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Ch. 8 – Molecular Structure. Molecular Geometry (p. 232 – 236). A. VSEPR Theory. V alence S hell E lectron P air R epulsion T heory Electron pairs orient themselves in order to minimize repulsive forces. Lone pairs repel more strongly than bonding pairs!!!. A. VSEPR Theory.

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Molecular Geometry (p. 232 – 236)


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    1. Ch. 8 – Molecular Structure Molecular Geometry(p. 232 – 236)

    2. A. VSEPR Theory • Valence Shell Electron Pair Repulsion Theory • Electron pairs orient themselves in order to minimize repulsive forces

    3. Lone pairs repel more strongly than bonding pairs!!! A. VSEPR Theory • Types of e- Pairs • Bonding pairs – form bonds • Lone pairs – nonbonding e- • Total e- pairs– bonding + lone pairs

    4. Bond Angle Bond Angle A. VSEPR Theory • Lone pairs reduce the bond angle between atoms

    5. Know the 13 common shapes & their bond angles! B. Determining Molecular Shape • Draw the Lewis Diagram • Tally up e- pairs on central atom (bonds + lone pairs) • double/triple bonds = ONE pair • Shape is determined by the # of bonding pairs and lone pairs

    6. BeH2 C. Common Molecular Shapes 2 total 2 bond 0 lone → Electronic Geometry = linear Hybridization = sp LINEAR 180°

    7. BF3 C. Common Molecular Shapes → Electronic Geometry = trigonal planar Hybridization = sp2 3 total 3 bond 0 lone TRIGONAL PLANAR 120°

    8. NO21- C. Common Molecular Shapes → Electronic Geometry = trigonal planar Hybridization = sp2 3 total 2 bond 1 lone BENT <120°

    9. CH4 C. Common Molecular Shapes → Electronic Geometry = tetrahedral Hybridization = sp3 4 total 4 bond 0 lone TETRAHEDRAL 109.5°

    10. NCl3 C. Common Molecular Shapes → Electronic Geometry = tetrahedral Hybridization = sp3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° <109.5°

    11. H2O C. Common Molecular Shapes → Electronic Geometry = tetrahedral Hybridization = sp3 4 total 2 bond 2 lone BENT 104.5° <109.5°

    12. PI5 C. Common Molecular Shapes → Electronic Geometry = trigonal bipyramidal Hybridization = dsp3 5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90°

    13. KrF4 C. Common Molecular Shapes → Electronic Geometry = trigonal bipyramidal Hybridization = dsp3 5 total 4 bond 1 lone SEESAW <120°/<90°

    14. ClF3 C. Common Molecular Shapes → Electronic Geometry = trigonal bipyramidal Hybridization = dsp3 5 total 3 bond 2 lone T-SHAPE <90°

    15. I31- C. Common Molecular Shapes → Electronic Geometry = trigonal bipyramidal Hybridization = dsp3 5 total 3 bond 2 lone LINEAR 180°

    16. SH6 C. Common Molecular Shapes → Electronic Geometry = octahedral Hybridization = d2sp3 6 total 6 bond 0 lone OCTAHEDRAL 90°

    17. IF5 C. Common Molecular Shapes → Electronic Geometry = octahedral Hybridization = d2sp3 6 total 5 bond 1 lone SQUARE PYRAMIDAL <90°

    18. SF4 C. Common Molecular Shapes → Electronic Geometry = octahedral Hybridization = d2sp3 6 total 4 bond 2 lone SQUARE PLANAR 90°

    19. O O Se O D. Examples • SeO3 3 total 3 bond 0 lone E.G. = TRIGONAL PLANAR M.G. = TRIGONAL PLANAR 120°

    20. H As H H D. Examples • AsH3 4 total 3 bond 1 lone E.G. = TETRAHEDRAL M.G. = TRIGONAL PYRAMIDAL 107° (<109.5°)

    21. E. Hybridization • Provides information about molecular bonding and molecular shape • Several atomic orbitals mix to form same total of equivalent hybrid orbitals

    22. E. Hybridization • Carbon is common example (orbital diagram) • One of 2s electrons is promoted to 2p • 4 identical orbitals form sp3 hybridization

    23. Remember the subscript is the orbital, not e- configuration! E. Hybridization • Other types of hybridization • Be – 2 ve- forms sp • Al – 3 ve- forms sp2 • Si – 4 ve- forms sp3 • Kr – 8 ve- forms dsp3 • S – 6 ve- forms d2sp3 exceptions

    24. F. Hybridization Example • Compare shapes and hybrid orbitals: PF3 PF5