A mole is the amount of substance that contains L particles of that substance. It is SI unit of the amount of substance since 1971. A mole of substance is defined as 6.02 x 1023 (also known as Avogadro constant and shown with L or NA).
A mole is found with a mass of substance divided by molar mass. Molar mass is the mass of one mole of any substances. For instance one mole of oxygen’s molar mass is 16 gmol-1.
The term of molar (from the Latin moles, meaning a large mass) was first introduced into chemistry by the German chemist, August Wilhelm Hofmann around 1865. It was originally intended to indicate any large macroscopic mass, in contrast to a submicroscopic or “molecular” mass. The term of mole which we use, was first introduced into chemistry by German physical chemist, Wilhelm Ostwald. However Ostwald used this term to mean “molecular weight in gram”. In 1900, Max Planck determined the value of Avogadro constant with this explanation: “To one gram molecule of a substance correspond 6,175 x 1023 molecules”.
Swedish chemist Jöns Jacob Berzelius 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. However in the 1960's the oxygen-16 definition was replaced with one based on carbon-12. And in 1971, the mole’s definition was changed by the International Bureau of Weights and Measures as “the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12.”. With this definition, one mole of pure 12C had a mass of exactly 12 grams. Also in 2011, the value of Avogadro constant was refined to 6.02 x 1023. In 2018, scientists decided that SI units would not be defined in terms of physical objects but they would be defined by constants. In order to that, the mole were redefined on the basis of a constant which is 6.02 x 1023 (the Avogadro constant) at the 26th General Conference on Weights and Measures (CGPM).
In chemistry, the concept of the mole helps to put quantitative information about what happens in a chemical equation on a macroscopic level. Also they use the mole to measure small amounts of atoms and molecules. However we do not use the mole only in chemistry. We use the mole for different areas and usages. In medicine, for example, we use the mole for figuring out how much of each ingredient is needed to react together while we are making drugs. Another example of using the mole is combustion. We use the mole in combustion to know how much air is needed, how much exhaust would be produced and how much heat is created. There are other examples such as plastics and battery production.
References
Jensen, William B. 2009. “Ask the Historian - The Origin of the Mole Concept.” Cincinnati: Department of Chemistry, University of Cincinnati.
Neuss, Geoffrey. 2014. “Stoichiometric Relationships.” Essay. In Chemistry for the IB Diploma: 2014 Edition., 2–2. Oxford: Oxford University Press.
Marquardt, Roberto. 2019. “The Mole and IUPAC: a Brief History.” De Gruyter. De Gruyter. July 1, 2019. https://www.degruyter.com/view/journals/ci/41/3/article-p50.xml?language=en.
“Mole (Unit).” 2020. Wikipedia. Wikimedia Foundation. September 14, 2020. https://en.wikipedia.org/wiki/Mole_(unit).
“The Mole.” 2019. LNE, Laboratoire National De Métrologie Et D'essais. 2019. https://www.lne.fr/en/learn-more/international-system-units/mole.
The Editors of Encyclopaedia Britannica. 2020. “Mole.” Encyclopædia Britannica. Encyclopædia Britannica, inc. August 14, 2020. https://www.britannica.com/science/mole-chemistry.
Al-Neaimi, Mouza, and Hamda Al-Mansoori. 2015. “How Do We Use Moles Everyday?” Prezi.
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