Structure and Bonding Chemistry^3 Mindmap
- Created by: Strelly22
- 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
- octet rule = atoms combine so in a molecule, each atom has 8 electrons in its valence 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
- 4e- pairs and 0LP = tetrahedral
- 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
- LP both at equatorial sites, giving T shaped 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-
- 3 assumptions:
- 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
- sp3 hybridization
- 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
- The Lewis Model
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