Chemical element of hydrogen

Properties, combustion, electron energy levels. Elemental molecular forms. Covalent, organic compounds. Hydrides, protons, acids, isotopes. Discovery and use. Role in quantum theory. Natural occurrence, production: laboratory, industrial, thermochemical.

Рубрика Химия
Вид реферат
Язык английский
Дата добавления 13.11.2009
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Summary

Subject: Hydrogen

Content

1. Introduction

2. Properties

2.1 Combustion

2.2 Electron energy levels

2.3 Elemental molecular forms

2.4 Compounds

2.4.1 Covalent and organic compounds

2.4.2 Hydrides

2.4.3 Protons and acids

2.5 Isotopes

3. History

3.1 Discovery and use

3.2 Role in quantum theory

4. Natural occurrence

5. Production

5.1 Laboratory

5.2 Industrial

5.3 Thermochemical

6. Conclusion

7. References

1. Introduction

Hydrogen is the most abundant chemical element, constituting roughly 75 % of the universe's elemental mass[3]. Stars in the main sequence are mainly composed of hydrogen in its plasma state. Naturally occuring elemental hydrogen is relatively rare on Earth.

The most common isotope of hydrogen is protium (name rarely used, symbol H) with a single proton and no neutrons. In ionic compounds it can take a negative charge (an anion known as a hydride and written as H?), or as a positively-charged species H+. The latter cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds always occur as more complex species. Hydrogen forms compounds with most elements and is present in water and most organic compounds. It plays a particularly important role in acid-base chemistry with many reactions exchanging protons between soluble molecules. As the simplest atom known, the hydrogen atom has been of theoretical use. For example, as the only neutral atom with an analytic solution to the Schrodinger equation, the study of the energetics and bonding of the hydrogen atom played a key role in the development of quantum mechanics.

Industrial production is mainly from the steam reforming of natural gas[4]. and less often from more energy-intensive hydrogen production methods like the electrolysis of water. Most hydrogen is employed near its production site, with the two largest uses being fossil fuel processing (e.g., hydrocracking) and ammonia production, mostly for the fertilizer market.

Hydrogen is important in metallurgy as it can embrittle many metals[5], complicating the design of pipelines and storage tanks[6]. Hydrogen has increasingly received attention as an energy-storage medium which burns in a less-polluting way than do fossil fuels.

2. Properties

2.1 Combustion

Figure. 0. The Space Shuttle Main Engine burns hydrogen with oxygen, producing a nearly-invisible flame at full thrust.

Hydrogen gas (dihydrogen[7]) is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume.[8] The enthalpy of combustion for hydrogen is ?286 kJ/mol:[9]

2 H2(g) + O2(g) > 2 H2O(l) + 572 kJ (286 kJ/mol)[note 1]

Hydrogen gas forms explosive mixtures with air in the concentration range 4-74% (volume per cent of hydrogen in air) and with chlorine in the range 5-95%. The mixtures spontaneously detonate by spark, heat or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C (932 °F).[10] Pure hydrogen-oxygen flames emit ultraviolet light and 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. The detection of a burning hydrogen leak may require a flame detector; such leaks can be very dangerous. The destruction of the Hindenburg airship was an infamous example of hydrogen combustion; the cause is debated, but the visible flames were the result of combustible materials in the ship's skin.[11] Because hydrogen is buoyant in air, hydrogen flames tend to ascend rapidly and cause less damage than hydrocarbon fires. Two-thirds of the Hindenburg passengers survived the fire, and many deaths were instead the result of falls or burning diesel fuel.[12]

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.[13]

2.2 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 wavelength.[14]

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.[15]

A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrodinger equation or the equivalent Feynman path integral formulation to calculate the probability density of the electron around the proton.[16]

Figure. 1. Depiction of a hydrogen atom showing the diameter as about twice the Bohr model radius (image not to scale).

2.3 Elemental molecular forms

There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei.[17] 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 (?+?); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 (?-?). 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".[18] 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 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.[19] 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.[20]

There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei.[17] 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 (?+?); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 (?-?). 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".[18] 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 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.[19] 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.[20]

A molecular form called protonated molecular hydrogen, or H+3, 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. 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.[24]

Figure. 2. First tracks observed in liquid hydrogen bubble chamber at the Bevatron

2.4 Compounds

2.4.1 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. 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.[25] 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.[26][27] 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.[28]

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;[29] the study of their properties is known as organic chemistry[30] and their study in the context of living organisms is known as biochemistry.[31] By some definitions, "organic" compounds are only required to contain carbon. 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.[29]

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.[32]

2.4.2 Hydrides

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 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.[33] 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 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.[34] Binary indium hydride has not yet been identified, although larger complexes exist.[35]

2.4.3 Protons and acids

Oxidation of hydrogen, in the sense of removing its electron, formally gives H+, containing 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" (H3O+). However, even in this case, such solvated hydrogen cations are thought more realistically physically to be organized into clusters that form species closer to H9O+4.[36] Other oxonium ions are found when water is in solution with other solvents.[37]

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.[38]

2.5 Isotopes

Figure. 3. 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.[39][40]

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.[41]

2H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. Essentially 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 1H-NMR spectroscopy.[42] Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.[43]

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.[32] 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.[44] It is used in nuclear fusion reactions,[45] as a tracer in isotope geochemistry,[46] and specialized in self-powered lighting devices.[47] Tritium has also been used in chemical and biological labeling experiments as a radiolabel.[48]

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.[49] 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.[50]

3. History

3.1 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.[51] 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.[52] 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 "flammable air" and further finding in 1781 that the gas produces water when burned. He is usually given credit for its discovery as an element.[53][54] In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek hydro meaning water and genes meaning creator)[55] when he and Laplace reproduced Cavendish's finding that water is produced when hydrogen is burned.[54]

Hydrogen was liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask.[54] He produced solid hydrogen the next year.[54] Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck.[53] Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932.[54] Francois Isaac de Rivaz built the first internal combustion engine powered by a mixture of hydrogen and oxygen in 1806. Edward Daniel Clarke invented the hydrogen gas blowpipe in 1819. The Dobereiner's lamp and limelight were invented in 1823.[54]

The first hydrogen-filled balloon was invented by Jacques Charles in 1783.[54] 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.[54] 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.[54] 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 May 6, 1937.[54] 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 aluminized fabric coating by static electricity. But the damage to hydrogen's reputation as a lifting gas was already done. 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,[56] 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).[57] For example, the ISS,[58] Mars Odyssey[59] and the Mars Global Surveyor[60] are equipped with nickel-hydrogen batteries. The Hubble Space Telescope, at the time its original batteries were finally changed in May 2009, more than 19 years after launch, led with the highest number of charge/discharge cycles.

3.2 Role in quantum theory

Figure. 4. Hydrogen emission spectrum lines in the visible range. These are the four visible lines of the Balmer series

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.[61] 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.[62]

4. Natural occurrence

Figure. 5. 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.[63] 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.[64]

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.[65]

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. However, hydrogen is the third most abundant element on the Earth's surface.[66] Most of the Earth's hydrogen is in the form of chemical compounds such as hydrocarbons and water.[32] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus. Methane is a hydrogen source of increasing importance.[67]

5. 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.

5.1 Laboratory

In the laboratory, H2 is usually prepared by the reaction of acids on metals such as zinc with Kipp's apparatus.

Zn + 2 H+ > Zn2+ + H2

Aluminium can also produce H2 upon treatment with bases:

2 Al + 6 H2O + 2 OH? > 2 Al(OH)4? + 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%.[68]

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.[69]

5.2 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.[70] At high temperatures (1000-1400 K, °C;700-1100 °C or 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 (2.0MPa, 20 atm or 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:[70]

CO + H2O > CO2 + H2

Other important methods for H2 production include partial oxidation of hydrocarbons:[71]

2 CH4 + O2 > 2 CO + 4 H2

and the coal reaction, which can serve as a prelude to the shift reaction above:[70]

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.[72] Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.[73]

5.3 Thermochemical

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.[74] 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.[75]

6. Conclusion

As an individual representative of the periodic table of chemical elements Dmitry Ivanovich Mendeleyev, the element has unique chemical and physical properties

Element is of great economic importance and plays a major role in world culture

7. References

"Dihydrogen". O=CHem Directory. University of Southern Maine. Retrieved 2009-04-06.

"Hydrogen Basics -- Production". Florida Solar Energy Center. 2007. Retrieved 2008-02-05.

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.

Christensen, C.H.; Norskov, J.K.; Johannessen, T. (9 July 2005). "Making society independent of fossil fuels -- Danish researchers reveal new technology". Technical University of Denmark. Retrieved 2008-03-28.

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

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 0309091632.

Dziadecki, J. (2005). "Hindenburg Hydrogen Fire". Retrieved 2007-01-16.

Kelly, M.. "The Hindenburg Disaster". About.com:American history. Retrieved 2009-08-08.

Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81th edition, CRC press.

Millar, Tom (December 10, 2003). "Lecture 7, Emission Lines -- Examples". PH-3009 (P507/P706/M324) Interstellar Physics. University of Manchester. Retrieved 2008-02-05.

Palmer, D. (13 September 1997). "Hydrogen in the Universe". NASA. Retrieved 2008-02-05.

Patnaik, P (2007). A comprehensive guide to the hazardous properties of chemical substances. Wiley-Interscience. p. 402. ISBN 0471714585.

Rogers, H.C. (1999). "Hydrogen Embrittlement of Metals". Science 159 (3819): 1057-1064. doi:10.1126/science.159.3819.1057. PMID 17775040.

Simpson, J.A.; Weiner, E.S.C. (1989). "Hydrogen". Oxford English Dictionary. 7 (2nd ed.). Clarendon Press. ISBN 0-19-861219-2.


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