Posts Tagged ‘single bond’

Hypervalence and octet-expansion in trimethylene-λ6-sulfane and related species.

Monday, November 27th, 2017

Previously: “Non-polar” species such as SeMe6, SMe6, ClMe3, ClMe5 all revealed interesting properties for the Se-C, S-C or Cl-C “single” bonds. The latter two examples in particular hinted at internal structures for these single bonds, as manifested by two ELF basins for some of the bonds. Here I take a look at the related molecule where a formal double bond between carbon and the central sulfur atom replacing the single-bond might also hint at octet expansions and hypervalence.

Starting with X=Y=Z=CH2,‡ the calculated (ωB97Xd/Def2-TZVPP) geometry has an interesting chiral D3-symmetric form. The density based ELF-basin centroids are shown below, with each formal C=S π-double bond represented by two ELF basins above and below the C-S axis and with each pair of ELF basins being twisted by 48° with respect to the other two pairs. The total valence shell count around the S is 10.98e and the octet is “expanded” (by ~3e).

The orbital-based NBO approach indicates little utilisation of higher (Rydberg) atomic orbital shells (S: [core]3S(1.13)3p(3.35)3d(0.11)4p(0.02); C: [core]2S(1.15)2p(3.77)3p(0.01)3d(0.01) ). Each S-C bond has a Wiberg bond order of 1.36 (significantly less than a double bond), and the central S has an overall bond index of 4.102. There is a real mis-match between the orbital partitioning (2*1.36 = 2.72e) and the ELF partitioning (2*1.83 = 3.66e) into the S-C bonds. The former indicates that ~two of the twelve valence electrons are entering into anti-bonding orbitals to reduce the total bond index from a possible six to just four, but that they still contribute to the electron-density based ELF disynaptic C-S basins. To cast light on this behaviour, successively one to three of the CH2 groups can be replaced by O.

For each “S=O” bond, we find the ELF basin population more or less halves and electrons instead populate the non-bonding O “lone pairs”. The S-C ELF populations in contrast remain approximately constant. These species therefore have “double” S=C bonds but just “single” S-O bonds. The Rydberg population increases slightly; S: [core]3S(1.06)3p(2.95)3d(0.16)4p(0.02)) and the S bond index is 4.18 for one oxygen and S: [core]3S(0.99)3p(2.67)3d(0.19)4p(0.02) and S bond index 4.16 for two oxygens.

Sulfur trioxide (below) seems best represented by S-O rather than S=O bonds. The Rydberg population is S: [core]3S(0.91)3p(2.41)3d(0.21)4p(0.03) and the S bond index is 4.32.

Just for good measure sulfur trisulfide S(S)3 shows rather lower lone pair population because of course it is less electronegative than oxygen, and hence has a slightly greater S-S ELF basin population. Rydberg, S: [core]3S(1.43)3p(3.73)3d(0.21)4p(0.03) and central S bond index 4.04.

It seems molecules where the electrons in a valence shell exceed the “octet” are only too happy to let the excess electrons leak out into adjacent electronegative atoms as lone pairs, where they are no longer classified as  “shared”. Trimethylene-λ6-sulfane does not have this option and the excess electrons remain in the region of the valence shell, but here they do not contribute to augmenting the bond index at the central atom.  In this specific interpretation, the octet is exceeded, but hypervalence is not induced. It is a slippery concept; one where general agreement about its properties may indeed be difficult to achieve!

‡The FAIR data DOI collection for this post is 10.14469/hpc/3316.

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Long C-C bonds.

Wednesday, November 30th, 2016

In an earlier post, I searched for small C-C-C angles, finding one example that was also accompanied by an apparently exceptionally long C-C bond (2.18Å). But this arose from highly unusual bonding giving rise not to a single bond order but one closer to one half! How long can a “normal” (i.e single) C-C bond get, a question which has long fascinated chemists.

A naive search of the CSD is not as straightforward as it seems. Using the simple sub-structure R3C-CR3 as the search query gives LIRPEI, DOI: 10.5517/CCQ043Y[1] an apparently unexceptional molecule with a very exceptional C-C distance of 1.87Å. With long bonds one has to be ultra-careful to look at the crystallographic analysis before drawing any conclusions. One class of molecule where this has been done by many groups is the system shown below (red = long bond), with 47 entries and for which the longest C-C bond emerges with the value of 1.79Å[2]

 

long-bonds

long-cc

You can view this structure at DOI: 10.5517/CCS0R6Q[3] and the authors go to some pains to assure us that it is still a closed shell single bond, and not a biradical. That does seem to be the current record holder, but of course we are only talking here about molecules whose crystal structure has been determined.

I will end with an open question; how SHORT could a “single” C-C bond get? Here, a search of the CSD is entirely dominated by crystallographic artefacts, and I am not sure what the value might be. 

References

  1. Zhang, Jian., Chen, Shumei., Valle, H.., Wong, M.., Austria, C.., Cruz, M.., and Bu, Xianhui., "CCDC 655529: Experimental Crystal Structure Determination", 2008. https://doi.org/10.5517/ccq043y
  2. T. Takeda, H. Kawai, R. Herges, E. Mucke, Y. Sawai, K. Murakoshi, K. Fujiwara, and T. Suzuki, "Negligible diradical character for the ultralong C–C bond in 1,1,2,2-tetraarylpyracene derivatives at room temperature", Tetrahedron Letters, vol. 50, pp. 3693-3697, 2009. https://doi.org/10.1016/j.tetlet.2009.03.202
  3. Takeda, T.., Kawai, H.., Herges, R.., Mucke, E.., Sawai, Y.., Murakoshi, K.., Fujiwara, K.., and Suzuki, T.., "CCDC 715703: Experimental Crystal Structure Determination", 2010. https://doi.org/10.5517/ccs0r6q

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A (very) short history of shared-electron bonds.

Tuesday, March 26th, 2013

The concept of a shared electron bond and its property of an order is almost 100 years old in modern form, when G. N. Lewis suggested a model for single and double bonds that involved sharing either 2 or 4 electrons between a pair of atoms[1]. We tend to think of such (even electron) bonds in terms of their formal bond order (an integer), recognising that the actual bond order (however defined) may not fulfil this value. I thought I would very (very) briefly review the history of such bonds.

  1. 1916: G. N. Lewis[1] proposed a model for carbon involving a cube with one electron at each corner, thus making an octet.‡ A single bond would be created by two atoms sharing a common edge (= 2 shared electrons), and a double bond by sharing a common face (= 4 shared electrons). The recognition that the formal bond order of two could be partitioned into one electron pair of σ symmetry and one of π was not achieved until ~1929 (by Hückel). It is also now recognised that whilst most bonds of order 1 are of type σ, a rare few can be π (these are called homo or “suspended” bonds).
  2. 1916: Lewis also speculates about a rather less well-known model comprising “eight electrons in which pairs are symmetrically placed about a center gives … the model of the tetrahedral carbon atom.” He then points out that two tetrahedra, attached by one, two or three corners each would represent the single, the double and the triple bond. The latter “represents the highest possible degree of union between two atoms“. He chooses acetylene as an example, representing it as H:C:::C:H and two “tautomers” (we would now call them valence bond isomers)† with lower bond orders, these being what we now call a bis-carbene and a biradical:Lewis
  3. 1965: It took a remarkable wait of 49 years (a span which encompasses the development and maturity of quantum mechanics) to extend the “highest possible degree of union” to the quadruple bond, identified by Cotton in the previously known compound [Re2Cl8]2-.[2].
    Click for 3D

    Click for 3D

    In fact, Mulliken had drawn a quadruple bond between the two carbons in C2 back in 1939[3] (see Table 1, p 779) but he probably thought of it as a very high energy excited state and that it did not merit further discussion. The latest thoughts are that C2  does indeed have (a weak) fourth bond[4] in its ground electronic state.

  4. 2005: Another 40 years elapsed before quintuple or “fivefold” bonding was discovered by Power[5] in ArCrCrAr. There has been a bit of a race since to discover the shortest example of this genre.
    Click for 3D

    Click for 3D

  5. 2013: Unlike the lower bond orders, where direct structural data for larger molecules is available, speculation about sextuple bonds is limited largely to theoreticians, who have been at it for quite a while. The latest thinking is summarised here[6] (also doi: 10.1039/C2CP43559D). The current best candidates for a sextuple bond include Mo2 and W2.
  6. What is the limit of the formal integer bond order? I do not believe anyone thinks that septuple or octuple bonds (formal or otherwise) will be discovered (or even speculated upon) any time soon, but there is no fundamental law which would prohibit them.[7] Quite possibly if we get beyond element 120 in the periodic table, examples might emerge!

‡A formula for predicting the filled electron shells is 2(N+1)2, which gives the values 2, 8, 18, 32[8],[9] 50. It is also, as it happens, the rule for 3D aromaticity in clusters.

†A bis-carbene form, whilst not appropriate for carbon, may indeed become more realistic as one proceeds down column 14 of the periodic table. Thus [10], where Ar-Sn≡Sn-Ar has a C-Sn-Sn bond angle of 125°.

Click for 3D.

Click for 3D.

Or perhaps an even better example[11] with a C-Sn-Sn angle of 98°. There is also an example of C-Pb-Pb[12] with an angle of 94°.

References

  1. G.N. Lewis, "THE ATOM AND THE MOLECULE.", Journal of the American Chemical Society, vol. 38, pp. 762-785, 1916. https://doi.org/10.1021/ja02261a002
  2. F.A. Cotton, "Metal-Metal Bonding in [Re<sub>2</sub>X<sub>8</sub>]<sup>2-</sup> Ions and Other Metal Atom Clusters", Inorganic Chemistry, vol. 4, pp. 334-336, 1965. https://doi.org/10.1021/ic50025a016
  3. R.S. Mulliken, "Note on Electronic States of Diatomic Carbon, and the Carbon-Carbon Bond", Physical Review, vol. 56, pp. 778-781, 1939. https://doi.org/10.1103/physrev.56.778
  4. S. Shaik, H.S. Rzepa, and R. Hoffmann, "One Molecule, Two Atoms, Three Views, Four Bonds?", Angewandte Chemie International Edition, vol. 52, pp. 3020-3033, 2013. https://doi.org/10.1002/anie.201208206
  5. T. Nguyen, A.D. Sutton, M. Brynda, J.C. Fettinger, G.J. Long, and P.P. Power, "Synthesis of a Stable Compound with Fivefold Bonding Between Two Chromium(I) Centers", Science, vol. 310, pp. 844-847, 2005. https://doi.org/10.1126/science.1116789
  6. F. Ruipérez, M. Piris, J.M. Ugalde, and J.M. Matxain, "The natural orbital functional theory of the bonding in Cr<sub>2</sub>, Mo<sub>2</sub>and W<sub>2</sub>", Phys. Chem. Chem. Phys., vol. 15, pp. 2055-2062, 2013. https://doi.org/10.1039/c2cp43559d
  7. G. Frenking, and R. Tonner, "The six-bond bound", Nature, vol. 446, pp. 276-277, 2007. https://doi.org/10.1038/446276a
  8. I. Langmuir, "Types of Valence", Science, vol. 54, pp. 59-67, 1921. https://doi.org/10.1126/science.54.1386.59
  9. J. Dognon, C. Clavaguéra, and P. Pyykkö, "Towards a 32‐Electron Principle: Pu@Pb<sub>12</sub> and Related Systems", Angewandte Chemie International Edition, vol. 46, pp. 1427-1430, 2007. https://doi.org/10.1002/anie.200604198
  10. A.D. Phillips, R.J. Wright, M.M. Olmstead, and P.P. Power, "Synthesis and Characterization of 2,6-Dipp<sub>2</sub>-H<sub>3</sub>C<sub>6</sub>SnSnC<sub>6</sub>H<sub>3</sub>-2,6-Dipp<sub>2</sub> (Dipp = C<sub>6</sub>H<sub>3</sub>-2,6-Pr<sup>i</sup><sub>2</sub>):  A Tin Analogue of an Alkyne", Journal of the American Chemical Society, vol. 124, pp. 5930-5931, 2002. https://doi.org/10.1021/ja0257164
  11. Y. Peng, R.C. Fischer, W.A. Merrill, J. Fischer, L. Pu, B.D. Ellis, J.C. Fettinger, R.H. Herber, and P.P. Power, "Substituent effects in ditetrel alkyne analogues: multiple vs. single bonded isomers", Chemical Science, vol. 1, pp. 461, 2010. https://doi.org/10.1039/c0sc00240b
  12. L. Pu, B. Twamley, and P.P. Power, "Synthesis and Characterization of 2,6-Trip<sub>2</sub>H<sub>3</sub>C<sub>6</sub>PbPbC<sub>6</sub>H<sub>3</sub>-2,6-Trip<sub>2</sub> (Trip = C<sub>6</sub>H<sub>2</sub>-2,4,6-<i>i</i>-Pr<sub>3</sub>):  A Stable Heavier Group 14 Element Analogue of an Alkyne", Journal of the American Chemical Society, vol. 122, pp. 3524-3525, 2000. https://doi.org/10.1021/ja993346m

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