Atomic mass

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The atomic mass , the mass of a single atom , can be specified like any mass in the SI unit kilogram (kg). Usually the mass of an atom is expressed in atomic mass units,


The atomic mass unit , formerly with ( atomic mass unit hereinafter) is one-twelfth the mass of an atom of the carbon - isotope 12 C . In biochemistry , and in the USA also in organic chemistry , the atomic mass unit is also known as the Dalton (unit symbol: Da ), named after the English naturalist John Dalton .

In the chemical is on the recommendation of IUPAC , the numerical value by itself, without unity, as relative atomic mass (Engl. Atomicweight ) designated and formally as a separate, dimensionless understood, namely as the mass ratio of each atom to an imaginary atomic mass . In contrast to this relative atomic mass of the specified in kg, g, or u mass is as absolute atomic mass (engl. Atomic mass ), respectively.

The atomic masses of the nuclides are approximate, but because of the different masses of proton and neutron and the mass defect, they are not exactly integral multiples of the mass of the hydrogen atom . In lists like Atomic Mass Adjustment 2012 and in interactive nuclide maps , the excess mass is often given instead of the atomic mass , sometimes both excess mass and atomic mass.

The mass ratios of the substances involved in a chemical reaction can be calculated from the atomic masses, the molecular masses that can be calculated from them and the molar mass derived from them .

The average atomic mass of a mixing element is calculated as the weighted arithmetic mean of the atomic masses of the isotopes with the natural abundances of the isotopes as weights. In chemistry, this average atomic mass is called the atomic weight of the element.


Table with atomic weights in Johann Samuel Traugott Gehler's physical dictionary 1840

The first table of relative atomic masses was published by John Dalton in 1805 . He obtained it based on the mass ratios in chemical reactions, where he chose the lightest atom, the hydrogen atom , as the "unit of mass" (see atomic unit of mass ).

Further relative atomic and molecular masses were calculated for gaseous elements and compounds on the basis of Avogadro's law , that is, by weighing and comparing known gas volumes, later also with the help of Faraday's law . Avogadro still called the smallest imaginable parts molecules. Berzelius then introduced the term atom for the smallest conceivable part of a substance. Arbitrarily he set the atomic weight of oxygen equal to 100. Later researchers chose the lightest substance, hydrogen, as the standard, but set the hydrogen molecule equal to 1. They then received the “ equivalent weight ” 6 for carbon and 8 for oxygen.

The real pioneer for correct atomic weights of elements was Jean Baptiste Dumas . He determined the atomic weights of 30 elements very precisely and found that 22 elements had atomic weights that are multiples of the atomic weight of hydrogen.

It was only Stanislao Cannizzaro who introduced the current distinction between atom and molecule in 1858. He assumed that a molecule of hydrogen consists of two atoms of hydrogen. For the individual hydrogen atom he arbitrarily set the atomic weight of 1, a hydrogen molecule consequently has a molecular mass of 2. In 1865, oxygen , whose atoms have on average approximately 16 times the mass of the hydrogen atom, was proposed by Jean Servais Stas as a reference element and mass was suggested to him 16.00 allocated.

In 1929, WF Giauque and HL Johnston discovered that oxygen has three isotopes. This prompted IUPAP to introduce a mass scale based on m ( 16 O), while IUPAC continued to use A r (O) = 16, i.e. oxygen in its natural isotopic composition.

In 1957, AO Nier and A. Ölander independently proposed that A r ( 12 C) and m ( 12 C) = 12 u should replace the old atomic mass units. IUPAP and IUPAC then agreed on this in the years 1959–1961. Up until then, physicists and chemists had two slightly different mass scales. In 1960 Everling, König, Josef Mattauch and Aaldert Wapstra published masses of nuclides

The carbon isotope 12 C with a mass of 12 u serves as the reference base to this day. The atomic mass indicates how many times greater the mass of the respective atom is than 1/12 of the mass of the 12 carbon atom. As mentioned above, the atomic masses of the nuclides are approximate, but not exactly, integer multiples of the mass of the hydrogen atom.

The following table shows some average ( see below ) relative atomic masses, i.e. atomic weights, depending on the four different reference masses:

element based on
nat H = 1 nat O = 16 16 O = 16 12 C = 12
nat H 01,000 01.008 01.008 01.008
nat Cl 35.175 35.457 35,464 35,453
nat O 15.872 16,000 16.004 15.999
nat N 13,896 14.008 14.011 14.007
nat C 11.916 12.011 12.015 12.011

Measurement, data collection

Exact atomic masses are determined today with mass spectrometers . The atomic masses of the individual isotopes result very precisely. To determine the atomic mass of the elements in their natural isotopic composition (atomic weights), the isotope ratio must then be determined. For the purposes of chemistry this average atomic mass of the natural isotope mixture in the earth's crust is given; In special cases, the origin of the isotope mixture must be taken into account.

Further examples of the relative atomic weights of some chemical elements:

An international group of experts founded by Aaldert Wapstra has been collecting measurement results of the atomic masses of all known nuclides from original publications and using them to form estimated (i.e. evaluated, professionally assessed) weighted mean values. The results were published in the journal Nuclear Physics A up to 2003 . Wapstra's co-author Georges Audi summarized the history of measuring the masses of nuclides and their estimates in 2006. His work also contains many references on this story. The group publishes the latest status of the estimated atomic masses about every ten years, most recently (status 2016) in 2012 under the name Ame2012 ( Atomic mass evaluation ) in the journal Chinese Physics . The data list for this evaluation is available free of charge from some servers.

For atomic weights in the chemical sense, a Microsoft Excel 97-2003 workbook from IUPAC with the title Table of Standard Atomic Weights 2013 can be downloaded from the network. For the mixing element iron z. B. is found there as the current best value of the average mass of a neutral atom (the number in brackets indicates the uncertainty of the last digit).

Web links

Wiktionary: atomic mass  - explanations of meanings, word origins, synonyms, translations

Individual evidence

  1. a b JR de Laeter et al. : Atomic weights of the elements: Review 2000 (IUPAC technical report) . In: Pure and applied chemistry . tape 75 , no. 6 , 2003, p. 683-800 ( online [PDF; accessed March 27, 2018]). P. 687 f: “When Tomas Batuecas , President of the Atomic Weight Committee, persuaded the authorities in the IUPAC Bureau to change the term to atomic mass in 1963 , traditional chemists revolted, atomic weight was retained and Edward Wichers , who was formerly President of the Commission has been tacitly reinstated as Chairman of the Atomic Weight Commission. "
  2. ^ A b G. Audi, M. Wang, AH Wapstra, FG Kondev, M. MacCormick, X. Xu, and B. Pfeiffer: The Ame2012 atomic mass evaluation. Chinese Physics C Volume 36 (2012), pages 1287-1602, Atomic Mass Adjustment 2012 .
  3. ^ Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten: Chemistry: Studying compact , 10th, updated edition, Munich 2011, p. 51 .
  4. G. Audi, The History of Nuclidic Masses and of their Evaluation, Int. J. Mass Spectr. Ion Process. 251 (2006) 85-94, arxiv
  5. ^ F. Everling, LA König, JME Mattauch, AH Wapstra: Relative nuclidic masses . In: Nucl. Phys. A . tape 18 , 1960, p. 529-569 .
  6. G. Audi, A. Wapstra: The 1993 atomic mass evaluation: (I) Atomic mass table. Nuclear Physics A , Vol. 565 (1993) pp. 1-65, doi : 10.1016 / 0375-9474 (93) 90024-R
  7. ^ Georges Audi: The history of nuclidic masses and of their evaluation . In: International Journal of Mass Spectrometry . tape 251 , no. 2–3 , 2006, pp. 85–94 , doi : 10.1016 / j.ijms.2006.01.048 ( online [PDF; accessed December 28, 2017]).
  8. ^ IUPAC, Standard Atomic Weights Revised 2013 .