Hydrogen Compounds

Type of physical science: Chemistry

Field of study: Chemical compounds

Hydrogen is the lightest and most abundant of all the chemical elements and forms the most compounds. It is impossible to exaggerate the importance of hydrogen compounds such as water, hydrocarbons, and carbohydrates in human life.

Overview

Hydrogen is the most abundant element in the universe and the ninth most abundant on Earth, where its most important compound, water, covers about three-fourths of the planet's surface. Only traces of uncombined hydrogen exist on Earth naturally, but the bodies of plants and animals of all sorts are made of compounds containing hydrogen combined with other light elements such as oxygen, nitrogen, and carbon. The human body is about 65 percent water by weight.

Hydrogen compounds are conveniently classified by the type of chemical bonds existing between hydrogen and the other elements in the compound. Thus, the compounds with C-H, O-H, and N-H bonds are classified as covalent hydrides, as are compounds of hydrogen with other nonmetallic elements. In the biologically important carbohydrates, fats, and proteins, hydrogen is held by a single, strong covalent bond in which a shared pair of electrons occupies the region between the hydrogen nucleus and the bonded atom.

Within the realm of covalent hydrides, the hydrogen atom can exist in a variety of subtly different bonding arrangements, all based on the shared electron pair. The differences between C-H bonds in hydrocarbons such as methane (a component of natural gas) and O-H bonds in water molecules arise partly because a carbon atom attracts the shared electrons less strongly than an oxygen atom. This effect tends to give the hydrogen atom in methane a less positive charge than the hydrogen atom in water. The O-H bond is described as more polar than the C-H bond. From the difference in bond polarity comes other differences. Less polar bonds lead to compounds that are less volatile. Thus, methane and other simple hydrocarbons are gases at room temperature, and water is liquid. The relatively negative (hydridic) hydrogen atoms in hydrocarbons are also more reactive toward oxygen, leading to the combustion reactions that are so useful. Relatively positive (protonic) hydrogen atoms, occurring in compounds like hydrogen chloride, are sometimes able to break off from the hydride molecule and form hydrogen ions (protons). The protons thus formed have their own unique chemical behavior, and can accept electrons from active metals such as zinc or iron, leading to production of hydrogen gas. Hydrides that are able to produce protons are called acids and function best in water solution, where the protons can be stabilized by water molecules clustering around them.

The chemist's proton is really a hydrated proton that includes a cluster of water molecules, and bears little resemblance to the proton of physics.

Hydrogen atoms usually form one strong covalent bond, but they can, under the right circumstances, bond somewhat more weakly to one additional atom of nitrogen or oxygen in a neighboring molecule, or even in a neighboring part of the same molecule. This weak form of bonding (about 10 percent as strong as a normal covalent bond) is called a hydrogen bond and can produce interesting changes in physical properties. For example, in the water molecule, an oxygen atom is strongly bonded to each of two hydrogen atoms. In water vapor at low pressure, the water molecules are far apart from one another, and do not interact. As the vapor is compressed and cooled, however, the molecules approach closely and begin to clump together, with the hydrogen of one molecule bonding via a hydrogen bond to the oxygen of the next molecule. The effect of the hydrogen bonding is to allow the water vapor to liquefy at a higher temperature than it otherwise might have. In comparison with other hydrogen compounds, water is much less volatile. As cooling continues, the water freezes, producing ice crystals, which possess a tetrahedral pattern of hydrogen atoms around each oxygen. This unique structure results in ice having a lower density than liquid water, so ice floats.

Hydrogen bonds are also responsible for many of the structural features of proteins and nucleic acids. These molecules would in many cases exist as long, floppy chains if it were not for hydrogen bonds. The deoxyribonucleic acid (DNA) in cells forms helical molecules in which the unique shape is held together by hydrogen bonds between one turn of the helix and the next.

High temperatures or harsh chemicals can cause the hydrogen bonds to break, leading to loss of the helical shape, and with it the vital ability to replicate and pass on genetic information.

Covalent hydrides, whether polar or not, can form only if the hydrogen shares a pair of electrons with the other element. Many metals fail to attract electrons well enough to be able to form covalent hydrides. For these metals, other bonding modes are available.

The most reactive metallic elements, such as sodium and potassium, combine with hydrogen to produce solid, saltlike hydrides containing a hydride ion. These hydrides react violently with water or with any substance capable of the slightest acidic behavior. Few metals, however, are active enough electron donors to be able to produce an ionic hydride of this sort. A more common metal-hydrogen interaction is the formation of metallic hydrides, in which small hydrogen atoms penetrate and expand the atomic structures of metals like titanium, and produce compounds with a range of compositions, often resembling solid solutions of hydrogen in metal.

Metals such as nickel, palladium, and platinum interact with molecular hydrogen and activate it toward chemical reactions. Most of the commercially important reactions of molecular hydrogen would be impractical without this type of catalysis.

Intact hydrogen molecules may bond to metals like chromium, tungsten, and iridium in suitable complex molecules. The hydrogen molecule is held in such a manner that the individual hydrogen atoms are equidistant from the metal atom. This arrangement of atoms may also occur when hydrogen molecules are adsorbed on a metal surface, and can be regarded as a preliminary step toward the activation of hydrogen molecules for participation in other chemical reactions.

Studies of structure and bonding in molecular hydrogen-metal compounds are helpful in understanding catalysis of hydrogenation reactions by metals.

Applications

The hydrides of carbon, nitrogen, and oxygen are of overwhelming importance in practical applications, and are used in enormous quantities. Hydrides of other elements have smaller-scale uses based on their unique properties as acids, bases, catalysts, or hydrogen carriers.

Any account must begin with water, which is the commonest and most useful of the hydrides, has been known from time immemorial, and was considered an element by the ancients. Although water is plentiful on earth, more than 90 percent of the supply is salty, and most of the fresh water exists in perpetually frozen condition at the poles. Agriculture, industry, and individual families in the United States consume water at a rate of approximately 250 billion gallons per day, creating the need for a purification and delivery system of great proportions.

Good sanitation and prevention of water-borne diseases such as cholera and typhoid fever require analysis of the water supply, in both chemical and microbiological terms, and application of suitable treatment (for example, chlorination) to eliminate bacteria.

Many uses of water are related to its high specific-heat capacity. Of all substances, water absorbs a unit of heat with the least rise in temperature. Water cooling is essential in power plants, air-conditioning, and a multitude of industrial operations. Water puts out fires by absorbing heat and reducing the temperature of the burning materials to the point where combustion is too slow to be self-perpetuating.

Water's unique solvent properties are involved in chemical processes such as the manufacture of paper, of rayon, and of soap, to mention only a few. The solvent action of rainwater on rocks is a major source of weathering and erosion, while washing of dishes, clothes, automobiles, and floors also depend on solvent properties of water.

In medicine, water is used in intravenous fluids as a solvent for salts, sugar, and possibly therapeutic agents. The water content of the body is exploited in magnetic resonance imaging (MRI), since the hydrogen atoms in water can absorb radio frequency energy in the presence of a strong magnetic field. The amount of water in the tissues and its type of binding affect the frequency and intensity of the absorption, and make imaging possible without the use of potentially damaging X rays.

Life processes are often dependent on the solvent power of water, or on chemical reactions of water. Blood is largely water, using its unique solvent powers to transfer proteins, sugars, oxygen, and carbon dioxide. Arguably the most vital chemical reaction in the body is the reaction of water with adenosine triphosphate (ATP), which provides the energy for other vital reactions.

Water is the most plentiful source of hydrogen and yields this element by electrolysis, or by treatment with iron or carbon at high temperatures. Hydrogen can then be used to make other hydrides such as ammonia, to hydrogenate fats for margarine, or as a fuel to propel spacecraft. In space, electricity is generated by tapping the energy from the hydrogen-oxygen reaction in a fuel cell.

Hydrocarbons derived from petroleum, natural gas, coal, tar, sands, or oil shale are hydrogen compounds second only to water in industrial importance. Hydrogen, of all substances, releases the most heat per pound when it burns. Hydrocarbon fuels for transportation and heating derive most of their energy value from their hydrogen content, and most of their problems from their carbon content, which burns to form toxic carbon monoxide, or heat-trapping carbon dioxide. Demand for hydrocarbons is so great as to require importation of close to 9 million barrels of oil per day into the United States: more than half the amount consumed.

Hydrocarbons also provide feedstocks for the petrochemical industry, from which flow plastic, synthetic detergents, solvents, artificial rubber, and many other products that have become a part of life over the past hundred years, since the first oil wells were drilled. Although fewer tons of hydrocarbons are used for petrochemicals than for fuels, the petrochemical use is a more subtle one than mere burning. When petroleum shortages develop, it may be easier to find alternative fuels than alternative feedstocks for petrochemicals.

Ammonia, a compound of hydrogen and nitrogen, is manufactured in huge quantities for use as a fertilizer and as a precursor for all other industrial nitrogen compounds. Some familiar products requiring the use of ammonia for their manufacture are: explosives such as TNT and nitroglycerin, fabrics like Orlon and nylon, polyurethane plastic, and household cleaning liquids. Ammonia itself is also used as a refrigerant in industrial freezers. There is reason to believe that the earth's atmosphere once contained ammonia, and it may have been ammonia that provided the nitrogen for the first proteins of life. Ammonia has vanished now from the earth's atmosphere, but still exists (sometimes in liquid or solid form) on the planet Jupiter, and in comets.

The greatest importance of ammonia in the twentieth century has been as a source of available nitrogen for fertilizing crops such as corn, wheat, and soybeans. The yields of these crops have been dramatically increased by heavy application of ammonia and ammonium salts to the soil. The ammonia for this purpose is made by reaction of hydrogen with nitrogen at temperatures near 750 Kelvins and pressures of 100 to 700 bar (Haber-Bosch process). Iron catalysts are used in this process to speed the reaction at as low a temperature as possible, since high temperatures result in poorer yields. Typical ammonia plants may have capacities of 500 to 1,000 tons per day.

Ionic and metallic hydrides derive their applications from their action as hydrogen carriers and activators. Hydrogen in a hydride exists at a density that would be difficult to achieve with molecular hydrogen. Alloys of lanthanum and nickel have been found that can reversibly take up 1 to 2 percent by weight of hydrogen at room temperature and 2 to 3 bar pressure. The hydrogen can be released when needed by mild heating at atmospheric pressure.

Ionic hydrides such as lithium hydride contain more hydrogen, but release it only when water is added, and the hydride cannot be regenerated. High-density hydrogen storage will be important if hydrogen is ever to be used for powering vehicles. Alternatives to metal hydride storage systems are cryogenic storage and high-pressure gas storage, which suffer from adverse safety factors.

Metal hydride reducing agents like lithium aluminum hydride and sodium borohydride, discovered in the 1940's, have revolutionized organic synthesis and are used routinely for introducing hydrogen atoms into organic structures in a controlled manner. The main impact has been in pharmaceuticals. These hydrides are referred to as "complex metal hydrides" because they contain two elements besides hydrogen. In inorganic chemistry, complex hydrides such as lithium aluminum hydride and sodium borohydride can be used to prepare hydrides of most of the other elements.

Context

Hydrogen and its compounds have always played an important role in chemical theory.

In 1766, Henry Cavendish (1731-1810) reported the discovery of hydrogen and showed that two volumes of hydrogen combine with one volume of oxygen to form water. It was not clear until much later that both hydrogen and oxygen are diatomic molecules. Cavendish also did some of the earliest experimental work on equivalent weights. He showed that if samples are taken of two different acids in such a manner that each sample neutralizes the same weight of a base (such as one gram of potash), then these two samples will also neutralize equal weights of other bases, such as marble. The ratio of the weights of the two acids was constant, independent of the base used. Cavendish failed to recognize hydrogen as the active principle of all acids, but Sir Humphry Davy (1778-1829) did so early in the nineteenth century, based on his studies of hydrochloric acid. Davy, while realizing that hydrogen was present in all known acids, could not show why only some hydrogen compounds are acids while others are not.

Svante Arrhenius (1859-1927) proposed in 1887 that the characteristic properties associated with acids could be explained by assuming that all acids dissociated in solution to form the hydrogen ion (hydrated proton). The hydrogen compounds that failed to act as acids simply were not able to dissociate into ions. Arrhenius' theory was very helpful in interpreting the functioning of acids as catalysts and in the understanding of electrical conductivities of acid solutions, but was somewhat awkward in its treatment of bases. In 1923, Johannes N. Brønsted (1879-1947) and Thomas M. Lowry (1874-1936) clarified the role of solvents in acid behavior and defined bases as proton acceptors, thus making it much easier to interpret the behavior of compounds, like ammonia, that do not function as direct donors of hydroxide ions (like metal hydroxides), but still act as bases.

Hydrogen, because it occurs in so many common compounds and is so versatile in its bonding arrangements, has often been found in the forefront of new discoveries in bonding.

Ordinary shared electron pair bonds had been the rule up until structural studies of boron hydrides and other "electron-deficient" molecules revealed that hydrogen could participate in two-electron, three-center bonds. In such bonds, the hydrogen atom shares an electron pair simultaneously with two boron atoms. The discovery of this bonding mode for hydrogen was followed by the realization that multicenter bonding is common and may involve atoms other than hydrogen.

Increasingly, in the twentieth century, developments in physics began to have a strong influence on the development of chemistry. Atomic structure theory based on quantum mechanics, the discovery and purification of isotopes, and certain discoveries in nuclear science had profound effects. Nuclear magnetic resonance is a case in point.

In 1945, physicists Edward M. Purcell (1912- ) and Felix Bloch (1905-1983) demonstrated the phenomenon of nuclear magnetic resonance for hydrogen nuclei in simple materials such as paraffin. In a strong magnetic field, hydrogen nuclei absorb radio frequency energy. The intensity of absorption is proportional to the number of hydrogen nuclei, and the frequency of absorption is shifted by the details of the chemical bonding of the hydrogen ("chemical shift"). The NMR method has become indispensable to chemists for detecting hydrogen atoms and determining their structural relationships in compounds. Since it can be used on liquid samples, nmr has opened up new frontiers of knowledge about chemical species in solution, but the greatest impact has come in the rapid analysis of complex organic molecules.

Synthetic organic chemistry can be done faster because of NMR and can benefit from previously unobtainable information about hydrogen atoms in molecules.

As chemistry develops in the years to come, many of the advances in knowledge of bonding will certainly involve hydrogen compounds because of the comparative simplicity of the hydrogen atom and its versatility in forming new types of compounds.

Principal terms

COVALENT BOND: a force of attraction resulting from the sharing of electrons between atoms

HYDRATED PROTON: often called a proton by chemists, it is actually a proton surrounded by a cluster of water molecules; responsible for the distinctive properties of acids

HYDRIDE ION: a hydrogen atom with an extra electron

HYDROGEN BOND: a bond formed by hydrogen beyond its normal bonding state of one covalent bond; this bond is weaker than a covalent bond, and its presence is felt mainly in changed physical properties, such as boiling point

HYDROGEN MOLECULE: two hydrogen atoms joined by a single covalent bond; gaseous hydrogen as it is normally encountered is made up of hydrogen molecules.

IONIC BOND: a force of attraction between oppositely charged atoms or groups of atoms (ions); responsible for the stability of solids like sodium chloride

PROTON: a hydrogen atom that has lost its electron; the smallest positively charged ion

Bibliography

Atkins, P. W. MOLECULES. San Francisco: W. H. Freeman, 1987. In this beautifully illustrated book, individual molecules are shown as colored pictures of molecular models. The discussion helps to bring out the spatial relationships in the molecular structures. Most of the compounds treated involve hydrogen atoms, and there are interesting cases of hydrogen bonding.

Behrman, A. S. WATER IS EVERYBODY'S BUSINESS. Garden City, N.Y.: Doubleday, 1968. A simple account of the varied methods of filtration, softening, chlorination, fluoridation, desalinization, and water quality testing. The concluding chapter concerns public policy for water resources. There is also a brief list of suggested readings.

Davis, K. S., and J. A. Day. WATER, THE MIRROR OF SCIENCE. Garden City, N.Y.: Doubleday, 1961. The unique physical and chemical properties of water are explained in layperson's terms. Many achievements of science have involved water in some way, making it the "mirror of science." Examples come from fields such as geology, cosmology, nuclear science, and meteorology.

Lauterbur, Paul C. "Magnetic Resonance Technology for Medical Studies." SCIENCE 226 (October 19, 1984): 288-298. A review article covering various uses of magnetic resonance in medicine. Includes magnetic resonance imaging.

Leicester, H. M., and H. S. Klickstein. A SOURCEBOOK IN CHEMISTRY. Cambridge, Mass.: Harvard University Press, 1952. A lengthy quotation from the published work of Henry Cavendish appears here, in which the experiments leading to the discovery of hydrogen are described. There are explanatory notes that help to place the work in context and to clear up some of the old-fashioned terminology that Cavendish uses.

Pake, George F. "Nuclear Magnetic Resonance." SCIENTIFIC AMERICAN 199 (August, 1958): 58-68. The theory and practice of nuclear magnetic resonance are treated, with emphasis on hydrogen compounds. Other uses of the nuclear magnetic resonance method are mentioned, including magnetometers for use in prospecting for minerals and the mapping of the earth's magnetic field.

Partington, J. R. A SHORT HISTORY OF CHEMISTRY. 2d ed. London: Macmillan, 1948. Contains a clear account of the contributions of Henry Cavendish and other chemists of the eighteenth and nineteenth centuries toward understanding the composition of water and other compounds. There are pictures of old-fashioned apparatuses and many references to the original literature.

Shaw, Bernard L. INORGANIC HYDRIDES. Oxford: Pergamon Press, 1967. Begins with a discussion of hydrogen itself and the major classes of hydrides. Individual groups of elements and their hydrides are described, and the book concludes with chapters on bonding in hydrides, and physical studies of hydrides by infrared and nuclear magnetic resonance spectroscopy.

Wendland, Ray T. PETROCHEMICALS: THE NEW WORLD OF SYNTHETICS. Garden City, N.Y.: Doubleday, 1969. Hydrocarbons from petroleum are used in making plastic, synthetic rubber, transportation fuels, and a host of other useful products. These applications are described from an industrial and historical perspective.

Williams, L. O. HYDROGEN POWER. Oxford: Pergamon Press, 1980. Generation and use of hydrogen for power are discussed. Use of solid metallic hydrides for hydrogen storage has been tried experimentally but seems impractical for use in vehicles. Considers many schemes for splitting water into hydrogen and oxygen without using electricity. Clear analysis of the variables involved.

VanderWerf, Calvin A. ACIDS, BASES, AND THE CHEMISTRY OF THE COVALENT BOND. New York: Reinhold, 1961. An elementary approach to the Brönsted/Lowry and Lewis concepts of acid and base. Factors affecting the strengths of acids are discussed, and it is shown how acid-base relationships can lead to a better understanding of many different reactions.

Acids and Bases

Chemical Formulas and Combinations

The Structure of Ice