Posts Tagged ‘double 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.

Thursday, December 1st, 2016

Following on from a search for long C-C bonds, here is the same repeated for C=C double bonds.

sq

The query restricts the search to each carbon having just two non-metallic substituents. To avoid conjugation with these, they each are 4-coordinated; the carbons themselves are three-coordinated. Further constraints are the usual no disorder, no errors and R < 0.1 and the C=C distance > 1.4Å (the standard value is ~1.32-1.34Å). The search query is deposited as DOI: 10.14469/hpc/1959[1]

c_c

The apparent longest example is LIRVEN, DOI: 10.5517/CC4R2MK[2] with a value of 1.589Å, longer than most C-C single bonds! Closer inspection reveals the presence of lithium cations, and so the molecule bearing the C=C bond must sustain two negative charges. So this apparent C=C bond is in fact anionic, with one electron going into each of the π* orbitals, thus lengthening the CC bond.‡ Not a true example of a neutral C=C bond[3] but it now becomes interesting for what its spin state might be. Is it a biradical or a triplet for example? One to be investigated further I fancy! Another example of this type is QUKCEE[4]

10-5517cc4r2mk-lirven

This next FAZWIM has a C=C length of 1.546Å. It is an old structure (1986), and comes without attached hydrogen atoms. Although drawn with no hydrogens on the central C=C bond, the length suggests this molecule is simply mis-assigned.†fazwin-diag fazwim

The final example I will highlight is pretty ordinary looking and published in 2016 as a private communication; ALOVOO, DOI: 10.5517/CCDC.CSD.CC1LJSWS[5] with a C=C length of 1.443Å. Again no obvious reason for the bond to be longer than normal.‡†

10-5517ccdc-csd-cc1ljsws-alovoo

In hunting for such unusual deviations from the norm, the most obvious explanation is normally some anomaly in the crystallographic analysis. Although the CSD (crystal structure database) is a very heavily curated resource, it seems unlikely that each deposition would be carefully inspected for its chemistry, and this must be our task here. But such anomalies can themselves point to interesting or unusual chemistry, which in  turn can be subjected to quantum computation to see if either the unusual value can be replicated or other reasons identified.  In this case, this exercise can been conducted by a human, but one can easily envisage the entire process being automated on a far larger scale.  The future?

‡ In fact the stoichiometry shows each “double bond” is actually a di-anion, with two electrons entering each of the the π* orbitals.

†A calculation on the singlet state for the structure as drawn (ωB97XD/Def2-TZVPP, DOI: 10.14469/hpc/1960) gives a bond length of 1.342Å, i.e. that expected for a double bond. The triplet state is similar in energy, but with a much longer central bond length of 1.476Å, DOI: 10.14469/hpc/1962 but the geometry at the carbons is planar and not bent as shown above. The quintet state is 1.45Å and is again planar, doi 10.14469/hpc/1963. So calculations on FAZWIM strongly suggest the structure as shown is an error.

‡†The computed value is 1.324Å, perfectly normal. DOI: 10.14469/hpc/1966[6]

References

  1. H. Rzepa, "Long C=C bonds", 2016. https://doi.org/10.14469/hpc/1959
  2. Matsuo, T.., Watanabe, H.., Ichinohe, M.., and Sekiguchi, A.., "CCDC 141348: Experimental Crystal Structure Determination", 2000. https://doi.org/10.5517/cc4r2mk
  3. T. Matsuo, H. Watanabe, M. Ichinohe, and A. Sekiguchi, "Reduction of the 1,4,5,8-tetrasila-1,4,5,8-tetrahydroanthracene derivative with lithium metal. Isolation and characterization of the tetralithium salt of a tetraanion, and observation of an Si–H⋯Li+ interaction", Inorganic Chemistry Communications, vol. 2, pp. 510-512, 1999. https://doi.org/10.1016/s1387-7003(99)00136-7
  4. T. Matsuo, H. Watanabe, and A. Sekiguchi, "A Novel Tetralithium Salt of a Tetraanion and a Dilithium Salt of a Dianion, Formed by the Reduction of the Tetrasilylethylene Moiety. Synthesis, Characterization, and Observation of an Si-H···Li+ Interaction", Bulletin of the Chemical Society of Japan, vol. 73, pp. 1461-1467, 2000. https://doi.org/10.1246/bcsj.73.1461
  5. M.E. Light, S. Bain, and J. Kilburn, "CCDC 1475906: Experimental Crystal Structure Determination", 2016. https://doi.org/10.5517/ccdc.csd.cc1ljsws
  6. H. Rzepa, "ALOVOO", 2016. https://doi.org/10.14469/hpc/1966

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Molecule orbitals as indicators of reactivity: bromoallene.

Thursday, September 1st, 2016

Bromoallene is a pretty simple molecule, with two non-equivalent double bonds. How might it react with an electrophile, say dimethyldioxirane (DMDO) to form an epoxide?[1] Here I explore the difference between two different and very simple approaches to predicting its reactivity. bromoallene

Both approaches rely on the properties of the reactant and use two types of molecule orbitals derived from its electronic wavefunction. The first of these is very well-known as the molecular orbital (MO), which has the property that it tends to delocalise over all the contributing atoms (the “molecule”). MOs are often used in this context; the highest energy occupied MO is thought of as being associated with the most nucleophilic (electron donating) regions of the molecule and so such a HOMO would be expected to predict the region of nucleophilic attack. The second is known as the natural bond orbital (NBO), which is evaluated in a manner that tends to localise it on bonds (the functional groups or reaction centres) and atom centres. What do these respective orbitals reveal for bromoallene? 

The MOs
HOMO, -0.3380 HOMO-1, -0.3692 au
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Click for 3D
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The NBOs
HONBO, -0.3769 HONBO-2, -0.3898
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The table above shows the energies (in Hartrees) of the four relevant orbitals. The less negative (less stable) the orbital, the more nucleophilic it is. The (heavily) delocalized HOMO is located on the C=C bond bond carrying the C-Br bond, Δ1,2 alkene, but it also has a large component on the Br. The more stable HOMO-1 is located on the C=C bond located away from the Br, the Δ2,3 alkene and also with a (different type of) component on the Br.

In contrast, the HONBO is located on the Δ2,3 alkene and it is the HONBO-2 that is on the Δ1,2 alkene. Both these orbitals have very little “leakage” onto other atoms, they are almost completely localised.

Well, now we have a problem since these two analyses lead to diametrically opposing predictions! So what does experiment say? A recent article[1] addresses this issue by isolating the initially formed epoxide from reaction with DMDO and characterising it using crystallography. But here comes the catch; such isolation only proved possible if the allene was also substituted with large sterically bulky groups such as t-butyl or adamantyl. And the isolated product was the Δ1,2 epoxide. So does that mean that the MO method was correct and the NBO method wrong? Well, not necessarily. Those large groups play an additional role via steric effects. To factor in such effects one has to look at the transition state model for the reaction rather than depending purely on the reactant properties. And the steric effects in this case appear to win out over the electronic ones.[1]

The Klopman[2]-Salem[3] equation (shown in very simplified, and original, form below for just the covalent term) casts some light on what is going on. This term is a double summation over occupied/unoccupied (donor-acceptor) orbital interactions, involving the coefficients of the orbitals (the overlap integrals in effect) in the numerator and the energy difference between the occupied/unoccupied orbital pair as denominator.

KS1

Performing such a double summation is rarely attempted; instead the equation is reduced to just one single term involving the donor of highest energy and the acceptor of lowest energy, ensuring the energy difference is a minimum and hence the term itself is (potentially) the largest in the summation. There is still the issue of the orbital coefficients, and here we get to the crux of the difference between the use of MOs and NBOs. You can see by inspection that the two π-MOs for bromoallene have different coefficients on the two atoms of interest, the two carbons of the double bond. One really has to evaluate the size of this term in the summation by using quantitative values for the respective coefficients and to very probably include the further terms in the summation for any other orbitals which also have significantly non-zero coefficients on these two atoms. But with the NBOs, the localisation procedure used to derive them has reduced the coefficients to just the carbon atoms and effectively no other atoms; all the other terms in the double summation in effect do drop out entirely. So with NBOs, the only number that matters is the energy difference between the occupied/empty orbitals (the denominator). But since the acceptor (the electrophile, DMDO in this case) is the same for both regiochemistries, things reduce even further to just comparing the donor energies for the two alternatives (Table above). The higher/less stable of these will have the greater contribution in the Klopman-Salem equation.

This little molecule teaches the important lesson that electronic and steric effects both play a role in directing reactions, and in this system they may well oppose each other. Simple interpretations based on reactant orbitals may give only a partial and even potentially misleading answer.

References

  1. D. Christopher Braddock, A. Mahtey, H.S. Rzepa, and A.J.P. White, "Stable bromoallene oxides", Chemical Communications, vol. 52, pp. 11219-11222, 2016. https://doi.org/10.1039/c6cc06395k
  2. G. Klopman, "Chemical reactivity and the concept of charge- and frontier-controlled reactions", Journal of the American Chemical Society, vol. 90, pp. 223-234, 1968. https://doi.org/10.1021/ja01004a002
  3. L. Salem, "Intermolecular orbital theory of the interaction between conjugated systems. I. General theory", Journal of the American Chemical Society, vol. 90, pp. 543-552, 1968. https://doi.org/10.1021/ja01005a001

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A wider look at π-complex metal-alkene (and alkyne) compounds.

Monday, June 13th, 2016

Previously, I looked at the historic origins of the so-called π-complex theory of metal-alkene complexes. Here I follow this up with some data mining of the crystal structure database for such structures.

Alkene-metal "π-complexes" have what might be called a representational problem; they do not happily fit into the standard Lewis model of using lines connecting atoms to represent electron pairs. Structure 1 was the original representation used by Dewar intending the meaning of partial back donation from a filled metal orbital to the empty π* of the alkene. At the other extreme these compounds can be called metallacyclopropanes (2) in which only single bonds feature (these can be thought of as representing full back bonding from metal to alkene and full forward bonding from alkene to metal). Representations 3 and 4 are a more fuzzy blend of these, implying some sort of partial bond order for the metal-carbon bonds. Taken together, they imply that the formal bond order of the C-C bond might vary between single to double. Structures 1 and 2 in particular imply that there might be two distinct ways in arranging the bonding and that π-complexes and metallacyclopropanes might therefore be distinct valence-bond isomers, each potentially capable of separate existence.

Why do these representations matter? Well, I am going to mine the crystal structure database for these species to try to see if there is any evidence for a bimodal distribution in the C-C lengths, perhaps indicating evidence of the isomerism suggested above. Such a structural database is indexed against atom-pair connectivity in the first instance and then bond type; one can specify the following types of bond connecting any two atoms: single, double, triple, quadruple, polymeric, delocalised, pi and any. It is not entirely obvious which if any of these types apply to structure 1 (it is not possible to draw a bond ending at the mid-point of another bond using the Conquest structure editor); the dashed lines in structures 3 and 4 could be classed as delocalised, pi, or most generally any. The search query can be constructed thus, where the two carbons carry R which can be either H or C and all four C-R bonds are specified as acyclic (to try to avoid complications by excluding compounds such as cyclic metallacenes). Because representation 1 cannot be constructed in the editor, I am going to specify that each carbon carries four bonds of any type in the first instance. The torsion specified is defined as R-C-C-M and the full queries can be found deposited here.[1]

If the metallacyclopropane representation 2 is defined with explicit single bonds, one gets only 22 hits (no errors, no disorder, R < 0.1). The distribution of C-C bond lengths is shown below. Already one sees a representational problem emerging. A true metallacyclopropane might be expected to show a C-C single bond length, say > ~1.5Å. But only one or two of these examples actually have this value, the most probable value being ~1.4Å.

Using representation 3, one gets 1861 hits, but as before one sees a maximum at ~1.4Å with a tail reaching to both single and double bond values for the C-C distance.‡

If the C-C bond is also specified as "any", the hits increase to 3948, but the bond length distribution is still very similar, with no sign of any bimodal distribution.

Such a distribution is however found if the torsions between the R-C bond vector and the C-M bond vector are plotted (for all types of bond). A large number of the complexes have a torsion <90°, which suggests that in fact the substituent R is probably interacting with the metal (even though this would lead to formal cyclicity, specifying R-C as acyclic does not detect this interaction). Could this be masking a bimodal distribution in the C-C lengths?

If the previous search is repeated, but this time specifying that all four torsions must lie in the range 90-180° (the range expected for a "classical" alkene-metal complex and selecting only the top right hand side cluster in the plot above) the reduced value of 1051 hits are obtained, but the monomodal distribution remains.

For this last set, here is a plot of the two C-metal bond length, with colour indicating the C-C bond length, indicating the two C-metal bonds are clearly linearly correlated.

One final variation;  the atom on either C can only be H or a 4-coordinate (sp3) carbon; 645 hits. Again, a monomodal distribution centered at 1.4Å.

So this foray through metal alkene complexes suggests that there is a continuum between the formal metallacyclopropane with a C-C single bond and the only slightly perturbed alkene-metal complex with a C=C double bond. Whilst this would not prevent any one of these compounds existing as two distinctly different valence-bond isomers, it makes it very unlikely. I had noted in an earlier post that for molecules of the type RX≡XR (X=Si, Ge, Sn, Pb) that there was indeed a clear bimodal distribution of the X-X lengths evident in the crystal structures (for a relatively small sample number). The structures 1-4 shown at the start of this post are all simply just variations in a continuum and not distinct isomers.

POSTSCRIPT:  I noted above the bimodel distribution in compounds involving formal triple bonds. So I repeated the search above for π-complex metal-alkyne complexes. Specifying an acyclic C-R bond, and any for the CC bond type, one gets the following.

There is now a tantalizing suggestion of two clusters, one at 1.3 and another at 1.4Å. The torsional distribution shows that the latter distance appears to be associated with much smaller torsions, whereas the top right cluster is associated with shorter lengths.

If the torsions are restricted to the range 90-180, then the histogram looses the smaller cluster, and perhaps gains a second cluster at 1.22Å?  As I said, all quite tantalizing!

‡The tail in all the histograms extends into the 1.1-1.3Å region, which seems unreasonable for a carbon where four bonds are specified. This region probably represents errors in the crystallographic analysis or reporting. But who knows, perhaps some very unusual compounds are lurking there!

 

References

  1. H. Rzepa, "A wider look at the π-complex theory of metal-alkene compounds.", 2016. https://doi.org/10.14469/hpc/642

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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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