MADISON PUBLIC SCHOOLS
Advanced Placement Chemistry
Authored by: Claire Miller
Reviewed by: Dr. Barbara Sargent
Assistant Superintendent for Curriculum and Instruction
Board of Education approval: July 2008
Members of the Board of Education:
Dr. Richard Noonan, Superintendent
Lisa Ellis, President
Patrick Rowe, Vice-President
David Arthur
Kevin Blair
Linda Gilbert
Shade Grahling
George Martin
James Novotny
Madison Public Schools
359 Woodland Road, Madison, NJ 07940
www.madisonpublicschools.org
I. OVERVIEW
Advanced Placement Chemistry (AP) at the high school level is a course designed to be the equivalent of a first year
of general college chemistry. Students in AP Chemistry usually have taken Honors Chemistry and Honors Physics as
prerequisites. They have elected to continue advanced study because of their interest, ability, past performance and teacher
recommendation. After successful completion of the course of study, students are required to take the Advanced Placement
Chemistry exam given by the College Board of Princeton, New Jersey.
The curriculum is centered around fundamental concepts in chemistry. These include the atomic structure,
periodicity, chemical bonding, molecular structure, the states of matter, the colligative properties of solutions, chemical
reactions, stoichiometry, kinetics, thermodynamics, equilibrium, acid-base chemistry, nuclear chemistry and
electrochemistry. Descriptive chemistry is also covered. It serves to illustrate the concepts and to provide real world
applications.
The class meets seven periods (45 minutes each) a week. Since there are two double periods, a significant portion
of class time is spent in the laboratory. Experiments are used to illustrate concepts, to reinforce theory and to develop
competence in dealing chemical problems. Emphasis is placed on the calculation and interpretation of quantitative data, as
well as the evaluation of experimental procedures. Students have the opportunity to work in a collaborative manner to
perform both microscale and macroscale experiments. In addition, the instructor makes use of chemical demonstrations
wherever possible to further illustrate the concepts that are studied. Both course and laboratory work are intended to aide in
the development of the students’ ability to think and communicate ideas with clarity and logic.
II.
RATIONALE
.
Students completing the AP Chemistry course will be prepared for the AP Chemistry exam. The content of this
course is guided by the College Board and represents a typical introductory college chemistry course. Successful completion
of the course and the subsequent exam may earn college credit.
III. COURSE REQUIREMENTS
In order to fulfill all the requirements set forth by the College Board, the class must move at a quick pace. Due to
the rigorous nature of this course, students are expected to be responsible, self-disciplined, and hardworking. They must
actively participate in completing laboratory experiments to enhance and reinforce learning. Students must work safely and
responsibly in the laboratory. Since the class is centered on problem solving sessions and group work in the laboratory,
teamwork is emphasized. Also, each student is expected to complete a summer assignment before entering AP Chemistry.
The assignment consists of problem sets from the first two chapter of the textbook, as well as the memorization of the
names/formulas/charges of the common ions and the solubility rules.
During the school year, time lines are provided to students every two weeks. The time lines list the class schedule
with all laboratory exercises and assignments including due dates for each assignment. Since students are given ample
notice, work is expected to be handed in on time with the exception of absence. About three weeks before the AP test, each
student is given a review folder. The folder consists of retired AP free response and multiple choice tests. In addition, there
are sheets with contain AP problems by topic. As each topic is reviewed in class, students are assigned 2 to 3 problems to
complete for homework. Problems are corrected and review the next day. Also, time is spent on predicting the products
and writing equations for reactions of various types.
IV. STUDENT OUTCOMES (Linked to N.J. Core Curriculum Standards)
I
Matter and Measurement (Summer Asssignment)
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Differentiate between pure substance and mixtures. (5.6)
2. Note the difference between physical and chemical properties, as well as physical and chemical change.
(5.6).
3. Use SI units for measured quantities. (5.3)
4. Apply the rules of significant figures to calculations involving experimental data. (5.3)
5. Use dimensional analysis to solve problems. (5.3)
B.
II
The Nature of Matter (Summer Assignment)
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Relate the experiments of Thomson, Milikan and Rutherford to the modern view of the atom. (5.2)
2. Describe the modern atomic mass scale. (5.3)
3. Determine the number of electrons, protons and neutrons given the appropriate information. (5.6)
4. Use the periodic table to show the major division of elements. (5.6)
5. Write the formula or name an ionic compound.
6. Write the formula or name of a binary molecular compound.
7. Write the formula or name an acid
B.
III
Laboratory Exercises and Activities:
1. The Language of Chemistry: Writing Formulas and Naming Compounds (Hague and Smith)
Stoichiometry
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Explain the importance of Avogadro’s number of particles and the mole. (5.3)
2. Determine an empirical formula form analysis data. (5.3)
3. Determine the molecular formula. (5.3)
4. Write a balanced equation of a reaction. (5.6)
5. Predict the mass of a product or a reactant given appropriate data. (5.3)
6. Relate the actual yield of the product to theoretical yield by calculating the percent yield for a reaction.
(5.3)
7. Determine the limiting reagent. (5.3)
B.
IV
Laboratory Exercises and Activities:
1. Safety in the Laboratory
2. “How Big Is An Atom” or the Determination of the diameter of a zinc atom on a piece of galvanized
iron-2 periods (Flinn)
3. “The Estimation of Avogadro’s Number” - 2 periods (Hollenberg,Postma, and Roberts)
Laboratory Exercises and Activities:
1. Determination of an Empirical Formula: “Synthesis of Manganese Chloride” -2 periods (George Gross)
2. Determination of Percent Yield of a Reaction: “The Decomposition of Chlorate” -2 periods (Flinn)
Solution Chemistry
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Determine the percent composition and or the molarity of a solution. (5.3)
2. Characterize reactions that occur in aqueous solution in terms of molecules and ions present. (5.6)
3. Classify methathesis reactions by the formation of an insoluble substance, a weak or nonelectrolyte, and
the formation of a gas. (5.6)
4. Use solubility rules to predict the formation of a precipitate.
5. Predict the products of an acid-base reaction. (5.6)
6. Recognize redox reactions that occur in solution.
7. Assign oxidation state numbers.
8. Write net ionic equations for a reaction. (5.6)
9. Extend the principles of stoichiometry to reactions in solution. (5.3)
10. Use titration data to determine volume, concentration, molecular mass or the composition of a solution.
(5.3)
B.
Laboratory Exercises and Activities:
1. Teacher Demonstation- “Using Conductivity to Find the Equivalence Point of Titration” (Vernier)
2. “Measuring the Conductivity of Solution” (Vernier)
3. Titration of Acid-Base – The Standardization of NaOH and The Determination of the Molecular Mass
of a Solid Acid-2 periods (Kemp and Nelson)
V
Thermodynamics
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Define energy and the First Law of Thermodynamics. (5.7)
2. Relate the change in internal energy to heat and work. (5.6)
3. Define the change in enthalpy and relate it to heat flow. (5.7)
4. Determine the change in enthalpy using calorimetry data or Hess’s Law. (5.3)
5. Determine the change in enthalpy using heats of formation. (5.3)
6. Define the Second Law of Thermodynamics. (5.7)
7. Use the standard molar entropy values to determine the change in entropy of a reaction.(5.3)
8. Determine the change in free energy using the free energies of formation. (5.3)
9. Use the Gibbs equation to calculate the change in free energy and predict the spontaneity of a reaction.
(5.3)
10. Determine the effect of temperature on the spontaneity of a reaction.
11. Relate the change in free energy to the equilibrium constant of a reaction.
B.
VII
Electronic Structure of Atoms
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Describe the electromagnetic radiation in terms of wavelength, frequency and energy. (5.7)
2. Calculate the energy of a wavelength of light. (5.3)
3. Describe the Bohr model of the hydrogen atom. (5.2)
4. Explain emission spectra in terms of the quantized energy state of an electron. (5.7)
5. Use the Bohr model to determine the wavelength of light that corresponds to an energy transition of an
electron in a hydrogen atom. (5.2)
6. Describe the quantum mechanical model of the atom. (5.2)
7. Tell the meaning of the four quantum designation for an electron.
8. Write an acceptable set of four quantum numbers for an electron in an atom.
9. Describe an orbital in terms of energy, shape and orientation.
10. Write an electron configuration of an atom or ion in its ground state. (5.6)
11. Expand the configuration to show the orbital and spin of the electron. (5.6)
12. Write the configuration in the abbreviated form. (5.6)
B.
VIII
Laboratory Exercises and Activities:
1. “Hess’s Law Laboratory” The Determination of the Heats of Solution and Neutralization-2 periods
(Vernier)
2. “Determination of the Heat of Combustion of Magnesium” -2 periods (Vernier)
3. Demonstration-“Entropy as a Driving Force” -.5 period (Bilash, Koob and Gross)
Laboratory Exercises and Activities:
1. Atomic Spectra and Atomic Structure -2 periods (Kemp and Nelson)
2. Determination of the Formula of a Hydrate -2 periods
Periodic Properties
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Relate the element’s position on the table to the electron configuration. (5.6)
2. Predict and explain the general trends on the table with respect to atomic size, ionization energy,
electron affinity, and metallic character. (5.6)
3. Arrange elements in order of relative size, IE, electron affinity and metallic character. (5.6)
4. Compare the chemical and physical properties of groups of metals and nonmetals. (5.6)
5. Use periodic position to predict the properties of the representative elements and their compounds. (5.6)
6. Differentiate between groups of metals and nonmetals on the basis of relative reactivity. (5.6)
B.
IX
Chemical Bonding
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Explain why ionic bonds form. (5.6)
2. Describe the energies involved in the formation of an ionic bond using the Born-Haber cycle.(5.6)
3. Relate the charges on ions to the electronic structure.
4. Predict and explain the trend ion ionic radius.
5. Explain the relationship between electronegativity and bond polarity.
6. Draw Lewis structures for molecules and polyatomic ions.
7. Use formal charge to determine the preferred Lewis structure.
8. Predict if a species will exhibit resonance and write resonance structures.
9. Relate bond strength to bond enthalpy. (5.6)
10. Estimate the change in enthalpy using bond energies. (5.3)
B.
X
Laboratory Exercises and Activities:
1. Drawing of Lewis Structures of Molecules and Polyatomic Ions -1 period
Molecular Geometry and Bond Theory
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Use VSEPR theory to predict geometry of molecules including those with expanded and incomplete
octets.
2. Predict the bond angles within molecules including the deviations due to lone pairs of electrons.
3. Predict the polarity of molecules.
4. Explain why hybridization is required to explain the bonding around the central atom of molecules.
5. Predict the type of hybridization around the central atom in a molecule.
6. Differentiate between a sigma and pi bond and count their relative number in a given structure.
7. Use Molecular Orbital theory to describe the bonding in simple diatomic molecules.
8. Use Molecular Orbital theory to determine the bond order and the magnetic properties of molecules.
9. Relate bond order to the relative number of bonds, bond length and bond strength.
B.
XI
Laboratory Exercises and Activities:
1. Analyzing the Properties of Representative Elements-The Chemical Properties of the Halogens- 2
periods (Microscale)
2. Teacher Demonstration- The Acidic/Basic Properties of Row Three Elements and Their Oxides -1
period (Hollenberg, Postma, and Roberts)
Gases
A.
Laboratory Exercises and Activities:
1. Construction of Molecular Shapes Using Models-1 period
2. Construction of Hybrid Orbitals Using Models -1 period
Objectives:
At the conclusion of this unit, the student will be able to:
1. Describe how the pressure of a gas is measured and define the units of pressure.
2. Use Boyle’s and Charles’s laws to quantitatively relate volume, pressure and temperature of a gas. (5.3)
3. Use Avogadro’s law to relate the number of moles of gas to the volume. (5.3)
4. Define molar volume of an ideal gas at STP.
5. Define the Ideal Gas Law.
6. Use the Ideal Gas Law to calculate the volume of a gas resulting from pressure changes, temperature
changes, and changes in the moles of gas. (5.3)
7. Use the Ideal Gas Law to calculate the density or molecular mass of a gas.(5.3)
8. Use the Ideal Gas Law and the principles of stoichiometry to determine the volume of gas involved in a
chemical reaction. (5.3)
9. Use Dalton’s Law of Partial Pressure to determine partial pressure or mole fraction of a gas in a
mixture. (5.3)
10. Use Graham’s Law to relate the molecular mass of a gas to its relative velocity.
11. Use the Kinetic Molecular Theory to explain the behavior of gases. (5.6)
12. Compare the behavior of a real gas to an ideal gas under certain conditions of pressure and temperature.
13. Describe how the van der Waals equation is used to predict the behavior of a real gas.
B.
XII
Liquids and Solids
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Use the particle model to explain the properties of solids and liquids. (5.6)
2. Differentiate between intermolecular forces of attraction: London dispersions, dipole-dipole forces and
hydrogen bonding.
3. Calculate the quantity of heat absorbed or released during a phase change. (5.7)
4. Define vapor pressure.
5. Describe the relationship between vapor pressure, volatility and boiling point of a liquid. (5.6)
6. Describe effect of intermolecular forces on physical properties such as viscosity, surface tension and
boiling point.. (5.6)
7. Draw a phase diagram given appropriate data.
8. Interpret a phase diagram. (5.6)
9. Determine the number of atoms per unit cell.
10. Given the type of unit cell for a solid, relate cell dimensions to the atomic radius and density. (5.3)
11. Distinguish between interparticle forces of attraction (ionic, covalent network, molecular, metallic
bonds) with regards to structure, holding forces, and the physical properties of a substance. (5.6)
12. Given the formula or the physical properties of a substance according to bond type. (5.6)
13. Predict the variation in the physical properties of a substance on the basis of bond type. (5.6)
B.
XIII
Laboratory Exercises and Activities:
1. “Determination of the Molar Volume of a Gas-Oxygen”- 2 periods (Hollenberg, Postma, and Roberts)
2. “Determination of the Molecular Mass of a Gas-Carbon Dioxide”-(Hollenberg, Postma, and Roberts)
3. Evaluation of R, the Universal Gas Constant Using Butane- 2 periods
(Flinn)
4. Teacher Demonstration- Graham’s Law of Diffusion- 1 period
Laboratory Exercises and Activities:
1. “Evaporation and Intermolecular Attractions” –The Determination of the Relative Strength of
Intermolecular Attractions Using the Rate of Evaporation-2 periods (Vernier)
2. The Packing of Particles in Crystals- Construction of Unit Cells and the Determination of the
Efficiency of Packing - 2 periods
Properties of Solutions
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Describe the solution process in terms of the energy and entropy changes that take place.
2. Predict the interaction between solute and solvent in the solution process. (5.6)
3. Define and compare ways of expressing concentration of solution: molarity, mass percent, mole
fraction and molality.
4. Calculate the molality and mole fraction given the composition of the solution. (5.3)
5. Distinguish between unsaturated, saturated and supersaturated solutions. (5.6)
6. Predict the effect that solute-solvent interaction, changes in temperature, and changes in pressure have
on the solubility.
7. Demonstrate how a solution’s vapor pressure is affected by the concentration and type of solute.
8. Use Raoult’s Law to determine the vapor pressure of the volatile components above a solution. (5.3)
9. Calculate the boiling point, freezing point and osmotic pressure of a solution given appropriate data.
(5.3)
10. Use osmotic pressure, freezing point depression, boiling point elevation and vapor pressure to determine
the molecular mass of the solute. (5.3)
11. Compare the colligative properties of electrolytes to those of non-electrolytes.
B.
Laboratory Exercises and Activities:
1. “The Determination of Molecular Mass of Benzoic Acid from Freezing Point Depression” -2 periods
(Vernier)
2. “Paper Chromatography: Separation of Cations”(Ni, Fe and Cu Ions)-2 periods (Kemp and Nelson)
XIV
Chemical Kinetics
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Define reaction rate and show how rate can be determined from experimental data. (5.3)
2. Determine the rate law for a reaction given appropriate data.(5.3)
3. Calculate the value of the specific rate constant and specify the units (5.3).
4. For first order reactions, calculate the concentration after a give time or the time required for the
concentration to drop to a certain level given appropriate data. (5.3)
5. Given data for concentration as a function of time, use graphical methods of analysis to determine the
reaction order with respect to a reactant. (5.3)
6. Define activation energy and describe its effect on the rate of reaction. (5.6)
7. Use the Arrhenius equation or graphical methods to determine the activation energy of a reaction. (5.3)
8. Use collision theory to discuss the role that temperature, catalyst and the nature of the reactants have on
the rate of reaction. (5.6)
9. Interpret energy diagrams in terms of the changes in energy that take place. (5.6)
10. Propose a mechanism for a reaction consistent with the rate expression and other experimental data.
11. Describe the role of a catalyst in a chemical reaction. (5.6)
12. Differentiate between a homogeneous and heterogeneous catalyst.
B.
XV
Chemical Equilibrium
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Describe how equilibrium is established in a reaction.
2. Define the Law of Mass Action.
3. Write the equilibrium expression for a given reaction.
4. Given equilibrium concentrations or partial pressures, calculate Kc or Kp. (5.3)
5. Given the magnitude of the equilibrium constant, predict the direction in which the system will move to
reach equilibrium.
6. Given the value of K and the initial concentrations, calculate equilibrium concentrations. (5.3)
7. Show how Kc and Kp are related. (5.3)
8. Use LeChatelier’s Principle to predict and explain the effect of changes in concentration, temperature
and pressure on the position of equilibrium.
9. Predict the effect of a catalyst on a system at equilibrium.
10. Write an equation for the equilibrium that exists in a saturated solution.
11. Write the Ksp expression for a saturated solution.
12. Calculate the value of K sp given solubility data.(5.3)
13. Use K sp to determine the molar solubility of a salt in water or in a solution that contains a common ion.
(5.3)
14. Use K sp to determine if a precipitate forms when two solutions are mixed.
B.
XVI
Laboratory Exercises and Activities:
1. Determination of the Order of the Reaction using Graphical Methods- 2 periods (Microscale)
2. Determination of the Activation Energy of a Clock Reaction-1 period
3. “Determining the Concentration (Ni +2) of a Solution: Beer’s Law”-2 periods (Vernier)
Laboratory Exercises and Activities:
1. Teacher Demonstration Chemical Equilibrium: “LeChatelier’s Principle” -1 period (Vonderbrink)
2. “The Determination of an Equilibrium Constant for the Formation of FeSCN+2-2 periods (Vernier)
3. “The Determination of the K sp of Ca(OH)2 Through Titration”-2 periods
Acid-Base Equilibria
A.
Objectives:
At the conclusion of this unit, the student will be able to:
1. Describe the Arrhenius theory of acid and base.
2. Define the dissociation constant for water Kw.
3. Describe the pH and the pOH scale.
4. Given one of the four quantities: [H+], [OH-1], pH or pOH, determine the other three. (5.3)
5. Classify species as acid or base using the Bronsted-Lowry theory.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
B.
XVII
A.
Identify a Bronsted-Lowry acid-base conjugate pair.
Differentiate between weak and strong acids; weak and strong bases.
Calculate the pH of a strong acid or base given the molar concentration. (5.3)
Calculate the ionization constant of a weak acid or base given the appropriate data. (5.3)
Describe a polyprotic acid and determine the pH of the solution.
Predict whether a given ion of a salt will react to form an acidic, basic or neutral solution.
Write the net ionic equation for the hydrolysis of a salt.
Determine the pH of a salt solution given the appropriate data. (5.3)
Show how bond strength and polarity affect the acidic-basic properties of ions or molecules.
Classify a species as an acid or base using the Lewis model.
Describe the characteristics and composition of a buffer solution.
Calculate the pH of a buffer solution. (5.3)
Determine the buffer capacity of a solution. (5.3)
Interpret the pH titration curves for an acid-base reaction.
Predict the acidic or basic properties of the solution formed at the end point of a titration.
Calculate the pH at any point in the titration process.(5.3)
Describe the composition of an indicator and explain how an indicator works.
Select the appropriate indicator for an acid-base titration.
Laboratory Exercises and Activities:
1. The Determination of the Ka of Acetic Acid-2 periods (Vonderbrink)
2. The Percent Hydrolysis of Salts-1 period
3. Teacher Demonstration-Construction of Titration Curves for Acid-Base Reactions-2 periods (Vernier)
4. The Determination of Buffer Capacity of the Acetic Acid/Sodium Acetate Buffer Using an Indicator-1
period (Microscale)
Electrochemistry
Objectives:
At the conclusion of this unit, the student will be able to:
1. Determine the oxidation state number of each atom in a compound.
2. Identify the oxidizing agent and the reducing agent in a redox reaction.
3. Balance redox reactions using the half-reaction method.
4. Balance redox reactions in acidic and basic solution.
5. Construct a voltaic cell, identify the oxidizing and reducing agent, label anode and cathode, and
indicate the flow of electrons and ions. (5.4)
6. Calculate the standard cell potential and use it to predict the direction of the reaction. (5.3)
7. Use the standard reduction potential to compare relative strength of oxidizing agents and reducing
agents.
8. Use the Nernst equation to calculate the cell potential at nonstandard state conditions.(5.3)
9. Use the Nernst equation to calculate the concentration of cell components given the voltage of the
cell. (5.3)
10. Use the Nernst equation to calculate the equilibrium constant for a redox reaction given the cell
potential. (5.3)
11. Calculate the change in free energy for an electrochemical process from the net cell potential. (5.3)
12. Differentiate between a voltaic cell and an electrolytic cell. (5.4)
13. Construct an electrolytic cell and label all parts. (5.4)
14. Predict the products that are form in an electrolytic cell.
15. Use Faraday’s law to determine the amount of product formed at the electrodes in an electrolytic cell.
(5.3)
B.
Laboratory Exercises and Activities:
1. Redox Titration- The Determination of the % of Hydrogen Peroxide in a Solution Using a Standard
Potassium Permanganate Solution-2 periods
2. Construction of Electrochemical Cells -2 periods (Hollenberg, Postma, and Roberts)
3. Electrolysis and Faraday’s Law-1 period (Microscale)
XVIII
A.
Nuclear Chemistry
Objectives:
At the conclusion of this unit, the student will be able to:
1. Classify types of radioactive decay.
2. Write balanced nuclear decay equations.
3. Establish the patterns of nuclear stability by relating stability to the neutron to proton ratio in the
isotope.
4. Predict the probable type of decay using “the band of stability.”
5. Describe the methods of induced nuclear transmutation.
6. Define half-life.
7. Use the first order rate law to determine the decay time, fraction that remains or the half-life, given
appropriate data. (5.3)
8. Explain how materials can be dated using radioactive decay.
9. Differentiate between nuclear fission and fusion and describe how both can be harnessed to produce
energy. (5.4)
10. Calculate the energy associated with nuclear transformations. (5.7)
11. Explain the biological effects of radiation. (5.10)
B.
XIX
A.
Laboratory Exercises and Activities:
1. Determination of the Half-Life of Ba-137-1 period
Chemistry of Coordination Compounds
Objectives:
At the conclusion of this unit, the student will be able to:
1. Describe the characteristics of transition metal cations that allow them to form complex ions.
2. Define the terms used in coordination chemistry.
3. Given the charge and composition of a complex ion, determine the charge and the coordination of the
central metal cation.
4. Describe the hybridization around the central metal cation.
5. Describe a chelating agent
6. Given the name or the formula of a complex ion or compound, determine the name or formula.
7. Identify and draw the various types of isomers that can be formed from the structure of a complex ion.
8. Explain how bonding theory accounts for the color and magnetism of a complex ion.
9. Write the net ionic equations for the reactions that involve complex ions.
10. Calculate the equilibrium concentrations of the species present given the formation constant for a
complex ion. (5.3)
11. Identify the amphoteric hydroxides.
B.
Laboratory Exercises and Activities:
1. Construction of Complex Ions Using Models -1 period
2. Teacher Demonstration-Reactions of Complex Ions-1 period (Hollenberg, Postma, and Roberts)
XX
AP Test Review
XXI
A.
B.
Organic Chemistry
Objectives:
At the conclusion of this unit, the student will be able to:
1. Predict the hybridization around the carbon atom in organic molecules.
2. Draw the structural formal of a hydrocarbon form the molecular formula.
3. Write the structural formula given the name.
4. Name the compound given the structural formula.
5. Draw all possible structural and geometric isomers given the formula or name.
6. Locate the chiral carbon in organic molecules in order to determine optical activity.
7. Predict the variation in physical properties associated with the isomeric
8. Classify organic compounds on the basis of their functional group.
9. Relate the physical properties of a molecule to the structure of the functional group. (5.6)
10. Identify common organic reactions and predict the products of the reactions.
11. Define a polymer and describe how polymers can be produced.
Laboratory Exercises and Activities:
1.
2.
3.
4.
V.
I
Construction of Hydrocarbons Using Molecular Models-1 period
Construction of Isomers Using Models-1 period
Construction of Functional Groups Using Models-1 period
Polymer Activity- The Exploration of the Physical Properties of Polymers-2 periods
XXII
A.
Qualitative Analysis
Objectives:
At the conclusion of this unit, the student will be able to:
1. Describe the method of selective precipitation to separate a mixture of ions in solution. (5.1)
2. Use a metallic ion’s ability to form a complex ion to change its solubility in solution.
3. Use the formation constant of a complex ion and the solubility product constant in order to determine if
a precipitate will dissolve.
4. Predict how the solubility of a substance is affected by pH.
5. Use laboratory data to develop a scheme of analysis to separate and test for the presence of certain ion
in a solution. (5.1)
6. After performing a laboratory analysis, state whether a certain ion is present or absent in an unknown
solution. (5.1)
7. Write net ionic equations for the reactions involved in qualitative analysis
B.
Laboratory Exercises and Activities:
1. “Qualitative Analysis of Silver, Mercury and Lead Ions”-3 periods (Hollenberg, Postma, and Roberts)
2. “Qualitative Analysis of Common Anions” -3 periods (Hollenberg, Postma, and Roberts)
3. “Qualitative Analysis of Aluminum, Iron and Zinc”-3 periods (Hollenberg, Postma, and Roberts)
ESSENTIAL QUESTIONS /SCOPE AND SEQUENCE
What is the nature of matter?
Matter and Measurement (Summer Assignment) (.5 week)
A. Classification and the Properties of Matter
B. Measurement: Units and Uncertainty
C. Dimensional Analysis
II
What is the meaning of the formula for an element or compound?
The Nature of Matter (Summer Assignment) (.5 week)
A. Atomic Theory
B. Periodic Table
C. Ionic Compounds and Binary Molecular Compounds
E. Acids
III
How can one describe the quantitative nature of a chemical formula?
What does a balanced equation reveal about the quantitative relationships between products and reactants in a
chemical reaction?
Stoichiometry (1.5 weeks)
A. Chemical Equations
B. The Mole
B. Quantitative Information from Chemical Formulas
1. Percentage Composition
2. Empirical Formulas and Molecular Formulas
C. Quantitative Information from Balanced Equations
1. Limiting Reagent
2. Theoretical Versus Actual Yield
IV
What characterizes reactions that occur in solution?
How is the composition of an aqueous solution determined?
Solution Chemistry (2 weeks)
A. Properties of Solutes in Aqueous Solutions
B. Metathesis Reactions
C. Redox Reactions
D. Net Ionic Equations
E. Molarity
F. Solution Stiochiometry and Titration
V
What is the relationship between energy changes and chemical reactions?
How can one determine in advance if a reaction occurs spontaneously?
Thermodynamics (2 weeks)
A. The Nature of Energy
B. First Law of Thermodynamics
C. Calorimetry
D. Hess’s Law
E. Enthalpy of Formation
F. Entropy and the Second Law of Thermodynamics
G. Determination of Entropy Change
H. Gibbs Free Energy
I. Free Energy and Temperature
J. Free Energy and the Equilibrium Constant
VII
What is the nature of electromagnetic radiation and it relationship to the atom?
How can one describe the probable energy states for an electron in an atom or ion?
Electronic Structure of Atoms (1.5 week)
A. Electromagnetic Radiation
B. Bright Line Spectra
C. Bohr Model for Hydrogen
D. Quantum Mechanical Model
E. Quantum Numbers
F. Electron Configuration
VIII
How is the periodic table used to organize chemical facts and predict chemical properties?
How do the properties of elements change as you move across a row or down a column on the periodic table?
Periodic Properties (.5 week)
A. Development of the Table
B. Electron Configuration and Position of the Table
C. Periodic Trends
1. Atomic/Ionic Radii
2. Ionization Energy
3. Electron Affinity
4. Metallic Character
D. Group Trends for Metals
E. Group Trends for Nonmetals
IX
What is the nature of a chemical bond?
What bonding models are useful in understanding the nature of ionic and covalent bonds?
Chemical Bonding (1.5 weeks)
A. Ionic Bonding
B. Covalent Bonding
C. Electronegativity and Bond Polarity
D. Lewis Structures
E. Resonance
F. Exceptions to the Octet Rule
G. Bond Energies
X
How can you predict the shape and polarity of a molecule?
Molecular Geometry and Bond Theory (1.5 weeks)
A. VSEPR Theory
B. Molecular Shape and Polarity
C. Hybridization
D. Multiple Bonds
E. Molecular Orbital Theory
XI
How can we predict the physical properties of gases quantitatively?
What do the physical properties of gases tell us about the behavior of molecules?
Gases (1.5 weeks)
A. Characteristics of Gases
B. Pressure
C. The Gas Laws: Boyle’s, Charles’s and Avogadro’s Laws
D. Ideal Gas Equation
E. Applications of Ideal Gas Equation
1. Gas Density and Molar Mass
2. Gas Stoichiometry
F. Dalton’s Law of Partial Pressure
G. Kinetic Molecular Theory
H. Graham’s Law
I. Real Gases
XII
What are the intermolecular forces of attraction that explain the physical properties of solids and liquids and the
phase changes that they undergo?
Liquids and Solids (2 weeks)
A. Intermolecular Forces of Attraction
B. Properties of Liquids
C. Phase Changes
D. Vapor Pressure
E. Phase Diagrams
F. Structure of Solids: Unit Cells
G. Bonding in Solids
XIII
What are forces involved in the interaction of a solute and solvent in a solution?
What are the colligative properties of a solution?
Properties of Solutions (2 weeks)
A. The Solution Process
B. Concentration Units: Mole Fraction and Molality
C. Factors Affecting Solubility
D. Colligative Properties
1. Vapor Pressure Lowering
2. Boiling Point Elevation and Freezing Point Depression
3. Osmotic Pressure
XIV
How can the speed of reaction be expressed?
What factors determine how rapidly a reaction occurs?
Chemical Kinetics (2 weeks)
A. Factors Affecting Rate
B.
C.
D.
E.
XV
Concentration and Rate: Orders of Reaction
Temperature and Rate
Reaction Mechanism
Catalysis
What is the nature of chemical equilibrium?
How can the equilibrium position of a reaction be expressed in quantitative terms?
What factors affect a system at equilibrium?
Chemical Equilibrium (2 weeks)
A. The Concept of Equilibrium
B. Equilibrium Expression
C. Equilibrium Constant: Kc and Kp
D. Applications of the Equilibrium Constant
E. Le Chatelier’s Principle
F. Solubility Equilibria: Ksp
G. Factors Affecting Solubility
H. Selective Precipitation of Ions
XVI
What are the properties that allow a substance to behave as an acid or base?
What characterizes the behavior of acids and bases in terms of the equilibrium in which they participate?
What are the common reactions of acids and bases?
Acid-Base Equilibria (2.5 weeks)
A. Dissociation Constant of Water, Kw
B. pH and pOH Scale
C. Bronsted-Lowry Acids and Bases
D. Acid-Base Strength
E. Ionization Constants, Ka and Kb
F. Hydrolysis
G. Acid-Base Behavior and Chemical Structure
H. Buffer Solutions
I. Acid-Base Titration and Titration Curves
XVII
How can chemical energy be converted to electrical energy in a electrochemical cell or battery?
How can electrical energy be used to make a nonspontaneous chemical process occur?
What are the products of a redox reaction?
Electrochemistry (2 weeks)
A. Oxidation States
B. Balancing Redox Reactions in Acidic or Basic Solutions
C. Voltaic Cells
D. Standard Cell Potential
F. Nernst Equation
G. Free Energy and Redox Reactions
H. Equilibrium Constants for Redox Reactions
I. Electrolysis and Faraday’s Law
XVIII
What characterizes nuclear reactions?
Nuclear Chemistry (1 week)
A. Types of Radioactive Decay
B. Patterns of Nuclear Stability
C. Nuclear Transmutation
D. Rates of Radioactive Decay
E. Energy Changes
F. Fission vs. Fusion
XIX
How does bonding theory explain the structure and properties of coordination compounds?
How does the formation of complex ions affect the reactivity and solubility of metallic ions?
Chemistry of Coordination Compounds (1 week)
A. Structure
B. Nomenclature
C. Isomerism
D. Reactions of Complex Ions/Amphoterism
XX
What are the fundamental concepts in chemistry?
AP Test Review (2 weeks)
XXI
How can bonding theory explain the existence of so many organic compounds?
What are the reactions that characterize organic chemistry?
Organic Chemistry (2 weeks)
A. Structure of Alkanes, Alkenes and Alkynes
B. Isomers
C. Nomenclature
D. Functional Groups
E. Reactions
F. Polymers
XXII
How can the principles of chemistry be used to analyze a solution in order to determine the identity of the ions it
contains.?
Qualitative Analysis (2 weeks)
A. Selective Precipitation of Ions
B. Formation of Complex Ions
C. Equilibrium Considerations
D. Amphoterism
E. Schemes of Analysis
F. Reactions and Equations
VI.
INSTRUCTIONAL STRATEGIES
The Advanced Placement Chemistry curriculum is centered on the major concepts in chemistry. Development of the
concepts takes place using strategies that include, but are not limited to the following:
A. Laboratory Activities, used to:
1. Demonstrate, apply and verify concepts
2. Evaluate models
3. Analyze data
4. Test hypothesis
B. Problem Solving Activities
1. Model building
2. Mathematical applications
3. Cooperative group work
4. Student discussion
C. Teacher-Guided Instruction
1. PowerPoint to introduce content
2 Questioning and discussion
3. Overhead transparencies
4. Handouts
5. Assignments
D. Internet used to research topics
E. Videodiscs
1. Chemistry, Coronet Videodisc
2. Chemistry at Work, Videodiscovery
3. Cosmic Chemistry, Optical Data Corporation
4. Physical Science, Optical Data Corporation
5. The World of Chemistry: Selected Demonstrations
F. Videos
1. Chem Study
2. Discovery
3. NOVA
4. The World of Chemistry: Annenberg Collection
VII. RESOURCES
Textbook:
Brown, Lemay and Bursten, Chemistry: The Central Science, 10th edition, Pearson Education, Inc. Upper
Saddle River, New Jersey, 2006.
Laboratory Resources:
There is no one standard lab manual for AP Chemistry. Experiments are taken from a variety of resources
and reflect the instructor’s teaching experience at the AP level. The experiments are all hands-on activities which
use both mircroscale and macroscale techniques. Each lab is adapted to meet the class time schedule, and
laboratory resources available. Currently, the high school laboratory has computers which enable the students to
use Vernier software and technology for data collection. Each student receives a lab sheet which contains an
introduction to the lab, the directions, safety concerns, a data sheet, as well as instructions for processing the data
and questions which analyze the lab results. Emphasis in placed on the evaluation of laboratory results with the
analysis of technique, comparison with the accepted values and the discussion of possible sources of error. Students
work in pairs in the laboratory. After completion, each student submits a laboratory report for evaluation. The
teacher places the graded lab report into a folder in order to create a portfolio of student laboratory work. The
portfolio is then given to the student upon successful completion of the course. The following resources are used as
a basis for the laboratory experiments, chemical reaction equations and chemical demonstrations.
Bilash, Gross and Koob, A Demo A Day, Flinn Scientific, Inc. Illinois, 1995.
Cesa, Flinn Scienfitic Chem Topic Labs, Flinn Scientific, Inc. Illinois, 2004.
Hague and Smith, The Ultimate Chemical Equations Handbook, Flinn Scientific, Inc. Illinois, 2001.
Holmquist and Volz, Chemistry with Computers, Vernier Software and Technology, Beaverton, Oregon,
2000.
Kemp and Nelson, Laboratory Experiments; Chemistry: The Central Science, Pearson Education, Inc.
Upper Saddle River, New Jersey, 2006.
Randall, Advanced Chemistry with Vernier, Vernier Software and Technology, Beaverton, Oregon, 2004.
Hollenberg, Postma, and Roberts, Chemistry in the Laboratory, 5th Edition, W. H Freeman and Company,
New York, 2000.
Vonderbrink, Sally Ann, Laboratory Experiments for Advanced Placement Chemistry, Flinn Scientific,
Illinois, 1995.
VIII. EVALUATION
Assessment may include:
A.
B.
C.
D.
E.
Laboratory Reports
Quizzes
Tests
Student Projects
Midterm and Final Exams