Moving between conscious calculation and hunch, between intuition and analysis, Mendeleev arrived within a few weeks at a tabulation of thirty-odd elements in order of ascending atomic weight, a tabulation that now suggested there was a recapitulation of properties with every eighth element. And on the night of February 16, 1869, it is said, he had a dream in which he saw almost all of the known elements arrayed in a grand table. The following morning, he committed this to paper.«43»
The logic and pattern of Mendeleev’s table were so clear that certain anomalies stood out at once. Certain elements seemed to be in the wrong places, while certain places had no elements. On the basis of his enormous chemical knowledge, he repositioned half a dozen elements, in defiance of their accepted valency and atomic weights. In doing this, he displayed an audacity that shocked some of his contemporaries (Lothar Meyer, for one, felt it was monstrous to change atomic weights simply because they did not ‘fit’).
In an act of supreme confidence, Mendeleev reserved several empty spaces in his table for elements ‘as yet unknown.’ He asserted that by extrapolating from the properties of the elements above and below (and also, to some extent, from those to either side) one might make a confident prediction as to what these unknown elements would be like. He did exactly this in his 1871 table, predicting in great detail a new element (‘eka-aluminium’) which would come below aluminium in Group III. Four years later just such an element was found, by the French chemist Lecoq de Boisbaudran, and named (either patriotically, or in sly reference to himself, gallus, the cock) gallium.
The exactness of Mendeleev’s prediction was astonishing: he predicted an atomic weight of 68 (Lecoq got 69.9) and a specific gravity of 5.9 (Lecoq got 5.94) and correctly guessed at a great number of gallium’s other physical and chemical properties – its fusibility, its oxides, its salts, its valency. There were some initial discrepancies between Lecoq’s observations and Mendeleev’s predictions, but all of these were rapidly resolved in favor of Mendeleev. Indeed, it was said that Mendeleev had a better grasp of the properties of gallium – an element he had never even seen – than the man who actually discovered it.
Suddenly Mendeleev was no longer seen as a mere speculator or dreamer, but as a man who had discovered a basic law of nature, and now the periodic table was transformed from a pretty but unproven scheme to an invaluable guide which could allow a vast amount of previously unconnected chemical information to be coordinated. It could also be used to suggest all sorts of research in the future, including a systematic search for ‘missing’ elements. ‘Before the promulgation of this law,’ Mendeleev was to say nearly twenty years later, ‘chemical elements were mere fragmentary, incidental facts in Nature; there was no special reason to expect the discovery of new elements.’
Now, with Mendeleev’s periodic table, one could not only expect their discovery, but predict their very properties. Mendeleev made two more equally detailed predictions, and these were also confirmed with the discovery of scandium and germanium a few years later.«44» Here, as with gallium, he made his predictions on the basis of analogy and linearity, guessing that the physical and chemical properties of these unknown elements, and their atomic weights, would be between those of the neighboring elements in their vertical groups.«45»
The keystone to the whole table, curiously, was not anticipated by Mendeleev, and perhaps could not have been, for this was not a question of a missing element, but of an entire family or group. When argon was discovered in 1894 – an element which did not seem to fit anywhere in the table – Mendeleev denied at first that it could be an element and thought it was a heavier form of nitrogen (N3, analogous to ozone, 03). But then it became apparent that there was a space for it, right between chlorine and potassium, and indeed, for a whole group coming between the halogens and the alkali metals in every period. This was realized by Lecoq, who went on to predict the atomic weights of the other yet-to-be-discovered gases – and these, indeed, were discovered in short order. With the discovery of helium, neon, krypton, and xenon, it was clear that these gases formed a perfect periodic group, a group so inert, so modest, so unobtrusive, as to have escaped for a century the chemist’s attention.«46» The inert gases were identical in their inability to form compounds; they had a valency, it seemed, of zero.«47»
The periodic table was incredibly beautiful, the most beautiful thing I had ever seen. I could never adequately analyze what I meant here by beauty – simplicity? coherence? rhythm? inevitability? Or perhaps it was the symmetry, the comprehensiveness of every element firmly locked into its place, with no gaps, no exceptions, everything implying everything else.
I was disturbed when one enormously erudite chemist, J.W. Mellor, whose vast treatise on inorganic chemistry I had started dipping into, spoke of the periodic table as ‘superficial’ and ‘illusory,’ no truer, no more fundamental than any other ad hoc classification. This threw me into a brief panic, made it imperative for me to see if the idea of periodicity was supported in any ways beyond chemical character and valency.
Exploring this took me away from my lab, took me to a new book that immediately became my bible, the CRC Handbook of Physics and Chemistry, a thick, almost cubical book of nearly three thousand pages, containing tables of every imaginable physical and chemical property, many of which, obsessively, I learned by heart.
I learned the densities, melting points, boiling points, refractive indices, solubilities, and crystalline forms of all the elements and hundreds of their compounds. I became consumed with graphing these, plotting atomic weights against every physical property I could think of. I became more and more excited, exuberant, the more I explored, for almost everything I looked at showed periodicity: not only density, melting point, boiling point, but conductivity for heat and electricity, crystalline form, hardness, volume changes with fusion, expansion by heat, electrode potentials, etc., etc. It was not just valency, then, it was physical properties, too. The power, the universality of the periodic table was increased for me by this confirmation.
There were exceptions to the trends shown in the periodic table, anomalies, too – some of them profound. Why, for example, was manganese such a bad conductor of electricity, when the elements on either side of it were reasonably good conductors? Why was strong magnetism confined to the iron metals? And yet these exceptions, I was somehow convinced, reflected special additional mechanisms at work, and in no sense invalidated the overall system.«48»
Using the periodic table, I tried my hand at prediction too, trying to predict the properties of a couple of still-unknown elements as Mendeleev had done for gallium and the others. I had observed, when I first saw the museum table, that there were four gaps in it. The last of the alkali metals, element 87, was still missing, as was the last of the halogens, element 85. Element 43, the one below manganese, was still missing, though this space read ‘?Masurium’ with no atomic weight.«49 »Finally there was a rare earth, element 61, missing too.
It was easy to predict the properties of the unknown alkali metal, for the alkali metals were all very similar, and one had only to extrapolate from the other elements in the group. 87, I reckoned, would be the heaviest, most fusible, most reactive of them all; it would be a liquid at room temperature, and like cesium have a golden sheen. Indeed, it might be salmon pink, like molten copper. It would be even more electropositive than cesium, and show an even stronger photoelectric effect. Like the other alkali metals, it would color flames a striking color – probably a bluish color, since the flame colors from lithium to cesium tended in this direction.
It was equally easy to predict the properties of the unknown halogen, for the halogens, too, were very similar, and the group showed simple, linear trends.
But predicting the properties of 43 and 61 would be trickier, for these were not ‘typical’ elements (in Mendeleev’s term). And it was precisely with such nontypical elements that Mendeleev had run into trouble, leading him to revise his original table. The transition metals had a sort of ho
mogeneity. They were all metals, all thirty of them, and most of them, like iron, were hard and tough, dense and infusible. This was especially so of the heavy transition elements, like the platinum metals and filament metals Uncle Dave had introduced me to. My interest in color brought home another fact, that where compounds of typical elements were usually colorless, like common salt, the compounds of transition metals often had vivid colors: the pink minerals and salts of manganese and cobalt, the green of nickel and copper salts, the many colors of vanadium; going with their many colors were their many valencies, too. All these properties showed me that the transition elements were a special sort of animal, different in nature from the typical elements.
Still, one might hazard a guess that element 43 would have some of the characteristics of manganese and rhenium, the other metals in its group (it would, for instance, have a maximum valency of 7, and form colored salts); but it would also be generically similar to the neighboring transition metals in its period – niobium and molybdenum to the left, and the light platinum metals to the right. So one could also predict that it would be a shining, hard, silvery metal with a density and melting point similar to theirs. It would be just the sort of metal Uncle Tungsten would love, and just the sort of metal which would have been discovered by Scheele in the 1770s – that is, if it existed in sensible amounts.
The hardest prediction, in any detail, would be for element 61, the missing rare earth metal, for these elements were in many ways the most baffling of all.
I think I first heard of the rare earths from my mother, who was a chain smoker and lit cigarette after cigarette with a small Ronson lighter. She showed me the ‘flint’ one day, pulling it out, and said it was not really flint, but a metal that produced sparks when it was scratched. This ‘mischmetal’ – cerium mostly – was a mishmash of half a dozen different metals, all of them very similar, all of them rare earths. This odd name, the rare earths, had a mythical or fairy-tale sound to it, and I imagined the rare earths as not only rare and precious, but as having special, secret qualities possessed by nothing else.
Later Uncle Dave told me of the extraordinary difficulty which chemists had had in separating the individual rare earths – there were a dozen or more – for they were astoundingly similar, at times indistinguishable in their physical and chemical properties. Their ores (which for some reason all seemed to come from Sweden) never contained a single rare-earth element, but a whole cluster of them, as if nature herself had trouble distinguishing them. Their analysis formed a whole saga in chemical history, a saga of passionate research (and frequently frustration) in the hundred years or more it took to identify them. The separation of the last few rare-earth elements, indeed, was beyond the powers of chemistry in the nineteenth century, and it was only with the use of physical methods such as spectroscopy and fractional crystallization that they were finally separated. No fewer than fifteen thousand fractional crystallizations, exploiting the infinitesimal differences in solubility between their salts, were needed to separate the final two, ytterbium and lutecium – an enterprise that occupied years.
Nonetheless there were chemists who were enthralled with the intransigent rare-earth elements and spent their entire lives trying to isolate them, sensing that their study might cast an unexpected light on all the elements and their periodicities:
The rare earths [wrote William Crookes] perplex us in our researches, baffle us in our speculations, and haunt us in our very dreams. They stretch like an unknown sea before us, mocking, mystifying, and murmuring strange revelations and possibilities.
If the rare-earth elements baffled, mocked, and haunted chemists, they positively maddened Mendeleev as he struggled to assign them a place in his periodic table. There were only five rare earths known when he constructed his first table in 1869, but then more and more were discovered in the decades that followed, and with each discovery the problem grew, because all of them, with their consecutive atomic weights, belonged (it seemed) in a single space in the table, crushed, as it were, between two adjoining elements in Period 6. Others, too, struggled with the placement of the maddeningly similar elements, further frustrated by a deep uncertainty as to how many rare-earth elements there might ultimately prove to be.
Many chemists, by the end of the nineteenth century, were inclined to put both the transition and the rare-earth elements into separate ‘blocks,’ for one needed a periodic table with more space, more dimensions, to accommodate these ‘extra’ elements that seemed to interrupt the basic eight groups of the table. I tried making different forms of periodic table myself to accommodate these blocks, experimenting with spiral ones and three-dimensional ones. Many others, I later found, had done the same: more than a hundred versions of the table appeared during Mendeleev’s lifetime.
All of the tables I made, all of the tables I saw, ended with uncertainty, ended with a question mark, centered around the ‘last’ element, uranium. I was intensely curious about this, about Period 7, which started with the as-yet-unknown alkali metal, element 87, but only got as far as uranium, element 92. Why, I wondered, should it stop here, after only six elements? Could there not be more elements, beyond uranium?
Uranium itself had been placed by Mendeleev under tungsten, the heaviest of the Group VI transition elements, for it was very much like tungsten, chemically. (Tungsten formed a volatile hexafluoride, a very dense vapor, and so did uranium – this compound, UF6, was used in the war to separate out the isotopes of uranium.) Uranium seemed like a transition metal, seemed like eka-tungsten – and yet, I felt somehow uncomfortable about this, and decided to do a little exploring, to examine the densities and melting points of all the transition metals. As soon as I did this I discovered an anomaly, for where the densities of the metals steadily increased through Periods 4, 5, and 6, they unexpectedly declined when one came to the elements in Period 7. Uranium was actually less dense than tungsten, though one would have expected it to be more so (thorium, similarly, was less dense than hafnium, not more so, as one would have expected). It was precisely the same with their melting points: these reached a maximum in Period 6, then suddenly declined.
I was excited about this; I felt I had made a discovery. Was it possible, despite all the similarities between uranium and tungsten, that uranium did not in fact belong in the same group, was not even a transition metal at all? Might this also be the case for the other Period 7 elements, thorium and protoactinium, and the (imaginary) elements beyond uranium? Could it be that these elements were instead the beginning of a second rare-earth series precisely analogous to the first one in Period 6? If this was the case, then eka-tungsten would not be uranium, but an as-yet-undiscovered element, which would appear only after the second rare-earth series had completed itself. In 1945, this was still unimaginable, the stuff of science fiction.
I was thrilled, soon after the war, to find that I had guessed right, when it was revealed that Glenn Seaborg and his coworkers in Berkeley had succeeded in making a number of transuranic elements – elements 93, 94, 95, and 96 – and found that these indeed were part of a second series of rare-earth elements (which, by analogy with the first rare-earth series, the lanthanides, he called the actinides).«50»
The number of elements in the second series of rare earths, Seaborg argued, by analogy with the first series, would also be fourteen, and after the fourteenth (element 103) one might expect ten transition elements, and only then the concluding elements of Period 7, ending with an inert gas at element 118. Beyond this, Seaborg suggested, a new period would start, beginning, like all the others, with an alkali metal, element 119.
It seemed that the periodic table might thus be extended to new elements far beyond uranium, elements that might not even exist in nature. Whether there was any limit to such transuranic elements was not clear: perhaps the atoms of such elements would become too big to hold together. But the principle of periodicity was fundamental, and could be extended, it seemed, indefinitely.
While Mendeleev saw the periodic ta
ble primarily as a tool for organizing and predicting the properties of the elements, he also felt it embodied a fundamental law, and he wondered on occasion about ‘the invisible world of chemical atoms.’ For the periodic table, it was clear, looked both ways: outward to the manifest properties of the elements, and inward to some as-yet-unknown atomic property which determined these.
In that first, long, rapt encounter in the Science Museum, I was convinced that the periodic table was neither arbitrary nor superficial, but a representation of truths which would never be overturned, but would, on the contrary, continually be confirmed, show new depths with new knowledge, because it was as deep and simple as nature itself. And the perception of this produced in my twelve-year-old self a sort of ecstasy, the sense (in Einstein’s words) that ‘a corner of the great veil had been lifted.’
CHAPTER SEVENTEEN
A Pocket Spectroscope
We had always celebrated Guy Fawkes Night, before the war, by setting off fireworks. Bengal lights, burning brilliantly green or red, were my favorites. The green, my mother had told me, was due to an element called barium, the red to strontium. I had no idea at that point what barium and strontium were, but their names, like their colors, stayed in my mind.
When my mother saw how enthralled I was by these lights, she showed me how, if one threw a pinch of salt on the stove, the gas flame suddenly flared and turned a brilliant yellow – this was due to the presence of another element, sodium (even the Romans, she said, had used it to give their fires and flares a richer color). So, in a sense, I was introduced to ‘flame tests’ even before the war, but it was only a few years later, in Uncle Dave’s lab, that I learned they were an essential part of chemical life, an instant way of detecting certain elements, even if present in minute amounts.