Manganese dioxide

Manganese dioxide
Names
IUPAC names
Manganese oxide
Manganese(IV) oxide
Other names
Pyrolusite, hyperoxide of manganese, black oxide of manganese
Identifiers
1313-13-9 YesY
3D model (Jmol) Interactive image
ChemSpider 14117 YesY
ECHA InfoCard 100.013.821
EC Number 215-202-6
PubChem 14801
RTECS number OP0350000
Properties
MnO
2
Molar mass 86.9368 g/mol
Appearance Brown-black solid
Density 5.026 g/cm3
Melting point 535 °C (995 °F; 808 K) (decomposes)
insoluble
Thermochemistry
53 J·mol−1·K−1[1]
−520 kJ·mol−1[1]
Hazards
Safety data sheet ICSC 0175
Harmful (Xn)
Oxidizer (O)
R-phrases R20/22
S-phrases (S2), S25
NFPA 704
Flash point 535 °C (995 °F; 808 K)
Related compounds
Other anions
 Manganese disulfide
Other cations
 Technetium dioxide
Rhenium dioxide
Manganese(II) oxide
Manganese(II,III) oxide
Manganese(III) oxide
Manganese heptoxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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 MnO
2 is for dry-cell batteries, such as the alkaline battery and the zinc-carbon battery.[2] MnO
2 is also used as a pigment and as a precursor to other manganese compounds, such as KMnO
4. It is used as a reagent in organic synthesis, for example, for the oxidation of allylic alcohols. MnO
2 in the α polymorph can incorporate a variety of atoms (as well as water molecules) in the "tunnels" or "channels" between the magnesium oxide octahedra. There is considerable interest in α-MnO
2 as a possible cathode for lithium ion batteries.[3][4]

Structure

Several polymorphs of MnO
2 are claimed, as well as a hydrated form. Like many other dioxides, MnO
2 crystallizes in the rutile crystal structure (this polymorph is called β-MnO
2), with three-coordinate oxide and octahedral metal centres.[2] 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. The α-polymorph of MnO
2 has a very open structure with "channels" which can accommodate metal atoms such as silver or barium. α-MnO
2 is often called Hollandite, after a closely related mineral.

Production

Naturally occurring manganese dioxide contains impurities and a considerable amount of manganese(III) oxide. Only a limited number of deposits contain the γ modification in purity sufficient for the battery industry.

Production of batteries and ferrite (two of the primary uses of manganese dioxide) requires high purity manganese dioxide. Batteries require "electrolytic manganese dioxide" while ferrites require "chemical manganese dioxide".[5]

Chemical manganese dioxide

One of method starts with natural manganese dioxide and converts it using dinitrogen tetroxide and water to a manganese(II) nitrate solution. Evaporation of the water, leaves the crystalline nitrate salt. At temperatures of 400 °C, the salt decomposes, releasing N
2O
4 and leaving a residue of purified manganese dioxide.[5] These two steps can be summarized as:

MnO
2 + N
2O
4 ⇌ Mn(NO
3)
2

In another process manganese dioxide is carbothermically reduced to manganese(II) oxide which is dissolved in sulfuric acid. The filtered solution is treated with ammonium carbonate to precipitate MnCO
3. 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.[5]

A third process involves manganese heptoxide and manganese monoxide. The two reagents combine with a 1:3 ratio to form manganese dioxide:

Mn
2O
7 + 3 MnO → 5 MnO
2

Lastly the action of potassium permanganate over manganese sulphate crystals produces the desired oxide.[6]

2 KMnO
4 + 3 MnSO
4 + 2 H
2O→ 5 MnO
2 + K
2SO
4 + 2 H
2SO
4

Electrolytic manganese dioxide

Electrolytic manganese dioxide (EMD) is used in zinc-carbon batteries together with zinc chloride and ammonium chloride. EMD is commonly used in zinc manganese dioxide rechargeable alkaline (Zn RAM) cells also. For these applications, purity is extremely important. EMD is produced in a similar fashion as electrolytic tough pitch (ETP) copper: The manganese dioxide is dissolved in sulfuric acid (sometimes mixed with manganese sulfate) and subjected to a current between two electrodes. The MnO2 dissolves, enters solution as the sulfate, and is deposited on the anode.

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 H
2O
2 → 2 H
2O + O
2

Manganese dioxide decomposes above about 530 °C to manganese(III) oxide and oxygen. At temperatures close to 1000 °C, the mixed-valence compound Mn
3O
4 forms. Higher temperatures give MnO.

Hot concentrated sulfuric acid reduces the MnO
2 to manganese(II) sulfate:[2]

2 MnO
2 + 2 H
2SO
4 → 2 MnSO
4 + O
2 + 2 H
2O

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

MnO
2 + 4 HCl → MnCl
2 + Cl
2 + 2 H
2O

As a source of hydrogen chloride, Scheele treated sodium chloride with concentrated sulfuric acid.[2]

Eo (MnO
2(s) + 4 H+
 + 2 e− ⇌ Mn2+ + 2 H
2O) = +1.23 V
Eo (Cl
2(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 MnO
2 in air gives green potassium manganate:

2 MnO
2 + 4 KOH + O
2 → 2 K
2MnO
4 + 2 H
2O

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

Applications

The predominant application of MnO
2 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.[7] 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 oxidant in organic synthesis.[8] 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.[9] The mineral pyrolusite makes a poor reagent. Usually, however, the reagent is generated in situ by treatment of an aqueous solution KMnO
4 with a Mn(II) salt, typically the sulfate. MnO
2 oxidizes allylic alcohols to the corresponding aldehydes or ketones:[10]

cis-RCH=CHCH
2OH + MnO
2 → cis-RCH=CHCHO + MnO + H
2O

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 MnO
2 to dialdehydes or diketones. Otherwise, the applications of MnO
2 are numerous, being applicable to many kinds of reactions including amine oxidation, aromatization, oxidative coupling, and thiol oxidation.

Pigment

Manganese dioxide, in the form of umber, 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. When dry avoid breathing in fine particles by wearing a simple medical mask or such to avoid damage to lungs.

See also

References

  1. 1 2 Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
  2. 1 2 3 4 Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 1218–20. ISBN 0-08-022057-6..
  3. ↑ Barbato, S (31 May 2001). "Hollandite cathodes for lithium ion batteries. 2. Thermodynamic and kinetics studies of lithium insertion into BaMMn7O16 (M=Mg, Mn, Fe, Ni)". Electrochimica Acta. 46 (18): 2767–2776. doi:10.1016/S0013-4686(01)00506-0.
  4. ↑ Tompsett, David A.; Islam, M. Saiful (25 June 2013). "Electrochemistry of Hollandite α-MnO : Li-Ion and Na-Ion Insertion and Li Incorporation". Chemistry of Materials. 25 (12): 2515–2526. doi:10.1021/cm400864n.
  5. 1 2 3 Preisler, Eberhard (1980), "Moderne Verfahren der Großchemie: Braunstein", Chemie in unserer Zeit, 14 (5): 137–48, doi:10.1002/ciuz.19800140502.
  6. ↑ Arthur Sutcliffe (1930) Practical Chemistry for Advanced Students (1949 Ed.), John Murray - London.
  7. ↑ 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.
  8. ↑ 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.
  9. ↑ Attenburrow, J.; Cameron, A. F. B.; Chapman, J. H.; Evans, R. M.; Hems, B. A.; Jansen, A. B. A.; Walker, T. (1952), "A synthesis of vitamin a from cyclohexanone", J. Chem. Soc.: 1094–1111, doi:10.1039/JR9520001094.
  10. ↑ Leo A. Paquette and Todd M. Heidelbaugh. "(4S)-(−)-tert-Butyldimethylsiloxy-2-cyclopen-1-one". Org. Synth.; Coll. Vol., 9, p. 136 (this procedure illustrates the use of MnO2 for the oxidation of an allylic alcohol.
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