Exceptions to the Octet Rule

The octet rule has three kinds of exceptions: molecules with an odd number of electrons (such as NO and NO₂), molecules in which the central atom has fewer than eight electrons (such as BF₃ and BeCl₂), and molecules in which the central atom has more than eight electrons, called expanded octets (such as PCl₅ and SF₆). In brief: boron in BF₃ has only six electrons, nitrogen in NO has seven, and sulfur in SF₆ has twelve; elements of the second period never exceed an octet, but elements of the third period and below can.

The octet rule works well for carbon, nitrogen, oxygen and fluorine, but many important compounds do not follow it. Knowing these exceptions is necessary to draw correct Lewis structures and to understand the reactivity of molecules such as boron trifluoride, nitrogen dioxide and sulfur hexafluoride. This article describes each type of exception, explains the reasons, summarizes examples in a table, works through the Lewis structures with electron counts and gives some history.

What are odd-electron molecules?

A molecule with an odd total number of valence electrons cannot have all of its atoms with complete octets, since electrons cannot all be paired. The unpaired electron makes the species a free radical, which is usually very reactive and paramagnetic. Examples include nitrogen monoxide (NO, 11 electrons), nitrogen dioxide (NO₂, 17 electrons) and chlorine dioxide (ClO₂, 19 electrons). NO₂ dimerizes at low temperatures to give N₂O₄, in which all electrons are paired. Many radicals are important in the atmosphere and in combustion, and the hydroxyl radical OH has a central role in atmospheric chemistry.

What are incomplete octets?

Some compounds of beryllium, boron and aluminum have a central atom with fewer than eight electrons. In BF₃, boron forms three bonds and has only six electrons. Forming a double bond with fluorine would put a positive formal charge on fluorine, the most electronegative element, which is unfavorable, so the structure with an incomplete octet is preferred. Such molecules are strong Lewis acids, which means they accept a lone pair from another molecule. BF₃ reacts with ammonia to form F₃B–NH₃, in which boron has an octet. Aluminum chloride exists as dimers, Al₂Cl₆, in which chlorine bridges supply the missing electrons. Beryllium chloride, BeCl₂, has only four electrons around beryllium in the gas phase. Hydrogen, with a duet, and lithium compounds are other examples of atoms that do not have an octet.

What are expanded octets?

Atoms of the third period and below (P, S, Cl, Br, I, Xe and others) are bigger and can accommodate more than eight electrons around them. They are called hypervalent. They form compounds such as PCl₅ (10 electrons), SF₆ (12 electrons), ClF₃ (10), XeF₄ (12) and the triiodide ion I₃⁻ (10). Expanded octets are found mostly with very electronegative atoms (F, O, Cl) bonded to the central atom. The atoms of the second period (C, N, O, F) cannot expand their octet, since there are no low-lying orbitals available for more than four pairs. This is why NF₅ does not exist but PF₅ does.

Summary of exceptions

Worked examples

1. Boron trifluoride, BF₃. Total = 3 + 3 × 7 = 24 electrons. Three B–F bonds use 6 electrons, and the three fluorine octets need 3 × 6 = 18 more electrons. The total is 24. Boron has 6 electrons. Formal charges: B 3 − 0 − 3 = 0, F 7 − 6 − 1 = 0. All zero, so the structure is accepted, though the octet of boron is incomplete.
2. Phosphorus pentachloride, PCl₅. Total = 5 + 35 = 40. Five P–Cl bonds use 10 electrons. Each chlorine needs 6 more electrons: 5 × 6 = 30. The total is 40. Phosphorus has 10 electrons.
3. Sulfur hexafluoride, SF₆. Total = 6 + 42 = 48. Six S–F bonds use 12 electrons. The six fluorine atoms use 6 × 6 = 36. The total is 48. Sulfur has 12 electrons, in an octahedral arrangement.
4. Xenon tetrafluoride, XeF₄. Total = 8 + 28 = 36. Four Xe–F bonds use 8, the fluorine lone pairs use 4 × 6 = 24, and the remaining 36 − 32 = 4 electrons form two lone pairs on xenon. Xenon has 4 bonding pairs and 2 lone pairs, 12 electrons. The molecule is square planar.
5. Nitrogen dioxide, NO₂. Total = 5 + 12 = 17 electrons. Two N–O bonds use 4 electrons and the oxygen octets use 12 more, which leaves 1 electron, placed on nitrogen. Nitrogen then has only 5 electrons, so a lone pair from one oxygen is turned into a second N–O bond. The result is O=N–O with one unpaired electron on nitrogen. Count: N=O (4) + N–O (2) + unpaired electron (1) + double-bonded O lone pairs (4) + single-bonded O lone pairs (6) = 17 ✓. Formal charges: N 5 − 1 − 3 = +1; the double-bonded O 6 − 4 − 2 = 0; the single-bonded O 6 − 6 − 1 = −1. The sum is 0. The double bond can be on either oxygen, so there are two resonance structures.
6. Number of electron pairs in ClF₃. Total = 7 + 21 = 28. Three Cl–F bonds use 6 electrons; the fluorine lone pairs use 18; the remaining 4 electrons form two lone pairs on chlorine. Chlorine has 3 bonding pairs and 2 lone pairs, which makes 10 electrons.

Why do the exceptions occur?

For atoms with incomplete octets, the electronegativity of the atoms bonded to the central atom prevents π donation, or the central atom is small and has few valence electrons. For odd-electron molecules, there is no way to pair every electron. For hypervalent molecules, there are two explanations. The older explanation used the d orbitals of the central atom to hold the additional electrons (“sp³d” and “sp³d²” hybridization). Modern calculations show that d orbitals contribute little, and the bonding can be explained by the polarity of the bonds and by three-center, four-electron bonds, which are described by molecular orbital theory. For teaching at an introductory level, the expanded octet remains a convenient way to describe the electron count and the shape of the molecule.

How has the understanding changed?

Lewis and Langmuir formulated the octet rule in 1916 to 1919, and chemists soon found exceptions such as PCl₅ and SF₆. Linus Pauling proposed hybridization with d orbitals to explain them in the 1930s. In 1962, Neil Bartlett’s discovery of the first noble gas compound was a challenge to the idea of the octet as a rigid rule. Since the 1990s, computational chemistry has shown that d orbital participation is minor, and the description of hypervalent bonding changed. The octet rule remains a powerful guide for the lighter elements.

Which mistakes are common?

Mistake 1: Forcing an octet on boron

BF₃ with a B=F double bond has a positive formal charge on fluorine, which is unfavorable. The six-electron structure is usually preferred.

Mistake 2: Expanding the octet of a second-period element

C, N, O and F never have more than eight electrons. There is no NF₅, and the compound OF₆ does not exist.

Mistake 3: Ignoring the electron count of odd-electron species

If the total number of valence electrons is odd, there will be an unpaired electron.

Mistake 4: Placing extra electrons on outer atoms instead of the central atom

After the outer octets are complete, put the remaining electrons on the central atom, which may then exceed eight if it is in period 3 or below.