College Bound & Advanced Chemistry Pacing Guide
Unit
Pacing
SOL**
Chapters
1. Introductory Concepts
14 days
(7 classes)
1.a-j; 2.h
CB: 1-3
ADV: 1-3
2. Atomic Structure
18 days
(9 classes)
2.a-g, i
CB: 4-6
ADV: 4, 11
3. Chemical Bonding
14 days
(7 classes)
3.d; 2.d-g
CB: 7-8
ADV: 12
4. Chemical Nomenclature
10 days
(5 classes)
3.a
CB: 9
ADV: 5
5. The Mole
12 days
(6 classes)
4.a
CB: 10
ADV: 8
6. Chemical Reactions
16 days
(8 classes)
3.b-c, e
CB: 11, 20
ADV: 6-7, 18
7. Basic Stoichiometry
14 days
(7 classes)
4.b
CB: 12
ADV: 9
8. Energy
12 days
(6 classes)
5.d-f
CB: 17
ADV: 10
9. Kinetics & Equilibrium
6 days
(3 classes)
3.f
CB: 18
ADV: 17
10. States of Matter
14 days
(7 classes)
2.h; 5.a-d
CB: 13-14
ADV: 13-14
11. Solutions
12 days
(6 classes)
2.h; 4.c; 5.cd, g
CB: 15-16
ADV: 15
12. Acids & Bases
12 days
(6 classes)
4.d
CB: 19
ADV: 16
13. Nuclear Chemistry
8 days
(4 classes)
2.b; 3.e
CB: 25
ADV: 19
Note: This schedule allows the SOLs to be covered before the Unit 13 test is administered and accounts for a semester exam
and/or for 9 weeks testing. This is a tight schedule that does not leave any room for missed or shortened days.
** Chemistry SOLs 1.j and 6.a-b are incorporated throughout the curriculum.
Unit One: Introductory Concepts (Chemistry SOL 1.a-j; 2.h)
CB: Chapters 1-3 & ADV: Chapters 1-3
A. Introduction to Chemistry
1. The scientific method
a. The place of the hypothesis
b. The difference between theory and law
2. Lab techniques and safety
B. Scientific measurement
1. Metric system (SI)
2. Accuracy and precision
3. Significant figures
4. Mass, volume, density, specific gravity
5. Heat and temperature
C. Problem solving skills
D. Matter and change
1. States of matter
2. Classification of matter
3. Physical and Chemical changes
The student will:
1. Demonstrate in practical situations how the terms experiment, hypothesis, theory, and law fit into
the scientific method.
2. Demonstrate safe laboratory practices, procedures, and techniques.
3. Understand Material Safety Data Sheet (MSDS) warnings including handling chemicals, hazards,
disposal, and chemical spill clean-up.
4. Identify, locate, and know how to use laboratory safety equipment.
5. Correctly name and identify laboratory equipment.
6. Design and perform controlled experiments to test predictions and predict the outcome when a
variable is changed.
7. Evaluate simple arithmetic equations and manipulate numbers in scientific notation without a
calculator.
8. Use a graphing calculator to solve chemistry problems.
9. Use appropriate technology for data collection and analysis.
10. Construct line and bar graphs utilizing the following: independent variable; dependent variable;
scale and unit of a graph; regression line.
11. Name common SI base units and list examples of SI combined units.
12. Name common metric units.
13. List and define the common metric prefixes.
14. Convert measurements using dimensional analysis.
15. Determine the mean of a set of measurements.
16. Distinguish between the accuracy and precision of a measurement.
17. Solve problems dealing with error and deviation, and evaluate for procedural errors.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
Perform calculations according to the rules for significant figures.
State the difference between the mass and weight of an object.
Calculate the density of an object.
Convert between Celsius and Kelvin temperature scales.
Differentiate among the three states of matter.
Differentiate between pure substances and mixtures, and between homogeneous and
heterogeneous mixtures.
Explain the difference between an element and a compound.
Classify substances as elements or compounds.
Distinguish between physical and chemical properties.
Describe what happens during a physical change and a chemical change.
Unit Two: Atomic Structure (Chemistry SOL 2.a-g, i)
CB: Chapters 4-6 & ADV: Chapters 4, 11
A. Atomic theory
1. Atomic number
2. Atomic mass
3. Mass number
4. Isotopes
B. Electronic structure
1. Atomic models
a.
Early models: Democritus, Dalton, Thomson
b.
Bohr atom
1) Quantized energy
2) Emission spectra
c.
Wave mechanical model
1) Quantum numbers
2) Electron configuration
3) Orbital notation
4) Lewis structures
C. Chemical Periodicity
1. Electron configuration and periodicity
2. Periodic trends
a. Atomic size
b. Ionization energy
c. Electron affinity
d. Electronegativity
e. Ionic size
1) Cations
2) Anions
3. Chemical families
The student will:
1. Describe the development of atomic theory from Democritus to quantum mechanical.
2. Identify the key contributions of the following scientists: Democritus, Dalton, Thomson,
Rutherford, Millikan, Bohr, Mendeleev, Moseley, Planck, Heisenberg, and de Broglie.
3. Summarize Dalton's atomic theory.
4. Differentiate between the major atom components (proton, neutron, and electron) in terms of
location, size, and charge.
5. Determine the atomic number, atomic mass, the number, of protons, and the number of electrons
of any atom of a particular element using the periodic table.
6. Use the location on the periodic table to predict common charges for ions in groups 1, 2, 15, 16,
17, and 18.
7. Use the charge on an ion to find the number of electrons in the ion.
8. Define an atomic mass unit.
9. Use the concept of isotopes to explain why the atomic masses of elements are not whole numbers.
10. Determine the number of neutrons in an isotope given its mass number.
11. Perform calculations to determine the “weighted” average atomic mass of an element from
isotope data.
12. Describe the contributions that Thomson and Rutherford made to the development of atomic
theory.
13. Explain how Bohr's model of the atom differed from its predecessors.
14. Compare the quantum mechanical model of the atom with previous models.
15. Describe the general shape of s, p, and d orbitals.
16. Use the Aufbau principle, the Pauli exclusion principle, and Hund's rule to write the electron
configurations of the elements.
17. Relate the position of an element on the periodic table to its electron configuration.
18. Determine the number of valence electrons from an element’s electron configuration.
19. Identify electrons as valence or core.
20. Write the electron configuration for any element on the periodic table.
21. Work problems involving wavelength, frequency, and energy of light.
22. Explain the origin of the atomic emission spectrum of an element.
23. Explain the origin of the periodic table.
24. Distinguish between a period and a group on the periodic table.
25. State the periodic law.
26. Recognize the demarcation of the periodic table into a s block, p block, d block, and f block.
27. Describe variations within a group and a period of atomic radii, ionization energy, electron
affinity, and electronegativity.
28. Compare an element’s reactivity to the reactivity of other elements in the table.
29. Explain how the shielding effect influences periodic trends.
30. Identify an element as an alkali metal, alkaline earth metal, transition metal, semimetal, halogen,
or noble gas.
31. Distinguish between physical and chemical properties of metals and non-metals.
Unit Three: Chemical Bonding (Chemistry SOL 3.d; 2.d-g)
CB: Chapters 7-8 & ADV: Chapter 12
A. Ionic bonds
B. Metallic bonds
C. Covalent bonds
1. Models of chemical bonding - Valence bond model
2. Bond dissociation
3. Intermolecular attractions
a. London forces
b. Dipole-dipole attraction
c. Hydrogen bonds
The student will:
1. Use the periodic table to find the number of valence electrons in an atom.
2. Draw electron dot formulas (Lewis dot diagrams) of the representative elements.
3. State the octet rule.
4. Describe the formation of a cation or an anion.
5. Give the characteristics of an ionic bond.
6. Explain the electrical conductivity of a melted and aqueous solution of ionic compounds.
7. Use the theory of metallic bonds to explain the physical properties of metals.
8. Describe the formation of a covalent bond between two nonmetallic elements.
9. Describe double and triple covalent bonds.
10. Draw electron dot formulas (Lewis dot diagrams) for simple covalent molecules.
11. Define resonance and draw resonance structures of compounds.
12. Show why some molecules that are exceptions to the octet rule may be paramagnetic.
13. Describes the features of sigma and pi bonds.
14. Use VSEPR theory to describe the shapes of simple covalently bonded molecules.
15. Predict, draw, and name molecular shapes (bent, linear, trigonal planar, tetrahedral, and
trigonal pyramidal).
16. Describe the shapes of simple molecules using orbital hybridization.
17. Determine whether a bond is non polar covalent, polar covalent, or ionic using
electronegativity values.
18. Show the relationship between polar covalent bonds and polar molecules.
19. Define bond dissociation energy.
20. Name and describe the weak attractive forces that hold molecules together.
21. Identify the characteristics of molecular substances.
Unit Four: Chemical Nomenclature (Chemistry SOL 3.a)
CB: Chapter 9 & ADV: Chapter 5
A. Writing chemical formulas using ion charges
B. Naming ionic compounds
C. Naming binary molecular compounds
D. Naming acids
E. Naming simple organic compounds
The student will:
1. Define the terms cation and anion and show how they are related to the terms metal and
nonmetal.
2. Use experimental data to show that different samples of the same compound obey the law of
definite proportions.
3. Differentiate between ionic compounds and molecular compounds, and between formula
units and molecules.
4. Use experimental data to show that two different compounds composed of the same two
elements obey the law of multiple proportions.
5. Use the periodic table to determine the charge on an ion.
6. Name binary ionic compounds using the Roman numeral system where appropriate.
7. Use polyatomic ions for naming and writing formulas of ionic compounds including
carbonate, sulfate, nitrate, hydroxide, phosphate, and ammonium.
8. Name binary covalent/molecular compounds.
9. Write the chemical formulas for certain common substances such as ammonia, water, carbon
dioxide, carbon monoxide, sulfur dioxide, and carbon tetrachloride.
10. Write the chemical formula of a compound, when given the name of the compound.
11. Name a compound when given the formula.
12. Name an acid when given the formula and vice versa.
13. Name a simple hydrocarbon when given its structure.
Unit Five: The Mole (Chemistry SOL 4.a)
CB: Chapter 10 & ADV: Chapter 8
A. Measuring matter
1. The mole
a. Molar mass
b. Avogadro’s number
c. Molar volume
2. Percent composition
3. Empirical formulas
4. Molecular formulas
The student will:
1. Identify the mole as the SI unit that measures the "amount of substance."
2. Define a mole as Avogadro's number of representative particles of a substance.
3. Calculate the number of moles in a given number of representative particles of any substance.
4. Distinguish between the terms gram atomic mass, gram molecular mass, gram formula mass,
and molar mass.
5. Calculate the mass of one mole of any substance.
6. Calculate the mass of a given number of moles of a substance.
7. Calculate the number of moles in a given mass of a substance.
8. Use the molar volume of a mole of gas (STP) to work mole-volume problems.
9. Calculate the gram molecular mass of a gas from density measurements of gases at STP.
10. Calculate the percentage composition of a substance from its chemical formula or
experimental data.
11. Distinguish between an empirical and a molecular formula.
12. Derive empirical and molecular formulas from appropriate experimental data.
Unit Six: Chemical Reactions (Chemistry SOL 3.b-c, e)
CB: Chapters 11, 20 & ADV: Chapters 6-7, 18
A. Chemical reactions
1. Writing and balancing chemical equations
2. Predicting chemical products
a. Combination/Synthesis
b. Decomposition
c. Single displacement —activity series
d. Double displacement
e. Combustion
3. Writing net ionic equations
4. Calculating oxidation states
The student will:
1. Identify the reactants and products in a chemical equation.
2. Transform word equations into chemical equations.
3. Write a balanced chemical equation when given the names or formulas of all the reactants
and products in a chemical reaction.
4. Classify a reaction as combination (synthesis), decomposition, single-replacement, double
replacement (metathetical), neutralization, and/or combustion.
5. Predict the products of simple combination and decomposition reactions.
6. Recognize equation for redox reactions and neutralization reactions.
7. Calculate oxidation states.
8. Define oxidation and reduction in terms of a change in oxidation number and identify atoms
being oxidized or reduced in redox reactions.
9. Identify the oxidizing agent and reducing agent in a redox reaction.
Unit Seven: Basic Stoichiometry (Chemistry SOL 4.b)
CB: Chapter 12 & ADV: Chapter 9
A. Stoichiometry
1. Mass-mass problems
2. Mass-volume problems
3. Volume-volume problems
4. Percent yield
5. Limiting reagents
The student will:
1. Determine the limiting reagent (reactant) in a chemical system.
2. Interpret a balanced chemical equation in terms of interacting moles, representative particles,
masses, and volumes of gases (STP).
3. Perform stoichiometric calculations with balanced equations using moles, mass,
representative particles, and volumes of gases (STP).
4. Knowing the limiting reagent in a reaction, calculate the maximum amount of product(s)
produced and the amount of any unreacted excess reagent.
Unit Eight: Energy (Chemistry SOL 5.d-f)
CB: Chapter 17 & ADV: Chapter 10
A. Heat Changes in chemical reactions
1. Calorimetry
2. Phase changes
3. Problems involving the 1st law of thermodynamics
B. Hess's Law
1. Heat of reaction
2. Standard molar heat of formation
3. Standard molar heat of combustion
C. Stoichiometry and heat
D. Factors which drive reactions
1. Enthalpy
2. Entropy
The student will:
1. Classify reactions as exothermic or endothermic.
2. Use the heating curve to solve calorimetry problems.
3. Distinguish between heat capacity and specific heat.
4. Understand the factors that drive chemical reactions: enthalpy and entropy.
5. State the laws of thermodynamics.
6. Use a calorimeter to determine the heat of combustion of a substance.
7. Define the enthalpy of a substance.
8. Use standard heats of formation to calculate the enthalpy change of a reaction.
9. Interpret and/or draw a potential energy diagram for both exothermic and endothermic
reactions correctly labeling the axes, the enthalpy of the reaction, and the activation energy.
10. Distinguish between standard molar heat of combustion, molar heat of reaction, and standard
molar heat of formation.
11. Solve simple stoichiometric energy problems.
12. Solve Hess’s Law problems.
Unit Nine: Kinetics & Equilibrium (Chemistry SOL 3.f)
CB: Chapter 18 & ADV: Chapter 17
A. Kinetics
1. Factors that affect rate
a. temperature
b. nature of reactants
c. state of subdivision
d. concentration
e. catalysts
2. Rate laws
B. Equilibrium
1. Le Chatelier's Principle
2. Law of mass action
The student will
1. Relate the ideas of activation energy and the activated complex to the rate of a reaction.
2. Understand the factors that influence the rate of a reaction.
3. Know the form for a rate law equation.
4. State and apply Le Chatelier's principle.
5. Write a law of mass action for a reversible reaction.
6. Understand the meaning of the Keq.
Unit Ten: States of Matter (Chemistry SOL 2.h; 5.a-d)
CB: Chapters 13-14 & ADV: Chapters 13-14
A. Kinetic molecular theory
B. Avogadro’s hypothesis
C. The nature of liquids
1. Equilibrium vapor pressure
2. Phase changes and phase diagrams
D. The nature of solids
1. Amorphous solids
2. Crystalline solids
a. Ionic solids
b. Molecular solids
c. Metallic solids
d. Covalent network solids
E. The nature of gases
1. Boyle’s law
2. Charles’ law
3. Law of Pressure/Temperature
4. Combined gas laws
5. Ideal gas law
6. Graham’s law
The student will:
1. Use the kinetic molecular theory to explain the motion, pressure, and temperature of a gas.
2. Relate temperature to the average kinetic energy of the particles in a substance.
3. Use the kinetic molecular theory to explain phenomena such as diffusion, compressibility, or
pressure of gases.
4. Explain the meaning of dynamic equilibrium in phase changes.
5. Describe what happens on a particle level at the melting point and boiling point of a
substance.
6. Relate the physical properties of solids to their type.
7. Identify the basic regions shown in a phase diagram. Locate each of the following on the
diagram: normal boiling point, normal freezing point, triple point, critical point.
8. Explain the significance of absolute zero, giving its value in degrees Celsius and Kelvin.
9. Convert between units of pressure--kPa, atm, and torr.
10. State the values of standard temperature and pressure.
11. State and explain the significance of Avogadro's hypothesis.
12. Use the gas laws to solve problems involving volume, temperature, pressure, and moles.
13. State and use Dalton's law of partial pressures.
14. State and use the ideal gas law.
15. Explain why real gases deviate from the gas laws.
16. State and use Graham's law of diffusion.
17. State and explain the major points of the kinetic molecular theory.
Unit Eleven: Solutions (Chemistry SOL 2.h; 4.c; 5.c-d,g)
CB: Chapters 15-16 & ADV: Chapter 15
A. Water
1. Properties of water
2. Aqueous solutions
a. electrolytes and non-electrolytes
b. colloids and suspensions
B. Solution formation
1. Solubility
2. Concentration units
3. Colligative properties
The student will:
1. Describe the hydrogen bonding that occurs in water based on the structure of the polar water
molecule.
2. Use the concept of hydrogen bonding to explain the following properties of water: high
surface tension, low vapor pressure, high specific heat, high heat of vaporization, and high
boiling point.
3. Define the terms solution, aqueous solution, solute, and solvent and give an example of each.
4. Use the rule that "like dissolves like" to predict the solubility of one substance in another.
5. Calculate the percent of water in a given hydrate.
6. Give the characteristics of colloids and suspensions that distinguish them from solutions.
7. Explain the difference between saturated, unsaturated, and supersaturated solutions.
8. Define and work problems involving the molarity and molality of a solution.
9. Describe how to prepare dilute solutions from concentrated solutions or solids.
10. Explain on a particle basis how the addition of a solute to a pure solvent causes an elevation
of the boiling point and a depression of the freezing point of the resultant solution.
Unit Twelve: Acids & Bases (Chemistry SOL 4.d)
CB: Chapter 19 & ADV: Chapter 16
A. Properties of acids and bases
B. Theories of acids and bases
1. Arrhenius
2. Bronsted-Lowry
3. Lewis
C. Strengths of acids and bases
D. Equilibrium of the ionization of water
1. pH of strong acids
2. pH of weak acids
3. pOH
4. Buffer systems
5. Hydrolysis of salts
E. Solubility product
The student will:
1. List the properties of acids and bases.
2. Write the equation for the self-ionization of water.
3. Calculate the pH of a solution given the [H+] or [OH-], or vice versa.
4. Define and give examples of Arrhenius, Bronsted-Lowry, and Lewis acids and bases.
5. Use the Bronsted-Lowry theory to classify substances as acids or bases; identify conjugate
acid-base pairs in reactions.
6. Use the extent of ionization and the Ka or Kb to distinguish between strong and weak acids
and bases. List the strong acids.
7. Calculate an acid dissociation constant from concentration and pH measurements.
8. Complete and balance a neutralization reaction.
9. Perform calculations involving acid-base reactions.
10. Explain the steps of a titration.
11. Use the concept of hydrolysis to explain why aqueous solutions of some salts are acidic or
basic.
12. Define a buffer, and show with equations how a buffer system works.
13. Use Le Chatelier's principle to explain the common ion effect.
Unit Thirteen: Nuclear Chemistry (Chemistry SOL 2.b; 3.e)
CB: Chapter 25 & ADV: Chapter 19
A. Nuclear
1. Radioactive decay
a. alpha
b. beta
c. gamma
2. Nuclear transformations
a. alpha bombardment
b. neutron bombardment
3. Detection of radioactivity
4. Half-life
5. Uses of radioactivity
6. Fission/Fusion
7. Nuclear Reactors
8. Effects of Radiation
The student will be able to:
1. Write the isotope symbols for alpha, beta, neutrons, and protons.
2. Complete reactions for nuclear decay and bombardment.
3. Explain the use of a Geiger counter and a dosimeter.
4. Work simple problems involving half-life.
5. List 5-6 uses for radioactivity, including power, dating, and medical applications.
6. Distinguish between fission and fusion.
7. Describe in general terms how a nuclear power plant works.
8. Identify the common effects of radiation.