7.2 Electron Pair Geometry versus Molecular Structure

Electron-Pair Geometry versus Molecular Structure

It is important to note that electron-pair geometry around a central atom is not the same thing as its molecular structure. The electron-pair geometries shown in Figure 7.2.3 describe all regions where electrons are located, bonds as well as lone pairs. Molecular structure describes the location of the atoms, not the electrons.

We differentiate between these two situations by naming the geometry that includes all electron pairs the electron-pair geometry. The structure that includes only the placement of the atoms in the molecule is called the molecular structure. The electron-pair geometries will be the same as the molecular structures when there are no lone electron pairs around the central atom, but they will be different when there are lone pairs present on the central atom.

VSEPR structures like the one in Figure 7.2.1 are often drawn using the wedge and dash notation, in which solid lines represent bonds in the plane of the page, solid wedges represent bonds coming up out of the plane, and dashed lines represent bonds going down into the plane. For example, the methane molecule, [latex]\ce{CH4}[/latex], which is the major component of natural gas, has four bonding pairs of electrons around the central carbon atom; the electron-pair geometry is tetrahedral, as is the molecular structure (Figure 7.2.4). On the other hand, the ammonia molecule, [latex]\ce{NH3}[/latex], also has four electron pairs associated with the nitrogen atom, and thus has a tetrahedral electron-pair geometry. One of these regions, however, is a lone pair, which is not included in the molecular structure, and this lone pair influences the shape of the molecule (Figure 7.2.2).

As seen in Figure 7.2.2, small distortions from the ideal angles in Figure 7.2.6 can result from differences in repulsion between various regions of electron density. VSEPR theory predicts these distortions by establishing an order of repulsions and an order of the amount of space occupied by different kinds of electron pairs. The order of electron-pair repulsions from greatest to least repulsion is:

lone pair-lone pair  >  lone pair-bonding pair  >  bonding pair-bonding pair

This order of repulsions determines the amount of space occupied by different regions of electrons. A lone pair of electrons occupies a larger region of space than the electrons in a triple bond; in turn, electrons in a triple bond occupy more space than those in a double bond, and so on. The order of sizes from largest to smallest is:

lone pair  >  triple bond  >  double bond  >  single bond

Consider formaldehyde, [latex]\ce{H2CO}[/latex], which is used as a preservative for biological and anatomical specimens (Figure 1). This molecule has regions of high electron density that consist of two single bonds and one double bond. The basic geometry is trigonal planar with 120° bond angles, but we see that the double bond causes slightly larger angles (121°), and the angle between the single bonds is slightly smaller (118°).

In the ammonia molecule, the three hydrogen atoms attached to the central nitrogen are not arranged in a flat, trigonal planar molecular structure, but rather in a three-dimensional trigonal pyramid (Figure 7.2.5) with the nitrogen atom at the apex and the three hydrogen atoms forming the base. The ideal bond angles in a trigonal pyramid are based on the tetrahedral electron pair geometry. Again, there are slight deviations from the ideal because lone pairs occupy larger regions of space than do bonding electrons. The [latex]\ce{H-N-H}[/latex] bond angles in [latex]\ce{NH3}[/latex] are slightly smaller than the 109.5° angle in a regular tetrahedron (Figure 7.2.3) because the lone pair-bonding pair repulsion is greater than the bonding pair-bonding pair repulsion (Figure 7.2.2). Figure 7.2.3 illustrates the ideal molecular structures, which are predicted based on the electron-pair geometries for various combinations of lone pairs and bonding pairs.
According to VSEPR theory, the terminal atom locations (Xs in Figure 7.2.3) are equivalent within the linear, trigonal planar, and tetrahedral electron-pair geometries (the first three rows of the table). It does not matter which X is replaced with a lone pair, because the molecules can be rotated to convert positions. For trigonal bipyramidal electron-pair geometries, however, there are two distinct X positions, as shown in Figure 7.2.4: an axial position (if we hold a model of a trigonal bipyramid by the two axial positions, we have an axis around which we can rotate the model) and an equatorial position (three positions form an equator around the middle of the molecule). As shown in Figure 7.2.6, the axial position is surrounded by bond angles of 90°, whereas the equatorial position has more space available because of the 120° bond angles. In a trigonal bipyramidal electron-pair geometry, lone pairs always occupy equatorial positions because these more spacious positions can more easily accommodate the larger lone pairs.

Theoretically, we can come up with three possible arrangements for the three bonds and two lone pairs for the [latex]\ce{ClF3}[/latex] molecule (Figure 7.2.4). The stable structure is the one that puts the lone pairs in equatorial locations, giving a T-shaped molecular structure.
Figure 7.2.4. (a) In a trigonal bipyramid, the two axial positions are located directly across from one another, whereas the three equatorial positions are located in a triangular arrangement. (b–d) The two lone pairs (red lines) in [latex]\ce{ClF3}[/latex] have several possible arrangements, but the T-shaped molecular structure (b) is the one actually observed, consistent with the larger lone pairs both occupying equatorial positions.

When a central atom has two lone electron pairs and four bonding regions, we have an octahedral electron-pair geometry. The two lone pairs are on opposite sides of the octahedron (180° apart), giving a square planar molecular structure that minimizes lone pair-lone pair repulsions (Figure 7.2.3).

Predicting Electron Pair Geometry and Molecular Structure

The following procedure uses VSEPR theory to determine the electron pair geometries and the molecular structures:

1. Write the Lewis structure of the molecule or polyatomic ion.
2. Count the number of regions of electron density (lone pairs and bonds) around the central atom. A single, double, or triple bond counts as one region of electron density.
3. Identify the electron-pair geometry based on the number of regions of electron density: linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral (Figure 7.2.6, first column).
4. Use the number of lone pairs to determine the molecular structure (Figure 7.2.6). If more than one arrangement of lone pairs and chemical bonds is possible, choose the one that will minimize repulsions, remembering that lone pairs occupy more space than multiple bonds, which occupy more space than single bonds. In trigonal bipyramidal arrangements, repulsion is minimized when every lone pair is in an equatorial position. In an octahedral arrangement with two lone pairs, repulsion is minimized when the lone pairs are on opposite sides of the central atom.

The following examples illustrate the use of VSEPR theory to predict the molecular structure of molecules or ions that have no lone pairs of electrons. In this case, the molecular structure is identical to the electron pair geometry.

Example 7.2.1: Predicting Electron-pair Geometry and Molecular Structure: [latex]\ce{CO2}[/latex] and [latex]\ce{BCl3}[/latex]

Predict the electron-pair geometry and molecular structure for each of the following:

1. carbon dioxide, [latex]\ce{CO2}[/latex], a molecule produced by the combustion of fossil fuels
2. boron trichloride, [latex]\ce{BCl3}[/latex], an important industrial chemical

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1. We write the Lewis structure of [latex]\ce{CO2}[/latex] as:
   
   This shows us two regions of high electron density around the carbon atom—each double bond counts as one region, and there are no lone pairs on the carbon atom. Using VSEPR theory, we predict that the two regions of electron density arrange themselves on opposite sides of the central atom with a bond angle of 180°. The electron-pair geometry and molecular structure are identical, and [latex]\ce{CO2}[/latex] molecules are linear.
2. We write the Lewis structure of [latex]\ce{BCl3}[/latex] as:
   
   Thus we see that [latex]\ce{BCl3}[/latex] contains three bonds, and there are no lone pairs of electrons on boron. The arrangement of three regions of high electron density gives a trigonal planar electron-pair geometry. The [latex]\ce{B-Cl}[/latex] bonds lie in a plane with 120° angles between them. [latex]\ce{BCl3}[/latex] also has a trigonal planar molecular structure (Figure 7.2.8).

Figure 7.2.5 shows the electron-pair geometry and molecular structure of [latex]\ce{BCl3}[/latex] are both trigonal planar. Note that the VSEPR geometry indicates the correct bond angles (120°), unlike the Lewis structure shown above.

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Example 7.2.2: Predicting Electron-pair Geometry and Molecular Structure: Ammonium

Two of the top 50 chemicals produced in the United States, ammonium nitrate and ammonium sulfate, both used as fertilizers, contain the ammonium ion. Predict the electron-pair geometry and molecular structure of the [latex]\ce{NH4+}[/latex] cation.

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We write the Lewis structure of [latex]\ce{NH4+}[/latex] as:

We can see that NH₄⁺ contains four bonds from the nitrogen atom to hydrogen atoms and no lone pairs. We expect the four regions of high electron density to arrange themselves so that they point to the corners of a tetrahedron with the central nitrogen atom in the middle (Figure 7.2.6). Therefore, the electron pair geometry of [latex]\ce{NH4+}[/latex] is tetrahedral, and the molecular structure is also tetrahedral (Figure 9).

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The next several examples illustrate the effect of lone pairs of electrons on molecular structure.

Example 7.2.3: Predicting Electron-pair Geometry and Molecular Structure: Lone Pairs on the Central Atom

Predict the electron-pair geometry and molecular structure of a water molecule.

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The Lewis structure of [latex]\ce{H2O}[/latex] indicates that there are four regions of high electron density around the oxygen atom: two lone pairs and two chemical bonds:

We predict that these four regions are arranged in a tetrahedral fashion (Figure 10), as indicated in Figure 7.2.6. Thus, the electron-pair geometry is tetrahedral and the molecular structure is bent with an angle slightly less than 109.5°. In fact, the bond angle is 104.5°.

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Example 7.2.4: Predicting Electron-pair Geometry and Molecular Structure: [latex]\ce{SF4}[/latex]

Sulfur tetrafluoride, [latex]\ce{SF4}[/latex], is extremely valuable for the preparation of fluorine-containing compounds used as herbicides (i.e., [latex]\ce{SF4}[/latex] is used as a fluorinating agent). Predict the electron-pair geometry and molecular structure of a [latex]\ce{SF4}[/latex] molecule.

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The Lewis structure of [latex]\ce{SF4}[/latex] indicates five regions of electron density around the sulfur atom: one lone pair and four bonding pairs:

We expect these five regions to adopt a trigonal bipyramidal electron-pair geometry. To minimize lone pair repulsions, the lone pair occupies one of the equatorial positions. The molecular structure (Figure 11) is that of a seesaw (Figure 7.2.6).

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Example 7.2.5: Predicting Electron-pair Geometry and Molecular Structure: [latex]\ce{XeF4}[/latex]

Of all the noble gases, xenon is the most reactive, frequently reacting with elements such as oxygen and fluorine. Predict the electron-pair geometry and molecular structure of the [latex]\ce{XeF4}[/latex] molecule.

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The Lewis structure of [latex]\ce{XeF4}[/latex] indicates six regions of high electron density around the xenon atom: two lone pairs and four bonds:

These six regions adopt an octahedral arrangement (Figure 7.2.6), which is the electron-pair geometry. To minimize repulsions, the lone pairs should be on opposite sides of the central atom (Figure 12). The five atoms are all in the same plane and have a square planar molecular structure.

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Glossary

axial position: location in a trigonal bipyramidal geometry in which there is another atom at a 180° angle and the equatorial positions are at a 90° angle

electron-pair geometry: arrangement around a central atom of all regions of electron density (bonds, lone pairs, or unpaired electrons)

equatorial position: one of the three positions in a trigonal bipyramidal geometry with 120° angles between them; the axial positions are located at a 90° angle

molecular structure: structure that includes only the placement of the atoms in the molecule