Manganese dioxide

Manganese dioxide
IUPAC name Manganese oxide Manganese(IV) oxide
Other names Pyrolusite
Identifiers
CAS number	1313-13-9 Y
PubChem	14801
ChemSpider	14117 Y
EC number	215-202-6
Jmol-3D images	Image 1
SMILES O=[Mn]=O
InChI InChI=1S/Mn.2O Y Key: NUJOXMJBOLGQSY-UHFFFAOYSA-N Y
Properties
Molecular formula	MnO2
Molar mass	86.9368 g/mol
Appearance	Brown-black solid
Density	5.026 g/cm3
Melting point	535 °C decomp
Solubility in water	insoluble
Thermochemistry
Std enthalpy of formation ΔfHo298	−520.9 kJ/mol
Standard molar entropy So298	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
Y (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Manganese(IV) oxide 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 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] MnO₂ is also used as a pigment and as a precursor to other manganese compounds, such as KMnO₄. It is used as a reagent in organic synthesis, for example, for the oxidation of allylic alcohols.

Structure

Several polymorphs of MnO₂ are claimed, as well as a hydrated form. Like many other dioxides, MnO₂ crystallizes in the rutile crystal structure (this polymorph is called β-MnO₂), with three-coordinate oxide and octahedral metal centres.^[1] MnO₂ is characteristically nonstoichiometric, being deficient in oxygen. The complicated solid-state chemistry of this material is relevant to the lore of "freshly prepared" MnO₂ 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 MnO₂ are associated with its redox, both oxidation and reduction.

Reduction

MnO₂ is the principal precursor to ferromanganese and related alloys, which are widely used in the steel industry. The conversions involve carbothermal reduction using coke:

MnO₂ + 2 C → Mn + 2 CO

The key reactions of MnO₂ in batteries is the one-electron reduction:

MnO₂ + e^- + H⁺ → MnO(OH)

MnO₂ catalyses several reactions that form O₂. 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]

E^o (MnO₂(s) + 4 H⁺ + 2 e⁻ Mn²⁺ + 2 H₂O) = +1.23 V

E^o (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.

This reaction is also a convenient way to remove the manganese dioxide precipitate from the ground glass joints after running a reaction (i. e., an oxidation with potassium permanganate).

Oxidation

Heating a mixture of KOH and MnO₂ in air gives green potassium manganate:

2 MnO₂ + 4 KOH + O₂ → 2 K₂MnO₄ + 2 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 MnO₂ as an inorganic pigment in ceramics and in glassmaking.

Organic synthesis

A specialized use of manganese dioxide is as 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, A. (1984). Chemistry of the Elements. Oxford: Pergamon. pp. 1218–20. ISBN 0-08-022057-6. .
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, 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., 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 .
6. ^ Leo A. Paquette and Todd M. Heidelbaugh, "(4S)-(−)-tert-Butyldimethylsiloxy-2-cyclopen-1-one", Org. Synth., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv9p0136 ; Coll. Vol. 9: 136  (this procedure illustrates the use of MnO2 for the oxidation of an allylic alcohol.

External links

• REACH Mn Consortium
• Index of Organic Synthesis procedures utilizing MnO₂
• Example Reactions with Mn(IV) oxide
• National Pollutant Inventory - Manganese and compounds Fact Sheet
• PubChem summary of MnO₂
• International Chemical Safety Card 0175
• Potters Manganese Toxicity by Elke Blodgett