Hydrogen(Redirected from Molecular hydrogen)
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of circa , hydrogen is the lightest element on the 1.008periodic table. Its monatomic form (H) is the most abundant chemical substance in the Universe, constituting roughly 75% of all baryonic mass.[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.
Purple glow in its plasma state
|Standard atomic weight (Ar, std)||[84, 1.00711] conventional: 1.0081.008|
|Hydrogen in the periodic table|
|Atomic number (Z)||1|
|Group, period||group 1, period 1|
|Element category||diatomic nonmetal|
Electrons per shell
|Phase (at STP)||gas|
|Melting point||13.99 K (−259.16 °C, −434.49 °F)|
|Boiling point||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)|
|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)|
|Oxidation states||−1, +1 (an amphoteric oxide)|
|Electronegativity||Pauling scale: 2.20|
|Covalent radius||31±5 pm|
|Van der Waals radius||120 pm|
|Speed of sound||1310 m/s (gas, 27 °C)|
|Thermal conductivity||0.1805 W/(m·K)|
|Magnetic susceptibility||−3.98·10−6 cm3/mol (298 K)|
|Discovery||Henry Cavendish (1766)|
|Named by||Antoine Lavoisier (1783)|
|Main isotopes of hydrogen|
The universal emergence of atomic hydrogen first occurred during the recombination epoch. 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, 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, 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. 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 a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas (dihydrogen or molecular hydrogen) is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol:
- 2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ (286 kJ/mol)[note 2]
Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% 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). 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.
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 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are also potentially dangerous acids.
Electron energy levels
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.
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. 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
There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1 ( 1⁄2+ 1⁄2); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 ( 1⁄2– 1⁄2). 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". 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. 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.
The uncatalyzed interconversion between para and ortho H2 increases with increasing temperature; thus rapidly condensed H2 contains large quantities of the high-energy ortho form that converts to the para form very slowly. 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 compounds, are used during hydrogen cooling.
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 (e.g., F, Cl, Br, I), or oxygen; in these compounds hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of medium-strength noncovalent bonding with the hydrogen of other similar molecules, a phenomenon called hydrogen bonding that is critical to the stability of many biological molecules. 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.
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. The study of their properties is known as organic chemistry and their study in the context of living organisms is known as biochemistry. 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 which 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. Millions of hydrocarbons are known, and they are usually formed by complicated synthetic pathways that seldom involve elementary hydrogen.
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 stoichiometry quantity of hydrogen at the anode. 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
2, which is polymeric. In lithium aluminium hydride, the AlH−
4 anion carries hydridic centers firmly attached to the Al(III).
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. Binary indium hydride has not yet been identified, although larger complexes exist.
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.
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 Bronsted-Lowry 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
4. Other oxonium ions are found when water is in acidic solution with other solvents.
H 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 descriptive but rarely used formal name protium.
H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in the nucleus. All deuterium in the universe is thought to have been produced at the time of the Big Bang, and has endured since that time. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1
H-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.
H is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into helium-3 through beta decay with a half-life of 12.32 years. It is so radioactive that it can be used in luminous paint, making it useful in such things as watches. The glass prevents the small amount of radiation from getting out. Small amounts of tritium are produced naturally by the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests. It is used in nuclear fusion reactions, as a tracer in isotope geochemistry, and in specialized self-powered lighting devices. Tritium has also been used in chemical and biological labeling experiments as a radiolabel.
Hydrogen is the only element that has different names for 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
H and 3
H) 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. In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry (IUPAC) allows any of D, T, 2
H, and 3
H to be used, although 2
H and 3
H are preferred.
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. Muonium was discovered in 1960. During the muon's µs lifetime, muonium can enter into compounds such as muonium chloride (MuCl) or sodium muonide (NaMu), analogous to 2.2 hydrogen chloride and sodium hydride respectively.
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. 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" 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. In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek ὑδρο- hydro meaning "water" and -γενής genes meaning "creator") when he and Laplace reproduced Cavendish's finding that water is produced when hydrogen is burned.
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. He produced solid hydrogen the next year. Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck. Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932. 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.
The first hydrogen-filled balloon was invented by Jacques Charles in 1783. 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. 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. 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. 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.; because of the thermal conductivity of hydrogen gas, this is the most common type in its field today.
The nickel hydrogen battery was used for the first time in 1977 aboard the U.S. Navy's Navigation technology satellite-2 (NTS-2). For example, the ISS, Mars Odyssey and the Mars Global Surveyor 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, more than 19 years after launch and 13 years beyond their design life.
Role in quantum theory
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. Furthermore, study of the corresponding simplicity of the hydrogen molecule and the corresponding cation H+
2 brought understanding of the nature of the chemical bond, which followed shortly after the quantum mechanical treatment of the hydrogen atom had been developed in the mid-1920s.
One of the first quantum effects to be explicitly noticed (but not understood at the time) was a Maxwell observation involving hydrogen, half a century before full quantum mechanical theory arrived. Maxwell observed that the specific heat capacity of H2 unaccountably departs from that of a diatomic gas below room temperature and begins to increasingly resemble that of a monatomic gas at cryogenic temperatures. According to quantum theory, this behavior arises from the spacing of the (quantized) rotational energy levels, which are particularly wide-spaced in H2 because of its low mass. These widely spaced levels inhibit equal partition of heat energy into rotational motion in hydrogen at low temperatures. Diatomic gases composed of heavier atoms do not have such widely spaced levels and do not exhibit the same effect.
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.) 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 and the CNO cycle nuclear fusion.
Throughout the universe, hydrogen is mostly found in the atomic and plasma states, with properties quite different 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.
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, mostly in the form of chemical compounds such as hydrocarbons and water. 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.
A molecular form called protonated molecular hydrogen (H+
3) is found in the interstellar medium, where it is generated by ionization of molecular hydrogen from cosmic rays. This charged 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+
3 is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium. Neutral triatomic hydrogen H3 can exist only in an excited form and is unstable. By contrast, the positive hydrogen molecular ion (H+
2) is a rare molecule in the universe.
2 is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions.
Hydrogen can be prepared in several different ways, but economically the most important processes involve removal of hydrogen from hydrocarbons, as about 95% of hydrogen production came from steam reforming around year 2000. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas. 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
4 + H
2O → 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
2 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:
4 → C + 2 H
Consequently, steam reforming typically employs an excess of H
2O. 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:
- CO + H
2O → CO
2 + H
Other important methods for H
2 production include partial oxidation of hydrocarbons:
- 2 CH
4 + O
2 → 2 CO + 4 H
and the coal reaction, which can serve as a prelude to the shift reaction above:
- C + H
2O → 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. Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.
- Zn + 2 H+
Aluminium can also produce H
2 upon treatment with bases:
- 2 Al + 6 H
2O + 2 OH−
→ 2 Al(OH)−
4 + 3 H
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 80–94%.
- 2 H
2O(l) → 2 H
2(g) + O
An alloy of aluminium and gallium in pellet form added to water can be used to generate hydrogen. The process also produces alumina, but the expensive gallium, which prevents the formation of an oxide skin on the pellets, can be re-used. This has important potential implications for a hydrogen economy, as hydrogen can be produced on-site and does not need to be transported.
There are more than 200 thermochemical cycles which can be used for water splitting, around a dozen 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 are under research and in testing phase to produce hydrogen and oxygen from water and heat without using electricity. 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.
Under anaerobic conditions, iron and steel alloys are slowly oxidized by the protons of water concomitantly reduced in molecular hydrogen (H
2). The anaerobic corrosion of iron leads first to the formation of ferrous hydroxide (green rust) and can be described by the following reaction:
- Fe + 2 H
2O → Fe(OH)
2 + H
In its turn, under anaerobic conditions, the ferrous hydroxide (Fe(OH)
2 ) can be oxidized by the protons of water to form magnetite and molecular hydrogen. This process is described by the Schikorr reaction:
- 3 Fe(OH)
2 → Fe
4 + 2 H
2O + H
- ferrous hydroxide → magnetite + water + hydrogen
The well crystallized magnetite (Fe
4) is thermodynamically more stable than the ferrous hydroxide (Fe(OH)
Geological occurrence: the serpentinization reaction
In the absence of atmospheric oxygen (O
2), in deep geological conditions prevailing far away from Earth atmosphere, hydrogen (H
2) is produced during the process of serpentinization by the anaerobic oxidation by the water protons (H+) of the ferrous (Fe2+) silicate present in the crystal lattice of the fayalite (Fe
4, the olivine iron-endmember). The corresponding reaction leading to the formation of magnetite (Fe
4), quartz (SiO
2) and hydrogen (H
2) is the following:
4 + 2 H
2O → 2 Fe
4 + 3 SiO
2 + 3 H
- fayalite + water → magnetite + quartz + hydrogen
Formation in transformers
From all the fault gases formed in power transformers, hydrogen is the most common and is generated under most fault conditions; thus, formation of hydrogen is an early indication of serious problems in the transformer's life cycle.
Consumption in processes
Large quantities of H
2 are needed in the petroleum and chemical industries. The largest application of H
2 is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H
2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H
2 has several other important uses. H
2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H
2 is also used as a reducing agent of metallic ores.
Hydrogen is highly soluble in many rare earth and transition metals and is soluble in both nanocrystalline and amorphous metals. Hydrogen solubility in metals is influenced by local distortions or impurities in the crystal lattice. 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, complicating the design of pipelines and storage tanks.
Apart from its use as a reactant, H
2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding. 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. Because H
2 is lighter than air, having a little more than 1⁄14 of the density of air, it was once widely used as a lifting gas in balloons and airships.
In more recent applications, hydrogen is used pure or mixed with nitrogen (sometimes called forming gas) as a tracer gas for minute leak detection. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries. Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties.
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. Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects. Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, as an isotopic label in the biosciences, and as a radiation source in luminous paints.
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.
Hydrogen is not an energy resource, except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development. The Sun's energy comes from nuclear fusion of hydrogen, but this process is difficult to achieve controllably on Earth. 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.
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. Nevertheless, elemental hydrogen has been widely discussed in the context of energy, as a possible future carrier of energy on an economy-wide scale. For example, CO
2 sequestration followed by carbon capture and storage could be conducted at the point of H
2 production from fossil fuels. Hydrogen used in transportation would burn relatively cleanly, with some NOx emissions, but without carbon emissions. However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial. Fuel cells can convert hydrogen and oxygen directly to electricity more efficiently than internal combustion engines.
Hydrogen is employed to saturate broken ("dangling") bonds of amorphous silicon and amorphous carbon that helps stabilizing material properties. It is also a potential electron donor in various oxide materials, including ZnO, SnO2, CdO, MgO, ZrO2, HfO2, La2O3, Y2O3, TiO2, SrTiO3, LaAlO3, SiO2, Al2O3, ZrSiO4, HfSiO4, and SrZrO3.
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. The natural cycle of hydrogen production and consumption by organisms is called the hydrogen cycle.
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. Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H2 gas even in the presence of oxygen. Efforts have also been undertaken with genetically modified alga in a bioreactor.
Safety and precautions
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. In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids. Hydrogen dissolves in many metals, and, in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement, leading to cracks and explosions. 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.
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.
- Simpson, J. A.; Weiner, E. S. C. (1989). "Hydrogen". Oxford English Dictionary. 7 (2nd ed.). Clarendon Press. ISBN 0-19-861219-2.
- Meija, J.; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure Appl. Chem. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
- Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 240. ISBN 0123526515.
- 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 0-8493-0486-5.
- Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
- "Hydrogen". Van Nostrand's Encyclopedia of Chemistry. Wylie-Interscience. 2005. pp. 797–799. ISBN 0-471-61525-0.
- Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 183–191. ISBN 0-19-850341-5.
- Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 16–21. ISBN 0-19-508083-1.
- Palmer, D. (13 September 1997). "Hydrogen in the Universe". NASA. Retrieved 5 February 2008.
- 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. Retrieved 4 May 2015.
- Presenter: Professor Jim Al-Khalili (21 January 2010). "Discovering the Elements". Chemistry: A Volatile History. 25:40 minutes in. BBC. BBC Four.
- "Hydrogen Basics – Production". Florida Solar Energy Center. 2007. Retrieved 5 February 2008.
- Rogers, H. C. (1999). "Hydrogen Embrittlement of Metals". Science. 159 (3819): 1057–1064. Bibcode:1968Sci...159.1057R. PMID 17775040. doi:10.1126/science.159.3819.1057.
- 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. Retrieved 19 May 2015.
- "Dihydrogen". O=CHem Directory. University of Southern Maine. Archived from the original on 13 February 2009. Retrieved 6 April 2009.
- 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.
- 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 0-309-09163-2.
- Patnaik, P. (2007). A Comprehensive Guide to the Hazardous Properties of Chemical Substances. Wiley-Interscience. p. 402. ISBN 0-471-71458-5.
- 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.
- Clayton, D. D. (2003). Handbook of Isotopes in the Cosmos: Hydrogen to Gallium. Cambridge University Press. ISBN 0-521-82381-1.
- NAAP Labs (2009). "Energy Levels". University of Nebraska Lincoln. Retrieved 20 May 2015.
- "photon wavelength 13.6 eV". Wolfram Alpha. 20 May 2015. Retrieved 20 May 2015.
- Stern, D. P. (16 May 2005). "The Atomic Nucleus and Bohr's Early Model of the Atom". NASA Goddard Space Flight Center (mirror). Retrieved 20 December 2007.
- Stern, D. P. (13 February 2005). "Wave Mechanics". NASA Goddard Space Flight Center. Retrieved 16 April 2008.
- Staff (2003). "Hydrogen (H2) Properties, Uses, Applications: Hydrogen Gas and Liquid Hydrogen". Universal Industrial Gases, Inc. Retrieved 5 February 2008.
- Tikhonov, V. I.; Volkov, A. A. (2002). "Separation of Water into Its Ortho and Para Isomers". Science. 296 (5577): 2363. PMID 12089435. doi:10.1126/science.1069513.
- Hritz, J. (March 2006). "CH. 6 – Hydrogen" (PDF). NASA Glenn Research Center Glenn Safety Manual, Document GRC-MQSA.001. NASA. Retrieved 5 February 2008.
- Shinitzky, M.; Elitzur, A. C. (2006). "Ortho-para spin isomers of the protons in the methylene group". Chirality. 18 (9): 754–756. PMID 16856167. doi:10.1002/chir.20319.
- 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.
- Amos, Wade A. (1 November 1998). "Costs of Storing and Transporting Hydrogen" (PDF). National Renewable Energy Laboratory. pp. 6–9. Retrieved 19 May 2015.
- 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.
- Clark, J. (2002). "The Acidity of the Hydrogen Halides". Chemguide. Retrieved 9 March 2008.
- Kimball, J. W. (7 August 2003). "Hydrogen". Kimball's Biology Pages. Retrieved 4 March 2008.
- IUPAC Compendium of Chemical Terminology, Electronic version, Hydrogen Bond
- Sandrock, G. (2 May 2002). "Metal-Hydrogen Systems". Sandia National Laboratories. Archived from the original on 24 February 2008. Retrieved 23 March 2008.
- "Structure and Nomenclature of Hydrocarbons". Purdue University. Retrieved 23 March 2008.
- "Organic Chemistry". Dictionary.com. Lexico Publishing Group. 2008. Retrieved 23 March 2008.
- "Biochemistry". Dictionary.com. Lexico Publishing Group. 2008. Retrieved 23 March 2008.
- 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.
- 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.
- 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.
- Miessler, G. L.; Tarr, D. A. (2003). Inorganic Chemistry (3rd ed.). Prentice Hall. ISBN 0-13-035471-6.
- 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.
- 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.
- 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.
- 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.
- 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.
- Urey, H. C.; Brickwedde, F. G.; Murphy, G. M. (1933). "Names for the Hydrogen Isotopes". Science. 78 (2035): 602–603. Bibcode:1933Sci....78..602U. PMID 17797765. doi:10.1126/science.78.2035.602.
- 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. PMID 1330130. doi:10.1007/BF01875525.
- Broad, W. J. (11 November 1991). "Breakthrough in Nuclear Fusion Offers Hope for Power of Future". The New York Times. Retrieved 12 February 2008.
- Traub, R. J.; Jensen, J. A. (June 1995). "Tritium radioluminescent devices, Health and Safety Manual" (PDF). International Atomic Energy Agency. p. 2.4. Retrieved 20 May 2015.
- Staff (15 November 2007). "Tritium". U.S. Environmental Protection Agency. Retrieved 12 February 2008.
- Nave, C. R. (2006). "Deuterium-Tritium Fusion". HyperPhysics. Georgia State University. Retrieved 8 March 2008.
- Kendall, C.; Caldwell, E. (1998). "Fundamentals of Isotope Geochemistry". US Geological Survey. Retrieved 8 March 2008.
- "The Tritium Laboratory". University of Miami. 2008. Retrieved 8 March 2008.
- 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. Lubbock: Texas Tech University. 25 (2): 97–107. doi:10.1023/A:1010655610575.
- van der Krogt, P. (5 May 2005). "Hydrogen". Elementymology & Elements Multidict. Retrieved 20 December 2010.
- § IR-3.3.2, Provisional Recommendations, Nomenclature of Inorganic Chemistry, Chemical Nomenclature and Structure Representation Division, IUPAC. Accessed on line 3 October 2007.
- IUPAC (1997). "Muonium". In A.D. McNaught, A. Wilkinson. Compendium of Chemical Terminology (2nd ed.). Blackwell Scientific Publications. ISBN 0-86542-684-8. doi:10.1351/goldbook.M04069.
- 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.
- 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.
- Boyle, R. (1672). "Tracts written by the Honourable Robert Boyle containing new experiments, touching the relation betwixt flame and air..." London.
- Winter, M. (2007). "Hydrogen: historical information". WebElements Ltd. Retrieved 5 February 2008.
- Musgrave, A. (1976). "Why did oxygen supplant phlogiston? Research programmes in the Chemical Revolution". In Howson, C. Method and appraisal in the physical sciences. The Critical Background to Modern Science, 1800–1905. Cambridge University Press. Retrieved 22 October 2011.
- Cavendish, Henry (12 May 1766). "Three Papers, Containing Experiments on Factitious Air, by the Hon. Henry Cavendish, F. R. S.". Philosophical Transactions. The Royal Society. 56: 141–184. Bibcode:1766RSPT...56..141C. JSTOR 105491.
- National Electrical Manufacturers Association (1946). A chronological history of electrical development from 600 B.C. p. 102.
- "NTS-2 Nickel-Hydrogen Battery Performance 31". Aiaa.org. Retrieved 6 April 2009.
- 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. ISBN 0-7803-7296-4. doi:10.1109/IECEC.2002.1391972. Retrieved 11 November 2011.
- Anderson, P. M.; Coyne, J. W. (2002). "A lightweight high reliability single battery power system for interplanetary spacecraft". Aerospace Conference Proceedings. 5: 5–2433. ISBN 0-7803-7231-X. doi:10.1109/AERO.2002.1035418.
- "Mars Global Surveyor". Astronautix.com. Archived from the original on 10 August 2009. Retrieved 6 April 2009.
- Lori Tyahla, ed. (7 May 2009). "Hubble servicing mission 4 essentials". NASA. Retrieved 19 May 2015.
- Hendrix, Susan (25 November 2008). Lori Tyahla, ed. "Extending Hubble's mission life with new batteries". NASA. Retrieved 19 May 2015.
- Crepeau, R. (1 January 2006). Niels Bohr: The Atomic Model. Great Scientific Minds. Great Neck Publishing. ISBN 1-4298-0723-7.
- 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.
- Charlton, Mike; Van Der Werf, Dirk Peter (1 March 2015). "Advances in antihydrogen physics". Science Progress. 98 (1): 34–62. doi:10.3184/003685015X14234978376369.
- Kellerbauer, Alban (29 January 2015). "Why Antimatter Matters". European Review. 23 (01): 45–56. doi:10.1017/S1062798714000532.
- Gagnon, S. "Hydrogen". Jefferson Lab. Retrieved 5 February 2008.
- Haubold, H.; Mathai, A. M. (15 November 2007). "Solar Thermonuclear Energy Generation". Columbia University. Archived from the original on 2011-12-11. Retrieved 12 February 2008.
- 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. Bibcode:2000ApJ...543..552S. arXiv: . doi:10.1086/317138.
- Dresselhaus, M.; et al. (15 May 2003). "Basic Research Needs for the Hydrogen Economy" (PDF). Argonne National Laboratory, U.S. Department of Energy, Office of Science Laboratory. Archived from the original (PDF) on 13 February 2008. Retrieved 5 February 2008.
- Berger, W. H. (15 November 2007). "The Future of Methane". University of California, San Diego. Retrieved 12 February 2008.
- McCall Group; Oka Group (22 April 2005). "H3+ Resource Center". Universities of Illinois and Chicago. Retrieved 5 February 2008.
- Helm, H.; et al. "Coupling of Bound States to Continuum States in Neutral Triatomic Hydrogen" (PDF). Department of Molecular and Optical Physics, University of Freiburg, Germany. Retrieved 25 November 2009.
- 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.
- Oxtoby, D. W. (2002). Principles of Modern Chemistry (5th ed.). Thomson Brooks/Cole. ISBN 0-03-035373-4.
- "Hydrogen Properties, Uses, Applications". Universal Industrial Gases, Inc. 2007. Retrieved 11 March 2008.
- 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.
- Lees, A. (2007). "Chemicals from salt". BBC. Archived from the original on 26 October 2007. Retrieved 11 March 2008.
- 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.
- Venere, E. (15 May 2007). "New process generates hydrogen from aluminum alloy to run engines, fuel cells". Purdue University. Retrieved 5 February 2008.
- Weimer, Al (25 May 2005). "Development of solar-powered thermochemical production of hydrogen from water" (PDF). Solar Thermochemical Hydrogen Generation Project.
- Perret, R. "Development of Solar-Powered Thermochemical Production of Hydrogen from Water, DOE Hydrogen Program, 2007" (PDF). Retrieved 17 May 2008.
- Hirschler, M. M. (2000). Electrical Insulating Materials: International Issues. ASTM International. pp. 89–. ISBN 978-0-8031-2613-8. Retrieved 13 July 2012.
- Chemistry Operations (15 December 2003). "Hydrogen". Los Alamos National Laboratory. Retrieved 5 February 2008.
- 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.
- Kirchheim, R.; Mutschele, T.; Kieninger, W.; Gleiter, H.; Birringer, R.; Koble, T. (1988). "Hydrogen in amorphous and nanocrystalline metals". Materials Science and Engineering. 99: 457–462. doi:10.1016/0025-5416(88)90377-1.
- 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.
- 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.
- "Atomic Hydrogen Welding". Specialty Welds. 2007. Archived from the original on 16 July 2011.
- 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.
- Almqvist, Ebbe (2003). History of industrial gases. New York, N.Y.: Kluwer Academic/Plenum Publishers. pp. 47–56. ISBN 0306472775. Retrieved 20 May 2015.
- Block, M. (3 September 2004). Hydrogen as Tracer Gas for Leak Detection. 16th WCNDT 2004. Montreal, Canada: Sensistor Technologies. Retrieved 25 March 2008.
- "Report from the Commission on Dietary Food Additive Intake" (PDF). European Union. Retrieved 5 February 2008.
- 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.
- Bergeron, K. D. (2004). "The Death of no-dual-use". Bulletin of the Atomic Scientists. Educational Foundation for Nuclear Science, Inc. 60 (1): 15. doi:10.2968/060001004.
- Quigg, C. T. (March 1984). "Tritium Warning". Bulletin of the Atomic Scientists. 40 (3): 56–57.
- International Temperature Scale of 1990 (PDF). Procès-Verbaux du Comité International des Poids et Mesures. 1989. pp. T23–T42. Retrieved 25 March 2008.
- McCarthy, J. (31 December 1995). "Hydrogen". Stanford University. Retrieved 14 March 2008.
- "Nuclear Fusion Power". World Nuclear Association. May 2007. Retrieved 16 March 2008.
- "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.
- "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.
- "Carbon Capture Strategy Could Lead to Emission-Free Cars" (Press release). Georgia Tech. 11 February 2008. Retrieved 16 March 2008.
- 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.
- Romm, J. J. (2004). The Hype About Hydrogen: Fact And Fiction in the Race To Save The Climate (1st ed.). Island Press. ISBN 1-55963-703-X.
- Garbak, John (2011). "VIII.0 Technology Validation Sub-Program Overview" (PDF). DOE Fuel Cell Technologies Program, FY 2010 Annual Progress Report. Retrieved 20 May 2015.
- 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.
- Van de Walle, C. G. (2000). "Hydrogen as a cause of doping in zinc oxide". Physical Review Letters. 85 (5): 1012–1015. Bibcode:2000PhRvL..85.1012V. PMID 10991462. doi:10.1103/PhysRevLett.85.1012.
- Janotti, A.; Van De Walle, C. G. (2007). "Hydrogen multicentre bonds". Nature Materials. 6 (1): 44–47. Bibcode:2007NatMa...6...44J. PMID 17143265. doi:10.1038/nmat1795.
- 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.
- 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.
- Cammack, R.; Robson, R. L. (2001). Hydrogen as a Fuel: Learning from Nature. Taylor & Francis Ltd. pp. 202–203. ISBN 0-415-24242-8.
- Rhee, T. S.; Brenninkmeijer, C. A. M.; Röckmann, T. (19 May 2006). "The overwhelming role of soils in the global atmospheric hydrogen cycle". Atmospheric Chemistry and Physics. 6 (6): 1611–1625. doi:10.5194/acp-6-1611-2006. Retrieved 20 May 2015.
- 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". The Journal of Biological Chemistry. 280 (40): 34170–7. PMID 16100118. doi:10.1074/jbc.M503840200.
- 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. Retrieved 6 August 2016.
- Williams, C. (24 February 2006). "Pond life: the future of energy". Science. The Register. Retrieved 24 March 2008.
- Brown, W. J.; et al. (1997). "Safety Standard for Hydrogen and Hydrogen Systems" (PDF). NASA. Retrieved 12 July 2017.
- "Liquid Hydrogen MSDS" (PDF). Praxair, Inc. September 2004. Archived from the original (PDF) on 27 May 2008. Retrieved 16 April 2008.
- "'Bugs' and hydrogen embrittlement". Science News. Washington, D.C. 128 (3): 41. 20 July 1985. JSTOR 3970088. doi:10.2307/3970088.
- Hayes, B. "Union Oil Amine Absorber Tower". TWI. Retrieved 29 January 2010.
- Walker, James L.; Waltrip, John S.; Zanker, Adam (1988). John J. McKetta; William Aaron Cunningham, eds. Lactic acid to magnesium supply-demand relationships. Encyclopedia of Chemical Processing and Design. 28. New York: Dekker. p. 186. ISBN 082472478X. Retrieved 20 May 2015.
- Chart of the Nuclides (17th ed.). Knolls Atomic Power Laboratory. 2010. ISBN 978-0-9843653-0-2.
- Ferreira-Aparicio, P.; Benito, M. J.; Sanz, J. L. (2005). "New Trends in Reforming Technologies: from Hydrogen Industrial Plants to Multifuel Microreformers". Catalysis Reviews. 47 (4): 491–588. doi:10.1080/01614940500364958.
- Newton, David E. (1994). The Chemical Elements. New York: Franklin Watts. ISBN 0-531-12501-7.
- Rigden, John S. (2002). Hydrogen: The Essential Element. Cambridge, Massachusetts: Harvard University Press. ISBN 0-531-12501-7.
- Romm, Joseph, J. (2004). The Hype about Hydrogen, Fact and Fiction in the Race to Save the Climate. Island Press. ISBN 1-55963-703-X.
- Scerri, Eric (2007). The Periodic System, Its Story and Its Significance. New York: Oxford University Press. ISBN 0-19-530573-6.