The scandium it is a transition metal whose chemical symbol is Sc. It is the first of the transition metals in the periodic table, but it is also one of the less common elements of the rare earths; although its properties may resemble those of lanthanides, not all authors approve of classifying it in such a way.
At the popular level, it is a chemical element that goes unnoticed. Its name, born from the rare earth minerals from Scandinavia, may current next to copper, iron or gold. However, it is still impressive, and the physical properties of its alloys can compete with those of titanium..
Also, more and more steps are being made in the world of technology, especially in terms of lighting and lasers. Anyone who has observed a lighthouse radiating a light similar to that of the sun, will have indirectly witnessed the existence of scandium. Otherwise it is a promising item for aircraft manufacturing.
The main problem facing the scandium market is that it is widely dispersed, and there are no minerals or rich sources of it; therefore, its extraction is expensive, even when it is not a metal with a low abundance in the earth's crust. In nature it is found as its oxide, a solid that cannot be easily reduced.
In a large part of its compounds, inorganic or organic, it participates in the bond with an oxidation number of +3; that is, assuming the presence of the Sc cation3+. Scandium is a relatively strong acid, and can form very stable coordination bonds with the oxygen atoms of organic molecules..
Article index
Scandium was recognized as a chemical element in 1879, by the Swiss chemist Lars F. Nilson. He worked with the minerals euxenite and gadolinite with the intention of obtaining the yttrium contained in them. He discovered that there was an unknown element in his traces thanks to the study of spectroscopic analysis (atomic emission spectrum).
From the minerals, he and his team managed to obtain the respective scandium oxide, a name received for having surely collected the samples from Scandinavia; minerals that by then were called rare earths.
However, eight years earlier, in 1871, Dmitri Mendeleev had predicted the existence of scandium; but with the name of ekaboro, which meant that its chemical properties were similar to those of boron.
And it was in fact the Swiss chemist Per Teodor Cleve who attributed scandium to ekaboro, thus being the same chemical element. Specifically, the one that begins the block of transition metals in the periodic table.
Many years passed when in 1937, Werner Fischer and his collaborators, managed to isolate metallic scandium (but impure), by means of the electrolysis of a mixture of potassium, lithium and scandium chlorides. It was not until 1960 that it could finally be obtained with a purity around 99%..
Elemental scandium (native and pure) can crystallize into two structures (allotropes): compact hexagonal (hcp) and the body-centered cubic (bcc). The first is usually referred to as the α phase, and the second the β phase..
The denser, hexagonal α phase is stable at ambient temperatures; while the less dense cubic β phase is stable above 1337 ºC. Thus, at this last temperature a transition occurs between both phases or allotropes (in the case of metals).
Note that although scandium normally crystallizes into an hcp solid, it does not mean that it is a very dense metal; at least, yes more than aluminum. From its electronic configuration it can be known which electrons normally participate in its metallic bond:
[Ar] 3d1 4stwo
Therefore, the three electrons of the 3d and 4s orbitals intervene in the way in which the Sc atoms are located in the crystal..
To compact into a hexagonal crystal, the attraction of its nuclei must be such that these three electrons, weakly shielded by the electrons of the inner shells, do not stray too far from the Sc atoms and, consequently, the distances between them are narrowed..
The α and β phases are associated with changes in temperature; However, there is a tetragonal phase, similar to that of the metal niobium, Nb, which results when the metallic scandium undergoes a pressure greater than 20 GPa.
Scandium can lose up to a maximum of its three valence electrons (3d14stwo). In theory, the first to "go" are those in the 4s orbital..
Thus, assuming the existence of the cation Sc+ in the compound, its oxidation number is +1; which is the same as saying that he lost an electron from the 4s orbital (3d14s1).
If it is the Sctwo+, your oxidation number will be +2, and you will have lost two electrons (3d14s0); and if it is the Sc3+, the most stable of these cations, will have an oxidation number of +3, and is isoelectronic to argon.
In short, their oxidation numbers are: +1, +2, and +3. For example, in the SctwoOR3 the oxidation number of scandium is +3 because the existence of Sc is assumed3+ (Sctwo3+OR3two-).
It is a silvery white metal in its pure and elemental form, with a soft and smooth texture. It acquires yellowish-pink tones when it begins to be covered with a layer of oxide (SctwoOR3).
44.955 g / mol.
1541 ºC.
2836 ºC.
25.52 J / (mol K).
14.1 kJ / mol.
332.7 kJ / mol.
66 μΩ cm at 20 ºC.
2.985 g / mL, solid, and 2.80 g / mL, liquid. Note that its solid state density is close to that of aluminum (2.70 g / mL), which means that both metals are very light; but scandium melts at a higher temperature (the melting point of aluminum is 660.3 ºC).
1.36 on the Pauling scale.
First: 633.1 kJ / mol (Sc+ gaseous).
Second: 1235.0 kJ / mol (Sctwo+ gaseous).
Third: 2388.6 kJ / mol (Sc3+ gaseous).
162 pm.
Paramagnetic.
Of all the isotopes of scandium, Four. FiveSc occupies almost 100% of the total abundance (this is reflected in its atomic weight very close to 45 u).
The rest consist of radioisotopes with different half-lives; As the 46Sc (t1/2 = 83.8 days), 47Sc (t1/2 = 3.35 days), 44Sc (t1/2 = 4 hours), and 48Sc (t1/2 = 43.7 hours). Other radioisotopes have t1/2 less than 4 hours.
The cation Sc3+ it is a relatively strong acid. For example, in water it can form the aqueous complex [Sc (HtwoOR)6]3+, which in turn can turn the pH to a value below 7, because it generates H ions3OR+ as a product of its hydrolysis:
[Sc (HtwoOR)6]3+(ac) + HtwoO (l) <=> [Sc (HtwoOR)5OH]two+(ac) + H3OR+(ac)
The acidity of scandium can also be interpreted according to the Lewis definition: it has a high tendency to accept electrons and, therefore, to form coordination complexes.
An important property of scandium is that its coordination number, both in most of its inorganic compounds, structures or organic crystals, is 6; that is, the Sc is surrounded by six neighbors (or forms six bonds). Above, the complex aqueous [Sc (HtwoOR)6]3+ is the simplest example of all.
In crystals, the centers of Sc are octahedral; either interacting with other ions (in ionic solids), or with covalently bonded neutral atoms (in covalent solids).
Example of the latter we have [Sc (OAc)3], which forms a chain structure with the AcO groups (acetyloxy or acetoxy) acting as bridges between the Sc atoms.
Because almost by default the oxidation number of scandium in most of its compounds is +3, it is considered unique and the nomenclature is therefore significantly simplified; very similar as it happens with alkali metals or aluminum itself.
For example, consider your rust, SctwoOR3. The same chemical formula indicates in advance the oxidation state of +3 for scandium. Thus, to call this compound scandium, and like others, the systematic, stock and traditional nomenclatures are used..
The SctwoOR3 It is then scandium oxide, according to the stock nomenclature, omitting (III) (although it is not its only possible oxidation state); scandic oxide, with the suffix -ico at the end of the name according to traditional nomenclature; and diescandium trioxide, obeying the rules of the Greek numerical prefixes of the systematic nomenclature.
Scandium, for the moment, lacks a defined biological role. That is, it is unknown how the body can accumulate or assimilate Sc ions3+; which specific enzymes can use it as a cofactor, if it exerts an influence on cells, albeit similar, to Ca ionstwo+ o Faith3+.
It is known, however, that Sc ions3+ exert antibacterial effects possibly by interfering with Fe ion metabolism3+.
Some statistical studies within medicine possibly link it to stomach disorders, obesity, diabetes, cerebral leptomeningitis, and other diseases; but without sufficiently enlightening results.
Likewise, plants do not usually accumulate appreciable amounts of scandium in their leaves or stems, but rather in their roots and nodules. Therefore, it can be argued that its concentration in biomass is poor, indicative of little participation in its physiological functions and, consequently, it ends up accumulating more in soils..
Scandium may not be as abundant as other chemical elements, but its presence in the earth's crust exceeds that of mercury and some precious metals. In fact, its abundance is close to that of cobalt and beryllium; for every ton of rocks, 22 grams of scandium can be extracted.
The problem is that its atoms are not located but scattered; that is, there are no minerals that are precisely rich in scandium in their mass composition. Therefore, it is said that it has no preference for any of the typical mineral-forming anions (such as carbonate, CO3two-, or sulfur, Stwo-).
It is not in its pure state. Nor is its most stable oxide, SctwoOR3, which combines with other metals or silicates to define minerals; such as thortveitite, euxenite, and gadolinite.
These three minerals (rare in themselves) represent the main natural sources of Scandium, and are found in regions of Norway, Iceland, Scandinavia and Madagascar..
Otherwise, the ions Sc3+ they can be incorporated as impurities in some gemstones, such as aquamarine, or in uranium mines. And in the sky, within the stars, this element ranks number 23 in abundance; quite high if the entire Cosmos is considered.
It has just been said that scandium can also be found as an impurity. For example, it is found in TiO pigmentstwo; in the waste from uranium processing, as well as its radioactive minerals; and in bauxite residues in the production of metallic aluminum.
It is also found in nickel and cobalt laterites, the latter being a promising source of scandium in the future..
The tremendous difficulties surrounding the extraction of scandium, and which took so long to obtain in the native or metallic state, were due to the SctwoOR3 it is hard to reduce; even more than TiOtwo, for showing the Sc3+ an affinity greater than that of Ti4+ towards the Otwo- (assuming 100% ionic character in their respective oxides).
That is, it is easier to remove oxygen from TiOtwo than to SctwoOR3 with a good reducing agent (typically carbon or alkali or alkaline earth metals). That is why the SctwoOR3 it is first transformed into a compound whose reduction is less problematic; such as scandium fluoride, ScF3. Next, the ScF3 is reduced with metallic calcium:
2ScF3(s) + 3Ca (s) => 2Sc (s) + 3CaFtwo(s)
The SctwoOR3 Either it comes from the minerals already mentioned, or it is a by-product of the extractions of other elements (such as uranium and iron). It is the commercial form of scandium, and its low annual production (15 tons) reflects the high costs of processing, in addition to its extraction from rocks..
Another method of producing scandium is to first obtain its chloride salt, ScCl3, and then subject it to electrolysis. Thus, metallic scandium is produced in one electrode (like a sponge), and chlorine gas is produced in the other.
Scandium not only shares with aluminum the characteristics of being light metals, but they are also amphoteric; that is, they behave like acids and bases.
For example, it reacts, like many other transition metals, with strong acids to produce salts and hydrogen gas:
2Sc (s) + 6HCl (aq) => 2ScCl3(aq) + 3Htwo(g)
In doing so, it behaves like a base (reacts with HCl). But, it reacts in the same way with strong bases, such as sodium hydroxide:
2Sc (s) + 6NaOH (aq) + 6HtwoO (l) => 2Na3Sc (OH)6(aq) + 3Htwo(g)
And now it behaves like an acid (reacts with NaOH), to form a scandate salt; that of sodium, Na3Sc (OH)6, with the scandate anion, Sc (OH)63-.
When exposed to air, scandium begins to oxidize to its respective oxide. The reaction is accelerated and autocatalyzed if a heat source is used. This reaction is represented by the following chemical equation:
4Sc (s) + 3Otwo(g) => 2SctwoOR3(s)
Scandium reacts with all halogens to form halides of the general chemical formula ScX3 (X = F, Cl, Br, etc.).
For example, it reacts with iodine according to the following equation:
2Sc (s) + 3Itwo(g) => 2ScI3(s)
In the same way it reacts with chlorine, bromine and fluorine.
Metallic scandium can dissolve in water to give rise to its respective hydroxide and hydrogen gas:
2Sc (s) + 6HtwoO (l) => 2Sc (OH)3(s) + Htwo(g)
The aqueous complexes [Sc (HtwoOR)6]3+ can be hydrolyzed in such a way that they end up forming Sc- (OH) -Sc bridges, until defining a cluster with three scandium atoms.
It is unknown, in addition to its biological role, what exactly are the physiological and toxicological effects of scandium.
In its elemental form it is believed to be non-toxic, unless its finely divided solid is inhaled, thereby causing damage to the lungs. Likewise, its compounds are attributed zero toxicity, so the intake of their salts in theory should not represent any risk; as long as the dose is not high (tested in rats).
However, the data regarding these aspects is very limited. Therefore, it cannot be assumed that any of the scandium compounds are truly non-toxic; even less if the metal can accumulate in soils and waters, then passing to plants, and to a lesser extent, to animals.
At the moment, scandium still does not represent a palpable risk compared to heavier metals; such as cadmium, mercury and lead.
Although the price of scandium is high compared to other metals such as titanium or yttrium itself, its applications end up being worth the efforts and investments. One of them is to use it as an additive for aluminum alloys..
In this way, Sc-Al alloys (and other metals) retain their lightness, but become even more resistant to corrosion, at high temperatures (they do not crack), and are as strong as titanium.
So much so is the effect that scandium has on these alloys, that it is enough to add it in trace amounts (less than 0.5% by mass) for its properties to improve drastically without observing an appreciable increase in its weight. It is said that, if used massively one day, it could reduce the weight of aircraft by 15-20%.
Likewise, scandium alloys have been used for the frames of revolvers, or for the manufacture of sporting goods, such as baseball bats, special bicycles, fishing rods, golf clubs, etc .; although titanium alloys tend to replace them because they are cheaper.
The best known of these alloys is AltwentyLitwentyMg10SctwentyYou30, which is as strong as titanium, as light as aluminum and as hard as ceramic.
Sc-Al alloys have been used to make metallic 3D prints, with the purpose of placing or adding layers of them on a preselected solid.
Scandium iodide, ScI3, it is added (along with sodium iodide) to mercury vapor lamps to create artificial lights that mimic the sun. That is why in stadiums or some sports fields, even at night, the lighting inside them is such that they provide the sensation of watching a game in broad daylight..
Similar effects have been intended for electrical devices such as digital cameras, television screens, or computer monitors. Also, headlights with such lamps from ScI3-Hg have been located in film and television studios.
SOFC, for its acronym in English (solid oxide fuel cell) use an oxide or ceramic as the electrolytic medium; in this case, a solid containing scandium ions. Its use in these devices is due to its great electrical conductivity and ability to stabilize temperature increases; so they work without getting too hot.
An example of one such solid oxide is scandium stabilized zirconite (in the form of SctwoOR3, again).
Scandium carbide and titanium make up a ceramic of exceptional hardness, only surpassed by that of diamonds. However, its use is restricted to materials with very advanced applications..
Sc ions3+ can coordinate with multiple organic ligands, especially if they are oxygenated molecules.
This is because the Sc-O bonds formed are very stable, and therefore end up building crystals with amazing structures, in whose pores chemical reactions can be triggered, behaving as heterogeneous catalysts; or to host neutral molecules, behaving as a solid storage.
Likewise, such organic scandium coordination crystals can be used to design sensory materials, molecular sieves, or ion conductors..
Yet No Comments