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 mass fraction (in commercial contexts often called weight fraction), by mole fraction (fraction of atoms by numerical count, or sometimes fraction of molecules in gases), or by 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 about 33% 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 during 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, despite their low atomic number, are rare because, although they are produced by nuclear fusion, they are destroyed by other reactions in the stars. [1] [2] Their natural occurrence is the result of cosmic ray spallation of carbon, nitrogen and oxygen in a type of nuclear fission reaction.The elements from carbon to iron are relatively more abundant in the universe because of the ease of making them in supernova nucleosynthesis. Elements of higher atomic number than iron (element 26) become progressively rarer in the universe, because they increasingly absorb stellar energy in their production. Also, elements with even atomic numbers are generally more common than their neighbors in the periodic table, due to favorable energetics of formation (see Oddo-Harkins rule), and among the lightest nuclides helium through sulfur the most abundant isotopes of equal number of protons and neutrons.
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 aluminium 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.
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 |
Total | 999,060 |
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 inductive reasoning based on observations, but they have not been directly observed and their nature is not well understood.
Most ordinary (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.
There are 80 known stable elements, and the lightest 16 comprise 99.9% of the ordinary matter of the universe. These same 16 elements, hydrogen through sulfur, fall on the initial linear portion of the Table of Nuclides (also called theSegrè plot), a plot of the proton versus neutron numbers of all matter both ordinary and exotic, containing hundreds of stable isotopes and thousands more that are unstable. The Segrè plot is initially linear because (aside from hydrogen) the vast majority of ordinary matter (99.4% in the Solar System [5] ) contains an equal number of protons and neutrons (Z=N). To be sure, 74% ordinary matter exists as mononucleonic protons (hydrogen). But when nucleons combine to form stable nuclides, they combine in a ratio of one part proton to one part neutron in 99.4% of ordinary matter. The structural basis of the equality of nucleon numbers in baryonic matter is one of the simplest and most profound unsolved mysteries of the atomic nucleus.
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. [6] [7] [8] 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 by large stars in the process of becoming supernovae, or by smaller stars in the process of dying. One type of Iron, Iron-56, is particularly common, since it is the most stable nuclide (in that it has the highest nuclear binding energy per nucleon) and 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.
The table shows the ten 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 apparently uniform throughout the universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.
As shown in the periodic table, the abundance of elements is in keeping with their origin. Very abundant hydrogen and helium are products of the Big Bang. The next three elements in the periodic table (lithium, beryllium, and boron) are rare, despite their low atomic number. They had little time to form in the Big Bang. They are produced in small quantities by nuclear fusion in dying stars or by breakup of heavier elements in interstellar dust, caused by cosmic ray spallation. In supernova stars, they are produced by nuclear fusion, but then destroyed by other reactions. [1]
Heavier elements, beginning with carbon, have been produced in dying or supernova stars by buildup from alpha particles (helium nuclei), contributing to 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.
The following graph (note log scale) shows abundance of elements in the Solar System.
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: | 3,616 | 172 | |
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 |
Loose correlations have been observed between estimated elemental abundances in the universe and the nuclear binding energy curve (also called the binding energy per nucleon). Roughly speaking, the relative stability of various atomic nuclides in withstanding the extremely energetic conditions of Big Bang Nucleosynthesis (BBN) 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 an explanation of 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 estimated abundances of the chemical elements in which the relative abundance of even atomic numbers is roughly 2 orders of magnitude greater than the relative abundance of odd atomic numbers (Oddo-Harkins rule). A similar alternation between even and odd atomic numbers can be observed in the nuclear binding energy curve in the neighborhood of carbon and oxygen, but here the loose correlation between relative abundance and binding energy ends. The binding energy for beryllium (an even atomic number), for example, is less than the binding energy for boron (an odd atomic number), as illustrated in the nuclear binding energy curve. Additionally, the alternation in the nuclear binding energy between even and odd atomic numbers resolves above oxygen as the graph increases steadily up to its peak at iron. 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. [10]
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 5.97×1024 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. [11]
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.
Atomic number | Name | Symbol | Mass fraction (ppm) | Atomic fraction (ppb) [12] |
---|---|---|---|---|
8 | oxygen | O | 297,000 | 482,000,000 |
12 | magnesium | Mg | 154,000 | 164,000,000 |
14 | silicon | Si | 161,000 | 150,000,000 |
26 | iron | Fe | 319,000 | 148,000,000 |
13 | aluminium | Al | 15,900 | 15,300,000 |
20 | calcium | Ca | 17,100 | 11,100,000 |
28 | nickel | Ni | 18,220 | 8,010,000 |
1 | hydrogen | H | 260 | 6,700,000 |
16 | sulfur | S | 6,350 | 5,150,000 |
24 | chromium | Cr | 4,700 | 2,300,000 |
11 | sodium | Na | 1,800 | 2,000,000 |
6 | carbon | C | 730 | 1,600,000 |
15 | phosphorus | P | 1,210 | 1,020,000 |
25 | manganese | Mn | 1,700 | 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 | caesium | 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 |
The mass-abundance of the nine most abundant elements in the Earth's crust is approximately: oxygen 46%, silicon 28%, aluminium 8.3%, 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:
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 their most stable isotopes 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 in the crust by decay chains from radioactive elements, 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. [13] Other cosmically common elements such as hydrogen, carbon and nitrogen form volatile compounds such as ammonia and methane that easily boil away into space from the heat of planetary formation and/or the Sun's light.
"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 stable 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 seven most abundant elements in the Earth's mantle is approximately: oxygen 44.3%, magnesium 22.3%, silicon 21.3%, iron 6.32%, calcium 2.48%, aluminium 2.29%, nickel 0.19%. [14]
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. [5]
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 (0.13%), sulfur (0.09%), calcium (0.04%), potassium (0.04%), bromine (0.007%), carbon (0.003%), and boron (0.0004%).
The order of elements by volume fraction (which is approximately molecular mole fraction) in the atmosphere is nitrogen (78.1%), oxygen (20.9%), [15] 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, 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.
For a complete list of the abundance of elements in urban soils, see Abundances of the elements (data page)#Urban soils.
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). 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. [16]
Below is a periodic table highlighting nutritional elements. [17]
Essential elements [18] [19] [20] [21] [22] | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
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 | * | Lu | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl | Pb | Bi | Po | At | Rn | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
Fr | Ra | ** | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
* | La | Ce | Pr | Nd | Pm | Sm | Eu | Gd | Tb | Dy | Ho | Er | Tm | Yb | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
** | Ac | Th | Pa | U | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | Md | No |
Legend: Quantity elements Essentiality or function in mammals debated No evidence for biological action in mammals, but essential in some lower organisms. (In the case of the lanthanides, the definition of an essential nutrient as being indispensable and irreplaceable is not completely applicable due to their extreme similarity. The stable early lanthanides La–Nd are known to stimulate the growth of various lanthanide-using organisms, and Sm–Gd show lesser effects for some such organisms. The later elements in the lanthanide series do not appear to have such effects.) [23] |
Atoms are the basic particles of the chemical elements. An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. The chemical elements are distinguished from each other by the number of protons that are in their atoms. For example, any atom that contains 11 protons is sodium, and any atom that contains 29 protons is copper. Atoms with the same number of protons but a different number of neutrons are called isotopes of the same element.
A chemical element is a chemical substance that cannot be broken down into other substances by chemical reactions. The basic particle that constitutes a chemical element is the atom. Chemical elements are identified by the number of protons in the nuclei of their atoms, known as the element's atomic number. For example, oxygen has an atomic number of 8, meaning that each oxygen atom has 8 protons in its nucleus. Two or more atoms of the same element can combine to form molecules, in contrast to chemical compounds or mixtures, which contain atoms of different elements. Atoms can be transformed into different elements in nuclear reactions, which change an atom's atomic number.
Neon is a chemical element; it has symbol Ne and atomic number 10. It is the second noble gas in the periodic table. Neon is a colorless, odorless, inert monatomic gas under standard conditions, with approximately two-thirds the density of air.
In physical cosmology, Big Bang nucleosynthesis is the production of nuclei other than those of the lightest isotope of hydrogen during the early phases of the universe. This type of nucleosynthesis is thought by most cosmologists to have occurred from 10 seconds to 20 minutes after the Big Bang. It is thought to be responsible for the formation of most of the universe's helium, along with small fractions of the hydrogen isotope deuterium, the helium isotope helium-3 (3He), and a very small fraction of the lithium isotope lithium-7 (7Li). In addition to these stable nuclei, two unstable or radioactive isotopes were produced: the heavy hydrogen isotope tritium and the beryllium isotope beryllium-7 (7Be). These unstable isotopes later decayed into 3He and 7Li, respectively, as above.
Nucleosynthesis is the process that creates new atomic nuclei from pre-existing nucleons and nuclei. According to current theories, the first nuclei were formed a few minutes after the Big Bang, through nuclear reactions in a process called Big Bang nucleosynthesis. After about 20 minutes, the universe had expanded and cooled to a point at which these high-energy collisions among nucleons ended, so only the fastest and simplest reactions occurred, leaving our universe containing hydrogen and helium. The rest is traces of other elements such as lithium and the hydrogen isotope deuterium. Nucleosynthesis in stars and their explosions later produced the variety of elements and isotopes that we have today, in a process called cosmic chemical evolution. The amounts of total mass in elements heavier than hydrogen and helium remains small, so that the universe still has approximately the same composition.
A period on the periodic table is a row of chemical elements. All elements in a row have the same number of electron shells. Each next element in a period has one more proton and is less metallic than its predecessor. Arranged this way, elements in the same group (column) have similar chemical and physical properties, reflecting the periodic law. For example, the halogens lie in the second-to-last group and share similar properties, such as high reactivity and the tendency to gain one electron to arrive at a noble-gas electronic configuration. As of 2022, a total of 118 elements have been discovered and confirmed.
The triple-alpha process is a set of nuclear fusion reactions by which three helium-4 nuclei are transformed into carbon.
Stellar nucleosynthesis is the creation (nucleosynthesis) of chemical elements by nuclear fusion reactions within stars. Stellar nucleosynthesis has occurred since the original creation of hydrogen, helium and lithium during the Big Bang. As a predictive theory, it yields accurate estimates of the observed abundances of the elements. It explains why the observed abundances of elements change over time and why some elements and their isotopes are much more abundant than others. The theory was initially proposed by Fred Hoyle in 1946, who later refined it in 1954. Further advances were made, especially to nucleosynthesis by neutron capture of the elements heavier than iron, by Margaret and Geoffrey Burbidge, William Alfred Fowler and Fred Hoyle in their famous 1957 B2FH paper, which became one of the most heavily cited papers in astrophysics history.
In astrophysics, silicon burning is a very brief sequence of nuclear fusion reactions that occur in massive stars with a minimum of about 8–11 solar masses. Silicon burning is the final stage of fusion for massive stars that have run out of the fuels that power them for their long lives in the main sequence on the Hertzsprung–Russell diagram. It follows the previous stages of hydrogen, helium, carbon, neon and oxygen burning processes.
Helium-4 is a stable isotope of the element helium. It is by far the more abundant of the two naturally occurring isotopes of helium, making up about 99.99986% of the helium on Earth. Its nucleus is identical to an alpha particle, and consists of two protons and two neutrons.
The Oddo–Harkins rule holds that an element with an even atomic number is more abundant than the elements with immediately adjacent atomic numbers. For example, carbon, with atomic number 6, is more abundant than boron (5) and nitrogen (7). Generally the relative abundance of an even atomic numbered element is roughly two orders of magnitude greater than the relative abundances of the immediately adjacent odd atomic numbered elements to either side. The Oddo-Harkins rule is true for all elements beginning with carbon produced by stellar nucleosynthesis but not true for the lightest elements below carbon produced by big bang nucleosynthesis and cosmic ray spallation. This pattern was first reported by Giuseppe Oddo in 1914 and William Draper Harkins in 1917.
Supernova nucleosynthesis is the nucleosynthesis of chemical elements in supernova explosions.
Cosmic ray spallation, also known as the x-process, is a set of naturally occurring nuclear reactions causing nucleosynthesis; it refers to the formation of chemical elements from the impact of cosmic rays on an object. Cosmic rays are highly energetic charged particles from beyond Earth, ranging from protons, alpha particles, and nuclei of many heavier elements. About 1% of cosmic rays also consist of free electrons.
Nuclear binding energy in experimental physics is the minimum energy that is required to disassemble the nucleus of an atom into its constituent protons and neutrons, known collectively as nucleons. The binding energy for stable nuclei is always a positive number, as the nucleus must gain energy for the nucleons to move apart from each other. Nucleons are attracted to each other by the strong nuclear force. In theoretical nuclear physics, the nuclear binding energy is considered a negative number. In this context it represents the energy of the nucleus relative to the energy of the constituent nucleons when they are infinitely far apart. Both the experimental and theoretical views are equivalent, with slightly different emphasis on what the binding energy means.
Nuclear astrophysics is an interdisciplinary part of both nuclear physics and astrophysics, involving close collaboration among researchers in various subfields of each of these fields. This includes, notably, nuclear reactions and their rates as they occur in cosmic environments, and modeling of astrophysical objects where these nuclear reactions may occur, but also considerations of cosmic evolution of isotopic and elemental composition (often called chemical evolution). Constraints from observations involve multiple messengers, all across the electromagnetic spectrum (nuclear gamma-rays, X-rays, optical, and radio/sub-mm astronomy), as well as isotopic measurements of solar-system materials such as meteorites and their stardust inclusions, cosmic rays, material deposits on Earth and Moon). Nuclear physics experiments address stability (i.e., lifetimes and masses) for atomic nuclei well beyond the regime of stable nuclides into the realm of radioactive/unstable nuclei, almost to the limits of bound nuclei (the drip lines), and under high density (up to neutron star matter) and high temperature (plasma temperatures up to 109 K). Theories and simulations are essential parts herein, as cosmic nuclear reaction environments cannot be realized, but at best partially approximated by experiments. In general terms, nuclear astrophysics aims to understand the origin of the chemical elements and isotopes, and the role of nuclear energy generation, in cosmic sources such as stars, supernovae, novae, and violent binary-star interactions.
The iron peak is a local maximum in the vicinity of Fe on the graph of the abundances of the chemical elements.
Isotopes are distinct nuclear species of the same chemical element. They have the same atomic number and position in the periodic table, but differ in nucleon numbers due to different numbers of neutrons in their nuclei. While all isotopes of a given element have almost the same chemical properties, they have different atomic masses and physical properties.
The atomic mass (ma or m) is the mass of an atom. Although the SI unit of mass is the kilogram (symbol: kg), atomic mass is often expressed in the non-SI unit dalton (symbol: Da) – equivalently, unified atomic mass unit (u). 1 Da is defined as 1⁄12 of the mass of a free carbon-12 atom at rest in its ground state. The protons and neutrons of the nucleus account for nearly all of the total mass of atoms, with the electrons and nuclear binding energy making minor contributions. Thus, the numeric value of the atomic mass when expressed in daltons has nearly the same value as the mass number. Conversion between mass in kilograms and mass in daltons can be done using the atomic mass constant .
In chemistry and physics, the iron group refers to elements that are in some way related to iron; mostly in period (row) 4 of the periodic table. The term has different meanings in different contexts.
Nuclear transmutation is the conversion of one chemical element or an isotope into another chemical element. Nuclear transmutation occurs in any process where the number of protons or neutrons in the nucleus of an atom is changed.
Correlations between abundance and nuclear binding energy [Subsection title]