CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
I.
II.
Distinguish between hypotheses, laws, theories, and models




Identify key factors that limit science’s ability to unambiguously
answer all questions
Define the desirability quotient and apply it in a given situation
Describe the steps in the FiLCHeRS approach to evaluating claims
Apply the FiLCHeRS approach to given situations


Define chemistry as the study of matter and the changes it undergoes
Identify physical and chemical properties

Identify physical and chemical changes

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
2. Heterogeneous
UNIT 3

UNIT 2
Desirability quotient – benefits/risks (1.5)
Critical thinking – FiLCHeRS approach ( a modification called
FLaReS is in Section 1.13)
Characteristics of Chemistry- Study of Matter and its Transformations
a. The study of chemistry (Introduction to Chapter 1)
b. Properties (1.8)
i. physical
ii. chemical
c. Transformations (1.8)
i. physical
ii. chemical
d. Classification of matter (1.9)
c.
d.

Unit
#
UNIT 1
Science and Its Methods
a. Basic characteristics of science (1.3)
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)
Learning Objectives
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
III.
c.
i. Review of exponential notation (Appendix A-1)
ii. Basic prefixes and their interconversions
Key derived units
d.
i. Volume (1.10)
ii. Density (1.11)
Temperature units and their interconversion (1.12)
e.
Energy units and their interconversion (1.12)
Atoms – History and Early Indications
a. Greek view of atoms (2.1)
b.
Fundamentals laws that led to early conclusions about atoms



Write numbers in exponential notation
Convert numbers from exponential notation to standard form
Convert metric units from one prefix to another



Convert area and volume units
Define density
Use the density definition to find missing information

Carry out conversions between the three temperature scales given the
mathematical relationships








Distinguish between energy and heat
Convert between various energy units
Describe the contrast in views between Aristotle and Democritus in
relation to matter
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
UNIT 6
i. Lavoisier: The law of conversation of mass (2.2)
ii. Proust: The law of definite proportions (2.3)
iii. Dalton: The law of multiple proportions (2.3)
List the seven SI units in use
Recognize the kilogram as the only prototype SI unit
UNIT 5
b.
i. SI units
ii. Current definitions of seven base units
The metric system (1.10)


Unit
#
UNIT 4
IV.
Measurement
a. The International System (SI) (1.10)
Learning Objectives
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
c.
John Dalton’s atomic theory of matter (2.4)
Learning Objectives




State the current number of chemical elements known (this semester)

Describe two of the electrical experiments, the workers involved, and
the conclusions drawn regarding atomic structure
UNIT 9
i. Electrolysis (Davy and Faraday) – atoms electrical in nature
ii. Cathode ray tubes (Crookes) –stream of particles leaves
cathode
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

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
UNIT 8
VI.
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. (2.5)
Historical Physical Evidence for Atoms
a. Relationship between electricity and the atom (3.1)

UNIT 7
V.
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)
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
Unit
#
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
Learning Objectives
The arrangement of particles in the nucleus
c.
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 electrons in
mostly empty space (Rutherford) (3.4)
The atomic nucleus (3.5)


List the three subatomic particles and their charges and relative masses
Describe Rutherford’s scattering experiment and its implications for
atomic structure




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
UNIT 10
b.
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 the
number of neutrons in the nucleus; mass number (A) is sum
of protons (Z) and neutrons
d.
X where X is the
chemical symbol
Electron arrangement – the Bohr model (3.6)
i. Relationship between wavelength and energy
ii. Bohr’s description of the relationship of colors emitted in
spectra to electron energy levels
Electron arrangement – the quantum model (3.7)
i. The concept of an orbital
ii. Writing electron configurations for atoms



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

Write electron configurations for elements in the first three periods
UNIT 12
e.
A
Z
UNIT 11

iii. Symbolic representation of isotopes -
Unit
#
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
f.
i. Identification of valence electrons with group numbers
ii. Identification of named groups and regions in the periodic
table
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




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

Determine the molecular shape of covalent compounds from their
Lewis structures
State whether a given molecule is nonpolar or polar
i.
ii.
iii.
iv.

i. Relating Lewis structure to molecular shape – The VSEPR
Theory (4.12)
ii.
Relate shape to polarity of molecules (4.13)
UNIT 16




UNIT 15
d.
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)
Relating shape and polarity to Lewis structure

Use the periodic table to identify the number of valence electrons in an
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
UNIT 14
c.
Lewis symbols for individual atoms (4.2)
Lewis forms for ionic compounds (4.3-4.4)
The octet rule (4.4)
Writing formulas for and naming binary ionic compounds
(4.5)
Lewis structures and covalent species

Unit
#
UNIT 13
VII.
Relationship between electron configurations and the periodic table
Learning Objectives
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
VIII.
Chemical Accounting
a. Chemical equations (5.1)
Learning Objectives






d.
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
Working with Avogadro’s number, mass, and the mole (5.4)



Give the significance of Avogadro’s number
Give Avogadro’s number
Relate formula mass to numbers of moles


Interconvert between grams, moles, and numbers of particles
Apply the mass and mole relationships to chemical reactions
UNIT 19
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
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
UNIT 18
c.
i. Law of combining volumes (Gay-Lussac)
ii. Avogadro’s hypothesis (Avogadro)
Introduction to Avogadro’s number and the mole (5.3)
Identify reactants and products in a chemical equation
Convert a verbal description of a reaction to a balanced chemical
equation
Balance chemical equations
UNIT 17
i. Identifying reactants and products
ii. Balancing chemical equations
b. Volume relationships in chemical equations (5.2)
Unit
#
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
Solutions – Qualitative (5.5)
f.
i. Definition of a solution
ii. Qualitative terms related to solutions
1. Unsaturated
2. Saturated
3. Supersaturated
4. Dilute
5. Concentrated
Solutions – Quantitative (5.5)
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
ii. Percent concentrations
1. Percent by volume
2. Percent by mass



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


Define the terms molarity, percent by mass, and percent by volume
Determine missing information in solution problems
Unit
#
UNIT 20
e.
Learning Objectives
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
IX.
Gases, Liquids, Solids, and Intermolecular Forces (Chapter 6)
a. Solids, liquids, and gases (6.1)
c.



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

State key physical differences between ionic and molecular compounds


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.
d.

Identify the key forces that determine the solubility of a solute in a
solvent
UNIT 22

Dipole-dipole forces
ii. Dispersion forces
iii. Hydrogen bonding
Forces in solution (between solute and solvent) (6.4)
Unit
#
UNIT 21
b.
i. Identification of molecular level differences
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
Comparison of properties of 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
Forces between molecules (intermolecular forces) (6.3)
Learning Objectives
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
f.
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
Simple gas laws (6.6)
g.
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
Ideal gas law (6.7)
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


State the five postulates of the kinetic molecular theory of gases
Describe how the kinetic molecular theory accounts for observed
properties of gases




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




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
Unit 24
Gases: The kinetic-molecular theory (6.5)
Unit
#
Unit 23
e.
Learning Objectives
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
X.
Acids and Bases (Chapter 7)
a. Experimental definitions of acids and bases (7.1)
b. Acids, bases, and salts (7.2)
Learning Objectives

Identify oxides as acidic or basic



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

Complete the chemical equation for a neutralization reaction given the
reactants
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
Identify conjugate acid-base pairs
e.
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)
f.
The pH scale (7.6)



g.
h.
Buffers and conjugate acid-base pairs (7.7)
Applications (7.8-7.10)
i. Acid rain
ii. Antacids
iii. Use in industry
iv. Use in health and disease issues

c.
UNIT 27
d.
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)
i. Nonmetal oxides – acidic anhydrides
ii. Metallic oxides – basic anhydrides
Strong and weak acids and bases (7.4)
UNIT 26
Give physical characteristics of acids and bases
State the Arrhenius theory of acids and bases
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
UNIT 25







Unit
#
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
XI.
Oxidation and Reduction (Chapter 8)
a. Three views of oxidation-reduction (8.1)
Learning Objectives



b.
c.
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
Oxidizing and reducing agents (8.2)
i. Oxidizing agent – substance that is reduced
ii. Reducing agent – substance that is oxidized
Electrochemistry: cells and batteries (8.3)

Identify oxidizing and reducing agents in a given chemical equation


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


Identify the role of oxidation-reduction in corrosion processes
Describe the function of a sacrificial anode
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
Give examples of hydrogen as a reducing agent in applications

UNIT 29
d.
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
Corrosion (8.4)
e.
i. The rusting of iron
ii. Protecting materials from corrosion
Common oxidizing agents (8.7)

Common reducing agents (8.8)




f.
UNIT 28

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
Unit
#
CHEM 1004
Descriptive Chemistry (Summer 2011)
Course Outline
(Numbers in parentheses refer to 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)
The links on the right (Unit 1, etc.) will bring you directly to the Unit PowerPoint slides.
Topic
XII.
Nuclear Chemistry (Chapter 11)
a. Exposure to natural radioactivity (11.1)
b. Writing nuclear equations (11.2)
Learning Objectives



d.
i. Describe constancy for a specimen
ii. Relate to radioisotopic dating (11.4)
Artificial transmutations (11.5)

e.
f.
Uses of radioisotopes (11.6)
Penetrating power of radiation (11.7)


g.
Energy from the nucleus (11.8)


h.
i.
i. E = mc2
ii. Binding energy
iii. Nuclear fission
iv. Conditions for chain reactions
Radioactive Fallout (11.10)
Nuclear power plants (11.11)


j.
Thermonuclear reactions (11.12)


Identify the three natural sources of radiation exposure
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
UNIT 30
c.
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
Half-life (11.3)
Unit
#


Use half-life information to determine quantities of material remaining
after specific lengths of time
Given sufficient information, determine the age of a specimen
Describe the reason for prolonged radioactive fallout from a release
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
Distinguish between nuclear fission and nuclear fission
UNIT 31
Recognize artificial transmutation as man-made process for producing
new elements
Identify three radioisotopes and describe their applications
Rate alpha particles, beta particles, and gamma rays in terms of their
penetrating power
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