I. General Information

1. Course Title:
Chemical Principles I

2. Course Prefix & Number:
CHEM 1424

3. Course Credits and Contact Hours:

Credits: 5

Lecture Hours: 4

Lab Hours: 3

Internship Hours: 0

4. Course Description:

This course includes a more rigorous collegiate treatment of topics in physical measurement, dimensional analysis, state of matter, nomenclature, chemical reactions, stoichiometry, gas laws, thermochemistry, atomic structure, and molecular bonding theory. MnTC Goal 3

5. Placement Tests Required:

Accuplacer (specify test):

Reading College Level CLC or Reading College Level, and Math Introductory College Level or Algebra College Level or Pre-Calculus College Level or Calculus College Level

Score:

6. Prerequisite Courses:

7. Other Prerequisites
This course requires an Accuplacer Next Gen Advanced Algebra Functions score of 250 or higher, or completion of MATH 0820 or MATH 1520.

9. Co-requisite Courses:

There are no corequisites for this course.

II. Transfer and Articulation

1. Course Equivalency - similar course from other regional institutions:

St. Cloud State University, CHEM 210 General Chemisry I, 4 credits
Bemidji State University, CHEM 1211 Principles of Chemistry I, 4 credits

2. Transfer - regional institutions with which this course has a written articulation agreement:

3. Prior Learning - the following prior learning methods are acceptable for this course:
Advanced Placement (AP)

III. Course Purpose

2. MN Transfer Curriculum (General Education) Courses - This course fulfills the following goal area(s) of the MN Transfer Curriculum:
Goal 3 – Natural Sciences

IV. Learning Outcomes

1. College-Wide Outcomes

College-Wide Outcomes/Competencies	Students will be able to:
Demonstrate written communication skills	Complete written reports using data collected from the laboratory
Apply abstract ideas to concrete situations	Predict physical properties of matter using atomic and molecular theories
Work as a team member to achieve shared goals	Work together as team members to successfully complete laboratory experiments

2. Course Specific Outcomes - Students will be able to achieve the following measurable goals upon completion of the course:

• Determine the number of significant figures in a number resulting from addition, subtraction, multiplication and division of measured values (Goal 3);
• Use dimensional analysis to convert from one unit or dimension of measurement to another (Goal 3);
• Describe and apply the scientific method used by scientists in solving problems (Goal 3);
• Construct and name chemical formulas for various combinations of representative metals, nonmetals and polyatomic ions (Goal 3);
• Describe electrons, protons, and neutrons, and the general structure of the atom (Goal 3);
• Define isotope and determine the atomic number, mass number, and number of neutrons for a specified isotope as well as identify atomic number and atomic mass for any element (Goal 3);
• Calculate the average atomic mass of an element from isotopic abundances and isotopic masses (Goal 3);
• Correlate wavelength, frequency, and energy of light with electron energy levels in the atom via the photoelectric effect and the Bohr model (Goal 3);
• Apply wave-particle duality and the uncertainty principle to describe properties of electrons (Goal 3);
• Apply the results of the Schrödinger quantum mechanical model of the atom to assign quantum numbers to electrons and write electron configurations of multi-electron atoms and ions (Goal 3);
• Predict patterns of ionization energies, electron affinities, electronegativities and sizes of atoms and ions within periods and groups of the periodic table using quantum theory (Goal 3);
• Perform various calculations involving grams, formula weight, moles, Avogadro’s number, percentage composition, empirical and molecular formulas (Goal 3);
• Balance chemical equations and use stoichiometric relationships to calculate product and reactant amounts, determine limiting reagents, and calculate percent yields (Goal 3);
• Solve various gravimetric and volumetric stoichiometry problems (Goal 3);
• Identify insoluble ionic compounds, strong and weak acids and bases, and oxidizing and reducing agents in precipitation, acid-base neutralization, and oxidation-reduction (redox) reactions and how to predict products, identify spectator ions, and write net ionic equations (Goal 3);
• Know the relationships between heat, work, energy, enthalpy as it relates to thermochemical equations, heat transfer involving temperature measurements, heat capacity or specific heats (Goal 3);
• Apply Hess’s Law and enthalpies of formation to determine enthalpies of reaction as well as reaction enthalpies using calorimetry data (Goal 3); and
• Describe bonding in pure covalent, polar covalent and ionic structures, using Lewis structures and Valence Bond Theory to predict shape, polarity and bonding (Goal 3).

V. Topical Outline

Listed below are major areas of content typically covered in this course.

1. Lecture Sessions

1. Basic Concepts of Chemistry
   
   • Hypotheses, Laws, and Theories
   • States of Matter and Kinetic Molecular Theory
   • Matter at the Macroscopic and Particulate Levels
   • Pure Substances
   • Mixtures: Homogeneous and Heterogeneous
   • Elements and Atoms
   • Compounds and Molecules
   • Physical Properties
   • Extensive and Intensive Properties
   • Physical and Chemical Changes
2. The Tools of Quantitative Chemistry
   
   • Units of Measurement
   • Temperature Scales
   • Experimental Error
   • Standard Deviation
   • Exponential of Scientific Notation
   • Significant Figures
   • Problem Solving by Dimensional Analysis
3. Basic Concepts of Chemistry
   
   • Hypotheses, Laws, and Theories
   • States of Matter and Kinetic Molecular Theory
   • Matter at the Macroscopic and Particulate Levels
   • Pure Substances
   • Mixtures: Homogeneous and Heterogeneous
   • Elements and Atoms
   • Compounds and Molecules
   • Physical Properties
   • Extensive and Intensive Properties
   • Physical and Chemical Changes
4. The Tools of Quantitative Chemistry
   
   • Units of Measurement
   • Temperature Scales
   • Experimental Error
   • Standard Deviation
   • Exponential of Scientific Notation
   • Significant Figures
   • Problem Solving by Dimensional Analysis
5. Atoms, Molecules, and Ions
   
   • Atomic Structure
   • Atomic Number
   • Atomic Weight and the Atomic Mass Unit
   • Mass Number
   • Isotope Abundance
   • Determining Atomic Mass and Isotope Abundance
   • Atomic Weight
   • The Periodic Table
   • Molecules, Compounds, and Formulas
   • Ionic Compounds
   • Formulas of Ionic Compounds
   • Names of Ionic Compounds
   • Atoms and Molar Mass
   • Percent Composition
   • Empirical and Molecular Formulas from Percent Composition
   • Determining a Formula from Mass Data
   • Hydrated Compound
6. Chemical Reactions
   
   • Balancing Chemical Equations
   • Chemical Reactions in Aqueous Solution
   • Ions and Molecules in Aqueous Solution
   • Predicting the Outcome of a Precipitation Reaction
   • Net Ionic Equations
   • Acids and Bases: The Arrhenius Definition
   • Acids and Bases: The Bronsted-Lowry Definition
   • Gas-Forming Reactions
   • Oxidation-Reduction Reactions
   • Classifying Reactions in Aqueous Solution
7. Stoichiometry: Quantitative Information about Chemical Reactions
   
   • Mass Relationships in Chemical Reactions: Stoichiometry
   • A Stoichiometry Calculation with a Limiting Reactant
   • Percent Yield
   • Chemical Equations and Chemical Analysis
   • Solution Concentration: Molarity
   • Preparing Solutions of Known Concentration
   • pH, a Concentration Scale for Acids and Bases
   • Solution Stoichiometry
   • Standardizing and Acid or Base
8. Principles of Chemical Reactivity: Energy and Chemical Reactions
   
   • Temperature and Heat
   • Systems and Surroundings
   • Energy Units
   • Specific Heat Capacity
   • Energy Changes of State
   • The First Law of Thermodynamics
   • Enthalpy
   • Enthalpy
   • State Functions
   • Enthalpy Changes for Chemical Reactions
   • Calorimetry
   • Hess’s Law
   • Standard Enthalpies of Formation
   • Enthalpy Change for a Reaction
9. The Structure of Atoms
   
   • Electromagnetic Radiation
   • Planck’s Equation
   • Einstein and the Photoelectric Effect
   • Energy and Chemistry: Using Planck’s Equation
   • The Bohr Model of the Hydrogen Atom
   • Particle-Wave Duality
   • Quantum Numbers and Orbitals
   • Shells and Subshells
   • The Shapes of Atomic Orbitals
   • Electron Spin
10. The Structure of Atoms and Periodic Trends
    
    • The Pauli Exclusion Principle
    • Atomic Subshell Energies and Electron Assignments
    • Effective Nuclear Charge, Z*
    • Electron Configurations of Atoms
    • Electron Configurations of Ions
    • Atomic Properties and Periodic Trends
    • Periodic Trends and Chemical Properties
11. Bonding and Molecular Structure
    
    • Covalent Bonding and Lewis Structures
    • Atom Formal Charges in Covalent Molecules and Ions
    • Resonance
    • Exceptions to the Octet Rule
    • Molecular Shapes
    • Bond Polarity and Electronegativity
    • Bond and Molecular Polarity
    • Bond Properties: Order, Length, Energy
12. Bonding and Molecular Structure: Orbital Hybridization and Molecular Orbitals
    
    • Orbitals and Theories of Chemical Bonding
    • Valence Bond Theory

2. Laboratory/Studio Sessions

1. Safety and Laboratory Techniques
2. Measurement of Density
3. Hydrate Analysis
4. Determination of a Formula
5. Chemical Nomenclature
6. Copper Cycle
7. Vitamin C Analysis
8. Standarization of HCl and NaOH
9. Gravimetric Analysis
10. Antacid Analysis
11. Enthalpy of Neutralization
12. Energy in Foods