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C121 11-3 lec28GL - Lecture 28 Bonding Specific Models Last...

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Lecture 28 November 16, 2011 Bonding: Specific Models Last Day: I Localized Electron Bonding Models (iii) Hybrid Atomic Orbitals (Chapter 14) Examples Comments on Multiple Bonds Today: II Molecular Orbital Theory (14.2-14.4)
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Orbitals used to form bonds in Ethylene
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Bonding in Benzene C 6 H 6
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σ - bonded network π - bonded network
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Recap: Lewis Structures/VSEPR Theory 2 centre-2 electron bonds predicts geometry in 3-D not informative about orbitals used in bonding Hybrid Atomic Orbital Theory 2 centre-2 electron bonds rationalizes molecular geometry and bonding using overlapping hybridized orbitals ( σ -bonding) and unhybridized p-orbitals (π –bonding)
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Limitations of these theories: they do not explain species with unpaired electrons unreliable in accurately describing energy levels of the e - For example: CH 4 according to hybrid orbital theory C sp 3 hybridized predicts electrons occupy four energetically equivalent orbitals Reality! Spectroscopy shows one pair of bonding electrons is lower in energy than the other three Molecular Orbital Theory explains the energetics of molecules very well
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Molecular Orbital Theory molecular orbital theory assumes that valence electrons occupy orbitals (molecular orbitals) that spread throughout the entire molecule (ie. they are delocalized) molecular orbitals (MO’s) are built by adding together (or superposing) valence atomic orbitals of similar energy and symmetry as a theory it successfully explains the energetics of many molecules ex. observe that liquid O 2 is attracted to the poles of a strong magnet ie. it possesses unpaired e - (paramagnetism) Lewis structure predicts only paired e - O::O MO theory (as we will see) correctly predicts paramagnetism : : : :
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MO theory can be quite complex! We will simplify by: (a) presenting a non-mathematical approach (b) considering only the simplest of molecules: first and second row homonuclear and heteronuclear diatomics same atoms ie. H 2 , Li 2 , N 2 + etc . different atoms ie. CO, CN - , NO etc.
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The H 2 Molecule (H A H B ) Recall that the wavefunction of an orbital describes the electron density as a function of the distance (R) from the nucleus for the 1s orbital Ψ(1s) ~ e - k R (spherical) Let’s construct the molecular orbitals for the H 2 molecule by combining the H 1s valence atomic orbitals on the two atoms Ψ R H A Ψ R H B bring the two atoms together
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when we superimpose the wavefunctions there are two possible atomic orbital combinations: (a) Ψ + = Ψ(1s) A + Ψ(1s) B Ψ R constructive overlap Ψ + region of enhanced electron density between the two nuclei bonding MO H A H B resulting molecular orbital Ψ
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(b) Ψ - = Ψ(1s) A - Ψ(1s) B H A H B resulting molecular orbital Ψ Ψ R destructive overlap Ψ - nodal plane (zero electron density) between the two nuclei anti-bonding MO
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