Atomic number

Last updated November 12, 2024

The atomic number or nuclear charge number (symbol Z) of a chemical element is the charge number of its atomic nucleus. For ordinary nuclei composed of protons and neutrons, this is equal to the proton number (n_p) or the number of protons found in the nucleus of every atom of that element. The atomic number can be used to uniquely identify ordinary chemical elements. In an ordinary uncharged atom, the atomic number is also equal to the number of electrons.

For an ordinary atom which contains protons, neutrons and electrons, the sum of the atomic number Z and the neutron number N gives the atom's atomic mass number A. Since protons and neutrons have approximately the same mass (and the mass of the electrons is negligible for many purposes) and the mass defect of the nucleon binding is always small compared to the nucleon mass, the atomic mass of any atom, when expressed in daltons (making a quantity called the "relative isotopic mass"), is within 1% of the whole number A.

Atoms with the same atomic number but different neutron numbers, and hence different mass numbers, are known as isotopes. A little more than three-quarters of naturally occurring elements exist as a mixture of isotopes (see monoisotopic 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 standard atomic 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 symbol Z comes from the German word Zahl 'number', which, before the modern synthesis of ideas from chemistry and physics, merely denoted an element's numerical place in the periodic 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 this Z number was also the nuclear charge and a physical characteristic of atoms, did the word Atomzahl (and its English equivalent atomic number) come into common use in this context.

The rules above do not always apply to exotic atoms which contain short-lived elementary particles other than protons, neutrons and electrons.

History

In the 19th century, the term "atomic number" typically meant the number of atoms in a given volume.^[1]^[2] Modern chemists prefer to use the concept of molar concentration.

In 1913, Antonius van den Broek proposed that the electric charge of an atomic nucleus, expressed as a multiplier of the elementary charge, was equal to the element's sequential position on the periodic 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.^[3]^[4] 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.^[5]^[6]

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.^[7]

The periodic table and a natural number for each element

Russian chemist Dmitri Mendeleev, creator of the periodic table. DIMendeleevCab.jpg — Russian chemist Dmitri Mendeleev, creator of the periodic table.

The periodic table of elements creates an ordering of the elements, and so they can be numbered in order.^[8]^: 222 Dmitri Mendeleev arranged his first periodic tables (first published on March 6, 1869) in order of atomic weight ("Atomgewicht").^[9] However, in consideration of the elements' observed chemical properties, he changed the order slightly and placed tellurium (atomic weight 127.6) ahead of iodine (atomic weight 126.9).^[9]^[10] 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 as argon and potassium, cobalt and nickel) 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.^[8]^: 222 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 (element 71) onward (hafnium was not known at this time).

The Rutherford-Bohr model and van den Broek

In 1911, Ernest Rutherford gave a model 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 element Z = 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 were exactly equal to its place in the periodic table (also known as element number, atomic number, and symbolized Z). This eventually proved to be the case.

Moseley's 1913 experiment

The experimental position improved dramatically after research by Henry Moseley in 1913.^[11] Moseley, after discussions with Bohr who was at the same lab (and who had used Van den Broek's hypothesis in his Bohr model of the atom), 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 (K and L lines) produced by the elements from aluminium (Z = 13) to gold (Z = 79) used as a series of movable anodic targets inside an x-ray tube.^[12] The square root of the frequency of these photons (x-rays) increased from one target to the next in an arithmetic 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 calculated electric charge of the nucleus, i.e. the element number Z. Among other things, Moseley demonstrated that the lanthanide series (from lanthanum to lutetium inclusive) must have 15 members—no fewer and no more—which was far from obvious from known chemistry at that time.

Missing elements

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 (with Z < 92) which were not found and therefore identified as still undiscovered, corresponding to atomic numbers 43, 61, 72, 75, 85, 87 and 91.^[13] From 1918 to 1947, all seven of these missing elements were discovered.^[14] By this time, the first four transuranium elements had also been discovered, so that the periodic table was complete with no gaps as far as curium (Z = 96).

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 (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 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,^[15] 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 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 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.

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 nuclear charge now was realized to come entirely from a content of 79 protons. Since Moseley had previously shown that the atomic number Z of an element equals this positive charge, it was now clear that Z is identical to the number of protons of its nuclei.

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 atomic number). 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 2024, all 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 of a nuclide becomes shorter as atomic number increases,^{[ citation needed ]} though undiscovered nuclides with certain "magic" numbers of protons and neutrons may have relatively longer half-lives and comprise an island of stability.

A hypothetical element composed only of neutrons, neutronium, has also been proposed and would have atomic number 0,^[16] but has never been observed.

Related Research Articles

Atoms are the basic particles of the chemical elements. An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. The chemical elements are distinguished from each other by the number of protons that are in their atoms. For example, any atom that contains 11 protons is sodium, and any atom that contains 29 protons is copper. Atoms with the same number of protons but a different number of neutrons are called isotopes of the same element.

Atomic theory is the scientific theory that matter is composed of particles called atoms. The definition of the word "atom" has changed over the years in response to scientific discoveries. Initially, it referred to a hypothetical concept of there being some fundamental particle of matter, too small to be seen by the naked eye, that could not be divided. Then the definition was refined to being the basic particles of the chemical elements, when chemists observed that elements seemed to combine with each other in ratios of small whole numbers. Then physicists discovered that these particles had an internal structure of their own and therefore perhaps did not deserve to be called "atoms", but renaming atoms would have been impractical by that point.

In atomic physics, the Bohr model or Rutherford–Bohr model was the first successful model of the atom. Developed from 1911 to 1918 by Niels Bohr and building on Ernest Rutherford's nuclear model, it supplanted the plum pudding model of J J Thomson only to be replaced by the quantum atomic model in the 1920s. It consists of a small, dense nucleus surrounded by orbiting electrons. It is analogous to the structure of the Solar System, but with attraction provided by electrostatic force rather than gravity, and with the electron energies quantized.

A chemical element is a chemical substance whose atoms all have the same number of protons. The number of protons is called the atomic number of that element. For example, oxygen has an atomic number of 8, meaning each oxygen atom has 8 protons in its nucleus. Atoms of the same element can have different numbers of neutrons in their nuclei, known as isotopes of the element. Two or more atoms can combine to form molecules. Some elements are formed from molecules of identical atoms, e. g. atoms of hydrogen (H) form diatomic molecules (H₂). Chemical compounds are substances made of atoms of different elements; they can have molecular or non-molecular structure. Mixtures are materials containing different chemical substances; that means (in case of molecular substances) that they contain different types of molecules. Atoms of one element can be transformed into atoms of a different element in nuclear reactions, which change an atom's atomic number.

Henry Gwyn Jeffreys Moseley was an English physicist, whose contribution to the science of physics was the justification from physical laws of the previous empirical and chemical concept of the atomic number. This stemmed from his development of Moseley's law in X-ray spectra.

<span class="mw-page-title-main">Nuclear physics</span> Field of physics that studies atomic nuclei

Nuclear physics is the field of physics that studies atomic nuclei and their constituents and interactions, in addition to the study of other forms of nuclear matter.

The periodic table, also known as the periodic table of the elements, is an ordered arrangement of the chemical elements into rows ("periods") and columns ("groups"). It is an icon of chemistry and is widely used in physics and other sciences. It is a depiction of the periodic law, which states that when the elements are arranged in order of their atomic numbers an approximate recurrence of their properties is evident. The table is divided into four roughly rectangular areas called blocks. Elements in the same group tend to show similar chemical characteristics.

A proton is a stable subatomic particle, symbol p, H⁺, or ¹H⁺ with a positive electric charge of +1 e (elementary charge). Its mass is slightly less than the mass of a neutron and approximately 1836 times the mass of an electron (the proton-to-electron mass ratio). Protons and neutrons, each with a mass of approximately one atomic mass unit, are jointly referred to as nucleons (particles present in atomic nuclei).

The atomic radius of a chemical element is a measure of the size of its atom, usually the mean or typical distance from the center of the nucleus to the outermost isolated electron. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Four widely used definitions of atomic radius are: Van der Waals radius, ionic radius, metallic radius and covalent radius. Typically, because of the difficulty to isolate atoms in order to measure their radii separately, atomic radius is measured in a chemically bonded state; however theoretical calculations are simpler when considering atoms in isolation. The dependencies on environment, probe, and state lead to a multiplicity of definitions.

<span class="mw-page-title-main">Nuclide</span> Atomic species

A nuclide is a class of atoms characterized by their number of protons, Z, their number of neutrons, N, and their nuclear energy state.

Prout's hypothesis was an early 19th-century attempt to explain the existence of the various chemical elements through a hypothesis regarding the internal structure of the atom. In 1815 and 1816, the English chemist William Prout published two papers in which he observed that the atomic weights that had been measured for the elements known at that time appeared to be whole multiples of the atomic weight of hydrogen. He then hypothesized that the hydrogen atom was the only truly fundamental object, which he called protyle, and that the atoms of other elements were actually groupings of various numbers of hydrogen atoms.

<span class="mw-page-title-main">Mass number</span> Number of heavy particles in the atomic nucleus

The mass number (symbol A, from the German word: Atomgewicht, "atomic weight"), also called atomic mass number or nucleon number, is the total number of protons and neutrons (together known as nucleons) in an atomic nucleus. It is approximately equal to the atomic (also known as isotopic) mass of the atom expressed in atomic mass units. Since protons and neutrons are both baryons, the mass number A is identical with the baryon number B of the nucleus (and also of the whole atom or ion). The mass number is different for each isotope of a given chemical element, and the difference between the mass number and the atomic number Z gives the number of neutrons (N) in the nucleus: N = A − Z.

<span class="mw-page-title-main">Antonius van den Broek</span> Dutch physicist (1870–1926)

Antonius Johannes van den Broek was a Dutch mathematical economist and amateur physicist, notable for being the first who realized that the position of an element in the periodic table corresponds to the charge of its atomic nucleus. This hypothesis was published in 1911 and inspired the experimental work of Henry Moseley, who found good experimental evidence for it by 1913.

<span class="mw-page-title-main">Nuclear binding energy</span> Minimum energy required to separate particles within a nucleus

Nuclear binding energy in experimental physics is the minimum energy that is required to disassemble the nucleus of an atom into its constituent protons and neutrons, known collectively as nucleons. The binding energy for stable nuclei is always a positive number, as the nucleus must gain energy for the nucleons to move apart from each other. Nucleons are attracted to each other by the strong nuclear force. In theoretical nuclear physics, the nuclear binding energy is considered a negative number. In this context it represents the energy of the nucleus relative to the energy of the constituent nucleons when they are infinitely far apart. Both the experimental and theoretical views are equivalent, with slightly different emphasis on what the binding energy means.

A table or chart of nuclides is a two-dimensional graph of isotopes of the elements, in which one axis represents the number of neutrons and the other represents the number of protons in the atomic nucleus. Each point plotted on the graph thus represents a nuclide of a known or hypothetical chemical element. This system of ordering nuclides can offer a greater insight into the characteristics of isotopes than the better-known periodic table, which shows only elements and not their isotopes. The chart of the nuclides is also known as the Segrè chart, after the Italian physicist Emilio Segrè.

<span class="mw-page-title-main">Isotope</span> Different atoms of the same element

Isotopes are distinct nuclear species of the same chemical element. They have the same atomic number and position in the periodic table, but different nucleon numbers due to different numbers of neutrons in their nuclei. While all isotopes of a given element have similar chemical properties, they have different atomic masses and physical properties.

The atomic nucleus is the small, dense region consisting of protons and neutrons at the center of an atom, discovered in 1911 by Ernest Rutherford based on the 1909 Geiger–Marsden gold foil experiment. After the discovery of the neutron in 1932, models for a nucleus composed of protons and neutrons were quickly developed by Dmitri Ivanenko and Werner Heisenberg. An atom is composed of a positively charged nucleus, with a cloud of negatively charged electrons surrounding it, bound together by electrostatic force. Almost all of the mass of an atom is located in the nucleus, with a very small contribution from the electron cloud. Protons and neutrons are bound together to form a nucleus by the nuclear force.

In chemistry and atomic physics, an electron shell may be thought of as an orbit that electrons follow around an atom's nucleus. The closest shell to the nucleus is called the "1 shell" (also called the "K shell"), followed by the "2 shell" (or "L shell"), then the "3 shell" (or "M shell"), and so on further and further from the nucleus. The shells correspond to the principal quantum numbers (n = 1, 2, 3, 4 ...) or are labeled alphabetically with the letters used in X-ray notation (K, L, M, ...). A useful guide when understanding electron shells in atoms is to note that each row on the conventional periodic table of elements represents an electron shell.

Chemistry: A Volatile History is a 2010 BBC documentary on the history of chemistry presented by Jim Al-Khalili. It was nominated for the 2010 British Academy Television Awards in the category Specialist Factual.

<span class="mw-page-title-main">Discovery of the neutron</span> Scientific background leading to the discovery of subatomic particles

The discovery of the neutron and its properties was central to the extraordinary developments in atomic physics in the first half of the 20th century. Early in the century, Ernest Rutherford developed a crude model of the atom, based on the gold foil experiment of Hans Geiger and Ernest Marsden. In this model, atoms had their mass and positive electric charge concentrated in a very small nucleus. By 1920, isotopes of chemical elements had been discovered, the atomic masses had been determined to be (approximately) integer multiples of the mass of the hydrogen atom, and the atomic number had been identified as the charge on the nucleus. Throughout the 1920s, the nucleus was viewed as composed of combinations of protons and electrons, the two elementary particles known at the time, but that model presented several experimental and theoretical contradictions.

References

↑ Leopold Gmelin (1848). Hand-book of Chemistry, p. 52: "...the specific gravity divided by the atomic weight gives the Atomic 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
1 2 Pais, Abraham (2002). Inward bound: of matter and forces in the physical world (Reprint ed.). Oxford: Clarendon Press [u.a.] ISBN 978-0-19-851997-3.
1 2 The Periodic Table of Elements Archived 18 August 2023 at the Wayback Machine , American Institute of Physics
↑ The Development of the Periodic Table Archived 26 July 2012 at the Wayback Machine , Royal Society of Chemistry
↑ Ordering the Elements in the Periodic Table Archived 4 March 2016 at the Wayback Machine , Royal Chemical Society
↑ Moseley, H.G.J. (1913). "XCIII.The high-frequency spectra of the elements". Philosophical Magazine. Series 6. 26 (156): 1024–1034. doi:10.1080/14786441308635052. Archived (PDF) from the original on 8 July 2023. Retrieved 12 December 2023.
↑ Eric Scerri, A tale of seven elements, (Oxford University Press 2013) ISBN 978-0-19-539131-2, p.47
↑ Scerri chaps. 3–9 (one chapter per element)
↑ Ernest Rutherford | NZHistory.net.nz, New Zealand history online Archived 1 December 2012 at the Wayback Machine . Nzhistory.net.nz (19 October 1937). Retrieved on 2011-01-26.
↑ von Antropoff, A. (1926). "Eine neue Form des periodischen Systems der Elementen". Zeitschrift für Angewandte Chemie (in German). 39 (23): 722–725. Bibcode:1926AngCh..39..722V. doi:10.1002/ange.19260392303.

This page is based on this Wikipedia article
Text is available under the CC BY-SA 4.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.

[1] Leopold Gmelin (1848). Hand-book of Chemistry, p. 52: "...the specific gravity divided by the atomic weight gives the Atomic number, that is to say, the number of atoms in a given volume.

[2] 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."

[3] 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.

[4] 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.

[5] Eric Scerri (2020). The Periodic Table: Its Story and Its Significance, p. 185

[6] Helge Kragh (2012). Niels Bohr and the Quantum Atom, p. 33

[7] Helge Kragh (2012). Niels Bohr and the Quantum Atom, p. 34

[PaisInward-8] 1 2 Pais, Abraham (2002). Inward bound: of matter and forces in the physical world (Reprint ed.). Oxford: Clarendon Press [u.a.] ISBN 978-0-19-851997-3.

[dm1869-9] 1 2 The Periodic Table of Elements Archived 18 August 2023 at the Wayback Machine , American Institute of Physics

[10] The Development of the Periodic Table Archived 26 July 2012 at the Wayback Machine , Royal Society of Chemistry

[11] Ordering the Elements in the Periodic Table Archived 4 March 2016 at the Wayback Machine , Royal Chemical Society

[12] Moseley, H.G.J. (1913). "XCIII.The high-frequency spectra of the elements". Philosophical Magazine. Series 6. 26 (156): 1024–1034. doi:10.1080/14786441308635052. Archived (PDF) from the original on 8 July 2023. Retrieved 12 December 2023.

[13] Eric Scerri, A tale of seven elements, (Oxford University Press 2013) ISBN 978-0-19-539131-2, p.47

[14] Scerri chaps. 3–9 (one chapter per element)

[15] Ernest Rutherford | NZHistory.net.nz, New Zealand history online Archived 1 December 2012 at the Wayback Machine . Nzhistory.net.nz (19 October 1937). Retrieved on 2011-01-26.

[Antropoff_1926-16] von Antropoff, A. (1926). "Eine neue Form des periodischen Systems der Elementen". Zeitschrift für Angewandte Chemie (in German). 39 (23): 722–725. Bibcode:1926AngCh..39..722V. doi:10.1002/ange.19260392303.

[1]

[2]

[3]

[4]

[5]

[6]

[7]

[8]

[9]

[10]

[11]

[12]

[13]

[14]

[15]

[16]