Periodic Law

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INTRODUCTION Periodic Law, in chemistry, law stating that many of the physical and chemical properties of the elements tend to recur in a systematic manner with increasing atomic number. Progressing from the lightest to the heaviest atoms, certain properties of the elements approximate those of precursors at regular intervals of 2, 8, 18, and 32. For example, the 2nd element (helium) is similar in its chemical behavior to the 10th (neon), as well as to the 18th (argon), the 36th (krypton), the 54th (xenon), and the 86th (radon). The chemical family called the halogens, composed of elements 9 (fluorine), 17 (chlorine), 35 (bromine), 53 (iodine), and 85 (astatine), is an extremely reactive family.


As a result of discoveries that firmly established the atomic theory of matter in the first quarter of the 19th century, scientists could determine the relative weights of atoms of the then known elements. The development of electrochemistry during this period by the British chemists Sir Humphry Davy and Michael Faraday led to the discovery of many additional elements. By 1829 a sufficient number of elements had been discovered to permit the German chemist Johann Wolfgang Dцbereiner to observe that certain elements with closely similar properties occur in triads, or groups of three, such as chlorine, bromine, and iodine; calcium, strontium, and barium; sulfur, selenium, and tellerium; and iron, cobalt, and manganese. Because of the limited number of known elements and the confusion that existed concerning the distinction between atomic weights and molecular weights, however, chemists were unable to grasp the significance of the Dцbereiner triads.

The development of the spectroscope in 1859 by the German physicists Robert Wilhelm Bunsen and Gustav Robert Kirchhoff made possible the discovery of many more elements (see Spectrum). In 1860, at the first international chemical congress ever held, the Italian chemist Stanislao Cannizzaro clarified the fact that some of the elements-for example, oxygen-have molecules containing two atoms. This realization finally enabled chemists to achieve a self-consistent listing of the elements.

These developments gave new impetus to the attempt to reveal interrelationships among the properties of the elements. In 1864 the British chemist John A. R. Newlands listed the elements in the order of increasing atomic weights and noted that a given set of properties recurs at every eighth place. He named this periodic repetition the law of octaves, by analogy with the musical scales. Newlands's discovery failed to impress his contemporaries, probably because the observed periodicity was limited to only a small number of the known elements.


The chemical law that the properties of all the elements are periodic functions of their atomic weights was developed independently by two chemists, in 1869 by the Russian Dmitry Mendeleyev and in 1870 by the German Julius Lothar Meyer. The key to the success of their efforts was the realization that previous attempts had failed because a number of elements were as yet undiscovered and that vacant places must be left for such elements in the classification. Thus, although no element then known had an atomic weight between those of calcium and titanium, Mendeleyev left a vacant space for it in his table. This place was later assigned to the element scandium, discovered in 1879, which has properties justifying its position in the sequence. The discovery of scandium proved to be one of a series of dramatic verifications of the predictions based on the periodic law, and validation of the law accelerated the development of inorganic chemistry.

The periodic law has undergone two principal elaborations since its original formulation by Mendeleyev and Meyer. The first revision involved extending the law to include a whole new family of elements, the existence of which was completely unsuspected in the 19th century. This group comprised the first three of the noble, or inert, gases (see Noble Gases) argon, helium, and neon, discovered in the atmosphere between 1894 and 1898 by the British physicist John William Strutt, 3rd Baron Rayleigh, and the British chemist Sir William Ramsay. The second development in the periodic law was the interpretation of the cause of the periodicity of the elements in terms of the Bohr theory (1913) of the electronic structure of the atom (see Atom).


The periodic law is most commonly expressed in chemistry in the form of a periodic table, or chart. The so-called short-form periodic table, based on Mendeleyev's table, with subsequent emendations and additions, is still in widespread use. In this table the elements are arranged in seven horizontal rows, called the periods, in order of increasing atomic weights, and in 18 vertical columns, called the groups. The first period, containing two elements, hydrogen and helium, and the next two periods, each containing eight elements, are called the short periods. The remaining periods, called the long periods, contain 18 elements, as in periods 4 and 5, or 32 elements, as in period 6. The long period 7 includes the actinide series, which has been filled in by the synthesis of radioactive nuclei through element 103, lawrencium. Heavier transuranium elements have also been synthesized.

The groups or vertical columns of the periodic table have traditionally been labeled from left to right using Roman numerals followed by the symbol a or b, the b referring to groups of transition elements. Another labeling scheme, which has been adopted by the International Union of Pure and Applied Chemistry (IUPAC), is gaining in popularity. This new system simply numbers the groups sequentially from 1 to 18 across the periodic table.

All the elements within a single group bear a considerable familial resemblance to one another and, in general, differ markedly from elements in other groups. For example, the elements of group 1 (or Ia), with the exception of hydrogen, are metals with chemical valence of +1; while those of group 17 (or VIIa), with the exception of astatine, are nonmetals, commonly forming compounds in which they have valences of -1.

ELECTRON SHELL THEORY In the periodic classification, noble gases, which in most cases are unreactive (valence = 0), are interposed between highly reactive metals that form compounds in which their valence is +1 on one side and highly reactive nonmetals forming compounds in which their valence is -1 on the other side. This phenomenon led to the theory that the periodicity of properties results from the arrangement of electrons in shells about the atomic nucleus. According to the same theory, the noble gases are normally inert because their electron shells are completely filled; other elements, therefore, may have some shells that are only partly filled, and their chemical reactivities involve the electrons in these incomplete shells. Thus, all the elements that occupy a position in the table preceding that of an inert gas have one electron less than the number necessary for completed shells and show a valence of -1, corresponding to the gain of one electron in reactions. Elements in the group following the inert gases in the table have one electron in excess of the completed shell structure and in reactions can lose that electron, thereby showing a valence of + 1.

An analysis of the periodic table, based on this theory, indicates that the first electron shell may contain a maximum of 2 electrons, the second builds up to a maximum of 8, the third to 18, and so on. The total number of elements in any one period corresponds to the number of electrons required to achieve a stable configuration. The distinction between the a and b subgroups of a given group also may be explained on the basis of the electron shell theory. Both subgroups have the same degree of incompleteness in the outermost shell but differ from each other with respect to the structures of the underlying shells. This model of the atom still provides a good explanation of chemical bonding.

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