Standard atomic weight

Example for copper in terrestrial sources. Two isotopes are present: copper-63 (62.93 u) and copper-65 (64.93 u), in different abundances (%). The standard atomic weight (Ar, standard) for copper is the average, taken their abundance into account, and then divided by the standardised  1⁄12 12C unit.[1] (Note: numbers in this example are simplified)

The standard atomic weight (Ar, standard) or atomic weight is a physical quantity for a chemical element, expressed as a relative atomic mass (Ar). It is specified by (restricted to) the IUPAC (CIAAW) definition of natural, stable, terrestrial sources.

Because of this practical definition, the value is widely used as 'the' atomic weight for real life substances. For example, in pharmaceuticals and scientific research.

Out of the 118 known chemical elements, 84 are stable and have this Earth-environment based value. Typically, such a value is, for example helium: Ar, standard(He) = 4.002602(2). The "(2)" indicates the uncertainty in the last digit shown, to read 4.002602 ±0.000002. IUPAC also publishes abridged values, rounded to five significant figures. For helium, Ar, abridged(He) = 4.0026.

For twelve elements the samples diverge on this value, because their sample sources have had a different decay history. For example, thallium (Tl) in sedimentary rocks has a different isotopic composition than in igneous rocks and volcanic gases. For these elements, the standard atomic weight is noted as an interval: Ar, standard(Tl) = [204.38, 204.39]. With such an interval, for less requiring situations, IUPAC also publishes an conventional value. For thallium, Ar, conventional(Tl) = 204.38.

Definition

Excerpt of an IUPAC Periodic Table showing the interval notation of the standard atomic weights of boron, carbon, and nitrogen (Chemistry International, IUPAC). Example: the pie chart for boron shows it to be composed of about 20% 10B and 80% 11B. This isotope mix causes the standard atomic weight of ordinary Earthly boron samples to be expected to fall within the interval 10.806 to 10.821. Boron samples from unusual sources, particularly non-terrestrial sources, might have measured atomic weights that fall outside this range. Atomic weight and relative atomic mass are synonyms.

The standard atomic weight is a more specific value of a relative atomic mass. It is defined as the relative atomic mass of a source in the local environment of the Earth's crust and atmosphere as determined by the IUPAC Commission on Atomic Weights and Isotopic Abundances. (CIAAW)[2] In general, values from different sources are subject to natural variation due to a different radioactive history of sources. By limiting the sources to terrestrial origin only, the CIAAW determined values have less variance, and are a more precise value for atomic masses actually found and used in worldly materials.

The CIAAW-published values are used and sometimes lawfully required in mass calculations. The values have an uncertainty (noted in brackets), or are an expectation interval (see example in illustration immediately above). This uncertainty reflects natural variability in isotopic distribution for an element, rather than uncertainty in measurement (which is much smaller with quality instruments).[3]

Although there is an attempt to cover the range of variability on Earth with standard atomic weight figures, there are known cases of mineral samples which contain elements with atomic weights that are outliers from the standard atomic weight range.[4]

For synthetic elements the isotope formed depends on the means of synthesis, so the concept of natural isotope abundance has no meaning. Therefore, for synthetic elements the total nucleon count of the most stable isotope (i.e., the isotope with the longest half-life) is listed in brackets, in place of the standard atomic weight.

When the term "atomic weight" is used in chemistry, usually it is the more specific standard atomic weight that is implied. It is standard atomic weights that are used in periodic tables and many standard references in ordinary terrestrial chemistry.

Lithium represents a unique case where the natural abundances of the isotopes have in some cases been found to have been perturbed by human isotopic separation activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources, such as rivers.

Terrestrial definition

An example of why “conventional terrestrial sources" must be specified in giving standard atomic weight values is the element argon. Between locations in the Solar System, the atomic weight of argon varies as much as 10%, due to extreme variance in isotopic composition. Where the major source of argon is the decay of 40
K in rocks, 40
Ar will be the dominant isotope. Such locations include the planets Mercury and Mars, and the moon Titan. On Earth the ratios of the three isotopes 36Ar : 38Ar : 40Ar are approximately 5 : 1 : 1600, giving terrestrial argon a standard atomic weight of 39.948(1). This atomic weight is larger than that of the next element potassium, causing confusion in the days when the places of elements in the periodic table was largely determined according to atomic weight.

However, such is not the case in the rest of the universe. Argon produced directly by stellar nucleosynthesis, is dominated by the alpha-process nuclide 36
Ar. Correspondingly, solar argon contains 84.6% 36
Ar (according to solar wind measurements),[5] and the ratio of the three isotopes 36Ar : 38Ar : 40Ar in the atmospheres of the outer planets is 8400 : 1600 : 1.[6]The atomic weight of argon in the Sun and most of the universe, therefore, would be only approximately 36.3.[7]

Causes of uncertainty on Earth

Famously, the published atomic weight value comes with an uncertainty. This uncertainty (and related: precision) follows from its definition, the source being "terrestrial and stable". Systematic causes for uncertainty are:

  1. Measurement limits. As always, the physical measurement is never finite. There is always more detail to be found and read. This applies to every single, pure isotope found. For example, today the mass of the main natural fluorine isotope can be measured to the accuracy of eleven decimal places: 18.998403163(6). But a still more precise measurement system could become available, producing more decimals.
  2. Imperfect mixtures of isotopes. In the samples taken and measured the mix (relative abundance) of those isotopes may vary. For example copper. While in general its two isotopes make out 69.15% and 30.85% each of all copper found, the natural sample being measured can have had an incomplete 'stirring' and so the percentages are different. The precision is improved by measuring more samples of course, but there remains this cause of uncertainty. (Example: lead samples vary so much, it can not be noted more precise than four figures: 207.2)
  3. Earthly sources with a different history. A source is the greater area being researched, for example 'ocean water' or 'volcanic rock' (as opposed to a 'sample': the single heap of material being investigated). It appears that some elements have a different isotopic mix per source. For example thallium in igneous rock has more lighter isotopes, while in sedimentary rock it has more heavy isotopes. There is no Earthly mean number. These elements show the interval notation: Ar, standard(Tl) = [204.38, 204.39]. For practical reasons, a simplified 'conventional' number is published too (for Tl: 204.38).

These three uncertainties are accumulative. The published value is a result of all these.

Abridged atomic weight

The abridged atomic weight, also published by CIAAW, is derived from the standard atomic weight reducing the numbers to five digits (five significant figures). The name does not say 'rounded'.

Interval borders are rounded downwards for the first (lowmost) border, and upwards for the upward (upmost) border. This way, the more precise original interval is fully covered.

Examples:

Conventional atomic weight

Twelve chemical elements have a standard atomic weight that is defined not as a single number, but as an interval. For example, hydrogen has Ar, standard(H) = [1.00 784, 1.00811]. This notation states that the various sources on Earth have substantially different isotopic constitutions, and uncertainties are incorporated in the two numbers. For these elements, there is not an 'Earth average' constitution, and the 'right' value is not its middle (that would be 1.007975 for hydrogen, with an uncertainty of (±0.000135) that would make it just cover the interval). However, for situations where a less precise value is acceptable, CIAAW has published a single-number conventional atomic weight that can be used for example in trade. For hydrogen, Ar, conventional(H) = 1.008. The twelve elements are: hydrogen, lithium, boron, carbon, nitrogen, oxygen, magnesium, silicon, sulfur, chlorine, bromine and thallium.

Naming controversy

The use of the name "atomic weight" has attracted a great deal of controversy among scientists.[8] Objectors to the name usually prefer the term "relative atomic mass" (not to be confused with atomic mass). The basic objection is that atomic weight is not a weight, that is the force exerted on an object in a gravitational field, measured in units of force such as the newton or poundal.

In reply, supporters of the term "atomic weight" point out (among other arguments)[8] that

It could be added that atomic weight is often not truly "atomic" either, as it does not correspond to the property of any individual atom. The same argument could be made against "relative atomic mass" used in this sense.

Determination of relative atomic mass

Modern relative atomic masses (a term specific to a given element sample) are calculated from measured values of atomic mass (for each nuclide) and isotopic composition of a sample. Highly accurate atomic masses are available[10][11] for virtually all non-radioactive nuclides, but isotopic compositions are both harder to measure to high precision and more subject to variation between samples.[12][13] For this reason, the relative atomic masses of the 22 mononuclidic elements (which are the same as the isotopic masses for each of the single naturally occurring nuclides of these elements) are known to especially high accuracy. For example, there is an uncertainty of only one part in 38 million for the relative atomic mass of fluorine, a precision which is greater than the current best value for the Avogadro constant (one part in 20 million).

Isotope Atomic mass[11] Abundance[12]
Standard Range
28Si 27.976 926 532 46(194) 92.2297(7)% 92.21–92.25%
29Si 28.976 494 700(22) 4.6832(5)% 4.67–4.69%
30Si 29.973 770 171(32) 3.0872(5)% 3.08–3.10%

The calculation is exemplified for silicon, whose relative atomic mass is especially important in metrology. Silicon exists in nature as a mixture of three isotopes: 28Si, 29Si and 30Si. The atomic masses of these nuclides are known to a precision of one part in 14 billion for 28Si and about one part in one billion for the others. However the range of natural abundance for the isotopes is such that the standard abundance can only be given to about ±0.001% (see table). The calculation is

Ar(Si) = (27.97693 × 0.922297) + (28.97649 × 0.046832) + (29.97377 × 0.030872) = 28.0854

The estimation of the uncertainty is complicated,[14] especially as the sample distribution is not necessarily symmetrical: the IUPAC standard relative atomic masses are quoted with estimated symmetrical uncertainties,[15] and the value for silicon is 28.0855(3). The relative standard uncertainty in this value is 1×10–5 or 10 ppm. To further reflect this natural variability, in 2010, IUPAC made the decision to list the relative atomic masses of 10 elements as an interval rather than a fixed number.[16]

Periodic table with standard atomic weights

Standard atomic weight (abridged, conventional*)[1]
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Group →
↓ Period
1 H 1.008*
He4.0026
2 Li6.94* Be9.0122
B 10.81* C 12.011* N 14.007* O 15.999* F 18.998 Ne20.180
3 Na22.990 Mg24.305*
Al26.982 Si28.085* P 30.974 S 32.06* Cl35.45* Ar39.948
4 K 39.098 Ca40.078(4) Sc44.956 Ti47.867 V 50.942 Cr51.996 Mn54.938 Fe55.845(2) Co58.933 Ni58.693 Cu63.546(3) Zn65.38(2) Ga69.723 Ge72.630(8) As74.922 Se78.971 Br79.904* Kr83.798(2)
5 Rb85.468 Sr87.62 Y 88.906 Zr91.224(2) Nb92.906 Mo95.95 Tc[98] Ru101.07(2) Rh102.91 Pd106.42 Ag107.87 Cd112.41 In114.82 Sn118.71 Sb121.76 Te127.60(3) I 126.90 Xe131.29
6 Cs132.91 Ba137.33 La138.91 1 asterisk Hf178.49(2) Ta180.95 W 183.84 Re186.21 Os190.23(3) Ir192.22 Pt195.08 Au196.97 Hg200.59 Tl204.38* Pb207.2 Bi208.98 Po[209] At[210] Rn[222]
7 Fr[223] Ra[226] Ac[227] 1 asterisk Rf[267] Db[268] Sg[269] Bh[270] Hs[277] Mt[278] Ds[281] Rg[282] Cn[285] Nh[286] Fl[289] Mc[290] Lv[293] Ts[294] Og[294]
1 asterisk Ce140.12 Pr140.91 Nd144.24 Pm[145] Sm150.36(2) Eu151.96 Gd157.25 Tb158.93 Dy162.50 Ho164.93 Er167.26 Tm168.93 Yb173.05 Lu174.97
1 asterisk Th232.04 Pa231.04 U 238.03 Np[237] Pu[244] Am[243] Cm[247] Bk[247] Cf[251] Es[252] Fm[257] Md[258] No[259] Lr[266]
 * Standard atomic weight The formal standard atomic weight may look like 4.002602(2) for helium, and [1.0078, 1.0082] for hydrogen. The "(n)" is the uncertainty.
 * Abridged The value is abridged to five significant figures. The ± uncertainty is noted as "(x)", or "(1)" when omitted. See abridged standard atomic weight.
 * Conventional When the formal standard atomic weight is an interval, like [1.0078, 1.0082], a simple, single number is published too. See conventional standard atomic weight.

See also

References

  1. 1 2 Meija, J.; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure Appl. Chem. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
  2. ↑ IUPAC Definition of Standard Atomic Weight
  3. ↑ ATOMIC WEIGHTS OF THE ELEMENTS 2005 (IUPAC TECHNICAL REPORT), M. E. WIESER Pure Appl. Chem., V.78, pp. 2051, 2006
  4. ↑ Definition of standard atomic weights: "Recommended values of relative atomic masses of the elements revised biennially by the IUPAC Commission on Atomic Weights and Isotopic Abundances and applicable to elements in any normal sample with a high level of confidence. A normal sample is any reasonably possible source of the element or its compounds in commerce for industry and science and has not been subject to significant modification of isotopic composition within a geologically brief period."
  5. ↑ Lodders, K. (2008). "The solar argon abundance". Astrophysical Journal. 674: 607–611. Bibcode:2008ApJ...674..607L. arXiv:0710.4523 Freely accessible. doi:10.1086/524725.
  6. ↑ Cameron, A. G. W. (1973). "Elemental and isotopic abundances of the volatile elements in the outer planets". Space Science Reviews. 14 (3–4): 392–400. Bibcode:1973SSRv...14..392C. doi:10.1007/BF00214750.
  7. ↑ This can be determined from the preceding figures per the definition of atomic weight and WP:CALC
  8. 1 2 de Bièvre, P.; Peiser, H. S. (1992). "'Atomic Weight'—The Name, Its History, Definition, and Units". Pure Appl. Chem. 64 (10): 1535–43. doi:10.1351/pac199264101535.
  9. ↑ Dalton, John (1808). A New System of Chemical Philosophy. Manchester.
  10. ↑ National Institute of Standards and Technology. Atomic Weights and Isotopic Compositions for All Elements.
  11. 1 2 Wapstra, A.H.; Audi, G.; Thibault, C. (2003), The AME2003 Atomic Mass Evaluation (Online ed.), National Nuclear Data Center. Based on:
  12. 1 2 Rosman, K. J. R.; Taylor, P. D. P. (1998), "Isotopic Compositions of the Elements 1997" (PDF), Pure and Applied Chemistry, 70 (1): 217–35, doi:10.1351/pac199870010217
  13. ↑ Coplen, T. B.; et al. (2002), "Isotopic Abundance Variations of Selected Elements" (PDF), Pure and Applied Chemistry, 74 (10): 1987–2017, doi:10.1351/pac200274101987
  14. ↑ Meija, Juris; Mester, Zoltán (2008). "Uncertainty propagation of atomic weight measurement results". Metrologia. 45: 53–62. Bibcode:2008Metro..45...53M. doi:10.1088/0026-1394/45/1/008.
  15. ↑ Holden, Norman E. (2004). "Atomic Weights and the International Committee—A Historical Review". Chemistry International. 26 (1): 4–7.
  16. ↑ IUPAC - International Union of Pure and Applied Chemistry: Atomic Weights of Ten Chemical Elements About to Change
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