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Atomic mass unit
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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.
precise definition is that it is one twelfth of the mass of an unbound atom of carbon-12 at rest and in its ground state.
- 1 u = 1/NA gram = 1/ (1000 NA) kg (where NA is Avogadro's number)
- 1 u = =
The atomic mass unit (amu) is an older name for a very similar scale (unified atomic mass unit, dalton, or universal mass unit).
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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.
Definition
The precise definition is that it is one twelfth of the mass of an unbound atom of carbon-12 at rest and in its ground state.
- 1 u = 1/NA gram = 1/ (1000 NA) kg (where NA is Avogadro's number)
- 1 u = =
The atomic mass unit (amu) is an older name for a very similar scale (unified atomic mass unit, dalton, or universal mass unit). The symbol amu for atomic mass unit is not a symbol for the unified atomic mass unit. Its use is an historical artifact (written during the time when the amu scales were used), an error (possibly deriving from confusion about historical usage), or correctly referring to the historical scales that used it (see History). Atomic masses are often written without any unit and then the unified atomic mass unit is implied.
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 International System of Units (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 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 0.01 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. For example, chlorine has an atomic weight of because it is composed of 76% and 24% .
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 will have a mass of precisely .
For example, the molecular mass of a water molecule containing one isotope and two 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.
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 of the mass of one atom of oxygen-16 as his unit.
Before 1961, the physical atomic mass unit (amu) was defined as of the mass of one atom of oxygen-16, while the chemical atomic mass unit (amu) was defined as 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.
See also
External links
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