AP Chem Syllabus

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AP® Chemistry
Text
Chemistry: The Central Science by Brown, LeMay, and Bursten, 9th ed., Prentice Hall, 2003.
ISBN: 0-13-038168-3
Additional References
Chemistry & Chemical Reactivity by Kotz and Purcell, 2nd ed., Saunders College Publishing, a
division of Holt, Rinehart and Winston, Inc., 1991. ISBN: 0-03-072569-0
The Ultimate Chemical Equations Handbook by Hague and Smith, Flinn Scientific, Inc., 2001.
ISBN: 1-877991-63-5
Laboratory Experiments for Advanced Placement Chemistry by Vonderbrink, Flinn Scientific,
Inc.,1995. ISBN: 1-877991-34-1
Laboratory Experiments for Advanced Placement Chemistry by Vonderbrink, 2nd ed., Flinn
Scientific, Inc.,2006. ISBN: 978-1-933709-03-1
Inquiry-Based Experiments in Chemistry by Lechtanski, American Chemical Society, 2000.
ISBN: 0-8412-3570-8
Molecular Origami by Hanson, University Science Books, 1995. ISBN: 0-935702-30-X
Overview
AP Chemistry meets daily for one 51-minute period. Students complete 3-5 problem sets, 2-3
quizzes, and a test for each chapter. A cumulative exam is also given at the end of each quarter.
The equivalent of one double-period per week is spent engaged in laboratory work resulting in
1 – 3 labs per chapter. Students work in pairs to record observations and data, perform analysis
calculations, and answer conclusion questions. Lab assessment is based on either formal or
abbreviated lab reports that are kept as a portfolio.
This is a second-year course so students have a strong foundation in:
 atomic theory
 atomic structure
 periodicity
 chemical bonding
 basic chemical reactions
 stoichiometry
 kinetic molecular theory
 gas laws
1
The text book used in the first-year course is Chemistry by Wilbraham, Staley, Matta, and
Waterman, Pearson Education Inc., publishing as Pearson Prentice Hall, 2008,
ISBN: 0-13-251211-4. Modern Chemistry by Davis, Metcalfe, Williams, and Castka, Holt,
Rinehart, and Winston, 1999, ISBN 0-03-051389-8, is also used as a reference. Additional
references are the ChemTopic™ books published by Flinn Scientific, Inc.
During the first-year course, the equivalent of one double-period per week is spent engaged in
laboratory work resulting in 1 – 3 labs per chapter. Students work in pairs to record observations
and data, perform analysis calculations, and answer conclusion questions. Lab assessment is
based on either formal or abbreviated lab reports that are kept as a portfolio.
First-Year Curriculum
Chapter 3 – Scientific Measurement (3 weeks)
The student will:
1. distinguish between qualitative and quantitative observations.
2.
identify the 7 SI base units and quantities.
3.
perform density calculations.
4.
make conversions between metric measurements by moving the decimal.
5.
make temperature conversions between the Kelvin and Celcius scales.
6.
distinguish between accuracy and precision.
7.
perform percent error calculations.
8.
determine the number of significant figures in measurements.
9.
perform mathematical operations involving significant figures.
10. convert measurements into scientific notation.
11. distinguish between inversely and directly proportional relationships.
Laboratories:
 Is Density an Intensive or Extensive Property? – An Inquiry Lab
 Determining the Densities of Pre- and Post-1982 Pennies
Chapter 2 – Matter and Change (2.5 weeks)
The student will:
1.
classify properties of matter as physical or chemical and as extensive or
intensive.
2.
classify changes of matter as physical or chemical.
2
3.
identify the reactants and products of a chemical reaction.
4.
identify four possible clues that a chemical change has taken place.
5.
explain the relationship between phase changes and energy.
6.
categorize phase changes as endothermic or exothermic.
7.
classify matter as an element, compound, homogeneous mixture, or
heterogeneous mixture.
Laboratories:
 Observation and Experimentation: Reaction of CaCl2, NaHCO3, and Phenol
Red
 Investigating Physical and Chemical Changes
 Separation of a Mixture
Chapter 4 – Atomic Structure (1.5 weeks)
The student will
1.
explain and apply the law of conservation of mass.
2.
recognize examples of the law of definite proportions and the law of
multiple proportions.
3.
summarize the essential points of Dalton’s atomic theory.
4.
summarize the experiments done with cathode rays and tell what was
concluded from these experiments.
5.
summarize the experiment carried out by Ernest Rutherford and tell what
was concluded from this experiment.
6.
identify and describe the three major subatomic particles.
7.
define the words atom and isotope.
8.
explain what is meant by atomic number and mass number, and describe
how they apply to isotopes.
9.
given the identity of a nuclide, determine its number of protons, neutrons,
and electrons.
10.
describe the relative scale used to determine atomic mass.
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11.
calculate the average atomic mass of an element using the mass of its
isotopes and their relative abundances.
12.
use a periodic table to write symbols and names of elements.
13.
use a periodic table to locate the group and period of elements.
14.
list the characteristics that distinguish metals, nonmetals, and metalloids.
Laboratory:
 Calculating Weighted Atomic Mass Using Beans
Chapter 5 – Electrons in Atoms (3 weeks)
The student will:
1.
explain the dual wave-particle nature of light.
2.
use the mathematical relationships between the velocity, wavelength,
frequency, and energy of electromagnetic radiation.
3.
describe how the line emission-spectrum of hydrogen contributed to
Bohr’s atomic model.
4.
describe the Bohr model of the hydrogen atom.
5.
compare and contrast the Bohr and the quantum models.
6.
list the four quantum numbers and explain what they represent.
7.
relate the number of orbitals and electrons in an energy level to the
principal quantum number.
8.
write orbital notations and electron configurations for atoms applying the
Aufbau principle, Pauli exclusion principle, and Hund’s rule.
9.
trace the evolution of the atomic model from Dalton’s model to the
quantum model recognizing the major developments of each model.
Laboratory:
 Flame Tests and Observation of Line Emission Spectra
Chapter 6 – The Periodic Table (3 weeks)
The student will:
1. explain the roles of Mendeleev and Moseley in the development of the
periodic table.
4
2.
state the modern periodic law.
3.
describe the relationship between the number of valence electrons and the
group numbers.
4.
locate and describe the general properties of the alkali metals, alkaline
earth metals, the halogens, and the noble gases.
5.
define atomic and ionic radius, ionization energy, and electronegativity.
6.
state, apply, and explain the periodic trends of atomic radius, ionization
energy, and electronegativity.
7.
correlate the pattern of reactivity of elements to the patterns of atomic
radius, ionization energy and electronegativity.
Laboratories:
 Reactivity of the Halogens
 Reactivity of the Alkaline Earth Metals
Chapters 7 and 8 – Ionic, Covalent, and Metallic Bonding (2 weeks)
The student will:
1. define the following terms: chemical bond, molecule, multiple bond.
2.
explain why atoms bond to other atoms.
3.
describe ionic and covalent bonding.
4.
explain the difference between polar covalent bonds and nonpolar
covalent bonds.
5.
classify bonding type according to electronegativity difference or by the
type of atoms bonded.
6.
state the octet rule.
7.
write dot diagrams for main group elements and their ions.
8.
draw Lewis structures for molecules.
9.
describe how ions are arranged in a crystal.
10. use the VSEPR theory to predict the geometry of simple molecules.
5
Laboratories:
 Properties of Solids: Structure and Bonding
 Investigation of Chemical Bonds
Chapter 9 – Chemical Names and Formulas (2 weeks)
The student will:
1. list, describe, and distinguish three kinds of chemical formulas.
2. distinguish between ionic and molecular compounds based on names and
formulas.
3. determine the charge of monatomic ions from the periodic table.
4. write formulas for ionic and molecular compounds from their names.
5. write names for ionic and molecular compounds from their formulas.
Laboratory:
 Naming Ionic Compounds: Cu+ or Cu2+, Fe2+ or Fe3+
Chapter 11 – Chemical Reactions (3 weeks)
The student will:
1. identify the reactants and products of a chemical reaction.
2. write formula equations from word equations.
3. balance chemical equations and explain how this satisfies the law of
conservation .
4. include energy as a reactant or product of a reaction.
5. recognize the symbols used to represent the phase of reactants or products.
6. identify and classify reactions as either synthesis, decomposition, single
replacement, or double replacement.
7. explain what a catalyst is.
8. describe the four types of reactions and predict the products of these types of
reactions using the activity series and solubility table.
9. define spectator ions and write net ionic equations.
Laboratory:
 Investigating Four Types of Chemical Reactions
6
Chapter 10 – Chemical Quantities (3 weeks)
The student will:
1.
determine or calculate the atomic/molar mass of an element and the
formula mass/molar mass of a compound.
2.
use conversion factors to change between mass and moles.
3.
calculate the percentage composition of a compound from the formula or
experimental data.
4.
calculate the empirical formula from mass data and percentage
composition.
5.
determine the molecular formula from the empirical formula and molar
mass.
Laboratories:
 What is the Mass Percent of Water in a Hydrate? – An Inquiry Lab
 The Empirical Formula of Magnesium Oxide
Chapter 12 – Stoichiometry (3 weeks)
The student will:
1.
describe the importance of the mole ratio in stoichiometric calculations.
2.
solve mole-mole problems using the mole ratio indicated by the
coefficients of a balanced chemical equation.
3.
solve mass-mole and mole-mass problems.
4.
solve mass-mass problems.
5.
distinguish between the limiting reactant and the excess reactant.
6.
determine the limiting reactant and excess reactant of a reaction.
7.
calculate the amount of product formed by the limiting reactant.
8.
calculate the amount of excess reactant leftover.
9.
distinguish between theoretical yield, actual yield, and percent yield.
10.
calculate the percent yield of a reaction.
Laboratories:
 Determining Percent Yield of a Chemical Reaction
7
Chapter 14 – The Behavior of Gases (3 weeks)
The student will:
1.
state the 5 assumptions of the kinetic theory.
2.
distinguish between an ideal gas and a real gas.
3.
explain how the kinetic theory accounts for the physical properties of
gases.
4.
state, explain, and apply Boyle’s law.
5.
state, explain, and apply Charles’ law.
6.
state, explain, and apply Gay-Lussac’s law.
7.
use the combined gas law to solve problems.
8.
state, explain, and apply Dalton’s law of partial pressure.
9.
use the ideal gas law to solve problems.
10.
explain Graham’s law of effusion and use it to solve problems.
Laboratories:
 Determining the Molar Volume of a Gas and The Ideal Gas Constant
Second-Year Curriculum
Chapter 13 – Properties of Solutions (5 weeks)
The student will:
1.
name and define the two parts of a solution.
2.
describe the factors affecting solubility.
3.
describe the factors affecting the rate of solution.
4.
explain the energy changes that occur during the solution process.
5.
interpret solubility curves.
6.
explain the difference between unsaturated, saturated, and supersaturated
solutions.
7.
make calculations involving the concentration of a solution using molarity,
molality, mole fraction, mass fraction (mass percentage), ppm, and ppb.
8
8.
make calculations to convert from one concentration unit to another
concentration unit.
9.
explain what colligative properties are and give two examples.
10.
make calculations involving Raoult’s law, freezing point depression, and
boiling point elevation.
Laboratories:
 Solubility of Polar and Non-polar Solutes
 Determining the Molarity of a Saturated Solution – An Inquiry Lab
 Determining Molar Mass by Freezing Point Depression of Butylated
Hydroxytoluene, BHT
Chapter 15 – Chemical Equilibrium (2 weeks)
The student will:
1. define chemical equilibrium and reversible reactions.
2.
explain why chemical equilibria are dynamic.
3.
write equilibrium constant expressions and calculate Keq values.
4.
determine Keq values for equations that have been manipulated.
5.
calculate Kp from Kc or vice versa.
6.
compare the extent of reactions bases on Keq values.
7.
calculate Q values and determine which reaction needs to be favored to reach
equilibrium.
8.
calculate Keq values or equilibrium concentrations when the other is known.
9.
state Le Chatelier’s Principle.
10. predict the reaction favored when temperature, pressure, or concentrations
are changed.
Laboratories:
 Determination of Equilibrium Constant for the Formation of FeSCN2+
 Equilibrium and Le Chatelier’s Principle
Chapter 16 – Acid – Base Equilibria (4 weeks)
The student will:
1. give Arrhenius definitions of acids and bases.
9
2. give Bronsted-Lowry definitions of acids and bases.
3. give Lewis definitions of acids and bases.
4. distinguish between dissociation and ionization.
5. identify the hydronium ion and describe how it is formed.
6. identify and label conjugate pairs and write conjugate partners for acids and
bases.
7. explain what monoprotic, polyprotic, amphoteric, and amphiprotic mean.
8. explain the difference between weak and strong acids and bases.
9. calculate pH, pOH, [H3O+], or [OH-] for weak or strong acids and bases.
10. predict the color an indicator will appear when an acid or base is added to it.
11. make calculations concerning the titration of strong acids and strong bases.
Laboratories:
 pH of Common Liquids
 Determiniation of the Dissociation Constants of Weak Acids
 Acid/Base Titration
Chapter 17 – Additional Aspects of Aqueous Equilibria (3 weeks)
The student will:
1. define solubility.
2. explain what precipitation reactions are.
3. write Ksp expressions for insoluble compounds.
4. calculate Ksp from solubilities.
5. calculate solubilities from Ksp values.
6. compare solubilities using Ksp values.
7. calculate Q values and predict if precipitation will occur.
8. calculate ion concentrations necessary for precipitation.
10
Laboratory:
 Determination of the Hardness of Water
Chapter 21 – Nuclear Chemistry (2 weeks)
The student will:
1. define radioactivity.
2. name and describe three types of radioactivity.
3. discuss some of the uses of radioactivity.
4. write balanced nuclear reactions for the various types of radioactivity.
5. explain fission, how it is used in nuclear reactors and some of the advantages
and disadvantages.
6. explain nuclear fusion and some of the advantages and disadvantages.
Laboratory:
 Simulation of Radioactive Decay
Chapters 8 & 9 – Basic Concepts of Chemical Bonding and Molecular Geometry (3 weeks)
The student will:
1. explain the role of valence electrons in chemical bonding and determine the
number of valence electrons for atoms and molecules.
2. explain why atoms bond and describe the two types of chemical bonds.
3. describe the difference between polar covalent and nonpolar covalent
bonding.
4. draw Lewis structures for molecules using the octet rule and use these Lewis
structures to determine the number of sigma and pi bonds in a molecule.
5. draw resonance structures.
6. determine the bond order of a bond.
7. use information about bond length and dissociation energy to make judgments
about the properties of compounds.
8. use a table of bond energies to determine the enthalpy change for a reaction.
9. predict the shape of molecules using the VSEPR theory.
10. predict the polarity of a molecule.
11
11. determine the hybridization of an atom in a molecule.
Laboratory:
 Origami Molecular Models
 Molar Mass of a Volatile Liquid
Chapter 11 – Intermolcular Forces, Liquids, and Solids (2.5 weeks)
The student will:
1. compare gases to liquids and solids in terms of the Kinetic Molecular Theory.
2. explain what intermolecular forces are.
3. explain four types of intermolecular forces and compare them by strength.
4. identify the type of force acting on various molecules.
5. state and explain the relationship between the strength of intermolecular forces
and the physical properties of liquids.
6. use the Clausius-Clapeyron equation to solve problems.
7. define lattice energy and explain how it affects the physical properties of
solids.
8. describe the structure of network covalent solids and give an example.
9. describe the structure of amorphus solids and give an example.
10.explain how hydrogen bonding is responsible for the unique properties of
water.
11.explain why temperature does not change during phase changes.
12.interpret information from a phase diagram.
13.define triple point.
Laboratories:
 Vapor Pressure and Enthalpy of Vaporization of Water
 Expansion of Water as It Freezes – An Inquiry Lab
Chapter 5 – Thermochemistry (3 weeks)
The student will:
1. define energy and list several forms of energy.
12
2. define thermochemistry.
3. define specific heat capacity, molar heat capacity, and solve related problems.
4. explain why phase changes are not accompanied by temperature changes and
solve related problems.
5. define change in enthalpy and classify a process as endothermic or exothermic
by knowing the sign of H.
6. state Hess’s Law and use it to determine the H for a reaction.
7. use tables of Hof values to calculate Horxn.
8. make calculations involving calorimetry.
Laboratories:
 Determination of Enthalpy of Solution Using Calorimetry
 Thermochemistry and Hess’s Law
Chapter 19 – Chemical Thermodynamics (1.5 weeks)
The student will:
1. define spontaneous and nonspontaneous reactions.
2. list two examples of spontaneous processes that are endothermic.
3. define entropy.
4. state the third law of thermodynamics.
5. calculate S for phase changes and other reactions.
6. determine the sign of S for processes or molecules.
7. state and explain the second law of thermodynamics.
8. use the Gibbs free energy equation to predict reaction spontaneity.
9. calculate Gorxn from Gof values.
10. distinguish between enthalpy and entropy driven reactions from the sign of
H and S.
11. determine the temperature at which a reaction changes spontaneity.
12. make calculations using the relationship between K and Go.
Chapter 14 – Chemical Kinetics (4 weeks)
The student will:
1. define reaction mechanism and chemical kinetics.
13
2. list four factors that affect the speed of a chemical process and tell how they
affect it.
3. define reaction rate in terms of concentration and time.
4. calculate the rate of a reaction during a particular time.
5. state what rate expressions show and write rate expressions for reactions.
6. state the meaning of the rate constant and reaction order and determine each
using experimental data and the rate equation.
7. make calculations involving concentration and time for first and second order
reactions.
8. calculate half-life or rate constant when one is known.
9. determine reaction orders and rate constants graphically.
10. define activated complex and activation energy.
11. compare the activation energy and energy released as products are formed for
endothermic and exothermic reactions.
12. explain the collision theory.
13. solve problems using the Arrhenius equation.
14. determine the order and rate expression from the reaction mechanism.
15. explain how catalysts increase the rate of reactions.
Laboratories:
 The Effect of Concentration and Temperature on the Rate of the Reaction
between S2O32- and H3O+
Chapter 20 – Electrochemistry (3 weeks)
The student will:
1. define oxidation and reduction
2. assign oxidation numbers
3. identify species that are oxidized, reduced, oxidizing agents, or reducing
agents
4. balance redox reactions
14
5. label the anode and cathode of a voltaic cell and show the movement of
electrons and ions of the salt bridge when given the cell reaction
o
6. define cell potential, E , and describe its relationship to spontaneity including
o
o
G = -nFE
7. design an electrochemical cell and calculate E
using an
o
E red
o
tot
given the species involved
table
8. use the Nernst equation to calculate cell voltage at nonstandard conditions
9. determine K from E
o
Laboratory:
 Electrochemical Cells
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