Читать книгу Foundations of Chemistry - Philippa B. Cranwell - Страница 101

When a central atom has four sets of electron density surrounding it, the areas of electron density are based upon a tetrahedron, but the final shape of the molecule depends on the number of bonded pairs. These shapes are called tetrahedral, pyramidal and, bent (or angular). Figure 2.16 shows examples of molecules that adopt each of the three shapes. When there are four single bonds to the central atom, e.g. in methane, the molecule adopts a tetrahedral shape (Figure 2.16a). When there are three single bonds to the central atom and one lone pair of electrons, e.g. in ammonia, NH₃, the molecule adopts a pyramidal shape (Figure 2.16b). When there are two bonds and two pairs of electrons around the central atom, e.g. water, H₂O, the molecule adopts a bent or angular shape (Figure 2.16c). The lone pairs have an impact on the bond angles between the bonded atoms and the central atoms because the lone pairs occupy more space than the bonding pairs. In a symmetrical tetrahedral molecule, the angle between each of the covalent bonds is 109.5°; in a pyramidal compound, the angle is 107°; and in a bent compound, it is 104.5°. This reflects the larger space that the lone pairs occupy, pushing the bonding electrons closer together; see Figure 2.16.

Figure 2.16 (a) Bonding angles in a tetrahedral bonding centre; (b) bonding angles in a pyramidal bonding centre; (c) bonding angles in a bent (or angular) bonding centre.

The rule of thumb for deciding the order of interactions between areas of electron density in a molecule is:

 Lone pair–lone pair > lone pair–bonded pair > bonded pair–bonded pair