AP Chemistry Course Syllabus

Textbook(s):

Brown, Theodore E., H. Eugene LeMay, and Bruce E. Bursten. Chemistry: The Central Science, 10^th Edition.
 Upper Saddle River, NJ: Prentice Hall.

Nelson, John H., Kenneth C. Kemp. Laboratory Experiments for Chemistry: The Central Science, 10^th Edition. Upper Saddle River, NJ: Prentice Hall.

Randall, Jack, Advanced Chemistry with Vernier, Beaverton, OR: Vernier.

Additional Materials:

Powerpoint Slides from the Instructor Resource Center accompanying Chemistry: The central Science.

Vernier LabPro and LoggerPro software, with numerous probes listed below.

Primary Source Readings:

Robert Boyle: Doubting the Four Elements (1661)

Marie Curie: Obtaining Radium (1923)

Herbert Butterfield: Chemistry Transformed (1965)

Isaac Asimov: Death in the Laboratory (1965)

Primo Levi: Carbon (1975)

Class Time:

M, W, F:	1:50-3:25 pm (Most formal labs will occur during these periods)
T, Th:	1:50-2:35 pm
*We will utilize 90 minutes per week, on average, for lab activities.

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Aims of Course:

AP Chemistry is an exciting, rigorous secondary Chemistry course. This course provides students a platform to develop an awareness, appreciation and understanding of the natural world that surrounds them. This course includes many of the key aspects of college level Chemistry, including a strong lab component.  The course is designed for motivated students who have an interest in knowing more about and pursuing post-secondary degrees in the sciences. Concepts covered in AP Chemistry will serve to prepare you for the AP Chemistry Exam given in the spring of the year.  Students are expected to take this exam and I will help you prepare for it.

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Objectives:

• Students mastering the material of AP Chemistry at our school will have a sufficient understanding and recall of facts, nomenclature, and concepts of chemistry to be successful on the AP Chemistry Exam.
• Students will demonstrate the ability to answer relevant questions in the format of those on the AP exam.
• Students will demonstrate basic problem-solving skills required for predicting chemical behavior including kinetics, equilibrium phenomena, and thermodynamics.
• Students will demonstrate competence in basic chemistry laboratory techniques and safety.
• Students will demonstrate competence in planning a laboratory procedure utilizing standard operational procedures and appropriate safety measures.
• Students will demonstrate competence in keeping a laboratory research notebook.

· In each laboratory experiment, students will physically manipulate equipment and materials in order to make relevant observations and collect data; use the collected data to form conclusions and verify hypotheses; and communicate and compare results and procedures (informally to other experimenters, and also in a formal, written report to the teacher)

• Students will demonstrate competence in writing research laboratory reports.

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Assessment Items:

Students should expect approximately 15-20 homework problems assigned and reviewed per chapter. The assignments will be either written from text or online from my website. Chapter exams consist of 20-30 multiple-choice questions and up to 4 free-response questions. These final four questions typically come from retired AP Released Exams. During the final quarter, I grade the completion of several practice AP Exams.

A student’s grade is a weighted average of the following:

· Tests 50%

· Labs 25%

· Quizzes 15%

· Homework 10%

Equipment and Stationery:

* This class is heavily supplemented with handouts in an outline form. Students will need a three-ring binder for these lecture notes and handouts. A divider may be used to separate lab materials from course notes. Powerpoint Slides with presentation material and practice problems will be used for lecture on a daily basis, and these slides will be available for handout periodically.

* Graph paper, pens, pencils, ruler, TI-83+, etc. are required on a daily basis. Do not come to class unprepared. Additional materials will be required as announced.

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Expectations:

Students are expected to keep a complete set of detailed notes based on class lectures, textbook and supplementary reading assignments. You are responsible for learning all reading topics covered in the textbook, and all material covered in class. Students are expected to have pertinent readings completed by the beginning of class. When specific homework is not assigned, students are still expected to average 45 minutes/day reading the upcoming material. Late assignments will be penalized or not accepted. Students are expected to arrive at class promptly. Do NOT miss a scheduled lab.

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Test Preparation:

Adequate notice (usually a minimum of 1 week) will be given for a change in the testing schedule. Pop quizzes will occur from time to time to ensure that the assigned readings are being completed. Students are not required to make up missed pop quizzes, but are required to make up announced quizzes. Be sure to regularly check the calendar for schedule.

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Extra Help:

I have an open door policy. Extra help will be available to any student upon request. I will be available from 3:45-4 on specific days to answer questions and provide assistance outside of class. Use our blog to post questions for the class and I am available by email. Please take advantage of this as required.

General Course Content

As required by the College Board, the following objectives are emphasized. The degree to which they are stressed corresponds to the level of assessment indicated in parentheses.

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I. Course Objective 1 : Structure of Matter (20%) These parenthetical percentages represent the amount of material on this topic that is on the actual AP exam

A. Atomic theory and atomic structure

1. Evidence for the atomic theory

2. Atomic masses; determination by chemical and physical means

3. Atomic number and mass number; isotopes

4. Electron energy levels: atomic spectra, quantum numbers, atomic orbitals

5. Periodic relationships including, for example, atomic radii, ionization energies, electron affinities, oxidation states

B. Chemical bonding

1. Binding forces

a. Types: ionic, covalent, metallic, hydrogen bonding, van der Waals (including

London dispersion forces)

b. Relationship to states, structure, and properties of matter

c. Polarity of bonds, electronegativities

2. Molecular models

a. Lewis structures

b. Valence bond: hybridization of orbitals, resonance, sigma and pi bonds

c. VSEPR

3. Geometry of molecules and ions, structural isomerism of simple organic molecules and coordination complexes; dipole moments of molecules; relation of properties to structure

C. Nuclear chemistry: nuclear equations, half-lives, and radioactivity; chemical applications

II. Course Objective 2 : States of Matter (20%)

A. Gases

1. Laws of ideal gases

a. Equation of state for an ideal gas

b. Partial pressures

2. Kinetic-molecular theory

a. Interpretation of ideal gas laws on the basis of this theory

b. Avogadro's hypothesis and the mole concept

c. Dependence of kinetic energy of molecules on temperature

d. Deviations from ideal gas laws

B. Liquids and solids

1. Liquids and solids from the kinetic-molecular viewpoint

2. Phase diagrams of one-component systems

3. Changes of state, including critical points and triple points

4. Structure of solids; lattice energies

C. Solutions

1. Types of solutions and factors affecting solubility

2. Methods of expressing concentration

3. Raoult's law and colligative properties (nonvolatile solutes); osmosis

4. Non-ideal behavior (qualitative aspects)

III. Course Objective 3 : Reactions (35-40%)

A. Reaction types

1. Acid-base reactions; concepts of Arrhenius, Bronsted-Lowry, and Lewis; coordination complexes; amphoterism

2. Precipitation reactions

3. Oxidation-reduction reactions

a. Oxidation number

b. The role of the electron in oxidation-reduction

c. Electrochemistry: electrolytic and galvanic cells; Faraday's laws; standard halfcell potentials; Nernst equation; prediction of the direction of redox reactions

B. Stoichiometry

1. Ionic and molecular species present in chemical systems: net ionic equations

2. Balancing of equations including those for redox reactions

3. Mass and volume relations with emphasis on the mole concept, including empirical formulas and limiting reactants

C. Equilibrium

1. Concept of dynamic equilibrium, physical and chemical; LeChatelier's principle; equilibrium constants

2. Quantitative treatment

a. Equilibrium constants for gaseous reaction: Kc, Kp

b. Equilibrium constants for reactions in solution

i. Constants for acids and bases; pK; pH

ii. Solubility product constants and their application to precipitation and the dissolution of slightly soluble compounds

iii. Common ion effect; buffers; hydrolysis

D. Kinetics

1. Concept of rate of reaction

2. Use of experimental data and graphical analysis to determine reactant order, rate constants, and rate laws

3. Effect of temperature change on rates

4. Energy of activation; the role of catalysts

5. The relationship between the rate-determining step and a mechanism

E. Thermodynamics

1. State functions

2. First law: change in enthalpy; heat of formation; heat of reaction; Hess's law; heats of vaporization and fusion; calorimetry

3. Second law: entropy; free energy of formation; free energy of reaction; dependence of change in free energy on enthalpy and entropy changes

4. Relationship of change in free energy to equilibrium constants and electrode potentials

IV. Course Objective 4 : Descriptive Chemistry (10-15%)

A. Chemical reactivity and products of chemical reactions

B. Relationships in the periodic table

C. Introduction to organic chemistry: hydrocarbons and functional groups (structure, nomenclature, chemical properties)

V. Course Objective 5 : Laboratory (5-10%) The AP Chemistry Examination includes some questions based on experiences and skills students acquire in the laboratory:

A. making observations of chemical reactions and substances

B. recording data

C. calculating and interpreting results based on the quantitative data obtained

D. communicating effectively the results of experimental work

VI. Course Objective 6 : Emphasis on chemical calculations and mathematical formulation of principles

These skills are integral to every unit we study, and practiced in assignments in each unit; they are also an active part of each laboratory investigation.

VII. Course Objective 7 : Lab Component comparable to a college-level introductory Chemistry Laboratory.

We spend two full periods weekly, on average, doing hands-on laboratory investigations. Students will gain practice in the manipulation of lab and computer equipment in order to make observations and collect data. These observations will be used to test hypotheses and make conclusions. The format for lab reports is outlined below.

Calendar Overview

Beginning	Week #	General Topic	Assignments Due
8/27	Week 1:	The Fundamentals Chapter 1: Introduction: Matter and Measurement Chapter 2: Atoms, Molecules, and Ions	Read Chp. 1-2, complete problem set
9/4	Week 2:	Chemical Reactions Chapter 3: Stoichiometry: Calculations with Chemical Formulas and Equations; Sections: 3.1–3.2 Chapter 4: Aqueous Reactions and Solution Stoichiometry; Sections: 4.1–4.4	Read Chp. 3.1-3.2 and 4.1-4.4, complete problem set Lab Report 1 Due
9/10	Weeks 3–4:	Stoichiometry Chapter 3: Stoichiometry: Calculations with Chemical Formulas and Equations; Sections: 3.3–3.7 Chapter 4: Aqueous Reactions and Solution Stoichiometry; Sections: 4.5–4.6	Read Chp. 3.3-3.7 and 4.5-4.6, complete problem set Lab 8 Report Due Lab 4 Report Due
9/24	Week 5:	Thermochemistry EXAM 1: Chp. 1-4 Chapter 5: Thermochemistry	Read Chp. 5, complete problem set Lab 6 Report Due
10/1	Weeks 6–7:	Electron Structure and Periodicity Chapter 6: Electron Structure of Atoms Chapter 7: Periodic Properties of the Elements	Read Chp. 6-7, complete problem set Lab 9 Report Due
10/15	Weeks 8–10:	Chemical Bonding and Molecular Geometry EXAM 2: Chp. 5-7 Chapter 8: Basic Concepts of Chemical Bonding Chapter 9: Molecular Geometry and Bonding Theories	Read Chp. 8-9, complete problem set Lab Assignment Due
11/12	Week 11:	Gases Chapter 10: Gases EXAM 3: Chp. 8-10	Read Chp. 10, complete problem set Lab 16 Report Due
11/26	Weeks 12–13:	Solids, Liquids, Changes in Phase, and Intermolecular Forces Chapter 11: Intermolecular Forces, Liquids, and Solids	Read Chp. 11 Lab 5 Report Due
12/10	Weeks 14–15:	Solution Properties Chapter 13: Properties of Solutions	Read Chp. 13 Lab 15 Report Due
1/7	Weeks 16–17:	Chemical Kinetics EXAM 4: Chp. 11,13 Chapter 14: Chemical Kinetics	Read Chp. 14 Lab 17 Report Due
1/28	Weeks 18–19:	Equilibrium Chapter 15: Chemical Equilibrium	Read Chp. 15 Lab 12 Report Due
2/11	Weeks 20–22:	Acid-Base Reactions and Solution Equilibria EXAM 5: Chp. 14-15 Chapter 16: Acid-Base Equilibria Chapter 17: Additional Aspects of Aqueous Equilibria	Read Chp. 16-17 Lab 10 Report Due Lab 7 Report Due Lab 11 Report Due
3/3	Weeks 23–24:	Chemical Thermodynamics EXAM 6: Chp. 16-17 Chapter 19: Chemical Thermodynamics	Read Chp. 19 Lab 19 Report Due
4/1	Weeks 25–26:	Redox Reactions and Electrochemistry Chapter 20: Electrochemistry	Read Chp. 20 Lab 13 Report Due
4/14	Weeks 27–28:	Descriptive, Organic, and Nuclear Chemistry	Assigned Reading Lab 20,21 Reports Due
4/28	Weeks 29–30:	EXAM 7: Chp. 19-20 Review and AP Exam Practice	Lab 22 Report Due

Detailed Scope and Sequence

By Week:	Objectives	Lab
Week 1:	The Fundamentals 1. Distinguish between physical and chemical properties and changes. 2. Understand the difference between elements, compounds, and mixtures. 3. Be familiar with the units of the metric system of measurement and the temperature scales. 4. Be able to convert measurements, especially within the metric system, by using dimensional analysis. 5. Determine the number of significant figures in a measurement and be able to express the results of a calculation with the proper number of significant figures. 6. Distinguish between protons, neutrons, and electrons, and be able to describe the composition of an atom of any particular element in terms of these subatomic particles. 7. Describe the basic anatomy of an atom and the ratio of the diameter of the nucleus to that of the atom. 8. Know the difference between an atom, an ion, and a molecule. 9. Have a basic knowledge of the periodic table, which includes being able to predict whether an element is a metal or a nonmetal, and what will be the probable charge of its ion. 10. Distinguish between empirical, molecular, and structural formulas. 11. Be able to write the correct name of an inorganic compound from its formula and vice versa. 12. Define hydrocarbon, alkane, and alcohol and be able to write the name from the formula and vice versa for simple alkanes and alcohols.	1,2,3,18
HW Problems	Exercises (pp. 30-35): 1.1, 1.3, 1.5, 1.9, 1.11, 1.13, 1.19, 1.23, 1.25, 1.27, 1.29, 1.33, 1.35, 1.37, 1.41, 1.45, 1.68, 1.71, 1.79 and Exercises (pp. 70-76): 2.1, 2.3, 2.5, 2.9, 2.17, 2.21, 2.29, 2.37, 2.39, 2.47, 2.49, 2.53, 2.59, 2.61, 2.63, 2.65, 2.81
Week 2:	Chemical Reactions [C3a] 1. Be able to balance chemical equations. 2. Write balanced chemical equations from word descriptions. 3. Predict the products of reactions based on the types presented, including combustion of compounds of C, H, and O. 4. Complete and balance these reactions. 5. Predict to some extent whether a substance will be a strong electrolyte, weak electrolyte, or nonelectrolyte. 6. Predict the ions that an electrolyte dissociates into. 7. Identify substances as acids, bases, and salts. 8. Using solubility rules, predict if a precipitate forms in a metathesis reaction, and thus predict its products and write a balanced equation. 9. Predict the products and write a balanced chemical equation for neutralization reactions. 10. After constructing molecular reactions for metathesis reactions, be able to identify spectator ions and write the net ionic equations. 11. Assign oxidation numbers to atoms. 12. Determine whether a reaction is Redox or not. 13. Use the activity series to predict whether a Redox (single replacement) reaction will occur, and be able to write the molecular and net ionic equations if it does.	8
HW Problems	Exercises (pp. 111-112, 158-159): 3.1, 3.5, 3.11, 3.15, 3.17 and 4.15, 4.18, 4.24, 4.31, 4.37, 4.41, 4.49, 4.51
Weeks 3–4:	Stoichiometry [C3b] 1. Calculate the atomic weight (average atomic mass) of an element from the relative abundances and masses of its naturally occurring isotopes. 2. Calculate the percentage composition of a compound from its formula. 3. Calculate the molar mass of a substance from its chemical formula. 4. Be able to interconvert between moles, mass, and number of particles of a substance. 5. Calculate the empirical formula of a compound from either elemental percent composition or quantity of CO2 and H2O produced from its combustion. 6. Calculate the molecular formula of a compound from the empirical formula and molecular weight. 7. Find the mass of any substance in a chemical reaction from the mass of one substance. 8. Determine the limiting reactant (limiting reagent) in a reaction and then calculate the amount of each product and the mass of the excess reactant left over. 9. Calculate theoretical yield. 10. Calculate moles of solute, volume of solution, or molarity of the solution from the other two. 11. Recognize and solve dilution problems. 12. Calculate the volume of a certain molarity solution required to react with another solution of known molarity. 13. Calculate the mass of a substance that would be required to react with a given volume of a solution of known molarity. 14. Calculate mass of solute or concentration of an unknown solution from titration data.	4,6
HW Problems	Exercises (pp. 112-115, 159-165): 3.23, 3.27, 3.29, 3.33, 3.37 3.43, 3.51, 3.59, 3.63, 3.71, 3.103 and 4.61, 4.71, 4.77, 4.83
Week 5:	Thermochemistry 1. Understand what the system, the surroundings, and the universe mean. 2. Be familiar with the units of energy. 3. Understand what the First Law of Thermodynamics means. 4. Be familiar with how the internal energy of a system is affected by exchanges of heat and work between the system and the surroundings. 5. Understand what a state function is. 6. Define enthalpy, and explain how heat transfer from or to the system at constant pressure changes it. 7. Know what the sign of the enthalpy indicates about the reaction. 8. Be able to sketch an enthalpy diagram for reactions given their enthalpy changes. 9. Be able to calculate the amount of heat released or absorbed by a reaction, knowing the quantity of the reactants and the enthalpy of the reaction on a mole basis. 10. Define heat capacity and specific heat (capacity). 11. Be able to work problems on calorimetry. 12. State and apply Hess's Law of Constant Heat Summation in calculating enthalpies of reaction from enthalpies of other reactions. 13. Know what the standard state of an element or compound is. 14. Define and illustrate what is meant by standard enthalpy of formation. 15. Calculate the enthalpy change of a reaction using a table of standard enthalpies of formation.	9
HW Problems	Exercises (pp. 204-213): 5.3, 5.17, 5.19, 5.23, 5.25, 5.27, 5.29, 5.31, 5.33, 5.37, 5.41, 5.43, 5.45, 5.49(c), 5.53, 5.55, 5.60, 5.61, 5.65, 5.67, 5.71, 5.75, 5.85,5.99,5.112
Weeks 6–7:	Electron Structure and Periodicity 1. Understand the relationships c = and E = h. 2. Understand the concept of a quantized atom and its relationship to a line spectra of atoms. 3. Explain the concept of ionization energy. 4. Describe the Uncertainty Principle and its affect on atomic theory. 5. Understand the relationship = h/mv and its affect on atomic theory. 6. Describe how quantum numbers define electron orbitals and their value limitations. 7. Describe the shapes of the orbital types. 8. Understand the concept of electron spin and how it relates to electron configuration. 9. Write the electron configuration both symbolically and as an orbital diagram for any element. 10. Be able to write electron configurations, especially valence configurations, for any element, using the periodic table with the knowledge of the s,p,d, and f blocks. 11. Describe the variations of atomic radii in the groups and periods on the periodic table and the underlying reasons for the variations. 12. Describe and explain the observed changes in successive ionization energies for a given atom. 13. Describe the variations in first ionization energies in the groups and periods on the periodic table and the underlying reasons for the variations. 14. Do the same with the electron affinities of the elements. 15. Describe the periodic trends in metallic and nonmetallic behavior and chemical activity.	Lab Activity to be announced
HW Problems	Exercises (pp. 251-259, pp. 292-299) 6.2, 6.4, 6.5, 6.7 - 6.17, 6.21 - 6.28, 6.32 - 6.42, 6.45 - 6.74, 6.90, 6.92, 6.97 and 7.7, 7.9 - 7.34, 7.35, 7.37 - 7.52, 7.52.
Weeks 8–10:	Chemical Bonding and Molecular Geometry 1. Be able to write the Lewis symbol for any atom. 2. Understand the energies involved in the formation of ionic bonds—ionization energy, electron affinity, and lattice energy. 3. Predict the formula of an ionic compound between representative elements using the octet rule, and the periodic table to predict an atom's probable valence. 4. Describe what happens to radius when an atom forms an ion. 5. Be able to explain the variation in size of an isoelectronic series. 6. Describe the nature of the covalent bond in terms of electron cloud overlap. 7. Be able to show covalent bond formation using Lewis symbols. 8. Be able to draw Lewis structures for bonds between atoms—single, double, and triple covalent. 9. Relate bond energies to bond order. 10. Explain electronegativity, how it varies on the periodic table, and its relationship to the nature of the bond between two atoms. 11. Predict the polarities of bonds between any two atoms from their electonegativities or their positions on the periodic table. 12. Write correct Lewis structures for any simple molecule or ion even when there is an exception to the octet rule. 13. Be able to write resonance structures when no one structure is adequate. 14. Relate the number of electron domains in the valence shell of an atom to the geometric arrangement of electrons around the atom. 15. Understand that the relative degree of repulsion between nonbonding pairs is greater than between bonding pairs of electrons. 16. Predict the molecular shape of a molecule or ion from its Lewis structure. 17. Predict, from its molecular shape and the electronegativities of the atoms involved, whether a molecule is polar (has a dipole). 18. Explain the types of hybridization. 19. Assign the type of hybridization on the basis of the electron geometry of the valence shell of an atom. 20. Describe the bonding between atoms in a molecule as sigma or pi. 21. Explain the concept of delocalization in bonds. 22. Describe how molecular orbitals are formed from atomic orbitals. 23. Explain the meaning of bonding and antibonding molecular orbitals. 24. Construct the molecular-orbital energy-level diagram for a diatomic molecule or ion predicting the bond order and the number of unpaired electrons.	16
HW Problems	Exercises (pp. 336-343, 388-397): 8.2 - 8.4, 8.7 - 8.21, 8.29 - 8.42, 8.45 - 8.64, 8.84 - 8.86 and 9.1 - 9.6, 9.8, 9.11 - 9.30, 9.32, 9.34, 9.35, 9.42, 9.43, 9.47, 9.50 - 9.53, 9.57, 9.58
Week 11:	Gases 1. Describe properties of gases compared to other physical states. 2. Define common units of gas pressure. 3. Describe how gases respond to changes in V, n, P, and T. 4. Be able to solve problems using combined and ideal gas equations. 5. Be able to calculate molar mass from gas density and vice versa. 6. Calculate the partial pressure of any gas from the composition of its mixture. 7. Understand the process and calculation of the pressure of a gas collected over water. 8. Calculate mole fraction from partial pressure. 9. Describe how the relative rates of diffusion and effusion of gases depends on their molar masses. 10. Understand the kinetic molecular theory. 11. Be able to work through gas stoichiometry problems. 12. Understand that real gases deviate from ideal gases especially at high pressure and/or low temperature. 13. Know the real gas equation, with corrections for particle attraction and size.	5
HW Problems	Exercises (pp.432-440): 10.1, 10.4, 10.9, 10.15, 10.21, 10.29, 10.32, 10.45, 10.47,10.54, 10.55, 0.57, 10.63, 10.69, 10.71, 10.73, 10.75, 10.81, 10.83, 10.89, 10.97, 10.102
Weeks 12–13:	Solids, Liquids, Changes in Phase, and Intermolecular Forces 1. Understand the kinetic molecular theory explanation of physical states. 2. Describe the types of intermolecular force and be able to state the type expected for a substance knowing its molecular structure. 3. Know the meaning of viscosity, surface tension, critical temperature, and critical pressure, and how they relate to the intermolecular force. 4. Understand how vapor pressure depends on intermolecular attraction and temperature. 5. Define boiling point. 6. From the heat capacities and enthalpies of state change needed, be able to calculate the amount of heat to change a substance from one temperature and state to another. 7. Predict the type of solid (ionic, molecular, metallic, or covalent network) a substance is and the properties it has because of this.	15
HW Problems	Exercises (pp. 478-484): 11.1, 11.2, 11.5, 11.13-11.24, 11.29, 11.31, 11.33, 11.45, 11.47, 11.51, 11.53, 11.55, 11.61(b), 11.65, 11.71-11.76, 11.88, 11.95
Weeks 14–15:	Solution Properties 1. Describe the energy changes associated with the formation of a solution and the role of entropy. 2. "Like dissolves like!" 3. Effects of temperature and pressure on solubility. 4. Define units of concentration, mass percent, ppm, mole fraction, molarity, molality, and be able to calculate each from appropriate data. 5. Be able to convert a concentration from one unit to the other. 6. Describe the effect of solute (or solvent) concentration on each colligative property—vapor pressure, boiling point, freezing point, osmotic pressure. Be able to calculate any of these effects from concentration data. 7. Calculate the concentration and molar mass of a nonvolatile, nonelectrolyte from its effect on a colligative property. 8. Explain the difference in magnitude of these effects caused by electrolytes compared to nonelectrolytes. Define the van't Hoff factor, i. 9. Become familiar with the types of colloids.	14,17
HW Problems	Exercises (pp. 564-573): 13.7, 13.13, 13.17, 13.23, 13.35, 13.45, 13.49, 13.57, 13.65, 13.67
Weeks 16–17:	Chemical Kinetics [C3d] 1. Express the rate of a reaction in terms of changes in the concentration of a reactant or a product per time. Understand how to change from one to the other. 2. Understand the difference graphically between average rate and instantaneous rate. Be able to calculate both. 3. Explain the meaning of the reaction rate law and the rate law constant. 4. Be able to determine a reaction rate law for a reaction from experimental data. 5. Calculate the rate law constant (including units) after finding the rate law constant from experimental data. After this, calculate the rate of another experiment not included in the data. 6. Understand what is meant by order in terms of a reactant as well as the overall order. 7. Explain graphically the concept of activation energy and how temperature affects reaction rate. 8. Understand how temperature affects the rate law constant for a reaction. 9. Be able to relate the collision model to all of the above. 10. Explain what is meant by a reaction mechanism and know the meaning of elementary steps, rate-determining step, and intermediate species. 11. Be able to explain and show how a rate law is derived from a certain reaction mechanism. 12. Describe the theory of how a catalyst works.	12
HW Problems	Exercises (pp. 617-627): 14.15, 14.19(b), 14.23, 14.29,14.33, 14.37, 14.39, 14.45, 14.51,14.59, 14.61, 14.63, 14.69, 14.71, 14.75,14.81, 14.87, 14.103
Weeks 18–19:	Equilibrium [C3c] 1. Understand the meaning of dynamic equilibrium. 2. Write the equilibrium expression for any chemical reaction. 3. Understand the meaning of the magnitude of the value of Keq. 4. Calculate Keq when given appropriate data. 5. Calculate Q, the reaction quotient, to determine if a reaction is at equilibrium and if not determine its direction. 6. Knowing the value of Keq and initial concentrations, calculate equilibrium concentrations. 7. Explain how an equilibrium is shifted by stresses (changes in temperature, pressure, or concentration)—Le Chatelier's Principle. 8. Explain how temperature changes the value of Keq. 9. Describe the effect of a catalyst on an equilibrium.	10
	Exercises (pp. 660-666): 15.2, 15.9, 15.11, 15.13, 15.15, 15.21, 15.27, 15.31, 15.35, 15.43, 15.48, 15.49, 15.51, 15.53, 15.65, 15.74
Weeks 20–22:	Acid-Base Reactions and Solution Equilibria [C3c] 1. List general properties that characterize acidic and basic solutions and the ions responsible. 2. Understand the Brönsted-Lowry theory and be able to identify conjugate acids and bases. 3. Explain the autoionization of water and write the KW expression. 4. Define pH and be able to interconvert between [H+], [OH–], pH, and pOH. 5. Understand what is meant by strength of an acid or a base. 6. Given the acid concentration, be able to interconvert between Ka and pH. Given the base concentration, be able to interconvert between Kb and pH. 7. Calculate the percent ionization from the Ka or the Kb, and vice versa. 8. Understand the relationship between the strength of an acid and the strength of its conjugate base; interconvert between Ka and Kb. 9. Predict whether the solution of a particular salt will be acidic, basic, or neutral. 10. Define an acid and a base in the Lewis sense. 11. Calculate the concentration of each species in a solution formed by mixing an acid and a base. 12. Describe how a buffer solution works and how one can be made at a particular pH. 13. Calculate the change in pH of a buffer upon the addition of a strong acid or a strong base. 14. Distinguish between the various titration curves. 15. Calculate the pH at any point in an acid-base titration. 16. Write a Ksp expression for a salt. 17. Interconvert between solubility and Ksp. 18. Calculate the effect of a common ion on the solubility of a slightly soluble salt. 19. Predict whether a precipitate will form when two solutions are mixed. 20. Understand the effect of pH on the solubility equilibrium of an acidic or basic ion.	7,11,19
HW Problems	Exercises (pp. 712-719, 760-767): 16.6, 16.15, 16.16, 16.18, 16.19, 16.21, 16.25, 16.29, 16.36, 16.43, 16.45, 16.49, 16.53, 16.55, 16.64, 16.73, 16.75, 16.82, 16.85, 16.86, 16.87, 16.94, 16.102, 16.103 and 17.2, 17.3, 17.5, 17.12, 17.16, 17.17, 17.20, 17.23, 17.31, 17.35, 17.41, 17.45, 17.47, 17.51, 17.57, 17.64, 17.70
Weeks 23–24:	Chemical Thermodynamics [C3e] 1. Define entropy in terms of randomness or disorder, and state the second law of thermodynamics. 2. Predict the sign of the entropy of a given process, and state the third law of thermodynamics. 3. Describe the effect of temperature and state changes on entropy. 4. Calculate S° for a reaction using a table of absolute entropies, S°. 5. Define free energy in terms of enthalpy and entropy and explain the relationship of the sign of G, and the spontaneity of a reaction. 6. Calculate G° for a reaction using a table of Gf° for the reactants and products. 7. Describe the conditions of "standard" state for standard free energy. 8. Interconvert G° and K for a reaction. 9. Describe the relationship between G and work. 10. Calculate the free energy change for a reaction at nonstandard conditions, G, knowing G°, T, and the data needed to calculate Q. 11. Predict how G changes with T, given the signs of H, and S. 12. Estimate G° at any given temperature, given H° and S°.	13
HW Problems	Exercises (pp. 836-845): 19.8, 19.21, 19.37, 19.41, 19.47, 19.53, 19.56, 19.61, 19.73, 19.75
Weeks 25–26:	Redox Reactions and Electrochemistry 1. Identify redox reactions, the species oxidized, reduced, the oxidizing agent, and the reducing agent. 2. Balance redox reactions by using oxidation number method and half-reactions method. 3. Diagram and label electrochemical cells, both voltaic and electrolytic. 4. Calculate emf of voltaic cell given electrode potentials. 5. Given electrode potentials predict if a reaction is spontaneous. 6. Interconvert E°, G°, and K for a redox reaction. 7. Be able to calculate any variable in the Nernst equation given the others. 8. Calculate time, current, or amount of a substance produced by electrolysis given the other two. 9. Calculate the maximum electrical work performed by a voltaic cell.	20,21
HW Problems	Exercises (pp. 890-899): 20.3, 20.12, 20.13, 20.15, 20.17, 20.20, 20.23, 20.25, 20.31, 20.34, 20.39, 20.47, 20.50, 20.58, 20.60, 20.69, 20.83, 20.85
Weeks 27–28:	Descriptive, Organic, and Nuclear Chemistry 1. Review Sections 3.2 and 4.2–4.4 and in the process try to establish overall rules for writing chemical reactions and for predicting the products given the reactants. 2. Go through sections 7.6–7.8 and 22.1 and add to these rules those you've established, be prepared to predict the products of reactions. 3. Predict the products of chemical reactions involving oxidation or combustion involving oxygen or proton transfer reactions. 4. Identify the types of hydrocarbons. 5. Understand and be able to identify structural isomers. 6. Know the major function groups. 7. Be able to write, balance and predict the products of nuclear reactions. 8. Understand the meaning of half-life.	22
HW Problems	Various Problems from the textbook will be assigned for this unit. Additional online assignments will likely accompany this unit.
Weeks 29–30:	Review and AP* Exam Practice
Exam:	Tuesday, May 15 2008

AP Chemistry Lab Report Guidelines

About Your Lab Book

Your Lab Book is a permanent record of the GRADED laboratory work that you have completed in AP Chemistry. I will grade this notebook at the end of each quarter. You should safeguard it, since some colleges require it as proof of completion of the lab component prior to extending credit for an AP Lab Science course such as this one. No white-out should be used on a completed lab.

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Required Elements

Each section should be marked with its name (Title, Date, Purpose, etc.,) as a header on the left edge of the page. Remember that you must always use complete sentences as well as correct spelling, punctuation and grammar. Include the name of your lab partner on the title page. It is essential that you collaborate with your partner(s) to divide labor ; for instance, one partner can collect data while the other takes the readings off the instrumentation. Each lab report should be typed in a standard 10pt font, with all graphs and tables being computer generated as well. The ONLY component of your lab that should be hand-written is your signature, along with the date of experiment at the bottom of each page. You will typically have one week to complete your lab report. When it is returned to you, place it in your Lab Book, the archive of all your AP Chemistry labs.

Pre-Lab Elements

Title

The title should be descriptive. “Experiment 3” is not a descriptive title. “Determination of the Molecular Weight of Oxygen from the Decomposition of Potassium Chlorate” is a descriptive title.

Date

This is the date (or dates) that you performed the experiment.

Objective

A brief statement of what you are attempting to do.

Procedure

Organize your procedure along a single column on the left half of the page in this section, leaving lots of room on the right side of the page for corresponding notes, observations, etc. You are not to copy the entire set of procedures from the lab manual! Do not include lengthy, detailed instructions; you will follow the detailed procedure given to you in the handout,

The most significant operations should be noted in a sentence or two; in this manner, you will have useful notes available to you when completing your write-up. You may use the class blog to share information with each other to share and compare data and complete the calculations for each given lab.

Data

Record all your data directly in your preliminary lab guide. Organize your data in a neat, orderly way. Label all data very clearly. Use correct significant digits, and always include proper units (g, mL, etc.). Space things out—don’t try to cram everything into a small space. Use tables where appropriate. Remember- YOU are responsible for collaborating to get data from partners!!

Calculations and Graphs

You should show how calculations are carried out. Always provide the equation used and then show how your values are substituted into it. Give the calculated values, with correct units and significant figures. If graphs are included, make the graphs an appropriate size and scale. Label all axes and give each graph a title. I am not responsible for the loss of any materials that are turned in “loose” in your Lab Book. Of course, if experiments are not quantitative, this section may be omitted.

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Conclusions

Your conclusion should have the following components:

1. Discussion of Theory

In this section, you should include information such as:

· What theory was demonstrated in this experiment?

· What do the calculations show?

· How was the purpose of the experiment fulfilled?

· Why does (or doesn’t) the experiment work?

*Refer back to the purpose of the lab to write this section.

2. Experimental Sources of Error

What are some specific sources of error, and how do they influence the data? Do they make the values obtained larger or smaller than they should be? Which measurement was the least precise? Instrumental error and human error exist in all experiments and should not be mentioned as a source of error unless they cause a significant fault. Significant digits and mistakes in calculations are NOT a valid source of error. In writing this section, it is sometimes helpful to ask yourself what you would do differently if you were to repeat the experiment and wanted to obtain better precision. If you can calculate a percent error or percent deviation, do so and include it in this section.

Sources of experimental error are just that: Experimental.

Calculations are not sources of experimental error!

Human error, scale inaccuracy, etc., are not sources of experimental error!

3. Questions

Answer any questions included in the lab directions.

4. Conclusion Statement

Finally, Make a simple statement concerning what you can conclude from the experiment.

Lab Descriptions:

#	Purpose of Experiment	Experiment Name and Objectives
1	Determination of the formula of a compound	Experiment 1: The Determination of a Chemical Formula Determine the water of hydration in a copper chloride hydrate sample. Conduct a reaction between a solution of copper chloride and solid aluminum. Use the results of the reaction to determine the mass and moles of Cu and Cl in the reaction. Calculate the empirical formula of the copper chloride compound.
2	Determination of the percentage of water in a hydrate	Experiment 2: The Determination of the Percent Water in a Compound Carefully heat a measured sample of a hygroscopic ionic compound. Determine the water of hydration of the compound. Complete the chemical formula of the compound.
4	Determination of molar mass by freezing-point depression	Experiment 4: Using Freezing-Point Depression to Find Molecular Weight Determine the freezing temperature of the pure solvent, lauric acid. Determine the freezing temperature of a mixture of lauric acid and benzoic acid. Calculate the freezing point depression of the mixture. Calculate the molecular weight of benzoic acid.
5	Determination of molar mass of a gas	Experiment 5: The Molar Volume of a Gas Measure the gas production of a chemical reaction by a pressure change. Determine the molar volume of the gas produced in the reaction. Calculate the molar volume of a gas at STP.
6	Standardization of a solution using a primary standard	Experiment 6: Standardizing a Solution of Sodium Hydroxide Prepare an aqueous solution of sodium hydroxide to a target molar concentration. Determine the concentration of your NaOH solution by titrating it with a solution of potassium hydrogen phthalate, abbreviated KHP, of precise molar concentration
7	Determination of concentration by acid-base titration, including a weak acid or weak base	Experiment 7: Acid-Base Titration Accurately conduct acid-base titrations. Determine the equivalence point of a strong acid - strong base titration. Determine the equivalence point of a weak acid - strong base titration. Calculate the molar concentrations of two acid solutions.
8	Determination of concentration by oxidation-reduction titration	Experiment 8: An Oxidation-Reduction Titration: The Reaction of Fe^2+ and Ce^4+ Conduct the potentiometric titration of the reaction between ferrous ammonium sulfate hexahydrate and ammonium cerium (IV) nitrate. Measure the potential change of the reaction. Determine the molar concentration of iron (II) ions in a sample of ferrous ammonium sulfate hexahydrate.
9	Determination of mass and mole relationship in a chemical reaction	Experiment 9: Determining the Mole Ratios in a Chemical Reaction Measure the enthalpy change of a series of reactions. Determine the stoichiometry of an oxidation-reduction reaction in which the reactants are known but the products are unknown.
10	Determination of the equilibrium constant for a chemical reaction	Experiment 10: The Determination of an Equilibrium Constant Prepare and test standard solutions of FeSCN^2+ in equilibrium. Test solutions of SCN^– of unknown molar concentration. Determine the molar concentrations of the ions present in an equilibrium system. Determine the value of the equilibrium constant, Keq.
11	Determination of appropriate indicators for various acid-base titrations; pH determination	Experiment 11: Investigating Indicators Conduct strong acid-strong base titrations using solutions of hydrochloric acid and sodium hydroxide, and three different indicator solutions. Select the proper indicator to use with a titration involving a weak acid or a weak base, based on your observations and measurements.
12	Determination of the rate of a reaction and its order	Experiment 12: The Decomposition of Hydrogen Peroxide Conduct the catalyzed decomposition of hydrogen peroxide under various conditions. Calculate the rate constant for the reaction. Determine the rate law for the reaction. Calculate the activation energy for the reaction.
13	Determination of enthalpy change associated with a reaction	Experiment 13: Determining the Enthalpy of a Chemical Reaction Use Hess’s Law to determine the enthalpy change of the reaction between aqueous ammonia and aqueous hydrochloric acid. Compare your calculated enthalpy change with the experimental results.
14	Separation and qualitative analysis of cations and anions	Experiment 14A: Separation and Qualitative Analysis of Cations Prepare and analyze a solution that contains ten selected cations. Analyze an unknown solution that contains a selection of cations. Experiment 14B: Separation and Qualitative Analysis of Anions Prepare and analyze a solution that contains six selected anions. Analyze an unknown solution that contains a selection of anions.
15	Synthesis of a coordination compound and its chemical analysis	Experiment 15A: The Synthesis of Alum Synthesize a sample of potassium aluminum sulfate dodecahydrate (alum). Observe and record the process of synthesizing a compound. Calculate the percent yield of your synthesis. Experiment 15B: The Analysis of Alum Determine the melting temperature of a sample of alum. Determine the water of hydration of a sample of alum. Determine the percent sulfate of a sample of alum. Verify the chemical formula of a sample of alum.
16	Analytical gravimetric determination	Experiment 16: Conductometric Titration and Gravimetric Determination of a Precipitate Measure the conductivity of the reaction between sulfuric acid and barium hydroxide. Use conductivity values as a means of determining the equivalence point of the reaction. Measure the mass of a product of the reaction as a means of determining the equivalence point of the reaction gravimetrically. Calculate the molar concentration of a barium hydroxide solution.
17	Colorimetric or spectrophotometric analysis	Experiment 17: Determining the Concentration of a Solution: Beer’s Law Prepare and test the absorbance of five standard copper (II) sulfate solutions. Calculate a standard curve from the test results of the standard solutions. Test the absorbance of a copper (II) sulfate solution of unknown molar concentration. Calculate the molar concentration of the unknown CuSO4 solution.

Labs, continued

18	Separation by chromatography	Experiment 18: Liquid Chromatography Conduct an liquid chromatographic separation. Conduct a step gradient chromatographic separation. Complete the necessary measurements and calculations to evaluate the components of a mixture that have been separated by liquid chromatography.
19	Preparation and properties of buffer solutions	Experiment 19: Buffers Evaluate a standard buffer solution. Prepare and test an acid buffer solution. Determine the buffer capacity of the standard buffer and the prepared buffer.
20	Determination of electrochemical series	Experiment 20: Electrochemistry: Voltaic Cells Prepare a Cu-Pb voltaic cell and measure its potential. Test two voltaic cells that use unknown metal electrodes to identify the metals. Prepare copper and lead concentration cells, observe, and measure their respective cell potentials. Use the Nernst equation to calculate the Ksp of PbI2.
21	Measurements using electrochemical cells and electroplating	Experiment 21: Electroplating Prepare and operate an electrochemical cell to plate copper onto a brass surface. Measure the amount of copper that was deposited in the electroplating process. Calculate the amount of energy used to complete the electroplating process.
22	Synthesis, purification, and analysis of an organic compound	Experiment 22: The Synthesis and Analysis of Aspirin Synthesize a sample of acetylsalicylic acid (aspirin). Calculate the percent yield of your synthesis. Measure the melting temperature of your aspirin sample. Conduct a colorimetric analysis of your aspirin sample.

All labs are hands-on and use combinations of Vernier LabPro Data Acquisition devices. We typically use LoggerPro software, but occasionally use DataMate software on Texas Instrument calculators. Occasionally we perform hands-on labs in microscale. The procedures will be slight modifications of the corresponding procedures in Advanced Chemistry with Vernier.

The following probes will be utilized:

• Stainless Steel Temperature Probes
• pH Sensors
• Drop Counter
• Colorimeter
• Gas Pressure Sensor
• Radiation Monitor
• ORP Sensor
• Conductivity Probe
• Current Probe
• Voltage Probe