Structure and Bonding Chemistry^3 Mindmap

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  • Created on: 17-12-22 17:39
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  • Structure and Bonding
    • The Lewis Model
      • octet rule = atoms combine so in a molecule, each atom has 8 electrons in its valence shell
        • non-bonding electrons are lone pairs
      • formal charge = (no. e- in valence shell of free atom) - (no. bonds to atom) - (no. unshared e-)
      • hypervalent = more than 8 electrons in its outer shell
    • Valence shell electron pair repulsion theory
      • 3 assumptions:
        • 1) atoms in a molecule are held together by pairs of e-'s = bonding pairs
        • 2) some molecules may have pairs of e-'s not involved with bonding = lone pairs
        • 3) electron pairs adopt positions as far apart as possible
      • VSEPR for tetrahedra and octahedra
        • 4e- pairs and 0LP = tetrahedral
          • 1LP = trigonal pyramidal
          • 2LP = bent
        • 6e- pairs and 0LP = octahedral
          • 1LP = square pyramidal
          • 2LP = square planar
      • structures based on trigonal bipyramid
        • eg. SF4 [most favourable is lone pair in equatorial]
        • eg. ClF3
          • LP both at equatorial sites, giving T shaped molecule
            • most favourable as has smallest no. repulsions involving LP
          • LP both at axial sites, giving trigonal planar molecule
          • 1LP at axial, 1LP at equatorial, giving pyramidal molecule
      • Compounds containing multiple bonds
        • POCl3 = tetrahedral
        • CO3 2- = trigonal planar
      • Limitations:
        • doesn't apply to d-block metal complexes
        • doesn't work for TeBr6 2-
    • Bond polarity and polar molecules
      • bonds between different atoms are normally polar and magnitude depends on electronegativity difference
      • whether a polyatomic molecule is polar or not depends on its shape and whether polarities of bonds cancel out
      • polarity of a solvent affects nature of compounds that dissolve in it
    • Valence bond theory for polyatomic molecules
      • sp3 hybridization
        • interactions between 2s orbital and 3 2p orbitals to form 4 sp3 hybrid orbitals
        • eg. methane. CH4 is tetrahedral as the geometry minimises the repulsions between e- pairs in C-H bonds
        • eg. ammonia. 1 of the sp3 orbitals contains a lone pair, N-H bond formed between a N sp3 hybrid orbital and a H 1s orbital
      • sp2 hybridization
        • 2s, 2px, 2py combine to give 3 2sp2 orbitals
        • the unhybridized pz orbital interact to form a pi bond
        • eg. ethene, BH3
      • sp hybridization
        • 2s and 2pz orbitals hybridize to form 2 sp hybrid orbitals
        • 1s orbital forms a C-H bond, unhybridized 2px and 2py form 2 pi bonds
        • eg. ethyne, BeH2
    • Resonance
      • explains bonding in molecules with bonds that appear to be different in a lewis structure but are the same experimentally
      • don't exist independently - actual structure is the average and has lower energy than any individual form
      • eg. benzene c6h6
      • hypervalency
        • in compounds of the 3rd period, is explained using hybridization schemes involving 1 or more d orbitals or assuming charge separation in the resonance forms
        • explanations why it never exceeds the octet rule: relatively small and the extra electrons would need to go into an empty orbital and are all too high in energy to be occupied


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