The materials to prepare students for

The materials to prepare students for practical lessons of inorganic chemistry

Lesson № 25

Theme: d-elements of the ІІІВ – VIIВ groups. Titanium, Vanadium. Actinides. Sub- group of Manganese.

 

Vanadium is a chemical element with the symbol V and atomic number 23. It is a hard, silvery gray, ductile and malleable transition metal. The element is found only in chemically combined form iature, but once isolated artificially, the formation of an oxide layer stabilizes the free metal somewhat against further oxidation.

Andrés Manuel del Río discovered compounds of vanadium in 1801 by analyzing a new lead-bearing mineral he called “brown lead,” and presumed its qualities were due to the presence of a new element, which he named erythronium (Greek for “red”) since, upon heating, most of its salts turned from their initial color to red. Four years later, however, he was convinced by other scientists that erythronium was identical to chromium. Chlorides of vanadium were generated in 1830 by Nils Gabriel Sefström who thereby proved that a new element was involved, which he named “vanadium” after the Germanic goddess of beauty and fertility, Vanadís (Freyja). Both names were attributed to the wide range of colors found in vanadium compounds. Del Rio’s lead mineral was later renamed vanadinite for its vanadium content. In 1867 Henry Enfield Roscoe obtained the pure element.

Vanadium occurs naturally in about 65 different minerals and in fossil fuel deposits. It is produced in China and Russia from steel smelter slag; other countries produce it either from the flue dust of heavy oil, or as a byproduct of uranium mining. It is mainly used to produce specialty steel alloys such as high speed tool steels. The most important industrial vanadium compound, vanadium pentoxide, is used as a catalyst for the production of sulfuric acid.

Large amounts of vanadium ions are found in a few organisms, possibly as a toxin. The oxide and some other salts of vanadium have moderate toxicity. Particularly in the ocean, vanadium is used by some life forms as an active center of enzymes, such as the vanadium bromoperoxidase of some ocean algae. Vanadium is probably a micronutrient in mammals, including humans, but its precise role in this regard is unknown.

The isolation of vanadium metal proved difficult. In 1831, Berzelius reported the production of the metal, but Henry Enfield Roscoe showed that Berzelius had in fact produced the nitride, vanadium nitride (VN). Roscoe eventually produced the metal in 1867 by reduction of vanadium(II) chloride, VCl₂, with hydrogen. In 1927, pure vanadium was produced by reducing vanadium pentoxide with calcium. The first large-scale industrial use of vanadium in steels was found in the chassis of the Ford Model T, inspired by French race cars. Vanadium steel allowed for reduced weight while simultaneously increasing tensile strength.

Characteristics

Vanadium is a hard, ductile, silver-gray metal. Some sources describe vanadium as “soft”, perhaps because it is ductile, malleable and not brittle. Vanadium is harder than most metals and steels (see Hardnesses of the elements and iron). It has good resistance to corrosion and it is stable against alkalis, sulfuric and hydrochloric acids. It is oxidized in air at about 933 K (660 °C, 1220 °F), although an oxide layer forms even at room temperature.

Isotopes

Naturally occurring vanadium is composed of one stable isotope ⁵¹V and one radioactive isotope ⁵⁰V. The latter has a half-life of 1.5×10¹⁷ years and a natural abundance 0.25%. ⁵¹V has a nuclear spin of 7/2 which is useful for NMR spectroscopy. 24 artificial radioisotopes have been characterized, ranging in mass number from 40 to 65. The most stable of these isotopes are ⁴⁹V with a half-life of 330 days, and ⁴⁸V with a half-life of 16.0 days. All of the remaining radioactive isotopes have half-lives shorter than an hour, most of which are below 10 seconds. At least 4 isotopes have metastable excited states. Electron capture is the main decay mode for isotopes lighter than the ⁵¹V. For the heavier ones, the most common mode is beta decay. The electron capture reactions lead to the formation of element 22 (titanium) isotopes, while for beta decay, it leads to element 24 (chromium) isotopes.

Chemistry and compounds

Oxidation states of vanadium, from left +2 (lilac), +3 (green), +4 (blue) and +5 (yellow).

The chemistry of vanadium is noteworthy for the accessibility of the four adjacent oxidation states 2-5. In aqueous solution the colours are lilac V²⁺(aq), green V³⁺(aq), blue VO²⁺(aq) and, at high pH, yellow VO₄^2-. Vanadium(II) compounds are reducing agents, and vanadium(V) compounds are oxidizing agents. Vanadium(IV) compounds often exist as vanadyl derivatives which contain the VO²⁺ center.

Ammonium vanadate(V) (NH₄VO₃) can be successively reduced with elemental zinc to obtain the different colors of vanadium in these four oxidation states. Lower oxidation states occur in compounds such as V(CO)₆, [V(CO)₆]⁻ and substituted derivatives.

The vanadium redox battery utilizes all four oxidation states; one electrode uses the +5/+4 couple and the other uses the +3/+2 couple. Conversion of these oxidation states is illustrated by the reduction of a strongly acidic solution of a vanadium(V) compound with zinc dust or amalgam. The initial yellow color characteristic of the pervanadyl ion [VO₂(H₂O)₄]⁺ is replaced by the blue color of [VO(H₂O)₅]²⁺, followed by the green color of [V(H₂O)₆]³⁺ and then the violet color due to [V(H₂O)₆]²⁺.

The most commercially important compound is vanadium pentoxide, which is used as a catalyst for the production of sulfuric acid. This compound oxidizes sulfur dioxide (SO₂) to the trioxide (SO₃). In this redox reaction, sulfur is oxidized from +4 to +6, and vanadium is reduced from +5 to +3:

V₂O₅ + 2 SO₂ → V₂O₃ + 2 SO₃

The catalyst is regenerated by oxidation with air:

V₂O₃ + O₂ → V₂O₅

 

Oxy-anions and cations the decavanadate structure

The oxyanion chemistry of vanadium(V) is complex: the predominance diagram for vanadates in aqueous solution shows at least 11 species to be predominant under specified conditions of pH and concentration. The tetrahedral vanadate ion, VO³⁻₄, is the principal species present at pH 12-14. On acidification, the monomer [HVO₄]^2- and dimer [V₂O₇]^– are formed, with the monomer predominant at vanadium concentration of less than ca. 10⁻²M (pV > 2; pV is equal to minus the logarithm of the total vanadium concentration/M). The formation of the divanadate ion is analogous to the formation of the dichromate ion. As the pH is reduced, further protonation and polymerization to polyvanadates occur: at pH 4-6 [H₂VO₄]^– is predominant at pV greater than ca. 4, while at higher concentrations trimers and tetramers are formed. Between pH 2-4 decavanadates predominate. In decavanadates there is a distorted octahedron of oxygen atoms around each vanadium atom. Vanadic acid, H₃VO₄ has a very low concentration because protonation of the tetrahedral species [H₂VO₄]^– results in the preferential formation of the octahedral [VO₂(H₂O)₄]⁺ species. In strongly acidic solutions, pH<2. [VO₂(H₂O)₄]⁺ is the predominant species, while the oxide V₂O₅ precipitates from solution at high concentrations. The oxide is formally the inorganic anhydride of vanadic acid. The structures of many of these and other vanadate ions have been determined by X-ray crystallography of crystalline compounds.

The acid dissociation constants for the vanadium and phosphorus series are remarkably similar. Chains, rings and clusters involving tetrahedral vanadium, analogous to the polyphosphates, are known. The correspondence between vanadate and phosphate chemistry can be attributed to the similarity in size and charge of phosphorus(V) and vanadium(V). Orthovanadate VO³⁻₄ is used in protein crystallography to study the biochemistry of phosphate.

The Pourbaix diagram for vanadium in water

The Pourbaix diagram for vanadium in water, which shows the redox potentials between various vanadium species in different oxidation states is also complex.

Vanadium(V) also forms various peroxo-complexes. The species VO(O)₂(H₂O)₄⁺ is present in acidic solutions. In alkaline solutions species with 2, 3 and 4 peroxide groups are present; the last forms violet crystals M₃V(O₂)₄ nH₂O (M=Li, Na, K, NH₄⁺), in which the vanadium has an 8-coordinate dodecahedral structure.

Chalcogenide and halide compounds

Vanadium forms a very large variety of binary compounds with sulfur, selenium and tellurium, often with complicated structures. The tetrahedral sulfa-anion [VS₄]^3-, analogous to the orthovanadate ion, is well-known, but there are no thio-analogues of the polymeric oxo-vanadates.

All four halides are known for oxidation states +2 and +3, but the iodide is not known for V(IV) and VF₅ is the only halide known for oxidation state 5. VCl₄ may be used as a catalyst for polymerization of dienes.

Examples of oxyhalides include.

· vanadium(V): VOF₃, VOX₃ and VO₂X (X=F, Cl)

· vanadium(IV): VOX₂ (X=F, Cl, Br)

· vanadium(III): VOX, (X=Cl,Br)

Coordination compounds

A ball-and-stick model of VO₅(C₅H₇)₂

Vanadium’s early position in the transition metal series lead to three rather unusual features of the coordination chemistry of vanadium. Firstly, metallic vanadium has the electronic configuration [Ar]3d³4s², so compounds of vanadium are relatively electron-poor. Consequently, most binary compounds are Lewis acids (electron pair acceptors); examples are all the halides forming octahedral adducts with the formula VX_nL_6−n (X = halide; L = other ligand). Secondly, the vanadium ion is rather large and can achieve coordinatioumbers higher than 6, as is the case in [V(CN)₇]⁴⁻. Thirdly, the vanadyl ion, VO²⁺, is featured in many complexes of vanadium(IV) such as vanadyl acetylacetonate (V(=O)₅(C₅H₇)₂). In this complex, the vanadium is 5-coordinate, square pyramidal, meaning that a sixth ligand, such as pyridine, may be attached, though the association constant of this process is small. Many 5-coordinate vanadyl complexes have a trigonal bypyramidal geometry, such as VOCl₂(NMe₃)₂.

Organometallic compounds

Organometallic chemistry of vanadium is well developed, but organometallic compounds are of minor commercial significance. Vanadocene dichloride is a versatile starting reagent and even finds minor applications in organic chemistry. Vanadium carbonyl, V(CO)₆, is a rare example of a metal carbonyl containing an unpaired electron, but which exists without dimerization. The addition of an electron yields V(CO)⁻₆ (isoelectronic with Cr(CO)₆), which may be further reduced with sodium in liquid ammonia to yield V(CO)³⁻₆ (isoelectronic with Fe(CO)₅).

Occurrence

Vanadinite

Metallic vanadium is not found iature, but is known to exist in about 65 different minerals. Economically significant examples include patronite (VS₄), vanadinite (Pb₅(VO₄)₃Cl), and carnotite (K₂(UO₂)₂(VO₄)₂·3H₂O). Much of the world’s vanadium production is sourced from vanadium-bearing magnetite found in ultramafic gabbro bodies. Vanadium is mined mostly in South Africa, north-western China, and eastern Russia. In 2010 these three countries mined more than 98% of the 56,000 tonnes of produced vanadium.

Vanadium is also present in bauxite and in fossil fuel deposits such as crude oil, coal, oil shale and tar sands. In crude oil, concentrations up to 1200 ppm have been reported. When such oil products are burned, the traces of vanadium may initiate corrosion in motors and boilers. An estimated 110,000 tonnes of vanadium per year are released into the atmosphere by burning fossil fuels. Vanadium has also been detected spectroscopically in light from the Sun and some other stars.

Production

Ferrovanadium chunks

Electrolytically refined vanadium dendritic crystals (99,9%)

Iodide refined vanadium

Most vanadium is used as an alloy called ferrovanadium as an additive to improve steels. Ferrovanadium is produced directly by reducing a mixture of vanadium oxide, iron oxides and iron in an electric furnace. Vanadium-bearing magnetite iron ore is the main source for the production of vanadium. The vanadium ends up in pig iron produced from vanadium bearing magnetite. During steel production, oxygen is blown into the pig iron, oxidizing the carbon and most of the other impurities, forming slag. Depending on the ore used, the slag contains up to 25% of vanadium.

Vanadium metal is obtained via a multistep process that begins with the roasting of crushed ore with NaCl or Na₂CO₃ at about 850 °C to give sodium metavanadate (NaVO₃). An aqueous extract of this solid is acidified to give “red cake”, a polyvanadate salt, which is reduced with calcium metal. As an alternative for small-scale production, vanadium pentoxide is reduced with hydrogen or magnesium. Many other methods are also in use, in all of which vanadium is produced as a byproduct of other processes. Purification of vanadium is possible by the crystal bar process developed by Anton Eduard van Arkel and Jan Hendrik de Boer in 1925. It involves the formation of the metal iodide, in this example vanadium(III) iodide, and the subsequent decomposition to yield pure metal.

2 V + 3 I₂ 2 VI₃

Applications

Tool made from vanadium steel

Alloys

Approximately 85% of vanadium produced is used as ferrovanadium or as a steel additive. The considerable increase of strength in steel containing small amounts of vanadium was discovered in the beginning of the 20th century. Vanadium forms stable nitrides and carbides, resulting in a significant increase in the strength of the steel. From that time on vanadium steel was used for applications in axles, bicycle frames, crankshafts, gears, and other critical components. There are two groups of vanadium containing steel alloy groups. Vanadium high-carbon steel alloys contain 0.15% to 0.25% vanadium and high speed tool steels (HSS) have a vanadium content of 1% to 5%. For high speed tool steels, a hardness above HRC 60 can be achieved. HSS steel is used in surgical instruments and tools.

Vanadium stabilizes the beta form of titanium and increases the strength and temperature stability of titanium. Mixed with aluminium in titanium alloys it is used in jet engines, high-speed airframes and dental implants. One of the common alloys is Titanium 6AL-4V, a titanium alloy with 6% aluminium and 4% vanadium.

Other uses

Vanadium(V) oxide is a catalyst in the contact process for producing sulfuric acid

Vanadium is compatible with iron and titanium, therefore vanadium foil is used in cladding titanium to steel. The moderate thermal neutron-capture cross-section and the short half-life of the isotopes produced by neutron capture makes vanadium a suitable material for the inner structure of a fusion reactor. Several vanadium alloys show superconducting behavior. The first A15 phase superconductor was a vanadium compound, V₃Si, which was discovered in 1952. Vanadium-gallium tape is used in superconducting magnets (17.5 teslas or 175,000 gauss). The structure of the superconducting A15 phase of V₃Ga is similar to that of the more common Nb₃Sn and Nb₃Ti.

The most common oxide of vanadium, vanadium pentoxide V₂O₅, is used as a catalyst in manufacturing sulfuric acid by the contact process and as an oxidizer in maleic anhydride production. Vanadium pentoxide is also used in making ceramics. Another oxide of vanadium, vanadium dioxide VO₂, is used in the production of glass coatings, which blocks infrared radiation (and not visible light) at a specific temperature. Vanadium oxide can be used to induce color centers in corundum to create simulated alexandrite jewelry, although alexandrite iature is a chrysoberyl. The possibility to use vanadium redox couples in both half-cells, thereby eliminating the problem of cross contamination by diffusion of ions across the membrane is the advantage of vanadium redox rechargeable batteries. Vanadate can be used for protecting steel against rust and corrosion by electrochemical conversion coating. Lithium vanadium oxide has been proposed for use as a high energy density anode for lithium ion batteries, at 745 Wh/L when paired with a lithium cobalt oxide cathode. It has been proposed by some researchers that a small amount, 40 to 270 ppm, of vanadium in Wootz steel and Damascus steel, significantly improves the strength of the material, although it is unclear what the source of the vanadium was.

Biological role

Vanadium plays a very limited role in biology, and is more important in ocean environments than on land.

Ascidiacea (sea squirts) contain vanadium as vanabin.

Tunicates such as this bluebell tunicate contain vanadium as vanabin.

Amanita muscaria contains amavadin.

Bromoperoxidases in algae

Organobromine compounds in a number of species of marine algae are generated by the action of a vanadium dependent bromoperoxidase. This is a haloperoxidase in algae which requires bromide and is an absolutely vanadium-dependent enzyme. Most organobromine compounds in the sea ultimately arise via the action of this vanadium bromoperoxidase.

Vanadium accumulation in tunicates and ascidians

German chemist Martin Henze discovered vanadium in the blood cells (or coelomic cells) of Ascidiacea (sea squirts) in 1911. It is essential to ascidians and tunicates, where it is stored in the highly acidified vacuoles of certain blood cell types, designated vanadocytes. Vanabins (vanadium binding proteins) have been identified in the cytoplasm of such cells. The concentration of vanadium in these ascidians’ blood is up to ten million times higher than the concentration of vanadium in the seawater around them. The function of this vanadium concentration system, and these vanadium-containing proteins, is still unknown but the vanadocytes are later deposited just under the outer surface of the tunic where their presence may deter predation.

Nitrogen fixation

A vanadium nitrogenase is used by some nitrogen-fixing micro-organisms, such as Azotobacter. In this role vanadium replaces more common molybdenum or iron, and gives the nitrogenase slightly different properties.

Fungi

Several species of macrofungi, namely Amanita muscaria and related species, accumulate vanadium (up to 500 mg/kg in dry weight). Vanadium is present in the coordination complex, amavadin, in fungal fruit-bodies. However, the biological importance of the accumulation process is unknown. Toxin functions or peroxidase enzyme functions have been suggested.

Mammals and birds

Rats and chickens are also known to require vanadium in very small amounts and deficiencies result in reduced growth and impaired reproduction. Vanadium is a relatively controversial dietary supplement, primarily for increasing insulin sensitivity and body-building. Whether it works for the latter purpose has not been proven, and there is some evidence that athletes who take it are merely experiencing a placebo effect. Vanadyl sulfate may improve glucose control in people with type 2 diabetes. In addition, decavanadate and oxovanadates are species that potentially have many biological activities and that have been successfully used as tools in the comprehension of several biochemical processes.

Safety

All vanadium compounds should be considered toxic. Tetravalent VOSO₄ has been reported to be over 5 times more toxic than trivalent V₂O₃. The Occupational Safety and Health Administration (OSHA) has set an exposure limit of 0.05 mg/m³ for vanadium pentoxide dust and 0.1 mg/m³ for vanadium pentoxide fumes in workplace air for an 8-hour workday, 40-hour work week. The National Institute for Occupational Safety and Health (NIOSH) has recommended that 35 mg/m³ of vanadium be considered immediately dangerous to life and health. This is the exposure level of a chemical that is likely to cause permanent health problems or death.

Vanadium compounds are poorly absorbed through the gastrointestinal system. Inhalation exposures to vanadium and vanadium compounds result primarily in adverse effects on the respiratory system. Quantitative data are, however, insufficient to derive a subchronic or chronic inhalation reference dose. Other effects have been reported after oral or inhalation exposures on blood parameters, on liver, oeurological development in rats, and other organs.

There is little evidence that vanadium or vanadium compounds are reproductive toxins or teratogens. Vanadium pentoxide was reported to be carcinogenic in male rats and male and female mice by inhalation in an NTP study, although the interpretation of the results has recently been disputed. Vanadium has not been classified as to carcinogenicity by the United States Environmental Protection Agency.

Vanadium traces in diesel fuels present a corrosion hazard; it is the main fuel component influencing high temperature corrosion. During combustion, it oxidizes and reacts with sodium and sulfur, yielding vanadate compounds with melting points down to 530 °C, which attack the passivation layer on steel, rendering it susceptible to corrosion. The solid vanadium compounds also cause abrasion of engine components.

TITANIUM

Titanium is a chemical element with the symbol Ti and atomic number 22. It has a low density and is a strong, lustrous, corrosion-resistant (including sea water, aqua regia and chlorine) transition metal with a silver color.

Titanium was discovered in Cornwall, Great Britain, by William Gregor in 1791 and named by Martin Heinrich Klaproth for the Titans of Greek mythology. The element occurs within a number of mineral deposits, principally rutile and ilmenite, which are widely distributed in the Earth’s crust and lithosphere, and it is found in almost all living things, rocks, water bodies, and soils. The metal is extracted from its principal mineral ores via the Kroll process or the Hunter process. Its most common compound, titanium dioxide, is a popular photocatalyst and is used in the manufacture of white pigments. Other compounds include titanium tetrachloride (TiCl₄), a component of smoke screens and catalysts; and titanium trichloride (TiCl₃), which is used as a catalyst in the production of polypropylene.

Titanium can be alloyed with iron, aluminium, vanadium, molybdenum, among other elements, to produce strong lightweight alloys for aerospace (jet engines, missiles, and spacecraft), military, industrial process (chemicals and petro-chemicals, desalination plants, pulp, and paper), automotive, agri-food, medical prostheses, orthopedic implants, dental and endodontic instruments and files, dental implants, sporting goods, jewelry, mobile phones, and other applications.

The two most useful properties of the metal form are corrosion resistance and the highest strength-to-weight ratio of any metal. In its unalloyed condition, titanium is as strong as some steels, but 45% lighter. There are two allotropic forms and five naturally occurring isotopes of this element, ⁴⁶Ti through ⁵⁰Ti, with ⁴⁸Ti being the most abundant (73.8%). Titanium’s properties are chemically and physically similar to zirconium, as both of them have the same number of valence electrons and are in the same group in the periodic table.

History

Titanium was discovered included in a mineral in Cornwall, Great Britain, in 1791 by the clergyman and amateur geologist William Gregor, then vicar of Creed parish. He recognized the presence of a new element in ilmenite when he found black sand by a stream in the nearby parish of Manaccan and noticed the sand was attracted by a magnet. Analysis of the sand determined the presence of two metal oxides; iron oxide (explaining the attraction to the magnet) and 45.25% of a white metallic oxide he could not identify. Gregor, realizing that the unidentified oxide contained a metal that did not match the properties of any known element, reported his findings to the Royal Geological Society of Cornwall and in the German science journal Crell’s Annalen.

Around the same time, Franz-Joseph Müller von Reichenstein produced a similar substance, but could not identify it. The oxide was independently rediscovered in 1795 by Prussian chemist Martin Heinrich Klaproth in rutile from Boinik (Germaame of unknown place) village of Hungary (Now in Slovakia). Klaproth found that it contained a new element and named it for the Titans of Greek mythology. After hearing about Gregor’s earlier discovery, he obtained a sample of manaccanite and confirmed it contained titanium.

The processes required to extract titanium from its various ores are laborious and costly; it is not possible to reduce the ore in the normal manner, by heating in the presence of carbon, as that produces titanium carbide. Pure metallic titanium (99.9%) was first prepared in 1910 by Matthew A. Hunter at Rensselaer Polytechnic Institute by heating TiCl₄ with sodium at 700–800 °C in the Hunter process. Titanium metal was not used outside the laboratory until 1932 when William Justin Kroll proved that it could be produced by reducing titanium tetrachloride (TiCl₄) with calcium. Eight years later he refined this process by using magnesium and even sodium in what became known as the Kroll process. Although research continues into more efficient and cheaper processes (e.g., FFC Cambridge), the Kroll process is still used for commercial production.

Titanium of very high purity was made in small quantities when Anton Eduard van Arkel and Jan Hendrik de Boer discovered the iodide, or crystal bar, process in 1925, by reacting with iodine and decomposing the formed vapors over a hot filament to pure metal.

In the 1950s and 1960s the Soviet Union pioneered the use of titanium in military and submarine applications (Alfa Class and Mike Class) as part of programs related to the Cold War. Starting in the early 1950s, titanium began to be used extensively for military aviation purposes, particularly in high-performance jets, starting with aircraft such as the F100 Super Sabre and Lockheed A-12.

In the USA, the Department of Defense realized the strategic importance of the metal and supported early efforts of commercialization. Throughout the period of the Cold War, titanium was considered a strategic material by the U.S. government, and a large stockpile of titanium sponge was maintained by the Defense National Stockpile Center, which was finally depleted in the 2000s. According to 2006 data, the world’s largest producer, Russian-based VSMPO-Avisma, was estimated to account for about 29% of the world market share. As of 2009, titanium sponge metal was produced in six countries: China, Japan, Russia, Kazakhstan, USA and Ukraine (in order of output).

In 2006, the U.S. Defense Agency awarded $5.7 million to a two-company consortium to develop a new process for making titanium metal powder. Under heat and pressure, the powder can be used to create strong, lightweight items ranging from armor plating to components for the aerospace, transport, and chemical processing industries.

2 FeTiO₃ + 7 Cl₂ + 6 C → 2 TiCl₄ + 2 FeCl₃ + 6 CO (900 °C)

TiCl₄ + 2 Mg → 2 MgCl₂ + Ti (1100 °C)

About 50 grades of titanium and titanium alloys are designated and currently used, although only a couple of dozen are readily available commercially. The ASTM International recognizes 31 Grades of titanium metal and alloys, of which Grades 1 through 4 are commercially pure (unalloyed). These four are distinguished by their varying degrees of tensile strength, as a function ofoxygen content, with Grade 1 being the most ductile (lowest tensile strength with an oxygen content of 0.18%), and Grade 4 the least (highest tensile strength with an oxygen content of 0.40%). The remaining grades are alloys, each designed for specific purposes, be it ductility, strength, hardness, electrical resistivity, creep resistance, resistance to corrosion from specific media, or a combination thereof.

The grades covered by ASTM and other alloys are also produced to meet Aerospace and Military specifications (SAE-AMS, MIL-T), ISO standards, and country-specific specifications, as well as proprietary end-user specifications for aerospace, military, medical, and industrial applications.

In terms of fabrication, all welding of titanium must be done in an inert atmosphere of argon or helium in order to shield it from contamination with atmospheric gases such as oxygen, nitrogen, orhydrogen. Contamination will cause a variety of conditions, such as embrittlement, which will reduce the integrity of the assembly welds and lead to joint failure. Commercially pure flat product (sheet, plate) can be formed readily, but processing must take into account the fact that the metal has a “memory” and tends to spring back. This is especially true of certain high-strength alloys. Titanium cannot be soldered without first pre-plating it in a metal that is solderable. The metal can be machined using the same equipment and via the same processes asstainless steel.

Physical properties

A metallic element, titanium is recognized for its high strength-to-weight ratio. It is a strong metal with low density that is quite ductile (especially in an oxygen-free environment), lustrous, and metallic-white in color. The relatively high melting point (more than 1,650 °C or 3,000 °F) makes it useful as a refractory metal. It is paramagnetic and has fairly low electrical and thermal conductivity.

Commercial (99.2% pure) grades of titanium have ultimate tensile strength of about 63,000 psi (434 MPa), equal to that of common, low-grade steel alloys, but are 45% lighter. Titanium is 60% more dense than aluminium, but more than twice as strong as the most commonly used 6061-T6 aluminium alloy. Certain titanium alloys (e.g., Beta C) achieve tensile strengths of over 200,000 psi (1,400 MPa). However, titanium loses strength when heated above 430 °C (806 °F).

Titanium is fairly hard (although not as hard as some grades of heat-treated steel), non-magnetic and a poor conductor of heat and electricity. Machining requires precautions, as the material will soften and gall if sharp tools and proper cooling methods are not used. Like those made from steel, titanium structures have a fatigue limit which guarantees longevity in some applications. Titanium alloys have lower specific stiffnesses than in many other structural materials such as aluminium alloys and carbon fiber.

The metal is a dimorphic allotrope whose hexagonal alpha form changes into a body-centered cubic (lattice) β form at 882 °C (1,620 °F). The specific heat of the alpha form increases dramatically as it is heated to this transition temperature but then falls and remains fairly constant for the β form regardless of temperature. Similar to zirconium and hafnium, an additional omega phase exists, which is thermodynamically stable at high pressures, but is metastable at ambient pressures. This phase is usually hexagonal (ideal) or trigonal (distorted) and can be viewed as being due to a soft longitudinal acoustic phonon of the β phase causing collapse of (111) planes of atoms.

Chemical properties

The Pourbaix diagram for titanium in pure water, perchloric acid or sodium hydroxide

Like aluminium and magnesium metal surfaces, titanium metal and alloy surfaces oxidize immediately when they are exposed to air. Titanium readily reacts with oxygen at 1,200 °C (2,190 °F) in air, and at 610 °C (1,130 °F) in pure oxygen, forming titanium dioxide. However, it is slow to react with water and air, as it forms a passive and protective oxide coating that protects it from further reaction. When it first forms, this protective layer is only 1–2 nm thick but continues to slowly grow; reaching a thickness of 25 nm in four years.

When exposed to nitrogen clean titanium metal is covered with a titanium nitride layer. The thin titanium dioxide and nitride layers on titanium surfaces are very hard and inert.

The most noted chemical property of titanium is its excellent resistance to corrosion; it is almost as resistant as platinum, capable of withstanding attack by dilute sulfuric acid and hydrochloric acid as well as chlorine gas, chloride solutions, and most organic acids. However, it is soluble in concentrated acids. The Pourbaix diagram in the image shows that titanium is actually thermodynamically a very reactive metal.

The metal cannot be melted in open air since it burns before the melting point is reached. Melting is only possible in an inert atmosphere or in a vacuum. At 550 °C (1,022 °F), it combines with chlorine. It also reacts with the other halogens and absorbs hydrogen.

Titanium is one of the few elements that burns in pure nitrogen gas, reacting at 800 °C (1,470 °F) to form titanium nitride, which causes embrittlement. Titanium powder can form an explosive suspension in air.

Because clean titanium metal is highly reactive and titanium binds strongly to oxygen, nitrogen and many other gases, titanium filaments are applied in titanium sublimation pumps. Titanium sublimation pumps are inexpensive and reliable devices which are applied as the last step to create extremely low pressures in ultra-high vacuum systems.

Experiments have shown that natural titanium becomes radioactive after it is bombarded with deuterons, emitting mainly positrons and hard gamma rays.

Compounds

TiN coated drill bit

The +4 oxidation state dominates titanium chemistry, but compounds in the +3 oxidation state are also common. Because of this high oxidation state, many titanium compounds have a high degree of covalent bonding.

Star sapphires and rubies get their asterism from the titanium dioxide impurities present in them. Titanates are compounds made with titanium dioxide. Barium titanate has piezoelectric properties, thus making it possible to use it as a transducer in the interconversion of sound and electricity. Esters of titanium are formed by the reaction of alcohols and titanium tetrachloride and are used to waterproof fabrics.

Titanium nitride (TiN), having a hardness equivalent to sapphire and carborundum (9.0 on the Mohs Scale), is often used to coat cutting tools, such as drill bits. It also finds use as a gold-colored decorative finish, and as a barrier metal in semiconductor fabrication.

Titanium tetrachloride (titanium(IV) chloride, TiCl₄) is a colorless liquid which is used as an intermediate in the manufacture of titanium dioxide for paint. It is widely used in organic chemistry as a Lewis acid, for example in the Mukaiyama aldol condensation. Titanium also forms a lower chloride, titanium(III) chloride (TiCl₃), which is used as a reducing agent.

Titanocene dichloride is an important catalyst for carbon-carbon bond formation. Titanium isopropoxide is used for Sharpless epoxidation. Other compounds include titanium bromide (used in metallurgy, superalloys, and common titanium alloys are made by reduction. For example, cuprotitanium (rutile with copper added is reduced), ferrocarbon titanium (ilmenite reduced with coke in an electric furnace), and manganotitanium (rutile with manganese or manganese oxides) are reduced.

2 FeTiO₃ + 7 Cl₂ + 6 C → 2 TiCl₄ + 2 FeCl₃ + 6 CO (900 °C)

TiCl₄ + 2 Mg → 2 MgCl₂ + Ti (1100 °C)

Applications

A titanium cylinder, “Grade 2” quality

Titanium is used in steel as an alloying element (ferro-titanium) to reduce grain size and as a deoxidizer, and in stainless steel to reduce carbon content. Titanium is often alloyed with aluminium (to refine grain size), vanadium, copper (to harden), iron, manganese, molybdenum, and with other metals. Applications for titanium mill products (sheet, plate, bar, wire, forgings, castings) can be found in industrial, aerospace, recreational, and emerging markets. Powdered titanium is used in pyrotechnics as a source of bright-burning particles.

Pigments, additives and coatings

Titanium dioxide is the most commonly used compound of titanium

About 95% of titanium ore extracted from the Earth is destined for refinement into titanium dioxide (TiO₂), an intensely white permanent pigment used in paints, paper, toothpaste, and plastics. It is also used in cement, in gemstones, as an optical opacifier in paper, and a strengthening agent in graphite composite fishing rods and golf clubs.

TiO₂ powder is chemically inert, resists fading in sunlight, and is very opaque: this allows it to impart a pure and brilliant white color to the brown or gray chemicals that form the majority of household plastics. Iature, this compound is found in the minerals anatase, brookite, and rutile. Paint made with titanium dioxide does well in severe temperatures, and stands up to marine environments. Pure titanium dioxide has a very high index of refraction and an optical dispersion higher than diamond. In addition to being a very important pigment, titanium dioxide is also used in sunscreens due to its ability to protect skin by itself. Recently, titanium oxide has been put to use in air purifiers (as a filter coating), or in film used to coat windows on buildings so that when titanium oxide becomes exposed to UV light (either solar or artificial) and moisture in the air, reactive redox species like hydroxyl radicals are produced so that they can purify the air or keep window surfaces clean.

Aerospace and marine

Due to their high tensile strength to density ratio, high corrosion resistance, fatigue resistance, high crack resistance, and ability to withstand moderately high temperatures without creeping, titanium alloys are used in aircraft, armor plating, naval ships, spacecraft, and missiles. For these applications titanium alloyed with aluminium, vanadium, and other elements is used for a variety of components including critical structural parts, fire walls, landing gear, exhaust ducts (helicopters), and hydraulic systems. In fact, about two thirds of all titanium metal produced is used in aircraft engines and frames. The SR-71 “Blackbird” was one of the first aircraft to make extensive use of titanium within its structure, paving the way for its use in modern military and commercial aircraft. An estimated 59 metric tons (130,000 pounds) are used in the Boeing 777, 45 in the Boeing 747, 18 in the Boeing 737, 32 in the Airbus A340, 18 in the Airbus A330, and 12 in the Airbus A320. The Airbus A380 may use 77 metric tons, including about 11 tons in the engines. In engine applications, titanium is used for rotors, compressor blades, hydraulic system components, and nacelles. The titanium 6AL-4V alloy accounts for almost 50% of all alloys used in aircraft applications.

Due to its high corrosion resistance to sea water, titanium is used to make propeller shafts and rigging and in the heat exchangers of desalination plants; in heater-chillers for salt water aquariums, fishing line and leader, and for divers’ knives. Titanium is used to manufacture the housings and other components of ocean-deployed surveillance and monitoring devices for scientific and military use. The former Soviet Union developed techniques for making submarines largely out of titanium.

Industrial

High-purity (99.999%) titanium with visible crystal structure

Welded titanium pipe and process equipment (heat exchangers, tanks, process vessels, valves) are used in the chemical and petrochemical industries primarily for corrosion resistance. Specific alloys are used in downhole and nickel hydrometallurgy applications due to their high strength (e. g.: titanium Beta C alloy), corrosion resistance, or combination of both. The pulp and paper industry uses titanium in process equipment exposed to corrosive media such as sodium hypochlorite or wet chlorine gas (in the bleachery). Other applications include: ultrasonic welding, wave soldering, and sputtering targets.

Titanium tetrachloride (TiCl₄), a colorless liquid, is important as an intermediate in the process of making TiO₂ and is also used to produce the Ziegler-Natta catalyst, and is used to iridize glass and because it fumes strongly in moist air it is also used to make smoke screens.

Consumer and architectural

Titanium metal is used in automotive applications, particularly in automobile or motorcycle racing, where weight reduction is critical while maintaining high strength and rigidity. The metal is generally too expensive to make it marketable to the general consumer market, other than high-end products, particularly for the racing/performance market. Late model Corvettes have been available with titanium exhausts.

Titanium is used in many sporting goods: tennis rackets, golf clubs, lacrosse stick shafts; cricket, hockey, lacrosse, and football helmet grills; and bicycle frames and components. Although not a mainstream material for bicycle production, titanium bikes have been used by race teams and adventure cyclists. Titanium alloys are also used in spectacle frames. This results in a rather expensive, but highly durable and long lasting frame which is light in weight and causes no skin allergies. Many backpackers use titanium equipment, including cookware, eating utensils, lanterns, and tent stakes. Though slightly more expensive than traditional steel or aluminium alternatives, these titanium products can be significantly lighter without compromising strength. Titanium is also favored for use by farriers, since it is lighter and more durable than steel when formed into horseshoes.

Titanium has occasionally been used in architectural applications: the 40 m (131 foot) memorial to Yuri Gagarin, the first man to travel in space, in Moscow, is made of titanium for the metal’s attractive color and association with rocketry. The Guggenheim Museum Bilbao and the Cerritos Millennium Library were the first buildings in Europe and North America, respectively, to be sheathed in titanium panels. Other construction uses of titanium sheathing include the Frederic C. Hamilton Building in Denver, Colorado and the 107 m (350 foot) Monument to the Conquerors of Space in Moscow.

Because of its superior strength and light weight when compared to other metals traditionally used in firearms (steel, stainless steel, and aluminium), and advances in metalworking techniques, the use of titanium has become more widespread in the manufacture of firearms. Primary uses include pistol frames and revolver cylinders. For these same reasons, it is also used in the body of laptop computers (for example, in Apple’s PowerBook line).

Some upmarket categories of tools made to be lightweight and corrosion-resistant, such as shovels and flashlights, are made of titanium or titanium alloys as well.

Jewelry

Because of its durability, titanium has become more popular for designer jewelry (particularly, titanium rings). Its inertness makes it a good choice for those with allergies or those who will be wearing the jewelry in environments such as swimming pools. Titanium is also alloyed with gold to produce an alloy that can be marketed as 24-carat gold, as the 1% of alloyed Ti is insufficient to require a lesser mark. The resulting alloy is roughly the hardness of 14-carat gold and thus is more durable than a pure 24-carat gold item would be.

Titanium’s durability, light weight, dent- and corrosion resistance makes it useful in the production of watch cases. Some artists work with titanium to produce artworks such as sculptures, decorative objects and furniture.

The inertness and ability to be attractively colored makes titanium a popular metal for use in body piercing. Titanium may be anodized to produce various colors, which varies the thickness of the surface oxide layer and causes interference fringes.

Medical

Because it is biocompatible (it is non-toxic and is not rejected by the body), titanium is used in a gamut of medical applications including surgical implements and implants, such as hip balls and sockets (joint replacement) that can stay in place for up to 20 years. The titanium is often alloyed with about 4% aluminium or 6% Al and 4% vanadium.

Titanium has the inherent ability to osseointegrate, enabling use in dental implants that can remain in place for over 30 years. This property is also useful for orthopedic implant applications. These benefit from titanium’s lower modulus of elasticity (Young’s modulus) to more closely match that of the bone that such devices are intended to repair. As a result, skeletal loads are more evenly shared between bone and implant, leading to a lower incidence of bone degradation due to stress shielding and periprosthetic bone fractures, which occur at the boundaries of orthopedic implants. However, titanium alloys’ stiffness is still more than twice that of bone, so adjacent bone bears a greatly reduced load and may deteriorate.

Since titanium is non-ferromagnetic, patients with titanium implants can be safely examined with magnetic resonance imaging (convenient for long-term implants). Preparing titanium for implantation in the body involves subjecting it to a high-temperature plasma arc which removes the surface atoms, exposing fresh titanium that is instantly oxidized.

Titanium is also used for the surgical instruments used in image-guided surgery, as well as wheelchairs, crutches, and any other products where high strength and low weight are desirable.

Precautions

Titanium is non-toxic even in large doses and does not play any natural role inside the human body. An estimated quantity of 0.8 milligrams of titanium is ingested by humans each day, but most passes through without being absorbed. It does, however, have a tendency to bio-accumulate in tissues that contain silica. One study indicates a possible connection between titanium and yellow nail syndrome. An unknown mechanism in plants may use titanium to stimulate the production of carbohydrates and encourage growth. This may explain why most plants contain about 1 part per million (ppm) of titanium, food plants have about 2 ppm, and horsetail and nettle contain up to 80 ppm.

As a powder or in the form of metal shavings, titanium metal poses a significant fire hazard and, when heated in air, an explosion hazard. Water and carbon dioxide–based methods to extinguish fires are ineffective on burning titanium; Class D dry powder fire fighting agents must be used instead.

When used in the production or handling of chlorine, care must be taken to use titanium only in locations where it will not be exposed to dry chlorine gas which can result in a titanium/chlorine fire. A fire hazard exists even when titanium is used in wet chlorine due to possible unexpected drying brought about by extreme weather conditions.

Titanium can catch fire when a fresh, non-oxidized surface comes in contact with liquid oxygen. Such surfaces can appear when the oxidized surface is struck with a hard object, or when a mechanical strain causes the emergence of a crack. This poses the possible limitation for its use in liquid oxygen systems, such as those found in the aerospace industry.

ACTINIDE

The actinide or actinoid series encompasses the 15 chemical elements with atomic numbers from 89 to 103, actinium to lawrencium. The actinide series derives its name from the group 3 element actinium; although actinoid (actinide) means “actinium-like” and therefore should exclude actinium, it is usually included in the series for the purpose of comparison. Only thorium and uranium occur in more than trace quantities in nature; the other actinides are synthetic elements. The actinides are usually considered to be f-block elements; they show much more variable valence than the lanthanides. All actinides are radioactive and release energy upon radioactive decay; uranium and plutonium, the most abundant actinides on Earth, are used iuclear reactors and nuclear weapons.

Properties

Actinides have similar properties to lanthanides. The 6d and 7s electronic shells are completed in actinium and thorium, and the 5f shell is being filled with further increase in atomic number; the 4f shell is filled in the lanthanides. The first experimental evidence for the filling of the 5f shell in actinides was obtained by McMillan and Abelson in 1940. As in lanthanides, the ionic radius of actinides monotonically decreases with atomic number.

 

Physical properties

Actinides are typical metals. All of them are soft and have a silvery color (but tarnish in air), relatively high density and plasticity. Some of them can be cut with a knife. Their electrical resistivity varies between 15 and 150 µOhm·cm. The hardness of thorium is similar to that of soft steel, so heated pure thorium can be rolled in sheets and pulled into wire. Thorium is nearly half as dense as uranium and plutonium, but is harder than either of them. All actinides are radioactive, paramagnetic, and, with the exception of actinium, have several crystalline phases: plutonium has seven, and uranium, neptunium and californium three. Crystal structures of protactinium, uranium, neptunium and plutonium do not have clear analogs among the lanthanide and are more similar to those of the 3d-transition metals.

All actinides are pyrophoric, especially when finely divided, that is, they spontaneously ignite upon reaction with air. The melting point of actinides does not have a clear dependence on the number of f-electrons. The unusually low melting point of neptunium and plutonium (~640 °C) is explained by hybridization of 5f and 6d orbitals and the formation of directional bonds in these metals.

Chemical properties

Like the lanthanides, all actinides are highly reactive with halogens and chalcogens; however, actinides react more easily. Actinides, especially those with a small number of 5f-electrons, are prone to hybridization. This is explained by the similarity of the electron energies at the 5f, 7s and 6d shells. Most actinides exhibit a larger variety of valence states, and the most stable are +6 for uranium, +5 for protactinium and neptunium, +4 for thorium and plutonium and +3 for actinium and other actinides.

Chemically, actinium is similar to lanthanum, which is explained by their similar ionic radii and electronic structure. Like lanthanum, actinium has oxidation of +3, but it is less reactive and has more pronounced basic properties. Among other trivalent actinides Ac³⁺ is least acidic, i.e. has the weakest tendency to hydrolyze in aqueous solutions.

Thorium is rather active chemically. Owing to lack of electrons on 6d and 5f orbitals, the tetravalent thorium compounds are colorless. At pH < 3, the solutions of thorium salts are dominated by the cations [Th(H₂O)₈]⁴⁺. The Th4+ ion is relatively large, and depending on the coordinatioumber can have a radius between 0.95 and 1.14 Å. As a result, thorium salts have a weak tendency to hydrolyse. The distinctive ability of thorium salts is their high solubility, not only in water, but also in polar organic solvents.

Protactinium exhibits two valence states; the +5 is stable, and the +4 state easily oxidizes to protactinium(V). Thus tetravalent protactinium in solutions is obtained by the action of strong reducing agents in a hydrogen atmosphere. Tetravalent protactinium is chemically similar to uranium(IV) and thorium(IV). Fluorides, phosphates, hypophosphate, iodate and phenylarsonates of protactinium(IV) are insoluble in water and dilute acids. Protactinium forms soluble carbonates. The hydrolytic properties of pentavalent protactinium are close to those of tantalum(V) and niobium(V). The complex chemical behavior of protactinium is a consequence of the start of the filling of the 5f shell in this element.

Uranium has a valence from 3 to 6, the last being most stable. In the hexavalent state, uranium is very similar to the group 6 elements. Many compounds of uranium(IV) and uranium(VI) are non-stoichiometric, i.e. have variable composition. For example, the actual chemical formula of uranium dioxide is UO2+x, where x varies between −0.4 and 0.32. Uranium(VI) compounds are weak oxidants. Most of them contain the linear “uranyl” group, UO₂+2. Between 4 to 6 ligands can be accommodated in an equatorial plane perpendicular to the uranyl group. The uranyl group acts as a hard acid and forms stronger complexes with oxygen-donor ligands than with nitrogen-donor ligands. NpO₂+2 and PuO₂+2 are also the common form of Np and Pu in the +6 oxidation state. Uranium(IV) compounds exhibit reducing properties, e.g., they are easily oxidized by atmospheric oxygen. Uranium(III) is a very strong reducing agent. Owing to the presence of d-shell, uranium (as well as many other actinides) forms organometallic compounds, such as UIII(C₅H₅)₃ and UIV(C₅H₅)₄.

Neptunium has valence states from 3 to 7, which can be simultaneously observed in solutions. The most stable state in solution is +5, but the valence +4 is preferred in solid neptunium compounds. Neptunium metal is very reactive. Ions of neptunium are prone to hydrolysis and formation of coordination compounds.

Plutonium also exhibits the valence between 3 and 7, and thus is chemically similar to neptunium and uranium. It is highly reactive, and quickly forms an oxide film in air. Plutonium reacts with hydrogen even at temperatures as low as 25–50 °C; it also easily forms halides and intermetallic compounds. Hydrolysis reactions of plutonium ions of different oxidation states are quite diverse. Plutonium(V) can enter polymerization reactions.

The largest chemical diversity among actinides is observed in americium, which can have valence between 2 and 6. Divalent americium is obtained only in dry compounds and non-aqueous solutions (acetonitrile). Oxidation states +3, +5 and +6 are typical for aqueous solutions, but also in the solid state. Tetravalent americium forms stable solid compounds (dioxide, fluoride and hydroxide) as well as complexes in aqueous solutions. It was reported that in alkaline solution americium can be oxidized to the heptavalent state, but these data proved erroneous. The most stable valence of americium is 3 in the aqueous solutions and 3 or 4 in solid compounds.

Valence 3 is dominant in all subsequent elements up to lawrencium (with the possible exception of nobelium). Curium can be tetravalent in solids (fluoride, dioxide). Berkelium, along with a valence of +3, also shows the valence of +4, more stable than that of curium; the valence 4 is observed in solid fluoride and dioxide. The stability of Bk4+ in aqueous solution is close to that of Ce4+. Only valence 3 was observed for californium, einsteinium and fermium. The divalent state is proven for mendelevium and nobelium, and iobelium it is more stable than the trivalent state. Lawrencium shows valence 3 both in solutions and solids.

The redox potential  increases from −0.32 V in uranium, through 0.34 V (Np) and 1.04 V (Pu) to 1.34 V in americium revealing the increasing reduction ability of the An4+ ion from americium to uranium. All actinides form AnH₃ hydrides of black color with salt-like properties. Actinides also produce carbides with the general formula of AnC or AnC2 (U2C3 for uranium) as well as sulfides An₂S₃ and AnS₂.

Like the lanthanides, the actinides form a family of elements with similar properties. Within the actinides, there are two overlapping groups: transuranium elements, which follow uranium in the periodic table—andtransplutonium elements, which follow plutonium. Compared to the lanthanides, which (except for promethium) are found iature in appreciable quantities, most actinides are rare. The most abundant, or easy to synthesize actinides are uranium and thorium, followed by plutonium, americium, actinium, protactinium and neptunium.

The existence of transuranium elements was suggested by Enrico Fermi based on his experiments in 1934. However, even though four actinides were known by that time, it was not yet understood that they formed a family similar to lanthanides. The prevailing view that dominated early research into transuranics was that they were regular elements in the 7th period, with thorium, protactinium and uranium corresponding to 6th-period hafnium, tantalum and tungsten, respectively. Synthesis of transuranics gradually undermined this point of view. By 1944 an observation that curium failed to exhibit oxidation states above 4 (whereas its supposed 6th period neighbor, platinum, can reach oxidation state of 7) prompted Glenn Seaborg to formulate a so-called “actinide hypothesis”. Studies of known actinides and discoveries of further transuranic elements provided more data in support of this point of view, but the phrase “actinide hypothesis” (the implication being that “hypothesis” is something that has not been decisively proven) remained in active use by scientists through the late 1950s.

At present, there are two major methods of producing isotopes of transplutonium elements: irradiation of the lighter elements with either neutrons or accelerated charged particles. The first method is most important for applications, as only neutron irradiation using nuclear reactors allows the production of sizeable amounts of synthetic actinides; however, it is limited to relatively light elements. The advantage of the second method is that elements heavier than plutonium, as well as neutron-deficient isotopes, can be obtained, which are not formed during neutron irradiation.

In 1962–1966, there were attempts in the United States to produce transplutonium isotopes using a series of six underground nuclear explosions. Small samples of rock were extracted from the blast area immediately after the test to study the explosion products, but no isotopes with mass number greater than 257 could be detected, despite predictions that such isotopes would have relatively long half-lives of α-decay. This inobservation was attributed to spontaneous fission owing to the large speed of the products and to other decay channels, such as neutron emission and nuclear fission.

From actinium to neptunium

Uranium and thorium were the first actinides discovered. Uranium was identified in 1789 by the German chemist Martin Heinrich Klaproth in pitchblende ore. He named it after the planet Uranus, which had been discovered only eight years earlier. Klaproth was able to precipitate a yellow compound (likely sodium diuranate) by dissolving pitchblende iitric acid and neutralizing the solution with sodium hydroxide. He then reduced the obtained yellow powder with charcoal, and extracted a black substance that he mistook for metal. Only 60 years later, the French scientist Eugène-Melchior Péligot identified it with uranium oxide. He also isolated the first sample of uranium metal by heating uranium tetrachloride with potassium. The atomic mass of uranium was then calculated as 120, but Dmitri Mendeleev in 1872 corrected it to 240 using his periodicity laws. This value was confirmed experimentally in 1882 by K. Zimmerman.

Thorium oxide was discovered by Friedrich Wöhler in the mineral, which was found in Norway (1827). Jöns Jacob Berzeliuscharacterized this material in more detail by in 1828. By reduction of thorium tetrachloride with potassium, he isolated the metal and named it thorium after the Norse god of thunder and lightning Thor. The same isolation method was later used by Péligot for uranium.

Actinium was discovered in 1899 by André-Louis Debierne, an assistant of Marie Curie, in the pitchblende waste left after removal of radium and polonium. He described the substance (in 1899) as similar to titanium and (in 1900) as similar to thorium. The discovery of actinium by Debierne was however questioned in 1971 and 2000, arguing that Debierne’s publications in 1904 contradicted his earlier work of 1899–1900. The name for word actinium comes from the Greek aktis, aktinos (ακτίς, ακτίνος), meaning beam or ray. This metal was discovered not by its own radiation but by the radiation of the daughter products. Owing to the close similarity of actinium and lanthanum and low abundance, pure actinium could only be produced in 1950. The term actinide was probably introduced by Victor Goldschmidt in 1937.

Protactinium was possibly isolated in 1900 by William Crookes. It was first identified in 1913, when Kasimir Fajans and Oswald Helmuth Göhring encountered the short-lived isotope ^234mPa (half-life 1.17 minutes) during their studies of the ²³⁸U decay. They named the new element brevium (from Latin brevis meaning brief); the name was changed to protoactinium (from Greek πρῶτος + ἀκτίς meaning “first beam element”) in 1918 when two groups of scientists, led by the Austrian Lise Meitner and Otto Hahn of Germany and Frederick Soddy and John Cranston of Great Britain, independently discovered ²³¹Pa. The name was shortened to Protactinium in 1949. This element was little characterized until 1960, when A. G. Maddock and co-workers in UK produced 130 grams of protactinium from 60 tonnes of waste left after extraction of uranium from its ore.

Neptunium (named for the planet Neptune, the next planet out from Uranus, after which uranium was named) was discovered by Edwin McMillan and Philip H. Abelson in 1940 in Berkeley, California. They produced the ²³⁹Np isotope (half-life 2.4 days) by bombarding uranium with slow neutrons. It was the first transuranium element produced synthetically.

Thirty-one isotopes of actinium and eight excited isomeric states of some of its nuclides were identified by 2010. Three isotopes, ²²⁵Ac, ²²⁷Ac and ²²⁸Ac, were found iature and the others were produced in the laboratory; only the three natural isotopes are used in applications. Actinium-225 is a member of radioactiveneptunium series; it was first discovered in 1947 as a fission product of uranium-233, it is an α-emitter with a half-life of 10 days. Actinium-225 is less available than actinium-228, but is more promising in radiotracer applications. Actinium-227 (half-life 21.77 years) occurs in all uranium ores, but in small quantities. One gram of uranium (in radioactive equilibrium) contains only 2×10⁻¹⁰ gram of ²²⁷Ac. Actinium-228 is a member of radioactive thorium series formed by the decay of ²²⁸Ra; it is a β^– emitter with a half-life of 6.15 hours. In one tonne of thorium there is 5×10⁻⁸ gram of ²²⁸Ac. It was discovered by Otto Hahn in 1906.

Twenty nine isotopes of protactinium are known with mass numbers 212–240 as well as three excitedisomeric states. Only ²³¹Pa and ²³⁴Pa have been found iature. All the isotopes have short lifetime, except for protactinium-231 (half-life 32,760 years). The most important isotopes are ²³¹Pa and ²³³Pa, which is an intermediate product in obtaining uranium-233 and is the most affordable among artificial isotopes of protactinium. ²³³Pa has convenient half-life and energy of γ-radiation, and thus was used in most studies of protactinium chemistry. Protactinium-233 is a β-emitter with a half-life of 26.97 days.

Uranium has the highest number (25) of both natural and synthetic isotopes. They have mass numbers of 217–242, and three of them, ²³⁴U, ²³⁵U and ²³⁸U, are present in appreciable quantities iature. Among others, the most important is ²³³U, which is a final product of transformations of ²³²Th irradiated by slow neutrons. ²³³U has a very higher fission efficiency by low-energy (thermal) neutrons, compared e.g. with ²³⁵U. Most uranium chemistry studies were carried out on uranium-238 owing to its long half-life of 4.4×10⁹years.

There are 19 isotopes of neptunium with mass numbers from 225 to 244; they are all highly radioactive. The most popular among scientists are long-lived ²³⁷Np (t_½ = 2.20×10⁶ years) and short-lived ²³⁹Np, ²³⁸Np (t_½ ~ 2 days).

Sixteen isotopes of americium are known with mass numbers from 232 to 248. The most important are ²⁴¹Am and ²⁴³Am, which are alpha-emitters and also emit soft, but intense γ-rays; both of them can be obtained in an isotopically pure form. Chemical properties of americium were first studied with ²⁴¹Am, but later shifted to ²⁴³Am, which is almost 20 times less radioactive. The disadvantage of ²⁴³Am is production of the short-lived daughter isotope ²³⁹Np, which has to be considered in the data analysis.

Among 19 isotopes of curium, the most accessible are ²⁴²Cm and ²⁴⁴Cm; they are α-emitters, but with much shorter lifetime than the americium isotopes. These isotopes emit almost no γ-radiation, but undergo spontaneous fission with the associated emission of neutrons. More long-lived isotopes of curium (^245–248Cm, all α-emitters) are formed as a mixture during neutron irradiation of plutonium or americium. Upon short irradiation, this mixture is dominated by curium-246, and then curium-248 begins to accumulate. Both of these isotopes, especially ²⁴⁸Cm, have a longer half-life (3.48×10⁵ years) and are much more convenient for carrying out chemical research than ²⁴²Cm and ²⁴⁴Cm, but they also have a rather high rate of spontaneous fission. ²⁴⁷Cm has the longest lifetime among isotopes of curium (1.56×10⁷ years), but is not formed in large quantities because of the strong fission induced by thermal neutrons.

Fourteen isotopes of berkelium were identified with mass numbers 238–252. Only ²⁴⁹Bk is available in large quantities; it has a relatively short half-life of 330 days and emits mostly soft β-particles, which are inconvenient for detection. Its alpha radiation is rather weak (1.45×10⁻³% with respect to β-radiation), but is sometimes used to detect this isotope. ²⁴⁷Bk is an alpha-emitter with a long half-life of 1,380 years, but it is hard to obtain in appreciable quantities; it is not formed upoeutron irradiation of plutonium because of the β-stability of isotopes of curium isotopes with mass number below 248.

Isotopes of californium with mass numbers 237–256 are formed iuclear reactors; californium-253 is a β-emitter and the rest are α-emitters. The isotopes with even mass numbers (²⁵⁰Cf, ²⁵²Cf and ²⁵⁴Cf) have a high rate of spontaneous fission, especially ²⁵⁴Cf of which 99.7% decays by spontaneous fission. Californium-249 has a relatively long half-life (352 years), weak spontaneous fission and strong γ-emission that facilitates its identification. ²⁴⁹Cf is not formed in large quantities in a nuclear reactor because of the slow β-decay of the parent isotope ²⁴⁹Bk and a large cross section of interaction with neutrons, but it can be accumulated in the isotopically pure form as the β-decay product of (pre-selected) ²⁴⁹Bk. Californium produced by reactor-irradiation of plutonium mostly consists of ²⁵⁰Cf and ²⁵²Cf, the latter being predominant for large neutron fluences, and its study is hindered by the strong neutron radiation.

Among the 16 known isotopes of einsteinium with mass numbers from 241 to 257 the most affordable is ²⁵³Es. It is an α-emitter with a half-life of 20.47 days, a relatively weak γ-emission and small spontaneous fission rate as compared with the isotopes of californium. Prolonged neutron irradiation also produces a long-lived isotope ²⁵⁴Es (t_½ = 275.5 days).

Nineteen isotopes of fermium are known with mass numbers of 242–260. ²⁵⁴Fm, ²⁵⁵Fm and²⁵⁶Fm are α-emitters with a short half-life (hours), which can be isolated in significant amounts.²⁵⁷Fm (t_½ = 100 days) can accumulate upon prolonged and strong irradiation. All these isotopes are characterized by high rates of spontaneous fission.

Among the 15 known isotopes of mendelevium (mass numbers from 245 to 260), the most studied is ²⁵⁶Md, which mainly decays through the electron capture (α-radiation is ≈10%) with the half-life of 77 minutes. Another alpha emitter, ²⁵⁸Md, has a half-life of 53 days. Both these isotopes are produced from rare einsteinium (²⁵³Es and ²⁵⁵Es respectively), that limits their so their availability.

Long-lived isotopes of nobelium and isotopes of lawrencium (and of heavier elements) have relatively small half-lives. For nobelium 11 isotopes are known with mass numbers 250–260 and 262. Chemical properties of nobelium and lawrencium were studied with ²⁵⁵No (t_½ = 3 min) and ²⁵⁶Lr (t_½ = 35 s). The longest-lived nobelium isotope ²⁵⁹No has a half-life of 1.5 hours.

Thorium and uranium are the most abundant actinides iature with the respective mass concentrations of 1.6×10⁻³% and 4×10⁻⁴%. Uranium mostly occurs in the Earth’s crust as a mixture of its oxides in the minerals uraninite, which is also called pitchblende because of its black color. There are several dozens of other uranium minerals such as carnotite (KUO₂VO₄·3H₂O) and autunite (Ca(UO₂)₂(PO₄)₂·nH₂O). The isotopic composition of natural uranium is ²³⁸U (relative abundance 99.2742%), ²³⁵U (0.7204%) and ²³⁴U (0.0054%); of these ²³⁸U has the largest half-life of 4.51×10⁹ years. The worldwide production of uranium in 2009 amounted to 50,572 tonnes, of which 27.3% was mined in Kazakhstan. Other important uranium mining countries are Canada (20.1%), Australia (15.7%), Namibia (9.1%), Russia (7.0%), andNiger (6.4%).

The most abundant thorium minerals are thorianite (ThO₂), thorite(ThSiO₄) and monazite, ((Th,Ca,Ce)PO₄). Most thorium minerals contain uranium and vice versa; and they all have significant fraction of lanthanides. Rich deposits of thorium minerals are located in the United States (440,000 tonnes), Australia and India (~300,000 tonnes each) and Canada (~100,000 tonnes).

The abundance of actinium in the Earth’s crust is only about 5×10⁻¹⁵%. Actinium is mostly present in uranium-containing, but also in other minerals, though in much smaller quantities. The content of actinium in most natural objects corresponds to the isotopic equilibrium of parent isotope ²³⁵U, and it is not affected by the weak Ac migration. Protactinium is more abundant (10⁻¹²%) in the Earth’s crust than actinium. It was discovered in the uranium ore in 1913 by Fajans and Göhring. As actinium, the distribution of protactinium follows that of ²³⁵U.

The half-life of the longest-lived isotope of neptunium, ²³⁷Np, is negligible compared to the age of the Earth. Thus neptunium is present in nature iegligible amounts produced as intermediate decay products of other isotopes. Traces of plutonium in uranium minerals were first found in 1942, and the more systematic results on ²³⁹Pu are summarized in the table (no other plutonium isotopes could be detected in those samples). The upper limit of abundance of the longest-living isotope of plutonium, ²⁴⁴Pu, is 3×10⁻²⁰%. Plutonium could not be detected in samples of lunar soil. Owing to its scarcity iature, most plutonium is produced synthetically. Negligible amounts of americium, curium, berkelium and californium are produced by neutron capture reactions and beta decay in very highly concentrated uranium-bearing deposits.

Owing to the low abundance of actinides, their extraction is a complex, multistep process. Fluorides of actinides are usually used because they are insoluble in water and can be easily separated with redox reactions. Fluorides are reduced with calcium,magnesium or barium.

Among the actinides, thorium and uranium are the easiest to isolate. Thorium is extracted mostly from monazite: thorium diphosphate (Th(PO₄)₂) is reacted with nitric acid, and the produced thorium nitrate treated with tributyl phosphate. Rare-earthimpurities are separated by increasing the pH in sulfate solution.

In another extraction method, monazite is decomposed with a 45% aqueous solution of sodium hydroxide at 140 °C. Mixed metal hydroxides are extracted first, filtered at 80 °C, washed with water and dissolved with concentrated hydrochloric acid. Next, the acidic solution is neutralized with hydroxides to pH = 5.8 that results in precipitation of thorium hydroxide (Th(OH)₄) contaminated with ~3% of rare-earth hydroxides; the rest of rare-earth hydroxides remains in solution. Thorium hydroxide is dissolved in an inorganic acid and then purified from the rare earth elements. An efficient method is the dissolution of thorium hydroxide iitric acid, because the resulting solution can be purified by extraction with organic solvents:

Th(OH)₄ + 4 HNO₃ → Th(NO₃)₄ + 4 H₂O

Metallic thorium is separated from the anhydrous oxide, chloride or fluoride by reacting it with calcium in an inert atmosphere:

ThO₂ + 2 Ca → 2 CaO + Th

Sometimes thorium is extracted by electrolysis of a fluoride in a mixture of sodium and potassium chloride at 700–800 °C in a graphite crucible. Highly pure thorium can be extracted from its iodide with the crystal bar process.

Uranium is extracted from its ores in various ways. In one method, the ore is burned and then reacted with nitric acid to convert uranium into a dissolved state. Treating the solution with a solution of tributyl phosphate (TBP) in kerosene transforms uranium into an organic form UO₂(NO₃)₂(TBP)₂. The insoluble impurities are filtered and the uranium is extracted by reaction with hydroxides as (NH₄)₂U₂O₇ or with hydrogen peroxide as UO₄·2H₂O.

When the uranium ore is rich in such minerals as dolomite, magnesite, etc., those minerals consume much acid. In this case, the carbonate method is used for uranium extraction. Its main component is an aqueous solution of sodium carbonate, which converts uranium into a complex [UO₂(CO₃)₃]^4–, which is stable in aqueous solutions at low concentrations of hydroxide ions. The advantages of the sodium carbonate method are that the chemicals have low corrosivity (compared to nitrates) and that most non-uranium metals precipitate from the solution. The disadvantage is that tetravalent uranium compounds precipitate as well. Therefore, the uranium ore is treated with sodium carbonate at elevated temperature and under oxygen pressure:

2 UO₂ + O₂ + 6 CO₂⁻³ → 2 [UO₂(CO₃)₃]^4–

This equation suggests that the best solvent for the uranium carbonate processing is a mixture of carbonate with bicarbonate. At high pH, this results in precipitation of diuranate, which is treated with hydrogen in the presence of nickel yielding an insoluble uranium tetracarbonate.

Another separation method uses polymeric resins as a polyelectrolyte. Ion exchange processes in the resins result in separation of uranium. Uranium from resins is washed with a solution of ammonium nitrate or nitric acid that yields uranyl nitrate, UO₂(NO₃)₂·6H₂O. When heated, it turns into UO₃, which is converted to UO₂ with hydrogen:

UO₃ + H₂ → UO₂ + H₂O

Reacting uranium dioxide with fluoric acid changes it to uranium tetrafluoride, which yields uranium metal upon reaction with magnesium metal:

4 HF + UO₂ → UF₄ + 2 H₂O

To extract plutonium, neutron-irradiated uranium is dissolved iitric acid, and a reducing agent (FeSO₄, or H₂O₂) is added to the resulting solution. This addition changes the oxidation state of plutonium from +6 to +4, while uranium remains in the form of uranyl nitrate (UO₂(NO₃)₂). The solution is treated with a reducing agent and neutralized withammonium carbonate to pH = 8 that results in precipitation of Pu⁴⁺ compounds.

In another method, Pu⁴⁺and UO₂+ 2 are first extracted with tributyl phosphate, then reacted with hydrazine washing out the recovered plutonium.

The major difficulty in separation of actinium is the similarity of its properties with those of lanthanum. Thus actinium is either synthesized iuclear reactions from isotopes of radium or separated using ion-exchange procedures.

Applications

While actinides have some established daily-life applications, such as in smoke detectors (americium) and gas mantles (thorium), they are mostly used iuclear weapons and use as a fuel iuclear reactors. The last two areas exploit the property of actinides to release enormous energy iuclear reactions, which under certain conditions may become self-sustaining chain reaction.

The most important isotope for nuclear power applications is uranium-235. It is used in the thermal reactor, and its concentration iatural uranium does not exceed 0.72%. This isotope strongly absorbs thermal neutrons releasing much energy. One fission act of 1 gram of 235U converts into about 1 MW·day. Of importance, is that 235U emits more neutrons than it absorbs; upon reaching the critical mass, 235U enters into a self-sustaining chain reaction. Typically, uranium nucleus is divided into two fragments with the release of 2–3 neutrons, for example:

Other promising actinide isotopes for nuclear power are thorium-232 and its product from the thorium fuel cycle, uranium-233.

The core of any nuclear reactor contains a set of hollow metal rods, usually made of zirconium alloys, filled with nuclear fuel cells – mostly oxide, carbide, nitride or monosulfide of uranium, plutonium or thorium, or their mixture (the so-called MOX fuel). The most common fuel is oxide of uranium-235. Nuclear reactor scheme

Fast neutrons are slowed by moderators, which contain water, carbon, deuterium, or beryllium, as thermal neutrons to increase the efficiency of their interaction with uranium-235. The rate of nuclear reaction is controlled by introducing additional rods made of boron or cadmium or a liquid absorbent, usually boric acid. Reactors for plutonium production are called breeder reactor or breeders; they have a different design and use fast neutrons.

Emission of neutrons during the fission of uranium is important not only for maintaining the nuclear chain reaction, but also for the synthesis of the heavier actinides. Uranium-239 converts via β-decay into plutonium-239, which, like uranium-235, is capable of spontaneous fission. The world’s first nuclear reactors were

About half of the produced thorium is used as the light-emitting material of gas mantles. Thorium is also added into multicomponent alloys of magnesium and zinc. So the Mg-Th alloys are light and strong, but also have high melting point and ductility and thus are widely used in the aviation industry and in the production of missiles. Thorium also has good electron emission properties, with long lifetime and low potential barrier for the emission. The relative content of thorium and uranium isotopes is widely used to estimate the age of various objects, including stars (see radiometric dating).

The major application of plutonium has been iuclear weapons, where the isotope plutonium-239 was a key component due to its ease of fission and availability. Plutonium-based designs allow reducing the critical mass to about a third of that for uranium-235. The “Fat Man”-type plutonium bombs produced during the Manhattan Project used explosive compression of plutonium to obtain significantly higher densities thaormal, combined with a central neutron source to begin the reaction and increase efficiency. Thus only 6.2 kg of plutonium was needed for an explosive yield equivalent to 20 kilotons of TNT. (See also Nuclear weapon design.) Hypothetically, as little as 4 kg of plutonium—and maybe even less—could be used to make a single atomic bomb using very sophisticated assembly designs.

Plutonium-238 is potentially more efficient isotope for nuclear reactors, since it has smaller critical mass than uranium-235, but it continues to release much thermal energy (0.56 W/g) by decay even when the fission chain reaction is stopped by control rods. Its application is limited by the high price (about 1000 USD/g). This isotope has been used in thermopiles and water distillation systems of some space satellites and stations. So Galileo and Apollo spacecraft (e.g. Apollo 14) had heaters powered by kilogram quantities of plutonium-238 oxide; this heat is also transformed into electricity with thermopiles. The decay of plutonium-238 produces relatively harmless alpha particles and is not accompanied by gamma-irradiation. Therefore this isotope (~160 mg) is used as the energy source in heart pacemakers where it lasts about 5 times longer than conventional batteries.

Actinium-227 is used as a neutron source. Its high specific energy (14.5 W/g) and the possibility of obtaining significant quantities of thermally stable compounds are attractive for use in long-lasting thermoelectric generators for remote use. 228Ac is used as an indicator of radioactivity in chemical research, as it emits high-energy electrons (2.18 MeV) that can be easily detected. 228Ac-228Ra mixtures are widely used as an intense gamma-source in industry and medicine.

Development of self-glowing actinide-doped materials with durable crystalline matrices is a new area of actinide utilization as the addition of alpha-emitting radionuclides to some glasses and crystals may confer luminescence.

Toxicity

Radioactive elements: the most stable isotope is very long-lived, with a half-life of over four million years.

Radioactive elements: the most stable isotope has half-life between 800 and 34.000 years.

Radioactive elements: the most stable isotope has half-life between one day and 103 years.

Highly radioactive elements: the most stable isotope has half-life between several minutes and one day.

Extremely radioactive elements: the most stable isotope has half-life less than several minutes. Very little is known about these elements due to their extreme instability and radioactivity.

Radioactive substances can harm human health via (i) local skin contamination, (ii) internal exposure due to ingestion of radioactive isotopes, and (iii) external overexposure by β-activity and γ-radiation. Together with radium and transuranium elements, actinium is one of the most dangerous radioactive poisons with high specific α-activity. The most important feature of actinium is its ability to accumulate and remain in the surface layer of skeletons. At the initial stage of poisoning, actinium accumulates in the liver. Another danger of actinium is that it undergoes radioactive decay faster than being excreted. Adsorption from the digestive tract is much smaller (~0.05%) for actinium than radium.

Protactinium in the body tends to accumulate in the kidneys and bones. The maximum safe dose of protactinium in the human body is 0.03 µCi that corresponds to 0.5 micrograms of 231Pa. This isotope, which might be present in the air as aerosol, is 2.5×108 times more toxic than hydrocyanic acid.

Plutonium, when entering the body through air, food or blood (e.g. a wound), mostly settles in the lungs, liver and bones with only about 10% going to other organs, and remains there for decades. The long residence time of plutonium in the body is partly explained by its poor solubility in water. Some isotopes of plutonium emit ionizing α-radiation, which damages the surrounding cells. The median lethal dose (LD50) for 30 days in dogs after intravenous injection of plutonium is 0.32 milligram per kg of body mass, and thus the lethal dose for humans is approximately 22 mg for a person weighing 70 kg; the amount for respiratory exposure should be approximately four times greater. Another estimate assumes that plutonium is 50 times less toxic than radium, and thus permissible content of plutonium in the body should be 5 µg or 0.3 µCi. Such amount is nearly invisible in under microscope. After trials on animals, this maximum permissible dose was reduced to 0.65 µg or 0.04 µCi. Studies on animals also revealed that the most dangerous plutonium exposure route is through inhalation, after which 5–25% of inhaled substances is retained in the body. Depending on the particle size and solubility of the plutonium compounds, plutonium is localized either in the lungs or in the lymphatic system, or is absorbed in the blood and then transported to the liver and bones. Contamination via food is the least likely way. In this case, only about 0.05% of soluble 0.01% insoluble compounds of plutonium absorbs into blood, and the rest is excreted. Exposure of damaged skin to plutonium would retaiearly 100% of it.

Using actinides iuclear fuel, sealed radioactive sources or advanced materials such as self-glowing crystals has many potential benefits. However, a serious concern is the extremely high radiotoxicity of actinides and their migration in the environment. Use of chemically unstable forms of actinides in MOX and sealed radioactive sources is not appropriate by modern safety standards. There is a challenge to develop stable and durable actinide-bearing materials, which provide safe storage, use and final disposal. A key need is application of actinide solid solutions in durable crystalline host phases.

d-elements of VIIВ group

Side subgroup VII group includes the following elements: Mn – Mn, Tc – Technetium and Re – rhenium. The main adverse representative subgroup VII group is manganese.

Manganese, technetium and rhenium – full electronic analogs of the valence electron configuration (n-1) d⁵s². They unite in a subgroup of manganese.

Atomic radii of rhenium and tehnetsiyu close because their properties are similar.

For typical manganese oxidation state +2, +4, +7.

Simple matter of manganese and its analogues are a silvery–white metal color.

Growth and melting temperature of boiling in the range Mn-Tc-Re explained by the increased covalent bond formed by (n-1) d-orbitals.

 

MANGANESE

Manganese is a chemical element, designated by the symbol Mn. It has the atomic number 25. It is found as a free element iature (often in combination with iron), and in many minerals. As a free element, manganese is a metal with important industrial metal alloy uses, particularly in stainless steels.

Manganese ions have various colors, depending on their oxidation state, and are used industrially as pigments. The permanganates of sodium, potassium and barium are powerful oxidisers. Manganese dioxide is used as the cathode (electron acceptor) material in standard and alkaline disposable dry cells and batteries.

Manganese(II) ions function as cofactors for a number of enzymes in higher organisms, where they are essential in detoxification of superoxide free radicals. The element is a required trace mineral for all known living organisms. In larger amounts, and apparently with far greater activity by inhalation, manganese can cause a poisoning syndrome in mammals, with neurological damage which is sometimes irreversible.

Notable chemical characteristics

Chemical properties

Manganese is a gray-white metal, resembling iron. It is a hard metal and is very brittle, fusible with difficulty, but easily oxidized. Manganese metal and its common ions are paramagnetic. While manganese metal does not form a permanent magnet, it does exhibit strong magnetic properties in the presence of an external magnetic field.

 

The +3 oxidation state is known in compounds like manganese(III) acetate, but these are quite powerful oxidizing agents and also prone to disproportionation in solution to manganese(II) and manganese(IV). Solid compounds of manganese(III) are characterized by their preference for distorted octahedral coordination due to the Jahn-Teller effectand its strong purple-red color. The oxidation state 5+ can be obtained if manganese dioxide is dissolved in molten sodium nitrite. Manganate (VI) salts can also be produced by dissolving Mn compounds, such as manganese dioxide, in molten alkali while exposed to air. Permanganate (+7 oxidation state) compounds are purple, and can give glass a violet color.Potassium permanganate, sodium permanganate and barium permanganate are all potent oxidizers. Potassium permanganate, also called Condy’s crystals, is a commonly used laboratory reagent because of its oxidizing properties and finds use as a topical medicine (for example, in the treatment of fish diseases). Solutions of potassium permanganate were among the first stains and fixatives to be used in the preparation of biological cells and tissues for electron microscopy.

The most common oxidation states of manganese are +2, +3, +4, +6 and +7, though oxidation states from +1 to +7 are observed. Mn²⁺ often competes with Mg²⁺ in biological systems, and manganese compounds where manganese is in oxidation state +7 are powerful oxidizing agents.

Mn²⁺ often competes with Mg²⁺ in biological systems. Manganese compounds where manganese is in oxidation state +7, which are restricted to the unstable oxide Mn₂O₇ and compounds of the intensely purple permanganate anion MnO₄⁻, are powerful oxidizing agents. Compounds with oxidation states +5 (blue) and +6 (green) are strong oxidizing agents and are vulnerable to disproportionation.

The most stable oxidation state for manganese is +2, which has a pale pink color, and many manganese(II) compounds are known, such as manganese(II) sulfate (MnSO₄) and manganese(II) chloride (MnCl₂). This oxidation state is also seen in the mineral rhodochrosite (manganese(II) carbonate). The +2 oxidation of Mn results from removal of the two 4s electrons, leaving a “high spin” ion in which all five of the 3d orbitals contain a single electron. Absorption of visible light by this ion is accomplished only by a spin-forbidden transition in which one of the d electrons must pair with another, to give the atom a change in spin of two units. The unlikeliness of such a transition is seen in the uniformly pale and almost colorless nature of Mn(II) compounds relative to other oxidation states of manganese.

Mn + Cl₂ = MnCl₂ ;

Mn + S = MnS ;

Mn + 2P = Mn₃P₂ ;

3Mn + N₂ = Mn₂N₂ ;

2Mn + Si = Mn₂Si .

Physical properties

Manganese is a silvery-gray metal that resembles iron. It is hard and very brittle, difficult to fuse, but easy to oxidize. Manganese metal and its common ions areparamagnetic. Manganese tarnishes slowly in air and “rusts” like iron, in water containing dissolved oxygen.

Isotopes

Naturally occurring manganese is composed of one stable isotope, ⁵⁵Mn. Eighteenradioisotopes have been characterized, with the most stable being ⁵³Mn with a half-lifeof 3.7 million years, ⁵⁴Mn with a half-life of 312.3 days, and ⁵²Mn with a half-life of 5.591 days. All of the remaining radioactive isotopes have half-lives that are less than three hours and the majority of these have half-lives that are less than one minute. This element also has three meta states. Manganese is part of the iron group of elements, which are thought to be synthesized in large stars shortly before the supernova explosion. ⁵³Mn decays to ⁵³Crwith a half-life of 3.7 million years. Because of its relatively short half-life, ⁵³Mn occurs only in tiny amounts due to the action ofcosmic rays on iron in rocks. Manganese isotopic contents are typically combined with chromium isotopic contents and have found application in isotope geology and radiometric dating. Mn–Cr isotopic ratios reinforce the evidence from ²⁶Al and ¹⁰⁷Pdfor the early history of the solar system. Variations in ⁵³Cr/⁵²Cr and Mn/Cr ratios from several meteorites indicate an initial⁵³Mn/⁵⁵Mn ratio that suggests Mn–Cr isotopic composition must result from in situ decay of ⁵³Mn in differentiated planetary bodies. Hence ⁵³Mn provides additional evidence for nucleosynthetic processes immediately before coalescence of the solar system. The isotopes of manganese range in atomic weight from 46 u (⁴⁶Mn) to 65 u (⁶⁵Mn). The primary decay modebefore the most abundant stable isotope, ⁵⁵Mn, is electron capture and the primary mode after is beta decay.

The origin of the name manganese is complex. In ancient times, two black minerals from Magnesia in what is now modern Greece, were both called magnes from their place of origin, but were thought to differ in gender. The male magnes attracted iron, and was the iron ore we now know as lodestone ormagnetite, and which probably gave us the term magnet. The female magnes ore did not attract iron, but was used to decolorize glass. This feminine magnes was later called magnesia, known now in modern times as pyrolusite or manganese dioxide. Neither this mineral nor manganese itself is magnetic. In the 16th century, manganese dioxide was called manganesum (note the two n’s instead of one) by glassmakers, possibly as a corruption and concatenation of two words, since alchemists and glassmakers eventually had to differentiate a magnesia negra (the black ore) from magnesia alba (a white ore, also from Magnesia, also useful in glassmaking). Michele Mercati called magnesia negra manganesa, and finally the metal isolated from it became known as manganese (German: Mangan). The namemagnesia eventually was then used to refer only to the white magnesia alba (magnesium oxide), which provided the name magnesium for that free element, when it was eventually isolated, much later.

Several oxides of manganese, for example manganese dioxide, are abundant iature, and owing to their color, these oxides have been used as since the Stone Age. The cave paintings in Gargas contain manganese as pigments and these cave paintings are 30,000 to 24,000 years old.

Manganese compounds were used by Egyptian and Roman glassmakers, to either remove color from glass or add color to it. The use as “glassmakers soap” continued through the Middle Ages until modern times and is evident in 14th-century glass fromVenice.

Because of the use in glassmaking, manganese dioxide was available to alchemists, the first chemists, and was used for experiments. Ignatius Gottfried Kaim (1770) and Johann Glauber (17th century) discovered that manganese dioxide could be converted topermanganate, a useful laboratory reagent. By the mid-18th century, the Swedish chemist Carl Wilhelm Scheele used manganese dioxide to produce chlorine. First,hydrochloric acid, or a mixture of dilute sulfuric acid and sodium chloride was made to react with manganese dioxide, later hydrochloric acid from the Leblanc process was used and the manganese dioxide was recycled by the Weldon process. The production of chlorine and hypochlorite containing bleaching agents was a large consumer of manganese ores.

Scheele and other chemists were aware that manganese dioxide contained a new element, but they were not able to isolate it. Johan Gottlieb Gahn was the first to isolate an impure sample of manganese metal in 1774, by reducing the dioxide with carbon.

The manganese content of some iron ores used in Greece led to the speculations that the steel produced from that ore contains inadvertent amounts of manganese, making the Spartan steel exceptionally hard. Around the beginning of the 19th century, manganese was used in steelmaking and several patents were granted. In 1816, it was noted that adding manganese to iron made it harder, without making it any more brittle. In 1837, British academic James Couper noted an association between heavy exposures to manganese in mines with a form of Parkinson’s disease. In 1912, manganese phosphating electrochemical conversion coatings for protecting firearms against rust and corrosion were patented in the United States, and have seen widespread use ever since.

The invention of the Leclanché cell in 1866 and the subsequent improvement of the batteries containing manganese dioxide as cathodic depolarizer increased the demand of manganese dioxide. Until the introduction of the nickel-cadmium battery and lithium-containing batteries, most batteries contained manganese. The zinc-carbon battery and thealkaline battery normally use industrially produced manganese dioxide, because natural occurring manganese dioxide contains impurities. In the 20th century, manganese dioxide has seen wide commercial use as the chief cathodic material for commercial disposable dry cells and dry batteries of both the standard (zinc-carbon) and alkaline types.

Manganese makes up about 1000 ppm (0.1%) of the Earth’s crust, making it the 12th most abundant element there. Soil contains 7–9000 ppm of manganese with an average of 440 ppm. Seawater has only 10 ppm manganese and the atmosphere contains 0.01 µg/m³. Manganese occurs principally as pyrolusite (MnO₂), braunite, (Mn²⁺Mn³⁺₆)(SiO₁₂), тpsilomelane (Ba,H₂O)₂Mn₅O₁₀, and to a lesser extent as rhodochrosite (MnCO₃).

The most important manganese ore is pyrolusite (MnO₂). Other economically important manganese ores usually show a close spatial relation to the iron ores. Land-based resources are large but irregularly distributed. About 80% of the known world manganese resources are found in South Africa; other important manganese deposits are in Ukraine, Australia, India, China, Gabon and Brazil. In 1978, 500 billion tons ofmanganese nodules were estimated to exist on the ocean floor. Attempts to find economically viable methods of harvesting manganese nodules were abandoned in the 1970s.

Manganese is mined in South Africa, Australia, China, Brazil, Gabon, Ukraine, India and Ghana and Kazakhstan. US Import Sources (1998–2001): Manganese ore: Gabon, 70%; South Africa, 10%; Australia, 9%; Mexico, 5%; and other, 6%. Ferromanganese: South Africa, 47%; France, 22%; Mexico, 8%; Australia, 8%; and other, 15%. Manganese contained in all manganese imports: South Africa, 31%; Gabon, 21%; Australia, 13%; Mexico, 8%; and other, 27%.

For the production of ferromanganese, the manganese ore is mixed with iron ore and carbon, and then reduced either in a blast furnace or in an electric arc furnace. The resulting ferromanganese has a manganese content of 30 to 80%. Pure manganese used for the production of iron-free alloys is produced by leaching manganese ore withsulfuric acid and a subsequent electrowinning process.

In 1972 the CIA’s Project Azorian, through billionaire Howard Hughes, commissioned the ship Hughes Glomar Explorer with the cover story of harvesting manganese nodules from the sea floor. That triggered a rush of activity to attempt to collect manganese nodules, which was not actually practical. The real mission of Hughes Glomar Explorer was to raise a sunken Soviet submarine, the K-129, with the goal of retrieving Soviet code books.

Industrially important compounds

Methylcyclopentadienyl manganese tricarbonyl is used as an additive in unleaded gasoline to boost octane rating and reduce engine knocking. The manganese in this unusual organometallic compound is in the +1 oxidation state.

The most stable oxidation state for manganese is +2, which has a pink to red color, and many manganese(II) compounds are known, such as manganese(II) sulfate (MnSO₄) and manganese(II) chloride (MnCl₂). This oxidation state is also seen in the mineral rhodochrosite, (manganese(II) carbonate). The +2 oxidation state is the state used in living organisms for essential functions; all of the other states are much more toxic.

Compounds with Mn(II) have reduction properties

6Mn(OH)₂ + O₂ = 2Mn₂MnO₄ + 6H₂O

3MnSO₄ + 2KClO₃ + 12KOH = 3K₂MnO₄ + 2KCl + 3K₂SO₄ + 6H₂O.

2MnSO₄ + 5PbO₂ + 6HNO₃ = 2HMnO₄ + 3Pb(NO₃)₂ + 2PbSO₄¯ + 2H₂O

4MnO₂ = 2Mn₂O₃ + O₂­

The +3 oxidation state is known, in compounds such as manganese(III) acetate, but these are quite powerful oxidizing agents.

Manganese(IV) oxide (manganese dioxide, MnO₂) is used as a reagent in organic chemistry for the oxidation of benzylic alcohols (i.e. adjacent to an aromatic ring). Manganese dioxide has been used since antiquity to oxidatively neutralize the greenish tinge in glass caused by trace amounts of iron contamination. MnO₂ is also used in the manufacture of oxygen and chlorine, and in drying black paints. In some preparations it is a brown pigment that can be used to make paint and is a constituent of natural umber.

2MnO₂ + 2H₂SO₄ = 2MnSO₄ + O₂­ + 2H₂O

4MnO₂ = 2Mn₂O₃ + O₂­

MnO₂ + CaO = CaMnO₃

MnO₂ + 4HCl = MnCl₂ + Cl + 2H₂O

3MnO₂ + KClO₃ + 6KOH ® 3K₂MnO₄ + KCl + 3H₂O

Manganese(IV) oxide was used in the original type of dry cell battery as an electron acceptor from zinc, and is the blackish material found when opening carbon-zinc type flashlight cells. The same material also functions iewer alkaline batteries (usually battery cells), which use the same basic reaction, but a different electrolyte mixture.

MnO₂ + 4H⁺ + 2e = Mn²⁺ + 2H₂O, E = 1.23V

Manganese phosphating is used as a treatment for rust and corrosion prevention on steel.

Permanganate (+7 oxidation state) manganese compounds are purple, and can color glass an amethyst color. Potassium permanganate, sodium permanganate and barium permanganate are all potent oxidizers. Potassium permanganate, also called Condy’s crystals, is a commonly used laboratory reagent because of its oxidizing properties and finds use as a topical medicine (for example, in the treatment of fish diseases). Solutions of potassium permanganate were among the first stains and fixatives to be used in the preparation of biological cells and tissues for electron microscopy.

Substitutes: Manganese has no satisfactory substitute in its major applications, which are related to metallurgical alloy use. In minor applications, (e.g., manganese phosphating), zinc and sometimes vanadium are viable substitutes. In disposable battery manufacture, standard and alkaline cells using manganese will probably eventually be mostly replaced with lithium battery technology.

The overall level and nature of manganese use in the United States is expected to remain about the same in the near term. No practical technologies exist for replacing manganese with other materials or for using domestic deposits or other accumulations to reduce the complete dependence of the United States on other countries for manganese ore.

Metal alloys

 Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties. Steelmaking, including its ironmaking component, has accounted for most manganese demand, presently in the range of 85% to 90% of the total demand. Among a variety of other uses, manganese is a key component of low-cost stainless steel formulations and certain widely used aluminium alloys.

Occurrence

Manganese ore

Manganese ore

Psilomelane (manganese ore)

Spiegeleisen is an iron alloy with a manganese content of approximately 15%

Manganese oxide dendrites on limestone from Solnhofen, Germany – a kind of pseudofossil. Scale is in mm

Mineral rhodochrosite (manganese(II) carbonate) in which the deep red color is due to impurities, not manganese

Manganese ore Psilomelane (manganese ore) Spiegeleisen is an iron alloy with a manganese content of approximately 15% Manganese oxide dendrites on limestone from Solnhofen, Germany – a kind of pseudofossil. Scale is in mm Mineral rhodochrosite (manganese(II) carbonate) in which the deep red color is due to impurities, not manganese.

Manganese occurs principally as pyrolusite (MnO₂), braunite, psilomelane, and to a lesser extent as rhodochrosite (MnCO₃). Land-based resources are large but irregularly distributed. Over 80% of the known world manganese resources are found in South Africa and Ukraine. Other important manganese deposits are in China, Australia, Brazil, Gabon, Ghana, India, and Mexico.

Psilomelane (manganese ore)

Manganese oxide dendrites on a limestone bedding plane from Solingen, Germany.

Isotopes

Naturally occurring manganese is composed of 1 stable isotope; 55Mn. 18 radioisotopes have been characterized with the most stable being 53Mn with a half-life of 3.7 million years, 54Mn with a half-life of 312.3 days, and 52Mn with a half-life of 5.591 days. All of the remaining radioactive isotopes have half lives that are less than 3 hours and the majority of these have half lives that are less than 1 minute. This element also has 3 meta states.

Manganese is part of the iron group of elements which are thought to be synthesized in large stars shortly before supernova explosion. 53Mn decays to 53Cr with a half-life of 3.7 million years. Because of its relatively short half-life, 53Mn is an extinct radionuclide. Manganese isotopic contents are typically combined with chromium isotopic contents and have found application in isotope geology and radiometric dating. Mn-Cr isotopic ratios reinforce the evidence from 26Al and 107Pd for the early history of the solar system. Variations in 53Cr/52Cr and Mn/Cr ratios from several meteorites indicate an initial 53Mn/55Mn ratio that suggests Mn-Cr isotopic systematics must result from in-situ decay of 53Mn in differentiated planetary bodies. Hence 53Mn provides additional evidence for nucleosynthetic processes immediately before coalescence of the solar system.

The isotopes of manganese range in atomic weight from 46 u (46Mn) to 65 u (65Mn). The primary decay mode before the most abundant stable isotope, 55Mn, is electron capture and the primary mode after is beta decay.

Precautions

Manganese compounds are less toxic than those of other widespread metals such as nickel and copper. Exposure to manganese dusts and fumes should not exceed the ceiling value of 5 mg/m³ even for short periods because of its toxicity level. Manganese poses a particular risk for children due to its propensity to bind to CH-7 receptors. Manganese poisoning has been linked to impaired motor skills and cognitive disorders.

Acidic permanganate solutions will oxidize any organic material they come into contact with. The oxidation process can generate enough heat to ignite some organic substances.

Potassium manganate K₂MnO₄

Potassium manganate is the chemical compound with the formula K₂MnO₄. This green salt is an intermediate in the industrial synthesis of potassium permanganate, a common chemical. Occasionally, potassium manganate and potassium permanganate are confused, but they are different compounds with distinctly different properties.

Structure

K₂MnO₄ is a salt, consisting of K⁺ cations and MnO₄^2- anions.

Synthesis

The industrial route entails treatment of MnO₂ with air:

MnO₂ + 2 KOH + ½ O₂ → K₂MnO₄ + H₂O

The transformation gives a green-colored melt. In fact, one can test an unknown substance for the presence of manganese by heating the sample in strong KOH in air. The production of a green coloration indicates the presence of Mn. This green color results from an intense absorption at 610 nm.

In laboratory, K₂MnO₄ can be synthesized by heating a solution of KMnO₄ in concentrated KOH solution followed by cooling to give green crystals:

4 KMnO₄ + 4 KOH → 4 K₂MnO₄ + O₂ + 2 H₂O

This reaction illustrates the relatively rare role of hydroxide as a reducing agent. Solutions of K₂MnO₄ are generated by allowing a solution of KMnO₄ in 5-10 M KOH to stir for a day at room temperature followed by removal of MnO₂, which is insoluble. The concentration of K₂MnO₄ in such solutions can be checked by measuring their absorbance at 610 nm.

The one-electron reduction of permanganate to manganate can also be effected using iodide as the reducing agent:

KMnO₄ + KI → K₂MnO₄ + ½ I₂

The conversion is signaled by the color change from purple, characteristic of permanganate, to the green color of manganate. This reaction also illustates the fact that manganate(VII) can serve as an electron acceptor in addition to its usual role as an oxygen-transfer reagent. Barium manganate, BaMnO₄, is generated by the reduction of KMnO₄ with iodide in the presence of barium chloride. Just like BaSO₄, BaMnO₄ exhibits low solubility in virtually all solvents.

3K₂MnO₄ + 2H₂O = 2KMnO₄ + MnO₂ + 4KOH

2K₂MnO₄ + Cl₂ = 2KMnO₄ + 2KCl.

Potassium permanganate KMnO₄

Potassium permanganate is an inorganic chemical compound with the formula KMnO₄. It is a salt consisting of K⁺ and MnO₄⁻ ions. Formerly known as permanganate of potash or Condy’s crystals, it is a strong oxidizing agent. It dissolves in water to give intense purple solutions, the evaporation of which gives prismatic purplish-black glistening crystals.

Potassium permanganate decomposes when exposed to light:

2 KMnO₄(s) → K₂MnO₄(s) + MnO₂(s)+O₂(g)

2KMnO₄ + H₂SO₄ = Mn₂O₇ + K₂SO₄ + H₂O

Structure and preparation

It forms orthorhombic crystals with constants: a = 910.5 pm, b = 572.0 pm, c = 742.5 pm. The Mn-O bond distances are 162.9 pm.

Potassium permanganate is produced industrially from manganese dioxide, which also occurs as the mineral pyrolusite. The MnO₂ is fused with potassium hydroxide and heated in air or with potassium nitrate (a source of oxygen). This process gives potassium manganate, which upon electrolytic oxidation in alkaline solution gives potassium permanganate.

2 MnO₂ + 4 KOH + O₂ → 2 K₂MnO₄ + 2 H₂O

2 MnO₄^2– + Cl₂ → 2 MnO₄^– + 2 Cl^–

Permanganates can also be generated by treating a solution of Mn²⁺ ions with strong oxidants such as lead dioxide (PbO₂), or sodium bismuthate (NaBiO₃). These reactions use the vivid violet colour of permanganate as a test for the presence of manganese.

Organic chemistry

Dilute solutions of KMnO₄ convert alkenes into diols (glycols). This behaviour is also used as a qualitative test for the presence of double or triple bonds in a molecule, since the reaction decolorizes the initially purple permanganate solution and generates a brown precipitate (MnO₂). It is sometimes referred to as Baeyer’s reagent. However, bromine serves better in measuring unsaturation (double or triple bonds) quantitatively, since KMnO₄, being a very strong oxidizing agent, can react with a variety of groups.

Under acidic conditions, the alkene double bond is cleaved to give the appropriate carboxylic acid:

CH₃(CH₂)₁₇CH=CH₂ + 2 KMnO₄ + 3 H₂SO₄ → CH₃(CH₂)₁₇COOH + CO₂ + 4 H₂O + K₂SO₄ + 2 MnSO₄

Potassium permanganate oxidizes aldehydes to carboxylic acids, such as the conversion of n-heptanal to heptanoic acid:

5 C₆H₁₃CHO + 2 KMnO₄ + 3 H₂SO₄ → 5 C₆H₁₃COOH + 3 H₂O + K₂SO₄ + 2 MnSO₄

Even an alkyl group (with a benzylic hydrogen) on an aromatic ring is oxidized, e.g. toluene to benzoic acid.

5 C₆H₅CH₃ + 6 KMnO₄ + 9 H₂SO₄ → 5 C₆H₅COOH + 14 H₂O + 3 K₂SO₄ + 6 MnSO₄

Glycols and polyols are highly reactive toward KMnO₄. For example, addition of potassium permanganate to an aqueous solution of sugar and sodium hydroxide produces the “chemical chameleon” reaction, which involves dramatic color changes associated with the various oxidation states of manganese. A related vigorous reaction is exploited as a fire starter in survival kits. For example, a mixture of potassium permanganate and glycerol or pulverized glucose ignites readily. Its sterilizing properties are another reason for inclusion of KMnO₄ in a survival kit.

By itself, potassium permanganate does not dissolve in many organic solvents. If an organic solution of permanganate is desired, “purple benzene” may be prepared, either by treating a two phase mixture of aqueous potassium permanganate and benzene with a quaternary ammonium salt, or by sequestering the potassium cation with a crown ether.

Reaction with acids

Concentrated sulfuric acid reacts with KMnO₄ to give Mn₂O₇, which can be explosive. Its reaction with concentratedhydrochloric acid gives chlorine. The Mn-containing products from redox reactions depend on the pH. Acidic solutions of permanganate are reduced to the faintly pink manganese(II) ion (Mn²⁺) and water. Ieutral solution, permanganate is only reduced by three electrons to give MnO₂, wherein Mn is in a +4 oxidation state. This is the material that stains one’s skin when handling KMnO₄. KMnO₄ spontaneously reduces in an alkaline solution to green K₂MnO₄, wherein manganese is in the +6 oxidation state.

A curious reaction occurs upon addition of concentrated sulfuric acid to potassium permanganate. Although no reaction may be apparent, the vapor over the mixture will ignite paper impregnated with alcohol. Potassium permanganate and sulfuric acid react to produce some ozone, which has a high oxidising power and rapidly oxidises the alcohol, causing it to combust. As the reaction also produces explosive Mn₂O₇, this should only be attempted with great care.

Photodecomposition

Potassium permanganate decomposes when exposed to light:

2 KMnO₄ → K₂MnO₄ + MnO₂(s) + O₂

In 1659, Johann Rudolf Glauber fused a mixture of the mineral pyrolusite (Manganese Dioxide: MnO₂) and potassium carbonateto obtain a material that, when dissolved in water, gave a green solution (potassium manganate) which slowly shifted to violet and then finally red. This report represents the first description of the production of potassium permanganate. Just under two hundred years later London chemist Henry Bollmann Condy had an interest in disinfectants, and marketed several products including ozonised water. He found that fusing pyrolusite with NaOH and dissolving it in water produced a solution with disinfectant properties. He patented this solution, and marketed it as Condy’s Fluid. Although effective, the solution was not very stable. This was overcome by using KOH rather than NaOH. This was more stable, and had the advantage of easy conversion to the equally effective potassium permanganate crystals. This crystalline material was known as Condy’s crystals or Condy’s powder. Potassium permanganate was comparatively easy to manufacture so Condy was subsequently forced to spend considerable time in litigation in order to stop competitors from marketing products similar to Condy’s Fluid or Condy’s Crystals.

Early photographers used it as a component of flash powder. It is now replaced with other oxidizers, due to the instability of permanganate mixtures. Aqueous solutions of KMnO₄ have been used together with T-Stoff (i.e. 80% hydrogen peroxide) as propellant for the rocket plane Messerschmitt Me 163. In this application, it was known as Z-Stoff. This combination of propellants is sometimes still used in torpedoes.

Almost all applications of potassium permanganate exploit its oxidizing properties. As a strong oxidant that does not generate toxic byproducts, KMnO₄ has many niche uses.

Potassium permanganate is one of the principal chemicals used in the film and television industries to “age” props and set dressings. Its oxidising effects create “hundred year old” or “ancient” looks on hessian cloth, ropes, timber and glass.

Water treatment and disinfection

As an oxidant, potassium permanganate can act as an antiseptic. For example, dilute solutions are used to treat canker sores (ulcers), disinfectant for the hands and treatment for mild pompholyx, dermatitis, and fungal infections of the hands or feet.  Potassium permanganate is used extensively in the water treatment industry. It is used as a regeneration chemical to remove iron and hydrogen sulfide (rotten egg smell) from well water via a “Manganese Greensand” Filter. “Pot-Perm” is also obtainable at pool supply stores, is used additionally to treat waste water. Historically it was used to disinfect drinking water. It currently finds application in the control of nuisance organisms such as Zebra mussels in fresh water collection and treatment systems.

Organic synthesis

Aside from its use in water treatment, the other major application of KMnO₄ is as a reagent for the synthesis of organic compounds. Significant amounts are required for the synthesis of ascorbic acid, chloramphenicol, saccharin, isonicotinic acid, and pyrazinoic acid.

Analytical use

Potassium permanganate can be used to quantitatively determine the total oxidisable organic material in an aqueous sample. The value determined is known as the permanganate value. In analytical chemistry, a standardized aqueous solution of KMnO₄ is sometimes used as an oxidizing titrant forredox titrations (permanganometry). In a related way, it is used as a reagent to determine the Kappa number of wood pulp. For the standardization of KMnO₄ solutions, reduction by oxalic acid is often used.

Aqueous, acidic solutions of KMnO₄ are used to collect gaseous mercury in flue gas during stationary source emissions testing.

In histology, potassium permanganate is used to bleach melanin which obscures tissue detail. Pre-treated with potassium permanganate, to obliterate Congo red reactivity, was thought to be definitive for AA amyloidosis; this is now generally considered to be unreliable.

Fruit preservation

Ethylene absorbents extend storage time of bananas even at high temperatures. This effect can be exploited by packing bananas in polyethylenewith potassium permanganate as an ethylene absorbent, that doubles a bananas lifespan up to 3–4 weeks without the need for refrigeration. Ethylene released by the bananas encourage it to ripen; by removing it via oxidation, the bananas’ ripening is delayed.

Survival kits

Potassium permanganate is typically included in survival kits: as a fire starter (mixed with antifreeze from a car radiator or glycerin), water sterilizer, and for creating distress signals on snow.

As an oxidizer that generates the dark brown product MnO₂, potassium permanganate rapidly stains virtually any organic material such as skin, paper, and clothing. Solid KMnO₄ is a strong oxidizer and thus should be kept separated from oxidizable substances. Reaction with concentrated sulfuric acid produces the highly explosivemanganese(VII) oxide (Mn₂O₇). When solid KMnO₄ is mixed with pure glycerol or other simple alcohols it will result in a violent combustion reaction.

3 C₃H₅(OH)₃ + 14 KMnO₄ → 14 MnO₂ + 7 K₂CO₃ + 2 CO₂ + 12 H₂O

 

Applications

Manganese has no satisfactory substitute in its major applications, which are related to metallurgical alloy use. In minor applications, (e.g., manganese phosphating), zinc and sometimes vanadium are viable substitutes. In disposable battery manufacture, standard and alkaline cells using manganese will be generally replaced in the future with lithium battery technology.

Steel

US Marine Corps steel helmet

Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties. Steelmaking, including its ironmaking component, has accounted for most manganese demand, presently in the range of 85% to 90% of the total demand. Among a variety of other uses, manganese is a key component of low-cost stainless steel formulations.

Small amounts of manganese improve the workability of steel at high temperatures, because it forms a high melting sulfide and therefore prevents the formation of a liquid iron sulfide at the grain boundaries. If the manganese content reaches 4%, the embrittlement of the steel becomes a dominant feature. The embrittlement decreases at higher manganese concentrations and reaches an acceptable level at 8%. Steel containing 8 to 15% of manganese can have a high tensile strength of up to 863 MPa. Steel with 12% manganese was used for British steel helmets. This steel composition was discovered in 1882 by Robert Hadfield and is still known as Hadfield steel.

Aluminium alloys

The second large application for manganese is as alloying agent for aluminium. Aluminium with a manganese content of roughly 1.5% has an increased resistance against corrosion due to the formation of grains absorbing impurities which would lead to galvanic corrosion. The corrosion-resistant aluminium alloys 3004 and 3104 with a manganese content of 0.8 to 1.5% are the alloys used for most of the beverage cans. Before year 2000, in excess of 1.6 million tonnes have been used of those alloys; with a content of 1% manganese, this amount would need 16,000 tonnes of manganese.

Other uses

Methylcyclopentadienyl manganese tricarbonyl is used as an additive in unleaded gasoline to boost octane rating and reduce engine knocking. The manganese in this unusual organometallic compound is in the +1 oxidation state.

Manganese(IV) oxide (manganese dioxide, MnO₂) is used as a reagent in organic chemistry for the oxidation of benzylic alcohols (i.e. adjacent to an aromatic ring). Manganese dioxide has been used since antiquity to oxidatively neutralize the greenish tinge in glass caused by trace amounts of iron contamination. MnO₂ is also used in the manufacture of oxygen and chlorine, and in drying black paints. In some preparations, it is a brown pigment that can be used to make paint and is a constituent of natural umber.

Manganese(IV) oxide was used in the original type of dry cell battery as an electron acceptor from zinc, and is the blackish material found when opening carbon–zinc type flashlight cells. The manganese dioxide is reduced to the manganese oxide-hydroxide MnO(OH) during discharging, preventing the formation of hydrogen at the anode of the battery.

MnO₂ + H₂O + -e → MnO(OH) + OH⁻

The same material also functions iewer alkaline batteries (usually battery cells), which use the same basic reaction, but a different electrolyte mixture. In 2002, more than 230,000 tons of manganese dioxide was used for this purpose.

World-War-II-era 5-cent coin (1942-5 identified by mint mark P,D or S above dome) made from a 56% copper-35% silver-9% manganese alloy

The metal is very occasionally used in coins; until 2000, the only United States coin to use manganese was the “wartime” nickel from 1942 to 1945. An alloy of 75% copper and 25% nickel was traditionally used for the production of nickel coins. However, because of shortage of nickel metal during the war, it was substituted by more available silver and manganese, thus resulting in an alloy of 56% copper, 35% silver and 9% manganese. Since 2000, dollar coins, for example the Sacagawea dollar and the Presidential $1 coins, are made from a brass containing 7% of manganese with a pure copper core. In both cases of nickel and dollar, the use of manganese in the coin was to duplicate the electromagnetic properties of a previous identically sized and valued coin, for vending purposes. In the case of the later U.S. dollar coins, the manganese alloy was an attempt to duplicate properties of the copper/nickel alloy used in the previous Susan B. Anthony dollar.

Manganese compounds have been used as pigments and for the coloring of ceramics and glass. The brown color of ceramic is sometimes based on manganese compounds. In the glass industry, manganese compounds are used for two effects. Manganese(III) reacts with iron(II) to induce a strong green color in glass by forming less-colored iron(III) and slightly pink manganese(II), compensating for the residual color of the iron(III). Larger amounts of manganese are used to produce pink colored glass.

Biological role

Reactive center of arginase with boronic acid inhibitor – the manganese atoms are shown in yellow

Biological role

Manganese is an essential trace nutrient in all forms of life. The classes of enzymes that have manganese cofactors are very broad, and include oxidoreductases, transferases, hydrolases, lyases, isomerases, ligases, lectins, and integrins. The reverse transcriptases of many retroviruses (though not lentiviruses such as HIV) contain manganese. The best-known manganese-containing polypeptides may be arginase, the diphtheria toxin, and Mn-containing superoxide dismutase (Mn-SOD).

Mn-SOD is the type of SOD present in eukaryotic mitochondria, and also in most bacteria (this fact is in keeping with the bacterial-origin theory of mitochondria). The Mn-SOD enzyme is probably one of the most ancient, for nearly all organisms living in the presence of oxygen use it to deal with the toxic effects of superoxide, formed from the 1-electron reduction of dioxygen. Exceptions include a few kinds of bacteria, such as Lactobacillus plantarum and related lactobacilli, which use a different nonenzymatic mechanism, involving manganese (Mn²⁺) ions complexed with polyphosphate directly for this task, indicating how this function possibly evolved in aerobic life.

The human body contains about 12 mg of manganese, which is stored mainly in the bones; in the tissue, it is mostly concentrated in the liver and kidneys. In the human brain, the manganese is bound to manganese metalloproteins, most notably glutamine synthetase in astrocytes.

Manganese is also important in photosynthetic oxygen evolution in chloroplasts in plants. The oxygen-evolving complex (OEC) is a part of photosystem II contained in the thylakoid membranes of chloroplasts; it is responsible for the terminal photooxidation of water during the light reactions of photosynthesis, and has a metalloenzyme core containing four atoms of manganese. For this reason, most broad-spectrum plant fertilizers contain manganese.

Precautions

Manganese compounds are less toxic than those of other widespread metals, such as nickel and copper. However, exposure to manganese dusts and fumes should not exceed the ceiling value of 5 mg/m³ even for short periods because of its toxicity level. Manganese poisoning has been linked to impaired motor skills and cognitive disorders.

The permanganate exhibits a higher toxicity than the manganese(II) compounds. The fatal dose is about 10 g, and several fatal intoxications have occurred. The strong oxidative effect leads to necrosis of the mucous membrane. For example, the esophagus is affected if the permanganate is swallowed. Only a limited amount is absorbed by the intestines, but this small amount shows severe effects on the kidneys and on the liver.

In 2005, a study suggested a possible link between manganese inhalation and central nervous system toxicity in rats.

Manganese exposure in United States is regulated by Occupational Safety and Health Administration.

Generally, exposure to ambient Mn air concentrations in excess of 5 μg Mn/m3 can lead to Mn-induced symptoms. Increased ferroportin protein expression in human embryonic kidney (HEK293) cells is associated with decreased intracellular Mn concentration and attenuated cytotoxicity, characterized by the reversal of Mn-reduced glutamate uptake and diminished lactate dehydrogenase leakage.

Environmental health concerns

Manganese in drinking water

Waterborne manganese has a greater bioavailability than dietary manganese. According to results from a 2010 study, higher levels of exposure to manganese in drinking water are associated with increased intellectual impairment and reduced intelligence quotients in school-age children. It is hypothesized that long-term exposure to the naturally occurring manganese in shower water puts up to 8.7 million Americans at risk.

Manganese in gasoline

Methylcyclopentadienyl manganese tricarbonyl (MMT) is a gasoline additive used to replace lead compounds for unleaded gasolines, to improve the octane number in low octane number petrol distillates. It functions as an antiknock agent by the action of the carbonyl groups. Fuels containing manganese tend to form manganese carbides, which damage exhaust valves. The need to use lead or manganese compounds is merely historic, as the availability of reformation processes which create high-octane rating fuels increased. The use of such fuels directly or in mixture with non-reformed distillates is universal in developed countries (EU, Japan, etc.). In USA the imperative to provide the lowest possible price per volume on motor fuels (low fuel taxation rate) and lax legistation of fuel content (before 2000) caused refineries to use MMT. Compared to 1953, levels of manganese in air have dropped. Many racing competitions specifically ban manganese compounds in racing fuel (cart, minibike). MMT contains 24.4–25.2% manganese. There is strong correlation between elevated atmospheric manganese concentrations and automobile traffic density.

Role ieurological disorders

Manganism

Manganese overexposure is most frequently associated with manganism, a rare neurological disorder associated with excessive manganese ingestion or inhalation. Historically, persons employed in the production or processing of manganese alloys have been at risk for developing manganism; however, current health and safety regulations protect workers in developed nations. The disorder was first described in 1837 by British academic John Couper, who studied two patients who were manganese grinders

Manganism is a biphasic disorder. In its early stages, an intoxicated person may experience depression, mood swings, compulsive behaviors, and psychosis. Early neurological symptoms give way to late-stage manganism, which resembles Parkinson’s disease. Symptoms include weakness, monotone and slowed speech, an expressionless face, tremor, forward-leaning gait, inability to walk backwards without falling, rigidity, and general problems with dexterity, gait and balance. Unlike Parkinson’s disease, manganism is not associated with loss of smell and patients are typically unresponsive to treatment with L-DOPA. Symptoms of late-stage manganism become more severe over time even if the source of exposure is removed and brain manganese levels return to normal.

Childhood developmental disorders

Several recent studies attempt to examine the effects of chronic low-dose manganese overexposure on development in children. The earliest study of this kind was conducted in the Chinese province of Shanxi. Drinking water there had been contaminated through improper sewage irrigation and contained 240–350 µg Mn/L. Although WMn concentrations at or below 300 µg Mn/L are considered safe by the US EPA and 400 µg Mn/L are considered safe by the WHO, the 92 children sampled (between 11 and 13 years of age) from this province displayed lower performance on tests of manual dexterity and rapidity, short-term memory, and visual identification when compared to children from an uncontaminated area. More recently, a study of 10-year-old children in Bangladesh showed a relationship between WMn concentration in well water and diminished IQ scores. A third study conducted in Quebec examined school children between the ages of 6 and 15 living in homes that received water from a well containing 610 µg Mn/L; controls lived in homes that received water from a 160 µg Mn/L well. Children in the experimental group showed increased hyperactive and oppositional behaviours.

Neurodegenerative diseases

Chronic low-dose manganese intoxication is strongly implicated in a number of neurodegenerative disorders, including Alzheimer’s disease, Parkinson’s disease, and amyotrophic lateral sclerosis. It may also play a role in the development of multiple sclerosis, restless leg syndrome, and Huntington’s disease. A protein called DMT1 is the major transporter involved in manganese absorption from the intestine, and may be the major transporter of manganese across the blood–brain barrier. DMT1 also transports inhaled manganese across the nasal epithelium. The putative mechanism of action is that manganese overexposure and/or dysregulation leads to oxidative stress, mitochondrial dysfunction, glutamate-mediated excitoxicity, and aggregates of protein.

 

RHENIUM

Rhenium is a chemical element with the symbol Re and atomic number 75. It is a silvery-white, heavy, third-row transition metal in group 7 of the periodic table. With an estimated average concentration of 1 part per billion (ppb), rhenium is one of the rarest elements in the Earth’s crust. The free element has the third-highest melting point and highest boiling point of any element. Rhenium resembles manganese chemically and is obtained as a by-product of molybdenum and copper ore’s extraction and refinement. Rhenium shows in its compounds a wide variety of oxidation states ranging from −1 to +7.

Discovered in 1925, rhenium was the last stable element to be discovered. It was named after the river Rhine in Europe.

Nickel-based superalloys of rhenium are used in the combustion chambers, turbine blades, and exhaust nozzles of jet engines, these alloys contain up to 6% rhenium, making jet engine construction the largest single use for the element, with the chemical industry’s catalytic uses being next-most important. Because of the low availability relative to demand, rhenium is among the most expensive of metals, with an average price of approximately US$4,575 per kilogram (US$142.30 per troy ounce) as of August 2011; it is also of critical strategic military importance, for its use in high performance military jet and rocket engines.

History

Rhenium (Latin: Rhenus meaning: “Rhine”) was the last element to be discovered having a stable isotope (other new radioactive elements have been discovered iature since then, such as neptunium and plutonium). The existence of a yet undiscovered element at this position in the periodic table had been first predicted by Dmitry Mendeleev. Other calculated information was obtained by Henry Moseley in 1914. It is generally considered to have been discovered by Walter Noddack, Ida Tacke, and Otto Berg in Germany. In 1925 they reported that they detected the element in platinum ore and in the mineral columbite. They also found rhenium in gadolinite and molybdenite. In 1928 they were able to extract 1 g of the element by processing 660 kg of molybdenite. It was estimated in 1968 that 75% of the rhenium metal in the United States was used for research and the development of refractory metal alloys. It took several years from that point on before the super alloys became widely used.

In 1908, Japanese chemist Masataka Ogawa announced that he discovered the 43rd element and named it nipponium (Np) after Japan (Nippon in Japanese). However, later analysis indicated the presence of rhenium (element 75), not element 43. The symbol Np was later used for the element neptunium.

Characteristics

Rhenium is a silvery-white metal with one of the highest melting points of all elements, exceeded by only tungsten and carbon. It is also one of the densest, exceeded only by platinum, iridium and osmium. Rhenium has a hexagonal close-packed crystal structure, with lattice parameters a = 276.1 pm and c = 445.6 pm.

Its usual commercial form is a powder, but this element can be consolidated by pressing and sintering in a vacuum or hydrogen atmosphere. This procedure yields a compact solid having a density above 90% of the density of the metal. When annealed this metal is very ductile and can be bent, coiled, or rolled. Rhenium-molybdenum alloys are superconductive at 10 K; tungsten-rhenium alloys are also superconductive around 4–8 K, depending on the alloy. Rhenium metal superconducts at 1.697 ± 0.006 K.

In bulk form and at room temperature and atmospheric pressure, the element resists alkalis, sulfuric acid, hydrochloric acid, dilute, but not concentrated nitric acid, and aqua regia.

Isotopes

Rhenium has a stable isotope, rhenium-185, which nevertheless occurs in minority abundance, a situation found only in two other elements (indium and tellurium). Naturally occurring rhenium is 37.4% ¹⁸⁵Re, which is stable, and 62.6% ¹⁸⁷Re, which is unstable but has a very long half-life (~10¹⁰ years). This lifetime is affected by the charge state of rhenium atom. The beta decay of ¹⁸⁷Re is used for rhenium-osmium dating of ores. The available energy for this beta decay (2.6 keV) is one of the lowest known among all radionuclides. There are twenty-six other recognized radioactive isotopes of rhenium.

Compounds

Rhenium compounds are known for all nine oxidation states between −1 and +7. The oxidation states +7, +6, +4, and +2 are the most common. Rhenium is most available commercially as salts of perrhenate, including sodium and ammonium perrhenates. These are white, water-soluble compounds.

Halides and oxyhalides

The most common rhenium chlorides are ReCl₆, ReCl₅, ReCl₄, and ReCl₃. The structures of these compounds often feature extensive Re-Re bonding, which is characteristic of this metal in oxidation states lower than VII. Salts of [Re₂Cl₈]^2- feature a quadruple metal-metal bond. Although the highest rhenium chloride features Re(VI), fluorine gives the d⁰ Re(VII) derivative rhenium heptafluoride. Bromides and iodides of rhenium are also well known.

Like tungsten and molybdenum, with which it shares chemical similarities, rhenium forms a variety of oxyhalides. The oxychlorides are most common, and include ReOCl₄, ReO₃Cl.

Oxides and sulfides

Perrhenic acid adopts an unconventional structure.

The most common oxide is the volatile colourless Re₂O₇, which adopts a molecular structure, unlike most metal oxides. The d1 species ReO₃ adopts a defect perovskite structure. Other oxides include Re₂O₅, ReO₂, and Re₂O₃. The sulfides are ReS₂ and Re₂S₇. Perrhenate salts can be converted to tetrathioperrhenate by the action of ammonium hydrosulfide.

Other compounds

Rhenium diboride (ReB₂) is a hard compound having the hardness similar to that of tungsten carbide, silicon carbide, titanium diboride or zirconium diboride.

Organorhenium compounds

Dirhenium decacarbonyl is the most common entry to organorhenium chemistry. Its reduction with sodium amalgam gives Na[Re(CO)₅] with rhenium in the formal oxidation state −1. Dirhenium decacarbonyl can be oxidised with bromine to bromopentacarbonylrhenium(I):

Re₂(CO)₁₀ + Br₂ → 2 Re(CO)₅Br

Reduction of this pentacarbonyl with zinc and acetic acid gives pentacarbonylhydridorhenium:

Re(CO)₅Br + Zn + HOAc → Re(CO)₅H + ZnBr(OAc)

Methylrhenium trioxide (“MTO”), CH₃ReO₃ is a volatile, colourless solid has been used as a catalyst in some laboratory experiments. It can be prepared by many routes, a typical method is the reaction of Re₂O₇ and tetramethyltin:

Re₂O₇ + (CH₃)₄Sn → CH₃ReO₃ + (CH₃)₃SnOReO₃

Analogous alkyl and aryl derivatives are known. MTO catalyses for the oxidations with hydrogen peroxide. Terminal alkynes yield the corresponding acid or ester, internal alkynes yield diketones, and alkenes give epoxides. MTO also catalyses the conversion of aldehydes and diazoalkanes into an alkene.

Nonahydridorhenate

Structure of ReH²⁻₉

A distinctive derivative of rhenium is nonahydridorhenate, originally thought to be the rhenide anion, Re−, but actually containing the ReH²⁻₉ anion in which the oxidation state of rhenium is +7.

Occurrence

Molybdenite

Rhenium is one of the rarest elements in Earth’s crust with an average concentration of 1 ppb; other sources quote the number of 0.5 ppb making it the 77th most abundant element in Earth’s crust. Rhenium is probably not found free iature (its possible natural occurrence is uncertain), but occurs in amounts up to 0.2% in the mineral molybdenite (which is primarily molybdenum disulfide), the major commercial source, although single molybdenite samples with up to 1.88% have been found. Chile has the world’s largest rhenium reserves, part of the copper ore deposits, and was the leading producer as of 2005. It was only recently that the first rhenium mineral was found and described (in 1994), a rhenium sulfide mineral (ReS₂) condensing from a fumarole on Russia’s Kudriavy volcano, Iturup island, in the Kurile Islands. Kudryavy discharges up to 20–60 kg rhenium per year mostly in the form of rhenium disulfide. Named rheniite, this rare mineral commands high prices among collectors.

Production

Ammonium perrhenate

Commercial rhenium is extracted from molybdenum roaster-flue gas obtained from copper-sulfide ores. Some molybdenum ores contain 0.001% to 0.2% rhenium. Rhenium(VII) oxide and perrhenic acid readily dissolve in water; they are leached from flue dusts and gasses and extracted by precipitating with potassium or ammonium chloride as the perrhenate salts, and purified by recrystallization. Total world production is between 40 and 50 tons/year; the main producers are in Chile, the United States, Peru, and Kazakhstan. Recycling of used Pt-Re catalyst and special alloys allow the recovery of another 10 tons per year. Prices for the metal rose rapidly in early 2008, from $1000–$2000 per kg in 2003–2006 to over $10,000 in February 2008. The metal form is prepared by reducing ammonium perrhenate with hydrogen at high temperatures:

2 NH₄ReO₄ + 7 H₂ → 2 Re + 8 H₂O + 2 NH₃

Applications

The Pratt & Whitney F-100 engine uses rhenium-containing second-generation superalloys

Rhenium is added to high-temperature superalloys that are used to make jet engine parts, using 70% of the worldwide rhenium production. Another major application is in platinum-rhenium catalysts, which are primarily used in making lead-free, high-octane gasoline.

Alloys

The nickel-based superalloys have improved creep strength with the addition of rhenium. The alloys normally contain 3% or 6% of rhenium. Second generation alloys contain 3%; these alloys were used in the engines of the F-16 and F-15, while the newer single-crystal third generation alloys contain 6% of rhenium; they are used in the F-22 and F-35 engines. Rhenium is also used in the superalloys, such as CMSX-4 (2nd gen) and CMSX-10 (3rd gen) that are used in industrial gas turbine engines like the GE 7FA. Rhenium can cause superalloys to become microstructurally unstable, forming undesirable TCP (topologically close packed) phases. In 4th and 5th generation superalloys, ruthenium is used to avoid this effect. Among others the new superalloys are EPM-102 (with 3% Ru) and TMS-162 (with 6% Ru), both containing 6% rhenium, as well as TMS-138 and TMS-174.

CFM International CFM56 jet engine still with blades made with 3% rhenium

For 2006, the consumption is given as 28% for General Electric, 28% Rolls-Royce plc and 12% Pratt & Whitney, all for superalloys, while the use for catalysts only accounts for 14% and the remaining applications use 18%. In 2006, 77% of the rhenium consumption in the United States was in alloys. The rising demand for military jet engines and the constant supply made it necessary to develop superalloys with a lower rhenium content. For example the newer CFM International CFM56 high-pressure turbine (HPT) blades will use Rene N515 with a rhenium content of 1.5% instead of Rene N5 with 3%.

Rhenium improves the properties of tungsten. Tungsten-rhenium alloys are more ductile at low temperature, allowing them to be more easily machined. The high-temperature stability is also improved. The effect increases with the rhenium concentration, and therefore tungsten alloys are produced with up to 27% of Re, which is the solubility limit. One application for the tungsten-rhenium alloys is X-ray sources. The high melting point of both compounds, together with the high atomic mass, makes them stable against the prolonged electron impact. Rhenium tungsten alloys are also applied as thermocouples to measure temperatures up to 2200 °C.

The high temperature stability, low vapor pressure, good wear resistance and ability to withstand arc corrosion of rhenium are useful in self-cleaning electrical contacts. In particular, the discharge occurring during the switching oxidizes the contacts. However, rhenium oxide Re₂O₇ has poor stability (sublimes at ~360 °C) and therefore is removed during the discharge.

Rhenium has a high melting point and a low vapor pressure similar to tantalum and tungsten. Therefore, rhenium filaments exhibit a higher stability if the filament is operated not in vacuum, but in oxygen-containing atmosphere. Those filaments are widely used in mass spectrometers, in ion gauges and in photoflash lamps in photography.

Catalysts

Rhenium in the form of rhenium-platinum alloy is used as catalyst for catalytic reforming, which is a chemical process to convert petroleum refinery naphthas with low octane ratings into high-octane liquid products. Worldwide, 30% of catalysts used for this process contain rhenium. The olefin metathesis is the other reaction for which rhenium is used as catalyst. Normally Re₂O₇ on alumina is used for this process. Rhenium catalysts are very resistant to chemical poisoning from nitrogen, sulfur and phosphorus, and so are used in certain kinds of hydrogenation reactions.

Other uses

¹⁸⁸Re and ¹⁸⁶Re isotopes are radioactive and are used for treatment of liver cancer. They both have similar penetration depth in tissue (5 mm for ¹⁸⁶Re and 11 mm for ¹⁸⁸Re), but ¹⁸⁶Re has advantage of longer lifetime (90 hours vs. 17 hours).

Related by periodic trends, rhenium has a similar chemistry with technetium; work done to label rhenium onto target compounds can often be translated to technetium. This is useful for radiopharmacy, where it is difficult to work with technetium – especially the 99m isotope used in medicine – due to its expense and short half-life.

High-temperature rhenium super-alloys are used to make jet engine parts. Platinum-rhenium catalysts are used in lead-free, high-octane, gasoline. Rhenium is also used as a filament for mass spectrographs and ion gauges. Because Rhenium has good wear resistance and withstands arc corrosion, it is used as an electrical contact material. Thermocouples made of Re-W are used for measuring temperatures up to 2200C, and rhenium wire is used in photographic flash lamps. Rhenium is available as metal and compounds with purities from 99% to 99.999% (ACS grade to ultra-high purity). High Purity (99.999%) Rhenium Oxide (ReO2) PowderElemental or metallic forms include pellets, rod, wire and granules for evaporation source material purposes. High Purity (99.999%) Rhenium (Re) Sputtering TargetRhenium nanoparticles and nanopowders provide ultra-high surface area which nanotechnology research and recent experiments demonstrate function to create new and unique properties and benefits. Oxides are available in powder and dense pellet form for such uses as optical coating and thin film applications. Oxides tend to be insoluble. Fluorides are another insoluble form for uses in which oxygen is undesirable such as metallurgy, chemical and physical vapor deposition and in some optical coatings. Rhenium is also available in soluble forms including chlorides, nitrates and acetates. These compounds can be manufactured as solutions at specified stoichiometries.

Very little is known about the toxicity of rhenium and its compounds. Safety data for Rhenium and its compounds can vary widely depending on the form. For potential hazard information, toxicity, and road, sea and air transportation limitations, such as DOT Hazard Class, DOT Number, EU Number, NFPA Health rating and RTECS Class, please see the specific material or compound referenced in the Products tab below.

 

Practical skills

Qualitative tests on VII group d-elements

Mn²⁺

2MnSO₄ + 16HNO₃(diluted) + 5NaBiO₃ ® 2HMnO₄ + 5Bi(NO₃)₃ + 2Na₂SO₄ + NaNO₃ + 7H₂O – purple violet solution

MnO₄^–

2KMnO₄ + 3H₂SO₄(diluted) + 5H₂O₂ ® 2MnSO₄ + 5O₂­ + 8H₂O + K₂SO₄

 

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

9. http://en.wikipedia.org/wiki

The following website shows the reaction of d-elements. It’s cool stuff! Check it out!

www.youtube.com/watch?v=K0FUH54U-XQ

www.youtube.com/watch?v=Xkv1v7pee7E

www.youtube.com/watch?v=X2dgpObsJkg

www.youtube.com/watch?v=yQcAfZMAUDQ

www.youtube.com/watch?v=xuV_Ule3NWo

www.youtube.com/watch?v=W5X9NqPIwcA

 

Prepared by PhD Falfushynska H.