Noble Gases

Type of physical science: Chemistry

Field of study: Chemistry of the elements

The noble gases—helium, neon, argon, krypton, xenon, and radon—are significant in theoretical chemistry because they possess highly stable, filled shells of electrons. Limited chemical reactions and compounds have been discovered for argon, krypton, xenon, and radon, but neon and helium are not yet known to be capable of chemical combination.

Overview

The noble gases (formerly called "rare" or "inert" gases) are usually included at the extreme right-hand side of the chemist's periodic table, just beyond the other groups of nonmetallic elements, the chalcogens and halogens. The noble gases differ remarkably from other nonmetals by having a very low degree of chemical reactivity, which tended to render them invisible to many of the research activities of early scientists. Dmitry Ivanovich Mendeleyev was unable to include a single noble gas element in his periodic table in 1871. During the years 1892 to 1902, the noble gases were discovered and added to the periodic table as a new "group zero": helium, neon, argon, krypton, xenon, and radon, in order of increasing atomic number.

The noble gas elements are extremely volatile, with helium (boiling point 4.3 Kelvins, or -268.9 degrees Celsius) exceeding all other elements in this respect. Even radon, the least volatile noble gas, boils at 211 Kelvins (-62 degrees Celsius). These elements exist as monatomic gases, which distinguishes them from the other elemental gases, such as hydrogen, nitrogen, and oxygen, which are all diatomic. Until 1962, the noble gases were not known to form any chemical bonds, although "clathrate" compounds were known, in which a solid "host" compound could crystallize, trapping noble gas atoms as "guests" in the crystal lattice. Numerous compounds are now known for krypton, xenon, argon, and radon.

Paradoxically, the noble gases, at their discovery, posed such problems for chemical theorists that some refused to recognize them as elements at all. The electronic theory of the atom permitted the view that stable, filled electron shells are not only the source of the inertness of the noble gases, but also the central factor in covalent bond formation by all the other elements.

During a solar eclipse in 1868, Pierre Janssen examined the solar spectrum and discovered a new line. This line was shown to be caused by a new element by Sir Joseph Norman Lockyer, who named the element for the Greek word for "sun."

William Ramsay and John W. Strutt (Lord Rayleigh) were the first to experiment with helium, using samples of the gas obtained from air or from uranium ore. Ramsay and Lord Rayleigh also discovered neon, argon, krypton, and xenon in air, and recognized them as new elements. The discovery of radon occurred in 1900, when Friedrich Dorn showed that a radioactive gas is produced by radium as it disintegrates. Ernest Rutherford (1871–1937) and Frederick Soddy (1877–1956) also studied radon, and liquefied it. Later, investigations in Ramsay's laboratory showed that radon was the densest gas known up to that time, and was the element below xenon in the periodic table.

Ramsay made many attempts to prepare compounds of argon, without success. The other gases of periodic group [0] were similarly "inert" to all chemical reagents tried, and so for many years were referred to as the "inert gases." Scientific interest in these gases was thus confined for many years to physical studies. As the electronic structure of atoms began to be understood, the inertness of the group [0] elements was interpreted as a consequence of their having filled shells of electrons: two electrons for helium, eight electrons for the others. The stability of the octet became a central feature of a theory of covalent chemical bonding put forth by Gilbert Newton Lewis (1875–1946) in 1916.

Helium attracted attention for other reasons. The difficulty of liquefying it raised fundamental questions about the gaseous state: some substances might be truly "permanent gases," incapable of existence as liquids. In the Netherlands, at the University of Leiden, Heike Kamerlingh Onnes (1853–1926) became a master practitioner of low-temperature technology, and succeeded in preparing liquid helium in 1908. The unprecedentedly low temperatures produced by cooling with liquid helium (boiling point 4 Kelvins) made possible crucial tests of the predictions of the quantum theory. Completely unexpected changes in mechanical and electrical properties of materials were found to occur at low temperatures, perhaps the most intriguing of which was the complete loss of electrical resistance in metals cooled in liquid helium. The samples of helium used by Kamerlingh Onnes had to be obtained from monazite, a mineral containing trapped helium originating from decay of radioactive elements. In 1907, the presence of helium in natural gas was reported at concentrations of 1 to 2 percent, making this a more significant source than monazite.

Liquid helium, when cooled to 2.2 Kelvins (called the λ point), undergoes a remarkable change in properties and enters what is sometimes called a fourth state of matter: the superfluid state. Superfluid helium exceeds all other substances in its ability to conduct heat and in its almost total lack of viscosity, or resistance to flow. Helium in the superfluid state can climb up over the sides of its container in defiance of gravity, and can seep with extraordinary rapidity through the tiniest of openings. These remarkable properties have posed a great challenge to theoretical physicists, and require sophisticated applications of quantum theory for their interpretation. At issue are the quantum theoretical idea of zero point energy, and the influence of symmetry. Isotopic helium 3, with a different nuclear spin from normal helium 4 and lighter atoms, does not become superfluid when cooled; but since it has a lower boiling point than helium 4, it can be used to produce even lower temperatures than the heavier isotope. Studies of low-temperature phenomena are essential in determining absolute entropies of substances and in verifying quantum theoretical predictions of the properties of solids and liquids.

Astrophysics also has benefited from the consideration of helium. The occurrence of helium in the sun was identified as an important factor in the thermonuclear reactions that produce solar heat and light and make life possible on earth. In 1938, Hans Bethe (1906–2005) and Carl Friedrich von Weizsaecker (1912–2007) independently worked out the series of thermonuclear reactions called the "carbon cycle." The reactions of the carbon cycle result in the fusion of hydrogen atoms to form helium atoms plus energy. Carbon and other elements are involved as preliminary stages to the formation of helium. The reactions of the carbon cycle can account for part of the energy emitted by the sun. Fusion reactions also take place in thermonuclear weapons: the so-called "hydrogen bomb." Such nuclear fusion under controlled conditions could be very desirable as an energy source, but has yet to be realized. Repeated reports of fusion reactions occurring at near room temperature during electrolysis experiments have remained controversial and have not resulted in practical energy sources.

The year 1962 was important in the history of the group [0] elements, marking their transition from "inertness" to "nobility." Neil Bartlett (1932–2008) at the University of British Columbia, Canada, did an experiment with xenon and platinum hexafluoride, a volatile substance with powerful oxidizing properties. Upon allowing the two gases to mix, Bartlett observed the formation of an orange, solid compound, the first instance of true chemical combination ever recorded for a group [0] element. The experiment was carefully planned, and compound formation was considered likely based on known physical properties of the reactants and by analogy with previously known reactions of the platinum compound. The group [0] gases now began to be referred to as "noble" rather than "inert." The use of "noble" in this way had been common since the Middle Ages in reference to metals, like gold, that resist the attack of all but the most powerful acids. The "base" metals, such as lead, are corroded and dissolved by nitric acid. The noble gases were the first nonmetals to be so designated. Suggestions to use names such as "argonons" or "aerogens" for the group [0] elements have not been widely followed.

As news of Bartlett's work spread, many chemists turned their experimental efforts toward the preparation of noble gas compounds. The fluorine chemistry experts at Argonne National Laboratory in Illinois soon reported that xenon will react with fluorine to produce xenon fluorides. Later, krypton and radon were also brought into combination with fluorine. By 1966, more than four hundred scientific papers had appeared on the subject of noble gas compounds. The catalog of known xenon compounds has grown very long, and includes compounds with carbon, nitrogen, oxygen, or fluorine bonded to the xenon atom. In 2005, it was first demonstrated that argon, too, can form compounds.

Applications

Air contains nearly 1 percent argon, 18 parts per million of neon, 5 parts per million of helium, and about 1 part per million each of krypton and xenon. Argon, neon, krypton, and xenon all are recovered for use by fractional distillation of liquid air. They are by-products of the huge liquid-air industry, and probably would be uneconomical to produce were it not for the demand that exists for liquid nitrogen and liquid oxygen. Helium is separated from natural gas, in which it exists at a concentration of several percent, rather than from air. Radon can be obtained only as a decay product of radium, which produces about 1 cubic centimeter of radon per month per gram of radium.

Many applications of the noble gases result from the low chemical reactivity of the elements. Argon and helium are used as inert, "blanketing" gases to protect hot metals from reacting with atmospheric oxygen or nitrogen. Blanketing can be important in welding metals and in preparation of reactive metals (titanium or magnesium, for example). Lamp filaments also survive much longer in an inert gas atmosphere than in air; argon and helium are used in lamps for this reason. Krypton may be substituted if it is desirable to have a higher operating temperature, which can produce from 7 to 20 percent increased light output from a bulb. Noble gas blanketing is also used to protect high-purity materials from contamination during assembly of integrated circuits.

The spectroscopic properties of noble gases have led to many applications. So-called "neon signs" are gas-filled tubes provided with a high-voltage discharge. Neon, argon, and helium, individually or in mixtures, are used to produce different colors. Lasers make use of the same gases to produce tight beams of coherent red or blue light that have been used for accurate distance measurements and many other purposes. The atomic spectrum of krypton is used as a standard of wavelength and forms the basis of the definition of the meter as a unit of length.

Xenon-krypton mixtures in flash lamps emit bright light flashes of extremely short duration when subjected to electrical excitation, useful in high-speed photography.

A different sort of application of noble gases in spectroscopy is their use as solvents, or matrices, for samples. Many short-lived chemical species can be trapped in frozen argon long enough to permit spectroscopic measurements. This is called the matrix-isolation method. A related technique employs liquid xenon as an inert solvent in which species of interest can be generated by photolysis and studied spectroscopically. The xenon is transparent to the radiation used in the spectroscopic experiment and thus does not interfere with the spectrum of the sample under study.

The low molecular weight and correspondingly low density of helium lead to several uses. Lighter-than-air craft acquire their buoyancy from helium, which is nonflammable and thus preferable to the even lighter gas, hydrogen. Light molecules have greater average velocities than heavier ones, making helium a very good conductor of heat and leading to its use as a coolant in nuclear reactors. The inertness of helium prevents corrosion of the reactor materials and permits higher operating temperatures. High thermal conductivity and inertness are also important in the use of helium as an eluant gas in vapor phase chromatography, the most prevalent method for chemical analysis of complex mixtures of volatile substances.

The low density of helium means that helium in the atmosphere is quickly lost to outer space. Since helium is a unique resource, and nonrenewable, it is desirable to try to conserve it, and not vent it to the atmosphere unless absolutely necessary. The United States government maintains a national committee on helium conservation.

Noble gases also have applications in relation to human health and well-being.

Helium-oxygen mixtures are breathed by deep-sea divers in preference to ordinary air because helium is less soluble in blood than nitrogen and can diffuse faster through the tissues. This prevents the distressing symptoms of nitrogen narcosis, which can afflict divers breathing ordinary air. Xenon, in contrast to helium, is found to possess anesthetic properties when breathed, and, although expensive, is nonflammable, unlike other anesthetics such as ether and cyclopropane. Radon, because of its radioactivity, can be used in radiation therapy, but if accidentally breathed can be a health hazard. Radon is constantly emitted by some rocks and can be released into the cellar of a building, and its water solubility allows it to enter buildings through the tap water. Unless the radon is efficiently ventilated, dangerous levels of exposure can occur.

The argon and helium on earth are thought to have arisen mainly from radioactive decay of other elements. Careful measurement of argon levels in rocks can be combined with a knowledge of the rate of its formation to yield evidence of the age of the rocks. This is the potassium-argon dating method. Methods of this sort can be applied to extraterrestrial objects, too. Meteorites contain traces of noble gas isotopes that can be studied to determine the age of the meteorite or the extent of its exposure to cosmic radiation.

The low boiling point of helium leads to the use of liquid helium in the production of extreme cold. The electrical conductors in magnet windings are made superconductive by liquid helium cooling. The powerful magnetic fields are then used for instruments such as particle accelerators and nuclear magnetic resonance spectrometers. One of the world's largest helium liquefaction plants is located at the site of the Fermi National Laboratory near Batavia, Illinois, where its output can be used to cool the magnet coils in the accelerator. Helium-cooled magnets are also used in magnetic levitation ("mag-lev") devices for transportation applications such as monorail trains.

The discovery of noble gas compounds such as the oxides and fluorides of xenon makes it possible to contemplate the possible use of these compounds as chemical reagents.

Xenon difluoride shows promise as a mild fluorinating agent and may find application in the synthesis of pharmaceuticals. Xenon trioxide, in water solution, is a powerful oxidizing agent and can replace periodates in the analytical procedure for manganese.

Context

When Lord Rayleigh began his studies of the densities of elemental gases late in the nineteenth century, isotopes had not yet been discovered, nor had mass spectrometers been developed for the accurate measurement of atomic masses. The motivation for Rayleigh's work was to test the hypothesis, called Prout's hypothesis, for William Prout (1785–1850), that all elements were built up from hydrogen, and had atomic masses equal to multiples of the mass of a hydrogen atom. Having dealt with oxygen and found a value close to 16 for its atomic mass, Rayleigh proceeded to nitrogen. His research led to the discovery of argon, which was found to have a molecular weight of 39.9 and to be a monatomic gas, the first elemental gas with this property. More controversial was the fact that its atomic weight lay between those of potassium and calcium, which were adjacent elements in the periodic table. Acceptance of argon as an element seemed to mean abandonment of increasing atomic weight as the ordering principle for the elements. The controversy gradually died down as the other group [0] elements became known and as knowledge of the atom improved.

Early in the twentieth century, momentous discoveries in atomic theory that clarified the electronic structure of atoms and the nature of the nucleus were made. Prout's hypothesis became a dead issue, even though it was proved partially true: Atomic nuclei were found to be composed of different numbers of hydrogen nuclei, together with neutrons. The subject of atomic masses turned out to be much more complex than Prout or anyone else had foreseen. Now the group [0] elements found new importance because of their stable, filled shells of electrons.

Lewis and Irving Langmuir (1881–1957) used the stability of the noble gas electron shell ("octet") to explain the different numbers of covalent bonds formed by different elements.

Elements early in the periodic table were limited to four electron pair bonds, but could form fewer than four if they retained nonbonding pairs of electrons.

In ionic compounds, the noble gas structure was found for many of the monatomic positive and negative ions formed by the elements. No element gained or lost more electrons than it needed to produce an octet. Most of the common salts could be thought of in this way.

Discovery of xenon compounds in 1962 showed decisively that the noble gas octet is not inviolable, and can be split up under highly oxidizing conditions.

The noble gases have always assumed an important role in theories of chemical bonding, and will continue to do so as new types of compounds are synthesized.

Principal terms

ATOMIC SPECTRUM: light of characteristic wavelengths emitted by a heated chemical element that serves as a means of detecting the element and distinguishing it from other elements

COVALENT BONDING: a force of attraction resulting from sharing of electrons between atoms; a major source of stability for chemical compounds

ISOTOPES: varieties of an element differing in the number of neutrons in the atomic nucleus

KELVIN: a unit of temperature named for Lord Kelvin (William Thomson); zero Kelvin, also called "absolute zero," is equivalent to -273.15 degrees Celsius

PERIODIC TABLE: a matrix arrangement of the chemical elements in which similar elements appear in the same column, ordered by increasing atomic number (number of electrons in the neutral atom)

QUANTUM THEORY: the idea that light energy has a smallest unit (a quantum) was proposed by Max Planck in 1900; this theory and its ramifications constitute the basic theory of matter and energy

SUPERCONDUCTIVITY: the property of conducting electrical current without resistance; many solids exhibit this property at low temperatures

ZERO POINT ENERGY: the energy retained by a substance even at the lowest temperatures; the existence of such an energy was predicted by quantum theory

Bibliography

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Fisher, David E. Much Ado about (Practically) Nothing: A History of the Noble Gases. Oxford UP, 2010.

Halford, Bethany. "To Get Noble Gases to Forge Bonds, Chemists Go To Extremes." Chemical and Engineering News, 29 May 2019, cen.acs.org/materials/inorganic-chemistry/IYPT-get-noble-gases-to-forge-bonds-chemists-go-to-extremes/97/i22. Accessed 6 Feb. 2025. 

Halka, Monica, and Brian Nordstrom. Halogens and Noble Gases. Facts on File, 2010.

"Noble Gases." GCSE Bitesize, BBC, 2014, www.bbc.co.uk/schools/gcsebitesize/science/aqa‗pre‗2011/oils/changesrev6.shtml. Accessed 6 Feb. 2025.

Schrobilgen, Gary J. "Noble Gas." Encyclopedia Britannica, 30 Jan. 2025, www.britannica.com/science/noble-gas. Accessed 6 Feb. 2025.