Periodic Law

I

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.

II

Historical Development

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 Döbereiner 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 tellurium; 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 Döbereiner 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.

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III

Mendeleyev and Meyer

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).

IV

Short-Form Periodic Table

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 102, nobelium. 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.

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