|SHAPES OF MOLECULES|
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?
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
|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°.|
|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:
These are both variations of the trigonal bipyramid. The
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|
(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).
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