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## Atomic mass
The The The The ## Additional recommended knowledge
## Mass defects in atomic massesThe pattern in the amounts the atomic masses deviate from their mass numbers is as follows: the deviation starts positive at hydrogen-1, becomes negative until a minimum is reached at iron-56, iron-58 and nickel-62, then increases to positive values in the heavy isotopes, with increasing atomic number. This corresponds to the following: nuclear fission in an element heavier than iron produces energy, and fission in any element lighter than iron requires energy. The opposite is true of nuclear fusion reactions: fusion in elements lighter than iron produces energy, and fusion in elements heavier than iron requires energy. ## Measurement of atomic massesDirect comparison and measurement of the masses of atoms is achieved with mass spectrometry. ## Conversion factor between atomic mass units and gramsThe standard scientific unit for dealing with atoms in macroscopic quantities is the mole (mol), which is defined arbitrarily as the amount of a substance with as many atoms or other units as there are in 12 grams of the carbon isotope C-12. The number of atoms in a mole is called Avogadro's number, the value of which is approximately 6.022 × 10 The formulaic conversion between atomic mass and SI mass in grams for a single atom is: where u is the atomic mass unit and ## Relationship between atomic and molecular massesSimilar definitions apply to molecules. One can compute the molecular mass of a compound by adding the atomic masses of its constituent atoms (nuclides). One can compute the molar mass of a compound by adding the relative atomic masses of the elements given in the chemical formula. In both cases the multiplicity of the atoms (the number of times it occurs) must be taken into account, usually by multiplication of each unique mass by its multiplicity. ## HistoryIn the history of chemistry the first scientists to determine atomic weights were John Dalton between 1803 and 1805 and Jöns Jakob Berzelius between 1808 and 1826. Atomic weight was originally defined relative to that of the lightest element hydrogen taken as 1.00. Stanislao Cannizzaro in the 1860's refined atomic weights by applying Avogadro's law. He formulated a law to determine atomic weights of elements: In the early twentieth century, up until the 1960's chemists and physicists used two different atomic mass scales. The chemists used a scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to the atomic mass of the most common oxygen isotope (containing eight protons and eight neutrons). However, because oxygen-17 and oxygen-18 are also present in natural oxygen this led to two different tables of atomic mass. The unified scale based on carbon-12, The term ## Table of standard atomic weights
## See also- Tutorial on the concept and measurement of atomic mass
- Atomic Weights and the International Committee — A Historical Review
## References**^**IUPAC Definition of Atomic Mass**^**IUPAC Definition of Relative Atomic Mass**^**IUPAC Definition of Standard Atomic Weight**^**ATOMIC WEIGHTS OF THE ELEMENTS 2005 (IUPAC TECHNICAL REPORT), M. E. WIESER Pure Appl. Chem., V.78, pp. 2051, 2006**^***Origin of the Formulas of Dihydrogen and Other Simple Molecules*Andrew Williams Vol. 84 No. 11 November 2007 • Journal of Chemical Education 1779**^**'ATOMIC WEIGHT' -THE NAME, ITS HISTORY, DEFINITION, AND UNITS, P. DE BIEVRE and H. S. PEISER Pure&App. Chem., 64, 1535, 1992
Categories: Chemical properties | Stoichiometry |
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This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Atomic_mass". A list of authors is available in Wikipedia. |