What is the best Lewis Structure of ICl4- when using a formal charge argument?

Most resources (such as this Yahoo! Answer: http://answers.yahoo.com/question/index?qid=200712… list IC4- as a 6 electron domain species. It’s worth being familiar with this question / answer before continuing.

Formal charge guidelines tell us to put the negative formal charge on the more electronegative element, yet (as the answer above describes) a 6 electron domain version of ICl4- puts the negative formal charge on the iodine.

Why not create a double bond (using one of the lone pairs on the iodine) between I and Cl, thus putting a negative formal charge on Cl, following the rules of formal charge, and resulting in a 5 electron domain, see saw atom arrangement.

My guess is that empirical data suggests that ICl4- is in fact square planar, but how can I explain this based on formal charge?

In response to the submitted answers by Spunk and Doctorwho.

Spunk: thanks for the detailed response. It turns out though that the octet ‘rule’ only applies for 2nd row nonmetals. As soon as you have access to d orbitals (3rd row and lower) the octet rule can be broken. For example, sulfurhexafluoride or xenon trioxided. Also a 5 electron domain species (by putting a double bond between Cl and I) would actually INCREASE the space between between the domains (a change from octahedral to trigonal bipyramidal) in the equatorial plane from 90 degrees to 120 degrees.

Doctorwho: ah…orbital overlap! Yeah, that must be it. I hadn’t thought about the diffuseness of the orbitals involved in the double bond I proposed.

4 Answers

  • I can’t answer all of this, but I can answer this part:

    “… Why not create a double bond (using one of the lone pairs on the iodine) between I and Cl, …”

    Very simple: the (most likely) p orbitals that would be involved in double bond formation, 5p on iodine and 3p on chlroine, would have very poor overlap and thus would not be able to form a stable double bond. (By the time you get to 5p orbitals, the space they occupy is pretty large, so their electron cloud is fairly diffuse.)

  • This Site Might Help You.

    RE:

    What is the best Lewis Structure of ICl4- when using a formal charge argument?

    Most resources (such as this Yahoo! Answer: http://answers.yahoo.com/question/index?qid=200712… list IC4- as a 6 electron domain species. It’s worth being familiar with this question / answer before continuing.

    Formal charge guidelines tell us to put the negative formal…

  • You can’t put a double bond on Cl because of the octet rule, which would be broken if you had a double bond. With the structure suggested in the answer, all Cl’s follow the octet rule. I does break the octet rule, however, it is very large and polarizable and it has access to d orbitals. Cl is less polarizable and does not have access to d orbitals.

    With those two considerations in mind… basically. the molecule is just more stable with the negative on the I than with it on the Cl.

    Plus, with the square planar shape, the molecules are the maximum distance apart. If you put the negative on the Cl (a double bond, at that), that would shove the other Cls VERY close to eachother, and all of that electron density so close is not a stable conformation.

  • It would be set up like this “:” means lone pair of electrons, ignore the “~” ~~~:~~:~~~~: H—O—Br===O: ~~~:~~: the formal charge would be as follows: H: 1-1/2(2)-0=0 O: 6-1/2(4)-4=0 Br: 7-1/2(6)-4=0 O: 6-1/2(4)-4=0 Br can hold extra electrons because it can leak those electrons over to the D orbital

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