The
alchemists, and especially the early
metallurgists, probably had some notion of amount of substance, but there are no surviving records of any generalization of the idea beyond a set of recipes. In 1758,
Mikhail Lomonosov questioned the idea that mass was the only measure of the quantity of matter, but he did so only in relation to his theories on
gravitation. The development of the concept of amount of substance was coincidental with, and vital to, the birth of modern chemistry. •
1777:
Wenzel publishes
Lessons on Affinity, in which he demonstrates that the proportions of the "base component" and the "acid component" (
cation and
anion in modern terminology) remain the same during reactions between two neutral
salts. •
1789:
Lavoisier publishes
Treatise of Elementary Chemistry, introducing the concept of a
chemical element and clarifying the
law of conservation of mass for chemical reactions. •
1792:
Richter publishes the first volume of
Stoichiometry or the Art of Measuring the Chemical Elements (publication of subsequent volumes continues until 1802). The term "
stoichiometry" is used for the first time. The first tables of
equivalent weights are published for
acid–base reactions. Richter also notes that, for a given acid, the equivalent mass of the acid is proportional to the mass of oxygen in the base. • :The concept of atoms raised the question of their weight. While many were skeptical about the reality of atoms, chemists quickly found atomic weights to be an invaluable tool in expressing stoichiometric relationships. •
1808: Publication of Dalton's
A New System of Chemical Philosophy, containing the first table of
atomic weights (based on H = 1). •
1809:
Gay-Lussac's law of combining volumes, stating an integer relationship between the volumes of reactants and products in the chemical reactions of gases. •
1811:
Avogadro hypothesizes that equal volumes of different gases (at same temperature and pressure) contain equal numbers of particles, now known as
Avogadro's law. •
1813/1814:
Berzelius publishes the first of several tables of atomic weights based on the scale of
m(O) = 100. •
1815:
Prout publishes his
hypothesis that all atomic weights are integer multiple of the atomic weight of hydrogen. The hypothesis is later abandoned given the observed atomic weight of
chlorine (approx. 35.5 relative to hydrogen). •
1819:
Dulong–Petit law relating the atomic weight of a solid element to its
specific heat capacity. •
1819:
Mitscherlich's work on
crystal isomorphism allows many
chemical formulae to be clarified, resolving several ambiguities in the calculation of atomic weights. • :The
ideal gas law was the first to be discovered of many relationships between the number of atoms or molecules in a system and other physical properties of the system, apart from its mass. However, this was not sufficient to convince all scientists of the existence of atoms and molecules, many considered it simply being a useful tool for calculation. •
1834:
Faraday states his
laws of electrolysis, in particular that "the chemical decomposing action of a current is
constant for a constant quantity of electricity". •
1856:
Krönig derives the ideal gas law from
kinetic theory.
Clausius publishes an independent derivation the following year. •
1860: The
Karlsruhe Congress debates the relation between "physical molecules", "chemical molecules" and atoms, without reaching consensus. •
1865:
Loschmidt makes the first estimate of the size of gas molecules and hence of number of molecules in a given volume of gas, now known as the
Loschmidt constant. •
1886:
van't Hoff demonstrates the similarities in behaviour between dilute solutions and ideal gases. •
1886:
Eugen Goldstein observes
discrete particle rays in gas discharges, laying the foundation of
mass spectrometry, a tool subsequently used to establish the masses of atoms and molecules. •
1887:
Arrhenius describes the dissociation of
electrolyte in solution, resolving one of the problems in the study of colligative properties. •
1893: First recorded use of the term
mole to describe a unit of amount of substance by
Ostwald in a university textbook. •
1897: First recorded use of the term
mole in English. • By the
turn of the twentieth century, the concept of atomic and molecular entities was generally accepted, but many questions remained, not least the size of atoms and their number in a given sample. The concurrent development of
mass spectrometry, starting in 1886, supported the concept of atomic and molecular mass and provided a tool of direct relative measurement. •
1905:
Einstein's paper on
Brownian motion dispels any last doubts on the physical reality of atoms, and opens the way for an accurate determination of their mass. •
1909:
Perrin coins the name
Avogadro constant and estimates its value. •
1913: Discovery of
isotopes of non-radioactive elements by
Soddy and
Thomson. •
1914:
Richards receives the Nobel Prize in Chemistry for "his determinations of the atomic weight of a large number of elements". •
1920:
Aston proposes the
whole number rule, an updated version of
Prout's hypothesis. •
1921: Soddy receives the Nobel Prize in Chemistry "for his work on the chemistry of radioactive substances and investigations into isotopes". •
1922: Aston receives the Nobel Prize in Chemistry "for his discovery of isotopes in a large number of non-radioactive elements, and for his whole-number rule". •
1926: Perrin receives the
Nobel Prize in Physics, in part for his work in measuring the Avogadro constant. •
1959/
1960: Unified atomic mass unit scale based on
m(C) = 12 u adopted by
IUPAP and
IUPAC. •
1968: The mole is recommended for inclusion in the
International System of Units (SI) by the
International Committee for Weights and Measures (CIPM). •
1972: The mole is approved as the
SI base unit of amount of substance. •
2019: The mole is redefined in the SI as "the amount of substance of a system that contains specified elementary entities". == See also ==