Dismiss Notice
Join Physics Forums Today!
The friendliest, high quality science and math community on the planet! Everyone who loves science is here!

Hybridization of SnCl3 -

  1. May 25, 2012 #1
    I just had my exam today, so I had this question that is, what is the type of hybridization which occurs in SnCl3 - , my answer is sp2 hybridization because I thought that the last electron left wouldn't be hybridized. Am I correct?
  2. jcsd
  3. May 25, 2012 #2


    User Avatar
    Science Advisor

    Hybridization is not an observable property of an atom, so this question makes little sense.
    More appropriate would be to ask with which kind of assumed hybridization would best fit the observed bond properties. Normally it makes little sense to consider hybridization beyond second row atoms as the size of s and p orbitals is becomes too different and the energy from bond formation doesn't make up for the propagation. So the best description would probably one where there is no hybridization, the free electron pair fills the s orbital and the bonding is due to the p orbitals, only.
  4. May 25, 2012 #3
    hmmm, let me quote the whole question to avoid any miscommunication.
    Tin (Sn) is an element of Group 14. What type of hybrid orbitals would tin use in SnCl3 (-1)

    A. sp
    B. sp2
    C. sp3
    D. sp3d

    I'm not sure if this changes anything but I hope there's an answer to this.

    PS: one more question
    For the hydrogen atom, the electron transition with the longest wavelength is

    A. n5 to n4
    B. n7 to n 6
    C. n6 to n1
    D. n2 to n1

    My friends say the answer should be B because the energy released when the electron moves from n=7 to n=6 is the smallest, thus the frequency is the smallest, and the wavelength is the longest.

    However I did a quick check on the internet, it seems most the answer given were D.

    Am I overlooking something important?
  5. May 25, 2012 #4
    I think it should be sp3. 2 paired electrons in the s orbital and 3 unpaired electron in p orbital used for bonding
  6. May 25, 2012 #5
    wouldn't that be 5 valence electrons when there's only 3 chlorine attached and the charge is negative 1
  7. May 25, 2012 #6


    User Avatar
    Science Advisor

    Answer to the second question is D as you corretly said. The first question doesn't exactly rise my confidence in the school system of your country.
  8. May 26, 2012 #7
    Yes, and I would like to add that the majority of my country's people would agree with you. The country is going through brain drain, there's no helping it.

    Anyhow, can you give a brief explanation as to why the answer is D? Because I remember the difference of frequency between two energy levels decreases as it goes higher up, so the frequency produced for answer B should much lower than D.
  9. May 27, 2012 #8


    User Avatar

    Staff: Mentor

    Strange, I would expect B to be correct as well (for the hydrogen question).
  10. May 27, 2012 #9
    Taking a look at a model of SnCl3- (http://ce.sysu.edu.cn/echemi/inocbx/ic3/Sn/SnCl3-.html) it looks like it can be described with an hybridization. If the free electron filled the s orbital and the bonding were due to the p orbitals, wouldn't the bond angles be closer to 90 degrees?
  11. May 29, 2012 #10


    User Avatar
    Science Advisor

    I am deeply sorry. I wanted to write "B is correct". Confounded the letters.
  12. May 29, 2012 #11


    User Avatar
    Science Advisor

    According to crystal structures of salts containing the anion (e.g. K[SnCl_3].KCl.H2O), the bond angle found is indeed very close to 90 degree, e.g. 87.7 and 90.8 degrees for the salt mentioned.
    See: http://dx.doi.org/10.1016/0022-1902(62)80247-4 [Broken]
    Similar conclusions are found from IR or Raman spectroscopy in the solution:
    Last edited by a moderator: May 6, 2017
  13. May 29, 2012 #12
    Is there any way to find this hybridization without any actual spectroscopy or such experiments? I mean, can something like the electronic configuration be used? After all, the OP was asked this for in a theoretical manner.
  14. May 29, 2012 #13


    User Avatar
    Science Advisor

    My general advice is never to invoke hybridization when you are not forced to do so, especially not in main group compounds beyond the second row.
    Keep it simple!
    I gave a theoretical justification in post #2.
    Even in compound like water and NH3 little is gained in assuming hybridized orbitals.
    Also bond angles aren't a good argument in general. After all, there is no physical law why orbitals should line up 100% for optimal bonding.
  15. May 29, 2012 #14
    From post #2...

    I'm not sure I understand that properly. You say, the energy from bond formation doesn't make up for the propagation, which propagation is that?

    Thanks for the reply :smile:

    But the last statement puts me into confusion again. Bond angles do indicate the structure where the molecule is in maximum stability which every molecule is trying to attain, right? So orbitals should pretty much line up like that, I believe. Wasn't hybridization brought in to explain this, in the first place?
  16. May 29, 2012 #15


    User Avatar
    Science Advisor

    I mean the energy necessary to lift e.g. an s electron to a p electron before hybridization. I confused promotion with propagation, sorry.
  17. May 29, 2012 #16
    Aha! Did some more reading into it, and I get it now. Thanks :smile:
  18. May 29, 2012 #17
    Yes, there is a way within a theoretical framework. Count all of the electrons, etc..., create your Lewis structure, and then count the number of electron domains, from there you'd get your molecular geometry. Within this case, there is a lone pair of electrons on the tin atom, but it is also a tetrahedral molecule (electron-pair geometry), however, its molecular geometry is trigonal pyramidal, so the hybridization for tetrahedral generally is sp^3.

    Sometimes, for instance in general chemistry, or organic chemistry, you can get the hybridization on that alone, however, there are instances where a molecules geometry won't exactly match up to its hybridization so knowing how to draw and deduce the sigma/pi - bonds, and knowings vesper theory would be more paramount than simply going on the "status-quo" so to speak. But within his course and this question, it is simply trying to test him on a more basic level of understanding.
  19. May 29, 2012 #18
    Gotcha. Thanks for the explanation!
  20. May 30, 2012 #19


    User Avatar
    Science Advisor

    up to here, that is plain VSEPR theory.

    This has not the least scientific basis. Can you give a serious reference (I don't mean a high school text book) for this claim?
  21. May 30, 2012 #20
    Last edited by a moderator: May 6, 2017
  22. May 30, 2012 #21


    User Avatar
    Science Advisor

    There is no proof by authority in science.

    From the first link you post:
    ouch... it hurts.

    No, but seriously: Don't take those introductory texts too seriously. It is not uncommon that they are written by people who themselves do not have a deep understanding of theory (just for reference: That does not mean that they are not good scientists; just that you cannot assume that a typical organics professor has any concept of what a "Slater determiant", a "molecular orbital" or a "covalent bond" actually is, and should take their theoretical explanations with a grain of salt).

    Regarding "tetrahedral implies sp3": For example, tetrahedral coordination is also found in coordination complexes and these are most definitely not sp3 hybridized. And even in covalent complexes, tetrahedral coordination can result for any number of reasons; but expecting a hybridization to occur in later than 2nd row atoms is normally not called for, for the reasons DrDu already explained earlier in this thread.

    To get a clearer picture of those issues, you might want to read Kutzeniggs "Chemical bonding in higher main group elements" (http://dx.doi.org/10.1002/anie.198402721 ).
    Last edited by a moderator: May 6, 2017
  23. May 31, 2012 #22
    It isn't that I am taking it too seriously, it is just that I am telling the other poster, for his general chemistry course what to expect. I know it is just a mathematical model trying to understand how certain molecules would interact/bond, but my point here isn't being solely focused on the uselessness of going past sp, or up to d orbitals, rather, as said before, if posed the question of it (previous post stated "theoretical framework") then you could use a simple way of getting what it may look like based on the mathematical framework developed by Pauling.

    As noted before: "This is basic general chemistry 101"..., not an overly advanced course. And, in the U.S. this is taught the exact way those links state, we aren't taught or some of us aren't told that going beyond the second row is useless, usually we are given a bunch of question of a bunch of chemical molecules and are told to figure out its hybridization... However, at least in organic chemistry it is rather different. The professor I took organic chemistry with only stressed that the sp2 hybridization was only valid and we didn't deal with anything beyond that. However, we did touch on sp3 of carbon-hydrogen bonds though, but that isn't what I am trying to discuss here.

    Thanks for the link though. I appreciate being more knowledgeable of chemistry. But, one thing, I can only view the abstract...
  24. Jun 1, 2012 #23


    User Avatar
    Science Advisor

    Kurokari doesn't seem to be from US. So I think it is quite irrelevant what may be taught in a specific class, what counts is a correct understanding of science.
    Having done my thesis in molecular physics I share the folling complaints with many theoretical chemists:

    1. Why are introductory chemistry students treated with mutilated theories of bonding they don't have the mathematical means to understand instead of being taught chemistry?

    2. Why do introductory texts perpetuate even those concepts from Paulings books (which are still apedagogical masterpieces and a source of scientific insight) from the first half of the 20th century which have been proven wrong or seriously modified since then when quantitative checks of the predictions became possible with the introduction of computers?

    I think the dominance of Pauling in introductory chemistry is - not too surprisingly - especially an US problem.

    As far as the article by Kutzelnigg (who wrote many pedagogical articles trying to correct some of the most serious theoretical misconceptions among chemists) I fear you either have to buy it or get it from your library.
Share this great discussion with others via Reddit, Google+, Twitter, or Facebook