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General Chemistry Test Study Guide

by: Samantha Weller

General Chemistry Test Study Guide CHEM 101 - 08

Marketplace > Radford University > Chemistry > CHEM 101 - 08 > General Chemistry Test Study Guide
Samantha Weller
GPA 3.2

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Study guide for a test on the final chapters of the course
General Chemistry
Timothy J Fuhrer
Study Guide
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This 15 page Study Guide was uploaded by Samantha Weller on Thursday December 10, 2015. The Study Guide belongs to CHEM 101 - 08 at Radford University taught by Timothy J Fuhrer in Fall 2015. Since its upload, it has received 30 views. For similar materials see General Chemistry in Chemistry at Radford University.


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Date Created: 12/10/15
FINAL Chem 101 Electromagnetic Radiation Frequency- # of wave peaks that pass by a given point in a certain time Hertz- The SI unit for frequency= s -1 Wavelength- distance form one peak of a wave to the next Amplitude- the height of a wave, from centerline to peak Balmer-Rhydberg m and n are whole numbers where n>m -2 -1 R is a constant= 1.097 * 10 nm E= hc/X h=6.626 * 10 -34J-S Heisenburg Schrodinger Y is the wave function or ortital H is the Hamilton and is a mathematical operator E is the energy of the electron being described 2 Y is the probability of finding the electron in a particular region of space Consequences of the Wave Function Principle quantum number- positive integer (n=1,2,3) on which the size and energy level (shell) of the orbital depend Angular Momentum Quantum Number- defines the 3D shape (subshell) of an orbital Magnetic Quantum Number- the special orientation Shielding 33As 1s 2s 2p 3s 3p 4s 3d 2 104p 3 nodes-1 less than it’s energy level number The Fourth Quantum Number Spin Quantuum number (m ) =s+1/2 or -1/2 Related to the magnetic field generated by a spinning charge Pauli exclusion principle- no two electrons in an atom can have the same 4 quantum numbers n l ml ms 1 0 0 +1/2, -1/2 Electron configuration Aufbau Principle- lowest energy orbitals are filled first. Orbitals can only hold two electrons, which must have opposite spin. Hund’s rule- if two or more orbitals with the same energy (degenerate) are available, one electron goes in each until each have one, before pairing up. Electrons in half filled orbitals will have the same spin numbers The Periodic Table Valence electrons- electrons in the highest principle energy level of a given atom 6s  valence electrons -All elements in a given group on table have similar valence- shell configurations z=atomic number as atomic number increases, the valence-shell electrons are attracted more strongly to the nucleus Ionic Compounds, Periodic Trends, and Bonding Theory Covalent bond- chemical bond that forms when two atoms share several electrons H ** H Molecule- two or more atoms joined by covalent bonds Diatomic element- exist as elements only in the form of diatomic molecules rather than atoms (H2, N2, O2, F2, Cl2, Br2, I2) Octet Rule- trying to get 8 valence electrons. (Metals do this easier) -nonmetals in the 2rd row or lower can bond with 6 atoms sometimes more than 8 Ionic bond- chemical bond resulting from the transfer of one or more electrons from one atom or another. Polyatomic ions- groups of atoms held together by covalent bonds, but with a zero-net charge *Cations are smaller than their neutral atoms, because the principle quantum number of the valence shell electron is smaller for cations because the atomic number is bigger Ionization energy- the energy necessary to completely remove one electron from a gaseous atom U= R 2 21 1 ___________ d Electron Affinity- (affinity means like, love) The energy change that occurs when an electron is added to an atom in the gaseous state (E ) ea The Octet rule Main group elements tend to undergo reactions that leave them with 8 outer-shell electrons 2 2 6 1 2 2 6 Ex: Na (1s , 2s 2p 3s )  Na+ (1s 2s 2p ) +e- ^8 electrons in outer shell Na Cl O 4 +8 +1 +7 Energetics of Ionic Bonding Sum of energy changes= net energy change equals the net energy for overall reaction. The born-haber cycle (forms loops) Chap Lattice Energy (pg 213) Coulombs law: F= K(z 1 )2vd Cation + anion energy Covalent bonding and electron-dot structure ethyl alcohol ether Molecule- the unit of matter held together by covalent bonds Electronegativity The ability of an atom in a molecule to attract the shared electrons in a covalent bond N, CL are hog electrons (hog the charge) Polar covalent bonds- shared but not equal Non polar bonds are equal and symmetrical Bond Polarity Na+ Cl Ionic H-Cl Polar covalent bond O-O equal Resonance When it is possible to draw two different valid structures for the same molecule moving only electrons and not atoms -Actual electric structure is an average of possibilities, called a resonance hybrid Formal Charge # of valence electrons on free atom minus the number of valence electrons assigned to bonded atom subtract original charge by bonds around central atom (# of valence electrons in free atom) -1/2 (# of bonding e-) – (# # of nonbonding e-) Formal charge of zero are more favorable when not all atoms can be zero, the most electronegative charge, and vice versa. SCN- = 15 *Carbon is usually the center atom Chapter 8- Covalent Bonding and electron-dot structures Molecular Shapes VSEPR- valence shell electron pair repulsion model # Charge Clouds Bond Angle Hybridization 2 180 sp 3 120 Sp 2 3 4 109 Sp 5 120/90/180 Sp d 6 90/180 Sp d 2 Valence Bond Theory Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite spin; the overlapping orbitals must be the same phrase Each bonded atom maintains its atomic orbitals, but electron pair in the overlapping orbitals are shared by both atoms. The greater the amount of orbitals overlap, the stronger the bond. (This causes directional character for the bond when other than s orbitals are involved.) Orbital Hybridization Valence bond theory as we know it so far, breaks down in describing many molecules, beginning with CH 4 sigma bond ------- single Double bond- pi bond Triple 1 is a sigma 2 pi bonds


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