Octet rule — Lewis structures, exceptions (incomplete octet, expanded octet, odd-electron species): Question bank

Intuition Picture the four exception "characters"

Before the traps, meet the shapes you'll be reasoning about. Each panel below is labelled with its own name, its electron count, and an arrow pointing at the feature that makes it special — so you read the picture directly, without mapping text to regions in your head.

Panel 1 — $C O_{2}$ (green): everyone reaches exactly 8, the happy default. No arrow, nothing broken.
Panel 2 — $B F_{3}$ (red): the red "gap!" arrow points to boron's missing pair — it sits at only 6 (incomplete octet, hungry Lewis acid).
Panel 3 — $S F_{6}$ (blue): the blue "bulge >8" arrow points to sulfur's overfull crowd of 12 (expanded octet, period-3 privilege).
Panel 4 — $N O$ (white atoms): the single red dot with the "no partner" arrow is the lonely unpaired electron of an odd-electron radical.

True or false — justify

The octet rule is a physical law that atoms must always obey.

False. It is a tendency toward noble-gas stability, not a law; boron settles for 6, sulfur holds 12, and radicals cannot pair at all.

Hydrogen tries to reach 8 electrons like every other atom.

False. H (and He) aim for a duet of 2, matching helium's full

1 s^{2}

shell — its valence shell only has one

s

orbital.

In Lewis counting, each atom of a double bond counts all 4 of the double-bond electrons toward its own octet.

True. A double bond is 4 shared electrons; both atoms count the same 4 electrons toward their own octet — that shared counting is exactly why bonding fills octets so efficiently.

Every atom in

C O_{2}

has exactly 8 electrons around it.

True. Carbon has two double bonds (

4 + 4 = 8

) and each oxygen has one double bond plus two lone pairs (

4 + 4 = 8

B F_{3}

the boron atom satisfies the octet by forming a B=F double bond.

False. A B=F double bond would put a

+ 1

formal charge on fluorine — the most electronegative atom — which is energetically absurd; boron stays with only 6 electrons.

Nitrogen can expand its octet to hold 10 electrons in molecules like

H N O_{3}

False. Period-2 elements have no accessible low-energy d-orbitals and are too small (see the energy-ladder figure above); N is capped at 4 bonds (8 electrons). Keep the formal charges instead.

S F_{6}

violates the octet rule because sulfur is unstable.

False. The opposite — sulfur (period 3) stably expands to 12 electrons; the molecule is very stable. "Exception" means "differs from 8," not "unstable."

Odd-electron species like

N O

can still give every atom a complete octet.

False. With 11 total valence electrons the number is odd, so at least one electron stays unpaired and one atom is left short of 8.

Adding a negative charge to an ion means subtracting electrons in Step A.

False. A negative charge means extra electrons, so you add them (see the Step A definition above); you subtract only for positive (cation) charges.

The best Lewis structure is always the one where every atom has zero formal charge (FC).

Partly false. The best structure minimizes the magnitude of formal charges — some charges may be unavoidable (as in

C O_{3}^{2 -}

, shown in the FC figure above); when charges exist, negative FC belongs on the most electronegative atom.

Spot the error

"I drew

B e C l_{2}

with Be having 8 electrons by making two Be=Cl double bonds — so it obeys the octet."

Error: double bonds put positive formal charge (FC) on chlorine and are not favored; Be is genuinely electron-deficient with only 4 electrons (incomplete octet), acting as a Lewis acid.

S O_{4}^{2 -}

has 30 valence electrons: S(6) + 4×O(6) = 30."

Error: the

2 -

charge adds 2 electrons in Step A, so the total is

6 + 24 + 2 = 32

, not 30. Forgetting the charge wrecks every later bond and formal-charge count.

"Oxygen is in period 2, so like sulfur it can form an expanded octet in

O_{3}

Error: period-2 elements never expand — no low-energy d-orbitals and too small (see the energy-ladder figure). Ozone is handled by resonance and formal charges, keeping every O at 8.

"NO has 12 valence electrons because N and O are both roughly 6."

Error: N is group 15 → 5 valence electrons, O is group 16 → 6, giving

5 + 6 = 11

— an odd number, which is exactly why NO is a radical.

"For

C O_{3}^{2 -}

I put a

- 1

formal charge (FC) on the carbon since the ion is negative."

Error: carbon has

F C = 4 - 0 - 4 = 0

(as annotated in the FC figure above); the two

- 1

charges sit on the single-bonded oxygens, the more electronegative atoms, summing to the ion's

- 2

P C l_{5}

can't exist because phosphorus would break the octet."

Error: P is period 3 and expands to 10 electrons (5 bonds);

P C l_{5}

is a real, stable molecule — see Hypervalency & d-orbital participation.

Why questions

Why does the octet rule target the number 8 specifically?

The valence shell has one

s

orbital (2 e⁻) and three

p

orbitals (6 e⁻); filling

s^{2} p^{6} = 8

electrons reproduces a noble-gas configuration, the lowest-energy arrangement.

Why do we place the least electronegative atom at the center of a Lewis structure?

The central atom shares its electrons with the most neighbors; a less electronegative atom holds its electrons loosely and is happy to share them out to several bonds.

Why does nature prefer Lewis structures with the smallest formal charges?

Large separated charges cost energy; a structure closer to overall neutrality (charge spread thinly, near the atoms that hold it best) sits lower in energy.

Why can sulfur expand its octet but nitrogen cannot, even though they are in the same column region?

They differ by period: S is in period 3 with accessible vacant d-orbitals just above its valence (small green gap in the energy-ladder figure) and a larger size; N is in period 2 with its

3 d

far above and out of reach — see Periodic trends.

Why are odd-electron species like

N O

and

C l O_{2}

so reactive?

The lone unpaired electron (the red "no partner" dot in the first figure) has no partner, so the molecule readily reacts to pair it up; such radicals are also paramagnetic (attracted to magnetic fields). See Free radicals.

Why does

B F_{3}

behave as a Lewis acid?

With only 6 electrons boron has an empty orbital, so it eagerly accepts an electron pair from a donor to complete its octet — the definition of a Lewis acid.

Why does drawing resonance structures for

C O_{3}^{2 -}

give a bond order of

4/3

The one double bond can sit on any of the three equivalent C–O positions; the real molecule is the average, spreading 4 bonds over 3 links →

4/3

each — this is exactly the resonance whose FC bookkeeping the figure above walks through.

Edge cases

Does

B e C l_{2}

obey the octet rule?

No — beryllium keeps only 4 electrons (two single bonds), an incomplete octet; it is the smallest, most extreme electron-deficient case.

What happens to the octet rule for a lone hydrogen atom aiming to bond?

It follows the duet rule instead, seeking just 2 electrons; treat H specially — it never carries lone pairs or reaches 8.

Can an atom have fewer than the electrons the shared-bond arithmetic suggests?

Yes — for electron-deficient centers (B, Be) the formula may propose extra bonds that would create absurd formal charges, so the real structure leaves the central atom short of 8.

Is a molecule with a formal charge automatically wrong or unstable?

No — formal charges are a bookkeeping tool, not real charge; ions and many neutral molecules legitimately carry them, and they can be the physically correct choice over an octet-forcing alternative.

What is the maximum number of bonds nitrogen can form, and why is this an edge case people miss?

Four bonds maximum (as in

N H_{4}^{+}

), giving 8 electrons; students wrongly add a fifth bond to cancel formal charges, but N cannot expand its octet.

If a species has an even total of valence electrons, can it still be a radical?

Rarely — an even count usually allows full pairing, but some biradicals (e.g.

O_{2}

has two unpaired electrons) still carry unpaired spins; the odd-electron rule guarantees a radical, but an even count does not guarantee its absence.

Recall One-line self-test

Cover every answer above and re-run the page. If you can justify each verdict (not just state it), you own the octet rule and its three exceptions — "I Expand Oddly."

Connections

Parent topic
Formal charge & resonance — the tool behind most "spot the error" items.
Electronegativity — decides where negative formal charge belongs.
Lewis acids and bases — incomplete-octet traps.
Hypervalency & d-orbital participation — expanded-octet edge cases.
Free radicals — odd-electron reactivity.
Periodic trends — the "why S not N" reasoning.
VSEPR theory — where correct Lewis structures lead next.