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CH 221 Pre-course Material

by: Virginia Brown

CH 221 Pre-course Material CH 221

Marketplace > North Carolina State University > Chemistry > CH 221 > CH 221 Pre course Material
Virginia Brown
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These notes cover the first three lectures posted on Moodle that should be watched before class begins. They are a basic overview of the concepts from Chemistry 101 that prove necessary for this co...
Organic Chemistry 1
Christopher Gorman
Class Notes
Organic Chemistry, Chemistry




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This 5 page Class Notes was uploaded by Virginia Brown on Tuesday August 16, 2016. The Class Notes belongs to CH 221 at North Carolina State University taught by Christopher Gorman in Fall 2016. Since its upload, it has received 4 views. For similar materials see Organic Chemistry 1 in Chemistry at North Carolina State University.

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Date Created: 08/16/16
CH 221; Sec. 004 Fall 2016 Prof. Christopher Gorman Note Taker- Virginia M Brown Pre Course Material: August 6, 2016 Electronic Structure:  Properties of organic compounds depend on structure  Found in nucleus- protons and neutrons  Charged- protons (+) and electrons (-)  Atomic number (Z): number of protons  Mass number (A): sum of neutrons (n) and protons (z)  Neutral Atoms: proton # = electron #  Electronic properties are one of the two main factors that dictate the reactivities of organic compounds  All about electrons:  Exhibit dual-wave particle nature (electrons behave as waves in an atom)  Schrödinger equation (wave equation)  Describes the energy of an electron in a hydrogen atom  Solutions = wave functions (Ψ)  Orbitals = probability distribution of an electron (Ψ )  4 Quantum numbers determine shell and energy of electrons  “n” (shell) determines shell and energy  “n” is the principle quantum number  1,2,3,4  1 = lowest energy  4 = highest energy  “l” (subshell) determines subshell, shape, and energy  “l” is the angular quantum number (determines angular momentum)  0, 1, 2, …, n-1  Ex- when n = 3; l = 0, 1, 2 (because of 3 subshells)  In Periodic Table: s = 0; p = 1; d = 2; f = 3  “m” (lrbital) designates the orbital within a subshell  “m”lis the magnetic quantum number (describes the orientation of orbitals in space)  Ex- when l = 2; m l 0, ±1, ±2 CH 221; Sec. 004 Fall 2016 Prof. Christopher Gorman Note Taker- Virginia M Brown  “m ” smagnetic spin) determines the orbital’s magnetic spin (either + or -)  Arrangement of electrons in an atom: Shells (n)  Subshells (l)  Orbitals (m)l  Principles/Rules to memorize:  Pauli Exclusion Principle- no two electrons in the same atom can have the same set of 4 quantum numbers; each has a unique combination of the 4  Aufbau’s Principle- electrons are generally assigned to orbitals of successively higher energy  Hund’s Rule- When placing electrons in a set of orbitals having the same energy, we place them singly for as long as possible  Electron Configurations:  Ex- H: 1s 1 2 1  Ex- Li: 1s 2s  Ex- B: 1s 2s 2p2 1 CH 221; Sec. 004 Fall 2016 Prof. Christopher Gorman Note Taker- Virginia M Brown  Node- (Ex- 2p electrons) an electron has a 0% probability of being found at the nucleus (shown on right)  A 2s electron has a significant probability of being found at the nucleus (shown on right)  Coulomb’s Law- the closer oppositely charged species are, the lower their energy of interaction  Ex- 2s has a lower energy than 2p  Valence electrons- in outer-most layer  Involved in chemical bonding and reactions  Involved in achieving noble gas configuration  Core electrons- Not involved in chemical bonding and reactions  Getting/maintaining an octet of electrons in octet shell:  Accomplished by:  Ionic Compounds:  Gaining valence electrons (nonmetals) and forming anions  Losing valence electrons (metals) and forming cations  Covalent Compounds:  Sharing valence electrons (nonmetal – nonmetal)  Bonding:  Ionic Bonding- electrostatic attraction between oppositely charged species  Ex- NaCl  Cation (loses electrons) Na: 1s 2s 2p 3s  Na : 1s 2s 2p2 2 6  Anion (gains electrons) Cl: 1s 2s 2p 3s 3p  Cl : 1s 2s 2p 3s 3p6 2 6  Ionic bonding is considered nondirectional bonding; it maximizes interactions and occurs in 3-dimensions  Covalent Bonding- nuclei od two nonmetals attracted to the electron density between them  Covalent binding is considered directional bonding because the attraction occurs between two nuclei Lewis Structures:  Lewis Symbols- Dots represent valence electrons (examples shown on right)  Lewis Structures- “Gotta have 8”  Nonmetals share electrons so that they have access to an octet of electrons  Depending on placement in the periodic table, diff. nonmetals form differing numbers of bonds to have octets within a molecule: CH 221; Sec. 004 Fall 2016 Prof. Christopher Gorman th Note Taker- Virginia M Brown  4 Family- 4 bonds; 0 lone pairs (Ex- CH4)  5 Family- 3 bonds; 1 lone pair (Ex- NH 3  6 Family- 2 bonds; 2 lone pairs (Ex- H O) th 2  7 Family- 1 bond; 3 lone pairs (Ex- HF)  8 Family- 0 bonds; 4 lone pairs (Ex- Ne)  Formal Charge- the charge that will reside on an atom if we assume that all bonding electrons are shared equally (assume covalent bond).  Family # – possession #  Completes the picture of the Lewis Structure and gives further information about the electronic character of bonded atoms Formal Charge on Atom = VE – [NB + ½ BE] VE = valence electrons NB = nonbonding electrons on an atom BE = bonding electrons around an atom Orbital Hybridization:  Lewis Structures and formal charge are good 2-D representations. However, directional nature (better shown in 3-D):  Dictates shape of molecules  Influences reactivity of molecules  VESPR- used to approximate the geometry of a molecule  Valence Shell Electron Pair Repulsion  General Formula: AX E m n  A = central atom  X = any atom (or ligand) surrounding A  E = a lone pair of the central atom A  CN = coordination number (# of atoms surrounding A or X ) m  SN = steric number (sum of m + n)  Deviation from idealized geometry is due to a non-homologous electron density around A  Lone pairs take up more room than bonding pairs  Therefore, place lone pairs in positions to minimize the small angle contacts  Leads to bond angle compression  Multiple bonds take up more room than single bonds  Larger atoms take up more room  Be able to determine: LDS  Steric #  shape  hybridization CH 221; Sec. 004 Fall 2016 Prof. Christopher Gorman Note Taker- Virginia M Brown SN Electron- pair Geometry Molecular Geometry Hybridization 2 Linear (180°) Linear AX 2 sp (2 sp hybrids) 3 Trigonal Planar (120°) Trigonal Planar AX 3 sp (3 sp Bent AX E hybrids) 2 3 4 Tetrahedral (109.5°) Tetrahedral AX 4 sp (4 sp Trigonal Pyramidal AX 3 hybrids) Bent AX 2 2 3 5 Trigonal Bipyramidal Trigonal Bipyramidal AX 5 sp d (5 hybrid (120° and 90°) See-saw AX 4 orbitals) T-shaped AX 3 2 Linear AX 2 3 3 2 6 Octahedral (90°) Octahedral AX 6 sp d (6 hybrid Square Pyramidal AX 5 orbitals Square Planar AX 4 2 T-Shaped AX 3 3 Linear AX 2 4 Examples: (Solid wedge bond = coming out towards you; dashed wedge bond = going behind) 1) AX E2 2 Bent (109°) 2) AX 3 Trigonal Planar (120°) 3) AX 3  Trigonal Pyramid (109°) 4) AX 4 Tetrahedral (109°)


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