We begin by writing the Lewis structure. The H atoms are terminal atoms. There are three central atoms and (3 x 1) + 4 + 6 + 4 + (3 x 1) = 20 valence electrons, or 10 pairs. A plausible Lewis structure is drawn at right. Each central atom is surrounded by four electron pairs, requiring sp3 hybridization. The valence-shell orbital diagrams for the atoms follow. [Pg.235]

Look again at Table 9-3 on the previous page. Notice that the number of atomic orbitals mixed to form the hybrid orbital equals the total number of pairs of electrons. In addition, the number of hybrid orbitals formed equals the number of atomic orbitals mixed. For example, AICI3 has a total of three pairs of electrons and VSEPR predicts a trigonal planar molecular shape. To have this shape, one s and two p orbitals on the central atom A1 must mix to form three identical sp hybrid orbitals. [Pg.261]

To see why allylic radicals are so stable, look at the orbital picture in Figure 10.3. The radical carbon atom with an unpaired electron can adopt sp2 hybridization, placing the unpaired electron in a p orbital and giving a structure that is electronically symmetrical. The p orbital on the central carbon can therefore overlap equally well with a p orbital on either of the two neighboring carbons. [Pg.341]

Long before it was. possible to perform MO calculations on even the simplest molecules, the equivalence of the bonds led to the development of a different conception of the bonding in AB molecules, in which nonequivalent AOs on the central atom are combined into hybrid orbitals. These hybrid orbitals provide a set of equivalent lobes directed at the set (or subset) of symmetry equivalent B atoms. It is therefore obvious that all A—B bonds to all equivalent B atoms will be equivalent. [Pg.222]

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