Standard atomic weight

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Example: copper in terrestrial sources. Two isotopes are present: copper-63 (62.9) and copper-65 (64.9), in abundances 69% + 31%. The standard atomic weight (Ardeg(Cu)) for copper is the average, weighted by their natural abundance, and then divided by the atomic mass constant mu. CIAAW 2013 - Standard atomic weight for cupper (29, Cu).svg
Example: copper in terrestrial sources. Two isotopes are present: copper-63 (62.9) and copper-65 (64.9), in abundances 69% + 31%. The standard atomic weight(Ar°(Cu)) for copper is the average, weighted by their natural abundance, and then divided by the atomic mass constant mu.

The standard atomic weight of a chemical element (symbol Ar°(E) for element "E") is the weighted arithmetic mean of the relative isotopic masses of all isotopes of that element weighted by each isotope's abundance on Earth. For example, isotope 63Cu (Ar = 62.929) constitutes 69% of the copper on Earth, the rest being 65Cu (Ar = 64.927), so

Contents

Because relative isotopic masses are dimensionless quantities, this weighted mean is also dimensionless. It can be converted into a measure of mass (with dimension M) by multiplying it with the dalton, also known as the atomic mass constant.

Among various variants of the notion of atomic weight (Ar, also known as relative atomic mass ) used by scientists, the standard atomic weight (Ar°) is the most common and practical. The standard atomic weight of each chemical element is determined and published by the Commission on Isotopic Abundances and Atomic Weights (CIAAW) of the International Union of Pure and Applied Chemistry (IUPAC) based on natural, stable, terrestrial sources of the element. The definition specifies the use of samples from many representative sources from the Earth, so that the value can widely be used as "the" atomic weight for substances as they are encountered in reality—for example, in pharmaceuticals and scientific research. Non-standardized atomic weights of an element are specific to sources and samples, such as the atomic weight of carbon in a particular bone from a particular archeological site. Standard atomic weight averages such values to the range of atomic weights that a chemist might expect to derive from many random samples from Earth. This range is the rationale for the interval notation given for some standard atomic weight values.

Of the 118 known chemical elements, 80 have stable isotopes and 84 have this Earth-environment based value. Typically, such a value is, for example helium: Ar°(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 fourteen 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°(Tl) =[204.38, 204.39]. With such an interval, for less demanding situations, IUPAC also publishes a 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% B and 80% B. This isotope mix causes the atomic weight of ordinary Earthly boron samples to be expected to fall within the interval 10.806 to 10.821. and this interval is the standard atomic weight. 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. IUPAC Periodic Table of the Elements 2011.jpg
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% B and 80% B. This isotope mix causes the atomic weight of ordinary Earthly boron samples to be expected to fall within the interval 10.806 to 10.821. and this interval is the standard atomic weight. 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 special value of the relative atomic mass. It is defined as the "recommended values" of relative atomic masses of sources 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. Thus, standard atomic weights are an expectation range of atomic weights from a range of samples or sources. By limiting the sources to terrestrial origin only, the CIAAW-determined values have less variance, and are a more precise value for relative atomic masses (atomic weights) 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. [2]

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.[ citation needed ][ dubious ]

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).

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), [4] and the ratio of the three isotopes 36Ar : 38Ar : 40Ar in the atmospheres of the outer planets is 8400 : 1600 : 1. [5] The atomic weight of argon in the Sun and most of the universe, therefore, would be only approximately 36.3. [6]

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 (fluorine-19) 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°(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.

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 [7] [8] for virtually all non-radioactive nuclides, but isotopic compositions are both harder to measure to high precision and more subject to variation between samples. [9] [10] 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.

IsotopeAtomic mass [8] Abundance [9]
StandardRange
28Si27.976 926 532 46(194)92.2297(7)%92.21–92.25%
29Si28.976 494 700(22)4.6832(5)%4.67–4.69%
30Si29.973 770 171(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, [11] especially as the sample distribution is not necessarily symmetrical: the IUPAC standard relative atomic masses are quoted with estimated symmetrical uncertainties, [12] 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. [13]

Naming controversy

The use of the name "atomic weight" has attracted a great deal of controversy among scientists. [14] 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.[ citation needed ]

In reply, supporters of the term "atomic weight" point out (among other arguments) [14] 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.

Published values

IUPAC publishes one formal value for each stable chemical element, called the standard atomic weight. [16] [1] :Table 1 Any updates are published biannually (in uneven years). In 2015, the atomic weight of ytterbium was updated. [16] Per 2017, 14 atomic weights were changed, including argon changing from single number to interval value. [17] [18]

The value published can have an uncertainty, like for neon: 20.1797(6), or can be an interval, like for boron: [10.806, 10.821].

Next to these 84 values, IUPAC also publishes abridged values (up to five digits per number only), and for the twelve interval values, conventional values (single number values).

Symbol Ar is a relative atomic mass, for example from a specific sample. To be specific, the standard atomic weight can be noted as Ar°(E), where (E) is the element symbol.

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. [1] :Table 2

Examples:

Conventional atomic weight

Fourteen chemical elements – hydrogen, lithium, boron, carbon, nitrogen, oxygen, magnesium, silicon, sulfur, chlorine, argon, bromine, thallium, and lead – have a standard atomic weight that is defined not as a single number, but as an interval. For example, hydrogen has Ar°(H) = [1.00 784, 1.00811]. This notation states that the various sources on Earth have substantially different isotopic constitutions, and that the uncertainties in all of them are just covered by the two numbers. For these elements, there is not an 'Earth average' constitution, and the 'right' value is not its middle (which 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, for example in trade, CIAAW has published a single-number conventional atomic weight. For hydrogen, Ar, conventional°(H) = 1.008. [1] :Table 3

A formal short atomic weight

By using the abridged value, and the conventional value for the fourteen interval values, a short IUPAC-defined value (5 digits plus uncertainty) can be given for all stable elements. In many situations, and in periodic tables, this may be sufficiently detailed. [1] :Tables 2 and 3

Element (E)
Ar°(E)
Value type
Ar°(E), abridged
or conventional
Mass number
[most stable isotope]
hydrogen 1H[1.00784, 1.00811]interval1.0080±0.0002
nitrogen 7N[14.00643, 14.00728]interval14.007±0.001
fluorine 9F18.998403162±0.000000005number ±uncertainty18.998±0.001
calcium 20Ca40.078±0.004number ±uncertainty40.078±0.004
technetium 43Tc(none)most stable isotope[97]

List of atomic weights

Standard atomic weight of the elements (IUPAC 20092021 [ref 1] )
ZSymbolNameAr, standardAbridgedYear changed
1H hydrogen [1.00784, 1.00811]1.0080±0.00022009
2He helium 4.002602±0.0000024.0026±0.00011983
3Li lithium [6.938, 6.997]6.94±0.062009
4Be beryllium 9.0121831±0.00000059.0122±0.00012013
5B boron [10.806, 10.821]10.81±0.022009
6C carbon [12.0096, 12.0116]12.011±0.0022009
7N nitrogen [14.00643, 14.00728]14.007±0.0012009
8O oxygen [15.99903, 15.99977]15.999±0.0012009
9F fluorine 18.998403162±0.00000000518.998±0.0012021
10Ne neon 20.1797±0.000620.180±0.0011985
11Na sodium 22.98976928±0.0000000222.990±0.0012005
12Mg magnesium [24.304, 24.307]24.305±0.0022011
13Al aluminium 26.9815384±0.000000326.982±0.0012017
14Si silicon [28.084, 28.086]28.085±0.0012009
15P phosphorus 30.973761998±0.00000000530.974±0.0012013
16S sulfur [32.059, 32.076]32.06±0.022009
17Cl chlorine [35.446, 35.457]35.45±0.012009
18Ar argon [39.792, 39.963]39.95±0.162017
19K potassium 39.0983±0.000139.098±0.0011979
20Ca calcium 40.078±0.00440.078±0.0041983
21Sc scandium 44.955907±0.00000444.956±0.0012021
22Ti titanium 47.867±0.00147.867±0.0011993
23V vanadium 50.9415±0.000150.942±0.0011977
24Cr chromium 51.9961±0.000651.996±0.0011983
25Mn manganese 54.938043±0.00000254.938±0.0012017
26Fe iron 55.845±0.00255.845±0.0021993
27Co cobalt 58.933194±0.00000358.933±0.0012017
28Ni nickel 58.6934±0.000458.693±0.0012007
29Cu copper 63.546±0.00363.546±0.0031969
30Zn zinc 65.38±0.0265.38±0.022007
31Ga gallium 69.723±0.00169.723±0.0011987
32Ge germanium 72.630±0.00872.630±0.0082009
33As arsenic 74.921595±0.00000674.922±0.0012013
34Se selenium 78.971±0.00878.971±0.0082013
35Br bromine [79.901, 79.907]79.904±0.0032011
36Kr krypton 83.798±0.00283.798±0.0022001
37Rb rubidium 85.4678±0.000385.468±0.0011969
38Sr strontium 87.62±0.0187.62±0.011969
39Y yttrium 88.905838±0.00000288.906±0.0012021
40Zr zirconium 91.224±0.00291.224±0.0021983
41Nb niobium 92.90637±0.0000192.906±0.0012017
42Mo molybdenum 95.95±0.0195.95±0.012013
43Tc technetium -
44Ru ruthenium 101.07±0.02101.07±0.021983
45Rh rhodium 102.90549±0.00002102.91±0.012017
46Pd palladium 106.42±0.01106.42±0.011979
47Ag silver 107.8682±0.0002107.87±0.011985
48Cd cadmium 112.414±0.004112.41±0.012013
49In indium 114.818±0.001114.82±0.012011
50Sn tin 118.710±0.007118.71±0.011983
51Sb antimony 121.760±0.001121.76±0.011993
52Te tellurium 127.60±0.03127.60±0.031969
53I iodine 126.90447±0.00003126.90±0.011985
54Xe xenon 131.293±0.006131.29±0.011999
55Cs caesium 132.90545196±0.00000006132.91±0.012013
56Ba barium 137.327±0.007137.33±0.011985
57La lanthanum 138.90547±0.00007138.91±0.012005
58Ce cerium 140.116±0.001140.12±0.011995
59Pr praseodymium 140.90766±0.00001140.91±0.012017
60Nd neodymium 144.242±0.003144.24±0.012005
61Pm promethium
62Sm samarium 150.36±0.02150.36±0.022005
63Eu europium 151.964±0.001151.96±0.011995
64Gd gadolinium 157.25±0.03157.25±0.031969
65Tb terbium 158.925354±0.000007158.93±0.012021
66Dy dysprosium 162.500±0.001162.50±0.012001
67Ho holmium 164.930329±0.000005164.93±0.012021
68Er erbium 167.259±0.003167.26±0.011999
69Tm thulium 168.934219±0.000005168.93±0.012021
70Yb ytterbium 173.045±0.010173.05±0.022015
71Lu lutetium 174.9668±0.0001174.97±0.012007
72Hf hafnium 178.486±0.006178.49±0.012019
73Ta tantalum 180.94788±0.00002180.95±0.012005
74W tungsten 183.84±0.01183.84±0.011991
75Re rhenium 186.207±0.001186.21±0.011973
76Os osmium 190.23±0.03190.23±0.031991
77Ir iridium 192.217±0.002192.22±0.012017
78Pt platinum 195.084±0.009195.08±0.022005
79Au gold 196.966570±0.000004196.97±0.012017
80Hg mercury 200.592±0.003200.59±0.012011
81Tl thallium [204.382, 204.385]204.38±0.012009
82Pb lead [206.14, 207.94]207.2±1.12020
83Bi bismuth 208.98040±0.00001208.98±0.012005
84Po polonium -
85At astatine -
86Rn radon -
87Fr francium -
88Ra radium -
89Ac actinium -
90Th thorium 232.0377±0.0004232.04±0.012013
91Pa protactinium 231.03588±0.00001231.04±0.012017
92U uranium 238.02891±0.00003238.03±0.011999
93Np neptunium -
94Pu plutonium -
95Am americium -
96Cm curium -
97Bk berkelium -
98Cf californium -
99Es einsteinium -
100Fm fermium -
101Md mendelevium -
102No nobelium -
103Lr lawrencium -
104Rf rutherfordium -
105Db dubnium -
106Sg seaborgium -
107Bh bohrium -
108Hs hassium -
109Mt meitnerium -
110Ds darmstadtium -
111Rg roentgenium -
112Cn copernicium -
113Nh nihonium -
114Fl flerovium -
115Mc moscovium -
116Lv livermorium -
117Ts tennessine -
118Og oganesson -
  1.  (This list:)
    CIAAW may publish changes to atomic weights (including its precision and derived values). Since 1947, any update this is done in odd years nominally; the actual date of publication may be some time later.
    • 2009 (introducing interval notation; Ge):
    "Atomic weights of the elements 2009 (IUPAC Technical Report)". Pure Appl. Chem. 83 (2): 359–396. 12 December 2010. doi:10.1351/PAC-REP-10-09-14.
    • 2011 (interval for Br, Mg):
    "Atomic weights of the elements 2011 (IUPAC Technical Report)". Pure Appl. Chem. 85 (5): 1047–1078. 29 April 2013. doi:10.1351/PAC-REP-13-03-02.
    • 2013 (all elements listed):
    Meija, Juris; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry . 88 (3): 265–91. doi: 10.1515/pac-2015-0305 .
    • 2015 (ytterbium changed):
    "Standard Atomic Weight of Ytterbium Revised". Chemistry International. 37 (5–6): 26. October 2015. doi: 10.1515/ci-2015-0512 . eISSN   0193-6484. ISSN   0193-6484.
    • 2017 (14 values changed):
    "Standard atomic weights of 14 chemical elements revised". CIAAW. 2018-06-05.
    * "2020" is an inconsistent year for change publication: CIAAW maintains that only odd years, changes are publicised.
    • 2021 (all elements listed); (4 values changed; introduced new symbol; merge "conventional" into "abridged" columns; change uncertainty notation (use "±")
    Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; et al. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN   1365-3075.
    Uncertainty handling
    About notation and handling of the uncertainty in the values, including those in [ ] range values:
    Outdated references
    See also: {{ Isotopes table/references }}

In the periodic table

Group 1 2   3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Hydrogen &
alkali metals
Alkaline earth metalsTrielsTetrelsPnicto­gensChal­co­gensHalo­gensNoble
gases
Period

1

Hydro­gen1H1.0080 He­lium2He4.0026
2 Lith­ium3Li6.94 Beryl­lium4Be9.0122 Boron5B10.81 Carbon6C12.011 Nitro­gen7N14.007 Oxy­gen8O15.999 Fluor­ine9F18.998 Neon10Ne20.180
3 So­dium11Na22.990 Magne­sium12Mg24.305 Alumin­ium13Al26.982 Sili­con14Si28.085 Phos­phorus15P30.974 Sulfur16S32.06 Chlor­ine17Cl35.45 Argon18Ar39.95
4 Potas­sium19K39.098 Cal­cium20Ca40.078 Scan­dium21Sc44.956 Tita­nium22Ti47.867 Vana­dium23V50.942 Chrom­ium24Cr51.996 Manga­nese25Mn54.938 Iron26Fe55.845 Cobalt27Co58.933 Nickel28Ni58.693 Copper29Cu63.546 Zinc30Zn65.38 Gallium31Ga69.723 Germa­nium32Ge72.630 Arsenic33As74.922 Sele­nium34Se78.971 Bromine35Br79.904 Kryp­ton36Kr83.798
5 Rubid­ium37Rb85.468 Stront­ium38Sr87.62 Yttrium39Y88.906 Zirco­nium40Zr91.224 Nio­bium41Nb92.906 Molyb­denum42Mo95.95 Tech­netium43Tc[97] Ruthe­nium44Ru101.07 Rho­dium45Rh102.91 Pallad­ium46Pd106.42 Silver47Ag107.87 Cad­mium48Cd112.41 Indium49In114.82 Tin50Sn118.71 Anti­mony51Sb121.76 Tellur­ium52Te127.60 Iodine53I126.90 Xenon54Xe131.29
6 Cae­sium55Cs132.91 Ba­rium56Ba137.33 Asterisks one.svg Lute­tium71Lu174.97 Haf­nium72Hf178.49 Tanta­lum73Ta180.95 Tung­sten74W183.84 Rhe­nium75Re186.21 Os­mium76Os190.23 Iridium77Ir192.22 Plat­inum78Pt195.08 Gold79Au196.97 Mer­cury80Hg200.59 Thallium81Tl204.38 Lead82Pb207.2 Bis­muth83Bi208.98 Polo­nium84Po[209] Asta­tine85At[210] Radon86Rn[222]
7 Fran­cium87Fr[223] Ra­dium88Ra[226] Asterisks 2 (vertical).svg Lawren­cium103Lr[266] Ruther­fordium104Rf[267] Dub­nium105Db[268] Sea­borgium106Sg[269] Bohr­ium107Bh[270] Has­sium108Hs[269] Meit­nerium109Mt[278] Darm­stadtium110Ds[281] Roent­genium111Rg[282] Coper­nicium112Cn[285] Nihon­ium113Nh[286] Flerov­ium114Fl[289] Moscov­ium115Mc[290] Liver­morium116Lv[293] Tenness­ine117Ts[294] Oga­nesson118Og[294]
Asterisks one.svg Lan­thanum57La138.91 Cerium58Ce140.12 Praseo­dymium59Pr140.91 Neo­dymium60Nd144.24 Prome­thium61Pm[145] Sama­rium62Sm150.36 Europ­ium63Eu151.96 Gadolin­ium64Gd157.25 Ter­bium65Tb158.93 Dyspro­sium66Dy162.50 Hol­mium67Ho164.93 Erbium68Er167.26 Thulium69Tm168.93 Ytter­bium70Yb173.05  
Asterisks 2 (vertical).svg Actin­ium89Ac[227] Thor­ium90Th232.04 Protac­tinium91Pa231.04 Ura­nium92U238.03 Neptu­nium93Np[237] Pluto­nium94Pu[244] Ameri­cium95Am[243] Curium96Cm[247] Berkel­ium97Bk[247] Califor­nium98Cf[251] Einstei­nium99Es[252] Fer­mium100Fm[257] Mende­levium101Md[258] Nobel­ium102No[259]

See also

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Gold (79Au) has one stable isotope, 197Au, and 37 radioisotopes, with 195Au being the most stable with a half-life of 186 days. Gold is currently considered the heaviest monoisotopic element. Bismuth formerly held that distinction until alpha-decay of the 209Bi isotope was observed. All isotopes of gold are either radioactive or, in the case of 197Au, observationally stable, meaning that 197Au is predicted to be radioactive but no actual decay has been observed.

Naturally occurring erbium (68Er) is composed of 6 stable isotopes, with 166Er being the most abundant. 39 radioisotopes have been characterized with between 74 and 112 neutrons, or 142 to 180 nucleons, with the most stable being 169Er with a half-life of 9.4 days, 172Er with a half-life of 49.3 hours, 160Er with a half-life of 28.58 hours, 165Er with a half-life of 10.36 hours, and 171Er with a half-life of 7.516 hours. All of the remaining radioactive isotopes have half-lives that are less than 3.5 hours, and the majority of these have half-lives that are less than 4 minutes. This element also has numerous meta states, with the most stable being 167mEr.

Naturally occurring terbium (65Tb) is composed of one stable isotope, 159Tb. Thirty-seven radioisotopes have been characterized, with the most stable being 158Tb with a half-life of 180 years, 157Tb with a half-life of 71 years, and 160Tb with a half-life of 72.3 days. All of the remaining radioactive isotopes have half-lives that are less than 6.907 days, and the majority of these have half-lives that are less than 24 seconds. This element also has 27 meta states, with the most stable being 156m1Tb, 154m2Tb and 154m1Tb.

Naturally occurring praseodymium (59Pr) is composed of one stable isotope, 141Pr. Thirty-eight radioisotopes have been characterized with the most stable being 143Pr, with a half-life of 13.57 days and 142Pr, with a half-life of 19.12 hours. All of the remaining radioactive isotopes have half-lives that are less than 5.985 hours and the majority of these have half-lives that are less than 33 seconds. This element also has 15 meta states with the most stable being 138mPr, 142mPr and 134mPr.

Naturally occurring cerium (58Ce) is composed of 4 stable isotopes: 136Ce, 138Ce, 140Ce, and 142Ce, with 140Ce being the most abundant and the only one theoretically stable; 136Ce, 138Ce, and 142Ce are predicted to undergo double beta decay but this process has never been observed. There are 35 radioisotopes that have been characterized, with the most stable being 144Ce, with a half-life of 284.893 days; 139Ce, with a half-life of 137.640 days and 141Ce, with a half-life of 32.501 days. All of the remaining radioactive isotopes have half-lives that are less than 4 days and the majority of these have half-lives that are less than 10 minutes. This element also has 10 meta states.

<span class="mw-page-title-main">Isotopes of lanthanum</span> Nuclides with atomic number of 57 but with different mass numbers

Naturally occurring lanthanum (57La) is composed of one stable (139La) and one radioactive (138La) isotope, with the stable isotope, 139La, being the most abundant (99.91% natural abundance). There are 39 radioisotopes that have been characterized, with the most stable being 138La, with a half-life of 1.02×1011 years; 137La, with a half-life of 60,000 years and 140La, with a half-life of 1.6781 days. The remaining radioactive isotopes have half-lives that are less than a day and the majority of these have half-lives that are less than 1 minute. This element also has 12 nuclear isomers, the longest-lived of which is 132mLa, with a half-life of 24.3 minutes. Lighter isotopes mostly decay to isotopes of barium and heavy ones mostly decay to isotopes of cerium. 138La can decay to both.

There are 39 known isotopes and 17 nuclear isomers of tellurium (52Te), with atomic masses that range from 104 to 142. These are listed in the table below.

Antimony (51Sb) occurs in two stable isotopes, 121Sb and 123Sb. There are 35 artificial radioactive isotopes, the longest-lived of which are 125Sb, with a half-life of 2.75856 years; 124Sb, with a half-life of 60.2 days; and 126Sb, with a half-life of 12.35 days. All other isotopes have half-lives less than 4 days, most less than an hour.

Indium (49In) consists of two primordial nuclides, with the most common (~ 95.7%) nuclide (115In) being measurably though weakly radioactive. Its spin-forbidden decay has a half-life of 4.41×1014 years, much longer than the currently accepted age of the Universe.

Naturally occurring silver (47Ag) is composed of the two stable isotopes 107Ag and 109Ag in almost equal proportions, with 107Ag being slightly more abundant. Notably, silver is the only element with all stable istopes having nuclear spins of 1/2. Thus both 107Ag and 109Ag nuclei produce narrow lines in nuclear magnetic resonance spectra.

Naturally occurring niobium (41Nb) is composed of one stable isotope (93Nb). The most stable radioisotope is 92Nb with a half-life of 34.7 million years. The next longest-lived niobium isotopes are 94Nb and 91Nb with a half-life of 680 years. There is also a meta state of 93Nb at 31 keV whose half-life is 16.13 years. Twenty-seven other radioisotopes have been characterized. Most of these have half-lives that are less than two hours, except 95Nb, 96Nb and 90Nb. The primary decay mode before stable 93Nb is electron capture and the primary mode after is beta emission with some neutron emission occurring in 104–110Nb.

Arsenic (33As) has 33 known isotopes and at least 10 isomers. Only one of these isotopes, 75As, is stable; as such, it is considered a monoisotopic element. The longest-lived radioisotope is 73As with a half-life of 80 days. Arsenic has been proposed as a "salting" material for nuclear weapons. A jacket of 75As, irradiated by the intense high-energy neutron flux from an exploding thermonuclear weapon, would transmute into the radioactive isotope 76As with a half-life of 1.0778 days and produce approximately 1.13 MeV gamma radiation, significantly increasing the radioactivity of the weapon's fallout for several hours. Such a weapon is not known to have ever been built, tested, or used.

Germanium (32Ge) has five naturally occurring isotopes, 70Ge, 72Ge, 73Ge, 74Ge, and 76Ge. Of these, 76Ge is very slightly radioactive, decaying by double beta decay with a half-life of 1.78 × 1021 years (130 billion times the age of the universe).

Naturally occurring zinc (30Zn) is composed of the 5 stable isotopes 64Zn, 66Zn, 67Zn, 68Zn, and 70Zn with 64Zn being the most abundant. Twenty-five radioisotopes have been characterised with the most abundant and stable being 65Zn with a half-life of 244.26 days, and 72Zn with a half-life of 46.5 hours. All of the remaining radioactive isotopes have half-lives that are less than 14 hours and the majority of these have half-lives that are less than 1 second. This element also has 10 meta states.

Naturally occurring chromium (24Cr) is composed of four stable isotopes; 50Cr, 52Cr, 53Cr, and 54Cr with 52Cr being the most abundant (83.789% natural abundance). 50Cr is suspected of decaying by β+β+ to 50Ti with a half-life of (more than) 1.8×1017 years. Twenty-two radioisotopes, all of which are entirely synthetic, have been characterized, the most stable being 51Cr with a half-life of 27.7 days. All of the remaining radioactive isotopes have half-lives that are less than 24 hours and the majority of these have half-lives that are less than 1 minute. This element also has two meta states, 45mCr, the more stable one, and 59mCr, the least stable isotope or isomer.

Naturally occurring vanadium (23V) is composed of one stable isotope 51V and one radioactive isotope 50V with a half-life of 2.71×1017 years. 24 artificial radioisotopes have been characterized (in the range of mass number between 40 and 65) with the most stable being 49V with a half-life of 330 days, and 48V with a half-life of 15.9735 days. All of the remaining radioactive isotopes have half-lives shorter than an hour, the majority of them below 10 seconds, the least stable being 42V with a half-life shorter than 55 nanoseconds, with all of the isotopes lighter than it, and none of the heavier, have unknown half-lives. In 4 isotopes, metastable excited states were found (including 2 metastable states for 60V), which adds up to 5 meta states.

Naturally occurring scandium (21Sc) is composed of one stable isotope, 45Sc. Twenty-five radioisotopes have been characterized, with the most stable being 46Sc with a half-life of 83.8 days, 47Sc with a half-life of 3.35 days, and 48Sc with a half-life of 43.7 hours and 44Sc with a half-life of 3.97 hours. All the remaining isotopes have half-lives that are less than four hours, and the majority of these have half-lives that are less than two minutes, the least stable being proton unbound 39Sc with a half-life shorter than 300 nanoseconds. This element also has 13 meta states with the most stable being 44m2Sc.

Naturally occurring titanium (22Ti) is composed of five stable isotopes; 46Ti, 47Ti, 48Ti, 49Ti and 50Ti with 48Ti being the most abundant. Twenty-one radioisotopes have been characterized, with the most stable being 44Ti with a half-life of 60 years, 45Ti with a half-life of 184.8 minutes, 51Ti with a half-life of 5.76 minutes, and 52Ti with a half-life of 1.7 minutes. All of the remaining radioactive isotopes have half-lives that are less than 33 seconds, and the majority of these have half-lives that are less than half a second.

<span class="mw-page-title-main">Commission on Isotopic Abundances and Atomic Weights</span> International scientific committee

The Commission on Isotopic Abundances and Atomic Weights (CIAAW) is an international scientific committee of the International Union of Pure and Applied Chemistry (IUPAC) under its Division of Inorganic Chemistry. Since 1899, it is entrusted with periodic critical evaluation of atomic weights of chemical elements and other cognate data, such as the isotopic composition of elements. The biennial CIAAW Standard Atomic Weights are accepted as the authoritative source in science and appear worldwide on the periodic table wall charts.

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