Vanadium was discovered in 1801 by Spanish-Mexican mineralogist Andrés Manuel del Río and named panchromium or erythronium. However, it was subsequently considered to be impure chromium.

The element was rediscovered in 1830 by Swedish chemist Nils Gabriel Sefström and named after Vanadis, the Scandinavian goddess of beauty and youth. The name was suggested because of the beautiful colors of vanadium compounds in solution.
English chemist Henry Enfield Roscoe first isolated the metal in 1867 by hydrogen reduction of vanadium dichloride.

American chemists John Wesley Marden and Malcolm N. Rich obtained it in 1925 in a purity of 99.7 percent by reducing vanadium pentoxide V₂O₅ with metallic calcium.

At the beginning of the 20th century, it was discovered that vanadium significantly increases the strength of steel. Henry Ford used it in the Model T (1908), which established vanadium's reputation as an alloying metal.

Today, vanadium is important for redox flow batteries, high-performance steels, and special alloys.

The most stable artificial isotopes are ⁴⁸V with a half-life of 16 days and ⁴⁹V with a half-life of 330 days. These are used as tracers. All other isotopes and nuclear isomers are very unstable and decay in minutes or seconds.

90 percent of vanadium demand comes from the steel industry, where vanadium is used in various steels and alloys for tools, axles, crankshafts, gears, and other critical components, as well as in jet engines and aircraft jets.

Vanadium is becoming increasingly important for vanadium redox flow batteries (VRFB). VRFBs are particularly suitable for large-scale renewable energy storage facilities, as they enable highly secure and environmentally friendly medium- and long-term energy storage. However, the limited availability of vanadium as a raw material is a disadvantage of VRFB technology.

Other applications include high-tech applications such as superconductors, nuclear reactors, catalysts, ceramics, and glass.

Vanadinit

China is the world leader in vanadium mining, with a 60 percent share of the global market. The most important company is Pangang in Sichuan, a major titanium ore producer.

In Russia, vanadium is produced as a by-product of steel and iron ore processing in the Urals. The EVRAZ Group is a key global player alongside VSMPO-AVISMA.

The Bushveld Complex in South Africa is another important source of vanadium mining. Brazil and Australia also mine vanadium.

Global annual production is around 100,000 tons.

Physical Properties

Vanadium is a non-magnetic, tough, malleable, and distinctly steel-blue heavy metal with a density of 6.11 g/cm³. Pure vanadium is relatively soft, but becomes harder when mixed with other elements and then has high mechanical strength. In most properties, it resembles its neighbor in the periodic table, titanium. The melting point of pure vanadium is 1910 °C, but this is significantly increased by impurities such as carbon. With a carbon content of 10%, it is around 2700 °C. Like chromium and niobium, vanadium crystallizes in a body-centered cubic crystal structure with the space group and the lattice parameter a = 302.4 pm, as well as two formula units per unit cell.

Below a transition temperature of 5.13 K, vanadium becomes a superconductor. Like pure vanadium, alloys of vanadium with gallium, niobium, and zirconium are also superconductive. At temperatures below 5.13 K, vanadium, like the vanadium group metals niobium and tantalum, exhibits a previously unexplained spontaneous electrical polarization in tiny clusters of up to 200 atoms, which is otherwise only found in non-metallic substances.

Chemical Properties

Vanadium is a base metal and is capable of reacting with many non-metals. When exposed to air, it retains its metallic luster for weeks. When observed over longer periods of time, green rust becomes clearly visible. If vanadium is to be preserved, it must be stored under argon. When heated, it is attacked by oxygen and oxidized to vanadium(V) oxide. While carbon and nitrogen only react with vanadium when white-hot, the reaction with fluorine and chlorine takes place even at low temperatures.

Vanadium is usually stable at room temperature when exposed to acids and bases due to a thin, passivating oxide layer. It is only attacked by hydrofluoric acid and strongly oxidizing acids such as hot nitric acid, concentrated sulfuric acid, and aqua regia.

A preliminary test is provided by the phosphorus salt bead, in which vanadium appears characteristically green in the reduction flame. The oxidation flame is faintly yellow and therefore too unspecific.

Qualitative detection of vanadium is based on the formation of peroxovanadium ions. For this purpose, an acidic solution containing vanadium in the oxidation state +5 is mixed with a small amount of hydrogen peroxide. This forms the reddish-brown [V(O₂)]³⁺ cation. This reacts with larger amounts of hydrogen peroxide to form the faint yellow peroxovanadium acid H₃[VO₂(O₂)₂].

Vanadium can be determined quantitatively by titration. To do this, a sulfuric acid solution containing vanadium is oxidized to pentavalent vanadium with potassium permanganate and then back-titrated with an iron(II) sulfate solution and diphenylamine as an indicator. It is also possible to reduce the pentavalent vanadium present with iron(II) sulfate to the tetravalent oxidation state and then perform potentiometric titration with potassium permanganate solution.

In modern analytics, vanadium can be detected using several methods. These include, for example, atomic absorption spectrometry at 318.5 nm and spectrophotometry with N-benzoyl-N-phenylhydroxylamine as a color reagent at 546 nm.

Biological significance: Vanadium compounds have various biological significance. Vanadium is characterized by the fact that it occurs both anionically as vanadate and cationically as VO₂⁺, VO²⁺, or V³⁺. Vanadates are very similar to phosphates and have correspondingly similar effects. Since vanadate binds more strongly to suitable enzymes than phosphate, it is able to block and thus control phosphorylation enzymes. This affects, for example, sodium-potassium ATPase, which controls the transport of sodium and potassium in cells. This blockage can be quickly reversed with desferrioxamine B, which forms a stable complex with vanadate. Vanadium also influences glucose uptake. It is able to stimulate glycolysis in the liver and inhibit the competing process of gluconeogenesis. This leads to a reduction in blood glucose levels. Therefore, research is being conducted to determine whether vanadium compounds are suitable for the treatment of type 2 diabetes mellitus. However, no clear results have been found yet. In addition, vanadium also stimulates the oxidation of phospholipids and suppresses the synthesis of cholesterol by inhibiting squalene synthase, a microsomal enzyme system in the liver. Consequently, a deficiency causes increased concentrations of cholesterol and triglycerides in the blood plasma.

In plants, vanadium plays a role in photosynthesis. It is able to catalyze the reaction for the formation of 5-aminolevulinic acid without enzymes. This is an important precursor for the formation of chlorophyll.

Vanadium-containing enzymes occur in some organisms; for example, some species of bacteria have vanadium-containing nitrogenases for nitrogen fixation. These include species of the genus Azotobacter and the cyanobacterium Anabaenavariabilis. However, these nitrogenases are not as efficient as the more common molybdenum nitrogenases and are therefore only activated in the event of molybdenum deficiency. Other vanadium-containing enzymes are found in brown algae and lichens. These possess vanadium-containing haloperoxidases, which they use to build chlorine, bromine, or iodine organic compounds.

The function of vanadium, which is present in large quantities in sea squirts as metalloproteins called vanabins, is not yet known. It was originally assumed that vanadium serves as an oxygen transporter similar to hemoglobin; however, this has been proven to be incorrect.

Hazards

Like other metal dusts, vanadium dust is flammable. Vanadium and its inorganic compounds have been shown to be carcinogenic in animal experiments. They are therefore classified as carcinogenic category 2. If vanadium dust is inhaled by workers in metal smelting over a long period of time, it can lead to a condition known as vanadism. This recognized occupational disease can manifest itself in mucous membrane irritation, green-black discoloration of the tongue, and chronic bronchial, lung, and intestinal diseases.

Compounds: Vanadium can occur in compounds in various oxidation states. The most common states are +5, +4, +3, and +2, while +1, 0, −1, and −3 are less common. The most important and stable oxidation states are +5 and +4.

Aqueous solution

In aqueous solution, vanadium can easily be converted to different oxidation states. Since the various vanadium ions have characteristic colors, this results in color changes.

In acidic solution, pentavalent vanadium forms colorless VO₂⁺ ions, which initially become blue tetravalent VO²⁺ ions upon reduction. The trivalent state with V³⁺ ions is green in color, while the lowest state achievable in aqueous solution, the divalent V²⁺ ion, is gray-violet.

Oxygen compounds

The most important and stable vanadium-oxygen compound is vanadium(V) oxide V₂O₅. This orange-colored compound is used in large quantities as a catalyst for sulfuric acid production. It acts as an oxygen carrier and is reduced during the reaction to another vanadium oxide, vanadium(IV) oxide VO₂. Other well-known vanadium oxides are vanadium(III) oxide V₂O₃ and vanadium(II) oxide VO.

In alkaline solution, vanadium(V) oxide forms vanadates, salts with the anion VO₄³⁻ . In contrast to the analogous phosphates, however, the vanadate ion is the most stable form; hydrogen and dihydrogen vanadates as well as free vanadic acid are unstable and are only known in dilute aqueous solutions. When basic vanadate solutions are acidified, polyvanadates are formed instead of hydrogen vanadates, in which up to ten vanadate units are combined. Vanadates are found in various minerals, examples being vanadinite, descloisite, and carnotite.

Halogen compounds

Vanadium forms a variety of compounds with the halogens fluorine, chlorine, bromine, and iodine. In the oxidation state +5, only one compound is known, vanadium(V) fluoride. In the oxidation states +4, +3, and +2, compounds exist with all halogens; only with iodine are compounds known in the states +2 and +3. However, only the chlorides vanadium(IV) chloride and vanadium(III) chloride are technically relevant. Among other things, they serve as catalysts for the production of ethylene-propylene-diene rubber.

Vanadium oxide chlorides

Vanadium also forms mixed salts with oxygen and chlorine, known as vanadium oxide chlorides. Vanadium(III) oxide chloride, VOCl, is a yellow-brown, water-soluble powder. Vanadium(IV) oxide chloride, VOCl₂, which is used in photography and as a textile dye, consists of green, hygroscopic crystal plates that dissolve in water with a blue color. Vanadium(V) oxide chloride, VOCl₃, is a yellow liquid that is very easily hydrolyzed by water. VOCl₃ serves as a catalyst component in low-pressure ethylene polymerization.

Other vanadium compounds

In organic vanadium compounds, vanadium reaches its lowest oxidation states of 0, −I, and −III. Metallocenes, known as vanadocenes, are particularly important here. These are used as catalysts for the polymerization of alkynes.

Vanadium carbide VC is used in powder form for plasma spraying and plasma powder welding, among other things. Vanadium carbide is also added to hard metals to reduce grain growth. This produces so-called cermets, which are particularly hard and wear-resistant.

General Information
Name, Symbol, Atomic Number	Vanadium, V, 23
Series	Übergangsmetalle
Group, Period, Block	5, 4, d-block
Appearance	Steel-gray metallic, with a bluish sheen
CAS Number	7440-62-2
Mass Fraction in Earth’s Crust	0,041 %
Atomic Properties
Atomic Mass	50,9415 u
Atomic Radius (calculated)	135 (171) pm
Covalent Radius	153 pm
Electron Configuration	[Ar] 3d3 4s2
1. Ionization Energy	650,9 kJ/mol
2. Ionization Energy	1414 kJ/mol
3. Ionization Energy	2830 kJ/mol
4. Ionization Energy	4507 kJ/mol
5. Ionization Energy	6298,7 kJ/mol
Physical Properties
State at 20 °C	Solid
Crystal Structure	Body-centered cubic (bcc)
Density	6,11 g/cm3 (at 20 °C)
Mohs Hardness	7,0
Magnetism	Paramagnetic ( = 3,8 · 10−4)
Melting Point	2183 K (1910 °C)
Boiling Point	3680 K (3407 °C)
Molar Volume	8,32 · 10−6 m3/mol
Heat of Vaporization	453 kJ/mol
Heat of Fusion	21,5[5] kJ/mol
Speed of Sound	4560 m/s at 293,15 K
Specific Heat Capacity	489 J/(kg · K)
Electrical Conductivity	5 · 106 A/(V · m)
Thermal Conductivity	31 W/(m · K)
Chemical Properties
Common Oxidation States	+5, +4 ,+3 ,+2
Electronegativity (Pauling scale)	1,63
Isotopes
Isotope NH t1/2 ZA ZE (MeV) ZP 48V {syn.} 15,9735 d ε 4,012 48Ti 49V {syn.} 330 d ε 0,602 49Ti 50V 0,25 % 1,5 · 1017 a ε 2,208 50Ti β− 1,037 50Cr 51V 99,75 % Stable
Safety Information
GHS Hazard Pictograms

H- and P-Statements H: No H-statements EUH: No EUH-statements P: No P-statements Hazard labeling (powder form)

Highly flammable	Irritant
(F)	(Xi)

R- and S-phrasesR: 17-36/37/38 (powder)S: 7-26-33-37-43-60 (powder)

A compound of titanium and oxygen was discovered in 1791 by the English chemist and mineralogist William Gregor and rediscovered and named independently in 1795 by the German chemist Martin Heinrich Klaproth.

The metal was not isolated in its pure form until 1910 by New Zealand metallurgist Matthew A. Hunter through the reduction of titanium tetrachloride with sodium in an airtight steel cylinder. However, the process known as the Hunter process proved to be very inefficient, which is why the element was not used industrially for a long time and was only used in laboratories.

In 1938, Luxembourg metallurgist William Justin Kroll revolutionized titanium production by reducing titanium(IV) chloride with magnesium. The Kroll process, named after him, is still used today and enabled the commercial exploitation of titanium.

The first pilot plant was built in Boulder City in the USA in 1944. The Soviet titanium industry was born in 1953 with the founding of VSMPO-AVISMA, which is now the world leader in this field.

More than 95 percent of the world's titanium ore raw materials are used to produce titanium dioxide for the manufacture of pigments, which are mainly used in paints, paper, and plastics.

The rest of the titanium ore is processed into titanium sponge, which is the starting material for titanium metal and titanium-iron alloys.
The most important industries for titanium metal and titanium alloys are aerospace, which accounts for around half of consumption. Titanium is used in NASA and SpaceX rockets, for example. Titanium alloys are also important in military jets and combat submarines.

Due to its biocompatibility, titanium is an important material for medical technology: from artificial hip and knee joints to dental implants and pacemaker housings to surgical instruments.

In the chemical industry, titanium is used in pipes and containers for aggressive chemicals due to its high corrosion resistance.
The automotive industry is another consumer of titanium materials, especially in high-performance engines and motorsports.

This lightweight but very strong material is also used in sporting goods, eyeglass frames, architecture, and jewelry.

The two most important commercial minerals are ilmenite and rutile. Natural rutile, with a 95 percent titanium dioxide content, is significantly purer than ilmenite, whose processing requires more steps and the use of environmentally harmful chemicals. Nevertheless, the extraction of ilmenite from heavy sand dominates titanium ore production, accounting for 90 percent of the total.

China is the largest producer of titanium ores from ilmenite, followed by Mozambique and South Africa.

Australia is the largest producer of rutile, followed by South Africa.

Over 95 percent of titanium mineral goes into titanium dioxide production. The rest is used for the production of alloys and titanium metal.

The production of pure titanium is difficult due to its reactivity. Titanium cannot be obtained by the usual method of reducing the oxide with carbon, as this easily produces a very stable carbide and the metal also reacts strongly with oxygen and nitrogen at elevated temperatures. Therefore, special processes were developed that transformed titanium from a laboratory curiosity into an important commercially produced structural metal after 1950.

Global annual production of titanium sponge, the precursor to titanium metal, amounts to around 320,000 tons, while annual titanium dioxide production reaches almost 10 million tons. China is the leader in titanium sponge production, accounting for 60 percent of the world market, followed by Russia, Japan, and Kazakhstan.

The largest single producer of titanium sponge is Pangang from Sichuan, China. The Japanese manufacturer Toho Titanium and the Kazakh company Ust-Kamnogorsk are other important players.

Saudi Arabia and India are not yet significant producers, but have plans to significantly expand their titanium sponge production.

Few materials offer the strength-to-weight ratio and corrosion resistance of titanium metal. In high-strength applications, titanium competes with aluminum, composite materials, intermetallic compounds, steel, and superalloys. Aluminum, nickel, special steels, and zirconium alloys can replace titanium in applications that require corrosion resistance.
Ground calcium carbonate, precipitated calcium carbonate, kaolin, and talc compete with titanium dioxide as white pigments.

Tellurium was isolated before it was known to be an element in its own right. Around 1782, Austrian mineralogist Franz Joseph Müller von Reichenstein examined an ore called Germanic gold. From this ore, he extracted a material that defied his analytical attempts and which he called Metallum problematikum. In 1798, Martin Heinrich Klaproth confirmed the elemental nature of the substance. He named the element after the human “celestial body” Tellus (Earth).

From the mid-19th century onwards, tellurium was used in metallurgy to refine steel and copper.

In the 1930s, tellurium was used in rubber vulcanization. In the 1950s, bismuth telluride was developed for use in thermoelectric cooling devices. In the 1960s, it began to be used in semiconductors and infrared detectors. In the 1980s, the first cadmium telluride (CdTe) solar cells were developed.

At the beginning of the millennium, the first CdTe thin-film solar cells were launched on the market.

Tellurium is mainly used in the production of cadmium telluride (CdTe) for thin-film solar cells. However, CdTe solar cells account for only five percent of the global photovoltaic market.

Another important area of application is the production of bismuth telluride (BiTe), which is used in thermoelectric devices for cooling and energy generation.
Metallurgical applications include its use as an alloying additive in steel to improve machining properties.

Anode slime from copper refining is the main source of tellurium, accounting for 90 percent of total production. China is the largest producer, with a 70 percent share of the global market, followed by Japan, Russia, and Sweden.

Within the EU, tellurium is mined as a by-product of gold extraction at the Kankberg mine in Sweden and refined at the nearby Rönnskär smelter owned by the Boliden Group.

Global annual production of tellurium amounts to between 900 and 1000 tons.

Several materials can replace tellurium in most applications, but usually with a loss of efficiency or product properties.
Amorphous silicon and copper indium gallium diselenide are the two main competitors to cadmium telluride in thin-film solar cells.
Bismuth selenide and organic polymers can replace BiTe in some thermoelectric components.
Bismuth, calcium, lead, phosphorus, selenium, and sulfur can replace tellurium in machine steel. Some of the chemical reactions catalyzed by tellurium can be carried out with other catalysts or by non-catalyzed processes.
In rubber compounds, sulfur or selenium can act as vulcanizing agents instead of tellurium.
The selenides and sulfides of niobium and tantalum can serve as electrically conductive solid lubricants instead of the tellurides of these metals.

In 1802, Swedish chemist Anders Gustaf Ekeberg discovered niobium. Due to the great chemical similarity between niobium and tantalum, determining the individual identity of the two elements posed a vexing problem. Therefore, chemists named the element after the mythological figure Tantalus, who suffered great torment.
It was not until 1844 that German chemist Heinrich Rose was able to prove that they were two different elements, naming the second element niobium after Niobe, the daughter of Tantalus.

From the beginning of the 20th century, tantalum was increasingly used in the electronics industry, especially for capacitors due to its ability to form stable oxide layers. From the 1940s onwards, tantalum developed into an important raw material for aerospace and medical technology.

Tantalum is an indispensable material for high-tech applications due to its extremely high corrosion resistance, biocompatibility, and excellent electrical performance.

Sixty to seventy percent is used in the electronics industry. The most important application for tantalum here is in capacitors. Tantalum capacitors offer high capacity in a small space, long service life, and stability, and are often used in miniaturized electrical circuits. They are installed in almost all smartphones, laptops, tablets, and consumer electronics. Tantalum is also used in the semiconductor industry, for example in chip coatings.

Its biocompatibility makes it interesting for use in medical implants and instruments.
In the chemical industry, its acid resistance is highly valued. It is used in acid-resistant equipment and in coatings for steel or titanium components.

Tantalum alloys are used in the aerospace industry in engine components that have to withstand extreme heat, as well as in satellites and heat shields.

In nuclear technology, tantalum absorbs neutrons. Tantalum is also used in some superconducting alloys.

Coltan (columbite-tantalite) is by far the most important tantalum ore. It is found mainly in the Democratic Republic of Congo, Rwanda, Brazil, and Australia. Tantalite is the purer tantalum variant of coltan. Microlite is often found in pegmatites and contains calcium in addition to tantalum. Wodginite is the ore found in the Wodgina mine in Australia.

The largest confirmed tantalum reserves are in Brazil and Australia, but the Congo dominates current production, which is characterized by small-scale mining. Rwanda also exports very large quantities of tantalum ore, but most of it is said to come from the conflict areas in eastern Congo.

The largest tantalum mine is Greenbushes in Australia, where lithium is also mined. Brazil is one of the top three tantalum producers in the world. Global Advanced Metals (GAM) processes tantalum from its mines in Brazil. The Australian company Pilbara Minerals operates the Pilgangoora lithium mine, where tantalum is produced as a by-product.

An important but controversial player is Mining Minerals Resources from Rwanda. In China, Ningxia Non-Ferrous Metals is a major processor of African and Chinese tantalum.
Global annual production of tantalum amounts to around 2,000 tons. However, official figures underestimate the actual quantities due to widespread illegal mining in Central Africa.

Tantalum can be replaced, but this may result in performance losses or higher costs.
Niobium and tungsten can be used in carbides, aluminum, ceramics, and niobium in electronic capacitors.
Glass, molybdenum, nickel, niobium, platinum, stainless steel, titanium, and zirconium replace tantalum in corrosion-resistant applications.
Hafnium, iridium, molybdenum, niobium, rhenium, and tungsten substitute tantalum in high-temperature applications.