by Manjit Kumar
While Rutherford was busy conducting his experiments, in Paris the Frenchman Henri Becquerel was trying to discover whether phosphorescent substances, which glow in the dark, could also emit X-rays. Instead he found that uranium compounds emitted radiation whether they were phosphorescent or not. Becquerel’s announcement of his ‘uranic rays’ aroused little scientific curiosity and no newspapers clamoured to report his discovery. Only a handful of physicists were interested in Becquerel’s rays for, like their discoverer, most believed that only uranium compounds emitted them. However, Rutherford decided to investigate the effects of ‘uranic rays’ on the electrical conductivity of gases. It was a decision he later described as the most important of his life.
Testing the penetration of uranium radiation using wafer-thin layers of ‘Dutch metal’, a copper-zinc alloy, Rutherford found that the amount of radiation detected depended on the number of layers used. At a certain point, adding further layers had little effect in reducing the intensity of radiation, but then surprisingly it began to fall once again as more layers were added. After repeating the experiment with different materials and finding the same general pattern, Rutherford could offer only one explanation. Two types of radiation were being emitted, and he called them alpha and beta rays.
When the German physicist Gerhard Schmidt announced that thorium and its compounds also emitted radiation, Rutherford compared it with alpha and beta rays. He found the thorium radiation to be more powerful and concluded that ‘rays of a more penetrative kind were present’.29 These were later called gamma rays.30 It was Marie Curie who introduced the term ‘radioactivity’ to describe the emission of radiation and who labelled substances that emitted ‘Becquerel rays’ as ‘radioactive’. She believed that since radioactivity was not confined to uranium alone, it must be an atomic phenomenon. It set her on the path to discovering, with her husband Pierre, the radioactive elements radium and polonium.
In April 1898, as Curie’s first paper was published in Paris, Rutherford learned that there was a vacant professorship at McGill University in Montreal, Canada. Although acknowledged as a pioneer in the new field of radioactivity, Rutherford put his name forward with little expectation of being appointed, despite a glowing letter of recommendation from Thomson. ‘I have never had a student with more enthusiasm or ability for original research than Mr Rutherford,’ wrote Thomson, ‘and I am sure if elected, he would establish a distinguished school of physics at Montreal.’31 He concluded: ‘I should consider any institution fortunate that secured the services of Mr Rutherford as professor of physics.’ After a stormy voyage, Rutherford, just turned 27, arrived in Montreal at the end of September and stayed for the next nine years.
Even before he left England he knew that he was ‘expected to do a lot of original work and to form a research school to knock the shine out of the Yankees!’32 He did just that, beginning with the discovery that the radioactivity of thorium decreased by half in one minute and then by half again in the next. After three minutes it had fallen to an eighth of its original value.33 Rutherford called this exponential reduction of radioactivity the ‘half-life’, the time taken for the intensity of radiation emitted to fall by half. Each radioactive element had its own characteristic half-life. Then came the discovery that would earn him the professorship in Manchester and a Nobel Prize.
In October 1901, Rutherford and Frederick Soddy, a 25-year-old British chemist at Montreal, began a joint study of thorium and its radiation and were soon faced with the possibility that it could be turning into another element. Soddy recalled how he stood stunned at the thought and let slip, ‘this is transmutation’. ‘For Mike’s sake, Soddy, don’t call it transmutation’, warned Rutherford. ‘They’ll have our heads off as alchemists.’34
The pair were soon convinced that radioactivity was the transformation of one element into another through the emission of radiation. Their heretical theory was met with widespread scepticism but the experimental evidence quickly proved decisive. Their critics had to discard long-cherished beliefs in the immutability of matter. No longer an alchemist’s dream, but a scientific fact: all radioactive elements did spontaneously transform into other elements, the half-life measuring the time it took for half the atoms to do so.
‘Youthful, energetic, boisterous, he suggested anything but the scientist’, is how Chaim Weizmann, later the first president of Israel but then a chemist at Manchester University, remembered Rutherford. ‘He talked readily and vigorously on any subject under the sun, often without knowing anything about it. Going down to the refectory for lunch, I would hear the loud, friendly voice rolling up the corridor.’35 Weizmann found Rutherford ‘devoid of any political knowledge or feelings, being entirely taken up with his epoch-making scientific work’.36 At the centre of that work lay his use of the alpha particle to probe the atom.
But what exactly was an alpha particle? It was a question that had long vexed Rutherford after he discovered that alpha rays were in fact particles with a positive charge that were deflected by strong magnetic fields. He believed that an alpha particle was a helium ion, a helium atom that had lost two electrons, but never said so publicly because the evidence was purely circumstantial. Now, almost ten years after discovering alpha rays, Rutherford hoped to find definitive proof of their true character. Beta rays had already been identified as fast-moving electrons. With the help of another young assistant, this time 25-year-old German Hans Geiger, Rutherford confirmed in the summer of 1908 what he had long suspected: an alpha particle was indeed a helium atom that had lost two electrons.
‘The scattering is the devil’, Rutherford had complained as he and Geiger tried to unmask the alpha particle.37 He had first noticed the effect two years earlier in Montreal when some alpha particles that had passed through a sheet of mica were slightly deflected from their straight-line trajectory, causing fuzziness on a photographic plate. Rutherford made a mental note to follow it up. Soon after arriving in Manchester, he had drawn up a list of potential research topics. Rutherford now asked Geiger to investigate one of those items – the scattering of alpha particles.
Together they devised a simple experiment that involved counting scintillations, tiny flashes of light produced by alpha particles when they strike a paper screen coated with zinc sulphide, after passing through a thin sheet of gold foil. Counting scintillations was an arduous task, with long hours spent in total darkness. Luckily, according to Rutherford, Geiger was ‘a demon at the work and could count at intervals for a whole night without disturbing his equanimity’.38 He found that alpha particles either passed straight through the gold foil or were deflected by one or two degrees. This was as expected. However, surprisingly, Geiger also reported finding a few alpha particles ‘deflected through quite an appreciable angle’.39
Before he could fully consider the implications, if any, of Geiger’s results, Rutherford was awarded the Nobel Prize for chemistry for discovering that radioactivity was the transformation of one element into another. For a man who regarded ‘all science as either physics or stamp collecting’, he appreciated the funny side of his own instant transmutation from physicist to chemist.40 After returning from Stockholm with his prize, Rutherford learnt to evaluate the probabilities associated with different degrees of alpha particle scattering. His calculations revealed that there was a very small chance, almost zero, that an alpha particle passing through gold foil would undergo multiple scatterings resulting in an overall large-angle deflection.
It was while Rutherford was preoccupied with these calculations that Geiger spoke to him about assigning a project to Ernest Marsden, a promising undergraduate. ‘Why not,’ said Rutherford, ‘let him see if any alpha particles can be scattered through a large angle?’41 He was surprised when Marsden did. As the search continued at ever-larger angles, there should have been none of the tell-tale flashes of light that Marsden had seen, signalling alpha particles crashing into the zinc sulphide screen.
As Rutherford struggled to make sense of ‘the natur
e of the huge electric or magnetic forces which could turn aside or scatter a beam of alpha particles’, he asked Marsden to check if any were reflected backwards.42 Not expecting him to find anything, he was utterly astonished when Marsden discovered alpha particles bouncing off the gold foil. ‘It was,’ Rutherford said, ‘almost as incredible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you.’43
Geiger and Marsden set about making comparative measurements using different metals. They found that gold scattered backwards almost twice as many alpha particles as silver and twenty times more than aluminium. Only one alpha particle in every 8,000 bounced off a sheet of platinum. When they published these and other results in June 1909, Geiger and Marsden simply recounted the experiments and stated the facts without further comment. A baffled Rutherford brooded for the next eighteen months as he tried to think his way through to an explanation.
The existence of atoms had been a matter of considerable scientific and philosophical debate throughout the nineteenth century, but by 1909 the reality of atoms had been established beyond any reasonable doubt. The critics of atomism were silenced by the sheer weight of evidence against them, two key pieces of which were Einstein’s explanation of Brownian motion and its confirmation, and Rutherford’s discovery of the radioactive transformation of elements. After decades of argument, in which many eminent physicists and chemists had denied its existence, the most favoured representation of the atom to emerge was the so-called ‘plum pudding’ model put forward by J.J. Thomson.
In 1903 Thomson suggested that the atom was a ball of massless, positive charge in which were embedded like plums in a pudding the negatively-charged electrons he had discovered six years earlier. The positive charge would neutralise the repulsive forces between the electrons that would otherwise tear the atom apart.44 For any given element, Thomson envisaged these atomic electrons to be uniquely arranged in a set of concentric rings. He argued that it was the different number and distribution of electrons in gold and lead atoms, for example, which distinguished the metals from one another. Since all the mass of a Thomson atom was due to the electrons it contained, it meant there were thousands in even the lightest atoms.
Exactly one hundred years earlier, in 1803, the English chemist John Dalton first put forward the idea that atoms of every element were uniquely characterised by their weight. With no direct way of measuring atomic weights, Dalton determined their relative weights by examining the proportions in which different elements combined to form various compounds. First he needed a benchmark. Hydrogen being the lightest known element, Dalton assigned it an atomic weight of one. The atomic weights of all the other elements were then fixed relative to that of hydrogen.
Thomson knew his model was wrong after studying the results of experiments involving the scattering of X-rays and beta particles by atoms. He had overestimated the number of electrons. According to his new calculations, an atom could not have more electrons than prescribed by its atomic weight. The precise number of electrons in the atoms of the different elements was unknown, but this upper limit was quickly accepted as a first step in the right direction. The hydrogen atom with an atomic weight of one could have only one electron. However, the helium atom with an atomic weight of four could have two, three, or even four electrons, and so on for the other elements.
This drastic reduction in electron numbers revealed that most of the weight of an atom was due to the diffuse sphere of positive charge. Suddenly, what Thomson had originally invoked as nothing more than a necessary artifice to produce a stable, neutral atom took on a reality of its own. But even this new, improved model could not explain alpha particle scattering and failed to pin down the exact number of electrons in a particular atom.
Rutherford believed that alpha particles were scattered by an enormously strong electric field within the atom. But inside J.J.’s atom, with its positive charge evenly distributed throughout, there was no such intense electric field. Thomson’s atom simply could not send alpha particles hurtling backwards. In December 1910, Rutherford finally managed to ‘devise an atom much superior to J.J.’s’.45 ‘Now,’ he told Geiger, ‘I know what the atom looks like!’46 It was nothing like Thomson’s.
Rutherford’s atom consisted of a tiny positively-charged central core, the nucleus, which contained virtually all the atom’s mass. It was 100,000 times smaller than the atom, occupying only a minute volume, ‘like a fly in a cathedral’.47 Rutherford knew that electrons inside an atom could not be responsible for the large deflection of alpha particles, so to determine their exact configuration around the nucleus was unnecessary. His atom was no longer the ‘nice hard fellow, red or grey in colour, according to taste’ that he once, tongue-in-cheek, said he had been brought up to believe in.48
Most alpha particles would pass straight through Rutherford’s atom in any ‘collision’, since they were too far from the tiny nucleus at its heart to suffer any deflection. Others would veer off course slightly as they encountered the electric field generated by the nucleus, resulting in a small deflection. The closer they passed to the nucleus, the stronger the effect of its electric field and the greater the deflection from their original path. But if an alpha particle approached the nucleus head-on, the repulsive force between the two would cause it to recoil straight back like a ball bouncing off a brick wall. As Geiger and Marsden had found, such direct hits were extremely rare. It was, Rutherford said, ‘like trying to shoot a gnat in the Albert Hall at night’.49
Rutherford’s model allowed him to make definite predictions, using a simple formula he had derived, about the fraction of scattered alpha particles to be found at any angle of deflection. He did not want to present his atomic model until it had been tested by a careful investigation of the angular distribution of scattered alpha particles. Geiger undertook the task and found alpha particle distribution to be in total agreement with Rutherford’s theoretical estimates.
On 7 March 1911, Rutherford announced his atomic model in a paper presented at a meeting of the Manchester Literary and Philosophical Society. Four days later, he received a letter from William Henry Bragg, the professor of physics at Leeds University, informing him that ‘about 5 or 6 years ago’ the Japanese physicist Hantaro Nagaoka had constructed an atom with ‘a big positive centre’.50 Unknown to Bragg, Nagaoka had visited Rutherford the previous summer as part of a grand tour of Europe’s leading physics laboratories. Less than two weeks after Bragg’s letter, Rutherford received one from Tokyo. Nagaoka wrote offering his gratitude ‘for the great kindness you showed me in Manchester’ and pointing out that in 1904 he had proposed a ‘Saturnian’ model of the atom.51 It consisted of a large heavy centre surrounded by rotating rings of electrons.52
‘You will notice that the structure assumed in my atom is somewhat similar to that suggested by you in your paper some years ago’, acknowledged Rutherford in his reply. Though alike in some respects, there were significant differences between the two models. In Nagaoka’s the central body was positively-charged, heavy and occupied most of the flat pancake-like atom. Whereas Rutherford’s spherical model had an incredibly tiny positively-charged core that contained most of the mass, leaving the atom largely empty. However, both models were fatally flawed and few physicists gave them a second thought.
An atom with stationary electrons positioned around a positive nucleus would be unstable, because the electrons with their negative charge would be irresistibly pulled towards it. If they moved around the nucleus, like planets orbiting the sun, the atom would still collapse. Newton had shown long ago that any object moving in a circle undergoes acceleration. According to Maxwell’s theory of electromagnetism, if it is a charged particle, like an electron, it will continuously lose energy in the form of electromagnetic radiation as it accelerates. An orbiting electron would spiral into the nucleus within a thousandth of a billionth of a second. The very existence of the material world was compelling evidence against Rutherford’s nuclear atom.
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p; He had long been aware of what appeared to be an intractable problem. ‘This necessary loss of energy from an accelerated electron,’ Rutherford wrote in his 1906 book Radioactive Transformations, ‘has been one of the greatest difficulties met with in endeavouring to deduce the constitution of a stable atom.’53 But in 1911 he chose to ignore the difficulty: ‘The question of the stability of the atom proposed need not be considered at this stage, for this will obviously depend upon the minute structure of the atom, and on the motion of the constituent charged part.’54
Geiger’s initial testing of Rutherford’s scattering formula had been quick and limited in scope. Marsden now joined him in spending most of the next year conducting a more thorough investigation. By July 1912 their results confirmed the scattering formula and the main conclusions of Rutherford’s theory.55 ‘The complete check,’ Marsden recalled years later, ‘was a laborious but exciting task.’56 In the process they also discovered that the charge of the nucleus, taking into account experimental error, was about half the atomic weight. With the exception of hydrogen, with an atomic weight of one, the number of electrons in all other atoms had to be approximately equal to half the atomic weight. It was now possible to nail down the number of electrons in a helium atom, for example, as two, where previously it could have been as many as four. However, this reduction in the number of electrons implied that Rutherford’s atom radiated energy even more strongly than had previously been suspected.
As Rutherford recounted tales from the first Solvay conference for Bohr’s benefit, he failed to mention that in Brussels neither he nor anyone else discussed his nuclear atom.
Back in Cambridge, the intellectual rapport that Bohr sought with Thomson never happened. Years later, Bohr identified one possible reason for the failure: ‘I had no great knowledge of English and therefore I did not know how to express myself. And I could say only that this is incorrect. And he was not interested in the accusation that it was not correct.’57 Infamous for neglecting papers and letters from students and colleagues alike, Thomson was also no longer actively engaged in electron physics.
