, 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
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 weights (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). The related concept of
equivalent mass had been in use at least a century earlier. In chemistry, it has been known since
Proust's law of definite proportions (1794) that knowledge of the mass of each of the components in a chemical
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
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.
Standardization Developments in
mass spectrometry led to the adoption of
oxygen-16 as the standard substance, in lieu of natural oxygen. 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 kilograms of carbon-12." Thus, by that definition, one mole of pure 12C had a mass of
exactly 12
g. The four different definitions were equivalent to within 1%. Because a
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
NA (the Avogadro constant) had to be determined experimentally. The experimental value adopted by
CODATA in 2010 is . In 2011 the measurement was refined to . The mole was made the seventh
SI base unit in 1971 by the 14th CGPM.
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
grams of
carbon-12 (the most common
isotope of carbon). The term
gram-molecule was formerly used to mean one mole of molecules, and
gram-atom for one mole of atoms. For example, 1 mole of
MgBr2 is 1 gram-molecule of MgBr2 but 3 gram-atoms of MgBr2. 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 constants that are, in their nature, exact. == Criticism ==