Unified atomic mass unit: Difference between revisions

Revision as of 05:01, 1 December 2007

The unified atomic mass unit (u), or dalton (Da), is a unit of atomic and molecular mass. By definition it is one twelfth of the mass of an unbound atom of the carbon-12 (¹²C) nuclide, at rest and in its ground state.

The relationship of the unified atomic mass unit to the macroscopic SI standard of mass, the kilogram, is given by Avogadro's number N_A. By the definition of Avogadro's number, the mass of N_A carbon-12 atoms, at rest and in their ground state, is 12 gram ( = 12×10⁻³ kg). From the latest value of N_A follows the latest value^[1] of one unified mass unit:

1 u ≈ 1.660538782(83) × 10⁻²⁷ kg ≈ 931.494028(23) MeV/c²

Future refinements in Avogadro's number by future improvements in counting large (on the order of 10²⁷) numbers of atoms, will give better accuracy of u.

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.

In the literature one still finds the obsolete unit amu (atomic mass unit). This is deplorable, not so much because it is not a supported SI unit, but because historically their are two different standard masses denoted by amu. There is the physics amu ( = 1/1.000 317 9 u) and there is the chemistry amu ( = 1/1.000 043 u). This difference arose from the fact that before 1960, 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. The chemist John Dalton introduced 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.

External links

↑ CODATA

[1] CODATA

[1]

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-The '''unified atomic mass unit''' ('''u'''), or '''dalton''' ('''Da'''), is a unit of atomic and molecular [[mass]]. By definition it is one twelfth of the mass of an unbound atom of the [[carbon-12]] [[nuclide]], at rest and in its ground state.
+The '''unified atomic mass unit''' ('''u'''), or '''dalton''' ('''Da'''), is a unit of atomic and molecular mass. By definition it is one twelfth of the mass of an unbound atom of the [[carbon]]-12 (<sup>12</sup>C) [[nuclide]], at rest and in its ground state.
+The relationship of the unified atomic mass unit to the macroscopic [[SI]] standard of mass, the [[kilogram]], is given by [[Avogadro's number]] ''N''<sub>A</sub>. By the definition of Avogadro's number, the mass of  ''N''<sub>A</sub>  carbon-12 atoms, at rest and in their ground state, is 12 gram ( = 12&times;10<sup>&minus;3</sup> kg). From the latest value of ''N''<sub>A</sub> follows the latest value<ref>CODATA</ref> of one unified mass unit:
+:1 u ≈ 1.660538782(83) × 10<sup>&minus;27</sup> kg ≈ 931.494028(23) MeV/c<sup>2</sup>
-<!--
+Future refinements in Avogadro's number by future improvements in counting large (on the order of 10<sup>27</sup>) numbers  of atoms, will give better accuracy of u.
-:1 u =  1/''N''<sub>A</sub> [[gram]] = 1/ (1000 ''N''<sub>A</sub>) [[kilogram|kg]] &nbsp;&nbsp;(where ''N''<sub>A</sub> is [[Avogadro's number]])
-:1 u ≈ 1.660538782(83) × 10<sup>−27</sup> kg ≈ 931.494028(23) MeV/c<sup>2</sup>
-See [[1 E-27 kg]] for a list of objects which have a mass of about 1 u.
+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'' [[proton]]s and [[neutron]]s will have a mass approximately equal to ''n'' u.
-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|History]]). Atomic masses are often written without any unit and then the unified atomic mass unit is implied.
+In the literature one still finds the obsolete unit amu (atomic mass unit). This is deplorable, not so much because it is not a supported [[SI]] unit, but because historically their are two different standard masses denoted by amu. There is the physics amu ( = 1/1.000 317 9 u) and there is the chemistry amu ( = 1/1.000 043 u). This difference arose from the fact that before 1960, 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 [[isotope]]s 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. The chemist [[John Dalton]] introduced 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.
-In [[biochemistry]] and [[molecular biology]] literature (particularly in reference to [[protein]]s), the term "dalton" is used, with the symbol '''Da'''.  Because proteins are large [[molecule]]s, they are typically referred to in kilodaltons, or "kDa", with one kilodalton being equal to 1000 daltons.
-The unified atomic mass unit, or dalton, is not an [[SI]] unit of mass, although 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'' [[proton]]s and [[neutron]]s will have a mass approximately equal to ''n'' u. (The reason is that a carbon-12 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 [[atomic nucleus|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.<ref>http://www.sisweb.com/referenc/source/exactmaa.htm</ref> For example, [[chlorine]] has an atomic weight of 35.45 u because it is composed of 76% <sup>35</sup>Cl (34.96 u) and 24% <sup>37</sup>Cl (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]] (''N''<sub>A</sub>) and the [[mole (unit)|mole]] are defined so that one mole of a substance with atomic or molecular mass 1 u will have a mass of precisely 1 [[gram]].
-For example, the molecular mass of a [[Water (molecule)|water molecule]] containing one <sup>16</sup>O isotope and two <sup>1</sup>H isotopes is 18.0106 u, and this means that one mole of this monoisotopic water has a mass of 18.0106 grams. 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 mass]]es 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 {{frac|1|16}} of the mass of one atom of [[oxygen]]-16 as his unit.
-Before [[1961]], the ''physical atomic mass unit'' (amu) was defined as {{frac|1|16}} of the mass of one atom of oxygen-16, while the ''chemical atomic mass unit'' (amu) was defined as {{frac|1|16}} of the ''average'' mass of an oxygen atom (taking the natural abundance of the different oxygen [[isotope]]s 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).
-== References ==
-{{reflist}}
-== See also ==
-* [[Molecular mass]]
 ==External links==
 *[http://www1.bipm.org/en/si/si_brochure/chapter4/table7.html SI website on acceptable non-SI units]
 *[http://physics.nist.gov/cgi-bin/cuu/Value?ukg Accepted value of 1u as of 2006]
--->
+[[Category: CZ Live]]
+[[Category: Chemistry Workgroup]]

Unified atomic mass unit: Difference between revisions

Revision as of 05:01, 1 December 2007

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