While Lewis dot structures can tell us how the atoms in molecules are bonded to each other, they don’t tell us the shape of the molecule. In this section, we’ll discuss the methods for predicting molecular shape. The most important thing to remember when attempting to predict the shape of a molecule based on its chemical formula and the basic premises of the VSPER model is that the molecule will assume the shape that most minimizes electron pair repulsions. In attempting to minimize electron pair repulsions, two types of electron sets must be considered: electrons can exist in bonding pairs, which are involved in creating a single or multiple covalent bond, or nonbonding pairs, which are pairs of electrons that are not involved in a bond, but are localized to a single atom.

The above table represents a single atom with all of the electrons that would be associated with it as a result of the bonds it forms with other atoms plus its lone electron pairs. However, since atoms in a molecule can never be considered alone, the shape of the actual molecule might be different from what you’d predict based on its structural pair geometry. You use the structural pair geometry to determine the molecular geometry by following these steps:

The table below shows all of the commonly occurring molecular geometries that are found for molecules with four or fewer bonding domains around their central atom.

As you can see from the table, atoms that have normal valence—meaning atoms that have no more than four structural electron pairs and obey the octet rule (and have no lone pairs)—are tetrahedral. For instance, look at methane, which is CH₄:

Ammonia (NH₃), which has three sigma bonds and a lone pair, however, is trigonal pyramidal:

Water (H₂O) has two lone pairs and its molecular geometry is “bent,” which is also called V shaped:

So as you can see, lone pairs have more repulsive force than do shared electron pairs, and thus they force the shared pairs to squeeze more closely together.

As a final note, you may remember that we mentioned before that only elements with a principal energy level of 3 or higher can expand their valence and violate the octet rule. This is because d electrons are necessary to make possible bonding to a fifth or sixth atom. In XeF₄, there are two lone pairs and four shared pairs surrounding Xe, and two possible arrangements exist:

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.

Draw the dot formula for SeF₄ and determine the hybridization at Se.

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

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.