Theatomic number ornuclear charge number (symbolZ) of achemical element is thecharge number of itsatomic nucleus. For ordinary nuclei composed ofprotons andneutrons, this is equal to theproton number (np) or the number of protons found in the nucleus of everyatom of that element. The atomic number can be used to uniquely identify ordinarychemical elements. In an ordinaryuncharged atom, the atomic number is also equal to the number ofelectrons.
For an ordinary atom which contains protons, neutrons andelectrons, the sum of the atomic numberZ and theneutron numberN gives the atom'satomic mass numberA. Since protons and neutrons have approximately the same mass (and the mass of the electrons is negligible for many purposes) and themass defect of thenucleon binding is always small compared to the nucleon mass, theatomic mass of any atom, when expressed indaltons (making a quantity called the "relative isotopic mass"), is within 1% of the whole numberA.
Atoms with the same atomic number but different neutron numbers, and hence different mass numbers, are known asisotopes. A little more than three-quarters of naturally occurring elements exist as a mixture of isotopes (seemonoisotopic elements), and the average isotopic mass of an isotopic mixture for an element (called the relative atomic mass) in a defined environment on Earth determines the element's standardatomic weight. Historically, it was these atomic weights of elements (in comparison to hydrogen) that were the quantities measurable by chemists in the 19th century.
The conventional symbolZ comes from the German wordZahl 'number', which, before the modern synthesis of ideas from chemistry and physics, merely denoted an element's numerical place in theperiodic table, whose order was then approximately, but not completely, consistent with the order of the elements by atomic weights. Only after 1915, with the suggestion and evidence that thisZ number was also the nuclear charge and a physical characteristic of atoms, did the wordAtomzahl (and its English equivalentatomic number) come into common use in this context.
The rules above do not always apply toexotic atoms which contain short-lived elementary particles other than protons, neutrons and electrons.
An explanation of the superscripts and subscripts seen inAZE notation. Atomic number is the number of protons, and therefore also the total positive charge, in the atomic nucleus.
The atomic number is used inAZE notation, (withA as themass number,Z the atomic number, and E forelement) to denote anisotope. When achemical symbol is used, e.g. "C" for carbon, standard notation uses asuperscript at the upper left of the chemical symbol for the mass number and indicates the atomic number with asubscript at the lower left (e.g.3 2He,4 2He,12 6C,14 6C,235 92U, and239 92U). Because the atomic number is given by the element symbol, it is common to state only the mass number in the superscript and leave out the atomic number subscript (e.g.3 He,4 He,12 C,14 C,235 U, and239 U).
The common pronunciation of the AZE notation is different from how it is written:4 2He is commonly pronounced as helium-four instead of four-two-helium, and235 92U as uranium two-thirty-five (American English) or uranium-two-three-five (British) instead of 235-92-uranium.Various notations appear in older sources were used, such asNe(22) in 1934,[1]: 226 Ne22 for neon-22 (1935)[2] or Pb210 for lead-210 (1933)[3]: 7
In the 19th century, the term "atomic number" typically meant the number of atoms in a given volume.[4][5] Modern chemists prefer to use the concept ofmolar concentration.
In 1913,Antonius van den Broek proposed that theelectric charge of an atomic nucleus, expressed as a multiplier of theelementary charge, was equal to the element's sequential position on theperiodic table.Ernest Rutherford, in various articles in which he discussed van den Broek's idea, used the term "atomic number" to refer to an element's position on the periodic table.[6][7] No writer before Rutherford is known to have used the term "atomic number" in this way, so it was probably he who established this definition.[8][9]
After Rutherford deduced the existence of the proton in 1920, "atomic number" customarily referred to the proton number of an atom. In 1921, the German Atomic Weight Commission based its new periodic table on the nuclear charge number and in 1923 the International Committee on Chemical Elements followed suit.[10]
The periodic table and a natural number for each element
Theperiodic table of elements creates an ordering of the elements, and so they can be numbered in order.[11]: 222 Dmitri Mendeleev arranged his first periodic tables (first published on March 6, 1869) in order ofatomic weight ("Atomgewicht").[12] However, in consideration of the elements' observed chemical properties, he changed the order slightly and placedtellurium (atomic weight 127.6) ahead ofiodine (atomic weight 126.9).[12][13] 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 atomic weight position was never entirely satisfactory. In addition to the case of iodine and tellurium, several other pairs of elements (such asargon andpotassium,cobalt andnickel) were later shown to have nearly identical or reversed atomic weights, thus requiring their placement in the periodic table to be determined by their chemical properties.[11]: 222 However the gradual identification of more and more chemically similarlanthanide elements, whose atomic number was not obvious, led to inconsistency and uncertainty in the periodic numbering of elements at least fromlutetium (element 71) onward (hafnium was not known at this time).
The Rutherford–Bohr model of thehydrogen atom (Z = 1) or a hydrogen-like ion (Z > 1). In this model, it is an essential feature that the photon energy (or frequency) of the electromagnetic radiation emitted (shown) when an electron jumps from one orbital to another be proportional to the mathematical square of atomic charge (Z2). Experimental measurements byHenry Moseley of this radiation for many elements (fromZ = 13 to 92) showed the results as predicted by Bohr. Both the concept of atomic number and the Bohr model were thereby given scientific credence.
In 1911,Ernest Rutherford gave amodel of the atom in which a central nucleus 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 (though it was almost 25% different from the atomic number of gold(Z = 79,A = 197), 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 (but was elementZ = 79 on the periodic table), 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 wereexactly equal to its place in the periodic table (also known as element number, atomic number, and symbolizedZ). This eventually proved to be the case.
The experimental position improved dramatically after research byHenry Moseley in 1913.[14] Moseley, after discussions with Bohr who was at the same lab (and who had used Van den Broek's hypothesis in hisBohr model of the atom), decided to test Van den Broek's and Bohr's hypothesis directly, by seeing ifspectral lines emitted from excited atoms fitted the Bohr theory's postulation that the frequency of the spectral lines be proportional to the square ofZ.
To do this, Moseley measured the wavelengths of the innermost photon transitions (K and L lines) produced by the elements fromaluminium (Z = 13) to gold (Z = 79) used as a series of movable anodic targets inside anx-ray tube.[15] The square root of the frequency of these photons(x-rays) increased from one target to the next in anarithmetic progression. This led to the conclusion (Moseley's law) that the atomic number does closely correspond (with an offset of one unit for K-lines, in Moseley's work) to the calculatedelectric charge of the nucleus, i.e. the element numberZ. Among other things, Moseley demonstrated that thelanthanide series (fromlanthanum tolutetium inclusive) must have 15 members—no fewer and no more—which was far from obvious from known chemistry at that time.
After Moseley's death in 1915, the atomic numbers of all known elements from hydrogen to uranium (Z = 92) were examined by his method. There were seven elements (withZ < 92) which were not found and therefore identified as still undiscovered, corresponding to atomic numbers 43, 61, 72, 75, 85, 87 and 91.[16] From 1918 to 1947, all seven of these missing elements were discovered.[17] By this time, the first fourtransuranium elements had also been discovered, so that the periodic table was complete with no gaps as far ascurium (Z = 96).
In 1915, the reason for nuclear charge being quantized in units ofZ, which were now recognized to be the same as the element number, was not understood. An old idea calledProut's hypothesis had postulated that the elements were all made of residues (or "protyles") 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 andThomas 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 anuclear reaction between alpha particles andnitrogen gas,[18] and believed he had proven Prout's law. He called the new heavy nuclear particles protons in 1920 (alternate names being proutons and protyles). 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 ofhydrogen nuclei, and thus there was required a hypothesis for the neutralization of the extraprotons presumed present in all heavy nuclei. A helium nucleus was presumed to have four protons plus two "nuclear electrons" (electrons bound inside the nucleus) to cancel two 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.
All consideration of nuclear electrons ended withJames Chadwick'sdiscovery of the neutron in 1932. An atom ofgold now was seen as containing 118 neutrons rather than 118 nuclear electrons, and its positive nuclear charge now was realized to come entirely from a content of 79 protons. Since Moseley had previously shown that the atomic numberZ of an element equals this positive charge, it was now clear thatZ is identical to the number of protons of its nuclei.
Each element has a specific set of chemical properties as a consequence of the number of electrons present in the neutral atom, which isZ (the atomic number). Theconfiguration of these electrons follows from the principles ofquantum mechanics. The number of electrons in each element'selectron shells, particularly the outermostvalence shell, is the primary factor in determining itschemical 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 ofany mixture of atoms with a given atomic number.
The quest for new elements is usually described using atomic numbers. As of 2025, all elements with atomic numbers 1 to 118have been observed. The most recent element discovered was number 117 (tennessine) in 2009. 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, thehalf-life of anuclide becomes shorter as atomic number increases,[citation needed] thoughundiscovered nuclides with certain "magic" numbers of protons and neutrons may have relatively longer half-lives and comprise anisland of stability.
A hypothetical element composed only of neutrons,neutronium, has also been proposed and would have atomic number 0,[19] but has never been observed.
^Leopold Gmelin (1848).Hand-book of Chemistry, p. 52: "...the specific gravity divided by the atomic weight gives theAtomic number, that is to say,the number of atoms in a given volume.
^James Curtis Booth, Campbell Morfit (1890).The Encyclopedia of Chemistry, Practical and Theoretical p.271: "The atomic number of a substance is its specific gravity, divided by its combining weight or equivalent. [...] the spec. grav. of a substance must be the number of atoms in a given volume, multiplied by their combining weight."
^Ernest Rutherford (March 1914)."The Structure of the Atom".Philosophical Magazine. 6.27:488–498.It is obvious from the consideration of the cases of hydrogen and helium, where hydrogen has one electron and helium two, that the number of electrons cannot be exactly half the atomic weight in all cases. This has led to an interesting suggestion by van den Broek that the number of units of charge on the nucleus, and consequently the number of external electrons, may be equal to the number of the elements when arranged in order of increasing atomic weight.
^Ernest Rutherford (11 December 1913). "The Structure of the Atom".Nature.92 (423).The original suggestion of van der Broek that the charge on the nucleus is equal to the atomic number and not to half the atomic weight seems to me very promising.
^Eric Scerri (2020).The Periodic Table: Its Story and Its Significance, p. 185
^Helge Kragh (2012).Niels Bohr and the Quantum Atom, p. 33
^Helge Kragh (2012).Niels Bohr and the Quantum Atom, p. 34
^abPais, Abraham (2002).Inward bound: of matter and forces in the physical world (Reprint ed.). Oxford: Clarendon Press [u.a.]ISBN978-0-19-851997-3.
