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Hydrogen, 1H
Hydrogen discharge tube.jpg
Purple glow in its plasma state
Appearancecolorless gas
Standard atomic weight Ar, std(H)[1.00784, 1.00811]conventional: 1.008
Hydrogen in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon


– ← hydrogenhelium
Atomic number (Z)1
Group 1: H and alkali metals
Period period 1
Block s-block
Element category   Reactive nonmetal
Electron configuration 1s1
Electrons per shell1
Physical properties
Phase at  STP gas
Melting point (H2) 13.99  K (−259.16 °C,−434.49 °F)
Boiling point (H2) 20.271 K(−252.879 °C,−423.182 °F)
Density (at STP)0.08988 g/L
when liquid (at m.p.)0.07 g/cm3(solid: 0.0763 g/cm3) [1]
when liquid (at b.p.)0.07099 g/cm3
Triple point 13.8033 K,7.041 kPa
Critical point 32.938 K, 1.2858 MPa
Heat of fusion (H2) 0.117  kJ/mol
Heat of vaporization (H2) 0.904 kJ/mol
Molar heat capacity (H2) 28.836 J/(mol·K)
Vapor pressure
P (Pa)1101001 k10 k100 k
at T (K)1520
Atomic properties
Oxidation states −1, +1 (an  amphoteric oxide)
Electronegativity Pauling scale: 2.20
Ionization energies
  • 1st: 1312.0 kJ/mol
Covalent radius 31±5  pm
Van der Waals radius 120 pm
Color lines in a spectral range Hydrogen spectrum visible.png
Color lines in a spectral range
Spectral lines of hydrogen
Other properties
Natural occurrence primordial
Crystal structure hexagonal
Speed of sound 1310 m/s (gas, 27 °C)
Thermal conductivity 0.1805 W/(m·K)
Magnetic ordering diamagnetic [2]
Magnetic susceptibility 3.98·10−6 cm3/mol(298 K) [3]
CAS Number 12385-13-6
1333-74-0 (H2)
Discovery Henry Cavendish [4] [5] (1766)
Named by Antoine Lavoisier [6] (1783)
Main isotopes of hydrogen
1H 99.98% stable
2H 0.02%stable
3H trace 12.32 y β 3He
Folder Hexagonal Icon.svg  Category: Hydrogen
| references

Hydrogen is the chemical element with the symbol H and atomic number  1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the universe, constituting roughly 75% of all baryonic mass. [7] [note 1] Non-remnant stars are mainly composed of hydrogen in the plasma state. The most common isotope of hydrogen, termed protium (name rarely used, symbol 1H), has one proton and no neutrons.


The universal emergence of atomic hydrogen first occurred during the recombination epoch (Big Bang). At standard temperature and pressure, hydrogen is a colorless, odorless, tasteless, non-toxic, nonmetallic, highly combustible diatomic gas with the molecular formula H2. Since hydrogen readily forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a particularly important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge (i.e., anion) when it is known as a hydride, or as a positively charged (i.e., cation) species denoted by the symbol H+. The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, [8] study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics.

Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, [9] and that it produces water when burned, the property for which it was later named: in Greek, hydrogen means "water-former".

Industrial production is mainly from steam reforming natural gas, and less often from more energy-intensive methods such as the electrolysis of water. [10] Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing (e.g., hydrocracking) and ammonia production, mostly for the fertilizer market. Hydrogen is problematic in metallurgy because it can embrittle many metals, [11] complicating the design of pipelines and storage tanks. [12]



Combustion of hydrogen with the oxygen in the air. When the bottom cap is removed, allowing air to enter at the bottom, the hydrogen in the container rises out of top and burns as it mixes with the air.
The Space Shuttle Main Engine burnt hydrogen with oxygen, producing a nearly invisible flame at full thrust. Shuttle Main Engine Test Firing cropped edited and reduced.jpg

Hydrogen gas (dihydrogen or molecular hydrogen, [13] is highly flammable:

2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ (286 kJ/mol) [note 2]

The enthalpy of combustion is −286 kJ/mol: [14]

Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% [15] and with chlorine at 5–95%. The explosive reactions may be triggered by spark, heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C (932 °F). [16]


Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the highly visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite. The detection of a burning hydrogen leak may require a flame detector; such leaks can be very dangerous. Hydrogen flames in other conditions are blue, resembling blue natural gas flames. [17] The destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.


H2 is relatively unreactive. The thermodynamic basis this low reactivity is the very strong H-H bond, with a bond dissociation energy of 435.7 kJ/mol. [18] The kinetic basis of the low reactivity is the nonpolar nature of H2 and its weak polarizability. It spontaneously reacts with chlorine and fluorine to form hydrogen chloride and hydrogen fluoride, respectively. [19] Molten sodium and potassium react with the gas to give the respective hydrides NaH and KH.[ citation needed ] The reactivity of H2 is strongly affected by the presence of metal catalysts. Thus, while H2 combusts readily, mixtures of H2 and O2 do not react in the absence of a catalyst.

Electron energy levels

Depiction of a hydrogen atom with size of central proton shown, and the atomic diameter shown as about twice the Bohr model radius (image not to scale) Hydrogen atom.svg

The ground state energy level of the electron in a hydrogen atom is −13.6  eV, [20] which is equivalent to an ultraviolet photon of roughly 91  nm wavelength. [21]

The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity. Because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, and therefore only certain allowed energies. [22]

A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or even the Feynman path integral formulation to calculate the probability density of the electron around the proton. [23] The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion.

Elemental molecular forms

First tracks observed in liquid hydrogen bubble chamber at the Bevatron Liquid hydrogen bubblechamber.jpg
First tracks observed in liquid hydrogen bubble chamber at the Bevatron

Molecular H2 exists a two spin isomers, i.e. compounds with two nuclear spin states. [24] In the orthohydrogen form, the spins of the two nuclei are parallel and form a triplet state with a molecular spin quantum number of 1 (12+12); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 (1212). At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form". [25] The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The liquid and gas phase thermal properties of pure parahydrogen differ significantly from those of the normal form because of differences in rotational heat capacities, as discussed more fully in spin isomers of hydrogen . [26] The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene, but is of little significance for their thermal properties. [27]

The ortho form that converts to the para form slowly at low temperatures. [28] The ortho/para ratio in condensed H2 is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate some of the hydrogen liquid, leading to loss of liquefied material. Catalysts for the ortho-para interconversion, such as ferric oxide, activated carbon, platinized asbestos, rare earth metals, uranium compounds, chromic oxide, or some nickel [29] compounds, are used during hydrogen cooling. [30]



Covalent and organic compounds

While H2 is not very reactive under standard conditions, it does form compounds with most elements. Hydrogen can form compounds with elements that are more electronegative, such as halogens (F, Cl, Br, I), or oxygen; in these compounds hydrogen takes on a partial positive charge. [31] When bonded to a more electronegative element, particularly fluorine, oxygen, or nitrogen, hydrogen can participate in a form of medium-strength noncovalent bonding with another electronegative element with a lone pair, a phenomenon called hydrogen bonding that is critical to the stability of many biological molecules. [32] [33] Hydrogen also forms compounds with less electronegative elements, such as metals and metalloids, where it takes on a partial negative charge. These compounds are often known as hydrides. [34]

Hydrogen forms a vast array of compounds with carbon called the hydrocarbons, and an even vaster array with heteroatoms that, because of their general association with living things, are called organic compounds. [35] The study of their properties is known as organic chemistry [36] and their study in the context of living organisms is known as biochemistry. [37] By some definitions, "organic" compounds are only required to contain carbon. However, most of them also contain hydrogen, and because it is the carbon-hydrogen bond that gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry. [35] Millions of hydrocarbons are known, and they are usually formed by complicated pathways that seldom involve elemental hydrogen.

Hydrogen is highly soluble in many rare earth and transition metals [38] and is soluble in both nanocrystalline and amorphous metals. [39] Hydrogen solubility in metals is influenced by local distortions or impurities in the crystal lattice. [40] These properties may be useful when hydrogen is purified by passage through hot palladium disks, but the gas's high solubility is a metallurgical problem, contributing to the embrittlement of many metals, [11] complicating the design of pipelines and storage tanks. [12]


Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. The term "hydride" suggests that the H atom has acquired a negative or anionic character, denoted H, and is used when hydrogen forms a compound with a more electropositive element. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group 1 and 2 salt-like hydrides, was demonstrated by Moers in 1920 by the electrolysis of molten lithium hydride (LiH), producing a stoichiometric quantity of hydrogen at the anode. [41] For hydrides other than group 1 and 2 metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group 2 hydrides is BeH
, which is polymeric. In lithium aluminium hydride, the AlH

Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, more than 100 binary borane hydrides are known, but only one binary aluminium hydride. [42] Binary indium hydride has not yet been identified, although larger complexes exist. [43]

In inorganic chemistry, hydrides can also serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminium complexes, as well as in clustered carboranes. [44]

Protons and acids

Oxidation of hydrogen removes its electron and gives H+, which contains no electrons and a nucleus which is usually composed of one proton. That is why H+
is often called a proton. This species is central to discussion of acids. Under the Brønsted–Lowry acid–base theory, acids are proton donors, while bases are proton acceptors.

A bare proton, H+
, cannot exist in solution or in ionic crystals because of its unstoppable attraction to other atoms or molecules with electrons. Except at the high temperatures associated with plasmas, such protons cannot be removed from the electron clouds of atoms and molecules, and will remain attached to them. However, the term 'proton' is sometimes used loosely and metaphorically to refer to positively charged or cationic hydrogen attached to other species in this fashion, and as such is denoted "H+
" without any implication that any single protons exist freely as a species.

To avoid the implication of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain a less unlikely fictitious species, termed the "hydronium ion" (H
). However, even in this case, such solvated hydrogen cations are more realistically conceived as being organized into clusters that form species closer to H
. [45] Other oxonium ions are found when water is in acidic solution with other solvents. [46]

Although exotic on Earth, one of the most common ions in the universe is the H+
ion, known as protonated molecular hydrogen or the trihydrogen cation. [47]

Atomic hydrogen

NASA has investigated the use of atomic hydrogen as a rocket propellant. It could be stored in liquid helium to prevent it from recombining into molecular hydrogen. When the helium is vaporized, the atomic hydrogen would be released and combine back to molecular hydrogen. The result would be an intensely hot stream of hydrogen and helium gas. The liftoff weight of rockets could be reduced by 50% by this method. [48]

Most interstellar hydrogen is in the form of atomic hydrogen because the atoms can seldom collide and combine. They are the source of the important 21 cm hydrogen line in astronomy at 1420 MHz. [49]


Hydrogen discharge (spectrum) tube Hydrogen discharge tube.jpg
Hydrogen discharge (spectrum) tube
Deuterium discharge (spectrum) tube Deuterium discharge tube.jpg
Deuterium discharge (spectrum) tube
Protium, the most common isotope of hydrogen, has 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). Protium.svg

Hydrogen has three naturally occurring isotopes, denoted 1
, 2
and 3
. Other, highly unstable nuclei (4
to 7
) have been synthesized in the laboratory but not observed in nature. [50] [51]

Unique among the elements, distinct names are assigned to its isotopes in common use today. During the early study of radioactivity, various heavy radioactive isotopes were given their own names, but such names are no longer used, except for deuterium and tritium. The symbols D and T (instead of 2
and 3
) are sometimes used for deuterium and tritium, but the corresponding symbol for protium, P, is already in use for phosphorus and thus is not available for protium. [61] In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry (IUPAC) allows any of D, T, 2
, and 3
to be used, although 2
and 3
are preferred. [62]

The exotic atom muonium (symbol Mu), composed of an antimuon and an electron, is also sometimes considered as a light radioisotope of hydrogen, due to the mass difference between the antimuon and the electron. [63] Muonium was discovered in 1960. [64] During the muon's 2.2  µs lifetime, muonium can enter into compounds such as muonium chloride (MuCl) or sodium muonide (NaMu), analogous to hydrogen chloride and sodium hydride respectively. [65]


Discovery and use

In 1671, Robert Boyle discovered and described the reaction between iron filings and dilute acids, which results in the production of hydrogen gas. [66] [67] In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by naming the gas from a metal-acid reaction "inflammable air". He speculated that "inflammable air" was in fact identical to the hypothetical substance called "phlogiston" [68] [69] and further finding in 1781 that the gas produces water when burned. He is usually given credit for the discovery of hydrogen as an element. [4] [5] In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek ὑδρο- hydro meaning "water" and -γενής genes meaning "creator") [70] when he and Laplace reproduced Cavendish's finding that water is produced when hydrogen is burned. [5]

Antoine-Laurent de Lavoisier Antoine-Laurent Lavoisier (by Louis Jean Desire Delaistre)RENEW.jpg
Antoine-Laurent de Lavoisier

Lavoisier produced hydrogen for his experiments on mass conservation by reacting a flux of steam with metallic iron through an incandescent iron tube heated in a fire. Anaerobic oxidation of iron by the protons of water at high temperature can be schematically represented by the set of following reactions:

   Fe +    H2O → FeO + H2
2 Fe + 3 H2O → Fe2O3 + 3 H2
3 Fe + 4 H2O → Fe3O4 + 4 H2

Many metals such as zirconium undergo a similar reaction with water leading to the production of hydrogen.

Hydrogen was liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask. [5] He produced solid hydrogen the next year. [5] Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck. [4] Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932. [5] François Isaac de Rivaz built the first de Rivaz engine, an internal combustion engine powered by a mixture of hydrogen and oxygen in 1806. Edward Daniel Clarke invented the hydrogen gas blowpipe in 1819. The Döbereiner's lamp and limelight were invented in 1823. [5]

The first hydrogen-filled balloon was invented by Jacques Charles in 1783. [5] Hydrogen provided the lift for the first reliable form of air-travel following the 1852 invention of the first hydrogen-lifted airship by Henri Giffard. [5] German count Ferdinand von Zeppelin promoted the idea of rigid airships lifted by hydrogen that later were called Zeppelins; the first of which had its maiden flight in 1900. [5] Regularly scheduled flights started in 1910 and by the outbreak of World War I in August 1914, they had carried 35,000 passengers without a serious incident. Hydrogen-lifted airships were used as observation platforms and bombers during the war.

The first non-stop transatlantic crossing was made by the British airship R34 in 1919. Regular passenger service resumed in the 1920s and the discovery of helium reserves in the United States promised increased safety, but the U.S. government refused to sell the gas for this purpose. Therefore, H2 was used in the Hindenburg airship, which was destroyed in a midair fire over New Jersey on 6 May 1937. [5] The incident was broadcast live on radio and filmed. Ignition of leaking hydrogen is widely assumed to be the cause, but later investigations pointed to the ignition of the aluminized fabric coating by static electricity. But the damage to hydrogen's reputation as a lifting gas was already done and commercial hydrogen airship travel ceased. Hydrogen is still used, in preference to non-flammable but more expensive helium, as a lifting gas for weather balloons.

In the same year, the first hydrogen-cooled turbogenerator went into service with gaseous hydrogen as a coolant in the rotor and the stator in 1937 at Dayton, Ohio, by the Dayton Power & Light Co.; [71] because of the thermal conductivity and very low viscosity of hydrogen gas, thus lower drag than air, this is the most common type in its field today for large generators (typically 60 MW and bigger; smaller generators are usually air-cooled).

The nickel hydrogen battery was used for the first time in 1977 aboard the U.S. Navy's Navigation technology satellite-2 (NTS-2). [72] For example, the ISS, [73] Mars Odyssey [74] and the Mars Global Surveyor [75] are equipped with nickel-hydrogen batteries. In the dark part of its orbit, the Hubble Space Telescope is also powered by nickel-hydrogen batteries, which were finally replaced in May 2009, [76] more than 19 years after launch and 13 years beyond their design life. [77]

Role in quantum theory

Hydrogen emission spectrum lines in the visible range. These are the four visible lines of the Balmer series Emission spectrum-H.svg

Because of its simple atomic structure, consisting only of a proton and an electron, the hydrogen atom, together with the spectrum of light produced from it or absorbed by it, has been central to the development of the theory of atomic structure. [78] Furthermore, study of the corresponding simplicity of the hydrogen molecule and the corresponding cation H+


Antihydrogen (
) is the antimatter counterpart to hydrogen. It consists of an antiproton with a positron. Antihydrogen is the only type of antimatter atom to have been produced as of 2015. [80] [81]

Cosmic prevalence and distribution

NGC 604, a giant region of ionized hydrogen in the Triangulum Galaxy Nursery of New Stars - GPN-2000-000972.jpg

Hydrogen, as atomic H, is the most abundant chemical element in the universe, making up 75% of normal matter by mass and more than 90% by number of atoms. (Most of the mass of the universe, however, is not in the form of chemical-element type matter, but rather is postulated to occur as yet-undetected forms of mass such as dark matter and dark energy. [82] ) This element is found in great abundance in stars and gas giant planets. Molecular clouds of H2 are associated with star formation. Hydrogen plays a vital role in powering stars through the proton-proton reaction in case of stars with very low to approximately 1 mass of the Sun and the CNO cycle of nuclear fusion in case of stars more massive than our Sun. [83]


Throughout the universe, hydrogen is mostly found in the atomic and plasma states, with properties quite distinct from those of molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity and high emissivity (producing the light from the Sun and other stars). The charged particles are highly influenced by magnetic and electric fields. For example, in the solar wind they interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora. Hydrogen is found in the neutral atomic state in the interstellar medium. The large amount of neutral hydrogen found in the damped Lyman-alpha systems is thought to dominate the cosmological baryonic density of the universe up to redshift z=4. [84]

Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H2. However, hydrogen gas is very rare in the Earth's atmosphere (1 ppm by volume) because of its light weight, which enables it to escape from Earth's gravity more easily than heavier gases. However, hydrogen is the third most abundant element on the Earth's surface, [85] mostly in the form of chemical compounds such as hydrocarbons and water. [44] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus, as is methane, itself a hydrogen source of increasing importance. [86]

A molecular form called protonated molecular hydrogen (H+
) is found in the interstellar medium, where it is generated by ionization of molecular hydrogen from cosmic rays. This ion has also been observed in the upper atmosphere of the planet Jupiter. The ion is relatively stable in the environment of outer space due to the low temperature and density. H+
is one of the most abundant ions in the universe, and it plays a notable role in the chemistry of the interstellar medium. [87] Neutral triatomic hydrogen H3 can exist only in an excited form and is unstable. [88] By contrast, the positive hydrogen molecular ion (H+
) is a rare molecule in the universe.



Electrolysis of water

The electrolysis of water is a simple method of producing hydrogen. A low voltage current is run through the water, and gaseous oxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or another inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used vs. energetic value of hydrogen produced) is in the range 88–94%. [89] [90]

2 H
(l) → 2 H
(g) + O

When determining the electrical efficiency of PEM (proton exchange membrane) electrolysis, the higher heat value (HHV) is used. [91] This is because the catalyst layer interacts with water as steam. As the process operates at 80 °C for PEM electrolysers the waste heat can be redirected through the system to create the steam, resulting in a higher overall electrical efficiency. The lower heat value (LHV) must be used for alkaline electrolysers as the process within these electrolysers requires water in liquid form and uses alkalinity to facilitate the breaking of the bond holding the hydrogen and oxygen atoms together. The lower heat value must also be used for fuel cells, as steam is the output rather than input.

Steam reforming (Industrial Method)

Hydrogen is often produced using natural gas, which involves the removal of hydrogen from hydrocarbons at very high temperatures, with about 95% of hydrogen production coming from steam reforming around year 2000. [92] Commercial bulk hydrogen is usually produced by the steam reforming of natural gas. [93] This method is also known as the Bosch process and is widely used for the industrial preparation of hydrogen.

At high temperatures (1000–1400 K, 700–1100 °C or 1300–2000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H

+ H
→ CO + 3 H

This reaction is favored at low pressures but is nonetheless conducted at high pressures (2.0  MPa, 20 atm or 600  inHg). This is because high-pressure H
is the most marketable product, and pressure swing adsorption (PSA) purification systems work better at higher pressures. The product mixture is known as "synthesis gas" because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon:

→ C + 2 H

Consequently, steam reforming typically employs an excess of H
. Additional hydrogen can be recovered from the steam by use of carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide: [93]

CO + H
+ H

Other important methods for H
production include partial oxidation of hydrocarbons: [94]

2 CH
+ O
→ 2 CO + 4 H

and the coal reaction, which can serve as a prelude to the shift reaction above: [93]

C + H
→ CO + H

Hydrogen is sometimes produced and consumed in the same industrial process, without being separated. In the Haber process for the production of ammonia, hydrogen is generated from natural gas. [95] Electrolysis of brine to yield chlorine also produces hydrogen as a co-product. [96]


Many metals react with water to produce H
, but the rate of hydrogen evolution depends on the metal, the pH, and the presence alloying agents. Most commonly, hydrogen evolution is induced by acids. The alkali and alkaline earth metals, aluminium, zinc, manganese, and iron react readily with aqueous acids. This reaction is the basis of the Kipp's apparatus, which once was used as a source of laboratory:

Zn + 2 H+
+ H

In the absence of acid, the evolution of H
is slower. Of technological significance because iron is widely used structural material, is its anaerobic corrosion:

Fe + 2 H
O → Fe(OH)
+ H

Many metals, e.g. aluminium, are slow to react with water because they form passivated coatings of oxides. An alloy of aluminium and gallium however does react with water. [97] } At high pH, aluminium can produce H

2 Al + 6 H
+ 2 OH
→ 2 Al(OH)
+ 3 H

Some metal-containing compounds react with acids to evolve H
. Under anaerobic conditions, ferrous hydroxide (Fe(OH)
) is be oxidized by the protons of water to form magnetite and H
. This process is described by the Schikorr reaction:

3 Fe(OH)
+ 2 H
O + H

This process occurs during the anaerobic corrosion of iron and steel in oxygen-free groundwater and in reducing soils below the water table.


More than 200 thermochemical cycles can be used for water splitting. Many of these cycles such as the iron oxide cycle, cerium(IV) oxide–cerium(III) oxide cycle, zinc zinc-oxide cycle, sulfur-iodine cycle, copper-chlorine cycle and hybrid sulfur cycle have been evaluated for their commercial potential to produce hydrogen and oxygen from water and heat without using electricity. [98] A number of laboratories (including in France, Germany, Greece, Japan, and the USA) are developing thermochemical methods to produce hydrogen from solar energy and water. [99]

Serpentinization reaction

In deep geological conditions prevailing far away from the Earth's atmosphere, hydrogen (H
) is produced during the process of serpentinization. In this process, water protons (H+) are reduced by ferrous (Fe2+) ions provided by fayalite (Fe
). The reaction forms magnetite (Fe
), quartz (SiO
), and hydrogen (H
): [100] [101]

+ 2 H
O → 2 Fe
+ 3 SiO
+ 3 H
fayalite + water → magnetite + quartz + hydrogen

This reaction closely resembles the Schikorr reaction observed in anaerobic oxidation of ferrous hydroxide in contact with water.


Petrochemical industry

Large quantities of H
are used in the "upgrading" of fossil fuels. Key consumers of H
include hydrodealkylation, hydrodesulfurization, and hydrocracking. Many of these reactions can be classified as hydrogenolysis, i.e., the cleavage of bonds to carbon. Illustrative is the separation of sulfur from liquid fossil fuels:

R-S-R + 2 H2 → H2S + 2 RH


Hydrogenation, the addition of H
to various substrates is conducted on a large scale. The hydrogenation of N2 to produce ammonia by the Haber-Bosch Process consumes a few percent of the energy budget in the entire industry. The resulting ammonia is used to supply the majority of the protein consumed by humans. [102] Hydrogenation is used to convert unsaturated fats and oils to saturated fats and oils. The major application is the production of margarine. Methanol is produced by hydrogenation of carbon dioxide. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H
is also used as a reducing agent for the conversion of some ores to the metals. [103]


Hydrogen is commonly used in power stations as a coolant in generators due to a number of favorable properties that are a direct result of its light diatomic molecules. These include low density, low viscosity, and the highest specific heat and thermal conductivity of all gases.

Energy carrier

Hydrogen is not an energy resource, [104] except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development. [105] The Sun's energy comes from nuclear fusion of hydrogen, but this process is difficult to achieve controllably on Earth. [106] Elemental hydrogen from solar, biological, or electrical sources requires more energy to make than is obtained by burning it, so in these cases hydrogen functions as an energy carrier, like a battery. Hydrogen may be obtained from fossil sources (such as methane), but these sources are unsustainable. [104]

The energy density per unit volume of both liquid hydrogen and compressed hydrogen gas at any practicable pressure is significantly less than that of traditional fuel sources, although the energy density per unit fuel mass is higher. [104] Nevertheless, elemental hydrogen has been widely discussed in the context of energy, as a possible future carrier of energy on an economy-wide scale. [107] For example, CO
sequestration followed by carbon capture and storage could be conducted at the point of H
production from fossil fuels. [108] Hydrogen used in transportation would burn relatively cleanly, with some NOx emissions, [109] but without carbon emissions. [108] However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial. [110] Fuel cells can convert hydrogen and oxygen directly to electricity more efficiently than internal combustion engines. [111]

Semiconductor industry

Hydrogen is employed to saturate broken ("dangling") bonds of amorphous silicon and amorphous carbon that helps stabilizing material properties. [112] It is also a potential electron donor in various oxide materials, including ZnO, [113] [114] SnO2, CdO, MgO, [115] ZrO2, HfO2, La2O3, Y2O3, TiO2, SrTiO3, LaAlO3, SiO2, Al2O3, ZrSiO4, HfSiO4, and SrZrO3. [116]

Niche and evolving uses

Apart from its use as a reactant, H
has a variety of smaller applications. It is used as a shielding gas in welding methods such as atomic hydrogen welding. [117] [118] H2 is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas. Liquid H2 is used in cryogenic research, including superconductivity studies. [119] Because H
is lighter than air, having a little more than 114 of the density of air, it was once widely used as a lifting gas in balloons and airships. [120]

Pure or mixed with nitrogen (sometimes called forming gas), hydrogen is a tracer gas for detection of minute leaks. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries. [121] Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties. [122]

Hydrogen's rarer isotopes also each have specific applications. Deuterium (hydrogen-2) is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions. [5] Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects. [123] Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, [124] as an isotopic label in the biosciences, [60] and as a radiation source in luminous paints. [125]

The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale at 13.8033  Kelvin. [126]

Biological reactions

H2 is a product of some types of anaerobic metabolism and is produced by several microorganisms, usually via reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H2 and its component two protons and two electrons. Creation of hydrogen gas occurs in the transfer of reducing equivalents produced during pyruvate fermentation to water. [127] The natural cycle of hydrogen production and consumption by organisms is called the hydrogen cycle. [128] parts per million (ppm) of H2 occurs in the breath of healthy humans. It results from the metabolic activity of hydrogenase-containing microorganisms in the large intestine. [129]

Water splitting, in which water is decomposed into its component protons, electrons, and oxygen, occurs in the light reactions in all photosynthetic organisms. Some such organisms, including the alga Chlamydomonas reinhardtii and cyanobacteria, have evolved a second step in the dark reactions in which protons and electrons are reduced to form H2 gas by specialized hydrogenases in the chloroplast. [130] Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H2 gas even in the presence of oxygen. [131] Efforts have also been undertaken with genetically modified alga in a bioreactor. [132]

Safety and precautions

GHS pictograms GHS-pictogram-flamme.svg
GHS Signal word Danger
P202, P210, P271, P403, P377, P381 [133]
NFPA 704 (fire diamond)
Flammability code 4: Will rapidly or completely vaporize at normal atmospheric pressure and temperature, or is readily dispersed in air and will burn readily. Flash point below 23 °C (73 °F). E.g. propaneHealth code 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g. sodium chlorideReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeHydrogen

Hydrogen poses a number of hazards to human safety, from potential detonations and fires when mixed with air to being an asphyxiant in its pure, oxygen-free form. [134] In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids. [135] Hydrogen dissolves in many metals and in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement, [136] leading to cracks and explosions. [137] Hydrogen gas leaking into external air may spontaneously ignite. Moreover, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to accidental burns. [138]

Even interpreting the hydrogen data (including safety data) is confounded by a number of phenomena. Many physical and chemical properties of hydrogen depend on the parahydrogen/orthohydrogen ratio (it often takes days or weeks at a given temperature to reach the equilibrium ratio, for which the data is usually given). Hydrogen detonation parameters, such as critical detonation pressure and temperature, strongly depend on the container geometry. [134]


  1. However, most of the universe's mass is not in the form of baryons or chemical elements. See dark matter and dark energy.
  2. 286 kJ/mol: energy per mole of the combustible material (molecular hydrogen).

Related Research Articles

Alkali metal Group of highly-reactive chemical elements

The alkali metals consist of the chemical elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). Together with hydrogen they constitute group 1, which lies in the s-block of the periodic table. All alkali metals have their outermost electron in an s-orbital: this shared electron configuration results in their having very similar characteristic properties. Indeed, the alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting well-characterised homologous behaviour. This family of elements is also known as the lithium family after its leading element.

Deuterium Isotope of hydrogen with 1 neutron

Deuterium is one of two stable isotopes of hydrogen. The nucleus of a deuterium atom, called a deuteron, contains one proton and one neutron, whereas the far more common protium has no neutrons 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% 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.

Heavy water Form of water

Heavy water is a form of water that contains only deuterium rather than the common hydrogen-1 isotope (1
or H, also called protium] that makes up most of the hydrogen in normal water. The presence of the heavier hydrogen isotope gives the water different nuclear properties, and the increase of mass gives it slightly different physical and chemical properties when compared to normal water.

Proton Nucleon (constituent of the nucleus of the atom) that has positive electric charge; symbol p or p+

A proton is a subatomic particle, symbol
, with a positive electric charge of +1e elementary charge and a mass slightly less than that of a neutron. Protons and neutrons, each with masses of approximately one atomic mass unit, are collectively referred to as "nucleons".

Tritium isotope of hydrogen with 2 neutrons

Tritium or hydrogen-3 is a rare and radioactive isotope of hydrogen. The nucleus of tritium 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.

In chemistry, a hydride is formally the anion of hydrogen, H. The term is applied loosely. At one extreme, all compounds containing covalently bound H atoms, are called hydrides: water is a hydride of oxygen, ammonia is a hydride of nitrogen, etc. For inorganic chemists, hydrides refer to compound and ions in which hydrogen is covalently attached to a less electronegative element. In such cases, the H centre has nucleophilic character, which contrasts with the protic character of acids. The hydride anion is very rarely observed.

A period 1 element is one of the chemical elements in the first row of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate periodic (recurring) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that analog elements fall into the same vertical columns. The first period contains fewer elements than any other row in the table, with only two: hydrogen and helium. This situation can be explained by modern theories of atomic structure. In a quantum mechanical description of atomic structure, this period corresponds to the filling of the 1s orbital. Period 1 elements obey the duet rule in that they need two electrons to complete their valence shell.

Neutron generator source of neutrons from linear particle accelerators

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 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 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.

Isotopes of hydrogen hydrogen with different numbers of neutrons


Lithium hydride chemical compound

Lithium hydride is an inorganic compound with the formula LiH. This alkali metal hydride is a colorless solid, although commercial samples are grey. Characteristic of a salt-like (ionic) hydride, it has a high melting point, and it is not soluble but reactive with all organic and protic solvents. It is soluble and nonreactive with certain molten salts such as lithium fluoride, lithium borohydride, and sodium hydride. With a molecular mass of slightly less than 8.0, it is the lightest ionic compound.

Trihydrogen cation chemical compound

The trihydrogen cation or protonated molecular hydrogen is a cation with formula H+
, consisting of three hydrogen nuclei (protons) sharing two electrons.

In chemistry, a hydron is the general name for a cationic form of atomic hydrogen, represented with the symbol H+
. However, virtually all chemists will call this species the "proton", which strictly speaking refers to the cation of protium, the most common isotope of hydrogen. The term "hydron", endorsed by the IUPAC, 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. Unless there is a need to explicitly address issues of isotopic composition, the term "hydron" is not used, and to refer to H+
as such is considered pedantic. Nevertheless, H+
will be referred to as the "hydron", for the sake of consistency with the article title.

Helium hydride ion Chemical compound

The helium hydride ion or hydridohelium(1+) ion or helonium is a cation (positively charged ion) with chemical formula HeH+. It consists of a helium atom bonded to a hydrogen atom, with one electron removed. It can also be viewed as protonated helium. It is the lightest heteronuclear ion, and is believed to be the first compound formed in the Universe after the Big Bang.

Transition metal hydrides are chemical compounds containing a transition metal bonded to hydrogen. Most transition metals form hydride complexes and some are significant in various catalytic and synthetic reactions. The term "hydride" is used loosely: some so-called hydrides are acidic (e.g., H2Fe(CO)4), whereas some others are hydridic, having H-like character (e.g., ZnH2).

Properties of water Physical and chemical properties of pure water

Water is a polar inorganic compound that is at room temperature a tasteless and odorless liquid, which is nearly colorless apart from an inherent hint of blue. It is by far the most studied chemical compound and is described as the "universal solvent" and the "solvent of life". It is the most abundant substance on Earth and the only common substance to exist as a solid, liquid, and gas on Earth's surface. It is also the third most abundant molecule in the universe.

Uranium hydride, also called uranium trihydride (UH3), is an inorganic compound and a hydride of uranium.

Chromium(I) hydride chemical compound

Chromium(I) hydride, systematically named chromium hydride, is an inorganic compound with the chemical formula (CrH)
. It occurs naturally in some kinds of stars where it has been detected by its spectrum. However, molecular chromium(I) hydride with the formula CrH has been isolated in solid gas matrices. The molecular hydride is very reactive. As such the compound is not well characterised, although many of its properties have been calculated via computational chemistry.

Iron(II) hydride, systematically named iron dihydride and poly(dihydridoiron) is solid inorganic compound with the chemical formula (FeH
. ). It is kinetically unstable at ambient temperature, and as such, little is known about its bulk properties. However, it known as a black, amorphous powder, which was synthesised for the first time in 2014.

Hydrogen chalcogenides are binary compounds of hydrogen with chalcogen atoms. Water, the first chemical compound in this series, contains one oxygen atom and two hydrogen atoms, and is the most common compound on the Earth's surface.

In chemistry, the decay technique is a method to generate chemical species such as radicals, carbocations, and other potentially unstable covalent structures by radioactive decay of other compounds. For example, decay of a tritium-labeled molecule yields an ionized helium atom, which might then break off to leave a cationic molecular fragment.


  1. Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 240. ISBN   978-0123526519.
  2. Lide, D. R., ed. (2005). "Magnetic susceptibility of the elements and inorganic compounds". CRC Handbook of Chemistry and Physics (PDF) (86th ed.). Boca Raton (FL): CRC Press. ISBN   978-0-8493-0486-6.
  3. Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN   978-0-8493-0464-4.
  4. 1 2 3 "Hydrogen". Van Nostrand's Encyclopedia of Chemistry. Wylie-Interscience. 2005. pp. 797–799. ISBN   978-0-471-61525-5.
  5. 1 2 3 4 5 6 7 8 9 10 11 12 Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 183–191. ISBN   978-0-19-850341-5.
  6. Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 16–21. ISBN   978-0-19-508083-4.
  7. Boyd, Padi (19 July 2014). "What is the chemical composition of stars?". NASA. Archived from the original on 15 January 2015. Retrieved 5 February 2008.
  8. Laursen, S.; Chang, J.; Medlin, W.; Gürmen, N.; Fogler, H. S. (27 July 2004). "An extremely brief introduction to computational quantum chemistry". Molecular Modeling in Chemical Engineering. University of Michigan. Archived from the original on 20 May 2015. Retrieved 4 May 2015.
  9. Presenter: Professor Jim Al-Khalili (21 January 2010). "Discovering the Elements". Chemistry: A Volatile History. 25:40 minutes in. BBC. BBC Four. Archived from the original on 25 January 2010. Retrieved 9 February 2010.
  10. "Hydrogen Basics – Production". Florida Solar Energy Center. 2007. Archived from the original on 18 February 2008. Retrieved 5 February 2008.
  11. 1 2 Rogers, H. C. (1999). "Hydrogen Embrittlement of Metals". Science . 159 (3819): 1057–1064. Bibcode:1968Sci...159.1057R. doi:10.1126/science.159.3819.1057. PMID   17775040.
  12. 1 2 Christensen, C. H.; Nørskov, J. K.; Johannessen, T. (9 July 2005). "Making society independent of fossil fuels – Danish researchers reveal new technology". Technical University of Denmark. Archived from the original on 21 May 2015. Retrieved 19 May 2015.
  13. "Dihydrogen". O=CHem Directory. University of Southern Maine. Archived from the original on 13 February 2009. Retrieved 6 April 2009.
  14. Committee on Alternatives and Strategies for Future Hydrogen Production and Use, US National Research Council, US National Academy of Engineering (2004). The Hydrogen Economy: Opportunities, Costs, Barriers, and R&D Needs. National Academies Press. p. 240. ISBN   978-0-309-09163-3.CS1 maint: multiple names: authors list (link)
  15. Carcassi, M. N.; Fineschi, F. (2005). "Deflagrations of H2–air and CH4–air lean mixtures in a vented multi-compartment environment". Energy. 30 (8): 1439–1451. doi:10.1016/j.energy.2004.02.012.
  16. Patnaik, P. (2007). A Comprehensive Guide to the Hazardous Properties of Chemical Substances. Wiley-Interscience. p. 402. ISBN   978-0-471-71458-3.
  17. Schefer, E. W.; Kulatilaka, W. D.; Patterson, B. D.; Settersten, T. B. (June 2009). "Visible emission of hydrogen flames". Combustion and Flame. 156 (6): 1234–1241. doi:10.1016/j.combustflame.2009.01.011.
  19. Clayton, D. D. (2003). Handbook of Isotopes in the Cosmos: Hydrogen to Gallium. Cambridge University Press. ISBN   978-0-521-82381-4.
  20. NAAP Labs (2009). "Energy Levels". University of Nebraska Lincoln. Archived from the original on 11 May 2015. Retrieved 20 May 2015.
  21. "photon wavelength 13.6 eV". Wolfram Alpha. 20 May 2015. Archived from the original on 12 May 2016. Retrieved 20 May 2015.
  22. Stern, D. P. (16 May 2005). "The Atomic Nucleus and Bohr's Early Model of the Atom". NASA Goddard Space Flight Center (mirror). Archived from the original on 17 October 2008. Retrieved 20 December 2007.
  23. Stern, D. P. (13 February 2005). "Wave Mechanics". NASA Goddard Space Flight Center. Archived from the original on 13 May 2008. Retrieved 16 April 2008.
  24. Staff (2003). "Hydrogen (H2) Properties, Uses, Applications: Hydrogen Gas and Liquid Hydrogen". Universal Industrial Gases, Inc. Archived from the original on 19 February 2008. Retrieved 5 February 2008.
  25. Tikhonov, V. I.; Volkov, A. A. (2002). "Separation of Water into Its Ortho and Para Isomers". Science. 296 (5577): 2363. doi:10.1126/science.1069513. PMID   12089435.
  26. Hritz, J. (March 2006). "CH. 6 – Hydrogen" (PDF). NASA Glenn Research Center Glenn Safety Manual, Document GRC-MQSA.001. NASA. Archived (PDF) from the original on 16 February 2008. Retrieved 5 February 2008.
  27. Shinitzky, M.; Elitzur, A. C. (2006). "Ortho-para spin isomers of the protons in the methylene group". Chirality. 18 (9): 754–756. doi:10.1002/chir.20319. PMID   16856167.
  28. Milenko, Yu. Ya.; Sibileva, R. M.; Strzhemechny, M. A. (1997). "Natural ortho-para conversion rate in liquid and gaseous hydrogen". Journal of Low Temperature Physics. 107 (1–2): 77–92. Bibcode:1997JLTP..107...77M. doi:10.1007/BF02396837.
  29. Amos, Wade A. (1 November 1998). "Costs of Storing and Transporting Hydrogen" (PDF). National Renewable Energy Laboratory. pp. 6–9. Archived (PDF) from the original on 26 December 2014. Retrieved 19 May 2015.
  30. Svadlenak, R. E.; Scott, A. B. (1957). "The Conversion of Ortho- to Parahydrogen on Iron Oxide-Zinc Oxide Catalysts". Journal of the American Chemical Society. 79 (20): 5385–5388. doi:10.1021/ja01577a013.
  31. Clark, J. (2002). "The Acidity of the Hydrogen Halides". Chemguide. Archived from the original on 20 February 2008. Retrieved 9 March 2008.
  32. Kimball, J. W. (7 August 2003). "Hydrogen". Kimball's Biology Pages. Archived from the original on 4 March 2008. Retrieved 4 March 2008.
  33. IUPAC Compendium of Chemical Terminology, Electronic version, Hydrogen Bond Archived 19 March 2008 at the Wayback Machine
  34. Sandrock, G. (2 May 2002). "Metal-Hydrogen Systems". Sandia National Laboratories. Archived from the original on 24 February 2008. Retrieved 23 March 2008.
  35. 1 2 "Structure and Nomenclature of Hydrocarbons". Purdue University. Archived from the original on 31 July 2012. Retrieved 23 March 2008.
  36. "Organic Chemistry". Dictionary.com. Lexico Publishing Group. 2008. Archived from the original on 18 April 2008. Retrieved 23 March 2008.
  37. "Biochemistry". Dictionary.com. Lexico Publishing Group. 2008. Archived from the original on 29 March 2008. Retrieved 23 March 2008.
  38. Takeshita, T.; Wallace, W. E.; Craig, R. S. (1974). "Hydrogen solubility in 1:5 compounds between yttrium or thorium and nickel or cobalt". Inorganic Chemistry . 13 (9): 2282–2283. doi:10.1021/ic50139a050.
  39. Kirchheim, R.; Mutschele, T.; Kieninger, W.; Gleiter, H.; Birringer, R.; Koble, T. (1988). "Hydrogen in amorphous and nanocrystalline metals". Materials Science and Engineering. 99 (1–2): 457–462. doi:10.1016/0025-5416(88)90377-1.
  40. Kirchheim, R. (1988). "Hydrogen solubility and diffusivity in defective and amorphous metals". Progress in Materials Science . 32 (4): 262–325. doi:10.1016/0079-6425(88)90010-2.
  41. Moers, K. (1920). "Investigations on the Salt Character of Lithium Hydride". Zeitschrift für Anorganische und Allgemeine Chemie. 113 (191): 179–228. doi:10.1002/zaac.19201130116. Archived (PDF) from the original on 24 August 2019. Retrieved 24 August 2019.
  42. Downs, A. J.; Pulham, C. R. (1994). "The hydrides of aluminium, gallium, indium, and thallium: a re-evaluation". Chemical Society Reviews. 23 (3): 175–184. doi:10.1039/CS9942300175.
  43. Hibbs, D. E.; Jones, C.; Smithies, N. A. (1999). "A remarkably stable indium trihydride complex: synthesis and characterisation of [InH3P(C6H11)3]". Chemical Communications (2): 185–186. doi:10.1039/a809279f.
  44. 1 2 3 Miessler, G. L.; Tarr, D. A. (2003). Inorganic Chemistry (3rd ed.). Prentice Hall. ISBN   978-0-13-035471-6.
  45. Okumura, A. M.; Yeh, L. I.; Myers, J. D.; Lee, Y. T. (1990). "Infrared spectra of the solvated hydronium ion: vibrational predissociation spectroscopy of mass-selected H3O+•(H2O)n•(H2)m". Journal of Physical Chemistry. 94 (9): 3416–3427. doi:10.1021/j100372a014.
  46. Perdoncin, G.; Scorrano, G. (1977). "Protonation Equilibria in Water at Several Temperatures of Alcohols, Ethers, Acetone, Dimethyl Sulfide, and Dimethyl Sulfoxide". Journal of the American Chemical Society. 99 (21): 6983–6986. doi:10.1021/ja00463a035.
  47. Carrington, A.; McNab, I. R. (1989). "The infrared predissociation spectrum of triatomic hydrogen cation (H3+)". Accounts of Chemical Research. 22 (6): 218–222. doi:10.1021/ar00162a004.
  48. "NASA/TM—2002-211915 : Solid Hydrogen Experiments for Atomic Propellants" (PDF). Archived from the original (PDF) on 27 September 2011. Retrieved 27 September 2011.
  49. "Hydrogen". mysite.du.edu. Archived from the original on 18 April 2009. Retrieved 20 April 2008.
  50. Gurov, Y. B.; Aleshkin, D. V.; Behr, M. N.; Lapushkin, S. V.; Morokhov, P. V.; Pechkurov, V. A.; Poroshin, N. O.; Sandukovsky, V. G.; Tel'kushev, M. V.; Chernyshev, B. A.; Tschurenkova, T. D. (2004). "Spectroscopy of superheavy hydrogen isotopes in stopped-pion absorption by nuclei". Physics of Atomic Nuclei. 68 (3): 491–97. Bibcode:2005PAN....68..491G. doi:10.1134/1.1891200.
  51. Korsheninnikov, A.; Nikolskii, E.; Kuzmin, E.; Ozawa, A.; Morimoto, K.; Tokanai, F.; Kanungo, R.; Tanihata, I.; 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. PMID   12633420.
  52. Urey, H. C.; Brickwedde, F. G.; Murphy, G. M. (1933). "Names for the Hydrogen Isotopes". Science. 78 (2035): 602–603. Bibcode:1933Sci....78..602U. doi:10.1126/science.78.2035.602. PMID   17797765.
  53. Oda, Y.; Nakamura, H.; Yamazaki, T.; Nagayama, K.; Yoshida, M.; Kanaya, S.; Ikehara, M. (1992). "1H NMR studies of deuterated ribonuclease HI selectively labeled with protonated amino acids". Journal of Biomolecular NMR . 2 (2): 137–47. doi:10.1007/BF01875525. PMID   1330130.
  54. Broad, W. J. (11 November 1991). "Breakthrough in Nuclear Fusion Offers Hope for Power of Future". The New York Times. Retrieved 12 February 2008.
  55. Traub, R. J.; Jensen, J. A. (June 1995). "Tritium radioluminescent devices, Health and Safety Manual" (PDF). International Atomic Energy Agency. p. 2.4. Archived (PDF) from the original on 6 September 2015. Retrieved 20 May 2015.
  56. Staff (15 November 2007). "Tritium". U.S. Environmental Protection Agency. Archived from the original on 2 January 2008. Retrieved 12 February 2008.
  57. Nave, C. R. (2006). "Deuterium-Tritium Fusion". HyperPhysics. Georgia State University. Archived from the original on 16 March 2008. Retrieved 8 March 2008.
  58. Kendall, C.; Caldwell, E. (1998). C. Kendall; J. J. McDonnell (eds.). "Chapter 2: Fundamentals of Isotope Geochemistry". Isotope Tracers in Catchment Hydrology. US Geological Survey. Archived from the original on 14 March 2008. Retrieved 8 March 2008.
  59. "The Tritium Laboratory". University of Miami. 2008. Archived from the original on 28 February 2008. Retrieved 8 March 2008.
  60. 1 2 Holte, A. E.; Houck, M. A.; Collie, N. L. (2004). "Potential Role of Parasitism in the Evolution of Mutualism in Astigmatid Mites". Experimental and Applied Acarology. 25 (2): 97–107. doi:10.1023/A:1010655610575. PMID   11513367.
  61. van der Krogt, P. (5 May 2005). "Hydrogen". Elementymology & Elements Multidict. Archived from the original on 23 January 2010. Retrieved 20 December 2010.
  62. § IR-3.3.2, Provisional Recommendations Archived 9 February 2016 at the Wayback Machine , Nomenclature of Inorganic Chemistry, Chemical Nomenclature and Structure Representation Division, IUPAC. Accessed on line 3 October 2007.
  63. IUPAC (1997). "Muonium". In A.D. McNaught, A. Wilkinson (ed.). Compendium of Chemical Terminology (2nd ed.). Blackwell Scientific Publications. doi:10.1351/goldbook.M04069. ISBN   978-0-86542-684-9.
  64. V.W. Hughes; et al. (1960). "Formation of Muonium and Observation of its Larmor Precession". Physical Review Letters . 5 (2): 63–65. Bibcode:1960PhRvL...5...63H. doi:10.1103/PhysRevLett.5.63.
  65. W.H. Koppenol; IUPAC (2001). "Names for muonium and hydrogen atoms and their ions" (PDF). Pure and Applied Chemistry . 73 (2): 377–380. doi:10.1351/pac200173020377. Archived (PDF) from the original on 14 May 2011. Retrieved 15 November 2016.
  66. Boyle, R. (1672). "Tracts written by the Honourable Robert Boyle containing new experiments, touching the relation betwixt flame and air..." London.
  67. Winter, M. (2007). "Hydrogen: historical information". WebElements Ltd. Archived from the original on 10 April 2008. Retrieved 5 February 2008.
  68. Musgrave, A. (1976). "Why did oxygen supplant phlogiston? Research programmes in the Chemical Revolution". In Howson, C. (ed.). Method and appraisal in the physical sciences . The Critical Background to Modern Science, 1800–1905. Cambridge University Press. doi:10.1017/CBO9780511760013. ISBN   9780521211109 . Retrieved 22 October 2011.
  69. Cavendish, Henry (12 May 1766). "Three Papers, Containing Experiments on Factitious Air, by the Hon. Henry Cavendish, F. R. S." Philosophical Transactions. 56: 141–184. Bibcode:1766RSPT...56..141C. doi: 10.1098/rstl.1766.0019 . JSTOR   105491.
  71. National Electrical Manufacturers Association (1946). A chronological history of electrical development from 600 B.C. New York, N.Y., National Electrical Manufacturers Association. p. 102. Archived from the original on 4 March 2016. Retrieved 9 February 2016.
  72. Stockel, J.F; j.d. Dunlop; Betz, F (1980). "NTS-2 Nickel-Hydrogen Battery Performance 31". Journal of Spacecraft and Rockets. 17: 31–34. Bibcode:1980JSpRo..17...31S. doi:10.2514/3.57704.
  73. Jannette, A. G.; Hojnicki, J. S.; McKissock, D. B.; Fincannon, J.; Kerslake, T. W.; Rodriguez, C. D. (July 2002). Validation of international space station electrical performance model via on-orbit telemetry (PDF). IECEC '02. 2002 37th Intersociety Energy Conversion Engineering Conference, 2002. pp. 45–50. doi:10.1109/IECEC.2002.1391972. ISBN   0-7803-7296-4. Archived (PDF) from the original on 14 May 2010. Retrieved 11 November 2011.
  74. Anderson, P. M.; Coyne, J. W. (2002). A lightweight high reliability single battery power system for interplanetary spacecraft. Aerospace Conference Proceedings. 5. pp. 5–2433. doi:10.1109/AERO.2002.1035418. ISBN   978-0-7803-7231-3.
  75. "Mars Global Surveyor". Astronautix.com. Archived from the original on 10 August 2009. Retrieved 6 April 2009.
  76. Lori Tyahla, ed. (7 May 2009). "Hubble servicing mission 4 essentials". NASA. Archived from the original on 13 March 2015. Retrieved 19 May 2015.
  77. Hendrix, Susan (25 November 2008). Lori Tyahla (ed.). "Extending Hubble's mission life with new batteries". NASA. Archived from the original on 5 March 2016. Retrieved 19 May 2015.
  78. Crepeau, R. (1 January 2006). Niels Bohr: The Atomic Model. Great Scientific Minds. ISBN   978-1-4298-0723-4.
  79. Berman, R.; Cooke, A. H.; Hill, R. W. (1956). "Cryogenics". Annual Review of Physical Chemistry. 7: 1–20. Bibcode:1956ARPC....7....1B. doi:10.1146/annurev.pc.07.100156.000245.
  80. Charlton, Mike; Van Der Werf, Dirk Peter (1 March 2015). "Advances in antihydrogen physics". Science Progress. 98 (1): 34–62. doi:10.3184/003685015X14234978376369. PMID   25942774.
  81. Kellerbauer, Alban (29 January 2015). "Why Antimatter Matters". European Review. 23 (1): 45–56. doi:10.1017/S1062798714000532.
  82. Gagnon, S. "Hydrogen". Jefferson Lab. Archived from the original on 10 April 2008. Retrieved 5 February 2008.
  83. Haubold, H.; Mathai, A. M. (15 November 2007). "Solar Thermonuclear Energy Generation". Columbia University. Archived from the original on 11 December 2011. Retrieved 12 February 2008.
  84. Storrie-Lombardi, L. J.; Wolfe, A. M. (2000). "Surveys for z > 3 Damped Lyman-alpha Absorption Systems: the Evolution of Neutral Gas". Astrophysical Journal. 543 (2): 552–576. arXiv: astro-ph/0006044 . Bibcode:2000ApJ...543..552S. doi:10.1086/317138.
  85. Dresselhaus, M.; et al. (15 May 2003). "Basic Research Needs for the Hydrogen Economy" (PDF). APS March Meeting Abstracts. Argonne National Laboratory, U.S. Department of Energy, Office of Science Laboratory. 2004: m1.001. Bibcode:2004APS..MAR.m1001D. Archived from the original (PDF) on 13 February 2008. Retrieved 5 February 2008.
  86. Berger, W. H. (15 November 2007). "The Future of Methane". University of California, San Diego. Archived from the original on 24 April 2008. Retrieved 12 February 2008.
  87. McCall Group; Oka Group (22 April 2005). "H3+ Resource Center". Universities of Illinois and Chicago. Archived from the original on 11 October 2007. Retrieved 5 February 2008.
  88. Helm, H.; et al. (2003), "Coupling of Bound States to Continuum States in Neutral Triatomic Hydrogen", Dissociative Recombination of Molecular Ions with Electrons, Department of Molecular and Optical Physics, University of Freiburg, Germany, pp. 275–288, doi:10.1007/978-1-4615-0083-4_27, ISBN   978-1-4613-4915-0
  89. Thomassen, Magnus. "Cost reduction and performance increase of PEM electrolysers" (PDF). fch.europa.eu/. FCH JU. Archived (PDF) from the original on 17 April 2018. Retrieved 22 April 2018.
  90. Kruse, B.; Grinna, S.; Buch, C. (2002). "Hydrogen Status og Muligheter" (PDF). Bellona. Archived from the original (PDF) on 16 February 2008. Retrieved 12 February 2008.
  91. Kruse, Bjørnar. "Hydrogen Status og muligheter" (PDF). bellona.org/. Bellona Norway. Archived (PDF) from the original on 22 April 2018. Retrieved 22 April 2018.
  92. Ogden, J. M. (1999). "Prospects for building a hydrogen energy infrastructure". Annual Review of Energy and the Environment. 24: 227–279. doi:10.1146/annurev.energy.24.1.227.
  93. 1 2 3 Oxtoby, D. W. (2002). Principles of Modern Chemistry (5th ed.). Thomson Brooks/Cole. ISBN   978-0-03-035373-4.
  94. "Hydrogen Properties, Uses, Applications". Universal Industrial Gases, Inc. 2007. Archived from the original on 27 March 2008. Retrieved 11 March 2008.
  95. Funderburg, E. (2008). "Why Are Nitrogen Prices So High?". The Samuel Roberts Noble Foundation. Archived from the original on 9 May 2001. Retrieved 11 March 2008.
  96. Lees, A. (2007). "Chemicals from salt". BBC. Archived from the original on 26 October 2007. Retrieved 11 March 2008.
  97. Parmuzina, A.V.; Kravchenko, O.V. (2008). "Activation of aluminium metal to evolve hydrogen from water". International Journal of Hydrogen Energy. 33 (12): 3073–3076. doi:10.1016/j.ijhydene.2008.02.025.
  98. Weimer, Al (25 May 2005). "Development of solar-powered thermochemical production of hydrogen from water" (PDF). Solar Thermochemical Hydrogen Generation Project. Archived (PDF) from the original on 17 April 2007. Retrieved 21 December 2008.
  99. Perret, R. "Development of Solar-Powered Thermochemical Production of Hydrogen from Water, DOE Hydrogen Program, 2007" (PDF). Archived (PDF) from the original on 27 May 2010. Retrieved 17 May 2008.
  100. Russell, M. J.; Hall, A. J.; Martin, W. (2010). "Serpentinization as a source of energy at the origin of life". Geobiology. 8 (5): 355–371. doi:10.1111/j.1472-4669.2010.00249.x. PMID   20572872.
  101. Schrenk, M. O.; Brazelton, W. J.; Lang, S. Q. (2013). "Serpentinization, Carbon, and Deep Life". Reviews in Mineralogy and Geochemistry. 75 (1): 575–606. Bibcode:2013RvMG...75..575S. doi:10.2138/rmg.2013.75.18.
  102. Smil, Vaclav (2004). Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production (1st ed.). Cambridge, MA: MIT. ISBN   9780262693134.
  103. Chemistry Operations (15 December 2003). "Hydrogen". Los Alamos National Laboratory. Archived from the original on 4 March 2011. Retrieved 5 February 2008.
  104. 1 2 3 McCarthy, J. (31 December 1995). "Hydrogen". Stanford University. Archived from the original on 14 March 2008. Retrieved 14 March 2008.
  105. "Nuclear Fusion Power". World Nuclear Association. May 2007. Archived from the original on 10 March 2008. Retrieved 16 March 2008.
  106. "Chapter 13: Nuclear Energy – Fission and Fusion". Energy Story. California Energy Commission. 2006. Archived from the original on 2 March 2008. Retrieved 14 March 2008.
  107. "DOE Seeks Applicants for Solicitation on the Employment Effects of a Transition to a Hydrogen Economy". Hydrogen Program (Press release). US Department of Energy. 22 March 2006. Archived from the original on 19 July 2011. Retrieved 16 March 2008.
  108. 1 2 "Carbon Capture Strategy Could Lead to Emission-Free Cars" (Press release). Georgia Tech. 11 February 2008. Archived from the original on 28 September 2013. Retrieved 16 March 2008.
  109. Heffel, J. W. (2002). "NOx emission and performance data for a hydrogen fueled internal combustion engine at 1500 rpm using exhaust gas recirculation". International Journal of Hydrogen Energy. 28 (8): 901–908. doi:10.1016/S0360-3199(02)00157-X.
  110. Romm, J. J. (2004). The Hype About Hydrogen: Fact And Fiction in the Race To Save The Climate (1st ed.). Island Press. ISBN   978-1-55963-703-9.
  111. Garbak, John (2011). "VIII.0 Technology Validation Sub-Program Overview" (PDF). DOE Fuel Cell Technologies Program, FY 2010 Annual Progress Report. Archived (PDF) from the original on 24 September 2015. Retrieved 20 May 2015.
  112. Le Comber, P. G.; Jones, D. I.; Spear, W. E. (1977). "Hall effect and impurity conduction in substitutionally doped amorphous silicon". Philosophical Magazine. 35 (5): 1173–1187. Bibcode:1977PMag...35.1173C. doi:10.1080/14786437708232943.
  113. Van de Walle, C. G. (2000). "Hydrogen as a cause of doping in zinc oxide" (PDF). Physical Review Letters. 85 (5): 1012–1015. Bibcode:2000PhRvL..85.1012V. doi:10.1103/PhysRevLett.85.1012. hdl:11858/00-001M-0000-0026-D0E6-E. PMID   10991462. Archived (PDF) from the original on 15 August 2017. Retrieved 1 August 2018.
  114. Janotti, A.; Van De Walle, C. G. (2007). "Hydrogen multicentre bonds". Nature Materials. 6 (1): 44–47. Bibcode:2007NatMa...6...44J. doi:10.1038/nmat1795. PMID   17143265.
  115. Kilic, C.; Zunger, Alex (2002). "n-type doping of oxides by hydrogen". Applied Physics Letters. 81 (1): 73–75. Bibcode:2002ApPhL..81...73K. doi:10.1063/1.1482783.
  116. Peacock, P. W.; Robertson, J. (2003). "Behavior of hydrogen in high dielectric constant oxide gate insulators". Applied Physics Letters. 83 (10): 2025–2027. Bibcode:2003ApPhL..83.2025P. doi:10.1063/1.1609245.
  117. Durgutlu, A. (2003). "Experimental investigation of the effect of hydrogen in argon as a shielding gas on TIG welding of austenitic stainless steel". Materials & Design. 25 (1): 19–23. doi:10.1016/j.matdes.2003.07.004.
  118. "Atomic Hydrogen Welding". Specialty Welds. 2007. Archived from the original on 16 July 2011.
  119. Hardy, W. N. (2003). "From H2 to cryogenic H masers to HiTc superconductors: An unlikely but rewarding path". Physica C: Superconductivity. 388–389: 1–6. Bibcode:2003PhyC..388....1H. doi:10.1016/S0921-4534(02)02591-1.
  120. Almqvist, Ebbe (2003). History of industrial gases. New York, N.Y.: Kluwer Academic/Plenum Publishers. pp. 47–56. ISBN   978-0306472770 . Retrieved 20 May 2015.
  121. Block, M. (3 September 2004). Hydrogen as Tracer Gas for Leak Detection. 16th WCNDT 2004. Montreal, Canada: Sensistor Technologies. Archived from the original on 8 January 2009. Retrieved 25 March 2008.
  122. "Report from the Commission on Dietary Food Additive Intake" (PDF). European Union. Archived (PDF) from the original on 16 February 2008. Retrieved 5 February 2008.
  123. Reinsch, J.; Katz, A.; Wean, J.; Aprahamian, G.; MacFarland, J. T. (1980). "The deuterium isotope effect upon the reaction of fatty acyl-CoA dehydrogenase and butyryl-CoA". J. Biol. Chem. 255 (19): 9093–97. PMID   7410413.
  124. Bergeron, K. D. (2004). "The Death of no-dual-use". Bulletin of the Atomic Scientists. 60 (1): 15–17. Bibcode:2004BuAtS..60a..15B. doi:10.2968/060001004. Archived from the original on 19 April 2008. Retrieved 13 April 2008.
  125. Quigg, C. T. (March 1984). "Tritium Warning". Bulletin of the Atomic Scientists. 40 (3): 56–57. doi:10.1080/00963402.1984.11459199.
  126. International Temperature Scale of 1990 (PDF). Procès-Verbaux du Comité International des Poids et Mesures. 1989. pp. T23–T42. Archived (PDF) from the original on 13 April 2008. Retrieved 25 March 2008.
  127. Cammack, R.; Robson, R. L. (2001). Hydrogen as a Fuel: Learning from Nature. Taylor & Francis Ltd. pp. 202–203. ISBN   978-0-415-24242-4.
  128. Rhee, T. S.; Brenninkmeijer, C. A. M.; Röckmann, T. (19 May 2006). "The overwhelming role of soils in the global atmospheric hydrogen cycle" (PDF). Atmospheric Chemistry and Physics. 6 (6): 1611–1625. doi:10.5194/acp-6-1611-2006. Archived (PDF) from the original on 24 August 2019. Retrieved 24 August 2019.
  129. Eisenmann, Alexander; Amann, Anton; Said, Michael; Datta, Bettina; Ledochowski, Maximilian (2008). "Implementation and interpretation of hydrogen breath tests". Journal of Breath Research. 2 (4): 046002. Bibcode:2008JBR.....2d6002E. doi:10.1088/1752-7155/2/4/046002. PMID   21386189.
  130. Kruse, O.; Rupprecht, J.; Bader, K.; Thomas-Hall, S.; Schenk, P. M.; Finazzi, G.; Hankamer, B. (2005). "Improved photobiological H2 production in engineered green algal cells" (PDF). The Journal of Biological Chemistry. 280 (40): 34170–7. doi:10.1074/jbc.M503840200. PMID   16100118.
  131. Smith, Hamilton O.; Xu, Qing (2005). "IV.E.6 Hydrogen from Water in a Novel Recombinant Oxygen-Tolerant Cyanobacteria System" (PDF). FY2005 Progress Report. United States Department of Energy. Archived (PDF) from the original on 29 December 2016. Retrieved 6 August 2016.
  132. Williams, C. (24 February 2006). "Pond life: the future of energy". Science. The Register. Archived from the original on 9 May 2011. Retrieved 24 March 2008.
  133. "MyChem: Chemical" (PDF). Archived (PDF) from the original on 1 October 2018. Retrieved 1 October 2018.
  134. 1 2 Brown, W. J.; et al. (1997). "Safety Standard for Hydrogen and Hydrogen Systems" (PDF). NASA. Archived (PDF) from the original on 1 May 2017. Retrieved 12 July 2017.
  135. "Liquid Hydrogen MSDS" (PDF). Praxair, Inc. September 2004. Archived from the original (PDF) on 27 May 2008. Retrieved 16 April 2008.
  136. "'Bugs' and hydrogen embrittlement". Science News. 128 (3): 41. 20 July 1985. doi:10.2307/3970088. JSTOR   3970088.
  137. Hayes, B. "Union Oil Amine Absorber Tower". TWI. Archived from the original on 20 November 2008. Retrieved 29 January 2010.
  138. Walker, James L.; Waltrip, John S.; Zanker, Adam (1988). "Lactic acid to magnesium supply-demand relationships". In John J. McKetta; William Aaron Cunningham (eds.). Encyclopedia of Chemical Processing and Design. 28. New York: Dekker. p. 186. ISBN   978-0824724788 . Retrieved 20 May 2015.

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