Mendeleev’s Garden
In honor of the late neurologist who charmed us with over a dozen books, a beloved essay from the archives.
This essay originally appeared in our Autumn 2001 issue.
In 1945 the Science Museum in London reopened (it had been closed for much of the war), and I—a boy of twelve with a passion for metals and numbers—first saw the giant periodic table displayed there. The table itself, covering a whole wall at the head of the stairs, was a cabinet made of dark wood with ninety-odd cubicles, each inscribed with the name, the atomic weight, and the chemical symbol of its element. And in each cubicle was a sample of the element itself (all of those elements, at least, that had been obtained in pure form, and that could be exhibited safely). It was labeled “The Periodic Classification of the Elements—after Mendeleeff.”
My first vision was of metals, dozens of them in every possible form: rods, lumps, cubes, wire, foil, discs, crystals. Most were gray or silver, some had hints of blue or rose. A few had burnished surfaces that shone a faint yellow, and then there were the rich colors of copper and gold.
In the upper right corner were the nonmetals—sulfur in spectacular yellow crystals and selenium in translucent red crystals; phosphorus, like pale beeswax, kept under water; and carbon, as tiny diamonds and shiny black graphite. There was boron, a brownish powder, and ridged crystalline silicon, with a rich black sheen like graphite or galena.
On the left were the alkali and alkaline earth metals, all (except magnesium) in protective baths of naphtha. I was struck by the lithium in the upper corner, for this, with its levity, was floating on the naphtha, and also by the cesium, lower down, which formed a glittering puddle beneath the naphtha. Cesium, I knew, had a very low melting point and it was a hot summer day. But I had not fully realized, from the tiny, partly oxidized lumps I had seen, that pure cesium was pale gold—it gave at first just a glint, a flash of gold, seeming to iridesce with a golden luster; then, from a lower angle, it was purely gold, and looked like a gilded sea, or golden mercury.
There were other elements that up to this point had been only names to me (or, almost equally abstract, names attached to some physical properties and atomic weights), and now for the first time I saw the range of their diversity and actuality. In this first, sensuous glance I saw the table as a gorgeous banquet, a huge table set with eighty-odd different dishes.
I had already become familiar with the properties of many elements and I knew they formed a number of natural families, such as the alkali metals, the alkaline earth metals, and the halogens. These families (Mendeleev called them “groups”) formed the verticals of the table—the alkali and alkaline earth metals to the left, the halogens and inert gases to the right, and everything else in four intermediate groups in between. The “groupishness” of these intermediate groups was somewhat less clear—thus in Group VI, I saw sulfur, selenium, and tellurium. I knew that these three (my “stinkogens”) were very similar, but what was oxygen doing heading the group? There must be some deeper principle at work—and indeed there was. This was printed at the top of the table, but in my impatience to look at the elements themselves, I had paid no attention to it at all. The deeper principle, I saw, was valency. The term valency was not to be found in my early Victorian books, for it had been properly developed only in the late 1850s, and Dmitri Mendeleev was one of the first to seize on it and use it as a basis for classification, to provide what had never been clear before: a rationale, a basis for the fact that elements seemed to form natural families, to have deep chemical and physical analogies with one another. Mendeleev now recognized eight such groups of elements in terms of their valencies.
Thus the elements in Group I, the alkali metals, had a valency of 1: one atom of these would combine with one atom of hydrogen, to form compounds such as LiH, NaH, KH, and so on. (Or with one atom of chlorine, to form compounds such as LiCl, NaCl, or KCL) The elements of Group II, the alkaline earth metals, had a valency of 2, and so would form compounds such as CaCl2, SrC12, BaC12, and so on. The elements of Group VIII had a maximum combining power of 8.
But while Mendeleev was organizing the elements in terms of valency, he was also fascinated by atomic weights and the fact that these were unique and specific to each element, that they were, in a sense, the atomic signature of each element. And if, mentally, he started to index the elements according to their valencies, he did this equally in terms of their atomic weights. And now, magically, the two came together. For if he arranged the elements, quite simply, in order of their atomic weights, in horizontal “periods,” as he called them, one could see recurrences of the same properties and valencies at regular intervals.
Every element echoed the properties of the one above, was a slightly heavier member of the same family. The same melody, so to speak, was played in each period—first an alkali metal, then an alkaline earth metal, then six more elements, each with its own valency or tone—but played in a different register (it was impossible to avoid thinking of octaves and scales here, for I lived in a musical house, and scales were the periodicity I heard daily).
It was eightness that dominated the periodic table before me, though one could also see, in the lower part of the table, that extra elements were interposed within the basic octets: ten extra elements apiece in Periods 4 and 5, and ten plus fourteen in Period 6.
So one went up, each period completing itself and leading to the next one in a series of dizzying loops—at least this is the form my imagination took, so that the sober, rectangular table before me was transformed, mentally, into spirals or loops. The table was a sort of cosmic staircase or a Jacob’s ladder, going up to, coming down from, a Pythagorean heaven.
I got a sudden, overwhelming sense of how startling the periodic table must have seemed to those who first saw it—chemists who were profoundly familiar with seven or eight chemical families but had never realized the basis of these families (valency), nor how all of them might be brought together into a single overarching scheme. I wondered if they had reacted as I did to this first revelation: “Of course! How obvious! Why didn’t I think of it myself?”
Whether one thought in terms of the verticals or in terms of the Horizontals—either way one arrived at the same grid. It was like a crossword puzzle that could be approached by either the “down” or the “across” clues, except that a crossword was arbitrary, a purely human construct, while the periodic table reflected a deep order in nature, for it showed all the elements arrayed in a fundamental relationship. I had the sense that it harbored a marvelous secret, but it was a cryptogram without a key—why was this relationship so?
I could scarcely sleep for excitement the night after seeing the periodic table—it seemed to me an incredible achievement to have brought the whole, vast, and seemingly chaotic universe of chemistry to an all-embracing order. The first great intellectual clarifications had occurred with Lavoisier’s defining of elements, with J. L. Proust’s finding that elements combined in discrete proportions only, and with Dalton’s notion that elements had atoms with unique atomic weights. With these, chemistry had come of age and had become the chemistry of the elements. But the elements themselves were not seen to come in any order; they could be listed only alphabetically (as Pepper did in his Playbook of Metals) or in terms of isolated local families or groups. Nothing beyond this was possible until Mendeleev’s achievement. To have perceived an overall organization, a superarching principle uniting and relating all the elements, had a quality of the miraculous, of genius. And this gave me, for the first time, a sense of the transcendent power of the human mind, and the fact that it might be equipped to discover or decipher the deepest secrets of nature, to read the mind of God.
I kept dreaming of the periodic table in the excited half-sleep of that night—I dreamed of it as a flashing, revolving pinwheel or Catherine wheel, and then as a great nebula, going from the first element to the last, and whirling beyond uranium, out to infinity. The next day I could hardly wait to get to the museum and dashed up to the top floor, where the table was, as soon as the doors were opened.
On this second visit I found myself looking at the table in almost geographic terms, as a realm, a kingdom, with different territories and boundaries. Seeing the table as a geographic realm allowed me to rise above the individual elements and to see certain general gradients and trends. Metals had long been recognized as a special category of elements, and now one could see, in a single synoptic glance, how they occupied three-quarters of the realm—all of the west side, most of the south—leaving only a smallish area, mostly in the northeast, for the nonmetals. A jagged line, like Hadrian’s Wall, separated the metals from the rest, with a few “semimetals,” metalloids—arsenic, selenium—straddling the wall. One could see the gradients of acid and base, how the oxides of the “western” elements reacted with water to form alkalis, the oxides of the “eastern” elements, mostly nonmetals, to form acids. One could see, again at a glance, how the elements on either border of the realm—the alkali metals and halogens, like sodium and chlorine, for example—showed the greatest avidity for each other and combined with explosive force, forming crystalline salts with high melting points, which dissolved to form electrolytes, while those in the middle formed a very different sort of compound: volatile liquids or gases that resisted electric currents. One could see, remembering how Volta and Davy and Berzelius ranked the elements into an electrical series, how the most strongly electropositive elements were all to the left, the most strongly electronegative to the right. Thus it was not just the placement of the individual elements, but trends of every sort that hit the eye when one looked at the table.
Seeing the table, “getting” it, altered my life. I took to visiting it as often as I could. I copied it into my exercise book and carried it everywhere; I got to know it so well—visually and conceptually—that I could mentally trace its paths in every direction, going up a group, then turning right on a period, stopping, going down one, yet always knowing where I was. It was like a garden, the garden of numbers I had loved as a child—but unlike that garden, it was real, a key to the universe. I spent hours now, totally absorbed, wandering, making discoveries, in the enchanted garden of Mendeleev.
There was a photograph of Mendeleev next to the periodic table in the museum; he looked like a cross between Fagin and Svengali, with a huge mass of hair and beard and piercing, hypnotic eyes. A wild, extravagant, barbaric figure—as romantic, in his way, as the Byronic Humphry Davy. I needed to know more of him, and to read his famous Principles of Chemistry, in which he had first published his periodic table.
His book, his life, did not disappoint me. He was a man of encyclopedic interests. He was also a music lover and a close friend of Borodin (who was also a chemist). And he was the author of the most delightful and vivid chemistry text ever published.1
Like my own parents, Mendeleev had come from a huge family—he was the youngest, I read, of fourteen children. His mother must have recognized his precocious intelligence, and when he reached fourteen, feeling that he would be lost without a proper education, she walked thousands of miles from Siberia with him—first to the University of Moscow (from which, as a Siberian, he was barred) and then to St. Petersburg, where he got a grant to train as a teacher. (She herself, apparently, nearing sixty at the time, died from exhaustion after this prodigious effort. Mendeleev, profoundly attached to her, was later to dedicate the Principles to her memory.)
Even as a student in St. Petersburg, Mendeleev showed not only an insatiable curiosity but a hunger for organizing principles of all kinds. Linnaeus, in the eighteenth century, had classified animals and plants, and (much less successfully) minerals, too. Dana, in the 1830s, had replaced the old physical classification of minerals with a chemical classification of a dozen or so main categories (native elements, oxides, sulfides, and so on). But there was no such classification for the elements themselves, and there were now some sixty elements known. Some elements, indeed, seemed almost impossible to categorize. Where did uranium go, or that puzzling, ultralight metal, beryllium? Some of the most recently discovered elements were particularly difficult—thallium, for example, discovered in 1862, was in some ways similar to lead, in others to silver, in others to aluminum, and in yet others to potassium.
It was nearly twenty years from Mendeleev’s first interest in classification to the emergence of his periodic table in 1869. This long pondering and incubation (so similar, in a way, to Darwin’s before he published On the Origin of Species) was perhaps the reason why, when Mendeleev finally published his Principles, he could bring a vastness of knowledge and insight far beyond any of his contemporaries—some of them also had a clear vision of periodicity, but none of them could marshal the overwhelming detail he could.
Mendeleev described how he would write the properties and atomic weights of the elements on cards and ponder and shuffle these constantly on his long railway journeys through Russia, playing a sort of patience or (as he called it) “chemical solitaire,” groping for an order, a system that might bring sense to all the elements, their properties and atomic weights.
There was another crucial factor. There had been considerable confusion, for decades, about the atomic weights of many elements. It was only when this was cleared up at the first international conference of chemists, at Karlsruhe, in 1860, that Mendeleev and others could even think of achieving a full taxonomy of the elements. Mendeleev had gone to Karlsruhe with Borodin (this was a musical as well as a chemical journey, for they stopped at many churches en route, trying out the local organs). With the old, pre-Karlsruhe atomic weights, one could get a sense of local triads or groups, but one could not see that there was a numerical relationship between the groups themselves.2 Only when Cannizzaro showed how reliable atomic weights could be obtained, and showed, for example, that the proper atomic weights for the alkaline earth metals (calcium, strontium, and barium) were 40, 88, and 137 (not 20, 44, and 68, as formerly believed) did it become clear how close these were to those of the alkali metals—potassium, rubidium, and cesium. It was this closeness, and in turn the closeness of the atomic weights of the halogens—chlorine, bromine, and iodine—that incited Mendeleev, in 1868, to make a small grid juxtaposing the three groups:
| Cl | 35.5 | K | 39 | Ca | 40 |
| Br | 80 | Rb | 85 | Sr | 88 |
| I | 127 | Cs | 133 | Ba | 137 |
And it was at this point, seeing that arranging the three groups of elements in order of atomic weight produced a repetitive pattern—a halogen followed by an alkali metal, followed by an alkaline earth metal—that Mendeleev, feeling this must be a fragment of a larger pattern, leapt to the idea of a periodicity governing all the elements: a Periodic Law.
Mendeleev’s first small table had to be filled in, and then extended in all directions, as if filling up a crossword puzzle; this in itself required some bold speculations. What element, he wondered, was chemically allied with the alkaline earth metals, yet followed lithium in atomic weight? No such element apparently existed—or could it be beryllium, usually considered to be trivalent, with an atomic weight of 14.5? What if it was bivalent instead, with an atomic weight, therefore, not of 14.5 but 9? Then it would follow lithium and fit into the vacant space perfectly.
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.3
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-aluminum”) that would come below aluminum 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 that 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.4 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.5
The keystone to the whole table, curiously, was not anticipated by Mendeleev, and perhaps could not have been, for this was a question not of a missing element but of an entire family or group. When argon was discovered in 1894—an element that 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.6 The inert gases were identical in their inability to form compounds; they had a valency, it seemed, of zero.7
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.8
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.9 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. Element 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 elements had a sort of homogeneity. 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 that I had been introduced to by my Uncle Dave, who manufactured light bulbs with tungsten filaments (we often called him Uncle Tungsten). 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 that 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 not only as rare and precious but as having special, secret qualities possessed by nothing else.
Later Uncle Dave told me of the extraordinary difficulty that chemists had had in separating the individual rare earths—there were a dozen or more, and 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) contained not just 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 of 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 a 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 on the “last” element, uranium. I was intensely curious about this, about Period 7, which started with an as-yet-unknown alkali metal, element 87, but got only 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 protactinium, 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 be not 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 co-workers 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).>10
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 table 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 that 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 that 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.”
Footnotes
1 In his very first footnote, in the preface, Mendeleev spoke of “how contented, free, and joyous is life in the realm of science”—and one could see, in every sentence, how true this was for him. The Principles grew like a living thing in Mendeleev’s lifetime, each edition larger, fuller, more mature than its predecessors, each filled with exuberating and spreading footnotes (footnotes that became so enormous that in the last editions they filled more pages than the text; indeed, some occupied nine-tenths of the page-I think my own love of footnotes, the excursions they allow, was partly determined by reading the Principles). ⇧
2 Mendeleev was not the first to see some significance in the atomic weights of elements. When the atomic weights of the alkaline earth metals were established by Berzelius, Dobereiner was struck by the fact that the atomic weight of strontium was just midway between that of calcium and barium. Was this an accident, as Berzelius thought, or an indication of something important and general? Berzelius himself had just discovered selenium in 1817 and at once realized that (in terms of chemical properties) it “belonged” between sulfur and tellurium. Dobereiner went further, and brought out a quantitative relationship too, for selenium’s atomic weight was just midway between theirs. And when lithium was discovered later that year (also in Berzelius’s kitchen lab), Dobereiner observed that it completed another triad, of alkali metals: lithium, sodium, and potassium. Feeling, moreover, that the gap in atomic weight between chlorine and iodine was too great, Dobereiner thought (as Davy had before him) that there must be a third element analogous to them, a halogen, with an atomic weight midway between theirs. (This element, bromine, was discovered a few years later.)
There were mixed reactions to Dobereiner’s “triads,” with their implication of a correlation between atomic weight and chemical character. Berzelius and Davy were doubtful of the significance of such “numerology,” as they saw it, but others were intrigued and wondered whether an obscure but fundamental significance was lurking in Dobereiner’s figures. ⇧
3 This, at least, is the accepted myth, and one that was later promulgated by Mendeleev himself, somewhat as Kekule was to describe his own discovery of the benzene ring years later, as the result of a dream of snakes biting their own tails. But if one looks at the actual table that Mendeleev sketched, one can see that it is full of transpositions, crossings-out, and calculations in the margins. It shows, in the most graphic way, the creative struggle for understanding, which was going on in his mind. Mendeleev did not wake from his dream with all the answers in place, but, more interestingly, perhaps, woke with a sense of revelation, so that within hours he was able to solve many of the questions that had occupied him for years. ⇧
4 In an 1889 footnote—even his lectures had footnotes, at least in their printed versions—he added: “I foresee some more new elements, but not with the same certitude as before.” Mendeleev was well aware of the gap between bismuth (with an atomic weight of 209) and thorium (232), and conceived that several elements must exist to fill it. He was most certain of the element immediately following bismuth—“an element analogous to tellurium, which we may call dvi-tellurium.” This element, polonium, was discovered by the Curies in 1898, and when finally isolated it had almost all the properties Mendeleev had predicted. (In 1899 Mendeleev visited the Curies in Paris and welcomed radium as his “eka-barium.”)
In the final edition of the Principles, Mendeleev made many other predictions—including two heavier analogs of manganese—an “eka-manganese” with an atomic weight of around 99, and a “trimanganese” with an atomic weight of 188; sadly, he never saw these. “Tri-manganese”—rhenium—was not discovered until 1925, the last of the naturally occurring elements to be found; while “ekamanganese,” technetium, was the first new element to be artificially made, in 1937.
He also envisaged, by analogy, some elements following uranium. ⇧
5 It is a remarkable example of synchronicity that in the decade following the Karlsruhe conference there emerged not one but six such classifications, all completely independent of one another: de Chancourtois’s in France, Odling’s and Newlands’s, both in England, Lothar Meyer’s in Germany, Hinrichs’s in America, and finally Mendeleev’s in Russia, all pointing toward a periodic law.
De Chancourtois, a French mineralogist, was the first to devise such a classification, and in 1862—just eighteen months after Karlsruhe—he inscribed the symbols of twenty-four elements spiraling around a vertical cylinder at heights proportional to their atomic weights, so that elements with similar properties fell one beneath another. Tellurium occupied the midpoint of the helix; hence he called it a “telluric screw,” a vis tellurique. But the Comptes Rendu, when they came to publish his paper, managed—grotesquely—to omit the crucial illustration, and this, among other problems, put paid to the whole enterprise, causing de Chancourtois’s ideas to be ignored.
Newlands, in England, was scarcely any luckier. He, too, arranged the known elements by increasing atomic weight, and seeing that every eighth element, apparently, was analogous to the first, he proposed a “Law of Octaves,” saying that “the eighth element, starting from a given one, is a kind of repetition of the first, like the 8th note in an octave of music.” (Had the inert gases been known at the time, it would, of course, have been every ninth element that resembled the first.) A too literal comparison to music, and the suggestion even that these octaves might be a sort of “cosmic music,” evoked a sarcastic response at the meeting of the Chemical Society at which Newlands presented his theory; it was said that he might have done as well to arrange the elements alphabetically.
There is no doubt that Newlands, even more than de Chancourtois, was very close to a periodic law. Like Mendeleev, Newlands had the courage to invert the order of certain elements when their atomic weight did not match what seemed to be their proper position in his table (though he failed to make any predictions of unknown elements, as Mendeleev did).
Lothar Meyer was also at the Karlsruhe conference and was one of the first to use the revised atomic weights published there in a periodic classification. In 1868 he came up with an elaborate sixteen-columned periodic table (but the publication of this was delayed until after. Mendeleev’s table had appeared). Lothar Meyer paid special attention to the physical properties of the elements and their relation to atomic weights, and in 1870 he published a famous graph plotting the atomic weights of the known elements against their “atomic volumes” (this being the ratio of atomic weight to density), a graph that showed high points for the alkali metals and low points for the dense, small atomed metals of Group VIII (the platinum and iron metals), with all the other elements falling nicely in between. This graph proved a most potent argument for a periodic law and did much to assist the acceptance of Mendeleev’s work.
But at the time of discovering his “Natural System,” Mendeleev was either ignorant of, or denied knowledge of, any attempts comparable to his own. Later, when his name and fame were established, he became more knowledgeable, perhaps more generous, less threatened by the notion of any co-discoverers or forerunners. When, in 1889, he was invited to give the Faraday Lecture in London, he paid a measured tribute to those who had come before him. ⇧
6 Cavendish, however, sparking the nitrogen and oxygen of air together, had observed in 1785 that a small amount (“not more than l/l20th part of the whole”) was totally resistant to combination, but no one paid any attention to this until the 1890s. ⇧
7 I think I identified at times with the inert gases and at other times anthropomorphized them, imagining them lonely, cut off, yearning to bond. Was bonding, bonding with other elements, absolutely impossible for them? Might not fluorine, the most active, the most outrageous of the halogens, so eager to combine that it had defeated efforts to isolate it for more than a century—might not fluorine, if given a chance, at least bond with xenon, the heaviest of the inert gases? I pored over tables of physical constants and decided that such a combination was, in principle, just possible.
In the early 1960s, I was overjoyed to hear (even though my mind at this time had moved on to other things) that Neil Bartlett had managed to prepare such a compound—a triple compound of platinum, fluorine, and xenon. Xenon fluorides and xenon oxides were subsequently made. Freeman Dyson has written to me describing his boyhood love of the periodic table and of the inert gases—he, too, saw them, in their bottles, in the Science Museum in South Kensington—and his excitement years later when he was shown a specimen of barium xenate, seeing the elusive, unreactive gas firmly and beautifully locked up in a crystal:
For me too, the periodic table was a passion … . As a boy, I stood in front of the display for hours, thinking how wonderful it was that each of these metal foils and jars of gas had its own distinct personality …. One of the memorable moments of my life was when Willard Libby came to Princeton with a little jar full of crystals of barium xenate. A stable compound, looking like common salt, but much heavier. This was the magic of chemistry, to see xenon trapped into a crystal. ⇧
8 A spectacular anomaly came up with the hydrides of the nonmetals—an ugly bunch, about as inimical to life as one could get. Arsenic and antimony hydrides were very poisonous and smelly; silicon and phosphorus hydrides were spontaneously inflammable. I had made in my lab the hydrides of sulfur (H2S), selenium (H2Se), and tellurium (H2Te), all Group VI elements, all dangerous and vile-smelling gases. The hydride of oxygen, the first Group VI element, one might predict by analogy, would be a foul smelling, poisonous, inflammable gas, too, condensing to a nasty liquid around -1002C. And instead it was water (H20)—stable, potable, odorless, benign, and with a host of special, indeed unique properties (its expansion when frozen, its great heat capacity, its capacity as an ionizing solvent, etc.) which made it indispensable to our watery planet, indispensable to life itself. What made it such an anomaly? The properties of water did not undermine for me the placement of oxygen in the periodic table, but they made me intensely curious as to why it was so different from its analogs. (This question, I found, had been resolved only recently, in the 1 930s, with Linus Pauling’s delineation of the hydrogen bond.)⇧
9 Ida Tacke Noddack was one of a team of German scientists who found element 75, rhenium, in 1925-26. Noddack also claimed to have found element 43, which she called masurium. But this claim could not be supported, and she was discredited. In 1934, when Fermi shot neutrons at uranium and thought he had made element 93, Noddack suggested that he was wrong, that he had in fact split the atom. But since she had been discredited with element 43, no one paid any attention to her. Had she been listened to, Germany would probably have had the atomic bomb and the history of the world would have been different. (This story was told by Glenn Seaborg when he was presenting his recollections at a conference in November 1997.) ⇧
10 Although elements 93 and 94, neptunium and plutonium, were created in 1940, their existence was not made public until after the war. They were given provisional names, when they were first made, of “extremium” and “ultimium,” because it was thought impossible that any heavier elements would ever be made. Elements 95 and 96, however, were created in 1944. Their discovery was not made public in the usual way—in a letter to Nature, or at a meeting of the Chemical Society—but on a children’s radio quiz show in November 1945, during which a twelve-year-old boy asked, “Mr. Seaborg, have you made any more elements lately?” ⇧
