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chromiummanganeseiron
-

Mn

Tc
Appearance
silvery metallic
General properties
Name, symbol, number manganese, Mn, 25
Element category transition metal
Group, period, block 74, d
Standard atomic weight 54.938045(5)g·mol−1
Electron configuration [Ar] 4s2 3d5
Electrons per shell 2, 8, 13, 2 (Image)
Physical properties
Phase solid
Density (near r.t.) 7.21 g·cm−3
Liquid density at m.p. 5.95 g·cm−3
Melting point 1519 K, 1246 °C, 2275 °F
Boiling point 2334 K, 2061 °C, 3742 °F
Heat of fusion 12.91 kJ·mol−1
Heat of vaporization 221 kJ·mol−1
Specific heat capacity (25 °C) 26.32 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 1228 1347 1493 1691 1955 2333
Atomic properties
Oxidation states 7, 6, 5, 4, 3, 2, 1, -1, -2, -3
(oxides: acidic, basic or amphoteric
depending on the oxidation state)
Electronegativity 1.55 (Pauling scale)
Ionization energies
(more)
1st: 717.3 kJ·mol−1
2nd: 1509.0 kJ·mol−1
3rd: 3248 kJ·mol−1
Atomic radius 127 pm
Covalent radius 139±5 (low spin), 161±8 (high spin) pm
Miscellanea
Crystal structure body-centered cubic
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 1.44 µΩ·m
Thermal conductivity (300 K) 7.81 W·m−1·K−1
Thermal expansion (25 °C) 21.7 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 5150 m/s
Young's modulus 198 GPa
Bulk modulus 120 GPa
Mohs hardness 6.0
Brinell hardness 196 MPa
CAS registry number 7439-96-5
Most stable isotopes
Main article: Isotopes of manganese
iso NA half-life DM DE (MeV) DP
52Mn syn 5.591 d ε - 52Cr
β+ 0.575 52Cr
γ 0.7, 0.9, 1.4 -
53Mn trace 3.74 ×106 y ε - 53Cr
54Mn syn 312.3 d ε 1.377 54Cr
γ 0.834 -
55Mn 100% 55Mn is stable with 30 neutrons

Manganese (pronounced /ˈmæŋɡəniːz/, MANG-gən-neez) is a chemical element, designated by the symbol Mn. It has the atomic number 25. It is found as a free element in nature (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 phosphating is used as a treatment for rust and corrosion prevention on steel. 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 oxidizers. 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.

Contents

Characteristics

Physical

Manganese is a gray–white metal, resembling iron. It is a hard metal and is very brittle, fusible with difficulty, but easily oxidized.[1] Manganese metal and its common ions are paramagnetic.[2]

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.[3]

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 53Cr with 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 of cosmic rays on iron in rocks [4]. 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.[3]

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.[3]

Chemical

Oxidation states
of manganese[note 1][5]
0 Mn2(CO)10
+1 K5[Mn(CN)6NO]
+2 MnCl2
+3 MnF3
+4 MnO2
+5 Na3MnO4
+6 K2MnO4
+7 KMnO4
Mineral rhodochrosite (manganese(II) carbonate)
Manganese(II) chloride
Aqueous solution of KMnO4

The most common oxidation states of manganese are +2, +3, +4, +6 and +7, though oxidation states from -3 to +7 are observed. Mn2+ often competes with Mg2+ in biological systems. Manganese compounds where manganese is in oxidation state +7, which are restricted to the oxide Mn2O7 and compounds of the intensely purple permanganate anion MnO4, are powerful oxidizing agents.[1] 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 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.

The +3 oxidation state is known, in compounds such as manganese(III) acetate, but these are quite powerful oxidizing agents and also disproportionate in solution to Mn(II) and Mn(IV). Solid compounds of Mn(III) are characterized by its preference for distorted octahedral coordination due to the Jahn-Teller effect and its strong purple-red color.

The oxidation state 5+ can be obtained if manganese dioxide is dissolved in molten sodium nitrite.[6] Manganate (VI) salts can also be produced by dissolving Mn compounds in alkaline melts in air.

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.[7]

History

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, but were thought to differ in gender. The male magnes attracted iron, and was the iron ore we now know 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 feminine magnes was later called magnesia, known now in modern times as pyrolusite or manganese dioxide. This mineral is never magnetic (although manganese itself is paramagnetic). In the 16th century, the latter compound was called manganesum (note the two n's instead of one) by glassmakers, possibly as a corruption 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 name magnesia 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.[8]

Some of the cave painting in Lascaux, France use manganese-based pigments.[9]

Several oxides of manganese, for example manganese dioxide, are abundant in nature and due to 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.[10]

Manganese compounds were used by Egyptian and Roman glassmakers, to either remove color from glass or add color to it.[11] The use as glassmakers soap continued through the middle ages until modern times and is evident in 14th century glass from Venice.[12]

Credit for first isolating of manganese is usually given to Johan Gottlieb Gahn

Due to 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 to permanganate, a useful laboratory reagent.[13] 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 reacted 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.[14] 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.[15] 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.[16]

With 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 of the batteries on the market contained manganese. The Zinc-carbon battery and the alkaline 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 (carbon–zinc) and alkaline type.[17]

Occurrence and production

Manganese makes up about 1000 ppm (0.1%) of the Earth's crust, making it the 12th most abundant element there.[18] Soil contains 7–9000 ppm of manganese with an average of 440 ppm.[18] Seawater has only 10 ppm manganese and the atmosphere contains 0.01 µg/m3.[18] Manganese occurs principally as pyrolusite (MnO2), braunite, (Mn2+Mn3+6)(SiO12),[19] psilomelane (Ba,H2O)2Mn5O10, and to a lesser extent as rhodochrosite (MnCO3).

Percentage of manganese output in 2006 by countries[20]

The most important manganese ore is pyrolusite (MnO2). Other economically important manganese ores usually show a close spatial relation to the iron ores.[1] 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 Australia, India, China, Gabon and Brazil.[20] In 1978 it was estimated that 500 billion tons of manganese nodules exist on the ocean floor.[21] Attempts to find economically viable methods of harvesting manganese nodules were abandoned in the 1970s.[22]

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%.[20][23]

For the production of ferromanganese the manganese ore are mixed with iron ore and carbon and then reduced either in a blast furnace or in an electric arc furnace.[24] The resulting ferromanganese has a manganese content of 30 to 80%.[1] Pure manganese used for the production of non-iron alloys is produced by leaching manganese ore with sulfuric acid and a subsequent electrowinning process.[25]

Applications

Manganese has no satisfactory substitute in its major applications, which are related to metallurgical alloy use.[20] 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.

Steel

British Brodie helmet

Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties. Steelmaking,[26] including its ironmaking component, has accounted for most manganese demand, presently in the range of 85% to 90% of the total demand.[25] Among a variety of other uses, manganese is a key component of low-cost stainless steel formulations.[23][27]

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%. The fact that steel containing 8 to 15% of manganese is cold hardening and can obtain a high tensile strength of up to 863  MPa,[28][29] steel with 12% manganese was used for the British steel helmets. This steel composition was discovered in 1882 by Robert Hadfield and is still known as Hadfield steel.[30]

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.[31] The corrosion resistant aluminium alloy 3004 and 3104 with a manganese content of 0.8 to 1.5% are the alloy used for most of the beverage cans.[32] Before year 2000, in excess of 1.6 million metric tons have been used of those alloys, with a content of 1% of manganese this amount would need 16,000 metric tons of manganese.[32]

Other use

Wartime nickel made from a copper-silver-manganese alloy

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.[33]

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.[12] 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.[34]

MnO2 + H2O + e → MnO(OH) + OH

The same material also functions in newer 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.[17][34]

The metal is very occasionally used in coins; the only United States coins to use manganese were the "wartime" nickel from 1942–1945.[35] Due to shortage of raw materials in the war the nickel in the alloy (75% copper and 25% nickel) used for the production of the nickel before was substituted by the less critical metals silver and manganese (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.[36]

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.[37] In the glass industry manganese compounds are used for two effects. Manganese(III) reacts with iron(II). The reaction induces a strong green color in glass by forming less-colored iron(III) and slightly pink manganese(II), compensating the residual color of the iron(III).[12] 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.

Manganese is an essential trace nutrient in all forms of life.[18] The classes of enzymes that have manganese cofactors are very broad and include such classes as 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).[38]

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 non-enzymatic 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 10 mg of manganese, which is stored mainly in the liver and kidneys. In the human brain the manganese is bound to manganese metalloproteins most notably glutamine synthetase in astrocytes.[39]

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.[40] 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.[41] 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.[42] 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.[43]

The permanganate exhibits a higher toxicity than the manganese(II) compounds. Several fatal intoxications have occurred, although the fatal dose is around 10 g. 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.[44][45]

In 2005, a study suggested a possible link between manganese inhalation and central nervous system toxicity in rats.[46] It is hypothesized that long-term exposure to the naturally occurring manganese in shower water puts up to 8.7 million Americans at risk.[46][47][48]

A form of neurodegeneration[49] similar to Parkinson's Disease called "manganism" has been linked to manganese exposure amongst miners and smelters since the early 19th century.[50] Allegations of inhalation-induced manganism have been made regarding the welding industry. Manganese exposure in United States is regulated by Occupational Safety and Health Administration.[51]

See also

Notes

  1. ^ Common oxidation states are in bold.

References

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  36. ^ Design of the Sacagawea dollar. United States Mint. http://www.usmint.gov/mint_programs/golden_dollar_coin/index.cfm?action=sacDesign. Retrieved 2009-05-04.  
  37. ^ Shepard, Anna Osler (1956). "Manganese and Iron–Manganese Paints". Ceramics for the archaeologist. Carnegie Institution of Washington. pp. 40–42. ISBN 9780872796201.  
  38. ^ Law, N. (1998). Manganese Redox Enzymes and Model Systems: Properties, Structures, and Reactivity. 46. p. 305. doi:10.1016/S0898-8838(08)60152-X.  
  39. ^ Takeda, A. (2003). "Manganese action in brain function". Brain Research Reviews 41: 79. doi:10.1016/S0165-0173(02)00234-5.  
  40. ^ Dismukes, G. Charles; Willigen, Rogier T. van (2006). "Manganese: The Oxygen-Evolving Complex & Models". Encyclopedia of Inorganic Chemistry. doi:10.1002/0470862106.ia128.  
  41. ^ Hasan, Heather (2008). Manganese. The Rosen Publishing Group. pp. 31. ISBN 9781404214088. http://books.google.com/books?id=nRmpEaudmTYC&pg=PA31.  
  42. ^ "Manganese Chemical Background". Metcalf Institute for Marine and Environmental Reporting University of Rhode Island. 2006-04. http://www.environmentwriter.org/resources/backissues/chemicals/manganese.htm. Retrieved 2008-04-30.  
  43. ^ "Risk Assessment Information System Toxicity Summary for Manganese". Oak Ridge National Laboratory. http://rais.ornl.gov/tox/profiles/mn.shtml. Retrieved 2008-04-23.  
  44. ^ Ong, K. L.; Tan; Cheung (1997). "Potassium permanganate poisoning--a rare cause of fatal self poisoning.". Emergency Medicine Journal 14 (1): 43. doi:10.1136/emj.14.1.43. PMID 9023625.  
  45. ^ Young, R.; Critchley; Young; Freebairn; Reynolds; Lolin (1996). "Fatal acute hepatorenal failure following potassium permanganate ingestion". Human & Experimental Toxicology 15 (3): 259. doi:10.1177/096032719601500313. PMID 8839216.  
  46. ^ a b Elsner, RJ; Spangler, JG (2005). "Neurotoxicity of inhaled manganese: Public health danger in the shower?". Medical Hypotheses 65 (3): 607–616. doi:10.1016/j.mehy.2005.01.043. PMID 15913899.  
  47. ^ Finley, John Weldon (1999). "Manganese deficiency and toxicity: Are high or low dietary amounts of manganese cause for concern?". BioFactors 10: 15. doi:10.1002/biof.5520100102.  
  48. ^ Barceloux, Donald (1999). "Manganese". Clinical Toxicology 37: 293. doi:10.1081/CLT-100102427.  
  49. ^ Normandin, Louise (2002). Metabolic Brain Disease 17: 375. doi:10.1023/A:1021970120965.  
  50. ^ Crossgrove, J; Zheng, W (2004). "Manganese toxicity upon overexposure.". NMR in biomedicine 17 (8): 544–553. doi:10.1002/nbm.931. ISSN 0952-3480. PMID 15617053.  
  51. ^ "Safety and Health Topics: Manganese Compounds (as Mn)". http://www.osha.gov/dts/chemicalsampling/data/CH_250190.html.  

External links


chromiummanganeseiron
-

Mn

Tc
25Mn
Appearance
silvery metallic
File:Mangan 1-crop.jpg
General properties
Name, symbol, number manganese, Mn, 25
Pronunciation /ˈmæŋɡənz/ MANG-gən-neez
Element category transition metal
Group, period, block 74, d
Standard atomic weight 54.938045(5)g·mol−1
Electron configuration [Ar] 4s2 3d5
Electrons per shell 2, 8, 13, 2 (Image)
Physical properties
Phase solid
Density (near r.t.) 7.21 g·cm−3
Liquid density at m.p. 5.95 g·cm−3
Melting point 1519 K, 1246 °C, 2275 °F
Boiling point 2334 K, 2061 °C, 3742 °F
Heat of fusion 12.91 kJ·mol−1
Heat of vaporization 221 kJ·mol−1
Specific heat capacity (25 °C) 26.32 J·mol−1·K−1
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 1228 1347 1493 1691 1955 2333
Atomic properties
Oxidation states 7, 6, 5, 4, 3, 2, 1, -1, -2, -3
(oxides: acidic, basic or amphoteric
depending on the oxidation state)
Electronegativity 1.55 (Pauling scale)
Ionization energies
(more)
1st: 717.3 kJ·mol−1
2nd: 1509.0 kJ·mol−1
3rd: 3248 kJ·mol−1
Atomic radius 127 pm
Covalent radius 139±5 (low spin), 161±8 (high spin) pm
Miscellanea
Crystal structure body-centered cubic
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 1.44 µΩ·m
Thermal conductivity (300 K) 7.81 W·m−1·K−1
Thermal expansion (25 °C) 21.7 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 5150 m/s
Young's modulus 198 GPa
Bulk modulus 120 GPa
Mohs hardness 6.0
Brinell hardness 196 MPa
CAS registry number 7439-96-5
Most stable isotopes
Main article: Isotopes of manganese
iso NA half-life DM DE (MeV) DP
52Mn syn 5.591 d ε - 52Cr
β+ 0.575 52Cr
γ 0.7, 0.9, 1.4 -
53Mn trace 3.74 ×106 y ε - 53Cr
54Mn syn 312.3 d ε 1.377 54Cr
γ 0.834 -
55Mn 100% 55Mn is stable with 30 neutrons

Manganese ( /ˈmæŋɡənz/, MANG-gən-neez) is a chemical element, designated by the symbol Mn. It has the atomic number 25. It is found as a free element in nature (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 phosphating is used as a treatment for rust and corrosion prevention on steel. Depending on their oxidation state, manganese ions have various colors and are used industrially as pigments. The permanganates of alkali and alkaline earth metals are powerful oxidizers. 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.

Contents

Characteristics

Physical properties

Manganese is a silvery-gray metal resembling iron. It is hard and very brittle, difficult to fuse, but easy to oxidize.[1] Manganese metal and its common ions are paramagnetic.[2]

Isotopes

Naturally occurring manganese is composed of 1 stable isotope, 55Mn. Eighteen 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.[3]

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 53Cr with 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 of cosmic rays on iron in rocks.[4] 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 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.[3]

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.[3]

Chemical properties

Oxidation states
of manganese[note 1][5]
0 Mn2(CO)10
+1 K5[Mn(CN)6NO]
+2 MnCl2
+3 MnF3
+4 MnO2
+5 Na3MnO4
+6 K2MnO4
+7 KMnO4

The most common oxidation states of manganese are +2, +3, +4, +6 and +7, though oxidation states from -3 to +7 are observed. 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.[1] 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 state is the state used in living organisms for essential functions; other states are toxic for the human body. 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.[6]

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 effect and its strong purple-red color.

The oxidation state 5+ can be obtained if manganese dioxide is dissolved in molten sodium nitrite.[7] 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.[8]

Mineral rhodochrosite (manganese(II) carbonate). The red color is due to impurities.  
Manganese(II) chloride  
Aqueous solution of KMnO4  

History

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, but were thought to differ in gender. The male magnes attracted iron, and was the iron ore we now know 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 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 name magnesia 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.[9]

[[File:|thumb| left |alt=A drawing of a left-facing bull, in black, on a cave wall |Some of the cave painting in Lascaux, France use manganese-based pigments.[10]]]

Several oxides of manganese, for example manganese dioxide, are abundant in nature and due to 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.[11]

Manganese compounds were used by Egyptian and Roman glassmakers, to either remove color from glass or add color to it.[12] The use as glassmakers soap continued through the middle ages until modern times and is evident in 14th century glass from Venice.[13]

File:Gahn Johan
Credit for first isolating manganese is usually given to Johan Gottlieb Gahn

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 to permanganate, a useful laboratory reagent.[14] 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 reacted 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.[15] 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.[16] 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.[17]

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 the alkaline 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.[18]

Occurrence and production

Manganese makes up about 1000 ppm (0.1%) of the Earth's crust, making it the 12th most abundant element there.[19] Soil contains 7–9000 ppm of manganese with an average of 440 ppm.[19] Seawater has only 10 ppm manganese and the atmosphere contains 0.01 µg/m3.[19] Manganese occurs principally as pyrolusite (MnO2), braunite, (Mn2+Mn3+6)(SiO12),[20] psilomelane (Ba,H2O)2Mn5O10, and to a lesser extent as rhodochrosite (MnCO3).

[[File:|center|border|175x120px|alt=|Manganese ore ]]
Manganese ore  
Psilomelane (manganese ore)  
[[File:|center|border|175x120px|alt=|Spiegeleisen is an iron alloy with a manganese content of approximately 15% ]]
Spiegeleisen is an iron alloy with a manganese content of approximately 15%  
[[File:|center|border|175x120px|alt=|Manganese oxide dendrites on a limestone bedding plane from Solnhofen, Germany—a kind of pseudofossil. Scale is in mm ]]
Manganese oxide dendrites on a limestone bedding plane from Solnhofen, Germany—a kind of pseudofossil. Scale is in mm  
File:World Manganese Production
Percentage of manganese output in 2006 by countries[21]

The most important manganese ore is pyrolusite (MnO2). Other economically important manganese ores usually show a close spatial relation to the iron ores.[1] 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.[21] In 1978 it was estimated that 500 billion tons of manganese nodules exist on the ocean floor.[22] Attempts to find economically viable methods of harvesting manganese nodules were abandoned in the 1970s.[23]

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%.[21][24]

For the production of ferromanganese, the manganese ore are mixed with iron ore and carbon and then reduced either in a blast furnace or in an electric arc furnace.[25] The resulting ferromanganese has a manganese content of 30 to 80%.[1] Pure manganese used for the production of non-iron alloys is produced by leaching manganese ore with sulfuric acid and a subsequent electrowinning process.[26]

Unexploited deposits includes Tamboa in Burkina Faso.

Applications

Manganese has no satisfactory substitute in its major applications, which are related to metallurgical alloy use.[21] 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.

Steel

[[File:| thumb|right|US Marine Corps steel helmet]] Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties. Steelmaking,[27] including its ironmaking component, has accounted for most manganese demand, presently in the range of 85% to 90% of the total demand.[26] Among a variety of other uses, manganese is a key component of low-cost stainless steel formulations.[24][28]

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 is cold hardening and can obtain a high tensile strength of up to 863 MPa.[29][30] Steel with 12% manganese was used for the British steel helmets. This steel composition was discovered in 1882 by Robert Hadfield and is still known as Hadfield steel.[31]

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.[32] The corrosion resistant aluminium alloy 3004 and 3104 with a manganese content of 0.8 to 1.5% are the alloy used for most of the beverage cans.[33] Before year 2000, in excess of 1.6 million metric tons have been used of those alloys, with a content of 1% of manganese this amount would need 16,000 metric tons of manganese.[33]

Other uses

File:War
World War II-time nickel made from a copper-silver-manganese alloy

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.[34]

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.[13] 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.[35]

MnO2 + H2O + e → MnO(OH) + OH

The same material also functions in newer 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.[18][35]

The metal is very occasionally used in coins; until 2000 the only United States coin to use manganese was the "wartime" nickel from 1942–1945.[36] 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.[37]

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.[38] In the glass industry manganese compounds are used for two effects. Manganese(III) reacts with iron(II). The reaction induces a strong green color in glass by forming less-colored iron(III) and slightly pink manganese(II), compensating the residual color of the iron(III).[13] Larger amounts of manganese are used to produce pink colored glass.

Biological role

[[File:|thumb|right|400px|Reactive center of arginase with boronic acid inhibitor. The manganese atoms are shown in yellow.]] Manganese is an essential trace nutrient in all forms of life.[19] 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).[39]

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 non-enzymatic 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 10 mg of manganese, which is stored mainly in the liver and kidneys. In the human brain the manganese is bound to manganese metalloproteins most notably glutamine synthetase in astrocytes.[40]

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.[41] 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.[42] 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.[43] Manganese poisoning has been linked to impaired motor skills and cognitive disorders.[44]

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.[45][46]

In 2005, a study suggested a possible link between manganese inhalation and central nervous system toxicity in rats.[47] It is hypothesized that long-term exposure to the naturally occurring manganese in shower water puts up to 8.7 million Americans at risk.[47][48][49]

A form of neurodegeneration[50] similar to Parkinson's Disease called "manganism" has been linked to manganese exposure amongst miners and smelters since the early 19th century.[51] Allegations of inhalation-induced manganism have been made regarding the welding industry. Manganese exposure in United States is regulated by Occupational Safety and Health Administration.[52]

Clinical toxicity

Manganism has occurred in persons employed in the production or processing of manganese alloys, patients receiving total parenteral nutrition, workers exposed to manganese-containing fungicides such as maneb, and abusers of drugs such as methcathinone made with potassium permanganate. Excessive exposure may be confirmed by measurement of blood or urine manganese concentrations.[53]

Chronic exposure to excessive Mn levels can lead to a variety of psychiatric and motor disturbances, termed manganism. 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 (LDH) leakage.[54]

See also

Notes

  1. ^ Common oxidation states are in bold.

References

  1. ^ a b c d Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985). "Mangan" (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1110–1117. ISBN 3-11-007511-3. 
  2. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics. CRC press. 2004. ISBN 0849304857. http://www-d0.fnal.gov/hardware/cal/lvps_info/engineering/elementmagn.pdf. 
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  5. ^ Schmidt, Max (1968). "VII. Nebengruppe" (in German). Anorganische Chemie II.. Wissenschaftsverlag. pp. 100–109. 
  6. ^ Descriptive Organic Chemistry Geoffrey Rayner-Canham, Tina Overton, Macmillan, 2003. p. 491
  7. ^ Temple, R. B.; Thickett, G. W. (1972). "The formation of manganese(v) in molten sodium nitrite". Australian Journal of Chemistry 25: 55. http://www.publish.csiro.au/?act=view_file&file_id=CH9720655.pdf. 
  8. ^ Luft, J. H. (1956). [Expression error: Unexpected < operator "Permanganate – a new fixative for electron microscopy"]. Journal of Biophysical and Biochemical Cytology 2 (6): 799–802. doi:10.1083/jcb.2.6.799. PMID 13398447. 
  9. ^ Calvert, J.B. (2003-01-24). "Chromium and Manganese". http://www.du.edu/~jcalvert/phys/chromang.htm. Retrieved 2009-04-30. 
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  11. ^ Chalmin, Y; Osuga, Y; Harada, M; Hirata, T; Koga, K; Morimoto, C; Hirota, Y; Yoshino, O et al.; Vignaud, C.; Salomon, H.; Farges, F.; Susini, J.; Menu, M. (2006). [Expression error: Unexpected < operator "Minerals discovered in paleolithic black pigments by transmission electron microscopy and micro-X-ray absorption near-edge structure"]. Applied Physics A 83 (12): 213–218. doi:10.1007/s00339-006-3510-7. PMID 16055459. 
  12. ^ Sayre, E. V.; Smith, R. W. (1961). [Expression error: Unexpected < operator "Compositional Categories of Ancient Glass."]. Science 133 (3467): 1824–1826. doi:10.1126/science.133.3467.1824. PMID 17818999. 
  13. ^ a b c Mccray, W. Patrick (1998). [Expression error: Unexpected < operator "Glassmaking in renaissance Italy: The innovation of venetian cristallo"]. Journal of the Minerals, Metals and Materials Society 50: 14. doi:10.1007/s11837-998-0024-0. 
  14. ^ Rancke-Madsen, E. (1975). [Expression error: Unexpected < operator "The Discovery of an Element"]. Centaurus 19 (4): 299–313. doi:10.1111/j.1600-0498.1975.tb00329.x. 
  15. ^ Alessio, L; Campagna, M; Lucchini, R (2007). [Expression error: Unexpected < operator "From lead to manganese through mercury: mythology, science, and lessons for prevention."]. American journal of industrial medicine 50 (11): 779–787. doi:10.1002/ajim.20524. PMID 17918211. 
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  54. ^ Yin, Z; Jiang, H; Lee, ES; Ni, M; Erikson, KM; Milatovic, D; Bowman, AB; Aschner, M (2010). "Ferroportin is a manganese-responsive protein that decreases manganese cytotoxicity and accumulation.". Journal of neurochemistry 112 (5): 1190–8. doi:10.1111/j.1471-4159.2009.06534.x. PMID 20002294. PMC 2819584. http://libres.uncg.edu/ir/uncg/f/K_Erickson_Ferroportin_2009.pdf. , and also: Cotzias et al. 1968; Olanow 2004; Aschner et al. 2007; Ellingsen et al. 2008

External links



1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

MANGANESE [[[symbol]] Mn; atomic weight, 54.93 (0 = t6)], a metallic chemical element. Its dioxide (pyrolusite) has been known from very early times, and was at first mistaken for a magnetic oxide of iron. In 1740 J. H. Pott showed that it did not contain iron and that it yielded a definite series of salts, whilst in 1774 C. Scheele proved that it was the oxide of a distinctive metal. Manganese is found widely distributed in nature, being generally found to a greater or less extent associated with the carbonates and silicates of iron, calcium and magnesium, and also as the minerals braunite, hausmannite, psilomelane, manganite, manganese spar and hauerite. It has also been recognized in the atmosphere of the sun (A. Cornu, Cornptes rendus, 1878, 86, pp. 3 1 5, 53 o), in sea water, and in many mineral waters.

The metal was isolated by J. G. Gahn in 1774, and in 1807 J. F. John (Gehlen's Jour. chem. phys., 1807, 3, p. 452) obtained an impure metal by reducing the carbonate at a high temperature with charcoal, mixed with a small quantity of oil. R. Bunsen prepared the metal by electrolysing manganese chloride in a porous cell surrounded by a carbon crucible containing hydrochloric acid. Various reduction methods have been employed for the isolation of the metal. C. Brunner (Pogg. Ann., 1857, 101, p. 264) reduced the fluoride by metallic sodium, and E. Glatzel (Ber., 1889, 22, p. 2857) the chloride by magnesium, H. Moissan (Ann. Claim. Phys., 1896 (7) 9, p. 286) reduced the oxide with carbon in the electric furnace; and H. Goldschmidt has prepared the metal from the oxide by means of his "thermite" process (see Chromium). W. H. Green and W. H. Wahl [German patent 70773 (1893)] prepare a 97% manganese from pyrolusite by heating it with 30% sulphuric acid, the product being then converted into manganous oxide by heating in a current of reducing gas at a dull red heat, cooled in a reducing atmosphere, and finally reduced by heating with granulated aluminium in a magnesia crucible with lime and fluorspar as a flux. A purer metal is obtained by reducing manganese amalgam by hydrogen (0. Prelinger, Monats., 18 94, 14, P 353).

Prelinger's manganese has a specific gravity of 7.42, and the variety obtained by distilling pure manganese amalgam in vacuo is pyrophoric (A. Guntz, Bull. Soc. [3], 7, 275), and burns when heated in a current of sulphur dioxide. The pure metal readily evolves hydrogen when acted upon by sulphuric and hydrochloric acids, and is readily attacked by dilute nitric acid. It precipitates many metals from solutions of their salts. It is employed commercially in the manufacture of special steels. (See Iron And Steel.) Compounds Manganese forms several oxides, the most important of which are manganous oxide, MnO, trimanganese tetroxide, Mn304, manganese sesquioxide, Mn203, manganese dioxide, Mn02, manganese trioxide, Mn03, and manganese heptoxide, Mn207.

Manganous oxide, MnO, is obtained by heating a mixture of anhydrous manganese chloride and sodium carbonate with a small quantity of ammonium chloride (J. v. Liebig and F. Wohler, Pogg. Ann., 1830, 21, p. 584); or by reducing the higher oxides with hydrogen or carbon monoxide. It is a dark coloured powder of specific gravity 5.09. Manganous hydroxide, Mn (OH) 2, is obtained as a white precipitate on adding a solution of a caustic alkali to a manganous salt. For the preparation of the crystalline variety identical with the mineral pyrochroite (see A. de Schulten, Comptes rend us, 1887, 105, p. 1265). It rapidly oxidizes on exposure to air and turns brown, going ultimately to the sesquioxide. Trimanganese tetroxide, Mn304, is produced more or less pure when the other oxides are heated. It may be obtained crystalline by heating manganese sulphate and potassium sulphate to a bright red heat (H. Debray, Comptes rendus, 1861, 52, p. 985). It is a reddish-brown powder, which when heated with hydrochloric acid yields chlorine. Manganese sesquioxide, Mn203, found native as the mineral braunite, may be obtained by igniting the other oxides in a mixture of nitrogen and oxygen, containing not more than .26% of the latter gas (W. Dittmar, Jour. Chem. Soc., 1864, 17, p.. 294). The hydrated form, found native as the mineral manganite, is produced by the spontaneous oxidation of manganous h y droxide. In the hydrated condition it is a dark brown powder which readily loses water at above too° C., it dissolves in hot nitric acid, giving manganous nitrate and manganese dioxide: 2MnO(OH) + 2HNO 3 = Mn(NO 3) 2 + MnO 2 + 2H 2 0. Manganese dioxide, or pyrolusite (q.v.), Mn02, the most important oxide, may be prepared by heating crystallized manganous nitrate until red fumes are given off, decanting the clear liquid, and heating to 150 0 to 160° C. for 40 to 60 hours (A. Gorgen, Bull. Soc., 1890 [3], 4, p. 16), or by heating manganese carbonate to 260° C. in the presence of air and washing the residue with very dilute cold hydrochloric acid. It is a hard black solid which readily loses oxygen when strongly heated, leaving a residue of Mn 3 0 4. When heated with concentrated hydrochloric acid it yields chlorine, and with concentrated sulphuric acid it yields oxygen. It is reduced to the monoxide when heated in a current of hydrogen. It is a strong oxidizing agent. It dissolves in cold concentrated hydrochloric acid, forming a dark brown solution which probably contains manganic chloride (see R. J. Meyer, Zeit. anorg. Chem., 1899, 22, p. 169; G. Neumann, Monats., 18 94, 1 5, p. 489). It is almost impossible to prepare a pure hydrated manganese dioxide owing to the readiness with which it loses oxygen, leaving residues of the type xMnO yMn0 2. Such mixtures are obtained by the action of alkaline hypochlorites on manganous salts, or by suspending manganous carbonate in water and passing chlorine through the mixture. The solid matter is filtered off, washed with water, and warmed with 10% nitric acid (A. Gorgen). It is a dark brown powder, which reddens litmus. Manganese dioxide combines with other basic oxides to form manganites, and on this property is based the Weldon process for the recovery of manganese from the waste liquors of the chlorine stills (see Chlorine). The manganites are amorphous brown solids, insoluble in water, and decomposed by hydrochloric acid with the evolution of chlorine. Manganese trioxide, Mn03, is obtained in small quantity as an unstable deliquescent red solid by dropping a solution of potassium permanganate in sulphuric acid on to dry sodium carbonate (B. Franke, Jour. prak. Chem., 1887 [2], 36, p. 31). Above 50° C. it decomposes into the dioxide and oxygen. It dissolves in water forming manganic acid, H 2 Mn0 4. Manganese heptoxide, Mn 2 0 7, prepared by adding pure potassium permanganate to well cooled, concentrated sulphuric acid, when the oxide separates as a dark oil (H. Aschoff, Pogg. Ann., 1860, 111, p. 217), is very unstable, continually giving off oxygen. It decomposes violently on heating, and explodes in contact with hydrogen, sulphur, phosphorus, &c. It dissolves in water to form a deep red solution which contains permanganic acid, HMnO 4. This acid is also formed by decomposing barium or lead permanganate with dilute sulphuric acid. It is only known in aqueous solution. This solution is of a deep violetred colour, and is somewhat fluorescent; it decomposes on exposure to light, or when heated. It is a monobasic acid, and a very powerful oxidizing agent (M. M. P. Muir, Jour. Chem. Soc., 1907, 91, p. 1485).

Table of contents

Manganous Salts

The anhydrous chloride, MnCl2, is obtained as a rose-red crystalline solid by passing hydrochloric acid gas over manganese carbonate, first in the cold and afterwards at a moderate red heat. The hydrated chloride, MnCl2.4H2O, is obtained in rose-red crystals by dissolving the metal or its carbonate in aqueous hydrochloric acid and concentrating the solution. It may be obtained in at least two different forms, one isomorphous with NaCl. 2H 2 O, by concentrating the solution between 15° C. and 20°C.; the other, isomorphous with FeCl 2.4H 2 0, by slow evaporation of the mother liquors from the former. It forms double salts with the chlorides of the alkali metals. The bromide MnBr2.4H20, iodide, Mn12, and fluoride, MnF2, are known.

Manganous Sulphate, MnSO 4, is prepared by strongly heating a paste of pyrolusite and concentrated sulphuric acid until acid fumes cease to be evolved. The ferric and aluminium sulphates present are thus converted into insoluble basic salts, and the residue yields manganous sulphate when extracted with water. The salt crystallizes with varying quantities of water, according to the temperature at which crystallization is effected: between - 4° C. and +6° C. with 7H 2 O, between 15° C. and 20° C. with 5H 2 O, and between 25° C. and 31° C, with 4H 2 O. It crystallizes in large pink crystals, the colour of which is probably due to the presence of a small quantity of manganic sulphate or of a cobalt sulphate. It combines with the sulphates of the alkali metals to form double salts.

Manganous Nitrate, Mn(NO 3) 2.6H 2 0, obtained by dissolving the carbonate in nitric acid and concentrating the solution, crystallizes from nitric acid solutions in long colourless needles, which melt at 25.8° C. and boil at 129.5° C. with some decomposition.

Manganous Carbonate, MnCO 3, found native as manganese spar, may be prepared as an amorphous powder by heating manganese chloride with sodium carbonate in a sealed tube to 150° C., or in the hydrated form as a white flocculent precipitate by adding sodium carbonate to a manganous salt. In the moist condition it rapidly turns brown on exposure to air.

Manganous Sulphide, MnS, found native as manganese glance, may be obtained by heating the monoxide or carbonate in a porcelain tube in a current of carbon bisulphide vapour. R. Schneider (Pogg. Ann., 18 74, 1 5 1, 449) obtained a crystalline variety by melting sulphur with anhydrous manganous sulphate and dry potassium carbonate, extracting the residue and drying it in a current of hydrogen. Four sulphides are known; the red and green are anhydrous, a grey variety contains much water, whilst the pink is a mixture of the grey and red (J. C. Olsen and W. S. Rapalje, Jour. Amer. Chem. Soc., 1904, 26, p. 1615). Ammonium sulphide alone gives incomplete precipitation of the sulphide. In the presence of ammonium salts the precipitate is dirt y white in colour, whilst in the presence of free ammonia it is a buff colour. This form of the sulphide is readily oxidized when exposed in the moist condition, and is easily decomposed by dilute mineral acids. Manganese Disulphide, MnS2, found native as hauerite, is formed as a red coloured powder by heating manganous sulphate with potassium polysulphide in a sealed tube at 160°-170° C. (H. v. Senarmont, Jour. prak. Chem., 18 5 0, 5 1, p. 385).

Manganic Salts

The sulphate, Mn2(S04)3, is prepared by gradually heating at 138° C. a mixture of concentrated sulphuric and manganese dioxide until the whole becomes of a dark green colour. The excess of acid is removed by spreading the mass on a porous plate, the residue stirred for some hours with nitric acid, again spread on a porous plate, and finally dried quickly at about 130° C. It is a dark green deliquescent powder which decomposes on heating or on exposure to moist air. It is readily decomposed by dilute acids. With potassium sulphate in the presence of sulphuric acid it forms potassium manganese alum, K2S04 Mn2(S04)3.24H20. A. Piccini (Zeit. anorg. Chem. 1898, 1 7, p. 355) has also obtained a manganese caesium alum. Manganic Fluoride, MnF3, a solid obtained by the action of fluorine on manganous chloride, is decomposed by heat into manganous fluoride and fluorine. By suspending the dioxide in carbon tetrachloride and passing in hydrochloric acid gas, W. B. Holmes (Abst. J.C.S., 1907, ii., p. 873) obtained a black trichloride and a reddish-brown tetrachloride.

Manganese Carbide, Mn 3 C, is prepared by heating manganous oxide with sugar charcoal in an electric furnace, or by fusing manganese chloride and calcium carbide. Water decomposes it, giving methane and hydrogen (H. Moissan); Mn 3 C+6H 2 O = 3Mn(OH)2+CH4+H2.

Manganates

These salts are derived from manganic acid H 2 Mn0 4. Those of the alkali metals are prepared by fusing manganese dioxide with sodium or potassium hydroxide in the presence of air or of some oxidizing agent (nitre, potassium chlorate, &c.); MnO 2 +2KHO+O=K 2 Mn0 4 +H 2 O. In the absence of air the reaction proceeds slightly differently, some manganese sesquioxide being formed; 3MnO 2 +2KHO = K 2 Mn0 4 +Mn 2 0 3 +H 2 O. The fused mass has a dark olive-green colour, and dissolves in a small quantity of cold water to a green solution, which is, however, only stable in the presence of an excess of alkali. The green solution is readily converted into a pink one of permanganate by a large dilution with water, or by passing carbon dioxide through it: 3K2Mn04+2C02= 2K2C03+2KMn04+Mn02.

Permanganates are the salts of permanganic acid, HMnO 4. The potassium salt, KMnO 4, may be prepared by passing chlorine or carbon dioxide through an aqueous solution of potassium manganate, or by the electrolytic oxidation of the manganate at the anode [German patent 101710 (1898)]. It crystallizes in dark purple-red prisms, isomorphous with potassium perchlorate. It acts as a powerful oxidizing agent, both in acid and alkaline solution; in the first case two molecules yield five atoms of available oxygen and in the second, three atoms: 2KMnO 4 +3H 2 SO 4 = K2S04+2MnS04+3H20+50; 2KMnO 4 +3H 2 O =2Mn02 H20+2KHO+30.

It completely decomposes hydrogen peroxide in sulphuric acid solution 2KMn04+5H202-I-3H2S04 = K2S04+2MnS04+8H20+502. It decomposes when heated to 200° - 240°C.: 2KMn04=K2Mn04+Mn02+02; and when warmed with hydrochloric acid it yields chlorine: 2 KM nO 4 + 16HC1= 2KC1 +2 MnC1 2 +8H 2 0 +5C12.

Sodium Permanganate, NaMn0 4.3H 2 O (?), may be prepared in a similar manner, or by precipitating the silver salt with sodium chloride. It crystallizes with great difficulty. A solution of the crude salt is used as a disinfectant under the name of "Condy's fluid." Ammonium Permanganate, NH 4 Mn0 4, explodes violently on rubbing, and its aqueous solution decomposes on boiling (W. Muthmann, Ber., 1893, 26, p. 1018); NH4 Mn04=Mn02+N2+2H20.

Barium Permanganate, BaMn 2 0 8, crystallizes in almost black needles, and is formed by passing carbon dioxide through water containing suspended barium manganate.

Detection

Manganese salts can be detected by the amethyst colour they impart to a borax-bead when heated in the Bunsen flame, and by the green mass formed when they are fused with a mixture of sodium carbonate and potassium nitrate. Manganese may be estimated quantitatively by precipitation as carbonate, this salt being then converted into the oxide, Mn 3 0 4 by ignition; or by precipitation as hydrated dioxide by means of ammonia and bromine water, followed by ignition to NIn 3 0 4. The valuation of pyrolusite is generally carried out by means of a distillation with hydrochloric acid, the liberated chlorine passing through a solution of potassium iodide, and the amount of iodine liberated being ascertained by means of a standard solution of sodium thiosulphate. The atomic weight of manganese has been frequently determined.

J. Berzelius, by analysis of the chloride, obtained the value 54.86; K. v. Hauer (Sitzb. Akad. Wien., 1857, 25, p. 132), by conversion of the sulphate into sulphide, obtained the value 54.78; J. Dewar and A. Scott (Chem. News, 188 3, 47, p. 98), by analysis of silver permanganate, obtained the value 55.038; J. M. Weeren (Stahl. u. Eisen, 18 93, 1 3, p. 559), by conversion of manganous oxide into the sulphate obtained the value 54.883, and of the sulphate into sulphide the value 54.876 (H = 1), and finally G. P. Baxter and Hines (Jour. Amer. Chem. Soc., 1906, 28, p. 1360), by analyses of the chloride and bromide, obtained 54.96 (0 = 16).


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Simple English

[[File:|thumb|Manganese slab]] Manganese is chemical element 25 on the periodic table. It symbol is Mn. (Some people get it confused with magnesium, the symbol is Mg). It has 25 protons. Its mass number is 54.94.

Contents

Properties

Manganese is a silvery-gray metal and is part of the group known as the transition metals. It is similar to iron. It is hard to melt, but easy to oxidize. Manganese forms chemical compounds in several oxidation states: +2, +4, and +7 are the most common. +2 compounds are pink or light brown. Manganese(II) chloride is a common example. +4 compounds are black and rarer. Manganese(IV) oxide is an example. They are oxidizing agents. +7 compounds are purple-black and powerful oxidizing agents. Potassium permanganate is an example. Manganese compounds can be black, brown, pink, red, green, blue, and purple.

Occurrence

Manganese is sometimes found alloyed with iron naturally. These rocks, called meteorites, came from space. Pyrolusite is one of the main sources of manganese. It also occurs as manganese carbonate. Some silicates have manganese in them.

Preparation

Manganese is normally made in an alloy with steel. This is made by mixing manganese ore and iron ore in a furnace and reducing it with carbon. This forms an alloy called ferromanganese. Pure manganese is made by reacting the manganese ore with sulfuric acid and electrolyzing it.

Uses

Manganese is used a lot in steel to make it stronger. This is the main use for manganese metal. Manganese compounds, particularly manganese(IV) oxide, are used in alkaline cells and Leclanche cells. Manganese metal is also alloyed with aluminium.

Our bodies and plants need manganese to work right. If we do not get enough manganese, we can get sick. We get manganese from our food and vitamins also have some manganese to make sure that we get enough.

Safety

Manganese dust can irritate lungs. Some manganese compounds cause toxicity when ingested. Manganese is less toxic than nickel or copper. Permanganates are the most toxic manganese compounds. When someone is exposed to manganese for a long time it can cause a problem with the nervous system.

See also

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