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Atomic mass unit
From Wikipedia, the free encyclopedia
The unified atomic mass unit or atomic mass unit (u), or dalton (Da) or, sometimes, universal mass unit (u), 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.
Values of 1 u
Contents
Units
1.660 538 782(83) × 10−24 g
1 Definition
2 History
3 See also
4 Notes
5 External links
1.660 538 782(83) × 10−27 kg
931.494 028(23) × 106
eV/c2
For details, see below.
Definition
The precise definition is that the atomic mass unit (u) is one twelfth of the mass of an isolated atom of carbon-12 (12C) at rest and in its ground
state.[1]
In other words, "A single atom of carbon-12 has a mass of 12 u exactly, by definition." Given that Mu is the Molar mass constant in and NA is
the Avogadro constant, the relationship between the atomic mass unit (u) and the gram (g) is
or
and
[2].
The atomic mass unit is also the name for two older units (both having the symbol amu) that are very similar to the modern atomic mass unit
(see History). The symbol amu is not a symbol for the unified atomic mass unit. Its use is a historical artifact (written during the time when the
amu scales were used), an error (possibly deriving from confusion about historical usage), or a correct reference to the historical scales that used
it. 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 mass of molecules, the term "dalton" is used, with the symbol Da. Because proteins
are large molecules, their masses are often 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.[3]
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 0.01 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 masses of chemical elements are averaged values of the various
stable isotope masses in the abundances which they naturally occur.[4] For example, chlorine has an atomic mass 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 atom and two 1H atoms 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 sorts 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 1⁄16 of the mass of one atom of oxygen-16 as his unit.
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Atomic mass unit - Wikipedia, the free encyclopedia
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Before 1961, the physical atomic mass unit (amu) was defined as 1⁄16 of the mass of one atom of oxygen-16, while the chemical atomic mass
unit (amu) was defined as 1⁄16 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.
See also
Mass-to-charge ratio
Notes
1.
2.
3.
4.
^ NIST CODATA "Unified atomic mass unit"
^ CODATA kilogram-atomic mass unit relationship
^ Non-SI units accepted for use with the SI, and units based on fundamental constants - International Bureau of Weights and Measures (BIPM)
^ Exact Masses and Isotopic Abundances of the Elements - Alphabet
External links
Accepted value of 1u as of 2006
atomic mass unit
Retrieved from "http://en.wikipedia.org/wiki/Atomic_mass_unit"
Categories: Nuclear chemistry | Units of mass
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