Posts Tagged ‘Hypervalent molecule’

Hypervalent or not? A fluxional triselenide.

Saturday, February 24th, 2018

Another post inspired by a comment on an earlier one; I had been discussing compounds of the type I.In (n=4,6) as possible candidates for hypervalency. The comment suggests the below as a similar analogue, deriving from observations made in 1989.[1]

This compound was investigated using 77Se NMR, with the following conclusions:

  1. The compound is fluxional, with the lines at room temperature broadened compared to those at -50°C.
  2. At -50°C the peaks are sharp enough to discern 1JSe-Se couplings, with multiplicities and integrations that suggest a central Se is surrounded by four equivalent further Se atoms, with shifts of 655.1 and 251.2 ppm.
  3. The magnitude of this 1JSe-Se coupling (391 Hz) leads to the suggestion of a considerable contribution of a resonance form with Se=Se bonds (structure 2 above).
  4. This was supported by 2J13C-77Se couplings which also imply a symmetrically coordinated central  Se.
  5. Thus the two resonance forms 1 or 2 above were suggested as the predominant form at -50°C, with an increasing incursion of the open chain isomer 3 at higher temperatures giving rise to the observed fluxional dynamic behaviour.
  6. One may surmise from these results that the central Se is certainly hypercoordinated and by the classical interpretations hypervalent.

Here are some calculations (R=H), at the ωB97XD/Def2-TZVPP/SCRF=chloroform level.‡ In red are the calculated Wiberg Se-Se bond orders, which give little indication of any Se=Se double bond character. 

The calculated 77Se shifts are shown in magenta, with the observed values being 655 and 255 ppm. The match is not good, the errors were 120 and 20.5 ppm.†  However calculated shifts for elements adjacent to e.g. Se or Br etc suffer from relativistic effects such as spin orbit coupling.[2] Thus the shift for the central Se, surrounded by four other Se atoms is likely to have a significant error, but the error for the four other Se atoms should be less. The reverse is true.

However, all the calculations of this species (up to Def2-TZVPPD basis set) showed this symmetric form of D2h symmetry to actually be a transition state, as per below.

There is a minimum with the structure below in which one pair of Se-Se lengths are longer than the other pair and for which the free energy is 6.5 kcal/mol lower. The Wiberg bond orders for the two sets of Se-Se bonds are now 0.16 and 0.86, which very much corresponds to structure 3 above.

Assuming that this compound is fluxional even at -50°C, the average of the pairs of Se atoms gives calculated shifts of 667 ppm (655 obs) whilst the central Se is 204.6 ppm (251 obs). The latter, influenced by two especially short Se-Se distances, is likely to have a very large spin-orbit coupling error, whilst for the former the error will be smaller (13C shifts adjacent to one Br typically have induced calculated errors of about 14 ppm[2]).

At this point I searched the Cambridge structure database for Se coordinated by four other Se atoms. A close analogue[3] has the structure shown below, in which pairs of Se-Se interactions have unequal bond lengths, the shorter being ~2.45Å. This matches the calculation above reasonably well.

Reconciling these various observations, we might assume that even at -50°C the fluxional behaviour has not been frozen out. Given that the fluxional barrier is only 6.5 kcal/mol, it is unlikely that the spectrum could be measured at a sufficiently low temperature to reveal not two sets of Se signals in the ratio 4:1 but three in the ratio 2:2:1. The spin-spin couplings reported presumably are a result of averaging a genuine 1JSe-Se coupling with a through space coupling.

So it appears that the analysis of the 77Se NMR reported in this article [1] may not be quite what it seems. A better interpretation is that structure 3 is the most realistic. This means no hypercoordination for the Se, never mind hypervalence!

‡FAIR data at DOI: 10.14469/hpc/3724. †The original reference, Me2Se was incorrectly calculated without solvation by chloroform. The values shown here are now corrected from those shown in the original post.

References

  1. Y. Mazaki, and K. Kobayashi, "Structure and intramolecular dynamics of bis(diisobutylselenocarbamoyl) triselenide as identified in solution by the 77Se-NMR spectroscopy", Tetrahedron Letters, vol. 30, pp. 2813-2816, 1989. https://doi.org/10.1016/s0040-4039(00)99132-9
  2. D.C. Braddock, and H.S. Rzepa, "Structural Reassignment of Obtusallenes V, VI, and VII by GIAO-Based Density Functional Prediction", Journal of Natural Products, vol. 71, pp. 728-730, 2008. https://doi.org/10.1021/np0705918
  3. R.O. Gould, C.L. Jones, W.J. Savage, and T.A. Stephenson, "Crystal and molecular structure of bis(NN-diethyldiselenocarbamato)-selenium(II)", Journal of the Chemical Society, Dalton Transactions, pp. 908, 1976. https://doi.org/10.1039/dt9760000908

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Hypervalent Helium – not!

Friday, February 16th, 2018

Last year, this article[1] attracted a lot of attention as the first example of molecular helium in the form of Na2He. In fact, the helium in this species has a calculated‡ bond index of only 0.15 and it is better classified as a sodium electride with the ionisation induced by pressure and the presence of helium atoms. The helium is neither valent, nor indeed hypervalent (the meanings are in fact equivalent for this element). In a separate blog posted in 2013, I noted a cobalt carbonyl complex containing a hexacoordinate hydrogen in the form of hydride, H–. A comment appended to this blog insightfully asked about the isoelectronic complex containing He instead of H–. Here, rather belatedly, I respond to this comment!

The complex [HCo6(CO)15]– has a calculated bond index at the hydrogen of 0.988 and a calculated NMR chemical shift of 21.6 ppm (ωB97XD/Def2-TZVPPD calculation) compared to a measured value of 23.2 ppm. Despite being six-coordinate, the hydride has a bond index that does not exceed one (it is not hypervalent).

So here is the neutral helium analogue. The He bond index emerges as 0.71 at the geometry of the hydride complex. Compare this with the bond index of 0.15 calculated for Na2He and it would be fair to say that at this geometry, the helium in [HeCo6(CO)15] would have a greater claim to be a molecular compound. Back in 2010, extrapolating from a series of posts here, I had speculated[2] about other molecular species of He, including the di-cation below. This has a He bond index of 0.54, rather less than that in [HeCo6(CO)15] but much more than in Na2He. It is also vibrationally stable.

But now, [HeCo6(CO)15] goes “pear-shaped” (why do pears have such a bad press?). I started a process of optimizing the geometry of this complex (ωB97Xd/Def2-TZVPPD). Slowly, the He started to creep out of the centre of the complex and emerge from the cavity. After about 100 steps it reached the geometry shown below, at which point the Wiberg bond index has dropped to 0.62 and still going down. I think it might take a few more steps to be completely expelled, but I have stopped the geometry optimisation at this stage.

So helium appears not to be valent in [HeCo6(CO)15]. However, I have yet to try Ne, which is both larger and softer. I will post results here.

‡All data at 10.14469/hpc/3587.

References

  1. X. Dong, A.R. Oganov, A.F. Goncharov, E. Stavrou, S. Lobanov, G. Saleh, G. Qian, Q. Zhu, C. Gatti, V.L. Deringer, R. Dronskowski, X. Zhou, V.B. Prakapenka, Z. Konôpková, I.A. Popov, A.I. Boldyrev, and H. Wang, "A stable compound of helium and sodium at high pressure", Nature Chemistry, vol. 9, pp. 440-445, 2017. https://doi.org/10.1038/nchem.2716
  2. H.S. Rzepa, "The rational design of helium bonds", Nature Chemistry, vol. 2, pp. 390-393, 2010. https://doi.org/10.1038/nchem.596

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Are diazomethanes hypervalent molecules? An attempt into more insight by more “tuning” with substituents.

Tuesday, December 26th, 2017

Recollect the suggestion that diazomethane has hypervalent character[1]. When I looked into this, I came to the conclusion that it probably was mildly hypervalent, but on carbon and not nitrogen. Here I try some variations with substituents to see what light if any this casts.

I have expanded the resonance forms of diazomethane by one structure from those shown in the previous two posts (a form by the way not considered in the original article[1]) to include a nitrene. This takes us back to an earlier suggestion on this blog that HC≡S≡CH is not a stable species but a higher order saddle point which distorts down to a bis-carbene, together with the suggestion that hypervalent triple bonds have the option of converting four of the six electrons into two carbene lone pairs, replacing the triple bond with a single bond. This in turn harks back to G. N. Lewis’ 101 year old idea for acetylene itself!

To explore this mode, I start by replacing the terminal ≡N in diazomethane with a ≡C-Me group, which cannot absorb electrons into lone-pairs in the manner that nitrogen can. A ωB97XD/Def2-TZVPP calculation‡ reveals that the linear form is a transition state for interconversion into a carbene. The IRC for the process (below) shows this carbene is ~10 kcal/mol lower than the linear “hypervalent” form. 

NBO analysis of this transition state reveals a similar orbital pattern to diazomethane itself, including a non-bonding orbital on the H2C carbon. The Wiberg carbon bond indices are 3.6764 and N 3.6454 and the bond orders C=N 1.1390 and N=CMe 1.6192.

ELF analysis of this transition state reveals the presence of two non-bonding pairs on the carbon atoms either side of the nitrogen but unshared with it, with populations of 1.19e and 1.37e (DFT). That nitrogen really does not like excess electrons! The four atoms C,N,C,C have ELF valence basins totalling 8.00, 6.94, 7.69 and 7.92e (DFT) or 8.07, 7.07 and 7.61e (CASSCF), suggesting that unlike diazomethane itself, the octet-excess induced hypervalence on carbon is slightly decreased.

Pumping even more electrons in by replacing the ≡C-Me group with ≡C-NH2 does not increase any hypervalence, but does induce more electrons to reside in “lone pairs”. Of the four atoms along the chain, three have “lone pairs” associated with them, a total of 4.83e that do not contribute to bonds (valence).

An electron withdrawing ≡C-CN group replacing the ≡C-NH2 reverses the effect of the latter, but this linear species is still a transition state for carbon isomerisation:

Finally, combining all we have learnt by adding in nitro groups on the first carbon. This is no longer a transition state but now a stable species; the sum of the ELF basin integrations around the carbon on the left reaches 8.95e, slightly higher than the dinitro-diazomethane discussed in the previous post. The numerical Wiberg atom bond indices are C 3.8713, N 3.6898, C 3.8503, C 3.9958 and N 3.0288 for the atoms along the chain, with the first nitrogen the “least-valent”.

So we see that “hypervalence”, or at least “octet-excess”, which is not exactly the same as hypervalence since it includes contributions from non-bonding electrons, is balanced on a knife-edge. Trying to increase the octet-excess by pumping electrons in turns the system into a transition state for carbene formation. Octet-excess is seen as a metastable property, to be relieved by geometric distortions where possible or localization of electrons into non-bonding lone pairs. And I remind yet again that no evidence has manifested in calculations of the molecules above that the central nitrogen of these diazomethane-like systems has any propensity for octet or valence-excess as implied by the formula C=N≡X.[1]

‡FAIR data for all calculations is available at DOI: 10.14469/hpc/3476

References

  1. M.C. Durrant, "A quantitative definition of hypervalency", Chemical Science, vol. 6, pp. 6614-6623, 2015. https://doi.org/10.1039/c5sc02076j

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Are diazomethanes hypervalent molecules? An attempt into more insight by more "tuning" with substituents.

Tuesday, December 26th, 2017

Recollect the suggestion that diazomethane has hypervalent character[1]. When I looked into this, I came to the conclusion that it probably was mildly hypervalent, but on carbon and not nitrogen. Here I try some variations with substituents to see what light if any this casts.

I have expanded the resonance forms of diazomethane by one structure from those shown in the previous two posts (a form by the way not considered in the original article[1]) to include a nitrene. This takes us back to an earlier suggestion on this blog that HC≡S≡CH is not a stable species but a higher order saddle point which distorts down to a bis-carbene, together with the suggestion that hypervalent triple bonds have the option of converting four of the six electrons into two carbene lone pairs, replacing the triple bond with a single bond. This in turn harks back to G. N. Lewis’ 101 year old idea for acetylene itself!

To explore this mode, I start by replacing the terminal ≡N in diazomethane with a ≡C-Me group, which cannot absorb electrons into lone-pairs in the manner that nitrogen can. A ωB97XD/Def2-TZVPP calculation‡ reveals that the linear form is a transition state for interconversion into a carbene. The IRC for the process (below) shows this carbene is ~10 kcal/mol lower than the linear “hypervalent” form. 

NBO analysis of this transition state reveals a similar orbital pattern to diazomethane itself, including a non-bonding orbital on the H2C carbon. The Wiberg carbon bond indices are 3.6764 and N 3.6454 and the bond orders C=N 1.1390 and N=CMe 1.6192.

ELF analysis of this transition state reveals the presence of two non-bonding pairs on the carbon atoms either side of the nitrogen but unshared with it, with populations of 1.19e and 1.37e (DFT). That nitrogen really does not like excess electrons! The four atoms C,N,C,C have ELF valence basins totalling 8.00, 6.94, 7.69 and 7.92e (DFT) or 8.07, 7.07 and 7.61e (CASSCF), suggesting that unlike diazomethane itself, the octet-excess induced hypervalence on carbon is slightly decreased.

Pumping even more electrons in by replacing the ≡C-Me group with ≡C-NH2 does not increase any hypervalence, but does induce more electrons to reside in “lone pairs”. Of the four atoms along the chain, three have “lone pairs” associated with them, a total of 4.83e that do not contribute to bonds (valence).

An electron withdrawing ≡C-CN group replacing the ≡C-NH2 reverses the effect of the latter, but this linear species is still a transition state for carbon isomerisation:

Finally, combining all we have learnt by adding in nitro groups on the first carbon. This is no longer a transition state but now a stable species; the sum of the ELF basin integrations around the carbon on the left reaches 8.95e, slightly higher than the dinitro-diazomethane discussed in the previous post. The numerical Wiberg atom bond indices are C 3.8713, N 3.6898, C 3.8503, C 3.9958 and N 3.0288 for the atoms along the chain, with the first nitrogen the “least-valent”.

So we see that “hypervalence”, or at least “octet-excess”, which is not exactly the same as hypervalence since it includes contributions from non-bonding electrons, is balanced on a knife-edge. Trying to increase the octet-excess by pumping electrons in turns the system into a transition state for carbene formation. Octet-excess is seen as a metastable property, to be relieved by geometric distortions where possible or localization of electrons into non-bonding lone pairs. And I remind yet again that no evidence has manifested in calculations of the molecules above that the central nitrogen of these diazomethane-like systems has any propensity for octet or valence-excess as implied by the formula C=N≡X.[1]

‡FAIR data for all calculations is available at DOI: 10.14469/hpc/3476

References

  1. M.C. Durrant, "A quantitative definition of hypervalency", Chemical Science, vol. 6, pp. 6614-6623, 2015. https://doi.org/10.1039/c5sc02076j

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Can any hypervalence in diazomethanes be amplified?

Saturday, December 23rd, 2017

In the previous post, I referred to a recently published review on hypervalency[1] which introduced a very simple way (the valence electron equivalent γ) of quantifying the effect. Diazomethane was cited as one example of a small molecule exhibiting hypervalency (on nitrogen) by this measure. Here I explore the effect of substituting diazomethane with cyano and nitro groups.‡

Firstly, dicyanodiazomethane. NBO analysis reveals the following atom bond indices; C, 3.810; N 3.834; N 2.971. Compare these values to diazomethane itself, C, 3.716; N 3.802; N 2.907 and you can see that the carbon bond index has increased slightly. The ELF basin integrations (below) which also take into account the “lone pair” on carbon are: C, 8.55, N, 6.65, N, 7.52 (DFT), again compared with diazomethane as C, 8.16; N, 6.59; N, 7.52. The CASSCF(14,14) result is very similar.

So the “γ(C)”† has increased from 8.2 to 8.55. Next, dinitrodiazomethane;

The NBO bond indices are C, 3.8203; N 3.8255; N 2.9802 and ELF integrations C, 8.82, N, 6.68, N, 7.49 (DFT).

So “γ(C)”†  increases along the series 8.16 → 8.55 → 8.82, whereas “γ(N)” changes as 6.59 → 6.65 → 6.68, a smaller effect. Whilst 8.82 is still some way off the value of γ(N)=10 quoted[1] for diazomethane, dinitrodiazomethane is still a pretty good candidate for hypervalent carbon. The question now is whether even larger values of “γ(C)” can be identified in other molecules. 

‡FAIR data for all calculations is available at DOI: 10.14469/hpc/3476. †The quotes in “γ(C)” indicate it is calculated here using ELF integrations rather than charge maps.

References

  1. M.C. Durrant, "A quantitative definition of hypervalency", Chemical Science, vol. 6, pp. 6614-6623, 2015. https://doi.org/10.1039/c5sc02076j

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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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Hypervalence revisited. The odd case of hexamethyl selenium.

Tuesday, November 7th, 2017

One thread that runs through this blog is that of hypervalency. It was therefore nice to come across a recent review of the concept[1] which revisits the topic, and where a helpful summary is given of the evolving meanings over time of the term hypervalent. The key phrase “it soon became clear that the two principles of the 2-centre-2-electron bond and the octet rule were sometimes in conflict” succinctly summarises the issue. Two molecules that are discussed in this review caught my eye, CLi6 and SeMe6. The former is stated as “anomalous in terms of the Lewis model“, but as I have shown in an earlier post, the carbon is in fact not anomalous in a Lewis sense because of a large degree of Li-Li bonding. When this is taken into account, the Lewis model of the carbon becomes more “normal”. Here I take a look at the other cited molecule, SeMe6.

I should start by summarising what I think are two fundamental ways in which electrons can be added to the valence shell of a main group element.

  • If an s/p basis only is used for the (main group) valence shell, then once eight electrons have populated the four bonding molecular orbitals constructed from this basis, additional electrons can then go into the four antibonding orbitals. This simple concept is often taught as an explanation for why the bond orders across the range N≡N, O=O, F-F and Ne…Ne decrease regularly (in fact one could also add CC to the left of this series, which is thought to have a weak quadruple bond). The Lewis octet is maintained throughout.
  • The second fundamental possibility is to expand the valence shell basis set to s/p/d or s/s/p/p (the second s or p-shell is the Rydberg level, i.e. 3s for carbon) or even s/s/p/p/d. That would be a true Lewis octet expansion, which depends on identifying significant Rydberg occupancy. This latter is in fact very rare, and few examples have been conclusively identified. One such as been discussed on this blog  and more examples were presented at the Aachen bond Slam in September 2017. Unlike the first mechanism (which reduces bond orders), this one actually increases bond orders and any bonds where the atomic orbital contributions have a significant Rydberg component can be considered as “Hyperbonds“.

One can then address the issue of hypervalency (and any octet expansion) by analysing the basis set contributions of the orbitals. These orbitals can be either canonical molecular orbitals or localised (NBO) orbitals. If a single determinant wavefunction is appropriate, then the orbitals would be doubly occupied (for closed shell species). If the molecule has multi-reference character, then of course fractional electron occupancy of these orbitals may be required (as would e.g. be the case for ozone, O3, another molecule asserted in the review as hypervalent[1]).

There are other ways of analysing the wavefunction. The one discussed at length in the review[1] is from an atomic charge map, but also mentioned is an ELF partitioning. This derives not from an orbital population but from the distribution of a function (ELF) calculated from the electron density itself. It was this latter method that was cited for SeMe6. The ELF method partitions electrons into so-called basins, which can be monosynaptic (lone pairs and ionic bonds), disynaptic (covalent bonds) and more rarely trisynaptic (3-centre bonds). Using this analysis, six disynaptic octahedrally-arranged ELF basins were located for  SeMe6  (“in which the Se–C bonds are relatively non-polar, can have electron populations exceeding 8 at the central atom”[1]) and for which the total integration cames to 11.34e (FAIR Data DOI for this calculation can be see at 10.14469/hpc/3219).

The key phrase is “non-polar”, since the Lewis concept relates to shared electron pair or covalent hypervalency. It was this aspect that I focused on seven years back in looking at whether e.g. IF7 was hypervalent (along with I(CN)7). These were too ionic to reveal disynaptic covalent ELF basins. So in an effort to reduce the polarity, I tried II7 and At.At7, on the grounds that these homonuclear molecules might be less polar. The seven I-I or At-At ELF disynaptic basins integrated to totals of 6.55 and 6.47e respectively; there was no evidence of “octet expansion” for either central halogen. Instead of course, the six electrons in the octet-excess needed to create seven I-I bonds actually populate I-I antibonding orbitals, as per method 1 above. Accordingly the I-I bond orders reduce from 1.0 to e.g. 0.47 for the axial substituents and 0.37 for the five equatorial groups.

One interesting property of the centroid of the ELF basins is that you can infer the polarity of the bond from its position along the bond axis. For II7, the centroids are displaced towards the central iodine, indicating it is more electronegative, and away from the terminal iodine, indicating it is the electropositive partner. I mention this since the ELF basins for SeMe6 show the centroids to be strongly displaced towards the carbon and away from the Se (0.38/0.62), indicating that this molecule is in fact polar and NOT non-polar as was asserted. 

To follow-up this latter observation, I did an NBO analysis of the wavefunction for SeMe6 (FAIR Data DOI: 10.14469/hpc/3220). This reveals the following properties.

  • Se populations: [core]4S( 1.30)4p( 3.04)4d( 0.06)5p( 0.01) of which Rydberg = 0.06516, natural charge on Se, 1.60224
  • C populations:  [core]2S( 1.20)2p( 3.65)3S( 0.01)3p( 0.01) of which Rydberg = 0.01838, natural charge on C -0.86670
  • H populations:  1S( 0.80)
  • Total Rydberg population: 0.20556
  • Wiberg Se-C bond orders:  0.6689
  • Wiberg bond indices:  Se 4.0431, C 3.8252,  H 0.9621

So according to this orbital-based analysis, SeMe6 is in effect a partially ionic compound with no evidence of significant Rydberg occupancies and hence no evidence of any octet expansion at Se. Thus we see two different interpretations emerging, depending on the analytical method used:

  1. SeMe6 is a polar molecule with no hypervalent attributes as judged using orbital analysis.
  2. As a polar molecule, it has six methyl carbanion-like substituents in which the carbon “lone pairs” all point towards the Se, manifesting as disynaptic ELF basins and indicating a total valence-basin octet-expanded population of 11.34e at Se. Even though this octet expansion originates mostly from the carbon atomic orbitals, the disynaptic nature of these valence basins means that the Se could indeed be defined as hypervalent.

Well, SeMe6 has turned out to be rather less clear-cut than implied by the assertion “in which the Se–C bonds are relatively non-polar”. There are however possible modifications to SeMe6 that might yet make it less polar. These may be the subject of a follow-up post.

References

  1. M.C. Durrant, "A quantitative definition of hypervalency", Chemical Science, vol. 6, pp. 6614-6623, 2015. https://doi.org/10.1039/c5sc02076j

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Real hypervalency in a small molecule.

Sunday, February 21st, 2016

Hypervalency is defined as a molecule that contains one or more main group elements formally bearing more than eight  electrons in their  valence shell. One example of a molecule so characterised was CLi6[1] where the description "“carbon can expand its octet of electrons to form this relatively stable molecule“ was used. Yet, in this latter case, the octet expansion is in fact an illusion, as indeed are many examples that are cited. The octet shell remains resolutely un-expanded. Here I will explore the tiny molecule CH3F2- where two extra electrons have been added to fluoromethane.

Two such electrons added to e.g. such a methane derivative can be in principle accommodated in two ways:

  1. The electrons on carbon could expand the octet shell by populating molecular orbitals constructed using 3s or 3p atomic orbitals (AOs) as well as the normal 2s and 2p shells. This is also the normal "explanation" for expanded octets, the assumption being that as one moves down the rows of the periodic table (e.g. P, S, Cl, etc) these shells become energetically more accessible (e.g. the 3d or 4s shell for P, S, Cl etc). In fact, for e.g. PF5, the occupancy of such  "Rydberg" shells is only ~0.2 electrons, not a significant octet expansion.
  2. The electrons can instead or as well as populate the antibonding molecular orbitals (MOs) formed from just the 2s/2p AOs. For a methane derivative, there are four bonding MOs (into which the octet of electrons are placed) and four anti-bonding MOs all constructed from the total of eight AOs. Well known examples of populating antibonding MOs are the series N≡N, O=O (singlet), F-F, Ne…Ne where the additional electrons are added to anti-bonding MOs and have the effect of reducing the bond orders from 3 to 2 to 1 to 0. And of course all core shells contain populated bonding and antibonding pairs.

Here are some ωB97XD/Def2-TZVPPD/scrf=water calculations. All these species are molecules with all-real vibrations, being stable toward dissociation to e.g. CH3– + H– or CH3– + F–.  A transition state for this latter dissocation with IRC[2] can be characterised. In all cases the energy of the highest occupied MO or NBO is -ve, meaning that the electrons are bound, at least in part due to the solvent field applied.

Molecule Wiberg CH order Wiberg CF order Natural Populations E HONBO, au dataDOI
CH42- 0.773 –

C:[core]2S(1.98)2p(3.82)3S( 0.15)4d( 0.01)

H:1S( 1.00)

-0.144CH4 [3]
CH3F2- 0.980 1.213

C:[core]2S(1.05)2p( 3.20)3S(1.26)4p( 0.01)4d( 0.01)

H:1S( 0.84)2S( 0.01)2p( 0.02)

F:[core]2S(1.88)2p( 5.61)3S( 0.30)3p( 0.04)3d( 0.01)4p( 0.01)

-0.068
Click for 3D

Click for 3D

 [4]
CH2F22- 0.871 0.897

C:[core]2S(1.60)2p( 2.64)3S(0.39)3p( 0.01)4d( 0.01)

H:1S(1.19)2S( 0.06)

F:[core]2S(1.86)2p( 5.52)3S( 0.01)3p( 0.01)4p( 0.01)

-0.281
Click for 3D

Click for 3D

 [5]
CF42- – 0.801

C:[core]2S(1.94)2p( 1.96)3S( 0.19)3p( 0.04)5d( 0.01)

F:[core]2S(1.89)2p( 5.54)3p( 0.01)3d( 0.02)

 

-0.148CF4 [6]
  1. CH42- shows only small Rydberg occupancy (< 0.2e), but a significantly reduced bond order for the four C-H bonds (each C-H bonding NBO also has some antibonding character for the other three CHs) and hence the molecule is not truly hypervalent.
  2. CH3F2- in contrast shows quite different behavour. The C-H bond order is almost 1 and the C-F bond order is actually >1. Of the two extra electrons, ~1.28 now occupy carbon Rydberg AOs and the fluorine also has significant Rydberg population (~0.36e). So this is a real hypervalent system, in which the total valencies exceed that expected from an octet.
  3. CH2F22- is somewhere inbetween the previous two systems. The carbon has modest Rydberg occupancy (~0.4e) but there is also significant occupation of the antibonding MOs. Both the C-H and C-F bond orders are <1.
  4. CF42- shows a further reduction in the C Rydberg occpancy (<0.2) and the C-F bond order is also reduced. This reduction in bond order is also seen in other so-called hypervalent systems such as PF5.

So of these systems, CH3F2- can be reasonably called hypervalent, whilst the others have much less such character. It does appear that there is a fine balance between placing extra electrons into Rydberg orbitals to expand the "octet" and hence valencies, and placing them in anti-bonding orbitals where the individual valencies are actually reduced. It seems that substituting methane with just one fluorine encourages population of the Rydberg orbitals, but that more fluorines encourage instead population of the antibonding orbitals. What is remarkable is that CH3F2- actually has a (small) barrier to dissociation. The challenge now is to try to design a system which has a significant Rydberg population, a low antibonding population AND is stable to dissociation; this will require some inspiration. So do not hold your breaths!

 

References

  1. H. Kudo, "Observation of hypervalent CLi6 by Knudsen-effusion mass spectrometry", Nature, vol. 355, pp. 432-434, 1992. https://doi.org/10.1038/355432a0
  2. https://doi.org/
  3. H.S. Rzepa, "C 1 H 4 -2", 2016. https://doi.org/10.14469/ch/191837
  4. H.S. Rzepa, "C 1 H 3 F 1 -2", 2016. https://doi.org/10.14469/ch/191919
  5. H.S. Rzepa, "C 1 H 2 F 2 -2", 2016. https://doi.org/10.14469/ch/191918
  6. H.S. Rzepa, "C 1 F 4 -2", 2016. https://doi.org/10.14469/ch/191916

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