Exceptions to the Octet Rule

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:

1. Odd-electron molecules have an odd number of valence electrons and therefore have an unpaired electron.
2. Electron-deficient molecules have a central atom with fewer electrons than needed for a noble gas configuration.
3. Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.

Odd-electron Molecules

Molecules that contain an odd number of electrons are called radicals. Nitric oxide, NO, is an example of an odd-electron molecule; it is produced in internal combustion engines when oxygen and nitrogen react at high temperatures.

To draw the Lewis structure for an odd-electron molecule like NO, the following steps are considered:

1. Determine the total number of valence (outer shell) electrons. The sum of the valence electrons is 5 (from N) + 6 (from O) = 11. The odd number indicates that it is a free radical, where not every atom has eight electrons in its valence shell.
2. Draw a skeleton structure of the molecule. A skeleton structure with an N–O single bond can easily be drawn.
3. Distribute the remaining electrons as lone pairs on the terminal atoms. In this case, there is no central atom, so the electrons are distributed around both atoms. Eight electrons are assigned to the more electronegative atom in these situations; thus, oxygen has the filled valence shell:
   
4. Place all remaining electrons on the central atom. Since there are no remaining electrons, this step does not apply.
5. Rearrange the electrons to make multiple bonds with the central atom in order to obtain octets wherever possible. Although an odd-electron molecule cannot have an octet for every atom, each atom should get electrons as close to an octet as possible. In this case, nitrogen has only five electrons around it. To move closer to an octet for nitrogen, one of the lone pairs from oxygen is utilized to form a NO double bond. (Another lone pair of electrons cannot be taken from oxygen to form a triple bond because nitrogen would then have nine electrons:)
   

Electron-deficient Molecules

Some molecules, however, contain central atoms that do not have a filled valence shell. Generally, these are molecules with central atoms from groups 2 and 13, outer atoms that are hydrogen, or other atoms that do not form multiple bonds. For example, in the Lewis structures of beryllium dihydride, BeH₂, and boron trifluoride, BF₃, the beryllium and boron atoms each have only four and six electrons, respectively. It is possible to draw a structure with a double bond between a boron atom and a fluorine atom in BF₃, satisfying the octet rule, but experimental evidence indicates the bond lengths are closer to that expected for B–F single bonds. This suggests the best Lewis structure has three B–F single bonds and electron-deficient boron. The reactivity of the compound is also consistent with electron-deficient boron. However, the B–F bonds are slightly shorter than what is actually expected for B–F single bonds, indicating that some double-bond character is found in the actual molecule.

An atom like the boron atom in BF₃, which does not have eight electrons, is very reactive. It readily combines with a molecule containing an atom with a lone pair of electrons. For example, NH₃ reacts with BF₃ because the lone pair on nitrogen can be shared with the boron atom:

Hypervalent Molecules

Elements in the second period of the periodic table (n = 2) can accommodate only eight electrons in their valence shell orbitals because they have only four valence orbitals (one 2s and three 2p orbitals). Elements in the third and higher periods (n ≥ 3) have more than four valence orbitals and can share more than four pairs of electrons with other atoms because they have empty d orbitals in the same shell. Molecules formed from these elements are sometimes called hypervalent molecules, such as PCl₅, and SF₆. In PCl₅, the central atom, phosphorus, shares five pairs of electrons. In SF₆, sulfur shares six pairs of electrons.

In some hypervalent molecules, such as IF₅ and XeF₄, some of the electrons in the outer shell of the central atom are lone pairs:

In the Lewis structures for these molecules, there are electrons left over after filling the valence shells of the outer atoms with eight electrons. These additional electrons must be assigned to the central atom.

This text is adapted from Openstax, Chemistry 2e, Section 7.3: Lewis Symbols and Structures.