Europium

Europium
Atomic number	63
Symbol	Eu
Boiling point	1529 °C
Discovery	Eugène-Anatole Demarçay (1901)
Electron configuration	[Xe] 4f7 6s2
Melting point	826 °C
Main isotopes	151Eu, 153Eu
Cas number	7440-53-1
Block	f
Phase STP	Solid
Oxidation states	+2, +3
Wikidata	Q1396

Europium (Eu, atomic number 63) is a silvery-white lanthanide metal. Like other rare-earth elements it is soft (hardness ≈ lead) and easily oxidized – in fact “Europium is the most chemically reactive … and least dense” of the lanthanides The metal crystallizes in a body-centered cubic lattice and has a low melting point (822 °C) and boiling point (1529 °C) for a metal At room temperature it is a solid (soft enough to cut with a knife with density about 5.24 g/cm^3 In its compounds Eu typically occurs in the +3 oxidation state (like its neighbors in the periodic table), but the +2 state is unusually stable for europium –Eu^2+ compounds are mild reductants. (Eu can even behave chemically like an alkali-earth in forming Eu^2+ salts.) Common compounds include #3 oxides and halides (e.g. Eu₂O₃, EuCl₃) and the corresponding +2 versions (EuO, EuCl₂, etc.) Europium has no known biological role and is relatively non-toxic (acute oral LD₅₀ for Eu salts is on the order of grams per kg) Its name comes from the continent of Europe, in honor of Demarçay’s 1901 naming of the element after the place of its discovery

Atomic Structure and Electron Configuration

Europium’s atomic number is 63, so its ground-state electron configuration is [Xe]4f^7 6s^2, with a half-filled 4f shell plus two 6s electrons In the trivalent state Eu^3+ it typically loses the two 6s electrons and one 4f electron (leaving 4f^6). The half-filled 4f^7 configuration of Eu^2+ is unusually stable among the lanthanides, which is why Eu^2+ salts are much easier to make than for most rare-earths Because of its f-electrons, europium is a large atom: the empirical atomic radius is about 185 pm Its ionic radius reflects the lanthanide contraction: for 6-coordinate ions Eu^2+ is roughly 131 pm while Eu^3+ is about 109 pm Analogous to its neighbors, europium is electropositive (Pauling χ ≈1.1–1.2) and has a relatively low first ionization energy (5.670 eV These trends are typical of a heavy f-block element: Eu is large, weakly electron-attracting, and loses electrons readily.

Isotopes and Nuclear Properties

Natural europium consists of two isotopes: ^153Eu (≈52.2%) and ^151Eu (≈47.8%) Although ^153Eu has so far shown no radioactivity, ^151Eu was found to undergo very slow α-decay with a half-life on the order of 4×10^18 years Practically speaking, both are treated as stable for chemistry. Hundreds of other radioisotopes of Eu can be made artificially; the longest-lived are ^150Eu (t½≈36.9 y), ^152Eu (13.5 y), ^154Eu (8.6 y) and ^155Eu (4.7 y) These heavier isotopes (and metastable nuclear isomers) decay by electron capture or β^– depending on mass. For example, ^155Eu is a minor fission product (half-life 4.7 y, emitting ~252 keV γ-rays) ^152Eu in particular is commonly used as a laboratory γ-ray calibration source due to its many emission lines. (Europium itself also has sharp atomic emission lines used in spectroscopy, but its main importance is in material luminescence – see below.) In geochemistry, europium’s variable valence leads to the “Europium anomaly”: under reducing conditions Eu^2+ preferentially enters plagioclase minerals, causing enriched or depleted Eu/other-REE ratios in rocks This anomaly is widely used as a tracer of magmatic processes. No long-lived radioisotope of Eu is used for radiometric dating or medical tracers, so its nuclear applications are mostly in research and calibration.

Allotropes and Typical Compounds

Europium has no special allotropes like carbon or sulfur. Pure europium metal has one normal crystalline form (the α–phase) at ambient pressures (body-centered cubic and it may transform to other phases only under extreme conditions. The chemistry of europium resembles that of the other lanthanides: its compounds are largely ionic, with Eu^3+ being the dominant cation. The oxide Eu₂O₃ is the most stable form of europium oxide; it is a white-to-pink solid with a high melting point obtained by decomposing Eu(NO₃)₃ or Eu(OH)₃ Under strong reducing conditions one can make mixed-valence Eu₃O₄ or EuO. Europium(II) oxide (EuO) is a deep-red rock-salt solid (ferromagnetic near 69 K) obtained by reducing Eu₂O₃; it and EuO can be prepared by hydrogen reduction or by decomposing Eu(OH)₂. In general EuO and other Eu^2+ compounds (like EuS) are known and have unusual magnetic and optical properties

Europium(III) halides (EuF₃, EuCl₃, EuBr₃, EuI₃) are readily formed by reacting europium metal with the halogens For example: 2 Eu + 3 Cl₂ → 2 EuCl₃ Under milder conditions (or by reduction of the trihalides), the divalent halides EuF₂, EuCl₂, EuBr₂, and EuI₂ also exist. In fact, the +2 state is quite stable in the halides: EuF₂ and EuI₂ are yellowish solids, while EuCl₂ and EuBr₂ are colorless, though the latter fluoresce bright blue under UV light Europium also forms a hydride EuH₂ (a black powder via hydrogenation of the metal) and a monomeric EuH₃ at very high H₂ pressures. Among the chalcogenides, EuS and EuSe exist (EuS can be prepared by sulfiding Eu₂O₃ and is ferromagnetic at low T). Many other Eu compounds are known (Eu₂(SO₄)₃, Eu(NO₃)₃, EuPO₄, etc.) that mirror typical Ln^3+ chemistry. Eu³⁺ is a hard (Lewis) acid and in water exists as the pale pink aquo-ion [Eu(H₂O)₉]³⁺ (coordination number ~9), which precipitates easily as Eu(OH)₃. In solution Eu^3+ behaves much like other rare-earth cations. A notable feature of europium compounds is their luminescence: many Eu³⁺ complexes fluoresce under UV excitation due to 4f–4f transitions Indeed, lipophilic β-diketonate ligands (e.g. Eu(fod)₃) are famous for producing sharp red emission.

Physical Properties

Europium metal is a good electrical conductor, like most metals: its resistivity at 25 °C is on the order of 90 ×10^−6 Ω·cm (less conductive than copper or iron). Its thermal conductivity is moderate (about 14 W/m·K at room temperature The metal is paramagnetic at ordinary temperatures (each Eu atom contributes a large f-electron moment), becoming magnetic only at very low T (Eu metal itself orders antiferromagnetically below ~90 K). Solid europium has an empirical atomic radius around 185 pm The 6s² electrons make it quite electropositive. Key physical constants include atomic weight 151.964 amu and standard atomic volume ~28.9 cm³/mol. Its specific heat is about 0.18 J/g·K and Debye temperature is relatively low, reflecting its soft lattice. Spectroscopically, europium and its ions have well-known lines: the Eu^2+ ion has broad emission bands (blue/green region) in many hosts, while Eu^3+ has sharp red emission near ~612 nm (the ^5D₀→^7F₂ “R-line”) which is exploited in lamps and lasers. (Eu atomic lines such as Eu II 6645 Å are also used by astronomers to measure r-process abundances in stars, though this is a niche application.)

Chemical Reactivity and Trends

Europium is extremely electropositive. The metal forms a dark oxide layer in air and burns readily if heated. For example, 4 Eu + 3 O₂ → 2 Eu₂O₃ (the Eu₂O₃ coating). It reacts slowly with cold water and more vigorously with hot water to give Eu(OH)₃ and hydrogen gas 2 Eu + 6 H₂O → 2 Eu(OH)₃ + 3 H₂ It also dissolves in dilute acids with evolution of hydrogen, forming Eu³⁺ solutions. (For instance, in H₂SO₄ one gets pale pink [Eu(H₂O)₉]³⁺ plus H₂ Europium reacts with all the halogens to make Eu³⁺ halides (see above) The +2 halides (EuX₂) and Eu(OH)₂ require more reducing conditions. In summary, Eu metal behaves like a very active alkaline-earth metal: it oxidizes in air and water, is readily attacked by acids, and must be stored under inert atmosphere. Its aqueous chemistry is dominated by Eu³⁺, which is a hard trivalent cation. Under basic conditions Eu(OH)₃ precipitates prominently; Eu²⁺ (whether dissolved or from Eu metal) is a strong reducing agent and oxidizes to Eu³⁺ upon exposure to oxygen or stronger oxidants.

Across the lanthanide series, europium sits in the middle of the “lanthanide contraction” trend. Its atomic and ionic sizes decrease only slightly from its nearest neighbors (e.g. Nd or Sm), but its half-filled 4f^7 configuration makes it unique chemically. Its first ionization energy (5.67 eV) is one of the lowest in the lanthanides, reflecting the stability of removing the two 6s electrons. Its electronegativity is low compared to transition metals. Thus Eu²⁺ salts (like EuCl₂) are more easily formed and more reducing than, say, Nd²⁺ (which virtually never stabilizes). Because of this, europium behaves more like a divalent earth metal than do most other rare earths. Corrosion- and passivation-wise, Eu metal quickly becomes Eu₂O₃-coated but this coating is not very protective; finely-divided Eu dust can even ignite spontaneously.

Occurrence and Production

Europium is very scarce in nature. Its crustal abundance is only on the order of 2 mg per kg (2 ppm) No europium-rich mineral is known; Eu is always a trace constituent in minerals of other rare earths. The most important ores containing europium are bastnäsite Ce,La)CO₃F] and monazite ((Ce,La)PO₄). These light-REE minerals typically contain only tenths of a percent Eu₂O₃. For example, bastnäsite from the old Mountain Pass mine (USA) had an Eu₂O₃ content around 0.1% by weight this was unusually high, roughly double a typical monazite deposit After mining such ore, europium is separated from the other rare earths by standard methods: the ore is dissolved (often by acid leaching or phosphate cracking), then solvent extraction or ion-exchange chromatography is used to isolate Eu from Ce, La, Nd, etc. The purified Eu usually comes out as Eu₂O₃ (europium oxide), which can be reduced to metal by calcium or other reductants.

Major europium producers are essentially the same as for other rare earths. China is by far the largest source (especially from Inner Mongolia bastnäsite deposits), with smaller production from countries like India (monazite sands), the USA, Australia, and others. The European Rare Earth industry has been nearly dormant, though new refineries (e.g. in France) are being planned to reduce reliance on imports. There is also growing interest in recycling europium from spent fluorescent lamps and electronics; for example, new processes aim to recover the Eu from used TV and CFL phosphors. Overall, europium supply is limited, and its cost is relatively high, reflecting its low abundance and complex extraction.

Applications and Technology

Europium’s signature feature is its luminance. Most of Eu’s commercial applications exploit phosphors or luminescent materials The breakthrough came in the early 1960s when Eu^3+-doped red phosphors were developed for color television. Before this, red glazes were dim, so TV picture cells had to reduce green/blue intensities to balance color. Introducing Eu-doped yttria or lanthanum oxysulfide gave a bright, pure red emission, revolutionizing the industry Today, europium-doped phosphors are used in CRT and plasma displays, in fluorescent lamps, and increasingly in LEDs. Two main classes of Eu phosphors produce red or blue light For example, Eu^2+-activated sulfides or silicates emit blue/green, while Eu^3+-activated oxides (like Y₂O₃:Eu or (Y,Gd)BO₃:Eu) give red. By combining Eu-doped red/blue phosphors with green/yellow ones, manufacturers obtain high-quality white LEDs or backlights. Europium is also used to dope optical glass and laser materials (e.g. Eu:Y₂O₃ or certain phosphors in laser cavities) for specialized optoelectronic applications

A notable modern use is in anti-counterfeiting and security. When the Euro currency was introduced, europium-doped fluorescent pigments were “fittingly chosen for use in … security features” of the banknotes Under UV light, these prints glow in characteristic red/orange colors that are very difficult to forge. (In fact, Europium is so closely associated with the Euro that it became a popular trivia question!). More generally, the sharp luminescence of Eu complexes is employed in fluorescent inks, paints, and sensors. On the nuclear side, ^152Eu sources are used to calibrate gamma detectors, and EuNitrides have been studied for potential spintronic devices because of their magnetic properties, but such uses remain niche.

Other uses of europium are minor. As a metal it has almost no structural or alloy applications (it is soft and reactive). It does find some use as an additive in specialized alloys or as an impurity absorber. In glassmaking, small Eu₂O₃ additions impart a red color or “flame test” distinguishers. Occasionally Eu compounds are used in phosphors for security inks and luminous paints. There are also research uses: for example, EuO and Eu-based perovskites are studied for novel magnetic or electronic effects. However, by tonnage the dominant demand for Eu is in luminescent materials for lighting and displays

Biology, Environment, and Safety

Europium has no known biological role. Like other lanthanides it is only minimally soluble in water at neutral pH, so it does not bioaccumulate in living organisms. Typical “rare-earth” exposures – say, breathing dust of Eu oxide – would behave similarly to other heavy metal particulates, potentially causing lung irritation in large doses, but Eu itself is not particularly poisonous. Toxicology tests indicate that europium salts have relatively high LD₅₀ values (low toxicity): for example, the intraperitoneal LD₅₀ of EuCl₃ in mice is about 550 mg/kg, and the oral LD₅₀ is about 5000 mg/kg These values are much higher (less toxic) than for many heavy metals. As [2] notes, there are “no clear indications that Europium is particularly toxic” Nevertheless, as with all powdered metals there are handling risks: europium metal or fine Eu dust should be protected from moisture, because it reacts and can ignite. The metal itself is pyrophoric in thin shavings. Safety standards for rare-earth elements generally suggest avoiding inhalation or ingestion of the dust. In the environment, released europium would remain largely as insoluble oxide or phosphate; it does not bio-magnify. No specific exposure limits for Eu have been set by agencies, but general heavy-metal precautions apply.

History and Etymology

Europium was isolated by the French chemist Eugène Demarçay in the late 19th century. He discovered it (around 1896) while working with samples of samarium oxide that showed unexplained spectral lines Initially he called it “Didymium Sigma” until more work confirmed it as a new element. In 1901 he formally named it “Europium” after Europe, marking a patriotic departure from the then-common practice of naming elements after discoverers or mythological concepts The naming proved prophetic: nearly a century later europium compounds were incorporated into the anti-counterfeiting features of the new Euro banknotes

Ethnographically, europium played a key role in mid-20th-century technology. The Mountain Pass deposit in California supplied much of the world’s europium in the 1960s. As mentioned, that ore famously contained only ~0.1% Eu₂O₃ by weight, but this small amount was critical for color television phosphors When color TVs became widespread, Eu-doped yttrium oxysulfide and other red phosphors dramatically improved the red component of displays. In fact, Eu’s impact on electronics and optics earned it nicknames like “the TV element.” Today its price and availability remain dominated by demand for phosphors.

Data Table – Key constants for europium:

Property	Value
Atomic number (Z)	63
Symbol	Eu
Element category	Rare-earth metal (lanthanide)
Period / Block	6 (f-block)
Most common oxidation states	+3, +2
Electron configuration	[Xe] 4f^7 6s^2
Appearance at STP	Silvery-white ductile solid
Density (20 °C)	5.24 g·cm⁻³
Melting point	822 °C (1095 K)
Boiling point	1529 °C (1802 K)
Crystal structure (metal)	Body-centered cubic (bcc)
Atomic radius (empirical)	≈185 pm
Ionic radius (Eu²⁺, CN=6)	≈131 pm
Ionic radius (Eu³⁺, CN=6)	≈109 pm
First ionization energy	5.6704 eV
Electronegativity (Pauling)	~1.2 (low)
Common isotopes (nat. occ.)	^153Eu (52.2%, stable); ^151Eu (47.8%, α-decays very slowly)
Abundance (Earth’s crust)	~2 ppm (very rare)