Atomic number

The atomic number of a chemical element (atomic species) is the number of protons in the nucleus of its atom. This number is conventionally denoted by Z and is by definition a non-negative natural number. For instance, the element carbon is characterized by Z = 6 and the element phosphorus by Z = 15. A proton having a positive charge e (the elementary charge), the atom of a chemical element has nuclear charge Ze. Thus, the carbon nucleus has charge 6e (contains 6 protons) and the phosphorus nucleus contains 15 protons and has charge 15e.

Since atoms are electrically neutral, Z gives simultaneously the number of negative particles (electrons) of charge −e orbiting the nucleus. The chemical and physical properties of an atom are solely determined by the number of its electrons and hence by its nuclear charge: the nuclear charge is a unique "fingerprint" of an element and Z labels the chemical elements uniquely.

All elements can be arranged in increasing atomic mass in rows, this gives the periodic system of elements. This arrangement was first thought of by the Russian chemist Dmitri Mendeléev, who discovered that chemically similar elements appear in the same columns of the periodic system. The sequence number of an element in this table is the atomic number (see "Short history" below). The first 94 elements, with atomic number 1 ≤ Z ≤ 94, appear in nature, although the elements with Z = 93, 94 appear only as minute traces. The elements with 95 ≤ Z ≤ 118 are man-made, and very short-lived.

A list of the elements sorted by atomic number Z can be found in this article.

Short history

At the end of the 18th century and the beginning of the 19th century, men like Antoine Lavoisier in France and John Dalton in England discovered that matter could not be decomposed indefinitely; at a certain point one ends up with indivisible particles—atoms. About the same time it was recognized that only a limited number of such different atoms exists, each type of atom constituting a so-called chemical element. (To avoid misunderstanding: the early 19th century chemists thought mistakenly atoms to be indivisible, now it is known that atoms can be broken up in a nucleus and electrons). In Dalton's days there were about forty different elements known; at present there are about 110.

In the early 19th century one was not yet able to determine atomic weights with great accuracy, in particular one could not yet distinguish the masses of different isotopes. This inaccuracy lead William Prout to surmise in 1815 that all atomic weights are integral multiples of the atomic weight of the lightest element, hydrogen. A century later, after the discovery of isotopes by Frederick Soddy in 1911, Prout's idea turned out to be very a good one, but during most of the 19th century it was believed to be false. When the weighing of atoms became more refined in the course of the 19th century, one discovered that atomic masses are not integral multiples of the hydrogen mass; now we understand that this is because isotopically averaged masses were determined.

Based upon atomic masses and chemical properties Dmitri Ivanovich Mendeléev arranged in 1869 the elements then known (ca. 60) in a very useful table, the Periodic Table of Elements. After he had done that, the atomic number Z could be defined as the sequence number of the element in the periodic table, counting rowwise from left to right and top to bottom. However, at that time the structure of the atom was not yet known, electrons were still to be discovered and the atomic number had no physical meaning whatever, other than as sequence number in Mendeléev's table.

In 1911 Ernest Rutherford discovered that atoms have a heavy, small, positively charged pit, which he called the nucleus of the atom. He observed that the charge is Z times the elementary charge e, where Z is the (already known) position of the element in the periodic table. Rutherford noticed that Z was often approximately half the atomic mass A. (We now know that many nuclei have the same number of neutrons as protons and that the masses of proton and neutron are almost equal, which explains why for many elements Z is to a good approximation equal to A/2.)

In 1913 Henry Gwyn Jeffreys Moseley performed X-ray measurements on a large number of elements and interpreted his results with a simple formula that contained the nuclear charge. For the first time he established unambiguously that the position of an element in Mendeléev's periodic table was solely determined by the nuclear charge of the atom.

It then took another twenty years before the structure of the nucleus was completely understood. In 1913 protons (which, after all, are nothing but atomic hydrogen cations) were well-known, but the neutron still had to be discovered. This was done by James Chadwick in 1932. As far as chemistry is concerned this discovery clinched the story of the atom. An atom consists of a positively charged nucleus with Z protons and A−Z neutrons; around it orbit Z electrons, each of charge −e.

References

• A. Pais, Inward Bound. Of Matter and Forces in the Physical World, Oxford University Press, New York (1986).