I.
UNIT 1
UNIT 2
II.
UNIT 3
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
Science and Its Methods
a. Basic characteristics of science (1.3)
 Distinguish between hypotheses, laws, theories, and models
i. Hypothesis – testable and reproducible
ii. Law – summary of many observations
iii. Theory – tentative attempt to explain observations
iv. Model – attempt to explain abstract ideas with more
familiar scale examples
b. The limitations of science (1.4)
 Identify key factors that limit science’s ability to unambiguously
answer all questions
c. Desirability quotient – benefits/risks (1.5)
 Define the desirability quotient and apply it in a given situation
d. Critical thinking – FiLCHeRS approach ( a modification called  Describe the steps in the FiLCHeRS approach to evaluating claims
FLaReS is in Section 1.13)
 Apply the FiLCHeRS approach to given situations
Characteristics of Chemistry- Study of Matter and its Transformations
a. The study of chemistry (Introduction to Chapter 1)
 Define chemistry as the study of matter and the changes it undergoes
b. Properties (1.8)
 Identify physical and chemical properties
i. physical
ii. chemical
c. Transformations (1.8)
States of Matter
 Identify physical and chemical changes
i. physical
ii. chemical
d. Classification of matter (1.9)
 Given a description of matter, classify it as either a pure substance or a
mixture
 If matter is a pure substance, further classify it as either and element or
a compound
 If matter is a mixture, further classify it as either a homogeneous or
heterogeneous mixture
i. States of matter
ii. Pure substance
1. element
2. compound
iii. Mixture
1. Homogeneous
III.
UNIT 4
UNIT 5
IV.
UNIT 6
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
2. Heterogeneous
Measurement
a. The International System (SI) (1.10)
 List the seven SI units in use
 Recognize the kilogram as the only prototype SI unit
i. SI units
ii. Current definitions of seven base units
b. The metric system (1.10)
 Write numbers in exponential notation
 Convert numbers from exponential notation to standard form
 Convert metric units from one prefix to another
i. Review of exponential notation (Appendix A-1)
ii. Basic prefixes and their interconversions
c. Key derived units
 Convert area and volume units
 Define density
 Use the density definition to find missing information
i. Volume (1.10)
ii. Density (1.11)
d. Temperature units and their interconversion (1.12)
 Carry out conversions between the three temperature scales given the
mathematical relationships

e. Energy units and their interconversion (1.12)
 Distinguish between energy and heat
 Convert between various energy units
↑↑↑↑↑WEEK 1 ↑↑↑↑↑
Atoms – History and Early Indications
a. Greek view of atoms (2.1)
 Describe the contrast in views between Aristotle and Democritus in
relation to matter
b. Fundamentals laws that led to early conclusions about atoms
 State the law of conservation of mass
 State the law of definite proportions
 State the law of multiple proportions
 Apply the law of multiple proportions in straightforward numerical
problems
i. Lavoisier: The law of conversation of mass (2.2)
ii. Proust: The law of definite proportions (2.3)
UNIT 8
VI.
UNIT 7
V.
UNIT 9
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
iii. Dalton: The law of multiple proportions (2.3)
c. John Dalton’s atomic theory of matter (2.4)
 State the four premises of Dalton’s atomic theory
 Suggest in a given situation which of Dalton’s premise may be capable
of explaining observed behavior
 Use proportions to carry out straightforward mass calculations for
chemical compounds
i. Four points of Dalton’s atomic theory (2.4)
ii. Physical phenomena explained by Dalton’s atomic
theory (2.4)
iii. Points in Dalton’s theory modified by more current
findings (2.4)
iv. Solving mass and atom ratios using Dalton’s approach
(2.4)
The Periodic Table
a. Historic development of the periodic table (2.5)
 Identify the primary individual associated with the development of
today’s periodic table
 State the historical driving force behind creating the periodic table in its
current arrangement
i. Dobereiner’s “Triads”
ii. De Chancourtois’s “Telluric Helix”
iii. Newland’s “Law of Octaves”
iv. Meyer’s system of elements
v. Mendeleev’s version
b. Current state – number of elements, physical arrangement, etc.
PeriodicTableLive!
 State the current number of chemical elements known (this semester)
(2.5)
↑↑↑↑↑WEEK 2 ↑↑↑↑↑
Historical Physical Evidence for Atoms
a. Relationship between electricity and the atom (3.1)
 Describe two of the electrical experiments, the workers involved, and
the conclusions drawn regarding atomic structure
i. Electrolysis (Davy and Faraday) – atoms electrical in
nature
ii. Cathode ray tubes (Crookes) –stream of particles leaves
cathode
UNIT 10
UNIT 11
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
iii. Determination of mass-to-charge ratio (Thomson) –
particles leaving cathode are negatively charged –
named the particles electrons
iv. Discovery of positive particles (Goldstein) – positive
particles also produced in cathode ray tubes when
negative particles are produced; positive particles are
called protons
v. Determination of electron charge (Millikan) – found
charge on electron in oil-drop experiment; also
provided mass because of Thomson’s mass-to-charge
ratio
b. The arrangement of particles in the nucleus
 List the three subatomic particles and their charges and relative masses
 Describe Rutherford’s scattering experiment and its implications for
atomic structure
i. Three fundamental particles (3.5)
1. Proton – positive, relatively massive
2. Neutron – neutral, relatively massive (slightly
more than the proton)
3. Electron – negative , relatively light (about
1/2000th of the other two particles)
ii. Rutherford - small positive center surrounded by
Build-an-atom
electrons in mostly empty space (Rutherford) (3.4)
Rutherford
scattering
c. The atomic nucleus (3.5)
 Identify an element by the number of protons in the nucleus
 Define the term isotope
 Define the term mass number
 Write the symbolic representation for an isotope given sufficient
information
 Given the symbolic representation for an isotope provide missing
information
i. Atomic number (Z) – identifies element and indicates
number of protons in the nucleus
ii. Isotopes – two nuclei of the same element that differ in
UNIT 12
UNIT 13
VII.
UNIT 14
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
the number of neutrons in the nucleus; mass number
(A) is sum of protons (Z) and neutrons
iii. Symbolic representation of isotopes - ZA X where X is
the chemical symbol
d. Electron arrangement – the Bohr model (3.6)
 Relate color of light to wavelength and energy
 Describe the electronic processes that lead to the color in line spectra
 Determine the maximum number of electrons in a given shell
i. Relationship between wavelength and energy
ii. Bohr’s description of the relationship of colors emitted
in spectra to electron energy levels
e. Electron arrangement – the quantum model (3.7)
 Write electron configurations for elements in the first three periods
i. The concept of an orbital
ii. Writing electron configurations for atoms
f. Relationship between electron configurations and the periodic
 Use the periodic table to identify the number of valence electrons in an
table
atom of an element
 Use the periodic table to describe the location of groups, periods, main
group elements, transitions elements, metals, nonmetals, alkali metals,
alkaline earth metals, and halogens
i. Identification of valence electrons with group numbers
ii. Identification of named groups and regions in the
periodic table
↑↑↑↑↑WEEK 3 ↑↑↑↑↑
Chemical Bonds (Chapter 4)
a. Identifying criteria for stable electron configurations (4.1)
b. Lewis symbols and ionic compounds
 Draw a Lewis symbol for a given atom or ion
 State the octet rule
 Write formulas for binary ionic compounds
 Name binary ionic compounds
i. Lewis symbols for individual atoms (4.2)
ii. Lewis forms for ionic compounds (4.3-4.4)
iii. The octet rule (4.4)
iv. Writing formulas for and naming binary ionic
UNIT 15
UNIT 16
VIII.
UNIT 17
U
NI
T
18
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
compounds (4.5)
c. Lewis structures and covalent species
 Draw Lewis structures for covalent molecules and ions
 Write formulas for a select set of polyatomic ions
 Name a select set of polyatomic ions
 Name covalent compounds
i. Writing Lewis structures for covalent bonds (4.6-4.8)
ii. Rules for writing Lewis structures (4.10)
iii. Identify and work with polyatomic ions (4.9)
iv. Name covalent compounds (4.6, 4.9)
d. Relating shape and polarity to Lewis structure
 Determine the molecular shape of covalent compounds from their
Lewis structures
 State whether a given molecule is nonpolar or polar
i. Relating Lewis structure to molecular shape – The
VSEPR Theory (4.12)
ii. Relate shape to polarity of molecules (4.13)
↑↑↑↑↑WEEK 4 ↑↑↑↑↑
Chemical Accounting
a. Chemical equations (5.1)
 Identify reactants and products in a chemical equation
 Convert a verbal description of a reaction to a balanced chemical
equation
 Balance chemical equations
i. Identifying reactants and products
ii. Balancing chemical equations
b. Volume relationships in chemical equations (5.2)
 State Gay-Lussac’s law of combining volumes
 Apply Gay-Lussac’s law of combining volumes to a chemical reaction
scenario
 State Avogadro’s hypothesis regarding the numbers of molecules in
equal volumes of gases
i. Law of combining volumes (Gay-Lussac)
ii. Avogadro’s hypothesis (Avogadro)
c. Introduction to Avogadro’s number and the mole (5.3)
 Give the significance of Avogadro’s number
 Give Avogadro’s number
UNIT 19
UNIT 20
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
 Relate formula mass to numbers of moles
i. Identification of Avogadro’s number and its
significance
ii. Definition of the term mole and its relationship to
Avogadro’s number
iii. Introduction to formula (or molecular mass for
molecules) mass and its relationship to Avogadro’s
number and the mole
d. Working with Avogadro’s number, mass, and the mole (5.4)
 Interconvert between grams, moles, and numbers of particles
 Apply the mass and mole relationships to chemical reactions
i. Mole-to-mass and mass-to-mole conversions
ii. Relationships in chemical equations
1. Mole relationships in chemical equations
2. Mass relationships in chemical equations
e. Solutions – Qualitative (5.5)
 Define the term solution
 Define the qualitative terms unsaturated, saturated, and supersaturated,
dilute, and concentrated
 Classify a given solution description according to the above terms
i. Definition of a solution
ii. Qualitative terms related to solutions
1. Unsaturated
2. Saturated
3. Supersaturated
4. Dilute
5. Concentrated
f. Solutions – Quantitative (5.5)
 Define the terms molarity, percent by mass, and percent by volume
 Determine missing information in solution problems
i. Definition of molarity and its applications
1. Relating molarity, moles, and volume of
solution
2. Relating molarity, mass, and volume of
solution
3. Preparation of solutions of given molarity
IX.
UNIT 21
UNIT 22
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
ii. Percent concentrations
1. Percent by volume
2. Percent by mass
↑↑↑↑↑WEEK 5 ↑↑↑↑↑
Gases, Liquids, Solids, and Intermolecular Forces (Chapter 6)
a. Solids, liquids, and gases (6.1)
 Identify molecular level differences between solids, liquids, and gases
 Define the terms associated with phase transitions – boiling point,
melting point, freezing point, condensation, freezing and sublimation
 Describe the physical meaning of term critical point
i. Identification of molecular level differences
States of Matter
ii. Identification of terminology associated with changes
from one phase to another
1. Melting (freezing)
2. Boiling (condensing)
3. Sublimation (deposition or condensation)
iii. Supercritical fluids
b. Comparison of properties of ionic and molecular compounds
 State key physical differences between ionic and molecular compounds
(6.2)
i.
Physical state at room temperature
ii.
Relative melting points
iii.
Energy required to melt
iv.
Conductivity in water solution
v.
Brittleness
c. Forces between molecules (intermolecular forces) (6.3)
 Define the three primary types of intermolecular forces
 Identify situations in which each of the intermolecular forces is
important
 Given a species identify the intermolecular forces which play a role in
that species
i. Dipole-dipole forces
ii. Dispersion forces
iii. Hydrogen bonding
d. Forces in solution (between solute and solvent) (6.4)
 Identify the key forces that determine the solubility of a solute in a
solvent
Unit 23
Unit 24
UNIT
25
X.
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
e. Gases: The kinetic-molecular theory (6.5)
Gas Properties
 State the five postulates of the kinetic molecular theory of gases
 Describe how the kinetic molecular theory accounts for observed
properties of gases
i. Atomic level particles in rapid, constant motion, and
move in straight lines
ii. Particles are tiny compared to volume of the container
iii. Very little attraction between particles
iv. Energy is conserved when particles collide
v. Temperature is a measure of the average kinetic energy
of gas molecules
f. Simple gas laws (6.6)
 State Boyle’s Law
 State Charles’s Law
 State the combined gas law
 Apply Boyle’s Law, Charles’s Law, and the combined gas law to gases
in a variety of situations
iv. Introduction to pressure measurement and units
v. Boyle’s Law: relates pressure and volume at constant
temperature and number of moles of gas
vi. Charles’s Law: relates temperature and volume at
constant pressure and number of moles of gas
g. Ideal gas law (6.7)
 State the ideal gas law
 Apply the ideal gas law to gases in a variety of situations
 State standard temperature and pressure conditions (STP)
 State the molar volume of an ideal gas at STP
i. Combined gas law: relates pressure, volume, and
temperature at constant number of moles of gas
ii. Ideal gas law: relates pressure, volume, temperature,
and number of moles of gas
iii. Molar volume at standard temperature and pressure
↑↑↑↑↑WEEK 6 ↑↑↑↑↑
Acids and Bases (Chapter 7)
a. Experimental definitions of acids and bases (7.1)
 Give physical characteristics of acids and bases
b. Acids, bases, and salts (7.2)
 State the Arrhenius theory of acids and bases
c.
d.
f.
UNIT 27
g.
h.
UNIT 26
e.
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
 Identify acids and bases using the Arrhenius theory
 State the Brønsted-Lowry theory of acids and bases
 Identify acids and bases using the Brønsted-Lowry theory
 Identify salts
 Identify the specific acids and bases that produce a given salt
i. Arrhenius theory – acid is proton donor in aqueous
solution, base is hydroxide donor in aqueous solution
ii. Brønsted-Lowry Acid-Base Theory – acid is proton
donor, base is proton acceptor
iii. Salt – product of neutralization reaction between acid
and base
Acidic and basic anhydrides (7.3)
 Identify oxides as acidic or basic
i. Nonmetal oxides – acidic anhydrides
ii. Metallic oxides – basic anhydrides
Strong and weak acids and bases (7.4)
 Describe the difference between strong and weak acids
 Describe the difference between strong and weak bases
 Identify hydrochloric, sulfuric, and nitric acid as strong acids
i. Define the terms strong and weak acids and bases
ii. Identify strong and weak acids and bases
Neutralization of acids and bases (7.5)
 Complete the chemical equation for a neutralization reaction given the
reactants
The pH scale (7.6)
pH Scale
 Recognize the pH scale is logarithmic
 Identify pH regions which are acidic, neutral, and basic
 Estimate pH from hydrogen concentration and hydrogen concentration
from pH
Buffers and conjugate acid-base pairs (7.7)
 Identify conjugate acid-base pairs
Applications (7.8-7.10)
i. Acid rain
ii. Antacids
iii. Use in industry
iv. Use in health and disease issues
↑↑↑↑↑WEEK 7 ↑↑↑↑↑
XI.
UNIT 28
UNIT 29
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
Oxidation and Reduction (Chapter 8)
a. Three views of oxidation-reduction (8.1)
 State the definitions of oxidation and reduction considering oxygen
atoms
 State the definitions of oxidation and reduction considering hydrogen
atoms
 State the definitions of oxidation and reduction considering the transfer
of electrons
 Apply the definitions of oxidation and reduction to identify species
oxidized and reduced in a given chemical equation
i. Oxidation – gain of oxygen atoms; reduction is the loss
of oxygen atoms
ii. Oxidation – loss of hydrogen atoms; reduction is the
gain of hydrogen atoms
iii. Oxidation – loss of electrons; reduction is the gain of
electrons
b. Oxidizing and reducing agents (8.2)
 Identify oxidizing and reducing agents in a given chemical equation
i. Oxidizing agent – substance that is reduced
ii. Reducing agent – substance that is oxidized
c. Electrochemistry: cells and batteries (8.3)
 Given a redox reaction break it into two half-reactions and identify each
as oxidation or reduction
 Given two unbalanced half-reactions, balance them and combine them
into one overall reaction
 For a given electrochemical cell identify the anode and the cathode
i. Separating reactions into two half-reactions
ii. Balancing half-reactions and combining into one
reaction
iii. Terminology: electrochemical cell, electrodes,
cathode, anode
iv. Descriptions of basic cells and batteries
d. Corrosion (8.4)
 Identify the role of oxidation-reduction in corrosion processes
 Describe the function of a sacrificial anode
i. The rusting of iron
ii. Protecting materials from corrosion
XII.
UNIT 30
UNIT 31
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
e. Common oxidizing agents (8.7)
 Write balanced chemical equations for the reaction of oxygen with
elements and compounds
 Identify ozone as an allotrope of oxygen
 Describe ozone’s dual role in environmental issues
 Describe some of the uses of chlorine as an oxidizing agent
f. Common reducing agents (8.8)
 Give examples of hydrogen as a reducing agent in applications
Nuclear Chemistry (Chapter 11)
a. Exposure to natural radioactivity (11.1)
 Identify the three natural sources of radiation exposure
b. Writing nuclear equations (11.2)
 State the three most common types of radiation and their characteristics
– alpha particles, beta particles, and gamma rays
 Complete given nuclear equations by supplying particles, mass
numbers, and/or atomic numbers
i. Identification of basic subatomic particles
ii. Review of writing atomic symbols with atomic
numbers and mass numbers
iii. Complete nuclear reactions by ensuring the sum of
atomic numbers and the sum of mass numbers is the
same on each side of the equation
c. Half-life (11.3)
 Use half-life information to determine quantities of material remaining
after specific lengths of time
 Given sufficient information, determine the age of a specimen
i. Describe constancy for a specimen
ii. Relate to radioisotopic dating (11.4)
Radioactive Dating
Game
d. Artificial transmutations (11.5)
 Recognize artificial transmutation as man-made process for producing
new elements
e. Uses of radioisotopes (11.6)
 Identify three radioisotopes and describe their applications
f. Penetrating power of radiation (11.7)
 Rate alpha particles, beta particles, and gamma rays in terms of their
penetrating power
g. Energy from the nucleus (11.8)
 Describe key aspects that lead to the nuclear fission process
 Describe three key technical issues that had to be overcome to build the
first nuclear weapon
2
i. E = mc
CHEM 1004
Descriptive Chemistry
Course Outline
(Numbers in parentheses are section numbers in the course textbook, Chemistry for Changing Times, John W. Hill, Terry W. McCreary, Doris K. Kolb, Prentice Hall, 2010
ISBN: 0-13-605449-8)
Underlined activities from PhET
Assignment
Topic
Learning Objectives
Activities
Topics
ii. Binding energy
iii. Nuclear fission
iv. Conditions for chain reactions
h. Radioactive Fallout (11.10)
 Describe the reason for prolonged radioactive fallout from a release
i. Nuclear power plants (11.11)
 State a key difference between the technology involved in nuclear
power plants compared to nuclear weapons
 State two key problems to be faced with nuclear power plants
j. Thermonuclear reactions (11.12)
 Distinguish between nuclear fission and nuclear fission
↑↑↑↑↑WEEK 8 ↑↑↑↑↑