Manganese dioxide

Last updated
Manganese dioxide
Manganese(IV) oxide.jpg
Rutile-unit-cell-3D-balls.png
Names
IUPAC names
Manganese dioxide
Manganese(IV) oxide
Other names
Pyrolusite, hyperoxide of manganese, black oxide of manganese, manganic oxide
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.013.821 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 215-202-6
PubChem CID
RTECS number
  • OP0350000
UNII
  • InChI=1S/Mn.2O Yes check.svgY
    Key: NUJOXMJBOLGQSY-UHFFFAOYSA-N Yes check.svgY
  • O=[Mn]=O
Properties
MnO
2
Molar mass 86.9368 g/mol
AppearanceBrown-black solid
Density 5.026 g/cm3
Melting point 535 °C (995 °F; 808 K) (decomposes)
Insoluble
+2280.0×10−6 cm3/mol [1]
Structure [2]
Tetragonal, tP6, No. 136
P42/mnm
a = 0.44008 nm, b = 0.44008 nm, c = 0.28745 nm
2
Thermochemistry [3]
54.1 J·mol−1·K−1
Std molar
entropy (S⦵298)
53.1 J·mol−1·K−1
−520.0 kJ·mol−1
−465.1 kJ·mol−1
Hazards
GHS labelling:
GHS-pictogram-exclam.svg
Warning
H302, H332
P261, P264, P270, P271, P301+P312, P304+P312, P304+P340, P312, P330, P501
NFPA 704 (fire diamond)
NFPA 704.svgHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 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
2
1
2
OX
Flash point 535 °C (995 °F; 808 K)
Safety data sheet (SDS) ICSC 0175
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
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)

Manganese dioxide 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. [4] 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 has an α-polymorph that can incorporate a variety of atoms (as well as water molecules) in the "tunnels" or "channels" between the manganese oxide octahedra. There is considerable interest in α-MnO
2 as a possible cathode for lithium-ion batteries. [5] [6]

Contents

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 pyrolusite or β-MnO
2), with three-coordinate oxide anions and octahedral metal centres. [4] 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. [7] The α-polymorph of MnO
2 has a very open structure with "channels", which can accommodate metal ions 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. 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". [8]

Chemical manganese dioxide

One 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. [8] 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. [8]

Lastly, the action of potassium permanganate over manganese sulfate crystals produces the desired oxide. [9]

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. [10]

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: [11]

MnO
2 + 2 C → Mn + 2 CO

The key redox 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, which is reduced only with difficulty. [11]

Hot concentrated sulfuric acid reduces MnO
2 to manganese(II) sulfate: [4]

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. [4]

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 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.

This reaction is also a convenient way to remove the manganese dioxide precipitate from the ground glass joints after running a reaction (for example, 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.

Occurrence and applications

The predominant application of MnO
2 is as a component of dry cell batteries: alkaline batteries and so called Leclanché cell, or zinc–carbon batteries. Approximately 500,000 tonnes are consumed for this application annually. [12] Other industrial applications include the use of MnO
2 as an inorganic pigment in ceramics and in glassmaking. It is also used in water treatment applications. [13]

Prehistory

Excavations at the Pech-de-l'Azé cave site in southwestern France have yielded blocks of manganese dioxide writing tools, which date back 50,000 years and have been attributed to Neanderthals . Scientists have conjectured that Neanderthals used this mineral for body decoration, but there are many other readily available minerals that are more suitable for that purpose. Heyes et al. (in 2016) determined that the manganese dioxide lowers the combustion temperatures for wood from above 650 °F to 480 °F, making fire making much easier and this is likely to be the purpose of the blocks. [14]

Organic synthesis

A specialized use of manganese dioxide is as oxidant in organic synthesis. [7] 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. [15] 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: [16]

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.

Microbiology

In Geobacteraceae sp., MnO2 functions as an electron acceptor coupled to the oxidation of organic compounds. This theme has implications for bioremediation. [17]

See also

Related Research Articles

<span class="mw-page-title-main">Carboxylic acid</span> Organic compound containing a –C(=O)OH group

In organic chemistry, a carboxylic acid is an organic acid that contains a carboxyl group attached to an R-group. The general formula of a carboxylic acid is often written as R−COOH or R−CO2H, sometimes as R−C(O)OH with R referring to an organyl group, or hydrogen, or other groups. Carboxylic acids occur widely. Important examples include the amino acids and fatty acids. Deprotonation of a carboxylic acid gives a carboxylate anion.

<span class="mw-page-title-main">Manganese</span> Chemical element with atomic number 25 (Mn)

Manganese is a chemical element; it has symbol Mn and atomic number 25. It is a hard, brittle, silvery metal, often found in minerals in combination with iron. Manganese was first isolated in the 1770s. It is a transition metal with a multifaceted array of industrial alloy uses, particularly in stainless steels. It improves strength, workability, and resistance to wear. Manganese oxide is used as an oxidising agent; as a rubber additive; and in glass making, fertilisers, and ceramics. Manganese sulfate can be used as a fungicide.

<span class="mw-page-title-main">Oxidizing agent</span> Chemical compound used to oxidize another substance in a chemical reaction

An oxidizing agent is a substance in a redox chemical reaction that gains or "accepts"/"receives" an electron from a reducing agent. In other words, an oxidizer is any substance that oxidizes another substance. The oxidation state, which describes the degree of loss of electrons, of the oxidizer decreases while that of the reductant increases; this is expressed by saying that oxidizers "undergo reduction" and "are reduced" while reducers "undergo oxidation" and "are oxidized". Common oxidizing agents are oxygen, hydrogen peroxide, and the halogens.

<span class="mw-page-title-main">Potassium permanganate</span> Chemical compound

Potassium permanganate is an inorganic compound with the chemical formula KMnO4. It is a purplish-black crystalline salt, that dissolves in water as K+ and MnO−
4, an intensely pink to purple solution.

In environmental chemistry, the chemical oxygen demand (COD) is an indicative measure of the amount of oxygen that can be consumed by reactions in a measured solution. It is commonly expressed in mass of oxygen consumed over volume of solution, which in SI units is milligrams per liter (mg/L). A COD test can be used to quickly quantify the amount of organics in water. The most common application of COD is in quantifying the amount of oxidizable pollutants found in surface water or wastewater. COD is useful in terms of water quality by providing a metric to determine the effect an effluent will have on the receiving body, much like biochemical oxygen demand (BOD).

<span class="mw-page-title-main">Manganese(II) chloride</span> Chemical compound

Manganese(II) chloride is the dichloride salt of manganese, MnCl2. This inorganic chemical exists in the anhydrous form, as well as the dihydrate (MnCl2·2H2O) and tetrahydrate (MnCl2·4H2O), with the tetrahydrate being the most common form. Like many Mn(II) species, these salts are pink, with the paleness of the color being characteristic of transition metal complexes with high spin d5 configurations.

<span class="mw-page-title-main">Permanganate</span> Chemical compound

A permanganate is a chemical compound with the manganate(VII) ion, MnO−
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4, also known as manganate(VI) because it contains manganese in the +6 oxidation state. Manganates are the only known manganese(VI) compounds.

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<span class="mw-page-title-main">Potassium manganate</span> Chemical compound

Potassium manganate is the inorganic compound with the formula K2MnO4. This green-colored salt is an intermediate in the industrial synthesis of potassium permanganate, a common chemical. Occasionally, potassium manganate and potassium permanganate are confused, but each compound's properties are distinct.

In chemistry, hypomanganate, also called manganate(V) or tetraoxidomanganate(3−), is a trivalent anion (negative ion) composed of manganese and oxygen, with formula MnO3−
4.

Potassium hypomanganate is the inorganic compound with the formula K3MnO4. Also known as potassium manganate(V), this bright blue solid is a rare example of a salt with the hypomanganate or manganate(V) anion, where the manganese atom is in the +5 oxidation state. It is an intermediate in the production of potassium permanganate and the industrially most important Mn(V) compound.

<span class="mw-page-title-main">Manganese(II) sulfate</span> Chemical compound

Manganese(II) sulfate usually refers to the inorganic compound with the formula MnSO4·H2O. This pale pink deliquescent solid is a commercially significant manganese(II) salt. Approximately 260,000 tonnes of manganese(II) sulfate were produced worldwide in 2005. It is the precursor to manganese metal and many other chemical compounds. Manganese-deficient soil is remediated with this salt.

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<span class="mw-page-title-main">Barium manganate</span> Chemical compound

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Barium permanganate is a chemical compound, with the formula Ba(MnO4)2. It forms violet to brown crystals that are sparingly soluble in water.

A lithium ion manganese oxide battery (LMO) is a lithium-ion cell that uses manganese dioxide, MnO
2, as the cathode material. They function through the same intercalation/de-intercalation mechanism as other commercialized secondary battery technologies, such as LiCoO
2. Cathodes based on manganese-oxide components are earth-abundant, inexpensive, non-toxic, and provide better thermal stability.

References

  1. ↑ Rumble, p. 4.71
  2. ↑ Haines, J.; Léger, J.M.; Hoyau, S. (1995). "Second-order rutile-type to CaCl2-type phase transition in β-MnO2 at high pressure". Journal of Physics and Chemistry of Solids. 56 (7): 965–973. Bibcode:1995JPCS...56..965H. doi:10.1016/0022-3697(95)00037-2.
  3. ↑ Rumble, p. 5.25
  4. 1 2 3 4 Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 1218–20. ISBN   978-0-08-022057-4..
  5. ↑ 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. hdl: 10533/173039 .
  6. ↑ 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. CiteSeerX   10.1.1.728.3867 . doi:10.1021/cm400864n.
  7. 1 2 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, pp. 1–16, doi:10.1002/047084289X.rm021.pub4, ISBN   9780470842898 .
  8. 1 2 3 Preisler, Eberhard (1980), "Moderne Verfahren der Großchemie: Braunstein", Chemie in unserer Zeit, 14 (5): 137–48, doi:10.1002/ciuz.19800140502 .
  9. ↑ Arthur Sutcliffe (1930) Practical Chemistry for Advanced Students (1949 Ed.), John Murray – London.
  10. ↑ Biswal, Avijit; Chandra Tripathy, Bankim; Sanjay, Kali; Subbaiah, Tondepu; Minakshi, Manickam (2015). "Electrolytic manganese dioxide (EMD): A perspective on worldwide production, reserves and its role in electrochemistry". RSC Advances. 5 (72): 58255–58283. doi:10.1039/C5RA05892A.
  11. 1 2 Wellbeloved, David B.; Craven, Peter M.; Waudby, John W. (2000). "Manganese and Manganese Alloys". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a16_077. ISBN   3527306730.
  12. ↑ 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   978-3-527-30385-4
  13. ↑ Ibrahim, Yazan; Wadi, Vijay S.; Ouda, Mariam; Naddeo, Vincenzo; Banat, Fawzi; Hasan, Shadi W. (15 January 2022). "Highly selective heavy metal ions membranes combining sulfonated polyethersulfone and self-assembled manganese oxide nanosheets on positively functionalized graphene oxide nanosheets". Chemical Engineering Journal. 428: 131267. doi:10.1016/j.cej.2021.131267. ISSN   1385-8947.
  14. ↑ "Neandertals may have used chemistry to start fires". www.science.org. Retrieved 2022-05-30.
  15. ↑ 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 .
  16. ↑ Paquette, Leo A. and Heidelbaugh, Todd M. "(4S)-(−)-tert-Butyldimethylsiloxy-2-cyclopen-1-one". Organic Syntheses {{cite journal}}: CS1 maint: multiple names: authors list (link); Collected Volumes, vol. 9, p. 136. (this procedure illustrates the use of MnO2 for the oxidation of an allylic alcohol)
  17. ↑ Lovley, Derek R.; Holmes, Dawn E.; Nevin, Kelly P. (2004). Dissimilatory Fe(III) and Mn(IV) Reduction. Advances in Microbial Physiology. Vol. 49. pp. 219–286. doi:10.1016/S0065-2911(04)49005-5. ISBN   9780120277490. PMID   15518832.

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