Uncle Tungsten: Memories of a Chemical Boyhood (2001) (34 page)

Utterly new principles had to be invoked, or invented, to come to terms with this impossibility. Learning of this was the third ecstasy of my life, at least of my ‘chemical’ life – the first having been learning of Dalton and atomic theory, and the second of Mendeleev and his periodic table. But the third, I think, was in some ways the most stunning of all, because it contravened (or seemed to) all the classical science I knew, and all I knew of rationality and causality.

It was Niels Bohr, also working in Rutherford’s lab in 1913, who bridged the impossible, by bringing together Rutherford’s atomic model with Planck’s quantum theory. The notion that energy was absorbed or emitted not continuously but in discrete packets, ‘quanta,’ had lain silently, like a time bomb, since Planck had suggested it in 1900. Einstein had made use of the idea in relation to photoelectric effects, but otherwise quantum theory and its revolutionary potential had been strangely neglected, until Bohr seized on it to bypass the impossibilities of the Rutherford atom. The classical view, the solar-system model, would allow electrons an infinity of orbits, all unstable, all crashing into the nucleus. Bohr postulated, by contrast, an atom that had a limited number of discrete orbits, each with a specific energy level or quantal state. The least energetic of these, the closest to the nucleus, Bohr called the ‘ground state’ – an electron could stay here, orbiting the nucleus, without emitting or losing any energy, forever. This was a postulate of startling, outrageous audacity, implying as it did that the classical theory of electromagnetism might be inapplicable in the minute realm of the atom.

There was, at the time, no evidence for this; it was a pure leap of inspiration, imagination – not unlike the leaps he now posited for the electrons themselves, as they jumped, without warning or intermediates, from one energy level to another. For, in addition to the electron’s ground state, Bohr postulated, there were higher-energy orbits, higher-energy ‘stationary states,’ to which electrons might be briefly translocated. Thus if energy of the right frequency was absorbed by an atom, an electron could move from its ground state into a higher-energy orbit, though sooner or later it would drop back to its original ground state, emitting energy of exactly the same frequency as it had absorbed – this is what happened in fluorescence or phosphorescence, and it explained the identity of spectral emission and absorption lines, which had been a mystery for more than fifty years.

Atoms, in Bohr’s vision, could not absorb or emit energy except by these quantum jumps – and the discrete lines of their spectra were simply the expression of the transitions between their stationary states. The increments between energy levels decreased with distance from the nucleus, and these intervals, Bohr calculated, corresponded exactly to the lines in the spectrum of hydrogen (and to Balmer’s formula for these). This coincidence of theory and reality was Bohr’s first great triumph. Einstein felt that Bohr’s work was ‘an
enormous
achievement,’ and, looking back thirty-five years later, he wrote, ‘ [it] appears to me as a miracle even today…This is the highest form of musicality in the sphere of thought.’ The spectrum of hydrogen – spectra in general – had been as beautiful and meaningless as the markings on butterflies’ wings, Bohr remarked; but now one could see that they reflected the energy states within the atom, the quantal orbits in which the electrons spun and sang. ‘The language of spectra,’ wrote the great spectroscopist Arnold Sommerfeld, ‘has been revealed as an atomic music of the spheres.’

Could quantum theory be extended to more complex, multi-electron atoms? Could it explain their chemical properties, explain the periodic table? This became Bohr’s focus as scientific life resumed after the First World War.«68»

As one moved up in atomic number, as the nuclear charge or number of protons in the nucleus increased, an equal number of electrons had to be added to preserve the neutrality of the atom. But the addition of these electrons to an atom, Bohr envisaged, was hierarchical and orderly. While he had concerned himself at first with the potential orbits of hydrogen’s lone electron, he now extended his notion to a hierarchy of orbits or shells for all the elements. These shells, he proposed, had definite and discrete energy levels of their own, so that if electrons were added one by one, they would first occupy the lowest-energy orbit available, and when that was full, the next-lowest orbit, then the next, and so on. Bohr’s shells corresponded to Mendeleev’s periods, so that the first, innermost shell, like Mendeleev’s first period, accommodated two elements, and two only. Once this shell was completed, with its two electrons, a second shell began, and this, like Mendeleev’s second period, could accommodate eight electrons and no more. Similarly for the third period or shell. By such a building-up, or
aufbau
₃
Bohr felt, all the elements could be systematically constructed, and would naturally fall into their proper places in the periodic table.

Thus the position of each element in the periodic table represented the number of electrons in its atoms, and each element’s reactivity and bonding could now be seen in electronic terms, in accordance with the filling of the outermost shell of electrons, the so-called valency electrons. The inert gases each had completed outer valency shells with a full complement of eight electrons, and this made them virtually unreactive. The alkali metals, in Group I, had only one electron in their outermost shell, and were intensely avid to get rid of this, to attain the stability of an inert-gas configuration; the halogens in Group VII, conversely, with seven electrons in their valency shell, were avid to acquire an extra electron and also achieve an inert-gas configuration. Thus when sodium came into contact with chlorine, there would be an immediate (indeed explosive) union, each sodium atom donating its extra electron, and each chlorine atom happily receiving it, both becoming ionized in the process.

The placement of the transition elements and the rare-earth elements in the periodic table had always given rise to special problems. Bohr now suggested an elegant and ingenious solution to this: the transition elements, he proposed, contained an additional shell of ten electrons each; the rare-earth elements an additional shell of fourteen. These inner shells, deeply buried in the case of the rare-earth elements, did not affect chemical character in nearly so extreme a way as the outer shells; hence the relative similarity of all the transition elements and the extreme similarity of all the rare-earth elements.

Bohr’s electronic periodic table, based on atomic structure, was essentially the same as Mendeleev’s empirical one based on chemical reactivity (and all but identical with the block tables devised in pre-electronic times, such as Thomsen’s pyramidal table and Werner’s ultralong table of 1905). Whether one inferred the periodic table from the chemical properties of the elements or from the electronic shells of their atoms, one arrived at exactly the same point.«69» Moseley and Bohr had made it absolutely clear that the periodic table was based on a fundamental numerical series that determined the number of elements in each period: two in the first period, eight each in the second and third, eighteen each in the fourth and fifth; thirty-two in the sixth and perhaps also the seventh. I repeated this series – 2, 8, 8, 18, 18, 32 – over and over to myself.

At this point I started to revisit the Science Museum and spend hours once again gazing at the giant periodic table there, this time concentrating on the atomic numbers inscribed in each cubicle in red. I would look at vanadium, for example – there was a shining nugget in its pigeonhole – and think of it as element 23, a 23 consisting of 5 + 18: five electrons in an outer shell around an argon ‘core’ of eighteen. Five electrons – hence its maximum valency of 5; but three of these formed an incomplete inner shell, and it was such an incomplete shell, I had now learned, that gave rise to vanadium’s characteristic colors and magnetic susceptibilities. This sense of the quantitative did not replace the concrete, the phenomenal sense of vanadium but heightened it, because I saw it now as a revelation, in atomic terms, of why vanadium had the properties it did. The qualitative and the quantitative had fused in my mind; the sense of ‘vanadiumness’ now could be approached from either end.

Between them, Bohr and Moseley had restored arithmetic to me, provided the essential, transparent arithmetic of the periodic table which had been intimated, though only in a muddy way, by atomic weights. The character and identity of the elements, much of it, anyhow, could now be inferred from their atomic numbers, which no longer just indicated nuclear charge but stood for the very architecture of each atom. It was all divinely beautiful, logical, simple, economical, God’s abacus at work.

What made metals metallic? Electronic structure explained why the metallic state seemed to be fundamental, so different in character from any other. Some of the mechanical properties of metals, their high densities and melting points, could now be explained in terms of the tightness with which electrons were bound to the nucleus. A very tightly bound atom, with a high ‘binding energy,’ seemed to go with unusual hardness and density, and high melting point. Thus it was that my favorite metals – tantalum, tungsten, rhenium, osmium: the filament metals – had the highest binding energies of any of the elements. (So there was, I was pleased to learn, an atomic justification for their exceptional qualities – and for my own preference.) The conductivity of metals was ascribed to a ‘gas’ of free and mobile electrons, easily detached from their parent atoms – this explained why an electric field could draw a current of mobile electrons through a wire. Such an ocean of free electrons, on the surface of a metal, could also explain its special luster, for oscillating violently with the impact of light, these would scatter or reflect any light back on its own path. The electron-gas theory carried the further implication that under extreme conditions of temperature and pressure, all the nonmetallic elements, all matter, could be brought into a metallic state. This had already been achieved with phosphorus in the 1920
s
, and it was predicted, in the 1930
s
, that at pressures in excess of a million atmospheres it might be achieved with hydrogen, too – there might be metallic hydrogen, it was speculated, at the heart of gas giants like Jupiter. The idea that
everything
could be ‘metallized’ I found deeply satisfying.«70»

I had long been puzzled by the peculiar powers of blue or violet light, short-wavelength light, as opposed to red or long-wavelength light. This was clear in the darkroom: one could have quite a bright ruby safelight that would not fog a developing film, whereas the least hint of white light, daylight (which of course contained blue), would fog it straightaway. It was clear, too, in the lab, where chlorine, for example, could be safely mixed with hydrogen in red light, but the mixture would explode in the presence of the least white light. And it was clear with Uncle Dave’s mineral cabinet, where one could induce phosphorescence or fluorescence with blue or violet light, but not with red or orange light. Finally, there were the photoelectric cells that Uncle Abe had in his house; these could be activated by the merest pencil of blue light, but would not respond to even a flood of red light. How could a huge amount of red light be less effective than a tiny amount of blue light? It was only after I had learned something of Bohr and Planck that I realized the answer to these apparent paradoxes must lie in the quantal nature of radiation and light, and the quantal states of the atom. Light or radiation came in minimum units or quanta, the energy of which depended on their frequency. A quantum of short-wavelength light – a blue quantum, so to speak – had more energy than a red one, and a quantum of X-rays or gamma rays had far more energy still. Each type of atom or molecule – whether of a silver salt in a photographic emulsion, or of hydrogen or chlorine in the lab, or of cesium or selenium in Uncle Abe’s photocells, or of calcium sulfide or tungstate in Uncle Dave’s mineral cabinet – required a certain specific level of energy to elicit a response; and this might be achieved by even a single high-energy quantum, where it could not be evoked by a thousand low-energy ones.

As a child I thought that light had form and size, the flower-like shapes of candle flames, like unopened magnolias, the luminous polygons in my uncle’s tungsten bulbs. It was only when Uncle Abe showed me his spinthariscope and I saw the individual sparkles in this that I started to realize that light, all light, came from atoms or molecules which had first been excited and then, returning to their ground state, relinquished their excess energy as visible radiation. With a heated solid, such as a white-hot filament, energies of many wavelengths were emitted; with an incandescent vapor, such as sodium in a sodium flame, only certain very specific wavelengths were emitted. (The blue light in a candle flame which had so fascinated me as a boy, I later learned, was generated by cooling dicarbon molecules as they emitted the energy they had absorbed when heated.)

But the sun, the stars, were like no lights on earth. They were of a brilliance, a whiteness, exceeding the hottest filament lamps (some, like Sirius, were almost blue). One could infer, from the radiation energy of the sun, a surface temperature of about 6,000 degrees. No one in his youth, Uncle Abe reminded me, had any idea what could allow the enormous incandescence and energy of the sun.
Incandescence
was scarcely the right word, for there was no burning, no combustion, in the ordinary sense – most chemical reactions, indeed, ceased above 1,000 degrees.

Could gravitational energy, the energy generated by a gigantic mass contracting, keep the sun going? This, too, it seemed, would be wholly inadequate to account for the blazing heat and energy of the sun and stars, undimmed for billions of years. Nor was radio-activity a plausible source of energy, because radioactive elements were not present in the stars in anywhere near the needed quantities, and their output of energy was too slow and unhurryable.