Valence Shell Electron Pair Repulsion theory (VSEPR) was developed by chemists to rationalize the results of experiments which were designed to provide information about the shapes of simple molecules. For example, when methane, CH4, is treated with dichlorine, Cl2, there is only one product formed that has the molecular formula CH3Cl. This result suggests that all the C-H bonds in methane are identical. This raises the question, "How can you arrange four hydrogen atoms around a carbon atom so that all the CH bonds are identical?" One possibilty is to place the carbon at the center of a square with a hydrogen atom at each of the corners: all five atoms lie in the same plane. Another option is to place the carbon at the center of a cube and to position the four hydrogen atoms at alternating corners of the cube. These two arrangements are called square planar and tetrahedral, respectively. Figure 1 illustrates both alternatives.
In Figure 1, the dashed lines are intended to emphasize the square and the cube, while the solid lines represent the bonds between the carbon atom and the hydrogens.
Both of the shapes depicted in Figure 1 are consistent with the experimental facts described above. Further experiments were required to differentiate between them. One such experiment involved conversion of CH4 into CH2ClBr. The experimental details of how this is done are less important than the results that were obtained: only one compound was isolated. This result eliminates the square planar geometry as an option for methane since you would expect two different structures for CH2ClBr if CH4 were square planar. Figure 2 illustrates these stuctures.
If methane were tetrahedral, then replacement of one H with Cl and another with Br would produce only one structure regardless of which hydrogens were replaced. You should make molecular models of "tetrahdral " CH2ClBr to convince yourself that this is true. If you would like to simulate experiments like those just described, go to How Do Chemists Know That Methane Is Tetrahedral?
These experiments, and many others, revealed that the shapes of countless molecules could be classified into eight groups: linear, bent, trigonal planar, square planar, tetrahedral, pyramidal, trigonal bipyramidal, and octahedral. These shapes are illustrated in Figure 3.
When chemists talk about the shape of a molecule, they are referring to the distribution of two or more atoms about a central atom. In Figure 3 the central atom is represented by the letter X. The atoms bonded to X are indicated by the symbol An, where n is used as a counter to keep track of the number of atoms attached to X. If you would like to examine three dimensional models of molecules with the shapes shown in Figure 3, go to the Interactive Molecular Model Kit created by Dr. Otis Rothenberger at Illinois State University.
Before you can use VSEPR theory to predict the shape of a molecule, you must be able to draw a proper Lewis structure for that molecule. Once you have done so, follow these rules:
If the number of atoms bonded to the central atom equals the number of regions of electron density around the central atom, then the geometry of the regions of electron density is the same as the shape of the molecule.
If the number of atoms bonded to the central atom is less than the number of regions of electron density around the central atom, then the shape of the molecule varies with the number of lone pairs of electrons on the central atom:
Notice the distinction between geometry and shape. For the purposes of this discussion the word geometry refers to the spatial distribution of regions of electron density around a central atom. The word shape describes the positions of the atoms that are bonded to the central atom. The distinction is necessitated by the fact that lone pairs of electrons occupy space even though they are not involved in bonding to other atoms.
Table 1 illustrates these rules for molecules containing central atoms which can have a maximum of 8 electrons in their valence shell.
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Central atom |
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Atoms bonded to central atom |
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Regions of electron density |
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Geometry of electron density |
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Number of lone pairs |
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Shape |
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In order to appreciate how and why VSEPR theory works, you must recall Coulomb's Law. In particular, you must recall that statement of Coulomb's Law which says "like charges repel". In terms of VSEPR theory, this means areas of electron density (i.e. bonds and non-bonding electron pairs) around a central atom will repel each other. They will get as far apart from each other as possible in order to minimize their interaction. Such a configuration has the lowest potential energy, i.e. is the most stable arrangement of elctron pairs around the central atom. To appreciate this idea, consider the alternative shapes for methane shown in Figure 1. In the square planar arrangement the H-C-H bond angles are all 90o. In the tetrahedral case, the H-C-H angles are 109.5o. In other words, the electron pairs are farther apart when they are arranged in a tetrahedral geometry than in a square planar configuration.
Another point that's worth reiterating: while chemists use the word shape to describe the positions of atoms around a central atom, it is the distribution of electron density around that central atom that determines those positions. Consider methane, CH4, and ammonia, NH3. The former molecule is tetrahedral while the latter is pyramidal. In both cases there are four pairs of electrons that have to be distributed around the central atom, but in the case of methane each pair forms a bond to a hydrogen atom. In the case of ammonia, three of the electron pairs form bonds to hydrogen atoms, but the fouth does not; it is a non-bonding pair. Even though there are only three atoms to be distributed around the nitrogen, there are four electron pairs. Figure 4 compares the two structures.
Now compare NH3 with BF3 as shown in Figure 5. In this case both central atoms are
bonded to three other atoms, but the nitrogen atom has four electron pairs to be distributed around it while the boron atom has only three. The F-B-F bond angles will be maximized when the three electron pairs connecting the boron atom to the fluorine atoms all lie in the same plane.
The first row elements B,C,N, O, and F, can accomodate a maximum of 8 electrons in their valence shell orbitals. This limits the geometries around these atoms to five of the eight possibilities shown in Figure 3; linear, bent, trigonal planar, pyramidal, and tetrahedral. In terms of orbital hybridization, the linear geometry corresponds to a situation in which the central atom is sp hybridized. When a molecule is trigonal planar, three pairs of electrons are distributed around the central atom and that atom is sp2 hybridized. The hybridization about the central atom is sp3 when that atom has four electron pairs around it. Note that the latter case applies to bent, pyramidal, and tetrahedral geometries. Figure 6 depicts three molecules in which the central atom uses sp3 hybridized orbitals to form bonds to other atoms.
As the top row of the figure indicates, the central atom has four pairs of electrons around it in all three molecules. Bonding pairs are shown as solid lines, while non-bonding pairs are represented by two dots. The shapes of these molecules become obvious when the orbitals are removed as shown in the bottom row of the figure.
Additional information about VSEPR theory may be found at:
http://www.shef.ac.uk/chemistry/vsepr/chime-false/vsepr.html
http://www2.gasou.edu/chemdept/general/molecule/tutorial/index.htm