How does #SO_4# have a charge of 2-?

Question

Naomi Aronson · Accepted Answer

We are agreed that sulfate anion derives from #H_2SO_4#, sulfuric acid, or from sodium sulfate, #Na_2SO_4#, which are both manifestly neutral entities.

A typical (if outdated) Lewis structure of sulfuric acid is: #(HO-)_2S(=O)_2# This Lewis structure is equivalent to: #(HO-)_2S^(2+)(-O^-)_2# For a neutral chalcogen atom (#"chalcogen = S or O"#), there must be 6 valence electrons. In the representation #(HO-)_2S(=O)_2# there are certainly 6 electrons associated with each sulfur or oxygen. Lone pairs are owned by the atom, and thus on neutral oxygen there are 2 electrons from the double bond, and 4 electrons in the lone pairs). Now of course both #H_2SO_4# and #HSO_4^-# are strong acids, and undergoes almost complete ionization in water: #H_2SO_4(aq) +2H_2O(l)rarr SO_4^(2-) + 2H_3O^+# Conservation of charge demands that the sulfate ion has 2 formal negative charges. Nitric acid has an even more problematic representation: #(O=)N^(+)(-O^(-))(-OH)#, where there is formal charge separation in even the neutral acid (6 electrons around nitrogen rather than 7; 9 electrons around oxygen rather than 8 ). See here for another example that assigns formal charge.

Kai Bowman · Answer

undefined The sulfate ion (SO₄²⁻) has a charge of 2- because it consists of one sulfur atom bonded to four oxygen atoms. Each oxygen atom contributes 2 electrons to the bond, and sulfur atom contributes 6 electrons. The sulfur atom also carries a formal charge of +2. Overall, the sulfate ion has 32 valence electrons and a total charge of -2.