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Molecular Geometry

>> Parts of this equation/concept include:

A -	Polarity

The geometry of a molecule is determined by arrangement of atoms around the central atom. There are two theories that explain this arrangement.

VSEPR theory predicts that groups of electrons will repel each other. Therefore these groups will arrange to be as far apart as possible. Of course, with fewer groups, the electrons can be further apart. A group of electrons is a bond (single, double, or triple doesn't matter) or a lone pair. The arrangement for different groups is listed in Table 7.2.

Because molecular geometry is the arrangement of atoms, lone pairs are not included in the molecular geometry. They do, however, influence the geometry in the basic arrangement of atoms, since they count as a repelling group and their strong repelling force pushes the bonds closer together than another bond would. Table 7.2 lists the appropriate names for the different arrangements and the angles between bonds.

Valence bond theory also predicts molecular geometry. In valence bond theory, the orbitals hybridize into a geometry that is the same as that predicted by VSEPR theory. Either theory gets the same result.

>> Example 1

What is the molecular geometry of the following?

1. XeF₄
2. PCl₅
3. I₃^–
4. CO₂
5. NO₃^–
6. SiCl₄
7. SF₆

Solution:

The Lewis structures for each of these atoms are shown in the previous examples (expanded octets and valence bond theory). You need the Lewis structure before you can make predictions about molecular geometry.

1. Xenon tetrafluoride has two lone pairs and four bonds around the central atom. That is a total of six groups. Therefore the arrangement of electrons is an octahedral. However, the lone pairs do not show up in the molecular geometry. Because lone pairs are particularly repulsive, they will get as far from each other as possible. In this example that is on opposite sides of the octahedral. Therefore the overall molecular geometry is square planar.

2. Phosphorus pentachloride has five bonds around the central atom. The electrons arrange in a trigonal bipyramid. Since there are no lone pairs, the molecular geometry is also a trigonal bipyramid.

3. Triiodide ion has three lone pairs and two bonds. The arrangement of electrons is a trigonal bipyramid. In this arrangement the lone pairs of electrons prefer equatorial positions to axial positions. This is because the 120° angle of the equatorial positions keeps the electrons further apart than the 90° angle of the axial position. With a lone pair taking each of the three equatorial positions, the remaining atoms have a linear geometry.

4. The carbon of carbon dioxide has two bonds. Its geometry (both electronic and molecular) is linear.

5. Nitrate ion has three bonds. (That this has resonance structures is irrelevant. Each structure gives the same geometry. It must be this way because the true structure is not any one of the three, but the average of them all.) With three bonds the molecular geometry is a planar triangle.

6. Silicon tetrachloride has four bonds and a tetrahedral molecular geometry.

7. Sulfur hexafluoride has six bonds and an octahedral geometry.

In a bond, the more electronegative atom attracts a bigger share of the shared electrons than the other. This separation of charge results in a bond dipole. When all the bond dipoles of a molecule are added, the net dipole is called the permanent dipole moment. When the bond dipoles cancel, the molecule's dipole moment is zero and the molecule is called nonpolar. If the bond dipoles do not cancel, the molecule is polar.

Lone pairs contribute to the polarity of a molecule. Since there is no positive nucleus to offset the negative charge of the lone pair, the end with the lone pair will always have a higher partial negative charge than any atom.

The key to determining whether or not dipoles cancel is to consider their direction as well as their magnitude. The direction is determined by the molecular geometry. The Lewis structure does not represent the molecular geometry.

A good clue is to look at the symmetry of the molecular geometry. Molecules that are not symmetric are polar.

Because ions already have a net charge, the partial charges due to a dipole aren't relevant. Consequently, no one bothers to classify ions as polar or nonpolar.

>> Example 2

Are the following molecules polar or nonpolar?

1. XeF₄
2. PCl₅
3. CO₂
4. SCl₂

Solution:

The Lewis structures of the molecules are shown in previous examples.

1. Because the lone pairs are exactly opposite of each other, they cancel. The bond dipoles of the xenon-fluorine bond are also exactly opposite each other. All bond dipoles cancel; the molecule is nonpolar.

2. This molecule has a trigonal bipyramid structure. The phosphorus-chlorine bond dipoles are all of the same magnitude. The axial bond dipoles cancel each other. The equatorial bond dipoles also cancel. Although the bond dipoles aren't exactly opposite, the angle of the two opposing bonds is sufficient to cancel that dipole. Thus the whole molecule is nonpolar. Any molecule where all terminal atoms are the same and with no lone pairs on the central atom is nonpolar.

3. This molecule has a linear structure. The carbon-oxygen double-bond dipoles cancel. The molecule is nonpolar.

4. This molecule has a bent structure, with two lone pairs on one side and the chlorine atoms on the other. Therefore it is a polar molecule, with the partial negative charge on the side with the lone pairs.

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