Exceptions to the Octet Rule: Expanded, Incomplete and Odd-Electron Molecules

The octet rule predicts the structures of thousands of compounds: every atom ends up with eight valence electrons, except hydrogen, which has two. But chemistry is full of important molecules that break it. Some atoms have fewer than eight electrons, some have more, and some molecules simply can’t give every atom an octet because they have an odd number of electrons. Knowing the exceptions, and why they occur, is essential for drawing correct Lewis structures and predicting shapes.

Type 1: Incomplete octets (fewer than eight)

Some atoms, especially small elements in groups 2 and 13, form stable molecules with fewer than eight electrons around them.

1. Boron trifluoride, BF₃

• Boron has 3 valence electrons and forms three single bonds to fluorine.
• Electrons around boron: 3 bonding pairs = 6.
• Giving boron an octet would require a B=F double bond, which would place a positive formal charge on the very electronegative fluorine. The structure with six electrons around boron is the better description.

Because boron is electron-deficient, BF₃ is a strong Lewis acid: it readily accepts a lone pair, for example from ammonia, to form F₃B–NH₃, in which boron does have an octet. See Lewis acids and bases.

2. Beryllium chloride, BeCl₂ (gas phase)

• Beryllium has 2 valence electrons and forms two bonds.
• Electrons around Be: 4.
• In the solid state, BeCl₂ forms chains in which chlorine atoms bridge between beryllium atoms, partly making up the deficiency.

3. Aluminium chloride, AlCl₃

• As an isolated molecule, aluminium has only 6 electrons.
• In the vapour at moderate temperatures, two AlCl₃ units join to form Al₂Cl₆, with two chlorine atoms bridging, giving each aluminium an octet. This is a neat example of an electron-deficient compound finding a way to complete its octets.

4. Boranes

Boron hydrides such as diborane (B₂H₆) have too few electrons for ordinary two-electron bonds everywhere. They use three-centre two-electron bonds, in which two electrons hold three atoms together (B–H–B bridges).

Type 2: Odd-electron molecules (radicals)

If a molecule has an odd total number of valence electrons, at least one electron must be unpaired, so not every atom can have an octet.

5. Nitrogen monoxide, NO

• Valence electrons: 5 (N) + 6 (O) = 11, an odd number.
• The best Lewis structure leaves nitrogen with 7 electrons.
• NO is a radical: reactive, and paramagnetic. It’s an important signalling molecule in the body (involved in blood vessel dilation) and a pollutant from car engines.

6. Nitrogen dioxide, NO₂

• Valence electrons: 5 + 6 + 6 = 17.
• One electron remains unpaired. Two NO₂ molecules readily pair up to form dinitrogen tetroxide, N₂O₄, pairing the odd electrons. That’s why brown NO₂ gas becomes paler when cooled: the equilibrium shifts towards colourless N₂O₄. See Le Chatelier’s principle.

7. Chlorine dioxide, ClO₂

• Valence electrons: 7 + 6 + 6 = 19.
• A radical used on a large scale to bleach wood pulp and disinfect water.

Odd-electron species are paramagnetic because of their unpaired electron. See paramagnetic vs diamagnetic and how to count unpaired electrons.

Type 3: Expanded octets (more than eight)

Elements in period 3 and below (P, S, Cl, Xe, I and others) can be surrounded by more than eight electrons in their compounds. These are often called hypervalent molecules.

8. Phosphorus pentachloride, PCl₅

• Phosphorus forms five P–Cl bonds: 10 electrons around phosphorus.
• Shape: trigonal bipyramidal.
• Compare with nitrogen, in the same group, which only forms NCl₃, never NCl₅.

9. Sulfur hexafluoride, SF₆

• Six S–F bonds: 12 electrons around sulfur.
• Shape: octahedral. SF₆ is extremely stable and unreactive, and it’s used as an insulating gas in high-voltage electrical equipment. It’s also a very powerful greenhouse gas, so its release is carefully controlled.

10. Xenon tetrafluoride, XeF₄

• Four Xe–F bonds plus two lone pairs: 12 electrons around xenon.
• Shape: square planar.
• Its existence (first made in 1962) overturned the belief that noble gases never form compounds. See noble gases.

11. Triiodide ion, I₃⁻

• The central iodine has two bonds and three lone pairs: 10 electrons.
• This ion forms when iodine dissolves in potassium iodide solution, and it’s the species responsible for the blue-black starch–iodine colour.

12. Sulfuric acid, H₂SO₄ (in some representations)

• Sulfuric acid is often drawn with two S=O double bonds, giving sulfur 12 electrons.
• It can also be drawn with only single bonds and formal charges (S²⁺ with O⁻ groups), keeping sulfur to an octet. Modern bonding calculations suggest the charge-separated picture is closer to reality, but both forms appear in textbooks.

Why can period 3 elements exceed eight?

The traditional explanation is that period 3 elements have empty 3d orbitals available to hold extra electrons. Modern calculations show that d orbitals play only a small role; the more important factors are:

• Size: larger central atoms can fit more atoms around them.
• Electronegative partners: hypervalent compounds almost always involve very electronegative atoms such as F, O or Cl, which draw electron density away from the central atom, making the bonds partly ionic.
• Delocalised bonding, such as three-centre four-electron bonds, which spread electrons over several atoms.

Period 2 elements (C, N, O, F) are too small and have no low-energy orbitals beyond 2s and 2p, so they never exceed an octet.

Type 4: Transition metal compounds

Transition metals use d orbitals in bonding and rarely follow the octet rule. For many metal complexes, an 18-electron rule is a better guide (s² p⁶ d¹⁰ = 18), especially for organometallic compounds, though it too has many exceptions. See transition metals.

Recognising exceptions in exam questions

1. Count the total valence electrons. If the total is odd, it’s an odd-electron species.
2. Look at the central atom. Boron, beryllium and aluminium often have incomplete octets.
3. Count the atoms attached. If a period 3 or lower central atom has more than four atoms (or bonds plus lone pairs) around it, expect an expanded octet.
4. Remember period 2 limits. C, N, O and F never exceed eight.

Summary table

Type	Examples	Electrons on central atom
Incomplete octet	BF₃, BeCl₂, AlCl₃ (monomer)	4 or 6
Odd-electron (radical)	NO, NO₂, ClO₂	odd number, e.g. 7
Expanded octet	PCl₅, SF₆, XeF₄, I₃⁻	10 or 12
Transition metal	many complexes	often 18

One final tip: when a structure seems to break the octet rule, first recount the valence electrons. Many apparent “exceptions” in homework turn out to be counting errors.

Key takeaways

• The octet rule is a useful guide, not a law; many stable molecules break it.
• Incomplete octets (BF₃, BeCl₂) make molecules electron-deficient and often Lewis acids.
• Odd-electron molecules (NO, NO₂) are radicals, paramagnetic and often reactive.
• Period 3 and heavier elements can have expanded octets (PCl₅, SF₆, XeF₄), especially with electronegative partners.
• Period 2 elements never exceed eight electrons; transition metals often follow an 18-electron guide instead.