Inventions in Chemistry

Oxygen (1770s) [pic] British educator and philosopher Joseph Priestley (1733 – 1804) discovered oxygen in experiments, isolated the gas, and described its function in combustion and respiration. He also invented soda or carbonated water by dissolving fixed air with water. Unaware of the significance of his discoveries and because of his stubborn refusal to abandon the phlogiston theory, he named the new gas “dephlogisticated air. However, it would be the French chemist Antoine Lavoisier (1743 – 1794) who gave the gas its present name, and was able to explain the nature of the element, accurately describing its role in combustion that totally discredit the phlogiston theory. In addition, Lavoisier collaborated with others to develop a systematic chemical nomenclature that facilitates dialogue among chemists and is still very much in use today. Who Discovered Oxygen? Everyone needs oxygen to survive – man and animals alike.

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Furthermore, oxygen is the third most abundant element in the universe and makes up nearly 21% of the earth’s atmosphere. Oxygen accounts for nearly half of the mass of the earth’s crust, two thirds of the mass of the human body and nine tenths of the mass of water. In this page we will try to outline the path to the discovery of this important substance. Oxygen was discovered for the first time by a Swedish Chemist, Carl Wilhelm Scheele, in 1772. Joseph Priestly, an English chemist, independently, discovered oxygen in 1774 and published his findings the same year, three years before Scheele published.

Antoine Lavoisier, a French chemist, also discovered oxygen in 1775, was the first to recognize it as an element, and coined its name “oxygen” – which comes from a Greek word that means “acid-former”. There is a historic dispute about who discovered oxygen. Most credit Priestly alone or Both Priestly and Scheele. To learn more about this dispute go to the link section, at the bottom of this page. Famous Experiments Oxygen from Minerals In 1772, Carl Wilhelm Scheele discovered that red-hot manganese oxide produces a gas. He called the gas “fire air” because of the brilliant sparks it produced when it came in contact with hot charcoal dust.

He repeated this experiment by heating potassium nitrate, mercury oxide, and many other materials and produced the same gas. He collected the gas in pure form using a small bag. He explained the properties of “fire air” using the phlogiston theory, which was soon discredited by Lavoisier. He carefully recorded his experiments in his notes, but waited several years before publishing them. In 1774, Priestley repeated Scheele’s experiments using a 12-inch-wide glass “burning lens”, he focused sunlight on a lump of reddish mercuric oxide in an inverted glass container.

The gas emitted, he found, was “five or six times as good as common air. ” (1) In succeeding tests, it caused a flame to burn intensely and kept a mouse alive about four times as long as a similar quantity of air. Priestley, a big supporter of the phlogiston theory, called his discovery “dephlogisticated air” on the theory that it supported combustion so well because it had no phlogiston in it, and hence could absorb the maximum amount during burning. Repeat Scheele’s and Priestley experiments:. Oxygen from Plants In August of 1771, Joseph Priestley, put a sprig of mint into a transparent losed space with a candle that burned out the air until it soon went out. After 27 days, he relit the extinguished candle again and it burned perfectly well in the air that previously would not support it. And how did Priestley light the candle if it was placed in a closed space? He focused sun light beams with a mirror onto the candle wick (Priestley had no bright source of light, and had to rely on the sun). Today, of course, we can use more sophisticated methods to light the candle like focusing light from a flood light through converging lens, or by an electrical spark.

So priestly proved that plants somehow change the composition of the air. In another celebrated Experiment from 1772, Priestley kept a mouse in a jar of air until it collapsed. He found that a mouse kept with a plant would survive. However, we do not recommend to repeat this experiment and hurt innocent animals. These kinds of observations led Priestley to offer an interesting hypothesis that plants restore to the air whatever breathing animals and burning candles remove – what was later coined by Lavoisier “oxygen”.

In these experiments, Priestly was the first to observe that plants release oxygen into the air – the process known to us as photosynthesis. Atomic Theory (1800s) [pic] [pic][pic][pic][pic][pic] John Dalton (1766 – 1844), English chemist and physicist, proposed the atomic theory, which states that: a. ) all elements are made up of tiny particles called atoms; b. ) all atoms of an element are identical; c. ) the atoms of dissimilar elements can be distinguished from one another by their corresponding relative weights; d. atoms of an element can be combined with atoms of another elements to form chemical compounds; and e. ) atoms cannot be created, broken down into smaller particles, nor destroyed in a chemical process. He also presented a way of associating invisible atoms with quantifiable amounts such as mass of a mineral or volume of a gas. Dalton’s theory has undergone modifications through the centuries, but it has as much significance for the future of the science as Lavoisier’s oxygen-based chemistry had been. Molecules are Made Up of Atoms (1810s – ) [pic] Read more in Chemistry How to Create a Cloud in a Bottle Learning About Chemistry: Introduction » At a time when the words “atom” and “molecule” were used interchangeably, Italian scientist Amedeo Avogadro (1776 – 1856) clarified that atoms combine to form molecules; and proposed his eponymous principle which asserts that “Equal volumes of ideal gases, at the same conditions of temperature and pressure, contain equal numbers of particles or molecules. ” [pic] Marie Curie was born on November 7 1867 and moved to Paris when she was 24 to study science and become a naturalized French citizen.

At the university, she met another instructor to whom she would eventually marry named Pierre Curie. The two worked together studying radioactive compounds and refining uranium pitchblende ore. Besides radium, Marie Curie also discovered a new substance polonium, which she named after her native homeland of Poland. In 1903, Marie and Pierre Curie and Henri Becquerel all received the Nobel Prize in physics for their research on radiation. Marie Curie was the first woman to ever receive the Nobel Prize. In 1911, Marie Curie received her second Nobel Prize, this time in chemistry for the discovery of radium and polonium.

Other than Marie Curie, Linus Pauling is the only other person to ever received two Nobel Prizes. Curie intentionally decided not to patent the process to extract and purify radium, leaving the door open to the scientific community to study the process unhindered. While being a brilliant scientist, Marie Curie also suffered from depression and kidney problems. During World War I, Marie Curie pushed for the use of radiography field units for the treatment of wounded soldiers. The units included tubes of radon gas that Curie purified herself.

Promptly after the war had started, Marie Curie don’t she and her husband’s gold Nobel Prize medals to help with the war efforts. On July 4, 1934, Marie Curie died of leukemia, which may have been caused by her exposure to radiation. Rumor Has It … Rumor has it that Marie Curie actually invented clog dancing and did so frequently in the bars and restaurants of Paris, much to the chagrin of her husband, Pierre. In fact, it was the clog dancing that contributed to his untimely death in the streets of Paris as he was racing to get away from the noise and clatter of the clogs at the time.

In a completed fabricated, yet entertaining bit of balderdash, it was rumored that after Pierre’s death, Marie Curie had begun an affair with a dwarf named Tito from the African north coast. Upon witnessing a drunken bit of consensual dwarf tossing at the local bar, Marie broke off the affair saying that her new love interest had offended her sense of decency and propriety. A recent discovery of an element has taken place for the periodic table. The chemical element known as WOMAN (symbol Wo) has an atomic weight of 120 (isotopic) and certain other characteristics.

It is found wherever man is. It can freeze at any time, melts when properly treated, and is bitter when not used well. It is seldom found in the free state. Wo is generally roundish in form and boils at 0 degrees Celcius. It is very active, possessing a great affinity for gold, silver, platinum, and precious stones. However, it undergoes violent reaction when left alone. It is able to absorb great quantities of food as well as energy. It turns green or red when placed beside a better looking specimen and seems to age rapidly. The element is highly ornamental and useful.

It equalizes the distribution of wealth and is probably the most powerful income agent known. This element should be respected as it is highly EXPLOSIVE when in inexperienced hands. Boyles law [pic] Boyle’s Law Toward the end of Van Helmont’s life, gases – air, in particular, since it was the most common gas – were attaining a new and dramatic importance. The Italian physicist Evangelista Torricelli (1608-1647) was able to prove, in 1643, that air exerted pressure. He showed that air could support a column of mercury thirty inches high and, in so doing, he invented the barometer.

Gases at once became less mysterious. They were matter, possessing weight as did the more easily studied liquids and solids. They differed from liquids and solids chiefly in their much lower density. The pressure exerted by the weight of the atmosphere was demonstrated in an astonishing manner by the German physicist Otto von Guericke (1602-1686). He invented an air pump with which he could pull the air out of containers, so that the air pressure on the outside was no longer equalized by air pressure on the inside. In 1654 Guericke prepared two metal hemispheres that fit together along a greased flange.

When the hemispheres were put together and the air within was removed by the air pump, air pressure from without held the hemispheres together. Teams of horses attached to each of the hemispheres and whipped into straining their utmost in opposite directions, could not pull the hemispheres apart. When air was allowed to re-enter the joined hemispheres, however, they fell apart of themselves. Demonstrations such as this roused great interest in the properties of air. In particular, the curiosity of an Irish chemist, Robert Boyle (1627-1691), was roused. He devised an air pump of his own that was even better than Guericke’s.

Then having, so to speak, pulled air apart in sucking it out of the container, he went on to try the opposite procedure of compressing it – that is, of pushing it together. In his experiments Boyle found that the volume of a sample of air varied with pressure according to a simple inverse relationship. He discovered this by dropping mercury into a very long, specially constructed tube and trapping a sample of air in the short, closed end which was fitted with a stopcock. By adding more mercury to the long open end he could increase the pressure on the trapped air.

If he added enough mercury to place the trapped air under doubled pressure ( a doubled weight of mercury), the volume of the trapped air was halved. If pressure was tripled the volume was reduced to a third. On the other hand, if pressure was eased off the volume expanded. This relationship whereby volume decreased in proportion as pressure increased was first published in 1660 and is still referred to as Boyle’s Law. This was the first attempt to apply exact measurement to changes in a substance of particular interest to chemists. It must be pointed out, though, that the change studied by Boyle was not a chemical one.

Air, however, it might be compressed or expanded, remains air. Such change in volume is a physical change. He was therefore involve in physical chemistry, the study of the physical changes of chemicals. This was not to come into its own for two centuries after the time of Boyle, but he laid the groundwork. Boyle did not specify that temperature must be held constant if Boyle’s law is to be valid. Probably he realized this and supposed it would be taken for granted. The French physicist Edme Mariotte (1630-1684), who discovered Boyle’s Law independently, about 1680, did specify that temperature must be held constant.

For this reason, Boyle’s Law is often referred to as Mariott’s law in continental Europe. Boyle’s experiments offered a focus for the gathering numbers of atomists. As stated earlier, Lucretius’s poem, introduced in a printed edition, had brought Greek views on atomism to the attention of European scholars. A French philosopher, Pierre Gassendi (1592-1655), was a convinced atomist as a result, and his writings impressed Boyle, who thereupon also became an atomist. As long as one concentrated on liquids and solids only, the evidence for atomism was no better in Boyle’s time than in that of Democritus.

Liquids and solids cannot be compressed by more than insignificant amounts. If they consist of atoms, those atoms must be in contact and cannot be pushed closer together. It is therefore hard to argue that liquids and solids must be made up of atoms, for if they were made up of continuous substances they would also be very difficult to compress. Why bother with atoms, then? Air as had been observe even in ancient times, and as Boyle had now made dramatically clear, can easily be compressed. How could this be unless it consisted of tiny atoms separated by empty space?

Compressing air simply would mean, from that point of view, the squeezing of empty space out of the volume, pushing the atoms closer together. If this view of gases is accepted it becomes easier to believe that liquids and solids are composed of atoms, too. For instance, water evaporates. How can that be unless it disappears tiny bit by tiny bit, and what could be simpler, then, than to suppose that it turns into vapor atom by atom? If water is heated it boils and vapor is visibly formed. The water vapor has the physical properties of an air-like substance and therefore, it is natural to suppose, is composed of atoms.

But if water is composed of atoms in its gaseous form, why not in its liquid form as well, and in its solid form of ice? And if this is true of water, why not of all matter? Arguments of this sort were impressive, and for the first time since atoms were first imagined two thousand years before, atomism began to win numerous converts. Newton, for instance, became an atomist. Nevertheless, atoms remained a misty concept. Nothing could be said about them except that if they were assumed to exist, it was easier to explain the behavior of gases. Another century and a half had to pass before atomism came into sharp focus.

The New View of Elements Boyle’s career marks the passing of the terms “alchemy” and “alchemist”. Boyle dropped the first syllable of the term in writing a book, The Sceptical Chymist, publish in 1661. From that time on, the science was chemistry and workers in the field were chemists. Boyle was “sceptical” because he was no longer willing to accept, blindly, the ancient conclusions that had been deduced from first principles. In particular, Boyle was dissatisfied with ancient attempts to identify the elements of the universe by mere reasoning. Instead, he defined elements in a matter-of-fact, practical way.

An element, it had been considered ever since Thales’ time, was one of the primal simple substances out of which the universe was composed. Well, then, a suspected element must be tested in order to see if it were really simple. If a substance could be broken into simpler substances it was not an element, but the simpler substances might be – until such time as chemists learned to break them down to still simpler substances. Furthermore, if two substances were each an element, they might be intimately combined to form a third substance called a compound.

If so, then that compound should lend itself to breakdown into the two original elements. The term “element” in this view, had only a practical meaning. A Substance such as quartz, for instance, could be considered an element until such time as experimental chemists discovered a way of converting it into two or more still simpler substances. In fact, no substance could ever be an element except in a provisional sense, according to this view, since one could never be certain when advancing knowledge might make it possible to devise a method for breaking down a supposed element into still simpler substances.

It was not until the coming of the twentieth century that the nature of elements could be defined in a non-provisional sense. The mere fact that Boyle wanted an experimental approach in defining elements (an approach that was adopted eventually) does not mean that he knew what the different elements were. It might have turned out, after all, that the experimental approach would indeed have proved Greek elements of fire, air, water, and earth to be elements. Boyle was convinced, for instance, of the validity of the alchemical viewpoint that the various metals were not elements and that one metal could be converted into another.

In 1689, he urged the British government to repeal the law against the alchemical manufacture of gold (they, too, feared the upset to the economy) because he felt that by forming gold out of base metal, chemists could help to prove the atomic view of matter. But Boyle was wrong there; the metals did prove to be elements. In fact, nine substances which we now recognize as elements had been known to the ancients; the seven metals (gold, silver, copper, iron, tin, lead, and mercury) and two non-metals (carbon and sulfur).

In addition, there were four substances now recognized as elements that had become familiar to the medieval alchemists: arsenic, antimony, bismuth, and zinc. Boyle, himself, came within a hair of being the discoverer of a new element. In 1680 he prepared phosphorus from urine. Some five to ten years before that, however, the feat had been accomplished by a German chemist, Hennig Brand (? -1692). Brand is sometimes called the “last of the alchemists”, and, indeed, his discovery came while he was searching for the philosopher’s stone which he thought he would find in (of all places) urine.

Brand was the first man to discover an element that had not been known, in a least some form, before the development of modern science. Phlogiston The seventeenth-century discoveries concerning air pressure and the unusual feat that one could perform by producing a vacuum and allowing air pressure to work, had important results. It occurred to several people that a vacuum might be formed without the use of an air pump. Suppose you boiled water and filled a chamber with steam, then cooled the chamber with cold water on the outside. The steam within the chamber would condense into a few drops of water, and a vacuum would exist in its place.

If one of the walls of the chamber were movable, air pressure on the other side would then drive that wall into the chamber. The movable wall could be pushed outward again if more steam were formed and allowed to enter the chamber, and then be pushed inward again if the steam were once more condensed. If you imagine the movable wall to be part of a piston, you can see that the piston will move in and out and that this in-an-out motion could be used to run a pump, for instance. By 1700, such a steam engine had actually been produced by an English engineer, Thomas Savery (c. 650-1715). It was a dangerous device because it used steam under high pressure at a time when high-pressure steam could not be safely controlled. However, another Englishman, Thomas Newcomen (1663-1729), working in partnership with Savery, devised a steam engine that would work on low-pressure steam. The device was improved and made really practical, toward the end of the eighteenth century, by the Scottish engineer James Watt (1736-1819). The result of these labors was that, for the first time, mankind was no longer dependent upon its own muscles or upon the muscles of animals.

Nor was man dependent upon the hit-or-miss force of the wind, or upon the spottily located energy of running water. Instead, he had a source of energy he could call upon at any time and in any place merely by boiling water over a wood or coal fire. This was the chief factor marking the start of the “Industrial Revolution”. The increasing interest from 1650 onward in the possibility of turning fire to new uses and, by way of the steam engine, making it do the heavy work of the world, brought to chemists a new awareness of fire. Why do some things burn and others not? What is the nature of combustion?

By old Greek notions something which could burn contained with itself the element of fire, and this something was released under the proper conditions. Alchemical notions were similar, except that a combustible was thought of as containing the principle of “sulfur” (though not necessarily actual sulfur). In 1669, a German chemist, Johann Joachim Becher (1635-1682), tried to rationalize this notion further, by introducing a new name. He imagined solids to be composed of three kinds of “earth”. One of these he called “terra pinguis” (“fatty earth”), and felt this to be the principle of inflammability.

A follower of Becher’s rather vague doctrines was the German physician and chemist Georg Ernest Stahl (1660-1734). He advanced a newer name still for the principle of inflammability, calling it phlogiston, from a Greek word meaning “to set on fire”. He went on to devise a scheme, involving phlogiston, that would explain combustion. Combustible objects, Stahl held, were rich in phlogiston, and the process of burning involved the loss of phlogiston to the air. What was left behind after combustion was without phlogiston and therefore could no longer burn.

Thus, wood possessed phlogiston, but ash did not. Stahl maintained further that the rusting of metals was analogous to the burning of wood, and he considered a metal to possess phlogiston while its rust (or “calx”) did not. This was an important insight, which made it possible to advance a reasonable explanation of the conversion of rocky ores into metals – the first great chemical discovery of civilized man. The explanation consisted of this: A rocky ore, poor in phlogiston, is heated with charcoal, which is very rich in phlogiston.

Phlogiston passes from the charcoal into the ore, so that the phlogiston-rich charcoal is turned into phlogiston-poor ash, while the phlogiston-poor ore is turned into phlogiston-rich metal. Air itself was considered by Stahl to be only indirectly useful to combustion, for it served only as a carrier, holding the phlogiston as it left the wood or metal and passing it on to something else (if something else were available). Stahl’s phlogiston theory met with opposition at first, notable from a Dutch physician, Hermann Boehaave (1668-1738), who argued that ordinary combustion and rusting could not be different versions of the same phenomenon.

To be sure, there is the presence of flame in one case and not in the other, but to Stahl the explanation was that in the combustion of substances, such as wood, phlogiston left so rapidly that its passage heated its surroundings and became visible as flame. In rusting the loss of phlogiston was slower and no flame appeared. Despite Boerhaave’s opposition, then, the phlogiston theory gained popularity throughout the eighteenth century. By 1780 it was almost universally accepted by chemists, since it seemed to explain so much so neatly. Yet a difficulty remained that neither Stahl nor any of his followers could explain.

Most combustible objects, such as wood, paper, and fat, seemed largely to disappear upon burning. The remaining soot or ash was much lighter than the original substance. This is to be expected, perhaps, since phlogiston had left that original substance. However, when metals rusted, they also lost phlogiston, according to Stahl’s theory, yet the rust was heavier than the original metal (a fact which had been noted by alchemists as early as 1490). Could phlogiston have negative weight, then, so that a substance that lost it was heavier than before, as some eighteenth-century chemists tried to maintain?

If so, why did wood lose weight in burning? Were there two kinds of phlogiston, one with weight and one with negative weight? This unanswered problem was not as serious in the eighteenth century as it seems to us today. We are used to measuring phenomena accurately, and an unexplained change in weight would disturb us. The eighteenth-century chemists, however, had not yet accepted the importance of accurate measurements, and they could shrug off the change in weight. As long as the phlogiston theory could explain changes in appearance and properties, changes in weight, they felt, could be ignored.

Mendeleev [pic] The Periodic Table [pic] Elements in Disorder There is a curious parallel in the histories of the organic chemistry and inorganic chemistry of the nineteenth century. The opening decades of the century saw a puzzling proliferation in the number of organic compounds, and also in the number of elements. The third quarter of the century saw the realm of organic compounds reduced to order, thanks to Kekule’s structural formula. It saw the realm of elements reduced to order also, and at least part of the credit for both changes goes to events at a particular nternational meeting of chemists. But let’s begin with the disorder at the beginning of the century. The discovery of elements over and above the nine known to the ancients and the four studied by medieval alchemists has been previously discussed. The gaseous elements, nitrogen, hydrogen, oxygen, and chlorine, had all been discovered in the eighteenth century. So had the metals, cobalt, platinum, nickel, manganese, tungsten, molybdenum, uranium, titanium, and chromium. In the first decade of the nineteenth century, no less than fourteen new elements were added to the list.

Among the chemists already mentioned in this work, Davy had isolated no fewer than six by means of electrolysis. Gay-Lussac and Thenard had isolated boron; Wollaston had isolated palladium and rhodium, while Berzelius had discovered cerium. Then, too, the English chemist Smithson Tennant (1761-1815), for whom Wollaston had worked as an assistant, discovered osmium and iridium. Another English chemist, Charles Hatchett (c. 1765-1847), isolated columbium (now officially called niobium), while a Swedish chemist, Anders Gustaf Ekebert (1767-1813), discovered tantalum.

The haul in succeeding decades was not quite as rich, but the number of elements continued to mount. Berzelius discovered four more elements: selenium, silicon, zirconium, and thorium. Louis Nicolas Vauquelin in 1797 discovered beryllium. By 1830, fifty-five different elements were recognized, a long step from the four elements of ancient theory. In fact, the number was too great for the comfort of chemists. The elements varied widely in properties and there seemed little order about them. Why were there so many? And how many more yet remained to be found? Ten? Fifty? A hundred? A thousand? An infinite number?

It was tempting to search for order in the list of elements already known. Perhaps in this manner some reason for the number of elements might be found and some way of accounting for the variation of properties that existed. The first to catch a glimmering of order was the German chemist Johann Wolfgang Dobereiner (1780-1849). In 1829, he noted that the element bromine, discovered three years earlier by the French chemist Antoine Jerome Balard (1802-1876), seemed just halfway in its properties between chlorine and iodine. (Iodine had been discovered by another French chemist, Bernard Courtois (1777-1838), in 1811. Not only did chlorine, bromine, and iodine show a smooth gradation in such properties as color and reactivity, but the atomic weight of bromine seemed to lie just midway between those of chlorine and iodine. Coincidence? Dobereiner went on to find two other groups of three elements exhibiting neat gradations of properties: calcium, strontium, and barium; and sulfur, selenium, and tellurium. In both groups the atomic weight of the element in the middle was about midway between those of the other two. Coincidence again? Dobereiner called these groups “triads”, and searched unsuccessfully for others.

The fact that five-sixths of the known elements could not be fitted into any triad arrangement made chemists decide that Dobereiner’s findings were merely coincidence. Furthermore, the manner in which atomic weights fit along with the chemical properties among the elements of Dobereiner’s triads did not impress chemists generally. In the first half of the nineteenth century, atomic weights tended to be underestimated. They were convenient in making chemical calculations, but there seemed no reason to use them in making lists of the elements. In was even doubtful that atomic weights were useful in making chemical calculations.

Some chemists did not distinguish carefully between atomic weight and molecular weight; some did not distinguish between atomic weight and equivalent weight. Thus, the equivalent weight of oxygen is 8, the atomic weight is 16, and the molecular weight is 32. In chemical calculations the equivalent weight, 8, is handiest; why then should the number 16 by used to determine the place of oxygen in the list of elements? This confusion among equivalent weight, atomic weight, and molecular weight spread its disorganizing influence not merely over the question of the list of elements but into the study of chemistry generally.

Disagreements over the relative weights to assign to different atoms led to disagreements over the number of atoms of particular elements within a given molecule. Kekule, shortly after he had published his suggestions leading to structural formulas, realized this concept would come to nothing if chemists could not agree, first of all, on empirical formulas. He therefore suggested a conference of important chemists from all over Europe to discuss the matter. As a result, an international scientific meeting was held for the first time in history.

It was called the First International Chemical Congress, and it met in 1860 in the town of Karlsruhe, in Germany. One hundred forty delegates attended, among them the Italian chemist Stanislao Cannizzaro (1826-1910). Two years earlier, Cannizzaro had come across the work of his countryman Avogadro. He saw how Avogadro’s hypothesis could be used to distinguish between the atomic weight and molecular weight of the important gaseous elements and how this distinction would serve to clarify the matter of atomic weights for the elements generally.

Furthermore, he saw the importance of distinguishing carefully between atomic weight and equivalent weight. At the Congress he made a strong speech on the subject and then distributed copies of a pamphlet in which he explained his points of view. Slowly and laboriously, he won over the chemical world to his views. From that time forward, the matter of atomic weight was clarified and the importance of berzelius’s table of atomic weights was appreciated. In organic chemistry this development meant that mean could now agree on empirical formulas and proceed onward to add detail in structural form, first in two dimensions, then in three.

In inorganic chemistry, the results were just as fruitful, for there was now at least one rational order in which to arrange the elements – in order of increasing atomic weight. Once that was done, chemists could look at the list with fresh eyes. Organizing the Elements In 1864, the English chemist John Alexander Reina Newlands (1837-1898) arranged the known elements in order of increasing atomic weights, and noted that this arrangement also placed the properties of the elements into at least a partial order. When he arranged his elements into vertical columns of seven, similar elements tended to fall into the same horizontal rows.

Thus, potassium fell next to the very similar sodium; selenium fell in the same row as the similar sulfur; calcium next to the similar magnesium, and so on. Indeed, each of Dobereiner’s three triads were to be found among the rows. Newlands called this the law of octaves (there are seven notes to an octave in music, the eighth note being almost a duplicate of the first note and beginning a new octave. ) Unfortunately, while some of the rows in his table did contain similar elements, other rows contained widely dissimilar elements.

It was felt by other chemists that what Newlands was demonstrating was coincidence rather than something of significance. He failed to get his work published. Two years earlier, a French geologist, Alexandre Emile Beguyer de Chancourtois (1820-1886) had also arranged elements in order of increasing atomic weight and had plotted them on a sort of cylindrical graph. Here, too, similar elements tended to fall into vertical columns. He published his paper, but not his graph, and his work went unnoticed, also. More successful was the German chemist Julius Lothar Meyer (1830-1895).

Meyer considered the volume taken up by certain fixed weights of the various elements. Under such conditions, each weight contained the same number of atoms of its particular element. This meant that the ratio of the volumes of the various elements was equal to the ratio of the volumes of single atoms of the various elements. Therefore, one could speak of atomic volumes. If the atomic volumes of the elements were plotted against the atomic weight, a series of waves was produced, rising to sharp peaks at the alkali metals: sodium, potassium, rubidium, and cesium. Each fall and rise to a peak corresponded to a period in the table of elements.

In each period a number of physical properties other than atomic volume also fell and rose. Hydrogen, the first in the list of elements (it has the lowest atomic weight) is a special case and can be considered as making up the first period all by itself. The second and third period in Meyer’s table included seven elements each, and duplicated Newlands’s law of octaves. However, the two waves following included more than seven elements, and this clearly showed where Newlands had made his mistake. One could not force the law of octaves to hold strictly throughout the table of elements, with seven elements in each row.

The later periods had to be longer than the earlier periods. Meyer published his work in 1870, but he was too late. The year before, a Russian chemist, Dmitri Ivanovich Mendeleev (1834-1907), had also discovered the change in length of the periods of elements, and then went on to demonstrate the consequences in a particularly dramatic fashion. Mendeleev was taking his graduate work in Germany at the time of the Karlsruhe Congress, and he was one of those who attended and heard Cannizzaro express his views on atomic weight.

After his return to Russia, he, too, began to study the list of elements in order of increasing atomic weight. Mendeleev tackled matters from the direction of valence. He noted that the earlier elements in the list showed a progressive change in valence. That is, hydrogen had a valence of 1, lithium of 1, beryllium of 2, boron of 3, carbon of 4, nitrogen of 3 (5), sulfur of 2 (6), fluorine of 1 (7), sodium of 1, magnesium of 2, aluminum of 3, silicon of 4, phosphorus of 3 (5), oxygen of 2 (6), chlorine of 1 (7), and so on.

Valence rose and fell, establishing periods; first, hydrogen itself; then two periods of seven elements each; then periods containing more than seven elements. Mendeleev used his information to prepare not merely a graph, as Meyer and Beguyer de Chancourtois, had, but a table like that of Newlands. Such a periodic table of the elements was clearer and more dramatic than a graph, and Mendeleev avoided Newlands’s mistake of insisting on equal periods throughout. Mendeleev published his table in 1869, the year before Meyer published his work.

However, the reason the lion’s share of the credit for the discovery of the periodic table is accorded to him over the other contributions is not a mere matter of priority of publication. It rests instead on the dramatic use to which Mendeleev put his table. In order to make the elements fit the requirements that those in a particular column all have the same valence, Mendeleev was forced in one or two cases to put an element of slightly higher atomic weight ahead of one of slightly lower atomic weight. This, tellurium (atomic weight 127. 6, valence 2) had to be put ahead of iodine (atomic weight 126. , valence 1) in order to keep tellurium in the valence-2 column and iodine in the valence-1 column. (His instinct in this respect led him in the correct direction, though the reason for it wasn’t made clear for nearly half a century) As if this were not enough, he also found it necessary to leave gaps altogether in his table. Rather than considering these gaps as imperfections in the table, Mendeleev seized upon them boldly as representing elements as yet undiscovered. In 1871, he pointed to three gaps in particular, those falling next to the elements boron, aluminum, and silicon in the table as modified in that year.

He went so far as to give names to the unknown elements that he insisted belonged in those gaps; eka-boron, eka-aluminum, and eka-silicon (“eka” is the Sanskrit word for “one”). He also predicted various properties of these missing elements, judging what these must be from the properties of the elements above and below the gaps in his table – thus following and completing the insight of Dobereiner. The world of chemistry remained skeptical and would perhaps have continued so if Mendeleev’s bold predictions had not been dramatically verified. That this happened was due, first of all, to use of a new chemical tool – the spectroscope.

Filling the Gaps In 1814, a German optician, Joseph von Fraunhofer (1787-1826), was testing the excellent prisms he manufactured. He allowed light to pass first through a slit and then through his triangular glass prisms. The light, he found, formed a spectrum of color that was crossed by a series of dark lines. He counted some six hundred of these lines, carefully noting their positions. These lines were made to yield startling information, in the late 1850’s, by the German physicist Gustav Robert Kirchhoff (1824-1887), working with the German chemist Robert Wilhelm Bunsen (1811-1899).

The basic source of light they used was a Bunsen burner, invented by Bunsen and known to every beginning student in a chemistry laboratory down to this day. This device burns a mixture of gas and air to produce a hot, scarcely luminous flame. When Kirchhoff placed crystals of various chemicals in the flame, it glowed with light of particular colors. If this light was passed through a prism it separated into bright lines. Each element, Kirchhoff showed, produced a characteristic pattern of bright lines when heated to incandescence, a pattern different from that of any other element.

Kirchhoff had this worked out a method of “fingerprinting” each element by the light it produced when heated. Once the elements had been fingerprinted, he could work backward and deduce the elements in an unknown crystal from the bright lines in its spectrum. The device used to analyze elements in this fashion was named the spectroscope. As we know today, light is produced as a result of certain events that occur within the atom. In each type of atom these events occur in a particular manner. Therefore, each element will emit light of certain wavelengths and no others.

Light falls upon vapor, those same events within the atoms of the vapor can be made to occur in reverse. Light of certain wavelengths is then absorbed rather then emitted. What’s more, since the same events are involved in either case (forward in one case, backward in the other), the wavelengths of light absorbed by vapor under one set of conditions are exactly the same as those that particular vapor would emit under another set of conditions. The dark lines in the spectrum of sunlight were produced, it seemed very likely, by absorption of the light of the glowing body of the sun by the gases of its relatively cool atmosphere.

The vapors in the sun’s atmosphere absorbed light, and from the position of the resulting dark lines in the spectrum one could tell what elements were present in the sun’s atmosphere. The spectroscope was used to show that the sun (and the stars) was made up of elements identical with those on the earth. This conclusion finally exploded Aristotle’s belief that the heavenly bodies consisted of substances distinct in nature from those making up the earth. The spectroscope offered a new and powerful method for detecting new elements.

If a mineral brought to incandescence should reveal spectral lines belonging to no known element, it seemed reasonable to suppose that an unknown element was involved. Bunsen and Kirchhoff proved this supposition handily when, in 1860, they tested a mineral with strange spectral lines and began to search it for a new element. They found the element and proved it to be an alkali metal, related in properties to sodium and potassium. They named it cesium, from a Lain word meaning “sky blue”, for the color of the most prominent line in its spectrum.

In 1861, they repeated their triumph by discovering still another alkali metal, rubidium, from a Latin word for red, again from the color of a spectral line. Other chemists began to make use of this new tool. One of them was the French chemist Paul Emile Lecoq de Boisbaudran (1838-1912), who spent fifteen years studying the minerals of his native Pyrenees by means of the spectroscope. In 1875, he tracked down some unknown lines and found a new element in zinc ore. He named it gallium, for Gaul (France).

Sometimes afterwards, he prepared enough of the new element to study its properties. Mendeleev read Lecoq de Boisbaudran’s report and at once pointed out that the new element was none other than his own eka-aluminum. Further investigation made the identification certain; Mendeleev’s prediction of the properties of eka-aluminum matched those of gallium in every respect. The other two elements predicted by Mendeleev were found by older techniques. In 1879, a Swedish chemist, Lars Fredrick Nilson (1840-1899), discovered a new element he called scandium (for Scandinavia).

When its properties were reported, one of Nilson’s colleagues, the Swedish chemist Per Theodor Cleve (1840-1905), at once pointed out its similarity to Mendeleev’s description of eka-boron. Finally, in 1886, a German chemist, Clemens Alexander Winkler (1838-1904), analyzing a silver ore, found that all the known elements it contained amounted to only 93 per cent of its weight. Tracking down the remaining 7 per cent, he found a new element he called germanium (for Germany). This turned out to be Mendeleev’s eka-silicon.

Thus, within fifteen years of Mendeleev’s description of three missing elements, all three had been discovered and found to match his descriptions with amazing closeness. No one could doubt thereafter the validity or usefulness of the periodic table. New Elements by Groups Mendeleev’s system had to withstand the impact of the discovery of still additional new elements, for which room might, or might not, be found in the periodic table. As far back as 1794 a Finnish chemist, Johan Gadolin (1760-1852), had discovered a new metallic oxide (or earth) in a mineral obtained from the Ytterby quarry near Stockholm, Sweden.

Because the new earth was much less common than such other earths as silica, lime, and magnesia, it was referred to as a rare earth. Gadolin named his oxide yttria after the quarry; fifty years later, it yielded the element yttrium. The rare earth minerals were analyzed during the mid-nineteenth century and were found to contain an entire group of new elements, the rare earth elements. The Swedish chemist Carl Gustav Mosander (1797-1858) discovered no fewer than four rare earth elements in the late 1830’s and early 1840’s. These were lanthanum, erbium, terbium, and didymium.

Actually, five were involved, for forty years later, in 1885, the Austrian chemist Carl Auer, Baron von Welsbach (1858-1929), found that didymium was a mixture of two elements, which he called praseodymium and neodymium. Lecoq de Boisbaudran discovered two others, samarium, in 1879, and dysprosium, in 1886. Cleve also discovered two: holmium and thulium, both in 1879. By 1907 when a French chemist, Georges Urbain (1872-1938), discovered the rare earth element lutetium, fourteen such elements in all had been discovered. The rare earths possessed very similar chemical properties, and all had a valence of 3.

One might suppose this meant they would all fall into a single column of the periodic table. Such an ordering was impossible. No column was long enough to hold fourteen elements. Besides, the fourteen rare earth elements had a very closely spaced set of atomic weights. On the basis of the atomic weights they all had to be placed in a single horizontal row – in one period, in other words. Room could be made for them in the sixth period provided that period were assumed to be longer than the fourth and fifth periods, just as those were longer than the second and third.

The similarity in properties of the rare earth elements went unexplained until the 1920’s. Until then, the lack of explanation cast a shadow over the periodic table. Another group of elements whose existence was completely unsuspected in Mendeleev’s time caused no such trouble. Indeed, they fit into the periodic table quite well. Knowledge concerning them began with the work of the English physicist John William Strutt, Lord Rayleigh (1842-1919), who, in the 1880’s, was working out with great care the atomic weights of oxygen, hydrogen, and nitrogen.

In the case of nitrogen he found that the atomic weight varied according to the source of the gas. Nitrogen from the air seemed to have a slightly higher atomic weight than nitrogen from chemicals in the soil. A Scottish chemist, William Ramsay (1852-1916), grew interested in this problem and recalled that Cavendish, in a long-neglected experiment, had tried to combine the nitrogen of the air with oxygen. He had found that a final bubble of gas was left over which could not be made to combine with oxygen in any circumstances.

That final bubble, then, could not have been nitrogen. Could it be that nitrogen, as ordinarily extracted from air, contained another gas, slightly denser then nitrogen, as an impurity, and that it was that gas which made nitrogen from air seem a little heavier than it ought to be? In 1894, Ramsay repeated Cavendish’s experiment and then applied an analytical instrument Cavendish had not possessed. Ramsay heated the final bubble of gas which would not react and studied the bright line of its spectrum. The strongest lines were in positions that fitted those of no known element.

The final bubble was a new gas denser than nitrogen and making up about 1 per cent of the volume of the atmosphere. It was chemically inert and could not be made to react with any other element, so it was named argon, from a Greek word meaning “inert”. Argon proved to have an atomic weight of just under 40. This meant that it would have to fit into the periodic table somewhere in the region of the following elements: sulfur (atomic weight 32), chlorine (atomic weight 35. 5), potassium (atomic weight 39), and calcium (atomic weight, just over 40).

If the atomic weight of argon were the only thing to be considered, the new element would have to go between potassium and calcium. However, Mendeleev had established the principle that valence was more important than atomic weight. Since argon combined with no element, it could be said to have a valence of 0. How did that fit? The valence of sulfur is 2, that of chlorine 1, that of potassium 1, and that of calcium 2. The progression of valence in that region of the periodic table is 2,1,1,2. A valence of 0 would fit neatly between the two 1’s: 2,1,0,1,2.

Therefore argon was placed between chlorine and potassium. However, if the periodic table was to be accepted as a guide, argon could not exist alone. It had to be one of a family of inert gases, each with a valence of 0. Such a family would fit neatly between the column containing the halogens (chlorine, bromine, iodine, etc. ) and that containing the alkali metals (sodium, potassium, etc. ), each with a valence of 1. Ramsay began the search. In 1895, he learned that in the United States samples of gas (that had been taken from nitrogen) had been obtained from a uranium mineral.

Ramsay repeated the work and found that the gas, when tested spectroscopic ally, showed lines that belonged neither to nitrogen nor argon. Instead, most astonishingly, they were the lines that had been observed in the solar spectrum by the French astronomer Pierre Jules Cesar Janssen (1824-1907) during a solar eclipse in 1868. At the time, the English astronomer Joseph Norman Lockyer (1836-1920) had attributed them to a new element which he had named helium, from a Greek work for sun.

On the whole, chemists had paid little attention at that time to a discovery of an unknown element in the sun based on evidence as fragile as a spectral line. But Ramsay’s work showed the same element to exist on the earth, and he retained Lockyer’s name. Helium is the lightest of the inert gases and, next to hydrogen, the element with the lowest atomic weight. In 1898, Ramsay carefully boiled liquid air, looking for samples of inert gases that he expected to bubble off first. He found three, which he named neon (“new”), krypton (“hidden”), and xenon (“stranger”).

The inert gases were at first considered mere curiosities, of interest only to the ivory-tower chemists. In researches beginning in 1910, however, the French chemist Georges Claude (1870-1960) showed that an electric current forced through certain gases such as neon produced a soft, colored light. Tubes filled with such gas could be twisted into multi-colored letters of the alphabet, words, and designs. By the 1940’s the incandescent light bulbs of new York City’s celebrated Great White Way and similar centers of festivity had been replaced with neon lights.

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