When Dalton published his atomic theory in 1808, the record of recognized elements was not large. By 1820 the number had reached fifty; some of these had been added by Berzelius. Not knowing how many additional elements were possible and not knowing the part of the world in which some new one might be found, the chemical pioneers were always on the alert. Each hoped to gain eternal fame by being the first to discover and name some element never known before. In most cases the discoveries followed the development of new tools for discovery.
In 1801, Humphry Davy, newly appointed professor of chemistry at the Royal Institution of London, set up a battery of electric cells to do some chemical experimenting. It was already established that this electrical device, invented by Alessandro Volta of Italy, was able to separate water into hydrogen and oxygen gases. Davy wanted to know what effect the electricity would have on other substances. He tried dilute sulfuric acid first, probably expecting to get sulfur. He got hydrogen and oxygen gases—and nothing else. The water of the dilute acid had been destroyed; the acid itself remained. Davy then tried a solution of potash lye. He got hydrogen and oxygen gases from the water of the solution, and nothing else.
By 1806 Davy had completed many experiments on the chemical effect of the electric current, but, over and over, he would get hydrogen and oxygen gases from any water that might be present. Yet, if water was omitted, the electricity from the battery of cells would not pass. He apparently needed his substances in liquid form but he could not use water to make that liquid. Then the idea came to him of melting the substance and, while it was melted, of trying to pass the current. He selected potash lye, which he knew would melt when heated, and touched the battery terminals to the outside of the partially melted solid. When the current was turned on there was a sputtering, then a flash of violet-colored flame. Whatever had formed had almost immediately burned up.
Taking all possible care, for there seemed to be some danger, he repeated the experiment on a larger scale. Here are some comments on the results as he reported them:
As soon as the current was turned on the solid melted both at the top where there was a conducting wire and at the bottom that rested on a disc of platinum connected to the negative side of the battery. Around the wire at the top there was a vigorous bubbling with gas being given off. At the lower surface no gas formed, but small globules, having a high metallic luster and looking much like mercury, appeared. Some of the globules burned with an explosion and bright flame as soon as they formed. Others remained and were merely tarnished and finally covered by a white film which formed on their surfaces.
At ordinary temperatures, if kept away from the air, the metal product is a soft and malleable solid, which has the luster of polished silver. In contact with a few drops of water it decomposes the water with great violence. An instantaneous explosion is produced and a brilliant flame; a solution of pure potash lye is formed.1
Davy, quite correctly, considered the new metal an element and named it "base of potash" or potassium. People crowded to a public lecture where he showed samples of this metal and performed experiments to demonstrate its strange properties. But the strain and excitement of weeks of tireless effort were too much for his frail health. He went to the hospital for several weeks—while all England prayed for his recovery. He was scarcely back on his feet again when he turned on the electric current in a new series of experiments. He was searching for another metal that he believed would be unearthed in soda lye. He found it, and named it "base of soda" or sodium. It was much like potassium but safer to handle. It burned with a strong yellow flame but did not catch on fire as readily as the other metal. Almost without resting, Davy, with the help of an assistant, went on searching for more new elements. Eventually he found calcium in lime, magnesium in white magnesia, barium in heavy baryta, and strontium in the Scottish earthy mineral strontia. The experimentation was difficult and the quantity of each metal produced was small. Then, while he was waiting for a few days for a larger and more powerful battery to be set up, a letter arrived from Sweden. It was Berzelius; he too had just discovered calcium and barium. Each wrote immediately congratulating the other. A little later, Davy recovered the element lithium from a Swedish mineral that Berzelius had found, purified, and sent to Davy in London. The discovery of other elements soon followed. Both Davy and Gay-Lussac produced boron from boric acid. Davy pointed out that chlorine was an element, not a compound of oxygen as Lavoisier had thought. He studied iodine, called it an element, and named it, though he did not discover it in the first place. He predicted that a new element would be found in the mineral fluorspar, gave it the name of fluorine, tried to prepare it, made some practical suggestions as to how it might be handled, but then gave up the task, for he found he was being poisoned by fluorine compounds.
All these discoveries gave Davy an excellent reputation as a chemist. They also gave him a great deal of publicity, even among those who knew little about chemistry. But a previous chemical "adventure" gave him almost too much publicity, and landed him in a hospital. It happened when he was twenty-two. Young as he was then, he had previously established a local reputation as an experimenter, especially in the preparation and handling of gases. So he was quite logically selected as "superintendent" of a special laboratory built and equipped by Dr. Thomas Beddoes, an English physician. The doctor, in rereading Priestley's account of the discovery of oxygen, had been struck by the experimenters' comment about the possible future use of the gas in lung and throat troubles. He also read over Priestley's account of his experiment on the gas that he called "diminished nitrous air" in which a mouse went soundly to sleep; it awoke drowsily when taken out of the gas. Dr. Beddoes wondered whether this gas would be used in cases of insomnia. He found in other reading that a breath of hydrogen gas affected the pitch of a person's speech. He wondered if that might not have a practical use in throat disorders. So, having the money to do it, he started on a planned schedule for trying the effect of these three gases upon humans. Davy volunteered to be a human patient.
The results of the breathing of oxygen and hydrogen revealed nothing not known before. But the use of diminished nitrous air, or nitrous oxide, as it is known today, was full of surprises. Up to this time only a mouse had taken a deep whiff of this gas; being a mouse, he could not put his reactions into words. Here are Robert Southey's remarks after he had filled his lungs with the gas, as reported by Davy:
My first definite sensation was a dizziness, such as to induce a fear of falling. This was momentary. When I took the mouthpiece furnishing the gas from my mouth, I immediately laughed. The laugh was involuntary, but highly pleasurable, accompanied by a thrill all through me. And a tingling in my toes and fingers, a sensation perfectly new and delightful.2
It was no wonder that people came from miles around to view the public demonstration of the new gas which they called "laughing gas". (It is used today as a mild anesthetic in certain surgical cases.) It was unfortunate that Dr. Beddoes and Davy did not stop their gas experiments at this point. Instead, they went to other gases. The inhaling of the gas that Priestley had called nitrous air was a grave error. Davy escaped with his life, but spent months in bed. Seven years later he suffered the severe illness that followed the strain and excitement of the discovery of potassium.
His scientific career was wonderfully brilliant, but short; he died at fifty-one in 1829. He had been knighted in 1811; after that, he was Sir Humphry Davy. But oddly enough the title did not honor him as a discoverer of elements. Scientists would have recognized the importance of that, of course. He was honored for having invented the Davy safety lamp, which had saved countless lives in the coal mines of England—he had granted the rights in the invention to the people of England—and because he had written in plain, terse English a scientific but practical treatise on English farming—a treatise that was to be used for half a century. Again, he had offered it, without compensation, to the people of England.
At the middle of the nineteenth century, that is, around the year 1850, there was a long list of known elements. Each of these responded to at least one chemical test that was distinctive. But no chemist could be certain that some undiscovered element might not also be able to respond to the tests of some known element, and so have missed detection. (A case like that did occur. In 1923 compounds of a new element hafnium were found, in small amounts, mixed with similar compounds of zirconium. The hafnium had passed all the special tests then used for zirconium, so its presence had never been suspected.) By 1859 the situation began to clear up. A way had been devised by which each element would have a sort of fingerprint of its own. After that, confused or erroneous identifications might always be checked. The idea had started in a rather indirect way. The compounds of a number of the metallic elements give a pronounced color to the flame of the alcohol burner. That certainty is now almost universally known. Anyone, for example, who lives by the sea can tell you that a nail of corroded copper in a slab of driftwood will give a green color to the fireplace flame as the wood slowly burns. The color given by strontium compounds is a brilliant red, useful for red flares as in those used by truckers and the highway patrol. Lithium compounds form a deep red, calcium an orange-red. Barium produces a light green, potassium a faint white-lavender, sodium an intense yellow. It might seem easy, then, to use the special flame color to detect the presence or absence of the compounds of a particular metal. But the situation becomes confused when two or more of the flame-giving elements are present in the mixture, since the shades of color will be mixed so completely that none will stand out clearly. The triangular glass prism used in science, can, however, do what the eye cannot. It can take a ray of colorless sunlight and, spreading it out, show that this ray is a mixture of shades of light that stretch across the spectrum from red to violet. In the chemistry laboratory that same glass prism can take a colored flame and sort it into the shades of red, orange, yellow, green, and lavender that are in it, putting each of these in its proper place on the spectrum screen. The chemist, looking at such a spread out array of colors, can then say with certainty, for example, "This substance being tested is a mixture. It contains compounds of both strontium and calcium; it does not contain those of lithium." Robert Bunsen, the chemist, and Gustav Kirchhoff, the physicist, turned the idea of the triangular glass prism into an accurate scientific instrument called a spectroscope. The device carried a scale of numbers. The yellow of the sodium flame gave a bright yellow line at a particular number location on this scale, and the same idea was true for other flame colors and other elements. So any colored line could be recorded by a number rather than by the description of the color shade, a detail that was to be highly important since color shades are difficult to describe. Two facts were determined almost immediately. Most of the elements giving colored flames had more than one bright line in their spectrums. No two elements had even a single bright line in common. One element could not, then, cover up or distort the lines of another element. Each had its own recognition lines.
Bunsen employed the spectroscope in the discovery of two metallic elements similar to sodium and potassium in their general nature, whose compounds exist in such small amounts in the earth's surface that they could scarcely have been discovered by the analytical chemist. This statement is from Bunsen's own account of the findings of the first of these elements:
If one brings into the flame of the spectroscope a drop of mother liquor from the Durkheim mineral water from which calcium, strontium, and magnesium have already been removed as compounds, one sees the bright lines of sodium, potassium, and lithium. And, in addition, two remarkable blue lines very close together. Now there had been no element known which would give two such rays. One might therefore conclude the certain existence of an unknown element belonging to the same chemical group as sodium, potassium, and lithium. We have proposed for this new metallic element the name cesium from the word caesius which the ancients used to designate the blue of the heavens.3
The second element, even more rare in its distribution in nature, had been found in the mineral lepidolite obtained from Saxony. Here is Bunsen's report on the final step without the elaborate details of analysis:
Two remarkable red lines appeared at the extreme red end of the solar spectrum. The color led us to give this new element of the sodium-potassium-lithium group the name rubidium, from rubidus, the color of the ruby.4
Only a month after the discovery of rubidium a report from William Crookes of England stated that he had discovered a new metallic element that gave a beautiful green flame. He named the element thallium, from the Greek for "green twig". A year later, indium, with its brilliant indigo-blue flame color, was added to the list of elements by Ferdinand Reich and Theodor Richter of Germany. Three more, all checked by the spectroscope, were added by Francois Lecoq de Boisbaudran of France. They were named gallium, samarium, and dysprosium. For these the discoverer had needed a higher source of heat for the flame test than the ordinary burner would give and had employed the heat of the electric spark. Since that time seven other elements have been identified by the electric-spark spectrum.
To the laboratory worker of today, Robert Bunsen's name is associated with laboratory burners. This is natural, since it was he who invented a burner that would successfully use a gas as a fuel. Previous burners had been for wood, charcoal, coal, or whale oil. In addition to the spectroscope he had invented also the Bunsen cell for battery current, the Bunsen filter pump, and the Bunsen valve for chemical apparatus.
To his students at the University of Heidelberg, Bunsen was a great teacher. To his advanced students he was a stimulating director of research. However, he had one peculiarity. No one in his laboratory was allowed to do any research on the compounds of carbon. The work was too dangerous. Had he not been nearly killed trying to find the chemical make-up of a compound of carbon that no one else had wanted to work with? It was a liquid with a fearful odor, its vapor was known to be very poisonous, and it readily formed compounds that might catch on fire without being ignited.
There were no hoods or ventilating fans in those early laboratories. For this dangerous experiment, Bunsen had developed an improvised gas mask with a large glass tube running out of a window through which to breathe. He seemed to be getting along moderately well when a drop of the material he had just produced fell on the heated portion of a glass flask. He reported that "the apparatus was demolished by an explosion, and a flame several feet high arose, covering the surrounding objects with a layer of black." He survived that accident unharmed. But a little later a second explosion cost him the sight of one eye and caused weeks of illness from the breathing of poisonous fumes. That was the last of his experiments on the make-up of compounds of carbon. "There are," he declared, "plenty of sensible, feasible problems in chemistry to undertake." Bunsen took his own advice—and lived to be eighty-eight.
Back in the 1700s, as we have seen, Jan van Helmont called the gas he was trying to collect "a wild spirit" because it had burst a flask and thrown pieces of glass in his face. That particular spirit was soon tamed. But the wildest of all the wild spirits of the chemical world would not be controlled or subdued until Henri Moissan of France did it in 1886.
This was the gas fluorine that Davy had thought would be found in fluorspar and had tried to obtain by electricity. It was this gas that he failed to produce and contain, being poisoned in the attempt. Two brothers, George and Thomas Knox of Ireland, tried to go on with Davy's experiments; one spent three years in bed as a result. P. Louyet of Belgium tried, knowing the danger, and died. The gas also killed Jerome Nickels of Nancy, France. George Gore of England was much more careful, and more fortunate. He escaped serious injury. His most difficult task was to find something in which to keep this fearfully wild spirit after he had captured it. Everything he used was soon riddled with holes.
Something should be said about the violent chemical behavior of this gas. It energetically attacks water, forming ozone and hydrogen fluoride gases. The hydrogen fluoride, so formed, can kill a person who inhales it. (The air has enough moisture in it to let this killer loose if fluorine leaks into the air.) In addition, hydrogen fluoride cannot be kept as a solution in a glass vessel, it destroys the glass. To keep fluorine, the gas, it was necessary, then, to avoid the presence of moisture. But the gas attacked metals, so metallic vessels could not hold it for more than a short time.
Gore, a skilled electrochemist, had no particular difficulty in getting the fluorine, for compounds of this element may readily be obtained. True, he could not pass the current he was using through a water solution of such a compound; the fluorine as it formed would attack the water. So he selected a fluorine compound that could be melted by heat, and passed electricity through the liquid. He got the gas. He used a small vial made of platinum to hold the collected gas, since platinum was not affected, in general, by chemicals. The expensive metal was completely destroyed. He tried palladium, another resistant metal. It was destroyed. He tried again, using gold. The metal was almost immediately changed to a powdery solid. He made a small jar out of carbon from an electric dry cell. In time the jar went to pieces, but before it was completely destroyed he collected some of the gas. He transferred it to a stoppered vial made of the mineral fluorspar, which the gas would not attack because it already was a fluorine compound. When a little of this gas was mixed with dry hydrogen gas the mixture exploded violently; no match was needed. Dangerous hydrogen fluoride was the product. Gore never made any more fluorine; he was satisfied.
In 1886 Henri Moissan carried on the production of fluorine along the general lines used by Gore and obtained it in sufficient quantity to make numerous experiments upon it. He was never able to prevent the gas from destroying his equipment, but he was able to cut down the rate of destruction. Here is a portion of his published report on the experiment:
I obtained the fluorine from a fluorine compound that had been added to a mineral having a low melting point and in which the fluorine compound dissolved readily. The use of electricity produced the fluorine at the positive terminal. Difficulty was experienced in getting any material for that terminal that would resist the chemical action of the gas. After some failures, and four interruptions of work caused by severe poisoning, the following arrangement of apparatus proved fairly satisfactory. Two electrodes were made from an alloy of platinum and iridium. These were sealed into a platinum U-tube closed with caps made from the mineral fluorspar, the caps being covered with a layer of gum-lac. The U-tube was chilled to 10 degrees below zero Fahrenheit to reduce the rate of the action of the fluorine on the platinum. The first test made with the gas was to bring it in contact with the element silicon. There was an immediate burst of flame, a gaseous product being formed.5
The announcement of his successful, though not inexpensive, production of fluorine made the headlines of the newspapers of the world. But Moissan was to be known for other difficult achievements also. He developed an electric furnace which produced a temperature so high that such uncommon but interesting metals as uranium, tungsten, vanadium, chromium, manganese, titanium, molybdenum, tantalum, and thorium were obtained from their compounds. The metals did not have complete purity but were produced in purer form than before. But the achievement that caused the greatest excitement at the time of announcement was Moissan's production of genuine diamonds out of carbon. Trying to duplicate nature's feat, he had used a combination of great pressure and high temperature. The plan did not have an upsetting effect upon the world's diamond market, however, since the largest gem he was able to get was one-thirty-second of an inch in length.
His interests and accomplishments might indicate that Moissan had been trained as an engineer or a metallurgist. Such was not the case. He had been an apothecary apprentice in Paris, and had gone on to become a registered pharmacist. Being able and ambitious, he had taken the time to study chemistry at the Ecole Polytechnique in Paris. Gaining fame in both pharmacy and chemistry, he had become professor of chemistry at the Ecole Superieure de Pharmacie in Paris. He had started his work on fluorine during vacation time. Moissan was a big, energetic man of great personality, who delighted in giving public addresses on scientific topics. His son Louis was his assistant in putting on the demonstrations that were an important feature of those addresses. (This son was killed on a battlefield of the First World War.) Moissan's father-in-law financed the work of the electric furnace and diamond production, and the entire family may have scarcely realized how much the work with fluorine was to shorten the experimenter's life. He died in 1907 at fifty-five. As an additional note we should state that not all compounds of fluorine are poisonous or dangerous. Electric refrigerators use a valuable nontoxic fluorine-containing fluid in the coils. The enamel of teeth is benefited by fluoridation.
Previous to 1869 a major portion of the elements in the earth's crust, in the oceans, and in the air had been found. But the chemist had not yet developed any plan by which he could know how many more there would be. The weights of the various kinds of atoms in comparison with the weight of hydrogen, the lightest of them all, had been worked out. But errors had been made in some of these weight values and there seemed no way to detect such errors. Then, almost quickly, the situation changed. The elements were found to fit into a natural system, with missing elements showing up because their locations remained vacant. Large errors in atomic weights could readily be detected. A spirit of romantic adventure invaded the study of chemistry.
In that year of 1869 a very thoughtful Russian chemist reported to the scientific world that the listed atomic weight for beryllium was off about two points, that the one given for indium was completely incorrect, and that the one for gold was heavier than that for platinum. He also stated quite positively that there were two elements with atomic weights lying between those of zinc and arsenic that were yet to be discovered. He acknowledged that he had not made a single experiment to prove any of these things, and the chemists of the world would have paid little attention to his remarks had not future work shown that these predictions were correct. As for the missing elements, he had gone much farther than merely predicting them. He had told what each would look like, how heavy it would be, what kinds of compounds it would form—and their formulas and colors. The metals were soon found and called gallium and germanium by the discoverers. The results were uncanny; every property he had announced they would possess was correct.
The name of that Russian chemist was Dmitri Mendelyeev. (His last name is also written in several other ways, such as Mendeleeff and Mendeljeff.) You might suppose that he had been playing a hunch in making his predictions. That was not the case. He had hit on a great truth—that the properties of the elements are periodic in their nature. There is, of course, nothing particularly unusual about movements or actions that come back, again and again, in periods or cycles. The truth that Mendelyeev had discovered was that the elements have their cyclic periods also. He had arranged the list of the elements in the order of their atomic weights, and had found that their properties kept repeating. At one place there would be an element that was a gas; spaced out beyond that would be an element that was also a gas; after that, spaced out in the same general way, would be a gas or a liquid that could be readily changed to a gas. Again, if there was a soft light metal at one point in the list there would be found, spaced out on either side, other soft light metals. These others would not be exactly as soft or as light as the first one, for the properties were never completely identical.
There were some chemists in later times who would liken the list of elements to the keys on a great piano. The comparison is interesting. As the hand moves along the keyboard each key when struck gives a distinctive tone, unlike that of any other key. The tones become lower in pitch as the weights of the wires that produce the tones grow heavier. But in these tones there is a cycle of sounds; the ear can hear it and the arrangement of keys will show it. Strike the key for middle C. Spread out across the keyboard, up and down, are other C's, for the C is a note in an octave and the octave is a repeating period or cycle. Select the F of the keyboard and there will be six other F's that may vibrate in tune with the one struck. The tone-giving properties of the piano keys are related in a simple periodic way to the position of the keys on the piano's keyboard. In describing the list of elements, Mendelyeev used almost the same wording as he wrote, "The properties of the elements are related in a simple periodic way to the position of the elements in the atomic-weight series." He called this the periodic law.
The piano is simpler in its periodicity than nature's arrangement of the elements. On that instrument every octave, up and down the keyboard, has the same number of notes. With the elements the first period has only two; then comes two periods of eight each; next comes two periods of eighteen each; after that a period of thirty-two, followed by an incomplete period. It was this variation in period size that kept earlier chemists from recognizing that a true periodicity existed. The plan of the piano was also simpler in the smooth transition from note to note along the octave. The arrangement for the elements showed a steady change at the beginning of the period and also at the close, with a marked break occurring as the elements moved on from metals to nonmetals—something that happened in every period but not at the same point in the period.
One thing more should be noted. The piano wires become heavier and heavier as the tones deepen in pitch. But the changes in heaviness do not in themselves produce the cyclic changes of the octaves; the overtones produced by vibrating wires do that. In the list of elements the changes in the atomic weights were not in themselves producing the cyclic changes of the periods; some factor, unknown at that time, was responsible.
One of Mendelyeev's great contributions to the understanding of periodicity among the elements was the construction of a periodic chart that would make the periods conspicuous. Each period formed a row on the chart, the rows being placed one above another. (The difficulty was in handling periods of so many different lengths.) Such a chart could indicate, close together, the elements that were most nearly alike in properties. One might show upon the chart, as a compact area, all elements that are gases; as another compact area, all metals used in alloy steels; as another area, the elements having colored compounds. Lines drawn on the chart can show the break between metals and nonmetals or point out elements of great hardness, or high melting points, or great heaviness. The use of the chart could never be a substitute of experimentation, but it was to give a sense of system and unity to the story of the elements such as had never been felt before. It is difficult to do justice to the life story of the author of the periodic law of the elements. The best approach to it may be through the remarkable story of a remarkable woman, his mother. The boy—twelve years younger than our General Grant of Civil War fame—was born in Tobolsk, Siberia. He was the fourteenth and youngest child. His father was an educated man, director of the college at Tobolsk. His mother came from a family that had pioneered in the manufacture of paper and glass. Her father was a printer—the printer of the first newspaper in Siberia. They were not political exiles. After the boy's birth the father gradually became blind from cataracts in both eyes, leaving the family with a government pension of about $500 a year. Needing additional funds, his mother reopened the glassmaking factory and acted personally as its manager. When her husband died she took her daughter Elizabeth and her son Dmitri, then fifteen, on the long and arduous journey to Moscow. She was determined that at least one of her children should be given a chance to study science intensively, for she saw industry's great need for it in Siberia. She felt that the boy was fully prepared for work at the University of Moscow. The officials turned the application down. She went on to St. Petersburg (Leningrad) to talk to an old friend of her husband. The boy was accepted for entrance into the physico-mathematical division of the college there. In addition, he was to receive a money-grant scholarship.
The rest of the story is the boy's. He made good at college, doing well in mathematics, physics, and chemistry. Just as he was completing his course he developed lung trouble. His mother died about the same time and his sister had married. He received, fortunately, a teaching position in southern Russia where, under favorable climatic conditions, he soon regained his health. Then he returned to St. Petersburg to teach chemistry at the college and to save his money. He wanted to go ahead with his mother's dream. He obtained approval from the minister of public education to leave Russia, and studied advanced chemistry in Paris under Robert Bunsen in Heidelberg. He returned to Russia with a doctor's degree in chemistry and was given a professorship, first at the Technological Institute and later at the University of St. Petersburg.
Mendelyeev remained in the St. Petersburg position for twenty-one years, then resigned. Always independent in his thinking, he had deeply resented having his classroom and laboratory work interfered with by officials, his research activities hedged about by political restrictions. A little later he was appointed director of the Bureau of Weights and Measures, and he kept that position until his death at the age of seventy-two.
As a person, Mendelyeev was quite a "character". He cut his hair but once a year—in the spring. He had an estate with servants and liked to be thought of as "the grand Russian of the province of Tver." He rarely wore ribbons and decorations, of which he had many. His wife, who was an artist, filled the home with sketches of the great chemists of history. Her influence apparently was responsible for his becoming a grand patron of art. He objected to his younger children going often to the theater. "It fills their minds with trifles and foolishness," was his comment.
In the brief period from 1894 to 1898 a new group of elements, five in number, was discovered and fitted into the periodic arrangement of the elements. Their presence had never been predicted; no one had sensed their omission. They were gases, and only Cavendish, and he a century before, was aware of the existence of even one of them. Cavendish had found it in the air, but had not thought of it as an element and had not named it. He had merely reported that the gas was chemically inert.
The finding of the first member of the group occurred in an indirect way. About the year 1890, an English physicist, Lord Rayleigh, had been officially checking the density values of the various gases, doing the task accurately and carefully. In the case of nitrogen the results were confusing and uncertain. Samples of this gas which he obtained from the air by removing the water vapor, carbon dioxide, and oxygen, leaving the nitrogen, proved to be about five parts in a thousand heavier than samples of nitrogen gas obtained by decomposing ammonia gas into hydrogen and nitrogen gases. Rayleigh reported his findings in a scientific journal and asked for readers' suggestions as to a possible reason for the observed differences of weight. Was some oxygen left with the nitrogen of the air? He had checked on that; the answer was "No". Was some hydrogen left with the nitrogen from ammonia? He had considered that; the answer, again, was "No". Did the nitrogen of the air have some molecules of extra-heavy weight—with some formula like N3--mixed in with the normal N2 molecules? That scarcely seemed possible. Then what could the explanation be?
William Ramsay, a Scotch chemist, was the only one to respond. He offered no solution, but he did have a suggestion. Why not use the spectroscope and the electric spark to get the spectrum lines of the gas? This could be used to find out if the nitrogen of the air had another element mixed in with it. With Rayleigh's cooperation Ramsay ran the test. The nitrogen lines appeared; there were also some red and green lines that were not those of nitrogen. In fact, those particular lines did not belong to the spectrum of any known element. A new element had been found. This spectroscope test was made in May 1894. At the August meeting of the British Association, Ramsay and Rayleigh gave their joint report on the discovery of this new element to be called argon, "the lazy one". The report listed the properties of the gas that they had found. To get the atomic weight in a chemical way, a compound of argon with one or more other elements was wanted. Ramsay reported that he had made dozens of attempts to prepare such a compound, and had failed. There was not a single element or group of elements with which this lazy element would combine. So they had assumed that each molecule of the gas contained one atom only; then by using the idea of Avogadro they had measured the atomic weight of argon as 39.88. This value put argon between chlorine, a chemically active nonmetal, and potassium, a chemically active metal, in the list of elements. The position of argon was understandable. It was neither nonmetal nor metal but an inert gas.
Helium was the next of the members of the inert-gas group to be obtained. Strangely enough, it had been named about a quarter of a century before—though the name did not fit the real substance very well. The story of the naming goes back to the eclipse of the sun in the year 1868. The spectroscope had been invented a few years before and there was some curiosity as to what the instrument would indicate about the composition of the sun's atmosphere. When the sun itself was darkened by the eclipse, the flaming corona of the sun's atmosphere would continue—dimly lighting up the earth. Would this corona show the same elements that the chemist had already found for this earth, or would there be new lines for special sun elements? Scientists watched carefully. As the dim coronal light took the place of the sun's direct rays, the spectrum changed abruptly from the full sun spectrum to the bright-line spectrum of the laboratory. Many earth elements were there, but one bright line in the yellow part of the coronal spectrum was strikingly new. It belonged to some element in the sun's atmosphere that was evidently quite abundant, for the line was clear and strong. Named by Joseph Norman Lockyer, the English astronomer, as helium, "metal of the sun," the element's properties could only be guessed at.
That was all that anyone knew of the "sun's element" for many years. About 1890, the American mineralogical chemist William F. Hillebrand reported an odd observation made on a certain uranium-containing mineral that had received the name of uraninite. In testing for its contents in a regular way he had dissolved the mineral in acid. The substance gave off a gas. But the gas was not carbon dioxide, for that is easily tested for, and it was not hydrogen, for it did not burn. He assumed that it was nitrogen but did not test the gas to see, since nitrogen is not readily tested for in any chemical way. It never occurred to him to ask for a spectroscope test using the electric spark.
In 1894 Ramsay was methodically going over the mineralogical journal, seeking clues as to possible argon compounds in nature, when his eye was caught by Hillebrand's comment about uraninite. He looked for that mineral in the English rock collections. Not finding any sample, he picked out a sample of a uranium-containing mineral from Sweden. In the acid test it gave out some gas which was neither carbon dioxide nor hydrogen. Hurriedly he set up the spectroscope with its electric-spark equipment. He found that the gas was a mixture. There were lines for nitrogen; there were lines for argon. But there were others—strange ones. Ramsay sent a sample of the gas to William Crookes, who had the best spectroscope in England. He sent another sample to Lockyer, who was an international authority on the spectra of the elements. Crookes reported that the gas sample showed something new. Lockyer was quite enthusiastic. He wrote back, "When I tested the gas, the glorious yellow effulgence of the capillary while the current was passing, was a sight to see."
Within a week it was shown that the main yellow line of the spectrum of the tested gas agreed in position with the main yellow line of the sun's element ascribed to helium. So the long-looked-for "metal of the sun" turned out to be not metal at all but an element as inert as argon. A search then showed that this very light, unburnable gas was likely to be present in small amounts in uranium-containing minerals. Natural gas from wells along the Rocky Mountains of the United States and Canada sometimes carries helium in commercial amounts. The atmosphere contains a mere trace. After the finding of argon and helium, Ramsay predicted three other members of this group. All should be gases, all should be chemically inert, not combining even with fluorine. The logical place to look for them was in the air. By this time the production of liquid air had developed into an important industry and Ramsay sought the aid of the men in this industry in his attempt to find the three inert-gas elements he was looking for. As the air was cooled intensely and kept under pressure, the nitrogen, oxygen, and argon changed from gases to liquids. A very small fraction of the air remained. It was in that fraction that Ramsay looked for, and found, the element neon, "new". When tested with the spectroscope and the electric spark it gave a surprisingly bright light that startled Ramsay and Morris Travers, his assistant. Travers Wrote:
The blaze of crimson light from the little tube containing a tiny quantity of neon told its own story, and it was a sight to dwell upon and never forget. It was worth the struggle of the previous years—and all the difficulties yet to be overcome before the research was finished. But for the moment when the electricity was turned on the actual spectrum of the gas did not matter, for nothing in the world gave such a glow as we have seen.6
That was the first neon light. Observed for the first time in June 1898, "the blaze of crimson light" is now a familiar one throughout commercial centers.
Ramsay predicted that the remaining two members of the group of inert gases would appear as tiny crystals in the liquid air. They were found. One was named krypton, "hidden". It required 120 tons of liquid air to get the other one in an amount that could be tested. It was named xenon, "stranger". Ramsay estimated that one molecule of xenon was present in 170,000,000 molecules of air! For the discovery and separation of the inert gases, and for other chemical work of high order, Ramsay was knighted in 1902. He was to be called, after that, Sir William Ramsay. He was awarded the 1904 Nobel prize in chemistry—a world honor.
These may seem great accomplishments for a boy who had studied no science until a senior at the University of Glasgow, Scotland. In the fall of that year he broke his leg playing football. While laid up, he picked up a chemistry textbook hoping to find how to make fireworks. The book was interesting, and he read it completely through. After that he decided to change his bedroom into a sort of laboratory. In that he had the active aid of a school friend. The friend has recorded some of the events of the months that followed.
There were a great many bottles always about—containing acids, salts, mercury and so on. We used to meet in the afternoons and do what practical work we could, making oxygen and hydrogen and various simple compounds such as oxalic acid from sugar. We made up nearly all the apparatus we used except flasks, retorts and beakers.7
After graduation in the spring Ramsay returned to the university for another year, this one given entirely to the study of physics and chemistry. He went to Germany for advanced study in chemistry, then returned first to Scotland and then to England as a teacher. After 1887 he held the chair of chemistry at University College, London. It was there that his research work on the inert gases was carried out. He was noted as a lecturer, having a rare sense of humor that greatly pleased his audiences. He was a great linguist. It is stated that he could lecture in perfect German before a cultured German audience, or in French before an assembly of French scientists. When, in 1913, he presided over the International Association of Chemical Societies, he addressed remarks in Italian to the Italian delegates, in French to the French group, in German to the German scientists, and in English in his main address—all with perfect skill and composure.
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