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Manganese dioxide

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Manganese dioxide
Manganese(IV) oxide
Names
IUPAC names
Manganese oxide
Manganese(IV) oxide
Other names
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.013.821 Edit this at Wikidata
EC Number
  • 215-202-6
  • O=[Mn]=O
Properties
MnO2
Molar mass 86.9368 g/mol
Appearance black solid
Density 5.026 g/cm3
Melting point 535 °C decomp
insoluble
Thermochemistry
53.1 J K−1 mol−1
−520.9 kJ/mol
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
1
1
2
OX
Flash point 535 °C
Related compounds
Other anions
 Manganese disulfide
Other cations
 Technetium dioxide
Rhenium dioxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Manganese(IV) oxide is the inorganic compound with the formula MnO
2. This blackish or brown solid occurs naturally as the mineral pyrolusite, which is the main ore of manganese and a component of manganese nodules. The principal use for MnO2 is for dry-cell batteries, such as the alkaline battery and the zinc-carbon battery.[1] MnO
2 is also used as a pigment and as a precursor to other manganese compounds, such as KMnO
4. It is used an reagent in organic synthesis, for example, for the oxidation of allylic alcohols.

Structure

Several polymorphs of MnO
2 are claimed, as well as a hydrated form. Like many other dioxides, MnO
2 crystallizes in the rutile motif (this polymorph is called β-MnO
2), with three-coordinate oxide and octahedral metal centres.[1] MnO
2 is characteristically nonstoichiometric, being deficient in oxygen. The complicated solid state chemistry of this material is relevant to the lore of "freshly prepared" MnO
2 in organic synthesis.

Production

Naturally occurring manganese dioxide contains impurities and a considerable amount of manganese in its 3+ oxidation state. Only a limited number of deposits contain the γ modification in purity sufficient for the battery industry. The production of ferrite also requires high purity manganese dioxide. 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, whereas EMD is used for the production of batteries.[2]

Chemical manganese dioxide

One of the two chemical methods starts from natural manganese dioxide and converts it using dinitrogen tetroxide and water to manganese(II) nitrate solution. Evaporation of the water, leaves the crystalline nitrate salt. At temperatures of 400 °C, the salt decomposes, releasing N2O4 and leaving a residue of purified manganese dioxide.[2] These two steps can be summarized as:

MnO2 + N2O4 → Mn(NO3)2
Mn(NO3)2 → MnO2 + N2O4

In the other chemical process, manganese dioxide ore is reduced by heating with oil or coal. The resulting manganese(II) oxide is dissolved in sulfuric acid, and the filtered solution is treated with ammonium carbonate to precipitate MnCO3. The carbonate is calcined in air to give a mixture of manganese(II) and manganese(IV) oxides. To complete the process, a suspension of this material in sulfuric acid is treated with sodium chlorate. Chloric acid, which forms in situ, converts any Mn(III) and Mn(II) oxides to the dioxide, releasing chlorine as a by-product.[2]

Electrolytical manganese dioxide

Manganese dioxide is used in zinc-carbon batteries together with zinc chloride and ammonium chloride.

Reactions

The important reactions of MnO
2 are associated with its redox, both oxidation and reduction.

Reduction

MnO
2 is the principal precursor to ferromanganese and related alloys, which are widely used in the steel industry. The conversions involve carbothermal reduction using coke:

MnO
2 + 2 C → Mn + 2 CO

The key reactions of MnO
2 in batteries is the one-electron reduction:

MnO
2 + e- + H+ → MnO(OH)

MnO
2 catalyses several reactions that form O
2. In a classical laboratory demonstration, heating a mixture of potassium chlorate and manganese dioxide produces oxygen gas. Manganese dioxide also catalyses the decomposition of hydrogen peroxide to oxygen and water:

2 H2O2 → 2 H2O + O2

Manganese dioxide decomposes above about 530 °C to manganese(III) oxide and oxygen. At temperatures close to 1000 °C, the mixed-valence compound Mn3O4 forms. Higher temperatures give MnO.

Hot concentrated sulfuric reduces the MnO2 to manganese(II):[1]

2 MnO2 + 2 H2SO4 → 2 MnSO4 + O2 + 2 H2O

The reaction of hydrogen chloride with MnO2 was used by Carl Wilhelm Scheele in the original isolation of chlorine gas in 1774:

MnO2 + 4 HCl → MnCl2 + Cl2 + 2 H2O

As a source of hydrogen chloride, Scheele treated sodium chloride with concentrated sulfuric acid.[1]

Eo (MnO2(s) + 4 H+ + 2 e− ⇌ Mn2+ + 2 H2O) = +1.23 V
Eo (Cl2(g) + 2 e− ⇌ 2 Cl−) = +1.36 V

The standard electrode potentials for the half reactions indicate that the reaction is endothermic at pH = 0 (1 M [H+]), but it is favoured by the lower pH as well as the evolution (and removal) of gaseous chlorine.

Oxidation

Heating a mixture of KOH and MnO
2 in air give green potassium manganate:

MnO2 + 2 KOH + ½ O2 → K2MnO4 + H2O

Potassium manganate is the precursor to potassium permanganate, a common oxidant.

Applications

The predominant application of MnO2 is as a component of dry cell batteries, so called Leclanché cell, or zinc–carbon batteries. Approximately 500,000 tonnes are consumed for this application annually.[3] Other industrial applications include the use of MnO
2 as an inorganic pigment in ceramics and in glassmaking.

Organic synthesis

A specialized use of manganese dioxide is as an oxidant in organic synthesis.[4] 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.[5] The mineral pyrolusite makes a poor reagent. Usually, however, the reagent is generated in situ by treatment of an aqueous solution KMnO4 with a Mn(II) salt, typically the sulfate. MnO2 oxidizes allylic alcohols to the corresponding aldehydes or ketones:[6]

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.

Pigment

Manganese dioxide was one of the earliest natural substances used by human ancestors. It was used as a pigment at least from the middle paleolithic. It was possibly used first for body painting, and later for cave painting. Some of the most famous early cave paintings in Europe were executed by means of manganese dioxide.

Hazards

Manganese dioxide can slightly stain human skin if it is damp or in a heterogeneous mixture, but the stains can be washed off quite easily with some rubbing.

References

  1. ^ a b c d Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 1218–20. ISBN 978-0-08-022057-4.. Cite error: The named reference "G&E" was defined multiple times with different content (see the help page).
  2. ^ a b c Preisler, Eberhard (1980), "Moderne Verfahren der Großchemie: Braunstein", Chemie in unserer Zeit, 14: 137–48, doi:10.1002/ciuz.19800140502.
  3. ^ Reidies, Arno H. (2002), "Manganese Compounds", Ullmann's Encyclopedia of Industrial Chemistry, vol. 20, Weinheim: Wiley-VCH, pp. 495–542, doi:10.1002/14356007.a16_123, ISBN 3-527-30385-5.
  4. ^ Cahiez, G.; Alami, M.; Taylor, R. J. K.; Reid, M.; Foot, J. S. (2004), "Manganese Dioxide", in Paquette, Leo A. (ed.), Encyclopedia of Reagents for Organic Synthesis, New York: J. Wiley & Sons.
  5. ^ 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 {{citation}}: Missing or empty |title= (help).
  6. ^ Leo A. Paquette and Todd M. Heidelbaugh. "(4S)-(−)-tert-Butyldimethylsiloxy-2-cyclopen-1-one". Organic Syntheses; Collected Volumes, vol. 9, p. 136. (this procedure illustrates the use of MnO2 for the oxidation of an allylic alcohol.
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