Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chemistry 107 - Fall 2009

Exam II Study Guide

F 2009 Version

Review

Quizzes (all quiz keys are posted on Moodle) and problem sets since Exam I. Review nomenclature so that you can read questions and understand them.

Periodic Table of the Elements
 IA IIA IIIA IVA VA VIA VIIA VIIIA
1 2 13 14 15 16 17 18
   H  He
Li Be    B C N O F Ne
Na Mg 3 4 5 6 7 8 9 10 11 12  Al Si P S Cl Ar
K Ca   Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
 Rb Sr                  Ag  Cd   Sn     I  Xe
 Cs  Ba        W        Pt Au  Hg    Pb        

Atoms and Atomic Structure

Bohr model for atom. What is wrong with the Bohr model? Geometry (shapes) of s- (spherical) and p- (two lobes along the x, y, or z axis with planar symmetry) orbitals for hydrogen. Define: orbital. s-orbital, p-orbital. Chemical bond, ionic bond, covalent bond.

Be able to write electronic configurations for any of the atoms (and their ions) on the periodic chart using the:

Chemical Periodicity

Read your text chapter on Periodicity.General properties of elements as exemplified on Periodic Table (which are most likely to lose electrons? gain them? across a period? within a group?). What charges do ions for groups I, II & VIIA usually have? ions for Al, O and S? (Remember our model for predicting ionic charges for elements: ion achieves electronic configuration of nearest Noble gas.) Where are metals located on the Periodic Table? non-metals? semi-metals (metalloids)? Which elements occur naturally in the pure state(e.g. Noble gases; C, O, N, S, Cu, Ag, Au,)? What are the formulas of the various gaseous elements, including the various halogens? Which elements occur as liquids at room temperature? Gases? What are allotropes of C (graphite and diamond)?

How do ionization energy, atomic radii and electronegativity vary on the Periodic chart. Know trends for these properties, elements with highest and lowest values. Which elements are considered to be metals? non-metals? metalloids (semi-metals)?

(Remember our model for predicting ionic charges for elements: ion achieves electronic configuration of nearest Noble gas.) Which elements occur naturally in the pure state (e.g. Noble gases; C, O, N, S, Cu, Ag, Au,)? lanthanide, actinide, s-block element, p-block element,

Bonding

What characterizes a covalent bond? an ionic bond? Predict whether a bond between two atoms is covalent or ionic. Predict the charges of ions of the representative elements. Explain why they have these charges (stable octets). Octet Rule. Valence electrons.

Lewis Dot Structures. What does the kernal of a Lewis Structure represent? Be able to write Lewis Structures for atoms, ions, and ionic compounds of the representative elements. How do you guess whether a compound is likely to be ionic or covalent? (hi, lo, intermediate rules for electronegativity). Predict the charges of ions of the representative elements. Explain why they have these charges (stable octets). Octet Rule. Valence electrons. Be able to draw correct Lewis Structures for molecules with multiple bonds and for molecules with resonace structures. (Many solved examples are in PS in lab manual, pg 27ff. Additional examples are available in problem Set A of the Supplemental Study Module) How do you guess whether a compound is likely to be ionic or covalent? (hi, lo, intermediate rules, for electronegativity; or if EN > 1.7 then ionic-Memorize the electronegativities of the Period 2 elements and hydrogen.). Don't forget charges for ions and brackets for negative ions!

Molecular Shapes and Polarity

VSEPR Theory. Steric number. Electronic shapes (linear, trigonal planar, tetrahedral) vs. molecular shapes (linear, bent, trigonal planar, trigonal pyramidal, tetrahedral). When and why do the various electronic shapes give rise to the different molecular shapes (electronic shape determined by connecting electron pairs and/or bonding electron clouds, molecular shape determined by connecting nuclei of attached atoms). Be able to predict the shapes of simple molecules using VSEPR Theory. Dipole. Dipole moment. Be able to predict the polarity of a covalent bond and the direction of its dipole moment. Practice problems are available in the VSEPR Theory and Molecular Geometry Supplemental Study Module.

Weak Bonds

What is a weak bond? What are the different types? (van der Waal's and Hydrogen bonds). What are van der Waal's bonds? How do they work? (generally contact only). When are they important?

Gases

Define/describe: pressure, barometer, manometer, Boyle's Law, Charles' Law, Standard Atmosphere, Avogadro's Principle, Ideal Gas, Perfect Gas, Dalton's Law of Partial Pressures, Kinetic Energy (=1/2 mv2), 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 gas law problems such as we have seen in class. 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! What is the volume of one mole of an ideal gas at STP? Practice problems are available in the Gas Law Problem Supplemental Study Module (1-5 ). Be able to do gas stoichiometry problems (Gas Law Module, #6 & 7) - Review simple equation balancing. What is the Kinetic Molecular Theory for gases? 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?

Liquids

Define/describe: weak forces, Hydrogen bonds, vapor pressure, phase change, boiling, solid, boiling point (bp), heat of vaporization/condensation, melting point (mp), heat of fusion/crystallization.

Be able to draw/interpret heating/cooling curves. (Remember that the cooling curve is just like the heating curve, but backwards - that is you start at the gas phase and remove heat.) Remember that the vapor pressure of a pure substance is dependent only on the nature of the substance and the temperature. Why is water's boiling point so high (vapor pressure so low)? How does it compare to other molecules? What are the other special properties of water and how do we explain them? Are there any molecules with similarly high boiling points? Why do liquids boil? When? Why are boiling points lower in the mountains?

Solutions - Review

Define/describe: solution, salt, solvent, solute, mass %, ppm, ppb, molarity (M, moles/liter solution) solution concentration problems using molarity (see Lecture 8 & Lecture 9 notes)

You may bring a 4" x 6" card (both sides) with any information you like to aid on the exam.

Organization is critical - you don't want to waste time finding stuff!!

You will be provided with a Periodic Chart.

 

refractometer icon

C107 Home

spectroscope icon

C107 Lecture Notes

© R Paselk

Last modified 13 November 2009