by Philip Ball
Until the Second World War began this was an international story, and as much collaborative as competitive. Discoveries made in Cambridge or Paris would be discussed within days in Berlin or Berkeley. In 1939 that ceased, and important discoveries in the nature and uses of nuclear power were regularly withheld from the science journals. It became a guessing game what one’s foreign peers were up to, what they knew, what they were trying to make. But everyone knew what the stakes were.
Invisible rays
Like quantum theory, nuclear physics began with a puzzle about radiation; indeed several puzzles, as one thing led to another. In the fin de siècle, invisible and intangible rays seemed to be everywhere. First there were cathode rays, studied by Lenard at Heidelberg and explained by J. J. Thomson in Cambridge. These led to the discovery of X-rays by Wilhelm Röntgen at the University of Würzburg, who noted that when cathode rays struck the glass wall of a cathode-ray tube, it was not just the glass that fluoresced—the influence extended further. Thomson and Lenard had already noticed that; in 1894 Thomson observed that glass tubing would glow even a few feet away from the discharge tube. That couldn’t be the work of the cathode rays themselves, which were stopped by the glass wall. Nor was it an effect of the fluorescent light from the glass: in 1895 Röntgen shielded the cathode-ray tube with thick black paper, yet still a phosphor screen beyond it glowed. If, however, he held up his hand between the tube and the screen, the phosphor image revealed the shadow of his bones. These mysterious penetrating rays, which Röntgen called X-rays, were apparently being produced when glass was stimulated into fluorescence by cathode rays. They could darken photographic plates, capturing the ghostly imprints permanently. Shown the X-ray of her skeletal hand, wedding ring and all, Röntgen’s wife exclaimed ‘I have seen my death!’
X-rays brought Röntgen a Nobel in 1901. Henri Becquerel in Paris heard about them in January 1896, and as an expert on fluorescence he wondered whether they might be emitted by fluorescing substances other than glass. One such was a salt of uranium, uranium potassium sulphate, which glowed after being exposed to sunlight.*1 Becquerel proposed to detect X-rays coming from the fluorescing uranium salt by their effect on photographic emulsion. He wrapped a photographic plate in black paper, masked a part of it with copper foil cut into a cross (which, he reasoned, should stop X-rays), scattered the salt on top and left it exposed to the sun to activate the effect. And indeed he found that when the plate was developed, the image of the cross was imprinted upon it.
But no X-rays had done this, as Becquerel subsequently realized in a famous stroke of serendipity. Having tried to repeat the experiment on an overcast February day, he put the plate, with the copper mask and uranium salt still on top, into a drawer for several days before deciding to develop it anyway, expecting to see at best only a weak imprint. Instead, the image of the copper cross was as clear as before. Sunlight wasn’t needed to activate the process, then; unlike Röntgen’s glass, Becquerel’s salt was giving off radiation spontaneously, without any apparent stimulation. This was something new.
Becquerel’s ‘uranic rays’ weren’t perceived as having the allure of X-rays, and the finding was not widely pursued. It did, however, spark the interest of Marie Curie, a Polish woman (née Maria Sklodowska) who had come to Paris to study science and mathematics at the Sorbonne. Marie married Pierre Curie, an instructor at the School of Chemistry and Physics, in 1895, and the birth of their daughter Irène in September 1897 prevented her from starting her doctorate until early the following year. She figured that Becquerel’s rays, which she named ‘radioactivity’, would make a good subject, ‘because the question was entirely new and nothing yet had been written upon it’.
Becquerel had noted that the uranic rays made air electrically conducting: they were capable of ejecting electrons from atoms, leaving the atoms charged (ionized). Pierre Curie had some years earlier invented an instrument for measuring electrical charge very accurately, called an electrometer, and the Curies now used this as a means of quantifying the activity of uranium. At first they used relatively pure uranium salts for their studies, but when Marie tested raw uranium ore (pitchblende, mined in Saxony), she found to her surprise that it was even more radioactive. She concluded that the ore must contain a second radioactive element with even greater activity than uranium itself. She had already discovered that uranium was not unique: the ‘uranic rays’ were also detected from the rare element thorium. But now she set about chemically analysing pitchblende to identify the source of the extra radioactivity. This meant using chemical separation techniques to extract the minor impurities that contained the additional radioactivity, repeating the process again and again on several tons of laboriously crushed, dirty brown pitchblende to collect as much of the trace material as possible. As they gradually made the solution of this impurity more concentrated, the Curies saw its activity increase. In July 1898 they presented their findings in a paper read by Becquerel to the Institut de France. Here they claimed to have identified a new radioactive element, which they named after Marie’s homeland: polonium.
There was a second radioactive impurity in pitchblende too, which showed different chemical behaviour and so could be separated from polonium. It was even more active—around a million times more so than uranium, enough for concentrated solutions to glow spontaneously as the ‘uranic rays’ excited fluorescence in the water. That accounted for the name that Pierre recorded in his lab book in December: radium.
Then radioactivity seemed to be everywhere. Electrometer measurements indicated that air itself was continually being ionized, and it was assumed that this was due to radiation streaming from natural radioactive elements in the ground or the air, such as the recently discovered inert gas radon (which is found in some types of rock). However, in 1912 the Austrian physicist Victor Hess in Vienna sent electrometers up in a balloon and discovered that the rate of ionization increased rather than decreased as the instruments were carried higher. The ionizing rays were coming from space: they were christened ‘cosmic rays’. Hess won the Nobel Prize in Physics in 1936, the same year that Peter Debye was rewarded in chemistry. Since Hess was a Jew, the National Socialists were unimpressed. Two years later he was arrested in Graz after the Anschluss because of his refusal to accept Nazi rule; he escaped and emigrated to the United States. Cosmic rays were of great interest to nuclear physicists, in part because they offered a source of very-high-energy particles as projectiles for nuclear-transmutation experiments, exceeding the energies that could be reached in particle accelerators. Werner Heisenberg devoted much attention to the subject from the late 1930s.
Atomic energy
What was radioactivity? It was unnerving, to say the least. X-rays were produced only in response to some energetic stimulation, such as bombardment with cathode rays. But uranium just went right on discharging its energetic rays, day after day, even in the dark and cold, no matter how the element itself might be transformed into different physical or chemical states. Unlike chemical energy, the energy of radioactivity seemed inexhaustible. Two German physics teachers, Hans Geitel and Julius Elster, took radioactive samples to the bottom of an 850-foot mineshaft in the Harz mountains of Saxony to test the hypothesis that they might be absorbing and then re-radiating rays that permeated all of space: they figured that a blanket of so much rock would attenuate any such influence. But it made no difference: the samples were as active as ever. The energy, they concluded, must be coming from the atom itself. It was atomic energy—and to understand it, one needed to understand the atom.
At the turn of the century, no one knew more about the atom than Ernest Rutherford, who arrived from New Zealand at Cambridge’s Cavendish Laboratory in the year X-rays were discovered. There he began experimenting on Becquerel’s uranic rays, and in 1899, shortly after moving to McGill University in Montreal as professor of physics, he reported that the rays are of two sorts: one, which he called alpha rays, are stopped by aluminium foil, whereas the second sort, called beta rays, are mo
re penetrating. Rutherford later deduced that both of these varieties are actually particles, not rays. The beta particles are equivalent to cathode rays, that is, electrons. And in 1908, after leaving McGill to work at the University of Manchester, Rutherford showed in a beautiful experiment that alpha particles are the positively charged nuclei of helium atoms.
How can helium nuclei come out of other atoms? Rutherford had already realized that this involved the transmutation of one element into another. Experiments on thorium radioactivity were complicated by the fact that they gave inconsistent results unless the thorium was enclosed in a metal box. In 1899 Rutherford realized that this was because a gas was escaping from thorium and would, unless prevented, carry some of the radioactivity away. What was this ‘thorium emanation’? That was a chemical question, and so Rutherford enlisted the help of a McGill chemist, an Englishman named Frederick Soddy. But Soddy’s answer seemed scarcely credible. The emanation didn’t undergo any chemical reactions at all, for it seemed to be nothing other than the inert gas argon discovered six years previously. ‘Rutherford,’ Soddy remembered stammering to his collaborator, ‘this is transmutation—the thorium is disintegrating.’ Rutherford boomed back that ‘they’ll have our heads off as alchemists!’
The thorium emanation wasn’t in fact argon, but the much heavier inert gas radon, into which thorium is transmuted by radioactive decay. Nonetheless, the principle of transmutation remained. Rutherford and Soddy realized it meant that radioactive emissions remove a part of the fabric of the atom, changing its chemical identity. As Rutherford’s student Henry Moseley showed in 1913, the ‘atomic number’ that defines the place of every chemical element in the sequence of the periodic table—from 1 for hydrogen to 92 for uranium—is not just an arbitrary label but quantifies the number of positive charges (relative to the hydrogen nucleus) in the atom’s nucleus. If a radioactive atom emits an alpha particle, this carries off two of those charges (the atomic number of helium is 2), and so reduces the atomic number by two, transmuting the element to that two places to the left in the periodic table. In this way, thorium, with atomic number 90, is transmuted to radium, with atomic number 88.*2 In other words, a radioactive atom ‘decays’ into another element as it emits radiation. (The emission of beta particles also induces transmutation, but that is a more complex matter which puzzled Rutherford and Marie Curie for years.) Some decays are so rapid as to be almost instantaneous, others are geologically slow. The rate is measured by the so-called half-life: the time taken for half of the atoms in a sample of a radioactive element to decay. This is always the same regardless of how much material you have. The half-life of uranium is about 4.5 billion years,†3 about the same as the age of the earth; that of thorium is just twenty-two minutes.
By radiating energetic particles (or high-energy photons called gamma rays in a third form of nuclear decay), radioactive atoms shed energy. In 1903 Rutherford and Soddy estimated the astonishing quantity of energy locked up in the atom, in comparison to which the energy released by any chemical process, such as the detonation of an explosive, is puny—at least 20,000 times less. Suppose, Rutherford mused, one could find a ‘detonator’ to expel all this atomic energy at once: then ‘some fool in a laboratory might blow up the universe unawares’.
Soddy returned to England, and in 1904 he delivered this sobering thought in a lecture. If this energy ‘could be tapped and controlled’, he said,
what an agent it would be in shaping the world’s destiny! The man who put his hand on the lever by which a parsimonious nature regulates so jealously the output of this store of energy would possess a weapon by which he could destroy the earth if he chose.
Soddy’s audience on that occasion was surely captivated by the thought: it was the Corps of Royal Engineers of the British Army.
The Curies had come to the same disturbing conclusion. In 1903 they and Becquerel were awarded the Nobel Prize in Physics for their work on radioactivity, but had been too busy and too ill to travel to Sweden for the ceremony (the kind of event that Pierre, in any case, loathed).*4 Marie gave birth to their second daughter Eve at the end of 1904, and so it was not until June 1905 that they finally came to Stockholm, where Pierre delivered the traditional Nobel lecture. ‘It can even be thought’, he said,
that radium could become very dangerous in criminal hands, and here the question can be raised whether mankind benefits from knowing the secrets of Nature, whether it is ready to profit from it or whether this knowledge will not be harmful for it.
Pierre was optimistic about that, but one would not have guessed it from his demeanour. He was prematurely aged, constantly tired and often depressed. It wasn’t clear why; he attributed his aches and pains to rheumatism. Marie too struggled with lethargy—she is, Pierre wrote to a correspondent in 1903, ‘always tired without being exactly ill’. Her second pregnancy had ended that summer in the premature birth of a baby who died soon after. The Curies noted that their experiments with radium could leave them with lesions and reddened skin on their fingers; Rutherford noted that Pierre’s hands were ‘in a very inflamed and painful state due to exposure to radium rays’. Yet for Pierre these ailments never took their likely course, as a year after his triumphant lecture in Stockholm he was killed in a road accident in Paris. Had he lived, he would surely have shared in the second Nobel awarded to Marie alone in 1911, this time in chemistry, for the discovery of radium and polonium.
The many faces of atoms
If you wanted to know about radiochemistry and the structure of the atom, Rutherford was the man to ask. The 26-year-old Otto Hahn went to Montreal in 1905 to work with him, and at Manchester Rutherford took on the German physicist Hans Geiger. Rutherford and Geiger developed a device that revealed the passage of an alpha particle by its ionization of air, triggering an electrical discharge between two electrodes with an audible click—this was the predecessor of Geiger’s famous counter. Geiger collaborated on the demonstration that alpha particles are ionized helium atoms in 1908, and the following year he, Rutherford and student Ernest Marsden conducted Rutherford’s most celebrated experiment in which he deduced that atoms have small, dense nuclei. They shot a stream of alpha particles at thin gold foil, expecting them to pass through with some slight deflection. Mostly that’s what happened; but to his perplexity, Marsden found that a very few of the particles bounced right back. ‘It was quite the most incredible event that has ever happened to me in my life’, Rutherford wrote later. ‘It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you.’ He concluded that this was only possible if atoms were not like little spheres—the ‘plum puddings’ proposed by J. J. Thomson, with electrons embedded in a suet of positive charge—but instead had most of the mass concentrated in a tiny, very dense and electrically charged centre, surrounded by a diffuse cloud of particles of the opposite charge. Mostly the atom was quite empty. Rutherford finally decided that the nucleus must have a positive charge, and the diffuse cloud contained electrons. This, he discovered, was much like the ‘Saturnian’ atom proposed in 1903 by the Japanese physicist Hantaro Nagaoka, in which rings of electrons orbited a positive (but by no means diminutive) core. As we saw earlier, this classical picture of a ‘planetary’ atom was soon reformulated in terms of quantum theory by Niels Bohr, who joined Rutherford at Manchester in 1912 ‘to get some experience in radioactive work’.
To understand the source of radioactive energy, one needed to understand the nucleus, which in Bohr’s quantum atom could be regarded as just a ball of positive charge. What was actually in there? In 1815 the chemist William Prout had proposed that the hydrogen nucleus (later recognized as a single positively charged particle called a proton) was the fundamental building block of all the other atoms. If so, then the masses of all atoms should be simple multiples of the mass of hydrogen. This was roughly true in general, but not exactly, and sometimes hardly at all. Carbon and oxygen atoms, for example, are about twelve and sixteen times the mass o
f hydrogen—but chlorine atoms seemed to be about 35.5 times that mass. So Prout’s law seemed highly inexact, and was largely ignored.
In 1919 Francis Aston at the Cavendish Laboratory, where Rutherford had recently been appointed director, devised an instrument for measuring atomic masses very accurately, which became known as a mass spectrometer. Aston was able to separate out atoms of different mass by removing electrons to make them electrically charged (ions), using an electric field to accelerate them down a channel evacuated of air, and then using a second field to bend their trajectories. If they have the same charge, then ions of different mass are deflected to different degrees and can be collected separately. Aston found that a pure element placed in the spectrometer could be separated into fractions with several different masses, each of them pretty much an exact multiple of the mass of hydrogen. Sulphur atoms, for example, could have masses of 32, 33 and 34, while chlorine atoms have masses of 35 and 37; their weighted average in a mixture of many billions of atoms resulted in the apparent ‘fractional masses’ that had seemed to flout Prout’s law.
The idea that atoms of an element could differ in mass wasn’t entirely new. Before the First World War interrupted Aston’s research by billeting him for military work with the Royal Air Force, he acted as an assistant to J. J. Thomson at the Cavendish and developed physical techniques to separate atoms or molecules of different mass, based on the fact that in gases such particles diffuse at different rates—lighter particles move more quickly, like children darting nimbly through a crowd. While investigating the inert gas neon, Aston found tantalizing evidence for two forms of the element with slightly different atomic masses of 20 and 22.