GENERAL CHEMISTRY - PART 2
CHEM 142
Spring 2009
Instructor: Dr. John D. Watts Conference Hours: MW 1 – 3 p.m.
Office: 505 JAPSB
Phone: 601 979 3488 Other times by appointment or if instructor is available.
E-mail: watts@chem.jsums.edu
Classroom: 209 JAPSB
Class schedule: MWF 11:00 – 11:50 a.m.
Instructor’s Webpage: http:/ /chem.jsums.edu/watts
Course Description
CHEM 142 is the se cond part of a two-semester course in General Chemistry. It is built upon the concepts and
theories presented in CHEM 141. The course covers to pics such as properties of gases, liquids and solids;
intermolecular interactions; properties of solutions; chemical kinetics; chemical equilibria; properties of acids and
bases; solubility and solubility product of slightly soluble salts; the second and third laws of thermodynamics;
electrochemistry.
Prerequisites:
1. The student must have taken CHEM 141 and earned a passing grade, preferably a 'C' or better. Alternatively, the
student must have been placed in the course with th e permission of the instructor.
2. The student must have officially registered with a valid fee statemen t.
3. It is stro ngly recommended that the st udent be registered in a CHML 142 laboratory class.
Course Objectives: After mastering the material from this course, a s tudent should be able to:
State and use the ideal gas equation (IGE) and its “sub-laws” (e.g. Boyle’s law, Charles’s law, Avogadro’s law, and
the combined gas law); calculate the density and molar mass of a gas; state and use Dalton’s law of partial
pressures; understand the basics of the kinetic th eory of gases, including the molecular origin of pressure and the
temperature dependence of molecular speeds; state and use Graham’s law of effusion; understand t he limitations
of the IGE.
State the charac teristics of the three states of matter and the various phase changes betwe en them; describe vapor
pressure and define boiling point in terms of it; construct and use pressure-temperature phase diagrams; describe
the different intermolecular forces, recognize the intermolecular forces present in diffe rent substances, and assess
their effect on properties such as melting and boiling points; compare the properties of different ty pes of solids.
Define, use, and interconvert between different conc entration measures (e.g. molarity, molality, mole fraction, mole
percent); use intermolecular forces to describe t he energetics of solution formation and predict miscibility; use
solubility concepts such as solubility curves; assess the effec t of temperature on solubility; state and use Henry’s
law on gas solubility; state and use Raoult’s law to calculate vapor pressures of mixtures; describe and do
calculations on colligative properties such as fr eezing point depression, boiling point elevation, and osmotic
pressure, using the appropriate van’t Hoff factor.
Define instantaneous and average reaction rate; determine rate laws from rate or concentration data ; perform
calculations with integrated first- and second-order rate equations; define half-life and relate it to rate constant;
know and use the Arrhenius equation; describe the tempe rature dependence of rate const ants; define reaction
profile, activation energy, transition state, and inter mediate; define reaction mechanism in terms of elementary
steps; predict the rate law from a proposed mechanism; explain the role of a catalyst.
Explain dynamic equilibrium; write the equilibrium constant expression for a chemical reaction; use the “algebra” of
equilibrium constants; define Kc and Kp and interconvert between them; distinguish between the equilibrium
constant and rea ction quotient; use the latter to predict the direction toward equilibrium; state and use Le Chatelier’s
principle to predict shifts in equilibria when a system at equilibrium is perturbed; calculate values of equilibrium
constants from equilibrium concentrations or partial pr essures; determine equilibrium concentrati ons or pressure
from given initial condition s.
Define acids and bases; identify conjugate acids and bases; wri te equilibrium constant expressions for acids and
bases (Ka and Kb) and water autoionizatio n (Kw); calculate pH, pOH, pK a, and pKb from concentrations and
equilibrium constants; calculate equilibrium concentrations from given initial concentrations and equilibrium
1
