Chapter Two

A Couple of Remaining Questions

By 1897 the end of the century was in sight, but the rash of increasingly fantastical scientific discoveries had only begun. The existence of atomic elements was established, and Henri Becquerel had found an odd property in the element uranium. It would leave a ghost image on a photographic plate, an effect still sought by spiritualists since the invention of photography. Instead of a filmy aura around a spirit medium or a glowing shape caught unaware in a darkened room, the ghost was somehow within the uranium atom. There was no clear explanation for Becquerel’s newly discovered rays, other than it was just an unusual property peculiar to uranium.

Back in England, at the Cavendish Laboratory, Sir Joseph John “J.J.” Thomson was involved with the ever popular vacuum-sealed glass tubes and high-voltage coils. He was going after what seemed the last frontier in science, to discover the nature of the “cathode rays.” With thousands of volts across two metal electrodes in the vacuum tube, rays would stream between the cathode and the anode. Röntgen had used them to produce his X-rays, but what exactly were they?

J.J. Thomson was born in 1856 in Cheetham Hill, a suburb of Manchester, England. His Scottish parents encouraged him to be practical and get an education in engineering, and he signed up for an apprenticeship with a local company. He was put on a two-year waiting list. Being naturally impatient, at the age of 14 he enrolled at the Owens College in Manchester, which had a gloriously impractical science faculty, concentrating on the whimsical frontiers without a nod to profitable industrial applications. On the recommendations of his professors he transferred to Trinity College in Cambridge in 1876 and four years later graduated with a BA in mathematics, Second Wrangler. In 1883 he obtained his MA in mathematical physics, became Cavendish Professor of Physics a year later, and in 1890 he married the daughter of Sir George Edward Paget, Regius Professor of Physics at Cambridge. Thomson now had it made, and thus he applied himself vigorously to the mystery of the cathode rays.

It was still the “sealing wax and string” era of physics, as red, Bank-of-England wax was used liberally to seal up leaking vacuum apparatus and delicate, blown-glass tubes were held up with strings. J.J. was somewhat fumble fingered, and his young assistants were known to strongly encourage him not to touch the equipment. He was, however, a genius at designing experiments and interpreting results. His three, sequential cathode-ray experiments would answer all his questions. For those in the field, for three experiments in a row to yield informative results is nothing short of a rarity.

The cathode rays involve a negative charge, because they originate at the cathode, or negative electrode, in the tube and vanish into the anode, or positive electrode. Thomson wished to know if this negative charge could be separated from the rays that caused the fluorescence on the glass, and for his first experiment he had a special tube built. It started out as a standard-type vacuum tube, with a metallic cathode, an empty flight path of a couple of inches, followed by an anode. The anode had a horizontal slit cut in it, so that rays would continue to fly out the positive end in a thin, wide beam. On the positive end the glass tube was extended. The beam would cross through the vacuum and hit a third electrode at the end. Thomson pumped down the tube, applied high voltage from a Ruhmkorff coil across the first two electrodes, and connected an electrometer between the first and third electrodes. As he had suspected, an electrical current flowed, proving that the rays at least had a component of pure electricity. He then applied a horse shoe magnet across the long flight path of the beam. The rays bent down in the magnetic field and missed the third electrode, hitting the glass and making it fluoresce. The electrical current stopped registering on his electrometer. He had now proven that the electrical charge and the rays were one and the same, inseparable from each other.

J.J. was pretty sure what the rays were, but he needed another experiment to nail it down. He had another tube built.¹² In this one he put a round hole in the positive electrode, to give him a pencil-thin beam of rays, and he eliminated the third electrode. Instead, he had the end of the tube lined with phosphorescent paint. In the flight path were mounted two parallel metal plates, with connecting wires sticking though the sides of the tube. If he could bend the beam with a voltage across the plates, it would further confirm that the rays were streams of electrical charge. This experiment had been tried many times before with no success, as tubes were being blown and wired up all over Europe, but Thomson thought he understood the failures. For this effect, the path had to be unusually long, and in an imperfect vacuum stray air molecules would get in the way of the beam. The Cavendish vacuum pumps were the best in the business, and there was plenty of sealing wax. A near-perfect vacuum was achieved in Thomson’s setup, thus alleviating the failures that had plagued other scientists.

He could see the ray beam from its glowing spot on the end of the tube.¹³ He connected a battery across the deflection plates, and instantly the spot moved toward the positive plate and away from the negative plate. J.J. was certain he knew the answer now, but there was still one more confirmation experiment necessary. He had a third tube built, just like the last one, only with no deflection plates. On this setup, two external magnetic coils replaced the internal plates.

Running the same experiment using a controlled magnetic field across the tube deflected the ray in a predictable manner. Having seen the results of both electrostatic and magnetic deflection of the beam, Thomson announced a bold conclusion: The cathode rays are composed of tiny, electrically charged particles he called “corpuscles.” How tiny? From precise observations of the angle with which a beam could be bent using a known voltage he calculated the charge-to-mass ratio. Not exactly knowing the charge, he could only guess at the mass of a single particle, but he could now illustrate that it was very, very small. The mass was almost negligible, and the charge on each and every particle was negative.

These particles were being ripped off atoms in the cathode and flung down the length of the tube, attracted to the positive charge on the anode. Atoms have zero residual charge, so what was left after the corpuscles were stripped away had to be a fairly heavy remnant, with a positive charge equal to that of the missing negative charge. From here Thomson skipped to the next conclusion. The atom, he theorized, was composed of a relatively heavy, large ball of positive electrical charge, with little negative corpuscles imbedded in the surface.¹⁴ It came to be called the “plum pudding” model of the atom, and it was revolutionary. It sent physics caroming off in a modified direction, like a startled school of fish. Thomson had proven that the atom, long thought to be indestructible, could be disassembled bit by bit, negative particle by negative particle. His charged corpuscles would eventually be named “electrons,” and he won the Nobel Prize in physics in 1906.¹⁵

By the application of pure genius and thorough lab work, physics was now detecting things smaller than the unseen atoms, using indirect implication. As the discoveries were becoming more abstract and farther removed from normal, human interactions with the physical world, the study of physics was becoming more organized, standardized, and systematic, with better peer review and cross-checking. Competition for finding the next discovery was becoming fierce, yet the collegial interlocking of findings was productive.¹⁶ At the end of the century another piece of the fantasy would lock into place, as the work of a woman from Poland intersected the physics time line.

Manya Skodowska was born in Warsaw, in the Tsarist Russian Empire sector of Poland, in 1867. Her father taught physics and her mother taught mathematics, and she was destined to achieve a higher education, but the Russian controls made it increasingly difficult for her. She was Polish and she was female, and these factors counted against her for admission to a college, and so she enrolled in an outlaw institution, the Floating University. Classes were only at night, and the campus was nonexistent. Manya studied mathematics and chemistry while acting as a governess for the children of wealthy families. Against all odds, with only this secondary education, she was accepted to the prestigious Sorbonne in Paris. In 1891 Manya moved to France, changed her name to Marie, and assimilated into French culture as best she could.

Marie overcame her spotty preparation and graduated first in her class with a master’s in physics in 1893, and a year later another master’s in mathematics, second in her class. With the celebrated Henri Becquerel on her thesis committee, she was accepted into the doctoral program.

In the late 19th century the industrial arts had come to appreciate the value of academic research, at least at a subsistence level. Funding from manufacturers dribbled down into the universities, and Marie secured grants for research on the magnetic properties of steel. Although she always missed Poland, she found the atmosphere in Paris conducive and even encouraging to science. She became laboratory chief at the Municipal School of Industrial Physics and Chemistry, where she shared an unusually dingy laboratory space with a like-minded French scientist named Pierre Curie. He was ten years her senior, and he had given up on love long before, but nature took its course, and Marie Skodowska became Madame Curie in July 1895 in a civil ceremony. To the chagrin of her family, the bride wore a dark blue ensemble, which for many years after served as a lab-coat.

Together the two scientists would share the excitement of discovery as they worked, with minimal funding, in a barely adequate, drafty structure fondly referred to as “the miserable shed.” It had been a medical school dissection room in better days. Water ran down the walls when it rained.

Marie grew interested in Becquerel’s new rays emanating from uranium. Röntgen’s X-rays were still the rage, and the medical applications for the electrically produced rays drew attention away from anything else. Marie was casting about for a Ph.D. thesis topic, and Becquerel’s rays were a perfect mystery to be solved. The concept of nuclear decay had not formed, because the nucleus was still an unknown structure, as all we still knew of atomic structures was the positive and negative charges, and there were no good theories as to the source of these invisible rays. A popular guess was that they were reflections of cosmic rays raining down from the cosmos. Madame Curie had an edge. Her new husband, Pierre, and his brother, Jacques, had developed an advanced electrometer, capable of detecting slight currents of electricity. With this impressive device, resembling a beer can sitting on a wooden bar-stool, she could detect and quantify the Becquerel rays from the air-ionization they produced.

Marie managed to wrangle a ton of pitchblende, a rich uranium ore, from an Austrian mining company and have it dumped at the miserable shed. She subjected it to the usual scientific tortures, and found that Becquerel was correct. The emanations were constant, regardless of whether the uranium was powdered or pulverized, pure or in compound, wet or dry, or subjected to heat or light. Minerals having the highest density of uranium emitted the most rays. Proceeding with appropriate intellectual caution, Marie began to form a hypothesis. Although it was difficult for her to consider in a scientific world in which the atom was thought to be indestructible, the rays seemed to come from the unique structure of the core of the uranium atom, as if the atom itself were deconstructing. Marie started referring to Becquerel rays as radiation, or radioactivity. The distinguished Henri Becquerel was not amused. Although he had been very helpful to the Curies advancing their careers, Becquerel was sensitive about his rays, and a tension developed.

In 1898 Marie discovered that another element emits radiation. It was thorium, and she was now firmly convinced that radioactivity was an atomic property, not something unique to uranium as previously thought. At this point Pierre realized that what Marie was doing was more fascinating than his work in crystal symmetries, so he dropped his research and joined in. Marie had just found something very interesting: Pitchblende, the ore from which uranium was extracted, was more radioactive than the pure uranium extracted from pitchblende. There was apparently something else in the ore.

Chemically processing barrel after barrel of pitchblende, the Curies were able, after considerable effort, to identify two new radioactive elements existing in the same mineral with uranium. The first element they discovered Marie named polonium, for her native Poland. The second she named radium, after her new term for the rays. Although they did not fully realize it, the Curies had discovered products of the decay chain of uranium. These are the heaviest naturally occurring elements, and are very unstable, and they slowly break down into lesser radioactive elements, the last of which is non-radioactive lead. With tremendous effort the Curies managed to refine the radium down to a sub-gram quantity of pure metal. In its undiluted form, it is an intensely radioactive material, and it produces an eerie blue glow as its radiation rips through nitrogen in the air. Marie carried around a couple of vials of it in her lab coat, just so she could look at the pretty color, and Pierre would place shallow dishes of radium compound around the courtyard outside the shed to impress visitors at night.

Her Ph.D. thesis of 1903, “Research on Radioactive Substances,” probably contributed more to the advancement of science than any previous piece of doctoral writing, and the same year she shared the Nobel Prize in physics with her husband and with her former mentor, Henri Becquerel, for the discovery of natural radiation.

The ultimate power and the danger in the blue glow was realized only dimly. Being in close, constant contact with highly radioactive materials significantly affected the Curies’ health. Their fingers and hands were burned, swollen, and peeling, and Pierre was particularly sickened by the radiation. He was in constant, severe pain.¹⁷ The connections between acute and continuous exposure to the radiation and sickness could not have been more obvious. What were they thinking? Even Röntgen noticed bad effects from his X-rays early on, and he started using lead shielding around his apparatus. Their zeal and mutual love of lab work and each other blinded them to any effects on their personal health.

Science was still in the fantasy era. Its practitioners studied invisible rays and were enthralled by god-like forces and effects. Reality had not really sunken in. With the speed of unfettered commerce and advertising, by 1904 a radiation industry sprang onto an excited public. Popular products ranged from “Thor-Radia,” a beauty cream based on thorium and radium, to “Doramad,” a thorium based toothpaste, both of which promised to kill germs and improve health. In retrospect, the concept of using these products is horrifying.¹⁸

Medical demand for radium for treating cancer and skin disorders was immediate and huge, keeping the prices for radioactive materials fabulously high. Fame and funding for the Curies was inevitable, and by 1919 Madame Curie was able to open a state-of-the-art laboratory at the Sorbonne, The Radium Institute. There would have to be a better understanding of the nature of the radiation before safety would become a serious issue. Clear, non-fantastic explanations were in order, but things were even more complicated than they seemed, and an understanding of the atomic structure would have to be dragged out of nature in small increments. The next handful of knowledge would be grabbed by a talented young man from New Zealand named Ernest Rutherford.

Ernest Rutherford, 1st Baron Rutherford of Nelson, answering to the name Lord Rutherford, was born somewhat near Nelson, New Zealand, in 1871. His parents had moved there from Perth, Scotland, to raise flax and children, and his early years were spent as a colonial at the rough edge of the empire. Somehow young Ernest showed academic potential, and he won scholarships to Nelson College and then to the University of New Zealand. At the age of twenty two he earned an M.A. with the double major of mathematics and physical science. He stayed on to work toward a B.Sc. with some academic research, and in 1893 became interested in studying Hertz’s electromagnetic waves and their influence on magnetized iron. For several months he held the record for the distance of a received radio transmission.

His work was original and ingenious, and an evaluation of it resulted in a scholarship for post-graduate work at Cambridge. A famous quote was his response when his mother shouted the news to him in the garden. He launched his spade away and exclaimed, “That’s the last potato I’ll ever dig!” His prediction proved accurate, as he went on to become the leading force in British science and the man who started nuclear physics. He showed up at the Cavendish Laboratory at Cambridge in September 1895, reporting directly to none other than J.J. Thomson.

Thomson announced his discovery of the electron in 1897 and Rutherford, seeing his work in radio as a dead end, jumped into the exciting new field of radioactivity research. The Curies in Paris had established that both uranium and thorium are radioactive, but they had not characterized the radiation itself. Rutherford began by studying the rays, and by methodical research he identified two types of distinct radiation. The first had little range. It was stopped by almost anything solid and by a short flight in air. He named it alpha. The second had greater penetrating power and better range in air. He named it beta.¹⁹

With this short triumph under his belt, Rutherford was invited to become professor of physics at McGill University in Montreal with a new, fully equipped laboratory, lavish funding, and a learned colleague in chemistry named Frederick Soddy. Rutherford accepted the position, and moved to Canada in 1899. Barely a year later he presented Soddy with a fascinating puzzle. It seemed as if some sort of gas was leaking out of radioactive thorium. Soddy jumped at the problem and offered to find its identity through chemical analysis.

The gas had no chemical properties whatsoever. Physics in 1899 was pure irony, and it had just given up another invisible secret, not by direct measurement, but by the absence of a measurement. The fact that there was no chemical response to the gas left one possibility: The thorium was turning into chemically inert argon. Stunned and excited, Soddy and Rutherford plunged immediately into a study of all known radioactive elements. They measured radioactivity by literally counting rays as they hit a phosphorescent screen, causing a slight flash visible in a microscope eyepiece. It was a terribly tedious measurement, but by timing a counting session with a watch they were able to characterize radiation intensity as the number of flashes per time interval. They found that different elements had different characteristic radiation rates, and that a given element’s rate would fall exponentially with time. The radioactivity of a substance would fall by half in a predictable period. Soddy named the interval “half-life.” The radioactivity would technically never disappear, but it would drop by half in a given length of time, then the remainder would drop by another half after the same interval, then by half again, and so on for all time.

Over the next two years, Rutherford and Soddy closely studied radioactive decay, concluding that that the argon gas emanation was indeed due to decay of the thorium atom. Upon emitting a radiation, an atom would down-shift or transmute to a lower, lighter element, and the product of the decay could be radioactive as well. Furthermore, one chemically identified element could have two or more sub-species, each with a characteristic radioactivity half-life. Soddy named the chemically identical sub-types “isotopes,” and there were now two ways to identify a radioactive substance.²⁰ A chemical analysis could identify the element, and a measurement of the half-life could identify the isotope of the element. The half-life of the only known radium isotope measured at 1,620 years. Two decay products of the thorium transmutation were in quantities too slight to be chemically analyzed, but one had a half-life of 27 days, and the other of 22 minutes. Some decay products disappeared almost instantly, with half-lives of less than one second. The list of decay product isotopes seemed endless. Rutherford and Soddy spent the next two years at McGill measuring radiation and recording results.

Rutherford had an uncanny intuition, and he suspected that beta rays were a naturally occurring form of the cathode rays characterized by J.J. Thomson. He was able to prove this hypothesis in a most elegant way, by using magnets and electrostatic fields to bend a collimated beam of beta rays in a vacuum, just as had been done with the electrons in Thomson’s cathode ray tubes. Beta rays are exactly cathode rays, or streams of electrons shot from deeper within the decaying atoms.

In 1903 Rutherford and Soddy collaborated on an important paper: “Radioactive Change.” In it they presented the first experimentally measured energy produced by the atomic decay of atoms. The numbers were astounding. They found that the energy released by the decay of one gram of radium was no less than 100,000,000 calories, and it was probably closer to 10,000,000,000 calories per gram. To put that measurement in perspective, consider the rocket fuel used in the upper stages of the Apollo moon rocket. It releases more energy per gram than any other known chemical reaction, and the fuel consists of liquid hydrogen burning in liquid oxygen, producing just 4,000 calories per gram. The obvious disparity in rocket fuel and radium energy release shows the difference between the electron binding force, as used in chemical reactions, and the force that keeps the atomic nucleus glued together. It was positively frightening. In an occasional off moment Rutherford contemplated the implications of the energy coiled up in radioactive elements, and the possibilities of artificially induced decay.

“Some fool in a laboratory,” he mused in his notes, “might blow up the universe unawares.”

Meanwhile, Rutherford was not the only one bucking for a Nobel in physics. Others were concocting wild but still rational theories, in keeping with the anything-goes attitude of nature in the small world of atoms. Philipp von Lénárd, a physicist at the University of Kiel in Germany, was still fiddling with Ruhmkorff coils, working on the project that Röntgen had abandoned back in ‘96 to bring cathode rays out of the tube and into the laboratory.²¹ Lénárd built a tube having a metal window thin enough to allow the rays to blow out into the air, and he was able to detect them using an external fluorescent screen. There was, however, an odd effect. He noticed that some of the rays were scattered when they went through the window. As a well informed scientist he was on board with J.J. Thomson’s theory about the cathode rays being tiny particles. Some particles would make it straight through the metal, but some seemed to hit something and be diverted or absorbed. If a particle could be absorbed by the metal window, then why weren’t they all absorbed? How could anything make it through untouched if the window was capable of stopping a particle in its tracks?

It was hard to break with Thomson’s well regarded model of the atom, involving a thick blob of “plum pudding.” A solid object, after all, was obviously hard, massive, opaque, continuous, and homogeneous. As Lord Rutherford once quipped, “I was brought up to look at the atom as a nice hard fellow, red or gray in color, according to taste.” Lénárd threw caution to the wind and announced an unacceptable conclusion based on his latest cathode ray experiments, that the atoms of which matter is built are composed of mostly vacuum. To make it clear, he added this example: A cubic meter of platinum is about as empty as outer space.

It was a busy year in physics. On the other side of the world, at the University of Tokyo, a physics professor named Hantaro Nagaoka was working on his own model of the atom, and it also did not involve plum pudding.²² Nagaoka had studied physics at universities in Vienna, Berlin, and Munich between 1892 and 1896, and he was particularly taken by James Clerk Maxwell’s first college essay from 1859, “On the Stability of Saturn’s Rings.” Maxwell had theorized that Saturn’s rings are composed of rocks, infinitesimally small compared to the massive planet, but still mechanically free and orbiting the planet as individual satellites. The weak gravitational coupling among rocks causes them to form rings around the equator as they orbit. If he modeled the electrons as rocks and the central positive charge as a planet, Nagaoka could visualize the atom as an invisibly small Saturn.

Nagaoka had a bold and interesting theory with a fatal flaw. Maxwell’s model of Saturn depended on there being a gravitational attraction between the planet and the rocks to hold them in orbit, and a weak gravity among the rocks to hold them in rings. The atom had electrostatic force between the positive center and the negative electrons to hold them in orbit, but there was a very strong repulsive force among the electrons, pushing them apart. The Nagaoka model would rip to pieces immediately, sending all electrons skipping out of orbit on their own initiative. Such behavior had not been observed in atoms, so what was holding them all together in the cohesive unit of “the atom?”

Rutherford was still at McGill in 1906, admitting that he could not find the flaw in Philip Lénárd’s pronouncement from three years before concerning the great void separating atoms. He was pondering it as he built another alpha particle experiment. This one would measure the magnetic deflection of alphas so he could confirm the charge on them. He scrounged up the biggest magnet he could find, set it up on a bench, and placed an alpha particle source behind a metal shield, with a thin slit to give him a narrow beam. The experiment was straightforward, just shooting alpha particles through the magnetic field. He used a photographic plate to record where the alphas landed after traversing his magnet. The best angle he could get was less than one degree. He then put a thin piece of mica over the slit in the metal shield, trying to improve the focus of his alpha beam. The alpha particles now made an odd, blurred image on the film as they blasted through the mica. To his amazement, one pathetic sliver of mica was deflecting alpha particles through an angle of two degrees, twice as much as he was getting with the best magnet he could find. He scribbled a calculation. To do that to the alpha particles would require an electrical field of 100,000,000 volts per centimeter of mica. There was something very intense in that flake of mica. He wasn’t sure what.

Back in England in 1907, Rutherford took the physics chair at the University of Manchester. By now he practically owned atomic physics, and he was not finished peeling back the veil. In 1908 he devised an experiment to prove another of his intuitions, that alpha rays are actually helium atoms with their electrons sheared off. His alpha experiment is noted in the history of physics as one of the most artistically perfect procedures ever performed. He had a glass tube made, with walls thin enough to allow alpha rays to pass through. He evacuated the tube, back-filled it with radon gas, a known alpha emitter, and flame-sealed the end. This tube then fit inside a larger tube with thick glass walls. He pumped all the air out of the outer tube and sealed it with flame. A light-spectrometer was aimed at the space between the thin, inner glass tube and the thick, outer glass tube.²³ It confirmed that there was nothing in the space but hard vacuum. Rutherford waited a few days, and then checked the spectrometer again. It registered helium. The alpha rays had escaped the first tube and had been confined by the second tube. Soddy had been half right in his evaluation of the gas generated by thorium. It was an inert gas, but it was not argon and, technically, it was not a gas escaping from a mineral. It was a chunk broken off the interior of an atom, which just happened to be indistinguishable from helium gas. The name of the radiation was adjusted from alpha rays to alpha particles, and Rutherford announced his newest finding to the audience in Stockholm as he accepted his Nobel Prize in Chemistry in 1908 for his work at McGill.²⁴

In 1910 Rutherford had one more experiment on his mind, and this one would shake physics like a 6.5 earthquake. It is forever enshrined in the pantheon of experimentalism as “The Gold Foil Experiment.”

He set up his colleague Hans Geiger, of Geiger counter fame, and an undergraduate, Ernest Marsden, to test out Lénárd’s troubling claim concerning the deflection of cathode rays. Rutherford decided to try the experiment using his heavy, energetic alpha particles, having a double positive charge. They built a tight beam of alphas using a speck of radium at the end of a long metal tube. All the alphas would be absorbed by the metal except those heading straight down the bore of the tube, like bullets out a gun barrel. They planned to aim it at very thin samples of several materials, such as aluminum, silver, and platinum, but first they decided to try gold, because it was the most malleable metal known, and could be flattened into very thin sheets. They set up a piece of gold foil vertically on the bench, with the alpha source aimed at a 45-degree angle. On the other side of the foil was positioned a glass plate painted with zinc sulfide. It would glow when an alpha particle hit it, and they would view it with the lights off using a microscope. It was an experiment design that could not fail, designed by the master himself. Geiger and Marsden turned off the lights and started counting alpha particles.

Something was wrong. Rutherford checked in, just to see how the setup was going. Marsden was embarrassed, and Geiger was annoyed.²⁵ They had predicted a spray pattern as big as four degrees wide, which is a fairly tight beam, but alphas were all over the place. Obviously, something was wrong with the alpha source. Not to worry. They would have it fixed shortly. They adjusted and fiddled with the alpha tube, but there was nothing wrong with it that could be adjusted. Rutherford came by again, checking on the experiment. Marsden reported that nothing was working right.

Rutherford had a hunch. They had logically been looking at the alpha particles as they blasted through the gold, looking for modest deflections. Rutherford told them to try looking on the other side of the foil, on the back. Marsden moved the microscope and the scintillation screen to the back of the foil. He slid a lead screen between the alpha tube and the screen, just to be certain that no stray particles out of the tube would contaminate his observations. He turned off the lights, let his retinas adjust to the dark, and looked through the eyepiece. He was thunderstruck by what he saw. The heavy, energetic alpha particles were hitting the gold foil and ricocheting backwards, through angles larger than 90 degrees. He scrambled up the steps leading to Rutherford’s private office, met him coming down, and broke the news.

Rutherford took it well.²⁶ The finding was incredible, but true. A piece of gold 0.00002 inches thick was deflecting alpha particles through an angle that no magnet on earth could manage. As Lord Rutherford recalled 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. On consideration I realized that this scattering backwards must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the atom was concentrated in a minute nucleus.

Lénárd had been absolutely correct in his loopy observations about the predominance of absolute vacuum in matter, and Rutherford tossed away Thomson’s plum pudding model. In Rutherford’s new model, mass was like a galaxy of tiny suns, and alpha particles were like comets. Send a comet through a matrix of suns, and if it gets too close to one, it will whip around and come back in nearly the same direction from which it approached, in a tight elliptical orbit, just like the alpha particles in the gold foil.²⁷ Around each sun he imagined tiny planets orbiting, called electrons, and in between was a lot of hard vacuum. There was a certain poetic beauty to his model, with the smallest structures in the universe, the atoms from which everything is made, having the same configuration as the largest structures, the solar systems. Rutherford had thus identified the fractal geometry of the cosmos. All of the power and the forces he had observed were concentrated into a very small and power-dense space at the center of each atom. He called it the nucleus, and a new form of physics was born.

As an occasion to announce his discovery Lord Rutherford chose a meeting of the Manchester Literary and Philosophical Society, a time-honored local organization, open to the public, on March 7, 1911. First up was a report from a Manchester fruit importer concerning a rare snake he had found in a shipment of bananas from Jamaica. He held up the snake to confirm his discovery. Next was Ernest Rutherford, to talk about the latest activity from the physics laboratory. To the young university students in the audience who understood what he was talking about, the lecture was a life-changing experience. It was like hearing the truth spoken aloud for the first time.

Rutherford’s concept of the atom would become an enduring visual icon. You have seen it a million times, a nucleus of balls stuck together, usually blue and yellow, with a few red balls, representing electrons, spinning around them. The red balls trace elliptical orbits behind them, with blurring tails indicating that the motion is fast. This bold diagram has been used everywhere from the seal of the Atomic Energy Commission to the logo of the American Nuclear Society. It is on everything from the flag of the International Atomic Energy Agency to the uniforms of the Albuquerque Isotopes. It is on my business card. It is in the symbol table in your computer, as Unicode number 269B, between the staff of Hermes and the fleur-de-lis.

With Rutherford’s announcement, science reached that rare plateau of nirvana in which everything was finally in balance. The universe was fully understood in 1911, from the motion of the planets to the structure of the atom. Using the equations of motion and gravity formulated by Isaac Newton, the positions of heavenly bodies could be calculated with precision. If you wanted to know the position of a moon of Jupiter the day and hour Galileo discovered it in 1610, it could be calculated on a sheet of paper using a pencil. Using the same equations, the orbital deflection of Haley’s comet around the Sun or an alpha particle around a nucleus could be calculated.

All was well. There were just a couple of remaining questions. Firstly, there was the small matter of a formula derived by a minor mathematician in Switzerland named Johann Jakob Balmer. Back in ‘85 Johann had cooked up a curious equation that seemed to predict the positions on the visible spectrum of the color bands coming from hydrogen gas excited by high voltage. There was no explanation, no theory, and no idea as to why this purely empirical formula worked so perfectly.

Secondly, there was the odd case of the planet Mercury. The motion of every single planet, satellite, comet, planetoid, or dust cloud in the sky could be predicted with great accuracy, except the smallest planet in the solar system. There was not the vaguest clue as to why Mercury seemed to show up slightly behind where it was supposed to be in the sky.²⁸

These anomalies would soon be cleared up. Experimental physics had come about as far as it could go without a blast of theory to explain what was going on. The trip into theory would be deeper than could be anticipated. The increasingly freakish science incubated in the previous century, as physics became more abstract and less descriptive of common reality, would be nothing compared to the coming development of purely mathematical, logical science.

Atop that, Lord Rutherford’s neat, visibly logical model of the atom would turn out to be wrong, and its continued use in letterheads and science books everywhere is the ghost of an old fragment of the fantasy. The atom does not look like that at all.²⁹ Rutherford’s little orbiting electrons would not work for the same reason Hantaro Nagaoka’s Saturn model was invalid, but the problem was even deeper. In the next 15 years it would be revealed that the verb “to look” had no meaning when applied to an atom.