Advanced Placement Chemistry Syllabus

2007 – 2008 School Year

Tara Griffith

 

General Introduction:
AP Chemistry is a course designed to prepare you for college chemistry. It is intended to be a very fast moving course, which hopefully will also prepare you for the AP Exam and give you the opportunity to earn college credit. I would encourage all students to work toward taking the AP Exam. Since it is a weighted course, it also has the possibility to boost your GPA.

 

Prerequisites:
In order to take AP Chemistry, as student MUST have completed one year of high school Honors Chemistry. The first eight chapters of AP Chemistry contain topics normally covered in a high school chemistry course. These topics will be discussed rather quickly. It is also suggested that a student have completed or be taking pre-calculus.

 

Student Responsibility:
Since this is an AP course, the student is expected to be self-motivated and be capable of reading the textbook and gleaning information about chemical principles from the text. It is the student’s responsibility to read all reading assignments faithfully. As you read through the text, be sure to look at diagrams and tables and read though the sample exercises. The green, blue, and purple colored sections in the text entitled "Chemistry and Life", “Chemistry and Work”, and “A Closer Look” are also deserving of the student’s attention. These make chemistry "real".

 

The students are also expected to work assigned questions and exercises. Students should have a 2 inch three ring binder. They will be required to keep all notes, questions and exercises, and labs in an organized manner from the beginning of the year.

 

Students will need to develop their own style of reviewing for tests and quizzes. Very little class time will be devoted for review and there will be no instructor generated chapter review sheets.

 

Methodology:

§         Lecture/Discussion

§         Student Presentations—particularly with homework problems and questions and review materials.

§         Group Work—most frequently used to check answers to homework, permitting us to focus as a large group only on those few areas where conceptual difficulties still remain.

§         Lab Work—where equipment or the number of hands needed is not an issue, students work alone, but when doing work that involves instrumentation, students work in pairs.

 

Grading: 

          Quizzes / Homework (10%)

·                     Every Friday I will give a brief 10-minute quiz that consists of one or two of the topics that were covered during the week. 

·                     I will also be giving FREQUENT reading quizzes, which students can use their notes that they took during reading (5 minutes quizzes).  Most of these quizzes will NOT be announced.

·                     There will be homework assignments almost every night with every few exceptions. Homework will graded according to the dice (explained in class).

          Labs (20%)

·                     I recommend a Student Lab Notebook with carbonless duplicate sets.

·                     More info to come during class time.

          Tests (70%)

·                     Unit tests are one-hour exams, which often include info from two chapters in chemistry text.  I will                  give a single unit test that covers both chapters.

·                     We will have one midterm exam, laboratory exam, and semester exam per semester.

·                     Exams will simulate the current AP Chemistry Exam.  Most exams consist of a “calculator-free” zone                  for the multiple-choice portion, equations, and essays, followed by a section where a calculator is                    allowed for problems.

 

Course Design: In an effort to avoid teaching some of the most challenging topics during the fourth quarter when there are numerous school-wide calendar interruptions and the students become mentally fatigued, the traditional course sequence will be altered.  Early in the year, when students are fresher, we will cover the more challenging topics.  At the end of the year, we will cover the chapters that are mostly review.

 

Teacher’s Schedule:

1st Hour

2nd Hour

3rd Hour

4th Hour

5th Hour

6th Hour

7th Hour

AP Chem

Chemistry

Chemistry

Chemistry

Chemistry

AP Chem

PLAN

 

Course Calendar:

First Quarter

I.       Chemical Calculations and General Stoichiometry        (1-3)  Aug 15 – Aug 24

A.                 Dealing with numerical information—significant digits in measurements and calculations, uncertainties, and percent uncertainties, graphing, dependent vs. independent variables, interpolation vs. extrapolation, analysis of solubility graphs.

B.                 Mole concept—atomic mass, molar masses, conversions from mass to moles to molecules to atoms and vice versa.

C.                 Subatomic particles—brief review, including isotope calculations

D.                Stoichiometry in chemical reactions

E.                 Equation balancing, oxidation numbers

F.                  Percent composition—empirical and molecular formulas, formulas from combustion data

G.                Laboratory experiments

1.                 Familiarization with laboratory equipment quiz

2.                 Paper Chromatography

 

III.     Solution Stoichiometry                (4 and 16)                       Aug 27 – Sep 7

A.                 Net ionic reactions—precipitation, acid-base (Arrhenius), oxidation-reduction

B.                 Solution concentration (molarity, molality, mass percent, mole fraction) and dilutions

C.                 Titrations, pH and pOH

D.                Limiting reagents

E.                 Laboratory experiments

1.                 Unknowns Lab

2.                 Strong Acid-Strong Base Titration (pH Meter)

3.                 Introduction to the Spec-20 (Beer’s Law [Cu2+] in an unknown)

 

III.     Bonding                 Chapters 9 and 10                               Sep 10 – Sep 26

A.                 For molecules or ions involving either octet bonding, nonoctet (expanded octet) bonding, or multiple-bonded systems.

          Determination of

·         Shapes (VSEPR Theory)

·         Polarities

·         Type of hybridization (sp3, sp3d2, etc.)

·         Relative bond stabilities

B.                 Resonance—what is it? When is it predicted? Resonance representations

C.                 Relative bond energies and bond lengths [ΔH calculations from bond energies are delayed until thermal chemistry.]

D.                Laboratory Experiments

1.                  Student-Work Constructing Molecular Models

2.                  Reaction of Cu Metal with Iodine à CuIx?

 

IV.     Gases                                        (5)                                  Oct 1 – Oct 17

A.                 Kinetic molecular theory, Avogadro’s hypothesis

B.                 Gas Laws—Boyle’s Law, Charles’s Law, Law of Partial Pressures, ideal gas law, Graham’s Law of Diffusion

C.                 Stoichiometry relationships in gaseous reaction

D.                Non-ideal gases—sources of deviations, pressure and volume consequences, reasons, Arrhenius equation

E.                 Laboratory Experiments

1.                 Determination of Molar Mass of Butane (via collection over water)

2.                 Molar Mass of a Volatile Liquid (via Dumas method)

3.                 Molar Mass of a Gas (via relative velocity-escape rates)

 

Oct 18 – Midterm Exam over Chapters 1-5, 9-10, 16

 

2nd Quarter

V.                 Energy Changes in Physical, Chemical, and Nuclear Reactions

(6 & 21)

A.                 Determination of heat—calories, Joules, endo- and exothermic systems, specific heat, heat capacity (Cp), heat of fusion, heat of vaporization

B.                 Heats of reaction—Hess’s Law, ΔHf, ΔHcombustion, ΔHrxn, ΔH from bond energies

C.                 Energy of nuclear reactions—fission, fusion, nuclear equations; alpha, beta, and positron decay; gamma radiation

D.                Laboratory Experiments

1.                 Heat of Fusion of Ice—Thermometer

2.                 Heat of Solution (dissolving ionic compounds in water)(CBLs)

3.                 Heat of Neutralization (CBLs)

4.                 Heat of Reaction (Mg and HCl) (CBLs)

 

VI.     Kinetics                                     (14)                               

A.                 Reaction rates and mechanisms, terminology, kinetic energy distributions

B.                 Factors affecting reaction rates

C.                 Reactant orders, overall reaction orders

D.                Potential energy diagrams—endothermic vs. exothermic, catalyzed and uncatalyzed, activation energy

E.                 Numerous demonstrations—light sticks in hot and cold water, catalysis

F.                 Laboratory experiments

1.                 Blue Bottle Reaction (determination of a possible mechanism)

2.                 Analysis of Water Flow through a Buret (Zero, 1st, or 2nd order?)

3.                 Rate Law for Crystal Violet with NaOH (via Spec-20 analysis)

 

VII.    Equilibrium—Gases and Ksp          (15 & 18)                        

A.                 Le Chatelier’s Principle—Qualitative shifts in equilibria

B.                 Mass action expression, Keq, Kp, Kc—[initial] vs. [equilib.]

C.                 Minimum energy and maximum randomness considerations

D.                Solubility guidelines, Ksp

E.                 Laboratory experiments

1.                 Keq of Fe(SCN)2+

2.                 Ksp of PbI2

First Semester Exam

 

3rd Quarter

VIII.   Acids and Bases                          (16 & 17)                        

A.                 Review of acid and base behaviors and theories—Arrhenius, Brønsted-Lowry, Lewis, amphoprotism

B.                 Kw pH-scale revisited, temperature effects on Kw and Ka values

C.                 Weak acids and weak bases—Ka, Kb, hydrolysis behaviors, neutralization, acid and base anhydrides, amphoterism

D.                Polyprotic acids and polybasic compounds

E.                 Laboratory Experiments

1.                 Behaviors of Buffered Solutions: Buffers and Equilibria

2.                 Determination of Ka and Molar mass (weak acid-strong base titration, CBLs, and pH Meter)

3.                 Ka’s of Acid-Base Indicators

 

IX.      Electrochemistry                         (20)                               

A.                 Quick review of terminology, redox reactions

B.                 Electrochemical cells and spontaneity

1.                 Galvanic cells, activity series, E˚cell, Ecell, nernst Equation, Keq values from cell voltages, batteries

2.                 Electrolytic reactions—fused salts vs. aqueous solutions, prediction of products and necessary voltages from experimental observations

C.                 Electroplating—Faraday constant, coulombs, determination of time or ion charges

D.                Laboratory experiments

1.                 Developmental of an Activity Series

2.                 Percent Cu in a Penny

3.                 Percent Fe via MnO4- Titration

4.                 Cell Voltages for Galvanic Cells and Factors Affecting Ecell

 

X.       Thermodynamics                        (19)                               

A.                 Terminology, standard state, first law of thermodynamics

B.                 Entropy and the second law of thermodynamics, ΔS, qualitative predictions, ΔS calculations

C.                 Gibbs’s free energy, ΔG˚, ΔG, and their meaning; qualitative predictions at low and high temperatures

D.                ΔG calculations involving

1.                 ΔGrxn = ΔGprod. – ΔGreact.

2.                 ΔG = ΔG˚ - RT lnK

3.                 ΔG˚ = - n F E˚

E.                 Laboratory experiments

1.                 Determination of ΔH, ΔG, ΔS, and Keq for Cu2+ and Zn Reaction

 

Unit Test—Acids and Bases, Electrochemistry, and Thermodynamics

 

4th Quarter

XI.              Solution Behaviors and Colligative Properties (12)               

A.                 Lattice energies, hydration energies, endo- and exothermic heats of dissolving

B.                 Electrolytes and nonelectrolytes, Debye-Huckel and concentration considerations, Born-Haber-type cycles, colloids

C.                 Applications involving freezing point depression, vapor pressure (Raoult’s Law), boiling point elevation

D.                Laboratory experiments

1.                 Molar Mass via Freezing Point Analysis

 

XII.    Solids and Liquids                       (11)                               

A.                 Intermolecular forces and factors that determine their relative strengths hydrogen bonding

B.                 Molecular, metallic, ionic, and network and planar covalent crystals and their properties, as well as allotropes

C.                 Amorphous solids, glasses, and plastics

D.                Laboratory experiment

1.                 ΔH of Vaporization of Water

 

XIII.   Atomic Structure and Periodicity  (7, 8, 23)                        

A.                 Atomic spectra, electrons in atoms, the Bohr atom, quantum mechanics in the hydrogen atom

B.                 Atomic orbital designations (s, p, d, f, as well, as n, m, ml, and spin), paramagnetism, diamagnetism

C.                 Transition—metal complexes

D.                Chemical families, physical and chemical properties, periodic trends, periods vs. rows, shielding, accounting for observed trends (and some anomalies)

E.                 Laboratory experiments

1.                 Synthesis and Analysis of a Transition Metal Complex Salt (Potassium iron oxalate)

2.                 Flame Tests and Ionization Energy

3.                 Qualitative Analysis—Anions

 

XIV.    Organic Chemistry                      (24)                               

A.                 Hydrocarbons—alkanes, alkenes, alkynes, cyclic structures

B.                 Functional groups—carbohydrates (alcohols, aldehydes, ketones, acids, esters, ethers) and nitrogen-bearing groups (amines, amides, nitro-s)