SCANDIUM: HISTORY, PROPERTIES, REACTIONS, RISKS AND USES - CHEMISTRY - 2025

The scandium is a transition metal whose chemical symbol is Sc is the first of the transition metals in the periodic table, but is also one of the least common rare earth elements.; 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.

Ultrapure elemental scandium sample. Source: Hi-Res Images of Chemical Elements

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; so its extraction is expensive, even when it is not a metal with 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 ³⁺ cation. Scandium is a relatively strong acid, and it can form very stable coordination bonds with the oxygen atoms of organic molecules.

History

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 their traces thanks to the study of spectroscopic analysis (atomic emission spectrum).

From the minerals, he and his team succeeded in obtaining 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%.

Structure and electronic configuration

Elemental scandium (native and pure) can crystallize into two structures (allotropes): the 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 make it 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:

3d ¹ 4s ²

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 their 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.

High pressure phase

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.

Oxidation numbers

Scandium can lose up to a maximum of its three valence electrons (3d ¹ 4s ²). In theory, the first to "go" are those in the 4s orbital.

Thus, assuming the existence of the Sc ⁺ cation in the compound, its oxidation number is +1; which is the same as saying that he lost an electron from the 4s orbital (3d ¹ 4s ¹).

If it is Sc ²⁺, its oxidation number will be +2, and it will have lost two electrons (3d ¹ 4s ⁰); and if it is Sc ³⁺, the most stable of these cations, it will have an oxidation number of +3, and it is isoelectronic to argon.

In short, their oxidation numbers are: +1, +2, and +3. For example, in Sc ₂ O ₃ the oxidation number of scandium is +3 because the existence of Sc ³⁺ (Sc ₂ ³⁺ O ₃ ^2-) is assumed.

Properties

Physical appearance

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 (Sc ₂ O ₃).

Molar mass

44.955 g / mol.

Melting point

1541 ° C.

Boiling point

2836 ° C.

Molar heat capacity

25.52 J / (mol · K).

Heat of fusion

14.1 kJ / mol.

Heat of vaporization

332.7 kJ / mol.

Thermal conductivity

66 µΩ · cm at 20 ° C.

Density

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).

Electronegativity

1.36 on the Pauling scale.

Ionization energies

First: 633.1 kJ / mol (Sc ⁺ gaseous).

Second: 1235.0 kJ / mol (Sc ²⁺ gaseous).

Third: 2388.6 kJ / mol (Sc ³⁺ gas).

Atomic radio

162 pm.

Magnetic order

Paramagnetic.

Isotopes

Of all the isotopes of scandium, ⁴⁵ Sc occupies almost 100% of the total abundance (this is reflected in its atomic weight very close to 45 u).

The others consist of radioisotopes with different half-lives; such as ⁴⁶ Sc (t _1/2 = 83.8 days), ⁴⁷ Sc (t _1/2 = 3.35 days), ⁴⁴ Sc (t _1/2 = 4 hours), and ⁴⁸ Sc (t _1/2 = 43.7 hours). Other radioisotopes have t _1/2 less than 4 hours.

Acidity

The Sc ³⁺ cation is a relatively strong acid. For example, in water it can form the aqueous complex ³⁺, which in turn can turn the pH to a value below 7, because it generates H ₃ O ⁺ ions as a product of its hydrolysis:

³⁺ (aq) + H ₂ O (l) <=> ²⁺ (aq) + H ₃ O ⁺ (aq)

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.

Coordination number

An important property of scandium is that its coordination number, in most of its inorganic compounds, structures or organic crystals, is 6; it means that the Sc is surrounded by six neighbors (or forms six bonds). Above, complex aqueous 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).

An example of the latter we have al, which forms a chain structure with the AcO groups (acetyloxy or acetoxy) acting as bridges between the Sc atoms.

Nomenclature

Because almost by default the oxidation number of scandium in most of its compounds is +3, it is considered as unique and the nomenclature is therefore significantly simplified; very similar as it happens with alkali metals or aluminum itself.

For example, consider its oxide, Sc ₂ O ₃. 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.

Sc ₂ O ₃ 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.

Biological role

Scandium, for the moment, lacks a defined biological role. That is, it is unknown how the body can accumulate or assimilate Sc ³⁺ ions; which specific enzymes can use it as a cofactor, if it exerts an influence on cells, albeit similar, to the Ca ²⁺ or Fe ³⁺ ions.

It is known, however, that Sc ³⁺ ions exert antibacterial effects possibly by interfering with the metabolism of Fe ³⁺ ions.

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 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.

Where to find and production

Minerals and stars

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 approximates that of cobalt and beryllium; For every ton of rocks, 22 grams of scandium can be extracted.

The problem is that their 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 to have no preference for any of the typical mineral-forming anions (such as carbonate, CO ₃ ^2-, or sulfide, S ^2-).

It is not in its pure state. Nor is its most stable oxide, Sc ₂ O ₃, 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, Sc ³⁺ ions may be incorporated as impurities in some gemstones, such as aquamarine, or in uranium mines. And in the sky, within the stars, this element is ranked number 23 in abundance; quite high if the entire Cosmos is considered.

Industrial waste and waste

It has just been said that scandium can also be found as an impurity. For example, it is found in TiO ₂ pigments; 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.

Metallurgical reduction

The tremendous difficulties surrounding the extraction of scandium, which took so long to obtain in the native or metallic state, were due to the fact that Sc ₂ O ₃ is difficult to reduce; even more than TiO ₂, since Sc ^{3+ shows} a greater affinity than Ti ⁴⁺ towards O ^2- (assuming 100% ionic character in their respective oxides).

That is, it is easier to de-oxygen TiO ₂ than Sc ₂ O ₃ with a good reducing agent (typically carbon or alkali or alkaline earth metals). That is why Sc ₂ O ₃ is first transformed into a compound whose reduction is less problematic; such as scandium fluoride, ScF ₃. Next, ScF ₃ is reduced with metallic calcium:

2ScF ₃ (s) + 3Ca (s) => 2Sc (s) + 3CaF ₂ (s)

Sc ₂ O ₃ either 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 the rocks.

Electrolysis

Another method to produce scandium is to first obtain its chloride salt, ScCl ₃, 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.

Reactions

Amphotericism

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) => 2ScCl ₃ (aq) + 3H ₂ (g)

In doing so, it behaves like a base (reacts with HCl). But, in the same way it reacts with strong bases, such as sodium hydroxide:

2Sc (s) + 6NaOH (aq) + 6H ₂ O (l) => 2Na ₃ Sc (OH) ₆ (aq) + 3H ₂ (g)

And now it behaves like an acid (reacts with NaOH), to form a scandate salt; that of sodium, Na ₃ Sc (OH) ₆, with the scandate anion, Sc (OH) ₆ ^3-.

Oxidation

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) + 3O ₂ (g) => 2Sc ₂ O ₃ (s)

Halides

Scandium reacts with all halogens to form halides of the general chemical formula ScX ₃ (X = F, Cl, Br, etc.).

For example, it reacts with iodine according to the following equation:

2Sc (s) + 3I ₂ (g) => 2ScI ₃ (s)

In the same way it reacts with chlorine, bromine and fluorine.

Hydroxide formation

Metallic scandium can dissolve in water to produce its respective hydroxide and hydrogen gas:

2Sc (s) + 6H ₂ O (l) => 2Sc (OH) ₃ (s) + H ₂ (g)

Acid hydrolysis

Aqueous ³⁺ complexes can be hydrolyzed in such a way that they end up forming Sc- (OH) -Sc bridges, until defining a cluster with three scandium atoms.

Risks

In addition to its biological role, the exact physiological and toxicological effects of scandium are unknown.

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 ingestion 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 so 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.

Applications

Alloys

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 articles, 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 Al ₂₀ Li ₂₀ Mg ₁₀ Sc ₂₀ Ti ₃₀, which is as strong as titanium, as light as aluminum, and hard as ceramic.

3D printing

Sc-Al alloys have been used to make metallic 3D prints, in order to place or add layers of them on a pre-selected solid.

Stadium illuminations

The lighthouses in the stadiums mimic the sunlight thanks to the action of scandium iodide together with mercury vapors. Source: Pexels.

Scandium iodide, ScI ₃, 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 used for electrical devices such as digital cameras, television screens, or computer monitors. Likewise, headlights with such _3- Hg ScI lamps have been located in film and television studios.

Solid oxide fuel cells

SOFC, for its acronym in English (solid oxide fuel cell) use an oxide or ceramic as electrolytic medium; in this case, a solid that contains scandium ions. Its use in these devices is due to its great electrical conductivity and ability to stabilize temperature increases; so they work without overheating.

An example of one such solid oxide is scandium stabilized zirconite (as Sc ₂ O ₃, again).

Ceramics

Scandium carbide and titanium make up a ceramic of exceptional hardness, second only to that of diamonds. However, its use is restricted to materials with very advanced applications.

Organic coordination crystals

Sc ³⁺ ions 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 like heterogeneous catalysts; or to house neutral molecules, behaving like a solid storage.

Likewise, such organic scandium coordination crystals can be used to design sensory materials, molecular sieves, or ion conductors.

References

1. Irina Shtangeeva. (2004). Scandium. Saint Petersburg State University Saint Petersburg. Recovered from: researchgate.net
2. Wikipedia. (2019). Scandium. Recovered from: en.wikipedia.org
3. The Editors of Encyclopaedia Britannica. (2019). Scandium. Encyclopædia Britannica. Recovered from: britannica.com
4. Dr. Doug Stewart. (2019). Scandium Element Facts. Chemicool. Recovered from: chemicool.com
5. Scale. (2018). Scandium. Recovered from: scale-project.eu
6. Helmenstine, Anne Marie, Ph.D. (July 03, 2019). An Overview of Scandium. Recovered from: thoughtco.com
7. Kist, AA, Zhuk, LI, Danilova, EA, & Makhmudov, EA (2012). On question of biological role of scandium. Recovered from: inis.iaea.org
8. WAGrosshans, YKVohra & WBHolzapfel. (1982). High pressure phase transformations in yttrium and scandium: Relation to rare earths and actinides crystal structures. Journal of Magnetism and Magnetic Materials Volume 29, Issues 1–3, Pages 282-286 doi.org/10.1016/0304-8853(82)90251-7
9. Marina O. Barsukova et al. (2018). Scandium-organic frameworks: progress and prospects. Russ. Chem. Rev. 87 1139.
10. Investing News Network. (November 11, 2014). Scandium Applications: An Overview. Dig Media Inc. Recovered from: investingnews.com