Interesting chemistry « Henry Rzepa's blog

Posts Tagged ‘Interesting chemistry’

Anatomy of an arrow-pushing tutorial: reducing a carboxylic acid.

Wednesday, December 1st, 2010

Arrow pushing (why never pulling?) is a technique learnt by all students of organic chemistry (inorganic chemistry seems exempt!). The rules are easily learnt (supposedly) and it can be used across a broad spectrum of mechanism. But, as one both becomes more experienced, and in time teaches the techniques oneself as a tutor, its subtle and nuanced character starts to dawn. An example of such a mechanism is illustrated below, and in this post I attempt to tease out some of these nuances.

The example chosen is the reduction of a carboxylic acid to an alcohol by borane (diborane). Lecture notes present this reaction as being specific to carboxylic acids, even in the presence of carboxylic esters. The tutor is then faced with how to explain this selectivity to students in a tutorial, using arrow pushing.

Scheme for reduction of a carboxylic acid by borane.

I start with grouping the arrow pushing into three sets:

The essential arrows (red). These will attempt to describe the key mechanistic step for which an answer is sought, in this example why the reaction is so selective for the carboxylic acid. A cruder, but perhaps pragmatic description is that these are the arrows needed to pass examinations in the subject (music to students’ ears).
The lazy arrows (blue). In this case, these arrows are essential to “prep the patient”, but they will not of themselves carry much insight into the operation of the mechanism.
The workup arrows (green). To continue the medical analogy, this is a post operative “closing the patient up” stage.

This tutorial actually starts with non-arrows. Process 1 involves converting the actual real structure of diborane (a bridged dimer) into its monomer, which is thought to be the active ingredient of this reagent. Because the bridging hydrogens are bound by three-centre-two-electron bonds, it is actually difficult to represent this process with conventional (two-centre-two-electron) arrows. So we do not even try!

Process 3 is the one which involves the essential (red) arrow pushing. It encapsulates the reason why only a carboxylic acid reacts in this process, and these arrows can be formalised by computing the transition state quantum mechanically (below). In fact there are two ways of illustrating this essential process. Process 2 involves first forming a O-B bond before the essential arrows. Process 4 involves another B-O bond AFTER the key transition state; the outcome of either process is identical. This illustrates another subtle behaviour in arrow pushing; the detailed timing or choreography of the arrows. In this example, the animated form of the reaction coordinate indicates relatively little B-O bond formation, so we will go with 2 and then 5 as the more realistic representation. In fact, the QM transition state is fascinating in its own right; note for example how one of the two extruding hydrogen atoms is moving far less (the hydridic one) than the other (the proton-like one; full details available at 10042/to-5725).

Key transition state? Click for 3D.

Process 6 is a lazy category. The preceding steps are simply repeated twice more to form a triacyloxyborane. There are many other forms of lazy arrow. Proton transfers are often thought of in this category, and double headed arrows involved in addition/elimination to e.g. carbonyl groups.

Process 7 and 8 are in an awkward category. Of themselves, they do not explain the selectivity of borane for this functional group, but they do represent another essential operation; namely the actual reduction of the carbonyl group. They are also somewhat speculative, and it is quite possible other routes could be devised.

Finally, with 9, we arrive at a resting phase which now requires workup (green). Thus 10 and 11 represent hydrolysis of the borate esters to the reduced alcohol and something starting to resemble boric acid. Clearly, more arrows are needed after 11, but few tutors (or examination graders) would begrudge a student if these were to be omitted. Step 11 also contains some lazy arrows, since a proton is transferred between oxygen atoms, but no arrows are shown for this process.

Clearly, there are plenty of nuances here, and it is perfectly possible that other arrow-pushers may even disagree with some of the ones I have shown above. But perhaps the above analysis might give you some ideas of your own on how to communicate the essential of reaction mechanisms to others.

Tags:acyloxyborane, arrow pushing, Interesting chemistry, post operative, tutor
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Gravitational fields and asymmetric synthesis

Saturday, November 20th, 2010

Our understanding of science mostly advances in small incremental and nuanced steps (which can nevertheless be controversial) but sometimes the steps can be much larger jumps into the unknown, and hence potentially more controversial as well. More accurately, it might be e.g. relatively unexplored territory for say a chemist, but more familiar stomping ground for say a physicist. Take the area of asymmetric synthesis, which synthetic chemists would like to feel they understand. But combine this with gravity, which is outside of their normal comfort zone, albeit one we presume is understood by physicists. Around 1980, chemists took such a large jump by combining the two, in an article spectacularly entitled Asymmetric synthesis in a confined vortex; Gravitational fields and asymmetric synthesis (DOI: 10.1021/ja00521a068). Their experiment was actually quite simple. They treated isophorone (a molecule with a plane of symmetry and hence achiral) with hydrogen peroxide and then measured the optical rotation.

Asymmetric synthesis of Isophorone oxide

Conventional wisdom is that the oxygen can be delivered with equal probability to either face of the alkene, resulting in a racemic (equal) mixture of the two enantiomers of the epoxide. But if one enantiomer is formed in slightly greater amount than the other, the reaction is said to proceed asymmetrically, and the product will exhibit an optical rotation. Normally, such asymmetry is induced by carrying out the reaction in the presence of a chiral molecule or catalyst. Light too can be chiral, but these brave chemists decided to use gravity. More specifically, the earth’s gravitational field. In the Northern Hemisphere. The reaction was conducted in a centrifuge in three ways. With the spun tube horizontally, and then vertically spinning clockwise or anticlockwise. The first of these produced product which exhibited no optical rotation within experimental error (e.g +0.2 ± 0.3 mdeg). The second gave results with a positive rotation (e.g. 12.8 ± 0.3 mdeg) and the third a negative rotation (-2.2 ± 0.2 mdeg). They considered that the reaction was occurring in e.g. a clockwise vortex constituting a P-helix (the vortex in other words was chiral) interacting with the earth’s coriolis force. They speculated (but did not do the experiment) that the reverse effect would be seen in the Southern Hemisphere. Their paper concluded with the grand speculation that prebiotic organic synthesis could have been partially asymmetric as a result of being conducted in a chiral gravitational field (nothing like aiming high!).

Shortly after this was published, a rebuttal appeared (DOI: 10.1021/ja00544a051) penned not by a synthetic chemist but a physicist. In truth, most of the four-paragraph article presents arguments few chemists are familiar with (and probably do not understand). Only one sentence, the very last, made the impact (and it sounds as if it was added as a throw-away afterthought). It simply stated that the magnitude of the earth’s field (in so-called natural units) suggested that a parameter ω related to the spinning velocity was ~10^-21 and that the corresponding value that would be required to induce the asymmetry observed (which can be computed from e.g. ΔG = -RT Ln K) was 10^-14. Put in rather plainer English, the earth’s magnetic field was seven orders of magnitude too weak to have the effect claimed for it. That last sentence on its own pretty much sunk the theory, and it is no longer thought that gravitational fields can induce asymmetry in reactions. I tell this story (which, as it happened thirty years ago is now largely forgotten), since seven orders of magnitude is quite a large mismatch! Chemists rarely have the opportunity to be so spectacularly wrong when they propose a theory. In reality, even two orders of magnitude is unusual.

It’s when you approach factors of say less than one order of magnitude (the nuances) that arguments about which interpretation is correct may break out. So which category might the subject of this post belong to? As I there noted, it’s all about whether the two carbon atoms separating carbon dioxide and cyclobutadiene in a crystalline lattice by a distance of 1.5Å are interacting by a covalent bond, or a van der Waals attraction. In terms of energies, most chemists agree that these differ by around two orders of magnitude. This, I suggest, does not come into the category of nuance!

For a nice review of asymmetric synthesis under physical fields, see here.

Tags:chemist, chiroptical, Interesting chemistry, physicist, synthetic chemist
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Can a cyclobutadiene and carbon dioxide co-exist in a calixarene cavity?

Friday, November 19th, 2010

On 8th August this year, I posted on a fascinating article that had just appeared in Science (DOI: 10.1126/science.1188002), in which the crystal structure was reported of two small molecules, 1,3-dimethyl cyclobutadiene and carbon dioxide, entrapped together inside a calixarene cavity. Other journals (e.g. Nature Chemistry, DOI: 10.1038/nchem.823) ran the article as a research highlight (where the purpose is not a critical analysis but more of an alerting service). A colleague, David Scheschkewitz, pointed me to the article. We both independently analyzed different aspects, and first David, and then I then submitted separate articles for publication describing what we had found. Science today published both David’s thoughts (DOI:10.1126/science.1195752) and also those of another independent group, Igor Alabugin and colleagues (DOI: 10.1126/science.1196188). The original authors have in turn responded (DOI: 10.1126/science.1195846). My own article on the topic will appear very shortly (DOI: 10.1039/C0CC04023A). You can see quite a hornet’s nest has been stirred up!

At issue is whether the two bonds (indicated with arrows below) are best described as normally covalent, or very much weaker van der Waals contacts, or essentially non-interacting atoms. The last two interpretations would sustain the claim that 1,3-dimethyl cyclobutadiene and carbon dioxide can co-exist as separate species inside the cavity. The first would argue that they have reacted to form a different molecule. You can inspect the 3D coordinates by clicking on the diagram below.

Reported X-Ray structure. Click for 3D

Barboiu et al originally argued that these two bonds were strong van der Waals contacts, with C-C and C-O distances of 1.5 and 1.6Å respectively, and with a OCO angle of 120°. The various responses to this claim tend to the view that these distances/angles clearly represent new covalent (or partially ionic-covalent) bonds, and that the combined species cannot be described as 1,3-dimethyl cyclobutadiene and carbon dioxide. There is obviously much more to it than that (including a detailed analysis of the errors present in a partially disordered crystal structure). So make your own minds up based on the articles cited above and if it helps, the original 3D coordinates, for your convenience made available above!

Tags:David Scheschkewitz, Igor Alabugin, Interesting chemistry, pericyclic, watoc11, X-ray
Posted in Interesting chemistry | 18 Comments »

A historical detective story: 120 year old crystals

Wednesday, November 17th, 2010

In 1890, chemists had to work hard to find out what the structures of their molecules were, given they had no access to the plethora of modern techniques we are used to in 2010. For example, how could they be sure what the structure of naphthalene was? Well, two such chemists, William Henry Armstrong (1847-1937) and his student William Palmer Wynne (1861-1950; I might note that despite working with toxic chemicals for years, both made it to the ripe old age of ~90!) set out on an epic 11-year journey to synthesize all possible mono, di, tri and tetra-substituted naphthalenes. Tabulating how many isomers they could make (we will call them AW here) would establish beyond doubt the basic connectivity of the naphthalene ring system. This was in fact very important, since many industrial dyes were based on this ring system, and patents depended on getting it correct! Amazingly, their collection of naphthalenes survives to this day. With the passage of 120 years, we can go back and check their assignments. The catalogued collection (located at Imperial College) comprises 263 specimens. Here the focus is on just one, specimen number number 22, which bears an original label of trichloronaphthalene [2:3:1] and for which was claimed a melting point of 109.5°C. What caught our attention is that a search for this compound in modern databases (Reaxys if you are interested, what used to be called Beilstein) reveals the compound to have a melting point of ~84°C. So, are alarm bells ringing? Did AW make a big error? Were many of the patented dyes not what they seemed?

1,2,3-trichloronaphthalene

The story starts to get murky when Reaxys reports the earliest literature for this compound as being 1941 (DOI: 10.1039/JR9410000243), the authority being Wynne himself (now a sprightly 80). The collection of 263 specimens was thought to go back to the 1890s, so how could it contain a compound only made about 50 years later? Time to do an X-ray determination. Remarkably, the 120 year old crystals of specimen 22 were still in good shape, but the determined structure held an initial surprise. The compound was in fact 1,6,7-trichloronaphthalene, quite a different species from the label.

1,6,7-trichloronaphthalene

So, did AW get things badly wrong, and were all those patents based on these structures potentially invalid? A little more detective work using Reaxys reveals that the 1,6,7 isomer melts at 109.5°C, the same as reported by AW in 1890 (Chem. News J. Ind. Sci., 1890 , 61, p. 273). So how did the 1,6,7-compound come to be mistaken for a 1,2,3,-isomer? The culprit turns out to be one prime (‘).

1,6,7 = 2:3:1' Click for 3D

Updated (see comment) Click for 3D

The numbering system in 1890 was different from what it is now. Then, primes were used to distinguish the numbering on each ring. When the collection was catalogued (in the 1990s), the 1′ was mistaken for 1 (you can see the prime on the original label). AW were correct all along, and the patent owners for all those naphthalene dyes can rest easy.

Sample 22 from AW collection

What this teaches us is that crystallography on 120 year old organic compounds is perfectly viable. Indeed, can anyone else claim to have solved the structure of such an old compound? And that those old chemists knew what they were doing, despite not having any instrumentation to help them. Oh, and a final comment. Precious few collections of molecules made by the original scientists exist nowadays. Many a collection has literally been skipped because of health and safety concerns. The AW collection itself was rescued from oblivion by the narrowest of margins. And we have permanently lost the opportunity for any detective work of the type described above. You can see that I am very upset by this. Heritage conservation should not just be old buildings, paintings etc, but the chemical heritage collections as well.

Thanks to Andrew White for the crystal structures (of this and three other samples, but their stories are for another day).

Tags:Andrew White, chemical heritage collections, detective, Historical, Imperial College, Interesting chemistry, News J. Ind., toxic chemicals, William Henry Armstrong, William Palmer Wynne, X-ray
Posted in Interesting chemistry | 11 Comments »

Rate enhancement of the Diels-Alder reaction inside a cavity

Saturday, October 30th, 2010

Reactions in cavities can adopt quite different characteristics from those in solvents. Thus first example of the catalysis of the Diels-Alder reaction inside an organic scaffold was reported by Endo, Koike, Sawaki, Hayashida, Masuda, and Aoyama (DOI: 10.1021/ja964198s), where the reaction shown below is speeded up very greatly in the presence of a crystalline lattice of the anthracene derivative shown below.

A Diels-Alder reaction. Click for animation.

Organic scaffold based on an anthracene derivative. Click for crystal structure.

Its difficult to be precise about how much faster, since the kinetics depend on reorganisation of the scaffold, the actual reaction kinetics, and diffusion of the products in and out of the cavity. It does however mean that a poor solution reaction (reflux, many hours, modest yield) can be accomplished in an hour or so at room temperature in high yield.

Some idea of what is going on can be probed using calculation. Because the host and the guest interact though van der Waals or dispersion forces, a new breed of density functional theory which takes these into account is used (ωB97XD). The basic assemblage comprises the reactants shown below, enclosed in a cage formed by four of the anthracene units. A total of 236 atoms. This is a pretty challenging size for a full-blown quantum mechanical calculation. Here, its been done using a reasonable basis set, 6-31G(d) and with a continuum solvation model applied (dichloromethane). If you are interested in this sort of thing, that is 2292 basis functions. I started the calculations in mid September, and its taken more than six weeks to optimise (on 8-processor computers).

Firstly, the results for a control calculation in dichloromethane. The energies of activation of the two isolated reactants coming together at the transition state are calculated as:
ΔG₂₉₈ 29.5, ΔH 15.5, T.ΔS -13.98 kcal mol^-1 (ΔS -46.9 cal K^-1mol^-1)

which are of course the various contributions to the equation ΔG = ΔH – T.ΔS. Note in particular how the last term increases the free energy barrier by ~14 kcal mol^-1! Using the equation Ln k/T = 23.76 – ΔG/RT, one can estimate a rate constant of ~4 x 10^-6 hour^-1 at 298K (i.e. very slow at room temperatures). If the unfavourable -T.ΔS term is ignored (ΔG = ΔH), the rate constant increases to ~9 x 10⁴ hour^-1 at 298K (i.e. fast), quite a difference. What about the values when the reactants and transition state are surrounded by the host?

ΔG₂₉₈ 20.0, ΔH 16.5, T.ΔS -3.49 kcal mol^-1 (ΔS -11.7 cal K^-1 mol^-1)

The key difference is that the last term is now much smaller, this reduces the free energy of activation and the estimated rate constant at 298K is now ~ 0.01 s^-1 (42.5 hour^-1). This magnitude of rate constant corresponds to a reasonably fast reaction at room temperatures.

Transition state for Diels Alder inside a cavity. Click for 3D.

This post demonstrates that the fascinating area of supermolecular chemistry can be just as amenable to computational exploration as the more conventional reaction.

Tags:animation, catalysis, free energy, free energy barrier, G/RT, Interesting chemistry, Organic scaffold, pericyclic
Posted in Interesting chemistry | 7 Comments »

The strongest bond in the universe!

Sunday, October 24th, 2010

The rather presumptious title assumes the laws and fundamental constants of physics are the same everywhere (they may not be). With this constraint (and without yet defining what is meant by strongest), consider the three molecules:

Property (CCSD/aug-cc-pVTZ)	N≡N	(H-N≡N)⁺	(H-N≡N-H)²⁺
NN length, Å	1.0967	1.0915	1.0795
NN stretch, cm^-1	2418.8	2356.4 2545.1^a/2451.5^b	2226.3/3024.0 2688.4^a/2567.7^b
ELF NN basin integration	3.57	4.31	4.59
QTAIM ρ(r)/∇²ρ(r)	0.714/-3.38	0.690/-3.07	0.700/-2.96
^aValue for hydrogen mass of 10,000 ^bValue for hydrogen mass of 0.001.

The series explores the effect of protonating dinitrogen (generally considered as strong as a diatomic bond gets).

Firstly, one notes that the N-N distance decreases with mono and then diprotonation, the second protonation having the greater effect. Is shorter stronger?
What about the NN stretching vibration? Here one encounters an annoying feature of vibrations; the modes are not always pure. Thus whilst in N₂ itself, there is only one normal mode, and it is as pure as they get, by the time we have di-protonated, we have three stretching modes, two involving H-N and one N-N. They mix and none can now be considered a pure N-N stretch. Thus in H₂N₂, the highest wavenumber mode of 3024 is a mixture of H-N and N-N, and likewise the 2226 mode, albeit in different proportions. So a trick has to be played. If the mass of each hydrogen is increased to 10,000, modes involving these super-heavy atoms no longer mix with any other mode. Now, the N-N mode becomes pure, and its value is 2688, a significant increase on nitrogen itself. The monoprotonated form also shows a lesser increase.
The ELF disynaptic basins for the three molecules also steadily increase their populations. Electrons that were previously in the terminal nitrogen lone pairs now creep into the N-N region instead, and hence make the bond stronger. The population does not reach six (the nominal value for a triple bond), since the H-N regions still contain more than 2 electrons. But ELF matches the previous two results.
QTAIM measures the electron density ρ(r) at the bond critical point. Here different behaviour is seen, with ρ(r) lower for the monprotonated, and the diprotonated form intermediate between the other two. Perhaps absolute electron densities measured at a single point do not measure bnd strengths after all. The Laplacian, ∇²ρ(r) steadily decreases along the series.

So is the NN bond in HNNH²⁺ the strongest bond in the universe? Almost certainly. OK, so bonds with higher formal bond orders (Cr₂ for example) exist, but they come nowhere near HN≡NH²⁺, which is crowned champion.

Oh, by the way, another article (DOI: 10.1063/1.1576756) claimed the title in 2003, but I make the claim for a stronger bond here!

Tags:Interesting chemistry
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(Almost) 100 years of Lewis structures: are they still fit for purpose?

Monday, September 27th, 2010

The molecule below was characterised in 1996 (DOI: 10.1246/cl.1996.489) and given the name tris(dithiolene)vanadium (IV). No attempt was made in the original article to give this molecule a Lewis structure using Lewis electron pair bonds. This blog will explore some of the issues that arise when this is attempted.¹

NAMPOG.

The name given to the molecule by the chemists who made it reflects the ligand used, which we can represent as cis-HS-CH=CH-SH (via its di-sodium salt and reaction with VCl₃). Its entry in the Cambridge crystal database is NAMPOG (which carries only the slightest of semantic or structural information). The chemical name however does carry some further information, namely the designation tris implies three fold symmetry (D_3h in this case), and hence that all three ligands are in fact identical (structurally).

A nominal first stab at a Lewis electron pair representation reflecting this symmetry might be as shown above. At this point we hit a logical problem with the final component of the assigned name; the formal oxidation state of the metal is designated IV. However, three moles of (-)S-CH=CH-S(-) imply the ligands carry a formal charge of 6-, and that therefore the metal must be 6+, or VI. Six however is not an oxidation state normally exhibited by vanadium. Why did the original discoverers designate it IV? Well, because careful electron counting reveals the system as a whole has 161 electrons, of which 71 are designated as valence electrons, and hence it must have one unpaired valence electron. In the representation above, that electron is shown resident on the V atom with a dot, and the ESR spectrum measured for the molecule turns out to be apparently characteristic of V(IV) systems (they do not mention whether they also compared the spectra with those derived from genuine examples of V(II), see below). This implies (as the authors note) that a total of only 4- must be delocalized over the three dithiolene ligands.

Returning to our electron counting, of the remaining 70 valence electrons, 24 electrons are implicit above as twelve sulfur lone pairs (which are sometimes shown as double dots, but their explicit inclusion here would cause clutter) and so we presume the remaining 46 electrons must be in Lewis-like electron pair bonds. Well, the structure above implies 24 such bonds (the six C-H lines, as well as the Hs are also omitted by convention, again to avoid clutter!). We can begin to see why the original article lacks a Lewis structure, since the one above contains too many electrons (48 rather than 46).

How might one proceed to rescue the situation? Because a great many possible Lewis structures could be drawn, we have to learn a little more about the molecule and seek recourse in the bond lengths measured for the system. The most obvious is the C-C length, which turns out to be 1.36Å, a value significantly longer than expected for a C=C double (i.e. a four electron) bond, but a little shorter than the 3-electron bond found in e.g. benzene.

A second attempt at a Lewis structure

The Lewis structure (one of three equivalent ones) now has 5 lines in the C-C region, or ~3.3 electrons per C-C bond averaged over three ligands, which seems to match the length a little better. It also has 25 lines representing nominal electron pairs and ten sulfur lone pairs, a total of 70 electrons. The net effect of this representation is to transfer two electrons from the sulfur lone pairs to the vanadium, and hence to reduce the formal charge at the metal from 6+ to 4+, or to V(IV). This sort of behaviour, where electrons can be borrowed from a ligand and used to reduce (or oxidise) the metal they are coordinated to is called non-innocent behaviour. The dithiolene ligand is notoriously non-innocent. It results in this case in our innocent assumptions that bonds are defined by an integer number of electrons [2,(3),4,(5) or 6 as in Lewis’ original classifications] are no longer adequate, and that non-integer descriptors must also be used.

There is still one counting rule we have not inspected. To complete its valence shell to reach Kr, V needs 18 valence electrons. The representation above gives it 13. So how about the following, which ends up with a valence shell of 17 electrons for vanadium (and an oxidation state of V(II))?

Third time lucky?

This implies that the V-S bonds might be a little shorter than normal. Well in NAMPOG its 2.35Å, perhaps slightly shorter than a typical V-S single bond of ~2.4Å, but in fact we are now down at the noise level, and its clear that we have probably reached (if not exceeded) the limit of semantic interpretation of the Lewis model. In this case, only three (of 100s of possible) Lewis structures have been discussed, and of course they were selected only because we had some experimental information to discriminate between them. And we must be aware that whilst Lewis structures are the simplest way of analysing the electron distribution in a molecule, far more sophisticated analyses are nowadays possible. The real question is which analysis can actually result in a greater insight into the molecule? But the least that can be said about molecule NAMPOG is that it causes one to think about the problems of representing bonding (I will draw the line however at using this example in my university admissions interviews!).

¹ I thank J. P. P. (Jimmy) Stewart for drawing this molecule on my blackboard and hence provoking this blog post.

Tags:Cambridge, chemical name, Historical, Interesting chemistry, metal, V
Posted in Interesting chemistry | 9 Comments »

Secrets of a university admissions interviewer

Sunday, September 19th, 2010

Many university chemistry departments, and mine is no exception, like to invite applicants to our courses to show them around. Part of the activities on the day is an “interview” in which the candidate is given a chance to shine. Over the years, I have evolved questions about chemistry which can form the basis of discussion, and I thought I would share one such question here. It starts by my drawing on the blackboard (yes, I really still use one!) the following molecular structure.

Mystery molecule.

The candidate is then invited to offer their initial impressions of this molecule, and shortly thereafter asked how they might make it (or where perhaps they might be able to buy it). This of course floors even the most confident of applicants! But after a moments thought, most students can derive not only a molecular formula, but an empirical formula. From which it becomes apparent that it is actually a trimer of carbon dioxide. In the previous post, I showed the structure of solid CO₂ and how an oxygen from one unit came fairly close to a carbon from another. So the next logical question might be to ask if this might lead to a molecule such as shown above. Why a trimer? Well, the aromatic core is also easily perceived, and one might expect some aromatic stabilization to result (which most of the candidates readily spot). Its also ionic, and here perhaps solvation may help stabilize. Finally, armed with Le Chatelier’s principle, one might conclude that pressure too would help. At this point of course, the realization normally dawns that possibly the purchase of a carbonated soft drink in a supermarket might offer perhaps a few molecules of the above. The discussion normally takes about 10 minutes, and is guaranteed to stimulate (and quite possibly exhaust) most students.

But here in this post, I would like to offer the denouement. What actually are the chances of forming this species? Enter a B3LYP/6-311G(d,p) calculation of the free energy. We can do this for various models:

The trimer energy evaluated in a continuum solvent (water). This works out at 83.4 kcal/mol higher than three monomers (in part due to the entropic requirements of coalescing three molecules into one). So, not many molecules in a fizzy drink then! (just as well perhaps, since e.g. benzene as an aromatic molecule would not be a pleasant additive).
How aromatic is the molecule? Well, a NICS(1) index of -2.3 ppm suggests little actual aromaticity. The C-O bond length (1.365Å) is certainly short enough. The Kekulé vibrational mode however is quite low (974 cm^-1) compared to benzene, which is ~1310 cm^-1 (remember, this mode represents how much energy it takes to distort an aromatic ring from a symmetric structure to the bond localized form).
If its not aromatic, then perhaps after all a better representation might be:
A better resonance structure

It is worth asking why even this perfectly reasonable form is so much higher in free energy than carbon dioxide itself.
Would solvating the structure with three explicit water molecules help (as per below)? Deciding quite how the hydrogen bonds will form is an interesting exercise in its own right!
Solvated trimer

But now the energy is +96.8 kcal/mol compared to the monomers. Its that entropy again!
Actually, oxygen is pretty poor at propagating aromaticity. Nitrogen is much better, so what about the following (historically, such s-triazines were in fact much better known than benzene itself in the first half of the 19th century).
Carbo-diimide trimer

This is now merely +35.8 kcal/mol higher in free energy compared to three momers. The Kekulé mode is up to 1355 cm^-1(discuss!).

There are many other facets of this that could be raised. But the main reason for introducing such a molecule for discussion is that just by looking at the structure, so many ideas can be explored. That, by and large, is how chemistry works.

Tags:energy, free energy, Interesting chemistry, trimer energy
Posted in Interesting chemistry | 3 Comments »

Solid carbon dioxide: hexacoordinate carbon?

Friday, September 17th, 2010

Carbon dioxide is much in the news, not least because its atmospheric concentration is on the increase. How to sequester it and save the planet is a hot topic. Here I ponder its solid state structure, as a hint to its possible reactivity, and hence perhaps for clues as to how it might be captured. The structure was determined (DOI 10.1103/PhysRevB.65.104103) as shown below.

The structure of solid carbon dioxide. Click for 3D

The two nominal double bond distances are 1.33Å, whilst a further four O…C contacts in the shape of a square complete the coordination (2.38Å each). All would probably agree that the central carbon is best described as hexa-coordinated. This is also a hot topic. For example, note the claim made recently to have created a hexa-coordinated carbon species by design (Synthesis and Structure of a Hexacoordinate Carbon Compound, DOI: 10.1021/ja710423d) based on a motif derived from an allene:

Designed hexacoordinate carbon. Click for 3D

This claim was supported by an unusual measured property, the electron density ρ(r) and its Laplacian in the putative O…C region. These two properties are one of those (relatively rare) meetings between experiment and quantum mechanics, and their usefulness has been noted in this blog on previous occasions. However, note that in this designed structure, the O…C distances are merely 2.65-2.7Å, significantly longer than in solid carbon dioxide! So carbon dioxide, in a form many of us are familiar with (solid), can certainly be justified as being described as having a hexacoordinate carbon (although we might draw the line at describing it as having hexavalent carbon).

If oxygen atoms can approach the carbon in CO₂ to within ~2.4Å, an interesting question can be posed. How close can another carbon get to CO₂ without actually reacting and forming a new molecule? C-C bonds, even weak ones, are so much more interesting than C-O bonds! It would have to be a particularly nucleophilic carbon, of course. A search of the August 2010 version of the Cambridge structural database (CSD) reveals no really close approaches of another carbon to CO₂. Only about 8 weak examples are found, and here the C-C distances are ~3.0-3.2Å, with the O=C=O angle in the CO₂ never less than 170°. In this context, there is an intriguing and very recent report (which has not yet made it into the searchable CSD) of the structure of CO₂ trapped in a cavity next to what was claimed to be a molecule of 1,3-dimethyl cyclobutadiene, or CBD (see 10.1126/science.1188002 and the discussion of this article in my earlier blog post). The focus in that report was on the “Mona Lisa of organic chemistry”, namely the CBD unit. One feels that the structure of the adjacent CO₂ was of lesser interest to the authors. According to a visual image of this system, the CBD and CO₂ pair show quite an intimate approach via their carbon atoms (a ghostly C-C bond is clearly represented). This raises the interesting question of whether the description of this pair should be of two intimate but nevertheless separate and relatively unperturbed molecules not connected by a covalent bond (“more indicative of a strong van der Waals contact than of covalent bonding“) or of a pair fully bound by a covalent C-C bond between them?

The issue of what is an interaction, and what is a bond continues to raise its often controversial head. And quantum theory continues to provide a multitude of interpretations as well.

Tags:13-dimethylcyclobutadiene, Borboiu, Cambridge, carbon dioxide, CBD and CO, CCDC, CDS, crystal structure, crystalline calixarene network, guest, host, Hypervalency, Interesting chemistry, often controversial head, van der Lee
Posted in Hypervalency, Interesting chemistry | 5 Comments »

The oldest reaction mechanism: updated!

Tuesday, September 14th, 2010

Unravelling reaction mechanisms is thought to be a 20th century phenomenon, coincident more or less with the development of electronic theories of chemistry. Hence electronic arrow pushing as a term. But here I argue that the true origin of this immensely powerful technique in chemistry goes back to the 19th century. In 1890, Henry Armstrong proposed what amounts to close to the modern mechanism for the process we now know as aromatic electrophilic substitution [1]. Beyond doubt, he invented what is now known as the Wheland Intermediate (about 50 years before Wheland wrote about it, and hence I argue here it should really be called the Armstrong/Wheland intermediate). This is illustrated (in modern style) along the top row of the diagram.

The mechanism of aromatic electrophilic substitution

In 1887, Armstrong had tabulated the well known ortho/meta/para directing properties of substituents already on the ring towards this reaction[2]. He even offered an explanation, which is not entirely wrong, given that in 1890, the electron had not yet been discovered! That did not stop Armstrong, who invented an entity he called the affinity for the purpose of developing his theories (in this theory, benzene had an inner circle of six affinities, which had a tendency to resist disruption). Armstrong’s description of the properties of the affinity matches that of the (yet to be discovered) electron very closely! But that is enough of history. The mechanism shown above emerged in its present representation (and naming) during the heyday of physical organic chemistry between 1926 – 1940, and of course is an absolute staple of all text books on organic chemistry. But, sacrilege, is it correct? Could what is referred to as an intermediate instead be a transition state? (shown in the bottom pathway of the scheme).

Consider instead the following, in which X is replaced by an acetic acid motif;

Transition state alternative to the Wheland

The two steps, a bond formation between the benzene and E, and the proton abstraction from the benzene by X, are now synchronized into a single step, and the intermediate is now transformed into a transition state. Time to put this theory to the test. X is going to be made trifluoroacetate (R=CF₃) and we are going to test it with E= NO⁺ and F⁺ (yes, trifluoroacetyl hypofluorite is a known chemical, and it really does fluorinate¹ aromatic compounds at -78C). Firstly, E= NO⁺. A B3LYP/6-311G(d,p) calculation[3] run in a solvent simulated as dichloromethane, reveals the mid point to indeed be a transition state and NOT an intermediate![4].

Wheland as a transition state. Click image for animation

There is one crucial aspect to this transition state that Armstrong himself made a point of. In the Wheland intermediate proper, the aromaticity of the benzene ring must be disrupted. As a transition state, it need not be (at least not completely). Thus the two bonds labeled as a have calculated lengths of ~1.415Å, only slightly longer than the aromatic length, and certainly not single bonds as implied by the Wheland intermediate! Notice also the significant motion by the hydrogen, which implies the reaction would be subject to a kinetic isotope effect (this would normally be interpreted in terms of the second stage of the stepwise reaction shown along the top a being rate limiting, but this result shows this need not be so). Thus, if the structure is favourable, this veritable old mechanism can be redesigned to give a new, 21st century look to a 19th century staple! By the way, the free energy of activation for this reaction is calculated as ~22 kcal/mol, a perfectly viable thermal reaction. No doubt, by suitable design of the group X, this might be reduced.

Now on to E=F⁺[5]. This looks a little different. F⁺ is now a much more voracious electrophile than the nitrosonium cation, and it therefore jumps ahead of the second mechanistic step, with no motion of the hydrogen being involved at this stage (one might also imagine making X a better base to swing things the other way).

Transition state E=F+ leading to Wheland Intermediate. Click for 3D model.

Genuine Wheland intermediate for E=F+ Click for 3D model

Now a full blown Armstrong/Wheland intermediate does indeed form (10042/to-5174); an intimate ion pair if you will, even in the relatively non polar dichloromethane as modelled solvent. The structure (shown above) is rather unexpected. This reaction has ΔG^† of ~5 kcal/mol, which is significantly lower than for the E=NO⁺ system.

Chemistry is full of surprises, and it is always a wonder how a slightly different take on even the oldest of reactions can reveal something new.

Reference.

p>1. Umemoto, T.; Mukono, T.. 1-Acylamido-2-fluoro-4-acylbenzenes. Jpn. Kokai Tokkyo Koho (1986), Patent number JP61246156.

References

"Proceedings of the Chemical Society, Vol. 6, No. 85", Proceedings of the Chemical Society (London), vol. 6, pp. 95, 1890. https://doi.org/10.1039/pl8900600095
H.E. Armstrong, "XXVIII.—An explanation of the laws which govern substitution in the case of benzenoid compounds", J. Chem. Soc., Trans., vol. 51, pp. 258-268, 1887. https://doi.org/10.1039/ct8875100258
"C 8 H 6 F 3 N 1 O 3", 2010. http://doi.org/10042/to-5172
S.R. Gwaltney, S.V. Rosokha, M. Head-Gordon, and J.K. Kochi, "Charge-Transfer Mechanism for Electrophilic Aromatic Nitration and Nitrosation via the Convergence of (ab Initio) Molecular-Orbital and Marcus−Hush Theories with Experiments", Journal of the American Chemical Society, vol. 125, pp. 3273-3283, 2003. https://doi.org/10.1021/ja021152s

Tags:acetic acid, animation, aromaticity, chemical, free energy, Henry Armstrong, Historical, Interesting chemistry, Wheland intermediate
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