Download Honors Chemistry

Course Content: Honors Chemistry
I.
MATTER, ENERGY, & CHANGE
a. Define Chemistry. S11.C.1.1.1
b. Identify applications of chemistry in everyday live in the different major areas of Chemistry.
S11.C.1.1.1
c. Explain the difference between mass and weight.
d. Differentiate between gaseous, liquid, and solid states in terms of particles. S11.C.2.1.2
e. Distinguish between physical and chemical properties. S11.C.1.1.2
f. Distinguish between physical and chemical changes. S11.C.1.1.2
g. Interpret the Law of Conservation of Mass. S11.C.2.1.2
h. Distinguish between endothermic and exothermic reactions. S11.C.2.1.2
i. Distinguish between mixtures and pure substances. S11.C.1.1.2
j. Distinguish between homogeneous and heterogeneous mixtures. S11.C.1.1.2
k. Distinguish between elements and compounds. S11.C.1.1.1
l. Identify common elements by name and symbol. S11.C.1.1.1
m. Interpret the Law of Definite Proportions and the Law of Multiple Proportions. S11.A.1.1.1,
S11.A.1.1.2
n. Calculate the percent composition (% by mass) of a given compound. S11.A.3.1.3
II.
MEASUREMENT AND PROBLEM SOLVING
a. Identify SI and other common units of measurement. S11.A.2.2.1
b. Perform Unit conversions using factor-label method. S11.A.2.2.1
c. Compute density of various materials. S11.A.2.2.1
d. Contrast accuracy and precision in data sets. S11.A.2.2.1
e. Calculate percent error for laboratory data. S11.A.2.2.1
f. Perform mathematical operations using proper significant figures and scientific notation.
S11.A.2.2.1
g. Differentiate proportional relationships both numerically and graphically. S11.A.3.1.3
III.
ATOMIC STRUCTURE, ISOTOPES, AND NUCLEAR CHEMISTRY
a. Explain the difference between Dalton’s Atomic Theory and the Modern Atomic Theory.
S11.A.1.1.1, S11.A.1.1.2
b. Summarize the experiments that led to the discovery of the nucleus and electrons. S11.C.1.1.1
c. Describe the properties of electrons, neutrons, and protons. S11.C.1.1.1
d. Define an isotope and employ the atomic number and mass number to identify the atomic
structure of a nuclide. S11.C.1.1.1
e. Calculate the average atomic mass of an element when the relative abundances of each isotope
are given. S11.C.1.1.1
f. Identify different types of radioactive materials. S11.C.1.1.2
g. Relate half-life to the dating processes. S11.C.1.1.2
h. Differentiate between fission and fusion processes. S11.A.1.3.2
i. Write balanced equations for nuclear reactions. S11.A.1.3.2
IV.
ELECTRON ARRANGEMENT AND PERIODIC LAW
a. Interpret the line emission spectrum of an element according to atomic structure. S11.C.2.2.1
b. Compare and contrast Bohr’s model of the atom and the Quantum model of the atom.
S11.A.1.1.1, S11.A.1.1.2
c. Describe the significance of the four quantum numbers relative to an atom’s electronic structure.
S11.C.1.1.4
d. Use the Aufbau Principle, Hund’s Rule, and the Pauli Exclusion Principle in describing an atoms
electronic arrangement using the orbital, electron, and electron dot notations. S11.C.1.1.4
e. Predict electron configurations of common monatomic ions. S11.C.1.1.3, S11.C.1.1.4
f. Relate the Periodic Law to the periodicity of the physical and chemical properties of the
elements. S11.C.1.1.4, S11.C.1.1.4, S11.A.3.3.1
g. Incorporate the sub-level blocks in the arrangement of the Periodic Table. S11.C.1.1.4
h. Relate group configurations and numbers to describe the location and properties of various
elements. S11.C.1.1.4, S11.A.3.3.1
i. Explain the periodic functions of atomic radii. S11.C.1.1.4, S11.A.3.3.1
j. Describe the differences between metals, nonmetals, and metalloids and their location on the
Periodic Table. S11.A.3.3.1
V.
CHEMICAL BONDING
a. Differentiate between an ionic and covalent bond. S11.C.1.1.2
b. Classify bonds according to electronegativity differences. S11.C.1.1.3
c. Explain the relationship of bond energy and bond length to bond formation. S11.C.1.1.3
d. Use the Octet Rule to determine the arrangement of electrons in a molecule during covalent
bonding. S11.C.1.1.3
e. Represent Ionic bonds using orbital notations, Lewis structures, and chemical formulas.
S11.C.1.1.3, S11.A.3.2.3
f. Represent Covalent bonds using orbital notations, Lewis structures, structural formulas, and
chemical formulas. S11.C.1.1.3, S11.A.3.2.3
g. Contrast the distinctive properties of ionic and covalent compounds. S11.C.1.1.2
h. Compare the size of ions to that of corresponding atoms explained by atomic structure.
S11.C.1.1.3
VI.
CHEMICAL NAMES & FORMULAS
a. Determine the formula of an ionic compound between any two given ions. S11.C.1.1.3,
S11A.3.2.3
b. Distinguish between the classical and Stock systems of nomenclature. S11.A.3.3.1
c. Produce the name of a common chemical compounds from its formula. S11.A.3.3.1
d. Develop the formula of a common chemical compound from its name. S11.A.3.3.1
e. Generalize the rules for assigning oxidation numbers. S11.A.3.3.1
f. Predict the oxidation number for each element in a given chemical compound. S11.A.3.3.1
g. Calculate the formula mass or molar mass of a given compound. S11.A.3.3.1
h. Convert between number of particles of a substance and mass and moles of a compound.
S11.A.3.1.3
i. Calculate the percent composition of a given compound. S11.A.3.1.3
j. Explain the relationship between the empirical formula and the molecular formula of a given
compound. S11.A.3.1.3
k. Determine the empirical formula from percent or mass composition. S11.A.3.1.3
l. Produce a molecular formula from an empirical formula. S11.A.3.1.3
VII.
CHEMICAL EQUATIONS AND REACTIONS
a. Formulate chemical equations into sentences. S11.A.1.3.2
b. Compose a chemical equation from the description of a chemical reaction. S11.A.1.3.2
c. Balance a formula equation by inspection. S11.A.1.3.2
d. Classify a reaction as synthesis, decomposition, single replacement, double displacement, and or
combustion. S11.A.1.3.2
e. Predict the products of simple reactions given the reactants. S11.A.1.3.2
f. Use the activity series and solubility chart to predict reactivity. S11.A.1.3.2, S11.A.3.3.1
g. Write appropriate net ionic equations for replacement reactions. S11.A.1.3.2
h. Write dissociation equations for soluble ionic compounds in water. S11.A.1.3.2
VIII.
HEAT AND ENERGY
a. Differentiate between heat and temperature. S11.C.2.1.2
b. Define energy and describe various types. S11.C.2.1.2
c. Interpret the Law of Conservation of Energy. S11.C.2.1.2
d. Distinguish between exothermic and endothermic chemical reactions. S11.c.2.1.2
e. Define and explain enthalpy. S11.C.2.1.2
f. Calculate heat changes for physical and chemical processes. S11.C.2.1.2, S11.A.1.3.1
g. Describe factors that influence reaction rates in terms of Collision Theory. S11.C.1.1.6
h. Explain the meaning of dynamic equilibrium as it applies to reversible chemical reactions.
S11.A.1.3.2
i. Use LeChatelier’s Principle to predict and explain the result of a stress placed on a chemical
reaction at equilibrium. S11.A.1.3.2, S11.A.1.3.1
IX.
STOICHIOMETRY
a. Define a mole in terms of Avogadro’s number. S11.A.1.1.2, S11.A.1.1.5
b. Solve problems involving mass, mole, molar mass, and number of atoms of an element.
S11.A.1.1.5
c. Apply mole ratios to stoichiometric calculations. S11.A.3.1.3
d. Calculate mole-mole, mole-mass, mass-mole, and mass-mass problems for given chemical
reactions. S11.A.3.1.3
e. Determine the limiting reactant and excess reactant in a chemical reaction. S11.A.3.1.3
f. Solve limiting reactant problems to determine the amount (moles and grams) of a substance
produced in a chemical reaction. S11.A.3.1.3
g. Calculate the percent yield in a chemical reaction. S11.A.3.1.3
X.
SOLUTIONS
a. Depict the structure of a water molecule. S11.C.1.1.2
b. Account for the various chemical and physical properties of liquid and solid water. S11.C.1.1.2
c. Generalize the energy processes contributing to endothermic and exothermic heat of solution of
an ionic compound. S11.C.2.1.2
d. Distinguish between electrolytes and nonelectrolytes. S11.C.1.1.2
e. Differentiate among saturated, unsaturated, and supersaturated solutions by using solubility
curves. S11.C.1.1.2
f. Compare the effects of temperature, pressure, and polarity on solubility. S11.C.1.1.2
g. Determine the concentration of a solution using molarity. S11.A.1.3.1
XI.
GASES
a. Use the five assumptions of Kinetic Molecular Theory to explain the characteristics of gases.
S11.C.1.1.5
b. Describe the conditions under which a real gas deviates from an ideal gas. S11.C.1.1.5
c. Describe the relationship between pressure, volume, temperature, and number of particles of a
gas. S11.C.1.1.5
d. Explain how pressure is measured and state standard conditions. S11.C.1.1.5
e. Relate absolute zero and standard temperature to the Kelvin scale. S11.C.1.1.5
f. Solve gas law problems employing Boyle’s, Charles’, Gay-Lussac’s, Dalton’s, and Combined
Gas Laws where appropriate. S11.C.1.1.5, S11.A.3.1.3
g. Incorporate Avogadro’s Principle in summarizing Gay-Lussac’s Law of Combining Volumes.
S11.C.1.1.5
h. Use standard molar volume of a gas to calculate gas mass, volume, molar mass, and density.
S11.C.1.1.5, S11.A.1.3.1
i. Construct the Ideal Gas Law from Boyle’s, Charles’, and Avagadro’s Principle. S11.C.1.1.5
j. Derive the ideal gas constant from the Ideal Gas Law. S11.C.1.1.5
k. Solve Ideal Gas Law problems from pressure, volume, temperature, moles, density, and molar
mass. S11.C.1.1.5, S11.A.1.3.1
l. Calculate volumes, masses, and moles using volume ratios, standard molar volume, and the gas
laws in gaseous chemical reactions. S11.C.1.1.5, S11.A.1.3.1
m. Determine the relative rates of effusion of diffusion of gases using Graham’s Law. S11.C.1.1.5,
S11.A.1.3.1
XII.
ACIDS/BASES
a. Define and give examples of traditional, Bronsted, and Lewis acids and bases. S11.C.1.1.2
b. Describe the general properties of aqueous acids and bases. S11.C.1.1.2
c. Differentiate between strong and weak acids and bases. S11.C.1.1.2
d. Explain and use the pH scale based upon the self ionization of water. S11.A.2.2.2
e. Solve problems for pH, pOH, concentration of hydronium ions and concentrations of hydroxide
ions in solutions. S11.A.1.3.1
f. Apply procedures involved in acid-base titration to calculate the molarity of a solution.
S11.A.1.3.1
Laboratory Program
Laboratory experiments are done throughout the course to reinforce classroom theory and concepts. All
experiments require students to use equipment, instrumentation, and materials to make observations
and/or collect data.
Students are required to submit a lab report after each experiment for evaluation. Some lab reports are
formal while others are completion of lab handouts supplied by the instructor.
Objectives of the laboratory programs include:
o Following directions
o Making accurate observations to collect qualitative data
o Taking accurate measurements to collect quantitative data
o Hypothesizing
o Organizing, classifying, critiquing, and analyzing data
o Predicting outcomes
o Inferring and drawing conclusions
o Creating visual, verbal, and mathematical models
o Demonstrating proper use of lab equipment
o Demonstrating proper lab techniques
o Applying laboratory safety rules
*Eligible Content addressed by these objectives include: S11.A.1.3.1, S11.A.1.3.2, S11.A.2.1.2, S11.A.2.1.3,
S11.A.2.1.4, S11.A.2.1.5, S11.A.2.2.1, and S11.A.2.2.2