Abundance of the chemical elements

The abundance of a chemical element measures how relatively common (or rare) the element is, or how much of the element is present in a given environment by comparison to all other elements. Abundance may be variously measured by the mass-fraction (the same as weight fraction), or mole-fraction (fraction of atoms by numerical count, or sometimes fraction of molecules in gases), or by volume-fraction. Measurement by volume-fraction is a common abundance measure in mixed gases such as planetary atmospheres, and is close to molecular mole-fraction for ideal gas mixtures (i.e., gas mixtures at relatively low densities and pressures).

For example, the mass-fraction abundance of oxygen in water is about 89%, because that is the fraction of water's mass which is oxygen. However, the mole-fraction abundance of oxygen in water is only 33% because only 1 atom of 3 in water is an oxygen atom. In the universe as a whole, and in the atmospheres of gas-giant planets such as Jupiter, the mass-fraction abundances of hydrogen and helium are about 74% and 23–25% respectively, while the (atomic) mole-fractions of these elements are closer to 92% and 8%. However, since hydrogen is diatomic while helium is not, in the conditions of Jupiter's outer atmosphere, the molecular mole-fraction (fraction of total gas molecules, or fraction of atmosphere by volume) of hydrogen in the outer atmosphere of Jupiter is about 86%, and for helium, 13%.

Most abundances in this article are given as mass-fraction abundances.

Contents

Abundance of elements in the Universe

Ten most common elements in the Milky Way Galaxy estimated spectroscopically[1]
Z Element Mass fraction in parts per million
1 Hydrogen 739,000 71 × mass of Oxygen (red bar)
2 Helium 240,000 23 × mass of Oxygen (red bar)
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
 
See also: Stellar population, Cosmochemistry, and Astrochemistry

The elements – that is, ordinary (baryonic) matter made out of protons and neutrons (as well as electrons) – are only a small part of the content of the Universe. Cosmological observations suggest that only 4.6% of the universe comprises the visible baryonic matter which constitutes stars, planets and living beings. The rest is made up of dark energy (72%) and dark matter (23%).[2] The latter are forms of matter and energy believed to exist on the basis of theory and observational deductions, but their details are still the subject of research. They have not yet been directly observed and are not well understood.

Most standard (baryonic) matter is found in stars and interstellar clouds, in the form of atoms or ions (plasma), although other unusual kinds of matter can be found in astrophysical settings, such as the high densities inside white dwarfs and neutron stars.

Hydrogen is the most abundant element in the known 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 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.[3][4][5] In astronomy, a "metal" is any element other than hydrogen, helium or lithium. This distinction is significant because hydrogen and helium (together with trace amounts of lithium) are the only elements that occur naturally without the nuclear fusion activity of stars. Thus, the metallicity of a galaxy or other object is an indication of past stellar activity.

The following graph (note log scale) shows abundance of elements in our solar system. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass.[1] 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. As physical laws and processes appear common throughout the visible universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.

The abundance of elements in the Solar System (see graph above) 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 quanties 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). In general, such elements up to iron are made in large stars in the process of becoming supernovas. 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 their atomic number.

Most abundant isotopes in the Solar System[6]
Isotope 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 5,920 477
Carbon-12 12 3,032 326
Nitrogen-14 14 1,105 102
Neon-20 20 1,548 100
Other elements: 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
Aluminum-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

Elemental abundance and 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 isotopes has exerted a strong influence on the relative abundance of elements formed in the Big Bang, and during the development of the universe thereafter. [7] 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 to odd atomic numbers, and is explained by the Pauli Exclusion Principle. [8] The semi-empirical mass formula (SEMF), also called Weizsäcker's formula or the Bethe-Weizsäcker mass formula, gives a theoretical explanation of the curve of nuclear binding energy. [9]

Abundance of elements in the Earth

See also: Earth#Chemical composition

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. 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. Due to mass segregation, the core region 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.[10]

Earth's bulk (total) elemental abundance

Click "show" at right, to show numerical values in full table:

An estimate[11] of the elemental abundances in the total mass of the Earth. Note that numbers are estimates, and they will vary depending on source and method of estimation. Order of magnitude of data can roughly be relied upon. ppb (atoms) is parts per billion, meaning that is the number of atoms of a given element in every billion atoms in the Earth.

Earth's crustal elemental abundance

Main article: Abundance of elements in Earth's crust

This graph illustrates the relative abundance of the chemical elements in Earth's upper continental crust, which 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.

Note that there are two breaks where the unstable elements technetium (atomic number: 43) and promethium (atomic number: 61) would be. These are both extremely rare, since on Earth they are only produced 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 of the first two of these elements have been identified spectroscopically in the atmospheres of stars, where they are produced by ongoing nucleosynthetic processes. There are also breaks 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 and are therefore extremely rare there. The six very rare, highly radioactive elements (polonium, astatine, francium, radium, actinium, and protactinium) 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.

Oxygen and silicon are notably quite common elements. They have frequently combined with each other to form common silicate minerals.

Rare-earth elemental abundance

"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 each similar in crustal concentration 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.

Oceanic elemental abundance

Earth's ocean water elemental abundance
Element Proportion (by mass)
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%
 
For a complete list of the abundance of elements in the ocean, see Abundances of the elements (data page)#Sea water.

Atmospheric elemental abundance

The order of elements by volume-fraction (which is approximately molecular mole-fraction) in the atmosphere is nitrogen (78.1%), oxygen (20.9%), 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.

Human body elemental abundance

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%
Main article: Chemical makeup of the human body

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: oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus. The next 0.75% is made up of the next five elements: potassium, sulfur, chlorine, sodium, and magnesium. 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 appear to be necessary to mammals in carefully-dust free conditions. Boron and silicon are notably necessary for plants but have uncertain roles in animals. The elements aluminum and silicon, although very common in the earth's crust, are conspicuously rare in the human body."[12]

Periodic table highlighting nutritional elements[13]

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 Uut Uuq Uup Uuh Uus Uuo
 
  * 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
The four organic basic elements Quantity elements Essential trace elements Suggested function from biochemistry and handling but no identified biological function in humans

See also

References

Footnotes

  1. ^ a b Croswell, Ken (February 1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. http://kencroswell.com/alchemy.html. 
  2. ^ WMAP- Content of the Universe
  3. ^ 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. 
  4. ^ Cameron, A.G.W. (1973). "Abundances of the elements in the solar system". Space Science Reviews 15. Bibcode 1973SSRv...15..121C. doi:10.1007/BF00172440. 
  5. ^ 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. 
  6. ^ Arnett, David (1996). Supernovae and Nucleosynthesis (First ed.). Princeton, New Jersey: Princeton University Press. ISBN 0-691-01147-8. OCLC 33162440. http://books.google.com/?id=PXGWGnPPo0gC&printsec=frontcove. 
  7. ^ Bell, Jerry A.; GenChem Editorial/Writing Team Jerry Bell (2005). "Chapter 3: Origin of Atoms". Chemistry: a project of the American Chemical Society. New York [u.a.]: Freeman. pp. 191–193. ISBN 9780716731269. "Correlations between abundance and nuclear binding energy [Subsection title]" 
  8. ^ Bell, Jerry A.; GenChem Editorial/Writing Team Jerry Bell (2005). "Chapter 3: Origin of Atoms". Chemistry: a project of the American Chemical Society. New York [u.a.]: Freeman. pp. 192. ISBN 9780716731269. "The higher abundance of elements with even atomic numbers [Subsection title]" 
  9. ^ Bailey, David. "Semi-empirical Nuclear Mass Formula". PHY357: Strings & Binding Energy. University of Toronto. http://www.upscale.utoronto.ca/GeneralInterest/DBailey/SubAtomic/Lectures/LectF25/Lect25.htm. Retrieved 2011-03-31. 
  10. ^ Morgan, J. W.; Anders, E. (1980). "Chemical composition of Earth, Venus, and Mercury". Proceedings of the National Academy of Science 71 (12): 6973–6977. Bibcode 1980PNAS...77.6973M. doi:10.1073/pnas.77.12.6973. PMC 350422. PMID 16592930. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=350422. 
  11. ^ William F McDonough The composition of the Earth. quake.mit.edu
  12. ^ Table data from Chang, Raymond (2007). Chemistry, Ninth Edition. McGraw-Hill. pp. 52. ISBN 0-07-110595-6. 
  13. ^ Ultratrace minerals. Authors: Nielsen, Forrest H. USDA, ARS Source: Modern nutrition in health and disease / editors, Maurice E. Shils ... et al.. Baltimore : Williams & Wilkins, c1999., p. 283-303. Issue Date: 1999 URI: [1]

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