CHE_106_Lecture22_2009

CHE_106_Lecture22_20 - Chemistry 106 Lecture 22 Topics Hybrid Orbitals Molecular Orbitals Chapter 9.5-9.8 Announcements NO LECTURES on Tuesday Nov

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Chemistry 106 Lecture 22 Topics: Hybrid Orbitals olecular rbitals Molecular Orbitals Chapter 9.5-9.8
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Announcements NO LECTURES on Tuesday Nov. 24 or Thursday Nov. 26 Next Lecture is Tuesday December 1 (Read Sections 10.1 - 10.5 ) Exam #3 is on Thursday, December 3 Exam #3 covers Chapters 7, 8, and 9
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Bond Theory In this chapter we will discuss the geometries of molecules in terms of their electronic structure. – We will also explore two theories of chemical bonding: valence bond theory and molecular orbital theory . Molecular geometry is the general shape of a molecule, as determined by the relative positions of the atomic nuclei.
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Valence Bond Theory Valence bond theory is an approximate theory to explain the covalent bond (sharing electrons) from a quantum mechanical view. – According to this theory, a bond forms between two atoms when the following conditions are met. 1. Two atomic orbitals “ overlap ”. 2. The total number of electrons contained in both overlapped orbitals is no more than two.
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Overlap and Bonding We think of covalent bonds forming through the sharing of electrons by adjacent atoms. In such an approach this can only occur when orbitals on the two atoms overlap.
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Hybrid Orbitals So far, we have just considered the overlap of atomic orbitals such as s and p. However, it’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from these “basic” atomic orbitals.
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Hybrid Orbitals One might expect the number of bonds formed by an atom would equal its unpaired electrons. – Chlorine, for example, with 1 unpaired electron generally forms 1 bond. (such as HCl) xygen with 2 unpaired electrons, usually – Oxygen, with 2 unpaired electrons, usually forms 2 bonds. (such as H 2 O) – However, carbon, with only 2 unpaired electrons, generally forms 4 bonds! For example, methane, CH 4 .
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Hybrid Orbitals The bonding in carbon (2s 2 2p 2 ) might be explained as follows: Idea : Four unpaired electrons are formed as an electron from the 2s orbital is promoted (excited) to the vacant 2p orbital. roblem Energy is required for this promotion and Problem : Energy is required for this promotion and therefore does not seem likely to happen. – However, more than enough energy is supplied (refunded) for this promotion from the formation of the two additional covalent bonds.
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s 2p s 2p Hybrid Orbitals C atom (ground state) 2s Energy 2s 1s 1s C atom (promoted)
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Hybrid Orbitals Using this new orbital diagram, one bond on carbon would form using the 2s orbital while the other three bonds would use the 2p orbitals. – This does not explain the fact that the four
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This note was uploaded on 02/09/2010 for the course CHE CHE 106 taught by Professor Korter during the Fall '09 term at Syracuse.

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CHE_106_Lecture22_20 - Chemistry 106 Lecture 22 Topics Hybrid Orbitals Molecular Orbitals Chapter 9.5-9.8 Announcements NO LECTURES on Tuesday Nov

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