Hydrogen



The Astrophysics of Hydrogen

Although hydrogen is the most abundant element in the universe, it was not created spontaneously during the explosion that began our universe 15 billion years ago. The big bang formed a chaotic mixture of matter, antimatter, and radiation. Antimatter meeting matter underwent mutually explosive annihilation, becoming energy that could be absorbed by the subatomic particles created in the blast. If the amount of antimatter had equaled the amount of matter, everything would have been annihilated within a tenth of a second. Fortunately, matter was a slightly larger portion of the stew, and the entire system cooled sufficiently that some of the matter could form nucleons—the collective name given to the neutrons and protons that form the cores of atoms. Several hundred thousand years had to pass before free-flying electrons, attracted to the positively charged protons, could remain attached and atoms were born. Elegant in its simplicity, hydrogen was the most easily formed and remains the dominant atomic species in the universe today.

But where is it? Hydrogen composes a minuscule portion of our atmosphere—only one part per million. Although it is bound in the molecules of our oceans and rivers, hydrogen does not exist in its pure molecular form in very many places on Earth.

To find where the majority of hydrogen is located, scientists have examined spectroscopic data from stars. In the early universe, currents and spirals formed from matter attracted to other matter via the gravitational force, initiating clouds of hydrogen atoms. This activity still goes on today. Once a cloud reaches a temperature around 5 million K and a density about 100 times that of water, the hydrogen nuclei begin to fuse into nuclei of helium. For each fusion event, about 5 × 10–12 joules of energy are released—a very small amount, but the huge number of fusion reactions occurring each second results in the Sun radiating energy at the astonishing rate of 4 × 1026 watts. It is this process of hydrogen fusing into helium in the Sun that is the source of light, heat, and ultimately all the fundamentally useable energy available on Earth.

The only way scientists know that the Sun is mostly hydrogen is from experiments performed here on Earth. Each element in the periodic table has a distinct signature, called its spectrum. Electrons do not stick to the nuclei in atoms, but surround the core in a fashion that scientists have modeled variously as orbits, clouds, and probability densities. More detail will unfold later in this chapter, but for the moment, the model of an atom to be pictured is that of an electron orbiting the nucleus like the Moon orbits Earth.

An electron, however, unlike the Moon, can have many different orbits if it absorbs the right amount of energy. That energy can come from collisions with other particles. The atom can also absorb light, allowing the electron to jump to a higher-energy orbit, but it does not tend to remain in that excited state. It will eventually relax back to its least energetic state. When the electron drops down from a higher to a lower energy state, it gives off energy in the form of light or, more precisely, electromagnetic radiation. The human eye cannot see all wave-lengths of electromagnetic radiation, but there are instruments called spectroscopes that can detect and measure the radiation. Hydrogen electrons emit different wavelengths than do helium electrons, which emit different wavelengths than, say, carbon. In fact, every element has a unique spectrum by which one can identify it.

This can be observed in the laboratory (even a high school physics laboratory) using easily obtainable tubes of atomic gases. Scientists look at the spectra from the heated tubes of gas in the lab and then compare these to spectra from stars. Most of our stars are moving away from us as the universe expands, so we need to include a redshift factor, but the patterns remain the same.

Hydrogen nuclei fuse to make helium nuclei in the core of the Sun, producing energy in the form of electromagnetic radiation. (Extreme Ultraviolet Imaging Telescope Consortium/NASA)


All the hydrogen there is—and all other naturally occurring elements—are produced in stars. Most Earth-based elements heavier than iron, however, must have been created in stellar supernova events.

DISCOVERY AND NAMING OF HYDROGEN

Hydrogen is the most abundant element in the universe and the tenth most abundant element in Earth’s crust. Hydrogen atoms make up 93 percent of all atoms in the universe. About 6.5 percent of the atoms in the universe are helium atoms. The remaining scant 0.5 percent of the universe consists of the atoms of all of the other elements, and yet some of those elements were isolated and identified by ancient people. Why then were other, less abundant, substances recognized as elements so much earlier than hydrogen was? The simplest answer is probably that hydrogen is a colorless, odorless gas. In ancient times, Greek philosophers thought there were only four elements—earth, air, fire, and water—and that all other substances were mixtures of those four elements. Scientists no longer classify any of these as elements. Later, even the elements that were known to medieval alchemists were a liquid (mercury) and solids (such as gold). Since hydrogen is a colorless and odorless gas, alchemists did not know to look for it, so it seems natural that the discovery of hydrogen—and gases like oxygen and nitrogen—came after the Middle Ages.

Suspicion of hydrogen’s existence dates to 1671, when the English natural philosopher Robert Boyle (1627–91) noted the flammability of the gas that results from the reaction of iron with hydrochloric acid. Boyle, however, did not identify the fumes he obtained as being those of a new element. Credit for the discovery of hydrogen goes to the English chemist Henry Cavendish (1731–1810). There were alchemists before Cavendish who had dissolved metals in acids and observed hydrogen and noted its flammability, but Cavendish, in 1766, was the first person to state that hydrogen was different from all other gases. He called the new gas “inflammable air from the metals.” He was also the first per-son to obtain pure samples of hydrogen and to describe its low density. Cavendish dissolved metals such as zinc, iron, and tin in hydrochloric and sulfuric acid to isolate hydrogen. He made the mistake, however, of thinking that hydrogen was a component of the metals, whereas shortly afterward scientists recognized that hydrogen gas came not from the metals, but from the acids used to dissolve the metals.

Following the discovery of oxygen in 1774, Cavendish combined hydrogen and oxygen to make water. This experiment showed that water is a compound, not an element, sounding the death knell to the Greek notion of the four elements earth, air, fire, and water. It was  the French chemist, Antoine-Laurent Lavoisier (1743–94), who gave  the name hydrogen to Cavendish’s “inflammable air” after hearing of Cavendish’s success at making water from hydrogen and oxygen. The word hydrogen means “water producer.”

A PLANETARY NOTION: THE BOHR MODEL

Since antiquity, natural philosophers have debated questions regarding the fundamental nature of matter. Is matter infinitely divisible, or does there exist a smallest possible indivisible piece of matter? Democritus, the ancient Greek philosopher, taught that matter consists of fundamental particles that cannot be further cut or subdivided. The word “atom” today stems from Democritus’s teaching: a = not, tom = cut. Aristotle (384–322 b.c.e.), however, believed that matter is infinitely divisible; there are no fundamental particles that cannot be further sub-divided. Since Aristotle was the most influential natural philosopher of ancient times, atomic theory was largely abandoned until the beginning of the 19th century.

In 1808, the English chemist John Dalton (1766–1844) published A New System of Chemical Philosophy, in which he revived the atomic theory of Democritus and showed that then-current knowledge in chemistry was consistent with the theory that matter is composed of atoms. Although atoms are too small to be seen by the naked eye—or even by optical microscopes—atomic theory came to be accepted during the 19th century as the best explanation for the fundamental composition of matter. In recognition of his contributions, Dalton is often referred to as the “father of modern atomic theory.”

Throughout the 19th century, however, little progress was made regarding the possible composition or structure of atoms. Imagined as 

Danish physicist Niels Bohr postulated that the electron in hydrogen travels around the nucleus in the manner in which a planet orbits the Sun. Hence his model was called the “planetary model.”

little hard spheres, the atoms of different elements were known to differ only in their relative masses. This view changed with the discovery of radioactivity in 1896 by the French physicist, Henri Becquerel (1852–1908), and the discovery of the electron in 1897 by the English physicist, J. J. Thomson (1856–1940). In addition, physicists soon began talking about atoms containing fundamental positively charged particles. After the discovery of atomic nuclei, the nucleus of the lightest isotope of hydrogen was called a proton. (Although scientists postulated the existence of electrically neutral particles in the early 1920s, the discovery of the neutron eluded physicists until 1932.) These discoveries suggested that atoms were not hard spheres; however, a more detailed model of the atom was needed. Being English, Thomson was well acquainted with the traditional yuletide dessert plum pudding, which consists mostly of raisins, plums, and spices. By way of analogy Thomson suggested a “plum-pudding model” of the atom in which the electrons resembled the relatively small raisins in the pudding, and the positively charged particles resembled the much larger plums. Thomson’s model was a beginning in the attempt to visualize the internal structure of atoms. Such a homogeneous model of the atom, however, offered little insight into any relationship between the structures of atoms and their proper-ties and was later replaced with more detailed models.
In 1911, the British physicist and Nobel laureate Ernest Ruther-ford (1871–1937) published the article “The Scattering of Alpha and Beta Particles by Matter and the Structure of the Atom” in Philosophical Magazine. In this article, Rutherford reported the results of an experiment that demonstrated that the protons and electrons in atoms are not distributed homogeneously. Instead, the protons are concentrated in a relatively tiny region Rutherford called the nucleus (from the Latin, meaning “kernel”). The electrons are extranuclear; electrons are located in a relatively much larger volume of space surrounding the nucleus. Rutherford’s discovery of the nucleus was immediately accepted within the scientific community. However, the relationship, if any, between atomic structure and properties was still unclear.
For several decades prior to 1910, scientists had been studying the spectra of elements and compounds. When white light passes through a sample of a cool gas or liquid, certain frequencies are absorbed, leading to an absorption spectrum consisting of dark lines embedded in the rainbow colors of the spectrum. When a gas is heated, the gas emits certain frequencies of light, leading to an emission spectrum consisting of bright lines against a dark background. (Similar results can be obtained using any frequency of electromagnetic radiation. Spectra outside the visible region must be measured by instruments sensitive to those particular regions of the spectrum.)
Scientists recognized that an element’s spectrum is like human fingerprints; just as each person has unique fingerprints, each element in the periodic table has a unique spectrum. The spectrum of a compound is simply a combination of the spectra of the elements in that compound. The uniqueness of spectra makes them powerful analytical tools in the identification of the elements found in minerals, in biological samples, in Earth’s atmosphere, and in the atmospheres of stars. Although much empirical data had been accumulated, there was no theoretical frame-work that explained or predicted spectra.
Finally, an explanation was provided in 1913 by Niels Bohr (1885–1962), a Danish physicist who tackled the problem of trying to under-stand fundamental atomic structure. Bohr postulated that the electron in hydrogen travels around the nucleus in the manner in which a planet orbits the Sun. Hence his model was called the “planetary model.” The important distinction between the orbit of a planet around the Sun and the orbit of an electron around a nucleus is that the distance of a planet from the Sun is arbitrary, whereas in the Bohr model an electron cannot exist at just any distance from a nucleus. An electron can orbit the nucleus only at particular fixed, or discrete, distances from the nucleus.
The energy of the electron in orbit around the nucleus is the sum of the electron’s kinetic and electrostatic potential energies. Since the potential energy depends on the distance the electron is from the nucleus, and in Bohr’s model distances can only be discrete quantities, the model dictates that the total energy of the electron can only take on discrete values. An electron seems to jump from one orbit to another, but it cannot exist in a stable state in the region between two orbits. Orbits close to the nucleus represent states of lower energy. As the distance of an orbit from the nucleus increases, the energy of the orbit increases. When an electron occupies the orbit closest to the nucleus, the electron is said to be in its ground state, or state of lowest energy. When an electron occupies an orbit farther from the nucleus, the electron is said to be in an excited state, or state of higher energy.
An electron can make the transition from a state of lower energy to a state of higher energy by absorbing a photon of light. However, it cannot absorb a photon possessing just any frequency (which, in the visible region, determines the color of the light). The energy of the photon is proportional to its frequency. This energy must correspond precisely to the difference between the electron’s higher and lower energy states. Because there are only a finite number of possible energy states, only photons with specific energies can be absorbed. Therefore, only very specific absorption lines are observed in an element’s or compound’s spectrum. (In the reverse process, when an electron makes the transition from higher to lower energy, a photon of light must be emitted. Various transitions provide a specific set of emission lines in an element’s or compound’s spectrum.)


Emission spectrum of hydrogen. When atoms of an element are excited (e.g., by heating), they return to their state of lowest energy by emitting radiation at specific wavelengths. If this radiation is passed through a spectrometer, a spectrum is produced that displays the element’s characteristic emission lines. The lines are a unique “fingerprint” of an element.
 


The planetary model of the hydrogen atom seemed at first to successfully explain the spectrum that is observed in the ultraviolet, visible, and infrared radiation regions. Qualitatively, Bohr’s planetary model provides a reasonable explanation for the origin of spectral lines for all elements. Quantitatively, however, Bohr’s model provides only approximate results for the hydrogen atom, but wildly incorrect results for all other atoms. These discrepancies were soon discovered. One, involving more precise measurements of the frequencies of hydrogen’s spectral lines, showed that the experimental frequencies did not exactly match the frequencies predicted by Bohr’s model. Another discrepancy was that Bohr’s model did not predict the effects of electric or magnetic fields on hydrogen’s spectrum. Nevertheless, because Bohr’s model incorporated quantum theory into its formulation, it was an extremely important step in the ultimate elucidation of atomic structure and it garnered Bohr the Nobel Prize in 1922. It was a few years later that a more detailed treatment of the atom using quantum theory was presented.

A QUANTUM SOLUTION

Although Bohr’s model was successful at predicting spectral lines, it only applied to one-electron atoms and provided no way to calculate transition rates. Transition rates were known to vary, as some spec-tral lines were considerably brighter than others, indicating that some energy jumps occurred much more frequently. The model also allowed only for circular paths, an unstable condition for an orbiting electron. In addition, Bohr’s model required that the classical orbital angular moment (a product of the linear momentum vector and the radius vector) be equal in magnitude to an integer multiple of Planck’s constant divided by the quantity 2π. This notion seemed to defy common sense, as it was unnatural in classical mechanics to have physical quantities that had to be whole-number multiples of another quantity.
Louis de Broglie, then a doctoral student at Paris University, knew that whole-number behavior in physics was commonly associated with periodicity in a system. Perhaps a periodic nature had to be a property of electrons in atoms. Since waves are the quintessential example of periodicity, he hypothesized that not only did light have a wave nature, but so must electrons. De Broglie pictured electrons as particles embedded in standing waves around a nucleus. If the two modes of existence were to mesh, he had somehow to relate wavelength to mass or momentum. His famous formula λ = h/p (where λ is the wavelength, p is the particle momentum, and h is Planck’s constant) had the right dimensions, but needed to be confirmed by experimental tests.
To find out if electrons present a diffraction pattern as light waves do, they had to be beamed through an opening of 10–10m—about 1/10,000 the radius of a human hair. It is impossible to machine such a narrow slit, but in 1927, physicists Clinton Davisson and Lester Germer at the Bell Laboratories in New Jersey were able to shoot electrons through a nickel crystal, whose structure is arranged by nature to have atomic planes just the right distance apart. They heated a metal filament to eject electrons that then traveled through the crystal to land in an electron detector. If electrons did not have a wave nature, the distribution of electron hits would have been uniform, but it was not. Instead, a distinct pattern of light and dark, like the crests and troughs of a wave, was observed, confirming de Broglie’s controversial idea.
De Broglie’s formula was useful for knowing a particle’s wavelength, but a more comprehensive understanding of the electron’s motion and location while bound to a nucleus was needed. In order to be able to 





Niels Bohr was a Danish physicist who tackled the problem of try-ing to understand fundamental atomic structure. (AP Photo/Alan Richard)



predict behavior such as the flux of electrons (rate of flow per unit area) in a scattering experiment like that of Davisson and Germer, one needs a mathematical function to describe electron motion. Austrian physicist Erwin Schrödinger took on the challenge by adapting the known wave equation for light (electromagnetic waves) to make it comply with the law of energy conservation, free particle motion, and de Broglie’s calculation for wavelength. (Eventually Schrödinger had to modify the equation to accommodate a spin nature to the electron as tested by observing atoms in a magnetic field.)
The solution to Schrödinger’s equation is a wave function. The solution provides a way to calculate the probability of finding the electron in a particular region near the nucleus for a given energy, angular momentum, and spin state as well as the probability for a jump between two states. It is important to note that an exact solution for Schrödinger’s equation is known only for one-electron atoms where no additional forces exist from further electrons. All solutions for multi-electron atoms require extensive approximation. For that rea-son, experimentation leads theory in the field of atomic physics.
Conceptualizing this quantum mechanical behavior can be a problem. Schrödinger’s equation implies that the electron cannot be found at a specific location; rather, there is some probability (less than 100 per-cent) that it may be found there. This uncertain idea of trying to figure out the odds of locating a particle led Einstein to state in a 1926 letter to Max Born, “I am convinced that He (God) does not play dice.” Wave mechanics, however, is the best description yet found for observed electron distributions in atoms.


To find out if electrons present a diffraction pattern as light waves do, physicists Clinton Davisson and Lester Germer at the Bell Laboratories in New Jersey in 1927 were able to “shoot” electrons through a nickel crystal, whose structure is arranged by nature to have atomic planes just the right distance apart. They heated a metal filament to eject electrons that then traveled through the crystal to land in an electron detector.




THE NEGATIVE HYDROGEN ION

The negative hydrogen ion (H–), or hydrogen with an extra electron attached, is a glaring example of the necessity of quantum mechanical theory to describe atomic behavior. Classical physics, considering only electrostatic forces among the three charged particles, predicts that this ion should not exist in a stable, bound state. Yet H– has been observed experimentally for decades.
It is only possible to understand how two electrons can be bound to one proton by considering the electron wave functions. In quantum mechanics, the electrons cannot be modeled as pointlike particles orbiting the nucleus, but must be pictured as fuzzy distributions of probability. In H–, the electrons are in close enough proximity that their probability distributions, or wave functions, overlap. This overlap induces a positive correlation that allows the bound state of the ion. This means that the electrons do not have simple individual independent wave functions, but share a different and more complicated wave function.
Electron correlation is a relatively new concept and has been studied by monitoring the doubly excited states of this simplest atomic three-body problem. Unlike the hydrogen atom, H– has no singly excited state where just one electron jumps to a higher level. Instead, when sufficient energy is introduced into the negative ion (via collision with other particles or excitation by light), both electrons are simultaneously excited to higher levels. These so-called doubly excited states are evidence of a rather alarming phenomenon—it has been demonstrated that the electrons can actually share the energy of a single photon. This took physicists by surprise because a photon was understood to be an indivisible packet of energy and should be absorbed as such. Work continues on doubly and even triply excited states of various atoms, mainly using electron storage rings—particle accelerators that produce synchrotron radiation to facilitate excitation of the electrons.

THE CHEMISTRY OF HYDROGEN
Hydrogen itself is easily obtainable from the electrolysis of water. In an electrolytic cell, two terminals—an anode and a cathode—are con-nected to a power supply. When an electrical current is sent through the cell, bubbles of hydrogen gas are formed at the cathode and bubbles of oxygen gas are formed at the anode. The net chemical equation for the reaction is

2H2O (l) 2H2 (g) + O2 (g)

Note that in water the hydrogen atoms are bonded covalently to an oxygen atom. In molecular hydrogen, the hydrogen atoms are bonded covalently to each other. Hydrogen gas itself is highly flammable; in the reverse of the reaction above, H2 and O2 recombine to form water, releasing energy, sometimes explosively.

Chemical bonding involves an atom’s electrons. Electrons can be transferred from one atom to another (making an ionic bond), or electrons can be shared between two atoms (making a covalent bond). When an atom loses one or more electrons, it is left with a net positive charge, and the resulting species is called a positive ion, or cation. When an atom gains one or more electrons, it assumes a net negative charge, and the resulting species is called a negative ion, or anion. When an atom more or less equally shares electrons with another atom, no ion is formed. As a general rule, the chemistry of metallic elements is deiminated by the tendency of the atoms of metals to form almost exclusively cations, whereas the chemistries of metalloids and nonmetals are deiminated by the tendency of their atoms to form either anions or covalent bonds, but rarely positive ions.

Hydrogen is more versatile than most elements in that hydrogen can form all three: cations, anions, and covalent bonds. The tendency of hydrogen to form positive ions, as metals do, is responsible for its usual placement in the periodic table at the top of the first column above lithium. Some periodic tables, however, show hydrogen at the top of the column of halogens, reflecting hydrogen’s ability also to form negative ions. Some periodic tables even place hydrogen in both positions.

There are various definitions of acids and bases. The one used here is attributed to a theory developed in 1923 independently by Johannes Brønsted (1879–1947), a Danish chemist, and Thomas Lowry (1874–1936), a British chemist. Recall that an atom of ordinary hydrogen has only a proton and an electron, and no neutrons. Therefore, a cation of ordinary hydrogen (H+) is just a proton. In the Brønsted-Lowry definition of acids and bases, an acid is a proton donor, that is, it can react with other compounds or ions by transferring one or more H+ ions to the other compounds or ions. A base is a proton acceptor: It can react with the H+ ions of compounds or ions that are acids. Some chemicals species, such as H2O, are said to be amphiprotic, that is, they are both donors and acceptors of protons; they are both an acid and a base. These definitions are illustrated in the following examples (“aq” means the reaction is taking place in aqueous solution):

HCl (aq) + H2O (l) Cl- (aq) + H3O+ (aq)

(ACID + BASE → BASE + ACID)

NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)

(BASE + ACID → ACID + BASE)

In the first equation, hydrochloric acid (HCl) is a proton donor and water is a proton acceptor. Therefore, in that equation, HCl is an acid and water is a base. In the second equation, ammonia (NH3 ) is a proton acceptor and water is a proton donor. Therefore, in that equation, NH3 is a base and water is an acid. The ability of water to behave as an acid in one reaction or as a base in a different reaction illustrates water’s amphiprotic nature. Notice that in an acid-base reaction, the reactant that is an acid is changed into a product that is a base; the reactant that is a base is changed into a product that is an acid. These combinations are called acid-base conjugate pairs. Thus, HCl and Cl– form an acid-base conjugate pair, and NH4+  and NH3 form a conjugate pair.
A slightly different but equivalent definition of acids and bases is suggested by these equations. The first reaction results in an increase in hydronium ions (H3O+). The second reaction results in an increase in hydroxide ions (OH-). Therefore, an acid can be defined as a chemicals substance that, when added to water, results in an increase in the concentration of hydronium ions. A base is a chemical substance that, when added to water, results in an increase in the concentration of hydroxide ions. Hydronium ions give acid solutions the properties we associate with acids—sour taste, corrosiveness, and the ability to turn blue litmus paper red. Hydroxide ions give basic (or alkaline) solutions the properties of feeling “soapy,” causticity, and the ability to turn red litmus paper blue.
Compounds of hydrogen with other elements are typically called hydrides. In metal hydrides, the hydrogen is a negative ion. With metalloids and nonmetals, the hydrides can involve covalent bonding or the formation of the positive hydrogen ion. When hydrogen combines with nitrogen, ammonia (NH3 ) is formed. Ammonia is a base. Hydrogen combines with phosphorus to form phosphine (PH3). Water (H2O) is both an acid and a base. The hydrides of the other elements in oxygen’s family are acids—H2S, H2Se, and H2Te. The hydrides of all the halogens are acids—HF, HCl, HBr, and HI. As a matter of convention, when the hydrogen is written first in a formula, the compound is an acid; when the hydrogen is written last, the compound is a base. In aqueous solutions, the molecules of HCl, HBr, and HI all completely dissociate into ions, so hydrogen forms hydrogen ions (H+). (How-ever, because a cation of ordinary hydrogen is just a proton, it is too small and too highly charged just to “float around” in water unattached. 

Hydrogen cations attach themselves to water molecules to form hydronium ions, (H3O+.)
Atoms of metals only form cations. Therefore, when atoms of hydro-gen combine with atoms of metals, hydrogen forms “hydride” ions (H–). Examples are lithium hydride (LiH) and sodium hydride (NaH).
An important property of hydrogen is its ability to form bridges, called hydrogen bonds, between two molecules. For example, in a glass of water, the water molecules are attracted to each other by the attraction of a hydrogen atom on one water molecule to the oxygen atom on another water molecule. This attraction, hydrogen bonding, is responsible for many of the properties we associate with water. For such a small molecule, water has an unusually high melting and boiling point. Under the normal atmospheric conditions found on Earth, with-out hydrogen bonding to hold molecules together more strongly, water would exist only in the gaseous state. Without hydrogen bonding, life on Earth would not exist.


FUEL CELLS: HYDROGEN AND THE ENERGY CRISIS 
Concerns about global warming and fears of an oil shortage resulting from recent growth in demand have boosted research into hydrogen fuel cells. Fuel cells can emit zero carbon dioxide—the only by-products being water vapor and heat—and, since hydrogen is the most abundant element in the universe, it seems we should never have a shortage.
Fuel cell technology is not new science. Although it did not produce enough energy to be useful, the first fuel cell was designed and built in 1839 by a Welsh physicist, Sir William Grove, and the technology has been progressing ever since. The basic design is battery-like, with negative and positive terminals (anode and cathode, respectively). At the anode, hydrogen molecules are split by a well-chosen catalyst into pro-tons and electrons. The resulting electron current can provide power to run any electrical device. The protons travel through an electrolyte to reach the cathode, where the protons and electrons combine with oxygen molecules to make water vapor and heat. Alkaline cells, whose electrolytic solutions are sodium hydroxide or potassium hydroxide, have been used successfully in Apollo spacecraft and the space shuttle as well as small submarines.
Major challenges remain, however, when it comes to developing models suitable for personal automobiles. The abiding problem is cost-effective storage of the hydrogen fuel. Storing it in gaseous form can be ruled out; unless the gas could be highly compressed, which is costly and hazardous, it would need too large a tank. Liquid fuel is a better choice for high-energy density, but its boiling point is –253°C, so it has to be cryogenically cooled. Daimler-Chrysler demonstrated the feasibility of using liquid hydrogen in its NECAR 4 as early as 1999, but cooling and insulation costs made it impractical for mass production.
Current research into metal hydrides, which store hydrogen at a fairly high density (depending on the material), could make fuel tank size practical as well as cost effective. In this type of material, individual hydrogen atoms are absorbed into the lattice of molecules, which, when heated, releases the hydrogen as needed to the anode. Metal hydrides are, however, heavy and difficult to catalyze at normal ambient temperatures.
Intense and competitive research and development in nearly every industrialized country continues, with intriguing results involving lightweight nonmetal hydrides, carbon nanotubes, and nanocrystalline metallic alloys. One of the greatest advances may turn out to be the development of miniature fuel cells to power cellular phones, estimated by researchers at Lawrence Berkeley Laboratory to currently use more than 600 million kW-h of energy per year in the United States alone.



Schematic of an alkaline hydrogen fuel cell


TECHNOLOGY AND CURRENT USES

Because it is the most abundant element, hydrogen has naturally become a crucial constituent in myriad processes useful to humankind. Hydro-gen is a chemical component of ammonia and hydrochloric acid and consumer products such as soaps and cleaners, as well as cosmetics. For cooking purposes, it converts oils to solid or nearly solid fats, which are more conducive to baking and do not spoil as easily as other fats such as lard. It is used in welding, reduction of metal ores, fixation of nitrogen from the air, hydrodesulfurization in natural gas refining, and hydrode-alkylation to break down aromatic hydrocarbons.
Hydrogen is also important in fuel production. Hydrocracking uses the partial pressure of hydrogen gas to break down complex organic molecules, and forms by-products such as ethane, aromatics, and jet fuels. Liquid hydrogen is also used as a rocket fuel. In fission-based nuclear reactors, heavy water (where deuterium replaces regular hydro-gen) is used as a neutron moderator.
Fusion research currently relies on the deuterium-tritium fusion reaction to produce useable power, a method that is well understood but still highly inefficient—much more energy must be put into the process than is produced. Fusion is not expected to be a viable source of power for humankind for at least the next 50 years.
For powering automobiles, hydrogen fuel cells may someday be commonly used. The basic design is battery-like with negative and posi-tive terminals (anode and cathode respectively). At the anode, hydrogen molecules are split by a well-chosen catalyst into protons and electrons. The resulting electron current can provide power to run any electrical device. The protons travel through an electrolyte to reach the cathode, where the protons and electrons combine with oxygen molecules to make water vapor and heat. Major challenges remain, however, when it comes to developing models suitable for personal automobiles, the most difficult being cost-effective storage of the hydrogen fuel. The slow pace of development in this area is the reason President Obama in May 2009 reduced federal funds for hydrogen fuel research by more than half.
One of the most marketable of hydrogen products is hydrogen per-oxide (H2O2), with applications ranging from household antiseptic to rocket propellant. Hydrogen is also an important element in all biologicals compounds. Clearly this element will continue to play a major role in research in every scientific discipline.














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