Molecular shape

Molecular Shape

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 non-bonding pairs, which are pairs of electrons that are not involved in a bond, but are localized to a single atom.

The VSPER Model—Determining Molecular Shape

Total number of single bonds, double bonds, and lone pairs on the central atom      Structural pair geometry       Shape

2        Linear

3        Trigonal planar

4        Tetrahedral

5        Trigonal bipyramidal

6        Octahedral

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:

Draw the Lewis dot structure for the molecule and count the total number of single bonds, multiple bonds, and unpaired electrons.

Determine the structural pair geometry for the molecule by arranging the electron pairs so that the repulsions are minimized (based on the table).

Use the table above to determine the molecular geometry.

For instance, look at methane, which is CH4:

simple m5

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:

simple m6

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

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