Whenever you see an image of phosphate, it's always shown having 5 bonds. 3 single bonds to anionic oxygen atoms and a double bond to an extra neutral oxygen.

Not only that, but they're shown being arranged in a tetrahedral manner. All of this additionally applies to sulfate, only that it has 2 double bonds instead of just one.

Now, from what I've been able to gather from the net, including this website is that this can be explained by invoking sp3 hybridization and a d orbital. The 3 anionic oxygens who already have a charge of -1 pair up with the 3 partially filled hybrid orbitals in phosphorous, while the remaining filled sp3 orbital on phosphorous donates it's electron pair to the empty orbital on an extra oxygen.

Once again, this is exactly the same for sulfate, except that it can only bond to 2 anionic oxygens while also being able to donate 2 electron pairs to other oxygen atoms.

This leaves us with 4 single bonds to oxygen atoms. This doesn't seem unusual at all. Nitrogen can easily take on 4 bonds just like phosphorous, and seeing how row 3 elements have a looser hold on their outer electrons, it makes sense that they can donate their complete electron pairs much easier than row two elements, seeing how their electron pairs would "stick out" more, making them much more available for bonding.

However, in Phosphate when you get to the 5th bond you have to explain it by being a little more creative. You could say that a p orbital on one of the oxygen atoms donates it's electron pair to an empty d orbital on phosphorous, creating and extra bond between phosphorous and one of the oxygens. The same thing happens in sulfate only that for some reason it can form two double bonds. (which I really wonder if that would be pushing the limits of stability at all. I mean wouldn't the two d orbitals kind of overlap at an awkward angle causing some kind of repulsion?)

Now, apparently from what I've heard, experimental evidence has shown that there is very little d orbital involvement in molecules like phosphate or sulfate, and they're practically all single bonds. So my question is, is this true? And if so, then why is the double bond picture still in practically every textbook, diagram, and piece of scientific literature that you come across? How did this picture of these ions emerge? I know it had something to do with linus pauling, but was this just accepted based on theory alone and no experimental evidence at all? Was there some experimental evidence for it in the beginning?

Sorry if it sounds like I'm asking too much, but I just find this whole situation so confusing. I'd appreciate your answers.

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    $\begingroup$ You can always go back to your question and edit it. Add more context, like linking the question or tags. Keep in mind that once you have obtained an answer it is unwise (and not very welcome here) to go back and change the question. But until then you can try to fine tune your question. $\endgroup$ Commented Oct 21, 2015 at 5:23

1 Answer 1


Recent experimental evidence shows that d orbitals are involved in the extra double bond of sulfate or phosphate ions. Therefore a stable sulfate or phosphate ion has 2 double bonds instead of only one to make the formal charge on sulfur or phosphorus zero. Any Lewis structure with expanded octet (periods 3 and higher) should have zero formal charge on the central atom.

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    $\begingroup$ It will be very helpful if you could quote those "recent experimental evidences". $\endgroup$ Commented Apr 20, 2019 at 17:24
  • $\begingroup$ Ok, I'll bite. How do you experimentally identify a $d$ orbital electron versus an $s$ or $p$ one? Oh boy this is going to be good! $\endgroup$ Commented May 20, 2019 at 21:02
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    $\begingroup$ This is completely untrue! There should be half a dozen downvotes here already. $\endgroup$
    – Mithoron
    Commented Jun 4, 2019 at 0:29

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