Course Outline: [CR2]

1st Semester – SC77W

  • (AP Chemistry Topic Covered*)

Ch 1 - 22 weeks

Chemical Foundations

Classification of matter, elements and compounds, review of metric system, uncertainty and significant figures, dimensional analysis

Atoms, Molecules, and Ions

Brief introduction to atomic theory and atomic structure, inorganic nomenclature

(BI 1 & 2)*

Students will:

  • Use the Factor-Label Method (Dimensional Analysis) in solving chemistry-related problems.
  • Use metric and SI systems of units in problems and in the lab.
  • Classify matter as elements, compounds or mixtures.
  • Write formulas for and give names of simple inorganic compounds.
  • Cite experimental evidence (including sketching a diagram of the apparatus) for models of the atom as they evolved.

Laboratory

Separation and Analysis by Chromatography1.5 hr

Ch 32.5 weeks

Stoichiometry – Introduction

Chemical mole, balancing equations, mass stoichiometry, including prediction of quantity of a reactant or product, % yield, limiting reactant, determination of empirical and molecular formulas. (Stoichiometry BI 3)*

Students will:

  • Complete and balance chemical equations.
  • Perform calculations using the mole concept of mass and number.
  • Solve problems involving the quantity of a reactant or product, % yield of a product and limiting reactants.
  • Determine the empirical and molecular formula from percentage composition or mass data.

Laboratory

Formula of a Hydrate1 hr

Synthesis of Aspirin1.5 hr

Ch 42 weeks

Types of Chemical Reactions and Solution Stoichiometry

Characteristics of strong, weak and non-electrolytes, molarity calculations, writing ionic equations for the reactions of acids, bases, and salts, stoichiometry of precipitation reactions

(Reaction Types & Stoichiometry BI 3)*

Students will:

  • Distinguish and characterize strong, weak and non electrolytes. Given the chemical formula
  • state whether the crystal structure of a substance in the solid state is ionic or molecular.
  • predict whether it exists in aqueous solution as ions or molecules.
  • use the terms strong or weak to describe the behavior of an electrolyte.
  • cite experimental evidence to support predictions
  • Describe how to prepare a solution of given concentration by specifying
  • the appropriate mass of solute for the volume of solution desired
  • how to dilute a more concentrated solution to the required concentration
  • Determine the concentration of an unknown solution or the quantity of reactant or product based on stoichiometric data from a precipitation reaction or a titration.
  • Write balanced (complete and net) ionic equations to describe reactions involving precipitation, formation of weak electrolyte, or the formation of a gas.
  • Use a table of solubility trends of common cations and anions to be able to predict which combinations are soluble and which are not.

Laboratory

Standardization of a Solution1.5 hr

Determination of % KHP in an Unknown1.5 hr

Ch 52 weeks

Gases

Properties of gases, ideal gas law equation, kinetic-molecular theory, Graham’s law, van der Waal’s equation.

(Gases, BI 1 & 2)*

Students will

  • Convert between various measures of pressure (atm, mm Hg, psi, kPa).
  • State the standard conditions of temperature and pressure.
  • State the relationships that exist between the number of molecules, pressure, volume and temperature for an ideal gas.
  • Predict how a change in any one of the above factors would change another of the factors.
  • Relate the partial pressure of a gas in a mixture to the mole fraction.
  • State the postulates of Kinetic Molecular Theory (KMT).
  • Relate the rate of gaseous diffusion or effusion to the molar mass of the gas.
  • Relate the kinetic energy (Ek) of molecules in a sample to the absolute temperature.

  • Explain the deviations from the ideal gas law in terms of KMT
  • use the van der Waals equation to calculate the real value of pressure or volume, or
  • relate the values of a or b to the strength of the attractive forces or relative size of the molecules.
  • State which gases are produced by automobiles and coal-fired power plants; write equations for the reactions in which products contributing to acid rain are produced.

Laboratory

Molar Volume of Hydrogen Gas and Determination of R (MBL)1 hr

Ch 62 weeks

Thermochemistry

Nature of energy, including storage and transfer modes, enthalpy and PV work, calorimetry, Hess’s Law, standard enthalpies of formation, present sources of energy, new energy sources.

(Thermodynamics, BI 5)*

Students will:

  • Distinguish between kinetic (thermal) and chemical potential energy.
  • Make diagrammatic representations (energy bar charts) to account for energy storage and transfer between system and surroundings
  • Distinguish between q and w; define enthalpy
  • Use Hess’s Law to determine H for a given reaction given either
  • a series of reactions, or
  • standard enthalpies of formation
  • Use calorimetric techniques to determine specific heat and heat of solution
  • Compare energy values of various fossil fuels

Laboratory

Determination of Specific Heats of Metals and Heats of Solution (MBL)2 hr

Ch 72 weeks

Atomic Structure and Periodicity

Line spectra, Bohr model, principles of quantum mechanics, quantum numbers and orbitals, Periodic Table and electron configurations, ionization energies, periodic trends.

(Atomic Theory & Atomic Structure, BI 1 & 2)*

Students will:

  • Use and to relate wavelengths of spectral lines to energies of emitted photons.
  • Construct a model of the hydrogen atom energy levels consistent with observed spectra.
  • Cite and give evidence for the tenets of the Bohr model of the atom.
  • Explain why we must use probability to discuss the location of electrons in an atom.
  • Relate orbital types to the quantum numbers n, l and m.
  • Sketch orbital diagrams and write electron configurations for many-electron atoms and ions.
  • Relate the structure of the Periodic Table to the filling of orbitals in many-electron atoms.
  • Relate periodic trends in ionization energy, atomic radii and electron affinity to the nucleus-electron attractions and electron-electron repulsions.
  • Explain trends in relative reactivity within the alkali metals and halogens to the electron structure of elements in these families.

Laboratory

Spectrum of Atomic Hydrogen2 hr

Ch 8 – 9 3.5 weeks

Bonding – General Concepts and Molecular Architecture

Chemical bonding in ionic compounds, lattice energy, bonding in covalent compounds, electronegativity and bond polarity, average bond energies and H for reactions, VSEPR and Lewis structures, exceptions to the octet rule, resonance, molecular structure and polarity.

Hybridization, localized electron model, Molecular Orbital model, bonding in homonuclear and heteronuclear diatomic molecules.

(Chemical Bonding, BI 1 & 2)*

Students will:

  • Use the Periodic Table to predict the bond type of binary compounds.
  • Use the Born-Haber cycle to calculate lattice energy for binary ionic compounds.
  • Predict relative ionic bond strength from ionic size and electron configurations
  • Use electronegativity to predict bond polarity in covalent bonds.
  • Sketch Lewis structures, including resonance, of various compounds and ions.
  • Use VSEPR to predict and sketch the shape of various compounds and ions.
  • Use the concept of formal charge to predict which, of possible Lewis structures, is most likely.
  • Match the hybridization of the atomic orbitals to the number of pairs of electrons on the central atom in molecules and ions.
  • Draw molecular orbital energy diagrams for various 2nd period homo- and hetero-nuclear molecules and ions. Use these to predict the order, relative length and stability of the bonds.

Laboratory

Making Molecular Models 1 hr

The 6-solution problem (intro to qualitative analysis)1 hr

Ch 10 – part 12 weeks

Solids and Liquids

Intermolecular forces, interactions in liquids, bonding and structures in metals, atomic and molecular solids and in ionic solids, types of unit cells vapor pressure and changes of state, phase diagrams.

(Liquids & Solids, BI 1 & 2)*

Students will:

  • Classify intermolecular forces in a given substance as ionic, covalent, London, dipole-dipole, hydrogen bonding, or metallic.
  • Classify a crystal as molecular, ionic, covalent or metallic.
  • Determine density from unit cell structure and atomic radii and vice-versa.
  • Use Fraunhofer and Bragg diffraction to determine spacing in crystals.

Laboratory

Simulating X-ray Diffraction Experiments (from ICE resources)2 hr

Total Time in Laboratory– 15.5 hours

This time does not include pre-lab discussion or post-lab analysis.

All of these labs are student-centered (hands-on).

**********

2nd Semester – SC77W

Ch 10 – part 21 week

Students will:

  • Relate vapor pressure to intermolecular forces and boiling point.
  • Use the Clausius-Clapeyron equation to determine Hv from vapor pressure and temperature data.

(Liquids & Solids, BI 1 & 2)*

Laboratory

Vapor Pressure Determination of Hv of Ethanol (MBL)2 hr

Ch 112 weeks

Properties of Solutions

Solution composition, energy and entropy considerations in solution formation, vapor pressure, colligative properties of electrolyte solutions, colloids.

(Solutions, BI 2)*

Students will:

  • Express the concentration of solutions in various ways (mass percent, mole fraction, molarity, molality).
  • Examine the steps in the formation of a liquid solution in terms of energy and entropy.
  • Use colligative properties to determine molar mass.
  • Relate the vapor pressure of a solution to its composition.

Laboratory

Using Freezing Point Depression to Determine Molar Mass (MBL)1.5 hr

Ch 12 2.5 weeks

Kinetics

Reaction rates, differential and integrated rate laws, reaction mechanisms, collision theory, activation energy and reaction rates, catalysis.

(Kinetics, BI 4)*

Students will:

  • Relate instantaneous rate to the slope of the tangent line to a [A] vs. time curve.
  • Use the method of initial rates to determine the order of reactants, and the value and units of k in a rate law.
  • Use graphical methods to determine the order of a reaction, relate slope to k and determine the half-life of a reaction.
  • Evaluate the likelihood of a reaction mechanism using experimental data.
  • Use the Arrhenius equation to relate rate constant, temperature and activation energy.
  • Explain catalysis in term of Ea, and the distribution of kinetic energies at a given temperature.

Laboratory

Rate and Order of a Chemical Reaction (decolorization of CV by OH–) (MBL) 1.5 hr

Ch 131.5 weeks

Chemical Equilibrium

The equilibrium condition, relationship between equilibrium constant and rate constants for opposing reactions, equilibrium expressions involving pressure and concentration, heterogeneous equilibria, solving equilibrium problems, LeChatelier's Principle.

(Equilibrium, BI 6)*

Students will:

  • Use simulations of opposing processes to show that the same equilibrium condition is reached whether the reactants or products initially predominate.
  • Write the expression for the equilibrium constant (in terms of pressures or concentrations) for reactions involving both homogeneous and heterogeneous equilibrium.
  • Use ICE tables to perform calculations to determine K or the concentration of a given species for an equilibrium system.
  • Use Le Chatelier’s Principle to predict the effect of changes imposed on equilibrium systems.

Laboratory

Colorimetric Determination of an Equilibrium Constant (MBL)2 hr

Ch 142 weeks

Acids & Bases

Nature of acids and bases, Arrhenius, Bronsted and Lewis models of acids and bases, acid strength, the pH scale, calculating the pH of strong and weak acids and bases, polyprotic acids, acid-base properties of salts, effect of structure on acid-base properties, acid-base properties of oxides, strategy for solving acid- base problems.

(Equilibrium, BI 6)*

Students will:

  • Distinguish features of Arrhenius, Bronsted and Lewis models of acids and bases.
  • Identify the relatively small number of strong acids and bases; relate acid strength to the value of the dissociation constant, Ka.
  • Explain the role of water in the Bronsted model of acids and bases.
  • Determine [H+], [OH–], and pH of a weak acid, given initial concentration and Ka.
  • Use pH to determine the Ka or Kb for a weak acid or base.
  • Use the phenomenon of hydrolysis to account for the acid/base properties of salts.
  • Account for acid strength in terms of structural considerations.
  • Use the Lewis model of an acid or base to explain the behavior of substances that B-L theory does not adequately explain

Laboratory

Relative Strengths of Acids1 hr

Hydrolysis of Salts1 hr

Ch 152 weeks

Applications of Aqueous Equilibria

Solutions of acids or bases containing a common ion, buffered solutions, titrations and pH curves, acid-base indicators, solubility equilibria and Ksp , precipitation and qualitative analysis, equilibria involving complex ions.

(Equilibrium, BI 6)*

Students will:

  • Explain how the presence of the salt of the conjugate of a weak acid or base affects the pH of the solution.
  • Describe how to prepare a buffered solution with a desired pH, and determine the buffering capacity, and describe how such a solution resists changes in pH.
  • Use the concept of hydrolysis to determine the pH at the endpoint of a titration of a weak acid or base.
  • Decide which indicator is appropriate for the titration of a weak acid or base.
  • Use Ksp to determine the solubility of a slightly soluble salt or calculate Ksp from solubility.
  • Determine whether a precipitate will form when solutions containing the ions of slightly soluble salts are mixed.
  • Explain how the formation of complex ions affects solubility equilibria.

Laboratory

Determination of the Ka of a Weak Acid (MBL)2 hr

Ch 162 weeks

Spontaneity, Entropy and Free Energy

Spontaneous processes and entropy, the Second Law of Thermodynamics, the effect of temperature on spontaneity, Gibbs Free Energy, entropy changes in chemical reactions, free energy and chemical reactions, free energy and equilibrium, free energy and work.

(Thermodynamics, BI 5)*

Students will:

  • Describe entropy in terms of the distribution of energy over available microstates.
  • View entropy as a way of determining what processes are probable.
  • Calculate changes in enthalpy, free energy, and entropy from appropriate thermodynamic data.
  • Relate Gibbs’ free energy and the value of the equilibrium constant.

Laboratory

None

Ch 172 weeks

Electrochemistry

Oxidation-reduction reactions, galvanic cells, standard reduction potentials; cell potential, electrical work and free energy, dependence of cell potential on concentration (Nernst equation), batteries, corrosion, electrolysis.

(Reaction Types, BI 3)*

Students will:

  • Complete and balance redox equations.
  • Use standard reduction potentials to determine:
  • whether a given redox reaction will be spontaneous
  • the cell potential if it is
  • Perform calculations relating Gibbs free energy and cell potential
  • Relate cell potential to the concentrations of the species
  • Describe the electrochemical reactions that occur in various kinds of batteries
  • Explain how sacrificial anodes protect against corrosion
  • Perform calculations involving electrolytic cells

Laboratory

Voltaic Cells1 hr

Determination of Avogadro’s Number from the Electrolysis of Water1 hr

Total Time in Laboratory– 13 hours

This time does not include pre-lab discussion or post-lab analysis.

All of these labs are student-centered (hands-on).

Materials to be Studied Outside Regular Class Time

Ch 21 Nuclear Chemistry

Nuclear stability and radioactive decay, kinetics of radioactive decay, detection and uses of radioactivity, radiocarbon dating.

(Nuclear Chemistry)

Ch 22 Organic Chemistry

Structure and nomenclature of alkanes, alkenes and alkynes, structural isomerism, aromatic hydrocarbons, functional groups and derivatives.

(Descriptive Chemistry, BI 2)*

AP Chemistry Exam Review (All)*

(BI 1) Big Ideal 1: Structure of Matter

(BI 2) Big Idea 2: Properties of matter, characteristics, states and forces of attraction.

(BI 3) Big Idea 3: Chemical Reactions

(BI 4) Big Idea 4: Rates of Reactions

(BI 5) Big Idea 5: Thermodynamics

(BI 6) Big Idea 6: Equilibrium

Laboratory

# / Recommended by The College Board / Performed in AP Chemistry or 1st year course at Dobson HS
1 / Determination of the formula of a compound / Empirical Formula of Zinc Chloride –
2 days / 1st year
2 / Determination of the percentage of water in a hydrate / Formula of a Hydrate
1 day / Ch 3
3 / Determination of molar mass by vapor density
4 / Determination of molar mass by freezing-point depression / Using Freezing Point Depression to Determine Molar Mass 1.5 days / Ch 11
5 / Determination of the molar volume of a gas / Molar Volume of Hydrogen Gas and Determination of R 1 day / Ch 5
6 / Standardization of a solution using a primary standard / Standardization of a Solution
1 day / Ch 4
7 / Determination of concentration by acid-base titration, including a weak acid or weak base / Determination of the Ka of a Weak Acid 2 days / Ch 15
8 / Determination of concentration by oxidation-reduction titration
9 / Determination of mass and mole relationship in a chemical reaction / Nail Lab
Cu – AgNO3 Reaction 2 days / 1st year
10 / Determination of the equilibrium constant for a chemical reaction / Colorimetric Determination of an Equilibrium Constant 1 day / Ch 13
11 / Determination of appropriate indicators for various acid-base titrations; pH determination / Relative Strength of Acids
1 day / Ch 14
12 / Determination of the rate of a reaction and its order / Rate and Order of a Chemical Reaction
1 day / Ch 12
13 / Determination of enthalpy change associated with a reaction / Determination of Specific Heats of Metals and Heats of Solution 1.5 days / Ch 6
14 / Separation and qualitative analysis of cations and anions / Qualitative Analysis of Group II Cations
2.5 days / 1st year
15 / Synthesis of a coordination compound and its chemical analysis
16 / Analytical gravimetric determination / Conservation of Mass in Chemical Change
2 days / 1st year
17 / Colorimetric or spectrophotometric analysis / Colorimetric Determination of an Equilibrium Constant
Rate and Order of a Chemical Reaction / Ch 13
Ch 12
18 / Separation by chromatography / Separation and Analysis by Chromatography 1.5 days / Ch 1-2
19 / Preparation and properties of buffer solutions
20 / Determination of electrochemical series / Activity Series / 1st year
21 / Measurements using electrochemical cells and electroplating / Determination of Avogadro’s Number
1 day / Ch 17
22 / Synthesis, purification, and analysis of an organic compound / Synthesis & Analysis of Aspirin
2 days / Ch 3

Note: There are 6 additional labs performed in AP Chemistry that are not prescribed by The College Board.

AP Chemistry Syllabus 2013 - 20141