The Foundations of Chemistry

D. Unshared Valence Electron Pairs in Octahedral Electronic Geometry

We can reason along the lines used in part D of Section 8-11 to predict the preferred
locations of lone pairs on the central atom in octahedral electronic geometry. Because of
the high symmetry of the octahedral arrangement, all six positions are equivalent, so it
does not matter in which position in the drawing we put the first lone pair. AB 5 U mole-
cules and ions are described as having square pyramidalmolecular geometry. When a
second lone pair is present, the most stable arrangement has the two lone pairs in two
octahedral positions at 180° angles from each other. This leads to a square planarmole-
cular geometry for AB 4 U 2 species. These arrangements are shown in Figure 8-3. Table
8-3 summarizes a great deal of information—study this table carefully.

Two lone pairs at 90° from each other

would be much more crowded.

8-12 Octahedral Electronic Geometry: AB 6 , AB 5 U, and AB 4 U 2 337

Problem-Solving Tip:Placing Lone Pairs on the Central Atom

Remember that lone pairs occupy more space than bonded pairs, so the lone pairs are

always put in positions where they will be least crowded.

If the Lewis formula for a molecule or ion shows only one lone pair:In linear, trigonal planar,

tetrahedral, or octahedral electronic geometry, all positions are equivalent, so it doesn’t

matter where we place the lone pair. In trigonal bipyramidal electronic geometry, place

the lone pair in the equatorial position where it is least crowded, and put the bonded

atoms in the other positions.

If the Lewis formula shows two lone pairs:In trigonal planar or tetrahedral electronic

geometry, we can place the lone pairs in any two positions and the bonded atoms in the

other position(s). In trigonal bipyramidal electronic geometry, place the two lone pairs

in two equatorial positions (120° apart) where they are least crowded, and put the bonded

atoms in the other positions. In octahedral electronic geometry, place the two lone pairs

in two positions across(180°) from each other, and put the bonded atoms in the other

positions.

Figure 8-3 Arrangements of bonded atoms and lone pairs (six regions of high electron
density—octahedral electronic geometry).

AB 5 U 5 bonded atoms (B)

1 lone pair (U)

Square pyramidal molecular geometry

Examples: IF 5 , BrF 5

AB 4 U 2 4 bonded atoms (B)

2 lone pairs (U)

Square planar molecular geometry

Examples: XeF 4 , IF 4 –