The SO2 Lewis structure places sulfur in the center, connects two oxygens, and distributes 18 valence electrons so that formal charges are minimized and both S–O bonds are equivalent through resonance.
Valence-electron accounting
- Total valence electrons: sulfur (6) + 2 × oxygen (6) = \(18\).
- Skeletal connectivity: O–S–O (oxygen atoms as terminal atoms; sulfur as the central atom).
- Single-bond framework: two S–O single bonds use \(2 \times 2 = 4\) electrons, leaving \(18 - 4 = 14\) electrons.
- Terminal oxygen octets: three lone pairs on each oxygen use \(2 \times 6 = 12\) electrons, leaving \(14 - 12 = 2\) electrons for sulfur.
- Central sulfur lone pair: the remaining 2 electrons become one lone pair on sulfur.
The single-bond skeleton with complete oxygen octets produces a valid electron count, but the resulting formal charges are not minimal; multiple bonding (and resonance) accounts for the observed equivalence and shorter S–O bond lengths.
Formal charges and preferred Lewis descriptions
Formal charge evaluates how well a Lewis structure assigns electrons: \[ FC = V - \left(N + \frac{B}{2}\right), \] where \(V\) is valence electrons for the free atom, \(N\) is nonbonding electrons on the atom, and \(B\) is bonding electrons in bonds to that atom.
Key Lewis candidates for SO2
| Representation | S–O bonding | Formal charges | Notes |
|---|---|---|---|
| All single bonds | O–S–O | \(FC(\text{S}) = +2\), \(FC(\text{each O}) = -1\) | Charge separation is large; not the best description. |
| Two resonance forms | One S=O and one S–O in each form | \(FC(\text{S}) = +1\), \(FC(\text{single-bond O}) = -1\), \(FC(\text{double-bond O}) = 0\) | Two equivalent resonance contributors; both S–O bonds become equivalent in the hybrid. |
| Expanded-octet depiction | O=S=O with one lone pair on S | \(FC(\text{S}) = 0\), \(FC(\text{each O}) = 0\) | Often drawn to show minimal formal charges; sulfur is a third-period element and can exceed an octet in Lewis bookkeeping. |
General-chemistry conventions commonly emphasize the resonance description because it explicitly shows electron delocalization and explains why the two S–O bonds are equivalent. The expanded-octet depiction is also widely accepted in introductory Lewis-structure practice and is consistent with minimized formal charge.
Resonance and bond order
The two resonance contributors are equivalent by symmetry: one places the negative charge on the left oxygen, the other on the right oxygen. The resonance hybrid has two equivalent S–O bonds with bond character between a single and a double bond, commonly summarized as an average bond order near \[ \text{bond order} \approx 1.5 \quad (\text{resonance model}). \]
Molecular geometry (VSEPR)
Around sulfur there are three electron domains: two S–O bonding regions and one lone pair. VSEPR predicts a trigonal-planar electron-domain arrangement and a bent molecular geometry for SO2. The O–S–O bond angle is close to \(120^\circ\) (often reported near \(119^\circ\) in typical conditions), and the molecule is polar because the bent shape prevents dipole cancellation.
Visualization: resonance contributors and bent geometry
Common pitfalls
- Electron count mismatch: SO2 has 18 valence electrons for the neutral molecule; charged species (SO2−, SO22−) require different totals.
- Octet overcorrection: oxygen remains an octet atom in standard Lewis structures; sulfur (period 3) can exceed an octet in Lewis bookkeeping when minimizing formal charge.
- Unequal-bond interpretation: resonance contributors are not separate molecules; experimentally the two S–O bonds are equivalent.
- Geometry confusion: two bonds plus one lone pair on sulfur implies a bent molecular shape, not linear.