One approach to reporting science which is perhaps better suited to the medium of a blog than a conventional journal article is the opportunity to follow ideas in unexpected, even unconventional directions. Thus my third attempt, like a dog worrying a bone, to explore hypervalency. I have, somewhat to my surprise, found myself contemplating the two molecules I8 and At8. Perhaps it might be better to write them as I(I)7 and At(At)7. This makes it easier to relate both to the known molecule I(F)7. What led to these (allotropes) of the halogens? Well, as I noted before, hypervalency is a concept rooted in covalency, albeit an excess of it! And bonds with the same atom at each end are less likely to be accused of ionicity. I earlier suggested that the nicely covalent IH7 was not hypervalent, with all the electrons which might contribute to hypervalency actually to be found in the H…H regions. The next candidate, I(CN)7 ultimately proved a little too ionic for comfort. So we arrive at II7. At the D5h geometry, it proves not to be a minimum, but a (degenerate) transition state for reductive elimination of I2 (I note parabolically that the 2010 Nobel prize for chemistry was awarded for reactions which involve similar reductive elimination of Pd and other metals to form covalent C-C bonds). Thus I8 is useful only as a thought experiment molecule, and not a species that could actually be made.
Posts Tagged ‘Hypervalency’
Hypervalency: Third time lucky?
Saturday, October 23rd, 2010Hypervalency: I(CN)7 is not hypervalent!
Sunday, October 17th, 2010In the last post, IH7 was examined to see if it might exhibit true hypervalency. The iodine, despite its high coordination, turned out not to be hypervalent, with its (s/p) valence shell not exceeding eight electrons (and its d-shell still with 10, and the 6s/6p shells largely unoccupied). Instead, the 14 valence electrons (7 from H, 7 from iodine) fled to the H…H regions. Well, perhaps H is special in its ability to absorb electrons into the H…H regions. So how about I(CN)7? (the species has not hitherto been reported in the literature according to CAS). The cyano group is often described as a pseudohalide, but the advantage of its use here is that it is about the same electronegativity as I itself, and hence the I-C bond is more likely to be covalent (than for example an I-F bond). As noted in the earlier blog, if the potentially hypervalent atom is very ionic, it can be difficult to know whether the electrons are truly associated with that atom, or whether they are in fact in lone pairs associated with the other electronegative atom (e.g. F). It is also important to avoid large substituents, otherwise steric interactions will cause problems around the equator.
Hypervalency: Is it real?
Saturday, October 16th, 2010The Wikipedia page on hypervalent compounds reveals that the concept is almost as old as that of normally valent compounds. The definition there, is “a molecule that contains one or more main group elements formally bearing more than eight electrons in their valence shells” (although it could equally apply to e.g. transition elements that would contain e.g. more than 18 electrons in their valence shell). The most extreme example would perhaps be of iodine (or perhaps xenon). The normal valency of iodine is one (to formally complete the octet in the valence shell) but of course compounds such as IF7 imply the valency might reach 7 (and by implication that the octet of electrons expands to 14). So what of IF7? Well, there is a problem due to the high electronegativity of the fluorine. One could argue that the bonds in this molecule are ionic, and hence that the valence electrons really reside in lone pairs on the F. Thus the apparently hypervalent PF5 could be written PF4+…F–, in which case the P is not really hypervalent after all. We need a compound with un-arguably covalent bonds. Well, what about IH7? One might probably still argue about ionicity (for example H+…IH6–) but that puts electrons on I and not H, and hence does not change any hypervalency on the iodine. Surely, if hypervalency is a real phenomenon, it should manifest in IH7?
Solid carbon dioxide: hexacoordinate carbon?
Friday, September 17th, 2010Carbon dioxide is much in the news, not least because its atmospheric concentration is on the increase. How to sequester it and save the planet is a hot topic. Here I ponder its solid state structure, as a hint to its possible reactivity, and hence perhaps for clues as to how it might be captured. The structure was determined (DOI 10.1103/PhysRevB.65.104103) as shown below.
The weirdest bond of all? Laplacian isosurfaces for [1.1.1]Propellane.
Wednesday, July 21st, 2010In this post, I will take a look at what must be the most extraordinary small molecule ever made (especially given that it is merely a hydrocarbon). Its peculiarity is the region indicated by the dashed line below. Is it a bond? If so, what kind, given that it would exist sandwiched between two inverted carbon atoms?
Quintuple bonds: part 2
Saturday, February 20th, 2010In the previous post, I ruminated about how chemists set themselves targets. Thus, having settled on describing regions between two (and sometimes three) atoms as bonds, they added a property of that bond called its order. The race was then on to find molecules which exhibit the highest order between any particular pair of atoms. The record is thus far five (six has been mooted but its a little less certain) for the molecule below
Quintuple bonds
Tuesday, February 16th, 2010Climbers scale Mt. Everest, because its there, and chemists have their own version of this. Ever since G. N. Lewis introduced the concept of the electron-pair bond in 1916, the idea of a bond as having a formal bond-order has been seen as a useful way of thinking about molecules. The initial menagerie of single, double and triple formal bond orders (with a few half sizes) was extended in the 1960s to four, and in 2005 to five. Since then, something of a race has developed to produce the compound with the shortest quintuple bond. One of the candidates for this honour is shown below (2008, DOI: 10.1002/anie.200803859) which is a crystalline species (a few diatomics which exist in the gas phase are also candidates; for other reviews of the topic see 10.1038/nchem.359, 10.1021/ja905035f and 10.1246/cl.2009.1122).
Blisteringly bent (quadruple) bonds
Saturday, January 23rd, 2010So ingrained is the habit to think of a bond as a simple straight line connecting two atoms, that we rarely ask ourselves if they are bent, and if so, by how much (and indeed, does it matter?). Well Hursthouse, Malik, and Sales, as long ago as 1978, asked just such a question about the unlikeliest of bonds, a quadruple Cr-Cr bond, found in the compound di-μ-trimethylsilylmethyl-bis-[(tri-methylphosphine) (trimethylsilylmethyI)chromium(II)(DOI: 10.1039/dt9780001314[1]). They arrived at this conclusion by looking very carefully at how the overlaps with the Cr d-orbitals might be achieved.
References
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M.B. Hursthouse, K.M.A. Malik, and K.D. Sales, "Crystal and molecular structure of di-µ-trimethylsilylmethyl-bis[(tri-methylphosphine)(trimethylsilylmethyl)chromium(
II )](4 Cr–Cr)", J. Chem. Soc., Dalton Trans., pp. 1314-1318, 1978. http://dx.doi.org/10.1039/dt9780001314
Contriving aromaticity from S≡C Triple bonds
Friday, January 1st, 2010In the previous post, the molecule F3S-C≡SF3 was found to exhibit a valence bond isomerism, one of the S-C bonds being single, the other triple, and with a large barrier (~31 kcal/mol, ν 284i cm-1) to interconversion of the two valence-bond forms. So an interesting extension of this phenomenon is shown below: