Europium (Eu)

Half of europium's electron behaviour follows from an arrangement of seven electrons. The 4f subshell holds fourteen when full, and a subshell exactly half full — one electron in each of its seven orbitals, all spins aligned — is unusually stable. Europium's neutral atom has precisely that.

The consequence is that europium is reluctant to give up a third electron the way the rest of its row does, because doing so would break the half-filled set. Taking that third electron out costs more energy for europium than for any other lanthanide except ytterbium, whose 4f shell is one short of full and which resists for the mirror reason. Those two are the only members of the series with a genuinely easy route out of the +3 state, and europium's is the more comfortable exit: its 2+ ion persists in ordinary aqueous solution, where the divalent ions of ytterbium and samarium reduce the water instead.

In the metal, the same reluctance means only two electrons per atom join the conduction band instead of three. Weaker metallic bonding follows, and the lattice sits loose: a mole of europium takes up around 29 cubic centimetres where its neighbours samarium and gadolinium manage on about 20, which is nearly half as much room again for an atom of essentially the same mass. That is where the density anomaly comes from. Europium is the least dense lanthanide by a wide margin — nothing else in the series falls below six grams per cubic centimetre — and it is the only one whose metal packs body-centred cubic at room temperature, an open arrangement, rather than into one of the close-packed structures the rest of the series adopts. It is also the most reactive of the fifteen, tarnishing instantly and reacting with water in the manner of a heavy alkaline earth rather than a rare earth.

The anomaly is worth stating precisely, because it is often overstated. Europium does not hold the series' melting-point record: cerium and ytterbium both freeze below it. Nor does it have the largest atoms — the radii tabulated here put ytterbium, lanthanum and praseodymium ahead of it. What divalency actually buys is volume and reactivity, and those two are where europium stands apart. Its boiling point is a second, quieter symptom: europium and ytterbium are the only lanthanides that boil below 2,000 K, because there is less metallic bonding to break.

The missing red

Early colour television had a specific, quantifiable problem: the red phosphor was bad.

A colour picture tube generates its image from three phosphors, and the white balance is set by the weakest of them. Through the 1950s the available reds — manganese-activated zinc phosphate and similar — were inefficient and broad-spectrum, so the green and blue had to be turned down to match. Sets were dim, and the red itself was muddy.

Europium fixed it. Yttrium orthovanadate activated with Eu³⁺, introduced in the mid-1960s and followed by yttrium oxide and oxysulfide versions, emits a narrow, intense line at around 611 nanometres — a saturated, unambiguous red at high efficiency. Picture brightness roughly doubled. Colour cathode-ray tubes built afterwards used a europium red almost without exception, and the technology's whole visual character, the reason 1970s television looks the way it does, rests on this one ion.

Why the colour is so pure

Lanthanide emission lines are narrow in a way that ordinary phosphors' are not, and the reason is structural. The 4f electrons doing the emitting sit inside the filled 5s and 5p shells, shielded from the crystal around them. In most luminescent materials the emitting electrons are in outer orbitals that the host lattice pushes and pulls on, smearing the emission into a broad band. Europium's 4f electrons barely notice the lattice, so they emit almost as a free atom would: sharp lines at fixed wavelengths, nearly independent of what material they are dissolved in.

That is why lanthanide phosphors give clean primary colours, and why the same europium red works in a picture tube, a fluorescent lamp and a security ink without retuning.

Both ends of a fluorescent lamp's spectrum

Compact fluorescent lighting depended on a three-band phosphor blend, and europium supplies two of the bands in two different oxidation states. Eu³⁺ gives the red. Eu²⁺, in barium magnesium aluminate, gives a strong blue — and Eu²⁺ emits a broad band rather than sharp lines, because its excited state involves a 5d electron that is exposed to the crystal field. The green comes from terbium.

Divalent europium also drives the modern glow-in-the-dark material. Strontium aluminate doped with Eu²⁺ and a co-dopant, commercialised in the early 1990s, is roughly ten times brighter and far longer-lasting than the zinc sulfide it replaced, and it is what allowed watch dials and emergency signage to abandon radioactive luminous paint entirely.

What a banknote knows

Security printing wants markers that are invisible in daylight, emit a specific and hard-to-copy colour under ultraviolet light, and cannot be bought in bulk. Narrow-line lanthanide emitters fit that description exactly, and euro banknotes fluoresce in a pattern that spectroscopic studies published shortly after the currency's introduction attributed to europium complexes — distinct emissions in different regions of the note.

The European Central Bank does not publish its formulations, so the specific compounds remain inference from measured spectra rather than confirmed fact. The general principle is not in doubt: europium and its neighbours appear widely in anti-counterfeiting inks, in document security and in product authentication, precisely because their emission lines are sharp enough to be verified by a machine and awkward to reproduce with dyes.

The europium anomaly

Geochemists plot the rare earth elements in order and expect a smooth curve, because chemically the series changes gradually. Europium routinely breaks the curve, and where it breaks tells you what happened to the rock.

Under reducing conditions europium becomes Eu²⁺, and the divalent ion is close in size and charge to Ca²⁺ and Sr²⁺. That lets it substitute into plagioclase feldspar readily, where the trivalent rare earths enter only grudgingly and in far smaller proportion. So plagioclase scavenges europium out of a melt: rocks made of accumulated plagioclase show a positive europium spike, and any melt left behind after plagioclase has crystallised shows a corresponding deficit.

The most consequential application of this is the Moon. The lunar highlands are anorthosite — almost pure plagioclase — and they carry a strong positive europium anomaly, while the potassium- and phosphorus-rich basalts known as KREEP carry a strong negative one. The two are complementary, which is exactly what would happen if the early Moon were largely molten and plagioclase floated to the top of that magma ocean while the residual liquid sank. The lunar magma ocean model rests substantially on europium arithmetic.

Discovery, and the isotope that is not quite stable

The element was glimpsed before it was obtained. William Crookes and then Boisbaudran noticed spectral lines in samarium-gadolinium fractions that belonged to neither. Eugène-Anatole Demarçay, a spectroscopist of exceptional skill, separated the responsible element in pure form in 1901 and named it after the continent.

Natural europium comes in two isotopes, one of which turned out in 2007 not to be stable after all. Europium-151 emits alpha particles on a half-life around 5 × 10¹⁸ years, so the element as tabulated is very slightly radioactive at a level no ordinary measurement could detect. Its reactor-produced relatives are more consequential: europium-152 emits gamma rays at a dozen or so well-spaced energies from about 120 keV to 1.4 MeV, which is exactly what is wanted for plotting the efficiency curve of a gamma spectrometer across its whole range from one source.

From shortage to glut in ten years

Europium's market history is an unusually sharp illustration of the rare earth industry's central awkwardness. Around 2010 to 2011, with compact fluorescent lamps selling in the billions and Chinese export quotas tightening, europium became one of the most expensive rare earths in the world and lighting manufacturers reformulated products to use less of it.

Then LEDs displaced fluorescent lighting almost completely, phosphor demand collapsed, and europium's price fell by an order of magnitude. Because europium can only be obtained by processing ore mined for the rare earths that are actually wanted, supply did not fall with it. Producers now have more europium than there are uses for — the mirror image of the dysprosium problem at the other end of the series, and the same underlying cause.