Noble gases: noble gases periodic table facts

What are the noble gases?

The noble gases are the column of elements at the far right edge of the periodic table. Under standard conditions (0 °C and one atmosphere of pressure) the first six noble gases exist as single, unbonded atoms rather than as molecules such as O₂ or N₂. They are nonflammable, they do not support combustion, and they have no taste, smell or color until an electric current makes them glow. Students counting electrons and atoms will find seaborgium atoms a tidy account of superheavy chemistry.

The group sits between two of the most reactive families in chemistry. To its left are the halogens, which are one electron short of a full shell and grab one aggressively; to its right, starting the next row, are the alkali metals, which carry one spare electron and give it away just as eagerly. The noble gas elements sit at the stable midpoint between those extremes, which is why noble gases behave so differently from the other nonmetals around them.

The noble gases, in order of increasing atomic number, are:

• Helium (He), the lightest noble gas and the second most abundant element in the universe
• Neon (Ne), the gas behind the classic red-orange sign glow
• Argon (Ar), by far the most common noble gas in the air you breathe
• Krypton (Kr) and xenon (Xe), heavier and rarer gases used in specialty lamps
• Radon (Rn), a radioactive noble gas produced underground
• Oganesson (Og), a man-made element whose chemistry is still mostly theory

Group 18 elements on the periodic table

Modern versions of the periodic table label the noble gases group 18, following the IUPAC numbering adopted in 1988. Older American tables called the column group VIIIA, and some European tables used group 0, a reminder of the period when chemists believed these elements had a combining power of zero. Each member closes a row of the periodic table: helium ends period 1, neon ends period 2, and so on down to oganesson at the end of period 7.

Their atomic number values (2, 10, 18, 36, 54, 86 and 118) are the "magic" totals at which an electron shell or subshell fills completely. When Mendeleev published his first periodic table in 1869, none of the noble gases was known, so his arrangement simply had no place for them; the column was added once argon and helium had been isolated.

Oganesson, the synthetic element

Like other superheavy elements, oganesson does not occur in nature. Only a handful of its atoms have ever been made, each surviving well under a millisecond, so no one has measured its properties directly.

Calculations suggest that relativistic effects, which become significant when inner electrons move at a large fraction of the speed of light, smear out its outer shell so much that oganesson would probably be a solid at room temperature and noticeably reactive, making it the one noble gas that may not behave like one. Its longest-lived known isotope has a half-life of roughly 0.7 milliseconds, far too short for practical study.

How the rare gases were discovered

Because noble gases react with almost nothing, they hid in plain sight for more than a century. In 1785 Henry Cavendish repeatedly sparked air with excess oxygen to remove its nitrogen and found a small bubble, about 1/120 of the original volume, that refused to react. He recorded it and moved on. For discovery history, the curium atomic number page covers Seaborg, James and Ghiorso, who made the first sample by bombarding plutonium with alpha particles.

The trail picked up in the 1890s. Lord Rayleigh noticed that nitrogen extracted from air was consistently heavier than nitrogen made from chemical compounds; the difference was tiny but reproducible. Working with William Ramsay, he showed that the atmospheric sample contained an unknown heavier gas, which they named argon in 1894.

Ramsay then isolated helium from a uranium mineral in 1895, matching it to spectral lines first spotted in the Sun's light during an 1868 eclipse, and realized that an entire column of elements was missing. By 1898 he and Morris Travers had pulled neon, krypton and xenon out of liquefied air. Radon, first seen as an "emanation" from radium in 1900, completed the set of natural noble gases.

The family name changed along with the chemistry. The first label, "rare gases," was misleading because argon makes up almost 1% of the atmosphere. "Inert" became standard once the electron-shell model explained their stability, while the alternative name aerogens never caught on.

The word noble, borrowed from the noble metals such as gold that resist corrosion, won out after 1962, when Neil Bartlett mixed xenon with platinum hexafluoride and produced an orange-yellow solid. That compound proved noble gases were reluctant to react, not incapable of it.

Why group 18 elements resist chemical reactions

Chemical reactions happen when elements gain, lose or share electrons to reach a lower-energy arrangement. A noble gas atom already has that arrangement.

Its outermost electrons are held tightly by the nucleus, it has no unpaired electron to share, and adding an extra electron would force it into a new, higher-energy shell. This built-in chemical stability, rather than any special reluctance, is the source of the inertness of noble gases.

Valence electrons and the octet rule

Helium has two valence electrons, which fill its only shell. Every other member of the noble gases has eight in its outer shell, an s²p⁶ arrangement. In 1916 Gilbert Lewis built that observation into the octet rule: atoms of other elements tend to react until their valence shell matches the eight-electron pattern of the nearest noble gas.

Sodium loses one electron to look like neon; chlorine gains one to look like argon. That is why so much introductory chemistry is framed around noble gases.

The same filled shells explain why noble gases have a slightly negative electron affinity. Unlike chlorine or oxygen, they release no energy by accepting an extra electron, so they never form stable negative ions.

Electron configuration shorthand

Because the noble gas cores are so stable, chemists use them as shorthand. Instead of writing every subshell, you write the symbol of the preceding noble gas in square brackets and add only the electrons beyond it.

This is called noble gas notation. For bromine (Z = 35), the long form and the short form of its electron configuration are:

$$\text{Br: } 1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,3d^{10}\,4s^{2}\,4p^{5} \;=\; [\text{Ar}]\,3d^{10}\,4s^{2}\,4p^{5}$$

Strontium (Z = 38) collapses even more neatly to \([\text{Kr}]\,5s^{2}\). Every element after helium can be written this way, because noble gases mark the end of each row. The bracketed core tells you at a glance which electrons are chemically inactive, so you can focus on the few that actually take part in bonding.

Ionization energy down the group

Each of the noble gases has the highest first ionization energy in its period, which is why stripping an electron from one is so hard. The value drops steadily as you move down the column, from 2,372 kJ/mol for helium to about 1,037 kJ/mol for radon, because the atomic radius grows and the outer electrons sit farther from the nucleus.

That trend divides the family in two: the lighter noble gases (helium and neon) have no known neutral compounds, while the heavier noble gases (krypton, xenon and radon) can be pushed into bonding with fluorine and oxygen. Xenon and radon even have measurable electronegativity values on scales built for bonded atoms, while electronegativity has no practical meaning for helium and neon, which form no bonds to measure.

Physical properties of the inert gases

With no chemical bonds to rely on, the only attraction holding noble gases together in the liquid or solid state is the weak London dispersion forces, a type of van der Waals forces that comes from momentary shifts in an atom's electron cloud. Bigger atoms have larger, more easily distorted clouds, so the attraction strengthens down the group. Compared with other elements of similar mass, noble gases boil at extraordinarily low temperatures, and that single trend drives the rise in melting and boiling temperature, enthalpy of vaporization and solubility in water that you see in the table below.

Boiling points down the group

Read down the boiling point column and every noble gas boils tens of kelvins higher than the one above it. Helium boils at 4.22 K, xenon at 165.05 K, and radon, the heaviest natural member, at about 211 K (roughly −62 °C).

Even radon is still a gas on the coldest day ever recorded on Earth. The melting points follow just a few kelvins below each boiling point, so the liquid range of each of the noble gases is remarkably narrow; argon, for example, is liquid only between 83.81 K and 87.30 K at one atmosphere.

Atomic mass climbs much faster than those weak attractions do, which is why the density column rises so steeply. A balloon of xenon is about 4.6 times heavier than the same balloon filled with air, while helium is roughly seven times lighter. Most noble gases also have several stable isotopes, and the small mass differences between those isotopes are what geologists exploit later in this article.

Helium near absolute zero

Helium breaks the pattern of the noble gases in ways no other element does. It is the only substance that will not freeze at atmospheric pressure, even at absolute zero; you need about 25 atmospheres of pressure to turn it solid. The reason comes from quantum theory: helium atoms are so light and attract each other so weakly that their zero point energy, the motion they keep even at the lowest possible temperature, is enough to stop them locking into a crystal.

Below 2.17 K, ordinary helium-4 becomes a superfluid that flows with no viscosity, creeps up the walls of its container, and conducts heat better than any metal. These effects made helium a cornerstone of low-temperature physics through the twentieth century.

Noble gas compounds and how they form

For decades textbooks said the noble gases formed no compounds at all. In 1933 Linus Pauling predicted that xenon and krypton should bond with fluorine, but early attempts failed, and the idea faded until Bartlett's 1962 result rewrote that corner of inorganic chemistry. Today chemists know of several hundred noble gas compounds, nearly all built around the heavier members and the most electronegative elements, fluorine and oxygen.

A separate class of structures traps noble gases without any true bond. In clathrates, argon, krypton or xenon sits locked inside cages in a host crystal such as ice or hydroquinone, held only by the cage's shape. Helium and neon are too small to stay trapped.

Xenon compounds and fluorides

Most known xenon compounds start with a direct reaction between xenon and fluorine gas. By controlling the ratio of gases and the heating conditions, chemists can steer the product toward one of three fluorides:

$$\mathrm{Xe} + \mathrm{F_2} \rightarrow \mathrm{XeF_2} \qquad \mathrm{Xe} + 2\,\mathrm{F_2} \rightarrow \mathrm{XeF_4} \qquad \mathrm{Xe} + 3\,\mathrm{F_2} \rightarrow \mathrm{XeF_6}$$

• Xenon difluoride (XeF₂) is a stable white solid sold commercially as a fluorinating agent for organic synthesis and for etching silicon.
• Xenon tetrafluoride (XeF₄) has a flat, square-planar shape with two lone pairs, a textbook example of VSEPR geometry.
• Xenon hexafluoride (XeF₆) is the most reactive of the three and hydrolyses with water to give the explosive oxide XeO₃, which in solution is a strong oxidizing agent.

In these fluorides xenon takes an oxidation state of +2, +4 or +6, and in perxenates it reaches +8. Because xenon appears to exceed the octet, chemists describe these species as hypervalent and explain the bonding with a three-center, four-electron model. Together these compounds show that the noble gases can bond whenever the partner is electronegative enough.

Krypton, radon and argon compounds

Krypton difluoride (KrF₂) is the best-known krypton compound and an even stronger fluorinating reagent than its xenon counterpart, though it slowly decomposes in storage. Radon should be more reactive than xenon, but its intense radioactivity makes study difficult, so radon compounds are known mostly from tracer experiments.

Argon held out until 2000, when chemists in Finland made argon fluorohydride (HArF), stable only below about 27 K. No neutral compound in which noble gas atoms of helium or neon are chemically bonded has been isolated, so those two remain the only truly inert noble gases.

Where group 0 gases come from

Each of the noble gases has its own origin story, and that origin controls its abundance and price. Argon is plentiful because the radioactive decay of potassium-40 in rocks has been releasing argon-40 for billions of years.

Helium also comes from decay, but it is so light that it escapes Earth's gravity once it reaches the atmosphere. Radon forms underground and decays within days.

Abundance in Earth's atmosphere

Dry air is about 0.93% argon by volume, making it the third most common gas after nitrogen and oxygen. The rest of the noble gases are measured in parts per million: roughly 18 ppm of neon, 5 ppm of helium, 1.1 ppm of krypton and under 0.1 ppm of xenon.

Xenon is so scarce that geochemists still debate where Earth's "missing" share went. That scarcity is why a liter of xenon costs hundreds of times more than a liter of argon.

Fractional distillation of liquid air

Neon, argon, krypton and xenon are by-products of the industrial plants that make liquid oxygen and nitrogen. Air is cooled and compressed until it condenses (liquefaction), then warmed slowly in tall columns so each component boils off at its own temperature. That process, fractional distillation, sends argon out between the nitrogen and oxygen fractions and leaves the denser krypton and xenon concentrated in the liquid oxygen.

Helium is the exception. It is extracted from natural gas fields where helium produced by the decay of the radioactive elements uranium and thorium has collected underground for millions of years, sometimes reaching several percent of the gas. That is why world helium supply follows natural gas production rather than air separation.

Radon as a health hazard

Radon is the only natural noble gas with no stable form, and its inertness is exactly what makes it dangerous. Because it bonds with nothing, soil and rock that contain radium and other radioactive elements can't hold on to the radon they produce, so it diffuses freely upward, enters buildings through foundation cracks and drains, and builds up in poorly ventilated basements.

Once inhaled, its short-lived decay products lodge in the lungs and emit alpha particles there, making radon the second leading cause of lung cancer after smoking. The U.S. Environmental Protection Agency advises action at 4 picocuries per liter, and since radon can't be seen or smelled, a test kit is the only way to know.

Industrial uses of each inert gas

Almost every application of these elements relies on one of three traits: noble gases don't react, they glow in distinctive colors when electrified, or (for helium) they stay liquid at temperatures nothing else can reach. Industry consumes far more argon than all the others combined, simply because it is the cheapest noble gas to produce.

Welding and inert atmospheres

When metals are hot enough to melt, oxygen and nitrogen in the air attack them instantly. Argon, often blended with helium, is used as a shielding gas in TIG and MIG welding to blanket the weld pool and prevent oxidation.

Argon's inertness and low cost make it the default wherever reactive surfaces are exposed: steelmakers anneal parts under it, titanium is refined in it, and silicon crystals for the semiconductor industry grow inside it. An inert atmosphere of high-purity argon keeps trace contamination out of the chips that become integrated circuits.

Neon signs, lamps and lasers

Run a high voltage through a low-pressure tube of gas and each element glows its own color. Pure neon gives the familiar red-orange; argon with a drop of mercury gives blue, and phosphor coatings supply the rest of the colors you see in neon signs. Lighting uses go well beyond signage:

• Argon fills incandescent light bulbs, slowing evaporation of the tungsten filament.
• Krypton allows brighter, longer-lived light bulbs because its heavy atoms slow filament loss even more.
• Xenon arc lamps produce a near-daylight spectrum for cinema projectors and car headlamps.
• Krypton and xenon flash lamps deliver the brief, intense bursts used in high-speed photography and to pump solid-state laser rods.
• Excimer lasers based on argon fluoride and krypton fluoride produce the ultraviolet light used to pattern microchips and reshape the cornea in eye surgery.

Liquid helium in cryogenic cooling

Liquid helium has a boiling point of 4.22 K, the coldest of all cryogenic refrigerants, and that is the temperature range where many metals lose all electrical resistance. Every MRI scanner relies on superconducting magnets bathed in liquid helium to keep their coils in that superconductivity state.

Particle accelerators, NMR spectrometers and quantum computers depend on the same coolant. Liquid neon is a less common alternative when you need to hold temperatures near 27 K.

Group VIIIA gases in medicine, diving and space

Outside the factory, the same properties of noble gases solve problems in medicine, exploration and research:

• Diving: Helium-oxygen breathing gases such as heliox and trimix replace some or all of the nitrogen divers would otherwise inhale. That avoids nitrogen narcosis at depth and reduces the risk of decompression sickness, because helium dissolves less in body tissues.
• Anesthesia: Xenon works as an anesthetic with few side effects and a fast recovery, but its price limits use to specialist cases.
• Imaging and therapy: Hyperpolarized xenon-129 and helium-3 light up lung airspaces in MRI scans, and radon was once used in radiotherapy before safer sources replaced it.
• Lift: Helium gives balloons and blimps their buoyancy without the fire risk of hydrogen.
• Spaceflight: Xenon is the propellant of choice for ion engines on satellites and deep-space probes, since a heavy, inert atom can be ionized and accelerated without corroding the thruster.
• Laboratory analysis: Helium is the standard carrier for gas chromatography, and frozen argon matrices hold unstable species still long enough for spectroscopy.

Across all of these uses in healthcare and research, what matters is that the gas does its physical job (diluting, carrying, lifting, glowing) without taking part in any reaction.

Aerogens as tracers in Earth science

Because the noble gases don't react with other elements, the isotopes trapped in rocks, groundwater and volcanic gas keep a clean record of where they came from. Geochemists read that record of noble gas isotopes in several ways:

• Dating rocks: potassium-argon dating measures how much argon-40 has built up from the decay of potassium-40, giving ages for volcanic rocks from thousands to billions of years old.
• Tracing the mantle: gas from deep mantle sources carries far more primordial helium-3 relative to helium-4 than gas from the crust, so the helium ratio in a hot spring shows how deep its source runs.
• Dating groundwater: dissolved krypton-81 and argon-39 act as clocks for water that has been underground for centuries to hundreds of thousands of years.

Sampling and measuring these gases takes care, since a single air bubble would swamp the signal, but the payoff is a direct view into processes no other tracer can reach.