In chemistry and physics, the atomic number of a chemical element is the number of protons
found in the nucleus of an atom of that element, and therefore identical to the charge number
of the nucleus. It is conventionally represented by the symbol Z. The atomic number uniquely
identifies a chemical element. In an uncharged atom, the atomic number is also equal to the
number of electrons. The atomic number, Z, should not be confused
with the mass number, A, which is the number of nucleons, the total number of protons and
neutrons in the nucleus of an atom. The number of neutrons, N, is known as the neutron number
of the atom; thus, A = Z + N. Since protons and neutrons have approximately the same mass
and the mass defect of nucleon binding is always small compared to the nucleon mass,
the atomic mass of any atom, when expressed in unified atomic mass units is roughly equal
to the whole number A. Atoms with the same atomic number Z but different
neutron numbers N, and hence different atomic masses, are known as isotopes. A little more
than three-quarters of naturally occurring elements exist as a mixture of isotopes, and
the average isotopic mass of an isotopic mixture for an element in a defined environment on
Earth, determines the element's standard atomic weight. Historically, it was these atomic
weights of elements that were the quantities measurable by chemists in the 19th century.
The conventional symbol Z comes from the German word Zahl meaning numberfigure, which, prior
to the modern synthesis of ideas from chemistry and physics, merely denoted an element's numerical
place in the periodic table, whose order is approximately, but not completely, consistent
with the order of the elements by atomic weights. Only after 1915, with the suggestion and evidence
that this Z number was also the nuclear charge and a physical characteristic of atoms, did
the word Atomzahl come into common use in this context.
History The periodic table and a natural number for
each element
Loosely speaking, the existence or construction of a periodic table of elements creates an
ordering of the elements, and so they can be numbered in order.
Dmitri Mendeleev claimed that he arranged his first periodic tables in order of atomic
weight. However, in consideration of the elements' observed chemical properties, he changed the
order slightly and placed tellurium ahead of iodine. This placement is consistent with
the modern practice of ordering the elements by proton number, Z, but that number was not
known or suspected at the time. A simple numbering based on periodic table
position was never entirely satisfactory, however. Besides the case of iodine and tellurium,
later several other pairs of elements were known to have nearly identical or reversed
atomic weights, thus requiring their placement in the periodic table to be determined by
their chemical properties. However the gradual identification of more and more chemically
similar lanthanide elements, whose atomic number was not obvious, led to inconsistency
and uncertainty in the periodic numbering of elements at least from lutetium onwards.
The Rutherford-Bohr model and van den Broek In 1911, Ernest Rutherford gave a model of
the atom in which a central core held most of the atom's mass and a positive charge which,
in units of the electron's charge, was to be approximately equal to half of the atom's
atomic weight, expressed in numbers of hydrogen atoms. This central charge would thus be approximately
half the atomic weight, the single element from which Rutherford made his guess). Nevertheless,
in spite of Rutherford's estimation that gold had a central charge of about 100, a month
after Rutherford's paper appeared, Antonius van den Broek first formally suggested that
the central charge and number of electrons in an atom was exactly equal to its place
in the periodic table. This proved eventually to be the case.
Moseley's 1913 experiment The experimental position improved dramatically
after research by Henry Moseley in 1913. Moseley, after discussions with Bohr who was at the
same lab, decided to test Van den Broek's and Bohr's hypothesis directly, by seeing
if spectral lines emitted from excited atoms fitted the Bohr theory's postulation that
the frequency of the spectral lines be proportional to the square of Z.
To do this, Moseley measured the wavelengths of the innermost photon transitions produced
by the elements from aluminum to gold used as a series of movable anodic targets inside
an x-ray tube. The square root of the frequency of these photons increased from one target
to the next in an arithmetic progression. This led to the conclusion that the atomic
number does closely correspond to the calculated electric charge of the nucleus, i.e. the element
number Z. Among other things, Moseley demonstrated that the lanthanide series must have 15 members—no
fewer and no more—which was far from obvious from the chemistry at that time.
The proton and the idea of nuclear electrons In 1915 the reason for nuclear charge being
quantized in units of Z, which were now recognized to be the same as the element number, was
not understood. An old idea called Prout's hypothesis had postulated that the elements
were all made of residues of the lightest element hydrogen, which in the Bohr-Rutherford
model had a single electron and a nuclear charge of one. However, as early as 1907 Rutherford
and Thomas Royds had shown that alpha particles, which had a charge of +2, were the nuclei
of helium atoms, which had a mass four times that of hydrogen, not two times. If Prout's
hypothesis were true, something had to be neutralizing some of the charge of the hydrogen
nuclei present in the nuclei of heavier atoms. In 1917 Rutherford succeeded in generating
hydrogen nuclei from a nuclear reaction between alpha particles and nitrogen gas, and believed
he had proven Prout's law. He called the new heavy nuclear particles protons in 1920. It
had been immediately apparent from the work of Moseley that the nuclei of heavy atoms
have more than twice as much mass as would be expected from their being made of hydrogen
nuclei, and thus there was required a hypothesis for the neutralization of the extra protons
presumed present in all heavy nuclei. A helium nucleus was presumed to be composed of four
protons plus two "nuclear electrons" to cancel two of the charges. At the other end of the
periodic table, a nucleus of gold with a mass 197 times that of hydrogen, was thought to
contain 118 nuclear electrons in the nucleus to give it a residual charge of + 79, consistent
with its atomic number. The discovery of the neutron makes Z the proton
number All consideration of nuclear electrons ended
with James Chadwick's discovery of the neutron in 1932. An atom of gold now was seen as containing
118 neutrons rather than 118 nuclear electrons, and its positive charge now was realized to
come entirely from a content of 79 protons. After 1932, therefore, an element's atomic
number Z was also realized to be identical to the proton number of its nuclei.
The symbol of Z The conventional symbol Z possibly comes from
the German word Atomzahl. However, prior to 1915, the word Zahl was used for an element's
assigned number in the periodic table. Chemical properties
Each element has a specific set of chemical properties as a consequence of the number
of electrons present in the neutral atom, which is Z. The configuration of these electrons
follows from the principles of quantum mechanics. The number of electrons in each element's
electron shells, particularly the outermost valence shell, is the primary factor in determining
its chemical bonding behavior. Hence, it is the atomic number alone that determines the
chemical properties of an element; and it is for this reason that an element can be
defined as consisting of any mixture of atoms with a given atomic number.
New elements The quest for new elements is usually described
using atomic numbers. As of 2010, elements with atomic numbers 1 to 118 have been observed.
Synthesis of new elements is accomplished by bombarding target atoms of heavy elements
with ions, such that the sum of the atomic numbers of the target and ion elements equals
the atomic number of the element being created. In general, the half-life becomes shorter
as atomic number increases, though an "island of stability" may exist for undiscovered isotopes
with certain numbers of protons and neutrons. See also
History of the periodic table Effective atomic number
Atomic theory Prout's hypothesis
References