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

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 / III. Stoichiometry
Scientific Method / Formula Stoichiometry
Matter / Mole Concept
Properties / Calculations – Concept Map
Chemical and Physical Properties / Mass Percent/Mass Ratio
Chemical and Physical Changes / Empirical/Molecular Formula Determination
Conservation of Mass / Reaction Stoichiometry
Classification Schemes / Reaction Equations
Physical States (Phases) / Writing/Balancing Equations
Composition / Calculations – Concept Map
Periodic Chart / Limiting Reagent & Yield
Measurement / Solution Stoichiometry
SI System (Units) / Terminology/Units/Preparation
Scientific Notation / Calculations – Dilution & Reaction
Significant Figures
Calculations / IV. Reaction Types
Temperature Conversions / Reaction Categories
Dimensional Analysis / Oxidation/Reduction (Redox)
Synthesis/Combination
II. Formulae & Nomenclature / Decomposition
Elements / Hydrocarbon Combustion
Atomic Theory & Structure / Single Replacement
Fundamental Laws / Double Replacement
Dalton to Rutherford / Reactions In Aqueous Solutions
Subatomic Particles / Strong/Weak/Non Electrolytes
Symbols/Formulas / Molecular/Ionic/Net Ionic Equations
Isotopes/Allotropes / Single Replacement
Compounds / Metal/Nonmetal Activity Series
Formula/Model Types / Double Replacement
Classification & Nomenclature / Precipitation Reactions/Solubility Rules
Ionic / Acid-Base Reactions
Binary Covalent / Strong/Weak Acids & Bases
Acids / Gas Production
Combination
Nonmetal Oxide + Water  Acid
Metal Oxide + Water  Base
Oxidation-Reduction Reactions
Concept/Terminology/States
Balancing – Total/Half Reaction Methods
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
VI. Chemical Bonding / Terminology and Units
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
X. Spontaneity and Thermodynamics / XII. Acids & Bases
Concepts and Terminology / Properties and Types
Spontaneity / Concepts and Terminology
Entropy / Acidity-Basicity Criteria
Free Energy / Acidic-Basic Salts
Laws of Thermodynamics / Theories
Calculations / Arrhenius
Chemical Changes / Bronsted-Lowry
S°, H°, and G° / Lewis
Temperature Range of Spontaneity / Self-Ionization of Water
Physical Changes / Equilibrium Relationships
S°, H°, and G° / Weak Acids/Bases
Boiling and Melting Points / Neutralization
Titration Curves/Indicators
XI. Reaction Equilibrium & Solubility / Indicators
Concepts and Terminology / Buffers
Dynamic and Static Equilibria / Calculations
Law of Mass Action/Reaction Quotient / [H+], [OH], pH, pOH
Le Chatelier's Principle / [Acid], [Base], Ka and Kb
Free Energy and Equilibrium / Percent Dissociation
Equilibrium Calculations / Molecular Weight
Equilibrium Constant (K) / Neutralizations and Titrations
Equilibrium Concentrations
Relationship between Kc and Kp / XIII. Chemical Kinetics
Le Chatelier's Principle / Reaction Rates
Temperature variation of K / Definitions and Terminology
Solubility Equilibria / Factors Affecting Reaction Rates
Concepts and Terminology / Rate Laws & Calculations
common-ion effect / Forms (Differential and Integrated)
fractional (selective) precipitation / Concentration Dependence
effect of pH / Temperature Dependence
Calculations / Determination From Data
Solubility / Molecular Visualization
solubility product / Collision Theory
common-ion effect / Transition State Theory
precipitate formation / 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

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

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)

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.

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
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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