Manganese

Manganese is a chemical element; it has symbol Mn and atomic number 25. It is a hard, brittle, silvery metal, often found in minerals in combination with iron. First isolated in the 1770s, manganese is a transition metal with many industrial alloy uses, particularly in stainless steels. It improves strength, workability, and resistance to wear. Manganese oxide is used as an oxidising agent, as a rubber additive, and in glass making, fertilizers, and ceramics. Manganese sulfate can be used as a fungicide.
Manganese is also an essential human dietary element, important in macronutrient metabolism, bone formation, and free radical defense systems. It is a critical component in dozens of proteins and enzymes. It is found mostly in the bones, but also the liver, kidneys, and brain. In the human brain, manganese is bound to manganese metalloproteins, most notably glutamine synthetase in astrocytes.
Manganese in the form of the deep violet salt potassium permanganate is commonly used in laboratories as an oxidizer. Potassium permanganate is also used as a biocide in water treatment.
It occurs at the active sites in some enzymes. Of particular interest is the use of a Mn–O cluster, the oxygen-evolving complex, in the production of oxygen by plants.

Characteristics

Physical properties

Manganese is a silvery-gray metal that resembles iron. It is hard and very brittle, difficult to melt, but oxidizes easily. Manganese and its common ions are paramagnetic. Manganese tarnishes slowly in air and oxidizes like iron in water containing dissolved oxygen.

Isotopes

Naturally occurring manganese is composed of one stable isotope, ⁵⁵Mn. Several radioisotopes have been isolated and described, ranging from ⁴⁶Mn to ⁷²Mn; the most stable ones are ⁵³Mn with a half-life of 3.7 million years, ⁵⁴Mn with a half-life of 312.08 days, and ⁵²Mn with a half-life of 5.591 days. All of the remaining radioactive isotopes have half-lives of less than three hours, and the majority of less than one minute. The primary decay mode in isotopes lighter than the most abundant stable isotope, ⁵⁵Mn, is electron capture, and the primary mode in heavier isotopes is beta decay. Manganese also has three meta states.
Manganese is part of the iron group of elements, which are thought to be synthesized in massive stars shortly before the supernova explosion. ⁵³Mn decays to ⁵³Cr with a half-life of 3.7 million years. Because of its short half-life, ⁵³Mn is relatively rare; it is produced by the impact of cosmic rays on iron.
Chromium and manganese are found together sufficiently for measurement of both to find application in isotope geology, and the Mn/Cr ratios here for dating the early Solar System. 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 suggest a non-zero initial ⁵³Mn/⁵⁵Mn ratio, which indicate that Cr isotopic composition variations must result from in situ decay of ⁵³Mn in differentiated planetary bodies. Hence, ⁵³Mn provides additional evidence for nucleosynthetic processes immediately before the coalescence of the Solar System.

Allotropes

Four allotropes of solid manganese are known, labeled α, β, γ and δ, and occur at successively higher temperatures. All are metallic, stable at standard pressure, and have a cubic crystal lattice, but they vary widely in their atomic structures.
Alpha manganese is the equilibrium phase at room temperature. It has a body-centered cubic lattice and is unusual among elemental metals in that it has a very complex unit cell, with 58 atoms per cell with manganese atoms in four different types of surroundings. It is paramagnetic at room temperature and antiferromagnetic at temperatures below.
Beta manganese forms when heated above the transition temperature of. It has a primitive cubic structure with 20 atoms per unit cell at two types of sites, which is as complex as that of any other elemental metal. It is easily obtained as a metastable phase at room temperature by rapid quenching of manganese at in ice water. It does not show magnetic ordering, remaining paramagnetic down to the lowest temperature measured.
Gamma manganese forms when heated above. It has a simple face-centered cubic structure. When quenched to room temperature it converts to β-Mn, but it can be stabilized at room temperature by alloying it with at least 5 percent of other elements. These solute-stabilized alloys distort into a face-centered tetragonal structure.
Delta manganese forms when heated above and is stable up to the manganese melting point of. It has a body-centered cubic structure.

Chemical compounds

Common oxidation states of manganese are +2, +3, +4, +6, and +7, although all oxidation states from −3 to +7 have been observed. Manganese in oxidation state +7 is represented by salts of the intensely purple permanganate anion. Potassium permanganate is a commonly used laboratory reagent because of its oxidizing properties; it is used as a topical medicine. 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.
Aside from various permanganate salts, Mn is represented by the unstable, volatile derivative Mn₂O₇. Oxyhalides are powerful oxidizing agents. The most prominent example of Mn in the +6 oxidation state is the green anion manganate, ²⁻. Manganate salts are intermediates in the extraction of manganese from its ores. Compounds with oxidation states +5 are somewhat elusive, and often found associated to an oxide or nitride ligand. One example is the blue anion hypomanganate ³⁻.
Mn is somewhat enigmatic because it is common in nature but far rarer in synthetic chemistry. The most common Mn ore, pyrolusite, is MnO₂. It is the dark brown pigment of many cave drawings and is also a common ingredient in dry cell batteries. Complexes of Mn, such as in manganese fluoride#Fluoromanganate complexes|K₂, are known but are rarer than those of manganese in the lower oxidation states. Mn-OH complexes are an intermediate in some enzymes, including the oxygen-evolving center in plants.
Simple derivatives of Mn³⁺ are rarely encountered but can be stabilized by suitably alkaline ligands. Manganese acetate is an oxidant useful in organic synthesis. Solid compounds of manganese are characterized by a strong purple-red color and a preference for distorted octahedral coordination resulting from the Jahn-Teller effect.
A particularly common oxidation state for manganese in aqueous solution is +2, which has a pale pink color. Many manganese compounds are known, such as the aquo complexes derived from manganese sulfate and manganese chloride. This oxidation state is also seen in the mineral rhodochrosite. Manganese commonly exists with a high-spin ground state, with 5 unpaired electrons, because of its high pairing energy. There are no spin-allowed d–d transitions in manganese, which explain its faint color.

Organomanganese compounds

Manganese forms a large variety of organometallic derivatives, i.e., compounds with Mn-C bonds. The organometallic derivatives include numerous examples of Mn in its lower oxidation states, i.e. Mn up through Mn. This area of organometallic chemistry is attractive because Mn is inexpensive and of relatively low toxicity.
Of greatest commercial interest is methylcyclopentadienyl manganese tricarbonyl, which is used as an anti-knock compound added to gasoline in some countries, featuring Mn. Consistent with other aspects of Mn chemistry, manganocene is high-spin. In contrast, its neighboring metal, iron, forms an air-stable, low-spin derivative in the form of ferrocene. When conducted under an atmosphere of carbon monoxide, reduction of Mn salts gives dimanganese decacarbonyl, an orange and volatile solid. The air-stability of this Mn compound reflects the powerful electron-acceptor properties of carbon monoxide. Many alkene complexes and alkyne complexes are derived from.
In Mn₂₂, Mn is low spin, which contrasts with the high spin character of its precursor, MnBr₂₂ ₂PCH₂CH₂P. Polyalkyl and polyaryl derivatives of manganese often exist in higher oxidation states, reflecting the electron-releasing properties of alkyl and aryl ligands. One example is ²⁻.

History

The origin of the name manganese is complex. In ancient times, two black minerals were identified from the regions of the Magnetes. They were both called magnes from their place of origin, but were considered to differ in sex. The male magnes attracted iron, and was the iron ore now known as lodestone or magnetite, and which probably gave us the term magnet. The female magnes ore did not attract iron, but was used to decolorize glass. This female magnes was later called magnesia, known now in modern times as pyrolusite or manganese dioxide. Neither this mineral nor elemental manganese is magnetic. In the 16th century, manganese dioxide was called manganesum by glassmakers, possibly as a corruption and concatenation of two words, since alchemists and glassmakers eventually had to differentiate a magnesia nigra from magnesia alba. Italian physician Michele Mercati called magnesia nigra manganesa, and finally the metal isolated from it became known as manganese. The name magnesia was eventually used to refer only to the white magnesia alba, which provided the name magnesium for the free element when it was isolated much later.
File:Lascaux painting.jpg|thumb| left |alt=A drawing of a left-facing bull, in black, on a cave wall |Some of the cave paintings in Lascaux, France, use manganese-based pigments.
Manganese dioxide, which is abundant in nature, has long been used as a pigment. The cave paintings in Gargas that are 30,000 to 24,000 years old are made from the mineral form of MnO₂ pigments.
Manganese compounds were used by Egyptian and Roman glassmakers, either to add to, or remove, color from glass. Use as "glassmakers soap" continued through the Middle Ages until modern times and is evident in 14th-century glass from Venice.
Because it was used in glassmaking, manganese dioxide was available for experiments by alchemists, the first chemists. Ignatius Gottfried Kaim and Johann Glauber discovered that manganese dioxide could be converted to permanganate, 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, and later hydrochloric acid from the Leblanc process was used and the manganese dioxide was recycled by the Weldon process.
Scheele and others were aware that pyrolusite contained a new element. Johan Gottlieb Gahn isolated an impure sample of manganese metal in 1774, which he did by reducing the dioxide with carbon. Ignatius Gottfried Kaim also may have reduced manganese dioxide to isolate the metal, but that is uncertain.
The manganese content of some iron ores used in Greece led to speculations that steel produced from that ore contains additional 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 documented that iron alloyed with manganese was harder but not more brittle. In 1837, British academic James Couper noted an association between miners' heavy exposure to manganese and a form of Parkinson's disease. In 1912, United States patents were granted for protecting firearms against rust and corrosion with manganese phosphate electrochemical conversion coatings, and the process has seen widespread use ever since.
The invention of the Leclanché cell in 1866 and the subsequent improvement of batteries containing manganese dioxide as cathodic depolarizer increased the demand for manganese dioxide. Until the development of batteries with nickel–cadmium and lithium, most batteries contained manganese. The zinc–carbon battery and the alkaline battery normally use industrially produced manganese dioxide because naturally occurring manganese dioxide contains impurities. In the 20th century, manganese dioxide was widely used as the cathode for commercial disposable dry batteries of both the standard and alkaline types.
Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties. This application was first recognized by the British metallurgist Robert Forester Mushet, who introduced the element to the steel manufacture process in 1856 in the form of spiegeleisen.