Lewis Structures & Octet Violations

Always begin by counting the total number of valence electrons present.
- Sum the valence electrons for each atom using periodic-table group numbers.
- Add electrons for negative charges; subtract for positive charges.
Assemble a skeletal structure:
- Place the least electronegative / most electro-positive atom in the center (exceptions: H and most halogens rarely go in the center).
Form single bonds between the central atom and each terminal atom (each bond = $2\,e^-$ ).
Complete the octets on terminal atoms first; place remaining electrons on the central atom.
When octets are not satisfied, employ multiple bonds or invoke an octet exception (electron deficiency, expanded octet, or odd-electron species).

Electron-deficient central atoms (e.g.
$\text{B},\,\text{Be}$ ) – stable with fewer than 8 electrons.
Expanded octets (period $n \ge 3$ elements that can access empty $d$ orbitals, e.g. S, P, Xe, I).
Odd-electron (radical) molecules – total valence electron count is odd; one atom must carry an unpaired “singlet” electron.

Valence-electron count
- Boron: group 13 $\Rightarrow 3\,e^-$ (instructor verbally said “13,” but he uses 3 valence electrons).
- Chlorine: $3\times7\,e^-$ .
- Total $=3+21 = 24\,e^-$ .
Skeleton: Boron central, three B–Cl single bonds (uses $6\,e^-$ → $18\,e^-$ left).
Octet completion
- Each Cl needs 6 more electrons (lone pairs): $3\times6 = 18\,e^-$ .
- No electrons remain for boron – it sits with only 6 electrons.
Key takeaway: Boron is satisfied at 6 electrons – textbook case of an electron-deficient octet violator.

Chemical context: Component of the classic starch indicator – turns deep purple in redox titrations.
Valence-electron count
- $3\times7\,e^- = 21\,e^-$ (iodine).
- Extra electron for the $-1$ charge → $22\,e^-$ total.
Skeleton: Must connect all three I atoms linearly (–I–I–I–).
- Two single bonds = $4\,e^-$ spent; $18\,e^-$ remain.
Octet completion on terminals
- Each terminal I requires 3 lone pairs → $16\,e^-$ .
- $2\,e^-$ left over.
Expanded octet step
- Central I (a 5th-period element) accepts the extra pair, expanding to 10 electrons (5 lone-pair equivalents when counting bonds).
Final drawing conventions
- Show the molecule in brackets with the overall $-1$ charge outside.
- Spread lone pairs symmetrically around the central I.
Why allowed? Iodine’s empty 5d orbitals provide spatial “storage,” legitimizing the expanded octet.

Environmental note: Brownish gas from automobile exhaust; monomer ( $\text{NO}<em>2$ ) eventually dimerizes into $\text{N}</em>2\text{O}_4$ .
Valence-electron count
- Nitrogen: $5\,e^-$ .
- Oxygen: $2\times6\,e^- = 12\,e^-$ .
- Total $=17\,e^-$ → odd → one unpaired “singlet.”
Choosing the central atom
- Rule: If not specified, place the more electropositive element in the center → nitrogen (appears left in the formula).
Skeleton: N with two single bonds to O (spends $4\,e^-$ → $13\,e^-$ remain).
Octet shortage analysis
- To satisfy both O atoms and N, $\approx16\,e^-$ would be needed → deficit.
Introduce a double bond (either O works – resonance later)
- Double bond consumes 2 extra bonding electrons (total $6\,e^-$ in bonds now).
Electron placement
- Double-bonded O gets 2 lone pairs (4 electrons) to complete 8.
- Single-bonded O gets 3 lone pairs (6 electrons).
- Only one electron remains for nitrogen → nitrogen holds the radical (7 electrons total).
Rationale for radical location
- Nitrogen is less electronegative → better site for electron deficiency.
- Analogy: “Bald-spot hair sharing” (hypochondriacion) – more electronegative atoms help spread charge; deficiency kept on the electropositive center.

Halogens almost never act as central atoms; exceptions arise with large halogens (Cl, Br, I) when forced by molecular composition.
The larger the halogen, the better it can accommodate charge & expanded octets.
Electropositive vs. electronegative placement
- Electropositive atoms in the center allow electronegative terminals to “spread out” electron density.
Deficiency (positive‐like) vs. surplus (negative‐like) mindset
- Instructor recommends thinking of a missing electron (singlet) as a positive defect; extra electrons behave as a negative surplus.

$\text{I}_3^-$ in starch solution gives the classic purple endpoint in redox titrations.
$\text{NO}_2$ is a pollutant; its behavior as a radical explains ease of dimerization and reactivity in the atmosphere.
Understanding octet violations is critical for:
- Predicting molecular geometry (VSEPR adjustments for expanded octets).
- Rationalizing stability of reactive intermediates in organic & inorganic reactions.

Octet violator – atom/molecule not adhering to the 8-electron rule.
Electron-deficient species – stable with fewer than 8 electrons (e.g.
$\text{BCl}_3$ ).
Expanded octet – central atom exceeds 8 electrons via access to $d$ orbitals (e.g.
$\text{I}_3^-$ ).
Odd-electron (radical) molecule – possesses one unpaired electron (e.g.
$\text{NO}_2$ ).
Singlet – the lone unpaired electron creating deficiency.
Hypochondriacion (classroom metaphor) – distributing electron density like combing a few hairs over a bald spot to hide a deficiency.

The lecture completes the catalogue of Lewis-structure types students need: normal, electron-deficient, expanded-octet, and odd-electron structures.
Next session promises to link these drawings to practical uses and further chemical insight.