hydrogen

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Dictionary: hy·dro·gen (hī'drə-jən) pronunciation

n. (Symbol H)

A colorless, highly flammable gaseous element, the lightest of all gases and the most abundant element in the universe, used in the production of synthetic ammonia and methanol, in petroleum refining, in the hydrogenation of organic materials, as a reducing atmosphere, in oxyhydrogen torches, and in rocket fuels. Atomic number 1; atomic weight 1.00794; melting point −259.14°C; boiling point −252.8°C; density at 0°C 0.08987 gram per liter; valence 1.

[French hydrogène : Greek hudro-, hydro- + French -gène, -gen.]

hydrogenous hy·drog'e·nous (-drōj'ə-nəs) adj.

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Science of Everyday Things: Hydrogen

Concept

First element on the periodic table, hydrogen is truly in a class by itself. It does not belong to any family of elements, and though it is a nonmetal, it appears on the left side of the periodic table with the metals. The other elements with it in Group 1 form the alkali metal family, but obviously, hydrogen does not belong with them. Indeed, if there is any element similar to hydrogen in simplicity and abundance, it is the only other one on the first row, or period, of the periodic table: helium. Together, these two elements make up 99.9% of all known matter in the entire universe, because hydrogen atoms in stars fuse to create helium. Yet whereas helium is a noble gas, and therefore chemically unreactive, hydrogen bonds with all sorts of other elements. In one such variety of bond, with carbon, hydrogen forms the backbone for a vast collection of organic molecules, known as hydrocarbons and their derivatives. Bonded with oxygen, hydrogen forms the single most important compound on Earth, and the most important complex substance other than air: water. Yet when it bonds with sulfur, it creates toxic hydrogen sulfide; and on its own, hydrogen is extremely flammable. The only element whose isotopes have names, hydrogen has long been considered as a potential source of power and transportation: once upon a time for airships, later as a component in nuclear reactions—and, perhaps in the future, as a source of abundant clean energy.

How It Works

The Essentials

The atomic number of hydrogen is 1, meaning that it has a single proton in its nucleus. With its single electron, hydrogen is the simplest element of all. Because it is such a basic elemental building block, figures for the mass of other elements were once based on hydrogen, but the standard today is set by ¹²C or carbon-12, the most common isotope of carbon.

Hydrogen has two stable isotopes—forms of the element that differ in mass. The first of these, protium, is simply hydrogen in its most common form, with no neutrons in its nucleus. Protium (the name is only used to distinguish it from the other isotopes) accounts for 99.985% of all the hydrogen that appears in nature. The second stable isotope, deuterium, has one neutron, and makes up 0.015% of all hydrogen atoms. Tritium, hydrogen's one radioactive isotope, will be discussed below.

The fact that hydrogen's isotopes have separate names, whereas all other isotopes are designated merely by element name and mass number (for example, "carbon-12") says something about the prominence of hydrogen as an element. Not only is its atomic number 1, but in many ways, it is like the number 1 itself—the essential piece from which all others are ultimately constructed. Indeed, nuclear fusion of hydrogen in the stars is the ultimate source for the 90-odd elements that occur in nature.

The mass of this number-one element is not, however, 1: it is 1.008 amu, reflecting the small quantities of deuterium, or "heavy hydrogen," present in a typical sample. A gas at ordinary temperatures, hydrogen turns to a liquid at −423.2°F (−252.9°C), and to a solid at −434.°F (−259.3°C). These figures are its boiling point and melting point respectively; only the figures for helium are lower. As noted earlier, these two elements make up all but 0.01% of the known elemental mass of the universe, and are the principal materials from which stars are formed.

Normally hydrogen is diatomic, meaning that its molecules are formed by two atoms. At the interior of a star, however, where the temperature is many millions of degrees, H₂ molecules are separated into atoms, and these atoms become ionized. In other words, the electron separates from the proton, resulting in an ion with a positive charge, along with a free electron. The positive ions experience fusion—that is, their nuclei bond, releasing enormous amounts of energy as they form new elements.

Because the principal isotopic form of helium has two protons in the nucleus, it is natural that helium is the element usually formed; yet it is nonetheless true—amazing as it may seem—that all the elements found on Earth were once formed in stars. On Earth, however, hydrogen ranks ninth in its percentage of the planet's known elemental mass: just 0.87%. In the human body, on the other hand, it is third, after oxygen and carbon, making up 10% of human elemental body mass.

Hydrogen and Bonding

Having just one electron, hydrogen can bond to other atoms in one of two ways. The first option is to combine its electron with one from the atom of a nonmetallic element to make a covalent bond, in which the two electrons are shared. Hydrogen is unusual in this regard, because most atoms conform to the octet rule, ending up with eight valence electrons. The bonding behavior of hydrogen follows the duet rule, resulting in just two electrons for bonding.

Examples of this first type of bond include water (H₂O), hydrogen sulfide (H₂S), and ammonia (NH₃), as well as the many organic compounds formed on a hydrogen-carbon backbone. But hydrogen can form a second type of bond, in which it gains an extra electron to become the negative ion H−, or hydride. It is then able to combine with a metallic positive ion to form an ionic bond. Ionic hydrides are convenient sources of hydrogen gas: for instance, calcium hydride, or CaH₂, is sold commercially, and provides a very convenient means of hydrogen generation. The hydrogen gas produced by the reaction of calcium hydride with water can be used to inflate life rafts.

The presence of hydrogen in certain types of molecules can also be a factor in intermolecular bonding. Intermolecular bonding is the attraction between molecules, as opposed to the bonding within molecules, which is usually what chemists mean when they talk about "bonding."

Hydrogen's Early History

Because it bonds so readily with other elements, hydrogen almost never appears in pure elemental form on Earth. Yet by the late fifteenth century, chemists recognized that by adding a metal to an acid, hydrogen was produced. Only in 1766, however, did English chemist and physicist Henry Cavendish (1731-1810) recognize hydrogen as a substance distinct from all other "airs," as gases then were called.

Seventeen years later, in 1783, French chemist Antoine Lavoisier (1743-1794) named the substance after two Greek words: hydro (water) and genes (born or formed). It was another two decades before English chemist John Dalton (1766-1844) formed his atomic theory of matter, and despite the great strides he made for science, Dalton remained convinced that hydrogen and oxygen in water formed "water atoms." Around the same time, however, Italian physicist Amedeo Avogadro (1776-1856) clarified the distinction between atoms and molecules, though this theory would not be generally accepted until the 1850s.

Contemporary to Dalton and Avogadro was Swedish chemist Jons Berzelius (1779-1848), who developed a system of comparing the mass of various atoms in relation to hydrogen. This method remained in use for more than a century, until the discovery of neutrons, protons, and isotopes pointed the way toward a means of making more accurate determinations of atomic mass. In 1931, American chemist and physicist Harold Urey (1893-1981) made the first separation of an isotope: deuterium, from ordinary water.

Real-Life Applications

Deuterium and Tritium

Designated as ²H, deuterium is a stable isotope, whereas tritium—³H—is unstable, or radioactive. Not only do these two have names; they even have chemical symbols (D and T, respectively), as though they were elements on the periodic table. Just as hydrogen represents the most basic proton-electron combination against which other atoms are compared, these two are respectively the most basic isotope containing a single neutron, and the most basic radioisotope, or radioactive isotope.

Deuterium is sometimes called "heavy hydrogen," and its nucleus is called a deuteron. In separating deuterium—an achievement for which he won the 1934 Nobel Prize—Urey collected a relatively large sample of liquid hydrogen: 4.2 qt (4 l). He then allowed the liquid to evaporate very slowly, predicting that the more abundant protium would evaporate more quickly than the heavier isotope. When all but 0.034 oz (1 ml) of the sample had evaporated, he submitted the remainder to a form of analysis called spectroscopy, adding a burst of energy to the atoms and then analyzing the light spectrum they emitted for evidence of differing varieties of atoms.

With an atomic mass of 2.014102 amu, deuterium is almost exactly twice as heavy as protium, which has an atomic mass of 1.007825. Its melting points and boiling points, respectively −426°F (−254°C) and −417°F (−249°C), are higher than for protium. Often, deuterium is applied as a tracer, an atom or group of atoms whose participation in a chemical, physical, or biological reaction can be easily observed.

Deuterium in War and Peace

In nuclear power plants, deuterium is combined with oxygen to form "heavy water" (D₂O), which likewise has higher boiling and melting points than ordinary water. Heavy water is often used in nuclear fission reactors to slow down the fission process, or the splitting of atoms. Deuterium is also present in nuclear fusion, both on the Sun and in laboratories.

During the period shortly after World War II, physicists developed a means of duplicating the thermonuclear fusion process. The result was the hydrogen bomb—more properly called a fusion bomb—whose detonating device was a compound of lithium and deuterium called lithium deuteride. Vastly more powerful than the "atomic" (that is, fission) bombs dropped by the United States over Japan (Nagaski and Hiroshima) in 1945, the hydrogen bomb greatly increased the threat of worldwide nuclear annihilation in the postwar years.

Yet the power that could destroy the world also has the potential to provide safe, abundant fusion energy from power plants—a dream as yet unrealized. Physicists studying nuclear fusion are attempting several approaches, including a process involving the fusion of two deuterons. This fusion would result in a triton, the nucleus of tritium, along with a single proton. Theoretically, the triton and deuteron would then be fused to create a helium nucleus, resulting in the production of vast amounts of energy.

Tritium

Whereas deuterium has a single neutron, tritium—as its mass number of 3 indicates—has two. And just as deuterium has approximately twice the mass of protium, tritium has about three times the mass, or 3.016 amu. Its melting and boiling points are higher still than those of deuterium: thus tritium heavy water (T₂O) melts at 40°F (4.5°C), as compared with 32°F (0°C) for H₂O.

Tritium has a half-life (the length of time it takes for half the radioisotopes in a sample to become stable) of 12.26 years. As it decays, its nucleus emits a low-energy beta particle, which is either an electron or a subatomic particle called a positron, resulting in the creation of the helium-3 isotope. Due to the low energy levels involved, the radioactive decay of tritium poses little danger to humans. In fact, there is always a small quantity of tritium in the atmosphere, and this quantity is constantly being replenished by cosmic rays.

Like deuterium, tritium is applied in nuclear fusion, but due to its scarcity, it is usually combined with deuterium. Sometimes it is released in small quantities into the groundwater as a means of monitoring subterranean water flow. It is also used as a tracer in biochemical processes, and as an ingredient in luminous paints.

Hydrogen and Oxygen

Water

Water, of course, is the most well-known compound involving hydrogen. Nonetheless, it is worthwhile to consider the interaction between hydrogen and oxygen, the two ingredients in water, which provides an interesting illustration of chemistry in action.

Chemically bonded as water, hydrogen and oxygen can put out any type of fire except an oil or electrical fire; as separate substances, however, hydrogen and oxygen are highly flammable. In an oxyhydrogen torch, the potentially explosive reaction between the two gases is controlled by a gradual feeding process, which produces combustion instead of the more violent explosion that sometimes occurs when hydrogen and oxygen come into contact.

Hydrogen Peroxide

Aside from water, another commonly used hydrogen-oxygen compound is hydrogen peroxide, or H₂O₂. A colorless liquid, hydrogen peroxide is chemically unstable (not "unstable" in the way that a radioisotope is), and decomposes slowly to form water and oxygen gas. In high concentrations, it can be used as rocket fuel.

By contrast, the hydrogen peroxide used in homes as a disinfectant and bleaching agent is only a 3% solution. The formation of oxygen gas molecules causes hydrogen peroxide to bubble, and this bubbling is quite rapid when the peroxide is placed on cuts, because the enzymes in blood act as a catalyst to speed up the reaction.

Hydrogen Chloride

Another significant compound involving hydrogen is hydrogen chloride, or HCl—in other words, one hydrogen atom bonded to chlorine, a member of the halogens family. Dissolved in water, it produces hydrochloric acid, used in laboratories for analyses involving other acids. Normally, hydrogen chloride is produced by the reaction of salt with sulfuric acid, though it can also be created by direct bonding of hydrogen and chlorine at temperatures above 428°F (250°C).

Hydrogen chloride and hydrochloric acid have numerous applications in metallurgy, as well as in the manufacture of pharmaceuticals, dyes, and synthetic rubber. They are used, for instance, in making pharmaceutical hydrochlorides, water-soluble drugs that dissolve when ingested. Other applications include the production of fertilizers, synthetic silk, paint pigments, soap, and numerous other products.

Not all hydrochloric acid is produced by industry, or by chemists in laboratories. Active volcanoes, as well as waters from volcanic mountain sources, contain traces of the acid. So, too, does the human body, which generates it during digestion. However, too much hydrochloric acid in the digestive system can cause the formation of gastric ulcers.

Hydrogen Sulfide

It may not be a pleasant subject, but hydrogen—in the form of hydrogen sulfide—is also present in intestinal gas. The fact that hydrogen sulfide is an extremely malodorous substance once again illustrates the strange things that happen when elements bond: neither hydrogen nor sulfur has any smell on its own, yet together they form an extremely noxious—and toxic—substance.

Pockets of hydrogen sulfide occur in nature. If a person were to breathe the vapors for very long, it could be fatal, but usually, the foul odor keeps people away. The May 2001 National Geographic included two stories relating to such natural hydrogen-sulfide deposits, on opposite sides of the Earth, and in both cases the presence of these toxic fumes created interesting results.

In southern Mexico is a system of caves known as Villa Luz, through which run some 20 underground springs, many of them carrying large quantities of hydrogen sulfide. The National Geographic Society's team had to enter the caves wearing gas masks, yet the area teems with strange varieties of life. Among these are fish that are red from high concentrations of hemoglobin, or red blood cells. The creatures need this extra dose of hemoglobin, necessary to move oxygen through the body, in order to survive on the scant oxygen supplies. The waters of the cave are further populated by microorganisms that oxidize the hydrogen sulfide and turn it into sulfuric acid, which dissolves the rock walls and continually enlarges the cave.

Thousands of miles away, in the Black Sea, explorers supported by a grant from the National Geographic Society examined evidence suggesting that there indeed had been a great ancient flood in the area, much like the one depicted in the Bible. In their efforts, they had an unlikely ally: hydrogen sulfide, which had formed at the bottom of the sea, and was covered by dense layers of salt water. Because the Black Sea lacks the temperature differences that cause water to circulate from the bottom upward, the hydrogen sulfide stayed at the bottom.

Under normal circumstances, the wreck of a 1,500-year-old wooden ship would not have been preserved; but because oxygen could not reach the bottom of the Black Sea—and thus wood-boring worms could not live in the toxic environment—the ship was left undisturbed. Thanks to the presence of hydrogen sulfide, explorers were able to study the ship, the first fully intact ancient shipwreck to be discovered.

Hydrocarbons

Together with carbon, hydrogen forms a huge array of organic materials known as hydrocarbons—chemical compounds whose molecules are made up of nothing but carbon and hydrogen atoms. Theoretically, there is no limit to the number of possible hydrocarbons. Not only does carbon form itself into seemingly limitless molecular shapes, but hydrogen is a particularly good partner. Because it has the smallest atom of any element on the periodic table, it can bond to one of carbon's valence electrons without getting in the way of the others.

Hydrocarbons may either be saturated or unsaturated. A saturated hydrocarbon is one in which the carbon atom is already bonded to four other atoms, and thus cannot bond to any others. In an unsaturated hydrocarbon, however, not all the valence electrons of the carbon atom are bonded to other atoms.

Hydrogenation is a term describing any chemical reaction in which hydrogen atoms are added to carbon multiple bonds. There are many applications of hydrogenation, but one that is particularly relevant to daily life involves its use in turning unsaturated hydrocarbons into saturated ones. When treated with hydrogen gas, unsaturated fats (fats are complex substances that involve hydrocarbons bonded to other molecules) become saturated fats, which are softer and more stable, and stand up better to the heat of frying. Many foods contain hydrogenated vegetable oil; however, saturated fats have been linked with a rise in blood cholesterol levels—and with an increased risk of heart disease.

Petrochemicals and Functional Groups

One important variety of hydrocarbons is described under the collective heading of petrochemicals—that is, derivatives of petroleum. These include natural gas; petroleum ether, a solvent; naphtha, a solvent (for example, paint thinner); gasoline; kerosene; fuel for heating and diesel fuel; lubricating oils; petroleum jelly; paraffin wax; and pitch, or tar. A host of other organic chemicals, including various drugs, plastics, paints, adhesives, fibers, detergents, synthetic rubber, and agricultural chemicals, owe their existence to petrochemicals.

Then there are the many hydrocarbon derivatives formed by the bonding of hydrocarbons to various functional groups—broad arrays of molecule types involving other elements. Among these are alcohols—both ethanol (the alcohol in beer and other drinks) and methanol, used in adhesives, fibers, and plastics, and as a fuel. Other functional groups include aldehydes, ketones, carboxylic acids, and esters. Products of these functional groups range from aspirin to butyric acid, which is in part responsible for the smell both of rancid butter and human sweat. Hydrocarbons also form the basis for polymer plastics such as Nylon and Teflon.

Hydrogen for Transportation and Power

We have already seen that hydrogen is a component of petroleum, and that hydrogen is used in creating nuclear power—both deadly and peaceful varieties. But hydrogen has been applied in many other ways in the transportation and power industries.

There are only three gases practical for lifting a balloon: hydrogen, helium, and hot air. Each is much less dense than ordinary air, and this gives them their buoyancy. Because hydrogen is the lightest known gas and is relatively cheap to produce, it initially seemed the ideal choice, particularly for airships, which made their debut near the end of the nineteenth century.

For a few decades in the early twentieth century, airships were widely used, first in warfare and later as the equivalent of luxury liners in the skies. One of the greatest such craft was Germany's Hindenburg, which used hydrogen to provide buoyancy. Then, on May 6, 1937, the Hindenburg caught fire while mooring at Lakehurst, New Jersey, and 36 people were killed—a tragic and dramatic event that effectively ended the use of hydrogen in airships.

Adding to the pathos of the Hindenburg crash was the voice of radio announcer Herb Morrison, whose audio report has become a classic of radio history. Morrison had come to Lakehurst to report on the landing of the famous airship, but ended up with the biggest—and most horrifying—story of his career. As the ship burst into flames, Morrison's voice broke, and he uttered words that have become famous:"Oh, the humanity!"

Half a century later, a hydrogen-related disaster destroyed a craft much more sophisticated than the Hindenburg, and this time, the medium of television provided an entire nation with a view of the ensuing horror. The event was the explosion of the space shuttle Challenger on January 28, 1986, and the cause was the failure of a rubber seal in the shuttle's fuel tanks. As a result, hydrogen gas flooded out of the craft and straight into the jet of flame behind the rocket. All seven astronauts aboard were killed.

The Future of Hydrogen Power

Despite the misfortunes that have occurred as a result of hydrogen's high flammability, the element nonetheless holds out the promise of cheap, safe power. Just as it made possible the fusion, or hydrogen, bomb—which fortunately has never been dropped in wartime, but is estimated to be many hundreds of times more lethal than the fission bombs dropped on Japan—hydrogen may be the key to the harnessing of nuclear fusion, which could make possible almost unlimited power.

A number of individuals and agencies advocate another form of hydrogen power, created by the controlled burning of hydrogen in air. Not only is hydrogen an incredibly clean fuel, producing no by-products other than water vapor, it is available in vast quantities from water. In order to separate it from the oxygen atoms, electrolysis would have to be applied—and this is one of the challenges that must be addressed before hydrogen fuel can become a reality.

Electrolysis requires enormous amounts of electricity, which would have to be produced before the benefits of hydrogen fuel could be realized. Furthermore, though the burning of hydrogen could be controlled, there are the dangers associated with transporting it across country in pipelines. Nonetheless, a number of advocacy groups—some of whose Web sites are listed below—continue to promote efforts toward realizing the dream of nonpolluting, virtually limitless, fuel.

Where to Learn More

American Hydrogen Association (Web site). <http://www.clean-air.org> (June 1, 2001).

Blashfield, Jean F. Hydrogen. Austin, TX: Raintree Steck-Vaughn, 1999.

Farndon, John. Hydrogen. New York: Benchmark Books, 2001.

"Hydrogen" (Web site). <http://pearl1.lanl.gov/periodic/elements/1.html> (June 1, 2001).

Hydrogen Energy Center (Web site). <http://www.h2eco.org/> (June 1, 2001).

Hydrogen Information Network (Web site). <http://www.eren.doe.gov/hydrogen/> (June 1, 2001).

Knapp, Brian J. Carbon Chemistry. Illustrated by David Woodroffe. Danbury, CT: Grolier Educational, 1998.

Knapp, Brian J. Elements. Illustrated by David Woodroffe and David Hardy. Danbury, CT: Grolier Educational, 1996.

National Hydrogen Association (Web site). <http://www.ttcorp.com/nha/> (June 1, 2001).

Uehling, Mark. The Story of Hydrogen. New York: Franklin Watts, 1995.

Sci-Tech Encyclopedia: Hydrogen

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The first chemical element in the periodic system. Under ordinary conditions it is a colorless, odorless, tasteless gas composed of diatomic molecules, H₂. The hydrogen atom, symbol H, consists of a nucleus of unit positive charge and a single electron. It has atomic number 1 and an atomic weight of 1.00797. The element is a major constituent of water and all organic matter, and is widely distributed not only on the Earth but throughout the universe. There are three isotopes of hydrogen: protium, mass 1, makes up 99.98% of the natural element; deuterium, mass 2, makes up about 0.02%; and tritium, mass 3, occurs in extremely small amounts in nature but may be produced artificially by various nuclear reactions. See also Deuterium; Periodic table; Tritium.

Physical properties

Ordinary hydrogen has an atomic weight of 1.00797, and a molecular weight of 2.01594. The gas has a density at 0°C (32°F) and 1 atm (10⁵ pascals) of 0.08987 g/liter (5.610 × 10⁻³ lb/ft³). Its specific gravity, compared to air, is 0.0695. The lightest substance known, it has a buoyancy in air of 1.203 g/liter (7.510 × 10⁻² lb/ft³) [see table].

Properties of hydrogen*
Property	Value
Melting point	−259.34°C
Boiling point at 1 atm	−252.87°C
Density of solid at −259.34°C	0.0858 g/cm³
Density of liquid at −252.87°C	0.0708 g/cm³
Critical temperature	−240.17°C
Critical pressure	12.8 atm
Critical density	0.0312 g/cm³
Specific heat at constant pressure
Gas at 25°C	3.42 cal/(g)(°C)
Liquid at −256°C	1.93 cal/(g)(°C)
Solid at −259.8°C	0.63 cal/(g)(°C)
Heat of fusion at −259.34°C	13.9 cal/g
Heat of vaporization at −252.87°C	107 cal/g
Thermal conductivity at 25°C	0.000444 cal/(cm)(s)(°C)
Viscosity at 25°C	0.00892 centipoise

*The low-temperature properties all refer to hydrogen which has the para-hydrogen concentration corresponding to its low-temperature equilibrium value.

Hydrogen dissolves in water to the extent of 0.0214 volume per volume of water at 0°C (32°F), 0.018 volume at 20°C (68°F), and 0.016 volume at 50°C (122°F). It is somewhat more soluble in organic solvents, and 0.078 volume dissolves in 1 volume of ethanol at 25°C (77°F). Many metals absorb hydrogen. Palladium is particularly notable in this respect, and dissolves about 1000 times its volume of the gas. The adsorption of hydrogen in steel may cause “hydrogen embrittlement,” which sometimes leads to the failure of chemical processing equipment.

The hydrogen atom has an ionization potential of 13.54 volts. The hydrogen nucleus (proton, mass 1) has a spin of ½ℏ and a magnetic moment of 2.79270 nuclear magnetons. Its absorption cross section for thermal neutrons is 0.332 × 10⁻²⁴ cm².

Chemical properties

At ordinary temperatures hydrogen is a comparatively unreactive substance unless it has been activated in some manner, for example, by a suitable catalyst. At elevated temperatures it is highly reactive.

Although ordinarily diatomic, molecular hydrogen dissociates at high temperatures into free atoms. Atomic hydrogen is a powerful reducing agent, even at room temperature. It reacts with the oxides and chlorides of many metals, including silver, copper, lead, bismuth, and mercury, to produce the free metals. It reduces some salts, such as nitrates, nitrites, and cyanides of sodium and potassium, to the metallic state. It reacts with a number of elements, both metals and nonmetals, to yield hydrides such as sodium hydride (NaH), potassium hydride (KH), and phosphorus hydride (PH₃). Sulfur forms a number of hydrides; the simplest is H₂S. With oxygen atomic hydrogen yields hydrogen peroxide, H₂O₂. With organic compounds atomic hydrogen reacts to produce a complex mixture of products. With ethylene, C₂H₄, for example, the products include ethane, C₂H₆, and butane C₄H₁₀. The heat liberated when hydrogen atoms recombine to form hydrogen molecules is used to obtain very high temperatures in atomic hydrogen welding.

Hydrogen reacts with oxygen to form water. At room temperature this reaction is immeasurably slow, but is accelerated by catalysts, such as platinum, or by an electric spark.

With nitrogen, hydrogen reacts to give ammonia. Hydrogen reacts at elevated temperatures with a number of metals, including lithium, sodium, potassium, calcium, strontium, and barium, to give hydrides. See also Metal hydrides.

Principal compounds

Hydrogen is a constituent of a very large number of compounds containing one or more other elements. Such compounds include water, acids, bases, most organic compounds, and many minerals. Compounds in which hydrogen is combined with a single other element are commonly referred to as hydrides. These may be divided into three general classes: the ionic or saltlike hydrides, the covalent or molecular hydrides, and the transition metal hydrides. See also Hydride.

Preparation

A large number of methods may be used to prepare hydrogen gas. The choice of method is determined by such factors as the quantity of hydrogen desired, the purity required, and the availability and cost of raw materials. Among the processes frequently used are the reactions of metals with water or acides, the electrolysis of water, the reaction of steam with hydrocarbons or other organic materials, and the thermal decomposition of hydrocarbons.

Uses

The largest single use of hydrogen is in the synthesis of ammonia. Ammonia plants are often built adjacent to petroleum refineries or coking plants to utilize by-product hydrogen that might otherwise be wasted. An important use for hydrogen is in petroleum-refining operations, such as hydrocracking and hydrogen treatment for removal of sulfur. Large quantities of hydrogen are consumed in the catalytic hydrogenation of unsaturated liquid vegetable oils to make solid fats. Hydrogenation is used in the manufacture of organic chemicals, such as alcohols from esters and glycerides, amines from nitriles, and cycloparaffins from aromatic hydrocarbons. Methanol is produced commercially by reaction of hydrogen with carbon monoxide. Reaction of hydrogen with chlorine is a major source of hydrochloric acid.

Liquid hydrogen

Liquid hydrogen, a clear colorless fluid which boils at −252.87°C (−423.17°F) and has the smallest boiling point density of any known liquid, was first produced in 1898. James Dewar was successful with this experiment because his development of the glass vacuum flask enabled him to retain this low-boiling-point liquid. For many years, hydrogen was liquefied by using high-pressure processes that are dangerous, since small air (oxygen) impurities in the hydrogen can result in an explosion. Modern liquefiers, however, generally use closed-circuit helium refrigeration cycles which condense the hydrogen at low pressure, and the hazards are considerably reduced.

The density of liquid hydrogen is roughly 790 times greater than that of the gas under normal conditions, so quantities of hydrogen in liquid form are transported conveniently and economically in relatively light-weight vacuum-jacketed containers that are open to the atmosphere. The alternative, to compress the gas at room temperature into heavy-walled steel containers, involves much less shipping efficiency. The space program uses large amounts of hydrogen for rocket fuels, in conjunction with oxygen or fluorine, and here the ability to transport hydrogen in liquid form with minimal insulation has great advantages. Liquid (and solid) hydrogen (and its heavier isotopes, deuterium and tritium) can be used for the production of electrical power by nuclear fusion.

Spin-polarized atomic hydrogen

Spin-polarized atomic hydrogen is a special preparation of the element which is expected to remain a gas down to the absolute zero of temperature. It is created by dissociating the hydrogen molecule (H₂), exposing the atoms to a high magnetic field which aligns the electronic spins, and storing them at temperatures so low that thermal effects are unlikely to disturb the spin alignment. Typically, spin-polarized hydrogen is stored at 0.3 K in a magnetic field of 8 tesla.

Spin-polarized hydrogen has been used to study the quantum theory of weakly interacting gases and has had applications in the areas of polarized proton sources for high-energy physics and cryogenic hydrogen masers. The ultimate goal is to observe the Bose-Einstein phase transition to a state in which a finite fraction of the atoms in the gas essentially come to rest. The gas in this state is also expected to be a superfluid. The transition occurs when the temperature of the gas is so slow that the quantum-mechanical wavelength associated with the particles becomes comparable to the interparticle spacing.

Dental Dictionary: hydrogen

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n
H

A gaseous, univalent element. Its atomic number is 1 and its atomic weight is 1.008. It is the simplest and lightest of the elements and is normally a colorless, odorless, highly flammable diatonic gas.

Britannica Concise Encyclopedia: hydrogen

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Lightest chemical element, chemical symbol H, atomic number 1. A colourless, odourless, tasteless, flammable gas, it occurs as the diatomic molecule H₂. Its atom consists of one proton (the nucleus) and one electron; the isotopes deuterium and tritium have an additional one and two nuclear neutrons, respectively. Though only the ninth most abundant element on Earth, it represents about 75% of all matter in the universe. Hydrogen was formerly used to fill airships; nonflammable helium has replaced it. It is used to synthesize ammonia, ethanol, aniline, and methanol; to treat petroleum fuels; as a reducing agent (see reduction) and to supply a reducing atmosphere; to make hydrogen chloride (see hydrochloric acid) and hydrogen bromide; and in hydrogenation (e.g., of fats). Liquid hydrogen (boiling point -423 °F [-252.8 °C]) is used in scientific and commercial applications to produce extremely low temperatures and as a rocket propellant and a fuel for fuel cells. Combustion of hydrogen with oxygen gives water as the sole product. The properties of most acids, especially in water solutions, arise from the hydrogen ion (H⁺, also referred to as the hydronium ion, H₃O⁺, the form in which H⁺ is found in a water environment). See also hydride; hydrocarbon.

For more information on hydrogen, visit Britannica.com.

Columbia Encyclopedia: hydrogen

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hydrogen (hī'drəjən) [Gr.,=water forming], gaseous chemical element; symbol H; at. no. 1; at. wt. 1.00794; m.p. −259.14°C; b.p. −252.87°C; density 0.08988 grams per liter at STP; valence usually +1.

The Isotopes and Forms

Atmospheric hydrogen is a mixture of three isotopes. The most common is called protium (mass no. 1, atomic mass 1.007822); the protium nucleus (protium ion) is a proton. A second isotope of hydrogen is deuterium (mass no. 2, atomic mass 2.0140), the so-called heavy hydrogen, often represented in chemical formulas by the symbol D. The deuterium nucleus, or ion, is called the deuteron; it consists of a proton plus a neutron. The two isotopes are found in atmospheric hydrogen in the proportion of about 1 atom of deuterium to every 6,700 atoms of protium. Protium and deuterium differ slightly in their chemical and physical properties; for example, the boiling point of deuterium is about 3°C lower than protium. The properties of compounds they form differ depending on the ratio of the two isotopes present.

Deuterium oxide (D₂O), the so-called heavy water, is present in ordinary water; the concentration of deuterium oxide is increased by electrolysis of the water. The melting point (3.79°C), boiling point (101.4°C), and specific gravity (1.107 at 25°C) of deuterium oxide are higher than those of ordinary water. Deuterium oxide is used as a moderator in nuclear reactors. Deuterium is also of importance because of the wide use it has found in scientific research; for example, chemical reaction mechanisms have been studied by the use of deuterium atoms as tracers (i.e., deuterium is substituted for atoms of ordinary hydrogen in compounds), making it possible to follow the course of individual molecules in a reaction.

Tritium (mass no. 3, atomic mass 3.016), a third hydrogen isotope, is a radioactive gas with a half-life of about 12¹/₄ years; it is often represented in chemical formulas by the symbol T. It is produced in nuclear reactors and occurs to a very limited extent in atmospheric hydrogen. It is used in the hydrogen bomb, in luminous paints, and as a tracer. The tritium nucleus, or ion, is called the triton; it consists of a proton plus two neutrons. Tritium oxide (T₂O) has a melting point (4.49°C) higher than that of deuterium oxide.

Besides being a mixture of three isotopes, hydrogen is a mixture of two forms, an ortho form and a para form, which differ in their electronic and nuclear spins. At room temperature atmospheric hydrogen is about ³/₄ ortho-hydrogen and ¹/₄ para-hydrogen. The two forms differ slightly in their physical properties.

Properties

Under ordinary conditions hydrogen is a colorless, odorless, tasteless gas that is only slightly soluble in water; it is the least dense gas known. It is the first element in Group 1 of the periodic table. Ordinary hydrogen gas is made up of diatomic molecules (H₂) that react with oxygen to form water (H₂O) and hydrogen peroxide (H₂O₂), usually as a result of combustion. A jet of hydrogen burns in air with a very hot blue flame. The flame produced by a mixture of oxygen and hydrogen gases (as in the oxyhydrogen blowpipe) is extremely hot and is used in welding and to melt quartz and certain glasses. Hydrogen gas must be used with caution because it is highly flammable; it forms easily ignited explosive mixtures with oxygen or with air (because of the oxygen in the air). At high temperatures hydrogen is a chemically active mixture of monohydrogen (atomic hydrogen) and the normal diatomic hydrogen (see allotropy).

Hydrogen has a great affinity for oxygen and is a powerful reducing agent (see oxidation and reduction). It reacts with nitrogen to form ammonia. With the halogens it forms compounds (hydrogen halides) that are strongly acidic in water solution. With sulfur it forms hydrogen sulfide (H₂S), a colorless gas with an odor like rotten eggs; with sulfur and oxygen it forms sulfuric acid. It combines with several metals to form metal hydrides such as calcium hydride. Combined with carbon (and usually other elements) it is a constituent of a great many organic compounds, such as hydrocarbons, carbohydrates, fats, oils, proteins, and organic acids and bases.

It is theoretically possible for hydrogen to exhibit the properties of a metal, such as electrical conductivity. Although researchers have been able to squeeze hydrogen into liquid and crystalline solid states through applications of intense heat, cold, and pressure, the metallic form eluded them until 1996. By compressing liquid hydrogen to nearly 2 million atmospheres pressure and a temperature of 4,400K, a team at the Lawrence Livermore National Laboratory created metallic hydrogen for a millionth of a second. While there is no practical application for the accomplishment, proof of the existence of a metallic form of hydrogen may have implications for theories of how Jupiter's magnetic field is produced.

Sources and Commercial Preparation

While hydrogen is only about one part per million in the atmosphere, it is the most abundant element in the universe. It is believed that hydrogen makes up about three quarters of the mass of the universe, or over 90% of the molecules. It is found in the sun and in other stars, where it is the major fuel in the fusion reactions (see nucleosynthesis) from which stars derive their energy.

Hydrogen is prepared commercially by catalytic reaction of steam with hydrocarbons, by the reaction of steam with hot coke (carbon), by the electrolysis of water, and by the reaction of mineral acids on metals. Millions of cubic feet of hydrogen gas are produced daily in the United States alone.

Uses

Hydrogen was formerly used for filling balloons, airships, and other lighter-than-air craft, a dangerous practice because of hydrogen's explosive flammability; there were disastrous fires, e.g., the immolation of the German airship Hindenburg at its mooring at Lakehurst, N.J., in 1937. Helium is preferable for use in lighter-than-air craft since it is not flammable. Hydrogen is used in the Haber process for the fixation of atmospheric nitrogen, in the production of methanol, and in hydrogenation of fats and oils. It is also important in low-temperature research. It can be liquefied under pressure and cooled; when the pressure is released, rapid evaporation takes place and some of the hydrogen solidifies.

Discovery of Hydrogen and Its Isotopes

Although hydrogen was prepared many years earlier, it was first recognized as a substance distinct from other flammable gases in 1766 by Henry Cavendish, who is credited with its discovery; it was named by A. L. Lavoisier in 1783. Deuterium was discovered by H. C. Urey, F. G. Brickwedde, and G. M. Murphy in 1932, although its existence had been suspected for some years. Deuterium oxide was also discovered by Urey and was first obtained in nearly pure form by G. N. Lewis. Tritium was synthesized by Ernest Rutherford, L. E. Oliphant, and Paul Harteck in 1935.

Science Dictionary: hydrogen

Top

The lightest chemical element; its symbol is H. Hydrogen normally consists of a single electron in orbit around a nucleus made up of a single proton. It is usually found as a gas and has several uses as a fuel.

Hydrogen atoms are combined to form helium atoms in fusion reactions in stars and in hydrogen bombs, which release huge amounts of energy. Hydrogen also burns rapidly, producing water as it combines with oxygen (see H and oxidation).

For a time, hydrogen was frequently used to fill blimps and dirigibles because of its extremely low weight. In 1937, however, the hydrogen in the dirigible Hindenburg caught fire, and many of the passengers and crew were killed. Since that time, helium has been widely preferred to hydrogen for use in airships; it is not as buoyant (see buoyancy) or cheap as hydrogen, but, being an inert gas, it does not burn.

Because there is so much hydrogen in stars, it is by far the most abundant element in the universe.

Veterinary Dictionary: hydrogen

Top

A chemical element, atomic number 1, atomic weight 1.00797, symbol H. It exists as the mass 1 isotope (protium, or light or ordinary hydrogen), mass 2 isotope (deuterium, heavy hydrogen), and mass 3 isotope (tritium).

h. bonding — weak electrostatic attraction between one electronegative atom and the hydrogen atom covalently linked to a second electronegative atom.
h. breath test — detects hydrogen production as a product of bacterial fermentation of carbohydrates, an indicator of inflammatory bowel disease or carbohydrate malabsorption.
h. cyanide — hydrocyanic acid.
heavy h. — hydrogen having double the mass of ordinary hydrogen; deuterium.
h. ion balance — see acid–base balance.
h. ion concentration — the degree of concentration of hydrogen ions (the acid element) in a solution. Its symbol is pH, and expresses the degree to which a solution is acidic or alkaline. The pH range extends from 0 to 14, pH 7 being neutral. A pH of less than 7 indicates acidity, above 7 indicates alkalinity. See also acid–base balance and ph.
h. peroxide — H₂O₂, used in solution as an antibacterial agent. A 3% solution foams on touching skin or mucous membrane and appears to have a mechanical cleansing action.
h. peroxide-based teat dips — see teat dip.
h. sulfide — an ill-smelling, colorless, poisonous gas, H₂S; much used as a chemical reagent. Hydrogen sulfide is often present in gases from oil wells and from manure vats under slatted floor barns. Poisoning of cattle causes diarrhea, dehydration, dyspnea and death in convulsions. The feces are black and the breath smells of hydrogen sulfide. Called also hydrosulfuric acid. See also manure pit gas poisoning.
h. swell — defective canned meat can. Can is distended due to production of hydrogen as a result of corrosion of the can surface.

Cosmic Lexicon: Hydrogen

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An element with atomic number 1; symbol: H. It is the most abundant element in the solar system, making up 90 percent of the Sun. Hydrogen, carbon, nitrogen, and oxygen are essential for life.

Word Tutor: hydrogen

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IN BRIEF: A gas that has no color or smell, burns very easily and is the lightest of all known substances.

Some scientists claim that hydrogen, because it is so plentiful, is the basic building block of the universe. — Frank Zappa (1940-1993)

Wikipedia: Hydrogen

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This article is about the chemistry of hydrogen. For the physics of atomic hydrogen, see Hydrogen atom. For other meanings, see Hydrogen (disambiguation).

neutronium ← hydrogen → helium

e⁻
↑
H
↓
Li

Periodic table

colorless gas

General

hydrogen, H, 1

1, 1, s

Standard atomic weight

1.00794 (7) g·mol⁻¹

Electron configuration

1s¹

Electrons per shell

1 (Image)

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 (-259°C), 7.042 kPa

Critical point

32.97 K, 1.293 MPa

Heat of fusion

(H₂) 0.117 kJ·mol⁻¹

Heat of vaporization

(H₂) 0.904 kJ·mol⁻¹

Specific heat capacity

(25 °C) (H₂)
28.836 J·mol⁻¹·K⁻¹

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)

2.20 (Pauling scale)

1st: 1312.0 kJ·mol⁻¹

31±5 pm

120 pm

Miscellaneous

Magnetic ordering

diamagnetic^[1]

Thermal conductivity

(300 K) 180.5 m W·m⁻¹·K⁻¹

Speed of sound

(gas, 27 °C) 1310 m/s

CAS registry number

1333-74-0

Most stable isotopes

Main article: Isotopes of hydrogen
iso	NA	half-life	DM	DE (MeV)	DP
¹H	99.985%	¹H is stable with 0 neutrons
²H	0.015%	²H is stable with 1 neutron
³H	trace	12.32 y	β⁻	0.01861	³He

References

Hydrogen (pronounced /ˈhaɪdrədʒən/^[2]) is the chemical element with atomic number 1. It is represented by the symbol H. At standard temperature and pressure, hydrogen is a colorless, odorless, nonmetallic, tasteless, highly flammable diatomic gas with the molecular formula H₂. With an atomic weight of 1.00794 u, hydrogen is the lightest element.

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. Elemental hydrogen is relatively rare on Earth. Industrial production is from hydrocarbons such as methane with most being used "captively" at the production site. The two largest uses are in fossil fuel processing (e.g., hydrocracking) and ammonia production mostly for the fertilizer market. Hydrogen may be produced from water by electrolysis at substantially greater cost than production from natural gas.^[4]

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 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 is important in metallurgy as it can embrittle many metals,^[5] complicating the design of pipelines and storage tanks.^[6] Hydrogen is highly soluble in many rare earth and transition metals^[7] and is soluble in both nanocrystalline and amorphous metals.^[8] Hydrogen solubility in metals is influenced by local distortions or impurities in the crystal lattice.^[9]

Combustion

The Space Shuttle Main Engine burns hydrogen with oxygen, producing a nearly invisible flame

Hydrogen gas (dihydrogen^[10]) is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume.^[11] The enthalpy of combustion for hydrogen is −286 kJ/mol:^[12]

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

Hydrogen-oxygen mixtures are explosive across a wide range of proportions. Its autoignition temperature, the temperature at which it ignites spontaneously in air, is 560 °C (1,040 °F).^[13] 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 explosion 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.^[14] 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.^[15]

H₂ reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding halides: hydrogen chloride and hydrogen fluoride.^[16]

Electron energy levels

Main article: Hydrogen atom

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.^[17]

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.^[18]

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.^[19]

Elemental molecular forms

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); in these compounds hydrogen takes on a partial positive charge.^[28] 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.^[29]^[30] 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.^[31]

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;^[32] the study of their properties is known as organic chemistry^[33] and their study in the context of living organisms is known as biochemistry.^[34] 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.^[32]

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.^[35]

Hydrides

Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. The term "hydride" implies 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.^[36] 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.^[37] Binary indium hydride has not yet been identified, although larger complexes exist.^[38]

Protons and acids

Further information: Acid-base reaction

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 H₉O₄⁺.^[39] Other oxonium ions are found when water is in solution with other solvents.^[40]

Although exotic on earth, one of the most common ions in the universe is the H₃⁺ ion, known as protonated molecular hydrogen or the triatomic hydrogen cation.^[41]

Isotopes

Main article: Isotopes of hydrogen

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.^[42]^[43]

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.^[44]
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.^[45] Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.^[46]
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.^[35] 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.^[47] It is used in nuclear fusion reactions,^[48] as a tracer in isotope geochemistry,^[49] and specialized in self-powered lighting devices.^[50] Tritium has also been used in chemical and biological labeling experiments as a radiolabel.^[51]

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.^[52] 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.^[53]

Natural occurrence

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.^[54] 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.^[55]

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.^[56]

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 (in chemically combined form) is the third most abundant element on the Earth's surface.^[57] Most of the Earth's hydrogen is in the form of chemical compounds such as hydrocarbons and water.^[35] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus. Methane is a hydrogen source of increasing importance.^[58]

History

Discovery and use

Main article: Timeline of hydrogen technologies

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.^[59] 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.^[60] In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by identifying the gas from a metal-acid reaction as "inflammable air" and further finding in 1781 that the gas produces water when burned. He is usually given credit for its discovery as an element.^[61]^[62] In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek hydro meaning water and genes meaning creator)^[63] when he and Laplace reproduced Cavendish's finding that water is produced when hydrogen is burned.^[62]

Hydrogen was liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask.^[62] He produced solid hydrogen the next year.^[62] Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck.^[61] Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932.^[62] François 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 Döbereiner's lamp and limelight were invented in 1823.^[62]

The first hydrogen-filled balloon was invented by Jacques Charles in 1783.^[62] 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.^[62] 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.^[62] 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.^[62] 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,^[64] 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).^[65] For example, the ISS,^[66] Mars Odyssey^[67] and the Mars Global Surveyor^[68] 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 of any NiH2 battery in low earth orbit.^[69]

Role in quantum theory

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.^[70] 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.^[71]

Production

For more details on this topic, see Hydrogen production.

H₂ is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions.

Laboratory

In the laboratory, H₂ is usually prepared by the reaction of acids on metals such as zinc 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%.^[72]

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.^[73]

Industrial

Main article: Hydrogen production

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.^[74] At high temperatures (700–1100 °C; 1,300–2,000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H₂.

CH₄ + H₂O → CO + 3 H₂

This reaction is favored at low pressures but is nonetheless conducted at high pressures (20 atm; 600 inHg) since high pressure H₂ is the most marketable product. The product mixture is known as "synthesis gas" because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon:

CH₄ → C + 2 H₂

Consequently, steam reforming typically employs an excess of H₂O. Additional hydrogen can be recovered from the steam by use of carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide:^[74]

CO + H₂O → CO₂ + H₂

Other important methods for H₂ production include partial oxidation of hydrocarbons:^[75]

2 CH₄ + O₂ → 2 CO + 4 H₂

and the coal reaction, which can serve as a prelude to the shift reaction above:^[74]

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.^[76] Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.^[77]

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.^[78] 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.^[79]

Applications

Large quantities of H₂ are needed in the petroleum and chemical industries. The largest application of H₂ is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H₂ in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H₂ has several other important uses. H₂ is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H₂ is also used as a reducing agent of metallic ores.^[80]

Apart from its use as a reactant, H₂ has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding.^[81]^[82] 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.^[83] 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 gas in balloons and airships.^[84]

In more recent applications, hydrogen is used pure or mixed with nitrogen (sometimes called forming gas) as a tracer gas for minute leak detection. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries.^[85] Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties.^[86]

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.^[62] Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects.^[87] Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs,^[88] as an isotopic label in the biosciences,^[51] and as a radiation source in luminous paints.^[89]

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

Energy carrier

Hydrogen is not an energy resource,^[91] except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development.^[92] The Sun's energy comes from nuclear fusion of hydrogen, but this process is difficult to achieve controllably on Earth.^[93] Elemental hydrogen from solar, biological, or electrical sources require more energy to make it than is obtained by burning it, so in these cases hydrogen functions as an energy carrier, like a battery. Hydrogen may be obtained from fossil sources (such as methane), but these sources are unsustainable.^[91]

The energy density per unit volume of both liquid hydrogen and compressed hydrogen gas at any practicable pressure is significantly less than that of traditional fuel sources, although the energy density per unit fuel mass is higher.^[91] Nevertheless, elemental hydrogen has been widely discussed in the context of energy, as a possible future carrier of energy on an economy-wide scale.^[94] For example, CO₂ sequestration followed by carbon capture and storage could be conducted at the point of H₂ production from fossil fuels.^[95] Hydrogen used in transportation would burn relatively cleanly, with some NOx emissions,^[96] but without carbon emissions.^[95] However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial.^[97]

Biological reactions

Further information: Biohydrogen and Biological hydrogen production

H₂ is a product of some types of anaerobic metabolism and is produced by several microorganisms, usually via reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H₂ and its component two protons and two electrons. Creation of hydrogen gas occurs in the transfer of reducing equivalents produced during pyruvate fermentation to water.^[98]

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.^[99] Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H₂ gas even in the presence of oxygen.^[100] Efforts have also been undertaken with genetically modified alga in a bioreactor.^[101]

Safety and precautions

Main article: Hydrogen safety

Inhalation of air with high concentration of hydrogen as with hydrox and hydreliox displaces oxygen and may cause the above symptoms as an asphyxant.^[102]

Hydrogen poses a number of hazards to human safety, from potential detonations and fires when mixed with air to being an asphyxant in its pure, oxygen-free form.^[103] In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids.^[104] Hydrogen dissolves in some metals, and, in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement.^[105] Hydrogen gas leaking into external air may spontaneously ignite. Moreover, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to accidental burns.^[106]

Even interpreting the hydrogen data (including safety data) is confounded by a number of phenomena. Many physical and chemical properties of hydrogen depend on the parahydrogen/orthohydrogen ratio (it often takes days or weeks at a given temperature to reach the equilibrium ratio, for which the data is usually given). Hydrogen detonation parameters, such as critical detonation pressure and temperature, strongly depend on the container geometry.^[103]

Notes

^ 286 kJ/mol: energy per mole of the combustible material (hydrogen)

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^ Heffel, James W. (2002-12-24). "NOx emission and performance data for a hydrogen fueled internal combustion engine at 1500 rpm using exhaust gas recirculation". International Journal of Hydrogen Energy (Riverside, CA: University of California) 28 (8): 901–908. doi:10.1016/S0360-3199(02)00157-X. http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6V3F-47HJVY6-8&_user=10&_rdoc=1&_fmt=&_orig=search&_sort=d&view=c&_acct=C000050221&_version=1&_urlVersion=0&_userid=10&md5=bbc8c5bce46f1d4ba3a814f5c828ee80. Retrieved on 2008-03-16.
^ See Romm, Joseph J. (2004). The Hype About Hydrogen: Fact And Fiction In The Race To Save The Climate (1st edition ed.). Island Press. ISBN 155963703X.
^ Cammack, Richard; Robson, R. L. (2001). Hydrogen as a Fuel: Learning from Nature. Taylor & Francis Ltd. ISBN 0415242428.
^ Kruse, O.; Rupprecht, J.; Bader, K.-P.; Thomas-Hall, S.; Schenk, P. M.; Finazzi, G.; Hankamer, B (2005). "Improved photobiological H₂ production in engineered green algal cells". The Journal of Biological Chemistry 280 (40): 34170–7. doi:10.1074/jbc.M503840200. PMID 16100118.
^ Smith, H. O.; Xu, Q (2005). "IV.E.6 Hydrogen from Water in a Novel Recombinant Oxygen-Tolerant Cyanobacteria System" (PDF). FY2005 Progress Report. United States Department of Energy. http://ec.europa.eu/food/fs/sfp/addit_flavor/flav15_en.pdf. Retrieved on 2008-02-05.
^ Williams, Chris (2006-02-24). "Pond life: the future of energy". Science (The Register). http://www.theregister.co.uk/2006/02/24/pond_scum_breakthrough/. Retrieved on 2008-03-24.
^ Lenntech: Health effects of hydrogen - Environmental effects of hydrogen Received on 11 February, 2009
^ ^a ^b Smith, H. O.; Xu, Q (1997). "Safety Standard for Hydrogen and Hydrogen Systems" (PDF). NASA. http://www.hq.nasa.gov/office/codeq/doctree/canceled/871916.pdf. Retrieved on 2008-02-05.
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E numbers

Colors (E100–199) • Preservatives (E200–299) • Antioxidants & Acidity regulators (E300–399) • Thickeners, stabilisers & emulsifiers (E400–499) • pH regulators & anti-caking agents (E500–599) • Flavour enhancers (E600–699) • Miscellaneous (E900–999) • Additional chemicals (E1100–1599)

Waxes (E900–909) • Synthetic glazes (E910–919) • Improving agents (E920–929) • Packaging gases (E930–949) • Sweeteners (E950–969) • Foaming agents (E990–999)

Calcium peroxide (E930) • Argon (E938) • Helium (E939) • Dichlorodifluoromethane (E940) • Nitrogen (E941) • Nitrous oxide (E942) • Butane (E943a) • Isobutane (E943b) • Propane (E944) • Oxygen (E948) • Hydrogen (E949)

Alkaline earth metals

Other metals

Other nonmetals

v • d • e Hydrogen compounds

HBr · HCN · HCl · HF · HI · HNC · HOF · H₂O · H₂O₂ · H₂S · H₂Se · H₂Te

v • d • e Diatomic chemical elements

Hydrogen H₂ \| Nitrogen N₂ \| Oxygen O₂ \| Fluorine F₂ \| Chlorine Cl₂ \| Bromine Br₂ \| Iodine I₂ \| Astatine At₂ \|

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Translations: Hydrogen

Top

Dansk (Danish)
n. - brint

idioms:

hydrogen bomb brintbombe
hydrogen peroxide brintoverilte

Nederlands (Dutch)
waterstof

Français (French)
n. - hydrogène

idioms:

hydrogen bomb bombe à hydrogène
hydrogen peroxide eau oxygénée

Deutsch (German)
n. - Wasserstoff

idioms:

hydrogen bomb Wasserstoffbombe
hydrogen peroxide (Chem.) Wasserstoffperoxyd

Ελληνική (Greek)
n. - (χημ.) υδρογόνο

idioms:

hydrogen bomb βόμβα υδρογόνου, υδρογονοβόμβα
hydrogen peroxide (χημ.) οξυζενέ, υπεροξείδιο υδρογόνου

Italiano (Italian)
idrogeno

idioms:

hydrogen bomb bomba a idrogeno
hydrogen peroxide acqua ossigenata

Português (Portuguese)
n. - hidrogênio (m) (Quím.)

idioms:

hydrogen bomb bomba (f) de hidrogênio (Eng. Nucl.)
hydrogen peroxide peróxido (m) de hidrogênio (Quím.)

Русский (Russian)
водород

idioms:

hydrogen bomb водородная бомба
hydrogen peroxide перекись водорода

Español (Spanish)
n. - hidrógeno

idioms:

hydrogen bomb bomba de hidrógeno
hydrogen peroxide peróxido de hidrógeno

Svenska (Swedish)
n. - väte

中文（简体）(Chinese (Simplified))
氢

idioms:

hydrogen bomb 氢弹
hydrogen peroxide 过氧化氢, 双氧水

中文（繁體）(Chinese (Traditional))
n. - 氫

idioms:

hydrogen bomb 氫彈
hydrogen peroxide 過氧化氫, 雙氧水

한국어 (Korean)
n. - 수소

日本語 (Japanese)
n. - 水素

idioms:

hydrogen bomb 水素爆弾
hydrogen peroxide 過酸化水素

العربيه (Arabic)
‏(الاسم) الهيدروجين أو الايدروجين‏

עברית (Hebrew)
n. - ‮מימן‬

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	Covalent Bonding
	Diatomic
	Duet Rule

	Which of the following substances contains molecules that will not form hydrogen bonds Hydrogen Hydrogen Fluoride Water or Ammonia? Read answer...
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	If 1 gram of hydrogen reacts with 19 grams of flourine to make hydrogen flouride how many grams of hydrogen flouride will form? Read answer...

Help us answer these

	Determine the theoretical and percent yield of hydrogen gas if 36.0g of water undergoes electrolysis to produce hydrogen and oxygen and 3.80g of hydrogen is collected?
	When a Hydrogen bond is formed there is a hydrogen bond donor and a hydrogen bond acceptor Which atoms does not allow hydrogen to behave as a good hydrogen donor?
	In a hydrogen gas experiment to collect hydrogen gas with a hydrogen generator why is it necessary to discard the first bottle of hydrogen you collected before continuing with the experiment?

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Copyrights:

		Dictionary. The American Heritage® Dictionary of the English Language, Fourth Edition Copyright © 2007, 2000 by Houghton Mifflin Company. Updated in 2007. Published by Houghton Mifflin Company. All rights reserved. Read more
		Science of Everyday Things. Science of Everyday Things. Copyright © 2002 by The Gale Group, Inc. All rights reserved. Read more

hydrogen

Contents

Combustion

Electron energy levels

Elemental molecular forms

Compounds

Covalent and organic compounds

Hydrides

Protons and acids

Isotopes

Natural occurrence

History

Discovery and use

Role in quantum theory

Production

Laboratory

Industrial

Thermochemical

Applications

Energy carrier

Biological reactions

Safety and precautions

See also

Notes

References

Further reading

External links