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TAMU CHEM 101 - Valence Bond Theory & Molecular Orbital Theory
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CHEM 101 1st Edition Lecture 22Outline of Last Lecture I. Covalent Bonds and Lewis StructuresII. Atom Formal ChargesIII. Bonding and Molecular StructureIV. Molecular ShapesV. Bond Polarity and ElectronegativityVI. Formal Charge and ElectronegativityOutline of Current Lecture I. Theories of Chemical BondingII. Valence Bond TheoryIII. Molecular Orbital TheoryCurrent LectureChapter 9: Orbital Hybridization and Molecular Orbitals- Theories of Chemical Bondingo Two common approaches to rationalize chemical bonding based on orbitals are the Valence Bond (VB) theory (Linus Pauling) and the Molecular Orbital (MO) theory (Robert S Mulliken).- Valence Bond Theory- close relation to the Lewis structures and the VSEPR (valence shell electron-pair repulsion) model; electrons are localized on a particular atom (in atomic orbitals)o Chemical bonds are described by valence electronso Half-filled atomic orbitals of bonding atoms overlap to form bondso Bonds are localized between atomso The molecular shape of the compound can be predicted.o Atom orbitals overlap to form a bond between two atomso Overlapping orbitals hold two electrons of opposite spino Usually, one electron is supplied by each of the two bonded atomso The bonding electrons are localized with a higher probability of being found within a region of space between the bonding nucleio Both electrons are simultaneously attracted to both nucleio When two vertical p orbitals overlap, it’s called a pi-bondo When two horizontal p orbitals overlap, it’s called a sigma-bondo Hybridization in the methane molecule: the 2s and 2p atomic orbitals of carbon ([He]2s^(2)2p^(2)) are mixed to four sp^(3) hybrid orbitalso The number of hybrid orbitals is equal to the number of atomic orbitals which are mixed to the hybrid orbital seto Hybrid orbitals are built from the combination of an s orbital with as many p (and d orbitals) to have enough hybrid orbitals to accommodate bond and lone pairs on the central atomo EX: Be: [He]2s^(2)  2 x spEx: B: [He] 2s^(2)2p^(1)  3 x sp^(2)o Hybridization in the ethylene molecule: The 2s and 2p atomic orbitals of one carbon ([He]2s^(2)2p^(2)) are mixed to three sp^(2) hybrid orbitals (that form three sigma bonds) The unhyridized p orbitals form one C – C pi bond- Molecular Orbital Theory- atomic orbitals become molecular orbitals which are delocalized over the whole moleculeo There are 4 principles of the MO theory: The total number of MO is always equal to the total number of atomic orbitals contributed by the combined atoms The bonding MOs are lower in energy than the parent MOs; the anti-bonding MOs are higher in energy than the parent MOs The electrons of the molecule are assigned to orbitals of successively higher energy according to the Pauli principle and Hund’s rule The atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar


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TAMU CHEM 101 - Valence Bond Theory & Molecular Orbital Theory

Type: Lecture Note
Pages: 2
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