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, 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.
68
By 1914, the scientists of Britain and France and Germany and Austria were all caught up, in various ways, in the First World War. Pure chemistry and physics were largely suspended for the duration, and applied science, war science, took its place. Rutherford ceased his fundamental research, and his lab was reorganized for work on submarine detection. Geiger and Marsden, who had observed the alpha-particle deflections that gave rise to Rutherford’s atom, found themselves at the Western Front, on different sides. Chadwick and Ellis, younger colleagues of Rutherford’s, were prisoners of war in Germany. And Moseley, aged twenty-eight, was killed by a bullet in the brain, at Gallipoli. My father often used to talk of the young poets, the intellectuals, the cream of a generation wiped out tragically in the Great War. Most of the names he mentioned were unknown to me, but Moseley’s was the one I knew, and the one I mourned most.
69
This gave Bohr predictive power too. Moseley had observed that element 72 was missing, but could not say whether it would be a rare-earth element or not (elements 57-71 were rare earths, and 73, tantalum, was a transition element, but no one was sure how many rare earths there would be). Bohr, with his clear idea of the numbers of electrons in each shell, was able to predict that element 72 would not be a rare-earth element, but a heavier analog of zirconium. He suggested that his colleagues in Denmark seek this new element in zirconium ores, and it was swiftly found (and named hafnium, after the old name for Copenhagen). This was the first time the existence and properties of an element were predicted not by chemical analogy, but on the purely theoretical basis of its electronic structure.