1 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)
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.
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
Competencies (Eligible
Content)
CHEM.A.1.1.1:
• Classify physical or
chemical changes within a
system in terms of matter
and/or energy.
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
Textbook
Pages
Duration
(in days)
The Nature of Chemistry
Analytical chemistry
Biochemistry
Chemistry
Inorganic chemistry
Organic chemistry
Physical chemistry
Science
Technology
L1: Ch 1 (p3-27)
2-5
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
L1 and L2:
Teacher and
department
generated
materials; Flinn
Scientific Safetly
Contract
2-5
1 CHEMISTRY
•
advance a solution.
Communicate and
defend a scientific
argument.
Evaporating dish
Funnel
Hot plate
Laboratory apron
Micropipettes
Pasteur pipettes
Petri dish
Ring stand
Rubber policeman
Rubber stoppers
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
L1: C1 (p11-16)
5-10
1 CHEMISTRY
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)
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)
1 CHEMISTRY
Matter & Energy
Topic(s)
Concepts
(with Standards)
Competencies
(Eligible Content)
Anchor Descriptions
Matter & Energy
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.
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.B2:
• Explain how the overall
energy flowing through a
system remains constant.
CHEM.A.1.1.3:
• Utilize significant figures to
communicate the uncertainty
in a quantitative observation.
3.2.12.A1:
• Compare and contrast
colligative properties of
mixtures.
• Compare and contrast the
unique properties of water
to other liquids.
CHEM.A.1.1.4:
• Relate the physical
properties of matter to its
atomic or molecular
structure.
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
CHEM.A.1.1.2:
• Classify observations as
qualitative and/or
quantitative.
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, dissociating).
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
Potential energy
Textbook
Pages
L1: Ch 2 (p20-41);
Ch 5 (p135-141),
and Ch 10 (p288293)
Duration
(in days)
8-15
1 CHEMISTRY
disorder or randomness
(entropy).
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.
Pure substance
Radiant energy
Solid
Solution
Sublimation
Suspension
Symbol
Triple point
Unit cell
Vapor
Vapor pressure
Vaporization
1 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.
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, distribution of
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
Textbook
Pages
Duration
(in days)
EMR
Absorption spectrum
Amplitude
Color
Electromagnetic radiation
Electromagnetic spectrum
Emission spectrum
Energy level
Excited state
Frequency (ν or or ν)
Ground state
Heisenberg uncertainty principle
Matter-wave
Photoelectric effect
Photon
Plank’s constant (h)
L1: Ch 11
(p324-327)
10
L1: Ch 3
(p46-78);
Ch 11
(p322-323)
6-7
Principle quantum number (n)
Quantized
Quantum
Spectrum
Speed of light (c)
Wavelength (λ)
People
Bohr, Neils
DeBroglie, Louis
Einstein, Albert
Heisenberg, Werner
Planck, Max
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
Compound
1 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.
electrons in orbitals,
shapes of orbitals).
CHEM.A.2.2.4:
• Relate the existence of
quantized energy levels to
atomic emission spectra.
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)
L1: Ch 19
(p606-627)
5-10
1 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
L1: Ch 11
(p328-352)
10-15
1 CHEMISTRY
Einstein, Albert
Hund, Friedrich
Pauli, Wolfgang
1 CHEMISTRY
The Periodic Table
Topic(s)
Concepts
(with Standards)
The Periodic Table
and Periodic
Trends
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 to the
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).
Anchor
Descriptions
CHEM.A.2.2:
• Describe the
behavior of
electrons in atoms.
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, distribution
of 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:
• Explain how
periodic trends in
the properties of
atoms allow for the
prediction of
physical and
chemical properties.
Vocabulary
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
People
Dobereiner, Johann
Erdmann, Hugo
Janssen, Pierre
Lockyer, Joseph Norman
Mendeleev, Dmitri
Meyer, Julius Lothar
Moseley, Henry
Newlands, John Alexander Reina
Textbook
Pages
L1: Ch 11 (p339352)
Duration
(in days)
5-8
1 CHEMISTRY
CHEM.A.2.3.2:
• Compare and/or
predict the properties
(e.g., electron affinity,
ionization energy,
chemical reactivity,
electronegativity,
atomic radius) of
selected elements by
using their locations
on the periodic table
and known trends.
Ramsey, William
Rayleigh, Lord (John William Strutt)
1 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.
3.2.12.A5:
• Use VSEPR theory to predict the
molecular geometry of simple
molecules.
3.2.12.B4:
• Describe conceptually the attractive
and repulsive forces between
objects relative to their charges and
the distance between them.
3.2.C.A1:
• Use electro-negativity to explain the
difference between polar and nonpolar covalent bonds.
3.2.C.A2:
• Explain how atoms combine to form
Competencies (Eligible
Content)
CHEM.B.1.3.1:
• Explain how atoms
combine to form
compounds through ionic
and covalent bonding.
Anchor
Descriptions
CHEM.B.1.3:
• Explain how
atoms form
chemical bonds.
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.
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, balland-stick models,
graphical models, solidsphere models, structural
formulas, skeletal
formulas, Lewis dot
structures).
CHEM.B.1.4.2:
• Utilize Lewis dot
structures to predict the
structure and bonding in
CHEM.B.1.4:
• Explain how
models can be
used to represent
bonding.
Vocabulary
Textbook
Pages
Duration
(in days)
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
L1: Ch 12 (p358382)
7-12
L1: Ch 12 (p358392)
5-10
People:
Dalton, John
Lewis, Gilbert
Molecular Geometry
Bent
Dipole
Hybrid orbitals
1 CHEMISTRY
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.
simple compounds.
CHEM.B.2.1.4:
• Predict products of
simple 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.
CHEM.B.2.1: Predict
what happens during a
chemical reaction.
Ionic compound
Linear
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
L1: Ch 14 (p449461)
7-12
L1: Ch 4 (p84106)
5-10
1 CHEMISTRY
Anion
B group
Base
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
monoditritetrapentahexaheptaoctanonadeca-
1 CHEMISTRY
Polyatomic Ions
Acetate
Ammonium
Bisulfate
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
L1: Ch 7, Ch 8
(p192-244)
10-12
1 CHEMISTRY
Synthesis reaction/ addition
reaction/ direct combination
reaction
1 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 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.
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:
• Apply the mole
concept to the
composition of
matter.
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
together in a compound
(e.g., computer models,
ball-and-stick models,
graphical models, solidsphere models, structural
formulas, skeletal formulas,
Lewis dot structures).
Vocabulary
Chemical Quantities: The Mole
and Stoichiometry
Actual yield/ experimental yield
Atomic mass unit
Avogadro’s number
Empirical formula
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
CHEM.B.1.4:
• Explain how models
can be used to
represent bonding.
Textbook
Pages
L1: Ch 6
(p152-185);
Ch9 (p251289)
Duration
(in days)
15-20
1 CHEMISTRY
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.
CHEM.B.2.1:
• Predict what
happens during a
chemical reaction.
1 CHEMISTRY
The Kinetic Molecular Theory
Topic(s)
Gas Laws
Concepts
(with Standards)
Competencies (Eligible
Content)
Anchor Descriptions
3.2.10.A3:
• Describe phases of matter
according to the kinetic
molecular theory.
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:
• Explain how the
kinetic molecular
theory relates to the
behavior of gases.
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.
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.
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
L1: Ch 13
(p398-433)
Duration
(in days)
10-12
1 CHEMISTRY
Solutions and Colligative Properties
Topic(s)
Solutions,
Concentrations,
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
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
L1: Ch 15
(p466-496)
Duration
(in days)
8-12
1 CHEMISTRY
heterogeneous and
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.
1 CHEMISTRY
Acids and Bases
Topic(s)
Acids and Bases,
pH, Neutralization
Reactions, and
Titrations
Concepts
(with Standards)
Competencies (Eligible
Content)
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.
CHEM.B.1.3.1:
• Explain how atoms combine
to form compounds through
ionic and covalent bonding.
3.2.10.A4:
• Describe chemical
reactions in terms of atomic
rearrangement and/or
electron transfer.
• Explain the difference
between endothermic and
exothermic reactions.
CHEM.B.1.4.2:
• Utilize Lewis dot structures to
predict the structure and
bonding in simple
compounds.
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.
Anchor Descriptions
CHEM.B.1.3:
• Explain how atoms
form chemical bonds.
CHEM.B.1.3.2:
• Classify a bond as being
polar covalent, non-polar
covalent, or ionic.
CHEM.B.2.1.4:
• Predict products of simple
chemical reactions (e.g.,
synthesis, decomposition,
single replacement, double
replacement, combustion).
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
Normality (N)
Textbook
Pages
L1: Ch 16
(p509-529)
Duration
(in days)
5-10
1 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.
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
1 CHEMISTRY
Thermochemistry
Topic(s)
Calorimetry and
Thermochemistry
Stoichiometry
Concepts
(with Standards)
Competencies (Eligible
Content)
Anchor Descriptions
Vocabulary
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.
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
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.
Textbook
Pages
L1: Ch 10
(p289-315)
Duration
(in days)
5-10
1 CHEMISTRY
Electrochemistry
Topic(s)
Concepts
(with Standards)
Competencies (Eligible
Content)
Anchor Descriptions
Oxidation-Reduction
(RedOx)
3.2.10.A4:
• Describe chemical
reactions in terms of
atomic rearrangement
and/or electron transfer.
CHEM.A.1.1.1:
• Classify physical or chemical
changes within a system in
terms of matter and/or
energy.
3.2.12.A4:
• Apply oxidation/reduction
principles to
electrochemical
reactions.
CHEM.B.2.1.1:
• Describe the roles of limiting
and excess reactants in
chemical reactions.
CHEM.A.1.1:
• Identify and describe
how observable and
measurable
properties can be
used to classify and
describe matter and
energy.
Electrochemistry
3.2.12.B4:
• Describe conceptually
the attractive and
repulsive forces between
objects relative to their
charges and the distance
between them.
3.2.C.A1:
• Use electro-negativity to
explain the difference
between polar and nonpolar covalent bonds
3.2.C.A2:
• Explain how atoms
combine to form
compounds through both
ionic and covalent
bonding.
• Predict the chemical
formulas for simple ionic
and molecular
compounds.
3.2.C.A4:
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
Oxidation- Reduction and
Electrochemistry
Anode
Battery
Cathode
Cell potential
Dry cell
Electrical cell
Electrical potential
Electrochemical cell
Electrochemical process
Electrode
Electrolysis
Fuel cells
Half cell
Half reactions
Half-reaction method
Oxidation
Oxidation number
Oxidation number change method
Oxidation-reduction reaction
Oxidizing agent
Red-ox reaction
Reducing agent
Reduction
Reduction potential
Salt bridge
Standard cell potential
Standard hydrogen electrode
Voltaic cell
Wet cell
Textbook
Pages
L1: Ch 18
(p574-600)
Duration
(in days)
5-15
1 CHEMISTRY
• Predict how
combinations of
substances can result in
physical and/or chemical
changes.
1 CHEMISTRY
Kinetics
Topic(s)
Chemical Kinetics,
Reaction Rates, and
Equilibrium
Concepts
(with Standards)
3.2.10.A4:
• Explain the difference
between endothermic
and exothermic
reactions.
• Identify the factors that
affect the rates of
reactions.
3.2.12.A5:
• Predict the shift in
equilibrium when a
system is subjected to a
stress.
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.A4:
• Predict how
combinations of
substances can result in
physical and/or chemical
changes.
Competencies (Eligible
Content)
Anchor Descriptions
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.
CHEM.A.1.2.3:
• Describe how factors (e.g.,
temperature, concentration,
surface area) can affect
solubility.
CHEM.B.1.3.1:
• Explain how atoms combine
to form compounds through
ionic and covalent bonding.
CHEM.B.2.1.4:
• Predict products of simple
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.
CHEM.A.1.2:
• Compare the
properties of
mixtures.
CHEM.B.1.3:
• Explain how atoms
form chemical
bonds.
CHEM.B.2.1:
• Predict what
happens during a
chemical reaction.
Vocabulary
Kinetics
Activated complex
Activation energy (Ea)
Biological inhibitor
Catalyst
Chemical equilibrium
Collision theory
Elementary reaction
Entropy (S)
Enzyme
Equilibrium constant (Keq)
Equilibrium position
First order reaction
(Gibbs) Free energy (G)
Inhibitor
Intermediate
LeChatelier’s principle
Non-spontaneous reaction
Rate
Rate law
Reaction mechanism
Reversible reaction
Second order reaction
Specific rate constant
Spontaneous reaction
Transition state (ŧ)
People
Gibbs, Josiah William
LeChatelier, Henry Louis
Braun, Karl Ferdinand
Textbook
Pages
L1: Ch 17
(p536-556)
Duration
(in days)
5-10