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Chapter 2.3

Chem 1100 Chapter 2.3 Review Notes.docx

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Western University
Chemistry 1027A/B
Felix Lee

2.3: Chemistry Review Valence Bond (VB) Theory – explains the discrepancies between VSEPR-predicted or experimentally observed shapes and those predicted by the geometries of the pure atomic orbitals 1. Covalent bonds between two nuclei are formed when each atom contributes one valence electron to a common orbital. This common orbital, which is an overlap of atomic orbitals, contains two electrons of opposite spin. 2. As a result of the overlap, the electrons are deemed to be localized. They are restricted to the area between the two respective nuclei. They cannot move around, or delocalize throughout the molecule. 3. Within the molecule, the atomic orbitals located on the central atom are not necessarily “pure” atomic orbitals. Bonds involving elements in the second or higher row of the periodic table involve combinations of atomic orbitals that form “hybrid” orbitals as the bonds are being made. Hybrid Orbitals – a mixture of pure atomic orbitals They can also contain non-bonding electrons. Example: Methane sp Hybridization: - Only one of the ‘p’ orbitals is used - These two remaining orbitals can be used to make double and triple bonds 2 sp Hybridization: - The remaining, unused ‘p’ orbital on each carbon atom may be used to make a double bond sp Hybridization: - There are no remaining ‘p’ orbitals, so they cannot form double or triple bonds sp d and sp d Hybridization: - Are not possible with elements in the second row because there are no d orbitals - Possible combinations: o One ‘s’ + three ‘p’ + one ‘d’ = five ‘sp d’ orbitals (four ‘d’ orbitals remain unused) 3 2 o One ‘s’ + three ‘p’ + two ‘d’ = six ‘spd d ’ orbitals (three ‘d’ orbitals remain unused) Summary of the Major Hybridizations: - Regions of e density Atomic Orbitals Used Hybrid Orbitals Formed Electronic Arrangement 2 One s, one p Two sp Linear 3 One s, two p Three sp 2 Trigonal planar 3 4 One s, three p Four sp Tetrahedral 5 One
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