Isotopes of hydrogen

Hydrogen (1H) has three naturally occurring isotopes, sometimes denoted 1H, 2H, and 3H. The first two of these are stable, while 3H has a half-life of 12.32 years. There are also heavier isotopes, which are all synthetic and have a half-life less than one zeptosecond (10−21 second). Of these, 5H is the most stable, and 7H is the least.[2][3]

Hydrogen is the only element whose individual isotopes have different names in common use today: the 2H (or hydrogen-2) isotope is deuterium[4] and the 3H (or hydrogen-3) isotope is tritium.[5] The symbols D and T are sometimes used for deuterium and tritium. The IUPAC accepts the D and T symbols, but recommends instead using standard isotopic symbols (2H and 3H) to avoid confusion in the alphabetic sorting of chemical formulas.[6] The ordinary isotope of hydrogen, with no neutrons, is sometimes called "protium".[7] (During the early study of radioactivity, some other heavy radioactive isotopes were given names, but such names are rarely used today.)

Main isotopes of hydrogen (1H)
Iso­tope Decay
abun­dance half-life (t1/2) mode pro­duct
1H 99.98% stable
2H 0.02% stable
3H trace 12.32 y β 3He
Standard atomic weight Ar, standard(H)
  • [1.007841.00811][1]
  • Conventional: 1.008
Hydrogen Deuterium Tritium Nuclei Schmatic-en
The three most stable isotopes of hydrogen: protium (A = 1), deuterium (A = 2), and tritium (A = 3).

Hydrogen-1 (protium)

Hydrogen
Protium, the most common isotope of hydrogen, consists of one proton and one electron. Unique among all stable isotopes, it has no neutrons. (see diproton for a discussion of why others do not exist)

1H (atomic mass 1.007825032241(94) u) is the most common hydrogen isotope with an abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the formal name protium.

The proton has never been observed to decay, and hydrogen-1 is therefore considered a stable isotope. Some grand unified theories proposed in the 1970s predict that proton decay can occur with a half-life between 1031 and 1036 years. If this prediction is found to be true, then hydrogen-1 (and indeed all nuclei now believed to be stable) are only observationally stable. To date, however, experiments have shown that the minimum proton half-life is in excess of 1034 years.

Hydrogen-2 (deuterium)

H-2 atom
A deuterium atom contains one proton, one neutron, and one electron

2H (atomic mass 2.01410177811(12) u), the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. The nucleus of deuterium is called a deuteron. Deuterium comprises 0.0026 – 0.0184% (by population, not by mass) of hydrogen samples on Earth, with the lower number tending to be found in samples of hydrogen gas and the higher enrichment (0.015% or 150 ppm) typical of ocean water. Deuterium on Earth has been enriched with respect to its initial concentration in the Big Bang and the outer solar system (about 27 ppm, by atom fraction) and its concentration in older parts of the Milky Way galaxy (about 23 ppm). Presumably the differential concentration of deuterium in the inner solar system is due to the lower volatility of deuterium gas and compounds, enriching deuterium fractions in comets and planets exposed to significant heat from the Sun over billions of years of solar system evolution.

Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of protium is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1H-NMR spectroscopy. Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.

Hydrogen-3 (tritium)

H-3 atom
A tritium atom contains one proton, two neutrons, and one electron

3H (atomic mass 3.01604928199(23) u) is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into helium-3 through β− decay with a half-life of 12.32 years.[8] Trace amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases. Tritium has also been released during nuclear weapons tests. It is used in thermonuclear fusion weapons, as a tracer in isotope geochemistry, and specialized in self-powered lighting devices.

The most common method of producing tritium is by bombarding a natural isotope of lithium, lithium-6, with neutrons in a nuclear reactor.

Tritium was once used routinely in chemical and biological labeling experiments as a radiolabel, which has become less common in recent times. D-T nuclear fusion uses tritium as its main reactant, along with deuterium, liberating energy through the loss of mass when the two nuclei collide and fuse at high temperatures.

Hydrogen-4

4H (atomic mass is 4.02643(11) u) contains one proton and three neutrons in its nucleus. It is a highly unstable isotope of hydrogen. It has been synthesised in the laboratory by bombarding tritium with fast-moving deuterium nuclei.[9] In this experiment, the tritium nucleus captured a neutron from the fast-moving deuterium nucleus. The presence of the hydrogen-4 was deduced by detecting the emitted protons. It decays through neutron emission into hydrogen-3 (tritium) with a half-life of about 139 ± 10 yoctoseconds, or (1.39 ± 0.10 × 10−22 seconds).[10]

In the 1955 satirical novel The Mouse That Roared, the name quadium was given to the hydrogen-4 isotope that powered the Q-bomb that the Duchy of Grand Fenwick captured from the United States.

Hydrogen-5

5H is a highly unstable isotope of hydrogen. The nucleus consists of a proton and four neutrons. It has been synthesised in the laboratory by bombarding tritium with fast-moving tritium nuclei.[9][11] In this experiment, one tritium nucleus captures two neutrons from the other, becoming a nucleus with one proton and four neutrons. The remaining proton may be detected, and the existence of hydrogen-5 deduced. It decays through double neutron emission into hydrogen-3 (tritium) and has a half-life of at least 910 yoctoseconds (9.1 × 10−22 seconds).[10]

Hydrogen-6

6H decays either through triple neutron emission into hydrogen-3 (tritium) or quadruple neutron emission into hydrogen-2 (deuterium) and has a half-life of 290 yoctoseconds (2.9 × 10−22 seconds).[10]

Hydrogen-7

7H consists of a proton and six neutrons. It was first synthesised in 2003 by a group of Russian, Japanese and French scientists at RIKEN's Radioactive Isotope Beam Factory by bombarding hydrogen with helium-8 atoms. In the resulting reaction, all six of the helium-8's neutrons were donated to the hydrogen's nucleus. The two remaining protons were detected by the "RIKEN telescope", a device composed of several layers of sensors, positioned behind the target of the RI Beam cyclotron.[3] Hydrogen-7 has a half life of 23 yoctoseconds (2.3×10−23 seconds).

List of isotopes

nuclide
symbol
Z(p) N(n) isotopic mass (u)[12] half-life
[resonance width]
decay
mode(s)[13]
daughter isotope(s)[n 1] nuclear
spin and
parity
representative
isotopic
composition
(mole fraction)[n 2]
range of natural
variation
(mole fraction)
1H 1 0 1.00782503224(9) Stable[n 3][n 4] 12+ 0.999885(70) 0.9998160.999974
2H (D)[n 5] 1 1 2.01410177811(12) Stable 1+ 0.000115(70)[n 6] 0.0000260.000184
3H (T)[n 7] 1 2 3.01604928199(23) 12.32(2) y β 3
He
12+ Trace[n 8]
4
H
1 3 4.02643(11) 1.39(10)×10−22 s
[3.28(23) MeV]
n 3
H
2
5
H
1 4 5.03531(10) > 9.1×10−22 s
[< 0.5 MeV]
2n 3
H
(​12+)
6
H
1 5 6.04496(27) 2.90(70)×10−22 s
[1.6(4) MeV]
3n 3
H
2#
4n 2
H
7
H
1 6 7.05275(108)# 5×10−22 s# 4n 3
H
12+#
  1. ^ Bold for stable isotopes.
  2. ^ Refers to that in water.
  3. ^ Unless proton decay occurs.
  4. ^ This and 3He are the only stable nuclides with more protons than neutrons.
  5. ^ Produced during Big Bang nucleosynthesis.
  6. ^ Tank hydrogen has a 2
    H
    abundance as low as 3.2×10−5 (mole fraction).
  7. ^ Produced during Big Bang nucleosynthesis, but not primordial, as all such atoms have since decayed to 3He.
  8. ^ Cosmogenic

Notes

  • Commercially available materials may have been subjected to an undisclosed or inadvertent isotopic fractionation. Substantial deviations from the given mass and composition can occur.
  • Values marked # are not purely derived from experimental data, but at least partly from systematic trends. Spins with weak assignment arguments are enclosed in parentheses.
  • Uncertainties are given in concise form in parentheses after the corresponding last digits. Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC, which use expanded uncertainties.
  • Isotope abundances are given by IUPAC Commission on Isotopic Abundances and Atomic Weights (CIAAW)

Decay chains

The majority of heavy hydrogen isotopes decay directly to 3H, which then decays to the stable isotope 3He. However, 6H has occasionally been observed to decay directly to stable 2H.

Note that the decay times are in yoctoseconds for all isotopes except 3H, which is expressed in years.

See also

References

Notes
  1. ^ Meija, Juris; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
  2. ^ Y. B. Gurov; et al. (2004). "Spectroscopy of superheavy hydrogen isotopes in stopped-pion absorption by nuclei". Physics of Atomic Nuclei. 68 (3): 491–497. Bibcode:2005PAN....68..491G. doi:10.1134/1.1891200.
  3. ^ a b A. A. Korsheninnikov; et al. (2003). "Experimental Evidence for the Existence of 7H and for a Specific Structure of 8He". Physical Review Letters. 90 (8): 082501. Bibcode:2003PhRvL..90h2501K. doi:10.1103/PhysRevLett.90.082501.
  4. ^ {{GoldBookRef|file=D01648|title=deuterium}
  5. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "tritium". doi:10.1351/goldbook.T06513
  6. ^ International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSCIUPAC. ISBN 0-85404-438-8. p. 48. Electronic version.
  7. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "protium". doi:10.1351/goldbook.P04903
  8. ^ G. L. Miessler; D. A. Tarr (2004). Inorganic Chemistry (3rd ed.). Pearson Prentice Hall. ISBN 978-0-13-035471-6.
  9. ^ a b G. M. Ter-Akopian; et al. (2002). "Hydrogen-4 and Hydrogen-5 from t+t and t+d transfer reactions studied with a 57.5-MeV triton beam". AIP Conference Proceedings. 610: 920–924. doi:10.1063/1.1470062.
  10. ^ a b c Audi, Georges; Wapstra, Aaldert Hendrik; Thibault, Catherne; Blachot, Jean; Bersillon, Olivier (2003). "The NUBASE evaluation of nuclear and decay properties" (PDF). Nuclear Physics A. 729 (1): 3–128. Bibcode:2003NuPhA.729....3A. CiteSeerX 10.1.1.692.8504. doi:10.1016/j.nuclphysa.2003.11.001. Archived from the original (PDF) on 2011-07-20.
  11. ^ A. A. Korsheninnikov; et al. (2001). "Superheavy Hydrogen 5H". Physical Review Letters. 87 (9): 92501. Bibcode:2001PhRvL..87i2501K. doi:10.1103/PhysRevLett.87.092501.
  12. ^ Wang, Meng; Audi, Georges; Kondev, Filip G.; Huang, Wen Jian; Naimi, Sarah; Xu, Xing (2017), "The AME2016 atomic mass evaluation (II). Tables, graphs, and references" (PDF), Chinese Physics C, 41 (3): 030003–1—030003–442, doi:10.1088/1674-1137/41/3/030003
  13. ^ Audi, Georges; Kondev, Filip G.; Wang, Meng; Huang, Wen Jia; Naimi, Sarah (2017), "The NUBASE2016 evaluation of nuclear properties" (PDF), Chinese Physics C, 41 (3): 030001–1—030001–138, Bibcode:2017ChPhC..41c0001A, doi:10.1088/1674-1137/41/3/030001
General references

Further reading

External links

1933 in science

The year 1933 in science and technology involved some significant events, listed below.

Cosmological lithium problem

In astronomy, the lithium problem, lithium discrepancy or lithium gap refers to the discrepancy between the observed abundance of lithium produced in Big Bang nucleosynthesis and the amount that should theoretically exist. Namely, the most widely-accepted models of the Big Bang suggest that three times as much primordial lithium, in particular lithium-7, should exist. This contrasts with the observed abundance of isotopes of hydrogen (1H and 2H) and helium (3He and 4He) that are consistent with predictions.

Deuterium

Deuterium (or hydrogen-2, symbol D or 2H, also known as heavy hydrogen) is one of two stable isotopes of hydrogen (the other being protium, or hydrogen-1). The nucleus of deuterium, called a deuteron, contains one proton and one neutron, whereas the far more common protium has no neutron in the nucleus. Deuterium has a natural abundance in Earth's oceans of about one atom in 6420 of hydrogen. Thus deuterium accounts for approximately 0.02% (or, on a mass basis, 0.03%) of all the naturally occurring hydrogen in the oceans, while protium accounts for more than 99.98%. The abundance of deuterium changes slightly from one kind of natural water to another (see Vienna Standard Mean Ocean Water).

The deuterium isotope's name is formed from the Greek deuteros, meaning "second", to denote the two particles composing the nucleus. Deuterium was discovered and named in 1931 by Harold Urey. When the neutron was discovered in 1932, this made the nuclear structure of deuterium obvious, and Urey won the Nobel Prize in 1934 “for his discovery of heavy hydrogen”. Soon after deuterium's discovery, Urey and others produced samples of "heavy water" in which the deuterium content had been highly concentrated.

Deuterium is destroyed in the interiors of stars faster than it is produced. Other natural processes are thought to produce only an insignificant amount of deuterium. Nearly all deuterium found in nature was produced in the Big Bang 13.8 billion years ago, as the basic or primordial ratio of hydrogen-1 to deuterium (about 26 atoms of deuterium per million hydrogen atoms) has its origin from that time. This is the ratio found in the gas giant planets, such as Jupiter. However, other astronomical bodies are found to have different ratios of deuterium to hydrogen-1. This is thought to be a result of natural isotope separation processes that occur from solar heating of ices in comets. Like the water cycle in Earth's weather, such heating processes may enrich deuterium with respect to protium. The analysis of deuterium/protium ratios in comets found results very similar to the mean ratio in Earth's oceans (156 atoms of deuterium per million hydrogens). This reinforces theories that much of Earth's ocean water is of cometary origin. The deuterium/protium ratio of the comet 67P/Churyumov-Gerasimenko, as measured by the Rosetta space probe, is about three times that of earth water. This figure is the highest yet measured in a comet.Deuterium/protium ratios thus continue to be an active topic of research in both astronomy and climatology.

François Robert

François Robert, born in Paris, France the 26th of January, 1951, is a French researcher specializing in isotope geochemistry and cosmochemistry. His work on the isotopes of hydrogen has enhanced the understanding of the origin of water and of organic matter in the solar system. He is famous for his work on lithium, beryllium and boron, light elements formed by the irradiation of interstellar matter. He received a Leonard Medal from the Meteoritical Society in 2011 for his work on the isotopic composition of stable nuclei.

Hydrogen (disambiguation)

Hydrogen is a chemical element with symbol H and atomic number 1.

Hydrogen may also refer to:

Hydrogen atom, about the physics of atomic hydrogen

Hydrogen ion

Hydrogen (software), drum machine software

Hydrogen vehicle

Isotopes of hydrogen

Hydrogen-2 (deuterium)

Hydrogen-3 (tritium)

Hydrogen-4

Hydrogen-5

Hydrogen (horse), a champion Australian Thoroughbred racehorse

Hydrogen atom

A hydrogen atom is an atom of the chemical element hydrogen. The electrically neutral atom contains a single positively charged proton and a single negatively charged electron bound to the nucleus by the Coulomb force. Atomic hydrogen constitutes about 75% of the baryonic mass of the universe.In everyday life on Earth, isolated hydrogen atoms (called "atomic hydrogen") are extremely rare. Instead, a hydrogen atom tends to combine with other atoms in compounds, or with another hydrogen atom to form ordinary (diatomic) hydrogen gas, H2. "Atomic hydrogen" and "hydrogen atom" in ordinary English use have overlapping, yet distinct, meanings. For example, a water molecule contains two hydrogen atoms, but does not contain atomic hydrogen (which would refer to isolated hydrogen atoms).

Atomic spectroscopy shows that there is a discrete infinite set of states in which a hydrogen (or any) atom can exist, contrary to the predictions of classical physics. Attempts to develop a theoretical understanding of the states of the hydrogen atom have been important to the history of quantum mechanics, since all other atoms can be roughly understood by knowing in detail about this simplest atomic structure.

Hydrogen deuteride

Hydrogen deuteride is a diatomic molecule substance or compound of the two isotopes of hydrogen: the majority isotope 1H protium and 2H deuterium. Its proper molecular formula is H2H but for simplification it is usually written as HD.

Hydron (chemistry)

In chemistry, a hydron is the general name for a cationic form of atomic hydrogen, represented with the symbol H+. However, this term is avoided and instead "proton" is used, which strictly speaking refers to the cation of protium, the most common isotope of hydrogen. The term "hydron" includes cations of hydrogen regardless of their isotopic composition: thus it refers collectively to protons (1H+) for the protium isotope, deuterons (2H+ or D+) for the deuterium isotope, and tritons (3H+ or T+) for the tritium isotope. Unlike most other ions, the hydron consists only of a bare atomic nucleus.

The negatively charged counterpart of the hydron is the hydride anion, H−.

Isotope

Isotopes are variants of a particular chemical element which differ in neutron number, and consequently in nucleon number. All isotopes of a given element have the same number of protons but different numbers of neutrons in each atom.The term isotope is formed from the Greek roots isos (ἴσος "equal") and topos (τόπος "place"), meaning "the same place"; thus, the meaning behind the name is that different isotopes of a single element occupy the same position on the periodic table. It was coined by a Scottish doctor and writer Margaret Todd in 1913 in a suggestion to chemist Frederick Soddy.

The number of protons within the atom's nucleus is called atomic number and is equal to the number of electrons in the neutral (non-ionized) atom. Each atomic number identifies a specific element, but not the isotope; an atom of a given element may have a wide range in its number of neutrons. The number of nucleons (both protons and neutrons) in the nucleus is the atom's mass number, and each isotope of a given element has a different mass number.

For example, carbon-12, carbon-13, and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13, and 14, respectively. The atomic number of carbon is 6, which means that every carbon atom has 6 protons, so that the neutron numbers of these isotopes are 6, 7, and 8 respectively.

Isotope analysis

Isotope analysis is the identification of isotopic signature, the abundance of certain stable isotopes and chemical elements within organic and inorganic compounds. Isotopic analysis can be used to understand the flow of energy through a food web, to reconstruct past environmental and climatic conditions, to investigate human and animal diets in the past, for food authentification, and a variety of other physical, geological, palaeontological and chemical processes. Stable isotope ratios are measured using mass spectrometry, which separates the different isotopes of an element on the basis of their mass-to-charge ratio.

Isotopologue

Isotopologues are molecules that differ only in their isotopic composition. They have the same chemical formula and bonding arrangement of atoms, but at least one atom has a different number of neutrons than the parent.

An example is water, where its hydrogen-related isotopologues are: "light water" (HOH or H2O), "semi-heavy water" with the deuterium isotope in equal proportion to protium (HDO or 1H2HO), "heavy water" with two deuterium isotopes of hydrogen per molecule (D2O or 2H2O), and "super-heavy water" or tritiated water (T2O or 3H2O, as well as HTO [1H3HO] and DTO [2H3HO], where some or all of the hydrogen atoms are replaced with tritium isotopes). Oxygen-related isotopologues of water include the commonly available form of heavy-oxygen water (H218O) and the more difficult to separate version with the 17O isotope. Both elements may be replaced by isotopes, for example in the doubly labeled water isotopologue D218O.

The atom(s) of the different isotope may be anywhere in a molecule, so the difference is in the net chemical formula. If a compound has several atoms of the same element, any one of them could be the altered one, and it would still be the same isotopologue. When considering the different locations of the same isotopically modified element, the term isotopomer, first proposed by Seeman and Paine in 1992, is used.

Isotopomerism is analogous to constitutional isomerism of different elements in a structure. Depending on the formula and the symmetry of the structure, there might be several isotopomers of one isotopologue. For example, ethanol has the molecular formula C2H6O. Mono-deuterated ethanol, C2H5DO, is an isotopologue of it. The structural formulas CH3−CH2−O−D and CH2D−CH2−O−H are two isotopomers of that isotopologue.

Neutron generator

Neutron generators are neutron source devices which contain compact linear particle accelerators and that produce neutrons by fusing isotopes of hydrogen together. The fusion reactions take place in these devices by accelerating either deuterium, tritium, or a mixture of these two isotopes into a metal hydride target which also contains deuterium, tritium or a mixture of these isotopes. Fusion of deuterium atoms (D + D) results in the formation of a He-3 ion and a neutron with a kinetic energy of approximately 2.5 MeV. Fusion of a deuterium and a tritium atom (D + T) results in the formation of a He-4 ion and a neutron with a kinetic energy of approximately 14.1 MeV. Neutron generators have applications in medicine, security, and materials analysis.The basic concept was first developed by Ernest Rutherford's team in the Cavendish Laboratory in the early 1930s. Using a linear accelerator driven by a Cockcroft–Walton generator, Mark Oliphant led an experiment that fired deuterium ions into a deuterium-infused metal foil and noticed that a small number of these particles gave off alpha particles. This was the first demonstration of nuclear fusion, as well as the first discovery of Helium-3 and tritium, created in these reactions. The introduction of new power sources has continually shrunk the size of these machines, from Oliphant's that filled the corner of the lab, to modern machines that are highly portable. Thousands of such small, relatively inexpensive systems have been built over the past five decades.

While neutron generators do produce fusion reactions, the number of accelerated ions that cause these reactions is very low. It can be easily demonstrated that the energy released by these reactions is many times lower than the energy needed to accelerate the ions, so there is no possibility of these machines being used to produce net fusion power. A related concept, colliding beam fusion, attempts to address this issue using two accelerators firing at each other.

Neutron source

A neutron source is any device that emits neutrons, irrespective of the mechanism used to produce the neutrons. Neutron sources are used in physics, engineering, medicine, nuclear weapons, petroleum exploration, biology, chemistry, and nuclear power.

Neutron source variables include the energy of the neutrons emitted by the source, the rate of neutrons emitted by the source, the size of the source, the cost of owning and maintaining the source, and government regulations related to the source.

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.

Pure fusion weapon

A pure fusion weapon is a hypothetical hydrogen bomb design that does not need a fission "primary" explosive to ignite the fusion of deuterium and tritium, two heavy isotopes of hydrogen (see thermonuclear weapon for more information about fission-fusion weapons). Such a weapon would require no fissile material and would therefore be much easier to develop in secret than existing weapons. The necessity of separating weapons grade uranium (U-235) or breeding plutonium (Pu-239) requires a substantial and difficult-to-conceal industrial investment, and blocking the sale and transfer of the needed machinery has been the primary mechanism to control nuclear proliferation to date.

Radioactive tracer

A radioactive tracer, radiotracer, or radioactive label, is a chemical compound in which one or more atoms have been replaced by a radionuclide so by virtue of its radioactive decay it can be used to explore the mechanism of chemical reactions by tracing the path that the radioisotope follows from reactants to products. Radiolabeling or radiotracing is thus the radioactive form of isotopic labeling.

Radioisotopes of hydrogen, carbon, phosphorus, sulfur, and iodine have been used extensively to trace the path of biochemical reactions. A radioactive tracer can also be used to track the distribution of a substance within a natural system such as a cell or tissue, or as a flow tracer to track fluid flow. Radioactive tracers are also used to determine the location of fractures created by hydraulic fracturing in natural gas production. Radioactive tracers form the basis of a variety of imaging systems, such as, PET scans, SPECT scans and technetium scans. Radiocarbon dating uses the naturally occurring carbon-14 isotope as an isotopic label.

Tritium

Tritium ( or ) or hydrogen-3 is a rare and radioactive isotope of hydrogen, with symbol T or 3H. The nucleus of tritium (sometimes called a triton) contains one proton and two neutrons, whereas the nucleus of the common isotope hydrogen-1 ("protium") contains just one proton, and that of hydrogen-2 ("deuterium") contains one proton and one neutron.

Naturally occurring tritium is extremely rare on Earth. The atmosphere has only trace amounts, formed by the interaction of its gases with cosmic rays. It can be produced by irradiating lithium metal or lithium-bearing ceramic pebbles in a nuclear reactor.

Tritium is used as a radioactive tracer, in radioluminescent light sources for watches and instruments, and, along with deuterium, as a fuel for nuclear fusion reactions with applications in energy generation and weapons.

The name of this isotope is derived from Greek τρίτος (trítos), meaning "third".

Vibrational bond

A vibrational bond is a chemical bond that happens between two very large atoms, like bromine, and a very small atom, like hydrogen, at very high energy states. Vibrational bonds only exist for a few milliseconds. This bond is detectable through modern analytic chemistry and is significant because it affects the rate at which other reactions can occur.

Wendell Mitchell Latimer

Wendell Mitchell Latimer (April 22, 1893 – July 6, 1955) was a prominent chemist notable for his description of oxidation states in his book "The Oxidation States of the Elements and Their Potentials in Aqueous Solution" (ASIN B000GRXLSA, first published 1938).

He received his Ph.D from the University of California, Berkeley for the work with George Ernest Gibson.

He earned many awards and honors for his scientific work.

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