Chapter C: E. Dutton, Critical observations on theories of the earth’s (5)

The best analytical work done in this country in the early part of our period was chiefly in connection with mineral analysis, and a large share of it was published in the Journal. Henry Seybert, of Philadelphia, in particular, showed remarkable skill in this direction, and published numerous analyses of silicates and other minerals, beginning in 1822. It was he who first detected boric acid in tourmaline (=6=, 155, 1822), and beryllium in chrysoberyl (=8=, 105, 1824). His methods for silicate analyses were very similar to those used at the present time.

J. Lawrence Smith in 1853 described his method for determining alkalies in minerals (=16=, 53), a method which in its final form (=1=, 269, 1871) is the best ever devised for the purpose. He also described (=15=, 94, 1853) a very useful method, still largely used in analytical work, for destroying ammonium salts by means of aqua regia. Carey Lea (=42=, 109, 1866) described the well-known test for iodides by means of potassium dichromate. F. W. Clarke (=49=, 48, 1870) showed that antimony and arsenic could be quantitatively separated from tin by the precipitation of the sulphides in the presence of oxalic acid. In 1864 Wolcott Gibbs (=37=, 346) began an important series of analytical notes from the Lawrence Scientific School, and he worked out later many difficult analytical problems, particularly in connection with his extensive researches upon the complex inorganic acids.

From 1850 on, Brush and his students made many important investigations upon minerals, and from 1877 Penfield (=13=, 425), beginning with an analysis of a new mineral from Branchville, Connecticut, described by Brush and E. S. Dana, displayed remarkable skill and industry in this kind of work. Both of the writers of this article were fortunate in being associated with Penfield in some of his researches upon minerals and one of us began as he did with the Branchville work. It is probably fair to say that Penfield did the most accurate work in mineral analysis that has ever been accomplished, and that he was similarly successful in crystallography and other physical branches of mineralogy.

The American analytical investigations that have been mentioned were all published in the Journal, with the exception of a part of Gibbs’s work. Many other American workers at mineral analysis might be alluded to here, but only the excellent work of a number of chemists in the United States Geological Survey will be mentioned. Among these Hillebrand deserves particular praise for the extent of his investigations and for his careful researches in improving the methods of rock analysis.

To our own Professor Gooch especial praise must be accorded for the very large number of analytical methods that have been devised, or critically studied, by him and his students, and for the excellent quality of this work. The publications in the Journal from his laboratory began in 1890 (=39=, 188), and the extraordinary extent of this work is shown by the fact that the three hundredth paper from the Kent Laboratory appeared in May, 1918. These very numerous and important investigations have been of great scientific and practical value, and they have formed a striking feature of the Journal for nearly 30 years. In 1912 Gooch published his “Methods in Chemical Analysis,” a book of over 500 pages, in which the work in the Kent Chemical Laboratory up to that time was concisely presented. Among the many workers who have assisted in these investigations, P. E. Browning, W. A. Drushel, F. S. Havens, D. A. Kreider, C. A. Peters, I. K. Phelps and R. G. Van Name are particularly prominent. Besides many other useful pieces of apparatus, the perforated filtering crucible was devised by Gooch, and this has brought his name into everyday use in all chemical laboratories.

Volumetric analysis was originated by Gay-Lussac, who described a method for chlorimetry in 1824, for alkalimetry in 1828, and for the determination of silver and chlorides in 1832. Margueritte devised titrations with potassium permanganate in 1846, while Bunsen, not far from the same time, introduced the use of iodine and sulphur dioxide solutions for the purpose of determining many oxidations and reductions. We owe to Mohr some improvements in apparatus and a German text-book on the subject, while Sutton wrote an excellent English work on volumetric analysis, of which many editions have appeared.

While volumetric analysis began to be used less than one hundred years ago, its applications have been gradually extended to a very great degree, and it is not only exceedingly important in investigations in pure chemistry, but its use is especially extensive in technical laboratories where large numbers of rapid analyses are required.

Not a few volumetric methods have been devised or improved in the United States, but mention will be made here only of Cooke’s important method for the determination of ferrous iron in insoluble silicates, published in the Journal (=44=, 347, 1867); to Penfield’s method for the determination of fluorine in 1878; and to the more recent general method of titration with an iodate in strong hydrochloric acid solutions, due to L. W. Andrews, a number of applications of which have been worked out in the Sheffield Laboratory.

A considerable amount of work with gases had been done by Priestley, Scheele, Cavendish, Lavoisier, Dalton, Gay-Lussac, and others before our hundred-year period began. Cavendish, about 1780, had analyzed atmospheric air with remarkable accuracy, and had even separated the argon from it and wondered what it was, and later Gay-Lussac had shown great skill in the study of gas reactions. During our period gas analysis has been further developed by many chemists. Bunsen, in particular, brought the art to a high degree of perfection in the course of a long period beginning about 1838, the last edition of his “Methods of Gas Analysis” having been published in 1877.

Important devices for the simplification of gas analysis in order that it might be used more conveniently for technical purposes have been introduced by Orsat in France and by Winkler, Hempel and Bunte in Germany.

It appears that our countryman Morley has surpassed all others in accurate work with gases in connection with his determinations of the combining weights and volumes of hydrogen and oxygen about the year 1891. Some of his publications have appeared in the Journal (=30=, 140, 1885; =41=, 220, 1891; and others).

Electrolytic analysis, involving the deposition of metals, or sometimes of oxides, usually upon a platinum electrode, was brought into use in 1865 by Wolcott Gibbs through an article published in the Journal (=39=, 58, 1865). He there described the electrolytic precipitation of copper and of nickel by the methods still in use. The application of the process has been extended to a number of other metals, and it has been largely employed, particularly in technical analyses. Important investigations and excellent books on this subject have been the contributions of Edgar F. Smith of the University of Pennsylvania, and the useful improvement, the rotating cathode, was devised by Gooch and described in the Journal (=15=, 320, 1903).

_General Inorganic Chemistry._

_The Chemical Symbols._—It is to Berzelius that we owe our symbols for the atoms, derived usually from their Latin names, such as C for carbon, Na for sodium, Cl for chlorine, Fe for iron, Ag for silver, and Au for gold. We owe to him also the use of small figures to show the number of atoms in a formula, as in N_{2}O_{5}. This was a marked improvement over the hieroglyphic symbols proposed by Dalton, which were set down as many times as the atoms were supposed to occur in formulas, forming groups of curious appearance, but in some respects not unlike some of our modern developed formulas. The advantages of Berzelius’s symbols were their simplicity, legibility, and the fact that they could be printed without the need of special type. It is true that at a later period Berzelius used certain symbols with horizontal lines crossing them to represent double atoms, and that these made some difficulty in printing. It should be mentioned also that Berzelius at one time made an effort to simplify formulas by placing dots over other symbols to represent oxygen, and commas to represent sulphur atoms. Examples of these are:

ĊaS⃛, calcium sulphate; F̋e, iron disulphide

This form of notation was quite extensively employed for a time, especially by mineralogists, but it was entirely abandoned later.

It is interesting to notice that Dalton, who lived until 1844, to reach the age of 78, differed from other chemists in refusing to accept the letter-symbols of Berzelius. In a letter written to Graham in 1837 he said: “Berzelius’s symbols are horrifying. A young student in chemistry might as soon learn Hebrew as to make himself acquainted with them. They appear like a chaos of atoms ... and to equally perplex the adepts of science, to discourage the learner, as well as to cloud the beauty and simplicity of the atomic theory.”

This forcibly expressed opinion was apparently tinged with self-esteem, but there is no doubt that Dalton was sincere in believing that the atoms were best represented by his circular symbols, because, as is well known, he thought that all the atoms were spherical in form, and it is evident that circles give the proper picture of spherical objects. At the present time some insight as to the structure of atoms is being gained, and it appears possible that the time may come when pictures of their external appearance that are not wholly imaginary may be made.

_Changes in Formulas._—Even before the year 1826, Berzelius displayed great skill in arriving at many formulas that agree with our present ones, for example, H_{2}O for water, ZnCl_{2} for zinc chloride, N_{2}O_{5} for nitric acid (anhydride), CaO for calcium oxide, CO and CO_{2} for the oxides of carbon, and many others. But at the same period other authorities, especially Gay-Lussac in France and Gmelin in Germany, on account of a lack of appreciation for Avogadro’s principle and for other reasons, such as the use of symbols to represent combining weights rather than atoms, were using different formulas for some of these compounds, such as HO, ZnCl and NO_{5}, so that their formulas for many of the compounds of hydrogen, chlorine, nitrogen and several other elements differed from those of Berzelius. The employment of different formulas involved the use of different atomic or combining weights. For example, with the formula H_{2}O for water the composition by weight requires the ratio 1 to 16 for the weights of the hydrogen and oxygen atoms, while with HO the ratio is 1 to 8.

Berzelius attempted to bring about greater uniformity in formulas and atomic weights by making changes in his table of atomic weights published in 1826. He practically doubled the relative atomic weights of hydrogen, chlorine, nitrogen, and of the other elements that gave twice as many atoms in his formulas as in those of others, and at the same time he wrote the symbols of these elements with a bar across them to indicate that they represented double atoms. For example, he wrote:

H̶O ZnC̶l N̶O_{5},

instead of

H_{2}O, ZnCl_{2} N_{2}O_{5}

This appears to have been an unfortunate concession to the views of others on the part of Berzelius, for the barred symbols were not generally adopted, partly on account of difficulties in printing, and the great achievement in theory made by him was lost sight of for a long period of time.

_The Law of Atomic Heats._—In 1819, Dulong and Petit of France, from experiments upon the specific heats of a number of solid elementary substances, came to the conclusion that the atoms of simple substances have equal capacities for heat, or in other words, that the specific heats of elements multiplied by their atomic weights give a constant called the atomic heat. For instance, the specific heats of sulphur, iron, and gold have been given as 0·2026, 0·110, and 0·0324, while their atomic weights are about 32, 56, and 197, respectively; hence the atomic heats obtained by multiplication are 6·483, 6·116, and 6·383.

Further investigations showed that the atomic heats display a considerable variation. Those of carbon, boron, beryllium, and silicon are very low at ordinary temperatures, although they increase and approach the usual values at higher temperatures. More recent work has shown, however, that the specific heats of other elements vary greatly with the temperature, almost disappearing at the temperature of liquid hydrogen, and hence possibly disappearing entirely at the absolute zero, where the electrical resistance of the metals appears to vanish likewise.

It has been found that most of the solid elements near ordinary temperatures give atomic heats that are approximately 6·4. Berzelius applied the law in fixing a number of atomic weights, and its importance for this purpose is still recognized.

It may be mentioned here that two well-known Yale men, W. O. Mixter and E. S. Dana, while students in Bunsen’s laboratory at Heidelberg in 1873, made determinations of the specific heats of boron, silicon, and zirconium. This was the first determination of this constant for zirconium, and it was consequently important in establishing the atomic weight of that element.

_Isomorphism and Polymorphism._—Mitscherlich observed in 1818 that certain phosphates and arsenates have the same crystalline form, and afterwards he reached the conclusion that identity in form indicates similarity in composition in connection with the number of atoms and their arrangement. This law of isomorphism was of much assistance in the establishment of correct formulas and consequently of atomic weights. For instance, since the carbonates of barium, strontium, and lead crystallize in the same form, the oxides of these metals must have analogous formulas. From such considerations Berzelius was able to make several improvements in his atomic weight table of 1826.

Mitscherlich was the first to observe two forms of sulphur crystals, and from this and other cases of dimorphism or of polymorphism it became evident that analogous compounds were not necessarily always isomorphous, a circumstance which has restricted the application of the law to some extent.

Besides its application in fixing analogous formulas, the law of isomorphism has come to be of much practical use in the understanding and simplification of the formulas for minerals, for these natural crystals very often contain several isomorphous compounds in varying proportions, and an understanding of this “isomorphous replacement,” as it is called, makes it possible to deduce simple general formulas for them.

In some cases isomorphism takes place to a greater or less extent between substances which are not chemically similar, and this brings about a variation in composition which at times has caused confusion. For instance, the mineral pyrrhotite has a composition which usually varies between Fe_{7}S_{8} and Fe_{11}S_{12}, and both these formulas have been assigned to it. It was recently shown by Allen, Crenshaw and Johnston in the Journal (=33=, 169, 1912) that this is a case where the compound FeS is capable of taking up various amounts of sulphur isomorphously.

The idea of solid solution was advanced by van’t Hoff to explain the crystallization of mixtures, including cases of evident isomorphism. This view has been widely accepted, and it has been particularly useful in cases where isomorphism is not evident. Solid solution between metals has been found to be exceedingly common, many alloys being of this character. A case of this kind was observed by Cooke and described in the Journal (=20=, 222, 1855). He prepared two well-crystallized compounds of zinc and antimony to which he gave the formulas Zn_{3}Sb and Zn_{2}Sb, but he observed that excellent crystals of each could be obtained which varied largely in composition from these formulas. As the two compounds were dissimilar in their formulas and crystalline forms, Cooke assumed that isomorphism was impossible and concluded “that it is due to an actual perturbation of the law of definite proportions, produced by the influence of mass.” We should now regard this as a case of solid solution.

_A Lack of Confidence in Avogadro’s Principle._—One reason why chemists were so slow in arriving at the correct atomic weights and formulas was a partial loss of confidence in Avogadro’s principle. About 1826 the young French chemist Dumas devised an excellent method for the determination of vapor densities at high temperatures, and his results and those of others showed some discrepancies in the expected densities. For example, the vapor density of sulphur was found to be about three times too great, that of phosphorus twice too great, that of mercury vapor and that of ammonium chloride only about half large enough to correspond to the values expected from analogy and other considerations. Thus, one volume of oxygen with two volumes of hydrogen make two volumes of steam, but only one third of a volume of sulphur vapor was found to unite with two volumes of hydrogen to make two volumes of hydrogen sulphide. Berzelius saw clearly that the results pointed to the existence of such molecules as S_{6}, P_{4}, and Hg_{1}, but it was not generally realized in those days that Avogadro’s rule is fundamentally reliable, and Berzelius himself appears to have lost confidence in it on account of these complications, for he did not apply Avogadro’s principle to decisions about atomic weights, except in the cases of substances gaseous at ordinary temperatures.

_Electro-chemical Theories._—The observation was made by Nicholson and Carlisle in 1800 that water was decomposed into its constituent gases by the electric current. Then in 1803 Berzelius and Hisinger found that salts were decomposed into their bases and acids by the same agency, and in 1807 Davy isolated potassium, sodium, and other metals afterwards, by a similar decomposition. Since those early times a vast amount of attention has been paid to the relation of electricity to chemical changes, a relation that is evidently of great importance from the fact that while electric currents decompose chemical compounds, these currents, on the other hand, are produced by chemical reactions.

Berzelius was particularly prominent in this direction, and in 1819 he published an elaborate electro-chemical theory. He believed that atoms were electrically polarized, and that this was the cause of their combination with one another. He extended this idea to groups of atoms, particularly to oxides, and regarded these groups as positive or negative, according to the excess of positive or negative electricity derived from their constituent atoms and remaining free. He thus arrived at his dualistic theory of chemical compounds, which attained great prominence and prevailed for a long time in chemical theory. According to this idea, each compound was supposed to be made up of a positive and a negative atom or group of atoms. For example, the formulas for potassium nitrate, calcium carbonate, and sulphuric acid corresponded to K_{2}O.N_{2}O_{5}, CaO.CO_{2} and H_{2}O.SO_{3} where we now write KNO_{3}, CaCO_{3} and H_{2}SO_{4}, and the theory was extended to embrace organic compounds also.

The eminent English chemist and physicist Faraday announced the important law of electro-chemical equivalents in 1834. This law shows that the quantities of elements set free by the passage of a given quantity of electricity through their solutions correspond to the chemical equivalents of those elements. Faraday made a table of the equivalents of a number of elements, regarding them important in connection with atomic weights, but at that time no sharp distinction was usually made between equivalents and atomic weights, and it was not fully realized that one atom of a given element may be the electrical equivalent of several atoms of another.

Faraday’s law, which is still regarded as fundamentally exact, has been of much practical use in the measurement of electric currents and in calculations connected with electro-chemical processes. In discussing his experiments, Faraday made use of several new terms, such as “electrolyte” for a substance which conducts electricity when in solution, and is thus “electrolyzed,” “electrode,” “anode,” and “cathode,” terms that have come into general use, and finally “ions” for the particles that were supposed to “wander” towards the electrodes to be set free there.

This term “ion” remained in comparative obscurity for more than half a century, when it was brought into great prominence among chemists by Arrhenius in connection with the ionic theory.

_Cannizzaro’s Ideas._—Up to about 1869 chaos reigned among the formulas used by different chemists. Various compound radicals and numerous type-formulas were employed, dualistic and unitary formulas of several kinds were in use, but the worst feature of the situation was the fact that more than one system of atomic weights was in vogue, so that water might be written

HO, H̶O, or H_{2}0

and similar discrepancies might appear in nearly all formulas containing elements of different valencies. In 1858, however, an article by the Italian chemist Cannizzaro appeared in which the outlines of a course in chemical philosophy were presented. This acquired wide circulation in the form of a pamphlet at a chemical convention somewhat later, and it dealt so clearly and ably with Avogadro’s principle, Dulong and Petit’s law, and other points in connection with formulas that it led to a rapid and almost universal reform among those who were using unsatisfactory formulas.

At about this time also the dualistic formulas of Berzelius were generally abandoned, and hydrogen came to be regarded as the characteristic element of all acids. For instance, CaO.SO_{3}, called “sulphate of lime,” came to be written CaSO_{4} and was called “calcium sulphate,” and while it had been shown as early as 1815 by Davy that “iodic acid,” I_{2}O_{5}, showed no acid reaction until it was combined with water, the accumulation of similar facts led to the formulation of sulphuric acid as H_{2}SO_{4} instead of SO_{3} or H_{2}O.SO_{3}, and that of other “oxygen acids” in a similar way. As a necessary consequence of this view of acids, the bases came to be regarded as compounds of the “hydroxyl” group, OH. Therefore the formula for caustic soda came to be written NaOH instead of Na_{2}O.H_{2}O, and so on.

_The Periodic System of the Elements._—The periodicity of the elements in connection with their atomic weights was roughly grasped by Newlands in England, who announced his “law of octaves” in 1863. This was at the time when the atomic weights were being modified and their numerical relations properly shown. The subject was worked out more fully by L. Meyer in Germany a little later, but it was most clearly and elaborately presented by the Russian chemist Mendeléeff in 1869.

In order that this subject may be explained to some extent Mendeléeff’s table is given here, with the addition of the recently discovered elements and some other modifications.

┌─────────┬────────────────┬───────────────────┬───────────────────┐ │ Groups │ I │ II │ III │ │ „ │ A B │ A B │ A B │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ Typical │ R_{2}O │ RO │ R_{2}O_{3} │ │Compounds│ │ │ │ │ „ │ RCl │ RCl_{2} │ RCl_{3} │ │ „ │ RH — │(RH_{2}) — │ — (RH_{3}) │ ╞═════════╪════════════════╪═══════════════════╪═══════════════════╡ │Series 1 │ │ │ │ │ │ │ │ │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 2 │ Lithium │ Beryllium│ _Boron_ │ │ │ 6·94 │ 9·1 │ 11·0 │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 3 │ Sodium │ Magnesium│ Aluminium│ │ │ 23·00 │ 24·32 │ 27·1 │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 4 │Potassium │ Calcium │Scandium │ │ │ 39·10 │ 40·07 │ 44·1 │ │ „ │ „ │ „ │ „ │ │ │ │ │ │ │ „ │ „ │ „ │ „ │ │ │ │ │ │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 5 │ Copper│ Zinc │ Gallium │ │ │ 53·57 │ 65·37 │ 69·9 │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 6 │Rubidium │Strontium │ Yttrium │ │ │ 85·43 │ 87·63 │ 89·0 │ │ „ │ „ │ „ │ „ │ │ │ │ │ │ │ „ │ „ │ „ │ „ │ │ │ │ │ │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 7 │ Silver│ Cadmium │ Indium │ │ │ 107·88│ 112·40 │ 114·8 │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ │ │ │Lanthanum │ │ 8 │ Cæsium │ Barium │139·0 to* │ │ │ 132·81 │ 137·87 │Lutecium │ │ │ │ │ 174.0 │ │ „ │ „ │ „ │ „ │ │ │ │ │ │ │ „ │ „ │ „ │ „ │ │ │ │ │ │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 9 │ Gold │ Mercury │ Thallium │ │ │ 197·2 │ 200·6 │ 204·0 │ ├─────────┼────────────────┼───────────────────┼───────────────────┤ │ 10 │ —— │ Radium │ —— │ │ │ │ 226·4 │ │ └─────────┴────────────────┴───────────────────┴───────────────────┘

┌─────────┬────────────────────┬─────────────────────┐ │ Groups │ IV │ V │ │ „ │ A B │ A B │ ├─────────┼────────────────────┼─────────────────────┤ │ Typical │ RO_{2} │ R_{2}O_{5} │ │Compounds│ │ │ │ „ │ RCl_{4} │ RCl_{3} │ │ „ │ — (RH_{4}) │ — RH_{3} │ ╞═════════╪════════════════════╪═════════════════════╡ │Series 1 │ │ │ │ │ │ │ ├─────────┼────────────────────┼─────────────────────┤ │ 2 │ _Carbon_ │ NITROGEN │ │ │ 12·00 │ 14·01 │ ├─────────┼────────────────────┼─────────────────────┤ │ 3 │ _Silicon_│ _Phosphorus_│ │ │ 28·3 │ 31·04 │ ├─────────┼────────────────────┼─────────────────────┤ │ 4 │ Titanium │Vanadium │ │ │ 48·1 │ 51·0 │ │ „ │ „ │ „ │ │ │ │ │ │ „ │ „ │ „ │ │ │ │ │ ├─────────┼────────────────────┼─────────────────────┤ │ 5 │ Germanium│ Arsenic │ │ │ 72·5 │ 74·96 │ ├─────────┼────────────────────┼─────────────────────┤ │ 6 │Zirconium │Niobium │ │ │ 90·6 │ 93·5 │ │ „ │ „ │ „ │ │ │ │ │ │ „ │ „ │ „ │ │ │ │ │ ├─────────┼────────────────────┼─────────────────────┤ │ 7 │ Tin 119·0│ Antimony │ │ │ │ 120·2 │ ├─────────┼────────────────────┼─────────────────────┤ │ │ (Cerium) │ │ │ 8 │ 140·25 │Tantalum │ │ │(Lutecium) │ 181·5 │ │ │ 174.0 │ │ │ „ │ „ │ „ │ │ │ │ │ │ „ │ „ │ „ │ │ │ │ │ ├─────────┼────────────────────┼─────────────────────┤ │ 9 │ Lead │ Bismuth │ │ │ 207·10 │ 208·0 │ ├─────────┼────────────────────┼─────────────────────┤ │ 10 │ Thorium │ —— │ │ │ 292·4 │ │ └─────────┴────────────────────┴─────────────────────┘

┌─────────┬─────────────────────┬───────────────────┬─────────────────┐ │ Groups │ VI │ VII │ VIII │ │ „ │ A B │ A B │ A B │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ Typical │ RO_{3} │ R_{2}O_{7} │(RO_{4}) — │ │Compounds│ │ │ │ │ „ │ RCl_{2} │ — RCl │ — R │ │ „ │ — RH_{2} │ — RH │ — R │ ╞═════════╪═════════════════════╪═══════════════════╪═════════════════╡ │Series 1 │ │ HYDROGEN │ HELIUM │ │ │ │ 1·008 │ 3·99 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 2 │ OXYGEN │ FLUORINE │ NEON │ │ │ 16·00 │ 19·0 │ 20·2 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 3 │ _Sulphur_ │ CHLORINE │ ARGON │ │ │ 32·07 │ 35·46 │ 39·88 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 4 │ Chromium │Manganese │ Iron │ │ │ 52·0 │ 54·93 │ 55·84 │ │ „ │ „ │ „ │ Cobalt │ │ │ │ │ 58·97 │ │ „ │ „ │ „ │ Nickel │ │ │ │ │ 58·68 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 5 │ _Selenium_│ _Bromine_│ KRYPTON│ │ │ 79·2 │ 79·92 │ 82·92 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 6 │Molybdenum │ —— 100 │Ruthenium │ │ │ 96·0 │ │ 101·7 │ │ „ │ „ │ „ │ Rhodium │ │ │ │ │ 102·9 │ │ „ │ „ │ „ │Palladium │ │ │ │ │ 106·7 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 7 │ Tellurium │ _Iodine_ │ XENON │ │ │ 127·5 │ 126·92 │ 130·2 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ │ │ │ │ │ 8 │ Tungsten │ —— 188 │ Osmium │ │ │ 184·0 │ │ 190·9 │ │ │ │ │ │ │ „ │ „ │ „ │ Iridium │ │ │ │ │ 193·1 │ │ „ │ „ │ „ │Platinum │ │ │ │ │ 195·2 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 9 │ —— │ —— │ NITON │ │ │ │ │ 222·4 │ ├─────────┼─────────────────────┼───────────────────┼─────────────────┤ │ 10 │ Uranium │ │ │ │ │ 238·5 │ │ │ └─────────┴─────────────────────┴───────────────────┴─────────────────┘

┌────────────────────────────────────────────────────────────────┐ │ * Lanthanum, Cerium, Praseodymium, Neodymium, ——, │ │Rare·Earth 139·0 140·25 140·6 144·3 │ │ Metals: │ │ „ Gadolinium, Terbium, Dysprosium, Holmium, Erbium,│ │ 157·3 159·2 162·5 163·5 167·7 │ └────────────────────────────────────────────────────────────────┘

┌─────────────────────────────────────────┐ │ * Samarium, Europeum, ——, │ │Rare·Earth 150·4 152·0 │ │ Metals: │ │ „ Thulium, Ytterbium, Lutecium,│ │ 168·5 172·0 174·0 │ └─────────────────────────────────────────┘

NOTE.—Distinctions in printing: GASEOUS ELEMENTS. _Other non-metallic
elements_, metallic elements. The heavy line encloses approximately
the acid-forming elements.

In this table the elements arranged in the order of their atomic weights fall into eight groups where the known oxides progress regularly, with the exception of two or three elements, from R_{2}O in Group I to R_{2}O_{7} in Group VII, while in Group VIII two oxides (of ruthenium and osmium) are known which carry the progression to RO_{4}.

It was pointed out by Mendeléeff that, with the exception of series 1 and 2 at the top of the table, the alternate members of the groups show particularly close relationships. These subordinate groups, marked A and B, in most cases show remarkable analogies and gradations in their properties, for example, in the alkali-metals from lithium to cæsium, and in the halogens from fluorine to iodine. The two divisions of a group do not usually show very close relations to each other, except in their valency, and they even display, in several instances, opposite gradations in chemical activity in the order of their atomic weights. For instance, cæsium stands at the electro-positive end, while gold stands at the electro-negative end of its subordinate group. The difference between the two divisions is very great in Groups VI and VII, but it is extreme in Group VIII, where heavy metals are on one side and inactive gases on the other. Many authorities separate these gases into a “Group O” by themselves at the left-hand side of the table, but this does not change their relative positions, and the plan may be objected to on the ground that many vacant places are thus left in the groups VIII and O.

The periodic law has been useful in rectifying certain atomic weights. At the outset Mendeléeff was obliged to change beryllium from 14·5 (assuming Be_{2}O_{3}) to 9 (assuming BeO), and later the atomic weights of indium and uranium were changed to make them fit the system. All of these changes have been confirmed by physical means.

Mendeléeff found a number of vacant places in his table, and was thus able to render further service to chemical science by predicting the properties of undiscovered elements, and his predictions were very closely confirmed by the later discovery of scandium, gallium, and germanium. The table indicates that there are still two undiscovered elements below manganese and probably two more among the rare-earth metals. The interesting observation has just recently been made by Soddy that the products of radioactive disintegration appear to pass in a symmetrical way through positions in the periodic system, giving off a helium molecule at alternate transformations until the place of lead is reached. It appears, therefore, that the five vacant places in the table above bismuth are probably occupied by these evanescent elements, and it is to be noticed that all of the elements that have been placed in this region of high atomic weights are radioactive.

There are some inconsistencies in the periodic system. The increments in the atomic weights are irregular, and there are three cases, argon and potassium, cobalt and nickel, and tellurium and iodine, where a higher atomic weight is placed before a lower one in order to bring these elements into their undoubtedly proper places. There is a peculiarity also in the heavy-metal division of Group VIII, where three similar elements occur in each of three places, and where the usual periodicity appears to be suspended, or nearly so, in comparison with most of the other elements. However, there seems to be a still more remarkable case of this kind in Group III, where fourteen metals of the rare earths have been placed. They are astonishingly similar in their chemical properties, hence it seems necessary to assume that periodicity is suspended here throughout the wide range of atomic weights from 139 to 174, where no elements save these have been found.

Several other interesting features of the table may be pointed out. The chlorides and hydrides, as indicated by the “typical compounds,” show a regular progression in both directions towards Group IV. (Where the type-formulas do not apply, as far as is known, to more than one or two elements, they have been placed in parentheses in the table given here.) It is a striking fact that the acid-forming elements occur together in a definite part of the table, and that the gases and other non-metallic elements, except the inactive gases of Group VIII, occur in the same region.

_Atomic Numbers._—As the result of a spectroscopic study of the wave lengths or frequencies of the X-rays produced when cathode rays strike upon anticathodes composed of different elements, Moseley in 1914 discovered that whole numbers in a simple series can be attributed to the atoms. These atomic numbers are: 1 for hydrogen, 2 for helium, 3 for lithium, 4 for beryllium, and so on, in the order in which the elements occur in Mendeléeff’s periodic table, and in the cases of argon and potassium, cobalt and nickel, and tellurium and iodine, they follow the correct chemical order, while the atomic weights do not. They appear to indicate, therefore, an even more fundamental relation between the atoms than that shown by the atomic weights.

These numbers are now available for every element up to lead, and they are particularly interesting in indicating, on account of missing numbers, the existence of two undiscovered elements in the manganese group, and two more among the rare-earth metals, in confirmation of the vacant places below lead in Mendeléeff’s table.

_The Isolation of Elements._—In the year 1818 about 53 elements were recognized, and since that time about 30 more have been discovered, but the elements already known comprised the more common ones, and nearly all of those which have been commercially important. A few of them, including beryllium, aluminium, silicon, magnesium, and fluorine, were then known only in their compounds, as they had not yet been isolated in the free condition.

Berzelius in 1823 prepared silicon, a non-metallic element resembling carbon in many respects. This element has recently been prepared on a rather large scale in electric furnaces at Niagara Falls, and has been used for certain purposes in the form of castings.

Wöhler created much sensation in 1827 by isolating aluminium and finding it to be a very light, strong and malleable metal, stable in the air, and of a silver-white color. For a long time this metal was a comparative rarity, being prepared by the reduction of aluminium chloride with metallic sodium; but about 25 years ago Hall, an American, devised a method of preparing it by electrolyzing aluminium oxide dissolved in fused cryolite. This process reduced the cost of aluminium to such an extent that it has now come into common use.

Wöhler and Bussy prepared beryllium in 1828, and Liebig and Bussy did the same service for magnesium in 1830. The latter metal has come to be of much practical importance, both as a very powerful reducing agent in chemical operations, and as an ingredient of flash-light powders and of mixtures used for fireworks. It is also used in making certain light alloys.

After almost innumerable attempts to isolate fluorine, during a period of nearly a century, this was finally accomplished in 1886 by Moissan in France by the electrolysis of anhydrous hydrogen fluoride. The free fluorine proved to be a gas of extraordinary chemical activity, decomposing water at once with the formation of hydrogen fluoride and ozonized oxygen. This fact explains the failure of many previous attempts to prepare it in the presence of water.

_Early Discoveries of New Elements._—The remarkable activity of chemical research at the beginning of our period is illustrated by the fact that three new elements were discovered in 1817. In that year Berzelius had discovered selenium, Arfvedson, working in Berzelius’s laboratory had discovered the important alkali-metal lithium, and Stromeyer had discovered cadmium.

In 1826 Ballard in France discovered bromine in the mother-liquor from the crystallization of common salt from sea water. Bromine proved to be an unusually interesting element, being the only non-metallic one that is liquid at ordinary temperatures, and being strikingly intermediate in its properties between chlorine and iodine. It has been obtained in large quantities from brines, and is produced extensively in the United States. The elementary substance and its compounds have found important applications in chemical operations, while the bromides have been found valuable in medicine and silver bromide is very extensively used in photography.

In 1828 Berzelius discovered thorium. The oxide of this metal has recently been employed extensively as the principal constituent of incandescent gas-mantles, and the element has acquired particular importance from the fact that, like uranium, it is radioactive, decomposing spontaneously into other elements.

Vanadium had been encountered as early as 1801 by Del Rio, who named it “erythronium,” but a little later it was thought to be identical with chromium and was lost sight of for a while. In 1830, however, it was re-discovered by, and received its present name from Sefström in Sweden. Berzelius immediately made an extensive study of vanadium compounds, but he gave them incorrect formulas and derived an incorrect atomic weight for the element, because he mistook a lower oxide for the element itself. Roscoe in England in 1867 isolated vanadium for the first time, found the right atomic weight, and gave correct formulas to its compounds. Vanadium is particularly interesting from the fact that it displays several valencies in its compounds, many of which are highly colored. It has found important use as an ingredient in very small proportions in certain “special steels” to which it imparts a high degree of resistance to rupture by repeated shocks.

Columbium was discovered early in the nineteenth century in the mineral columbite from Connecticut by Hatchett, an Englishman, who did not, however, obtain the pure oxide. It was afterwards obtained by Rose who named it niobium. Both names for the element are in use, but the former has priority. Attention was called to this fact by an article in the Journal by Connell, an Englishman (=18=, 392, 1854).

_The Platinum Group of Metals._—In 1854 a new member of the platinum group of metals, ruthenium, was discovered by Claus. Platinum had been discovered about the middle of the eighteenth century, while its other rarer associates, iridium, osmium, palladium, and rhodium, had been recognized in the very early years of the nineteenth century. It was during the latter period that platinum ware began to be employed to a considerable extent in chemical operations, and this use was greatly extended as time went on. The discovery was made by Phillips in 1831 that finely divided platinum by contact would bring about the combination of sulphur dioxide with atmospheric oxygen, and this application during the past 20 years has become enormously important in the sulphuric acid industry, while other important applications of platinum as a “catalytic agent” have also been made. Wolcott Gibbs and Carey Lea have contributed perhaps more than any other recent chemists to a knowledge of the platinum metals. Carey Lea (=38=, 81, 248, 1864) dealt chiefly with the separation of the metals from each other, while Gibbs’s work (=31=, 63, 1861; =34=, 341, 1862) included investigations of many of the compounds.

It may be mentioned that while platinum and its associates were formerly known only in the uncombined condition in nature, the arsenide sperrylite, PtAs_{2}, was described by the late S. L. Penfield, and the senior writer of this chapter, in articles published in the Journal (=37=, 67, 71, 1889).

_Applications of the Spectroscope._—The discovery in certain mineral waters of the rare alkali-metals rubidium and cæsium by Bunsen and Kirchoff in 1861 was in consequence of the application of spectroscopy by these same scientists a short time previously to the identification of elements imparting colors to the flame. Since that time the employment of the spectroscope for chemical purposes has been much extended, as it has been used in the examination of light from electric sparks and arcs, as well as from Geissler tube discharges and from colored solutions.

The metals rubidium and cæsium are interesting in being closely analogous to potassium and in standing at the extreme electro-positive end of the series of known metals. It should be noticed here that Johnson and Allen of our Sheffield Laboratory, having obtained a good supply of rubidium and cæsium material from the lepidolite of Hebron, Maine, made some important researches upon these elements, accounts of which were published in the Journal (=34=, 367, 1862; =35=, 94, 1863). They established the atomic weight of cæsium, thus correcting Bunsen’s determination which was unsatisfactory on account of the small quantity and impurity of his material. Pollucite, a mineral rich in cæsium, which had been found in very small amount on the Island of Elba, has more recently been obtained in large quantities—hundreds of pounds—at Paris, Maine, and its vicinity. This American pollucite was first analyzed and identified by the senior writer of this article (=41=, 213, 1891), and later (=43=, 17, 1892 _et seq._) the results of many investigations on cæsium and rubidium compounds, in which the junior writer played an important part, carried out in Sheffield Laboratory, were published in the Journal.

The application of the spectroscope led to the discovery of thallium in 1861 by Crookes of England, and to that of indium in 1863 by Reich and Richter in Germany. Both of these metals are extremely rare, but they are of considerable theoretical interest. Thallium is particularly remarkable in showing resemblances in its different compounds to several groups of metals.

The spectroscope was employed again in connection with the discovery of gallium in 1875 by Boisbaudran. It is in the same periodic group as thallium and indium, and it has a remarkably low melting point, just above ordinary room-temperature. It has been among the rarest of the rare elements, but within two or three years a source of it has been found in the United States in certain residues from the refining of commercial zinc. The recent issues of the Journal (=41=, 351, 1916; =42=, 389, 1916) show that Browning and Uhler of Yale have availed themselves of this new material in order to make important chemical and physical researches upon this metal.

_Germanium._—The discovery of germanium in the mineral argyrodite in 1886 by Winkler revealed a curious metal which gives a white sulphide that may be easily mistaken for sulphur and which is volatilized completely when its hydrochloric acid solution is evaporated, so that it is evasive in analytical operations. This element had been predicted with much accuracy by Mendeléeff, and it is rather closely related to tin.

A few years after the discovery of germanium, Penfield published in the Journal (=46=, 107, 1893; =47=, 451, 1894) some analyses of argyrodite, correcting the formula given by Winkler to the mineral; also he described canfieldite, an analogous mineral from Bolivia, in which a large part of the germanium was replaced by tin.

_The Rare Earths._—Before the year 1818 two rare earths, the oxides of yttrium and cerium, were known in an impure condition. Since that time about fourteen others have been discovered as associates of the first two. The rare earths are peculiar from the fact that many of them are always found mixed together in the minerals containing them, and also from the circumstance that most of them are remarkably similar in their chemical reactions and consequently exceedingly difficult to separate from each other. In many cases multitudes of fractional precipitations or crystallizations are needed to obtain pure salts of a number of these metals. The solutions of the salts of several of these elements give characteristic absorption bands when examined spectroscopically by the use of transmitted light.

No important practical application has been found for any of these earthy oxides, except that about one per cent of cerium oxide is mixed with thorium oxide in incandescent gas-mantles in order to obtain greatly increased luminosity.

_The Inactive Gases._—As long ago as 1785, Cavendish, that remarkable Englishman who first weighed the world and first discovered the composition of water, actually obtained a little argon in a pure condition by sparking atmospheric nitrogen with oxygen converting it into nitric acid (another discovery of his) and absorbing the excess of oxygen. The volume of this residual gas as estimated by him corresponds very closely to the volume of argon in the atmosphere, as now known.

It was more than a century later, in 1894, that Rayleigh and Ramsay discovered argon in the air. Lord Rayleigh had found that atmospheric nitrogen was about one-half per cent heavier than chemical nitrogen, a fact which led to the investigation. It was only necessary to repeat Cavendish’s experiment on a large scale, or to absorb oxygen with hot copper and nitrogen with hot magnesium, in order to obtain argon. The gas attracted much attention, both on account of having but a single atom in its molecule, and particularly because it failed to enter into chemical combination of any kind. This gas has been used of late for filling the bulbs of incandescent electric lamps in cases where a gas pressure without chemical action is desired.

In 1890 and 1891, Hillebrand published in the Journal (=40=, 384, 1890: =42=, 390, 1891) a series of analyses of the mineral uraninite and reported in some samples of the mineral as much as 2·5 per cent of an inactive gas. Hillebrand examined the gas spectroscopically but, just missing an important discovery, he detected only the spectrum lines of nitrogen. Ramsay, in searching for argon in some sort of natural combination, and doubtless remembering Hillebrand’s work, heated some cleveite, a variety of uraninite, and obtained, not argon, but a new gas. This gave a yellow spectrum-line corresponding to a line previously observed in the light of the sun’s corona and attributed to an element in the sun called helium. Helium, therefore, in 1895 had been found on the earth. This gas is a constant constituent of uranium minerals, as it is produced by the breaking down of radioactive elements. It has been found in very small quantity in the atmosphere, and is the most difficult of all known gases to liquefy, as its boiling point, as shown by Onnes in 1908, is only 4° above the absolute zero. It has not yet been solidified.

In 1898 Ramsay and Travers, by the use of ingenious methods of fractional distillation and absorption by charcoal, obtained three other much rarer inactive gases from the atmosphere which they called neon, krypton and xenon.

The inactive gases are all colorless, and as they form no chemical compounds they are characterized by their densities, which give their atomic weights, by their boiling points, and by their characteristic Geissler-tube spectra.

The gaseous radium emanation, or niton, belongs also to the inactive group, and it was also collected and studied by Ramsay who was compelled to work with only 0·0001 cc. of it, as the volume obtained by heating radium salts is very small. It is an evanescent element, disappearing within a few days on account of radioactive disintegration. Meanwhile it glows brilliantly when liquefied and cooled to the temperature of liquid air. It has an atomic weight of 222, four units below that of radium, and the difference is considered as due to the loss by radium of an atom of helium in passing into the emanation.

_The Radioactive Elements._—The discovery of radium in 1898 by Madame Curie, and the study of that and other radioactive elements has produced a profound effect upon chemical theory. It was found that the two elements of the highest atomic weights, uranium and thorium, are always spontaneously decomposing into other elements at a fixed rate of speed which can be controlled by no artificial means, and that the elements resulting from these decompositions likewise undergo spontaneous changes into still other elements at greatly varying rates of speed, forming in each case a remarkable series of temporary elements. These transformations are accompanied by the emission at enormous velocities of three kinds of rays, one variety of which has been shown to consist of helium atoms. The greater number of the elements formed in these transformations have not as yet been obtained in a pure condition, and they are known only in connection with their radioactivity, volatility, etc.; but radium and niton, two of these products, have been obtained in a pure condition, so that their atomic weights and their places in the periodic system have been fixed.

We owe much of our knowledge of the radioactive transformations to the researches of Rutherford and of Soddy, and of their co-workers, but one of the important products of the transformation of uranium, an element which he called ionium, was characterized by Boltwood of Yale (=25=, 365, 1908).

Radium and niton, apart from their radioactive properties, resemble barium and the inert gases of the atmosphere, respectively. The rates at which their progenitors produce them, and the rates at which they themselves decompose, bring about a state of equilibrium after a time. Therefore a given amount of uranium, which decomposes exceedingly slowly, can yield even after thousands of years only a very small proportional quantity of undecomposed radium, one-half of which disappears in about 2500 years, because the amount decomposed must eventually be equal to the amount produced. The first conclusive evidence that radium is a product of the decomposition of uranium was given by Boltwood in the Journal (=18=, 97, 1904). He found that all uranium minerals contain radium; and the amount of radium present is always proportional to the amount of uranium, which shows the genetic relation between the two.

In the case of niton, which is produced by radium, and is called also the radium emanation, the rate of decay is rapid, so that if the gas is expelled from radium by heating, equilibrium is reached after a few days, with the accumulation of the largest possible amount of niton.

The conclusion has been reached by Rutherford and others that the final product besides helium, in the radioactive transformations, is lead, or at least an element or elements resembling lead to such a degree that no separation of them by chemical means is possible. Atomic weight determinations by Richards and others have shown that specimens of lead found in radioactive minerals give distinctly different atomic weights from that of ordinary lead. This fact has led to the view that possibly the atoms of the elements are not all of the same weight, but vary within certain limits—a view that is contrary to previous conclusions derived from the uniformity in atomic weights obtained with material from many different sources.

The results of the investigations upon radioactivity have led to modified views in regard to the stability of the elements in general. There has been little or no proof obtained that any artificial transmutation of the elements is possible, but the spontaneous transformation of the radioactive elements brings forward the possibility that other elements are changing imperceptibly, and that a state of evolution exists among them. All of the radioactive changes that we know proceed from higher to lower atomic weights, and we are entirely ignorant of the process by which uranium and thorium must have been produced originally.

Since radioactive changes have been found to be accompanied by the release of vast amounts of energy, compared with which the energy of chemical reactions is trivial, a new aspect in regard to the structure of atoms has arisen,—they must be complex in structure, the seats of enormous energy.

The determination of the amount of radium in the earth’s crust has indicated that the heat produced by it is amply sufficient to supply the loss of heat due to radiation, and this source of heat is regarded by many as the cause of volcanic action. The sun’s radiant heat also has been supposed to be supplied by radioactive action, so that the older views regarding the limitation of the age of the earth and the solar system on account of loss of heat have been considerably modified by our knowledge of radioactivity.

_Physical Chemistry._

The application of physical methods as aids to chemical science began in early times, and some of these, such as the determinations of gas and vapor densities, specific heats, and crystalline forms have been mentioned already in this article. Within recent times physical chemistry has greatly developed and a few of its important achievements will now be described.

_Molecular Weight Determinations._—Gas and vapor densities in connection with Avogadro’s principle, formed the only basis for molecular weight determinations until comparatively recent times. The early methods of Gay-Lussac and Dumas for vapor density were supplemented in 1868 by the method of Hofmann, whereby vapors were measured under diminished pressure over mercury. In 1878 Victor Meyer introduced a simpler method depending upon the displacement of air or other gas by the vapor in a heated tube. As refractory tubes, such as those of porcelain or even iridium, could be used in this method, molecular weights at extremely high temperatures were determined with interesting results. For instance, it was found that iodine vapor, which shows the molecule I_{2} at lower temperatures, gradually becomes monatomic with rise in temperature, that sulphur vapor dissociates from S_{8} to S_{2} under similar conditions, and that most of the metals, including silver, have monatomic vapors.