Manganese dioxide
| Manganese dioxide | |
|---|---|
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| |
| IUPAC name | Manganese dioxide |
| Other names | Pyrolusite Manganese(IV) oxide |
| Identifiers | |
| InChI | InChI=1/Mn.2O/rMnO2/c2-1-3 |
| InChIKey | NUJOXMJBOLGQSY-MQJJFXMEAL |
| Standard InChI | InChI=1S/Mn.2O |
| Standard InChIKey | NUJOXMJBOLGQSY-UHFFFAOYSA-N |
| CAS number | [] |
| EC number | |
| RTECS | OP0350000 |
| PubChem | |
| Properties[1] | |
| Chemical formula | MnO2 |
| Molar mass | 86.937 g/mol |
| Appearance | black solid |
| Density | 5.026 g/cm3 |
| Melting point |
535 °C decomp. |
| Solubility in water | insoluble |
| Thermochemistry[2] | |
| Std enthalpy of formation ΔfH |
−519.5 kJ/mol |
| Standard molar entropy S |
53.0 J K−1 mol−1 |
| Hazards[3] | |
| Material safety data sheet (MSDS) | ICSC |
| EU index number | 025-001-00-3 |
| GHS pictograms | 
|
| GHS signal word | WARNING |
| GHS hazard statements | H302, H332 |
| GHS precautionary statements | P264, P270, P261, P271, P301+312, P330, P304+340, P501 |
| Flash point | non-flammable |
| Related compounds | |
| Other anions | Manganese disulfide |
| Other cations | Technetium dioxide Rhenium dioxide |
| Other manganese oxides | Manganese(II) oxide Manganese(II,III) oxide Manganese(III) oxide Manganese heptoxide |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) | |
Manganese dioxide is the inorganic compound with the formula MnO2. This blackish or brown solid occurs naturally as the mineral pyrolusite, which is the main ore of manganese. It is also present in manganese nodules. The principal use for MnO2 is for dry-cell batteries, such as the alkaline battery and the zinc–carbon battery.[4] In 1976 this application accounted for 500,000 tonnes of pyrolusite. MnO2 is also used for production of MnO4−. It is used extensively as an oxidizing agent in organic synthesis, for example, for the oxidation of allylic alcohols.
Contents
Structure
A number of polymorphs of MnO2 have been identified. The most common is β-MnO2,(pyrolusite) which has the TiO2, rutile structure.[5] Others are α-MnO2, γ-MnO2, ε-MnO2, and λ-MnO2.[5][6] The structures of all of these forms are closely related and are built from linked MnO6 octahedra and "tunnels" capable of holding cations are formed in all except β-MnO2.[5][6] Interest in these polymorphs is due to their use in lithium ion batteries.
History and chemical reactions
The chemical properties of MnO2 have been well examined since relatively pure materials have long been readily found or prepared. MnO2 was used for production of chlorine in the eighteenth century, before being displaced by electrolytic methods. The manganese dioxide was subsequently recovered by the Weldon process.
MnO2 catalyses several reactions associated with the formation of O2, reminiscent of the role of Mn in the oxygen evolving center. In a classical laboratory demonstration, a mixture of potassium chlorate and manganese dioxide is heated and the oxygen gas collected over water. Manganese dioxide also catalyses the decomposition of hydrogen peroxide to oxygen and water:
- 2H2O2 → O2 + 2H2O
In molten KOH in the presence of oxygen, MnO2 reacts to give the manganate anion.
Manganese dioxide decomposes above about 530 ºC to give manganese(III) oxide and oxygen, while hot concentrated sulfuric or hydrochloric acid reduce the MnIV to manganese(II).[4]
- 2MnO2 + 2H2SO4 → 2MnSO4 + O2 + 2H2O
- MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O
This latter reaction forms the basis for Scheele's first isolation of chlorine gas in 1774: Scheele used nascent hydrogen chloride, prepared from the reaction of sodium chloride with concentrated sulfuric acid.[7]
- MnO2 + 4NaCl + 2H2SO4 → 2Na2SO4 + MnCl2 + Cl2 + 2H2O
- E
o(MnO2(s) + 4 H+ + 2 e− ⇌ Mn2+ + 2H2O) = +1.23 V - E
o(Cl2(g) + 2 e− ⇌ 2 Cl−) = +1.36 V
- E
The reaction would not be expected to proceed, based on the standard electrode potentials, but is favoured by the extremely high acidity and the evolution (and removal) of gaseous chlorine.
Production
The natural occuring manganese dioxide contains impurities and a considerable amount of manganese(III) ions. Only a limited number of deposits contain the γ-modification in high enough purity, which is the prefered material for the battery industry. The ferrite production also needs manganese dioxide with only small amounts of impurities. Therfore the production of synthetic manganese dioxide is important. Two groups of methods are used, yielding chemical manganese dioxide (CMD) and electrolytical manganese dioxide (EMD). The CMD is mostly used for the production of ferrites, while the EMD is used for the production of batteries.[8]
Chemical manganese dioxide
One of the two chemical methods starts from natural manganese dioxide and converts it with dinitrogen tetroxide (N2O4) and water to manganese(II) nitrate solution, which is purified and after evaporation of the water a crystaline solid forms. At temperatures of 400 °C the reverse reaction releases the N2O4 and manganese dioxide is formed.[8]
- MnO2 + N2O4 → Mn(NO3)2
- Mn(NO3)2 → MnO2 + N2O4
In the other chemical process the natural manganeses dioxide ore is reduced with oil or coal to the managanese oxide. The manganese(II) oxide is dissolved in sulfuric acid and after purification the manganese(II) is precipitated as carbonate by adding ammonium carbonate. The carbonate is heated with air and a manganese oxide forms, which is a mixture between manganese(II) and manganese(IV) oxide. To complete the reaction the oxide is suspended in sulfuric acid and sodium chlorate is added. The chloric acid oxidizes the manganese oxide and chlorine is formed as by-product.[8]
- 2MnO2 + C → 2MnO + CO2
- MnO + H2SO4 → MnSO4 + H2O
- MnSO4 + (NH4)2CO3 → MnCO3↓ + (NH4)2SO4
- MnCO3 + O2 → MnO1.85 + CO2
- MnO1.85 + 2n HClO3 → MnO1.85 + nCl2 + nH2O
Electrolytical manganese dioxide
The electrochemical method starts from the manganese(II) sulfate solution, which is obtained by similar processes mentioned in the previous method, but some producers use hydrogen instead of coal and oil for the reduction. Iron is oxidized by addition of manganese dioxide and precipitates at a pH of 5 to 6, while other ions like cobalt, nickel and copper are precipitated as sulfides. Electrolysis is than used to oxidize the manganese. During that process the manganese dioxide forms a layer on the anode, which is made from titanium, lead or graphite.[8]
:MnO2 + H2 → MnO + H2O
- MnO + H2SO4 → MnSO4 + H2O
- Mn2+ + H2O → MnO2 + 2e− + 2H+
Applications
Ancient cave painters and later the ancient Egyptians used MnO2 as a black or brown pigment. One of its first uses was in glass making. Egyptian and Roman glasswares have been shown by analyses to contain over 2% manganous oxide. Pliny mentions the use of manganese oxide under the name of "magnes" among the Romans for decolorizing glass. He considered it a variety of lodestone or magnetic iron ore.[9] Eighteenth-century British chemists referred to MnO2 simply as manganese.[ref. needed]
In modern times, the predominant application of MnO2 is as a component of dry cell batteries, so called zinc–carbon batteries. Approximately 500,000 tonnes are consumed for this application annually. Two distinct synthetic forms of the dioxide are used for batteries, chemical manganese dioxide (CMD) and electrolytic manganese dioxide (EMD).[10]
MnO2 in organic synthesis
A specialized use of manganese dioxide is as an oxidant in organic synthesis.[11] The effectiveness of the reagent depends on the method of preparation, a problem that is typical for other heterogeneous reagents where surface area, among other variables, is a significant factor.[12] The mineral pyrolusite makes a poor reagent. Usually, however, the reagent is generated by treatment of an aqueous solution KMnO4 with a Mn(II) salt, typically the sulfate. MnO2 oxidizes allylic alcohols to the corresponding aldehydes:
- cis-RCH=CHCH2OH + MnO2 → cis-RCH=CHCHO + “MnO” + H2O
The configuration of the double bond is conserved in the reaction. The corresponding acetylenic alcohols are also suitable substrates, although the resulting propargylic aldehydes can be quite reactive. Benzylic and even unactivated alcohols are also good substrates. 1,2-Diols are cleaved by MnO2 to dialdehydes or diketones. Otherwise, the applications of MnO2 are numerous, being applicable to many kinds of reactions including amine oxidation, aromatization, oxidative coupling, and thiol oxidation.
Precursor to permanganate
The green salt potassium manganate is obtained in minutes when a solution of MnO2 in molten KOH or NaOH is treated with oxidizing agents such as potassium nitrate (KNO3), potassium perchlorate (KClO4), or even oxygen gas.[10]
- 2MnO2 + 4OH− + O2 → 2MnO42− + 2H2O
Potassium manganate converts into purple potassium permanganate in aqueous acidic solution:
- 3MnO42− + 4H+ → 2MnO4− + MnO2 + 2H2O
References
- ↑ CRC Handbook of Chemistry and Physics, 62nd ed.; Weast, Robert C., Ed.; CRC Press: Boca Raton, FL, 1981; p B-118. ISBN 0-8493-0462-8.
- ↑ Jolly, William L. Modern Inorganic Chemistry, 2nd ed.; McGraw-Hill: New York, 1991; p 586. ISBN 0-07-032768-8.
- ↑ Index no. 025-001-00-3 of Annex VI, Part 3, to Regulation (EC) No 1272/2008 of the European Parliament and of the Council of 16 December 2008 on classification, labelling and packaging of substances and mixtures, amending and repealing Directives 67/548/EEC and 1999/45/EC, and amending Regulation (EC) No 1907/2006. OJEU L353, 31.12.2008, pp 1–1355 at p 420.
- ↑ 4.0 4.1 Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984; pp 1218–20. ISBN 0-08-022057-6.
- ↑ 5.0 5.1 5.2 Wells, A. F. Structural Inorganic Chemistry, 5th ed.; Clarendon Press: Oxford, 1984. ISBN 0-19-855370-6.
- ↑ 6.0 6.1 Fong, G. C.; Kennedy, B. J.; Elcombe, M. M. A powder neutron diffraction study of λ- and γ-manganese dioxide and of LiMn2O4. Z. Kristallogr. 1994, 209 (12), 941–45.
- ↑ Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984; p 923. ISBN 0-08-022057-6.
- ↑ 8.0 8.1 8.2 8.3 Preisler, Eberhard Moderne Verfahren der Großchemie: Braunstein. Chemie in unserer Zeit 1980, 14, 137–48. DOI: 10.1002/ciuz.19800140502.
- ↑ Harder, Edmund Cecil Manganese Deposits of the United States with Sections on Foreign Deposits, Chemistry, and Uses; Bulletin of the United States Geological Survey 427; Government Printing Office: Washington, DC, 1910, <http://www.archive.org/details/manganesedeposi00hardgoog>.
- ↑ 10.0 10.1 Reidies, Arno H. Manganese Compounds. In Ullmann's Encyclopedia of Industrial Chemistry; Wiley-VCH: Weinheim, 2002; Vol. 20, pp 495–542. ISBN 3-527-30385-5. DOI: 10.1002/14356007.a16_123.
- ↑ Cahiez, G.; Alami, M.; Taylor, R. J. K.; Reid, M.; Foot, J. S. Manganese Dioxide. In Encyclopedia of Reagents for Organic Synthesis; Paquette, Leo A., Ed.; John Wiley: New York, 2004.
- ↑ Attenburrow, J.; Cameron, A. F. B.; Chapman, J. H.; Evans, R. M.; Hems, B. A.; Jansen, A. B. A.; Walker, T. A synthesis of vitamin a from cyclohexanone. J. Chem. Soc. 1952, 1094–11. DOI: 10.1039/JR9520001094.
External links
- International Chemical Safety Card 0175
- Preparations from Organic Syntheses in which manganese dioxide appears
- Entry for "Manganese and compounds" on the Australian National Pollutant Inventory
- REACH Mn Consortium
- Example Reactions with Mn(IV) oxide
- Potters Manganese Toxicity by Elke Blodgett
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