Hybridization and molecular shape

The shapes of these molecules cannot result from bonding between simple s and p atomic orbitals. Although s and p orbitals have the lowest energies for isolated atoms in space, they are not the best for forming bonds. To explain the shapes of common organic molecules, we assume that the s and p orbitals combine to form hybrid atomic orbitals that separate the electron pairs more widely in space and place more electron density in the bonding region between the nuclei. 

Common bond angles. Bond angles in organic compounds are usually close to 109°, 120°, or 180°. 

Formation of a pair of sp hybrid atomic orbitals. Addition of an, v orbital to a p orbital gives an sp hybrid atomic orbital, with most of its electron density on one side of the nucleus. Adding the p orbital with opposite phase gives the other sp hybrid orbital, with most of its electron density on the opposite side of the nucleus from the first hybrid. 

Draw the Lewis structure for beryllium hydride, BeH2. Draw the orbitals that overlap in the bonding of BeH2, and label the hybridization of each orbital. Predict the H—Be — H bond angle. 

There are only four valence electrons in BeH2 (two from Be and one from each H), so the Be atom cannot have an octet. The bonding must involve orbitals on Be that give the strongest bonds (the most electron density in the bonding region) and also allow the two pairs of electrons to be separated as far as possible. 

Structure of a double bond Tbe first pair of electrons forms a sigma bond. Tbe second pair forms a pi bond. 

Tbe pi bond bas its electron density centered in two lobes, above and below tbe sigma bond. Together, tbe two lobes of tbe pi bonding molecular orbital constitute one bond. 

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Although orbital hybridizations and molecular shapes for hypovalent metal hydrides of the early transition metals and the normal-valent later transition metals are similar, the M—H bonds of the early metals are distinctly more polar. For example, metal-atom natural charges for YH3 (+1.70), HfH4 (+1.75), and TaHs (+1.23) are all significantly more positive than those (ranging from +0.352 to —0.178) for the homoleptic hydrides from groups 6-10. Indeed, the empirical chemistry of early transition-metal hydrides commonly reveals greater hydricity than does that of the later transition-metal hydrides. 

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