Humboldt State University ® Department of Chemistry

Richard A. Paselk

SA 560B; 826-5719

Hour Exam 2 Study Guide

Spring 2015 version


Review nomenclature so that you can read questions and understand them. Review concepts and information from Exam I that we continue to use in looking at additional system etc.

Problem sets in text (on schedule), discussion modules, Sapling and appropriate lab calculations.

Be familiar with the HSU Chemistry Periodic Table and "Constants and Equations etc." sheet you will have on the exam!

Be able to:

convert numbers to scientific notation and use numbers expressed in scientific notation; do all calculations with proper significant figures; make all conversions within the metric system (SI)

Stoichiometry (Chapter 3)

Stoichiometry Problems

Be able to solve problems involving stoichiometric relationships in chemical reactions. Be able to do both problems involving reactants in excess (e.g. combustion of carbon in open air) and problems with a limiting reagent. (Lecture examples, and don't forget problems in text & Reaction Stoichiometry module)

Solutions (Chapter 4)

Define/describe: solution, salt, solvent, solute, saturated solution (Textbook), unsaturated solution, super saturated solution (Textbook), mass %, molarity, neutralization, equivalent (in chemical reactions). Why do some substances dissolve in each other? Why do others not? ("Like Dissolves Like"). Be able to solve problems involving concentrations as we have seen in class: find mass %, molarity, molality of solutions given their components. Be able to find concentrations of solutions after dilution or mixing with other components. Be able to do assigned problems in text. Be able to determine molarity. Be able to find number of moles given molar concentration. Be able to do dilution problems. (Recall examples of all of these problem types in lecture notes.) Be able to do assigned problems in text.

Solution Reactions

Electrolytes. Dissociation. What is meant by "strong" and "weak" in reference to electrolytes, acids and bases? Degree of dissociation. Aqueous Solution Reactions: Be able to define acids and bases by the Arrhenius definition (acids provide hydrogen ion, bases provide hydroxide ion). What are the meanings of "strong" and "weak" for acids and bases by these definitions? What are precipitation and complexation reactions?

Be able to write and balance net ionic equations! Keep in mind the charges on ions! (Net Ionic Equations module) For elemental ions look at the Roman Numeral Group number for a first guess (Group I = +1, etc. Group VII = 7-8 = -1 etc.). For transition metals and molecular (polyatomic) ions look at the on-line Table of Common Ions. (Note that the Ionic Reactions and Net Ionic Reactions lab exercise is a great source of examples.) Be familiar with classification scheme for reactions (text). Know solubility rules (Lecture 11) and strengths of electrolytes for common ions (those on our lists in lecture 13). Keep in mind that most ions in formulae are independent unless you know otherwise. Thus CaCl2 consists of one Ca2+ ion and two Cl1- ions NOT one Cl22- ion!


What is oxidation? What is reduction? Redox. Be able to recognize oxidizers/reducers, oxidants/reductants in chemical reactions. Know rules for determining oxidation numbers (Know Table of Oxidation Number Rules in lecture 13/module). Be able to assign oxidation numbers to atoms in compounds and molecular ions (e.g. sulfate or acetate ion). A set of examples, with some detailed solutions, are available on my Oxidation Number module.

Be able to write half-reactions. Be able to balance ionic redox reactions by the half-reaction method (Redox Balancing–Acid module, Redox Balancing–Basic module).

Gases (Chapter 5)


pressure, barometer, manometer, Boyle's Law, Charles' Law, Standard Atmosphere, Avogadro's Principle, Ideal Gas, Perfect Gas, Dalton's Law of Partial Pressures, Ideal Gas Law (PV=nRT). What does absolute zero represent? What is the rationale behind this concept? (can't have negative volume).

Be able to solve problems based on materials below:

Ideal Gas Law and gas problems such as we have seen in class, text homework & Gas Laws module. Partial pressure(Dalton's Law) problems. Remember: all temperatures must be in K (absolute temperature!), and if you are using R, pressures must be in atmospheres, volumes in liters and quantity in moles for our value of R! What is the volume of one mole of an ideal gas at STP?

What is the Kinetic Molecular Theory for gases? Kinetic Energy (KE=1/2 mv2). What are its postulates? What is the meaning of temperature (what is it a measure of)? What relationship is there between temperature and pressure (in microscopic terms- what are the particles doing)? Temperature and volume? Root Mean Square Velocity, urms = (3RT/M)1/2 (Textbook). Be able to do gas stoichiometry problems. Graham's Laws of effusion and diffusion - know and be able to solve problems. Module 10 (Gas law problems 6 & 7). Real Gases– know the van der Waals equations, its postulates, and how to use it: [Pobs + a(n/V)2] x (V – nb) = nRT (Textbook).

Thermochemistry (Chapter 6)


exothermic and endothermic. What are exothermic and endothermic reactions. What is heat? work? energy? enthalpy? DeltaE = q + w. DeltaH= DeltaE - w. When are enthalpy and energy equal?

Be able to solve:

heat capacity problems, calorimeter problems. Understand thermochemical equations (what happens to DeltaH if the reaction is rewritten in reverse? What is a standard state? Module 11 part 1 (Calorimeter Problems)


Gas constant, R= 0.082 1 L*atm*mol-1*K-1. Avogadro's number = 6.02 x 1023. Mass of one amu in grams (one gram divided by Avogadro's number).

You will be provided with a Periodic Table and "Exam 2 Constants and Equations etc.", available on Moodle


Syllabus / Schedule

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© R A Paselk

Last modified 5 March 2015