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From Citizendium, the Citizens' Compendium
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.
Before it was known that atoms have a charged nucleus, the atomic number was defined as the position (the sequence number Z) of the atom in the periodic system of elements (see "Short history" below). In the periodic system, all elements are arranged in increasing atomic mass in rows; the determination of atomic masses was possible long before the structure of atoms was known. The arrangement of atoms by mass was thought of by the Russian chemist Dmitri Mendeléev, who discovered that chemically similar elements appear in the same columns of the periodic system (see history of periodic table). 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 here.
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 twenty to thirty 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 Joseph Louis Proust 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, Proust'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, it was 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 in a very useful table, the Periodic Table of Elements. After he had done that, the atomic number Z was defined as the sequence number of the element in the periodic table, counting row wise 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.
After the discovery of the electron at the end of the 19th century, it was assumed that atoms formed a source of electrons, but how many electrons an atom contained was not known at all.
In 1911 Ernest Rutherford proposed that neutral atoms have a heavy small "pit" of positive charge N|e| that neutralizes the charge of the (negative) electrons of the atom; two years later he called this "pit" the nucleus of the atom. In his 1911 paper Rutherford noticed that N is often proportional to the atomic mass A and three years later he observed that for many nuclei N ≈ A/2. (We now know that many elements 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 the number of protons N is to a good approximation equal to A/2.) Since the position of the elements in the periodic system is determined by A, Rutherford had in 1911 a glimmer of insight in the relationship between this position and nuclear charge.
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 Z of an element in Mendeléev's periodic table, the atomic number (in the old sense), was solely determined by the nuclear charge N of the atom (the atomic number in the modern sense). Since in fact he established that N≡Z it is now common to use only the symbol Z.
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.
- ↑ E. Rutherford, Philosophical Magazine, vol. 21, p. 669 (1911)
- ↑ E. Rutherford, Radioactive substances and their radiation, Cambridge University Press, (1913). p. 184
- ↑ E. Rutherford, Philosophical Magazine, vol. 27, p. 488 (1914)
- A. Pais, Inward Bound. Of Matter and Forces in the Physical World, Oxford University Press, New York (1986).