Europium – facts and properties

Position in the periodic table

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Occurrence and extraction

Europium makes up only a small share of the ores monazite and bastnaesite, yet it is still more common in the Earth's crust than silver. The basics of the rare earths are covered on the lanthanum page.

One peculiarity makes it easier to separate: europium is easily reduced to doubly charged ions. Eu²⁺ is about as large as strontium and, like strontium, forms a sparingly soluble sulfate, while its triply charged neighbours stay dissolved. That lets chemists pull europium selectively out of a mixture of its neighbours.

Geologists exploit the same behaviour. Because Eu²⁺ can take the place of calcium in feldspar, many rocks contain more or less europium than their neighbours in the series would suggest. This europium anomaly shows whether feldspar crystallised from a melt or was removed from it, in rocks on Earth as well as in samples from the Moon.

The French chemist Eugène-Anatole Demarçay had already noticed spectral lines in samarium samples in the 1890s that did not belong to samarium. After countless crystallisations he managed to separate the new element in 1901 and named it after Europe.

Properties and reactions

Europium is the seventh lanthanide and sits in period 6, between samarium on its left and gadolinium on its right. Its electron configuration, [Xe] 4f⁷ 6s², ends in a half-filled 4f subshell, which is especially stable, so in the metal each atom gives up only two electrons instead of three like most lanthanides. As a result the atoms are packed loosely. The density of 5.24 g/cm³ is the lowest of any lanthanide, and the metal is about as soft as lead. The melting point is 822 °C and the boiling point 1529 °C.

Europium is the most reactive lanthanide. In air it quickly covers itself with oxide, turnings ignite at roughly 150 to 180 degrees Celsius, and with cold water it reacts much like calcium, giving hydrogen and europium hydroxide. It is stored under mineral oil or in argon.

Both ions occur in compounds. Eu³⁺ is the more common and is colourless to pale pink. Eu²⁺ is a mild reducing agent that air oxidises to Eu³⁺ in solution, although it is very stable in some solids.

The glow of the two ions has different causes. Eu³⁺ emits sharp red lines near 610 nanometres, because electrons jump within the shielded 4f subshell, so the colour hardly depends on the host material. In Eu²⁺ an electron changes between 4f and 5d states. The light is broad and its colour depends on the crystal that holds the ion; in lamp phosphors it is blue.

Uses in everyday life and technology

Red light was long the weak point of picture tubes. Europium phosphors made the red in the colour televisions of the 1960s much brighter. Fluorescent and low-energy lamps mix three phosphors: a red yttrium oxide doped with europium, a green terbium material and a blue europium material. The ultraviolet radiation of the mercury vapour in the tube excites them, and the three colours add up to white light. Since 2023 the EU has phased most such lamps out in favour of LEDs. White LEDs and the backlights of screens also contain europium phosphors that convert blue LED light into other colours.

A second field is protection against forgery. Euro banknotes contain fibres and inks with europium compounds that glow under UV light. This glow is one of the features that shop staff check with a UV lamp at the till. Such phosphors are hard to imitate with ordinary inks.

In biochemistry, europium complexes serve as labels. Their glow lasts milliseconds, far longer than the natural glow of biological samples. Measuring with a short delay, once the interfering light has died away, therefore gives a clean signal.

Europium and its compounds have low toxicity. For the metal, easy ignition is the main risk.

Quick question

What is the atomic number of europium (Eu)?

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