periodic law

periodic law statement of a periodic recurrence of chemical and physical properties of the elements when the elements are arranged in order of increasing atomic number . Such an arrangement in the form of a table in which the groupings of elements having similar properties are easily identified is called the periodic system or the periodic table .

Pioneering Periodic Arrangements of the Elements

Laws of Triads and of Octaves

Early in the 19th cent., a number of chemists had noticed certain relationships between the properties of elements and their atomic weight. In 1829 J. W. Döbereiner stated that there existed some three-element groups, or triads, in which the atomic weight of the middle element was the average of the other two and the properties of this element lay between those of the other two. For example, calcium, strontium, and barium form a triad; lithium, sodium, and potassium, another. The English chemist J. A. Newlands found (1863-65) that if the elements are listed according to atomic weight starting with the second, the 8th element following any given element has similar chemical properties, and so does the 16th. This became known as the law of octaves. About the same time, A. E. de Chancourtois arranged the elements according to increasing atomic weight in the form of a vertical helix with eight elements in a turn, so that elements having similar properties fell along vertical lines.

The Periodic Table

D. I. Mendeleev was the first to state the periodic law close to its present form. He proposed in 1869 that the properties of elements are periodic functions of the atomic weight and grouped the elements accordingly in a periodic system. Working independently and not aware of Mendeleev's work, Lothar Meyer arrived at a similar system, publishing his results about a year after Mendeleev's. When Mendeleev devised his periodic table a number of positions could not be fitted by any of the then known elements. Mendeleev suggested that these empty spaces represented undiscovered elements and by means of his system accurately predicted their general properties and atomic weights.

Introduction of Atomic Numbers

The work (1913-14) of H. G. Moseley on the X-ray spectra of elements (see X ray ) led to the present form of the periodic law. He found that the wavelength of the X-radiation of elements decreased with increasing atomic weight. However, the relationship was not a strict one. He assigned a new set of numbers, called atomic numbers, to the elements he had studied, so that there was a relation between the wavelength and the atomic number. The atomic number is the number of positive charges, or protons, contained in the atomic nucleus (see atom ) or, equivalently, the number of negative charges, or electrons, outside the nucleus in a neutral atom. The periodic law can be explained on the basis of the electronic structure of the atom, which is believed to be the main factor underlying the chemical properties and many of the physical properties of the elements. In turn, the electronic structures of atoms have been successfully accounted for by the quantum theory .

In spite of its great success, the periodic system that had been introduced by Mendeleev had some discrepancies. Arranged strictly according to atomic weight, not all elements fell into their proper groups. Better arrangement could be made if the positions of certain neighboring couples were interchanged. For example, to suit the chemical order of the table, the inert gas argon (at. wt. 39.948) should come before the chemically active metal potassium (at. wt. 39.0983). Through Moseley's work, it was found that although the atomic number of an element is roughly half its atomic weight, the atomic weight does not always increase with increasing atomic number. The discrepancies occur just for those elements where Mendeleev's law failed. Based on atomic number, the periodic law now has no exceptions. Although all the missing elements in the periodic table have been found (with the aid of the periodic table itself), the table retains its usefulness to the chemist as a reliable check for disputed or uncertain data concerning some of the known elements.