Periodic Table

I

Introduction

Periodic Table, a chart of all the chemical elements arranged in a manner that reflects the properties of the elements and the structures of their atoms. The elements are arranged generally (though not entirely) in order of their relative atomic mass (atomic weight). They are placed in seven horizontal rows, called periods, and in 18 vertical columns, called 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. Period 7 does not have a definite length, because only the first six elements occur naturally; further elements are created artificially. The heaviest atoms exist only fleetingly, and new ones are continually being produced (see Transuranic Elements).

The groups have traditionally been labelled from left to right using Roman numerals followed by the symbol a or b. Another labelling scheme is gaining in popularity, which simply numbers the groups sequentially (left to right) from 1 to 18 across the periodic table.

The table is arranged so that all the elements within a single group (column) bear a considerable family 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 a chemical valence of +1 (that is tending to lose an electron). Those of group 17 (or VIIa) are non-metals (though the heaviest, astatine, shows some metallic properties). These elements commonly form compounds in which they have valences of -1 (that is, they tend to gain an electron.)

II

Electron Shell Theory

The periodicity of properties results from the arrangement of electrons in shells about the atomic nucleus. The noble gases (group 0 or 18) are normally inert, because their electron shells are completely filled; other elements have some shells that are only partly filled, and their chemical reactivities involve the electrons in these incomplete shells. Thus, all the elements in group 17 (or VIIb), which is to the left of the inert gases, have one electron less than the number necessary for completed shells; they show a valence of -1. Elements in group 1 (or Ia), following the inert gases, have one electron in excess of the completed shell structure and in reactions can lose that electron, thereby showing a valence of + 1.

Moving across each period from left to right, the outer shells of the atoms fill up with electrons. The first electron shell may contain 1 electron (hydrogen) or 2 (helium). The second builds up from 1 (lithium, Li) to a maximum of 8 (neon, Ne). The third also builds up to 8, and so on. The total number of elements in any one period corresponds to the number of electrons required to achieve a stable configuration.

An element’s numerical position in the periodic table (counting from left to right and top to bottom) is called its atomic number. It is equal to the total number of electrons present in the normal, electrically neutral atom.

Beyond the inert gas argon (Ar, atomic number 18), the electrons do not fill the outermost shells in a simple and regular way. Potassium (K, atomic number 19) has 1 electron in its outermost subshell, or orbital, and calcium (Ca, atomic number 20) has 2; but in the next 10 elements new electrons are added to the next subshell in, where they have very little effect on the atom’s chemical properties. Thus the 10 elements from scandium (Sc, atomic number 21) to zinc (Zn, atomic number 30) have very similar properties. They are placed between calcium and gallium (Ga, atomic number 31), and are called transition elements. Their presence makes period 4 into a long period. Another series of transition elements occurs in period 5, running from yttrium (Y, atomic number 39) to cadmium (Cd, atomic number 48).

Periods 6 and 7 are made even longer by the fact that within each of their transition series is embedded an inner transition series in which a still deeper subshell is filled with electrons. The inner transition series in period 5, together with the chemically similar preceding element, lanthanum (La, atomic number 57), are called the lanthanides, lanthanoids, or rare earth elements. In period 7, the elements beginning with actinium (Ac, atomic number 89) are called the actinides, or actinoids.

III

Quantum Theory

The first successful versions of the periodic table were drawn up independently in the late 19th century by Dmitry Mendeleyev and Julius Lothar Meyer (see Periodic Law). With the development of quantum theory and its application to atomic structure by the Danish physicist Niels Bohr and other scientists, most of the detailed features of the periodic table have found a ready explanation. Every electron is characterized by four quantum numbers that designate its orbital motion in space. By means of the selection rules governing these quantum numbers and the exclusion principle of Wolfgang Pauli, which states that no two electrons in the same atom can have the same quantum numbers, physicists can determine theoretically the maximum number of electrons required to complete each shell. These results explain the electron structures inferred from experiment and expressed in the periodic table.

Further development of the quantum theory revealed why some elements have only one incomplete shell (namely, the outermost, or valence, shell), whereas others may have incomplete underlying shells as well.