The Periodic Table – History of Its Arrangement

The Periodic Table – History of Its Arrangement

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 international 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 126.9, 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 spectroscopically, 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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