Editing Dalton (unit) (section)

== History ==

=== Origin of the concept ===

[[File:Jean Perrin 1926.jpg|right|thumb|Jean Perrin in 1926]]
The interpretation of the [[law of definite proportions]] in terms of the [[atomic theory of matter]] implied that the masses of atoms of various elements had definite ratios that depended on the elements.  While the actual masses were unknown, the relative masses could be deduced from that law.  In 1803 [[John Dalton]] proposed to use the (still unknown) atomic mass of the lightest atom, hydrogen, as the natural unit of atomic mass.  This was the basis of the [[standard atomic weight|atomic weight scale]].<ref name=petley1989/>

For technical reasons, in 1898, chemist [[Wilhelm Ostwald]] and others proposed to redefine the unit of atomic mass as {{sfrac|1|16}} the mass of an oxygen atom.<ref name=hold2004/> That proposal was formally adopted by the [[Commission on Isotopic Abundances and Atomic Weights|International Committee on Atomic Weights]] (ICAW) in 1903. That was approximately the mass of one hydrogen atom, but oxygen was more amenable to experimental determination. This suggestion was made before the discovery of isotopes in 1912.<ref name=petley1989/> Physicist [[Jean Baptiste Perrin|Jean Perrin]] had adopted the same definition in 1909 during his experiments to determine the atomic masses and the [[Avogadro constant]].<ref name=perrin1909/> This definition remained unchanged until 1961.<ref name=chang2005/><ref name=kelt2008/> Perrin also defined the "mole" as an amount of a compound that contained as many molecules as 32 grams of oxygen ({{chem|O|2}}). He called that number the [[Avogadro number]] in honor of physicist [[Amedeo Avogadro]].

=== Isotopic variation ===
The discovery of isotopes of oxygen in 1929 required a more precise definition of the unit. Two distinct definitions came into use. Chemists choose to define the AMU as {{sfrac|1|16}} of the average mass of an oxygen atom as found in nature; that is, the average of the masses of the known isotopes, weighted by their natural abundance. Physicists, on the other hand, defined it as {{sfrac|1|16}} of the mass of an atom of the isotope oxygen-16 (<sup>16</sup>O).<ref name=hold2004/>

=== Joint definition by IUPAP and IUPAC ===

The existence of two distinct units with the same name was confusing, and the difference (about {{val|1.000282}} in relative terms) was large enough to affect high-precision measurements. Moreover, it was discovered that the isotopes of oxygen had different natural abundances in water and in air. 

In April 1957 [[Alfred O. C. Nier]] suggested to [[Josef Mattauch]] that the [[carbon-12]] be adopted as mass scale because of carbon's use as a secondary standard in [[mass spectrometry]]. Also, carbon-12 implied acceptable relative changes in the atomic weight scale, i.e., 42 parts-per-million (ppm) compared to 275 ppm for [[oxygen-16]] which would not acceptable to chemists. 

Following the approval of the [[International Union of Pure and Applied Physics]] (IUPAP) General Assembly at Ottawa, Canada in 1960 and the [[International Union of Pure and Applied Chemistry]] (IUPAC) General Assembly at Montreal, Canada in 1961, the atomic weights were officially given on the carbon-12 scale for the first time.<ref name=petley1989/><ref name=hold2004/>

The new unit was named the "unified atomic mass unit" and given a new symbol "u", to replace the old "amu" that had been used for the oxygen-based unit.<ref name=goldbUnAtMaUn/> However, the old symbol "amu" has sometimes been used, after 1961, to refer to the new unit, particularly in lay and preparatory contexts.

With this new definition, the [[standard atomic weight]] of [[carbon]] is about {{val|12.011|u=Da}}, and that of oxygen is about {{val|15.999|u=Da}}. These values, generally used in chemistry, are based on averages of many samples from [[Earth's crust]], its [[atmosphere]], and [[organic materials]].

=== Adoption by BIPM ===

The IUPAC 1961 definition of the unified atomic mass unit, with that name and symbol "u", was adopted by the [[International Bureau for Weights and Measures]] (BIPM) in 1971 as a [[non-SI unit accepted for use with the SI]].<ref name=bipm1971/>

=== Unit name ===
In 1993, the IUPAC proposed the shorter name "dalton" (with symbol "Da") for the unified atomic mass unit.<ref name=mills1993/><ref name=goldbDa/> As with other unit names such as watt and newton, "dalton" is not capitalized in English, but its symbol, "Da", is capitalized. The name was endorsed by the [[International Union of Pure and Applied Physics]] (IUPAP) in 2005.<ref name=iupap2005/>

In 2003 the name was recommended to the BIPM by the [[Consultative Committee for Units]], part of the [[CIPM]], as it "is shorter and works better with [SI] prefixes".<ref name=bipm-ccu15/> In 2006, the BIPM included the dalton in its 8th edition of the [[SI]] brochure of formal definitions as a [[non-SI units accepted for use with the SI|non-SI unit accepted for use with the SI]].<ref name=bipm8th/> The name was also listed as an alternative to "unified atomic mass unit" by the [[International Organization for Standardization]] in 2009.<ref name=ISO1/><ref name=ISO10/> It is now recommended by several scientific publishers,<ref name=oxsty/> and some of them consider "atomic mass unit" and "amu" deprecated.<ref name=rapsty/> In 2019, the BIPM retained the dalton in its 9th edition of the [[SI]] brochure, while dropping the unified atomic mass unit from its table of non-SI units accepted for use with the SI, but secondarily notes that the dalton&nbsp;(Da) and the unified atomic mass unit&nbsp;(u) are alternative names (and symbols) for the same unit.<ref name=bipm9th/>

=== 2019 revision of the SI ===
The definition of the dalton was not affected by the [[2019 revision of the SI]],<ref name=cipm106/><ref name=cgpm26/><ref name=bipm9th/> that is, 1&nbsp;Da in the SI is still {{sfrac|1|12}} of the mass of a carbon-12 atom, a quantity that must be determined experimentally in terms of SI units. However, the definition of a mole was changed to be the amount of substance consisting of exactly {{physconst|NA|unit=no|ref=no}} entities and the definition of the kilogram was changed as well. As a consequence, the [[molar mass constant]] remains close to but no longer exactly 1&nbsp;g/mol, meaning that the mass in grams of one mole of any substance remains nearly but no longer exactly numerically equal to its average molecular mass in daltons,<ref name=lehm2016/> although the relative standard uncertainty of {{val|4.5|e=-10}} at the time of the redefinition is insignificant for all practical purposes.<ref name="bipm9th" />  One entity, symbol ent, is the smallest amount of any substance (retaining its chemical properties). One mole is an aggregate of an Avogadro number of entities, 1 mol = ''N''<sub>0</sub> ent. This means that the appropriate atomic-scale unit for molar mass is dalton per entity, Da/ent = ''M''<sub>u</sub>, very nearly equal to 1 g/mol. For Da/ent to be exactly equal to g/mol, the dalton would need to be redefined exactly in terms of the (fixed-''h'') kilogram. Then, in addition to the identity g = (g/Da) Da, we would have the parallel relationship mol = (g/Da) ent = ''N''<sub>0</sub> ent, conforming to the original mole concept—that the Avogadro number is the gram-to-dalton mass unit ratio.