by Isaac Asimov
Seeking to explain the various properties of substances, the alchemists attached these properties to certain controlling elements that they added to the list. They identified mercury as the element that imparted metallic properties to substances, and sulfur as the element that imparted the property of flammability. One of the last and best of the alchemists, the sixteenth-century Swiss physician Theophrastus Bombastus von Hohenheim, better known as Paracelsus, added salt as the element that imparted resistance to heat.
The alchemists reasoned that one substance could be changed into another by merely adding and subtracting elements in the proper proportions. A metal such as lead, for instance, might be changed into gold by adding the right amount of mercury to the lead. The search for the precise technique of converting base metal to gold went on for centuries. In the process, the alchemists discovered substances vastly more important than gold—such as the mineral acids and phosphorus.
The mineral acids—nitric acid, hydrochloric acid, and, particularly, sulfuric acid (first prepared about 1300)—introduced a virtual revolution in alchemical experiments. These substances were much stronger acids than the strongest previously known (the acetic acid of vinegar); and with them, substances could be decomposed without the use of high temperatures and long waits. Even today, the mineral acids, particularly sulfuric acid, are of vital use in industry. It is said that the extent of the industrialization of a nation can be judged by its annual consumption of sulfuric acid.
Nevertheless, few alchemists allowed themselves to be diverted by these important side issues from what they considered the main quest. Unscrupulous members of the craft indulged in outright fakery, producing gold by sleight-of-hand, to win what we would call today “research grants” from rich patrons. This brought the profession into such disrepute that the very word alchemist had to be abandoned. By the seventeenth century, alchemist had become chemist, and alchemy had graduated to a science called chemistry.
In the bright birth of science, one of the first of the new chemists was Robert Boyle, the author of Boyle’s law of gases (see chapter 5) In his The Sceptical Chymist, published in 1661, Boyle first laid down the specific modern criterion of an element: a basic substance that can be combined with other elements to form compounds, and that, conversely, cannot be broken down to any simpler substance after it is isolated from a compound.
Boyle retained a medieval view about what the actual elements were, however. For instance, he believed that gold was not an element and could be formed in some way from other metals. So, in fact, did his contemporary Isaac Newton, who devoted a great deal of time to alchemy. (Indeed, Emperor Francis Joseph of Austria-Hungary subsidized experiments for making gold as late as 1867.)
In the century after Boyle, practical chemical work began to make clear which substances could be broken down into simpler substances and which could not. Henry Cavendish showed that hydrogen would combine with oxygen to form water, so water could not be an element. Later Lavoisier resolved the supposed element air into oxygen and nitrogen. It became plain that none of the Greek elements was an element by Boyle’s criterion.
As for the elements of the alchemists, mercury and sulfur did indeed turn out to be elements “according to Boyle.” But so did iron, tin, lead, copper, silver, gold, and such nonmetals as phosphorus, carbon, and arsenic. And Paracelsus’s “element” salt eventually was broken down into two simpler substances.
Of course, the definition of elements depended on the chemistry of the time. As long as a substance could not be broken down by the chemical techniques of the day, it could still be considered an element. For instance, Lavoisier’s list of thirty-three elements included such items as lime and magnesia. But fourteen years after Lavoisier’s death on the guillotine in the French Revolution, the English chemist Humphry Davy, using an electric current to split the substances, divided lime into oxygen and a new element he called calcium, and similarly split magnesia into oxygen and another new element he named magnesium.
On the other hand, Davy was able to show that a green gas that the Swedish chemist Carl Wilhelm Scheele had made from hydrochloric acid was not a compound of hydrochloric acid and oxygen, as had been thought, but a true element, and he named it chlorine (from the Greek word for “green”).
ATOMIC THEORY
At the beginning of the nineteenth century, there developed a radically new way of looking at elements which harked back to some of the Greeks, who had, after all, contributed what has turned out to be perhaps the most important single concept in the understanding of matter.
The Greeks argued about whether matter was continuous or discrete: that is, whether it could be divided and subdivided indefinitely into ever finer dust or would be found in the end to consist of indivisible particles. Leucippus of Miletus and his pupil Democritus of Abdera insisted, about 450 B.C., that the latter was the case. Democritus, in fact, gave the particles a name: he called them atoms (meaning “nondivisible”). He even suggested that different substances were composed of different atoms or combinations of atoms, and that one substance could be converted into another by rearranging the atoms. Considering that all this was only an intelligent guess, one is thunderstruck by the correctness of his intuition. Although the idea may seem obvious today, it was so far from obvious at the time that Plato and Aristotle rejected it out of hand.
It survived, however, in the teachings of Epicurus of Samos, who wrote about 300 B.C., and in the philosophic school (Epicureanism) to which he gave rise. An important Epicurean was the Roman philosopher Lucretius, who, about 60 B.C., embodied atomic notions in a long poem On the Nature of Things. One battered copy of Lucretius’s poem survived through the Middle Ages, and the poem was one of the earliest works to be printed once that technique had been invented.
The notion of atoms thus never entirely passed out of the consciousness of Western scholarship. Prominent among the atomists in the dawn of modern science were the Italian philosopher Giordano Bruno and the French philosopher Pierre Gassendi. Bruno had many unorthodox scientific views, such as a belief in an infinite universe with the stars distant suns about which planets revolved, and expressed himself boldly. He was burned as a heretic in 1600—the outstanding martyr to science of the Scientific Revolution. The Russians have named a crater on the other side of the moon in his honor.
Gassendi’s views impressed Boyle, whose own experiments showing that gases could easily be compressed and expanded seemed to show that these gases must be composed of widely spaced particles. Both Boyle and Newton were therefore among the convinced atomists of the seventeenth century.
In 1799, the French chemist Joseph Louis Proust showed that copper carbonate contained definite proportions by weight of copper, carbon, and oxygen, however it might be prepared. The proportions were in the ratio of small whole numbers: 5 to 4 to 1. He went on to show a similar situation for a number of other compounds.
That situation could best be explained by assuming that compounds are formed by the union of small numbers of bits of each element that could combine only as intact objects. The English chemist John Dalton pointed this out in 1803 and, in 1808, published a book in which all the new chemical information gathered in the past century and a half was shown to make sense if all matter were supposed to be composed of indivisible atoms. (Dalton kept the old Greek word as a tribute to the ancient thinkers.) It did not take long for this atomic theory to persuade most chemists.
According to Dalton, each element possesses a particular kind of atom, and any quantity of the element is made up of identical atoms of this kind. What distinguishes one element from another is the nature of its atoms. And the basic physical difference between atoms is in their weight. Thus sulfur atoms are heavier than oxygen atoms, which in turn are heavier than nitrogen atoms; they, in turn, heavier than carbon atoms; and these, in turn, heavier than hydrogen atoms.
The Italian chemist Amedeo Avogadro applied the atomic theory to gases in such a way as to show that it makes sense to suppose that equal vo
lumes of gas (of whatever nature) are made up of equal numbers of particles. This is Avogadro’s hypothesis. These particles were at first assumed to be atoms but eventually were shown to be composed, in most cases, of small groups of atoms called molecules. If a molecule contains atoms of different kinds (like the water molecule, which consists of an oxygen atom and two hydrogen atoms), it is a molecule of a chemical compound.
Naturally it became important to measure the relative weights of different atoms—to find the atomic weights of the elements, so to speak. The tiny atoms themselves were hopelessly beyond the reach of nineteenth-century weighing techniques. But by weighing the quantity of each element separated from a compound, and making deductions from an element’s chemical behavior, it was possible to work out the relative weights of the atoms. The first to go about this systematically was the Swedish chemist Fins Jacob Berzelius. In 1828, he published a list of atomic weights based on two standards—one giving the atomic weight of oxygen the arbitrary value of 100, the other taking the atomic weight of hydrogen as equal to 1.
Berzelius’s system did not catch on at once; but in 1860, at the first International Chemical Congress in Karlsruhe, Germany, the Italian chemist Stanislao Cannizzaro presented new methods for determining atomic weights, making use of Avogadro’s hypothesis, which had hitherto been neglected. Cannizzaro described his views so forcefully that the world of chemistry was won over.
The weight of oxygen rather than hydrogen was adopted as the standard at that time, because oxygen can more easily be brought into combination with various elements (and combination with other elements was the key step in the usual method of determining atomic weights). Oxygen’s atomic weight was arbitrarily taken by the Belgian chemist Jean Servais Stas, in 1850, as exactly 16, so that the atomic weight of hydrogen, the lightest known element, would be just about 1—1.0080, to be exact.
Ever since Cannizzaro’s time, chemists have sought to work out atomic weights with ever greater accuracy. This reached a climax, as far as purely chemical methods were concerned, in the work of the American chemist Theodore William Richards, who, in 1904 and thereafter, determined the atomic weights with an accuracy previously unapproached. For this he received the Nobel Prize in chemistry in 1914. On the basis of later discoveries about the physical constitution of atoms, Richards’s figures have since been corrected to still more refined values.
Throughout the nineteenth century, although much work was done on atoms and molecules, and scientists generally were convinced of their reality, there existed no direct evidence that they were anything more than convenient abstractions. Some prominent scientists, such as the German chemist Wilhelm Ostwald, refused to accept them in any other way. To him, they were useful but not “real.”
The reality of molecules was made clear by Brownian motion. This was first observed in 1827 by the Scottish botanist Robert Brown, who noted that pollen grains suspended in water jiggled erratically. At first it was thought that the jiggling was because of the life in the pollen grains, but equally small particles of completely inanimate dyes also showed the motion.
In 1863, it was first suggested that the movement was due to unequal bombardment of the particles by surrounding water molecules. For large objects, a slight inequality in the number of molecules striking from left and from right would not matter. For microscopic objects, bombarded by perhaps only a few hundred molecules per second, a few in excess—this side or that—can induce a perceptible jiggle. The random movement of the tiny particles is almost visible proof of the graininess of water, and of matter generally.
Einstein worked out a theoretical analysis of this view of Brownian motion and showed how one could work out the size of the water molecules from the extent of the little jiggling movements of the dye particles. In 1908, the French physicist Jean Perrin studied the manner in which particles settle downward through water under the influence of gravity. The settling is opposed by molecular collisions from below, so that a Brownian movement is opposing gravitational pull. Perrin used this finding to calculate the size of the water molecules by means of the equation Einstein had worked out, and even Ostwald had to give in. For his investigations Perrin received the Nobel Prize for physics in 1926.
So atoms have steadily been translated from semimystical abstractions into almost tangible objects. Indeed, today we can say that we have at last “seen” the atom. This is accomplished with the field ion microscope, invented in 1955 by Erwin W. Mueller of Pennsylvania State University. His device strips positively charged ions off an extremely fine needle tip and shoots them to a fluorescent screen in such a wayas to produce a 5 million-fold magnified image of the needle tip. This image actually makes the individual atoms composing the tip visible as bright little dots. The technique was improved to the point where images of single atoms could be obtained. The American physicist Albert Victor Crewe reported the detection of individual atoms of uranium and thorium by means of a scanning electron-microscope in 1970.
MENDELEEV’S PERIODIC TABLE
As the list of elements grew in the nineteenth century, chemists began to feel as if they were becoming entangled in a thickening jungle. Every element had different properties, and they could see no underlying order in the list. Since the essence of science is to try to find order in apparent disorder, scientists hunted for some sort of pattern in the properties of the elements.
In 1862, after Cannizzaro had established atomic weight as one of the important working tools of chemistry, a French geologist, Alexandre Emile Beguyer de Chancourtois, found that he could arrange the elements in the order of increasing atomic weight in a tabular form, so that elements with similar properties fell in the same vertical column. Two years later, a British chemist, John Alexander Reina Newlands, independently arrived at the same arrangement. But both scientists were ignored or ridiculed. Neither could get his suggestions properly published at the time. Many years later, after the importance of the periodic table had become universally recognized, their papers were published at last. Newlands even got a medal.
It was the Russian chemist Dmitri Ivanovich Mendeleev who got the credit for finally bringing order into the jungle of the elements. In 1869, he and the German chemist Julius Lothar Meyer proposed tables of the elements, making essentially the same point that de Chancourtois and Newlands had already made. But Mendeleev received the recognition because he had the courage and confidence to push the idea farther than the others.
In the first place, Mendeleev’s periodic table (so called because it showed the periodic recurrence of similar chemical properties) was more complicated than that of Newlands and nearer what we now believe to be correct (see table 6.1). Second, where the properties of an element placed it out of order according to its atomic weight, Mendeleev boldly switched the order, on the ground that the properties are more important than the atomic weight. He was eventually proved correct, as we shall see later in this chapter. For instance, tellurium, with an atomic weight of 127.61, should, on the weight basis, come after iodine, whose atomic weight is 126.91. But in the columnar table, putting tellurium ahead of iodine places it under selenium, which it closely resembles, and similarly puts iodine under its cousin bromine.
Table 6.1. The periodic table of the elements. The shaded areas of the table represent the two rare-earth series: the lanthanides and the actinides, named after their respective first members. The number in the lower right-hand corner of each box indicates the atomic weight of the element. An asterisk marks elements that are radioactive. Each element’s atomic number appears at top center of its box.
Finally, and most important, where Mendeleev could find no other way to make his arrangement work, he did not hesitate to leave holes in the table and to announce, with what seemed infinite gall, that elements must be discovered that belonged in those holes. He went farther. For three of the holes, he described the element that would fit each, utilizing as his guide the properties of the elements above and below the hole in the table. And here Mendeleev had a stroke o
f luck. Each of his three predicted elements was found in his own lifetime, so that he witnessed the triumph of his system. In 1875, the French chemist Lecoq de Boisbaudran discovered the first of these missing elements and named it gallium (after the Latin name for France). In 1879, the Swedish chemist Lars Fredrik Nilson found the second and named it scandium (after Scandinavia). And in 1886, the German chemist Clemens Alexander Winkler isolated the third and named it germanium (after Germany, of course). All three elements had almost precisely the properties predicted by Mendeleev.
ATOMIC NUMBERS
With the discovery of X rays by Roentgen, a new era opened in the history of the periodic table. In 1911, the British physicist Charles Glover Barkla discovered that when X rays are scattered by a metal, the scattered rays have a sharply defined penetrating power, depending on the metal; in other words, each element produces its own characteristic X rays. For this discovery Barkla was awarded the Nobel Prize in physics for 1917.
There was some question whether X rays were streams of tiny particles or consisted of wavelike radiations after the manner of light. One way of check ing was to see whether X rays could be diffracted (that is, forced to change direction) by a diffraction grating consisting of a series of fine scratches However, for proper diffraction, the distance between the scratches must be roughly equal to the size of the waves in the radiation. The most finely spaced scratches that could be prepared sufficed for ordinary light, but the penetrating power of X rays made it likely that, if X rays were wavelike, the waves would have to be much smaller than those of light. Therefore, no ordinary diffraction gratings would suffice to diffract X rays.
