11 central or interior atoms which form the bonds have the least tendency to

11 central or interior atoms which form the bonds

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11. central or interior atoms, which form the bonds, have the least tendency to hybridize. 12. we focus on the hybridization of the interior atoms and assume that all terminal atoms, those bonding to only one other atom, are hybridization: a) the number of standard atomic orbitals added together always equals the number of hybrid orbitals formed. (1) the total number of orbitals is conserved. b) the particular combinations of standard atomic orbitals added together determines the shapes and energies of the hybrid orbitals formed. c) the particular type of hybridization that occurs is the one that yields the lowest overall energy for the molecule. (1) since the actual energy calculations are beyond the scope of this book, we will use electron geometries by VSEPR theory to predict the type of hybridization. B. sp 3 Hybridization 1. the notation sp 3 indicates that the hybrid orbitals are mixtures of one s orbital and three p orbitals. a) notice that the hybrid orbitals all have the same energy, they are degenerate. 2. the four hybrid orbitals are arranged in a tetrahedral geometry with 109.5 angles between them. 3. hybridized orbital readily form chemical bonds because they tend to maximize overlap with other orbitals. a) however, if the central atom of a molecule contains lone pairs, hybrid orbitals can also accommodate them. C. sp 2 Hybridization and Double Bonds 1. hybridization of one s and two p orbitals results in three sp 2 hybrids and one leftover unhybridized p orbital. 2. notice that the three hybrid orbitals have a trigonal planar geometry with 120 degree angles between them. a) the unhybridized p orbitals is oriented perpendicular to the three hybridized orbitals. 3. each of the sp 2 orbitals is half-filled. a) the remaining electron occupies the leftover p orbital, even though it is slightly higher in energy. 4. when p orbitals overlap side by side, the resulting bond is called a pi ( ) bond, and the π electron density is above and below the intermolecular axis. 5. when orbitals overlap end to end, a in all of the rest of the bonds in the molecule, the resulting bond is call a sigma ( ) bond. σ 6. even though we represent the two electrons in a pi bond as two arrows in the upper lobe, they are actually spread out over both the upper and lower lobes (this is one of the limitations we encounter when we try to represent electrons with arrows). 7. we can now label all the bonds in the molecule using a notation that specifies the type of bond as well as the type of overlapping orbitals. 8. notice the correspondence between valence bond theory and Lewis theory.
Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory and Molecular Orbital Theory 7 a) double bonds in Lewis theory always correspond to one sigma and one pi bond in valence bond theory.

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