Pauling on periodic law

American theoretical physical chemist Linus Carl Pauling (1901–94) is the only person to have won two unshared Nobel Prizes. His first, a Nobel Prize for Chemistry (1954), was awarded for research into the nature of the chemical bond and its use in elucidating molecular structure; the second, a Nobel Prize for Peace (1962), recognized his efforts to ban the testing of nuclear weapons. Before his Nobel fame Pauling also wrote for the Encyclopædia Britannica, penning in 1948 the 14th edition’s article on “The Periodic Law,” highlighted below.

Periodic Law

It was discovered by Dmitri I. Mendeléyev in the mid-19th century that the chemical elements show a periodic recurrence of properties when they are arranged in a certain order, which is approximately the order of increasing atomic weight. The statement of this fact, which has been of inestimable value in the development of the science of chemistry, is called the periodic law, and a tabular arrangement of the elements which brings those with similar properties into juxtaposition is called the periodic table or periodic system of the elements.

It was recognized during the second decade of the 20th century that the order of elements in the periodic system is that of their atomic numbers, the integers which are equal to the positive electrical charges of the atomic nuclei expressed in electronic units; and in recent years great progress has been made in explaining the periodic law in terms of the electronic structure of atoms and molecules. This clarification has increased the value of the law, which is used as much today as it was 50 years ago, when it expressed the only known relationship among the elements.

History of the periodic law

The classification of the chemical elements is a comparatively recent problem, for it was not until about the beginning of the 19th century that the two kinds of chemical substances, elements and compounds, were clearly differentiated. The early years of the century witnessed a rapid development in analytical chemistry—the art of distinguishing different chemical substances—and the consequent building up of a vast body of knowledge of the chemical and physical properties of both elements and compounds. This rapid expansion of chemical knowledge soon necessitated classification, for on the classification of chemical knowledge are based not only the systematized literature of chemistry but also the laboratory arts by which chemistry is passed on as a living science from one generation of chemists to another. The elementary substances are few in number, about 100, with distinctive properties, whereas compound substances are extremely numerous, of the order of 1,000,000, with properties correspondingly diversified and often overlapping. It thus occurred that the classification of elements lagged many years behind that of compounds, and, in fact, no general agreement had been reached among chemists as to the classification of elements for nearly half a century after the systems of classification of compounds had become established in general use.

J.W. Döbereiner in 1817 showed that the combining weight of strontium lies midway between those of calcium and barium, and some years later he showed that other such “triads” exist (chlorine, bromine and iodine; and lithium, sodium and potassium). J.B.A. Dumas, L. Gmelin, E. Lenssen, Max von Pettenkofer and J.P. Cooke expanded Döbereiner’s suggestions between 1828 and 1854 by showing that similar relationships were more extensive than among triads of elements, fluorine being added to the halogens, and magnesium to the alkaline-earth metals, while oxygen, sulfur, selenium and tellurium were classed as one family, and nitrogen, phosphorus, arsenic, antimony and bismuth as another family of elements.

Attempts were subsequently made to show that the atomic weights of the elements could be expressed by an arithmetical function, and in 1862 A.E.B. de Chancourtois proposed a classification of the elements based on the new values of atomic weights given by Stanislao Cannizzaro’s system of 1858. De Chancourtois plotted the atomic weights on a helical curve such that corresponding points differed by 16, the atomic weight of oxygen. This curve brought closely related elements onto corresponding points, and he suggested in consequence that “the properties of the elements are the properties of numbers,” a remarkable prediction in the light of modern knowledge. In the following year, 1863, J.A.R. Newlands proposed a system of classification of the elements in the order of atomic weights, the elements being assigned ordinal numbers from unity upward and divided into seven groups having properties closely related to the first seven known elements, hydrogen, lithium, beryllium, boron, carbon, nitrogen and oxygen. This relationship was termed the law of octaves, by analogy with the seven intervals of the musical scale. Then in 1869 Mendeléyev, as a result of an extensive correlation of the properties and the atomic weights of the elements, with special attention to valency, proposed the period law, by which “the elements arranged according to the magnitude of atomic weights show a periodic change of properties.” Lothar Meyer had independently reached a similar conclusion, which he published after the appearance of Mendeléyev’s Russian paper and of an account of it in German.

Mendeléyev’s periodic table of 1869 contained 17 columns, with two nearly complete periods (sequences) of elements, from potassium to bromine and rubidium to iodine, preceded by two partial periods of seven elements each (lithium to fluorine and sodium to chlorine), and followed by three incomplete periods. In his 1871 paper Mendeléyev presented a revision of the 17-group table, the principal improvement being the correct repositioning of 17 elements. He, as well as Lothar Meyer, also proposed a table with eight columns, obtained by splitting each of the long periods into a period of seven, an eighth group containing the three central elements (such as Fe, Co, Ni; Mendeléyev also included Cu, instead of placing it in group I), and a second period of seven. The first and second periods of seven were later distinguished by use of the letters a and b attached to the group symbols, which were the Roman numerals.

With the discovery of the inert gases helium, neon, argon, etc., by Lord Rayleigh and Sir William Ramsay in 1894 and the following years, Mendeléyev and others proposed that a new “zero” group to accommodate them be added to the periodic table. The “short-period” form of the period law, with groups 0, I, II,…VIII, became popular, and remained in general use until recent years. A new table, devised in 1895 by J. Thomsen, was interpreted in terms of the electronic structure of atoms by Niels Bohr in 1922. In this table, there are periods of increasing length, each extending from one noble gas to the next noble gas; the table thus contains a period of 2 elements, two periods of 8 elements, two periods of 18 elements, a period of 32 elements and an incomplete period. The elements in each period are connected by tie-lines with corresponding elements in the following period. The principal disadvantage of this periodic table is the large space required by the period of 32 elements and the difficulty of tracing a sequence of closely similar elements along the slanting tie-lines. A very satisfactory compromise is to compress the period of 32 elements into 18 spaces by omitting the 14 rare earths. This form of the periodic table is the work of T. Bayley (1882).

The great value of the periodic law was made evident by Mendeléyev’s success in 1871 in finding that, by changing 17 elements from the positions indicated by the atomic weights which had been accepted for them into new positions, their properties could be correlated with those of other elements. This change indicated that there were small errors in the previously accepted atomic weights of several of the elements, and large errors for several others, for which wrong multiples of the combining weights had been used as atomic weights. Mendeléyev was also able to predict the existence and many of the properties of the undiscovered elements eka-boron, eka-aluminum, eka-silicon, eka-manganese, dui-manganese and eka-tantalum, now identified with the elements scandium, gallium, germanium, technetium (formerly masurium), rhenium and polonium, respectively. Similarly after the discovery of helium and argon the periodic law permitted the prediction of the existence of neon, krypton, xenon and radon. In recent years Niels Bohr pointed out that element 72 would be expected, from its position in the periodic system, to be similar to zirconium in its properties rather than to the rare earths; this observation led G. de Hevesy and D. Coster in 1922 to examine zirconium ores and to discover the missing element, which they named hafnium.

Many of the elements in the Mendeléyev and Lothar Meyer periodic tables of 1871 were required by their properties to be put in positions somewhat out of the order of atomic weights. Most of these displacements have been rectified by redeterminations of the atomic weights; three, however, remain, and are real. In the three pairs of elements argon and potassium, cobalt and nickel, and tellurium and iodine the first element of each pair has the greater atomic weight, but the earlier position in the periodic system. The solution of this difficulty had been indicated by Newlands, who had assigned the successive ordinal numbers to the elements arranged approximately according to atomic weight. After the discovery of the electron by J.J. Thomson and the development of the theory of the nuclear atom by Sir Ernest Rutherford, who had found by experiment that the electrical charge of the nuclei of heavier atoms, in units equal to the charge of the electron with changed sign, is roughly one-half the atomic weight of the atoms, it was suggested in 1911 by A. van den Broek that this quantity, the atomic number, might be equal to the ordinal number of the element in the periodic system. This suggestion was brilliantly confirmed in 1913 by H.G.J. Moseley’s measurements of the wave lengths of the characteristic X-ray spectral lines of many elements. There is no longer any uncertainty about the position of any element in the ordered series of the periodic system. All of the elements from 1 to 96 are now known, and it is to be expected that new elements beyond 96 will be made by nuclear reactions.

The small significance of the exact atomic weight of an element to its position in the periodic system is shown by the existence of isotopes of every element, atoms which have the same atomic number but different atomic weights. The chemical properties of the different isotopes of an element are essentially the same, and all of the isotopes of an element occupy the same place in the periodic system.

The detailed understanding of the periodic system has developed along with the quantum theory of spectra and the electronic structure of atoms, beginning with the work of Niels Bohr in 1913. Important forward steps were the formulation of the general rules of the old quantum theory by W. Wilson and A. Sommerfeld in 1917, the discovery of the exclusion principle by W. Pauli, Jr., in 1925, the discovery of the spin of the electron by G.E. Uhlenbeck and S. Goudsmit in 1925, and the development of quantum mechanics by W. Heisenberg and E. Schrödinger during the same year. The development of the electronic theory of valency and molecular structure, beginning with the discovery of the shared electron pair by G.N. Lewis in 1916, also played a very important part in explaining the periodic law.

Mendeléyev stated that a natural law acquires scientific importance only when it yields practical results in elucidating unexplained phenomena, in disclosing unrecognized occurrences and in evoking verifiable predictions. The periodic law has been found to be very useful indeed in the simplification of the science of chemistry through the classification of the elements, in the determination of the correct multiples of equivalent weights to be taken as the atomic weights of elements, in indicating errors in atomic weights, in the prediction of the existence and properties of new elements and especially as a basis for chemical reasoning and prediction in the planning of experimental work. Its value has increased with our increasing understanding of the law and its basis in the electronic structure of atoms, and it has now become one of the firm foundation stones of the great edifice of science.

The uranium elements

The only important general question about the periodic system to which the answer remains doubtful is that as to the nature of the last period. Because thorium is more closely similar to hafnium than to cerium in many of its properties, and uranium is more closely similar to tungsten than to neodymium, it has been thought that the uranium elements constitute the beginning of a transition period similar to the two long periods, and involving the occupancy of the 6 d subshell. However, G.T. Seaborg has pointed out that the properties of the trans-uranium elements suggest that these elements are members of a second rare-earth series, with occupancy of the 5 f orbitals beginning at thorium. It is true that uranium is unlike the rare-earth elements in that the +6 oxidation state is the most stable of its oxidation states (+6, +5, +4, +3, +2); but neptunium, instead of being similar to rhenium, is reported not to have a stable +7 state; instead, neptunium and plutonium are said to have the same oxidation states as uranium, but with a general shift in stability toward the lower oxidation states.

An answer to this question can be given by the consideration of the stability and nature of the electronic orbitals involved. Although d orbitals contribute to the formation of strong covalent bonds, inner f orbitals are not useful in this way, as is shown by the small tendency of the rare-earth elements to form covalent compounds. The fact that uranium does form covalent compounds shows that the 6 d orbitals are stable enough to be used for this purpose; and the stability of the +6 oxidation state shows that the 5 f orbitals of uranium are not sufficiently stable to attract any of the 6 valence electrons very strongly. Hence the uranium period differs from the immediately preceding period in that the 5 f orbitals are less stable relative to the 6 d orbitals than are the 4 f relative to the 5 d orbitals. However, with increase in atomic number the 5 f orbitals will increase in stability more rapidly than the 6 d orbitals, and ultimately the 5 f orbitals will be filled. It is likely that this process has begun at neptunium; and presumably it would be completed at element 103, which would be similar to lutetium in structure and properties, with probably a greater tendency to the formation of covalent compounds. The 6 d, 7 s and 7 p subshells would then be successively occupied, with element 118 as the next noble gas, though probably able to form some chemical compounds.

Thus this great generalization, the periodic law, which first showed its value through Mendeléyev’s prediction of the existence and properties of undiscovered elements, is again being put to the same use, the undiscovered elements of present interest being those which have greater atomic number than the known elements, and which will be sought not in nature’s minerals, but in preparations made by man.