Atomic mass: Difference between revisions
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Atomic mass is numerically equal to '''relative atomic mass''', denoted by ''A''<sub>r</sub>( X), where X is the [[isotope]] of which the mass is indicated. The difference between atomic mass and relative atomic mass is that the former has a dimension (u), while the latter is dimensionless. The relative atomic mass is the ratio of atomic mass to one twelfth of the mass of the nuclide <sup>12</sup>C at rest in its nuclear and electronic ground state. | Atomic mass is numerically equal to '''relative atomic mass''', denoted by ''A''<sub>r</sub>( X), where X is the [[isotope]] of which the mass is indicated. The difference between atomic mass and relative atomic mass is that the former has a dimension (u), while the latter is dimensionless. The relative atomic mass is the ratio of atomic mass to one twelfth of the mass of the nuclide <sup>12</sup>C at rest in its nuclear and electronic ground state. | ||
Recall that different isotopes of an atom have different numbers of neutrons and the same number (the [[atomic number]] ''Z'') of protons. So, different isotopes of a given atom have the same charge but differ in mass. | In practice there are two ways of dealing with the different masses of [[isotope]]s. Recall first that different isotopes of an atom have different numbers of neutrons and the same number (the [[atomic number]] ''Z'') of protons. So, different isotopes of a given atom have the same charge but differ in mass. As an example we look at the [[carbon]] atom (atomic number ''Z'' = 6, i.e., 6 protons). It has two stable isotopes and one radioactive—but long-lived—isotope. The respective atomic masses are, <sup>12</sup>C: 12 u (six neutrons), <sup>13</sup>C: 13.0033548378 u (seven neutrons), and <sup>14</sup>C: 14.003241988 u (eight neutrons). The ''relative'' atomic mass of e.g. the isotope <sup>13</sup>C is the dimensionless number 13.0033548378. | ||
In [[high resolution spectroscopy]] and [[mass spectrometry]] the masses of the isotopes are observed in the spectra. That is, one can distinguish the spectral peaks arising from the different [[isotopologue]]s, (same molecule, different isotopic composition) in the sample. In these fields it is common to consider the samples as mixtures of different compounds (i.e., mixtures of different isotopologues). | |||
In [[ | In most of practical [[chemistry]] one does not notice the different isotopologues that are always present in "off-the-shelf" chemicals. In these ordinary chemicals the concentration of different isotopes is determined by the terrestrial ''natural abundance'' of the isotope. | ||
In | |||
Take the element [[chlorine]] as an example. It has two stable isotopes: <sup>35</sup>Cl (with a mass of 34.96885271 u) and <sup>37</sup>Cl (with a mass of 36.96590260 u). Of all the chlorine atoms occurring on earth 75.78 % is of the lighter kind, while 24.22 % is the heavier isotope. The average mass of the Cl atom is | Take the element [[chlorine]] as an example. It has two stable isotopes: <sup>35</sup>Cl (with a mass of 34.96885271 u) and <sup>37</sup>Cl (with a mass of 36.96590260 u). Of all the chlorine atoms occurring on earth 75.78 % is of the lighter kind, while 24.22 % is the heavier isotope. The average mass of the Cl atom is | ||
: (34.969×75.78 + 36.966×24.22)/100 = 35.453 u. | : (34.969×75.78 + 36.966×24.22)/100 = 35.453 u. | ||
The atomic mass averaged over isotopic abundances is called the '''standard atomic weight'''. (For historical reasons the term "weight" is still used here.) | The atomic mass averaged over isotopic abundances is called the '''standard atomic weight'''. (For historical reasons the term "weight" is still used here.) In most of practical chemistry these standard weights are used in computations as "the" mass of the element, and in this manner one accounts for the different masses of the isotopes found in nature. | ||
==Note on nomenclature== | ==Note on nomenclature== | ||
The concept of "relative atomic mass" is in principle a simple one, yet there is some confusion about its definition. We followed [[NIST]], see the [http://physics.nist.gov/PhysRefData/Compositions/notes.html NIST web site], where clearly and unambiguously the ''relative mass'' is defined of an ''isotope''. The site states: | |||
<blockquote> | <blockquote> | ||
'''Relative Atomic Mass (of the isotope):''' ''A''<sub>r</sub>(X), where X is an isotope | '''Relative Atomic Mass (of the isotope):''' ''A''<sub>r</sub>(X), where X is an isotope | ||
Revision as of 02:16, 21 January 2008
In chemistry and physics, atomic mass (formerly known as atomic weight) is the mass of an atom expressed in unified atomic mass units (u).
Atomic mass is numerically equal to relative atomic mass, denoted by Ar( X), where X is the isotope of which the mass is indicated. The difference between atomic mass and relative atomic mass is that the former has a dimension (u), while the latter is dimensionless. The relative atomic mass is the ratio of atomic mass to one twelfth of the mass of the nuclide 12C at rest in its nuclear and electronic ground state.
In practice there are two ways of dealing with the different masses of isotopes. Recall first that different isotopes of an atom have different numbers of neutrons and the same number (the atomic number Z) of protons. So, different isotopes of a given atom have the same charge but differ in mass. As an example we look at the carbon atom (atomic number Z = 6, i.e., 6 protons). It has two stable isotopes and one radioactive—but long-lived—isotope. The respective atomic masses are, 12C: 12 u (six neutrons), 13C: 13.0033548378 u (seven neutrons), and 14C: 14.003241988 u (eight neutrons). The relative atomic mass of e.g. the isotope 13C is the dimensionless number 13.0033548378.
In high resolution spectroscopy and mass spectrometry the masses of the isotopes are observed in the spectra. That is, one can distinguish the spectral peaks arising from the different isotopologues, (same molecule, different isotopic composition) in the sample. In these fields it is common to consider the samples as mixtures of different compounds (i.e., mixtures of different isotopologues).
In most of practical chemistry one does not notice the different isotopologues that are always present in "off-the-shelf" chemicals. In these ordinary chemicals the concentration of different isotopes is determined by the terrestrial natural abundance of the isotope.
Take the element chlorine as an example. It has two stable isotopes: 35Cl (with a mass of 34.96885271 u) and 37Cl (with a mass of 36.96590260 u). Of all the chlorine atoms occurring on earth 75.78 % is of the lighter kind, while 24.22 % is the heavier isotope. The average mass of the Cl atom is
- (34.969×75.78 + 36.966×24.22)/100 = 35.453 u.
The atomic mass averaged over isotopic abundances is called the standard atomic weight. (For historical reasons the term "weight" is still used here.) In most of practical chemistry these standard weights are used in computations as "the" mass of the element, and in this manner one accounts for the different masses of the isotopes found in nature.
Note on nomenclature
The concept of "relative atomic mass" is in principle a simple one, yet there is some confusion about its definition. We followed NIST, see the NIST web site, where clearly and unambiguously the relative mass is defined of an isotope. The site states:
Relative Atomic Mass (of the isotope): Ar(X), where X is an isotope
This usage is followed by Mohr and Taylor[1] who state that (the atomic mass constant mu is a twelfth of the mass of 12C):
The relative atomic mass Ar(X) of an elementary particle, atom, or more generally an entity X, is defined by Ar(X) = m(X) /mu, where m(X) is the mass of X. Thus Ar(X) is the numerical value of m(X) when m(X) is expressed in u, and evidently Ar(12C)=12.
On the other hand, the official IUPAC publication, IUPAC Goldbook, defines:
relative atomic mass (atomic weight), Ar
The ratio of the average mass of the atom to the unified atomic mass unit
Although it is not explicitly stated in the Goldbook what the average mass is, it is likely and plausible that the averaging is over different isotopes weighted by terrestrial isotopic abundance. Hence, acccording to IUPAC's definition, the relative atomic mass is almost synonymous with the standard atomic weight defined above.
IUPAC also defines standard atomic weight, but adds recommended to its definition, that is, IUPAC defines standard atomic weight as recommended relative atomic mass, implying that the value may change in the future.
From reading Ref. [2] it becomes clear that the confusion is created by too many international comittees addressing this, basically very simple, problem of definition.
Standard Atomic Weights of the Elements
The following table [3] lists the standard atomic weights. The uncertainties in the last given decimal are in parentheses. Square brackets [ ] indicate the mass number of the most stable isotope. CS stands for chemical symbol. Z is the atomic number. See this article for a list of the full names of the elements.
| Z | CS | Mass | Z | CS | Mass | Z | CS | Mass | |||
|---|---|---|---|---|---|---|---|---|---|---|---|
| 1 | H | 1.00794(7) | 38 | Sr | 87.62(1) | 75 | Re | 186.207(1) | |||
| 2 | He | 4.002602(2) | 39 | Y | 88.90585(2) | 76 | Os | 190.23(3) | |||
| 3 | Li | 6.941(2) | 40 | Zr | 91.224(2) | 77 | Ir | 192.217(3) | |||
| 4 | Be | 9.012182(3) | 41 | Nb | 92.90638(2) | 78 | Pt | 195.078(2) | |||
| 5 | B | 10.811(7) | 42 | Mo | 95.94(2) | 79 | Au | 196.96655(2) | |||
| 6 | C | 12.0107(8) | 43 | Tc | [98] | 80 | Hg | 200.59(2) | |||
| 7 | N | 14.0067(2) | 44 | Ru | 101.07(2) | 81 | Tl | 204.3833(2) | |||
| 8 | O | 15.9994(3) | 45 | Rh | 102.90550(2) | 82 | Pb | 207.2(1) | |||
| 9 | F | 18.9984032(5) | 46 | Pd | 106.42(1) | 83 | Bi | 208.98038(2) | |||
| 10 | Ne | 20.1797(6) | 47 | Ag | 107.8682(2) | 84 | Po | [209] | |||
| 11 | Na | 22.989770(2) | 48 | Cd | 112.411(8) | 85 | At | [210] | |||
| 12 | Mg | 24.3050(6) | 49 | In | 114.818(3) | 86 | Rn | [222] | |||
| 13 | Al | 26.981538(2) | 50 | Sn | 118.710(7) | 87 | Fr | [223] | |||
| 14 | Si | 28.0855(3) | 51 | Sb | 121.760(1) | 88 | Ra | [226] | |||
| 15 | P | 30.973761(2) | 52 | Te | 127.60(3) | 89 | Ac | [227] | |||
| 16 | S | 32.065(5) | 53 | I | 126.90447(3) | 90 | Th | 232.0381(1) | |||
| 17 | Cl | 35.453(2) | 54 | Xe | 131.293(6) | 91 | Pa | 231.03588(2) | |||
| 18 | Ar | 39.948(1) | 55 | Cs | 132.90545(2) | 92 | U | 238.02891(3) | |||
| 19 | K | 39.0983(1) | 56 | Ba | 137.327(7) | 93 | Np | [237] | |||
| 20 | Ca | 40.078(4) | 57 | La | 138.9055(2) | 94 | Pu | [244] | |||
| 21 | Sc | 44.955910(8) | 58 | Ce | 140.116(1) | 95 | Am | [243] | |||
| 22 | Ti | 47.867(1) | 59 | Pr | 140.90765(2) | 96 | Cm | [247] | |||
| 23 | V | 50.9415(1) | 60 | Nd | 144.24(3) | 97 | Bk | [247] | |||
| 24 | Cr | 51.9961(6) | 61 | Pm | [145] | 98 | Cf | [251] | |||
| 25 | Mn | 54.938049(9) | 62 | Sm | 150.36(3) | 99 | Es | [252] | |||
| 26 | Fe | 55.845(2) | 63 | Eu | 151.964(1) | 100 | Fm | [257] | |||
| 27 | Co | 58.933200(9) | 64 | Gd | 157.25(3) | 101 | Md | [258] | |||
| 28 | Ni | 58.6934(2) | 65 | Tb | 158.92534(2) | 102 | No | [259] | |||
| 29 | Cu | 63.546(3) | 66 | Dy | 162.500(1) | 103 | Lr | [262] | |||
| 30 | Zn | 65.409(4) | 67 | Ho | 164.93032(2) | 104 | Rf | [261] | |||
| 31 | Ga | 69.723(1) | 68 | Er | 167.259(3) | 105 | Db | [262] | |||
| 32 | Ge | 72.64(1) | 69 | Tm | 168.93421(2) | 106 | Sg | [266] | |||
| 33 | As | 74.92160(2) | 70 | Yb | 173.04(3) | 107 | Bh | [264] | |||
| 34 | Se | 78.96(3) | 71 | Lu | 174.967(1) | 108 | Hs | [277] | |||
| 35 | Br | 79.904(1) | 72 | Hf | 178.49(2) | 109 | Mt | [268] | |||
| 36 | Kr | 83.798(2) | 73 | Ta | 180.9479(1) | 110 | Ds | [281] | |||
| 37 | Rb | 85.4678(3) | 74 | W | 183.84(1) | 111 | Rg | [272] | |||
Notes
- ↑ P. J. Mohr and B. N. Taylor, Reviews of Modern Physics, vol. 77, p. 1 (2005)
- ↑ Article about Atomic Weights
- ↑ Physical Reference Data. The numbers in this table are taken from the web site of NIST on December 2, 2007.