Posts Tagged ‘Alan Dronsfield’

The history of Alizarin (and madder).

Thursday, October 18th, 2018

The Royal Society of Chemistry historical group (of which I am a member) organises two or three one day meetings a year. Yesterday the October meeting covered (amongst other themes) the fascinating history of madder and its approximately synthetic equivalent alizarin. Here I add a little to the talk given by Alan Dronsfield on the synthesis of alizarin and the impact this had on the entire industry.

Although William Perkin famously (and accidentally) produced the first synthetic chemical dye in 1856 (Mauveine), the industry at that time was both large and dominated by dyes from natural products. Mauve was something of a niche colour; far more important was alizarin, both as a red dye (for cotton) and a red pigment (in painting) and up to 1869 it was sourced from the roots of the madder plant (which was difficult to farm) and from insects (which could be farmed). It was nonetheless expensive to produce it from either and so a race started to create it synthetically. Famously, two groups submitted patents for such a synthesis in 1869, William Perkin himself and two scientists working in BASF, Carl Graebe and Carl Liebermann.[1],[2] The latter were the winners (by one day) and they are now famed for their work‡ (what a difference one day can make; Perkin is known for his other work, but not as much for the synthesis of alizarin). As with mauveine, the structures of these dyes were not known with certainty (or for mauveine even approximately) at the time, but Graebe and Liebermann had managed to prove that alizarin was derived from anthracene by reducing the former to the latter using zinc dust. Trouble was, the structure of anthracene itself was not certain in 1869! There were two probable candidates, (a) and (b) below.

Alan told us how Graebe and Liebermann favoured structure (a), now known as phenanthrene, rather than (b), which we recognize as anthracene. A full story is told in this PhD thesis, written in 1919 and published in 1921[3] and I can only tell a tiny bit of it here. Essentially (a) was preferred over (b) because the former could sustain three aromatic (benzene-like) rings, whereas the latter only two (p 3 of the thesis above). Years later in 1972, this concept emerged as the Clar π-sextet rule, but the idea was already more than 100 years old by then! And indeed thermodynamically, phenanthrene is more stable than anthracene. By 1872, circumstantial evidence was accumulating that in fact alizarin was derived from (b), largely via attempts to synthesize the molecule by various reactions. These often were performed at high temperatures (red-hot tubes), and we now know that many complex rearrangements can occur at such temperatures. In 1889[4], Armstrong was quoting the structure of anthracene with no doubts about its structure. However, it took another 30 years or so for an entirely unambiguous total synthesis of anthracene to be devised.[3] Also around that time the first structures based on crystallography were emerging (by William Bragg) that supported this hypothesis. Even so, the first modern crystal structure had to wait until 1950.[5]

We learn from this story that many chemical structures established during the 19th century were largely based on (admittedly a large) body of circumstantial evidence. A wonderful example of how a systematic rather than a circumstantial proof of the structure of naphthalene was established using chemical synthesis and degradations alone can be found here in the work by Armstrong. Evidence obtained from instruments was largely restricted to techniques such as thermochemistry and polarimetry in the 19th century and for the first twenty years of the 20th to e.g. infra-red spectroscopy.[6] It is remarkable then that actually, most 19th century structures have stood the test of time. Moreover, not knowing the precise structure did not prevent the processes for making them to be patented. Nowadays of course, a simple crystal structure can often be solved in a few minutes and NMR spectroscopy takes a similar amount of time. We are no longer used to waiting for years or indeed decades for structural proof!

‡This synthesis proved to be very expensive (requiring a step using bromine and then a second step to remove it). But shortly after, a much more efficient synthesis which dispensed with the bromine brought the cost of the dye down dramatically. The madder industry never really recovered from this blow.

References

  1. C. Graebe, and C. Liebermann, "Ueber künstliche Bildung von Alizarin", Berichte der deutschen chemischen Gesellschaft, vol. 2, pp. 14-14, 1869. https://doi.org/10.1002/cber.18690020106
  2. C. Graebe, and C. Liebermann, "Ueber künstliches Alizarin", Berichte der deutschen chemischen Gesellschaft, vol. 2, pp. 332-334, 1869. https://doi.org/10.1002/cber.186900201141
  3. C.W. Colver, and W.A. Noyes, "SYNTHESIS OF ANTHRACENE<sup>1</sup> FROM NAPHTHALENE.", Journal of the American Chemical Society, vol. 43, pp. 898-905, 1921. https://doi.org/10.1021/ja01437a023
  4. "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
  5. A. McL Mathieson, J.M. Robertson, and V.C. Sinclair, "The crystal and molecular structure of anthracene. I. X-ray measurements", Acta Crystallographica, vol. 3, pp. 245-250, 1950. https://doi.org/10.1107/s0365110x50000641
  6. W.W. Coblentz, "Infra-red Absorption Spectra: I. Gases", Physical Review (Series I), vol. 20, pp. 273-291, 1905. https://doi.org/10.1103/physrevseriesi.20.273

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The atom and the molecule: A one-day symposium on 23 March, 2016 celebrating Gilbert N. Lewis.

Friday, December 11th, 2015

You might have noticed the occasional reference here to the upcoming centenary of the publication of Gilbert N. Lewis’ famous article entitled “The atom and the molecule“.[1] A symposium exploring his scientific impact and legacy will be held in London on March 23, 2016, exactly 70 years to the day since his death. A list of the speakers and their titles is shown below; there is no attendance fee, but you must register as per the instructions below.

Royal Society of Chemistry Historical Group Meeting on 23th March 2016, Burlington House, Piccadilly, London: The atom and the molecule: A symposium celebrating Gilbert N. Lewis.

  • Dr Patrick Coffey (Berkeley, USA): Does Personality Influence Scientific Credit? Simultaneous Priority Disputes: Lewis vs. Langmuir and Langmuir vs. Harkins
  • Professor Robin Hendry (Durham, UK): Lewis on Structure and the Chemical Bond
  • Professor Alan Dronsfield (UK): An organic chemist reflects on the Lewis two-electron bond
  • Dr Julia Contreras-García (UPMC, France): Do bonds need a name?
  • Professor Nick Greeves (Liverpool, UK): The influence of Lewis on organic chemistry teaching, textbooks and beyond
  • Professor Clark Landis (UWM, USA): Lewis and Lewis-like Structures in the Quantum Era
  • Professor Michael Mingos (Oxford, UK): The Inorganic dimension to Lewis and Kossel’s landmark contributions
  • Dr Patrick Coffey (Berkeley, USA): Lewis’ Life, Death, and Missing Nobel Prize

Prior registration is essential. Please email your name and address to Professor John Nicholson,  jwnicholson01 @ gmail.com

References

  1. G.N. Lewis, "THE ATOM AND THE MOLECULE.", Journal of the American Chemical Society, vol. 38, pp. 762-785, 1916. https://doi.org/10.1021/ja02261a002

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How is the bromination of alkenes best represented?

Sunday, October 14th, 2012

I occasionally delve into the past I try to understand how we got to our present understanding of chemistry. Thus curly arrow mechanistic notation can be traced back to around 1924, with style that bifurcated into two common types used nowadays (on which I have commented and about which further historical light at the end of this post). Here I try to combine these themes with some analysis of wavefunctions for a particularly troublesome reaction to represent, the dibromination of an alkene, which I represented in the previous post as shown below.

However, that first step in the mechanism (red box above) has been represented in two other ways in the past. The green box used to be common, but is now being superseded by the magenta style (as for example here). 

I should first explain (what I think is) the intent of these three schemes (this is not really declared very formally, and so I am interpreting rather than stating):

  1. The green scheme, showing two arrows, attempts to illustrate a nucleophile (the alkene pπ electrons) interacting with an electrophilic acceptor (the Br-Br σ* bond) in a filled/empty orbital sense.
  2. The magenta scheme, showing three arrows, adopts a notation where a nucleophilic electron pair attacks an electrophilic atom centre to form a new bond between the two atoms.
  3. The red scheme has the nucleophilic electron pair terminating at the (approximate) bond centroid of the newly forming bond rather than at the atom.†

I set out to see what light quantum mechanics might be able to cast. To do so, I subjected the wavefunction of the system shown in the blue box above to analysis, which represents the species being formed as a result of the curly arrows pushed in the schemes above.

Firstly, QTAIM, which determines the topology of the electron density in the molecule. The green dots in this diagram are what are called “bond-critical-points”, or BCPs. The numerical values associated with each green dot are the electron density ρ(r) at that point.

Now, I have pointed out elsewhere that the existence of such a topological feature does not necessarily coincide with what we think of as a bond. With that caveat in mind, one can see that the BCPs reveal a cyclobromonium ring, with one C-C single bond (from the value of  ρ(r) ), two C-Br bonds, and a ring point (shown in red). The slightly weaker bond (again from the value of  ρ(r) ) is the one about to cleave as a result of the tribromide anion attacking the carbon, with inversion of configuration at that carbon. 

This picture does seem to correspond to our intuitive thoughts about mechanism. It offers a way of interpreting the arrow pushing scheme shown in red above. In this sense, an arrow would start at a BCP of a nucleophilic (donor) bond in the reactant, and terminate at the BCP of the acceptor bond formed in the product. Should we need to do so, we could derive precise 3D coordinates of the relevant BCPs, and ensure that our curly arrows either start or end precisely at those coordinates. This method would not allow for example the magenta scheme, since the terminating point of what is after all an electron arrow cannot be at a nucleus but needs to be at a bond (critical point). There is however one aspect un-answered. Both the red and magenta schemes have one arrow starting at an electron “lone pair”, and QTAIM does not give us coordinates for these! I will deal with this aspect last.

Yet another way of looking at it is to interpret the wavefunction in terms of pairs of (doubly) occupied and empty localised orbitals so that a donor-acceptor interaction energy can be derived. On to the NBO technique (natural bond orbitals), which tells us about the donor/acceptor interactions within a molecule. This relates to how the green box above might be viewed. When this is done, two sets of NBO orbital pairs are especially pertinent; each is a (doubly occupied) donor originating from the alkene (purple and orange) and an (empty) acceptor corresponding to the Br-Br cleaving bond (red and blue). In the overlay of two NBOs below, the purple (donor) and blue (acceptor) densities are overlapping in-phase to form a C-Br bond (arrow 1) Equally, overlap of an orange (donor) originating in part from a lone pair on bromine and the red acceptor (from the B-Br anti-bond) are forming the second C-Br bond (arrow 2). This actually corresponds more closely to the magenta than the green box. 

NBO Analysis. Click for 3D.

NBO Analysis. Click for 3D.

This brings us back to a deficiency of QTAIM; it does not tell us what kind of bonds we are dealing with. We might have presumed that the formed C-Br bonds of the bromonium ring (e.g. blue box) were single and hence two-electron bonds.

We have one more way of looking at bonds, and this method also allows us to count the electrons in the bond. Remember, a bond does not have to contain integer numbers of electrons! It can just as easily be fractional, as for example in PF5. ELF (the electron localisation function) is a way of looking at the properties of the electron density to identify localised spin-paired probabilities. The ELF technique partitions the function into discrete basins, and these can then be integrated for the total number of electrons defined by the ρ(r), the electron density. The centroids of these basins then give us something actually quite similar to the bond-critical-points from QTAIM theory, but carry two additional benefits. Firstly, the total electron count associated with each basin. Secondly, it also gives us the centroids of any lone pairs (which we identified as something helpful for defining either a start point or an end point of a curly arrow in arrow pushing). I show below the ELF analysis of the ion-pair intermediate of this bromination (i.e.the outcome of the arrow pushing in the red or magenta boxes). The red dots are basin centroids; there are lots of them but I concentrate only on the two marked with black arrows. They are the result of the donor-acceptor orbital overlaps, the principle one of which was shown above. These two ELF basins each have electron integrations of ~1e. Each C-Br “bond” contains only one shared (i.e. covalent) electron.

Click for 3D.

So which of the three schemes above is the most realistic? Well, the green scheme uses only one curly arrow in the carbon-bromine region, and so it carries the message that the bonds in this region only involve two electrons. The red scheme corresponds closely to the topology of the electron density involved in the reaction, but clearly, its arrows are NOT simple two-electron arrows. Neither are those of the magenta scheme, which seems rather to fall between two stools; it is not accurate topologically, but neither are its arrows simple two-electron arrows.

My conclusion must be a reminder that when we push curly arrows, we may not necessarily be able to associate these arrows with simple pairs of electrons. This is quite a subversive statement to make. But then again,surely the concept of curly arrow pushing, dating as it does from 1924, is overdue a makeover?

† Alan Dronsfield has contacted me with some information about when styles 2 and 3 might have bifurcated. Two particularly influential early text-books on mechanism, one published by Gould in 1959 and another by Sykes in 1961, appear to have adopted respectively the magenta and red schemes.

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The first ever curly arrows.

Friday, July 20th, 2012

I was first taught curly arrow pushing in 1968, and have myself taught it to many a generation of student since. But the other day, I learnt something new. Nick Greeves was kind enough to send me this link to the origin of curly arrow pushing in organic chemistry, where the following diagram is shown and Alan Dronsfield sent me two articles he co-wrote on the topic (T. M. Brown, A. T. Dronsfield and P. J. T Morris, Education in Chemistry, 2001, 38, 102-104, 107 and 2003, 40, 129-134); thanks to both of them.

This diagram dates from 1924, and is to be found in an article published by Robert Robinson (J. Soc. Chem. Ind., 1924, 43, 1297, a journal difficult to get hold of nowadays). Here, Robinson was trying to explain why the nitroso group is o/p-directing in aromatic electrophilic substitution. Whilst the notation is remarkably modern, some aspects do need explaining. 

  1. Robinson shows the nitrogen lone pair (arrow 1) as a line, and not as we now do, a double dot.
  2. Similarly, he shows arrow 3 ending at a line. We now do not show this in the starting structure, but reveal it in the final result, as above on the right, and again shown as a double dot.
  3. Similarly, he shows a + charge on the nitrogen at the start, whereas we now show it as the outcome of the process.
  4. If Robinson intends to create a +ve charge, then he really should balance that by showing the creation of a negative charge in the p-position of the ring. He does not balance his charges! 
  5. As was the custom at the time, the benzene ring itself is not represented in the Kekule mode (which of course should have been well known in 1924) but as what looks to us now as cyclohexane. It must have been the case in 1924 (and for several decades after) that cyclohexane itself was not regarded as an interesting system, and hence there must have been little confusion about drawing benzene as (modern) cyclohexane. The implied semantic of showing such a ring was that it represented benzene.
    1. But this way of drawing it leads to really difficult issues. Thus Robinson’s arrow 2 departs from what looks to us like a single bond, in which case no bond would be left. Robinson of course means implicitly that arrow 2 reduces the bond order by one, and if we start with a double bond from a Kekule structure, that the bond is reduced to 1, not zero, as is shown in the modern notation above.
    2. Likewise, the destination of arrow 2 in Robinson’s notation clearly creates a double bond. Which again is an issue, since he is not showing the double bonds. The trouble really arises because Robinson does not illustrate the outcome of his process.
    3. Finally, whereas arrow 1 starts at a line representing a lone pair, that line is disconnected from the N. However, the destination of arrow 3 appears to create a bond, not a lone pair.

Now that we have clarified Robinson’s meaning, what else can we say about Robinson’s structure.

  1. It is important to realise that in 1924, the 3D characteristics of electrons (their wavefunction) were not known. Looking at the modern version of the diagram, chemists realise that when a double line is drawn, the two are not the same. One line represents a σ-bond, the other a π-bond. We recognise that the two have different spatial characteristics. Hückel it was who showed that in planar aromatics, the two sets are in fact orthogonal, and do not mix. At which point we need to sort out what the three arrows in Robinson’s diagram represent. Arrows 2 and 3 we recognise as π-arrows. But what of arrow 1? I decided to do a search of the Cambridge data base for nitrosobenzenes, finding 22 sets of coordinates. In all except one, the two atoms of the nitroso group were co-planar with the six of the benzene ring. We now know of course that this places the nitrogen lone pair firmly in the plane of the eight atoms, and hence of a σ-type. Strictly therefore, it is orthogonal to the π-arrows and cannot be mixed with them. The solution of course is to first rotate the nitroso group by 90° to bring the nitrogen lone pair into conjugation with the π-system, whereupon Robinson’s arrows now “work”. 
  2. On a more minor point, we recognise that the nitrogen lone pair occupies a trigonal position, and so we draw the C-NO group as bent, rather than linear as Robinson did.
  3. If the co-planarity of the nitroso and benzene rings is retained, then the only way to draw the arrows is in the opposite direction to Robinson, resulting in the creation of a -ve charge on the oxygen and a +ve charge on the p-carbon. This of course is the resonance we now show for the nitro group, and implies m-direction, not o/p. 
  4. Which raises the fascinating question. Why, if the structure of nitrosobenzenes appears to be planar and not rotated, is the nitroso group nevertheless observed to be an o/p director? The answer of course must be in looking at the properties of the transition state, and not the starting material itself. But in 1924, the concept of a transition state itself was not yet recognised.

So this little blast-from-the-past example still gives us lots to think about!

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