study guide ch 1 pt 2 - Professor M. J. Krische Chemistry...

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Professor M. J. Krische Chemistry 318M, Fall 2008 Chapter 1-Study Guide (Part 2 ) I. The Electronic Structure of Organic Molecules (Chapters 1) i. Historical Perspective ii. Lewis Structures iii. Molecular Shape (VSEPR) iv. Resonance Structures v. Atomic Orbitals vi. Molecular Orbitals vii. Hybrid Atomic Orbitals viii. Functional Groups v. Atomic Orbitals Lewis structures are helpful in providing an accurate accounting of electrons. Nevertheless, they are symbolic representations of atomic and molecular electronic structures. Electrons do not stand still in octets, but exist in atomic orbitals of specific energies that are defined as volumes of electron density. Each orbital is defined by 4 quantum numbers, n, l, m and s. n = Principle Quantum Number : Describes the shell in which the electron resides (i.e., which row of the periodic table). n has integer values of 1, 2, 3, etc. . The energy of the shell increases with the value of n. l = Angular Momentum Quantum Number : Describes the “ subshell ” or shape of the orbital . The possible integer values for l are 0 . .. n - 1 . The quantum number l is referred to as a letter: l = 0 is "s", l = 1 is "p" and l = 2 is "d". The energy of the subshell increases with the value of l. m = Magnetic Quantum Number: Describes the orientation of the orbital in space. The possible integer values for m are -l. .. 0. ..+l . The energies of these orientations are equivalent. s = Spin Quantum Number : Describes electron spin . May have values of ±1/2 only . This arises from the Pauli exclusion principle , which states that no more than two electrons (which must be of opposite spin) can occupy a given orbital. The simplest example to consider is Helium (He). The electronic configuration of He is 1s
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This note was uploaded on 02/09/2009 for the course CH 318M taught by Professor Bocknack during the Fall '08 term at University of Texas at Austin.

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study guide ch 1 pt 2 - Professor M. J. Krische Chemistry...

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