SAT Subject Test Chemistry






VSEPR—Electrostatic Repulsion

Properties of molecules depend not only on the bonding of atoms but also on the molecular geometry—the three-dimensional arrangement of the molecule”s atoms in space. The combination of the polarity of the bonds and the geometry of the molecule determine the molecular polarity. This can be defined as the uneven distribution of the molecular charge. The chemical formula reveals little information about a molecule”s geometry. It is only after doing many tests designed to reveal the shapes of the various molecules that chemists developed two different yet equally successful theories to explain certain aspects of their findings. One theory accounts structurally for molecular bond angles. The other is used to describe changes in the orbitals that contain the valence electrons of a molecule”s atoms. The structural theory that deals with the bond angles is called the VSEPR theory, whereas the one that describes changes in the orbitals that contain the valence electrons is called the hybridization theory. (VSEPR represents Valence Shell Electron Pair Repulsion.)


Two theories explain molecular structure: VSEPR theory uses valence shell electron pair repulsion. Hybridization theory uses changes in the orbitals of the valence electrons.

VSEPR uses as its basis the fact that like charges will orient themselves in such a way as to diminish the repulsion between them.

1. Mutual repulsion of two electron clouds forces them to the opposite sides of a sphere. This is called a linear arrangement.

EXAMPLE: BeF2, berylium fluoride

2. Minimum repulsion between three electron pairs occurs when the pairs are at the vertices of an equilateral triangle inscribed in a sphere. This arrangement is called a trigonal-planar arrangement.


These basic arrangements are important to learn!

EXAMPLE: BF3, boron trifluoride

3. Four electron pairs are farthest apart at the vertices of a tetrahedron inscribed in a sphere. This arrangement is called a tetrahedral-shaped distribution of electron pairs.


Configurations often appear as questions on the SAT test.

EXAMPLE: CH4, methane

4. Mutual repulsion of six identical electron clouds directs them to the corners of an inscribed regular octahedron. This is said to have an octahedral arrangement.

EXAMPLE: SF6, sulfur hexafluoride

VSEPR and Unshared Electron Pairs

Ammonia, NH3, and water, H2O, are examples of molecules in which the central atom has both shared and unshared electron pairs. Here is how the VSEPR theory accounts for the geometries of these molecules.

The Lewis structure of ammonia shows that, in addition to the three electron pairs the central nitrogen atom shares with the three hydrogen atoms, it also has one unshared pair of electrons:

VSEPR theory postulates that the lone pair occupies space around the nitrogen atom just as the bonding pairs do. Thus, as in the methane molecule shown in the preceding section, the electron pairs maximize their separation by assuming the four corners of a tetrahedron. Lone pairs do occupy space, but our description of the observed shape of a molecule refers to the positions of atoms only. Consequently, as shown in the drawing below, the molecular geometry of an ammonia molecule is that of a pyramid with a triangular base. The general VSEPR formula for a molecule such as ammonia (NH3) is AB3E, where A replaces N, B replaces H, and E represents the unshared electron pair.

A water molecule has two unshared electron pairs and can be represented as an AB2E2 molecule. Here, the oxygen atom is at the center of a tetrahedron, with two corners occupied by hydrogen atoms and two by the unshared pairs, as shown below. Again, VSEPR theory states that the lone pairs occupy space around the central atom but that the actual shape of the molecule is determined only by the positions of the atoms. In the case of water, this results in a “bent,” or angular, molecule.

VSEPR and Molecular Geometry

The following table summarizes the molecular shapes associated with particular types of molecules. Notice that, in VSEPR theory, double and triple bonds are treated in the same way as single bonds. It is helpful to use the Lewis structures and this table together to predict the shapes of molecules with double and triple bonds, as well as the shapes of polyatomic ions.


Know these molecular and Lewis structures.


The molecular configurations derived by VSEPR can also be arrived at through the concept of hybridization. Briefly stated, this means that chemists envision that two or more pure atomic orbitals (usually s, p, and d ) can be mixed to form two or more new hybrid atomic orbitals that are identical and conform to the known shapes of molecules. Hybridization can be illustrated as follows:

1. sp Hybrid Orbitals

Spectroscopic measurements of beryllium fluoride, BeF2, reveal a bond angle of 180° and equal bond lengths.

The ground state of beryllium is:

To accommodate the experimental data, we theorize that a 2s electron is excited to a 2p orbital; then the two orbitals hybridize to yield two identical orbitals called sp orbitals. Each contains one electron but is capable of holding two electrons.

2. sp 2 Hybrid Orbitals
Boron trifluoride, BF3, has bond angles of 120° of equal strength. To accommodate these data, the boron atom hybridizes from its ground state of 1s22s22p 1 to:

3. sp 3 Hybrid Orbitals
Methane, CH4, can be used to illustrate this hybridization. Carbon has a ground state of 1s 22s 22p 2. One 2s electron is excited to a 2p orbital, and the four involved orbitals then form four new identical sp 3 orbitals.

In some compounds where only certain sp 3 orbitals are involved in bonding, distortion in the bond angle occurs because of unbonded electron repulsion. Examples:

a. Water, H2O

b. Ammonia, NH3

4. sp3d2 Hybrid Orbitals
These orbitals are formed from the hybridization of an s and a p electron promoted to d orbital s and transformed into six equal sp3d2 orbitals. The spatial form is shown below. Sulfur hexafluoride, SF6, illustrates this hybridization.

The concept of hybridization is summarized in the accompanying table.


Know these hybrid orbitals designations and their corresponding shapes.