Exceptions to the Octet Rule in Chemistry

Instructor: Richard Craven

Dr. Craven (Doc Rico) has taught the majority of disciplines in chemistry as well as environmental and life sciences. He has a doctorate in Chemistry.

In this lesson, we will study molecular situations that do not obey the octet rule and use Lewis structures and formal charge calculations to determine the better molecular structure for these situations. Afterwards, you can test your understanding using a short quiz.

Identifying the Violators

The main-group elements from Period 2 onward on the Periodic Table, for the most part, follow the octet rule, which is the rule that states that atoms try to achieve greater energy stability by having a completely filled a valence shell of eight electrons when molecular bonds or often when forming single atom ions. However, there are exceptions to this tried and true rule. These exceptions fall into three categories:

  • Molecules with incomplete octets
  • Molecules with an odd number of valence-shell electrons
  • Molecules with an expanded octet on a central atom

Molecules with Incomplete Octets

There are times when a possible, valid Lewis structure (a 2D representation of the bonding between atoms in a molecule) will not give a valence-shell octet to every atom in the structure once all valence-shell electrons are totaled. The central atoms with these incomplete octets most often encountered are Be, B, and Al. An example of such a molecule is boron trifluoride, BF3. A typical Lewis structure for this molecule leaves the boron deficient in valence electrons:

Lewis and 3D structures for BF3

This structure is a legitimate structure accounting for all 24 valence electrons of the atoms and having a zero formal charge on each atom:

Formal Charge on an atom (FC) = (total # of valence electrons in free atom) - (total # of nonbonding electrons) - ½(total # of electrons in bonds to atom)

FC (B) = 3 - 0 - ½(6) = 0

FC (F) = 7 - 6 - ½(2) = 0

There is also another Lewis structure (actually three resonance structures) that does allow this molecule to follow the octet rule:

Resonance Lewis Structures, BF3

These structures give an octet to each atom but non-zero formal charges on some of the atoms:

FC (B) = 3 - 0 - ½(8) = -1

FC (F) = 7 - 4 - ½(4) = +1 (Fluorine with double bond to boron)

FC (F) = 7 - 6 - ½(2) = 0 (Fluorines with single bonds to boron)

Even though it's frowned upon to have a positive formal charge on the more electronegative element, it is allowed. Of the four possible structures shown above for boron trifluoride, experimental results have shown that a hybrid combination of all four of them is the best structure for this molecule.

Molecules with an Odd Number of Valence-shell Electrons

Not all molecules have an even number of valence-shell electrons with which to form the molecule. This typically results with one atom not satisfying the octet rule by having an odd number of valence-shell electrons in the final structure. Three such examples are nitrogen monoxide (NO), nitrogen dioxide (NO2), and chlorine dioxide (ClO2) (all gases with two having resonance Lewis structures):

Lewis Structure for NO

Resonance Structures for NO2

Resonance Structures for ClO2

If we were to look at the formal charges for each example, we would see that NO has zero formal charges whereas the others have two atoms with balanced, opposite non-zero formal charges.

>NO

FC (N) = 5 - 3 - ½(4) = 0

FC (O) = 6 - 4 - ½(4) = 0

>NO2

FC (N) = 5 - 1 - ½(6) = +1

FC (O) = 6 - 4 - ½(4) = 0 (Oxygen with double bond)

FC (O) = 6 - 6 - ½(2) = -1 (Oxygen with single bond)

>ClO2

FC (Cl) = 7 - 4 - ½(4) = +1

FC (O) = 6 - 5 - ½(2) = 0 (Oxygen with odd # nonbonding electrons)

FC (O) = 6 - 6 - ½(2) = -1 (Oxygen with even # nonbonding electrons)

There is also another group of molecules one might encounter that also have odd number of valence-shell electrons. This group is comprised of free radicals which are molecules or atoms that are mostly fragments of other molecules that exist for only a brief period of time during chemical reactions and are very highly reactive which explains their being short-lived in chemical reactions.

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