Hydrogen: Difference between revisions

Hydrogen, ₁H
	Purple glow in its plasma state
Hydrogen
Appearance	Colorless gas
Standard atomic weight Ar°(H)
	[1.00784, 1.00811]; 1.0080±0.0002 (abridged);
Hydrogen in the periodic table
	(none) ← hydrogen → helium
Atomic number (Z)	1
Group	group 1: hydrogen and alkali metals
Period	period 1
Block	s-block
Electron configuration	1s1
Electrons per shell	1
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)
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)
Atomic properties
Oxidation states	−1, 0, +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
	Spectral lines of hydrogen
Other properties
Natural occurrence	primordial
Crystal structure	hexagonal (hP4)
Lattice constants	a = 378.97 pm; c = 618.31 pm (at triple point)
Thermal conductivity	0.1805 W/(m⋅K)
Magnetic ordering	diamagnetic
Molar magnetic susceptibility	−3.98×10−6 cm3/mol (298 K)
Speed of sound	1310 m/s (gas, 27 °C)
CAS Number	12385-13-6 ; 1333-74-0 (H2)
History
Discovery	Henry Cavendish (1766)
Named by	Louis-Bernard Guyton de Morveau; Antoine Lavoisier (1787)
Isotopes of hydrogen v; e;
	Category: Hydrogen; view; talk; edit; \| references

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Hydrogen (Template:PronEng^[11]) is the chemical element with atomic number 1. It is represented by the symbol H. At standard temperature and pressure, hydrogen is a colorless, odorless, nonmetallic, tasteless, highly flammable diatomic gas with the molecular formula H₂. With an atomic mass of 1.00794 amu, hydrogen is the lightest element.

Hydrogen is the most abundant of the chemical elements, constituting roughly 75% of the universe's elemental mass.^[12] Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth, and is industrially produced from hydrocarbons such as methane, after which most elemental hydrogen is used "captively" (meaning locally at the production site), with the largest markets about equally divided between fossil fuel upgrading (e.g., hydrocracking) and ammonia production (mostly for the fertilizer market). Hydrogen may be produced from water using the process of electrolysis, but this process is presently significantly more expensive commercially than hydrogen production from natural gas.^[13]

The most common naturally occurring isotope of hydrogen, known as protium, has a single proton and no neutrons. In ionic compounds it can take on either a positive charge (becoming a cation composed of a bare proton) or a negative charge (becoming an anion known as a hydride). Hydrogen can form compounds with most elements and is present in water and most organic compounds. It plays a particularly important role in acid-base chemistry, in which many reactions involve the exchange of protons between soluble molecules. 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.

Chemistry and characteristics

The solubility and characteristics of hydrogen with various metals are very important in metallurgy (as many metals can suffer hydrogen embrittlement^[14]) and in developing safe ways to store it for use as a fuel.^[15] Hydrogen is highly soluble in many compounds composed of rare earth metals and transition metals^[16] and can be dissolved in both crystalline and amorphous metals.^[17] Hydrogen solubility in metals is influenced by local distortions or impurities in the metal crystal lattice.^[18]

Combustion

Hydrogen gas is highly flammable and will burn at concentrations as low as 4% H₂ in air.^[19] The enthalpy of combustion for hydrogen is −286 kJ/mol;^[20] it burns according to the following balanced equation.

2 H₂(g) + O₂(g) → 2 H₂O(l) + 572 kJ (286 kJ/mol)^[21]

When mixed with oxygen across a wide range of proportions, hydrogen explodes upon ignition. Hydrogen burns violently in air. It ignites automatically at a temperature of 560 °C.^[22] Pure hydrogen-oxygen flames burn in the ultraviolet color range and are nearly invisible to the naked eye, as illustrated by the faintness of flame from the main Space Shuttle engines (as opposed to the easily visible flames from the SRBs). Thus it is difficult to visually detect if a hydrogen leak is burning. The explosion of the Hindenburg airship was an infamous case of hydrogen combustion (pictured); the cause is debated, but combustible materials in the ship's skin were responsible for the coloring of the flames.^[23] Another characteristic of hydrogen fires is that the flames tend to ascend rapidly with the gas in air, as illustrated by the Hindenburg flames, causing less damage than hydrocarbon fires. Two-thirds of the Hindenburg passengers survived the fire, and many of the deaths which occurred were from falling or from diesel fuel burns.^[24]

H₂ reacts directly with other oxidizing elements. A violent and spontaneous reaction can occur at room temperature with chlorine and fluorine, forming the corresponding hydrogen halides: hydrogen chloride and hydrogen fluoride.^[25]

Electron energy levels

The ground state energy level of the electron in a hydrogen atom is −13.6 eV, which is equivalent to an ultraviolet photon of roughly 92 nm.^[26]

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 electromagnetic force attracts electrons and protons to one another, while planets and celestial objects are attracted to each other 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.^[27]

A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation or the equivalent Feynman path integral formulation to calculate the probability density of the electron around the proton.^[28]

Elemental molecular forms

There are two different types of diatomic hydrogen molecules that differ by the relative spin of their nuclei.^[29] In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state; in the parahydrogen form the spins are antiparallel and form a singlet. 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".^[30] The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but since 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 physical properties of pure parahydrogen differ slightly from those of the normal form.^[31] The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene.^[32]

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 convert to the para form very slowly.^[33] 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 the hydrogen liquid, leading to loss of the liquefied material. Catalysts for the ortho-para interconversion, such as iron compounds, are used during hydrogen cooling.^[34]

A molecular form called protonated molecular hydrogen, or H₃⁺, is found in the interstellar medium (ISM), where it is generated by ionization of molecular hydrogen from cosmic rays. It has also been observed in the upper atmosphere of the planet Jupiter. This molecule is relatively stable in the environment of outer space due to the low temperature and density. H₃⁺ is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium.^[35]

Compounds

Covalent and organic compounds

While H₂ is not very reactive under standard conditions, it does form compounds with most elements. Millions of hydrocarbons are known, but they are not formed by the direct reaction of elementary hydrogen and carbon (although synthesis gas production followed by the Fischer-Tropsch process to make hydrocarbons comes close to being an exception, as this begins with coal and the elemental hydrogen is generated in situ).^{[citation needed]} Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I); in these compounds hydrogen takes on a partial positive charge.^[36] When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of strong noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules.^[37]^[38] Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.^[39]

Hydrogen forms a vast array of compounds with carbon. Because of their general association with living things, these compounds came to be called organic compounds;^[40] the study of their properties is known as organic chemistry^[41] and their study in the context of living organisms is known as biochemistry.^[42] By some definitions, "organic" compounds are only required to contain carbon (as a classic historical example, urea). However, most of them also contain hydrogen, and since 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.^[40]

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.^[43]

Hydrides

Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. To chemists, the term "hydride" usually implies that the H atom has acquired a negative or anionic character, denoted H⁻. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 with the electrolysis of molten lithium hydride (LiH), that produced a stoichiometric quantity of hydrogen at the anode.^[44] For hydrides other than group I and II metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH₂, which is polymeric. In lithium aluminium hydride, the AlH₄⁻ 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, there are over 100 binary borane hydrides known, but only one binary aluminium hydride.^[45] Binary indium hydride has not yet been identified, although larger complexes exist.^[46]

"Protons" and acids

Oxidation of H₂ formally gives the proton, H⁺. This species is central to discussion of acids, though the term proton is used loosely to refer to positively charged or cationic hydrogen, denoted H⁺. A bare proton H⁺ cannot exist in solution because of its strong tendency to attach itself to atoms or molecules with electrons. To avoid the convenient fiction of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain the hydronium ion (H₃O⁺) organized into clusters to form H₉O₄⁺.^[47] Other oxonium ions are found when water is in solution with other solvents.^[48]

Although exotic on earth, one of the most common ions in the universe is the H₃⁺ ion, known as protonated molecular hydrogen or the triatomic hydrogen cation.^[49]

Isotopes

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

¹H 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.^[52]
²H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. 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 ¹H-NMR spectroscopy.^[53] Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.^[54]
³H 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.^[43] Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests.^[55] It is used in nuclear fusion reactions,^[56] as a tracer in isotope geochemistry,^[57] and specialized in self-powered lighting devices.^[58] Tritium has also been used in chemical and biological labeling experiments as a radiolabel.^[59]

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 names, but such names are no longer used). The symbols D and T (instead of ²H and ³H) are sometimes used for deuterium and tritium, but the corresponding symbol P is already in use for phosphorus and thus is not available for protium.^[60] In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry allows any of D, T, ²H, and ³H to be used, although ²H and ³H are preferred.^[61]

Natural occurrence

Hydrogen is the most abundant element in the universe, making up 75% of normal matter by mass and over 90% by number of atoms.^[62] 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 proton-proton reaction and CNO cycle nuclear fusion.^[63]

Throughout the universe, hydrogen is mostly found in the atomic and plasma states whose properties are quite different from 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.^[64]

Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H₂ (for data see table). 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. Still, hydrogen is the third most abundant element on the Earth's surface.^[65] Most of the Earth's hydrogen is in the form of chemical compounds such as hydrocarbons and water.^[43] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus. Methane is a hydrogen source of increasing importance.^[66]

History

Discovery and use

Hydrogen gas, H₂, was first artificially produced and formally described by T. Von Hohenheim (also known as Paracelsus, 1493–1541) via the mixing of metals with strong acids.^[67] He was unaware that the flammable gas produced by this chemical reaction was a new chemical element. In 1671, Robert Boyle rediscovered and described the reaction between iron filings and dilute acids, which results in the production of hydrogen gas.^[68] In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by identifying the gas from a metal-acid reaction as "inflammable air" and further finding in 1781 that the gas produces water when burned. He is usually given credit for its discovery as an element.^[7]^[8] In 1783, Antoine Lavoisier gave the element the name of hydrogen (from the Greek hydro meaning water and genes meaning creator)^[10] when he and Laplace reproduced Cavendish's finding that water is produced when hydrogen is burned.^[8]

Hydrogen was liquefied for the first time by James Dewar in 1898 by using his invention, the vacuum flask.^[8] He produced solid hydrogen the next year.^[8] Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck.^[7] Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932.^[8] One of the first uses of H₂ was for limelight.^[8]

The first hydrogen-filled balloon was invented by Jacques Charles in 1783.^[8] 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.^[8] 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.^[8] 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.^[8] The incident was broadcast live on radio and filmed. Ignition of leaking hydrogen as widely assumed to be the cause but later investigations pointed to ignition of the aluminumized fabric coating by static electricity. But the damage to hydrogen's reputation as a lifting gas was already done.

Role in quantum theory

Hydrogen emission spectrum.

Because of its relatively 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.^[69] Furthermore, the corresponding simplicity of the hydrogen molecule and the corresponding cation H₂⁺ allowed fuller 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.^[70]

Production

H₂ 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.

Laboratory

In the laboratory, H₂ is usually prepared by the reaction of acids on metals such as zinc.

Zn + 2 H⁺ → Zn²⁺ + H₂

Aluminium produces H₂ upon treatment with acids but also with base:

2 Al + 6 H₂O → 2 Al(OH)₃ + 3 H₂

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 between 80–94%.^[71]

2H₂O(aq) → 2H₂(g) + O₂(g)

In 2007, it was discovered that an alloy of aluminium and gallium in pellet form added to water could be used to generate hydrogen. The process also creates 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, since hydrogen can be produced on-site and does not need to be transported.^[72]

Industrial

Hydrogen can be prepared in several different ways, but economically the most important processes involve removal of hydrogen from hydrocarbons. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas.^[73] At high temperatures (700–1100 °C; 1,300–2,000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H₂.

CH₄ + H₂O → CO + 3 H₂

This reaction is favored at low pressures but is nonetheless conducted at high pressures (20 atm; 600 inHg) since high pressure H₂ is the most marketable product. 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:

CH₄ → C + 2 H₂

Consequently, steam reforming typically employs an excess of H₂O. 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:^[73]

CO + H₂O → CO₂ + H₂

Other important methods for H₂ production include partial oxidation of hydrocarbons:^[74]

2 CH₄ + O₂ → 2 CO + 4 H₂

and the coal reaction, which can serve as a prelude to the shift reaction above:^[73]

C + H₂O → CO + H₂

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.^[75] Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.^[76]

Applications

Large quantities of H₂ are needed in the petroleum and chemical industries. The largest application of H₂ is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H₂ in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H₂ has several other important uses. H₂ 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. H₂ is also used as a reducing agent of metallic ores.^[77]

Apart from its use as a reactant, H₂ has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding.^[78]^[79] 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.^[80] Since H₂ is lighter than air, having a little more than 1/15th of the density of air, it was once widely used as a lifting agent in balloons and airships.^[81]

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.^[82] Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties.^[83]

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.^[8] Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects.^[84] Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs,^[85] as an isotopic label in the biosciences,^[59] and as a radiation source in luminous paints.^[86]

The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale at 13.8033 kelvins.^[87]

Energy carrier

Hydrogen is not an energy source,^[88] except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development.^[89] The Sun's energy comes from nuclear fusion of hydrogen, but this process is difficult to achieve controllably on Earth.^[90] Elemental hydrogen from solar, biological, or electrical sources costs more in energy to make than is obtained by burning it. Hydrogen may be obtained from fossil sources (such as methane) for less energy than required to make it, but these sources are unsustainable, and are also themselves direct energy sources.^[88]

The energy density per unit volume of both liquid hydrogen and 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.^[88] Nevertheless, elemental hydrogen has been widely discussed in the context of energy, as a possible future carrier of energy on an economy-wide scale.^[91] For example, CO₂ sequestration followed by carbon capture and storage could be conducted at the point of H₂ production from fossil fuels.^[92] Hydrogen used in transportation would burn relatively cleanly, with some NOx emissions,^[93] but without carbon emissions.^[92] However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial.^[94]

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.^[95]

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.^[96] Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H₂ gas even in the presence of oxygen.^[97] Efforts have also been undertaken with genetically modified alga in a bioreactor.^[98]

Safety and precautions

Hydrogen poses a number of hazards to human safety, from potential detonations and fires when mixed with air to being an asphyxant in its pure, oxygen-free form.^[99] In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids. Hydrogen dissolves in some metals, and, in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement.^[100] Hydrogen gas leaking into external air may spontaneously ignite. However, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to accidental burns.^[101]

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.^[99]

References

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^ Palmer, David (November 13, 1997). "Hydrogen in the Universe". NASA. Retrieved 2008-02-05. {{cite web}}: Check date values in: |date= (help)
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^ Kirchheim, R. (1988). "Hydrogen in amorphous and nanocrystalline metals". Materials Science and Engineering. 99: 457–462. doi:10.1016/0025-5416(88)90377-1. {{cite journal}}: |access-date= requires |url= (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)
^ 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. {{cite journal}}: |access-date= requires |url= (help)
^ Carcassi, M. N. (June 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. {{cite journal}}: |access-date= requires |url= (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)
^ National Academy of Engineering, National Academy of Sciences (2004). The Hydrogen Economy: Opportunities, Costs,. National Academies Press. pp. p. 240. ISBN 0309091632. {{cite book}}: |pages= has extra text (help)
^ Energy is per mole of the combustible material, Hydrogen. 286 kJ/mol
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^ Dziadecki, John (2005). "Hindenburg Hydrogen Fire". Retrieved 2007-01-16.
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^ Clayton, Donald D. (2003). Handbook of Isotopes in the Cosmos: Hydrogen to Gallium. Cambridge University Press. ISBN 0521823811.
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^ Brown, Kevin (2007). "The Hydrogen Atom". MathPages. Retrieved 2007-12-21.
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^ Milenko, Yu. Ya. (1997). "Natural ortho-para conversion rate in liquid and gaseous hydrogen". Journal of Low Temperature Physics. 107 (1–2): 77–92. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
^ Svadlenak, R. Eldo (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. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
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^ Hibbs, David E. (1999). "A remarkably stable indium trihydride complex: synthesis and characterisation of [InH3{P(C6H11)3}]". Chemical Communications: 185–186. doi:10.1039/a809279f. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
^ Okumura, Anthony M. (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. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
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[1] "Standard Atomic Weights: Hydrogen". CIAAW. 2009.

[CIAAW2021-2] Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.

[3] Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 240. ISBN 978-0123526519.

[Arblaster_2018-4] Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.

[5] 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.

[6] Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 978-0-8493-0464-4.

[Nostrand-7] "Hydrogen". Van Nostrand's Encyclopedia of Chemistry. Wylie-Interscience. 2005. pp. 797–799. ISBN 978-0-471-61525-5. Cite error: The named reference "Nostrand" was defined multiple times with different content (see the help page).

[nbb-8] ^ ^a ^b ^c ^d ^e ^f ^g ^h ⁱ ^j ^k ^l Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 183–191. ISBN 978-0-19-850341-5. Cite error: The named reference "nbb" was defined multiple times with different content (see the help page).

[9] Miśkowiec, Paweł (April 2023). "Name game: The naming history of the chemical elements—part 1—from antiquity till the end of 18th century". Foundations of Chemistry. 25 (1): 29–51. doi:10.1007/s10698-022-09448-5.

[Stwertka-10] Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 16–21. ISBN 978-0-19-508083-4. Cite error: The named reference "Stwertka" was defined multiple times with different content (see the help page).

[11] Hydrogen, entry in the Oxford English Dictionary, prepared by J. A. Simpson and E. S. C. Weiner, vol. 7, second edition, Oxford: Clarendon Press, 1989. ISBN 0-19-861219-2 (vol. 7), ISBN 0-19-861186-2 (set.)

[12] Palmer, David (November 13, 1997). "Hydrogen in the Universe". NASA. Retrieved 2008-02-05. {{cite web}}: Check date values in: |date= (help)

[13] Staff (2007). "Hydrogen Basics — Production". Florida Solar Energy Center. Retrieved 2008-02-05.

[14] Rogers, H. C. (1999). "Hydrogen Embrittlement of Metals". Science. 159 (3819): 1057–1064. doi:10.1126/science.159.3819.1057. {{cite journal}}: |access-date= requires |url= (help)

[15] Christensen, C. H. (July 9, 2005). "Making society independent of fossil fuels — Danish researchers reveal new technology". Technical University of Denmark. Retrieved 2008-03-28. {{cite news}}: Check date values in: |date= (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Takeshita-16] Takeshita, T. (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. {{cite journal}}: |access-date= requires |url= (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Kirchheim1-17] Kirchheim, R. (1988). "Hydrogen in amorphous and nanocrystalline metals". Materials Science and Engineering. 99: 457–462. doi:10.1016/0025-5416(88)90377-1. {{cite journal}}: |access-date= requires |url= (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Kirchheim2-18] 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. {{cite journal}}: |access-date= requires |url= (help)

[19] Carcassi, M. N. (June 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. {{cite journal}}: |access-date= requires |url= (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)

[20] National Academy of Engineering, National Academy of Sciences (2004). The Hydrogen Economy: Opportunities, Costs,. National Academies Press. pp. p. 240. ISBN 0309091632. {{cite book}}: |pages= has extra text (help)

[21] Energy is per mole of the combustible material, Hydrogen. 286 kJ/mol

[22] Staff (September 10, 2005). "Safety data for hydrogen". Chemical and Other Safety Information. The Physical and Theoretical Chemistry Laboratory, Oxford University. Retrieved 2008-02-05. {{cite web}}: Check date values in: |date= (help)

[23] Dziadecki, John (2005). "Hindenburg Hydrogen Fire". Retrieved 2007-01-16.

[24] Werthmüller, Andreas. "The Hindenburg Disaster". Swiss Hydrogen Association. Retrieved 2008-02-05.

[25] Clayton, Donald D. (2003). Handbook of Isotopes in the Cosmos: Hydrogen to Gallium. Cambridge University Press. ISBN 0521823811.

[26] Millar, Tom (December 10, 2003). "Lecture 7, Emission Lines — Examples". PH-3009 (P507/P706/M324) Interstellar Physics. University of Manchester. Retrieved 2008-02-05. {{cite web}}: Check date values in: |date= (help)

[27] Stern, David P. (May 16, 2005). "The Atomic Nucleus and Bohr's Early Model of the Atom". NASA Goddard Space Flight Center. Retrieved 2007-12-20. {{cite web}}: Check date values in: |date= (help)

[28] Brown, Kevin (2007). "The Hydrogen Atom". MathPages. Retrieved 2007-12-21.

[uigi-29] Staff (2003). "Hydrogen (H₂) Properties, Uses, Applications: Hydrogen Gas and Liquid Hydrogen". Universal Industrial Gases, Inc. Retrieved 2008-02-05.

[Tikhonov-30] Tikhonov, Vladimir I. (2002). "Separation of Water into Its Ortho and Para Isomers". Science. 296 (5577): 2363. doi:10.1126/science.1069513. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[NASA-31] Hritz, James (March 2006). "CH. 6 - Hydrogen" (PDF). NASA Glenn Research Center Glenn Safety Manual, Document GRC-MQSA.001. NASA. Retrieved 2008-02-05.

[32] Shinitzky, Meir; Elitzur, Avshalom C. (2006-05-30). "Ortho-para spin isomers of the protons in the methylene group". Chirality. 18 (9). Rehovot, Israel: Weizmann Institute of Science: 754–756. doi:10.1002/chir.20319. Retrieved 2008-03-25.

[33] Milenko, Yu. Ya. (1997). "Natural ortho-para conversion rate in liquid and gaseous hydrogen". Journal of Low Temperature Physics. 107 (1–2): 77–92. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Svadlenak-34] Svadlenak, R. Eldo (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. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[35] McCall Group, Oka Group (April 22, 2005). "H3+ Resource Center". Universities of Illinois and Chicago. Retrieved 2008-02-05. {{cite web}}: Check date values in: |date= (help)

[36] Clark, Jim (2002). "The Acidity of the Hydrogen Halides". Chemguide. Retrieved 2008-03-09.

[37] Kimball, John W. (2003-08-07). "Hydrogen". Kimball's Biology Pages. Retrieved 2008-03-04.

[38] IUPAC Compendium of Chemical Terminology, Electronic version, Hydrogen Bond

[39] Sandrock, Gary (2002-05-02). "Metal-Hydrogen Systems". Sandia National Laboratories. Retrieved 2008-03-23.

[hydrocarbon-40] "Structure and Nomenclature of Hydrocarbons". Purdue University. Retrieved 2008-03-23.

[41] "Organic Chemistry". Dictionary.com. Lexico Publishing Group. 2008. Retrieved 2008-03-23.

[42] "Biochemistry". Dictionary.com. Lexico Publishing Group. 2008. Retrieved 2008-03-23.

[Miessler-43] Miessler, Gary L. (2003). Inorganic Chemistry (3rd edition ed.). Prentice Hall. ISBN 0130354716. {{cite book}}: |edition= has extra text (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Moers-44] Moers, Kurt (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-45] Downs, Anthony J. (1994). "The hydrides of aluminium, gallium, indium, and thallium: a re-evaluation". Chemical Society Reviews. 23: 175–184. doi:10.1039/CS9942300175. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Hibbs-46] Hibbs, David E. (1999). "A remarkably stable indium trihydride complex: synthesis and characterisation of [InH3{P(C6H11)3}]". Chemical Communications: 185–186. doi:10.1039/a809279f. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Okumura-47] Okumura, Anthony M. (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. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Perdoncin-48] Perdoncin, Giulio (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. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Carrington-49] Carrington, Alan (1989). "The infrared predissociation spectrum of triatomic hydrogen cation (H3+)". Accounts of Chemical Research. 22 (6): 218–222. doi:10.1021/ar00162a004. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

[Gurov-50] Gurov, Yu. B. (2004). "Spectroscopy of superheavy hydrogen isotopes in stopped-pion absorption by nuclei". Physics of Atomic Nuclei. 68 (3): 491–497. doi:10.1134/1.1891200. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

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v t e Diatomic chemical elements
Common	H₂ N₂ O₂ F₂ Cl₂ Br₂ I₂
Other	He₂ Ar₂ C₂ P₂ S₂ Li₂ Rb₂