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Gen Chem Exam 3 Study Guide

by: Nick Manning

Gen Chem Exam 3 Study Guide CHEM - 10060 - 001

Marketplace > Kent State University > Chemistry > CHEM - 10060 - 001 > Gen Chem Exam 3 Study Guide
Nick Manning
GPA 4.0

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These notes cover everything that will be on our next exam, including Chapter 8, 9, and 10.
Study Guide
lewis, Dot, structure, Bonding, periodic, Table
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This 6 page Study Guide was uploaded by Nick Manning on Tuesday April 19, 2016. The Study Guide belongs to CHEM - 10060 - 001 at Kent State University taught by TBA in Fall 2015. Since its upload, it has received 32 views. For similar materials see GENERAL CHEMISTRY I in Chemistry at Kent State University.


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Date Created: 04/19/16
COMPREHENSIVE CHEM EXAM STUDY GUIDE THIS IS IT. The last Gen Chem I exam is upon us. You can do it. I will be going in-depth in most subjects and covering and explaining everything on Dr. Leslie’s study guide in order. You can check out my weekly notes if you need help in one area! Chapter 8 (Periodic Table) Important things to know from this chapter:  How to draw an Electron Configuration Diagram! o When making an electron configuration diagram, the spsn (m ) is represented as arrows, either ↑ (+1/2) for clockwise or ↓ (-1/2) for counter clockwise. A Ground State Electron Configuration Diagram shows the arrangement of electrons in their orbitals for an atom in its most stable (ground) state and is filled with up arrows in each box first, then filed from the beginning with the down. Orbitals are displayed with boxes and in the boxes have the arrows, with a MAX of 2 e- per box. o EXAMPLE) Electron Config. of Carbon ↑ |↑ | ↑↓ ↑↓ 1s 2s 2p o Pauli’s Exclusion Principle: each orbital can hold max of 2e- if they have opposite signs. o Hund’s Rule: if there are not enough e- to fill a whole orbital, add ↑ arrows first to all boxes.  How to use this diagram to find out.. o # of Core e- / inner e-: in filled E lvels (+ e- in filled d or f levels) o # of Outer e-: e- in highest n level o # of Valence e-: outer electrons + unfilled d or f shells o # of Unpaired e-: electrons with only an up arrow and no down to match  Unpaired electrons also result in a substance being paramagnetic, while paired electrons results in a diamagnetic substance.  Know how to do the shortened version also! o write the nearest noble gas in parentheses then list the electrons in their groups after. For example, changing Carbon’s from  1s2 2s2 2p2 to  [He] 2s2 2p2  Know how to use the chart I used in the beginning to find an element based on the listed electron configuration.  Know the important trends in Electron Affinity, Ionization Energy, Atomic Size and Metallic Behavior from the Periodic Table. (Excuse the horrible periodic tables, I tried) Ionization Energy Atomic Size & & Electron Affinity Metallic Behavior & Acidity Chapter 9 (Bonding) You should know:  How to use the two elements in a bond to know what type of bond (metallic (2 metals), covalent (2 nonmetals), or ionic (metal & nonmetal))  How to draw a Lewis dot diagram o Use the Group number (aka # of valence electrons) and put the dots around the element, pairing them on each side and putting one on every side before you begin pairing. Usually a max of 8.  The Octet Rule: elements strive to get 8 valence electrons, and the when the charges reflect 8 electrons total, it is called a pseudo-noble gas configuration  The different formations of an ionic bond, and how to identify the change in energy. o When using a Lewis Dot, just draw lines to connect the electrons that will be transferring. o When using a partial Orbital Diagram or Electron Configuration, Just show where the electrons transferred, either by writing in the arrow or by writing the number.  Lattice Energy: measure of how strongly ions attract each other  Lattice Energy increases when size decreases and charge increases, because of less shielding and more protons pulling on the electrons.  Ionic Compound Properties: -Array of ions in a crystal - Hard but brittle -Conduct electricity in solution or molten forms - Nonconductors in solid forms - No free moving ions when solid - Ions are free to move when melted  In a covalent bond, both nuclei of bonded atoms attract shared e- - not a static bond, it’s more of a dynamic pulsing bond bw the two atoms  There are specific partners in covalent bonds and not specific in ionic bonds  How to predict the covalent bond strength: - in terms of atomic size- strength increases when size decreases - in terms of bond type (single, double, triple) The more bonds, the stronger the bond and the shorter the bond length  A polar covalent bond is when the electrons are not shared evenly, and the more electronegative element will “take more control” of the electrons. - If the difference between the elements is less than .4, it is a nonpolar covalent, between .4 and 1.7 equates a polar covalent, and over 1.7 means it is an ionic bond.  Metallic Bonds “pool” their valence electrons together into an electron sea. This makes metals malleable and still hard. They sort of share their electrons between all of the atoms.  Valence e- are free to move independently, allowing for great heat and electricity conduction  Metals are strong, but can be shaped & worked w/o breaking CHAPTER 10 LEWIS STRUCTURES You should know:  WHAT LEWIS STRUCTURES ARE. They are the dots and lines thing we have been talking about. They show covalent bonds between atoms. To draw one, you draw the elements Lewis Dot diagrams, and draw a line between the two elements for each two shared electrons, if there were four electrons, then two bonds, if six, then three bonds.  Resonance hybrid structures must be used when more than one Lewis structures could correctly be made; but three rules are used to find the truest structure o Usually prefers a full octet o Fewest nonzero charges o Prefers to carry negative formal charge and an atom with a higher Electronegativity  Expanded octets are used when elements above period 3 have more than 8 valence electrons.  Formal charge must also be shown with these structures o Compares the valence e- of unbonded atoms to “e- owned” by bonded atoms o Always 0 when bonded in preferred pattern  e- owned = all the non-bonding e- (dots; the lone pairs and single e-) o Also the lines = ½ the bonding e-  Formal charge is calculated by: val e- - e- owned  A formal charge of 0 mean the atom is making its preferred bonding.  Predict Polarity by taking the difference of the Electronegativity and study these molecular shapes  VSEPR = Valence Shell Electron Pair Repulsion  Wedge equals poking forward in 3D, and dotted lines equal going backwards I hope this all helps, use this to your advantage!!


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