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Why do some polyatomic ions violate the octet rule?

  1. Nov 7, 2011 #1
    Why do some polyatomic ions violate the octet rule? I'm trying to figure out the lewis diagrams of some ions by myself, and sometimes I see things like these:

    [PLAIN]http://www.chemprofessor.com/bonding_files/image035.jpg [Broken]

    [PLAIN]https://www2.bc.edu/~teeter/Courses/Phosphate.GIF [Broken]

    These things have 10 electrons for some of the atoms instead of 8. I did some searching on Google, and came upon reasons like formal charges, odd number of electrons...I don't understand!
    Last edited by a moderator: May 5, 2017
  2. jcsd
  3. Nov 7, 2011 #2
    This is a good question, I can't answer it fully, but always look at it this way, electrons are always "searching" for a lower state of potential energy given by the field of the nucleus. Simply because once bound in a state of lower potential, it takes energy to return the electron to a state of higher potential.

    In the case of normal polyatomic ions, an electron is gained to facilitate the general reduction in energy for other electrons in the molecule.

    The octet rule shows that a full principle energy shell is naturally lower in energy due to the pairing of electrons in all orbitals http://en.wikipedia.org/wiki/Exchange_interaction

    This lowers the potential of the system, thus making it a more stable compound. There are also other effects that lower its potential when an octet is formed, such as increased electronegativity decreasing the potential of electrons and such. In order for an atom to gain electrons in higher orbitals beyond an octet and such, it must decrease the potential beyond the point that the energy of the new electrons increases it.

    As for the sulfite ion, perhaps the increased electrical neutrality (addition of electron to sulfur, and subsequent removal of electron from an oxygen) offsets the electron density around sulfur, making the entire molecule more stable, even if it compromises sulfur's energy.

    I hope that helped, I wish I could help more:/ , but honestly someone else on here help because I'm curious for the exact reason myself :D
    Last edited: Nov 7, 2011
  4. Nov 7, 2011 #3
    Sorry, this is like the second year I'm studying chemistry, so all I know are the basics...I don't really understand electronegativity, potential or exchange interaction...
  5. Nov 7, 2011 #4
    Okay, here's what I know:

    The guy tells me to change the formal charges. I discovered that if I follow what he did, two atoms will have zero formal charges, while the other two will still have negative one formal charges. But what's the use of this? Why can't you stick with the original version (without the double bond in the sulfite ion)?
  6. Nov 7, 2011 #5
    Don't worry! Hahaha, if your interested you will get far. Electronegativity is really easy to understand, although a lot of teachers can try to make it seem more complicated (perhaps I am oversimplifying here). But, electronegativity is basically the difference in charges "focused" on certain electrons, since the inner electrons are closer to the nucleus, the charge on them is high, the actual difference in electronegativity occurs when you add subsequently higher orbitals with electrons, the outer electrons actually repel the inner electrons, pushing the inner electrons to a closer proximity relative to the nucleus, and pushing the outer electrons away from the nucleus, basically shifting the charge from acting on all electrons equally to acting with greater force on the inner electrons and leaving the outer electrons with little positive charge. If that made sense lol.

    And as for an electron's potential, it is the same basic concept as newtonian physics, as the distance an object is from a certain field gives its potential energy, or energy it could transfer to kinetic energy if it moved closer to the origin of the field. http://en.wikipedia.org/wiki/Potential_energy Except that electron orbitals can only exist in certain energies/distances from the nucleus, because other intermediate orbitals are unstable for the electron.

    As for exchange reaction I don't know that many reasons behind it either, other than the tested fact that it lowers the potential of the system by pairing opposite spins... thats particle physics and I'm only going to delve into it superficially hahahaha...yet.

    Ill watch the video now... should have before I wrote that lol
  7. Nov 7, 2011 #6
    Actually I knew about electronegativity last year. I tried to get ahead of the class. But after the holidays I forgot everything about it. I find chemistry quite daunting, it's not the same with physics. At least I still haven't seen violations of physics laws yet. (Or maybe I'm just a beginner).

    By the way, your explanation above is quite easy to understand. Thanks! :D
  8. Nov 7, 2011 #7
    Okay, What I am thinking is that the delocalized electron (resonate electron) increases the stability of the molecule, http://en.wikipedia.org/wiki/Resonance_(chemistry [Broken]) As for the formal charge of the atom, it seems (from my novice opinion) having not seen formal charge for a while that it is simply a way to arithmetically deduce the charge of an atom, so perhaps the instability of the first (octet obeying) sulfate is reduced by removing the charges from the oxygen and sulfur, neutralizing them, but I can't be sure.
    Last edited by a moderator: May 5, 2017
  9. Nov 7, 2011 #8
    Yeah, at least physics (classical at least) is simple... if I can say that. hahaha, chemistry is sort of its application, I am doing simple honors chemistry this year, (I'm 16) but I tried to "get ahead" aswell... worst ...or best decision I ever had, still not sure. Still understanding doesn't gain you college credit :/ hahaha Anyways at least my explanation is understandable :) What class are you taking this year?
  10. Nov 7, 2011 #9
    I'm doing pre-IB. Gotta get myself familiar with the basics first. I've managed to get through the local system with good grades by just memorizing. Not gonna depend on that this year.

    I guess this video explains a lot to me:
    Last edited by a moderator: Sep 25, 2014
  11. Nov 7, 2011 #10


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    Staff: Mentor

    Octet rule is - at best - only a "rule of thumb", a simplification of the reality, no wonder it is often violated. Don't treat it too seriously.
  12. Nov 7, 2011 #11


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    For sulfite, the first structure containing only single bonds is entirely correct while the second one is wrong although - as you well observed - the contrary is often to be found in some (either old or dubious) books. The reason these incorrect structures are still perpetuated lies in the fact that in former times it was thought that the d orbitals of Sulfur and Phosphorus could participate in bonding. However, when detailed calculations became possible, this turned out to be incorrect.
    A correct description of the bonding in higher main group compounds involves the high ionicity and polarizability of the bonds.
  13. Nov 7, 2011 #12
    I have another question:
    Peroxide has a formula of O2 with negative 2 charge. I looked at the lewis structure, and each oxygen atom has a negative 1 formal charge. I thought that atoms with same charges repel each other? So...why is there peroxide?
  14. Nov 8, 2011 #13


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    Staff: Mentor

    It can be explained in terms of MO (molecular orbitals) theory. There are enough electrons on bonding orbitals to make O22- stable in some environments (although it is highly reactive in general).
  15. Nov 8, 2011 #14


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    Multiply charged anions are usually only stable either in a crystal, where the positive counter charges are nearby or as a solution in polar solvents which tend to screen the electric field. E.g. in gas phase O2- isn't stable either.
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