So4 Lewis Structure !!top!! »
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 initial structure (Structure A) looks like this: so4 lewis structure
Sulfur is less electronegative than oxygen. Therefore, sulfur is the central atom. The four oxygen atoms surround it in a tetrahedral arrangement (though we draw it in 2D with S in the middle and O’s at the four cardinal points). We started with 32 electrons
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. To satisfy the octet rule, each oxygen needs
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).
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.
Formal Charge = (Valence electrons) - (Non-bonding electrons) - ½(Bonding electrons)
