Manganese dioxide

Source: Wikipedia, the free encyclopedia.
Manganese dioxide
Manganese(IV) oxideMn4O2
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 Edit this at Wikidata
EC Number
  • 215-202-6
RTECS number
  • OP0350000
UNII
  • InChI=1S/Mn.2O checkY
    Key: NUJOXMJBOLGQSY-UHFFFAOYSA-N checkY
  • O=[Mn]=O
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
+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
53.1 J·mol−1·K−1
Std enthalpy of
formation
(ΔfH⦵298)
 −520.0 kJ·mol−1
−465.1 kJ·mol−1
Hazards
GHS labelling:
GHS07: Exclamation mark
Warning
H302, H332
P261, P264, P270, P271, P301+P312, P304+P312, P304+P340, P312, P330, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 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
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).
☒N verify (what is checkY☒N ?)

Manganese dioxide is the

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]

Structure

Several

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

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

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

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

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

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

water
:

2 H
2O
2 → 2 H
2O + O
2

Manganese dioxide decomposes above about 530 °C to

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

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 (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

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

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

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

References

  1. ^ Rumble, p. 4.71
  2. doi:10.1016/0022-3697(95)00037-2
    .
  3. ^ Rumble, p. 5.25
  4. ^
    ISBN 978-0-08-022057-4
    ..
  5. hdl:10533/173039
    .
  6. doi:10.1021/cm400864n
    .
  7. ^
    ISBN 9780470842898
    .
  8. ^
    doi:10.1002/ciuz.19800140502
    .
  9. ^ Arthur Sutcliffe (1930) Practical Chemistry for Advanced Students (1949 Ed.), John Murray – London.
  10. doi:10.1039/C5RA05892A
    .
  11. ^
    ISBN 3527306730
    .
  12. ISBN 978-3-527-30385-4
  13. ISSN 1385-8947
    .
  14. ^ "Neandertals may have used chemistry to start fires". www.science.org. Retrieved 2022-05-30.
  15. 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. PMID 15518832
    .

Cited sources

External links

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