by Harry Cliff
The next morning a triumphant Rutherford bounded into the lab, his face beaming, and went straight to tell Geiger that he finally knew what the atom looked like. That very day Geiger got going on experiments to test the rough predictions of Rutherford’s model. After a few more weeks of bombarding gold foils with alpha particles, Geiger found that the angles they came ricocheting off at agreed neatly with Rutherford’s predictions. In March 1911 Rutherford felt ready to show his atom to the world. Appropriately, he chose the Manchester Literary and Philosophical Society, where John Dalton had discussed his own atomic theory a century earlier.
What Rutherford revealed that day was more than just the discovery of the atomic nucleus; he had seen, for the very first time, the shape of the subatomic world. The nucleus contains almost all of an atom’s mass squeezed into a tiny space, tens of thousands of times smaller than the atom itself, orbited by insubstantial clouds of electrons. If you were to blow up an atom to the size of something more familiar, say a football stadium, the nucleus would be about the size of a marble sitting at the center of the field, with the electrons whizzing around somewhere up in the stands.
However, there was a serious problem with Rutherford’s atom. If the atom was really like a tiny solar system, then it couldn’t possibly be stable. It’s a well-established law that when a charged particle is accelerated, in a circle for instance, it gives off electromagnetic radiation. This implied that the electrons should be constantly emitting light, losing a bit more energy with each orbit, until they eventually spiral into the nucleus. Previous attempts to propose solar-system-like atoms had failed for this very reason. J.J.’s plum-pudding metaphor for the atom had been largely motivated by an attempt to find a theoretical arrangement of electrons that was stable against collapse.
The solution to this paradox came from a brilliant young Danish physicist named Niels Bohr, who enlisted a strange new idea known as the “quantum.” In the early twentieth century, Albert Einstein and Max Planck had put forward the idea that light comes in discrete little lumps, or quanta. Taking inspiration from this idea, Bohr argued that electrons could only move around the nucleus in certain fixed orbits, emitting quanta of light as they jumped from one level to another. And as the electrons were restricted to only these levels, like trains moving on circular tracks, it was impossible for them to fall into the nucleus. Bohr’s marriage of quantum theory with Rutherford’s nuclear atom was a triumph, explaining a whole host of phenomena, in particular the peculiar fact that atoms of different chemical elements all give off and absorb characteristic wavelengths of light. In time, Bohr’s theory would lead inexorably to a revolutionary new description of the subatomic world: quantum mechanics. (Much more on that later.)
Augmented by Bohr, Rutherford’s atomic model finally allowed physicists to solve the riddle of the periodic table. In the months following Rutherford’s first paper on the nuclear atom, a young research student working under Rutherford in the Manchester lab, Henry Moseley, made another profound discovery. When Mendeleev had created the periodic table, he had given each element a label known as the “atomic number,” which simply recorded the order in which the elements were listed. Hydrogen, the lightest element, sat at position 1, helium the next heaviest at number 2, all the way up to uranium at position 92. This number seemed to be closely connected to the masses of the different elements, as in general the masses of the elements increased as you moved through the table. However, this wasn’t always the case. There were a few instances where the chemical properties of the elements had guided Mendeleev to place a heavier element in front of a lighter element. For example, cobalt (atomic number 27) came before nickel (atomic number 28), despite cobalt having a bigger atomic mass. This atomic number was thought to be nothing more than a useful label with no physical significance, but Mosely found that the frequencies of X-rays emitted by different chemical elements depended directly on the atomic number, not the atomic mass. The atomic number was more than just a label after all; it was actually the number of positive charges in the nucleus! So hydrogen contained one positive charge and uranium ninety-two positive charges. These positive charges were balanced by an equal number of negatively charged electrons in orbit around the nucleus, leaving the atom neutral overall. The patterns Mendeleev had first spotted playing cards on those long train journeys through Russia were all to do with the way these different numbers of electrons arranged themselves around the nucleus.
But all this raised a question: What was the nucleus made from? Was Prout’s idea that all atoms were made of hydrogen right after all? The fact that the charge of the nucleus was always a whole number multiple of the charge of the hydrogen atom seemed to suggest so, but then, radioactive decay released helium nuclei and electrons, so maybe the nucleus was made of them as well? While much of the physics community embraced the weird and wonderful new world of quantum theory, Rutherford led a relatively small group of physicists on a new journey of exploration. This time their goal was to unlock the makeup of the atomic nucleus itself.
Skip Notes
*1 He insisted on only getting his hair cut and beard trimmed once a year, cultivating a look that was a striking blend of Gandalf the Grey, Leonardo da Vinci, and Fagin.
*2 The same term used by Isaac Newton to describe his hypothetical particles of light.
*3 A third type of even more penetrating ray, named “gamma,” was described by Paul Villard in 1900.
*4 This detector was the forerunner of the modern Geiger counter, which is still used today to measure radiation levels, emitting an ominous clicking sound beloved by directors of disaster movies.
*5 Grandson of the famous naturalist Charles Darwin, author of the theory of evolution by natural selection
*6 Actually, at the time Rutherford wasn’t sure whether the center of the atom had a positive or negative charge—this was only decided on a few years later.
CHAPTER 4
Smashed Nuclei
How are we doing in our quest to make an apple pie from scratch? Well, we’ve certainly got a better understanding of the basic ingredients. The carbon, oxygen, and hydrogen that we extracted at the outset are made of distinct atoms, and every atom has the same essential structure: an unimaginably tiny nucleus containing almost all of the atom’s mass, with much lighter subatomic particles called electrons buzzing around it. The whole thing is held together by powerful electrical forces, which bind the negatively charged electrons to the positively charged nucleus, while some mysterious quantum wizardry stops the whole thing from collapsing in on itself and taking all matter in the universe with it.
We’ve also figured out what makes a carbon atom different from say an oxygen atom—it’s all down to the number of positive charges in the nucleus, which attract an equal number of negative electrons to leave the atom neutral overall. Hydrogen, we now know, is the simplest atom of all—its nucleus has a charge of +1 and is orbited by a single electron. Carbon, on the other hand, has six positive charges in the nucleus and six electrons, while oxygen has eight. Moseley discovered that the number of positive charges in the nucleus turns out to be exactly the same as the atomic number, which chemists used to think was just a label telling you where to find a given element in the periodic table. Since the chemical properties of the elements vary in a regular way as you move through the table, this tells us the chemical properties of an atom must be completely determined by the number of positive charges in the nucleus.
Now that’s all bloody marvelous, but despite the discovery of the electron and the nucleus, we still don’t know how to make hydrogen, carbon, oxygen, or indeed any of the other elements in an apple pie. If it’s the number of positive charges in the nucleus that determines whether an atom is carbon or uranium, we need to figure out what lies inside the atomic nucleus if we’re going to find recipes for all the elements in the periodic table.
When the Rutherford-Bohr model
of the atom was established around 1913, the nucleus remained shrouded in mystery. However, it was pretty clear that the nucleus wasn’t just a new version of the indivisible atom, shrunk tens of thousands of times in size. Marie Curie and Ernest Rutherford were both convinced that alpha, beta, and gamma radiation emerged from the nucleus itself, which meant that the nucleus must be made of even smaller things. The question was, what?
Since the alpha and beta particles that came flying out of the nucleus in radioactive decay were just helium nuclei and electrons, it was natural to assume that the nucleus contained helium nuclei and electrons. Lingering in the background was William Prout’s old idea that hydrogen atoms were the basic building blocks of all the heavier elements, but that idea had hit the rocks thanks to awkward elements like chlorine, whose atomic masses weren’t whole number multiples of the mass of hydrogen.
It was all a bit of a muddle. Before physicists could make any progress, they were going to need some new experimental clues, but getting information out of the nucleus was going to be no easy task. It would take two more truly heroic experiments, the first of which was carried out by the same scientist who discovered the nucleus, that boisterous, booming force of nature, Ernest Rutherford.
THE NUCLEUS SPLINTERS
In 1914, as the outbreak of war brought scientific research to a grinding halt across Europe, Rutherford abandoned his radioactive experiments to work on submarine detection for the Admiralty. However, even a world war couldn’t keep him away from his one true love for long. Although now in his midforties and recently elevated to Sir Ernest Rutherford, his curiosity was as strong as ever. The discovery of the nucleus had opened up a brand-new frontier and he was itching to explore it.
His keen scientific nose had already caught the scent: a nagging problem that could be traced back to that Sunday evening in 1910 when he had first shared his vision of the nucleus with the young Charles Galton Darwin. During their after-dinner chat, Darwin had pointed out that Rutherford’s idea implied that if you fired alpha particles into gases of light elements like hydrogen, then occasionally the much lighter hydrogen nuclei should be knocked out of the gas, like a snooker struck by the cue ball.
Just before the war broke out, Ernest Marsden, the young researcher who had first seen alpha particles bouncing backward off gold atoms, had followed up on Darwin’s suggestion by firing alpha particles into ordinary air. Now air contains a certain amount of water vapor (H2O), which of course contains hydrogen atoms, and just as Darwin had predicted, Marsden saw hydrogen nuclei being kicked out of the gas by the alpha particles. However, he was puzzled to find way more hydrogen nuclei come flying out than he expected based on the amount of water vapor in air. Somewhat stumped, Marsden eventually made the rather unconvincing suggestion that the radium atoms that produced the alpha particles must also be shooting out hydrogen nuclei.
Rutherford wasn’t persuaded. Unfortunately, Marsden had left Manchester for a university post in Wellington, New Zealand, in 1915, only to return to Europe to fight for the British Army in France. After writing to ask for Marsden’s permission, Rutherford picked up where his former student had left off, gradually spending more and more time on the experiments as the war ground on. The once bustling Manchester lab was now more or less deserted, and Rutherford found himself working down in a dark basement, accompanied only by the laboratory steward, William Kay.
The apparatus he worked with was an absolute classic of Rutherfordian simplicity: a battered brass box about 10 centimeters long with a radioactive lump of radium at one end and some pipes that could be used to feed in various gases. At the end farthest from the radium was a small window covered in a thin metal foil, which blocked alpha particles emitted by the radium but allowed the more penetrating hydrogen nuclei to escape. Just outside the window was a zinc-sulfide screen, which produced characteristic flashes of light when it was hit by the escaping hydrogen nuclei.
Again, the observations were made in almost total darkness, peering down a microscope at the zinc-sulfide screen. It was eye-straining work: the flickers of light produced by hydrogen were much fainter than those from alpha particles, and an observer could only keep counting for a couple of minutes before their eyesight became unreliable. It was even possible to fool yourself into thinking you’d seen a hydrogen flash if you watched for too long. Rutherford and Kay worked in shifts of around two minutes each, one counting while the other rested their eyes. Rutherford’s notebooks from the time tell a story of numerous experimental difficulties, from stray light reflecting off the metal foils to suspected contamination in the gas supplies, including frequent remarks like “No observations because of poor eyesight.”
For a long time, he struggled to make sense of what he was seeing. Were the hydrogen nuclei coming from contamination in the gas? Perhaps they were somehow being produced when the alpha particles hit the metal foil at the end of the brass box? Or maybe they came from the radium itself as Marsden had suggested? Again, he was forced to put his work on hiatus to go on a mission to the United States in the summer of 1917, but it turned out to be one of those useful breaks when stepping away from a problem lets your mind slowly work out the solution in the background. When Rutherford got back to the lab in September, he had the answer—the hydrogen nuclei weren’t already in the gas, they were being created when alpha particles collided with atomic nuclei in the gas.
In a burst of intense work in October and November, Rutherford tried out a variety of different gases, from ordinary air to pure carbon dioxide, nitrogen, and oxygen. When alphas were fired into air, the screen lit up with the flickers of hydrogen nuclei, but when he used pure carbon dioxide or oxygen, he saw almost no flashes at all. Pure nitrogen, on the other hand, sent an even larger torrent of hydrogen nuclei careering into the screen than ordinary air. Having eliminated every other possibility, Rutherford was forced to a staggering conclusion: the alpha particles were smashing the nitrogen nuclei apart, sending hydrogen nuclei flying out like shrapnel from an explosion. Well aware of just how momentous this discovery was, he completely absented himself from his submarine work, writing to his superiors at the Admiralty, “If, as I have reason to believe, I have disintegrated the nucleus of the atom, this is of greater significance than the war.”
After a year of checking and rechecking his results, he felt ready to draw his final, dramatic conclusion: “The hydrogen atom which is liberated formed a constituent part of the nitrogen nucleus.” Rutherford had finally found the first convincing evidence that the chemical elements were all ultimately made from hydrogen. He would later give the hydrogen nucleus a new name, confirming its status alongside the electron as one of the fundamental building blocks of all atoms, calling it the “proton.”*1 It was the climax of a long story that goes all the way back to John Dalton’s measurements of the relative masses of different atoms, which in turn had inspired William Prout to imagine that all the chemical elements might be built from hydrogen. Not only had Rutherford resurrected Prout’s hypothesis, he had opened a door to understanding the ultimate origins of the chemical elements. Armed with the proton and the electron, physicists could at last begin to imagine how the elements might be built up one by one, from helium all the way to uranium. And he had achieved all this despite wartime shortages, working down in the basement of a deserted laboratory with nothing more than a battered brass box, a few crumbs of radium, and the loyal assistance of William Kay.
But there was a dirty great fly in the ointment: the riddle of elements like chlorine. If all the elements were really made of hydrogen, why did chlorine have an atomic mass 35.5 times that of hydrogen? In fact, a possible way out of this conundrum had already been found by Rutherford’s old colleague from his days at McGill, Frederick Soddy. In 1913, Soddy made the curious discovery of a number of new radioactive elements that appeared to be chemically indistinguishable from other well-known nonradioactive elements. Among them was a radioactive form of lead, the garden
variety of which isn’t radioactive at all. That must mean that there could be multiple versions of the same chemical element, occupying the same slot in the periodic table but differing in their radioactivity. Soddy dubbed these chemical copies “isotopes.”
Now this all raised a tantalizing possibility: What if Soddy’s isotopes had the same nuclear charge, making them the same chemical element, but different atomic masses? Perhaps there were really two different isotopes of chlorine, one with a mass of 35 and another with a mass of 36, that when mixed together made it look as though chlorine had an atomic mass in between the two. It was an appealing idea, but it was hard to imagine how you could separate two different isotopes and measure their atomic masses. After all, isotopes were, by definition, chemically indistinguishable.
However, there is a way to do it. Back at the Cavendish Laboratory, the chemist Francis Aston had been beavering away down in a gloomy cellar on a brand-new instrument capable of weighing atoms with unheralded precision: the mass spectrograph. Aston’s new invention could be used to fire ions of different elements through a sort of lens made from electric and magnetic fields, focusing them onto different locations on a photographic strip depending on their masses.
The spectrograph was a revelation, and Aston soon used it to show that chlorine, the element that had scuppered the idea that hydrogen was the building block of all atoms, was actually a mixture of two isotopes: roughly three parts chlorine-35 to one part chlorine-37, giving an average mass of 35.5. By 1922 he had discovered forty-eight different isotopes of twenty-seven different elements, including six different isotopes of xenon alone. All of the elements that Aston weighed turned out to have masses that were whole number multiples of hydrogen’s,*2 a spectacular result that combined with Rutherford’s coup of knocking protons out of the nucleus all but confirmed that protons were the building blocks of the nucleus.