Manganese

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Not to be confused with Magnesium.
For the ghost town in Minnesota, see Manganese, Minnesota.
chromium ← manganese → iron
-
↑
Mn
↓
Tc
Appearance
silvery metallic
250px
General properties
Name, symbol, number manganese, Mn, 25
Pronunciation /ˈmæŋɡəniːz/ MANG-gən-neez
Element category transition metal
Group, period, block 7, 4, 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 (11px /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. 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.

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

Main article: Isotopes of manganese

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.
Mineral rhodochrosite (manganese(II) carbonate). The red color is due to impurities.  
Manganese(II) chloride
Manganese(II) chloride  
Aqueous solution of KMnO4
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:Lascaux-aurochs.jpg
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 Gottlieb.jpg
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

See also: Category:Manganese minerals

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).

Manganese ore
Manganese ore  
Psilomelane (manganese ore)
Psilomelane (manganese ore)  
Spiegeleisen is an iron alloy with a manganese content of approximately 15%
Spiegeleisen is an iron alloy with a manganese content of approximately 15%  
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 2006.svg
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]

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:M1917helmet.jpg
British Brodie 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

Main article: Aluminium alloy

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 Nickle.jpg
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:Arginase.jpeg
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

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External links

 Periodic table
H   He
Li Be   B C N O F Ne
Na Mg   Al Si P S Cl Ar
K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo
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  1. ↑ 1.0 1.1 1.2 1.3 Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  2. ↑ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics (PDF). CRC press. 2004. ISBN 0849304857.  |first1= missing |last1= in Authors list (help)
  3. ↑ 3.0 3.1 3.2 Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A. Atomic Mass Data Center. 729: 3–128. doi:10.1016/j.nuclphysa.2003.11.001. 
  4. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  5. ↑ Schmidt, Max (1968). "VII. Nebengruppe". Anorganische Chemie II (in German). Wissenschaftsverlag. pp. 100–109. 
  6. ↑ Descriptive Organic Chemistry Geoffrey Rayner-Canham, Tina Overton, Macmillan, 2003. p. 491
  7. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  8. ↑ Luft, J. H. (1956). "Permanganate – a new fixative for electron microscopy". Journal of Biophysical and Biochemical Cytology. 2 (6): 799–802. doi:10.1083/jcb.2.6.799. PMC 2224005 Freely accessible. PMID 13398447. 
  9. ↑ Calvert, J.B. (2003-01-24). "Chromium and Manganese". Retrieved 2009-04-30. 
  10. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  11. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  12. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  13. ↑ 13.0 13.1 13.2 Mccray, W. Patrick (1998). "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). "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). "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.  More than one of |author1= and |last1= specified (help); More than one of |author2= and |last2= specified (help); More than one of |author3= and |last3= specified (help)
  16. ↑ Couper, J. (1837). "On the effects of black oxide of manganese when inhaled into the lungs". Br. Ann. Med. Pharmacol. 1: 41–42. 
  17. ↑ Olsen, Sverre E.; Tangstad, Merete; Lindstad, Tor (2007). "History of manganese". Production of Manganese Ferroalloys. Tapir Academic Press. pp. 11–12. ISBN 9788251921916. 
  18. ↑ 18.0 18.1 Preisler, Eberhard (1980). "Moderne Verfahren der Großchemie: Braunstein". Chemie in unserer Zeit (in German). 14: 137–148. doi:10.1002/ciuz.19800140502. 
  19. ↑ 19.0 19.1 19.2 19.3 Emsley, John (2001). "Manganese". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, UK: Oxford University Press. pp. 249–253. ISBN 0-19-850340-7. 
  20. ↑ Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  21. ↑ 21.0 21.1 21.2 21.3 Corathers, Lisa A. (2009). "Mineral Commodity Summaries 2009: Manganese" (PDF). United States Geological Survey. Retrieved 2009-04-30. 
  22. ↑ Wang, X; Schröder, Hc; Wiens, M; Schlossmacher, U; Müller, We (2009). "Manganese/polymetallic nodules: micro-structural characterization of exolithobiontic- and endolithobiontic microbial biofilms by scanning electron microscopy". Micron (Oxford, England : 1993). 40 (3): 350–358. doi:10.1016/j.micron.2008.10.005. ISSN 0968-4328. PMID 19027306. 
  23. ↑ United Nations Ocean Economics and Technology Office, Technology Branch, United Nations (1978). Manganese Nodules: Dimensions and Perspectives. Springer. ISBN 9789027705006. 
  24. ↑ 24.0 24.1 Corathers, Lisa A. (June 2008). "2006 Minerals Yearbook: Manganese" (PDF). Washington, D.C.: United States Geological Survey. Retrieved 2009-04-30. 
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  27. ↑ Verhoeven., John D. (2007). Steel metallurgy for the non-metallurgist. Materials Park, Ohio: ASM International. pp. 56–57. ISBN 9780871708588. 
  28. ↑ Dastur, Y. N.; Leslie, W. C. (1981). "Mechanism of work hardening in Hadfield manganese steel". Metallurgical Transactions A. 12: 749. doi:10.1007/BF02648339. 
  29. ↑ Stansbie, John Henry (2007). Iron and Steel. Read Books. pp. 351–352. ISBN 9781408626160. 
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  31. ↑ Tweedale, Geoffrey (1985). "Sir Robert Abbott Hadfield F.R.S. (1858-1940), and the Discovery of Manganese Steel Geoffrey Tweedale". Notes and Records of the Royal Society of London. 40 (1): 63–74. doi:10.1098/rsnr.1985.0004. 
  32. ↑ "chemical properties of 2024 aluminum allow". Metal Suppliers Online, LLC. Retrieved 2009-04-30. 
  33. ↑ 33.0 33.1 Kaufman, John Gilbert (2000). "Applications for Aluminium Alloys and Tempers". Introduction to aluminum alloys and tempers. ASM International. pp. 93–94. ISBN 9780871706898. 
  34. ↑ Graham, L. A.; et al. (2005). "Manganese(I) poly(mercaptoimidazolyl)borate complexes: spectroscopic and structural characterization of MnH–B interactions in solution and in the solid state". Dalton Trans (1): 171–180. doi:10.1039/b412280a. PMID 15605161. 
  35. ↑ 35.0 35.1 Dell, R. M. (2000). "Batteries fifty years of materials development". Solid State Ionics. 134: 139–158. doi:10.1016/S0167-2738(00)00722-0. 
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  37. ↑ "Design of the Sacagawea dollar". United States Mint. Retrieved 2009-05-04. 
  38. ↑ Shepard, Anna Osler (1956). "Manganese and Iron–Manganese Paints". Ceramics for the archaeologist. Carnegie Institution of Washington. pp. 40–42. ISBN 9780872796201. 
  39. ↑ Law, N.; Caudle, M; Pecoraro, V (1998). "Manganese Redox Enzymes and Model Systems: Properties, Structures, and Reactivity". 46: 305. doi:10.1016/S0898-8838(08)60152-X. 
  40. ↑ Takeda, A. (2003). "Manganese action in brain function". Brain Research Reviews. 41 (1): 79. doi:10.1016/S0165-0173(02)00234-5. PMID 12505649. 
  41. ↑ Dismukes, G. Charles; Willigen, Rogier T. van (2006). "Manganese: The Oxygen-Evolving Complex & Models". Encyclopedia of Inorganic Chemistry. doi:10.1002/0470862106.ia128. 
  42. ↑ Hasan, Heather (2008). Manganese. The Rosen Publishing Group. p. 31. ISBN 9781404214088. 
  43. ↑ "Manganese Chemical Background". Metcalf Institute for Marine and Environmental Reporting University of Rhode Island. 2006-04. Retrieved 2008-04-30.  Check date values in: |date= (help)
  44. ↑ "Risk Assessment Information System Toxicity Summary for Manganese". Oak Ridge National Laboratory. Retrieved 2008-04-23. 
  45. ↑ Ong, K. L.; Tan, TH; Cheung, WL (1997). "Potassium permanganate poisoning--a rare cause of fatal self poisoning". Emergency Medicine Journal. 14 (1): 43. doi:10.1136/emj.14.1.43. PMC 1342846 Freely accessible. PMID 9023625. 
  46. ↑ Young, R.; Critchley, JA; Young, KK; Freebairn, RC; Reynolds, AP; Lolin, YI (1996). "Fatal acute hepatorenal failure following potassium permanganate ingestion". Human & Experimental Toxicology. 15 (3): 259. doi:10.1177/096032719601500313. PMID 8839216. 
  47. ↑ 47.0 47.1 Lua error in package.lua at line 80: module 'Module:Citation/CS1/Suggestions' not found.
  48. ↑ Finley, John Weldon; Davis, Cindy D. (1999). "Manganese deficiency and toxicity: Are high or low dietary amounts of manganese cause for concern?". BioFactors. 10: 15. doi:10.1002/biof.5520100102. 
  49. ↑ Barceloux, Donald; Barceloux, Donald (1999). "Manganese". Clinical Toxicology. 37: 293. doi:10.1081/CLT-100102427. 
  50. ↑ Normandin, Louise; Hazell, AS (2002). "Manganese neurotoxicity: an update of pathophysiologic mechanisms". Metabolic Brain Disease. 17 (4): 375. doi:10.1023/A:1021970120965. PMID 12602514. 
  51. ↑ 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. 
  52. ↑ "Safety and Health Topics: Manganese Compounds (as Mn)". 
  53. ↑ R. Baselt, Disposition of Toxic Drugs and Chemicals in Man, 8th edition, Biomedical Publications, Foster City, CA, 2008, pp. 883-886.
  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" (PDF). Journal of neurochemistry. 112 (5): 1190–8. doi:10.1111/j.1471-4159.2009.06534.x. PMC 2819584 Freely accessible. PMID 20002294. , and also: Cotzias et al. 1968; Olanow 2004; Aschner et al. 2007; Ellingsen et al. 2008
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