2 CHEMISTRY
The Nature of Chemistry
Topic(s)
The Scientific
Method
Laboratory
Procedure(s)
and Safety
The Metric
System,
Significant
Figures and
Measurement
Concepts
(with Standards)
Competencies (Eligible
Content)
3.2.10.A6:
• Compare and contrast
scientific theories.
• Know that both direct and
indirect observations are
used by scientists to study
the natural world and
universe.
• Identify questions and
concepts that guide scientific
investigations.
• Formulate and revise
explanations and models
using logic and evidence.
• Recognize and analyze
alternative explanations and
models.
• Explain the importance of
accuracy and precision in
making valid measurements.
CHEM.A.1.1.1:
• Classify physical or
chemical changes within
a system in terms of
matter and/or energy.
3.2.C.A6:
• Examine the status of
existing theories.
• Evaluate experimental
information for relevance and
adherence to science
processes.
• Judge that conclusions are
consistent and logical with
experimental conditions.
• Interpret results of
experimental research to
predict new information,
propose additional
investigable questions, or
advance a solution.
• Communicate and defend a
CHEM.A.1.1.2:
• Classify observations as
qualitative and/or
quantitative.
CHEM.A.1.1.3:
• Utilize significant figures
to communicate the
uncertainty in a
quantitative observation.
Anchor Descriptions
CHEM.A.1.1:
• Identify and describe how
observable and measurable
properties can be used to
classify and describe matter
and energy.
Vocabulary
The Nature of Chemistry
Analytical chemistry
Biochemistry
Chemistry
Inorganic chemistry
Organic chemistry
Physical chemistry
Science
Technology
Scientific Methods and Laboratory Safety
& Procedures
Control
Control experiment
Dependent variable
Erlenmeyer flask (E-flask)
Graduated cylinder
Hypothesis
Independent variable
Laboratory balance
Laboratory burner
Law
Observation
Science
Scientific method
Theory
Variable
Laboratory equipment
Beaker
Beaker tongs
Buret
Buret clamp
Chemical splash googles
Control
Control experiment
Crucible and cover
Crucible tongs
Dropper bottles
Evaporating dish
Funnel
Hot plate
Laboratory apron
Micropipettes
Pasteur pipettes
Petri dish
Ring stand
Rubber policeman
Rubber stoppers
Textbook
Pages
L2: Ch 1 (p216)
L1 and L2:
Teacher and
department
generated
materials;
Flinn
Scientific
Safely
Contract
Duration
(in days)
2-5
2-5
2 CHEMISTRY
scientific argument.
Scoopula
Spatula
Test tube
Test tube brush
Test tube holder/ test tube clamp
Test tube tongs
Thermometer
Utility clamp
Watch glass
Well plate/ spot plate
Wire gauze
Math- Metric system, Significant Figures,
Measurement
Absolute zero
Accepted value
Accuracy
Celcius temperature scale
Dimensional analysis
Error
Experimental value
Heat
Human error
Hydrometer
Kelvin temperature scale
Mass
Matter
Metric system
Percent error
Precision
Qualitative
Quantitative
Random error
Rounding
Scientific notation
Significant digits (significant figures)
Specific gravity
System Internacional (SI)
Systemic error
Temperature
Unit
Weight
Metric prefixes
Giga- (G, 109)
Mega- (M, 106)
Kilo- (K, 103)
deci- (d, 10-1)
centi- (c, 10-2)
milli- (m, 10-3)
micro- (µ, 10-6)
nano- (n, 10-9)
pico- (p, 10-12)
Metric Units
Atmosphere (atm)
L2: Ch 1 (p1520)
5-10
2 CHEMISTRY
Cubic meter (m3)
Degrees Celcius (°C)
Density (d)
Gram (g)
Joule (J)
Kelvins (K)
Liter (L)
Mole (mol)
Volume (V)
Newton (N)
Pascal (Pa)
Hertz (Hz or 1/sec or sec-1)
2 CHEMISTRY
Properties and Classification of Matter
Topic(s)
Matter &
Energy
Concepts
(with Standards)
Competencies
(Eligible Content)
3.2.10.A1:
• Identify properties of matter
that depend on sample size.
• Explain the unique properties
of water (polarity, high boiling
point, forms hydrogen bonds,
high specific heat) that
support life on Earth.
CHEM.A.1.1.1:
• Classify physical or
chemical changes within
a system in terms of
matter and/or energy.
3.2.10.B2:
• Explain how the overall
energy flowing through a
system remains constant.
3.2.12.A1:
• Compare and contrast
colligative properties of
mixtures.
• Compare and contrast the
unique properties of water to
other liquids.
3.2.C.A1:
• Differentiate between
physical properties and
chemical properties.
• Differentiate between pure
substances and mixtures;
differentiate between
heterogeneous and
homogeneous mixtures.
3.2.C.B2:
• Explore the natural tendency
for systems to move in a
direction of disorder or
randomness (entropy).
CHEM.A.1.1.2:
• Classify observations as
qualitative and/or
quantitative.
CHEM.A.1.1.3:
• Utilize significant figures
to communicate the
uncertainty in a
quantitative observation.
CHEM.A.1.1.4:
• Relate the physical
properties of matter to its
atomic or molecular
structure.
CHEM.A.1.1.5:
• Apply a systematic set of
rules (IUPAC) for naming
compounds and writing
chemical formulas (e.g.,
binary covalent, binary
ionic, ionic compounds
containing polyatomic
ions).
CHEM.A.1.2.1:
• Compare properties of
solutions containing ionic
or molecular solutes
(e.g., dissolving,
Anchor Descriptions
CHEM.A.1.1:
• Identify and describe how
observable and measurable
properties can be used to
classify and describe matter
and energy.
CHEM.A.1.2:
• Compare the properties of
mixtures.
Vocabulary
Matter and Energy
Allotrope
Alloy
Amorphous solid
Atom
Boiling
Boiling point
Chemical change
Chemical property
Chromatography
Colloid
Compound
Condensation
Crystal
Distillation
Electrolysis
Element
Energy
Evaporation
Freezing
Freezing point
Fusion
Gas
Glass
Heterogeneous mixture
Homogeneous mixture
Kinetic energy
Liquid
Matter
Melting
Melting point
Mixture
Phase
Phase change
Phase diagram
Physical change
Physical property
Plasma
Textbook
Pages
L2: Ch 2 (p2845), Ch 3 (p7476), Ch 10
(p274-287)
Duration
(in days)
8-15
2 CHEMISTRY
dissociating).
3.2.10.B3:
• Explain how heat energy will
move from a higher
temperature to a lower
temperature until equilibrium
is reached.
CHEM.A.1.2.2:
• Differentiate between
homogeneous and
heterogeneous mixtures
(e.g., how such mixtures
can be separated).
CHEM.A.1.2.3:
• Describe how factors
(e.g., temperature,
concentration, surface
area) can affect
solubility.
CHEM.A.1.2.4:
• Describe various ways
that concentration can be
expressed and
calculated (e.g., molarity,
percent by mass, percent
by volume).
CHEM.A.1.2.5:
• Describe how chemical
bonding can affect
whether a substance
dissolves in a given
liquid.
Potential energy
Pure substance
Radiant energy
Solid
Solution
Sublimation
Suspension
Symbol
Triple point
Unit cell
Vapor
Vapor pressure
Vaporization
2 CHEMISTRY
Atomic Theory and Structure
Topic(s)
EMR
Atomic Structure
and History
Nuclear Chemistry
and Radioactivity
Electron
Configurations
Concepts
(with Standards)
Competencies
(Eligible Content)
3.2.10.A5:
• Describe the historical
development of models
of the atom and how
they contributed to
modern atomic theory.
CHEM.A.2.1.1:
• Describe the evolution of
atomic theory leading to
the current model of the
atom based on the works
of Dalton, Thomson,
Rutherford, and Bohr.
3.2.12.A2:
• Distinguish among the
isotopic forms of
elements.
• Explain the probabilistic
nature of radioactive
decay based on
subatomic
rearrangement in the
atomic nucleus.
• Explain how light is
absorbed or emitted by
electron orbital
transitions.
3.2.12.A3:
• Explain how matter is
transformed into energy
in nuclear reactions
according to the
equation E=mc2.
3.2.C.A2:
• Compare the electron
configurations for the
first twenty elements of
the periodic table.
3.2.C.A3:
• Identify the three main
types of radioactive
CHEM.A.2.1.2:
• Differentiate between the
mass number of an
isotope and the average
atomic mass of an
element.
CHEM.A.2.2.1:
• Predict the ground state
electronic configuration
and/or orbital diagram for
a given atom or ion.
Anchor Descriptions
CHEM.A.2.1:
• Explain how atomic theory
serves as the basis for the
study of matter.
CHEM.A.2.2:
• Describe the behavior of
electrons in atoms.
Vocabulary
EMR
Absorption spectrum
Amplitude
Color
Electromagnetic radiation
Electromagnetic spectrum
Emission spectrum
Energy level
Excited state
Frequency ( or or ν)
Ground state
Heisenberg uncertainty principle
Matter-wave
Photoelectric effect
Photon
Plank’s constant (h)
Textbook
Pages
Duration
(in days)
L2: Ch 13
(p372- 384)
10
L2: Ch 5 (p106122)
6-7
Principle quantum number (n)
Quantized
Quantum
Spectrum
Speed of light (c)
Wavelength (λ)
CHEM.A.2.2.2:
• Predict characteristics of
an atom or an ion based
on its location on the
periodic table (e.g.,
number of valence
electrons, potential types
of bonds, reactivity).
People
Bohr, Neils
DeBroglie, Louis
Einstein, Albert
Heisenberg, Werner
Planck, Max
CHEM.A.2.2.3:
• Explain the relationship
between the electron
configuration and the
atomic structure of a
given atom or ion (e.g.,
energy levels and/or
orbitals with electrons,
Atomic Structure & History
Atom
Atomic mass number (atomic
mass, or mass number: A)
Atomic mass unit (amu)
Atomic number (Z)
Average atomic mass
Cathode ray
2 CHEMISTRY
decay and compare
their properties.
• Describe the process of
radioactive decay by
using nuclear equations
and explain the concept
of half-life for an
isotope.
• Compare and contrast
nuclear fission and
nuclear fusion.
3.2.C.A5:
• Recognize discoveries
from Dalton (atomic
theory), Thomson (the
electron), Rutherford (the
nucleus), and Bohr
(planetary model of atom),
and understand how each
discovery leads to modern
theory.
• Describe Rutherford’s
“gold foil” experiment that
led to the discovery of the
nuclear atom. Identify the
major components
(protons, neutrons, and
electrons) of the nuclear
atom and explain how
they interact.
distribution of electrons
in orbitals, shapes of
orbitals).
CHEM.A.2.2.4:
• Relate the existence of
quantized energy levels
to atomic emission
spectra.
Compound
Electron (e-)
Electrostatic Attraction
Element
Ion
Isotope
Molecule
Neutron (n0)
Nucleus
Nuclide
Proton (p+)
Strong (Nuclear) Force
People
Ancient Greeks: Aristotle,
Democritus, Leuccipus
Bequerel, Henri
Bohr, Neils
Chadwick, James
Curie, Marie
Curie, Pierre
Dalton, John
DeBroglie, Louis
Eintein, Albert
Heisenberg, Werner
Millikan, Robert Andrews
Planck, Max
Rutherford, Ernest
Thomson, Sir Joseph John
Nuclear Chemistry and
Radiation
Alpha decay
Alpha particle (α)
Alpha radiation
Band of stability
Beta decay
Beta particle (β)
Beta radiation
Electron capture
Film badge
Fission
Fusion
Gamma radiation
Gamma ray (Γ)
L2 : Ch 28
(p840-862)
5-10
2 CHEMISTRY
Geiger counter
Half life
Ionizing radiation
Neutron absorption
Neutron modification
Nuclide
Positron
Positron emission
Radiation
Radioactive decay
Radioactivity
Radioisotope
Scintillation counter
Strong (Nuclear) Force
Transmutation
Transuranium elements
People
Bequerel, Henri
Curie, Marie
Curie, Pierre
Geiger, Hans
Röentgen, Wilhelm
Electron configuration
Angular momentum (l)
Atomic orbital
Aufbau principle
Electron configuration
Energy level
Hund’s rule
Magnetic quantum number (ml)
Orbital
Orbital diagram
Pauli exclusion principle
Principle quantum number (n)
Quantum
Quantum mechanical model
Quantum number
Spin quantum number (s)
Sublevel
People
Bohr, Neils
Einstein, Albert
L2: Ch 13
(p361-384)
10-15
2 CHEMISTRY
Hund, Friedrich
Pauli, Wolfgang
2 CHEMISTRY
The Periodic Table
Topic(s)
The Periodic
Table and
Periodic
Trends
Concepts
(with Standards)
3.2.10.A1:
• Predict properties of elements
using trends of the periodic
table
3.2.C.A1:
• Explain the relationship of an
element’s position on the
periodic table to its atomic
number, ionization energy,
electro-negativity, atomic
size, and classification of
elements.
3.2.C.A2:
• Compare the electron
configurations for the first
twenty elements of the
periodic table.
• Relate the position of an
element on the periodic table
to its electron configuration
and compare its reactivity of
other elements in the table.
• Predict chemical formulas
based on the number of
valence electrons.
Competencies
(Eligible Content)
CHEM.A.2.2.2:
• Predict characteristics of
an atom or an ion based
on its location on the
periodic table (e.g.,
number of valence
electrons, potential types
of bonds, reactivity).
CHEM.A.2.2.3:
• Explain the relationship
between the electron
configuration and the
atomic structure of a
given atom or ion (e.g.,
energy levels and/or
orbitals with electrons in
orbitals, shapes of
orbitals).
CHEM.A.2.2.4:
• Relate the existence of
quantized energy levels
to atomic emission
spectra.
CHEM.A.2.3.1:
• Explain how the
periodicity of chemical
properties led to the
arrangement of elements
on the periodic table.
CHEM.A.2.3.2:
• Compare and/or predict
the properties (e.g.,
electron affinity,
ionization energy,
Anchor Descriptions
CHEM.A.2.2:
• Describe the behavior of
electrons in atoms.
CHEM.A.2.3:
• Explain how periodic trends in
the properties of atoms allow
for the prediction of physical
and chemical properties.
CHEM.A.1.2:
• Compare the properties of
mixtures.
Vocabulary
Textbook
Pages
The Periodic Table and Periodic
Trends
Alkali metal
Alkaline earth metal
Atomic radii
Boron group
Carbon group
Electron affinity
Electronegativity
Family
Group
Halides (halogens)
Inner transition metals
Ionic radii
Ionization energy
Law of octaves
Metal
Metalloid
Nitrogen group
Noble gases
Nonmetal
Octet rule
Oxygen group
Period
Representative elements
Second ionization energy
Semi-metal
Successive ionization energy
Transition metals
Triads
L2: Ch 5 (p123127); Ch 14
(p390-407)
People
Dobereiner, Johann
Erdmann, Hugo
Janssen, Pierre
Lockyer, Joseph Norman
Mendeleev, Dmitri
Meyer, Julius Lothar
Duration
(in days)
5-8
2 CHEMISTRY
chemical reactivity,
electronegativity, atomic
radius) of selected
elements by using their
locations on the periodic
table and known trends.
Moseley, Henry
Newlands, John Alexander Reina
Ramsey, William
Rayleigh, Lord (John William
Strutt)
2 CHEMISTRY
Chemical Bonding
Chemical Relationships and Reactions
Topic(s)
Lewis
Structures
and Bonding
Molecular
Geometry
Intermolecular
Forces (IMFs)
Nomenclature
Reaction
Types,
Predicting,
and Writing
Products
Concepts
(with Standards)
3.2.10.A2:
• Compare and contrast different
bond types that result in the
formation of molecules and
compounds.
• Explain why compounds are
composed of integer ratios of
elements.
3.2.10.A4:
• Describe chemical reactions in
terms of atomic rearrangement
and/or electron transfer.
• Explain the difference between
endothermic and exothermic
reactions.
Competencies
(Eligible Content)
CHEM.B.1.3.1:
• Explain how atoms
combine to form
compounds through ionic
and covalent bonding.
CHEM.B.1.3.2:
• Classify a bond as being
polar covalent, non-polar
covalent, or ionic.
CHEM.B.1.3.3:
• Use illustrations to predict
the polarity of a molecule.
3.2.12.B4:
• Describe conceptually the
attractive and repulsive forces
between objects relative to
their charges and the distance
between them.
CHEM.B.1.4.1.:
• Recognize and describe
different types of models
that can be used to
illustrate the bonds that
hold atoms together in a
compound (e.g., computer
models, ball-and-stick
models,, graphical models,
solid-sphere models,
structural formulas,
skeletal formulas, Lewis
dot structures).
3.2.C.A1:
• Use electro-negativity to
explain the difference between
polar and non-polar covalent
bonds.
CHEM.B.1.4.2:
• Utilize Lewis dot structures
to predict the structure and
bonding in simple
compounds.
3.2.C.A2:
• Explain how atoms combine to
CHEM.B.2.1.4:
• Predict products of simple
3.2.12.A5:
• Use VSEPR theory to predict
the molecular geometry of
simple molecules.
Anchor Descriptions
CHEM.B.1.3:
Explain how atoms form
chemical bonds.
CHEM.B.1.4:
Explain how models can be
used to represent bonding.
CHEM.B.2.1:
Predict what happens during a
chemical reaction.
Vocabulary
Lewis Structures and Bonding
Anion
Cation
Chemical formula
Compound
Coordinate covalent bond
Formal charge
Formula unit
Ion
Ionic compound
Isomer
Law of definite proportions
Law of multiple proportions
Lewis structure
Lone pairs of electron
Molecular compound
Molecular formula
Molecule
Monatomic ion
Non-bonded pairs
Octet rule
Polyatomic ion
Resonance
Steric number
Structural formula
Valence electrons
Textbook
Pages
Duration
(in days)
L2: Ch 6 (p133160); Ch 15
(p418-452)
7-12
L2: Ch 16 (p358392)
5-10
People:
Dalton, John
Lewis, Gilbert
Molecular Geometry
Bent
Dipole
Hybrid orbitals
Ionic compound
Linear
2 CHEMISTRY
form compounds through both
ionic and covalent bonding.
• Draw Lewis dot structures for
simple molecules and ionic
compounds.
• Predict the chemical formulas
for simple ionic and molecular
compounds.
3.2.C.A4:
• Predict how combinations of
substances can result in
physical and/or chemical
changes.
• Interpret and apply the laws of
conservation of mass,
constant composition (definite
proportions), and multiple
proportions.
• Balance chemical equations by
applying the laws of
conservation of mass.
• Classify chemical reactions as
synthesis (combination),
decomposition, single
displacement (replacement),
double displacement, and
combustion.
chemical reactions (e.g.,
synthesis, decomposition,
single replacement, double
replacement, combustion).
CHEM.B.2.1.5:
• Balance chemical
equations by applying the
Law of Conservation of
Matter.
Molecular orbital theory
Non-polar covalent compound
Octahedral
Parent geometry
Polar covalent compound
See-saw (Teeter- totter)
Square planar
Square pyramidal
Tetrahedral
Trigonal bipyramidal
Trigonal planar
Trigonal pyramidal
T-shaped
VSEPR theory
Intermolecular Forces (IMFs),
Solids and Liquids
Boiling
Capillary action
Emulsify
Hydrogen bond (H bond)
Hydrogen- dipole interactions
Hydrophilic
Hydrophobic
Induced dipole
Intermolecular forces (IMF)
Intramolecular force
Ion- dipole interactions
London dispersion forces
Micelle
Surface tension
Surface tension
Van der Waals forces (Dipoledipole forces or dipole-dipole
interactions)
Vapor pressure
Viscosity
Nomenclature
Acid
Alkali metal
Alkaline earth metal
Anion
B group
Base
L2: Ch 16 (p460468); Ch 17
(p475-481)
7-12
L2: Ch 6 (p149164)
5-10
2 CHEMISTRY
Binary compound
C group
Cation
Conductor
Covalent compound
Ductile
Electrolyte
Electronegativity
Formula unit
Halogens (halides)
-ide
Inner transition metals
Ion
Ionic compound
Malleable
Metal
Metallic bond
Molecular compound
N group
Noble gases
Non-metal
O group
Oxidation number
Salt
Semi-metal (metalloid)
Ternary compound
Transition metals
Prefixes for naming binary
covalent compounds
monoditritetrapentahexaheptaoctanonadecaPolyatomic Ions
Acetate
Ammonium
Bisulfate
2 CHEMISTRY
Bisulfate
Carbonate
Chlorate
Chlorite
Chromate
Hydronium
Hypochlorite
Monohydrogen phosphate
Nitrate
Nitrite
Perchlorate
Permanganate
Phosphate
Phosphite
Sulfate
Sulfite
Reactions
Activity series
Balanced equation
Catalyst
Chemical equation
Coefficient
Combustion reaction
Complete ionic equation
Decomposition reaction
Double replacement reaction
/double displacement reaction
Incomplete combustion reaction
Law of conservation of matter/ law
of conservation of mass
Net ionic equation
Precipitate
Product
Reactant/ Reagent
Salt
Single replacement reaction
/single displacement reaction
Solubility table
Synthesis reaction/ addition
reaction/ direct combination
reaction
L2: Ch 8 (p202278)
10-12
2 CHEMISTRY
The Mole
Topic(s)
The Mole
Empirical &
Molecular
Formulas,
Percent
Composition
Stiochiometry:
Limiting Reactant,
Percent Yield,
and Theoretical
Yield
Concepts
(with Standards)
3.2.10.A4:
• Predict the amounts of
products and reactants in a
chemical reaction using mole
relationships.
3.2.10.A5:
• Apply the mole concept to
determine number of particles
and molar mass for elements
and compounds.
3.2.C.A2:
• Explain how atoms combine to
form compounds through both
ionic and covalent bonding.
• Use the mole concept to
determine number of particles
and molar mass for elements
and compounds.
• Determine percent
compositions, empirical
formulas, and molecular
formulas.
3.2.C.A4:
• Interpret apply the laws of
conservation of mass, constant
composition (definite
proportions), and multiple
proportions.
• Balance chemical equations by
applying the laws of
conservation of mass.
• Use stoichiometry to predict
quantitative relationships in a
chemical reaction.
Competencies
(Eligible Content)
Anchor Descriptions
CHEM.B.1.1.1:
• Apply the mole
concept to
representative
particles (e.g.,
counting, determining
mass of atoms, ions,
molecules, and/or
formula units).
CHEM.B.1.1:
• Explain how the mole is a
fundamental unit of
chemistry.
CHEM.B.1.2.1:
• Determine the
empirical and
molecular formulas of
compounds.
CHEM.B.1.2.2:
• Apply the law of
definite proportions to
the classification of
elements and
compounds as pure
substances.
CHEM.B.1.2.3:
• Relate the percent
composition and mass
of each element
present in a
compound.
CHEM.B.1.4.1:
• Recognize and
describe different
types of models that
can be used to
illustrate the bonds
that hold atoms
CHEM.B.1.2:
• Apply the mole concept to the
composition of matter.
CHEM.B.1.4:
• Explain how models can be
used to represent bonding.
CHEM.B.2.1:
• Predict what happens during
a chemical reaction.
Vocabulary
Chemical Quantities: The Mole
and Stoichiometry
Actual yield/experimental yield
Atomic mass unit
Avogadro’s number
Empirical Theory
Excess reagent/excess reactant
Formula mass/gram formula mass
Formula unit
Limiting reagent/limiting reactant
Molar mass/gram molar mass/gram
molecular mass
Molar volume
Mole
Molecular formula
Percent composition
Percent yield
Representative particle
Stoichiometry
STP (standard temperature and
pressure)
Theoretical yield
People
Avogadro, Amadeo
Textbook
Pages
L2: Ch7 (p170196); Ch9
(p251-289)
Duration
(in days)
15-20
2 CHEMISTRY
together in a
compound (e.g.,
computer models, balland-stick models,
graphical models,
solid-sphere models,
structural formulas,
skeletal formulas,
Lewis dot structures).
CHEM.B.2.1.1:
• Describe the roles of
limiting and excess
reactants in chemical
reactions.
CHEM.B.2.1.2:
• Use stoichiometric
relationships to
calculate the amounts
of reactants and
products involved in a
chemical reaction.
CHEM.B.2.1.3:
• Classify reactions as
synthesis,
decomposition, single
replacement, double
replacement, or
combustion.
CHEM.B.2.1.5:
• Balance chemical
equations by applying
the Law of
Conservation of
Matter.
2 CHEMISTRY
The Kinetic Molecular Theory
Topic(s)
Gas Laws
Concepts
(with Standards)
3.2.10.A3:
• Describe phases of matter
according to the kinetic
molecular theory.
3.2.12.B3:
• Describe the relationship
between the average kinetic
molecular energy,
temperature, and phase
changes.
3.2.C.A3:
• Describe the three normal
states of matter in terms of
energy, particle motion, and
phase transitions.
Competencies
(Eligible Content)
CHEM.B.2.2.1:
• Utilize mathematical
relationships to predict
changes in the number
of particles, the
temperature, the
pressure, and the
volume in a gaseous
system (i.e., Boyle’s law,
Charles’s law, Dalton’s
law of partial pressures,
the combined gas law,
and the ideal gas law).
CHEM.B.2.2.2:
• Predict the amounts of
reactants and products
involved in a chemical
reaction using molar
volume of a gas at STP.
Anchor Descriptions
CHEM.B.2.2:
• Explain how the kinetic
molecular theory relates to
the behavior of gases.
Vocabulary
Gas Behavior and Gas Laws
Atmosphere (atm)
Atmospheric pressure
Barometer
Boyle’s law
Charles’s law
Combined gas law
Compressibility
Dalton’s law of partial pressures
Diffusion
Effusion
Gas pressure
Graham’s law of effusion
Guy-Lussac’s law
Ideal gas
Ideal gas law
Kilopascal (kPa)
Kinetic energy
Kinetic molecular theory
Manometer
Millimeters of Mercury (mmHg)
Pascal (Pa)
Pounds per square inch (psi)
Pressure
STP (standard temperature and
pressure)
Vacuum
People
Avogadro, Amadeo
Boyle, Robert
Charles, Jacques
Dalton, John
Gay-Lussac, Joseph Louis
Graham, Thomas
Textbook
Pages
L2: Ch 10
(p266-274);
Ch 12 (p326354)
Duration
(in days)
10-12
2 CHEMISTRY
Solutions and Colligative Properties
Topic(s)
Solutions,
Concentra-tions,
and Colligative
Properties
Concepts
(with Standards)
3.2.10.A2:
• Compare and contrast
different bond types that
result in the formation of
molecules and compounds.
• Explain why compounds
are composed of integer
ratios of elements.
3.2.10.A4:
• Predict the amounts of
products and reactants in a
chemical reaction using
mole relationships.
3.2.10.A5:
• Apply the mole concept to
determine number of
particles and molar mass
for elements and
compounds.
3.2.12.A1:
• Compare and contrast
colligative properties of
mixtures.
• Compare and contrast the
unique properties of water
to other liquids.
3.2.C.A1:
• Differentiate between
physical properties and
chemical properties.
• Differentiate between pure
substances and mixtures;
differentiate between
heterogeneous and
Competencies
(Eligible Content)
CHEM.A.1.2.1:
• Compare properties of
solutions containing
ionic or molecular
solutes (e.g.,
dissolving,
dissociating).
CHEM.A.1.2.2:
• Differentiate between
homogeneous and
heterogeneous
mixtures (e.g., how
such mixtures can be
separated).
CHEM.A.1.2.3:
• Describe how factors
(e.g., temperature,
concentration, surface
area) can affect
solubility.
CHEM.A.1.2.4:
• Describe various ways
that concentration can
be expressed and
calculated (e.g.,
molarity, percent by
mass, percent by
volume).
CHEM.A.1.2.5:
• Describe how chemical
bonding can affect
whether a substance
dissolves in a given
liquid.
Anchor Descriptions
CHEM.A.1.2:
• Compare the properties of
mixtures.
Vocabulary
Solutions and Colligative
Properties
Alloy
Boiling point elevation
Colligative property
Concentration
Freezing point depression
Henry’s Law
Immiscible
Miscible
Molality (M, molal)
Molarity (M, molar)
Saturated solution
Solute
Solution
Solvent
Supersaturated solution
Vapor pressure reduction
People
Henry, William
Textbook
Pages
L2: Ch 18
(p500-526)
Duration
(in days)
8-12
2 CHEMISTRY
homogeneous mixtures.
3.2.C.A2:
• Explain how atoms
combine to form
compounds through both
ionic and covalent
bonding.
3.2.C.A4:
• Predict how combinations
of substances can result in
physical and/or chemical
changes.
2 CHEMISTRY
Acids and Bases
Topic(s)
Acids and Bases,
pH, Neutralization
Reactions, and
Titrations
Concepts
(with Standards)
3.2.10.A2:
• Compare and contrast
different bond types that
result in the formation of
molecules and
compounds.
• Explain why compounds
are composed of integer
ratios of elements.
3.2.10.A4:
• Describe chemical
reactions in terms of
atomic rearrangement
and/or electron transfer.
• Explain the difference
between endothermic and
exothermic reactions.
3.2.12.A4:
• Describe the interactions
between acids and bases.
3.2.C.A2:
• Explain how atoms
combine to form
compounds through both
ionic and covalent
bonding.
• Use the mole concept to
determine number of
particles and molar mass
for elements and
compounds.
• Determine percent
compositions, empirical
formulas, and molecular
formulas.
Competencies
(Eligible Content)
CHEM.B.1.3.1:
• Explain how atoms
combine to form
compounds through
ionic and covalent
bonding.
CHEM.B.1.3.2:
• Classify a bond as
being polar covalent,
non-polar covalent, or
ionic.
CHEM.B.1.4.2:
• Utilize Lewis dot
structures to predict
the structure and
bonding in simple
compounds.
CHEM.B.2.1.4:
• Predict products of
simple chemical
reactions (e.g.,
synthesis,
decomposition, single
replacement, double
replacement,
combustion).
Anchor Descriptions
CHEM.B.1.3:
• Explain how atoms form
chemical bonds.
CHEM.B.1.4:
• Explain how models can be
used to represent bonding.
CHEM.B.2.1:
• Predict what happens during
a chemical reaction.
Vocabulary
Acid- Base Chemistry
Acid
Acid dissociation constant (Ka)
Acid-base indicator solution
Acidic solution
Alkaline
Amphoteric
Arrhenius acid
Arrhenius base
Base
Base dissociation constant (Kb)
Basic solution
BrØnsted-Lowrey Acid
BrØnsted-Lowrey Base
Buffer
Buffering capacity
Common ion
Common ion effect
Concentration
Conjugate acid
Conjugate acid-base pair
Conjugate base
Diprotic acid
End point/ equivalence point
Gram equivalent mass (gem)
H+ acceptor
H+ donor
Hydrogen ion (H+)
Hydronium ion (H3O+)
Hydroxide ion (OH-)
Ion-product constant of water (Kw)
Lewis acid
Lewis base
Molarity (M, molar)
Monoprotic acid
Neutral solution
Neutralization reaction
Neutralize
Textbook
Pages
L2: Ch 20, 21
(p577-638)
Duration
(in days)
5-10
2 CHEMISTRY
3.2.C.A4:
• Interpret and apply the laws
of conservation of mass,
constant composition
(definite proportions), and
multiple proportions.
• Balance chemical
equations by applying the
laws of conservation of
mass.
• Use stoichiometry to
predict quantitative
relationships in a chemical
reaction.
Normality (N)
pH
pOH
Polyprotic acid
Salt
Salt hydrolysis
Self ionization
Solubility constant (Ksp)
Standard solution
Strong acid
Strong base
Titration
Triprotic acid
Weak acid
Weak base
2 CHEMISTRY
Thermochemistry
Topic(s)
Calorimetry and
Thermochemistry
Stoichiometry
Concepts
(with Standards)
Competencies (Eligible
Content)
Anchor Descriptions
3.2.10.A4:
• Describe chemical
reactions in terms of
atomic rearrangement
and/or electron
transfer.
• Explain the difference
between endothermic
and exothermic
reactions.
• Predict the amounts of
products and reactants
in a chemical reaction
using mole
relationships.
CHEM.A.1.1.1:
• Classify physical or
chemical changes
within a system in terms
of matter and/or energy.
CHEM.A.1.1:
• Identify and describe how
observable and measurable
properties can be used to
classify and describe matter
and energy.
3.2.10.B3:
• Explain how heat
energy will move from a
higher temperature to a
lower temperature until
equilibrium is reached.
• Analyze the processes
of convection,
conduction, and
radiation between
objects or regions that
are at different
temperatures.
3.2.C.B2:
• Explore the natural
tendency for systems
to move in a direction
of disorder or
randomness (entropy).
CHEM.B.2.1.1:
• Describe the roles of
limiting and excess
reactants in chemical
reactions.
CHEM.B.2.1.2:
• Use stoichiometric
relationships to
calculate the amounts
of reactants and
products involved in a
chemical reaction.
CHEM.B.2.1.5:
• Balance chemical
equations by applying
the Law of
Conservation of Matter.
CHEM.B.2.1:
• Predict what happens during
a chemical reaction.
Vocabulary
Thermochemistry
Calorie (Cal) and calorie (cal)
Calorimetry
Chemical potential energy
Endothermic
Energy
Enthalpy (H)
Entropy (S)
Exothermic
Free energy (G)
Heat
Heat capacity
Heat of reaction
Hess’s Law
Joule (J)
Law of conservation of energy
(Molar) Heat of combustion (ΔHcomb)
(Molar) Heat of condensation
(ΔHcond)
(Molar) Heat of formation (ΔH°f)
(Molar) Heat of fusion (ΔHfus)
(Molar) Heat of solidification (ΔHsol)
(Molar) Heat of solution (ΔHsoln)
(Molar) Heat of vaporization (ΔHvap)
Potential energy
Specific heat capacity (c)
Standard heat of formation
Surroundings
System
Thermochemical equation
Thermochemistry
Textbook
Pages
L2: Ch 11
(p292-318)
Duration
(in days)
5-10