Editing Mole (unit) (section)

== History ==
[[File:Amadeo Avogadro.png|thumb|upright|[[Amedeo Avogadro|Avogadro]], who inspired the Avogadro constant]]
The history of the mole is intertwined with that of units of [[molecular mass]], and the [[Avogadro constant]].

The first table of [[standard atomic weight]] was published by [[John Dalton]] (1766–1844) in 1805, based on a system in which the relative atomic mass of [[hydrogen]] was defined as 1. These relative atomic masses were based on the [[Stoichiometry|stoichiometric]] proportions of chemical reaction and compounds, a fact that greatly aided their acceptance: It was not necessary for a chemist to subscribe to [[atomic theory]] (an unproven hypothesis at the time) to make practical use of the tables. This would lead to some confusion between atomic masses (promoted by proponents of atomic theory) and [[equivalent weight]]s (promoted by its opponents and which sometimes differed from relative atomic masses by an integer factor), which would last throughout much of the nineteenth century.

[[Jöns Jacob Berzelius]] (1779–1848) was instrumental in the determination of relative atomic masses to ever-increasing accuracy. He was also the first chemist to use [[oxygen]] as the standard to which other masses were referred. Oxygen is a useful standard, as, unlike hydrogen, it forms compounds with most other elements, especially [[metals]]. However, he chose to fix the atomic mass of oxygen as 100, which did not catch on.

[[Charles Frédéric Gerhardt]] (1816–56), [[Henri Victor Regnault]] (1810–78) and [[Stanislao Cannizzaro]] (1826–1910) expanded on Berzelius' works, resolving many of the problems of unknown stoichiometry of compounds, and the use of atomic masses attracted a large consensus by the time of the [[Karlsruhe Congress]] (1860). The convention had reverted to defining the atomic mass of hydrogen as 1, although at the level of precision of measurements at that time – relative uncertainties of around 1% – this was numerically equivalent to the later standard of oxygen = 16. However the chemical convenience of having oxygen as the primary atomic mass standard became ever more evident with advances in analytical chemistry and the need for ever more accurate atomic mass determinations.

The name ''mole'' is an 1897 translation of the German unit ''Mol'', coined by the [[chemist]] [[Wilhelm Ostwald]] in 1894 from the German word ''Molekül'' ([[molecule]]).<ref>
{{Cite book
|last=Helm |first=Georg
|year=1897
|title=The Principles of Mathematical Chemistry: The Energetics of Chemical Phenomena
|url=https://archive.org/details/principlesmathe00helmgoog
|others=transl. by Livingston, J.; Morgan, R.
|place=New York
|publisher=Wiley
|page=[https://archive.org/details/principlesmathe00helmgoog/page/n20 6]
}}</ref><ref>Some sources place the date of first usage in English as 1902. [[Merriam–Webster]] [http://www.merriam-webster.com/dictionary/mole%5B5%5D proposes] {{webarchive|url=https://web.archive.org/web/20111102181728/http://www.merriam-webster.com/dictionary/mole%5B5%5D |date=2011-11-02 }} an etymology from ''Molekulärgewicht'' ([[molecular weight]]).</ref><ref>{{cite book
|last=Ostwald |first=Wilhelm |author-link=Wilhelm Ostwald
|year=1893
|publisher=Wilhelm Engelmann
|location=Leipzig, Germany
|title=Hand- und Hilfsbuch zur Ausführung Physiko-Chemischer Messungen
|trans-title=Handbook and Auxiliary Book for Conducting Physico-Chemical Measurements
|page=119
|url=https://babel.hathitrust.org/cgi/pt?id=uc1.b4584562;view=1up;seq=131
}} From p. 119: ''"Nennen wir allgemein das Gewicht in Grammen, welches dem Molekulargewicht eines gegebenen Stoffes numerisch gleich ist, ein Mol, so ... "''  (If we call in general the weight in grams, which is numerically equal to the molecular weight of a given substance, a "mol", then ... )</ref> The related concept of [[equivalent mass]] had been in use at least a century earlier.<ref>'''mole''', '''''n.{{sup|8}}''''', [[Oxford English Dictionary]], Draft Revision Dec. 2008</ref>

In chemistry, it has been known since [[Joseph Proust|Proust's]] [[law of definite proportions]] (1794) that knowledge of the mass of each of the components in a chemical [[system (thermodynamics)|system]] is not sufficient to define the system. Amount of substance can be described as mass divided by Proust's "definite proportions", and contains information that is missing from the measurement of mass alone. As demonstrated by [[John Dalton|Dalton's]] [[law of partial pressures]] (1803), a measurement of mass is not even necessary to measure the amount of substance (although in practice it is usual). There are many physical relationships between amount of substance and other physical quantities, the most notable one being the [[ideal gas law]] (where the relationship was first demonstrated in 1857). The term "mole" was first used in a textbook describing these [[colligative properties]].<ref>{{cite book |last=Ostwald |first=Wilhelm |others=McGowan, George (transl.) |url=https://openlibrary.org/books/OL7204743M/The_scientific_foundations_of_analytical_chemistry | ol=7204743M | title=The Scientific Foundations of Analytical Chemistry: Treated in an Elementary Manner | year=1900 |edition=Second English |location=London | publisher=Macmillan and Co. }}</ref>

=== Standardization ===
Developments in [[mass spectrometry]] led to the adoption of [[oxygen-16]] as the standard substance, in lieu of natural oxygen.<ref>{{cite journal |last1=Busch |first1=Kenneth |title=Units in Mass Spectrometry |journal=Current Trends in Mass Spectrometry |date=May 2, 2003 |volume=18 |issue=5S |pages=S32-S34 [S33] |url=https://cdn.sanity.io/files/0vv8moc6/spectroscopy/cfb6f4cb3d02243b516bce3b11dc3584733be2b1.pdf/article-55961.pdf |access-date=29 April 2023}}</ref>

The oxygen-16 definition was replaced with one based on carbon-12 during the 1960s. The International Bureau of Weights and Measures defined the mole as "the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012&nbsp;kilograms of carbon-12." Thus, by that definition, one mole of pure <sup>12</sup>C had a mass of ''exactly'' 12&nbsp;[[Gram|g]].<ref name="SI114-15" /><ref name="IUPAChist">{{AtomicWeightHistory}}</ref> The four different definitions were equivalent to within 1%.

{|class="wikitable" align="center" style="margin:.5em;"
! Scale basis
! Scale basis<br />relative to {{sup|12}}C = 12
! Relative deviation<br />from the {{sup|12}}C = 12 scale
|-
| Atomic mass of hydrogen = 1
| 1.00794(7)
| align="center" | −0.788%
|-
| Atomic mass of oxygen = 16
| {{val|15.9994|(3)}}
| align="center" | +0.00375%
|-
| Relative atomic mass of {{sup|16}}O = 16
| {{val|15.9949146221|(15)}}
| align="center" | +0.0318%
|-
|}

Because a [[Dalton (unit)|dalton]], a unit commonly used to measure [[atomic mass]], is exactly 1/12 of the mass of a carbon-12 atom, this definition of the mole entailed that the mass of one mole of a compound or element in grams was numerically equal to the average mass of one molecule or atom of the substance in daltons, and that the number of daltons in a gram was equal to the number of elementary entities in a mole. Because the mass of a [[nucleon]] (i.e. a [[proton]] or [[neutron]]) is approximately 1 dalton and the nucleons in an atom's nucleus make up the overwhelming majority of its mass, this definition also entailed that the mass of one mole of a substance was roughly equivalent to the number of nucleons in one atom or molecule of that substance.

Since the definition of the gram was not mathematically tied to that of the dalton, the number of molecules per mole ''N''<sub>A</sub> (the Avogadro constant) had to be determined experimentally. The experimental value adopted by [[Committee on Data for Science and Technology|CODATA]] in 2010 is {{nowrap|1=''N''<sub>A</sub> = {{val|6.02214129|(27)|e=23|u=mol-1}}}}.<ref>[http://physics.nist.gov/cgi-bin/cuu/Value?na physics.nist.gov/] {{webarchive|url=https://web.archive.org/web/20150629063615/http://physics.nist.gov/cgi-bin/cuu/Value?na |date=2015-06-29 }} Fundamental Physical Constants: Avogadro Constant</ref>
In 2011 the measurement was refined to {{val|6.02214078|(18)|e=23|u=mol-1}}.<ref>{{cite journal | first = Birk | last = Andreas | title = Determination of the Avogadro Constant by Counting the Atoms in a <sup>28</sup>Si Crystal | journal=Physical Review Letters | volume = 106 | issue = 3 | year=2011 | pages = 30801 | doi=10.1103/PhysRevLett.106.030801 | pmid = 21405263 | bibcode=2011PhRvL.106c0801A|arxiv = 1010.2317 | s2cid = 18291648 |display-authors=etal}}</ref>

The mole was made the seventh [[SI base unit]] in 1971 by the 14th CGPM.<ref>{{cite web|url=http://www.bipm.org/en/CGPM/db/14/3/|title=BIPM – Resolution 3 of the 14th CGPM|website=www.bipm.org|access-date=1 May 2018|url-status=dead|archive-url=https://web.archive.org/web/20171009112117/http://www.bipm.org/en/CGPM/db/14/3|archive-date=9 October 2017}}</ref>

=== 2019 revision of the SI ===

Before the [[2019 revision of the SI]], the mole was defined as the amount of substance of a system that contains as many elementary entities as there are atoms in 12 [[gram]]s of [[carbon-12]] (the most common [[isotopes of carbon|isotope of carbon]]).<ref name=SI8>{{SIbrochure8th}}</ref>
The term ''gram-molecule'' was formerly used to mean one mole of molecules, and ''gram-atom'' for one mole of atoms.<ref name="SI114-15">{{SIbrochure8th|pages=114–15}}</ref> For example, 1 mole of [[MgBr2|MgBr<sub>2</sub>]] is 1&nbsp;gram-molecule of MgBr<sub>2</sub> but 3&nbsp;gram-atoms of MgBr<sub>2</sub>.<ref>
{{cite journal
| doi=10.1088/0953-8984/15/6/315
| last1=Wang | first1=Yuxing
| last2=Bouquet | first2= Frédéric
| last3=Sheikin | first3=Ilya
| last4=Toulemonde | first4=Pierre
| last5=Revaz | first5=Bernard
| last6=Eisterer | first6=Michael
| last7=Weber | first7=Harald W.
| last8=Hinderer | first8=Joerg
| last9=Junod | first9=Alain
| display-authors=etal
| title=Specific heat of MgB<sub>2</sub> after irradiation
| journal=Journal of Physics: Condensed Matter | year=2003
| volume=15
| issue=6 | pages=883–893|arxiv = cond-mat/0208169 |bibcode = 2003JPCM...15..883W| s2cid=16981008 }}</ref><ref>{{cite journal
| doi=10.1103/PhysRevB.72.024547
| last1=Lortz | first1=R.
| last2=Wang | first2=Y.
| last3=Abe | first3=S.
| last4=Meingast | first4=C.
| last5=Paderno | first5=Yu.
| last6=Filippov | first6=V.
| last7=Junod | first7=A.
| display-authors=etal
| title=Specific heat, magnetic susceptibility, resistivity and thermal expansion of the superconductor ZrB<sub>12</sub>
| journal=Phys. Rev. B | year=2005
| volume=72
| issue=2 | pages=024547
| arxiv = cond-mat/0502193
| bibcode = 2005PhRvB..72b4547L
| s2cid=38571250
}}</ref>

In 2011, the 24th meeting of the [[General Conference on Weights and Measures]] (CGPM) agreed to a plan for a possible revision of the [[SI base unit]] definitions at an undetermined date.

On 16 November 2018, after a meeting of scientists from more than 60 countries at the CGPM in Versailles, France, all SI base units were defined in terms of physical constants. This meant that each SI unit, including the mole, would not be defined in terms of any physical objects but rather they would be defined by [[physical constant]]s that are, in their nature, exact.<ref name="IUPACrev" />

Such changes officially came into effect on 20 May 2019. Following such changes, "one mole" of a substance was redefined as containing "exactly {{val|6.02214076|e=23}} elementary entities" of that substance.<ref>[https://www.bipm.org/utils/en/pdf/CIPM/CIPM2017-EN.pdf?page=23 CIPM Report of 106th Meeting] {{webarchive|url=https://web.archive.org/web/20180127202612/https://www.bipm.org/utils/en/pdf/CIPM/CIPM2017-EN.pdf?page=23 |date=2018-01-27 }} Retrieved 7 April 2018</ref><ref>{{cite journal |title=Redefining the Mole |url=https://www.nist.gov/si-redefinition/redefining-mole |journal=NIST |access-date=24 October 2018|date=2018-10-23 }}</ref>