Is science sometimes in danger of getting tunnel vision? Recently published ebook author, Ian Miller, looks at other possible theories arising from data that we think we understand. Can looking problems in a different light give scientists a different perspective?

Share this |

Share to Facebook Share to Twitter Share to Linked More...

Latest Posts

My alternative explanations for planetary formation survived a further two months. The reason for not giving a November update was not that I was hiding something, but there was word that a "huge announcement" would be made from the Curiosity team on December 5 so I waited. Could this fulfill one of my predictions? Er, no! The announcement was a bit of a squib: the last bit of equipment was working. Yes, this is a minor miracle, but it says nothing about planetary formation.
Obviously, a good number of papers were published, but very few had anything relevant to say about this theory. One of the more interesting came from Tobin et al., (Nature 492: 83-85). The protostar L1527 IRS has about 0.2 solar masses, while it is surrounded by a rotationally supported disk containing at least seven Jupiter masses, and further surrounded by an envelope containing about 1 solar mass. This is obviously a star in the early stages of formation, and as far as can be seen, it follows standard theory quite nicely. If we wait for about 3 million years we might see newly formed planets!
A paper by Crida and Charnoz (Science 338: 1196 – 1199) proposes that satellites form from massive rings around planets. A case is made that such rings form, then spread out beyond the Roche limit, accrete into larger bodies, and are moved out by tidal forces. It is not entirely clear, at least to me, how this works with gas accreting inwards, as gas drag should drag bodies inwards. The model is claimed to give good agreement with Neptune's inner minor satellites, the Uranian system, and Saturn's inner system. The Jovian system does not fit, and Titan is not really a good fit. One problem is that such rings must be extremely massive. The proposal differs from what I proposed, although my compositional proposal might still explain the rings, if they existed.
For those wondering how my theory differs from standard theory, the latter is essentially purely physical, with accretion being due to gravity and driven by gas flows, turbulence, etc that lead to collisions. I agree that these are important, and the main drivers once a certain size is reached, however standard theory cannot explain how the starting position, a distribution of planetesimals, form, because bodies up to tens of kilometers in size have negligible gravity, and collisions do not lead to binding strength. My major difference is that the initial stages are driven by chemistry, and our system is typical of such a system where the star blows out the disk within 1 My after forming. If this does not happen, planets keep accreting, and now become gravitationally unstable, which leads to a variety of different, smaller system types. By being driven by chemistry, the governing variable during the initial accretion stages is temperature.
Just before posting this, a new system was claimed: Tau ceti is claimed to have planets with semimajor axes at 0.1, 0.195, 0.37, 0.55 and 1.35 A.U.  I made approximate predictions for the outer part of this system, and the outcome is mixed. In my theory, the rocky planets are governed by two temperature profiles, the initial accretion, and that prior to the final removal of disk material, while the outer planets are primarily determined by the primary temperature distribution. If we interpret the planet at 1.35 A.U. as the accretor due to water ice (i.e. the Jupiter equivalent, irrespective of size), then the outer planets are about twice as close to the star as I suggested, which could arise if Tau Ceti accreted more slowly than our star, or because Tau Ceti has a lower metallicity, the accretion disk may have radiated heat better, by being somewhat more transparent. If so, and if these distances are real (see caveat below) there will be three further planets at about 2.1 A.U, 4.7 A.U., and 7.8 A.U. The one at 4.7 A.U. would be the smallest, while the one at 2.1 A.U would have formed an atmosphere similar to that of Titan. (I had predicted the second outermost one would have an atmosphere of nitrogen, if it were big enough to have an atmosphere, but because the planet is twice as close to the star, it too might have volatile methane.) The inner planets are further from the star than simple proportion, as would be expected because the second temperature profile is significantly due to stellar radiation. So is that confirmation or falsification? It may be neither, because a caveat must be noted: These planets are small; the so-called Jupiter equivalent is only 2% that of our Jupiter (although I predicted they would be small, due to the lower metallicity) and were detected at the very limit of stellar wobble. They may not be real, or may not be quite as reported.
Finally, I wish you all to enjoy the festive season. This will be my last blog until next year, when I shall return with some postscripts to my PhD project.
Posted by Ian Miller on Dec 21, 2012 10:13 PM GMT
The Roman festival of the Saturnalia, on the winter solstice, was where order was turned on its head. The question then is, do followers of this blog have the nerve to have their deepest beliefs challenged? For those who wish to exercise the brain over the next period, on December 21-23, there will be a free download of Aristotelian Methodology in the Physical Sciences from Amazon. The concept behind this ebook is that when Galileo threw out the bath water of Aristotle's physics, he also threw out the baby, namely Aristotle's methods for forming theories. Aristotle made at least two gigantic errors in physics but as this book shows, they arose because Aristotle forgot to use his own methodology (or he had yet to develop it; Physica was one of his earliest works.)
The first four chapters of Part 1 introduce the reader to my interpretation of what Aristotle was trying to say, then there are several examples of where failures to follow this advice have led to problems. One recent failure is the case of the non-classical carbenium ion, where two Nobel-prizewinners went head to head and failed to reach a conclusion. Part two then covers physics that I think a chemist should know about, which includes mechanics, waves, electromagnetic theory, quantum mechanics, and some other physics of general interest, then chemistry that I thought physicists should know about. That latter part is probably the weakest part of the book, and is largely based on explanations I have had to give to physicists I have worked with. A number of you may be able to offer suggestions for a revised edition.
It is the third part that is intended to offer the intellectual challenges, and it comprises 72 problems, in which the reader has to offer an alternative theory to . . .   Then, just to show it can be done, I offer answers, not all of which are beyond criticism because the second challenge is for the reader to accept or falsify my answers. Are you up to it? Examples include:
How could Priestley have ensured that the phlogiston theory prevailed?
Isaac Newton spun a bucket of water, and noticed centrifugal forces forced the water to the edges. Similarly, you can have artificial gravity in a spinning space ship. The problem is, how does the water know the bucket (or the astronaut, the space ship) is spinning? Isaac Newton appeared to fail on this, and a number of modern physics books mention the issue but provide no answer. Can you do better?
Yes, there is that non-classical carbenium ion. Surely you will not be put off by two Nobel-prizewinners failing to solve it?
I could not resist some of my own work, so one question starts off asking you to formulate an alternative interpretation of quantum mechanics. In the spirit of Christmas, a clue: Following Aristotle, either there is a wave of there is not. If there is, either it travels at the same velocity as the particle or it does not. That should get you going.
Finally, the interpretation I have come up with is non-local only to one wavelength, or one quantum of action. But Alain Aspect showed that entangled photons do not comply with Bell's inequality. If so, my interpretation is wrong. Your job is to falsify his claim (explained in sufficient detail earlier). Then, your last problem is to falsify my answer, which is based on the requirement that Bell's inequality must be followed if energy is conserved, and if the associative law of sets holds. As an aside, attempts to publish this argument in journals led to return by editors without peer review, except once when I was informed "This is wrong; the maths are trivial." Unfortunately, no clues as to where it is wrong, so you can help (or get help from a physicist.) Mind you, I have shown it to two professors of mathematical physics and they ducked for cover, as did a professor of theoretical physics. I sent it to a professor of theoretical physics who had written a book on Bell's inequality, he had promised to tell me where I was wrong but he never answered. My guess is you will fail on that one.
Posted by Ian Miller on Dec 15, 2012 1:57 AM GMT
In a previous blog, I discussed some misadventure, but there were other incidents of somewhat higher quality. (There had to be!)  To get to my lab from the main stairs, I had to go through the main lab. This had quite a number of benches, together with a couple of side rooms which seemed to be of more value for light entertainment when bored than anything else. However, when entering the main lab, some previous student, presumably inspired by Dante Alighieri, had ensured that one of the first things a visitor saw was a bun suspended from the ceiling by a wire. This bun was completely devoid of mould or any fungal growth! Enter not herein, for this laboratory is totally unsuitable for life!
One day, when I left my lab and entered the main lab, there to my right was something that I shall never forget. A friend was trying to purify a nitroazetidine by distillation, and was using a microburner on a pear-shaped flask. As I watched, suddenly there was a "pop", the top of the distillation equipment went upwards, and there was the poor victim sitting disconsolately holding the microburner while a light cloud of soot descended on him.
Then there was the fire alarm. A fire in a chemistry lab is a nightmare, although to be honest, I have never seen one, which says something about chemists ability to handle flammable materials. Anyway, the rules were clear. To make it easier to evacuate, those who could used the fire escape external to the building. Eventually, everybody assembled in the quad, and I received a few remarks, for I had come down the fire escape holding a lab bench drawer. In the drawer were all my precious samples and my lab book, and I was the only one who had thought to take that precaution.
Finally, an embarrassing scene. About the end of my second year, the Department introduced multiple choice questions for the first time, and as you might guess, the PhD students couldn't wait to get their hands on one of the papers. "What idiot set this question?" I asked, and looked up, and it was fairly clear the "idiot" was at the far end of the lab! Nevertheless, I stood by it. The question was, you are determining the molecular weight of benzene by the Victor Meyer method. If the benzene is wet, will the result be too high, too low, or about the same? I could justify each answer! Unless the experimentalist is a clod and does not introduce drops of water, the answer is, not much difference because water is essentially insoluble in benzene. If there is a drop of water, and the water stays in the vapour phase, the answer is, too low, but if the water condenses somewhere, the answer is, too high (because there is weight that does not give rise to vapour). I really hate these questions where the student cannot explain why the option was taken.
Slightly off the topic, but many years later I saw one of those Olympiad questions they inflict on students, and the question was, which has the higher boiling point: methyl cyclopropane or cyclobutane? When I saw this question, I thought it unreasonable, because while the former has a small dipole moment, the other has more degrees of vibrational freedom. I mentally picked the cyclobutane, but the answer the students were supposed to give was methyl cyclopropane, because of the dipole moment. The problem is, observation shows that answer is just plain wrong! People who set questions like that ought to be taken out and identified, and made to wear placards saying, "check the literature!"
Posted by Ian Miller on Dec 8, 2012 12:55 AM GMT
When writing up, I had to explain why my results were contradicting the emerging paradigm. How would you enjoy that, when future employment depends on success? My first step was to define some terms, and I interpreted "conjugation" to mean a system where a common wave function extended over more than two atoms. In particular, if cyclopropyl (Cy) conjugated with unit A, there was at least one two-electron wave function in the system Cy-A. Now, suppose there was a unit B; if Cy conjugated, then there was at least one two-electron wave function in the system Cy-B.  Thus in hexatriene, there are wave functions extending over the entire system. Now, I reasoned that if Cy-A was a delocalized system, through the virial theorem and the fact that a solution consistent with the stationary state Schrodinger equation can have only one energy, the wave function extending over the system had a common potential. The same applied to Cy-B. Therefore, if Cy conjugates, within this definition if substitution changed the potential at B, there must be a change at A. I had shown there was no mesomeric change, therefore there was no linking wave function. Now, at the very least, I had to explain how cyclopropane stabilized adjacent positive charge, why did cyclopropane give significant bathochromic shifts to certain UV spectra, and why was the dipole moment of cyclopropyl chloride reduced compared with other alkyl chlorides? Fortunately, if the first was answered, so was the second. The reason is, from Maxwell’s electromagnetic theory, light can only be absorbed if there is a change of electric moment, and if cyclopropyl stabilizes adjacent positive charge, it will stabilize the excited state when positive charge is adjacent, and by doing so, it will also increase the extinction coefficient. (In this respect, I think Pauling’s canonical structures/VB approach make these effects so much easier to understand at a lower level than MO theory.)
It was then that I saw what I believed was the answer. Coulson and Moffatt had proposed bent bonds. If we consider the carbon atoms to have orbitals in the sp3 configuration that overlap, in a bond, the forces from the opposing nucleus pull the orbital closer to the C-C line, the movement being opposed by the remaining four ring electrons, and an equilibrium was reached whereby the charge density of each bond was moved some distance towards the centroid of the ring. This movement of electrons might be symmetrical in cyclopropane, but it was not with respect to a substituent. Four lobes on the two distal carbon atoms moved more or less directly towards the substituent, but the two lobes on the geminal carbon atom changed their angles, but did not significantly change the electron distance to the substituent.
There were two ways of viewing this, and I probably chose the harder of the two, although in the event it was the way with the greatest ability to explain a wider range of observations. Both depend on the obvious: the potential energy is stored in electric fields. (That is due to Maxwell.) The first is to note that the repulsion energy on the four distal lobes will be overridden by the positive charge on the substituent. If so, the stabilization energy of a carbenium ion should not exceed 2/3 the strain energy, together with a standard polarization energy of alkane bonds. The second way of looking at it is that the four distal lobes move towards the positive charge and stabilize it. In short, at least qualitatively, the stabilization, in excess of that given by standard alkyl groups, of adjacent positive charge by the cyclopropyl group is required by Maxwell’s electromagnetic theory, and there is no necessity for special quantum effects.
I proposed to explain the electric moment of cyclopropyl chloride this way. Assume the linear combination of atomic wave functions. (Quantum mechanics is a linear theory. Waves always combine in a linear fashion.)  The reason alkyl chlorides have a dipole moment (as I saw it then) is that on wave interference, each orbital tries to increase its electron density by a given proportion, but the electron density is much higher around the chlorine atom before wave interference, so electron density has to move towards the chlorine after it. The higher electron density in the cyclopropane ring, which I referred to as a monopole, partly offsets that. So I had my explanations. I could write up the conclusions.
However, I had a minor problem: I had nobody to beta test what I was writing, because supervisor had gone to North America. This was a serious problem because I had nobody to check whether I might be going wrong somewhere, and nobody to check whether what I was writing was readily comprehensible. I was effectively saying everyone else was wrong. What should have happened was that supervisor should have sat me down in front of a physicist and straightened out what I was thinking so I could put it in a more orthodox form. That did not happen.
Posted by Ian Miller on Dec 1, 2012 2:09 AM GMT
At this point I had difficulties, but since I had made some new amines in order to measure their acid-base equilibrium constants, I did just that. Then, I measured some rate constants for the reaction of the amines with dinitrochlorobenzene. I then had two Hammett lines, but none with mesomeric withdrawing substituents because the para cyano substituent had also refused to make the amine. Part of the reason was because the very aggressive conditions needed to stop urea formation also tended to hydrolyse the cyano group. I was now back to the Hammett rho factor. We knew the rho value for anilines and benzylamines, so we could estimate that for 2-phenylethylamines. The cyclopropyl group enhanced the rho value by about 30% over the ethyl group, which was consistent with my concept of conductance and two routes. This indicated (to me, anyway) an absence of conjugation, but it was hardly a definitive result. Meanwhile, those wretched styrenes had still not put in an appearance, but of course without the mesomeric withdrawing groups it probably did not matter. I had a rho value, and it appeared that it would be impossible to get such amines consistent with having mesomeric withdrawing substituents. (I have no doubt that a more skilled synthetic chemist may have had more success here, but not with the para nitro derivative. I found it interesting that somebody tried to make this a decade later, and reported the same trouble.)
What now? Suddenly, supervisor put in a cameo appearance. (The 9 month sabbatical had extended to over 16 months, by my calculations.) He had an idea: take the acids, which I had already made and measure the rates of reaction with diphenyldiazomethane. Nobody measured rates of reaction of acids with this material, so get data, pad out the thesis! I was a bit skeptical about this comment, but then again, getting a PhD did have a certain attraction, so off I went. Since apparently nobody had bothered to carry out this with any other acids, I also had to measure the reaction rates with benzoic acids, so I could make some estimate of what the rho values should be, and of course to get an idea of how much mesomeric effect could be expected from substituents. The benzoic acids were available, so this meant plenty of measurements. If nothing else, thesis padding would ensue!
The reaction of the acids with diphenyldiazomethane depends on the acidity, so I expected this to merely reproduce existing   data on the phenylcyclopropanecarboxylic acids, but suddenly I realized that supervisor had made a key advance. The reactions were carried out in toluene, and this amplified the differences in acidity. The net result was that all the rho values were much bigger than anybody else had managed to get, and this meant that experimental errors did not have such significance with respect to sigma values. The rho value for the cyclopropane ring was again roughly where I expected it, but now I had some mesomeric donors, the para methoxy and the para fluoro substituent, and furthermore, the para nitro and para cyano substituted acids were available to anchor the line at a greater distance from the rest of the substituents.
Joy to the world! These results were quite unambiguous: the para substitution offered no mesomeric donating properties. I had a conclusion! Thank you, supervisor! Unfortunately, this brought a major problem: the scientific community was coming to the conclusion that cyclopropane conjugated with adjacent unsaturation, and here was me with results saying it did not. The reason for the conclusion that it did were:
(a) Hammett rho values for 2-phenylcyclopropyl-X were about 30% higher than those for 2-phenylethyl-X
(b) Cyclopropyl stabilized adjacent positive charge
(c) Cyclopropyl adjacent to a chromophore such as a benzene ring gave a bathochromic shift to UV signals.
(d) MO theory, and in particular CNDO/2 computations, said it conjugated.
(e) Dipole moments of molecules such as cyclopropyl chloride were about 0.3 Debye less than, say, isopropyl chloride.
All was not bad, because there were also data that indicated that cyclopropane did not conjugate. There were also some almost confusing options, thus the fact that the C-Cl bond in cyclopropyl chloride was slightly shorter than alkyl chlorides could be explained in terms of changed hybridization, but what did that mean in relation to conjugation? So, I had a further problem: I had to write up. No problem with much of the thesis, until the conclusions, where I had to deal with (a) to (e). My choices were simple: I had to show that my results were wrong (most undesirable because that was guaranteed failure, and anyway, I backed them) or irrelevant (again, undesirable, as why had I chosen this project?) or I had to find an alternative explanation for (a) to (e). Welcome to the start of an unusual career! I was happy that I had (a) under control, but . . . To add to my problems, supervisor disappeared again, apparently seeking a better paying job in North America.
How many of other PhD students had to take on an essentially emerging scientific consensus amongst those who had influence, essentially on their own? The problem was, how to do it?
Posted by Ian Miller on Nov 23, 2012 6:01 AM GMT
Ever had trouble making something? Yes, I had further problems. Recall, I needed a sequence of substituted trans-2-phenylcyclopropylamines, and I would get the amine through the carboxylic acid. My aim was to tell whether the para nitro substituent would give the sigma constant typical of conjugation, or no conjugation, which meant I needed the para nitro substituent, and I would prefer the meta nitro substituent to "anchor" the Hammett line.
The para nitro compound would be a straightforward nitration, or so I thought, of the carboxylic acid. The literature suggested concentrated nitric acid at 30 degrees C, to get an awful yield, so I tried fuming nitric in an ice bath. The product melted at 170 degrees C, not the required 198 degrees, and was shown to be the then unknown dinitrophenyl cyclopropanecarboxylic acid. Nitric acid diluted with acetic acid kept below 5 degrees for five hours gave a 75%  yield of the right stuff. So far, so good.
The meta nitro compound was obtainable, in principle, through my standard synthesis, except that after the Friedel Crafts reaction of succininc anhydride with benzene, I nitrated the resultant ketone, and ended up with the meta nitro substituted keto acid. Easy! Then all went well until the cyclization: a zero yield of cyclopropanecarboxylic acid, and a 100 % yield of shog, aka dirty brown rubbish.
It was around about now that supervisor went overseas on a 9 month sabbatical. No advice about what to do. Anyway, I thought I had better make the amines, because these wretched styrenes that I had ordered some time ago refused to turn up. Amine conversion went well, until, disaster, the para nitro compound refused to behave. The Curtius reaction went well enough, and I made the hydrochloride of trans-2-p-nitrophenylcyclopropylamine in good yield. If I dumped this into strong alkali, I got the amine (at least I believe so) but under any conditions near neutral pH, where there could be some of both protonated and base forms present, there was an immediate production of a reddish brown condensed product. By definition, the measurement of a dissociation constant requires both to be present. Using it as a nucleophile to carry out a displacement generally permits some of the protonated form to be present, as you cannot carry out kinetics of a nucleophilic displacement with molar NaOH present. Now what?
My nearest piece of inspiration was to risk some of my p-bromo substituted acid and try to carry out a nucleophilic substitution with cyanide. In this reaction, the cyclopropyl ring retards the reaction, but it worked. The yield may not have been magic, but it was adequate. Whew! Not so fast! Conversion to the amine was again, troublesome. I was in trouble. I was also on my own; supervisor remained absent, not even mail, and nobody else was offering help.
One final piece of amusement, that could have been in the previous blog. As a reward for supervising senior organic undergraduate labs, I was permitted to get some non-critical synthesis done for me, so I got a student to try the Friedel Crafts reaction of succinic anhydride on ethyl benzene. The student worked well (I supervised the first part very closely!) and eventually the student proudly announced he had a good yield of white material. I suggested he do a melting point, so we could see how pure it was. Some time later, he was still measuring the melting point, so I asked him what was taking so long. "It isn't melting," he replied. A quick flame test told me all I needed to know: he had thrown out the product and kept the aluminium oxide. It would take some time to measure the melting point of that!
Posted by Ian Miller on Nov 17, 2012 2:29 AM GMT
As most people who have undertaken a PhD in organic chemistry know, synthesis can take a long time, even when the objective is not primarily synthesis, so to get a sense of proportion in this account, this seems to be time for an interlude, wherein some various tales of, well, you can choose your word, can be told. So, here are some things that should not have happened.
The prequel, and warning that not all is as it should be, happened in the summer when I had been hired for my first taste of research. I came back from lunch and saw water on the floor from the bench behind me. A hose had come off a small condenser, and a small "river" of water was snaking across the bench. By itself, a nuisance, but in two dry bits, enveloped almost by the winding "river", were two sizable lumps of sodium sitting there! Fortunately, I knew where some tongs were, so tongs, then tap, then when the perpetrator returned, some language not befitting an output of the RSC!
I started the PhD in a small “temporary” building called “The Armoury”, a sort of overflow lab because the main chemistry building was “full”. One advantage was it was near the Students’ Union cafeteria, so coffee was at hand. Soon after starting, my supervisor announced he was to do some lab work, a statement not greeted with unqualified enthusiasm by the senior PhD students. His first move was to purify a very large amount of acetic anhydride by distillation, but after a short period into this exercise, there was a yell, an emergency wash,  and a very strong smell of acetic anhydride. No boiling chip or stirring, and not a lot of work for the rest of the afternoon. There was no repeat of this enthusiasm for lab work! It was also suggestive that good advice for lab work might not be forthcoming.
One morning I came in and saw a remarkable sight. There was a small room at one end directly facing the path to the Student’s Union, and on a bench on the far side from the window, someone had carried out a sealed tube pressure reaction. The top had blown, there was a 2 mm hole drilled through the safety shield, and in a direct straight line from the tube, hole in shield, there was a further 2 mm hole in the window about 3 meters away. That did not put me off pressure reactions (I have done quite a number through my career) but I always used steel.
After a year I was moved to the top floor of the chemistry building, and old stone building with basically one entrance. At some point, a Colombo Plan student was brought in and put at the bench behind me. Since nobody knew what his skill level was, I was asked to offer assistance. His first synthesis involved reacting a ketone with a Grignard reagent, so I discussed what would happen and set him off. I had to help him get it started, but soon the magnesium was reacting well. About two hours later, I heard the bleat, "I've got two layers." I pointed out that was expected when he poured his ether into the acidified ice-water. "But I haven't done that yet." Somewhat annoyed, I had to go around the other side of the island to see what he had done, and there were two layers! Had I tried, I could not have done that in a hundred years! He had managed to add the ketone as an ethereal layer on top of the Grignard! His hand shuddered, and as I dived for the floor there was a roar like a rocket motor (OK- slight exaggeration). Fortunately, there was no gas lit in the room. So, I gave him a lecture about the need for stirring or refluxing, and he was at it next day. About three in the afternoon, another bleat, "I've still got two layers!" I succeeded in getting him removed from the lab.  More next week.
Posted by Ian Miller on Nov 9, 2012 7:44 AM GMT
My supervisor finally reappeared, fresh from prolonged summer holidays, and for some reason I could not understand, my brilliant project was not greeted with unqualified enthusiasm. Just maybe he saw problems that eluded my somewhat innocent/inexperienced state but somehow he accepted and so I was off to the bench. The good news was that trans 2-phenylcyclopropylamine was made commercially; it is a monoamine oxidase inhibitor. (Memo to self: emphasize hygiene and clean working! Recalling that baggage, I did not want my brain messed up.) The simplest route was through the Curtius reaction from acids (acid chloride, nucleophilic substitution with azide, thermally decompose azide, acid hydrolysis of isocyanate). So, the problem now was to make the carboxylic acids. The most obvious route was rather long and only gave para substitution: Friedel Crafts with succinic anhydride, borohydride reduction to give the lactone, thionyl chloride to give the 4-(substituted phenyl)-4-chlorobutyroyl chloride, which was then esterified with ethanolic HCl, then cyclization with sodium t-amylate, purified by oxidizing out the olefinic material with alkaline KMnO4. Long-winded, an interesting filtration problem, but seemingly free of big problems.
There were two reasons to look for something else. The first was to get something quicker, while the second was to get some meta substitution. The literature said that ethyl diazoacetate on styrene gave a 5% yield of the desired compound. That did not impress me. I tried the zinc-copper couple/di-iodomethane on ethyl cinnamate, but no significant yield of desired product resulted. This was not unexpected, because that reaction was known to be difficult when the olefin was electron poor. The reaction of the ylide from trimethylsulphoxonium iodide on ethyl cinnamate did not work either, although a little help here from supervisor could have come in handy. The reference was from Corey, in a Tet. Lett. It was only much later, in a J. Org Chem., did I learn that the reaction only really works at minus 80 degrees. Silly me, I never thought of that, but a more experienced chemist might have asked, what were those bubbles – try cooling it until they stop. However, I did get a reasonable yield with alkali, triethylphosphonoacetate and styrene oxide. So, I immediately ordered some meta substituted styrenes. The basic problem with this project, of course, was that the electron withdrawing substituent that could show mesomeric effects would be para nitro, and its sigma value was some distance from the others in the line, which meant that the line had to be well-anchored so that the extrapolated "non-mesomeric" value had as little error range as possible.
The best way to avoid this difficulty was to make the meta nitro substitution, so while making a number of compounds by the "slow" route while waiting for styrenes, I made a lot of the simple reaction of succinic anhydride with benzene, and nitrated it. That gave me a lot of meta nitro substituted acid, so then I took that down the route to the cyclopropanecarboxylic acid, but when trying to form the cyclopropane ring, disaster! There was essentially a 100% yield of dirty brown rubbish. Supervisor had no suggestions, and it appeared that I was down to extrapolation, or doing something spectacular with one of the styrenes. Which raised the question, a year later, just where were those styrenes? It was around about now I started to get nervous. It was also somewhere about now that supervisor went off on a 9-month sabbatical. More soon!
Posted by Ian Miller on Nov 2, 2012 11:44 PM GMT
Again, my theories on planetary formation survive another month, although if I were rewriting I would amend the literature survey a little. There were three papers of particular significance. 
Paniello (Nature, 490: 376) examined zinc isotope evidence from lunar rocks and concluded that the moon had an additional condensation step than Terran rocks. This is in accord with the generally accepted theory that the Moon was formed by the condensate from a collision of a body called Theia with Earth. By itself, this is unexceptional, however there was also a comment by Elliot (Nature, 490: 346) in which he notes that collisional models suggest the Moon should be made predominantly from material originating from Theia, in which case isotope distributions of non-volatile elements should match Theia, but they are essentially identical to those of Earth. Elliot suggests that this requires an extremely rapidly rotating Earth prior to collision, which seems unlikely. Elliot overlooked the option that Theia accreted at an Earth-sun Lagrange point (Belbruno and Gott. 2005. Astron. J. 129: 1724) either L4 or L5, in which case Theia would have the same isotopes. I favour that interpretation, mainly because my theory argues that Earth forms at the most favourable temperature for rocky accretion.
Shcheka and Keppler (Nature 490: 531) did me a favour. One big problem in accounting for earth's atmosphere is why is xenon depleted compared with argon, and to a lesser extent, krypton. One answer is that the initial atmosphere suffered strong hydrodynamic escape to space, which led to an enhancement of heavy xenon isotopes, but that should remove almost all the argon. These authors found that perovskite, which makes up much of the Earth's mantle, can dissolve up to 1% argon. The reason is, anomalies in perovskite (MgSiO3) arise through elements like aluminium getting in, which create holes that are roughly the same size as argon. Xenon, being bigger, does not fit. Under this scenario, the early hydrodynamic escape (powered by intense solar radiation on an atmosphere of retained accretion disk gases) of hydrogen and helium drags off most of the other elements, and subsequent argon is released, without heavy isotope enhancement, by volcanic degassing.
Cassata et al. (Icarus 221: 461) determined isotope ratios of trapped argon from the Martian meteorite ALH84001 and concluded that the atmospheric pressure on Mars at 4.16 Gy BP was < 400 mbar, and accordingly a CO2 atmosphere could not have had sufficient pressure to have sufficient greenhouse gas to permit water to flow. This strongly supports my theory, in which it was ammonia that dissolved in water and lowered the melting point. At first I got excited, because these authors used a C/N ratio that only made sense if the nitrogen ended up underground. That would strongly support my theory, but unfortunately it did not. Closer reading showed they assumed the C/N ratio.
Finally, you may wonder why I got involved with planetary formation theory. In the early 1990s, thanks to a persistent economic downturn, I had some spare time, so I wrote a science-fiction book about the colonization of Mars, in which the bad people were intending to get rich by floating junk shares/stock on Earth. (There were plenty of examples of fraud to learn from in the previous decade!) To expose the fraud, I needed an unexpected discovery, and in my background, I had published one paper arising from the idea that the CO2 atmosphere beloved of geologists would lead to basalt weathering and giving ferrous ions, which, in turn, could photochemically reduce CO2. We did some experiments and it does, but then I became concerned; what happened to the ferric ions? A few experiments showed that ferric ions are very aggressive at photochemically attacking carbohydrates and amino acids. Somehow, CO2 being part of the origin of life became much less attractive.
Accordingly, for my required "unexpected discovery", I tried for a reduced atmosphere, which would lead to massive underground deposits of urea. (Yes, I know it might go further, but . . .) My agent managed to persuade the editor of a major publisher to look at it, but the editor died. The replacement cleared the desk, and rejected my novel on the grounds it was too implausible. I was fairly confident "colonization of Mars" was not too bad for SciFi, with big money around, fraud is hardly implausible, so that left the reduced atmosphere. How dare a literary editor trash my chemistry! So I became involved. The experts say that thanks to UV radiation, ammonia would only last decades and so is irrelevant, nevertheless the only available sample of ocean from about 3.2 Gy BP has levels of ammonia in it approaching those of potassium. I back observations over "experts", even if the observers did not realize the significance of what they found. Incidentally, if anyone is interested, the offending book, Red Gold is now available as an ebook on Amazon, and I have included a précis of the theory in an appendix.
Posted by Ian Miller on Oct 27, 2012 12:24 AM BST
Mid December, and time to select a supervisor and project. That was easy; I selected my supervisor on the grounds he was enthusiastic over a project. I did my background reading, got some glassware to my bench, and was all ready to start, but with two days before time to go away for Christmas, I headed for the library. Then, for a Christmas present, I gave my supervisor the results of the project: it was neatly written up in the latest JACS. Oops! Worse, there was no safety net; apart from these measurements there was nowhere to go with the compounds.
Back from Christmas, my supervisor gave me two new projects. One was to examine solvolyses of a number of substituted acetylacetonates of various metals, to see the degree of electronic effects that could be related to various metals. That seemed a great project, until the literature indicated that the rate constants of at least some were zero. Not promising! The alternative project involved fusing a cyclopropanecarboxylic acid to the 9-10 position of phenanthrene, and examine the possibility of conjugative effects of substitution on the phenanthrene on acidity. At the time, there was a debate as to whether cyclopropane transmitted conjugative effects.
I saw difficulties. To get a starting material, one took 3 kg of refluxing phenanthrene and carefully added about 700 g of ethyl diazoacetate while avoiding using scratched glassware, etc. If things went wrong, bang! The report indicated that it was possible to extract about 25 g of product from the 3.6 kg of resultant tar. As if that were not bad enough, how did one get substitution? Trying to make 3 kg of a substituted phenanthrene did not seem to be particularly attractive, especially since the most reactive carbons of phenanthrene itself are 9 and 10. Substituting the basic product was possible, but there was a risk the 2-position (para to the other phenyl) would be the most reactive. Accordingly, I was fairly dejected when the Head of Department met me walking towards the department. When I explained, he smiled and said, “Why don’t you select your own project?” That was probably just to get rid of me, but I took him up on this offer. (Note to self at time: supervisor not functioning on all cylinders, and most of time so far, he was somewhere else! Little did I know I was getting into the groove!)
At that time, two results had come in from the dissociation of a sequence of 2-phenycyclopropane carboxylic acids, employing the Hammett equation. Recall my baggage? I knew about that. Now the interesting thing about this was the two determinations, one in water and one in aqueous ethanol, came to opposite conclusions! Further, the conclusions were based on the Hammett rho values, which were known to attenuate substitution effects by about a third for every saturated carbon in the chain, but full conjugation, e.g. a double bond, showed little attenuation. The determination in water had substitution effects attenuated by the same as two methylene units, i.e. no conjugation, while the aqueous ethanol results had the rho value about 30%  higher, which led to the conclusion that cyclopropane conjugated about 30-50% as well as a vinyl, as in cinnamic acids. What do you think about that?
I regarded this as hopelessly naïve. Suppose whatever induction measured acted like a fluid. An example might be pushing electron charge, or, dare I think it at the time, enhancing probability? Cyclopropane has two routes to go through, therefore, when you add the second route, it should be 30% higher anyway. More interesting were the sigma values; these represent the ability of the substituent to "push or pull" charge/potential, and each has an inductive value, which is enhanced if the substituent can participate in a mesomeric effect. A quick graphing of the carboxylic acid results showed the rho values were too small to be sure, but if we discarded the para nitro substitution, the water results were also consistent, but with worse scatter, with the rho value being 30% higher. Discarding inconvenient results, I here you say. Well, yes, but there was a reason. The paranitro compound is almost totally insoluble in water, the others are not a lot better, so I considered that it was doubtful if true equilibrium was reached. The ethanol results would be better. The basic problem with considering sigma values was that rho was too small that experimental uncertainty made a conclusion unreliable, but what you could see indicated, based on para methoxy, that there was very little or no conjugation.
So, I could see a project: measure the dissociation constants of the 2-phenylcyclopropylamines. Amines, with one less atom in the road, have higher rho values, which I knew from my previous summer work. Baggage at work for me! So, all I had to do was wait for supervisor to reappear. More to come.
Posted by Ian Miller on Oct 19, 2012 11:07 PM BST
< Prev    1 2 3 4