Humboldt State University ® Department of Chemistry

Richard A. Paselk

SA 560B; 826-5719

Hour Exam 1 Study Guide

Spring 2013 version


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

Be familiar with the HSU Chemistrty Periodic Table you will have on the exam!

Introduction (Chapter 1)

What is Chemistry? Two "definitions" for science. Model vs. Theory.


Know the elements, ions, and common acids on the Chemistry Department Supplements on the web. Be able to name compounds using "Stock" (systematic) nomenclature as in your text and lab manual appendix. Be able to write formulae for compounds named systematically. Be able to recognize compounds of copper, iron, lead, mercury and tin using the -ous -ic system. (Nomenclature module)

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), including liters to m3 etc., know prefixes in metric system in notes (tera-, giga, mega-, kilo-, deci-, centi-, milli-, micro-, nano-, pico-, femto-.), make approximate conversions between American and metric systems (one inch = 2.54 cm exactly; a yard is about a meter; a quart is about a liter); be able to convert between Farenheit and Celcius (centigrade) temperatures give the formulae; be able to convert between Celsius and Kelvin temperatures; solve density problems and other problems using dimensional analysis.

Atoms (Chapter 2)

Be able to describe/discuss the basic premises of the Dalton atom model. Be able to describe/discuss the basic premises of the Thompson and Rutherford atom models, and know what key problem/experiment led to the rejection of each. What is the make-up of an atom according to the modern picture? Be able to fill out tables as we did in class for isotopes (for example: given A & Z find numbers of protons, symbol, etc. Atoms/Isotopes module). What is an isotope? How do isotopes differ from each other? Are there significant chemical differences between isotopes of the same element? Be able to determine the % composition of elements in terms of their isotopes (for an example see the sample Final Exam, problem II.3). Be able to determine the atomic weight of an element given its isotopic % composition.

Stoichiometry (Chapter 3)

Moles and formulae

What is a mole? (One mole = number of atoms in exactly 12 grams of 12C = 6.022 x 1023 particles). What is a formula unit? What is the mass of one atom in grams? What is the mass of one amu? Be able to solve mole problems. (e.g. How many moles of _ in _ ? How many atoms of _ in _ ? How many molecules of _ in _ ? How many grams of _ in _ ? etc. Moles etc. module). Be able to find formula weights of compounds. Be able to find % composition of compounds given formulae. Be able to find formulae of compounds given % composition. Be able to find molecular formulae given % composition and MW. 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.

Chemical Equations

Be able to balance by inspection simple chemical equations such as we have done in class (What are the products of complete combustion of compounds containing only elements such as C, H and possibly O?). Keep in mind our fundamental assumption that mass is conserved, which is manifested by conservation of atom type and number! A series of examples, including a method to keep track of atoms, may be found on my on-line supplements, Balancing etc. module. Note that these examples do NOT represent reactions in solution, so do not involve net ionic equations!

Solutions (Chapter 4)

Define/describe: solution, salt, solvent, solute, saturated solution, unsaturated solution, super saturated solution, 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.

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 and strengths of electrolytes for common ions (those on our lists in lecture 11). 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).

Stoichiometry Problems

What is stoichiometry? 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)

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

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?


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).


Syllabus / Schedule

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© R A Paselk

Last modified 22 February 2013