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Localized Electron Models - Lecture Slides | CHEM 162, Study notes of Chemistry

Material Type: Notes; Class: GENERAL CHEMISTRY; Subject: Chemistry; University: University of Washington - Seattle; Term: Autumn 2008;

Typology: Study notes

Pre 2010

Uploaded on 03/18/2009

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Download Localized Electron Models - Lecture Slides | CHEM 162 and more Study notes Chemistry in PDF only on Docsity! Lecture 2 Hybridization, σ and π Bonding Localized Electron Model (LEM) Fatal shortcomings of LEM Molecular Orbital Model σ bonding and π bonding Read pages 653-664 of Zumdahl • Two modes of bonding are important for 1st and 2nd row elements: σ bonding and π bonding • These two differ in their relationship to the internuclear axis: σ bonds have electron density ON the axis π bonds have electron density ABOVE AND BELOW the axis What did we learn so far? (1) Molecular orbitals are combinations of atomic orbitals (2) Atomic orbitals are “hybridized” to satisfy bonding in molecules (3) Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for molecular geometry The Localized Electron Model is very powerful for explaining geometries and basic features of bonding in molecules, but it is just a model. Example: O2 - Lewis dot structure O=O - All electrons are paired Contradicts Experiment! .. .. .. .. Major limitation of the LE model: Assumes electrons are highly localized between the nuclei - sometimes requires resonance structures - does not deal with excited states - incorrectly predicts physical properties in some cases The Molecular Orbital Model Basic premise: When atomic orbitals interact to form a bond, the result is the formation of new molecular orbitals HΨ = EΨ Important features of molecular orbitals: 1. Atomic Orbitals are solutions of the Schrödinger equation for atoms. Molecular orbitals are the solutions of the same Schrödinger equation applied to the molecule. Bonding is driven by stabilization of electrons • Electrons are negatively charged • Nuclei are positively charged The bonding combination puts electron density between the two nuclei - stabilization The antibonding combination moves electron density away from the nuclei - destabilization = = nucleus+ σ* M.O. is raised in energy σ M.O. is lowered in energy H atom: (1s)1 electron configuration H2 molecule: (σ1s)2 electron configuration Figure 13.1 Same as previous description of bonding (Ch 13): σ σ∗ Bond Order Bond Order = # bonding #antibonding electrons electrons 2 The bond order is an indication of bond strength Greater bond order Greater bond strength Bond Order: Examples H2 Bond order = (2-0)/2 = 1 Single bond Stable molecule (436 kJ/mol bond) He2 Bond order = (2-2)/2 = 0 No bond! Unstable molecule (0kJ/mol bond) He2+ H2+ Bond order = (2-1)/2 = 1/2 Half of a single bond Can be made, but its not very stable (250kJ/mol bond) Bond order = (1-0)/2 = 1/2 Half of a single bond Can be made, but its not very stable (255kJ/mol bond) For diatomics of atoms with valence electrons in the p orbitals, we must consider three possible bonding interactions: σ−typeπ−type π−type = nucleus Figure 14.36 + + + + + + + + – – – – – – – – – – + +– (+) constructive mixing (–) destructive mixing Example: the O2 Diatomic Oxygen has the 2s22p4 valence configuration Bond Order = (8-4)/2 = 2 O2 is stable (498 kJ/mol bond strength) (σ2s)2(σ2s*)2(σ2p)2(π2p)4(π2p*)2
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