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 VSPER Model—Determining Molecular Shape
|Total number of single bonds, double bonds,
and lone pairs on the central atom
||Structural pair geometry
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
the structural pair geometry for the molecule by arranging the electron
pairs so that the repulsions are minimized (based on the table).
the table above to determine the molecular geometry.
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
Three areas of electron density: trigonal
Four areas of electron density: tetrahedral
Five areas of electron density: trigonal
Six areas of electron density: octahedral
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