The octet rule is a bonding guideline for main-group elements: stable Lewis structures often place eight valence electrons around nonmetal atoms (a noble-gas configuration), counted as shared bonding pairs plus lone pairs.
Valence-shell stability and the octet rule
The octet rule describes a common electron-distribution pattern in covalent and ionic bonding. Atoms in the second period (C, N, O, F) are especially well described by the octet rule because their valence shell consists of the \(2s\) and \(2p\) orbitals and cannot exceed eight electrons. Many Lewis structures can be assessed by checking whether typical atoms reach an octet (or a duet for hydrogen) while maintaining a reasonable total electron count.
Electron count around an atom in a Lewis structure includes:
- Lone-pair electrons shown as dots on that atom.
- Bonding electrons counted as \(2\) for each bond to that atom (each bond represents a shared pair).
An octet corresponds to \(8\) electrons around the atom by this counting convention.
Lewis structures and electron bookkeeping
A Lewis structure begins with the total valence electron count from the periodic table. For a neutral molecule, this is the sum of group-valence electrons for all atoms. For ions, electrons are added for negative charge and removed for positive charge.
\[ V_{\text{total}}=\sum V_i \;\; \pm \;\; (\text{charge}) \]
Formal charge is a compact way to evaluate whether a proposed Lewis structure distributes electrons plausibly. It favors structures with small magnitudes of charge and, when charges are unavoidable, negative formal charge on the more electronegative atoms.
\[ \text{FC} = V - \left(N + \frac{B}{2}\right) \]
Here \(V\) is the number of valence electrons for the free atom, \(N\) is the number of nonbonding (lone-pair) electrons assigned to that atom, and \(B\) is the total number of bonding electrons in bonds to that atom.
Common exceptions and boundaries
The octet rule is reliable for many stable, main-group compounds, yet several well-known categories sit outside it. These are not “violations” in the sense of incorrect chemistry; they reflect orbital capacity, electron deficiency, or unpaired electrons.
| Category | Valence-shell outcome | Typical examples | Chemical reason (general chemistry level) |
|---|---|---|---|
| Duet rule (hydrogen) | \(2\) electrons around H | H2, HCl, CH4 | Hydrogen has only the \(1s\) valence orbital, which holds at most \(2\) electrons. |
| Incomplete octet | Fewer than \(8\) electrons on a central atom | BF3, BeCl2, AlCl3 | Electron-deficient centers form stable compounds because full octets are not always required for bonding stability in these cases. |
| Odd-electron species | An atom carries \(7\) (or another odd count) due to an unpaired electron | NO, NO2, ClO2 | Total valence electrons are odd, so at least one unpaired electron remains; strict octets for every atom are impossible. |
| Expanded octet (period \(\ge 3\)) | More than \(8\) electrons around a central atom | PCl5, SF6, XeF4 | Heavier main-group atoms can accommodate more electron density in bonding arrangements; second-period atoms cannot exceed an octet. |
| Resonance | Same connectivity, multiple valid electron placements | O3, NO3−, CO32− | Delocalized \(\pi\) bonding yields multiple Lewis structures; the actual electron distribution is a resonance hybrid. |
Illustrative electron counts
CO2 as a standard octet-rule structure
Carbon dioxide has \(4\) valence electrons from C and \(6\) from each O, giving a total of \(16\) valence electrons:
\[ V_{\text{total}} = 4 + 2(6) = 16 \]
The Lewis structure \( \mathrm{O{=}C{=}O} \) places an octet on carbon (two double bonds count as \(8\) electrons around C) and an octet on each oxygen (two lone pairs plus the double bond).
BF3 as an incomplete-octet example
Boron trifluoride is electron-deficient at boron in its simplest Lewis structure. Boron forms three single bonds, giving boron \(6\) electrons around it by bonding-pair counting. This incomplete octet is a standard, stable pattern for boron compounds and underpins Lewis-acid behavior.
Visualization of octet-rule patterns
Common pitfalls
- Electron totals that do not match the computed valence electron count for the entire species.
- Second-period limits that exceed eight electrons on C, N, O, or F; expanded octets belong to period \(3\) and beyond.
- Formal-charge placement that assigns large charges when a lower-charge resonance form exists.
- Hydrogen counts that treat H as an octet element; hydrogen follows the duet rule.
Summary statement
The octet rule is a central organizing idea for Lewis structures in general chemistry: electron bookkeeping aims for octets on many main-group atoms, with well-defined exceptions that include the duet rule for hydrogen, incomplete octets for electron-deficient centers, odd-electron species, resonance, and expanded octets for heavier central atoms.