Dictionary:

hydrogen

  (hī'drə-jən)

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.]

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.

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].

*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.

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.

For more information on hydrogen, visit Britannica.com.

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.

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.

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.