In the axial arrangement, shared pairs are situated “top and bottom.” In the equatorial arrangement, shared pairs surround Xe. The equatorial arrangement is more stable since the lone pairs are 180˚ apart and this minimizes their repulsion. In both molecular arrangements, the electronic geometry is octahedral, with 90˚ angles. The top figure has a molecular geometry known as “seesaw,” while the bottom figure has a molecular geometry that is more stable, known as square planar.
Example
Draw the dot formula for SeF4 and determine the hybridization at Se.
Explanation
First determine the number of valence electrons this molecule has: SeF4 has 6 + 4(7) = 34 valence electrons, which is equal to 17 pairs of electrons.
Selenium is surrounded by four fluorines and a lone pair of electrons. That’s five sites of electron density, which translates into sp3d hybridization. Se is from the fourth period, so it may have an expanded octet.
So, to recap, focus on the number of binding “sites” or areas of concentrated electron density:
Two areas of electron density: linear, planar molecule
Three areas of electron density: trigonal planar molecule
Four areas of electron density: tetrahedral molecule
Five areas of electron density: trigonal bipyramidal molecule
Six areas of electron density: octahedral molecule
Molecular Polarity
In chemical bonds, polarity refers to an uneven distribution of electron pairs between the two bonded atoms—in this case, one of the atoms is slightly more negative than the other. But molecules can be polar too, and when they are polar, they are called dipoles. Dipoles are molecules that have a slightly positive charge on one end and a slightly negative charge on the other. Look at the water molecule. The two lone electron pairs on the oxygen atom establish a negative pole on this bent molecule, while the bound hydrogen atoms constitute a positive pole. In fact, this polarity of water accounts for most of water’s unique physical properties. However, molecules can also contain polar bonds and not be polar. Carbon dioxide is a perfect example. Both of the C—O bonds in carbon dioxide are polar, but they’re oriented such that they cancel each other out, and the molecule itself is not polar.
