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)
I.
II.
III.
Science and Its Methods
a. Science compared to technology (1.1)
b. 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
c. The limitations of science (1.4)
d. Critical thinking – FLaRES approach (1.13)
e. Desirability quotient – benefits/risks (1.5)
Characteristics of Chemistry- Study of Matter and its Transformations
a. Properties (1.8)
i. physical
ii. chemical
b. Transformations (1.8)
i. physical
ii. chemical
c. Classification of matter (1.9)
i. States of matter
ii. Pure substance
1. element
2. compound
iii. Mixture
1. Homogeneous
2. Heterogeneous
iv. Atoms and molecules
Measurement
a. The International System (SI) (1.10)
i. SI units
ii. Current definitions of seven base units
b. The metric system (1.10)
i. Review of exponential notation
ii. Basic prefixes and their interconversions
c. Key Derived units
i. Volume (1.10)
ii. Density (1.11)
d. Temperature units and their interconversion (1.12)
e. Energy units and their interconversion (1.12)
IV.
V.
VI.
Atoms – History and Early Indications
a. Greek view of atoms (2.1)
b. Fundamentals laws that led to early conclusions about atoms
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)
c. John Dalton’s atomic theory of matter (2.4)
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)
i. Dobereiner’s “Triads”
ii. De Chancourtois’s “Telluric Helix”
iii. Newland’s “Law of Octaves”
iv. Meyer’s system of elements
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)
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-tocharge ratio
b. The arrangement of particles in the nucleus
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)
c. The atomic nucleus (3.5)
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
iii. Symbolic representation of isotopes - ZA X where X is the chemical symbol
d. Electron arrangement – the Bohr model (3.6)
i. Relationship between wavelength and energy
ii. Bohr’s description of the relationship to colors emitted in spectra to
electron energy levels
iii. Determining the maximum number of electrons in a given energy level
e. Electron arrangement – the quantum model (3.7)
i. The concept of an orbital
ii. Writing electron configurations for atoms
f. Relationship between electron configurations and the periodic table
i. Identification of valence electrons with group numbers
ii. Identification of named groups and regions in the periodic table
VII.
VIII.
Chemical Bonds (Chapter 4)
a. Identifying criteria for stable electron configurations (4.1)
b. Lewis symbols
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 compounds (4.5)
v. Writing Lewis structures for covalent bonds (4.6-4.8)
vi. Identify and work with polyatomic ions (4.9)
vii. Name covalent compounds (4.6, 4.9)
viii. Rules for writing Lewis structures (4.10)
ix. Relating Lewis structure to molecular shape – The VSEPR Theory (4.12)
x. Relate shape to polarity of molecules (4.13)
Chemical Accounting
a. Chemical equations (5.1)
i. Identifying reactants and products
ii. Balancing chemical equations
b. Volume relationships in chemical equations (5.2)
i. Law of combining volumes (Gay-Lussac)
ii. Avogadro’s hypothesis (Avogadro)
c. Introduction to Avogadro’s number and the mole (5.3)
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)
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 (5.5)
i. Definition of a solution
ii. Qualitative terms related to solutions
1. unsaturated
2. saturated
3. supersaturated
iii. 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
iv. Percent concentrations
1. Percent by volume
2. Percent by mass
IX.
Gases, Liquids, Solids, and Intermolecular Forces (Chapter 6)
a. Solids, liquids, and gases (6.1)
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 condensating)
iii. Supercritical fluids
b. 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
c. Forces between molecules (intermolecular forces) (6.3)
i. Dipole-dipole forces
ii. Dispersion forces
iii. Hydrogen bonding
d. Forces in solution (between solute and solution) (6.4)
e. Gases: The kinetic-molecular theory (6.5)
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)
i. Introduction to pressure measurement and units
ii. Boyle’s Law: relates pressure and volume at constant temperature and
number of moles of gas
iii. Charles’s Law: relates temperature and volume at constant pressure and
number of moles of gas
iv. Molar volume at standard temperature and pressure
g. 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
X.
Acids and Bases (Chapter 7)
a. Experimental definitions of acids and bases (7.1)
b. Acids, bases, and salts (7.2)
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
c. Acidic and basic anhydrides (7.3)
i. Nonmetal oxides – acidic anhydrides
ii. Metallic oxides – basic anhydrides
d. Strong and weak acids and bases (7.4)
i. Define the terms strong and weak acids and bases
ii. Identify strong and weak acids and bases
e. Neutralization of acids and bases (7.5)
f. The pH scale (7.6)
g. Buffers and conjugate acid-base pairs (7.7)
h. Applications (7.8-7.10)
i. Acid rain
ii. Antacids
iii. Use in industry
iv. Use in health and disease
XI.
Oxidation and Reduction (Chapter 8)
a. Three views of oxidation-reduction (8.1)
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)
i. Oxidizing agent – substance that is reduced
ii. Reducing agent – substance that is oxidized
c. Electrochemistry: cells and batteries (8.3)
i. Terminology: electrochemical cell, electrodes, cathode, anode
ii. Descriptions of basic cells and batteries
d. Corrosion (8.4)
i. The rusting of iron
ii. Protecting materials from corrosion
e. Common oxidizing agents (8.7)
f. Common reducing agents (8.8)
XII.
Nuclear Chemistry (Chapter 11)
a. Exposure to natural radioactivity (11.1)
b. Writing nuclear equations (11.2)
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)
i. Describe constancy for a specimen
ii. Relate to radioisotopic dating (11.4)
d. Artificial transmutations (11.5)
e. Uses of radioisotopes (11.6)
f. Penetrating power of radiation (11.7)
g. Energy from the nucleus (11.8)
i. E = mc2
ii. Binding energy
iii. Nuclear fission
iv. Conditions for chain reactions
h. Nuclear power plants (11.11)
i. Thermonuclear reactions (11.12)