| Manganese dioxide | |
|---|---|
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|
| IUPAC name |
Manganese oxide
Manganese(IV) oxide |
| Other names | Pyrolusite |
| Identifiers | |
| CAS number | 1313-13-9  |
| PubChem | 14801 |
| EC number | 215-202-6 |
| Properties | |
| Molecular formula | MnO2 |
| Molar mass | 86.9368 g/mol |
| Appearance | black solid |
| Density | 5.026 g/cm3 |
| Melting point |
535 °C decomp |
| Solubility in water | insoluble |
| Thermochemistry | |
| Std enthalpy of formation ΔfH |
−520.9 kJ/mol |
| Standard molar entropy S |
53.1 J K−1 mol−1 |
| Hazards | |
| MSDS | ICSC 0175 |
| EU Index | 025-001-00-3 |
| EU classification | Harmful (Xn) Oxidizer (O) |
| R-phrases | R20/22 |
| S-phrases | (S2), S25 |
| NFPA 704 |
 1
1
2
OX
|
| Flash point | 535 °C |
| Related compounds | |
| Other anions | Manganese disulfide |
| Other cations | Technetium dioxide Rhenium dioxide |
| Related manganese oxides | Manganese(II) oxide Manganese(II,III) oxide Manganese(III) oxide Manganese heptoxide |
| (what is this?) (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
|
| Infobox references | |
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.[1] In 1976 this application accounted for 500,000 tonnes of pyrolusite. MnO2 is also used for production of MnO−4. It is used extensively as an oxidizing agent in organic synthesis, for example, for the oxidation of allylic alcohols.
Contents |
Structure
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 °C to give manganese(III) oxide and oxygen, while hot concentrated sulfuric or hydrochloric acid reduce the MnIV to manganese(II).[1]
- 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.[2]
- 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.
==
The natural occurring 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 preferred material for the battery industry. The ferrite production also needs manganese dioxide with only small amounts of impurities. Therefore 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.[3]
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 crystalline solid forms. At temperatures of 400 °C the reverse reaction releases the N2O4 and manganese dioxide is formed.[3]
- MnO2 + N2O4 → Mn(NO3)2
- Mn(NO3)2 → MnO2 + N2O4
In the other chemical process, the natural manganese dioxide ore is reduced with oil or coal to the manganese 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.[3]
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).[4]
MnO2 in organic synthesis
A specialized use of manganese dioxide is as an oxidant in organic synthesis.[5] 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.[6] 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.
References
- ^ a b Greenwood, Norman N.; Earnshaw, A. (1984), Chemistry of the Elements, Oxford: Pergamon, pp. 1218–20, ISBN 0-08-022057-6.
- ^ Greenwood, Norman N.; Earnshaw, A. (1984), Chemistry of the Elements, Oxford: Pergamon, p. 923, ISBN 0-08-022057-6.
- ^ a b c Preisler, Eberhard (1980), "Moderne Verfahren der Großchemie: Braunstein", Chemie in unserer Zeit 14: 137–48, doi:.
- ^ Reidies, Arno H. (2002), "Manganese Compounds", Ullmann's Encyclopedia of Industrial Chemistry, 20, Weinheim: Wiley-VCH, pp. 495–542, doi:, ISBN 3-527-30385-5.
- ^ Cahiez, G.; Alami, M.; Taylor, R. J. K.; Reid, M.; Foot, J. S. (2004), "Manganese Dioxide", in Paquette, Leo A., Encyclopedia of Reagents for Organic Synthesis, New York: J. Wiley & Sons.
- ^ Attenburrow, J.; Cameron, A. F. B.; Chapman, J. H.; Evans, R. M.; Hems, B. A.; Jansen, A. B. A.; Walker, T. (1952), J. Chem. Soc.: 1094.
Sources
- Oosterhoeks Encyclopedie (Dutch)
External links
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