So4 Lewis Structure 〈ESSENTIAL〉

We represent this by drawing all significant resonance structures connected by double-headed arrows, or more commonly, by drawing a single structure with dashed lines or a circle to indicate delocalized bonding, though this is less precise. The above resonance model (using two double bonds) is excellent for explaining formal charge and bond equivalence. However, it violates a subtle but important rule: in the two-double-bond structure, sulfur has 10 electrons around it (four from each of two double bonds and two from each of two single bonds = 4+4+2+2 = 12? Wait, recalc carefully).

Formal Charge = (Valence electrons) - (Non-bonding electrons) - ½(Bonding electrons) so4 lewis structure

The initial structure (Structure A) looks like this: We represent this by drawing all significant resonance

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. Wait, recalc carefully)

We started with 32 electrons. After using 8 for bonds, we have (32 - 8 = 24) electrons left (or 12 lone pairs). Oxygen atoms are greedy for electrons. To satisfy the octet rule, each oxygen needs 6 more electrons (3 lone pairs) around it. (4 \text oxygens \times 6 \text electrons = 24) electrons. Perfect.

The actual sulfate ion is a resonance hybrid of multiple equivalent structures. In one resonance form, the double bonds are on the top and left oxygens. In another, they are on the top and right. In a third, on the bottom and left, and so on. The true ion is the average of all these forms, where each S–O bond has a bond order of 1.5 (halfway between single and double) and each oxygen carries a formal charge of -0.5.

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.