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22-AUGUST-2008 06:13:22 - Atomic mass unit The introduction to this article may be too long. Please help by moving some material from it into the body of the article. Read the layout guide and 's lead section guidelines for more information. Discuss this issue on the talk page The unified atomic mass unit u, or dalton Da or, sometimes, universal mass unit, is a unit of mass used to express atomic and molecular masses. It is the approximate mass of a hydrogen atom, a proton, or a neutron. The precise definition is that it is one twelfth of the mass of an unbound atom of carbon-12 12C at rest and in its ground state. 1 u = 1/NA gram = 1/ 1000 NA kg where NA is Avogadro's number 1 u = 1.66053878283×10-27 kg = 931.49402723 MeV/c2 The atomic mass unit amu is an older name for the same thing, which differs slightly in definition, and differs in value by one part in 1000. In biochemistry and molecular biology, when talking about proteins, the term kilodalton is used, with the symbol kDa. Because proteins are large molecules, their masses are in kilodaltons, where one kilodalton is 1000 daltons. The unified atomic mass unit, or dalton, is not an SI unit of mass, but it is accepted for use with SI under either name. The unit is convenient because one hydrogen atom has a mass of approximately 1 u, and more generally an atom or molecule that contains n protons and neutrons will have a mass approximately equal to n u. The reason is that a 12C atom contains 6 protons, 6 neutrons and 6 electrons, with the protons and neutrons having about the same mass and the electron mass being negligible in comparison.The mass of the electron is approximately 1/1836 of the mass of the proton. This is an approximation, since it does not account for the mass contained in the binding energy of an atom's nucleus; this binding energy mass is not a fixed fraction of an atom's total mass. The differences which result from nuclear binding are generally less than 0.1 u, however. Chemical element masses, as expressed in u, would therefore all be close to whole number values within 2% and usually within 1% were it not for the fact that atomic weights of chemical elements are averaged values of the various stable isotope masses in the abundances which they naturally occur.1 For example, chlorine has an atomic weight of 35.45 u because it is composed of 76% 35Cl 34.96 u and 24% 37Cl 36.97 u. Another reason the unit is used is that it is experimentally much easier and more precise to compare masses of atoms and molecules determine relative masses than to measure their absolute masses. Masses are compared with a mass spectrometer see below. Avogadro's number NA and the mole are defined so that one mole of a substance with atomic or molecular mass 1 u will have a mass of precisely 1 g. For example, the molecular mass of a water molecule containing one 16O isotope and two 1H isotopes is 18.0106 u, and this means that one mole of this monoisotopic water has a mass of 18.0106 g. Water and most molecules consist of a mixture of molecular masses due to naturally occurring isotopes. For this reason these sort of comparisons are more meaningful and practical using molar masses which are generally expressed in g/mol, not u. In other words the one-to-one relationship between daltons and g/mol is true but in order to be used accurately for any practical purpose any calculations must be with isotopically pure substances or involve much more complicated statistical averaging of multiple isotopic compositions. Contents 1 History 2 References 3 See also 4 External links History The chemist John Dalton was the first to suggest the mass of one atom of hydrogen as the atomic mass unit. Francis Aston, inventor of the mass spectrometer, later used 1â?„16 of the mass of one atom of oxygen-16 as his unit. Before 1961, the physical atomic mass unit amu was defined as 1â?„16 of the mass of one atom of oxygen-16, while the chemical atomic mass unit amu was defined as 1â?„16 of the average mass of an oxygen atom taking the natural abundance of the different oxygen isotopes into account. Both units are slightly smaller than the unified atomic mass unit, which was adopted by the International Union of Pure and Applied Physics in 1960 and by the International Union of Pure and Applied Chemistry in 1961. Hence, before 1961 physicists as well as chemists used the symbol amu for their respective and slightly different atomic mass units. One still sometimes finds this usage in the scientific literature today. However, the accepted standard is now the unified atomic mass unit symbol u, with: 1 u = 1.000 317 9 amu physical scale = 1.000 043 amu chemical scale. Since 1961, by definition the unified atomic mass unit is equal to one-twelfth of the mass of a carbon-12 atom. References ^ Exact Masses and Isotopic Abundances of the Elements - Alphabet See also Molecular mass Mass-to-charge ratio External links SI website on acceptable non-SI units Accepted value of 1u as of 2006 atomic mass unit Retrieved from http://en..org/wiki/Atomic_mass_unit Categories: Nuclear chemistry | Units of massHidden categories: introduction cleanup | All pages needing cleanup Views Article Discussion this page History Personal tools Log in / create account Navigation Main page Contents Featured content Current events Random article Search Go Search Interaction Community portal Recent changes Contact Donate to Help Toolbox What links here Related changes Upload file Special pages Printable version Permanent link Cite this page Languages العربية Asturianu Bosanski Brezhoneg БългарÑ?ки Català Česky Dansk Deutsch Eesti Español Esperanto Français í•œêµì–´ हिनà¥?दी Hrvatski Bahasa Indonesia Italiano עברית Lietuvių Magyar МакедонÑ?ки Nederlands 日本語 ‪Norsk bokmÃ¥l‬ ‪Norsk nynorsk‬ O'zbek Plattdüütsch Polski Português Română РуÑ?Ñ?кий Simple English SlovenÄ?ina SlovenÅ¡Ä?ina СрпÑ?ки / Srpski Srpskohrvatski / СрпÑ?кохрватÑ?ки Suomi Svenska ไทย Tiếng Việt Türkçe УкраїнÑ?ька ä¸æ–‡ This page was last modified on 21 August 2008, at 23:53. of the GNU Free Documentation License. ® , Inc., a U.S.
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