Download AP Chemistry Syllabus

Advanced Placement Chemistry
Student Syllabus
Revised 2007
Course Description
The science of Chemistry seeks to understand the structure and composition of matter and the changes
that it undergoes. Advanced Placement Chemistry examines the fundamental principles of the science
of Chemistry from both macroscopic (descriptive and quantitative) and microscopic viewpoints. Topics
include: matter, nomenclature, chemical stoichiometry and reactions, atomic theory and electronic
structure, chemical bonding and molecular geometry, kinetic molecular theory, thermochemistry,
thermodynamics, chemical equilibria, acids and bases, kinetics, and electrochemistry. Laboratory
experiments provide experience in conducting quantitative chemical measurements and illustrate the
principles discussed in class. The subject matter, laboratory skills, and expected level of understanding
are designed to be roughly equivalent to those in the initial two introductory chemistry courses taken by
chemistry or science majors in college. Students enrolling in the course should be responsible, well
organized, disciplined, focused academically, and have good time-management skills. Mathematics is
used extensively throughout the course.
Prerequisites: completion of Chemistry and Physics; concurrent enrollment in Trigonometry or
higher. (Juniors taking the course are expected to take Physics concurrently).
Course Goals and Student Expectations
Course Goals
· develop an understanding of the knowledge, fundamental principles and concepts of Chemistry
· comprehend the mathematical formulations of physical/chemical principles and recognize the
conditions for which each expression is applicable
Student Expectations
· perform and present results of laboratory experiments (individually and in groups)
· physically manipulate laboratory equipment and apparatus and perform basic lab procedures
· make/record quantitative and qualitative observations of physical/chemical properties and
chemical reactions.
· solve problems algebraically and graphically
· communicate orally and in writing
· describe, explain, and apply conceptual models
· interpret, manipulate, analyze, and evaluate actual and hypothetical data
· raise questions and learn from mistakes
· be an independent learner/thinker
· think analytically
· seek assistance from the instructor and/or other resources and materials as needed
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Course Content
The facts, ideas, inferences, rationalizations, models/theories, and mathematical formulations that make
up our understanding of Chemistry and the process of observation, experimentation, and analysis that
are the basis of this understanding are the dual themes of AP Chemistry. The first is the focus of the
Unit Content presentations/discussions whereas the second is the focus of the laboratory program.
Unit Content
I. Fundamentals
Scientific Method
Matter
Properties
Chemical and Physical Properties
Chemical and Physical Changes
Conservation of Mass
Classification Schemes
Physical States (Phases)
Composition
Periodic Chart
Measurement
SI System (Units)
Scientific Notation
Significant Figures
Calculations
Temperature Conversions
Dimensional Analysis
II. Formulae & Nomenclature
Elements
Atomic Theory & Structure
Fundamental Laws
Dalton to Rutherford
Subatomic Particles
Symbols/Formulas
Isotopes/Allotropes
Compounds
Formula/Model Types
Classification & Nomenclature
Ionic
Binary Covalent
Acids
III. Stoichiometry
Formula Stoichiometry
Mole Concept
Calculations – Concept Map
Mass Percent/Mass Ratio
Empirical/Molecular Formula Determination
Reaction Stoichiometry
Reaction Equations
Writing/Balancing Equations
Calculations – Concept Map
Limiting Reagent & Yield
Solution Stoichiometry
Terminology/Units/Preparation
Calculations – Dilution & Reaction
IV. Reaction Types
Reaction Categories
Oxidation/Reduction (Redox)
Synthesis/Combination
Decomposition
Hydrocarbon Combustion
Single Replacement
Double Replacement
Reactions In Aqueous Solutions
Strong/Weak/Non Electrolytes
Molecular/Ionic/Net Ionic Equations
Single Replacement
Metal/Nonmetal Activity Series
Double Replacement
Precipitation Reactions/Solubility Rules
Acid-Base Reactions
Strong/Weak Acids & Bases
Gas Production
Combination
Nonmetal Oxide + Water  Acid
Metal Oxide + Water  Base
Oxidation-Reduction Reactions
Concept/Terminology/States
Balancing – Total/Half Reaction Methods
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V. Atomic Structure & Periodicity
VII. Gases
Interactions of Light and Matter
Pressure Measurement & Units
Blackbody Radiation
Behavior/Calculations
Photoelectric Effect
Empirical Laws (P,V,T Relationships)
Line Spectra
Ideal Gas Law
Wave-Particle Nature of Light
Density and Molecular Weight
Energy, Frequency
Stoichiometry
Frequency, Wavelength, Speed of Light
Mixtures and Partial Pressures
Bohr Model
Effusion/Diffusion
Concept
Kinetic Molecular Theory of Gases
Atomic Spectra
Real Gases
Quantum Theory and Electronic Structure
Deviations From Ideal
Concepts
Van der Waals Equation
Quantum Numbers
Energy Levels/Sublevels
Orbitals/Orbital Shapes
VIII. Liquids, Solids, Solutions
Electron Configuration/Orbital Diagram
Kinetic Molecular Theory of Liquids and Solids
Periodicity
Intermolecular Forces
Organization of the Periodic Table
Types
Trends & Rationalizations
Relationship to Physical States
Atom/Ion Size
Boiling Point
Ionization Energy
Melting Point
Electron Affinity
Vapor Pressure
Electronegativity
Phase Diagrams
Solutions
Terminology and Units
VI. Chemical Bonding
Bond Types and Role of Electrons
Factors Influencing Solubility/Dissolution
Covalent Bonding
Henry's Law (Gas Solubility)
Lewis Dot Representations
Colligative Properties
Terminology & Octet Rule
Boiling Point Elevation
Resonance & Formal Charges
Freezing Point Depression
Bond Strength/Bond Length
Vapor Pressure Lowering
Molecular Geometry
Electrolytes/Non-Electrolytes
VSEPR Theory
Strong/Weak Electrolytes
Polarity – Bond & Molecule
Bonding Theories
IX. Thermochemistry
Valence Bond Theory
Terminology and Units
Concepts and Terminology
Physical Changes
Hybridization & Bond Types
Temperature Change (q = m·c·T)
Molecular Orbital Theory
Phase Changes (q = n·H)
Concepts and Terminology
Heating/Cooling Curves
Energy Level Diagram (Diatomics)
Chemical Changes
Ionic Bonding
Enthalpies of Reaction/Formation
Lattice Arrangement of Atoms
Stoichiometry
Bond Strength
Solution/Bomb Calorimetry
Lewis Dot Representations
Hess's Law
Bond Energies
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X. Spontaneity and Thermodynamics
Concepts and Terminology
Spontaneity
Entropy
Free Energy
Laws of Thermodynamics
Calculations
Chemical Changes
S°, H°, and G°
Temperature Range of Spontaneity
Physical Changes
S°, H°, and G°
Boiling and Melting Points
XI. Reaction Equilibrium & Solubility
Concepts and Terminology
Dynamic and Static Equilibria
Law of Mass Action/Reaction Quotient
Le Chatelier's Principle
Free Energy and Equilibrium
Equilibrium Calculations
Equilibrium Constant (K)
Equilibrium Concentrations
Relationship between Kc and Kp
Le Chatelier's Principle
Temperature variation of K
Solubility Equilibria
Concepts and Terminology
common-ion effect
fractional (selective) precipitation
effect of pH
Calculations
Solubility
solubility product
common-ion effect
precipitate formation
XII. Acids & Bases
Properties and Types
Concepts and Terminology
Acidity-Basicity Criteria
Acidic-Basic Salts
Theories
Arrhenius
Bronsted-Lowry
Lewis
Self-Ionization of Water
Equilibrium Relationships
Weak Acids/Bases
Neutralization
Titration Curves/Indicators
Indicators
Buffers
Calculations
[H+], [OH], pH, pOH
[Acid], [Base], Ka and Kb
Percent Dissociation
Molecular Weight
Neutralizations and Titrations
XIII. Chemical Kinetics
Reaction Rates
Definitions and Terminology
Factors Affecting Reaction Rates
Rate Laws & Calculations
Forms (Differential and Integrated)
Concentration Dependence
Temperature Dependence
Determination From Data
Molecular Visualization
Collision Theory
Transition State Theory
Potential Energy Diagrams
Reaction Mechanisms
Elementary Reactions
Molecularity
Slow and Fast Steps
Relationship To Rate Law
XIV. Electrochemistry
Electrochemical Cells
Terminology/Cell Diagram
Electromotive Force
Electrode/Cell potential-free energy relationship
Nernst Equation
Faraday’s Law
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Laboratory Experience
The laboratory program consists of investigations where good results require (1) the proper use and
application of laboratory equipment and procedures, (2) accurate quantitative and/or qualitative
data/observations, and (3) the manipulation/evaluation of data and/or the application of conceptual
models. College-level experiments form the basis of the laboratory experience, see Table 1.
Collaborative groups are used to perform, analyze, and report several of the more involved, or lengthy,
experiments. The repertoire of skills/techniques developed in the first-year Chemistry laboratory, see
Table 2 for selected experiments, are utilized and expanded on in the AP Chemistry course.
Table 1. AP Chemistry Experiments
Separation and Gravimetric Analysis – Composition of a Three-Component Mixture
Volumetric Analysis – Acetic Acid Content in Vinegar
Reactions in Aqueous Solutions – Double Replacement Reactions
Atomic Spectroscopy – Line Spectrum of Hydrogen
Spectrophotometric Determination of Cu2+ Concentration – Absorbance Spectrum/Beer’s Law
Gas Laws – Boyle’s Law/Average Molar Mass of Air/O2-N2 Ratio in Air
Intermolecular Forces
Thermochemistry – Heats of Reaction/Hess’s Law
Qualitative Equilibria – Le Chatelier's Principle
Solubility Product Determination
Titration/pH Curves – Determination of Ka and Molar Mass of a Weak Diprotic Acid
Kinetics – Determination of the Rate Expression for an Iodine Clock Reaction
Table 2. Selected First-Year Chemistry Experiments
Penny Analysis – Gravimetric Analysis & Percent Composition
White Powder – Comparison of Physical/Chemical Properties
Density – Identification of Unknown Solids and Liquids
Identification by Chemical Change – Identification of Six Solutions by Their Pair-Wise Reactions
KClO3 Decomposition – Percent Composition of Oxygen/Percent Error
Preparation of a Molar Solution
Empirical Formula Determination – Magnesium Chloride or Hydrate
Evidence of a Chemical Reaction – Characteristics of Chemical Reactions
Decomposition of NaHCO3 – Product Identification and Reaction Equation Determination
Activity of Metals – Determination of an Activity Series/Ionic and Net Ionic Equations
Ten Solutions – Identification of Precipitates in Double Replacement Reactions
Preparation of a Paint Pigment – Quantitative Precipitation and Filtration/Yield Determinations
Standardization of a NaOH Solution – Volumetric Analysis
Cation Flame Test
Qualitative Analysis of Cations – Pb2+, Ag+, Hg22+
Qualitative Analysis of Anions– SO42, CO32, Cl, I
Lewis Structures and Molecular Geometry/Models
Calorimetry – Heat of Combustion or Heat of Solidification or Temperature of a Flame
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Student Evaluation/Assessment
AP Chemistry is a full year course designed to be completed prior to the AP exam at approximately day
165. Students participating in this course meet seven periods a week, with two days consisting of
consecutive double periods. The double periods provide additional time for performing and analyzing
laboratory experiments. Including pre- and post- lab work/analysis, 15 – 20 percent of the available
time is spent on these investigations.
Each six weeks student’s will be evaluated on the basis of performance on assignments, written/oral lab
reports, quizzes, and tests. The six grading periods constitute 90% of the course grade, with two
semester (1/2 year) exams contributing the remaining 10%.
The course takes advantage of students’ first-year chemistry experience to move quickly through the
first several units.
AP courses are weighted courses. Students receive weighted credit only if the grade is an “A” or a “B.”
If an “A normally yields four points n a non-AP course, an “A” in an AP course yields five points. This
ultimately affects the student QPA calculation.
Sample questions (and answers)
1) A sample of dolomitic limestone containing only CaCO3 and MgCO3 was analyzed. When heated,
the limestone decomposes producing CO2 gas and a solid residue.
a) Write the equation for the decomposition of calcium carbonate as described above.
b) When a 0.2800 sample of this limestone was decomposed, it was found to contain 0.0488 g
of calcium. What percent of the limestone by mass was CaCO3?
Answers
a) CaCO3 (s)  CaO (s) + CO2 (g)
b)
0.0448 g Ca
1 mol Ca 1 mol CaCO3 100.0 g CaCO3
 0.112 g CaCO3



1
40.08g Ca
1 mol CaCO3
1 mol Ca
0.112 g CaCO3
 100  40.0 %
0.2800 g sample
2) The reaction H2 (g) + I2 (g)  2 HI (g) is exothermic at 298 K and is first order with respect to both
hydrogen and iodine. Predict the effects of each of the following changes on the initial rate of the
reaction and explain your prediction.
a) Addition of hydrogen gas at constant temperature and volume.
b) Increase in temperature.
Answers
a) Addition of hydrogen gas increases the initial rate of reaction. At constant temperature and
volume, increasing the amount of hydrogen in the container increases the concentration of
hydrogen, and since the reaction is first order with respect to hydrogen, the rate of reaction
increases.
b) The initial rate of reaction will increase. Increasing the temperature of the system shifts the
energy distribution of the molecules toward higher energies. This increases the fraction of
molecules having sufficient energy to overcome the reaction’s activation energy, thus
increasing the rate of reaction.
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Primary Course Materials
Texts
Ebbing, D.D. and S.D. Gammon. General Chemistry, 6th ed., Boston: Houghton Mifflin, 1999
A published laboratory text is not used; handouts are prepared for each laboratory experiment.
Laboratory Equipment
Ordinary equipment for handling of chemicals (beakers, flasks, test tubes burners, funnels, etc.) and
measuring properties or quantities of chemicals (single pan and analytical balances, burets, volumetric
pipets/flasks, pH meters, spectrophotometers, etc.)
Supplemental Materials and Suggested Reading List
Internet
http://antoine.frostburg.edu/chem/senese/101/index.shtml
http://library.thinkquest.org/3659/
http://library.thinkquest.org/10429/
http://library.thinkquest.org/3310/user/index.html
http://www.chem1.com/acad/webtext/virtualtextbook.html
http://preparatorychemistry.com/
http://dbhs.wvusd.k12.ca.us/webdocs/ChemTeamIndex.html
References/Resources
Weast, R.C. Ed., CRC Handbook of Chemistry and Physics, 61st. Ed. Boca Raton, CRC Press, 1981
Windholz, M. Ed., The Merck Index, 9th Ed. Rahway, Merck, 1976
Other college level textbooks and lab manuals.
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