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CHMY 211 Week 1

by: Bronwyn

CHMY 211 Week 1 CHMY 211-001

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About this Document

Book notes from Chapter 1. Lecture notes from 8/31-9/2 covering sections 1-3 of Chapter 1. (1.3 is incomplete and will be finished in class on 9/6)
Elements of Organic Chemistry
Kristian H. Schlick
Class Notes
General Chemistry, Organic Chemistry




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This 5 page Class Notes was uploaded by Bronwyn on Monday September 5, 2016. The Class Notes belongs to CHMY 211-001 at Montana State University taught by Kristian H. Schlick in Fall 2016. Since its upload, it has received 113 views. For similar materials see Elements of Organic Chemistry in Organic Chemistry at Montana State University.

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Date Created: 09/05/16
CH 1: ​ ovalent Bonding and Shapes of Molecules  ● 1.1 ­ Electronic Structure  ­​ ○ ground state electron configuration = e​ arranged at lowest energy  ○ Rules:   Orbitals fill from low to high energy  Each orbital hold 2 e​   Orbitals of equal energy fill before e​ pair up  ○ Lewis Structures ~ show valence electrons  ● 1.2 ­ Lewis Model  ○ octet rule = full outer shell  ○ anion = negatively charged  ○ cation = positively charged  ○ ionic bond = bond between anion and cation  ○ covalent bond = bond sharing e​   ­ ■ nonpolar = equally shared e​  polar = unequally shared e​  ○ electronegativity = the strength of an electron’s attraction to an atom  ○ bond length = distance between nuclei  ○ dipole  ○ formal charge = the charge on an atom in a molecule  ● 1.3 ­ Predicting Bond Angles  ○ VSEPR = valence electrons can be involved in single, double, or triple bonds or  unshared, this creates areas of electron density that want to repel each other  ● 1.4 ­ Polar or Nonpolar  ○ Polar and dipole describe a covalent bond in which atoms don’t have an equal  charge  ○ determined by electronegativity and molecular geometry  ○ nonpolar = atom with no polar bonds or has a vector sum of the dipoles of zero  ● 1.5 ­ Resonance  ○ resonance structures = molecule representations that differ only in the  arrangement of valence e​  ● 1.6 ­ Orbital Overlap  ○ shapes ~ s  is spherical, p is dumbbell  ○ sigma bond = covalent bond formed by orbital overlap  ○ hybridization = atomic orbitals that combine to decrease energy  ● 1.7 ­ Functional Groups  ○ the part of a molecule that undergoes chemical reactions = functional group  ○ review table 1.9  ○ alcohols = ­­OH group  ○ amine = bond to N, NH, or NH​   2 ○ aldehyde = O=CH (w/ an R group bonded to the Carbon)  ○ ketone = O=C (w/ 2 R groups bonded to the Carbon)  ○ carboxylic acid = O=C­OH (w/ an R group bonded to the Carbon)  ○ ester = O=C­O (w/ an R group on both the Carbon and the single bonded  Oxygen)  ○ amide = O=C­N (w/ an R group on the Carbon and 2 R groups on the Nitrogen)   8/29:  No notes  ​ 8/30: ​ 1.1 ­ Electronic Structure  ● Structural Theory  ○ Isomers  ○ Bonds  ■ C = 4 bonds  ■ N = 3  ■ O = 2  ■ Halides = 1  ● Reactions  ○ Collision → bonds break/ form by rearranging e​ ­  ● Atoms  ○ Nucleus = protons + neutrons = most of mass  ○ Extranuclear space = electrons = most of space  ● Covalent bond  ○ Electron sharing  ○ Stable bond = low energy ← at ideal intermolecular distance  ○ Ex: H­­H  ­​ ● e​ structure  ○ Shells hold 2n​  electrons  st​ ­ ■ 1​  shell = 2e​  ■ 2​  = 8  ■ 3​  = 18  th​ ■ 4​  = 32  ○ Valence electrons = outermost shell, involved in rxns, highest energy  ­ ■ Group # = # of valence e​   ● Lewis Structures  ­​ ○ Draw atoms → arrange atoms to form e​ pairs  ○ Octet rule = complete shell  ○ Ex: C2​H6​O    ○ Ex: why is Borane (BH​ ) 3​stable?  ■ It’s unstable because the octet rule isn’t met    ● Formal Charge  ­​ ­​ ○ (# e​ supposed to be there) ­ (# e​ actually there)  ● Electronegativity  ­ ○ How strongly an atom attracts an e​  ○ Increases moving up and moving right on periodic table  ­​ ○ Nonpolar covalent = same e​ attraction, same atom or no net dipole  ○ Polar covalent = different e​ attraction, different atom  ­​ ○ Induction = orbital e​ shift → dipole  ● Intermolecular forces  ○ H bonding  ■ Compounds with atoms capable of H bonds = protic  ○ Dipole­dipole interaction = lined up polars  ­​ ○ Dispersion forces = a temporary dipole from constant e​ movement  9/1: ​1.2 ­ VSEPR, Hybridization  ● Atomic Orbital  ○ s orbital = 2 e​; p orbital = 6 e​; d orbital = 10 e​  ○ Shells:  ■ 1​  shell: s orbital  nd​ ■ 2​ : s and p  ■ 3​ : s, p, and d  o​ ■ 3 p per shell (x, y, z), 90​  from each other  ○ Each orbital holds 2 e​, paired by spins  ○ Energy level principles  ■ Aufbau = fill orbitals from low to high energy  ■ Pauli = max 2 e​ w/ paired spins  ­​ ■ Hund = 1 e​ in each orbital of equal energy before pairing  ○ Hybridized orbitals  ­​ ■ Promotion of e​ through excitation  ■ Carbon uses this to have 4 bonds instead of 2  3​ ■ 1s + 3p = 4sp​  (CH​ 4​methane)  1s + 2p = 3sp​  (ethene)  1s + 1p = 1sp (acetylene)  ■ σ bonds form with overlapped hybrids  π bonds form with leftover orbitals overlapping  ­​ ● VSEPR = atomic regions of e​ density want to repel each other and have max distance  between them  ● Molecular Geometry  ○ Steric # = # of regions of e​ density  o ■ Steric = 4 → tetrahedral = 109.5​    3 → trigonal planar = 120​  o o 2 → linear = 180​         ​


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