Difference between revisions of "Europium"
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| − | '''Europium''' is a [[lanthanide]] with the symbol '''Eu''' and atomic number 63. It is a steel-gray metal about as reactive as calcium. Although difficult to find and rather expensive, | + | '''Europium''' is a [[lanthanide]] with the symbol '''Eu''' and atomic number 63. It is a steel-gray metal about as reactive as [[calcium]]. Although difficult to find and rather expensive, this metal and its salts have very interesting properties that make it an excellent addition to the amateur chemist's lab. Among these properties are multi-colored [[fluorescence]], [[redox|redox chemistry]] and [[paramagnetism]], brought about by the element's half-filled f-shell. Europium can exist in a +2 state in a reducing environment, which can be an excellent exercise in preparing reduced compounds, as the reaction Eu<sup>2+</sup> → Eu<sup>3+</sup> + e<sup>-</sup> is even less favored than [[Chromium|Cr]]<sup>2+</sup> → Cr<sup>3+</sup> + e<sup>-</sup>, which is a standard exercise in the lab. |
==Properties== | ==Properties== | ||
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===Physical=== | ===Physical=== | ||
| − | Freshly cut europium is grayish, but quickly develops a thick layer of yellow and graphite-colored oxides. It is a relatively soft metal, and can be cut with a cutting tool. Europium(II) compounds exhibit a blue fluorescence, and europium(III) compounds exhibit a red fluorescence. The fluorescence may be enhanced by the complexation of [[Dipicolinic acid|dipicolinate]] or other planar ligands to a europium atom. | + | Freshly cut europium is grayish, but quickly develops a thick layer of yellow and graphite-colored oxides. It is a relatively soft metal, softer than other lanthanides, and can be cut with a cutting tool. Europium(II) compounds exhibit a blue fluorescence (though this varies significantly with the coordinating ligands), and europium(III) compounds exhibit a red fluorescence. The fluorescence may be enhanced by the complexation of [[Dipicolinic acid|dipicolinate]] or other planar ligands to a europium atom. |
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| + | ===Chemical=== | ||
| + | Europium is by far the most reactive lanthanide, quickly corroding in air. The highly favorable reaction Eu → Eu<sup>2+</sup> + e<sup>-</sup> and the increased stability of europium(II) accelerates the corrosion rate of the metal such that a small piece exposed to dry air will corrode within a month. The resulting yellow powder, nearly the color of mustard, slowly fades as it turns from europium(II) to europium(III), the most stable form of europium. [[Europium(II,III) oxide|A mixed oxide of europium(II) and europium(III)]] has been reported, as have [[europium(II) sulfide]], [[europium(II) chloride]], and [[europium(II) sulfate]], which is very similar to [[calcium sulfate|calcium]] and [[barium sulfate]] in that it is highly insoluble in water. | ||
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| + | Europium metal burns in air with a bright red flame, identical to that of [[samarium]] to form the oxide. When europium contacts water, it will react with water about as quickly as [[calcium]], and will form the yellow [[Europium(II) hydroxide|divalent hydroxide]], which converts to the [[Europium(III) hydroxide|trivalent hydroxide]] on further exposure to water. When dry, it absorbs carbon dioxide from the air to make [[europium(III) carbonate]]. | ||
==Availability== | ==Availability== | ||
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*Producing fluorescent crystals (blue and red) | *Producing fluorescent crystals (blue and red) | ||
*Producing triboluminescent europium tetrakis(dibenzoylmethide)triethylammonium | *Producing triboluminescent europium tetrakis(dibenzoylmethide)triethylammonium | ||
| + | *Element collecting | ||
==Safety== | ==Safety== | ||
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===Storage=== | ===Storage=== | ||
| − | Europium requires storage under an inert atmosphere or mineral oil to prevent corrosion. When exposed to air, europium corrodes very easily, and it is not possible to reduce europium compounds to europium metal without an aprotic solvent. | + | Europium requires storage under an inert atmosphere or mineral oil to prevent corrosion. To remove mineral oil, rinse it off with toluene or xylene. When exposed to air, europium corrodes very easily, and it is not possible to reduce europium compounds back to europium metal without an aprotic solvent or powerful reductants like potassium. Keep europium away from oxidizers and acids. |
===Disposal=== | ===Disposal=== | ||
| − | Since europium expensive | + | Since europium is expensive, it's best to try to recycle it. |
==References== | ==References== | ||
Latest revision as of 21:15, 11 July 2021
| General properties | |||||
|---|---|---|---|---|---|
| Name, symbol | Europium, Eu | ||||
| Appearance | Silvery white | ||||
| Europium in the periodic table | |||||
| |||||
| Atomic number | 63 | ||||
| Standard atomic weight (Ar) | 151.964(1) | ||||
| Group, block | , f-block | ||||
| Period | period 6 | ||||
| Electron configuration | [Xe] 4f7 6s2 | ||||
per shell | 2, 8, 18, 25, 8, 2 | ||||
| Physical properties | |||||
| Silvery-white | |||||
| Phase | Solid | ||||
| Melting point | 1099 K (826 °C, 1519 °F) | ||||
| Boiling point | 1802 K (1529 °C, 2784 °F) | ||||
| Density near r.t. | 5.264 g/cm3 | ||||
| when liquid, at | 5.13 g/cm3 | ||||
| Heat of fusion | 9.21 kJ/mol | ||||
| Heat of | 176 kJ/mol | ||||
| Molar heat capacity | 27.66 J/(mol·K) | ||||
| pressure | |||||
| Atomic properties | |||||
| Oxidation states | 3, 2, 1 (a mildly basic oxide) | ||||
| Electronegativity | Pauling scale: 1.2 | ||||
| energies |
1st: 547.1 kJ/mol 2nd: 1085 kJ/mol 3rd: 2404 kJ/mol | ||||
| Atomic radius | empirical: 180 pm | ||||
| Covalent radius | 198±6 pm | ||||
| Miscellanea | |||||
| Crystal structure | Body-centered cubic (bcc) | ||||
| Thermal expansion | 35.0 µm/(m·K) (poly) | ||||
| Thermal conductivity | 13.9 W/(m·K) | ||||
| Electrical resistivity | 9·10-7 Ω·m (poly) | ||||
| Magnetic ordering | Paramagnetic | ||||
| Young's modulus | 18.2 GPa | ||||
| Shear modulus | 7.9 GPa | ||||
| Bulk modulus | 8.3 GPa | ||||
| Poisson ratio | 0.152 | ||||
| Vickers hardness | 165–200 MPa | ||||
| CAS Registry Number | 7440-53-1 | ||||
| History | |||||
| Naming | After Europe | ||||
| Discovery and first isolation | Eugène-Anatole Demarçay (1896, 1901) | ||||
Europium is a lanthanide with the symbol Eu and atomic number 63. It is a steel-gray metal about as reactive as calcium. Although difficult to find and rather expensive, this metal and its salts have very interesting properties that make it an excellent addition to the amateur chemist's lab. Among these properties are multi-colored fluorescence, redox chemistry and paramagnetism, brought about by the element's half-filled f-shell. Europium can exist in a +2 state in a reducing environment, which can be an excellent exercise in preparing reduced compounds, as the reaction Eu2+ → Eu3+ + e- is even less favored than Cr2+ → Cr3+ + e-, which is a standard exercise in the lab.
Contents
Properties
Physical
Freshly cut europium is grayish, but quickly develops a thick layer of yellow and graphite-colored oxides. It is a relatively soft metal, softer than other lanthanides, and can be cut with a cutting tool. Europium(II) compounds exhibit a blue fluorescence (though this varies significantly with the coordinating ligands), and europium(III) compounds exhibit a red fluorescence. The fluorescence may be enhanced by the complexation of dipicolinate or other planar ligands to a europium atom.
Chemical
Europium is by far the most reactive lanthanide, quickly corroding in air. The highly favorable reaction Eu → Eu2+ + e- and the increased stability of europium(II) accelerates the corrosion rate of the metal such that a small piece exposed to dry air will corrode within a month. The resulting yellow powder, nearly the color of mustard, slowly fades as it turns from europium(II) to europium(III), the most stable form of europium. A mixed oxide of europium(II) and europium(III) has been reported, as have europium(II) sulfide, europium(II) chloride, and europium(II) sulfate, which is very similar to calcium and barium sulfate in that it is highly insoluble in water.
Europium metal burns in air with a bright red flame, identical to that of samarium to form the oxide. When europium contacts water, it will react with water about as quickly as calcium, and will form the yellow divalent hydroxide, which converts to the trivalent hydroxide on further exposure to water. When dry, it absorbs carbon dioxide from the air to make europium(III) carbonate.
Availability
Europium is more common than iodine on Earth, but it is hard to find and extremely expensive. Places like Sigma-Aldrich charge $1000 for five measly grams.[1] One source for europium, as well as other rare earth metals, is Metallium. It is sold in 5 gram (thankfully only $50) and 25 gram sizes, as well as rods, ampoules, and coins. Metallium also takes custom orders. Europium and its compounds may be occasionally found on eBay. Other places such as NewMet will sell europium rods, sheets, foils at any size and no minimum order, though the price is on request.
Projects
- Preparing reduced europium compounds
- Producing fluorescent crystals (blue and red)
- Producing triboluminescent europium tetrakis(dibenzoylmethide)triethylammonium
- Element collecting
Safety
Safety
Europium metal, especially as a dust, should always be kept away from water and open flames. Europium fires can be identified by their bright red flames. Never use water to put out europium fire, as this will aggravate it. Class D fire extinguishers are recommended for this type of fire.
Storage
Europium requires storage under an inert atmosphere or mineral oil to prevent corrosion. To remove mineral oil, rinse it off with toluene or xylene. When exposed to air, europium corrodes very easily, and it is not possible to reduce europium compounds back to europium metal without an aprotic solvent or powerful reductants like potassium. Keep europium away from oxidizers and acids.
Disposal
Since europium is expensive, it's best to try to recycle it.
