by Bill Bryson
Mendeleyev’s Periodic Table, still as useful today as when he invented it in 1869. Mendeleyev was said to have modelled the table on the card game solitaire. (credit 7.8)
As is often the way in science, the principle had actually been anticipated three years previously by an amateur chemist in England named John Newlands. He suggested that when elements were arranged by weight they appeared to repeat certain properties—in a sense to harmonize—at every eighth place along the scale. Slightly unwisely, for this was an idea whose time had not quite yet come, Newlands called it the Law of Octaves and likened the arrangement to the octaves on a piano keyboard. Perhaps there was something in Newlands’ manner of presentation, but the idea was considered fundamentally preposterous and widely mocked. At gatherings, droller members of the audience would sometimes ask him if he could get his elements to play them a little tune. Discouraged, Newlands gave up pushing the idea and soon dropped out of sight altogether.
Mendeleyev used a slightly different approach, placing his elements into groups of seven, but employed fundamentally the same premise. Suddenly the idea seemed brilliant and wondrously perceptive. Because the properties repeated themselves periodically, the invention became known as the Periodic Table.
Mendeleyev was said to have been inspired by the card game known as solitaire in North America and patience elsewhere, wherein cards are arranged by suit horizontally and by number vertically. Using a broadly similar concept, he arranged the elements in horizontal rows called periods and vertical columns called groups. This instantly showed one set of relationships when read up and down and another when read from side to side. Specifically, the vertical columns put together chemicals that have similar properties. Thus copper sits on top of silver and silver sits on top of gold because of their chemical affinities as metals, while helium, neon and argon are in a column made up of gases. (The actual, formal determinant in the ordering is something called their electron valences, and if you want to understand them you will have to enrol in evening classes.) The horizontal rows, meanwhile, arrange the chemicals in ascending order by the number of protons in their nuclei—what is known as their atomic number.
The structure of atoms and the significance of protons will come in a following chapter; for the moment, all that is necessary is to appreciate the organizing principle: hydrogen has just one proton and so it has an atomic number of 1 and comes first on the chart; uranium has 92 protons and so it comes near the end and has an atomic number of 92. In this sense, as Philip Ball has pointed out, chemistry really is just a matter of counting. (Atomic number, incidentally, is not to be confused with atomic weight, which is the number of protons plus the number of neutrons in a given element.)
The Periodic Table of Chemical Elements, re-imagined as a landscape, with the peaks and slopes representing properties of the elements, such as atomic weights and numbers. This work is described as “a view from the northeast across the transition metals towards the peaks of Osmium and Iridium.” (credit 7.8a)
There was still a great deal that wasn’t known or understood. Hydrogen is the most common element in the universe, and yet no-one would guess as much for another thirty years. Helium, the second most abundant element, had only been found the year before—its existence hadn’t even been suspected before that—and then not on the Earth, but in the Sun, where it was found with a spectroscope during a solar eclipse, which is why it honours the Greek sun god Helios. It wouldn’t be isolated until 1895. Even so, thanks to Mendeleyev’s invention, chemistry was now on a firm footing.
For most of us, the Periodic Table is a thing of beauty in the abstract, but for chemists it established an immediate orderliness and clarity that can hardly be overstated. “Without a doubt, the Periodic Table of the Chemical Elements is the most elegant organizational chart ever devised,” wrote Robert E. Krebs in The History and Use of Our Earth’s Chemical Elements—and you can find similar sentiments in virtually every history of chemistry in print.
Today we have “120 or so” known elements—92 naturally occurring ones plus a couple of dozen that have been created in labs. The actual number is slightly contentious because the heavy, synthesized elements exist for only millionths of seconds and chemists sometimes argue over whether they have really been detected or not. In Mendeleyev’s day just sixty-three elements were known, but part of his cleverness was to realize that the elements as then known didn’t make a complete picture, that many pieces were missing. His table predicted, with pleasing accuracy, where new elements would slot in when they were found.
No-one knows, incidentally, how high the number of elements might go, though anything beyond 168 as an atomic weight is considered “purely speculative”; but what is certain is that anything that is found will fit neatly into Mendeleyev’s great scheme.
The nineteenth century held one last important surprise for chemists. It began in 1896 when Henri Becquerel in Paris carelessly left a packet of uranium salts on a wrapped photographic plate in a drawer. When he took the plate out some time later, he was surprised to discover that the salts had burned an impression in it, just as if the plate had been exposed to light. The salts were emitting rays of some sort.
Considering the importance of what he had found, Becquerel did a very strange thing: he turned the matter over to a graduate student for investigation. Fortunately the student was a recent émigré from Poland named Marie Curie. Working with her new husband, Pierre, Curie found that certain kinds of rocks poured out constant and extraordinary amounts of energy, yet without diminishing in size or changing in any detectable way. What she and her husband couldn’t know—what no-one could know until Einstein explained things the following decade—was that the rocks were converting mass into energy in an exceedingly efficient way. Marie Curie dubbed the effect “radioactivity.” In the process of their work, the Curies also found two new elements—polonium, which they named after her native country, and radium. In 1903 the Curies and Becquerel were jointly awarded the Nobel Prize in physics. (Marie Curie would win a second prize, in chemistry, in 1911; the only person to win in both chemistry and physics.)
At McGill University in Montreal the young New Zealand-born Ernest Rutherford became interested in the new radioactive materials. With a colleague named Frederick Soddy he discovered that immense reserves of energy were bound up in these small amounts of matter, and that the radioactive decay of these reserves could account for most of the Earth’s warmth. They also discovered that radioactive elements decayed into other elements—that one day you had an atom of uranium, say, and the next you had an atom of lead. This was truly extraordinary. It was alchemy pure and simple; no-one had ever imagined that such a thing could happen naturally and spontaneously.
The smudges left accidentally on a photographic plate in Henri Becquerel’s desk drawer that led to the discovery of radioactivity. (credit 7.9)
Ever the pragmatist, Rutherford was the first to see that there could be a valuable practical application in this. He noticed that in any sample of radioactive material, it always took the same amount of time for half the sample to decay—the celebrated half-life3—and that this steady, reliable rate of decay could be used as a kind of clock. By calculating backwards from how much radiation a material had now and how swiftly it was decaying, you could work out its age. He tested a piece of pitchblende, the principal ore of uranium, and found it to be 700 million years old—very much older than the age most people were prepared to grant the Earth.
In the spring of 1904, Rutherford travelled to London to give a lecture at the Royal Institution—the august organization founded by Count von Rumford only 105 years before, though that powdery and periwigged age now seemed a distant aeon compared with the roll-your-sleeves-up robustness of the late Victorians. Rutherford was there to talk about his new disintegration theory of radioactivity, as part of which he brought out his piece of pitchblende. Tactfully— for the ageing Kelvin was present, if not always fully awake—Rutherford noted that Kelvin himself had suggested t
hat the discovery of some other source of heat would throw his calculations out. Rutherford had found that other source. Thanks to radioactivity the Earth could be—and self-evidently was—much older than the 24 million years Kelvin’s final calculations allowed.
Kelvin beamed at Rutherford’s respectful presentation, but was in fact unmoved. He never accepted the revised figures and to his dying day believed his work on the age of the Earth his most astute and important contribution to science—far greater than his work on thermodynamics.
As with most scientific revolutions, Rutherford’s new findings were not universally welcomed. John Joly of Dublin strenuously insisted well into the 1930s that the Earth was no more than 89 million years old, and was stopped only then by his own death. Others began to worry that Rutherford had now given them too much time. But even with radiometric dating, as decay measurements became known, it would be decades before we got within a billion years or so of the Earth’s actual age. Science was on the right track, but still way out.
Pierre and Marie Curie photographed in Paris. The Curies were jointly awarded the Nobel Prize in physics with Henri Becquerel in 1903 for their work on radioactivity. (credit 7.10)
Kelvin died in 1907. That year also saw the death of Dmitri Mendeleyev. Like Kelvin, his productive work was far behind him, but his declining years were notably less serene. As he aged, Mendeleyev became increasingly eccentric—he refused to acknowledge the existence of radiation or the electron or anything else much that was new—and difficult. His final decades were spent mostly storming out of labs and lecture halls all across Europe. In 1955, element 101 was named mendelevium in his honour. “Appropriately,” notes Paul Strathern, “it is an unstable element.”
Radiation, of course, went on and on, literally and in ways nobody expected. In the early 1900s Pierre Curie began to experience clear signs of radiation sickness—notably dull aches in his bones and chronic feelings of malaise—which doubtless would have progressed unpleasantly. We shall never know for certain because in 1906 he was fatally run over by a carriage while crossing a Paris street.
Marie Curie spent the rest of her life working with distinction in the field, helping to found the celebrated Radium Institute of the University of Paris in 1914. Despite her two Nobel Prizes, she was never elected to the Academy of Sciences, in large part because after the death of Pierre she conducted an affair with a married physicist sufficiently indiscreet to scandalize even the French—or at least the old men who ran the academy, which is perhaps another matter.
For a long time it was assumed that anything so miraculously energetic as radioactivity must be beneficial. For years, manufacturers of toothpaste and laxatives put radioactive thorium in their products, and at least until the late 1920s the Glen Springs Hotel in the Finger Lakes region of New York (and doubtless others as well) featured with pride the therapeutic effects of its “Radio-active mineral springs.” It wasn’t banned in consumer products until 1938. By this time it was much too late for Mme Curie, who died of leukaemia in 1934. Radiation, in fact, is so pernicious and long-lasting that even now her papers from the 1890s—even her cookbooks—are too dangerous to handle. Her lab books are kept in lead-lined boxes and those who wish to see them must don protective clothing.
Never has the promise of glowing skin been more dangerously apt than in the early years of the twentieth century when radium was commonly used as a featured ingredient in beauty products. (credit 7.11)
Thanks to the devoted and unwittingly high-risk work of the first atomic scientists, by the early years of the twentieth century it was becoming clear that the Earth was unquestionably venerable, though another half-century of science would have to be done before anyone could confidently say quite how venerable. Science, meanwhile, was about to get a new age of its own— the atomic one.
Marie Curie, towards the end of her life, with Albert Einstein: their brilliant discoveries started scientists on the road that led to the atomic bomb. (credit 7.12)
1 The confusion over the aluminum/aluminium spelling arose because of some uncharacteristic indecisiveness on Davy’s part. When he first isolated the element in 1808, he called it alumium. For some reason he thought better of that and changed it to aluminum four years later. Americans dutifully adopted the new term, but many British users disliked aluminum, pointing out that it disrupted the -ium pattern established by sodium, calcium and strontium, so they added a vowel and syllable. Among his other achievements, Davy also invented the miner’s safety lamp.
2 The principle led to the much later adoption of Avogadro’s Number, a basic unit of measure in chemistry, which was named for Avogadro long after his death. It is the number of molecules found in 2.016 grams of hydrogen gas (or an equal volume of any other gas). Its value is placed at 6.0221367 × 1023, which is an enormously large number. Chemistry students have long amused themselves by computing just how large a number it is, so I can report that it is equivalent to the number of popcorn kernels needed to cover the United States to a depth of nine miles, or cupfuls of water in the Pacific Ocean, or soft-drink cans that would, evenly stacked, cover the Earth to a depth of two hundred miles. An equivalent number of American pennies would be enough to make every person on Earth a dollar trillionaire. It is a big number.
3 If you have ever wondered how the atoms determine which 50 per cent will die and which 50 per cent will survive for the next session, the answer is that the half-life is really just a statistical convenience—a kind of actuarial table for elemental things. Imagine you had a sample of material with a half-life of 30 seconds. It isn’t that every atom in the sample will exist for exactly 30 seconds or 60 seconds or 90 seconds or some other tidily ordained period. Each atom will in fact survive for an entirely random length of time that has nothing to do with multiples of 30; it might last until two seconds from now or it might oscillate away for years or decades or centuries to come. No-one can say. But what we can say is that for the sample as a whole the rate of disappearance will be such that half the atoms will disappear every 30 seconds. It’s an average rate, in other words, and you can apply it to any large sampling. Someone once worked out, for instance, that American dimes have a half-life of about thirty years.
A technician makes adjustments to a recently installed atom smasher at the University of Notre Dame in Indiana in 1941. The 8 million volts generated by the machine helped scientists experiment with atomic structure as well as with the production of radioactive metals. (credit p3.1)
EINSTEIN’S UNIVERSE
As the nineteenth century drew to a close, scientists could reflect with satisfaction that they had pinned down most of the mysteries of the physical world: electricity, magnetism, gases, optics, acoustics, kinetics and statistical mechanics, to name just a few, had all fallen into order before them. They had discovered the X-ray, the cathode ray, the electron and radioactivity, invented the ohm, the watt, the kelvin, the joule, the amp and the little erg.
If a thing could be oscillated, accelerated, perturbed, distilled, combined, weighed or made gaseous they had done it, and in the process produced a body of universal laws so weighty and majestic that we still tend to write them out in capitals: the Electromagnetic Field Theory of Light, Richter’s Law of Reciprocal Proportions, Charles’s Law of Gases, the Law of Combining Volumes, the Zeroth Law, the Valence Concept, the Laws of Mass Actions, and others beyond counting. The whole world clanged and chuffed with the machinery and instruments that their ingenuity had produced. Many wise people believed that there was nothing much left for science to do.
In 1875, when a young German in Kiel named Max Planck was deciding whether to devote his life to mathematics or to physics, he was urged most heartily not to choose physics because the breakthroughs had all been made there. The coming century, he was assured, would be one of consolidation and refinement, not revolution. Planck didn’t listen. He studied theoretical physics and threw himself body and soul into work on entropy, a process at the heart of thermodynamics, which seemed to hold much pr
omise for an ambitious young man.1 In 1891 he produced his results and learned to his dismay that the important work on entropy had in fact been done already, in this instance by a retiring scholar at Yale University named J. Willard Gibbs.
Gibbs is perhaps the most brilliant person most people have never heard of. Modest to the point of near-invisibility, he passed virtually the whole of his life, apart from three years spent studying in Europe, within a three-block area bounded by his house and the Yale campus in New Haven, Connecticut. For his first ten years at Yale he didn’t even bother to draw a salary. (He had independent means.) From 1871, when he joined the university as a professor, to his death in 1903, his courses attracted an average of slightly over one student a semester. His written work was difficult to follow and employed a private form of notation that many found incomprehensible. But buried among his arcane formulations were insights of the loftiest brilliance.
In 1875–8, Gibbs produced a series of papers, collectively titled On the Equilibrium of Heterogeneous Substances, which dazzlingly elucidated the thermodynamic principles of, well, nearly everything—“gases, mixtures, surfaces, solids, phase changes…chemical reactions, electrochemical cells, sedimentation, and osmosis,” to quote William H. Cropper. In essence, what Gibbs did was show that thermodynamics didn’t apply simply to heat and energy at the sort of large and noisy scale of the steam engine, but was also present and influential at the atomic level of chemical reactions. Gibbs’s Equilibrium has been called “the Principia of thermodynamics,” but for reasons that defy speculation Gibbs chose to publish these landmark observations in the Transactions of the Connecticut Academy of Arts and Sciences, a journal that managed to be obscure even in Connecticut, which is why Planck did not hear of him until too late.
