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Hydrogen is the chemical element with atomic number 1. It is represented by the symbol H. With an atomic weight of 1.00794 u, hydrogen is the lightest and most abundant chemical element, constituting roughly 75 % of the Universe's elemental mass.[1] Stars in the main sequence are mainly composed of hydrogen in its plasma state. Naturally occurring elemental hydrogen is relatively rare on Earth.

Hydrogens mest almindelige isotop er protium (navnet bliver sjældent benyttet, symbolet er 1H) med en enkelt proton og ingen neutroner. I ioniske forbindelser kan det tage form af en negativ ladning (en anion kendt som hydrid og skrives H), eller som en positivt ladet specie H+. Den sidstnævnte kation

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 Schrödinger equation, the study of the energetics and bonding of the hydrogen atom played a key role in the development of quantum mechanics.

Hydrogen gas (now known to be H2) was first artificially produced in the early 16th century, via the mixing of metals with strong acids. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance,[2] and that it produces water when burned, a property which later gave it its name, which in Greek means "water-former". At standard temperature and pressure, hydrogen is a colorless, odorless, nonmetallic, tasteless, highly combustible diatomic gas with the molecular formula H2.

Industrial production is mainly from the steam reforming of natural gas, and less often from more energy-intensive hydrogen production methods like the electrolysis of water.[3] 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 a concern in metallurgy as it can embrittle many metals,[4] complicating the design of pipelines and storage tanks.[5]

Properties[redigér | redigér wikikode]

Forbrænding[redigér | redigér wikikode]

A black cup-like object hanging by its bottom with blue glow coming out of its opening.
The Space Shuttle Main Engine burns hydrogen with oxygen, producing a nearly-invisible flame at full thrust.

Brint, H2, (dihydrogen gas) er yderst brandfarlig og vil brænder i luft ved en meget bred vifte af koncentrationer, nemlig mellem 4% og 75% (volumen procent).[6] The enthalpy of combustion for hydrogen er −286 kJ/mol:[7]

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

Brint danner eksplosive blandinger med luft i konentrationsintervallet fra 4-74% (volumenprocent af brint i luften) og med chlor i intervallet 5-95%. Blandinger eksploderer spontant ved gnister, varme eller sollys. Brints selvantændelsestemperatur, dvs. temperaturen hvor selvantændelse sker spontant, er 500 °C.[8]

Rene brint-ilt (H2-O2) flammer udsender ultraviolet lys og er næsten usynlige for det blotte øje, hvilket er illustreret ved den svage røgfane af Space Shuttle hovedmotoren i forhold til den meget synlige røgfane af Space Shuttle Solid Rocket Booster. Detektering af brint læk kræver en flam detector; sådanne lækager kan være ekstremt farlige. Ekplosionen på Hindenburg luftskibet er et berygtet eksempel på brints forbrænding; årsagen til eksplosionen er stadig ikke fastlagt, men de synlige flammer var resultatet af brændbart materialer i skibets konstruktion. [9]

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.[10]

Brint, H2 reagerer med alle oxiderende grundstoffer. Brint kan reagerer spontant og voldsomt med chlor og fluor ved stuetemperatur hvorved dannes de korreksponderende hydrogenhalider, saltsyre (HCl) og hydrogenfluorid (HF), som er potentielt farlige syrer.[11]

Electron energy levels[redigér | redigér wikikode]

Nuvola apps download manager2-70%.svg Hovedartikel: Hydrogen atom.
Drawing of a light-gray large sphere with a cut off quarter and a black small sphere and numbers 1.6 and 1.7x10-5 illustrating their relative diameters.
Depiction of a hydrogen atom showing the diameter as about twice the Bohr model radius (image not to scale).

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.[12]

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.[13]

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.[14]

Elemental molecular forms[redigér | redigér wikikode]

Two bright circles on dark background, both contain numerous thin black lines inside.
First tracks observed in liquid hydrogen bubble chamber at the Bevatron

There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei.[15] 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 (½+½); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 (½-½). 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".[16] 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.[17] 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.[18]

The uncatalyzed interconversion between para and ortho H2 increases with increasing temperature; thus rapidly condensed H2 contains large quantities of the high-energy ortho form that converts to the para form very slowly.[19] The ortho/para ratio in condensed H2 is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate some of the hydrogen liquid, leading to loss of liquefied material. Catalysts for the ortho-para interconversion, such as ferric oxide, activated carbon, platinized asbestos, rare earth metals, uranium compounds, chromic oxide, or some nickel[20] compounds, are used during hydrogen cooling.[21]

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.[22] Neutral triatomic hydrogen H3 can only exist in an excited from and is unstable.[23]

Compounds[redigér | redigér wikikode]

Yderligere information: [[Hydrogen compounds]]

Covalent and organic compounds[redigér | redigér wikikode]

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.[24] 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.[25][26] 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.[27]

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;[28] the study of their properties is known as organic chemistry[29] and their study in the context of living organisms is known as biochemistry.[30] 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.[28]

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.[31]

Hydrider[redigér | redigér wikikode]

Forbindelser af hydrogen kaldes ofte for hydrider, en term som bruges meget løst. Termen "hydrid" antyder at hydrogenatomet har fået en negativ eller anionis karakter, hvilket skrives H-, og bruges når hydrogen danner en forbindelse med et mere elektropositivt grundstof. Eksistensen af hydridanioner blev foreslået af Gilbert N. Lewis i 1916 for gruppe I og II salt-lignende hydrider, og blev eftervist af Moers i 1920 ved elektrolyse af smeltet lithiumhydrid (LiH), som dannede en støkiometrisk mængde af hydrigen ved anoden.[32] 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.[33] Binære indiumhydrider er endnu ikke blevet identificeret, selvom større komplekser eksisterer.[34]

Protons and acids[redigér | redigér wikikode]

Skabelon:See 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.[35] Other oxonium ions are found when water is in solution with other solvents.[36]

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.[37]

Isotopes[redigér | redigér wikikode]

Referencer[redigér | redigér wikikode]

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  3. ^ Hydrogen Basics — Production. Florida Solar Energy Center. 2007. Hentet 2008-02-05. 
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  7. ^ 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. s. 240. ISBN 0309091632. 
  8. ^ Patnaik, P (2007). A comprehensive guide to the hazardous properties of chemical substances. Wiley-Interscience. s. 402. ISBN 0471714585. 
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  10. ^ Kelly, M.. The Hindenburg Disaster. history. Hentet 2009-08-08. 
  11. ^ Clayton, D.D. (2003). Handbook of Isotopes in the Cosmos: Hydrogen to Gallium. Cambridge University Press. ISBN 0521823811. 
  12. ^ Millar, Tom (December 10, 2003). "Lecture 7, Emission Lines — Examples". PH-3009 (P507/P706/M324) Interstellar Physics (University of Manchester). Hentet 2008-02-05. 
  13. ^ Stern, David P. (2005-05-16). The Atomic Nucleus and Bohr's Early Model of the Atom. NASA Goddard Space Flight Center (mirror). Hentet 2007-12-20. 
  14. ^ Stern, David P. (2005-02-13). Wave Mechanics. NASA Goddard Space Flight Center. Hentet 2008-04-16. 
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Fodnotefejl: <ref>-tags eksisterer for en gruppe betegnet "note", men der blev ikke fundet et tilsvarende {{reflist|group="note"}}, eller et afsluttende </ref>-tag mangler