# Why does sulfate have this structure?

Look at the structure for sulfate:

Why does sulfur form so many covalent bonds. Its valency is only $-2$, it only needs two electrons, yet here it's getting $6$.

The same thing happens with phosphate.

Phosphorus has a valency of $-3$, yet, it forms $5$ covalent bonds.

• I think the key thing is that Sulphur can have a range of valencies. It just usually has an oxidation number of -2. In this case each Oxygen is -2 and the Sulphur is +6. – Nick Aug 22 '13 at 15:13
• Why does sulphur have multiple valencies? – Gerard Aug 22 '13 at 15:14
• Although it seems that current research favours a 3c,4e bond model with the D orbital contribution being small. I can't read the papers though as they are paywalled. – Nick Aug 22 '13 at 15:25

I think an important point to mention here is that Lewis dot structures, and the octet rule, are simply models that describe experimental observables. For example, the concept of valency works well with organic molecules but was challenged by Alfred Werner in the development of bonding models that adequately described coordination compounds.

The electron-dot model of bonding is so meaningful because it can be adapted to explain various experimental observables. You can draw two perfectly reasonable electron dot structures for $\ce{SO4^{2-}}$: one with all $\ce{S-O}$ single bonds and one with two $\ce{S=O}$ double bonds (and the additional resonance structures). The question is, which electron-dot structure best represents the real structure? We can use the concept of formal charges to predict that the structure with double bonds is most likely closer to reality.

We then look at the experimental data (which is reported on the sulfate wikipedia page) for S-O single and double bond lengths. The S-O bond length in $\ce{SO4^{2-}}$ is 149 pm, which is shorter than that observed in sulfuric acid (157 pm) and very close to the gas-phase bond length of sulfur monoxide which is 148 pm.

So, at the end of the day, an electron dot structure is a type of model that describes bonding. It has some additional features such as expanded octets and formal charge that help broaden its applicability to more situations; however, models such as this one don't tell us what an atom "needs"; rather, it provides an explanation for how we observe atoms "behaving".

Ultimately those structures are wrong. Most of the community of chemistry educators teaches that formal charge reduction creates hypervalency. All modern research and models shows this not to be the case. You can't hybridize a d orbital and even get the proper tetrahedral geometry here. Reducing formal charges to 0 makes no real world sense, when sulfur has a measured empirical charge of around 1.7 in sulfate.

Sulfate is 4 single bonds. The bond length difference is due to the fact that the sulfur carries extra charge, as does the oxygens. This ionic like bond is of course going to be shorter than more neutral forms of S-O. Further, why would two double and two single hybrid bonds average to be the same length as a double bond?

Essentially any general chemistry solution that invokes d-orbitals as it's solution, outside of transition metal chemistry, is wrong. Sulfur hexafluoride is not in fact 6 single bonds.

Unfortunately, biochemists are not going to update their picture of phosphate anytime in the next fifty years because no one cares.

The fact that basically every science major in the world learns this wrong is abhorrent.

• Why blame "the biochemists"? In biochemistry textbooks, the formula for phosphate is often $\ce{P_i}$, sometimes even with a nice circle around it. – Karsten Theis May 27 at 21:48

It's a highly debated topic and seems that none has won. I thought it to be fully explainable through MOT. However, while searching I found this link. It cleared few of my misunderstanding.