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(Hyper)activating the chemistry journal.

Monday, September 7th, 2009

The science journal is generally acknowledged as first appearing around 1665 with the Philosophical Transactions of the Royal Society in London and (simultaneously) the French Academy of Sciences in Paris. By the turn of the millennium, around 10,000 science and medical journals were estimated to exist. By then, the Web had been around for a decade, and most journals had responded to this new medium by re-inventing themselves for it. For most part, they adopted a format which emulated paper (Acrobat), with a few embellishments (such as making the text fully searchable) and then used the Web to deliver this new reformulation of the journal. Otherwise, Robert Hooke would have easily recognized the medium he helped found in the 17th century.

In 1994, a small group of us thought that one could, and indeed should go further than emulated paper. We argued [1] that journals should be activated by delivering not merely the logic of a scientific argument, but also the data on which it might have been based. Of course, we encountered the usual problem; doing this might cost publishers more in production resources, and in the absence of a market prepared to pay the extra, the business model did not make sense (to the publishers). Well, 15 years later, and most publishers are indeed now thinking about how their journals can be enhanced. A number of interesting projects (the RSC’s Project Prospect is one which strives to bring science alive) have emerged. Another is the topic of this blog; the activation of the journal with molecular coordinates and data using the Jmol applet.

Initially (~2005), this project met with resistance from publishers, and the issue really amounted to what the definitive version of a scientific article should be. Should that definitive version be printable? That model, after all had served the community well for more than 300 years! And journals from the very beginning are still as readable now as when first published. In other words, print lasts! But print is pretty limiting after all. For a start, it is limited to 2D static representations. Molecules, by and large, do their magic in a dynamic three dimensions (4D in an Einsteinian sense). But print is also expensive; not merely to produce, but to transport paper around the world.

From the turn of the millennium, a number of publishers, amongst them the American Chemical Society, started to evolve the scientific article such that the pre-eminent version would now be considered to be the HTML form (perhaps as a prelude to phasing out print entirely? See an interesting commentary by a journal editor) and perhaps a digital Acrobat form which would be deemed to loose some of its functionality once printed (again see here for how Acrobat can be used to enhance things). Again however, a chicken-and-egg scenario resulted. To enhance the articles with extra functionality (such as data), they would need to find authors prepared to put the extra work into preparing the material. In fact, most authors already do that, but they call it supporting information. This is often highly data rich, covering materials such as spectra, coordinates and other information nowadays provided to researchers for analysis. Unfortunately, what has been missing is the education of authors to provide this information in a proper digital form which can be easily re-used by others, and on a Web page, converted automatically to nice interactive models. Most spectra which form part of the supporting information are in fact still scanned versions of printed spectra!

Enter computational chemists. Nowadays, they live in a world that truly does not need printing! Almost all of their data is already suitably digital. So perhaps it is no surprise to find that when enhanced journal articles started appearing around 2005, many were produced by this group of chemists. By now perhaps you are wondering what such an article might look like. Well, the remainder of this blog will be devoted to listing some examples. You will also notice that they come exclusively from our own publications. Perhaps someone will find the time to collect a far more representative set to better illustrate the diversity of this form, and how it is evolving. Meanwhile, you might wish to take a look at the following.

Part 1: The early days: 1994 onwards

These examples all relied on a browser plugin called Chime, which is no longer with us! Hence the pages designed to invoke it no longer display properly. But the data associated with the articles is still there!

  1. An early 1994 example of (hyper)activating a journal article can be seen here as the preliminary communication and
  2. in 1995 here as the final full article. I am told that this was the article that actually inspired the developers of Chime to enhance (Netscape) with a chemical plugin.
  3. This one from 1998 illustrates how articles can decay in functionality when Chime is no longer available.
  4. An ab initio and MNDO-d SCF-MO Computational Study of Stereoelectronic Control in Extrusion Reactions of R2I-F Iodine (III) Intermediates, M. A. Carroll, S. Martin-Santamaria, V. W. Pike, H. S. Rzepa and D. A. Widdowson, Perkin Trans. 2, 1999, 2707-2714 with the supporting information here.
  5. Huckel and Mobius Aromaticity and Trimerous transition state behaviour in the Pericyclic Reactions of [10], [14], [16] and [18] Annulenes. Sonsoles Martên-Santamarêa, Balasundaram Lavan and H. S. Rzepa, J. Chem. Soc., Perkin Trans 2, 2000, 1415. with the supporting information here.
  6. Peter Murray-Rust, H. S. Rzepa and Michael Wright, “Development of Chemical Markup Language (CML) as a System for Handling Complex Chemical Content”, New J. Chem., 2001, 618-634. DOI: 10.1039/b008780g. This article broke new ground in that the supporting information was something of a misnomer. It was expressed entirely in XML, including all the chemistry data, and used XSLT transforms on the fly to regenerate the article. In that sense, it was actually a superset of the published article. It would be fair to say that this article was rather ahead of its time (although it does seem appropriate to publish it in a new journal!).
  7. M. Jakt, L. Johannissen, H. S. Rzepa, D. A. Widdowson and R. Wilhelm, “A Computational Study of the Mechanism of Palladium Insertion into Alkynyl and Aryl Carbon-Fluorine bonds”, Perkin Trans. 2, 2002, 576-581 and supporting information.
  8. P. Murray-Rust and H. S. Rzepa, chapter in “Handbook of Chemoinformatics. Part 2. Advanced Topics.”, ed. J. Gasteiger and T. Engel, 2003, Vol 1, was not enhanced per se, but did lay out the principles of how it might/should be done.
  9. K. P. Tellmann, M. J. Humphries, H. S. Rzepa and V. C. Gibson, “An experimental and computational study of β-H transfer between organocobalt complexes and 1-alkenes”, Organometallics, 2004, 23, 5503-5513. DOI: 10.1021/om049581h and supporting information.

Part 2: 2005.

These four examples all now invoke Jmol, which downloads upon request and hence does not rely on the presence of any browser plugin. The four articles were submited with supporting information in the form of HTML. These were associated with the main article, but were not formal part of that article. In that sense, they represent an incarnation of the traditional model, with all the data firmly resident in the supporting information.

  1. Gibson, Vernon C.; Marshall, Edward L.; Rzepa, H. S. ” A computational study on the ring-opening polymerization of lactide initiated by β-diketiminate metal alkoxides: The origin of heterotactic stereocontrol”, J. Am. Chem. Soc., 2005, 127, 6048-6051. DOI: 10.1021/ja043819b and supporting information.
  2. H. S. Rzepa, Mobius aromaticity and delocalization”, Chem. Rev., 2005, 105, 3697 – 3715. DOI: 10.1021/cr030092l and supporting information.
  3. H. S. Rzepa, “Double-twist Mšbius Aromaticity in a 4n+2 Electron Electrocyclic Reaction”, 2005, Chem Comm, 5220-5222. DOI: 10.1039/b510508k The supporting information is also available directly.
  4. H. S. Rzepa, “A Double-twist Mobius-aromatic conformation of [14]annulene”, Org. Lett., 2005, 7, 637 – 4639. DOI: 10.1021/ol0518333 and supporting information.

Part 3: 2006 onwards

The supporting information has now been assimilated into the main body of the article proper, and within these confines contribute components such as enhanced figures or tables (i.e. enhanced with data)

  1. A. P. Dove, V. C. Gibson, E. L. Marshall, H. S. Rzepa, A. J. P. White and D. J. Williams, “Synthetic, Structural, Mechanistic and Computational Studies on Single-Site β-Diketiminate Tin(II) Initiators for the Polymerization of rac-Lactide”, J. Am. Chem. Soc., 2006,128, 9834-9843. DOI: 10.1021/ja061400a The enhancement can be seen in Figure 11.
  2. O. Casher and H. S. Rzepa, “SemanticEye: A Semantic Web Application to Rationalise and Enhance Chemical Electronic Publishing”, J. Chem. Inf. Mod., 2006, 46, 2396-2411. DOI: 10.1021/ci060139e
  3. H S. Rzepa and M. E. Cass, “A Computational Study of the Nondissociative Mechanisms that Interchange Apical and Equatorial Atoms in Square Pyramidal Molecules”, Inorg. Chem., 2006, 45, 3958–3963. DOI 10.1021/ic0519988. Interactive table at 10.1021/ic0519988/ic0519988.html
  4. M. E. Cass and H. S. Rzepa, “In Search of The Bailar Twist and Ray-Dutt mechanisms that racemize chiral tris-chelates: A computational study of Sc(III), V(III), Co(III), Zn(II) and Ga(III) complexes of a ligand analog of acetylacetonate”, Inorg. Chem., 2007, 49, 8024-8031. DOI: 10.1021/ic062473y The enhancement can be seen in Figure 2
  5. H. S. Rzepa, “Lemniscular Hexaphyrins as examples of aromatic and antiaromatic Double-Twist Möbius Molecules”, Org. Lett., 2008, 10, 949-952.DOI:10.1021/ol703129z The enhancement can be seen in Web Table 1.
  6. D. C. Braddock and H. S. Rzepa, “Structural Reassignment of Obtusallenes V, VI and VII by GIAO-based Density functional prediction”, J. Nat. Prod., 2008, DOI: 10.1021/np0705918 and WEO1.
  7. S. M. Rappaport and H S. Rzepa, “Intrinsically Chiral Aromaticity. Rules Incorporating Linking Number, Twist, and Writhe for Higher-Twist Möbius Annulenes”, J. Am. Chem. Soc., 2008, 130,, 7613-7619. DOI: 10.1021/ja710438j and WEO1 to 4
  8. C. S. M. Allan and H. S. Rzepa, “AIM and ELF Critical point and NICS Magnetic analyses of Möbius-type Aromaticity and Homoaromaticity in Lemniscular Annulenes and Hexaphyrins”, J. Org. Chem., 2008, 73, 6615-6622. DOI: 10.1021/jo801022b and WEO1
  9. C. S. M. Allan and H. S. Rzepa, “Chiral aromaticities. Möbius Homoaromaticity”, J. Chem. Theory. Comp., 2008, 4, 1841-1848. DOI: 10.1021/ct8001915 and WEO1
  10. C. S. M Allan and H. S. Rzepa, “The structure of Polythiocyanogen: A Computational investigation”, Dalton Trans., 2008, 6925 – 6932. DOI: 10.1039/b810147g and enhanced Table
  11. H. S. Rzepa, “Wormholes in Chemical Space connecting Torus Knot and Torus Link π-electron density topologies”, Phys. Chem. Chem. Phys., 2009, 1340-1345. DOI: 10.1039/b810301a and enhanced Table.
  12. H. S. Rzepa, “The Chiro-optical properties of a Lemniscular Octaphyrin”, Org. Lett., 2009, 11, 3088-3091. DOI: 10.1021/ol901172g
  13. C. S. Wannere, H. S. Rzepa, B. C. Rinderspacher, A. Paul, H. F. Schaefer III, P. v. R. Schleyer and C. S. M. Allan, “The geometry and electronic topology of higher-order Möbius charged Annulenes”, J. Phys. Chem., 2009, DOI: 10.1021/jp902176a and enhanced table
  14. H. S. Rzepa, “The distortivity of π-electrons in conjugated Boron rings.”, Phys. Chem. Chem. Phys., 2009, DOI: 10.1039/B911817A and enhanced table.
  15. H. S. Rzepa, “The importance of being bonded”, Nature Chem., 2009, DOI: 10.1038/nchem.373 and the exploratorium.
  16. King Kuok Hii, J.L.Arbour, H.S.Rzepa, A.J.P.White, “Unusual Regiodivergence in Metal-Catalysed Intramolecular Cyclisation of γ-Allenols”, Chem. Commun, 2009, DOI: 10.1039/b913295c and enhanced table.
  17. L. F. V. Pinto, P. M. C. Glória, M. J. S. Gomes, H. S. Rzepa, S. Prabhakar, A. M. Lobo. “A Dramatic Effect of Double Bond Configuration in N-Oxy-3-aza Cope Rearrangements – A simple synthesis of functionalised allenes”, Tet. Lett., 2009, 50, 3446-3449. DOI: 10.1016/j.tetlet.2009.02.228 and interactive table.
  18. H. S. Rzepa and C. S. M. Allan, “Racemization of isobornyl chloride via carbocations: a non-classical look at a classic mechanism”, J. Chem. Educ., 2010, DOI: 10.1021/ed800058c and interactive table.
  19. K. Abersfelder, A. J. P. White, H. S. Rzepa, and D. Scheschkewitz “A Tricyclic Aromatic Isomer of Hexasilabenzene”, Science, 2010, DOI: 10.1126/science.1181771 and interactive table.
  20. A. C. Spivey, L. Laraia, A. R. Bayly, H. S. Rzepa and A. J. P. White “Stereoselective Synthesis of cis- and trans-2,3-Disubstituted Tetrahydrofurans via Oxonium−Prins Cyclization: Access to the Cordigol Ring System”, Org. Lett., 2010, DOI 10.1021/ol9024259 and interactive table.
  21. J. Kong, P. v. R. Schleyer and H. S. Rzepa, “Successful Computational Modeling of Iso-bornyl Chloride Ion-Pair Mechanisms”, J. Org. Chem., 2010, DOI: 10.1021/jo100920e and interactive table.
  22. A. Smith, H. S. Rzepa, A. White, D. Billen, K. K. Hii, “Delineating Origins of Stereocontrol in Asymmetric Pd-Catalyzed α-Hydroxylation of 1,3-Ketoesters”, J. Org. Chem., 2010, 75, 3085-3096. DOI: 10.1021/jo1002906 and interactive table.
  23. H. S. Rzepa “The rational design of helium bonds”, Nature Chem., 2010, 2, 390-393. DOI: 10.1038/NCHEM.596 and web enhanced table.
  24. P. Rivera-Fuentes, J. Lorenzo Alonso-Gómez, A. G. Petrovic, P. Seiler, F. Santoro, N. Harada, N. Berova, H. S. Rzepa, and F. Diederich, “Enantiomerically Pure Alleno–Acetylenic Macrocycles: Synthesis, Solid-State Structures, Chiroptical Properties, and Electron Localization Function Analysis”, Chem. Eur. J., 2010, DOI: 10.1002/chem.201001087 and interactive figure
  25. H. S. Rzepa, “The Nature of the Carbon-Sulfur bond in the species H-CS-OH”, J. Chem. Theory. Comput., 2010, 49, DOI: 10.1021/ct100470g and interactive table.
  26. H. S. Rzepa, “Can 1,3-dimethylcyclobutadiene and carbon dioxide co-exist inside a supramolecular cavity?”, Chem. Commun., 2010, DOI: 10.1039/C0CC04023A and interactive table
  27. M. R. Crittall, H. S. Rzepa, and D. R. Carbery, “Design, Synthesis, and Evaluation of a Helicenoidal DMAP Lewis Base Catalyst”, Org. Lett., 2011, DOI: 10.1021/ol2001705 and interactive table
  28. H. S. Rzepa, “The past, present and future of Scientific discourse”, J. Cheminformatics, 2011, 3, 46. DOI: 10.1186/1758-2946-3-46 and interactive figure 3, figure 4 and figure 5.
  29. H. S. Rzepa, “A computational evaluation of the evidence for the synthesis of 1,3-dimethylcyclobutadiene in the solid state and aqueous solution”, Chem. Euro. J., 2012, in press.
  30. J. L. Arbour, H. S. Rzepa, L. A. Adrio, E. M. Barreiro, P. G. Pringle and K. K. (Mimi) Hii, “Silver-catalysed enantioselective additions of O-H and N-H to C=C bonds: Non-covalent interactions in stereoselective processes”, Chem. Euro. J., 2012, in press, Web table 1 and Web table 2.
  31. H. S. Rzepa, “Chemical datuments as scientific enablers”, J. Chemoinformatics, submitted.
  32. A. P. Buchard, F. Jutz, F. M. R. Kember, H. S. Rzepa, C. K. Williams, C.K., “Experimental and Computational Investigation of the Mechanism of Carbon Dioxide/Cyclohexene Oxide Copolymerization Using A Dizinc Catalyst”, in press. Interactivity box
  33. D. C. Braddock, D. Roy, D. Lenoir, E. Moore, H. S. Rzepa, J. I-Chia Wu and P. von R. Schleyer, “Verification of Stereospecific Dyotropic Racemisation of Enantiopure d and l-1,2-Dibromo-1,2-diphenylethane in Non-polar Media”, Chem. Comm., 2012, just published. DOI: 10.1039/C2CC33676F and interactivity box.
  34. K. Leszczyńska, K. Abersfelder, M. Majumdar, B. Neumann, H.-G. Stammler, H. S. Rzepa, P. Jutzi and D. Scheschkewitz, “The Cp*Si+ Cation as a Stoichiometric Source of Silicon, Chem. Comm., 2012, 48, 7820-7822. DOI: 10.1039/c2cc33911k. Cites links to 10042/to-13974, 10042/to-13982, 10042/to-13969, 10042/20028, 10042/to-13973, 10042/to-13985
  35. H. S. Rzepa, “A computational evaluation of the evidence for the synthesis of 1,3-dimethylcyclobutadiene in the solid state and aqueous solution”, Chem. Euro. J., 2013, 4932-4937. DOI: 10.1002/chem.201102942 and WebTable
  36. H. S. Rzepa, “Chemical datuments as scientific enablers”, J. Chemoinformatics, 2013, 4, DOI: 10.1186/1758-2946-5-6. The interactivity box is integrated into the body of the article.
  37. M. J. Cowley, V. Huch, H. S. Rzepa, D. Scheschkewitz, “A Silicon Version of the Vinylcarbene – Cyclopropene Equilibrium: Isolation of a Base-Stabilized Disilenyl Silylene”, 2013, Nature Chem., in press and Webtable.
  38. M. J. S. Gomes, L. F. V. Pinto, H. S. Rzepa, S. Prabhakar, A. M. Lobo, “N-Heteroatom Substitution Effects in 3-Aza-Cope Rearrangements”, Chemistry Central, 2013, 7:94. doi:10.1186/1752-153X-7-94 and Table.
  39. H. S. Rzepa and C. Wentrup, “Mechanistic Diversity in Thermal Fragmentation Reactions: a Computational Exploration of CO and CO2 Extrusions from Five-Membered Rings”, J. Org. Chem., DOI: 10.1021/jo401146k and Table.
  40. D. C. Braddock, J. Clarke and H. S. Rzepa “Epoxidation of Bromoallenes Connects Red Algae Metabolites by an Intersecting Bromoallene Oxide – Favorskii Manifold”, Chem. Comm., 2013, DOI: 10.1039/C3CC46720A and Table.
  41. M. J. Fuchter, Ya-Pei Lo and H. S. Rzepa, “Mechanistic and chiroptical studies on the desulfurization of epidithiodioxopiperazines reveal universal retention of configuration at the bridgehead carbon atoms”, J. Org. Chem., 2013, in press. doi: 10.1021/jo401316a and table.

References

  1. H.S. Rzepa, B.J. Whitaker, and M.J. Winter, "Chemical applications of the World-Wide-Web system", Journal of the Chemical Society, Chemical Communications, pp. 1907, 1994. https://doi.org/10.1039/c39940001907

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Towards the ultimate bond!

Monday, August 24th, 2009


The 100th anniversary of G. N. Lewis’ famous electron pair theory of bonding is rapidly approaching in 2016 (DOI: 10.1021/ja02261a002). He set out a theory of bond types ranging from 1-6 electrons. The strongest bond recognized by this theory was the 6-electron triple bond, a good example of which occurs in dinitrogen, N2. In terms of valence electrons, nitrogen has an atomic configuration of 2s2, 2p3. Each atom has five electrons in total, some or all of which in principle could be used for forming bonds. An exploration of this motif across the entire periodic table is presented in part one of this blog.
Elements in Groups 5/15 of the Periodic Table.
Nitrogen is in the main group 15, and the element at the bottom of this group is Bismuth (also with the same atomic configuration). We can then move to the corresponding column of the transition series, this time occupying group 5. The first examplar in this set, Vanadium has an atomic configuration of 3d3, 4s2, again five valence electrons, but now utilizing the d- rather than the p-shell of valence atomic orbitals (AOs). The final forage across the period table would land us with Pr and Pa, which occupy the lanthanide and actinide series respectively, and which have atomic configurations of 4f3, 6s2 and 5f2, 6d1 and 7s2 respectively. You can now see the theme developing; how does the bonding develop between two atoms that between them have ten valence electrons occupying molecular orbitals constructed from s, and then either p, d or f atomic orbitals. The next in that series, g atomic orbitals, are thought unlikely to have any chemical significance in the presently known periodic table.

We are in fact going to study the diatomic molecule comprising two atoms of each element, and further we are going to protonate this species on one of these atoms resulting in the molecule HX2+. So let us start with the two systems HN2+ (Figure 1, 1a-1e) and HBi2+ (Figure 1, 2a-2e).

Figure 1. The molecular orbitals of HN2(+) and HBi2(+)

Figure 1. The molecular orbitals of HN2(+) and HBi2(+)

The two most stable valence molecular orbitals (MOs) for each system (1a and 2a) are the symmetric and antisymmetric combinations of (assumed pure) s-AOs, each populated with two electrons. For Bi (2a/2b) it is fairly clear that this bonding and anti-bonding combination cancel almost exactly. The bond order resulting from these four valence electrons is therefore close to zero. But N is in fact rather odder (and in part, the reason for protonating the systems was to tease out this oddity)! The apparently antibonding 2s-2s combination (1b) actually has electron density along the N-N bond, and the node occurs not along the bond, but at the actual nitrogen atom. So for N, the bonding and anti-bonding σ-bond combinations do not cancel, and the sum of these two may actually lead to a non-zero bond order.

With Bi, the next most stable MO results from the overlap of two 6p-AOs end on (2c); the anti-bonding combination of these AOs is in fact the 2e (with no occupying electrons). Result: bond order of +1, and this bond is called a σ-bond. The final MO shown (2d) is in fact one of a pair of MOs (only one of which is shown here), resulting from parallel overlap of two 6p-AOs. Result: bond order of +2, and we call these two p-π-bonds. The total bond order is +3, comprising 1 σ-bond and two p-π-bonds. No surprises here yet! For N, the relative order of the π- (1c) and the σ-bonds (1d) is swapped compared to Bi (which might be due to relativistic effects, see below), but one can again argue that these three orbitals together contribute 1 σ-bond and two p-π-bonds to the total bonding. So, to summarize, HBi2+ exhibits a classical triple Bi…Bi bond; HN2+ in contrast may actually exceed the bond order of +3, and a case could be made for arguing it is an abnormally strong bond (there is evidence that the N…N stretching frequency in the protonated species is significantly higher than the simple diatomic nitrogen gas).

Let us now move to the combination HV2+ (3) and HTa2+ (4, Figure 2). There are two major differences for V compared to N. Firstly, a 3d rather than 2p -AO is used for the bonding. Secondly, the 3d-AO is actually lower in energy than the 4s-AO, the reverse of the 2s/2p order for eg nitrogen. So what effect does this have on the resulting molecular orbitals?


Figure 2. Valence molecular orbitals for HV2(+) and HTa2(+)

Figure 2. Valence molecular orbitals for HV2(+) and HTa2(+). Click image for 3D of Delta orbital


The difference between N and V turns out to be spectacular, in several regards! The most stable MO (3a) now turns out to be composed of end-on overlaps of two 3d-AOs. There are in fact two of these orbitals (only one is shown in Figure 2), and together they form two 3d-π bonds. The next MO (3b) involves the end-on overlap of a 3dz2-AO, this time forming one 3d-σ-bond. Finally, the 4s-AOs get in on the act, forming now one bonding s-σ-bond (3c). So far, eight of the ten valence electrons have been consumed; two more to go. But now we have a problem. The next MO is formed by parallel overlap of two 3d-AOs (Figure 2, 3d), and in fact there is a pair of these combinations (again, only one is shown in the figure), which are in fact equal (degenerate) in energy. Because they are equal in energy, they must both be populated by an equal number of electrons, but since there are only two valence electrons left, we end up with one electron in each of these two orbitals, resulting in a triplet spin state. Exactly the same phenomenon is responsible for diatomic oxygen also adopting a triplet rather than singlet spin state. This parallel overlap of two 3d-AOs is said to form a 3d-δ-bond. The total bond order in molecule 3 therefore comprises two 3d-π-bonds, one 3d-σ-bond, one 4s-σ and two half 3d-δ-bonds, i.e. five in total and enforces a triplet spin state. So in the sense of formal bond order, the V-V bond in 3 is greater (and perhaps even stronger) than in 1. What a difference from nitrogen!

How about Ta? This has an atomic electronic configuration of 5d3, 6s2. The molecular orbitals are shown in Figure 2 (right). The significant difference in this region of the periodic table is that so-called relativistic effects start to influence the relative ordering of the atomic orbitals. This so-called relativistic contraction impacts upon s-AOs far more than p, or d or f. Thus orbital 4a (Figure 2, right) comprising a symmetric combination of Ta 5s-AOs is more compact than the corresponding V orbital (3c), and relatively more stable. Next come the end-on overlaps of two 5d-AOs (4b), of which there are two (degenerate) combinations (only one is shown in Figure 2). MOs 4c and 4d are again best described as originating from the 6s Ta AO, but with significant contributions from the 5d-AOs (sd-hybrids if you will). Significantly, the relativistic effect means that the δ-bond formed by parallel overlap of two 5d-AOs (4e) is left vacant. Thus the bonding in 4 comprises an 6s-σ-bond, two 5d-π-bonds, and two sd-σ-bonds (one of which does look rather anaemic) but NO 5d-δ-bond! The consequences of the relativistic effect is the relegation of the δ-bond to unoccupied status and the formation of a singlet rather than a triplet spin ground state.

Given the big differences in bonding which occur upon changing a 2p-valence atomic orbital to 3-5d-AOs, one wonders what will happen with 4-5f-AOs? A π-bond can be formed by the parallel overlap of two p-AOs, or end-on overlap of two d-AOs. A δ-bond can be formed by the parallel overlap of two d-AOs. Could then a δ-bond also be formed by the end-on overlap of two f-AOs? Would a φ-bond be formed by the parallel overlap of two f-AOs? The latter might look something like that shown in Figure 3, shown in two different orientations (these diagrams were obtained by inspecting the unfilled MOs in the V/Ta examples shown here; notice that although the φ bond appears to be bonding, the system has chosen not to occupy it with any electrons!).

Figure 3. The "phi" bond.

Figure 3. Two views of a φ bonding MO. Click image for 3D
Figure 3. The "phi" antibond.

Figure 4. A φ* antibonding MO. Click image for 3D

This blog will end by posing the question “can any molecule be devised which supports one or more φ-bonds”, or will the relativistic contraction always scupper such efforts by depriving such bonds of electrons (e.g. Ta above)? Are the systems reported in DOI: 10.1021/ja067281g examples of a 5f-φ bond for uranium (the claim is made in a very low key manner)? This will be investigated in the follow up to this post!

(For serious geeks/computational chemists only, N was computed at the B3LYP/6-31G(d) level, V at the ROHF/6-31G(d) level, and Bi/Ta at a triple-ζ-pseudopotential level (which incorporates some of the relativistic effects).

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Molecular toys: Tetrahedral cavities

Saturday, July 4th, 2009


An earlier post described how a (spherical) halide anion fitted snugly into a cavity generated by the simple molecule propanone, itself assembled by sodium cations coordinating to the oxygen. A recent elaboration of this theme, reminiscent of the children’s toys where objects have to be fitted into the only cavity that matches their shape, Nitschke and co-workers report the creation of a molecule with a tetrahedral rather than a spherical cavity (DOI: 10.1126/science.1175313 ), into which another but much smaller tetrahedral molecule is fitted.  The small molecule is P4, in which each of the three valencies of the P atom is directed to a corner of the tetrahedron. The large molecule  comprises four Fe atoms. These are each octahedrally coordinated with six ligand sites, three of which mimic the P atoms in also being directed towards the remaining three vertices of a tetrahedron.

P4 inside a Tetrahedral cavity.

Needless to say, the properties of the P4 molecule when entrained into this larger container are nothing like that of the free molecule. Now it is quite inert, but this is due purely to the snug fit. For example, the normal reaction of this molecule is to oxidize in air. But such oxidation would now produce a molecule too big to fit into the cavity. Hence no reaction!

So, now the search is on for a cubic container to include a cubic molecule!

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Longer is stronger.

Saturday, June 6th, 2009

The iconic diagram below represents a cornerstone of organic chemistry. Generations of chemists have learnt early on in their studies of the subject that these two representations of where the electron pairs in benzene might be located (formally called electronic resonance or valence bond forms) each contribute ~50% to the overall wavefunction, and that the real electronic description is in effect an average of these two (that is the implied meaning of the double headed arrow). This means that the six C-C bonds in benzene must all be of equal length. The diagrams, everyone knows, do not mean that benzene has three short and three long C-C bonds.

The Kekule structures of benzene.

The Kekulé structures of benzene. Click for 3D.

The diagram has much other implied semantics. Thus there is no explicit three dimensional information; the molecule looks (and is) flat, and it is tempting to conclude that the electrons are flat and two dimensional as well. Indeed, up to around 1930 (some 105 years after its first discovery), the electrons in benzene were always represented as all lying in the plane of the molecule. This changed when Hückel announced the principle of σ/π separation. These were the labels he gave to two different symmetries of electrons (actually derived for ethene), one set which did genuinely occupy the plane of the molecule, and a second (π) set for which this plane represented a node (a region of zero probability for the electron density). The π electrons could instead be regarded as occupying the space above and below that plane. Hückel went on to develop a quantum mechanical theory for benzene based purely on those π-electrons, of which there are six. This (now called Hückel) theory predicted that the averaged structure noted above emerged naturally, along with another concept known as π-electron resonance energy. This is the difference in energy between the symmetric form of benzene and a structure in which the six π electrons do not interact as a whole, but which are localized into three pairs located in the regions of the double bonds. Most people interpret this latter as being equivalent to the two Kekulé forms shown above. Symmetrizing the structure (from D3h to the higher D6h symmetry) is accompanied by reducing the π-energy of the system by that resonance term (often estimated as around -152 kJ/mol of stabilization). For benzene in other words, this is the difference in energy between the symmetric species and a (hypothetical) bond localized cyclohexatriene.

With such a focus on the π-electrons, it seemed natural to accept that the reason why benzene has six equal C-C lengths is because of the resonance energy gained by the π-electrons when adopting the six-fold symmetric form. Prior to around 1961, no-one would have dissented from that point of view. The first to do so was Berry (see DOI: 10.1063/1.1732256 ), but his was a lone voice at that time. But mysterious and inexplicable observations started to come to light. Perhaps the most direct was a study of the excited state of benzene, in which one π-electron is promoted from a bonding to a higher energy and antibonding π-orbital (known as a π-π* excitation, see DOI 10.1063/1.435193). A schematic illustration of this process is shown below.

The Hückel Molecular orbital picture for benzene

The Hückel Molecular orbital picture for ground and excited states of benzene

Diagram (a) shows the normal population of electrons in the (three lowest) energy levels derived using Hückel’s theory. Diagram (b) shows how this changes in the first excited singlet state, which would be expected to have weaker π-bonds. The vibrational spectrum of a molecule is one way of measuring how strong the bonds in a molecule are. Berry had already implied that one particular vibrational mode, the so-called Kekulé mode (also known as the b2u mode using group theory) seemed unusually low in frequency. In other words, this distorsion was easier than it should have been, and this Berry attributed to the (then almost heretical view) that the π-electrons did not in fact promote a hexagonal form of benzene. This was instead induced by the σ electrons, which occupy the plane of the molecule. This effect prevailed over the π-electrons, which were in fact trying to get benzene to adopt a bond-alternating geometry (managing instead only to lower the energy of the b2u mode). When the vibrational spectrum of the excited state of benzene was analyzed in 1977, it appeared to spectacularly vindicate Berry (DOI 10.1063/1.435193). The Kekulé mode has a value of 1309 cm-1 for the normal ground state of benzene, but an exalted value of 1570 cm-1 in the excited state. This means that as the bonding due to the π-electrons is weakened by placing one of them in an antibonding orbital, their overall ability to distort the geometry is also weakened. As a result, the resistance to such distorsion (the Kekulé mode) is in turn strengthened by an amount corresponding to +261 cm-1. It was evidence such as this, and much else besides that Shaik and his co-workers used to promote the idea of π-distortivity in benzene (DOI: 10.1021/cr990363l). Despite such advocacy, the idea that all the six bonds in benzene are equal despite rather than because of the π-electrons is still rarely taught in introductory organic chemistry.

But the story of excited benzene is not yet quite finished! In 2006, Blancafort and Sola (DOI: 10.1021/jp064885y) reminded us that the (1B2u) excited state of benzene exhibits a type of geometrical distorsion known as pseudo Jahn-Teller (PJT), the origins of which have nothing to do with any of the previous arguments. The effect instead arises because the promoted electron emerges from a so-called energy degenerate orbital, and jumps into another degenerate orbital (Figure b above). The exaltation of the b2 vibrational mode is in fact strongly coupled with this PJT effect, which complicates disentangling the two effects (PJT and π-distortivity).

So another excited state is here proposed which is not susceptible to the PJT effect. Figure (c) above shows a π-quintet state in which two electrons are both promoted to anti-bonding orbitals. Now the π-electron bonding has been well and truly weakened! When the vibrational modes are calculated for the (D6h-symmetric) geometry of benzene at the same level of theory (B3LYP/aug-cc-pvtz) for both singlet ground state and quintet excited state one finds that the b2u vibrational mode has the value of 1332 cm-1 for the former and 1524 cm-1 for the latter. Significantly, the former mode shows a contribution to the motion from the hydrogen atoms. These, being light, tend to increase the wavenumber of the vibrational mode. The same mode in the quintet state however shows motion of the carbon atoms only (Click on the diagram below to view the b2u mode for the quintet state of benzene, and note how little motion of the hydrogen atoms there is). It is a pure Kekulé mode, whereas that for the ground state is not! If the motion of the hydrogens in the ground state of benzene is suppressed by artificially changing the atomic weight of the hydrogen in the mass-weighting scheme to a large value, the calculated b2u vibrational mode drops to around 1317 cm-1. This means the quintet state mode of benzene is exalted by 207 cm-1, and being PJT-free, it is a truer reflection of the effect of the π-electrons. Thus the effect first speculated upon by Berry, and championed by Shaik is spectacularly vindicated (again!).

The b2u modes in benzene

The b2u modes in benzene for (a) ground state and (b) quintet state. Click for 3D.

But what of the title for this post? Well, the C-C length in the singlet ground state of benzene is 1.391Å. In the quintet state, it becomes longer at 1.454Å (which is almost exactly the value that Berry originally suggested should be used for the hypothetical cyclohexatriene geometry). Despite this lengthening, the Kekulé mode clearly gets stronger. Why is this noteworthy? Well, it is almost always assumed that if a bond is shorter, it means stronger. In this case, we have an example of six bonds each getting shorter and weaker (at least as measured by the b2u mode of vibration), or as the title states, longer and stronger in the quintet state of benzene. Oh, and what about that π-resonance energy which we started with? Does it play no role after in the symmetric structure of benzene? Well, in fact it does! The answer is that the π-resonance energy is still at its maximum stabilization at the hexagonal structure of benzene, but it is the total π-energy that achieves its maximum stability at the non-hexagonal structure. These two energies are quite different beasts, and they each prefer a different geometry!

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The mystery of the Finkelstein reaction

Saturday, May 16th, 2009

This story starts with an organic chemistry tutorial, when a student asked for clarification of the  Finkelstein reaction. This is a simple SN2 type displacement of an alkyl chloride or bromide, using sodium iodide in acetone solution, and resulting in an alkyl iodide. What was the driving force for this reaction he asked? It seemed as if the relatively strong carbon-chlorine bond was being replaced with a rather weaker carbon-iodine bond. But its difficult to compare bond strengths of discrete covalent molecules with energies of ionic lattices. Was a simple explanation even possible?

All is not as it seems however. The traditional explanation, found by the quick Google search linked above, is that the reaction illustrates Le Chatelier’s principle, whereby an equilibrium is driven over to completion by removal of one of the products (in this case sodium chloride or sodium bromide, which crystallize out of solution). Well, we have replaced one possible (and probably complicated) explanation based on bond strengths and ionic lattices by another based on the solubilities of an ionic material in a moderately polar solvent. But all we have done is ask a different question, which now becomes why is sodium iodide highly soluble in acetone, whereas sodium chloride and bromide are not? The answer to this is less easily found using Google!

A good start would be the crystal structure of any complex formed between acetone and sodium iodide. Fortunately, one such does exist, and it is shown below (sodium=yellow, iodine=purple).

(Acetone)3. NaI

(Acetone)3. NaI. Click for 3D.

The formula shows three acetone molecules for each sodium iodide. The carbonyl oxygen has two lone pairs of electrons, and each of these is used to coordinate a (different) sodium cation. This allows each sodium to be coordinated by a total of six lone pairs, giving it octahedral coordination. This sets up what in fact is quite a rigid scaffold, with the unusual feature of an approximately triangular shaped channel running down the lattice (two such are shown above). The size of this hole is determined by the methyl groups of the acetone, and it is into this cavity that the halide ion must fit.

As it happens, the iodide anion is exactly the right size to produce a perfectly snug fit up against those methyl groups (click on the image above to view this). If a chloride or bromide anion were to be fitted into the cavity, there would be empty space surrounding it. The cavity itself is too rigid to collapse around the halide anion to absorb this space. This means these halide anions are further away from the positively charged sodium than they would like to be such that they minimize their ionic lattice energies. Instead they avoid fraternizing with the acetone at all, and form a pure sodium chloride or sodium bromide lattice (where the two oppositely charged ions CAN approach at optimal distances). The result is that sodium chloride crystallizes out of solution, and the Finkelstein reaction proceeds to completion!

Acetone. NaI in spacefill mode

Acetone. NaI in spacefill mode. Click for 3D.

But that is not quite the end of the story. If you view the acetone.NaI lattice sideways (click on the diagram above to view this aspect), you will find that in fact there is still space in the scaffold after all! Each iodide anion has room above or below it, with space for exactly one more iodine atom to fit without having to change the shape of the scaffold. And indeed such a molecule has been reported, but it is an odd one! The stoichiometry is now (acetone)3.NaI2, which implies that the iodide anion has been joined by an iodine atom. I2(-) is called a radical anion, and as such has an unpaired electron. Just like two iodine atoms can couple their unpaired electrons to form a covalent bond, so can two I2 radical anions, forming I42- [or I3–.I–] or on to infinity as a linear iodine polymer, of formula n[I42-], with all the I…I distances equal at 3.224Å (a system with no Peierls distortion).  Straight rod-like polymeric chains of a single element might appear highly unusual, but curiously, another class of elements that exhibits this behaviour is Cu/Ag/Au and Ga (DOI: 10.1002/anie.200601726, the ultimate in thin wires!).

Acetone. NaI2

Acetone. NaI2. Click for 3D.

Finally, it is worth noting that the same phenomenon occurs with the dimethylformamide.NaI complex. In this example, only the NaI and not the NaI2complex has been reported.

DMF. NaI

DMF. NaI. Click for 3D.

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The Chirality of Lemniscular Octaphyrins

Tuesday, April 28th, 2009

In the previous post,  it was noted that  Möbius annulenes are intrinsically chiral, and should therefore in principle be capable of resolution into enantiomers. The synthesis of such an annulene by Herges and co-workers was a racemic one; no attempt was reported at any resolution into such enantiomers. Here theory can help, since calculating the optical rotation [α]D is nowadays a relatively reliable process for rigid molecules. The rotation (in °) calculated for that Möbius annulene was relatively large compared to that normally measured for most small molecules.

Recently, quite a number of cyclopolypyrroles, more commonly called phyrins, have been reported. The conventional number of pyrrole rings in many biological systems is of course four (chlorophyll, haemoglobin, etc), but these extended porphyrins can have anywhere between five and sixteen such rings comprising a larger macrocycle. For those with six-eight such rings, a commonly adopted geometric motif is found to be a figure-eight, or more properly a lemniscular one. Such shapes have recently (10.1021/ol703129z) been recognized as also being Möbius systems, albeit this time with two half twists in the π-electron cycle rather than just the single twist synthesized by Herges. As such, they also follow a simple electronic selection rule, being aromatic if 4n+2 π-electrons circulate around the ring.

One such molecule is shown below (10.1039/b502327k), albeit with four of the pyrroles replaced by a thiophene ring.

A 34-Octaphyrin. Click to see molecule

A 34-Octaphyrin. Click for 3D.

Just as with the Herges syntheses, most of these phyrins are also made as racemates. There appears to be only one report of such octaphyrin actually being separated into enantiomers (10.1002/(SICI)1521-3773(19991216)38:24<3650::AID-ANIE3650>3.0.CO;2-F) but no optical rotation could be measured due to its intense colour (in other words, so much light is absorbed by the system that too little remains to measure its rotation). So no-one knows what the magnitudes of the optical rotation values for these figure-eight or lemniscular molecules actually are.

Here again, theory can come to the rescue. The octaphyrin shown above for example (simplified such that Ar=H), [α]D has the stupendous value of -25517° (See 10042/to-2185). Values above 10,000 are common for this type of molecule! So these relatively small and simple class of molecules are currently easily the record holders for the size of their optical rotations. OK, the latter are merely predictions, but it certainly should serve as an encouragement for experimental measurements of this property.

Oh, by the way, if you click on the graphic above, you will get to see a molecular orbital calculated for the molecule. It is the most stable of the π-type of MOs, and shows the characteristic features of the lemniscate, namely the π-electrons take the form of a torus link (10.1039/b810301a).

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The chirality of Möbius annulenes

Wednesday, April 22nd, 2009

Much like climbing Mt. Everest because its there,  some hypothetical molecules are just too tantalizing for chemists to resist attempting a synthesis. Thus in 1964, Edgar Heilbronner  speculated on whether a conjugated annulene ring might be twistable into a  Möbius strip. It was essentially a fun thing to try to do, rather than the effort being based on some anticipated  (and useful) property it might have. If you read the original article (rumour has it the idea arose during a lunchtime conversation, and the manuscript was completed by the next day), you will notice one aspect of these molecules that is curious by its absence. There is no mention (10.1016/S0040-4039(01)89474-0) that such Möbius systems will be chiral. By their nature, they have only axes of symmetry, and no planes of symmetry, and such molecules therefore cannot be superimposed upon their mirror image; as is required of a chiral system (for a discussion of the origins and etymology of the term, see 10.1002/chir.20699).

The 16-annulene synthesized by Herges and his team.

The 16-annulene synthesized by Herges and his team. Click for 3D.

Heilbronner’s little vignette had little overt effect on the synthetic community until around 2003, when Rainer Herges announced that a crystalline annulene following this recipe had been rationally synthesized (10.1038/nature02224). This time, the chemical community really sat up and took notice. The synthesis was hailed as a major achievement, ranking (chemically) as similar to climbing Everest. But if you read Herges’ article carefully, yet again you will note the absence of any discussion of the chirality of their molecule. Their synthesis was of course racemic, in other words an equal proportion of both enantiomers was made. Indeed, it is not obvious how a non-racemic synthesis could be carried out, although resolution of the product might be an easier task. So in the absence of any pure enantiomer of this molecule, can one speculate on its chiral properties? One obvious such property is the optical rotation, and in particular the [α]D value in chloroform. Most optically pure molecules with molecular weights of < 500 Daltons  tend to have rotations also < 500°. Few molecules have values > 1000°. Now it should be said at the outset that a molecule with a large optical rotation is not more chiral than a molecule with a smaller value; indeed it seems generally agreed that the question “how chiral is this molecule” is either fairly, or even completely meaningless. But it seems a useful task of having a value to hand which is at least approximately accurate, so that some idea of whether any attempted resolution of the enantiomers has produced optically pure product or not. Fortunately, in the last decade or so, computing a value for [α]D has been entirely viable using the standard programs (see 10.1002/chir.20466 and 10.1021/jo070806i for a discussion). This is also useful for two reasons:

 

  1. If the magnitude of the rotation is > 100°, then the sign of this rotation can be very reliably matched to either enantiomer. This allows the absolute configuration to be assigned with a lot of confidence, and probably much more easily than trying to do it by other methods.
  2. The magnitude itself can be reliably predicted to within 10% of the true value if the molecule is conformationally rigid. However, if it has any rotatable groups (and that even includes e.g. OH groups), then the result can be enormously sensitive to that conformation (or Boltzmann mixture of conformations). Put the other way, calculating the optical rotation could be regarded as a very sensitive way of determining conformations!

So what of the 16-annulene synthesized by Herges and co-workers. Well at the B3LYP/6-311G(2df,2pd) and SCRF(CPCM,solvent=chloroform) level of theory (which is reasonably accurate, although one can do better of course), the enantiomer shown by clicking on the graphic above is predicted to have a rotation of -1355° (for the digital repository entry for the calculation, see 10042/to-2176). That is indeed a large value for such a relatively small molecule, and is probably more reliable because of the lack of conformational ambiguity. Well, you saw the prediction here! Anyone up for testing it experimentally?

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A molecule with an identity crisis: Aromatic or anti-aromatic?

Monday, April 13th, 2009

In 1988, Wilke (DOI: 10.1002/anie.198801851) reported molecule 1

A [24] annulene. Click on image for model.

A 24-annulene. Click for 3D.


It was a highly unexpected outcome of a nickel-catalyzed reaction and was described as a 24-annulene with an unusual 3D shape. Little attention has been paid to this molecule since its original report, but the focus has now returned! The reason is that a 24- annulene belongs formally to a class of molecule with 4n (n=6) π-electrons, and which makes it antiaromatic according to the (extended) Hückel rule. This is a select class of molecule, of which the first two members are cyclobutadiene and cyclo-octatetraene. The first of these is exceptionally reactive and unstable and is the archetypal anti-aromatic molecule. The second is not actually unstable, but it is reactive and conventional wisdom has it that it avoids the undesirable antiaromaticity by adopting a highly non-planar tub shape and hence instead adopts reactive non-aromaticity. Both these examples have localized double bonds, a great contrast with the molecule which sandwiches them, cyclo-hexatriene (i.e. benzene). The reason for the resurgent interest is that a number of crystalline, apparently stable, antiaromatic molecules have recently been discovered, and ostensibly, molecule 1 belongs to this select class!

So is 1 actually anti-aromatic? Let us look at some of the ways in which this might be estimated.

  1. One can inspect the bond lengths, measured from X-ray analysis. The longest is 1.463Å, labelled a above, and it corresponds to a single bond (value from the crystal structure).
  2. If the molecule had a bond alternating structure, the adjacent bonds would be expected to be much shorter, in the region of 1.32Å. In fact, they are rather longer, at 1.37Å. Indeed, in the cycloheptatriene part of the molecule, the alternation is much less than one might expect of an anti-aromatic molecule, oscillating between 1.37 and 1.43Å.
  3. One can also inspect aromaticity via a variety of magnetic indices. The simplest of these is the NICS probe. Placed at a ring centroid, a negative value of this index of around -10ppm indicates aromaticity (this is the value for benzene), whilst a strongly positive value (of up to +20 ppm) indicates anti-aromaticity. Molecule 1 has two potential centroids, one placed at the absolute centre of the system, and one placed at the centroid of the ~6-membered ring completed using bond b (in reality, the centroids were computed from the positions of ring critical points obtained from an AIM analysis). The NICS values at these two positions are both ~-4.4 ppm (See DOI: 10042/to-2156 for details of the calculation). These values does not indicate antiaromaticity! They could even be described as mildly aromatic. So what is going on?
  4. What about the chemical shifts of the other protons? All the hydrogens attached to sp2 carbons are predicted to resonate at around 6.7ppm (unfortunately Wilke does not report the experimental spectrum), which is typical of an aromatic system (anti-aromatic systems have high upfield shifts for such protons, at around 2ppm or even -2 ppm, see DOI: 10.1021/ol703129z for examples). The two protons of the methylene bridges are also quite different; 2.9 and 0.8 ppm. The latter is the proton endo to the cycloheptatrienyl ring, and is typical of a proton placed in the anisotropic magnetic shielding region of e.g. benzene. Thus the cycloheptatrienyl ring is itself behaving as if it were aromatic, whereas the overarching 24-annulene ring is certainly not behaving as if it were antiaromatic.

One possible explanation involves a concept known as homoaromaticity. The bond marked as b could be regarded as completing the 6π-electron local aromaticity of that ring (it would be formally considered as a 1π-electron bond, with no underlying σ-framework, see 10.1021/ct8001915 for further detail). Well characterised examples of such neutral homoaromatics are in fact very rare indeed; the phenomenon is thought to manifest mostly through cationic homoaromatics (i.e. the homotropylium cation, see 10.1021/jo801022b for discussion). So has the case been made for 1 being the first clear cut example of a neutral homoaromatic molecule, containing no less than four rings exhibiting this type of aromaticity?

There is one further concept that can be introduced. Clar (for a discussion, see DOI: 10.1021/cr0300946) proposed that benzenoid 6π-electron local aromaticity is preferred to less local or more extended cyclic conjugations, if the two compete. Many examples in a type of compound known as polybenzenoid aromatics are known where the most favourable resonance structure is that which maximises the number of Clar rings. More recently, quite a few ostensibly antiaromatic molecules have been shown to attenuate this unfavourable effect by forming instead groups of aromatic Clar islands containing delocalized benzene like rings (discussion of this point can be found at DOI: 10.1039/b810147g). In molecule 1, we could have a new phenomenon; a homoClar ring, formed to avoid antiaromaticity.

For further discussion, see the comment posted to Steve Bachrach’s blog.

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Conformational analysis and enzyme activity: models for amide hydrolysis.

Sunday, April 12th, 2009

The diagram below summarizes an interesting result recently reported by Hanson and co-workers (DOI: 10.1021/jo800706y. At ~neutral pH, compound 13 hydrolyses with a half life of 21 minutes, whereas 14 takes 840 minutes. Understanding this difference in reactivity may allow us to understand why some enzymes can catalyze the hydrolysis of peptides with an acceleration of up to twelve orders of magnitude.

Models for peptide cleavage.

Models for peptide cleavage.

The secret to understanding this behaviour lies in a technique known as conformational analysis, for which Derek Barton was awarded a Nobel prize. Indeed, the very molecules for which he first developed his technique were the decalins, of which molecule  13 is an example of a cis-decalin and 14 a trans-decalin. Barton’s insight was to recognize that both types of ring prefer to exist in chair conformations rather than the alternative boat shape.

The technique pioneered by Barton for estimating the energies of these various conformations is called Molecular Mechanics, and can be used to explain the difference in reactivity. Considering first molecule 13, one can calculate its molecular mechanics energy for two conformations, differing in whether the N-alkyl sidechain is equatorial (left) or axial (right).

Cis amide

Cis amide. Click for Equatorial 3D.

The equatorial form (green box) comes out about 5 kcal/mol lower in energy than the axial (red box). One can also calculate the energy of the product, which arises from the OH attacking the carbon of the amide (dashed lines), evicting ammonia, and forming a cyclic lactone. Here, the most stable product (by ~10 kcal/mol) is again that resulting from the green bond forming. From the simple relationship ΔG = -RT Ln K (where K describes the position of the equatorial/axial equilibrium), one can conclude that the ratio equatorial/axialis ~4000, i.e. the favoured reaction arises from the most abundant reactant.

Trans amide

Trans amide. Click for 3D.

With the trans amide, the equatorial conformation (green box) is around 3 kcal/mol lower than the axial (red box), but now the most stable lactone product (by ~ 3 kcal/mol) arises (green bond) from the less stableaxial reactant. For reaction to occur, the equatorial reactant has to first isomerise to the axial, which imposes a ~3 kcal/mol penalty on the reaction. This is enough to slow the rate of the reaction significantly compared to the un-penalised cis-decalin reaction.

 

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How do molecules interact with each other?

Sunday, April 12th, 2009

Understanding how molecules interact (bind) with each other when in close proximity is essential in all areas of chemistry. One specific example of this need is for the molecule shown below.

The Pirkle reagent

The Pirkle reagent

This is the so-called Pirkle Reagent and is much used to help resolve the two enantiomers of a racemic mixture, particularly drug molecules. The reagent binds to each enantiomer of a racemic drug differently, and this difference can be exploited by using e.g. column chromatography to separate the two forms. The conventional wisdom is that such chiral recognition occurs via a three-point binding model. In other words, at least three different interactions must occur between the Pirkle reagent and the drug to allow such chiral recognition.

So how do we identify where these bindings might occur? A good place to start is to look at the self-binding of the Pirkle reagent, in other words, how does it interact with itself in the crystal state? An X-ray structure of the pure enantiomer of the Pirkle reagent shows that it binds with itself to form a loose dimer. We are now in a position to analyze exactly how this binding occurs. To do this, we are going to invoke a technique known as Atoms-in-molecules or AIM. This effectively looks at the curvature of the electron density in the dimer, and from the characteristics of this curvature, identifies a series of so called critical points, or regions where the first derivative of the electron density (referred to as ρ(r) ) with respect to the geometry is zero. These critical points come in four varieties only;

  1. A nuclear critical point, which almost exactly corresponds to where the nuclei are
  2. A bond critical point, which is the key to understanding not only where actual bonds are in the molecule, but also a range of weaker interactions which are conventionally not graced with the term bond, but which nevertheless can be essential in understanding how to molecules interact weakly with each other.
  3. The remaining two types of critical point relate to rings and cages, and we will not be concerned further with them here.

The electron density required for this analysis could in principle come from the X-ray measurements themselves, but it is not easy to acquire this to the required accuracy (although it can be done). In this case, it is easier (and probably no less accurate) to calculate the density from a DFT-based quantum mechanical calculation. The result of this is shown below.

Pirkle dimer. Click on image to obtain model

Pirkle dimer. Click for 3D.


The light blue spheres show the position of selected bond critical points or BCPs in the AIM analysis. So what do they tell us about how two molecules of Pirkle molecule interact with each other? Three different points labelled 1-3 are highlighted for discussion.

  1. Points 1 connect the hydrogen of the OH group with the carbons of the π-face of the anthracene ring (the left ring of the molecule as shown above). This is an unusual type of interaction known as a π-facial hydrogen bond, and it has only been recognized as such in the last 30 years. Note that this interaction is not restricted to occur just between a pair of atoms, but can involve more (in this case almost a whole benzene ring). By finding the value of the electron density ρ(r) at this BCP, one can estimate the energy of interaction resulting from its formation. In this case, ρ(r) ~ 0.014 au, and comparison with other types of hydrogen bond suggests that this value corresponds to an interaction energy of around 2.5 kcal/mol. This is a little weaker than a conventional OH…O hydrogen bond, but is still quite significant. Two of these interactions occur in this Pirkle dimer.
  2. Points 2 are equally unexpected. They connect the oxygen of the same OH group involved in the previous interaction, and one of the ring C-H groups. Again, that C-H…O groups can interact has only been recognized relatively recently. The value of ρ(r) of ~ 0.018 indicates a hydrogen bond strength of ~3 kcal/mol, again hardly insignificant.
  3. There are four specific interactions of the final type 3. These occur in the region of overlap of the two anthracene rings, and these are referred to as π-π stacking interactions. Again, the ρ(r) of ~ 0.005, calibrated against known systems, suggests that each is individually worth around 1 kcal/mol.

So adding up all eight interactions indicates that the two molecules of the Pirkle reagent have an interaction energy of around 15 kcal/mol resulting just from these weak bonds (there are other types of interactions between two molecules known as dispersion forces, which also contribute), and which together provide more than enough free energy to overcome the entropy required to bring the two molecules together.

Armed with tools such as AIM, one can now be more confident in analyzing the various terms that contribute to two molecules interacting with each other, and in the case of chiral molecules, how these interactions may result in chiral recognitions.

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