a curriculum document based on the sas framework might ... · pdf filequalitative and/or...
TRANSCRIPT
2 CHEMISTRY The Nature of Chemistry
Topic(s) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
The Scientific Method Laboratory Procedure(s) and Safety The Metric System, Significant Figures and Measurement
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 advance a solution.
• Communicate and defend a
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.
CHEM.A.1.1: • Identify and describe how
observable and measurable properties can be used to classify and describe matter and energy.
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
L2: Ch 1 (p2-
16) L1 and L2:
Teacher and department generated materials; Flinn Scientific Safely Contract
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 (p15-
20)
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) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
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.
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.1.1.1: • 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.
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,
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.
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
L2: Ch 2 (p28-
45), Ch 3 (p74-76), Ch 10 (p274-287)
8-15
2 CHEMISTRY 3.2.10.B3: • Explain how heat energy will
move from a higher temperature to a lower temperature until equilibrium is reached.
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.
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) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
EMR Atomic Structure and History Nuclear Chemistry and Radioactivity Electron Configurations
3.2.10.A5: • Describe the historical
development of models of the atom and how they contributed to modern atomic theory.
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.2.1.1: • Describe the evolution of
atomic theory leading to the current model of the atom based on the works of Dalton, Thomson, Rutherford, and Bohr.
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,
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.
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) 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
L2: Ch 13
(p372- 384) L2: Ch 5 (p106-
122)
10 6-7
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) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
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 of other elements in the table.
• Predict chemical formulas based on the number of valence electrons.
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,
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.
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
L2: Ch 5 (p123-
127); Ch 14 (p390-407)
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) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
Lewis Structures and Bonding Molecular Geometry Intermolecular Forces (IMFs) Nomenclature Reaction Types, Predicting, and Writing Products
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 non-polar covalent bonds.
3.2.C.A2:
• Explain how atoms combine to
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.
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).
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
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.
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 People: Dalton, John Lewis, Gilbert Molecular Geometry Bent Dipole Hybrid orbitals Ionic compound Linear
L2: Ch 6 (p133-
160); Ch 15 (p418-452)
L2: Ch 16 (p358-
392)
7-12 5-10
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 (Dipole-
dipole forces or dipole-dipole interactions)
Vapor pressure Viscosity Nomenclature Acid Alkali metal Alkaline earth metal Anion B group Base
L2: Ch 16 (p460-
468); Ch 17 (p475-481)
L2: Ch 6 (p149-
164)
7-12 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 mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- Polyatomic 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 (p202-
278)
10-12
2 CHEMISTRY The Mole
Topic(s) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
The Mole Empirical & Molecular Formulas, Percent Composition Stiochiometry: Limiting Reactant, Percent Yield, and Theoretical Yield
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.
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.1.2.1: • 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.1.1: • Explain how the mole is a
fundamental unit of chemistry.
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.
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
L2: Ch7 (p170- 196); Ch9 (p251-289)
15-20
2 CHEMISTRY together in a
compound (e.g., computer models, ball-and-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) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
Gas Laws
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.
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.
CHEM.B.2.2: • Explain how the kinetic
molecular theory relates to the behavior of gases.
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
L2: Ch 10
(p266-274); Ch 12 (p326-354)
10-12
2 CHEMISTRY Solutions and Colligative Properties
Topic(s) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
Solutions, Concentra-tions, and Colligative Properties
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
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.
CHEM.A.1.2: • Compare the properties of
mixtures.
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
L2: Ch 18
(p500-526)
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) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
Acids and Bases, pH, Neutralization Reactions, and Titrations
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.
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).
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.
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
L2: Ch 20, 21
(p577-638)
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) Concepts (with Standards)
Competencies (Eligible Content)
Anchor Descriptions Vocabulary Textbook Pages
Duration (in days)
Calorimetry and Thermochemistry Stoichiometry
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.
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.A.1.1.1: • Classify physical or
chemical changes within a system in terms of matter and/or energy.
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.A.1.1: • Identify and describe how
observable and measurable properties can be used to classify and describe matter and energy.
CHEM.B.2.1: • Predict what happens during
a chemical reaction.
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
L2: Ch 11
(p292-318)
5-10