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Hydrogen

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                 1               (none) ← hydrogen → helium
                 -
                ↑
                H
                ↓
                Li

                                  Periodic Table - Extended Periodic Table

                                                                   General
                                       Name, Symbol, Number hydrogen, H, 1
                                                 Chemical series nonmetals
                                              Group, Period, Block 1, 1, s
                                                      Appearance colorless
                                             Atomic mass 1.00794 (7) g/mol
                                               Electron configuration 1s^1
                                                     Electrons per shell 1
                                                       Physical properties
                                                                 Phase gas
                                              Density (0 °C, 101.325 kPa)
                                                               0.08988 g/L
                                                    Melting point 14.01  K
                                          (−259.14 ° C, −434.45 ° F)
                                                     Boiling point 20.28 K
                                          (−252.87 ° C, −423.17 ° F)
                                         Triple point 13.8033 K, 7.042 kPa
                                         Critical point 32.97 K, 1.293 MPa
                                  Heat of fusion (H[2]) 0.117 kJ·mol^−1
                            Heat of vaporization (H[2]) 0.904 kJ·mol^−1
                                             Heat capacity (25 °C) (H[2])
                                                28.836 J·mol^−1·K^−1

   CAPTION: Vapor pressure

                                             P/Pa  1 10 100 1 k 10 k 100 k
                                            at T/K               15   20

                                                         Atomic properties
                                               Crystal structure hexagonal
                                                  Oxidation states 1, −1
                                                       ( amphoteric oxide)
                                    Electronegativity 2.20 (Pauling scale)
                                    Ionization energies 1st: 1312.0 kJ/mol
                                                       Atomic radius 25 pm
                                Atomic radius (calc.) 53 pm ( Bohr radius)
                                                     Covalent radius 37 pm
                                               Van der Waals radius 120 pm
                                                             Miscellaneous
                     Thermal conductivity (300 K) 180.5 mW·m^−1·K^−1
                                     Speed of sound (gas, 27 °C) 1310 m/s
                                             CAS registry number 1333-74-0
                                                         Selected isotopes

                 CAPTION: Main article: Isotopes of hydrogen

                                  iso   NA    half-life DM  DE ( MeV)  DP
                                  ^1H 99.985% H is stable with 0 neutrons
                                  ^2H 0.0115% H is stable with 1 neutron
                                  ^3H trace   12.32 y   β^− 0.019     ^3He

                                                                References

   Hydrogen ( IPA: /ˈhaɪdrə(ʊ)dʒən/, Latin: 'hydrogenium', from Ancient
   Greek ὕδωρ (hudor): "water" and Ancient Greek γείνομαι (geinomai): "to
   beget or sire") is a chemical element that, in the periodic table, has
   the symbol H and an atomic number of 1. At standard temperature and
   pressure it is a colorless, odorless, nonmetallic, tasteless, highly
   flammable diatomic gas (H[2]). With an atomic mass of 1.00794 g/ mol,
   hydrogen is the lightest element. It is also the most abundant,
   constituting roughly 75% of the universe's elemental mass. 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, after which most free 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 in ammonia production (mostly for the fertilizer
   market). However, hydrogen can easily be produced from water using the
   process of electrolysis.

   The most common naturally occurring isotope of hydrogen 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.

Nomenclature

   The word "hydrogen" has several different meanings:
    1. the name of an element.
    2. an atom, sometimes called "H dot", that is abundant in space but
       essentially absent on earth, because it dimerizes.
    3. a diatomic molecule that occurs naturally in trace amounts in the
       Earth's atmosphere; chemists increasingly refer to H[2] as
       dihydrogen to distinguish this molecule from atomic hydrogen and
       hydrogen found in other compounds.
    4. the atomic constituent within all organic compounds, water, and
       many other chemical compounds.

   The elemental forms of hydrogen should not be confused with hydrogen as
   it appears in chemical compounds.

History

Discovery of H[2]

   Hydrogen gas, H[2], 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. 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. 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 that the
   gas produces water when burned. Cavendish had stumbled on hydrogen when
   experimenting with acids and mercury. Although he wrongly assumed that
   hydrogen was a liberated component of the mercury rather than the acid,
   he was still able to accurately describe several key properties of
   hydrogen. He is usually given credit for its discovery as an element.
   In 1783, Antoine Lavoisier gave the element the name of hydrogen when
   he (with Laplace) reproduced Cavendish's finding that water is produced
   when hydrogen is burned. Lavoisier's name for the gas won out.

   One of the first uses of H[2] was for balloons. The H[2] was obtained
   by reacting sulphuric acid and metallic iron. Infamously, H[2] was used
   in the Hindenburg airship that was destroyed in a midair fire.

Role in history of quantum theory

   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. Furthermore, the
   corresponding simplicity of the hydrogen molecule and the corresponding
   cation H[2]^+ 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 Maxwell's observation, half a century
   before full quantum mechanical theory arrived. He observed that the
   specific heat capacity of H[2] 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 behaviour arises from the spacing of the (quantized)
   rotational energy levels, which are particularly wide-spaced in H[2]
   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.

Natural occurrence

   NGC 604, a giant region of ionized hydrogen in the Triangulum Galaxy
   Enlarge
   NGC 604, a giant region of ionized hydrogen in the Triangulum Galaxy

   Hydrogen is the most abundant element in the universe, making up 75% of
   normal matter by mass and over 90% by number of atoms. This element is
   found in great abundance in stars and gas giant planets. Molecular
   clouds of H[2] are associated with star formation. Hydrogen plays a
   vital role in powering stars through proton-proton reaction nuclear
   fusion.

   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.

   Under ordinary conditions on Earth, elemental hydrogen exists as the
   diatomic gas, H[2] (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. Although H atoms and H[2] molecules are abundant in
   interstellar space, they are difficult to generate, concentrate, and
   purify on Earth. Most of the Earth's hydrogen is 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. Methane is a hydrogen source of increasing importance.

The hydrogen atom

Electron energy levels

   Depiction of a hydrogen-1 atom, or protium, showing the Van der Waals
   radius and the proton nucleus
   Enlarge
   Depiction of a hydrogen-1 atom, or protium, showing the Van der Waals
   radius and the proton nucleus

   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.

   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. A more accurate description of
   the hydrogen atom comes from a purely quantum mechanical treatment that
   uses the Schrödinger equation to calculate the probability density of
   the electron around the proton. Treating the electron as a matter wave
   reproduces chemical results such as shape of the hydrogen atom more
   naturally than the particle-based Bohr model, although the energy and
   spectral results are the same. Modeling the system fully using the
   reduced mass of nucleus and electron (as one would do in the two-body
   problem in celestial mechanics) yields an even better formula for the
   hydrogen spectra, and also the correct spectral shifts for the isotopes
   deuterium and tritium. Very small adjustments in energy levels in the
   hydrogen atom, which correspond to actual spectral effects, may be
   determined by using a full quantum mechanical theory which corrects for
   the effects of special relativity (see Dirac equation), and by
   accounting for quantum effects arising from production of virtual
   particles in the vacuum and as a result of electric fields (see quantum
   electrodynamics).

   In hydrogen gas, the electronic ground state energy level is split into
   hyperfine structure levels because of magnetic effects of the quantum
   mechanical spin of the electron and proton. The energy of the atom when
   the proton and electron spins are aligned is higher than when they are
   not aligned. The transition between these two states can occur through
   emission of a photon through a magnetic dipole transition. Radio
   telescopes can detect the radiation produced in this process, which is
   used to map the distribution of hydrogen in the galaxy.

Isotopes

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

   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.
     * ^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.

     * ^2H, the other stable hydrogen isotope, is known as deuterium and
       contains one proton and one neutron in its nucleus. Deuterium
       comprises 0.0026–0.0184% of all hydrogen on Earth. It 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. Heavy water is used as a
       neutron moderator and coolant for nuclear reactors. Deuterium is
       also a potential fuel for commercial nuclear fusion.

     * ^3H is known as tritium and contains one proton and two neutrons in
       its nucleus. It is radioactive, decays through beta decay with a
       half-life of 12.32 years. 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. It is
       used in nuclear fusion reactions, as a tracer in isotope
       geochemistry, and specialized in self-powered lighting devices.
       Tritium was once routinely used in chemical and biological labeling
       experiments as a radiolabel (this has become less common).

   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. IUPAC states that while this use is common it is not
   preferred.

Elemental molecular forms

   First tracks observed in liquid hydrogen bubble chamber
   Enlarge
   First tracks observed in liquid hydrogen bubble chamber

   There are two different types of diatomic hydrogen 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; 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". 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. The
   ortho/para distinction also occurs in other hydrogen-containing
   molecules or functional groups, such as water and methylene.

   The uncatalyzed interconversion between para and ortho H[2] increases
   with increasing temperature; thus rapidly condensed H[2] contains large
   quantities of the high-energy ortho form that convert to the para form
   very slowly. The ortho/para ratio in condensed H[2] 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.

Chemical and physical properties

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

Combustion

   Hydrogen can combust rapidly in air, and was blamed for the disaster
   with Hindenburg on May 6, 1937
   Enlarge
   Hydrogen can combust rapidly in air, and was blamed for the disaster
   with Hindenburg on May 6, 1937

   Hydrogen gas is highly flammable and will burn at concentrations as low
   as 4% H[2] in air. The enthalpy of combustion for hydrogen is –286
   kJ/mol; it combusts according to the following balanced equation.

          2 H[2](g) + O[2](g) → 2 H[2]O(l) + 572 kJ

   When mixed with oxygen across a wide range of proportions, hydrogen
   explodes upon ignition. Hydrogen burns violently in air.
   Hydrogen-oxygen flames 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 shuttle
   boosters). Thus it is difficult to visually detect if a hydrogen leak
   is burning. The Hindenburg zeppelin flames seen in the adjacent picture
   are from the covering skin of the zeppelin which contained carbon and
   pyrophoric aluminium powder that may have started the fire. Another
   characteristic of hydrogen fires is that the flames tend to ascend
   rapidly with the gas in air, causing less damage than hydrocarbon
   fires. Two-thirds of the Hindenburg passengers survived and deaths were
   from falling or from gasoline burns.

   H[2] 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.

Compounds

Covalent and organic compounds

   While H[2] 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. Hydrogen can form compounds with elements that are more
   electronegative, such as halogens (e.g., F, Cl, Br, I) and chalcogens
   (O, S, Se); in these compounds hydrogen takes on a partial positive
   charge. 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. 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.

   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; 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 (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. (This latter
   definition is not perfect, however, as in this definition urea would
   not be included as an organic compound).

   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.

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 G.N. Lewis in 1916 for
   group I and II salt-like hydrides, was demonstated by Moers in 1920
   with the electrolysis of molten lithium hydride (LiH), that produced a
   stoichiometric quantity of hydrogen at the anode. 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 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, there are over 100 binary borane hydrides known,
   but only one binary aluminium hydride. Binary indium hydride has not
   yet been identified, although larger complexes exist.

"Protons" and acids

   Oxidation of H[2] 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
   (H[3]O^+) organized into clusters to form H[9]O[4]^+. Other oxonium
   ions are found when water is in solution with other solvents.

   Although exotic on earth, one of the most common ions in the universe
   is the H[3]^+ ion, known as protonated molecular hydrogen or the
   triatomic hydrogen cation.

Production

   H[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.

Laboratory syntheses

   In the laboratory, H[2] is usually prepared by the reaction of acids on
   metals such as zinc.

          Zn + 2 H^+ → Zn^2+ + H[2]

   Aluminium produces H[2] upon treatment with acids but also with base:

          2 Al + 6 H[2]O → 2 Al(OH)[3] + 3 H[2]

   The electrolysis of water is a simple method of producing hydrogen,
   although the resulting hydrogen necessarily has less energy content
   than was required to produce it. 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%. Bellona
   Report on Hydrogen

          2H[2]O(aq) → 2H[2](g) + O[2](g)

Industrial syntheses

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

          CH[4] + H[2]O → CO + 3 H[2]

   This reaction is favored at low pressures but is nonetheless conducted
   at high pressures (20 atm; 600  inHg) since high pressure H[2] 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:

          CH[4] → C + 2 H[2]

   Consequently, steam reforming typically employs an excess of H[2]O.

   Additional hydrogen from steam reforming can be recovered from the
   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[2]O → CO[2] + H[2]

   Other important methods for H[2] production include partial oxidation
   of hydrocarbons:

          CH[4] + 0.5 O[2] → CO + 2 H[2]

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

          C + H[2]O → CO + H[2]

   NB. Hydrogen is sometimes produced and consumed in the same industrial
   process, without being separated. In the Haber process for the
   production of ammonia (the world's fifth most produced industrial
   compound), hydrogen is generated from natural gas.

Biological syntheses

   H[2] 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 H[2] and its component two
   protons and two electrons. Evolution of hydrogen gas occurs in the
   transfer of reducing equivalents produced during pyruvate fermentation
   to water.

   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 H[2] gas by specialized hydrogenases in the chloroplast.
   Efforts have been undertaken to genetically modify cyanobacterial
   hydrogenases to efficiently synthesize H[2] gas even in the presence of
   oxygen.

   Other rarer but mechanistically interesting routes to H[2] production
   also exist in nature. Nitrogenase produces approximately one equivalent
   of H[2] for each equivalent of N[2] reduced to ammonia. Some
   phosphatases reduce phosphite to H[2].

Applications

   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.

   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. H[2] is used as the rotor coolant in
   electrical generators at power stations, because it has the highest
   thermal conductivity of any gas. Liquid H[2] is used in cryogenic
   research, including superconductivity studies. Since H[2] 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.
   However, this use was curtailed after the Hindenburg disaster convinced
   the public that the gas was too dangerous for this purpose.

   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.

   The triple point temperature of equilibrium hydrogen is a defining
   fixed point on the ITS-90 temperature scale.

Hydrogen as an energy carrier

   Having been used as an ingredient in some rocket fuels for several
   decades, hydrogen, or more specifically H[2], is now widely discussed
   in the context of energy. Hydrogen is not an energy source, since it is
   not an abundant natural resource and more energy is used to produce it
   than can be ultimately extracted from it. However, it could become
   useful as a carrier of energy, as elucidated in the United States
   Department of Energy's 2003 report, "Among the various alternative
   energy strategies, building an energy infrastructure that uses hydrogen
   — the third most abundant element on the earth's surface — as the
   primary carrier that connects a host of energy sources to diverse end
   uses may enable a secure and clean energy future for the Nation." The
   hydrogen would then locally be converted into usable energy either via
   combustion of fossil fuels or by electrochemical conversion into
   electricity in a fuel cell.

   One theoretical advantage of using H[2] as a carrier is the
   localization and concentration of environmentally unwelcome aspects of
   hydrogen manufacture. For example, CO[2] sequestration could be
   conducted at the point of H[2] production from methane. Hydrogen could
   also be produced using the electrolysis of water method; however, this
   is currently three to six times as expensive as production from natural
   gas. High-temperature electrolysis, which promises greater efficiency,
   is being investigated. Currently, however hydrogen production is
   expensive relative to other energy storage chemicals, and the bulk of
   hydrogen is now produced by the least expensive method, which (as
   noted) employs methane and which, as currently practiced, creates
   greenhouse gas emissions.

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