* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download AP Chemistry Syllabus
Chemical potential wikipedia , lookup
Stability constants of complexes wikipedia , lookup
Ultraviolet–visible spectroscopy wikipedia , lookup
Detailed balance wikipedia , lookup
Heat transfer physics wikipedia , lookup
Chemical bond wikipedia , lookup
Acid–base reaction wikipedia , lookup
Thermodynamics wikipedia , lookup
Electrochemistry wikipedia , lookup
Woodward–Hoffmann rules wikipedia , lookup
Marcus theory wikipedia , lookup
Rate equation wikipedia , lookup
Equilibrium chemistry wikipedia , lookup
Chemical equilibrium wikipedia , lookup
Reaction progress kinetic analysis wikipedia , lookup
Enzyme catalysis wikipedia , lookup
Hydrogen-bond catalysis wikipedia , lookup
George S. Hammond wikipedia , lookup
Chemical thermodynamics wikipedia , lookup
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 1 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 2 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 3 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 4 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 5 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. 6 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. 7