Octet rule and its limits
The octet rule is a compact description of a strong main-group pattern: many atoms form bonds that give them a noble-gas-like valence-shell count. It is especially useful for second-period elements, but it is not a fundamental law and it has well-defined limits.
Where the octet rule comes from
For C, N, O and F, the valence shell contains one 2s and three 2p orbitals: four orbitals holding at most eight electrons. Bonding often fills these states in ways that resemble the closed-shell configuration of neon.
Hydrogen follows a duet rule because its 1s shell holds only two electrons. Transition metals and many heavier main-group compounds require models beyond a simple octet count.
Electron-deficient molecules
BeH₂ and BF₃ are classic electron-deficient molecules. Forcing an octet onto boron by drawing unnecessary B=F double bonds can worsen the Lewis description because the experimental bonding and reactivity are better captured by electron-deficient B.
Electron deficiency often creates Lewis acidity: an empty orbital can accept a lone pair from a donor.
Odd-electron molecules and radicals
If the total valence-electron count is odd, every atom cannot have a paired-electron octet. NO has 11 valence electrons; NO₂ has 17. These species contain an unpaired electron and are radicals.
Radicals are not automatically unstable enough to be nonexistent. NO is a stable gas under ordinary conditions; many radical intermediates are detectable and central to combustion, atmospheric chemistry and polymerization.
Hypervalent Lewis structures
Third-period and heavier atoms can appear with more than eight electrons around the central atom in Lewis structures, as in PCl₅, SF₆ or XeF₄. Older explanations often invoked extensive use of low-lying d orbitals.
Modern bonding descriptions show that simple “expanded d-orbital octets” are generally not the right microscopic picture. Delocalized molecular orbitals, ionic resonance contributions and three-center bonding provide better explanations.
Formal charge, resonance and the better structure
The octet rule is only one criterion when choosing Lewis structures. Formal charges should be minimized when possible, negative formal charge is usually favored on more electronegative atoms, and equivalent resonance structures may distribute charge and bond order.
A Lewis structure is a bookkeeping model of valence electrons. It does not uniquely specify electron density, bond order or geometry.
A decision path
Count electrons
Odd total? Expect at least one unpaired electron.
Second-period center
C, N, O and F cannot exceed an octet in ordinary Lewis structures.
B or Be center
Electron deficiency may be the correct answer rather than a forced multiple bond.
Period 3 or heavier
Hypervalent Lewis counts are possible, but interpret them with modern delocalized bonding models.
Lewis structures before the octet test
The octet rule is most useful inside a disciplined Lewis-structure workflow. First count all valence electrons, choose a plausible skeleton, connect atoms with single bonds, complete terminal-atom shells, then place remaining electrons on the central atom. Only after that should multiple bonds or formal-charge alternatives be considered.
For a polyatomic ion, add one electron for each negative charge and subtract one for each positive charge. Brackets and the overall charge belong to the complete Lewis structure.
| Species | Total valence electrons | Key feature |
|---|---|---|
| CO₂ | 16 | Two C=O double bonds give octets with zero formal charge. |
| BF₃ | 24 | B remains electron-deficient with three B–F single bonds. |
| NO | 11 | Odd total requires an unpaired electron. |
| SO₄²⁻ | 32 | Several resonance/formal-charge drawings are possible. |
Hypervalence without the old d-orbital shortcut
For species such as XeF₂, a useful modern picture is a three-center four-electron interaction. Orbitals on F–Xe–F combine into bonding, nonbonding and antibonding molecular orbitals spread over all three atoms. Four electrons occupy the bonding and nonbonding combinations.
This delocalized picture explains how a Lewis structure can place more than eight electrons around the central atom without requiring a set of strongly localized “expanded-octet” d hybrids.
The exact bonding differs across hypervalent molecules, but the general lesson is stable: Lewis electron counts are bookkeeping; the microscopic wavefunction is delocalized.
Transition metals follow different electron-counting ideas
The octet rule was built for main-group valence shells. Transition-metal chemistry instead involves ns and (n−1)d electrons, ligand fields and metal–ligand covalency. The 18-electron rule is sometimes a useful analogue because one s, three p and five d orbitals can hold 18 electrons, but it also has many exceptions.
Applying the octet rule mechanically to Fe, Co, Ni or their complexes therefore produces misleading Lewis pictures. The model should match the electronic structure being studied.
Exercises
BF₃
How many electrons surround B in the simplest Lewis structure?
Solution
Three B–F bonds place 6 electrons around boron.
NO
Why can NO not give both atoms conventional paired-electron octets?
Solution
It has an odd total of 11 valence electrons, so one electron remains unpaired.
SF₆
Does a Lewis structure with 12 electrons around S prove sulfur uses six localized sp³d² hybrids?
Solution
No. The Lewis count is useful bookkeeping; modern bonding is better described with delocalized molecular orbitals and substantial ionic character.