I. General Information

1. Course Title:
Fundamentals of Chemistry

2. Course Prefix & Number:
CHEM 1414

3. Course Credits and Contact Hours:

Credits: 4

Lecture Hours: 3

Lab Hours: 3

Internship Hours: 0

4. Course Description:

This course involves the study of general laws of chemistry, periodicity, atomic and molecular structure, physical and chemical changes. MnTC Goal 3

5. Placement Tests Required:

Accuplacer (specify test):

Next Gen Math

Score:

235

6. Prerequisite Courses:

9. Co-requisite Courses:

There are no corequisites for this course.

II. Transfer and Articulation

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

Bemidji State University, CHEM 1111 General Chemistry I, 4 credits
St. Cloud State University, CHEM 160 Preparatory Chemistry, 4 credits

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

III. Course Purpose

1. Program-Applicable Courses – This course is required for the following program(s):

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 oral communication skills	Make a presentation of lab group results and/or current events involving chemistry.
Analyze and follow a sequence of operations	Follow a given lab protocol to successfully complete a lab experiment in the time allowed.
Apply abstract ideas to concrete situations	Apply theories concerning particles which are not visible to construct atomic structure and calculate gram amounts to use in an experiment.
Work as a team member to achieve shared goals	Interact with lab partners to successfully complete a lab experiment in the time allowed.

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

• Demonstrate understanding of scientific theories and the ways in which scientists develop, express, and question theories in the field of chemistry (MnTC Goal 3);
• Formulate and test hypothesis by performing chemistry experiments requiring collection of data, its statistical and/or graphical analysis, and an appreciation of uncertainty and sources of error (MnTC Goal 3);
• Communicate their findings, analyses, and interpretations with other students and the instructor orally and in writing (MnTC Goal 3);
• Demonstrate the use of dimensional analysis to convert from one unit or dimension of measurement to another (MnTC Goal 3);
• Demonstrate how to determine the number of significant figures in a number resulting from addition, subtraction, multiplication and division of measured values (MnTC Goal 3);
• Describe the models and list the properties for the three states of matter (MnTC Goal 3);
• Classify changes in matter as either physical or chemical (MnTC Goal 3);
• Draw a Bohr model for any representative element and predict its chemical reactivity (MnTC Goal 3);
• Solve various gas law problems involving the combined gas law (MnTC Goal 3);
• Compare and contrast ionic and covalent bonds (MnTC Goal 3);
• Draw Lewis electron dot structures for a given molecular formula (MnTC Goal 3);
• Construct and name chemical formulas for various combinations of representative metals, nonmetals and polyatomic ions (MnTC Goal 3);
• Balance chemical equations (MnTC Goal 3);
• Perform various calculations involving grams, formula weight, moles, Avogadro’s number, percentage composition, empirical and molecular formulas (MnTC Goal 3); and
• Calculate the stoichiometric amounts in grams or moles for a given chemical equation when give an initial amount of a reactant (MnTC Goal 3).

V. Topical Outline

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

1. Lecture Sessions

1. Matter and Measurement
   • State the steps of the scientific method.
   • Explain the difference between a fact and an inference.
   • Recognize qualitative verses quantitative data.
   • Define the term scientific law.
   • Explain the difference between a hypothesis and a theory.
   • State the two properties matter must possess.
   • Define and recognize a physical verses a chemical change.
   • Recognize the three states of matter.
   • Define the terms homogeneous verses heterogeneous.
   • Describe how matter can be classified by attempting to separate it by physical changes.
   • Describe how matter can be further classified by attempting to separate it by chemical changes once it has been completely separated by physical changes.
   • Express decimal numbers in scientific notation and vice versa.
   • Express the results of multiplication and division of scientific numbers in scientific notation.
   • Perform multiplication and division of scientific numbers on a scientific calculator.
   • Write unit paths and solve any problem that involves converting from one unit of measurement to another—given a table of conversion factors.
   • Interpret the meaning of proportionality and of a proportionality constant.
   • Solve problems that involve the use of a proportionality constant as a conversion factor.
   • Convert Fahrenheit temperatures to Celsius and vice versa.
   • Convert Celsius temperatures to Kelvin and vice versa.
   • Distinguish between units of the English system and Metric system.
   • List the standard metric units of measurement and their abbreviations for length, mass and volume.
   • Write equivalencies, with their correct abbreviations, between any standard metric unit and a prefixed metric unit of kilo, centi, or milli.
   • Write equivalencies between metric units when utilizing a table of prefixed metric units.
   • Determine the number of significant figures in a number by having memorized the rules for counting significant numbers.
   • Determine the uncertainty in a given measured value.
   • Determine the correct number of significant figures after adding and subtracting measured values.
   • Determine the correct number of significant figures after multiplying and dividing measured values.
2. Atoms and Elements
   • Identify three subatomic particles by name, relative mass, and charge, and describe the structure of the atom.
   • Write and name a given isotope and determine the number of subatomic particles it contains.
   • Define an atomic mass unit, amu.
   • Calculate the atomic mass of an element given a set of isotopes, masses and abundances.
   • Distinguish between periods and families or groups in the periodic table.
   • Describe why groups or families of elements have similar chemical and physical properties.
   • Locate the alkali metals, halogens, and noble gases in the periodic table.
   • Predict the relative sizes of atoms within a group or family and period.
   • Identify the metals, nonmetals, and metalloids by their positions in the periodic table.
   • Identify representative and transition elements by their positions in the periodic table.
3. Molecules, Ions and Compounds
   • Name molecules.
   • Recognize the seven elements that form stable diatomic molecules.
   • Name and know the charges of the polyatomic ions asked to memorize.
   • Name the two types of ionic formulas that contain polyatomic ions.
   • Describe how scientists arrived at Avogadro’s number.
   • Define a mole in terms of mass and Avogadro’s number.
   • Convert grams of a given substance to moles and vice versa.
   • Convert grams of a given substance to numbers of atoms, molecules, or formula units or vice versa.
   • Calculate the percentage by mass of each element in a given formula.
   • Determine the empirical formula from percentage data.
   • Determine the true molecular or ionic formula from percentage data and formula weight.
4. Chemical Equations and Stoichiometry
   • Balance chemical equations.
   • Identify the abbreviations used in chemical equations for solids, liquids and gases.
   • Classify balanced equations into one of the six types discussed in class.
   • Identify common acids and bases.
   • Solve the four types of possible stoichiometry problems.
   • State the definitions of endothermic and exothermic reactions.
   • Draw energy level diagrams for endothermic and exothermic reactions.
   • Calculate percentage yield for a given stoichiometric reaction.
   • Atomic Structure
   • Describe the wave nature of electromagnetic radiation and how wavelength plays a role in its character.
   • Describe how observations of a discharge tube can be used to make of model of electron energy levels.
   • List the two types of energy an electron possesses.
   • Describe how the total energy of an electron changes with increased average distance from the nucleus.
5. Atomic Electron Configuration
   • Describe the shapes of the s and p orbitals.
   • List the various orbitals, sublevels, and principal energy levels and their capacities for electrons.
   • Write the electron configuration for any element in the periodic table using the short-hand notation.
   • Define valence electrons.
   • Write electron-dot structures for any representative element.
6. Bonding and Molecular Structure
   • Define the terms cations and anions.
   • Predict the number of electrons representative metals and nonmetals tend to loss or gain so as to acquire a noble gas electron configuration when they are involved in an ionic bond.
   • Predict the charge the representative metals and nonmetal acquire when they are involved in an ionic bond.
   • Use electron-dot structures to draw the electron transfers necessary between metals and nonmetals to form ionic bonds.
   • Predict ionic formulas from those metals and nonmetals that have only one charge.
   • Calculate the charge a transition metal when contained in an ionic bond.
   • Name the nonmetal anions.
   • Name ionic formulas that contain metals that have only one oxidation state and those that contain metals that can have more than one oxidation state.
   • Draw electron-dot structures for covalently bonded substances.
   • Draw electron-dot structures for polyatomic ions.
   • Define electronegativity and apply it to determine if a formula contains ionic or covalent bonds.
   • Define polar, nonpolar or essentially nonpolar covalent bonds and be able to use electronegativities to determine the type found in a formula.
   • State the elements which disobey the octet rule.
   • Predict the geometry of a molecule.
   • Determine if a molecule is polar or nonpolar.
7. Gases and Their Properties
   • State the properties of gases.
   • Describe the model for ideal gases.
   • Describe a barometer and open ended monometer and how to make pressure measurements from each.
   • Convert one pressure measurement to another given the relationship between the pressure measurements.
   • State Boyle’s and apply it to solve applicable problems.
   • State Gay-Lassac’s Law, Charles’ Law, and the combined gas law and apply them to solve applicable problems.
   • State Avogadro’s Law.
   • State the ideal gas law and apply it to solve applicable problems.
   • State or derive the gas density equation from the ideal gas equation and apply it to solve applicable problems.
   • State or derive the molar volume equation from the ideal gas equation and apply it to solve applicable problems.
   • Define STP.
8. Intermolecular Forces, Liquids, Solids
   • Predict how vapor pressure, molar heat of vaporization, boiling point, viscosity, and surface tension varies with intermolecular forces of attraction.
   • Identify the intermolecular forces of attraction in a substance from its formula.
   • Draw an kinetic energy distribution curve for molecules in a liquid.
   • Explain how escape energy of molecules in a liquid is related to the intermolecular forces of attraction and how its value relates to the liquids vapor pressure.
   • Explain how the kinetic energy distribution curve changes with temperature and how that affects the vapor pressure of the liquid.
   • Write a reversible equation to represent the equilibrium between a liquid and its vapor state.
   • Define boiling point of a liquid in terms of the liquids vapor pressure and the total gaseous pressure exerted on the liquid and explain how the boiling point of a liquid can be changed.
   • Identify the five types of solids and their respective physical properties.
   • Define heats of vaporization, condensation, fusion, and solidification.
   • Calculate the amount of energy required to vaporize a given mass of liquid at its boiling point or condense the same amount of gas at its condensation temperature.
   • Calculate the amount of energy required to melt a given mass of solid at its melting point or freeze the same amount at its freezing point.
   • Identify melting, freezing, boiling, or condensation as endothermic or exothermic processes.
   • Define specific heat.
   • Calculate the amount of energy required to be added or removed to change the temperature of a given amount of solid, liquid, or gas.
   • Calculate the amount of energy required to be added or removed to make any variation in temperature including any combination of phase changes.
   • Solutions and Their Behavior
   • Explain the electric properties of polar and nonpolar molecules.
   • Know what hydrogen bonding is and when it occurs.
   • Know the definitions of homogeneous and heterogeneous mixtures, solutions, solute, and solvent.
   • Know how polar and nonpolar molecules effect the mixing process between solute and solvent.
   • Know how to calculate % wt., ppm, and molarity of a solute in a solution.
   • Know how to use molarity as a conversion factor to solve the four types of molarity problems discussed in class.
   • Know how to calculate dilution problems.

2. Laboratory/Studio Sessions

1. Laboratory Techiques
2. Introduction to Quantitative Measurement
3. Effect of Temperature on Solubility of a Salt
4. Separation Techniques
5. Freezing and Melting of Water
6. Chemical and Physical Changes
7. Properties of Solutions: Electrolytes and Non-Electrolytes
8. Chemical Changes and Equations
9. Endothermic and Exothermic Reactions
10. Determining the Concentration of a Solution: Beer’s Law
11. Acid-Base Titration
12. Gravimetric Determination
13. Boyle’s Law
14. Preparation of Soap