3.7 Electronegativity, polar bonds, & Vsepr theory  (Page 3/27)

Consider formaldehyde, H ₂ CO, which is used as a preservative for biological and anatomical specimens ( [link] ). This molecule has regions of high electron density that consist of two single bonds and one double bond. The basic geometry is trigonal planar with 120° bond angles, but we see that the double bond causes slightly larger angles (121°), and the angle between the single bonds is slightly smaller (118°).

In the ammonia molecule, the three hydrogen atoms attached to the central nitrogen are not arranged in a flat, trigonal planar molecular structure, but rather in a three-dimensional trigonal pyramid ( [link] ) with the nitrogen atom at the apex and the three hydrogen atoms forming the base. The ideal bond angles in a trigonal pyramid are based on the tetrahedral electron pair geometry. Again, there are slight deviations from the ideal because lone pairs occupy larger regions of space than do bonding electrons. The H–N–H bond angles in NH ₃ are slightly smaller than the 109.5° angle in a regular tetrahedron ( [link] ) because the lone pair-bonding pair repulsion is greater than the bonding pair-bonding pair repulsion ( [link] ). [link] illustrates the ideal molecular structures, which are predicted based on the electron-pair geometries for various combinations of lone pairs and bonding pairs.

According to VSEPR theory, the terminal atom locations (Xs in [link] ) are equivalent within the linear, trigonal planar, and tetrahedral electron-pair geometries (the first three rows of the table). It does not matter which X is replaced with a lone pair because the molecules can be rotated to convert positions. For trigonal bipyramidal electron-pair geometries, however, there are two distinct X positions, as shown in [link] : an axial position    (if we hold a model of a trigonal bipyramid by the two axial positions, we have an axis around which we can rotate the model) and an equatorial position    (three positions form an equator around the middle of the molecule). As shown in [link] , the axial position is surrounded by bond angles of 90°, whereas the equatorial position has more space available because of the 120° bond angles. In a trigonal bipyramidal electron-pair geometry, lone pairs always occupy equatorial positions because these more spacious positions can more easily accommodate the larger lone pairs.

Theoretically, we can come up with three possible arrangements for the three bonds and two lone pairs for the ClF ₃ molecule ( [link] ). The stable structure is the one that puts the lone pairs in equatorial locations, giving a T-shaped molecular structure.