Xenon

Xenon
Atomic number	54
Symbol	Xe
Group	18 (noble gases)
Boiling point	−108.1 °C
Electron configuration	[Kr] 4d10 5s2 5p6
Melting point	−111.8 °C
Period	5
Main isotopes	129Xe, 131Xe, 132Xe
Phase STP	Gas
Block	p
Oxidation states	0, +2, +6
Wikidata	Q1106

Xenon is a heavy, colorless, odorless noble gas with atomic number 54. It lies in group 18 (the noble gases) of the periodic table (period 5, p-block) and is monatomic under normal conditions. At standard temperature and pressure (STP: 0 °C, 1 atm), xenon is a gas; its first impressions are that of an inert, “stranger” in air. The element’s symbol is Xe, and its atomic weight is about 131.3 amu. Because its outer electron shell is full (valence shell 5s²5p⁶), xenon has an octet configuration (eight valence electrons) that makes it generally chemically inert. Its common oxidation state is 0 (the elemental form), but it can be forced into positive states such as +2, +4, +6, and +8 in compounds with highly electronegative elements (especially fluorine and oxygen). Only under extreme laboratory conditions does xenon form stable compounds; most of the time it remains elemental. In this section and the ones that follow, we give an organized overview of xenon’s properties, isotopes, chemistry, uses, and history.

Atomic Structure and Electron Configuration

A xenon atom has 54 protons in its nucleus and, when neutral, also 54 electrons. Its electron configuration is \[Xe\] 4d¹⁰ 5s² 5p⁶, where “[Xe]” denotes the filled inner shells (1s through 4p, corresponding to krypton’s electron configuration). In other words, xenon’s valence shell is the 5th shell with 5s²5p⁶, a closed shell of eight electrons. This filled octet in the 5p subshell is why xenon behaves like an inert gas under ordinary conditions.

In terms of periodic trends, xenon has a quite large atomic radius (ion + van der Waals). Its calculated atomic (covalent) radius is about 108 picometers (pm), and its van der Waals radius is about 216 pm. Being a heavy atom (mass number around 131) with many electron shells, xenon’s electrons are relatively far from the nucleus on average. Its large radius and high polarizability make xenon more chemically reactive than the lighter noble gases (neon, argon) when confronted with strong oxidizers. Xenon’s electronegativity on the Pauling scale is about 2.6. To interpret this number, recall that smaller atoms (like fluorine) have higher electronegativities; xenon’s value 2.6 is lower than that of halogens (2.5–4.0) but higher than typical values for true metals or the very light noble gases. Xenon’s first ionization energy (the energy needed to remove one electron to form Xe⁺) is about 1170 kJ/mol (12.1298 eV). This is less than argon’s (1521 kJ/mol) – reflecting xenon’s relatively loose outer electrons – but it is still large compared to many metals. The second ionization energy jumps to about 2046 kJ/mol, showing that once one electron is removed, the next is much harder to take. In short, xenon’s electrons are held strongly enough that it rarely forms ordinary ionic or covalent bonds without a very powerful oxidizer.

Isotopes and Nuclear Properties

Naturally occurring xenon on Earth is a mixture of nine isotopes ranging from mass numbers 124 to 136. These are: ¹²⁴Xe, ¹²⁶Xe, ¹²⁸Xe, ¹²⁹Xe, ¹³⁰Xe, ¹³¹Xe, ¹³²Xe, ¹³⁴Xe, and ¹³⁶Xe. (The isotopes ¹²⁴Xe and ¹³⁶Xe can undergo double nuclear decay modes—¹²⁴Xe can double-electron-capture and ¹³⁶Xe double-beta-decay—but these processes have extremely long half-lives and have not been observed directly in a laboratory, so those isotopes are often treated as effectively stable.) The most abundant xenon isotopes by natural abundance are ¹³²Xe (~27%), ¹²⁹Xe (~26%), ¹³¹Xe (~21%), and ¹³⁴Xe (~10%). Lesser amounts of ¹³⁰Xe (~4%) and ¹³⁶Xe (~9%) are present, and isotopes ¹²⁸Xe, ¹²⁶Xe, and ¹²⁴Xe each account for less than a few percent of natural xenon.

These isotopes differ in nuclear spin and stability. Notably, ¹²⁹Xe has nuclear spin 1/2 and ¹³¹Xe has spin 3/2. All the even-mass isotopes (including ¹³⁴Xe, ¹³⁶Xe, etc.) have nuclear spin 0. Because ¹²⁹Xe and ¹³¹Xe have unpaired nuclear spins, samples of xenon gas can be used in nuclear magnetic resonance (NMR) and magnetic resonance imaging (MRI) experiments after hyperpolarization. The most well-known use of hyperpolarized xenon is in lung MRI: inhaled ¹²⁹Xe (when spin-polarized by exposure to a laser-pumped alkali vapor) yields enhanced NMR signals in lung tissue and blood, allowing imaging of pulmonary structure and gas exchange.

Radioactive isotopes of xenon occur as well. Xenon-133 (with half-life 5.25 days) and xenon-135 (9.14 hours) are produced as fission products of uranium or plutonium. Xenon-133 is used in medical diagnostics, particularly as an inhaled tracer for lung ventilation scans. Xenon-135 is famous in nuclear reactor physics: it is a very strong neutron absorber (with an enormous capture cross-section) and builds up during fission, acting as a poisonous "neutron poison" that can shut down or alter reactor power levels if not managed. This “xenon poisoning” is a crucial factor in reactor design and operation. After reactor shutdown, the decay of iodine fission products leads to a temporary rise in Xe-135 levels, and operators must wait for Xe-135 to decay (with its ~9 h half-life) before restarting the reactor safely.

In planetary science, xenon isotopes are keys to understanding early solar system history. For example, ¹²⁹Xe in Earth’s atmosphere partly comes from the long-ago decay of extinct ¹²⁹I (iodine-129, half-life 16 million years). This has been used to date events like the formation of Earth’s atmosphere and the differentiation of the Moon. Likewise, the isotopic composition of xenon in Mars’s atmosphere suggests significant atmospheric loss over time. In short, xenon isotopes are tracers for nuclear processes both natural (like radioactive decay chains) and human-made (nuclear reactors and tests).

Allotropes and Typical Compounds

Xenon, being a noble gas, exists as monatomic particles and has no allotropes in the way carbon does (graphite/diamond) or oxygen (O₂/O₃). At very low temperatures and sufficient pressure, xenon does solidify into crystals, but those solid phases (face-centered cubic, then hexagonal close-packed under higher pressure) are just states of the elemental form, not chemically different allotropes. One interesting “excited-state” species is the xenon excimer or exciplex, Xe₂, which forms only in excited plasmas or lasers: two xenon atoms can temporarily bond in an electronically excited state and emit ultraviolet (for example, vapor lamps used for photolithography use Xe₂ emission at 172 nm). Once the electrons relax, the Xe₂ dissociates again to two Xe atoms.

Although elemental xenon is chemically inactive under normal conditions, it does form compounds with strongly electronegative elements. The first-principles are that xenon’s filled valence shell can be perturbed by very high-energy reagents. The most important compounds are xenon fluorides and oxides. Three xenon fluorides are well-characterized: xenon difluoride (XeF₂), xenon tetrafluoride (XeF₄), and xenon hexafluoride (XeF₆). These are all covalent molecules with xenon in the center surrounded by fluorine atoms. Xenon difluoride is linear (Xe–F bonds at 180°), XeF₄ is square-planar, and XeF₆ is a distorted octahedron roughly (often treated as a fluxional octahedron). These fluorides are colorless or white crystalline solids at room temperature (XeF₂ melts at 111 °C, XeF₄ at 117 °C, and XeF₆ sublimes at about 60 °C because it is highly volatile). They react vigorously with water: for instance, hydrolysis of XeF₂ or XeF₄ yields xenon oxide and hydrogen fluoride.

Xenon oxides are more exotic. Xenon trioxide (XeO₃) is a colorless crystalline solid (m.p. 156 °C) analogous to osmium tetroxide in many ways: it is a strong oxidizer and explosive, decomposing to xenon and oxygen (XeO₃ → XeO₂ + ½O₂). Xenon tetroxide (XeO₄) is even more unstable and decomposes explosively at room temperature (in fact it explosively decomposes above –35 °C). Salts and acids can exist: there is xenic acid (H₂XeO₄) and xenates (like Na₂XeO₄), though these generally are only stable in solution or under cryogenic isolation. The oxidation state of xenon in these oxides can reach +6 (XeO₃) and +8 (XeO₄).

Aside from fluorine and oxygen, no other stable binary xenon compounds are known at ambient conditions. Xenon does not form true hydrides (no Xe–H bonds) or nitrides; attempts to make Xe–Cl or Xe–Br bonds yield only instantaneous excimer molecules or van der Waals complexes but no isolable solids under normal conditions. However, chemists have found a few very specialized xenon compounds with less electronegative partners if the environment is forcing. For example, xenon can bind to certain carbons or boron atoms when attached to strongly electron-withdrawing groups (though such compounds typically require elaborate synthetic methods and are highly reactive). Xenon also forms clathrate compounds: for instance, water ice under pressure can trap xenon atoms in its lattice (clathrate hydrate structure), but this is a physical entrapment, not a xenonallotrope. All true xenon “compounds” are usually xenon in a positive oxidation state bonded to electronegative ligands (F, O, or certain polyatomic anions).

Physical Properties

At room temperature and atmospheric pressure, xenon is a colorless gas. It is quite dense for a gas – about 5.9 kg/m³ at 0 °C (for comparison, air is 1.28 kg/m³). In other words, one liter of xenon gas at STP weighs about 5.9 grams. The gas is odorless and inert, and it is not flammable (since it is non-reactive, it cannot support combustion and will not burn). Xenon gas is used in certain lamps and lasers where it emits a bright blue-white glow: under electrical discharge it can emit strong lines in the blue part of the spectrum, giving a pale lavender or blue flash.

Xenon liquefies at −108.1 °C (165.1 K) under 1 atm pressure, and freezes at −111.8 °C (161.4 K). Near its triple point (161.4 K, 0.818 atm), liquid xenon has a density of up to about 3.1 g/cm³, higher than water. Solid xenon, formed just below this temperature, is even denser – around 3.64 g/cm³ at the triple point. This makes solid xenon even denser than most common minerals (for instance, granite is ~2.7 g/cm³). As a solid, xenon adopts a face-centered cubic (fcc) crystal structure (space group Fm3m) at low pressures. Under very high pressures (tens of gigapascals), xenon undergoes a phase change to a hexagonal close-packed (hcp) structure and eventually becomes metallic (around 140–155 GPa its conduction bands overlap). In fact, compressed xenon appears sky-blue and conducts electricity much like a metal. These extreme phases are of interest only at very high pressures (e.g. in diamond anvils or planetary cores) and do not occur naturally on Earth’s surface.

Xenon’s thermal conductivity as a gas is quite low: about 0.0056 W/(m·K) at STP (roughly one-fifth that of argon). The low thermal conductivity and high molecular weight make xenon useful for thermal insulation (for example in specialized thermally-insulating windows or cryogenic systems). Xenon has no electrical conductivity (being a nonpolar gas), but in a gas discharge it readily ionizes and conducts (hence its use in lamps and plasma devices).

Regarding spectroscopy, xenon has many atomic emission lines across the UV, visible, and infrared. In the vacuum ultraviolet, xenon excimers emit around 147 nm and 172 nm. In the visible and near-infrared, strong emission lines occur at wavelengths such as 462 nm (blue), 484 nm (blue-green), 656 nm (red), 828 nm (near IR), among others. In particular, emission near 828 nm is very intense and is often seen in discharge spectra. These lines (especially in xenon arc lamps) give off a broad, sun-like spectrum with peaks in the blue, which is why xenon arc lamps are used in photography and projection to mimic daylight. The precise positions of xenon lines are used for calibration in spectroscopy.

Chemical Reactivity and Trends

Xenon’s chemical reactivity is best described as “mostly inert.” In its elemental form, it does not oxidize or react with most substances. It does not combine with oxygen, water, acids or bases under normal conditions. Compared to many common elements, xenon will not corrode metals or form conventional salts. In terms of a “reactivity series,” xenon sits at the bottom (like all noble gases), resisting reactions. However, xenon is the most reactive of all noble gases (after radon), because its large, polarizable electron cloud can be distorted by the most powerful oxidizers.

In strong redox chemistry, xenon can behave as a reducing agent (giving up electrons) when meeting extreme oxidizers, but only in specialized conditions. For example, xenon will react explosively with elemental fluorine gas (the strongest oxidizer) at elevated temperatures or with a spark, yielding xenon fluorides. It can also be oxidized by oxygen difluoride (OF₂) or by ozone under specific conditions to form xenon oxides. Thus, xenon can have positive oxidation states up to +8 when truly forced. Conversely, xenon has no native negative oxidation states nor does it easily accept electrons as a reducing agent.

As for acids and bases, xenon itself forms neither. It does not dissolve in water to form a xenon acid. The “xenic acid” H₂XeO₄ (if formed via hydrolysis of XeF₂) is a strong oxidizing acid, but it is unstable outside of solution. Xenon does not act as a Lewis base (it won’t easily donate electron pairs to metal cations) nor a Lewis acid (it won’t accept electron pairs from donors). Complex formation with xenon is generally very weak unless one counts physical inclusion complexes (xenon atoms incarcerated in cavities of other molecules, as in clathrates or fullerenes). Essentially, xenon’s redox chemistry is limited to it being oxidized (Xe → Xe⁺/XeF compounds) under extreme conditions.

One more note on reactivity context: xenon’s ability to react increases down the group of noble gases, so it is more reactive than krypton and argon. This is why xenon compounds were discovered first (in the 1960s) long before any krypton or argon compounds were made. Even so, compared to typical metals or nonmetals, xenon is still extraordinarily sluggish: most chemists encounter it only as an inert background gas, unless they deliberately engage in high-energy chemistry.

Occurrence and Production

On Earth, xenon is very rare. It is found as a trace component of the atmosphere at roughly 0.087 parts per million by volume (about 87 nanoliters per liter of air). In mass terms, Earth’s atmosphere contains on the order of 2 gigatonnes of xenon, but spread over the globe that equals only about 5–10 liters of xenon at STP per square kilometer of Earth’s surface. Xenon is also released in small amounts from some mineral springs and as a fission product in nuclear reactors.

Xenon’s cosmic abundance is also low. It is formed almost exclusively in the rapid neutron-capture processes (the r-process) during supernova explosions. In the Sun and meteorites, xenon constitutes about 1 part in 630,000 by mass. Planetary atmospheres vary: Jupiter’s atmosphere contains about 2.6 times as much xenon (relative to hydrogen) as the Sun does, for reasons not fully understood. Mars has a xenon abundance similar to Earth’s but with a higher fraction of the Xe-129 isotope, hinting that Mars’s atmosphere has lost lighter elements over time while leaving heavier isotopes behind. Earth’s atmosphere itself shows an interesting “xenon paradox”: geochemical evidence suggests that Earth’s initial xenon inventory was partly bound in minerals and then lost to space or sequestered, leaving the atmosphere enriched in heavier xenon isotopes.

Xenon does not occur naturally in mineral ores like many heavier elements do. Instead, commercial xenon comes from the air. Xenon is produced as a byproduct of the cryogenic fractional distillation of liquefied air. In a typical air-separation plant, air is cooled until it liquefies; the liquid is then distilled to separate nitrogen, oxygen, argon, krypton, and finally xenon and krypton. Since xenon is so rare, it typically ends up concentrated (to around 0.1–0.2% by volume) in a liquid oxygen product or in a krypton stream. Further distillation and adsorption steps (for example, using silica gel) isolate xenon gas. By convention, extracting one liter of xenon at STP from the atmosphere requires on the order of 200–250 watt-hours of energy. This high cost of separation is why xenon is expensive compared to more common gases.

Major industrial suppliers of xenon include large gas companies (for example, Air Liquide, Air Products, Linde, and Praxair) operating in regions like Europe, North America, and China. Global production of xenon is only on the order of 30–50 metric tons per year. Prices vary with purity and quantity, but xenon typically costs on the order of tens to hundreds of dollars per liter of liquid at STP. (By contrast, argon and nitrogen are so much more abundant that they cost pennies per liter.) Xenon used in specialized batteries or certain electronics (e.g. plasma TVs) may be recycled, and spent nuclear fuel can also produce xenon, although this route is not the primary industrial source.

Applications and Technology

Despite its rarity and cost, xenon has found many high-value applications due to its unique properties.

One of the largest uses of xenon is in lighting. Xenon gas in high-pressure lamps produces a brilliant white light. Xenon arc lamps (also called xenon flash lamps or strobe lamps) are used in photographic flash bulbs, cinematography (studio lighting), and in endoscopes and microscopy as intense light sources. Automobile high-intensity discharge (HID) headlights often use xenon. The reason is that xenon’s spectrum under electrical excitation is broad and bright, resembling natural sunlight, with good output across the visible range. For example, commercial xenon headlamp bulbs burn around 35–50 watts of electrical power and output a daylight-like beam. Xenon is also used in mercury-free fluorescent lamps and as a filler gas in some incandescent bulbs to reduce filament evaporation (xenon slows tungsten evaporation, prolonging bulb life).

Another major application is in laser and plasma technology. Xenon is used in excimer and exciplex lasers. In an excimer laser, a noble-gas halide (like XeCl or XeF) is excited in a gas discharge. The short-lived excited dimer (e.g. XeCl, xenon chloride excimer) emits a specific ultraviolet wavelength (308 nm for XeCl, 172 nm for XeF) and then dissociates. Excimer lasers are widely used in semiconductor photolithography (etching integrated circuits), medical eye surgery (LASIK uses ArF, KrF, and partially XeCl lasers), and micromachining. Xenon is sometimes mixed with other gases (like helium or krypton) in “xenon ion lasers” that stimulated emission at visible and infrared wavelengths; the first gas lasers (in 1962) used xenon plasmas. Xenon also emits efficiently in vacuum-UV, so it is used in specialized lamps for UV curing of adhesives and in solar simulators.

In aerospace, xenon is the propellant of choice for many electric propulsion (ion thruster) systems. Because xenon has a large atomic mass, each ionized xenon atom carries substantial momentum when expelled from an ion engine, yielding high thrust per ion. Its inertness ensures it will not corrode or contaminate the propulsion hardware. Xenon ion drives have been used on satellites and deep-space probes (for example, NASA’s Deep Space 1 test mission and the Dawn asteroid mission used xenon thrusters). New satellites, including communications constellations and space stations, use xenon-fueled Hall-effect thrusters for fine maneuvering. The fuel efficiency (specific impulse) of xenon ion engines is much higher than chemical rockets, so xenon-as-propellant is key to long-duration missions.

Xenon’s medical and biological uses stem from both its inertness and its effects on the human body at high concentrations. Medically, xenon is a potent anesthetic when inhaled at high pressure; it produces anesthesia similar to nitrous oxide or halogenated ethers but with minimal side effects and rapid recovery (since xenon is not metabolized, it is quickly exhaled). Because it causes no ozone depletion or greenhouse effect and is non-toxic, xenon was investigated as an anesthetic gas in early clinical trials (first used in 1951). However, its high cost has limited widespread adoption. Xenon-133 (radioactive) is used in nuclear medicine for lung ventilation scans and thyroid studies. Xenon isotopes are also used as tracers in various medical and environmental studies.

In analytical sciences, xenon plays roles in high-precision instruments. For example, xenon lamps are used as calibration sources in spectrometers and as light sources in mass spectrometry. Xenon’s unique spectral lines make it a convenient reference. Hyperpolarized ¹²⁹Xe gas is used for advanced MRI of lung function or even brain perfusion. In physics, liquid xenon is a popular detector medium for nuclear and particle experiments (e.g. dark matter searches like the XENON1T experiment use tons of ultra-pure liquid xenon to detect rare particle interactions). These detectors use xenon’s scintillation (light emission when struck by radiation) and ionization properties to sense particles.

Other applications include: xenon gas discharge as a non-toxic filling in plasma or discharge panels (some high-end plasma TVs used xenon/krypton mixtures), as an environmental tracer (released in small amounts to study atmospheric transport, since it has no chemical sink), and even as a doping agent in sports (xenon inhalation has been used illicitly to increase erythropoietin production, though this is mostly anecdotal and now banned by agencies like WADA). Xenon is also found in some specialized chemical processes; for instance, xenon difluoride (XeF₂) is an effective fluoride donor or etchant in microfabrication (it can etch silicon and silicon dioxide smoothly at low temperatures).

Biology, Environment, and Safety

Xenon has no known biological role. In nature it is physiologically inert and virtually non-reactive, so organisms have not evolved to use it metabolically. If inhaled by humans, xenon behaves as an anaesthetic rather than a nutrient. (Indeed, at modest pressures xenon produces euphoria or dizziness, and at higher concentrations it induces full surgical anesthesia; these effects are due to xenon’s interaction with neural receptors, particularly as an NMDA antagonist, even though xenon itself is chemically inert.) Xenon is not toxic in the conventional sense; breathing pure xenon in place of air would simply cause asphyxiation (by displacing oxygen), not poisoning or chemical injury. Safety data for xenon treat it as a simple asphyxiant gas – chronic exposure has no special toxicity, but oxygen deficiency from even high concentrations of xenon can be dangerous.

In environmental terms, xenon is benign. It does not react or accumulate; it does not contribute to smog or ozone depletion, and its greenhouse effect is negligible due to its inertness and low concentration. Once released, xenon atoms remain in the atmosphere for geologic timescales; they can slowly leak to space from the upper atmosphere (like helium and hydrogen, but at a far slower rate because xenon is heavy). The natural cycling of xenon is minimal. In some closed-support systems (like submarines or spacecraft), xenon could accumulate without proper filtration, but in Earth's open environment it simply dilutes back into air. Low natural abundance means that any anthropogenic release is extremely small compared to the ocean of nitrogen and oxygen.

Safe handling of xenon focuses on conventional gas safety. Being inert, it is not flammable or corrosive. The main hazards are its density and its cryogenic forms: heavy xenon gas can displace oxygen at low-lying areas, creating an asphyxiation hazard in poorly ventilated spaces. (Industrial plants avoid this by having oxygen monitors and ventilation.) Liquid or solid xenon are extremely cold; contact can cause frostbite. Cylinders of high-pressure xenon must be handled like any compressed gas (risk of explosion if ruptured). In medical or laboratory use, xenon is sometimes pressurized, requiring careful handling. Otherwise, xenon is considered one of the safer industrial gases: it has no occupational exposure limits for chemical toxicity (only for asphyxiation).

History and Etymology

Xenon’s story begins in the late 19th century, in the era of discovering the noble gases. After Sir William Ramsay discovered argon in 1894, he and collaborator Morris Travers set out to find the other inert gases. In 1898, by partially distilling liquid air, Ramsay and Travers isolated a new gaseous element in the fraction remaining after removing oxygen, nitrogen, neon, and krypton. On July 12, 1898, they announced the discovery of xenon. Ramsay gave it the name xenon from the Greek word *xénon (ξένον) meaning “stranger” or “guest,” reflecting its previous obscurity in Earth’s atmosphere. Ramsay and Travers immediately published its spectrum and noted its noble-gas character.

Early research established that xenon was quite unreactive; for many decades it was a novelty with no known chemistry. In 1962, this belief changed dramatically when British chemist Neil Bartlett found that xenon could be made to react with powerful oxidizers. Bartlett was working with platinum hexafluoride (PtF₆) when he hypothesized that it might oxidize xenon (since PtF₆ had earlier been found to oxidize oxygen). When he mixed PtF₆ with xenon gas, an orange-yellow solid formed — the first recognized compound of any noble gas. It was initially formulated as Xe⁺[PtF₆]⁻ (today it’s understood as an adduct [XeF][PtF₅. This landmark experiment showed that xenon could indeed enter into normal chemical compounds. Within months, other xenon fluorides (XeF₂, XeF₄, XeF₆) and oxides were synthesized by various researchers, launching the field of noble-gas chemistry. Bartlett’s discovery earned him fame (though not a Nobel Prize) and revolutionized how chemists viewed the periodic table’s “inert” elements.

In parallel with chemistry, xenon found technological uses in the mid-20th century. In 1939, Harold Edgerton pioneered xenon flash photography by sending high-voltage discharges through excited xenon gas, creating very bright flashes; the technique became a staple of high-speed photography. In 1951, an Iowa anesthesiologist (Stuart Cullen) became the first to use xenon as a surgical anesthetic in humans. Research in the 1950s and 1960s on xenon’s pharmacological effects led to clinical studies of xenon anesthesia, valued for its safety profile. In 1960–1962, Bell Labs discovered laser action in xenon-doped plasmas, giving rise to xenon ion lasers for research.

Throughout the late 20th century, xenon’s applications grew. The development of high-pressure xenon arc lamps (from the 1960s onward) brought xenon lighting into everything from searchlights to cinema projectors. In the 1980s and 1990s, xenon’s role in rocketry began: the Space Electric Propulsion community adopted xenon as fuel for Hall-effect and ion thrusters, used on satellites and space probes. Most recently, discoveries like the utilization of hyperpolarized ¹²⁹Xe in MRI and xenon-based detectors in particle physics (e.g., liquid xenon detectors for dark matter) have shown that this rare gas still has untapped potential in cutting-edge technology.

Throughout its history, xenon’s identity as a “stranger” element has been echoed in its name and its biography: almost too scarce and stable to matter – until chemists and engineers found ways to make it extraordinary.

Data Table

Property	Value
Symbol, Name	Xe, xenon
Atomic number (Z)	54
Atomic weight (standard)	131.293 u
Group, Period, Block	18, 5, p
Category	Noble gas
Phase at STP	Colorless gas
Density (gas, 0 °C, 1 atm)	5.9 kg/m³ (≈0.0059 g/cm³)
Melting point	–111.8 °C (161.4 K)
Boiling point	–107.1 °C (166.0 K)
Crystal structure (solid)	Face-centered cubic (fcc)
Atomic radius (calc.)	108 pm
Van der Waals radius	216 pm
Electron configuration	[Kr] 4d¹⁰ 5s² 5p⁶
Valence electrons (outer)	5s² 5p⁶ (8 electrons)
Electronegativity	2.6 (Pauling scale)
First ionization energy	1170 kJ/mol (12.1298 eV)
Oxidation states	0 (elemental); compounds with +2, +4, +6, +8 (typical)
Stable isotopes (mass)	124, 126, 128, 129, 130, 131, 132, 134, 136
Electrons per shell	2, 8, 18, 18, 8
Notable spectral lines	Emission lines in UV–Vis; strong lines at 828 nm, 484 nm, 462 nm (etc.)