SHAPES OF MOLECULES

Understanding the shapes of molecules is an important first step in being able to discuss and predict chemical properties. We shall discuss here some "simple" molecules. This topic, however, has important applications in understanding the behavior of much larger molecules. Much of biochemistry is now being discussed based on how macromolecules are shaped, and how different molecules "fit" together.

It is observed that the SF2 molecule is bent, the F-S-F bond angle being 98°. The BeCl2 molecule, however, is linear. Why are these two AX2 type molecules so different?
 
SF2 BeCl2

To understand any molecule, one must first complete a Lewis dot structure. It is then possible to predict the molecular shape using TWO BASIC PRINCIPLES:

1. The shapes of molecules are determined by the repulsion between electron pairs in the outer shell of the central atom. Both bond pairs (electron pairs shared by two atoms) and lone pairs (those located on a central atom but not shared) must be considered.

2. Lone pairs repel more than bond pairs.

The application of these principles is best seen by referring to specific examples. We shall start looking only at molecules with single bonds (for simplicity).
 

The BeCl2 molecule has a Lewis dot structure as shown. The central Be atom has two bond pairs in its outer shell. Repulsion between these two pairs (first principle above) causes the molecule to be linear (see above). If the molecule were bent in any direction, the two bond pairs would be brought closer together, increasing the repulsion.
 
 
The molecule BF3 has a dot symbol as: BF3
Here the B atom has three bond pairs in its outer shell. Minimizing the repulsion causes this molecule to have a trigonal planar shape, with the F atoms forming an equilateral triangle about the B atom. The F-B-F bond angles are all 120°, and all the atoms are in the same plane.

B. Central atoms with octet (noble gas) configurations:
 
The CH4 molecule has a dot structure as shown. The shape of this molecule, however, is not planar, as is suggested by the way we draw this dot structure.  CH4
Carbon has 4 bond pairs. The four H atoms are arranged about the C atom in a tetrahedral shape . This shape minimizes the repulsion between the bond pairs. The 109.5° angle is the same for all H-C-H bond angles and is called the tetrahedral angle
There are many molecules that have four bond pairs and this regular tetrahedral shape; CCl4, SiF4, and SnCl4 are just a few examples.
The molecule NH3 has a dot symbol much like that for BF3 (see above). Now, however, there is a lone pair in the outer shell of the central N atom.  NH3
In NH3 the N has 3 bond pairs and 1 lone pair, (4 total pairs). The shape is called trigonal pyramidal (approximately tetrahedral minus one atom).

WHACHACALLIT: A shape name is based on what is experimentally observed - the location of atomic nuclei. In NH3 the N, with 4 e pairs, will have a tetrahedral electron pair orientation. (The total number of e pairs determines the electron pair orientation.) The lone pair occupies one corner of the tetrahedron. It is difficult to "see" lone pairs experimentally. Looking only at the atoms, we see a short, rather distorted tetrahedron. This is called a pyramid. The pyramids of Egypt have square bases. The NH3 pyramid has a triangular base. Hence the shape is called trigonal pyramidal.

LONE PAIR DISTORTIONS: Due to the greater repelling character of lone pairs, (second principle above) the H atoms in NH3 are bent closer together than the normal tetrahedral angle of 109.5°. In NH3 the observed angle is 107.3°. Other molecules with this one-lone, three-bond-pair configuration (:NCl3 and :PCl3) have this same trigonal pyramidal shape, slightly different bond angles, but all less than 109.5°.
 
The H2O molecule has this dot structure: H2O
The O in H2O has 2 bond pairs and 2 lone pairs (again, 4 total pairs). The electron pair orientation around O is tetrahedral. Two corners of the tetrahedron are "missing" because they are occupied by lone pairs, not atoms.  The shape is called bent. The H-O-H bond angle is 104.4°. This angle is less than that in NH3, due in part to the greater repulsions felt with two lone pairs
Other  molecules with 2 bond plus 2 lone pairs include OF2, H2S, and SF2.  Bond angles vary, but all are significantly less than 109.5°

A number of molecules have more electrons in the outer shell than in a noble gas configuration.  This involves use of one or more d orbital (so as to not violate the Pauli Exclusion Principle). We will not worry, for now, about the types or shapes of the orbitals involved in the bonding, but will only consider electron pair repulsions.
 
The PCl5 molecule has 5 bond pairs in the outer shell of P. This molecule has a symmetrical trigonal bipyramidal shape.  PCl5
Note that the Cl atoms occupy two types of positions.  The two Cl atoms which are on a straight line which passes through the P nucleus are said to occupy axial positions.  The other three Cl are in equatorial positions. 
The axial Cl to P bond distance is slightly longer than the equatorial Cl to P bond distance.  Why? (See below.)

Examples of molecules with 5 electron pairs, but with lone pairs, are SF4 and ClF3:
 
SF4
"See-Saw"
ClF3
Distorted "T"

These are both variations of the trigonal bipyramid. The lone pairs always occupy equatorial positions. Equatorial pairs have fewer 90° repulsions, and thus are at lower potential energy. This is consistent with the longer axial bond lengths seen for PCl5. The greater repulsion of these lone pairs distort (bend) the both the axial and equatorial Cl atoms slightly away, to the opposite side of the molecule.
 
The SF6 molecule is an example of a molecule with 6 bond pairs. The least repelling shape for this is octahedral. The 6 F atoms are located at the corners of the octahedron. SF6
In Summary: To predict the shape of a molecule:

Study problems: Work out the shapes and bond angles (including distortions) for:  IF5, I3-, and XeF4.

Click here for information About The Files accessed in this document.

Click here to E-mail comments to the author, Gardiner H. Myers: gmyers@chem.ufl.edu.