Expanded Octets and Formal Charge Trade-offs

With four S-O single bonds, sulfur has an octet but carries a formal charge of +2 while each oxygen carries -1. Converting two of those bonds to double bonds gives sulfur twelve electrons and drops its formal charge to zero, leaving only two oxygens at -1.

Sulfate has 6 + 4(6) + 2 = 32 valence electrons. The all-single-bond drawing uses eight in bonds and twenty-four in oxygen lone pairs. In the two-double-bond drawing, twelve are in bonds and twenty in oxygen lone pairs: the budget is still thirty-two. Each new S=O component comes from an oxygen lone pair, not from additional electrons supplied to sulfur. The sulfur formal charge changes from +2 to 0 because its bond-order sum rises from four to six.

An expanded Lewis drawing of sulfate can reduce formal-charge magnitudes; an all-single-bond drawing keeps an octet at sulfur. Neither drawing alone is the physical electron density. Modern bonding descriptions do not require literal d-orbital hybridization, and this does not uniquely certify one Lewis drawing or assign numerical resonance weights. Compare the bookkeeping conventions explicitly and follow the representation requested in the question.

Neither single drawing implies two permanent short S–O bonds and two permanent long ones in an isolated symmetric sulfate ion. A family of equivalent expanded-valence contributors distributes the double-bond positions. Those drawings remain a model; they do not by themselves determine a literal integer bond order or prove that sulfur uses d hybrids.

SO₂ with one single and one double bond gives sulfur +1 and the single bonded oxygen -1. Drawing both bonds double gives every atom a formal charge of zero at the cost of ten electrons on sulfur. The same argument applies, and the same two answers appear in different sources.

SO₂ has eighteen valence electrons. In the zero-formal-charge double/double drawing, the two S=O bonds use eight electrons, the oxygens' four lone pairs use eight, and one sulfur lone pair uses two. Sulfur therefore has ten surrounding electrons: four bonding pairs plus one lone pair. Omitting that sulfur lone pair would leave only sixteen electrons and give the wrong formal charge. In an octet-only drawing, one S=O and one S–O plus the sulfur lone pair gives +1 on S and −1 on the single-bonded O; there are two equivalent placements.

Read the requested convention before drawing. If asked to minimize formal charges while allowing expanded octets, use that criterion; if an octet-only representation is requested, obey that constraint. Show total electrons, lone pairs and formal charges, with the sum equal to the ionic charge. Naming a criterion does not guarantee credit if the prompt requires a different convention.

Before answering, state whether the task requires an octet-only structure or allows expanded-valence drawings chosen to reduce formal-charge magnitudes. Then show lone pairs, local formal charges, bracket charge for ions, and the total electron count. Naming an alternative convention is useful explanation, but it does not override an explicit representation requested by a question.

P, S, Cl and heavier central atoms can appear with more than eight surrounding electrons in useful Lewis drawings. Ordinary C, N, O and F centers cannot use that bookkeeping shortcut. Period number therefore constrains which drawings are available; it does not require all compounds of period-3 elements to expand or justify adding extra bonds without checking the budget.

Modern descriptions of hypervalent bonding use delocalization and other orbital models rather than requiring literal occupation of d hybrids. The Lewis drawing remains useful for counting electrons and comparing representations. Keep that purpose separate from a claim about the exact physical distribution of bonding electrons.