a spectrophotometer to measure percent transmission or absorbance of a
