Academic Chemistry Syllabus

  • Academic Chemistry Semesters I & II. Curriculum Both Semesters

    I. Course Length:

    1 Semester (Two 9-weeks, 5 Class periods / week, 2 Lab periods / II. Grade Level: 10

    III. Credits: 0.7 credit per semester (1.4 credits per school year)

    IV. Course Description:

    This is an introduction course to chemistry designed for students planning on attending 4 YEAR COLLEGE ONLY. Students use their advanced math skills and problem solving techniques to learn the basics of chemistry. This course also includes a lab period two days per week. Prerequisites: C or better in the following course - Biology, Algebra I, Geometry.

    V. Course Outline:

    Semester I.

    Unit 1:
    Objectives: After successful completion of this unit the student will be able to:

    1. Understand the difference between precision and accuracy.

    2. Convert between Standard and Scientific notation.

    3. Perform mathematic operations involving significant figures.

    4. Understand the fundamental principles of the metric and SI systems.

    5. Solve problems utilizing dimensional analysis and factor labeling.

    A. Measurements in chemistry

    1. Precision

    2. Accuracy

    3. Standard notation

    4. Scientific notation

    5. Converting

      a. Standard notation to Scientific notation b. Scientific notation to standard notation

    B. Significant figures

    1. Defined

    2. Why we need them?

    3. Adding / Subtracting with

    4. Multiplying and dividing with

    5. Combined operations

    C. SIsystem

    1. Length

    2. mass

    3. volume

    4. time

    5. current

    6. temperature

    7. amount of substance

    D. SI system Derived units

    1. Area
    2. Density 3. Pressure

    E. Chemistry Problem solving

    1. Dimensional Analysis

    2. Factor label method

    3. Handout problem solving strategies

    Unit 2:
    Objectives: After successful completion of this unit the student will be able to:

    1. Differentiate between the various branches of chemistry.

    2. Understand what pure and applied chemistry are.

    3. Apply the Scientific method to solve problems.

    F. Chemistry intro

    1. Defined

    2. Father of Chem.

    3. Chem. All around us

    4. Chem. Contribution to other sciences

    G. Branches of Chemistry

    1. Biochemistry

    2. Organic Chemistry

    3. Inorganic Chemistry

    4. Physical Chemistry

    5. Analytical Chemistry

      a. Qualitative b. Quantitative

    H. Pure vs. Applied Chemistry

    I. Scientific Method

    1. Intro (VCR TAPE ANAOLGY)

    2. Definition

    3. Observation

    4. Hypothesis

    5. Experiment

    6. Conclusion

    7. Theory

    8. Law

    9. Theory vs. Law

    Unit 3:
    Objectives: After successful completion of this unit the student will be able to:

    1. Explain the principle theories of atomic structure and electron notation.

    2. Define the various subatomic particles.

    3. Use the periodic table to find atomic mass, atomic number and so on.

    4. Identify uses for various isotopes.

    5. Calculate formula mass.

    6. Explain the laws of “Conservation of Matter and Energy”

    7. Identify the various types of matter (e.g., homogeneous,

      heterogeneous) and methods for their separation.

    8. Identify physical and chemical properties of matter.

    9. Determine the difference between a fluid and a liquid, a gas and a

      vapor.

    10. Separate a mixture by physical means (e.g., filtration, distillation)

    J. Matter

    1. Defined

      1. Mass

      2. Weight

    2. Heterogeneous vs. Homogeneous

    3. Heterogeneous

    a. Mixtures (heterogeneous)

    4. Homogeneous

    1. Solutions

    2. Pure Substances

    3. Elements

    1. Atoms, Intro

      1. Defined 2. Theories

    2. Atoms

    a. Dalton
    b. Bohr
    c. Rutherford

    1. Subatomic particles
    a. Electrons

    b. Protons
    c. Neutrons
    d. Relative electrical charge

    M. Structure of the Atom

    1. Nucleus
    2. Rutherford’s Foil experiment

    N. Atomic Number

    1. Name

    2. Symbol

    3. Atomic Number

    4. Protons

    5. Neutrons

    6. Mass Number

    7. Finding Numbers of Protons, Neutrons, Electrons

    O. Mass
    1. Define

    1. Isotopes of the Elements

      1. Defined 2. Examples

      e. Hydrogen f. C

    2. Atomic Mass

    1. Atomic Mass Unit
    a. Defined

    b. Nature, Elements, and Isotopes

    R. Calculating the Atomic Mass of an Element 1. Examples

    Number

    Unit 4:
    Objectives: After successful completion of this unit the student will be able to:

    1. Read the periodic table.

    2. Explain the relationship between the electronic structure of an

      element and its position on the Periodic table.

    3. Predict the relationship between the electronic structure of an element

      and its position on the Periodic table.

    4. Apply the concepts of the Pauli Exclusion Principle, Aufbau principle,

      and Hund’s Rule to draw orbital diagrams and electron

      configurations.

    5. Write electron configurations of various elements.

    S. Intro

    1. Remember
    a. Location of electrons

    b. Valence
    c. Core
    d. Representation of valence and core electrons

    T. Periodic “Short Cuts”

    1. Families/Groups, vertical columns and valence electrons

    2. Series/Periods, horizontal row and energy levels

    U. Energylevels

    1. Location/ Arrangement on chart

    2. Defined

    3. Drawing of Ground State vs. Excited state of electron

    4. Ground state

    5. Excited state

    6. “Fluorescent Light Demo.”

    a. Conservation of Energy

    V. “The Diagonal Rule” (The order of electron fill.)

    1. Copy from over head

      1. Aufbau Principal (“Building up Order”)

      2. Electron Configuration

      3. Sublevel

      4. Space Orbital

      5. Orbital Diagram

    2. 1st Series

    3. 2nd Series

    a. Hund’s Rule

    4. 5. 6.

    3rd Series 4th –5th

    6th –7th

    1. Pauli Exclusion Principal

    2. Electron Dot notation

    1. Nobel Gas Core

    2. Halogens

    1. Chromium

    2. Copper

    SEMESTER I & II OVERLAP

    Unit 5:
    Objectives: After successful completion of this unit the student will be able to:

    1. Draw Lewis structures of various elements and their ions.

    2. Explain the mole concept.

    3. Interconvert moles and mass.

    4. Use standard molar volume (22.4 l = 1 mole) to convert moles to

      volume of gas at standard conditions.

    5. Use moles and interconvert between grams, liters of a gas, and

      number of particles.

    6. Understand the fundamental principles of electrochemistry. (e.g.,

      electrolytic and voltaic cells, cathodes and anodes.)

    7. Classify 1⁄2 cells reactions and reduction or oxidation.

    A. Ions (Polyatomic)

    1. Cation

    2. Anion

    3. Polyatomic ions

    a. 4 steps to drawing poly ions

    1. count “4 key numbers”

    2. draw skeleton

    3. give everybody 8 electrons

    4. check and adjust as necessary

    B. Mole

    1. Defined

    2. Avagadro’s number

    3. gram atomic weight

    C. Molarity

    1. defined

    2. relation to grams per unit volume

    3. solving mole/ Molarity problems (dilutions, concentrations)

    1. DIMO

      1. Acronym defined

      2. use in converting

        1. grams, liters, # of particles to mole

        2. # of moles to grams, liters, # of particles

    2. Electrochemistry

      1. Defined

      2. electrolytic cell

      3. voltaic/ galvanic cell

      4. cathodes/ anodes

    3. Redox reactions in electrochemical cells

      1. Oxidation

      2. Reduction

      3. 1⁄2 cell reactions

      4. cell potentials

        SEMESTER II.

    Unit 6:
    Objectives: After successful completion of this unit the student will be able to:

    1. Identify and differentiate the basic types of chemical bonds: ionic and covalent.

    2. Understand the relationship between bonding and physical properties.

    3. Indicate how electrons are shared and transferred to form bonds.

    4. Identify the types of chemical bonds in common substances (e.g.,

      Sodium chloride, Carbon dioxide, Methane)

    5. Show how bonding determines physical properties.

    6. Name chemical compounds from formulas (e.g., NaCl is Sodium

      chloride)

    7. Write chemical formula for compounds (e.g., Sodium chloride is

      NaCl)

    8. Identify the type of electron distribution in a given substance (e.g.,

      polar, ionic)

    9. Apply the concepts of basic chemical nomenclature.

    G. Chemical Bonds (General Overview)

    1. Chemical Bonds Defined

    2. Two Types of Bonds

      1. Ionic

      2. Covalent

    3. Specific Bonds form when...

      1. Metal to Non-Metal

      2. Non-Metal to Non-Metal

      3. Metals ≠ DO NOT bond with Metals

      4. Chart for indications of specific bonds

    4. Electronegativity

      1. Defined

      2. Trend with periodic chart

      3. Use in determining bonds

    H. Ionic Bonding

    1. Remember definition and when they occur.

    2. Ionic Compounds defined

    3. Ionic Bonds use valence electrons

    4. Refresher: Difference between Metals and Non-Metals

      1. Location on Periodic Chart

      2. Number of valence electrons, and the number desired

      3. “Gain” or “Give”

    5. Cation - Ion Formation

    6. Stoichiometry (“Balancing”)

    7. Nomenclature (“Naming Compounds”)

    Unit 7:
    Objective: After successful completion of this unit the student will be able to:

    1. Draw Lewis structures for resonance.

    2. Determine the various types of intermolecular forces.

    3. Identify/determine the various types of electron distributions in

      chemical bonding.

    4. Classify and identify the various types of molecular geometry and

      VSEPR theory.

    5. Explain the basis of the VSEPR theory.

    6. Relate molecular slope to electron structure.

    7. Predict the physical state and properties of a substance at a given

      temperature by consideration of intermolecular forces.

    I. Covalent Bonding

    1. Remember definition and when they occur.

    2. Covalent Compounds defined

    3. Covalent Bonds use various electrons

      4. Refresher: Difference between Metals and Non-Metals a. Location on Periodic Chart

    b. Number of valence electrons, and the number desired c. Everyone needs. Know one “Gains” or “Gives”. ALL

    SHARE.
    5. Stoichiometry (“Balancing”)

    6. Nomenclature (“Naming Compounds”)

    J. Resonance

    1. Defined

    2. Reasons for

    3. Examples

    K. Exceptions to Octet Rule

    1. Paramagnetic

    2. Diamagnetic

    3. Examples

    L. Intermolecular Forces

    1. Van der Waals

    1. Dispersion forces

    2. Dipole interactions

    i. Hydrogen bonds

    M. VSEPR Theory

    1. Defined 2. Why
    3. Examples

    Unit 8:
    Objectives: After successful completion of this unit the student will be able to:

    1.

    2. 3.

    4.

    N. Chemical

    1. 2. 3.

    Identify and differentiate the various symbols used in a chemical equation.
    Apply the principles of Stoichiometry.
    Write balanced chemical equation show that atoms are conserved in chemical reactions.

    Perform titrations to determine the concentration of an unknown acid or base.

    Equation

    Defined
    Basic Terms
    Frequently used Symbols

    O. Balancing Chemical Equations

    1. Remember “Nomenclature” and “Stoichiometry”

    2. Basic Rules

    3. Setting up the “Chart”

    4. Balancing Practice and Examples

    Unit 9:
    Objectives: After successful completion of this unit the student will be able to:

    1. Identify the basic types of chemical reactions.

    2. Define and differentiate the types of chemical reactions (e.g., single

      and double replacement, combination/synthesis, decomposition)

    3. Use the activity series to predict reactivity.

    4. Predict mass and/or volume of a species in chemical reactions.

    5. Calculate percent yields and identify limiting reagent.

    6. Perform calculations involving concentrations of solutions. (e.g.,

      molarity, actual yield)

    7. Classify and predict what type of reaction will occur. (e.g.,

      decomposition, combination, single replacement, double replacement,

      reversible, and redox)

    8. Identify and determine products and write balanced equations for

      reactions.

    9. Explain the heat change in a chemical reaction.

    P. Types

    of Reactions

    1. Endothermic

    2. Exothermic

    3. Combustion

    4. Synthesis

    5. Decomposition

    6. Single Replacement

    7. Activity Series

    8. Double Replacement

    9. Redox. (Oxidation / Reduction)

    Q. Calculations from Balanced Reactions

    1. Moles

    2. Grams

    3. Liters of Gas

    4. Number of Particles

    5. Limiting Reagents / Excess

    Unit 10:
    Objectives: After successful completion of this unit the student will be able to:

    1. Describe the various relationships among the variables that describe gases (e.g., pressure, temperature, volume, and moles.)

    2. Use standard molar volume (22.4 liters = 1 mole) to convert moles to volume of gas a standard conditions.

    3. Define the Universal Gas Constant.

    4. Predict the physical state and properties of a substance at a given

      temperature by consideration of pressure, volume, and temperature.

    5. Understand the concept of an Ideal Gas.

    6. Differentiate between a gas and a vapor.

    7. Use “The Gas Laws” to predict the pressure, volume, temperature

      and amount of a gas at specific conditions.

    8. Explain how the observed properties of gases (e.g., expansion,

      pressure, low density, and diffusion) relate to the physical and chemical properties of a gas.

    R. Intro

    1. Remember:
    a. 3 states of matter

    S. Basic

    Terms

    b. Differences in energy and atom/molecule proximity

    1. Pressure

    2. volume

    3. Mole

    4. Universal gas constant

    5. Temperature

    6. STP

    7. Ideal gas

    T. Observed Properties of Gases

    1. Expansion

    2. Pressure

    3. Low Density

    4. Diffusion

    U. Kinetic Theory (Description of a Gas)

    1. Defined

    2. Specific assumptions

    3. Attractive forces (Effect)

    1. Boyle’s Law

      1. Definition 2. Examples

    2. Charles’s Law

      1. Definition 2. Examples

    3. Combined Gas Law

      1. Definition 2. Derivation 3. Examples

    4. Ideal Gas Law

      1. Definition 2. Examples

    5. Dalton’s Law of Partial Pressures

      1. Definition

      2. Partial pressure

      3. Total pressure

      4. Mole fraction

      5. Examples

    VI. Assessment:

    33% Homework, Labs, class work, maintenance of lab drawer

    67% Test and quizzes