Even though many molecules seem to obey the VSEPR Theory and provide an accurate assumption of its shape, there are quite a few molecules that fail to agree with the representation provided by VSEPR. A few examples will be given and illustrated in depth.

Let's take a look at the sulfur dioxide molecule. We know that both oxygen and sulfur have six valence electrons each. Furthermore, there must be 6 + 6 + 6 = 18 / 2 = 9 pairs of electrons. Since sulfur is the central atom and since it has a pair of unshared electrons, the sulfur dioxide molecule could be represented in two different ways.

Now lets look at some molecules which do not obey the octet rule at all. For example, boron trihydride is one case. Remember that boron has three sp² hybrid orbitals; thus, only three hydrogen atoms can fit in. Because of this, boron can only accommodate three pairs of electrons, rather than four pairs predicted by the octet rule. The shape given to all boron compounds, and most elements in Group 13, is a trigonal planar shape.

In Phosphorus pentachloride, there are five chlorine atoms bonded to a single phosphorus atom. Since each chlorine atom shares one pair of electrons, it would mean that the phosphorus atom has a valence shell containing 10 electrons, rather than eight as dictated by the octet rule. How is this possible? The five electrons each have parallel spins because of dsp³ hybridization and thus, a d orbital was used. Such molecules with five orbitals used are shaped like a trigonal bipyramidal.

Likewise, in sulfur hexahydride (SH₆) and iodine hepatahydride (IH₇), there are twelve and fourteen shared electrons, respectively. The iodine heptahydride has a shape of a pentagonal bipyramid, while sulfur hexahydride has a shape of an octahedron. Only sulfur hexahydride is shown to the right. How are these electron configurations possible? It is a more complex hybridization, known as dsp³ and d²sp³ hybridization. Both types of hybridization are due to the s, p, and d orbitals mixing to produce equal energy orbitals. These hybridizations will only be discussed in the advanced topics.

One final case is with molecules with an odd number of electrons. For example, nitrogen(II) oxide has seven electrons in the nitrogen atom and eight in the oxygen atom, resulting in 15 electrons; eleven of these are valence electrons. However, when diagramming nitrogen(II) oxide, there are three possible dot diagrams as shown to the left. For such molecules with odd number of electrons, there exists no dot diagrams that obey the octet rule. In this case, resonance structures are used.

We are nearing the end of our quest! Only a few more sections remain before we finally understand the nature of covalent bonds. The next section discusses what distinguishes covalent molecules from ionic compounds. Press 'Next Section' to continue.