|Name, symbol, number||scandium, Sc, 21|
|Element category||transition metal|
|Group, period, block||3, 4, d|
|Standard atomic weight||44.955912(6) g·mol−1|
|Electron configuration||[Ar] 3d1 4s2|
|Electrons per shell||2, 8, 9, 2 (Image)|
|Density (near r.t.)||2.985 g·cm−3|
|Liquid density at m.p.||2.80 g·cm−3|
|Melting point||1814 K, 1541 °C, 2806 °F|
|Boiling point||3109 K, 2836 °C, 5136 °F|
|Heat of fusion||14.1 kJ·mol−1|
|Heat of vaporization||332.7 kJ·mol−1|
|Specific heat capacity||(25 °C) 25.52 J·mol−1·K−1|
|Oxidation states||3, 2,
(weakly basic oxide)
|Electronegativity||1.36 (Pauling scale)|
|1st: 633.1 kJ·mol−1|
|2nd: 1235.0 kJ·mol−1|
|3rd: 2388.6 kJ·mol−1|
|Atomic radius||162 pm|
|Covalent radius||170±7 pm|
|Van der Waals radius||211 pm|
calc. 562 nΩ·m
|Thermal conductivity||(300 K) 15.8 W·m−1·K−1|
|Young's modulus||74.4 GPa|
|Shear modulus||29.1 GPa|
|Bulk modulus||56.6 GPa|
|Brinell hardness||750 MPa|
|CAS registry number||7440-20-2|
|Most stable isotopes|
|Main article: Isotopes of scandium|
Scandium (pronounced /ˈskændiəm/, SKAN-dee-əm) is a chemical element with symbol Sc and atomic number 21. A silvery-white metallic transition metal, it has historically been sometimes classified as a rare earth element, together with yttrium and the lanthanoids. In 1879, Lars Fredrik Nilson and his team, found a new element with spectral analysis, in the minerals euxenite and gadolinite from Scandinavia.
Scandium is present in most of the rare earth element and uranium deposits, but it is extracted from these ores in only a few mines worldwide. Due to the low availability and the difficulties in the preparation of metallic scandium, which was first done in 1937, it took until the 1970s before applications for scandium were developed. The positive effects of scandium on aluminium alloys were discovered in the 1970s, and its use in such alloys remains the only major application of scandium.
Dmitri Mendeleev predicted the existence of an element that he called ekaboron, with an atomic mass between 40 and 48 in 1869. Ten years later Lars Fredrik Nilson found a new element in the minerals euxenite and gadolinite from Scandinavia. He was able to prepare 2 gram of scandium oxide of high purity. He named it scandium, from the Latin Scandia meaning "Scandinavia". Nilson was apparently unaware of Mendeleev's prediction, but Per Teodor Cleve recognized the correspondence and notified Mendeleev.
Metallic scandium was produced for the first time in 1937 by electrolysis of a eutectic mixture, at 700–800 °C, of potassium, lithium, and scandium chlorides. The first pound of 99% pure scandium metal was produced in 1960. The use for aluminium alloys began in 1971, following a US patent. Aluminium-scandium alloys were also developed in the USSR.
Groups 1 to 3 of the periodic table could be written as follows:
This grouping is consistent with Mendeleev's prediction for scandium as "eka-boron". It shows that the properties of Sc will be intermediate between the properties of Al and Y, in the same way that the properties of Ca are intermediate between those of Mg and Sr. It also shows that there will be a diagonal relationship between Mg and Sc, just as there is between Be and Al.
However, in the standard periodic table boron and aluminium are placed in group 13, where the relationships above are less obvious. As to the rest of group 3, there has been controversy as to whether yttrium is in the same group as lanthanum or as lutetium. In the chemical compounds of the elements shown as group 3, above, the predominant oxidation state is +3. The ions M3+ will all have the electronic configuration of a noble gas, so it is reasonable that they should be in the same group of the periodic table. Most modern text-books place Sc, Y, La and Ac in the same periodic group.
Scandium does not have a particularly low abundance in the earth's crust. Estimates vary from 18 to 25 ppm, which is comparable to the abundance of cobalt (20–30 ppm). However, scandium is distributed sparsely and occurs in trace amounts in many minerals. Rare minerals from Scandinavia and Madagascar such as thortveitite, euxenite, and gadolinite are the only known concentrated sources of this element. Thortveitite can contain up to 45%, as scandium(III) oxide.
Scandium is more common in the sun and certain stars than on Earth. Scandium is only the 50th most common element on earth (35th most abundant in the Earth's crust), but it is the 23rd most common element in the sun.
World production of scandium is in the order of 2 tonnes per year as scandium oxide. The primary production is 400 kg while the rest is from stockpiles of Russia created during the Cold War. In 2003 only three mines produced scandium: the uranium and iron mines in Zhovti Vody in Ukraine, the rare earth mines in Bayan Obo, China and the apatite mines in the Kola peninsula, Russia. In each case scandium is a byproduct from the extraction of other elements. and is sold as scandium oxide. The production of metallic scandium is in the order of 10 kg per year. The oxide is converted to scandium fluoride and reduced with metallic calcium.
Madagascar and Iveland-Evje Region in Norway have the only deposits of minerals with high scandium content, thortveitite (Y,Sc)2(Si2O7) and kolbeckite ScPO4·2H2O, but these are not being exploited. Other scandium sources include the nickel and cobalt mines at Syerston and Lake Innes, New South Wales, Australia, iron, tin, and tungsten deposits in China and uranium deposits in Russia and Kazakhstan. As of 2003, scandium was not being extracted from the tailings at any of these mines, but some scandium extraction may be started if there is sufficient demand. There is currently no primary production of scandium in the Americas, Europe, or Australia.
Naturally occurring scandium is composed of one stable isotope 45Sc with a nuclear spin of 7/2. 13 radioisotopes have been characterized with the most stable being 46Sc with a half-life of 83.8 days, 47Sc with a half-life of 3.35 days, and 48Sc with a half-life of 43.7 hours. All of the remaining radioactive isotopes have half lives that are less than 4 hours, and the majority of these have half-lives that are less than 2 minutes. This element also has 5 meta states with the most stable being 44mSc (t½ 58.6 h).
The isotopes of scandium range in atomic weight from 40 u (40Sc) to 54 u (54Sc). The primary decay mode at masses lower than the only stable isotope, 45Sc, is electron capture, and the primary mode at masses above it is beta emission. The primary decay products at atomic weights below 45Sc are calcium isotopes and the primary products from higher atomic weights are titanium isotopes.
Scandium metal is hard and has a silvery appearance. It develops a slightly yellowish or pinkish cast when exposed to air. It is not resistant to weathering and dissolves slowly in most dilute acids. It does not react with a 1:1 mixture of nitric acid (HNO3) and hydrofluoric acid, HF, presumably due to the formation of an impermeable passive layer on the surface of the metal. In the compounds ScB and ScC, boron and carbon are incorporated non-stoichiometrically into the lattice of the scandium.
The radii of M3+ ions in the following table
indicate why the chemistry of scandium is more closely related to that of yttrium than that of aluminium and explains why scandium has been classified as a lanthanide-like element. The oxide Sc2O3 is weakly acidic and the hydroxide Sc(OH)3 is amphoteric:
The halides ScX3 (X = Cl, Br, I) are very soluble in water, but ScF3 is insoluble. In all four halides the scandium is 6-coordinate. The halides are Lewis acids; for example, ScF3 dissolves a solution containing excess fluoride to form [ScF6]3−. This is a typical example of a complex of Sc(III) in which the coordination number is 6. In the larger Y and La ions 8- and 9- coordination are often found.
There are a few compounds known in which the oxidation state is less than 3. The cluster [Sc6Cl12]3− has a similar structure to that of the Nb6Cl12 cluster in which chlorine atoms bridge the 12 edges of an octahedron of metal atoms. Other sub-halides are known. The nature of the hydride ScH2 is not yet fully understood. It appears not to be a saline hydride of Sc(II), but may be a compound of Sc(III) with two hydrides and an electron which is delocalized in a kind of metallic structure. ScH can be observed spectroscopically at high temperatures in the gas phase.
Scandium forms a series of organometallic compounds with cyclopentadienyl, based on the Sc(Cp)2 motif. The chlorine-bridged dimer, [Sc(Cp)2Cl]2 is the starting point for the preparation of many compounds by replacement of the chlorine.
The addition of scandium to aluminium limits the excessive grain growth that occurs in the heat-affected zone of welded aluminium components. This has two beneficial effects: the precipitated Al3Sc forms smaller crystals than are formed in other aluminium alloys and the volume of precipitate-free zones that normally exist at the grain boundaries of age-hardening aluminium alloys is reduced. Both of these effects increase the usefulness of the alloy. However, titanium alloys, which are similar in lightness and strength, are cheaper and much more widely used.
The main application of scandium by weight is in aluminium-scandium alloys for minor aerospace industry components. These alloys contain between 0.1% and 0.5% of scandium. They were used in the Russian military aircraft Mig 21 and Mig 29.
Some items of sports equipment, which rely on high performance materials, have been made with scandium-aluminium alloys, including baseball bats, lacrosse sticks, as well as bicycle frames and components. U.S. gunmaker Smith & Wesson produces revolvers with frames composed of scandium alloy and cylinders of titanium.
Approximately 20 kg (as Sc2O3) of scandium is used annually in the United States to make high-intensity discharge lamps. Scandium iodide, along with sodium iodide, when added to a modified form of mercury-vapor lamp, produces a form of metal halide lamp, an artificial light source which produce a very white light with high color rendering index that sufficiently resembles sunlight to allow good color-reproduction with TV cameras. About 80 kg of scandium is used in metal halide lamps/light bulbs globally per year. The first scandium based metal halide lamps were patented by General Electric and initially made in North America, although they are now produced in all major industrialized countries. The radioactive isotope 46Sc is used in oil refineries as a tracing agent. Scandium triflate is a catalytic Lewis acid used in organic chemistry.
Elemental scandium is not considered to be toxic. Little animal testing of scandium compounds has been done. The median lethal dose (LD50) levels for scandium(III) chloride for rats have been determined and were intraperitoneal 4 mg/kg and oral 755 mg/kg. In the light of these results compounds of scandium should be handled as compounds of moderate toxicity.
SCANDIUM [[[symbol]] Sc, atomic weight 44.1 (0=16)1, one of the rare earth metals. It was isolated in 1879 by L. F. Nilson and was shown by Cleve to be identical with the ekaboron predicted by D. Mendeleeff. The separation of scandium from wolframite (which contains 0.34-0.16% of rare earths) is given by R. J. Meyer (Zeit, anorg. Chem. 1908, 60, p. 134), but it seems impossible to obtain a perfectly pure specimen of the oxide. The salts of scandium are all colourless, the chloride and bromide corresponding in composition to Sc 2 X 6.12H 2 0; the fluoride is anhydrous. The sulphate combines with the alkaline sulphates to form double salts of the type Sc 2 (SO 4) 3.3K 2 SO 4. A large number of salts, both of inorganic and organic acids, have been described by Sir W. Crookes (Phil. Trans. 3908, 209, A. p. 3s); those of the fatty acids are in most cases more soluble in cold than in hot water.
Scandium is a metal in a group known as the transition metals. It is also a rare earth metal. What this means is that there is not very much scandium found in the earth. Because of this, the pure metal can be expensive. The pure metal is very reactive, and will react with other elements like oxygen. The metal turns from shiny to gray.
Scandium is not very dangerous because there is not much of it on Earth, so there is not enough of it to harm us. It does not have many uses. Its main use is perhaps as a component in Mercury-vapor lamps. Such lamps are used to light Stadiums.