Abundance of the chemical elements

The abundance of the chemical elements is a measure of the occurrence of the chemical elements relative to all other elements in a given environment. Abundance is measured in one of three ways: by the mass-fraction (the same as weight fraction); by the mole-fraction (fraction of atoms by numerical count, or sometimes fraction of molecules in gases); or by the volume-fraction. Volume-fraction is a common abundance measure in mixed gases such as planetary atmospheres, and is similar in value to molecular mole-fraction for gas mixtures at relatively low densities and pressures, and ideal gas mixtures. Most abundance values in this article are given as mass-fractions.

For example, the abundance of oxygen in pure water can be measured in two ways: the mass fraction is about 89%, because that is the fraction of water's mass which is oxygen. However, the mole-fraction is 33.3333...% because only 1 atom of 3 in water, H2O, is oxygen. As another example, looking at the mass-fraction abundance of hydrogen and helium in both the Universe as a whole and in the atmospheres of gas-giant planets such as Jupiter, it is 74% for hydrogen and 23–25% for helium; while the (atomic) mole-fraction for hydrogen is 92%, and for helium is 8%, in these environments. Changing the given environment to Jupiter's outer atmosphere, where hydrogen is diatomic while helium is not, changes the molecular mole-fraction (fraction of total gas molecules), as well as the fraction of atmosphere by volume, of hydrogen to about 86%, and of helium to 13%.[Note 1]

The abundance of chemical elements in the universe is dominated by the large amounts of hydrogen and helium which were produced in the Big Bang. Remaining elements, making up only about 2% of the universe, were largely produced by supernovae and certain red giant stars. Lithium, beryllium and boron are rare because although they are produced by nuclear fusion, they are then destroyed by other reactions in the stars.[1][2] The elements from carbon to iron are relatively more common in the universe because of the ease of making them in supernova nucleosynthesis. Elements of higher atomic number than iron (element 26) become progressively more rare in the universe, because they increasingly absorb stellar energy in being produced. Elements with even atomic numbers are generally more common than their neighbors in the periodic table, also due to favorable energetics of formation.

The abundance of elements in the Sun and outer planets is similar to that in the universe. Due to solar heating, the elements of Earth and the inner rocky planets of the Solar System have undergone an additional depletion of volatile hydrogen, helium, neon, nitrogen, and carbon (which volatilizes as methane). The crust, mantle, and core of the Earth show evidence of chemical segregation plus some sequestration by density. Lighter silicates of aluminum are found in the crust, with more magnesium silicate in the mantle, while metallic iron and nickel compose the core. The abundance of elements in specialized environments, such as atmospheres, or oceans, or the human body, are primarily a product of chemical interactions with the medium in which they reside.


Ten most common elements in the Milky Way Galaxy estimated spectroscopically[3]
Z Element Mass fraction (ppm)
1 Hydrogen 739,000
2 Helium 240,000
8 Oxygen 10,400
6 Carbon 4,600
10 Neon 1,340
26 Iron 1,090
7 Nitrogen 960
14 Silicon 650
12 Magnesium 580
16 Sulfur 440

The elements – that is, ordinary (baryonic) matter made of protons, neutrons, and electrons, are only a small part of the content of the Universe. Cosmological observations suggest that only 4.6% of the universe's energy (including the mass contributed by energy, E = mc² ↔ m = E / c²) comprises the visible baryonic matter that constitutes stars, planets, and living beings. The rest is thought to be made up of dark energy (68%) and dark matter (27%).[4] These are forms of matter and energy believed to exist on the basis of scientific theory and observational deductions, but they have not been directly observed and their nature is not well understood.

Most standard (baryonic) matter is found in intergalactic gas, stars, and interstellar clouds, in the form of atoms or ions (plasma), although it can be found in degenerate forms in extreme astrophysical settings, such as the high densities inside white dwarfs and neutron stars.

Hydrogen is the most abundant element in the Universe; helium is second. However, after this, the rank of abundance does not continue to correspond to the atomic number; oxygen has abundance rank 3, but atomic number 8. All others are substantially less common.

The abundance of the lightest elements is well predicted by the standard cosmological model, since they were mostly produced shortly (i.e., within a few hundred seconds) after the Big Bang, in a process known as Big Bang nucleosynthesis. Heavier elements were mostly produced much later, inside of stars.

Hydrogen and helium are estimated to make up roughly 74% and 24% of all baryonic matter in the universe respectively. Despite comprising only a very small fraction of the universe, the remaining "heavy elements" can greatly influence astronomical phenomena. Only about 2% (by mass) of the Milky Way galaxy's disk is composed of heavy elements.

These other elements are generated by stellar processes.[5][6][7] In astronomy, a "metal" is any element other than hydrogen or helium. This distinction is significant because hydrogen and helium are the only elements that were produced in significant quantities in the Big Bang. Thus, the metallicity of a galaxy or other object is an indication of stellar activity, after the Big Bang.

In general, elements up to iron are made in large stars in the process of becoming supernovae. Iron-56 is particularly common, since it is the most stable element that can easily be made from alpha particles (being a product of decay of radioactive nickel-56, ultimately made from 14 helium nuclei). Elements heavier than iron are made in energy-absorbing processes in large stars, and their abundance in the universe (and on Earth) generally decreases with increasing atomic number.

Periodic table showing the cosmogenic origin of each element.
Periodic table showing the cosmogenic origin of each element.

Solar system

Most abundant nuclides
in the Solar System[8]
Nuclide A Mass fraction in parts per million Atom fraction in parts per million
Hydrogen-1 1 705,700 909,964
Helium-4 4 275,200 88,714
Oxygen-16 16 9,592 477
Carbon-12 12 3,032 326
Nitrogen-14 14 1,105 102
Neon-20 20 1,548 100
Other nuclides: 3,879 149
Silicon-28 28 653 30
Magnesium-24 24 513 28
Iron-56 56 1,169 27
Sulfur-32 32 396 16
Helium-3 3 35 15
Hydrogen-2 2 23 15
Neon-22 22 208 12
Magnesium-26 26 79 4
Carbon-13 13 37 4
Magnesium-25 25 69 4
Aluminium-27 27 58 3
Argon-36 36 77 3
Calcium-40 40 60 2
Sodium-23 23 33 2
Iron-54 54 72 2
Silicon-29 29 34 2
Nickel-58 58 49 1
Silicon-30 30 23 1
Iron-57 57 28 1

The following graph (note log scale) shows abundance of elements in the Solar System. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass.[3] Nearby galaxies that have evolved along similar lines have a corresponding enrichment of elements heavier than hydrogen and helium. The more distant galaxies are being viewed as they appeared in the past, so their abundances of elements appear closer to the primordial mixture. Since physical laws and processes are uniform throughout the universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.

The abundance of elements is in keeping with their origin from the Big Bang and nucleosynthesis in a number of progenitor supernova stars. Very abundant hydrogen and helium are products of the Big Bang, while the next three elements are rare since they had little time to form in the Big Bang and are not made in stars (they are, however, produced in small quantities by breakup of heavier elements in interstellar dust, as a result of impact by cosmic rays).

Beginning with carbon, elements have been produced in stars by buildup from alpha particles (helium nuclei), resulting in an alternatingly larger abundance of elements with even atomic numbers (these are also more stable). The effect of odd-numbered chemical elements generally being more rare in the universe was empirically noticed in 1914, and is known as the Oddo-Harkins rule.

Estimated abundances of the chemical elements in the Solar System (logarithmic scale).
Estimated abundances of the chemical elements in the Solar System (logarithmic scale).

Relation to nuclear binding energy

Loose correlations have been observed between estimated elemental abundances in the universe and the nuclear binding energy curve. Roughly speaking, the relative stability of various atomic nuclides has exerted a strong influence on the relative abundance of elements formed in the Big Bang, and during the development of the universe thereafter.[9] See the article about nucleosynthesis for the explanation on how certain nuclear fusion processes in stars (such as carbon burning, etc.) create the elements heavier than hydrogen and helium.

A further observed peculiarity is the jagged alternation between relative abundance and scarcity of adjacent atomic numbers in the elemental abundance curve, and a similar pattern of energy levels in the nuclear binding energy curve. This alternation is caused by the higher relative binding energy (corresponding to relative stability) of even atomic numbers compared with odd atomic numbers and is explained by the Pauli Exclusion Principle.[10] The semi-empirical mass formula (SEMF), also called Weizsäcker's formula or the Bethe-Weizsäcker mass formula, gives a theoretical explanation of the overall shape of the curve of nuclear binding energy.[11]


The Earth formed from the same cloud of matter that formed the Sun, but the planets acquired different compositions during the formation and evolution of the solar system. In turn, the natural history of the Earth caused parts of this planet to have differing concentrations of the elements.

The mass of the Earth is approximately 5.98×1024 kg. In bulk, by mass, it is composed mostly of iron (32.1%), oxygen (30.1%), silicon (15.1%), magnesium (13.9%), sulfur (2.9%), nickel (1.8%), calcium (1.5%), and aluminium (1.4%); with the remaining 1.2% consisting of trace amounts of other elements.[12]

The bulk composition of the Earth by elemental-mass is roughly similar to the gross composition of the solar system, with the major differences being that Earth is missing a great deal of the volatile elements hydrogen, helium, neon, and nitrogen, as well as carbon which has been lost as volatile hydrocarbons. The remaining elemental composition is roughly typical of the "rocky" inner planets, which formed in the thermal zone where solar heat drove volatile compounds into space. The Earth retains oxygen as the second-largest component of its mass (and largest atomic-fraction), mainly from this element being retained in silicate minerals which have a very high melting point and low vapor pressure.

Estimated abundances of chemical elements in the Earth.[13] The right two columns give the fraction of the mass in parts per million (ppm) and the fraction by number of atoms in parts per billion (ppb).
Atomic Number Name Symbol Mass fraction (ppm) Atomic fraction (ppb)[13]
8 oxygen O 297000 482,000,000
12 magnesium Mg 154000 164,000,000
14 silicon Si 161000 150,000,000
26 iron Fe 319000 148,000,000
13 aluminum Al 15900 15,300,000
20 calcium Ca 17100 11,100,000
28 nickel Ni 18220 8,010,000
1 hydrogen H 260 6,700,000
16 sulfur S 6350 5,150,000
24 chromium Cr 4700 2,300,000
11 sodium Na 1800 2,000,000
6 carbon C 730 1,600,000
15 phosphorus P 1210 1,020,000
25 manganese Mn 1700 800,000
22 titanium Ti 810 440,000
27 cobalt Co 880 390,000
19 potassium K 160 110,000
17 chlorine Cl 76 56,000
23 vanadium V 105 53,600
7 nitrogen N 25 46,000
29 copper Cu 60 25,000
30 zinc Zn 40 16,000
9 fluorine F 10 14,000
21 scandium Sc 11 6,300
3 lithium Li 1.10 4,100
38 strontium Sr 13 3,900
32 germanium Ge 7.00 2,500
40 zirconium Zr 7.10 2,000
31 gallium Ga 3.00 1,000
34 selenium Se 2.70 890
56 barium Ba 4.50 850
39 yttrium Y 2.90 850
33 arsenic As 1.70 590
5 boron B 0.20 480
42 molybdenum Mo 1.70 460
44 ruthenium Ru 1.30 330
78 platinum Pt 1.90 250
46 palladium Pd 1.00 240
58 cerium Ce 1.13 210
60 neodymium Nd 0.84 150
4 beryllium Be 0.05 140
41 niobium Nb 0.44 120
76 osmium Os 0.90 120
77 iridium Ir 0.90 120
37 rubidium Rb 0.40 120
35 bromine Br 0.30 97
57 lanthanum La 0.44 82
66 dysprosium Dy 0.46 74
64 gadolinium Gd 0.37 61
52 tellurium Te 0.30 61
45 rhodium Rh 0.24 61
50 tin Sn 0.25 55
62 samarium Sm 0.27 47
68 erbium Er 0.30 47
70 ytterbium Yb 0.30 45
59 praseodymium Pr 0.17 31
82 lead Pb 0.23 29
72 hafnium Hf 0.19 28
74 tungsten W 0.17 24
79 gold Au 0.16 21
48 cadmium Cd 0.08 18
63 europium Eu 0.10 17
67 holmium Ho 0.10 16
47 silver Ag 0.05 12
65 terbium Tb 0.07 11
51 antimony Sb 0.05 11
75 rhenium Re 0.08 10
53 iodine I 0.05 10
69 thulium Tm 0.05 7
55 cesium Cs 0.04 7
71 lutetium Lu 0.05 7
90 thorium Th 0.06 6
73 tantalum Ta 0.03 4
80 mercury Hg 0.02 3
92 uranium U 0.02 2
49 indium In 0.01 2
81 thallium Tl 0.01 2
83 bismuth Bi 0.01 1


Elemental abundances
Abundance (atom fraction) of the chemical elements in Earth's upper continental crust as a function of atomic number. The rarest elements in the crust (shown in yellow) are rare due to a combination of factors: all but 1 of them are the densest siderophiles (iron-loving) elements in the Goldschmidt classification of elements, meaning they have a tendency to mix well with iron, depleting them by being relocated deeper into the Earth's core. Their abundance in meteoroids is higher. Additionally, tellurium has been depleted from the crust due to formation of volatile hydrides.

The mass-abundance of the nine most abundant elements in the Earth's crust is approximately: oxygen 46%, silicon 28%, aluminum 8.2%, iron 5.6%, calcium 4.2%, sodium 2.5%, magnesium 2.4%, potassium 2.0%, and titanium 0.61%. Other elements occur at less than 0.15%. For a complete list, see abundance of elements in Earth's crust.

The graph at right illustrates the relative atomic-abundance of the chemical elements in Earth's upper continental crust— the part that is relatively accessible for measurements and estimation.

Many of the elements shown in the graph are classified into (partially overlapping) categories:

  1. rock-forming elements (major elements in green field, and minor elements in light green field);
  2. rare earth elements (lanthanides, La-Lu, and Y; labeled in blue);
  3. major industrial metals (global production >~3×107 kg/year; labeled in red);
  4. precious metals (labeled in purple);
  5. the nine rarest "metals" — the six platinum group elements plus Au, Re, and Te (a metalloid) — in the yellow field. These are rare in the crust from being soluble in iron and thus concentrated in the Earth's core.

Note that there are two breaks where the unstable (radioactive) elements technetium (atomic number 43) and promethium (atomic number 61) would be. These elements are surrounded by stable elements, yet both have relatively short half lives (~ 4 million years and ~ 18 years respectively). These are thus extremely rare, since any primordial initial fractions of these in pre-Solar System materials have long since decayed. These two elements are now only produced naturally through the spontaneous fission of very heavy radioactive elements (for example, uranium, thorium, or the trace amounts of plutonium that exist in uranium ores), or by the interaction of certain other elements with cosmic rays. Both technetium and promethium have been identified spectroscopically in the atmospheres of stars, where they are produced by ongoing nucleosynthetic processes.

There are also breaks in the abundance graph where the six noble gases would be, since they are not chemically bound in the Earth's crust, and they are only generated by decay chains from radioactive elements in the crust, and are therefore extremely rare there.

The eight naturally occurring very rare, highly radioactive elements (polonium, astatine, francium, radium, actinium, protactinium, neptunium, and plutonium) are not included, since any of these elements that were present at the formation of the Earth have decayed away eons ago, and their quantity today is negligible and is only produced from the radioactive decay of uranium and thorium.

Oxygen and silicon are notably the most common elements in the crust. On Earth and in rocky planets in general, silicon and oxygen are far more common than their cosmic abundance. The reason is that they combine with each other to form silicate minerals. In this way, they are the lightest of all of the two-percent "astronomical metals" (i.e., non-hydrogen and helium elements) to form a solid that is refractory to the Sun's heat, and thus cannot boil away into space. All elements lighter than oxygen have been removed from the crust in this way.

Rare-earth elements

"Rare" earth elements is a historical misnomer. The persistence of the term reflects unfamiliarity rather than true rarity. The more abundant rare earth elements are similarly concentrated in the crust compared to commonplace industrial metals such as chromium, nickel, copper, zinc, molybdenum, tin, tungsten, or lead. The two least abundant rare earth elements (thulium and lutetium) are nearly 200 times more common than gold. However, in contrast to the ordinary base and precious metals, rare earth elements have very little tendency to become concentrated in exploitable ore deposits. Consequently, most of the world's supply of rare earth elements comes from only a handful of sources. Furthermore, the rare earth metals are all quite chemically similar to each other, and they are thus quite difficult to separate into quantities of the pure elements.

Differences in abundances of individual rare earth elements in the upper continental crust of the Earth represent the superposition of two effects, one nuclear and one geochemical. First, the rare earth elements with even atomic numbers (58Ce, 60Nd, ...) have greater cosmic and terrestrial abundances than the adjacent rare earth elements with odd atomic numbers (57La, 59Pr, ...). Second, the lighter rare earth elements are more incompatible (because they have larger ionic radii) and therefore more strongly concentrated in the continental crust than the heavier rare earth elements. In most rare earth ore deposits, the first four rare earth elements – lanthanum, cerium, praseodymium, and neodymium – constitute 80% to 99% of the total amount of rare earth metal that can be found in the ore.


The mass-abundance of the eight most abundant elements in the Earth's mantle (see main article above) is approximately: oxygen 45%, magnesium 23%, silicon 22%, iron 5.8%, calcium 2.3%, aluminum 2.2%, sodium 0.3%, potassium 0.3%.

The mantle differs in elemental composition from the crust in having a great deal more magnesium and significantly more iron, while having much less aluminum and sodium.


Due to mass segregation, the core of the Earth is believed to be primarily composed of iron (88.8%), with smaller amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace elements.[12]


The most abundant elements in the ocean by proportion of mass in percent are oxygen (85.84), hydrogen (10.82), chlorine (1.94), sodium (1.08), magnesium (0.1292), sulfur (0.091), calcium (0.04), potassium (0.04), bromine (0.0067), carbon (0.0028), and boron (0.00043).


The order of elements by volume-fraction (which is approximately molecular mole-fraction) in the atmosphere is nitrogen (78.1%), oxygen (20.9%),[14] argon (0.96%), followed by (in uncertain order) carbon and hydrogen because water vapor and carbon dioxide, which represent most of these two elements in the air, are variable components. Sulfur, phosphorus, and all other elements are present in significantly lower proportions.

According to the abundance curve graph (above right), argon, a significant if not major component of the atmosphere, does not appear in the crust at all. This is because the atmosphere has a far smaller mass than the crust, so argon remaining in the crust contributes little to mass-fraction there, while at the same time buildup of argon in the atmosphere has become large enough to be significant.

Urban soils

For a complete list of the abundance of elements in urban soils, see Abundances of the elements (data page)#Urban soils.

Human body

Elemental abundance in the human body
Element Proportion (by mass)
Oxygen 65
Carbon 18
Hydrogen 10
Nitrogen 3
Calcium 1.5
Phosphorus 1.2
Potassium 0.2
Sulfur 0.2
Chlorine 0.2
Sodium 0.1
Magnesium 0.05
Iron < 0.05
Cobalt < 0.05
Copper < 0.05
Zinc < 0.05
Iodine < 0.05
Selenium < 0.01

By mass, human cells consist of 65–90% water (H2O), and a significant portion of the remainder is composed of carbon-containing organic molecules. Oxygen therefore contributes a majority of a human body's mass, followed by carbon. Almost 99% of the mass of the human body is made up of six elements: hydrogen (H), carbon (C), nitrogen (N), oxygen (O), calcium (Ca), and phosphorus (P). The next 0.75% is made up of the next five elements: potassium (K), sulfur (S), chlorine (Cl), sodium (Na), and magnesium (Mg). CHNOPS for short. Only 17 elements are known for certain to be necessary to human life, with one additional element (fluorine) thought to be helpful for tooth enamel strength. A few more trace elements may play some role in the health of mammals. Boron and silicon are notably necessary for plants but have uncertain roles in animals. The elements aluminium and silicon, although very common in the earth's crust, are conspicuously rare in the human body.[15]

Below is a periodic table highlighting nutritional elements.[16]

Nutritional elements in the periodic table
H   He
Li Be   B C N O F Ne
Na Mg   Al Si P S Cl Ar
K Ca Sc   Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y   Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba La * Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac ** Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
  * Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
  ** Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
  Quantity elements
  Essential trace elements
  Deemed essential trace element by U.S., not by European Union
  Suggested function from deprivation effects or active metabolic handling, but no clearly-identified biochemical function in humans
  Limited circumstantial evidence for trace benefits or biological action in mammals
  No evidence for biological action in mammals, but essential in some lower organisms.
(In the case of lanthanum, the definition of an essential nutrient as being indispensable and irreplaceable is not completely applicable due to the extreme similarity of the lanthanides. Thus Ce, Pr, and Nd may be substituted for La without ill effects for organisms using La, and the smaller Sm, Eu, and Gd may also be similarly substituted but cause slower growth.)

See also



  1. ^ Vangioni-Flam, Elisabeth; Cassé, Michel (2012). Spite, Monique, ed. Galaxy Evolution: Connecting the Distant Universe with the Local Fossil Record. Springer Science & Business Media. pp. 77–86. ISBN 9401142130.
  2. ^ Trimble, Virginia (1996). "The Origin and Evolution of the Chemical Elements". In Malkan, Matthew A.; Zuckerman, Ben. The origin and evolution of the universe. Sudbury, Mass.: Jones and Bartlett Publishers. p. 101. ISBN 0-7637-0030-4.
  3. ^ a b Croswell, Ken (February 1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. Archived from the original on 2011-05-13.
  4. ^ What is Dark Energy? Archived 2016-01-15 at the Wayback Machine, Space.com, 1 May 2013
  5. ^ Suess, Hans; Urey, Harold (1956). "Abundances of the Elements". Reviews of Modern Physics. 28: 53. Bibcode:1956RvMP...28...53S. doi:10.1103/RevModPhys.28.53.
  6. ^ Cameron, A.G.W. (1973). "Abundances of the elements in the solar system". Space Science Reviews. 15: 121. Bibcode:1973SSRv...15..121C. doi:10.1007/BF00172440.
  7. ^ Anders, E; Ebihara, M (1982). "Solar-system abundances of the elements". Geochimica et Cosmochimica Acta. 46 (11): 2363. Bibcode:1982GeCoA..46.2363A. doi:10.1016/0016-7037(82)90208-3.
  8. ^ Arnett, David (1996). Supernovae and Nucleosynthesis (First ed.). Princeton, New Jersey: Princeton University Press. p. 11. ISBN 0-691-01147-8. OCLC 33162440.
  9. ^ Bell, Jerry A.; GenChem Editorial/Writing Team (2005). "Chapter 3: Origin of Atoms". Chemistry: a project of the American Chemical Society. New York [u.a.]: Freeman. pp. 191–193. ISBN 978-0-7167-3126-9. Correlations between abundance and nuclear binding energy [Subsection title]
  10. ^ Bell, Jerry A.; GenChem Editorial/Writing Team (2005). "Chapter 3: Origin of Atoms". Chemistry: a project of the American Chemical Society. New York [u.a.]: Freeman. p. 192. ISBN 978-0-7167-3126-9. The higher abundance of elements with even atomic numbers [Subsection title]
  11. ^ Bailey, David. "Semi-empirical Nuclear Mass Formula". PHY357: Strings & Binding Energy. University of Toronto. Archived from the original on 2011-07-24. Retrieved 2011-03-31.
  12. ^ a b Morgan, J. W.; Anders, E. (1980). "Chemical composition of Earth, Venus, and Mercury". Proceedings of the National Academy of Sciences. 77 (12): 6973–6977. Bibcode:1980PNAS...77.6973M. doi:10.1073/pnas.77.12.6973. PMC 350422. PMID 16592930.
  13. ^ a b William F McDonough The composition of the Earth. quake.mit.edu, archived by the Internet Archive Wayback Machine.
  14. ^ Zimmer, Carl (3 October 2013). "Earth's Oxygen: A Mystery Easy to Take for Granted". The New York Times. Archived from the original on 3 October 2013. Retrieved 3 October 2013.
  15. ^ Table data from Chang, Raymond (2007). Chemistry (Ninth ed.). McGraw-Hill. p. 52. ISBN 0-07-110595-6.
  16. ^ Nielsen, Forrest H. (1998). "Ultratrace minerals.". In Maurice E. Shils; James A. Olsen; Moshe Shine; A. Catharine Ross. Modern nutrition in health and disease. Baltimore: Lippincott Williams & Wilkins. pp. 283–303. ISBN 978-0683307696.


  1. ^ Below Jupiter's outer atmosphere, volume fractions are significantly different from mole fractions due to high temperatures (ionization and disproportionation) and high density where the Ideal Gas Law is inapplicable.


External links


Abundance may refer to:

In science and technology:

Abundance (economics), the opposite of scarcities

Abundance (ecology), the relative representation of a species in a community

Abundance (programming language), a Forth-like computer programming language

Abundance, a property of abundant numbers

In chemistry:

Abundance (chemistry), when a substance in a reaction is present in high quantities

Abundance of the chemical elements, a measure of how common elements are

Natural abundance, the natural prevalence of different isotopes of an element on Earth

Abundance of elements in Earth's crustIn literature:

Abundance (play), a 1990 stage play written by Beth Henley

Al-Kawthar ("Abundance"), the 108th sura of the Qur'an

Abundance: The Future Is Better Than You Think, a 2012 book by Peter Diamandis and Steven KotlerOther usesAbundance Generation, a renewable energy investment platform

Fountain de la Abundancia, a former fountain in Madrid

Abundance, Royal Abundance and Abundance Declared, bids in the card game Solo whist; sometimes spelled "abondance"


CHON is a mnemonic acronym for the four most common elements in living organisms: carbon, hydrogen, oxygen, and nitrogen.

The acronym CHNOPS, which stands for carbon, hydrogen, nitrogen, oxygen, phosphorus, sulfur, represents the six most important chemical elements whose covalent combinations make up most biological molecules on Earth. Most of the elements are nonmetals.

Sulfur is contained in the amino acids cysteine and methionine.Phosphorus is contained in phospholipids, a class of lipids that are a major component of all cell membranes, as they can form lipid bilayers, which keep ions, proteins, and other molecules where they are needed for cell function, and prevent them from diffusing into areas where they should not be. Phosphate groups are also an essential component of the backbone of nucleic acids and are required to form ATP – the main molecule used as energy powering the cell in all living creatures.Carbonaceous asteroids are rich in CHON elements.

These asteroids are the most common type, and frequently collide with Earth as meteorites. Such collisions were especially common early in Earth's history, and these impactors may have been crucial in the formation of the planet's oceans.The simplest compounds to contain all of the CHON elements are fulminic acid and isocyanic acid (the latter of which is much more stable), having one of each atom.

IC 4651

IC 4651 is an open cluster of stars located about 2,900 light years distant in the constellation Ara. It was first catalogued by John Louis Emil Dreyer in his 1895 version of the Index Catalogue. This is an intermediate age cluster that is 1.2 ± 0.2 billion years old. Compared to the Sun, the members of this cluster have a higher abundance of the chemical elements other than hydrogen and helium. The combined mass of the active stars in this cluster is about 630 times the mass of the Sun.The currently known active stars in this cluster form only about 7% of the cluster's original mass. Of the remainder, about 35% of the mass consists of stars that have evolved into white dwarfs or other stellar remnants. The remainder of lost mass consists of stars that have migrated away from the main body of the cluster or have been lost completely.

Molar ionization energies of the elements

These tables list values of molar ionization energies, measured in kJ mol−1. This is the energy per mole necessary to remove electrons from gaseous atoms or atomic ions. The first molar ionization energy applies to the neutral atoms. The second, third, etc., molar ionization energy applies to the further removal of an electron from a singly, doubly, etc., charged ion. For ionization energies measured in the unit eV, see Ionization energies of the elements (data page). All data from rutherfordium onwards is predicted.

Natural abundance

In physics, natural abundance (NA) refers to the abundance of isotopes of a chemical element as naturally found on a planet. The relative atomic mass (a weighted average, weighted by mole-fraction abundance figures) of these isotopes is the atomic weight listed for the element in the periodic table. The abundance of an isotope varies from planet to planet, and even from place to place on the Earth, but remains relatively constant in time (on a short-term scale).

As an example, uranium has three naturally occurring isotopes: 238U, 235U and 234U. Their respective natural mole-fraction abundances are 99.2739–99.2752%, 0.7198–0.7202%, and 0.0050–0.0059%. For example, if 100,000 uranium atoms were analyzed, one would expect to find approximately 99,274 238U atoms, approximately 720 235U atoms, and very few (most likely 5 or 6) 234U atoms. This is because 238U is much more stable than 235U or 234U, as the half-life of each isotope reveals: 4.468 × 109 years for 238U compared with 7.038 × 108 years for 235U and 245,500 years for 234U.

Exactly because the different uranium isotopes have different half-lives, when the Earth was younger, the isotopic composition of uranium was different. As an example, 1.7×109 years ago the NA of 235U was 3.1% compared with today's 0.7%, and for that reason a natural nuclear fission reactor was able to form, something that cannot happen today.

However, the natural abundance of a given isotope is also affected by the probability of its creation in nucleosynthesis (as in the case of samarium; radioactive 147Sm and 148Sm are much more abundant than stable 144Sm) and by production of a given isotope as a daughter of natural radioactive isotopes (as in the case of radiogenic isotopes of lead).

Neutron capture nucleosynthesis

Neutron capture nucleosynthesis describes two nucleosynthesis pathways: the r-process and the s-process, for rapid and slow neutron captures, respectively. R-process describes neutron capture in a region of high neutron flux, such as during supernova nucleosynthesis after core-collapse, and yields neutron-rich nuclides. S-process describes neutron capture that is slow relative to the rate of beta decay, as for stellar nucleosynthesis in some stars, and yields nuclei with stable nuclear shells. Each process is responsible for roughly half of the observed abundances of elements heavier than iron. The importance of neutron capture to the observed abundance of the chemical elements was first described in 1957 in the B2FH paper.

Nuclear astrophysics

Nuclear astrophysics is an interdisciplinary branch of physics involving close collaboration among researchers in various subfields of nuclear physics and astrophysics: notably stellar modeling; measurement and theoretical estimation of nuclear reaction rates; physical cosmology and cosmochemistry; gamma ray, optical and X-ray astronomy; and extending our knowledge about nuclear lifetimes and masses. In general terms, nuclear astrophysics aims to understand the origin of the chemical elements and the energy generation in stars.

Oddo–Harkins rule

The Oddo–Harkins rule holds that an element with an even atomic number (such as carbon: element 6) is more abundant than both elements with the adjacently larger and smaller odd atomic numbers (such as boron: element 5 and nitrogen: element 7, respectively for the carbon). This tendency of the abundance of the chemical elements was first reported by Giuseppe Oddo in 1914 and William Draper Harkins in 1917.

Speeds of sound of the elements

The speed of sound in any chemical element in the fluid phase has one temperature-dependent value. In the solid phase, different types of sound wave may be propagated, each with its own speed: among these types of wave are longitudinal (as in fluids), transversal, and (along a surface or plate) extensional.

Stellar chemistry

Stellar chemistry is the study of the chemical composition of astronomical objects; stars in particular, hence the name stellar chemistry. The significance of stellar chemical composition is an open ended question at this point. Some research asserts that a greater abundance of certain elements (such as carbon, sodium, silicon, and magnesium) in the stellar mass are necessary for a star's inner solar system to be habitable over long periods of time. The hypothesis being that the "abundance of these elements make the star cooler and cause it to evolve more slowly, thereby giving planets in its habitable zone more time to develop life as we know it." Stellar abundance of oxygen also appears to be critical to the length of time newly developed planets exist in a habitable zone around their host star. Researchers postulate that if our own sun had a lower abundance of oxygen, the Earth would have ceased to "live" in a habitable zone a billion years ago, long before complex organisms had the opportunity to evolve.

Surveyor 5

Surveyor 5 was the fifth lunar lander of the American unmanned Surveyor program sent to explore the surface of the Moon. Surveyor 5 landed on Mare Tranquillitatis. A total of 19,049 images were transmitted to Earth.

Surveyor 6

Surveyor 6 was the sixth lunar lander of the American unmanned Surveyor program that reached the surface of the Moon. Surveyor 6 landed on the Sinus Medii. A total of 30,027 images were transmitted to Earth.

This spacecraft was the fourth of the Surveyor series to successfully achieve a soft landing on the Moon, obtain post landing television pictures, determine the abundance of the chemical elements in the lunar soil, obtain touchdown dynamics data, obtain thermal and radar reflectivity data, and conduct a Vernier engine erosion experiment. Virtually identical to Surveyor 5, this spacecraft carried a television camera, a small bar magnet attached to one footpad, and an alpha-scattering instrument as well as the necessary engineering equipment. It landed on November 10, 1967, in Sinus Medii, 0.49 deg in latitude and 1.40 deg w longitude (selenographic coordinates)–the center of the Moon's visible hemisphere. The spacecraft accomplished all planned objectives. The successful completion of this mission satisfied the Surveyor program's obligation to the Apollo project. On November 24, 1967, the spacecraft was shut down for the two-week lunar night. Contact was made on December 14, 1967, but no useful data was obtained.

Lunar soil surveys were completed using photographic and alpha particle backscattering methods. A similar instrument, the APXS, was used onboard several Mars missions.In a further test of space technology, Surveyor 6's engines were restarted and burned for 2.5 seconds in the first lunar liftoff on November 17 at 10:32 UTC. This created 150 lbf (700 N) of thrust and lifted the vehicle 12 feet (4 m) from the lunar surface. After moving west eight feet, (2.5 m) the spacecraft once again successfully soft landed and continued functioning as designed.

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