3.3: Breaking the octet rule with higher electron counts (hypervalent atoms)

The octet rule applies well to atoms in the second row of the periodic table, where a full valence shell includes eight electrons with an electron configuration of \(s^2p^6\). Even elements in the third and fourth row are known to follow this rule sometimes, but not always. In larger atoms, where \(n\geq3\) the valence shell contains additional subshells: the \(d, f, g...\) subshells. Therefore, atoms with \(n\geq3\) can have higher valence shell counts by "expanding" into these additional subshells. When atoms contain more than eight electrons in their valence shell, they are said to be hypervalent. Hypervalency allows atoms with \(n\geq3\) to break the octet rule by having more than eight electrons. This also means they can have five or more bonds; something that is nearly unheard of for atoms with \(n\leq2\). Complete the exercises below to see examples of molecules containing hypervalent atoms.

Exercise \(\PageIndex{1}\)

Draw the Lewis structures for SF₆.

Answer

Exercise \(\PageIndex{1}\)

Draw the Lewis structure for ClF₃.

Answer

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Is hypervalency real? Not exactly. Hypervalency is a concept associated with hybrid orbital theory and Lewis theory. It's useful for some simple things, like predicting how atoms are connected and predicting molecular shape. But the idea that the d-orbitals are involved in bonding isn't accurate according to wave mechanics.

For main group molecules, chemists (like Pauling) thought a long time ago that hypervalence is due to expanded s²p⁶ octets. The consensus is now clear that d orbitals are NOT involved in bonding in molecules like SF₆ any more than they are in SF₄ and SF₂. In all three cases, there is a small and roughly identical participation of d-orbitals in the wavefunctions. This has been established in both MO and VB theory. However using Hybrid orbitals with d-orbital contributions equips us with a language which can pragmatically describe the geometries of highly coordinated substances.

Hybrid orbitals are a powerful tool to describe the geometries and shape of molecules and metal complexes. However, in "real" molecules, their significance may be debated. Often a more realistically (and more complicated) molecular orbitals approach is needed. However, from an epistemologically simple point of view, bonding theories can only be judged by their predictions. To the extent that hybridization can explain the shapes of PF₅ and SF₆, valence bond theory is a perfectly good theory. To the extent that if you write out the valence bond wavefunction using hybridized orbitals and calculate energies and other properties à la Pauling (i.e., ionization energy and electron affinities) and find them to be off from experimental results (by tens of kcals/mol), then valence bond theory is not accurate.

Bonding theories can only be judged by their predictions.

A simple explanation that can be given is that molecular wavefunctions constructed out of hybridized atomic orbitals is accurate enough to predict some things, but not others. Predictions of any theory must be compared with empirical evidence to assess when they work and when they fail. When a theory gives the wrong answer, at least one assumption must not hold. In this case, the valence bond wavefunction is not accurate enough to capture some important features of a system's electronic structure. It may not be the most intellectually satisfying answer, but to say more would result in a much more complicated answer and certainly far beyond the level reasonably expected from general chemistry discussions.