In chemistry, we aim for a structure where formal charges are as close to zero as possible. Because sulfur is in the third period of the periodic table, it can have an (more than 8 electrons). To minimize formal charges:

Valence = 6. Non-bonding = 4 (two lone pairs). Bonding = 4 (the double bond). (FC = 6 - 4 - ½(4) = 6 - 4 - 2 = 0)

The Sulfur becomes 0, the double-bonded Oxygens become 0, and the two single-bonded Oxygens remain -1negative 1 each. This adds up to the overall -2negative 2 charge of the ion. Geometry and Hybridization

Connect each oxygen to the sulfur with a single bond (a line representing 2 electrons). This uses up (4 \text bonds \times 2 \text electrons = 8) electrons.

Valence = 6. Non-bonding = 0. Bonding = 8 (four bonds). (FC = 6 - 0 - ½(8) = 6 - 0 - 4 = +2)

Sum = (-2 + 0 + 0 + 0 + 0 = -2). Formal charges: S (-2), all O (0). While all oxygens are satisfied, sulfur now bears a -2 formal charge. Since oxygen is more electronegative than sulfur, it is more stable for the negative charges to reside on oxygen, not sulfur. Structure D is less favorable than Structure B.

Valence = 6. Non-bonding = 6 (three lone pairs). Bonding = 2 (one single bond). (FC = 6 - 6 - ½(2) = 6 - 6 - 1 = -1)

The Lewis structure for SO4^2- is:

Move a lone pair from two different oxygen atoms to form with sulfur.

Our goal is to distribute these 32 electrons as bonding pairs (lines) and lone pairs (dots) to satisfy the octet rule for as many atoms as possible.