General College Chemistry I

I.     Course Prefix/Number: CHM 121

       Course Name: General College Chemistry I

       Credits: 4 (3 lecture; 3 lab)

II.    Prerequisite

MAT 110 or MAT 120, with minimum grade of C, and one year of high school chemistry or CHM 101 or CHM 105, with minimum grade of C, or consent of instructor. MAT 140, or MAT 149, is recommended.

III.   Course (Catalog) Description

Course is first of two semester sequence (CHM 121 and CHM122). Content includes the periodic table of elements, atomic structure, basic concepts of quantum theory, stoichiometry of compounds and reactions, thermochemistry, molecular structure, bonding, intermolecular interactions, the gaseous state, and solutions. Weekly hands-on lab activities. Intended for students enrolled in science and/or pre-professional curricula.

IV.   Learning Objectives

  1. General Education.
    1. Identify, define, analyze, interpret, and evaluate: ideas, concepts, information, problems, solutions, and consequences. This includes the ability to compute and comprehend quantitative information and to engage in the scientific process.
  2. Lecture.
    1. Explain observations and laws using models, and apply the scientific approach to knowledge.
    2. Classify matter according to its state, composition, and properties.
    3. Differentiate between physical and chemical changes and between physical and chemical properties.
    4. Evaluate the reliability of measured and calculated quantities and report those values using standard SI units, rules of significant figures and rules of scientific notation.
    5. Utilize units (dimensional analysis) as a guide to problem solving.
    6. Predict properties of atoms, elements and compounds using the Periodic Table of Elements and modern atomic theory.
    7. Represent compounds using their names, chemical formulas, and models.
    8. Apply the concepts of mole and stoichiometry.
    9. Represent solutes and reactions in aqueous solutions, and write molecular, complete ionic and net ionic equations.
    10. Quantify solubility, solution concentration and titration equivalence points.
    11. Identify redox reactions.
    12. Correlate the tenets of kinetic molecular theory with ideal gas behavior and rationalize nonideal gas behavior.
    13. Summarize the nature of energy and enthalpy, and quantify energy and enthalpy changes.
    14. Discuss wave-particle duality of light and electrons.
    15. Apply the Bohr and quantum mechanical models of the atom.
    16. Apply Lewis theory, the VSEPR model, and valence bond theory.
    17. Correlate the physical properties of substances with intermolecular forces.
    18. Interpret phase diagrams and heating curves.
    19. Predict relative solubility of molecular substances in liquids.
    20. Correlate the properties of crystalline solids with the basic units that comprise them.
    21. Quantify colligative properties of solutions.
  3. Laboratory.
    1. Minimize risk to self and others by adhering to documented and verbalized laboratory safety policies.
    2. Safely demonstrate basic laboratory skills including filtration, titration, observation and testing of properties of various unknowns, as well as use of a Bunsen burner, high-precision balances, and other basic equipment and glassware.
    3. Perform routine laboratory measurements including mass, volume, and temperature, and record them with proper precision and units; distinguish between measured quantities (raw data) and calculated quantities.

V.    Academic Integrity and Student Conduct

Students and employees at Oakton Community College are required to demonstrate academic integrity and follow Oakton's Code of Academic Conduct. This code prohibits:

• cheating,
• plagiarism (turning in work not written by you, or lacking proper citation),
• falsification and fabrication (lying or distorting the truth),
• helping others to cheat,
• unauthorized changes on official documents,
• pretending to be someone else or having someone else pretend to be you,
• making or accepting bribes, special favors, or threats, and
• any other behavior that violates academic integrity.

There are serious consequences to violations of the academic integrity policy. Oakton's policies and procedures provide students a fair hearing if a complaint is made against you. If you are found to have violated the policy, the minimum penalty is failure on the assignment and, a disciplinary record will be established and kept on file in the office of the Vice President for Student Affairs for a period of 3 years.

Please review the Code of Academic Conduct and the Code of Student Conduct, both located online at

VI.   Sequence of Topics

  1. Scientific Method. Distinction Between Description and Explanation
    1. Observation
    2. Hypothesis
    3. Scientific Law
    4. Scientific Theory/Model
  2. Basic Properties and Classification of Matter
    1. Classification of Matter—Operational Definitions (Macroscopic)
      1. States of matter: Solid, Liquid and Gas
      2. Pure substance vs. Mixture; Element vs. Compound; Homogeneous vs. Heterogeneous Mixtures
    2. Classification of Matter—Conceptual Definitions and Nanoscopic Pictures
      1. States of matter: Solid, Liquid and Gas
      2. Pure substance vs. Mixture; Element vs. Compound; Molecular vs. Monatomic Element; Molecular vs. Ionic Compound; Homogeneous vs. Heterogeneous Mixtures
    3. Physical and Chemical Properties of Matter. Physical and Chemical Changes of Matter
  3. Measurements and Calculations in Chemistry
    1. Physical Quantities and Units
      1. Standard Units in SI System: Measurement of Length, Mass, Time, and Temperature
      2. SI Prefixes and Conversions; Conversions between SI and Other Systems
      3. Derived Quantities and Units (with focus on Volume and Density)
      4. Scientific Notation
    2. Uncertainty in Measured and Calculated Quantities; Precision vs. Accuracy.
      1. Meaning of Uncertainty and Relation to Significant Figures Concept
      2. Significant Figure Rules For Estimating Uncertainty in Calculated Results
      3. Precision vs. Accuracy and Relation to Random and Systematic Errors
    3. General Problem Solving Strategy
      1. Dimensional Analysis in Conversions and General Calculations
      2. Rearrangement and Usage of Mathematical Equations
  4. Early Models of Atoms (and Historical Development of)
    1. Laws of Conservation of Mass; Definite Proportions; Multiple Proportions
    2. Explanation of Laws: Dalton’s Atomic Theory
    3. Charge and Coulomb’s Law
    4. Subatomic Particles: Protons, Neutrons and Electrons
    5. Rutherford’s Experiment and Nuclear Model of Atom
    6. Atomic Number. Chemical Symbols. Mass Number, Isotopes, and Isotopic Notation.
    7. Cations and Anions
  5. Periodic Law and Periodic Table of the Elements (PTE)
    1. Mendeleev and Development of PTE; Periods vs. Groups in PTE
    2. Metals, Nonmetals, and Metalloids
    3. Atomic Mass and Abundance of Isotopes. Atomic Mass Unit (amu)
  6. Representing Chemical Substances: Chemical Formulas
    1. Relation to Basic Units and Nanoscopic Representations
      1. Monatomic vs. Molecular Elements (Atoms vs. Molecules)
      2. Ionic vs. Molecular Compounds (Ionic Lattice vs. Molecules)
      3. Molecular (or Formula) Mass
    2. Types of Formulae: Empirical, Molecular, Structural, and 3D
  7. Macroscopic Amounts of Matter: The Mole Concept/Unit
    1. Avogadro’s Number
    2. Application to Atoms, Molecules, or Formula Units (& More); Generality
    3. Importance: Mole Ratios of “Things” Equals Ratio of the “Things”
    4. Relation to Mass: Molar Mass of Monatomic Element, Molecular Element, or Compound
  8. Quantitative Aspects of Chemical Compounds
    1. Mass Percent Composition (by Element)
    2. Determination of Empirical Formula From Experimental Data (Mass Data, % Mass Data, Combustion Analysis Data); Molecular vs. Empirical Formula Determination
  9. IUPAC Nomenclature of Ionic Compounds, Binary Molecular Compounds, and Acids
    1. Monatomic and Polyatomic Anions and Cations (Type I and II)
    2. Formula Units and the “Neutrality Principle”; Meaning of Subscripts
    3. Binary Molecular Compounds (Definitions and IUPAC Rules)
    4. Acids: How to Recognize and Name
  10. Representing Chemical Reaction: Balanced Chemical Equations and How to Use Them to Calculate Quantities Involved in Chemical Change (Stoichiometry)
    1. Meaning of a Balanced Equation; How to Balance a Chemical Equation
    2. Mole-Mole, Mole-Mass, and Mass-Mass Calculations
    3. Limiting Reactant Situations
    4. Actual Yield, Theoretical Yield, and Percent Yield
  11. Substances in Aqueous Solution: Nature of and Reactions involving
    1. Dissolution vs. Lack of Dissolution (Operational and Conceptual Definitions)
    2. Electrolyte vs. Nonelectrolyte vs. Insoluble Substances (Operational and Conceptual Definitions)
    3. Strong vs. Weak Electrolytes (Including Acids); Nanoscopic Pictures
    4. Solubility Rules for Ionic Compounds (and Identifying Ionic Strong Electrolytes)
    5. Solution Concentration: Molarity (Including Ions of Electrolyte Solutes)
    6. Exchange Reactions (Acid-Base, Precipitation)
    7. Writing Molecular, Complete Ionic, and Net Ionic Equations
    8. Oxidation-Reduction (Redox) Reactions
      1. Definition of Oxidation/Reduction, Oxidized/Reduced, to Oxidize/Reduce, and Oxidizing Agent/Reducing Agent
      2. Oxidation Numbers (States): Definition of, Assignment of, and Use in Identifying Redox Reactions
    9. Solution Stoichiometry and Titrations
  12. Gases
    1. Pressure, Temperature, and Volume; Units, STP
    2. Boyle’s, Charles’, and Avogadro’s Laws; Absolute Zero of Temperature
    3. Ideal Gas Law (and Derivation of “All” Relations/Laws Relating P, V, n, and/or T)
    4. Application of Ideal Gas Law: Density of a Gas, Molar Mass of a Gas
    5. Gas Mixtures: Partial Pressure and Dalton’s Law
    6. Nonideal Gas Behavior (Conditions that Lead to Nonideal Behavior)
    7. Kinetic Molecular Theory (KMT); Use of to Explain Gas Laws/Gas Behavior (Including Nonideal Behavior Under Conditions Where Postulates Not Valid)
      1. Includes Velocity Distribution Curves (Dependence on T and Particle Mass)
    8. Stoichiometry of Reactions in Gaseous Phase
    9. Diffusion and Effusion of Gases; Explanation using KMT.
  13. Thermochemistry
    1. Forms of Energy; Temperature vs. Heat; System vs. Surroundings; Internal Energy (E)
    2. First Law of Thermodynamics; Work, Esys = q + w); Enthalpy (H) and H’s
    3. Specific Heat Capacity; Calorimetry to Determine H¬¬sys for Processes
    4. Exo- vs. Endothermic Processes (Predictions for Phase Changes, Bond Break/Making)
    5. Thermochemical Equations—Stoichiometry with Enthalpy
    6. Hess’s Law, Standard Enthalpies of Formation, Determination of Hsys from Data
  14. Modern (Quantum Mechanical) Model of the Atom (& Development of)
    1. Waves and the Wave (Classical) Nature of Light
      1. Wavelength (λ), Frequency (ν), Speed of Light (c), c = , Amplitude
      2. Electromagnetic Spectrum
    2. Quantization and a New Model of Light (Planck, Einstein)
      1. Photoelectric Effect Observations and Einstein’s Photon Model of Light (Explains)
      2. Wave-particle Duality of Light
    3. Bohr Model of Hydrogen Atom
      1. Line Spectrum Explained With Assumption of Quantization of Electron Energies / Electron Orbits and Electronic Transitions Between Them
      2. Absorption and Emission; Ground and Excited Electronic States of Atoms
    4. Quantum Mechanical (Wave) Model of the Atom
      1. Wave Particle Duality of Matter; Standing Wave Analogy
      2. Schrodinger Equation and Orbitals—Shapes, Types, and “Nature” (Probability)
      3. Ground State Electron Configurations for Multielectron Atoms from Pauli Principle and Hund’s Rule (Aufbau Principle); Core vs. Valence Electrons
        1. Position of Element in Table Relates to Valence Electron Configuration
        2. Explains Organization of Periodic Table (s-block, p-block, etc.)
        3. Explains Periodicity of Chemical Reactivity: Same Valence Configuration = Similar Chemical Reactivity
      4. Energy Level/”Shell” Model of Electronic Structure
        1. Including Effective Nuclear Charge
  15. Periodic Properties and Explanation of them Using “Shell” Model
    1. Periodic Trends in: Atomic Radius, 1st Ionization Energy, Higher Ionization Energy Patterns, Relative Sizes of Cations and Anions, Electron Affinities, Metallic Character
    2. Effective Nuclear Charge and Average Distance Determine the Force of Attraction of an Electron To the Nucleus (Coulomb’s Law Applied)
  16. Chemical Bonding
    1. Types of Chemical Bonds: Ionic, Covalent, and Metallic
    2. Lewis Structures of Atoms and Monoatomic Ions
    3. Ionic Compounds
      1. Lattice Energy; Relation to Ion Charges and Radii (by Coulomb’s Law)
    4. Covalent Bonding
      1. Bonding (Shared) Electron Pairs
      2. Electronegativity and Relation to Polar and Nonpolar Covalent Bonds
    5. Lewis Structures of Molecular Compounds and Polyatomic Ions
      1. The Octet Rule (a Starting Point)
      2. Resonance Structures
      3. Formal Charges (Needed to Help Assess Non-equivalent Resonance Structures
      4. Exception to the Octet Rule
    6. Bond Energy and Bond Length. Estimating Reaction Energies Using Bond Energies.
    7. Metallic Bond: Electron Sea Model
  17. Geometries Around Central Atoms—VSEPR Theory
    1. Electron Groups. Electron Group Geometry
    2. Molecular Shape (Geometry); Effect of Lone Pairs.
    3. Polyatomic Species with More than One Central Atom
    4. Molecular Polarity (or Not) Depends on Molecular Geometry as Well as Bond Polarity
  18. Valence Bond Theory
    1. Atomic Orbital (AO) Overlap.
    2. Using Hybridization of AO’s to Explain Many Observed Molecular Geometries
      1. sp, sp2, and sp3 Hybridization
      2. σ and π bonds; Rotation Restricted Around a π bond
      3. sp3d, and sp3d2 Hybridization
  19. Intermolecular Forces (IMFs) and Properties of Liquids and Solids
    1. Types of Intermolecular Forces
      1. Dispersion (London) Forces
      2. Dipole-Dipole Forces
      3. Hydrogen Bonding
      4. Ion-Dipole Forces
    2. Properties Dependent on IMFs
      1. Melting Point, Boiling Point, Heat of Fusion, Heat of Vaporization, Vapor Pressure, Surface Tension, Viscosity, Capillary Action
    3. Vapor Pressure, Boiling, and Boiling Point
      1. Temperature Dependence of Vapor Pressure
      2. Boiling and Boiling Point—Relation to External Pressure
    4. Sublimation and Fusion
    5. Heating Curve
    6. Phase Diagram: Vaporization, Fusion, and Sublimation Curve; Triple Point; Normal Melting and Boiling Point; Supercritical Temperature and Pressure
    7. Unusual Properties of Water
    8. Types of Crystalline Solids: Molecular, Ionic, Metallic, and Covalent Network
  20. Solutions
    1. Solubility: Solute-Solvent Intermolecular Interactions; Polarity of Solvent and Solute
    2. Energetics of Solution Formation; Heat of Hydration
    3. Factors Affecting Solubility
      1. Effect of Temperature on Solubility of Solids
      2. Solubility of Gases: Effect of Temperature and Pressure
    4. Solution Concentration
      1. Molarity
      2. Molality
      3. Mass Percent; ppm and ppb
      4. Mole Fraction and Mole Percent
    5. Colligative Properties of Solutions of Nonvolatile Nonelectrolytes
      1. Vapor Pressure Lowering (effect only)
      2. Boiling Point Elevation (quantitative)
      3. Freezing Point Depression (quantitative)
      4. Osmotic Pressure (effect only)
    6. Colligative Properties of Electrolyte Solutions (van’t Hoff factor)

Laboratory Activities

The laboratory activities will include a safety overview including the location and demonstration of the use of laboratory safety equipment. Weekly hands-on activities will include 12-14 of the specific activities listed below.

  1. Resolution of Matter into Pure Substances – Paper Chromatography
  2. Determination of the Density of Liquids and Solids
  3. Properties of Hydrates
  4. Determination of a Chemical Formula
  5. Standardization of a Basic Solution and Determination of the Molar Mass of an Acid
  6. Determination of Iron by Reaction with Permanganate—A Redox Titration
  7. Verifying the Absolute Zero of Temperature
  8. The “Air Bag” Problem—Application of Stoichiometry and the Ideal Gas Equation
  9. Heat Effects and Calorimetry
  10. Alkaline Earths and Halogens
  11. What’s in These Bottles?—A Nonstandard Qualitative Analysis “Puzzle”
  12. Geometric Structure of Molecules—An Experiment Using Molecular Models
  13. Classification of Chemical Substances
  14. Molar Mass Determination by Depression of the Freezing Point

VII.  Methods of Instruction

  • Lectures, which may be supplemented with classroom discussion, use of molecular models, use of multimedia, and/or use of computer based materials at the discretion of the instructor.
  • Hands-On Laboratory Activities
  • Individual and/or Group Problem Solving

Course may be taught as face-to-face, hybrid or online course.

VIII. Course Practices Required

  • Participation in class and/or small group discussions
  • Problem solving to include basic algebraic manipulations
  • Hands-on laboratory activities
  • Course may be taught as face-to-face, hybrid or online course.


IX.   Instructional Materials

Note: Current textbook information for each course and section is available on Oakton's Schedule of Classes.

The course textbook is Chemistry: A Molecular Approach, 2nd ed. (2011) by Nivaldo Tro (ISBN: 978-0-3216-5178-5) OR Oakton’s custom version of this text for CHM 121 (ISBN: 978-0-5588-7433-9), or comparable text. 

The laboratory text is Chemistry 121 Lab Manual (ISBN:  978-0-4958-3138-9), which is a custom version of Chemical Principles in the Laboratory, 8th edition by Emil Slowinski, Wayne Wolsey, and William Masterton, (2005), or comparable text.

Beginning with the Spring 2007 semester, students will be required to purchase their own Chemical Safety/Splash Goggles.  These goggles must meet the following criteria:

  • Fit snuggly against the forehead and face, protecting against splashes
  • Be impact resistant; ANSI rating of Z87 or higher
  • Include only indirect venting

Two varieties of such goggles compliant with the above criteria are available for purchase in the bookstore.  Students may also elect to find an alternative source for purchase, as long as the goggles meet the above criteria and are approved by the instructor.

X.    Methods of Evaluating Student Progress

Depending upon the instructor, any combination of the following assessments may be used to evaluate student progress and determine the course grade.

  • Quizzes, tests, and/or examinations which may include essay, short answer, multiple choice, true/false, and/or problem solving questions
  • Laboratory assignments, reports, results and/or practicals
  • Individual and/or group written reports
  • Individual and/or group oral reports
  • Individual and/or group problem solutions

XI.   Other Course Information

  1. Reading of the text and laboratory material ahead of the class or laboratory session is expected.
  2. Regular attendance at all sessions is expected.  Missed laboratory sessions are not able to be made up.
  3. Class policies on make-up of exams and accepting of late work will be determined by the individual instructor.
  4. Students will be required to review and sign off on their review, understanding, and adherence to basic laboratory safety regulations.
  5. Support services include the availability of open computer laboratories, the college library, and the availability of free tutoring through the Learning Center and/or office hours with the course instructor.

If you have a documented learning, psychological, or physical disability you may be entitled to reasonable academic accommodations or services. To request accommodations or services, contact the Access and Disability Resource Center at the Des Plaines or Skokie campus. All students are expected to fulfill essential course requirements. The College will not waive any essential skill or requirement of a course or degree program.

Oakton Community College is committed to maintaining a campus environment emphasizing the dignity and worth of all members of the community, and complies with all federal and state Title IX requirements.

Resources and support for
  • pregnancy-related and parenting accommodations; and
  • victims of sexual misconduct
can be found at

Resources and support for LGBTQ+ students can be found at

Electronic video and/or audio recording is not permitted during class unless the student obtains written permission from the instructor. In cases where recordings are allowed, such content is restricted to personal use only. Any distribution of such recordings is strictly prohibited. Personal use is defined as use by an individual student for the purpose of studying or completing course assignments.

For students who have been approved for audio and/or video recording of lectures and other classroom activities as a reasonable accommodation by Oakton’s Access Disabilities Resource Center (ADRC), applicable federal law requires instructors to permit those recordings. Such recordings are also limited to personal use. Any distribution of such recordings is strictly prohibited.

Violation of this policy will result in disciplinary action through the Code of Student Conduct.