要AP的话：
I. Structure of Matter (20%)
Atomic theory and atomic structure
Evidence for the atomic theory
Atomic masses; determination by chemical and physical means
Atomic number and mass number; isotopes
Electron energy levels: atomic spectra, quantum numbers, atomic orbitals
Periodic relationships including, for example, atomic radii, ionization energies, electron affinities, oxidation states
Chemical bonding
Binding forces
Types: ionic, covalent, metallic, hydrogen bonding, van der Waals (including London dispersion forces)
Relationships to states, structure, and properties of matter
Polarity of bonds, electronegativities
Molecular models
Lewis structures
Valence bond: hybridization of orbitals, resonance, sigma and pi bonds
VSEPR
Geometry of molecules and ions, structural isomerism of simple organic molecules and coordination complexes; dipole moments of molecules; relation of properties to structure
Nuclear chemistry: nuclear equations, half-lives, and radioactivity; chemical applications
II. States of Matter (20%)
Gases
Laws of ideal gases
Equation of state for an ideal gas
Partial pressures
Kinetic-molecular theory
Interpretation of ideal gas laws on the basis of this theory
Avogadro's hypothesis and the mole concept
Dependence of kinetic energy of molecules on temperature
Deviations from ideal gas laws
Liquids and solids
Liquids and solids from the kinetic-molecular viewpoint
Phase diagrams of one-component systems
Changes of state, including critical points and triple points
Structure of solids; lattice energies
Solutions
Types of solutions and factors affecting solubility
Methods of expressing concentration (The use of normalities is not tested.)
Raoult's law and colligative properties (nonvolatile solutes); osmosis
Non-ideal behavior (qualitative aspects)
III.Reactions (35-40%)
Reaction types
Acid-base reactions; concepts of Arrhenius, Brönsted-Lowry, and Lewis; coordination complexes; amphoterism
Precipitation reactions
Oxidation-reduction reactions
Oxidation number
The role of the electron in oxidation-reduction
Electrochemistry: electrolytic and galvanic cells; Faraday's laws; standard half-cell potentials; Nernst equation; prediction of the direction of redox reactions
Stoichiometry
Ionic and molecular species present in chemical systems: net ionic equations
Balancing of equations including those for redox reactions
Mass and volume relations with emphasis on the mole concept, including empirical formulas and limiting reactants
Equilibrium
Concept of dynamic equilibrium, physical and chemical; Le Chatelier's principle; equilibrium constants
Quantitative treatment
Equilibrium constants for gaseous reactions: Kp, Kc
Equilibrium constants for reactions in solution
Constants for acids and bases; pK; pH
Solubility product constants and their application to precipitation and the dissolution of slightly soluble compounds
Common ion effect; buffers; hydrolysis
Kinetics
Concept of rate of reaction
Use of experimental data and graphical analysis to determine reactant order, rate constants, and reaction rate laws
Effect of temperature change on rates
Energy of activation; the role of catalysts
The relationship between the rate-determining step and a mechanism
Thermodynamics
State functions
First law: change in enthalpy; heat of formation; heat of reaction; Hess's law; heats of vaporization and fusion; calorimetry
Second law: entropy; free energy of formation; free energy of reaction; dependence of change in free energy on enthalpy and entropy changes
Relationship of change in free energy to equilibrium constants and electrode potentials
IV. Descriptive Chemistry (10-15%)
Knowledge of specific facts of chemistry is essential for an understanding of principles and concepts. These descriptive facts, including the chemistry involved in environmental and societal issues, should not be isolated from the principles being studied but should be taught throughout the course to illustrate and illuminate the principles. The following areas should be covered:
Chemical reactivity and products of chemical reactions
Relationships in the periodic table: horizontal, vertical, and diagonal with examples from alkali metals, alkaline earth metals, halogens, and the first series of transition elements
Introduction to organic chemistry: hydrocarbons and functional groups (structure, nomenclature, chemical properties). Physical and chemical properties of simple organic compounds should also be included as exemplary material for the study of other areas such as bonding, equilibria involving weak acids, kinetics, colligative properties, and stoichiometric determinations of empirical and molecular formulas.
V. Laboratory (5-10%)
The differences between college chemistry and the usual secondary school chemistry course are especially evident in the laboratory work. The AP Chemistry Exam includes some questions based on experiences and skills students acquire in the laboratory: making observations of chemical reactions and substances; recording data; calculating and interpreting results based on the quantitative data obtained; and communicating effectively the results of experimental work.
Colleges have reported that some AP candidates, while doing well on the exam, have been at a serious disadvantage because of inadequate laboratory experience. Meaningful laboratory work is important in fulfilling the requirements of a college-level course of a laboratory science and in preparing a student for sophomore-level chemistry courses in college.
Because chemistry professors at some institutions ask to see a record of the laboratory work done by an AP student before making a decision about granting credit, placement, or both, in the chemistry program, students should keep reports of their laboratory work that can be readily reviewed.
Chemical Calculations
The following list summarizes types of problems either explicitly or implicitly included in the topic outline. Attention should be given to significant figures, precision of measured values, and the use of logarithmic and exponential relationships. Critical analysis of the reasonableness of results is to be encouraged.
Percentage composition
Empirical and molecular formulas from experimental data
Molar masses from gas density, freezing-point, and boiling-point measurements
Gas laws, including the ideal gas law, Dalton's law, and Graham's law
Stoichiometric relations using the concept of the mole; titration calculations
Mole fractions; molar and molal solutions
Faraday's law of electrolysis
Equilibrium constants and their applications, including their use for simultaneous equilibria
Standard electrode potentials and their use; Nernst equation
Thermodynamic and thermochemical calculations
Kinetics calculations