A | B | C | D | E | F | G | H | CH | I | J | K | L | M | N | O | P | Q | R | S | T | U | V | W | X | Y | Z | 0 | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9
 Pure manganese cube and oxidized manganese chips | |||||||||||||||||||||||||||||||||||
| Manganese | |||||||||||||||||||||||||||||||||||
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| Pronunciation | /ˈmæŋɡəniːz/ | ||||||||||||||||||||||||||||||||||
| Appearance | silvery metallic | ||||||||||||||||||||||||||||||||||
| Standard atomic weight Ar°(Mn) | |||||||||||||||||||||||||||||||||||
| Manganese in the periodic table | |||||||||||||||||||||||||||||||||||
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| Atomic number (Z) | 25 | ||||||||||||||||||||||||||||||||||
| Group | group 7 | ||||||||||||||||||||||||||||||||||
| Period | period 4 | ||||||||||||||||||||||||||||||||||
| Block | d-block | ||||||||||||||||||||||||||||||||||
| Electron configuration | [Ar] 3d5 4s2 | ||||||||||||||||||||||||||||||||||
| Electrons per shell | 2, 8, 13, 2 | ||||||||||||||||||||||||||||||||||
| Physical properties | |||||||||||||||||||||||||||||||||||
| Phase at STP | solid | ||||||||||||||||||||||||||||||||||
| Melting point | 1519 K (1246 °C, 2275 °F) | ||||||||||||||||||||||||||||||||||
| Boiling point | 2334 K (2061 °C, 3742 °F) | ||||||||||||||||||||||||||||||||||
| Density (at 20° C) | 7.476 g/cm3 [3] | ||||||||||||||||||||||||||||||||||
| when liquid (at m.p.) | 5.95 g/cm3 | ||||||||||||||||||||||||||||||||||
| Heat of fusion | 12.91 kJ/mol | ||||||||||||||||||||||||||||||||||
| Heat of vaporization | 221 kJ/mol | ||||||||||||||||||||||||||||||||||
| Molar heat capacity | 26.32 J/(mol·K) | ||||||||||||||||||||||||||||||||||
Vapor pressure
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| Atomic properties | |||||||||||||||||||||||||||||||||||
| Oxidation states | −3, −1, 0, +1, +2, +3, +4, +5, +6, +7 (depending on the oxidation state, an acidic, basic, or amphoteric oxide) | ||||||||||||||||||||||||||||||||||
| Electronegativity | Pauling scale: 1.55 | ||||||||||||||||||||||||||||||||||
| Ionization energies |
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| Atomic radius | empirical: 127 pm | ||||||||||||||||||||||||||||||||||
| Covalent radius | Low spin: 139±5 pm High spin: 161±8 pm | ||||||||||||||||||||||||||||||||||
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| Other properties | |||||||||||||||||||||||||||||||||||
| Natural occurrence | primordial | ||||||||||||||||||||||||||||||||||
| Crystal structure | α-Mn: body-centered cubic (bcc) (cI58) | ||||||||||||||||||||||||||||||||||
| Lattice constant |  | ||||||||||||||||||||||||||||||||||
| Thermal expansion | 23.61×10−6/K (at 20 °C)[3] | ||||||||||||||||||||||||||||||||||
| Thermal conductivity | 7.81 W/(m⋅K) | ||||||||||||||||||||||||||||||||||
| Electrical resistivity | 1.44 µΩ⋅m (at 20 °C) | ||||||||||||||||||||||||||||||||||
| Magnetic ordering | paramagnetic | ||||||||||||||||||||||||||||||||||
| Molar magnetic susceptibility | (α) +529.0×10−6 cm3/mol (293 K)[4] | ||||||||||||||||||||||||||||||||||
| Young's modulus | 198 GPa | ||||||||||||||||||||||||||||||||||
| Bulk modulus | 120 GPa | ||||||||||||||||||||||||||||||||||
| Speed of sound thin rod | 5150 m/s (at 20 °C) | ||||||||||||||||||||||||||||||||||
| Mohs hardness | 6.0 | ||||||||||||||||||||||||||||||||||
| Brinell hardness | 196 MPa | ||||||||||||||||||||||||||||||||||
| CAS Number | 7439-96-5 | ||||||||||||||||||||||||||||||||||
| History | |||||||||||||||||||||||||||||||||||
| Discovery | Carl Wilhelm Scheele (1774) | ||||||||||||||||||||||||||||||||||
| First isolation | Johann Gottlieb Gahn (1774) | ||||||||||||||||||||||||||||||||||
| Isotopes of manganese | |||||||||||||||||||||||||||||||||||
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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. Manganese was first isolated in the 1770s. Manganese is a transition metal with a multifaceted array of 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, fertilisers, 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.[6] It is found mostly in the bones, but also the liver, kidneys, and brain.[7] In the human brain, the manganese is bound to manganese metalloproteins, most notably glutamine synthetase in astrocytes.
It is familiar in the laboratory in the form of the deep violet salt potassium permanganate. It occurs at the active sites in some enzymes.[8] 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 fuse, but easy to oxidize.[9] Manganese metal and its common ions are paramagnetic.[10] Manganese tarnishes slowly in air and oxidizes ("rusts") like iron in water containing dissolved oxygen.[11]
Isotopes
Naturally occurring manganese is composed of one stable isotope, 55Mn. Several radioisotopes have been isolated and described, ranging in atomic weight from 46 u (46Mn) to 72 u (72Mn). The most stable are 53Mn with a half-life of 3.7 million years, 54Mn with a half-life of 312.2 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.[12] Manganese also has three meta states.[12]
Manganese is part of the iron group of elements, which are thought to be synthesized in large stars shortly before the supernova explosion.[13] 53Mn decays to 53Cr with a half-life of 3.7 million years. Because of its relatively short half-life, 53Mn is relatively rare, produced by cosmic rays impact on iron.[14] 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 suggest an initial 53Mn/55Mn ratio, which indicate that 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.[15][16][17][18]
Allotropes
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Four allotropes (structural forms) of solid manganese are known, labeled α, β, γ and δ, and occurring 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.[19][20][21]
Alpha manganese (α-Mn) is the equilibrium phase at room temperature. It has a body-centered cubic lattice and is unusual among elemental metals in having a very complex unit cell, with 58 atoms per cell (29 atoms per primitive unit cell) in four different types of site.[22][19] It is paramagnetic at room temperature and antiferromagnetic at temperatures below 95 K (−178 °C).[23]
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Beta manganese (β-Mn) forms when heated above the transition temperature of 973 K (700 °C; 1,290 °F). 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.[24] It is easily obtained as a metastable phase at room temperature by rapid quenching. It does not show magnetic ordering, remaining paramagnetic down to the lowest temperature measured (1.1 K).[24][25][26]
Gamma manganese (γ-Mn) forms when heated above 1,370 K (1,100 °C; 2,010 °F). It has a simple face-centered cubic structure (four atoms per unit cell). 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 (such as C, Fe, Ni, Cu, Pd or Au), and these solute-stabilized alloys distort into a face-centered tetragonal structure.[27][26]
Delta manganese (δ-Mn) forms when heated above 1,406 K (1,130 °C; 2,070 °F) and is stable up to the manganese melting point of 1,519 K (1,250 °C; 2,270 °F). It has a body-centered cubic structure (two atoms per cubic unit cell).[20][26]
Chemical compounds
Common oxidation states of manganese are +2, +3, +4, +6, and +7, although all oxidation states from −3 to +7 except –2 have been observed. Manganese in oxidation state +7 is represented by salts of the intensely purple permanganate anion MnO4−. Potassium permanganate is a commonly used laboratory reagent because of its oxidizing properties; it is used 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.[29]
Aside from various permanganate salts, Mn(VII) is represented by the unstable, volatile derivative Mn2O7. Oxyhalides (MnO3F and MnO3Cl) are powerful oxidizing agents.[9] 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 (O2−) or nitride (N3−) ligand.[30][31] One example is the blue anion hypomanganate 3−.
Mn(IV) 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 but is also a common ingredient in dry cell batteries. Complexes of Mn(IV) are well known, but they require elaborate ligands. Mn(IV)-OH complexes are an intermediate in some enzymes, including the oxygen evolving center (OEC) in plants.[32]
Simple derivatives Mn+3 are rarely encountered but can be stabilized by suitably basic ligands. Manganese(III) acetate is an oxidant useful in organic synthesis. Solid compounds of manganese(III) are characterized by its strong purple-red color and a preference for distorted octahedral coordination resulting from the Jahn-Teller effect.[citation needed]
A particularly common oxidation state for manganese in aqueous solution is +2, which has a pale pink color. Many manganese(II) compounds are known, such as the aquo complexes derived from manganese(II) sulfate (MnSO4) and manganese(II) chloride (MnCl2). This oxidation state is also seen in the mineral rhodochrosite (manganese(II) carbonate). Manganese(II) commonly exists with a high spin, S = 5/2 ground state because of the high pairing energy for manganese(II). There are no spin-allowed d–d transitions in manganese(II), which explain its faint color.[33]
| Oxidation states of manganese[34] | |
|---|---|
| −3 | Mn(CO)(NO) 3 |
| −1 | HMn(CO) 5 |
| 0 | Mn 2(CO) 10 |
| +1 | MnC 5H 4CH 3(CO) 3 |
| +2 | MnCl 2, MnCO 3, MnO |
| +3 | MnF 3, Mn(OAc) 3, Mn 2O 3 |
| +4 | MnO 2 |
| +5 | K 3MnO 4 |
| +6 | K 2MnO 4 |
| +7 | KMnO 4, Mn 2O 7 |
| Common oxidation states are in bold. | |
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(−III) up through Mn(I). This area of organometallic chemistry is attractive because Mn is inexpensive and of relatively low toxicity.[35]
Of greatest commercial interest is "MMT", methylcyclopentadienyl manganese tricarbonyl, which is used as an anti-knock compound added to gasoline (petrol) in some countries. It features Mn(I). Consistent with other aspects of Mn(II) chemistry, manganocene (Mn(C5H5)2) is high-spin. In contrast, its neighboring metal iron forms an air-stable, low-spin derivative in the form of ferrocene (Fe(C5H5)2). When conducted under an atmosphere of carbon monoxide, reduction of Mn(II) salts gives dimanganese decacarbonyl Mn2(CO)10, an orange and volatile solid. The air-stability of this Mn(0) compound (and its many derivatives) reflects the powerful electron-acceptor properties of carbon monoxide. Many alkene complexes and alkyne complexes are derived from Mn2(CO)10.[citation needed]
In Mn(CH3)2(dmpe)2, Mn(II) is low spin, which contrasts with the high spin character of its precursor, MnBr2(dmpe)2 (dmpe = (CH3)2PCH2CH2P(CH3)2).[36] 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−.[citation needed]
History
The origin of the name manganese is complex. In ancient times, two black minerals were identified from the regions of the Magnetes (either Magnesia, located within modern Greece, or Magnesia ad Sipylum, located within modern Turkey).[37] 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.[38] Neither this mineral nor elemental manganese is magnetic. In the 16th century, manganese dioxide was called manganesum (note the two Ns instead of one) by glassmakers, possibly as a corruption and concatenation of two words, since alchemists and glassmakers eventually had to differentiate a magnesia nigra (the black ore) from magnesia alba (a white ore, also from Magnesia, also useful in glassmaking). Michele Mercati called magnesia nigra manganesa, and finally the metal isolated from it became known as manganese (German: Mangan). The name magnesia eventually was then used to refer only to the white magnesia alba (magnesium oxide), which provided the name magnesium for the free element when it was isolated much later.[39]
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.[41]
Manganese compounds were used by Egyptian and Roman glassmakers, either to add to, or remove, color from glass.[42] Use as "glassmakers soap" continued through the Middle Ages until modern times and is evident in 14th-century glass from Venice.[43]
Because it was used in glassmaking, manganese dioxide was available for experiments by alchemists, the first chemists. Ignatius Gottfried Kaim (1770) and Johann Glauber (17th century) discovered that manganese dioxide could be converted to permanganate, a useful laboratory reagent.[44] Kaim also may have reduced manganese dioxide to isolate the metal, but that is uncertain.[45] 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. The production of chlorine and hypochlorite bleaching agents was a large consumer of manganese ores.[citation needed]
Scheele and others were aware that pyrolusite (mineral form of manganese dioxide) 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.[citation needed]
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.[46] 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.[47] 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.[48]
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 cathodic for commercial disposable dry batteries of both the standard (zinc–carbon) and alkaline types.[49]
Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties.[50] This application was first recognized by the British metallurgist Robert Forester Mushet (1811–1891) who, in 1856, introduced the element, in the form of Spiegeleisen.
Occurrence
Manganese comprises about 1000 ppm (0.1%) of the Earth's crust, the 12th most abundant of the crust's elements.[7] Soil contains 7–9000 ppm of manganese with an average of 440 ppm.[7] The atmosphere contains 0.01 μg/m3.[7] Manganese occurs principally as pyrolusite (MnO2), braunite (Mn2+Mn3+6)SiO12),[51] psilomelane (Ba,H2O)2Mn5O10, and to a lesser extent as rhodochrosite (MnCO3).
