Hydrogen

This article is about the chemistry of hydrogen. For the physics of atomic hydrogen, see hydrogen atom.

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Hydrogen ([hydrogenium] Error: {{Lang-xx}}: text has italic markup (help), from [hydro] Error: {{Lang-xx}}: text has italic markup (help): "water" and [genes] Error: {{Lang}}: text has italic markup (help): "forming") is a chemical element in the periodic table that has the symbol H and atomic number 1. At standard temperature and pressure it is a colorless, odorless, nonmetallic, univalent, tasteless, highly flammable diatomic gas (H₂). With an atomic mass of just 1.00794 g/mol, hydrogen is the lightest element. It is also the most abundant, constituting roughly 75% of the universe's elemental matter. Stars in their main sequence are overwhelmingly composed of hydrogen in its plasma state. Elemental hydrogen is industrially produced from hydrocarbons and is currently used primarily in fossil fuel upgrading but has a variety of other applications in both the energy and other sectors of the world's economy.

The most naturally common isotope of hydrogen contains one electron and an atomic nucleus of one proton. In ionic compounds it can take on either a positive charge (becoming a cation, formally a bare proton) or a negative charge (becoming an anion known as a hydride). Hydrogen can form compounds with most other elements and is present in water and all 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 element for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key historical and theoretical role in the development of quantum mechanics.

Different meanings of "hydrogen"

The word "hydrogen" has several different meanings that are important to distinguish. Possible uses:

• Hydrogen is the name of an element.
• Hydrogen is an atom, sometimes called "H dot" that is abundant in space but essentially absent on earth, because it dimerizes.
• Hydrogen is a diatomic molecule that occurs naturally in trace amounts in the Earth's atmosphere; chemists increasingly refer to H₂ as dihydrogen to distinguish this molecule from atomic hydrogen and hydrogen found in other compounds.
• Hydrogen is atomic constituent within all organic compounds, water, and many other chemical compounds.

It is especially important not to confuse elemental forms of hydrogen with hydrogen as it appears in chemical compounds.

History

Discovery of H₂

Hydrogen gas, H₂, was first artificially produced and formally described by Theophrastus Bombastus von Hohenheim (1493–1541)—also known as Paracelsus— 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 "flammable," and further finding that the gas produces water when burned in air. 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, including the fact that it produced water when burned. In 1783 Antoine Lavoisier gave the element its name and (with Laplace) reported that pure water is produced by burning hydrogen and oxygen.

One of the first uses of H₂ was for balloons. The H₂ was obtained by reacting sulfuric acid and metallic iron. Famously, H₂ was used in the Hindenburg airship that was destroyed in a midair fire.

Role in history of quantum theory

Has the ability to make a human explode when too much gets into the skin cells. some question this fact of life. they are dumbases. 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 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 Maxwell's observation, half a century before full quantum mechanical theory arrived, that the specific heat capacity of H₂ unaccountably resembles that of a monatomic gas below room temperature. According to quantum theory, this behavior arises from the spacing of the (quantized) rotational energy levels, which are particularly wide-spaced in H₂ due to 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 atomic 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 H₂ 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 plasma state 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 emmissivity (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.

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) due to its light weight, which enables it to escape from Earth's gravity more easily than heavier gases. Thus hydrogen is both ubiquitous in the universe and difficult to produce in concentrated form on Earth. Although H atoms and H₂ 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, from which it cannot be extracted into its elemental form without significant input of energy. Most of the Earth's hydrogen is in the form of water.. Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus.

The hydrogen atom

Electron energy levels

Main article: Hydrogen atom

depiction of a hydrogen-1 atom 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, electrons and protons are attracted to one another by the electromagnetic force, while planets and celestial objects are attracted to each other by gravity. Due to the discretization of energy inherent in quantum mechanics, the electron in the Bohr model can only occupy certain allowed distances from the proton. 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 experimental results such as the energy levels and hydrogen spectrum more accurately than the particle-based Bohr model. Finally, 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.

The electronic ground state energy level is split into hyperfine structure levels because of magnetic effects due to 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

Main article: Isotopes of hydrogen

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. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium is used in nuclear fusion reactions, as a radiolabel in biochemistry, and as a solvent in ¹H-NMR spectroscopy.

• 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 due to 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.

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.

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. The two forms have slightly different physical properties; for example, the melting and boiling points of parahydrogen are about 0.1 K lower than those of orthohydrogen. 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 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 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 The ortho/para ratio in condensed H₂ is an important consideration in the preparation and storage of liquid hydrogen, since the ortho-para conversion is exothermic and produces enough heat to evaporate the hydrogen liquid, which causes hydrogen loss after liquefying. Catalysts for the ortho-para interconversion, such as iron filings, are used during hydrogen cooling.

Chemical and physical properties

H₂ is less soluble in water, alcohol, or ether than oxygen is. Its solubility and adsorption characteristics 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.

Combustion

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

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

2 H₂(g) + O₂(g) → 2 H₂O(l) + 572 kJ

When mixed with oxygen across a wide range of proportions, hydrogen explodes upon ignition. Uniquely, 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. Although it is widely believed that the Hindenburg zeppelin burned due to the hydrogen gas it contained, the flames seen at right are actually from the covering skin of the blimp that contained carbon and pyrophoric aluminium powder. 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, partly for this reason.

H₂ reacts directly with other oxidizing elements. A violent reaction can occur with chlorine and fluorine, forming the corresponding hydrogen halides, HCl and HF.

Compounds

Further information: [[:Hydrogen 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. 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 association with living things, these compounds are called organic compounds; the study of their properties is known as organic chemistry and their study in the context of living organisms is known as biochemistry. (By some definitions "organic" compounds are only required to contain carbon; however most of them also contain hydrogen, and it is addition of hydrogen that gives them their particular chemical characteristics).

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 aluminum 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 hydride anion is a convenient bookkeeping tool but does not exist per se - alkali metal hydrides, e.g. sodium hydride (NaH), are polymeric and have no solution chemistry. In fact in 1920, K. Moers demonstrated that electrolysis of molten lithium hydride (LiH, m.p. 692 °C) produced a stoichiometric quantity of hydrogen at the anode. In lithium aluminum hydride, the AlH₄⁻ anion carries hydridic centers firmly attached to the Al(III). Palladium hydride contains interstitial hydrogen atoms, i.e. the H atoms are bonded to multiple Pd atoms without perturbing the overall Pd framework. Hydrogen forms hydrides with all main group elements with the exception of the noble gases and indium and thallium.

"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 due to 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⁺). Although even this representation is inadequate, the more accurate H₉O₄⁺ is rarely used in discussions of acids because it makes balancing reactions clunky and tedious. The basicity of water guarantees that the oxonium is the dominant form in water. In anhydrous acids, other forms are found.

Although exotic on earth, one of the most common ions in the universe is the H₃⁺ ion.

H₂ reacts with oxygen to form water, H₂O. Considerable energy is released in this process. At room temperature no reaction occurs between H₂ and O₂ in the absence of a catalyst.

Production

H₂ is produced in chemistry and biology laboratories, often as a byproduct 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, H₂ is usually prepared by the reaction of acids on metals such as zinc.

Zn + 2 H⁺ → Zn²⁺ + H₂

Aluminum 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 but expensive method of producing hydrogen. Typically the cathode electrode is made from platinum.

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), 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) 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 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 + H₂O → CO₂ + H₂

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

CH₄ + 0.5 O₂ → CO + 2 H₂

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

C + H₂O → CO + H₂

Note: 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 in situ from natural gas.

Biological syntheses

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. 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 H₂ gas by specialized hydrogenases in the chloroplast. Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H₂ gas even in the presence of oxygen

Other rarer but mechanistically interesting routes to H₂ production also exist in nature. Nitrogenase produces approximately one equivalent of H₂ for each equivalent of N₂ reduced to ammonia. Some phosphatases reduce phosphite to H₂.

Applications

Large quantities of H₂ are needed in the petroleum and chemical industries. By far the largest application of H₂ is for the processing ("upgrading") of fossil fuels. The key consumers of H₂ in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H₂ has several other important uses.

• used in the hydrogenation of fats and oils (found in items such as margarine), and in the production of methanol.
• H₂ is used in the manufacture of hydrochloric acid
• H₂ is used in certain welding methods
• H₂ is used in the reduction of metallic ores.
• H₂ is an ingredient in some rocket fuels.
• 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.
• The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale.
• 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. However, this use was curtailed after the Hindenburg disaster convinced the public that the gas was too dangerous for this purpose.
• Deuterium, an isotope of hydrogen (hydrogen-2), is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions. Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects.
• Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, as an isotopic label in the biosciences, and as a radiation source in luminous paints.

Hydrogen as an energy carrier

Main article: hydrogen economy

Hydrogen, or more specifically H₂, is 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.” One theoretical advantage of using H₂ as a carrier, is the localization and concentration of environmentally unwelcome aspects of hydrogen manufacture. For example, CO₂ sequestration could be conducted at the point of H₂ production.

See also

Antihydrogen Biofuel Deuterium Electric vehicle Fuel cell Hybrid vehicle Hydrogen atom

Hydrogen bomb Hydrogen bond Hydrogen vehicle Hydrogen cycle Hydrogen economy The Hype about Hydrogen Hydrogen leak testing Hydrogen line

Hydrogen molecule Hydrogen spectral series Hydrogen station Liquid hydrogen Natural gas Photohydrogen Tritium

General references

• "Chart of the Nuclides". Fourteenth Edition. General Electric Company. 1989. {{cite journal}}: Cite journal requires |journal= (help)
• Ferreira-Aparicio, P (2005). "New Trends in Reforming Technologies: from Hydrogen Industrial Plants to Multifuel Microreformers". Catalysis Reviews. 47: 491–588. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
• Krebs, Robert E. (1998). The History and Use of Our Earth's Chemical Elements : A Reference Guide. Westport, Conn.: Greenwood Press. ISBN 0-313-30123-9.
• Newton, David E. (1994). The Chemical Elements. New York, NY: Franklin Watts. ISBN 0-531-12501-7.
• Rigden, John S. (2002). Hydrogen : The Essential Element. Cambridge, MA: Harvard University Press. ISBN 0-531-12501-7.
• Romm, Joseph, J. (2004). The Hype about Hydrogen, Fact and Fiction in the Race to Save the Climate. Island Press. ISBN 1-55963-703-X.{{cite book}}: CS1 maint: multiple names: authors list (link) Author interview at Global Public Media.
• Stwertka, Albert (2002). A Guide to the Elements. New York, NY: Oxford University Press. ISBN 0-19-515027-9.

External links

• Basic Hydrogen Calculations of Quantum Mechanics
• Biohydrogen
• Computational Chemistry Wiki
• Hydrogen phase diagram
• RIKEN Beam Science Laboratory, Japan — Heavy hydrogen research
• Wavefunction of hydrogen