Shapes of Molecules

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).

A. Central atoms with less than octet configurations:

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 NH3the 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 NH3are 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°.

C. Central atoms with expanded octet configurations:

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:

Lewis dot structure

(2) Count the number of bond pairs and lone pairs around the central atom.

(3) Decide on the electron pair orientation based on the total number of electron pairs (4 = tetrahedral, 5 = trigonal bipyramidal).

(4) Consider the placement of lone pairs and any distortions from "regular" shapes.

(5) Name the shape based on the location of atoms (nuclei).

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

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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).

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°.