_oxide.jpg) | |
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| 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 |
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| ECHA InfoCard | 100.013.821 |
| EC Number |
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| RTECS number |
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| UNII | |
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| 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 | |
Formula units (Z) | 2 |
| Thermochemistry[3] | |
| 54.1 J·mol−1·K−1 | |
Std molar entropy(S⦵298) | 53.1 J·mol−1·K−1 |
Std enthalpy of formation(ΔfH⦵298) | −520.0 kJ·mol−1 |
Gibbs free energy(ΔfG⦵) | −465.1 kJ·mol−1 |
| Hazards | |
| GHS labelling: | |
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| Warning | |
| H302,H332 | |
| P261,P264,P270,P271,P301+P312,P304+P312,P304+P340,P312,P330,P501 | |
| NFPA 704 (fire diamond) | |
| Flash point | 535 °C (995 °F; 808 K) |
| Safety data sheet (SDS) | ICSC 0175 |
| Related compounds | |
Otheranions | Manganese disulfide |
Othercations | 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 theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Manganese dioxide is theinorganic compound with theformulaMnO
2. This blackish or brown solid occurs naturally as the mineralpyrolusite, which is the main ore ofmanganese and a component ofmanganese nodules. The principal use forMnO
2 is for dry-cellbatteries, such as thealkaline battery and thezinc–carbon battery, although it is also used for other battery chemistries such asaqueouszinc-ion batteries.[4][5]MnO
2 is also used as apigment and as a precursor to other manganese compounds, such aspotassium permanganate (KMnO4). It is used as areagent inorganic synthesis, for example, for the oxidation ofallylicalcohols.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 inα-MnO
2 as a possible cathode forlithium-ion batteries.[6][7]
Severalpolymorphs ofMnO
2 are claimed, as well as a hydrated form. Like many other dioxides,MnO
2 crystallizes in therutilecrystal structure (this polymorph is calledpyrolusite orβ-MnO
2), with three-coordinate oxide anions and octahedral metal centres.[4]MnO
2 is characteristicallynonstoichiometric, being deficient in oxygen. The complicatedsolid-state chemistry of this material is relevant to the lore of "freshly prepared"MnO
2 inorganic synthesis.[8] The α-polymorph ofMnO
2 has a very open structure with "channels", which can accommodate metal ions such as silver or barium.α-MnO
2 is often calledhollandite, after a closely related mineral. Two other polymorphs, Todorokite and RomanechiteMnO
2, have a similar structure toα-MnO
2 but with larger channels.δ-MnO
2 exhibits a layered structure more akin to that ofgraphite.[5]
Naturally occurring manganese dioxide contains impurities and a considerable amount ofmanganese(III) oxide. Production ofbatteries andferrite (two of the primary uses of manganese dioxide) requires high purity manganese dioxide. Batteries require "electrolytic manganese dioxide" while ferrites require "chemical manganese dioxide".[9]
One method starts with natural manganese dioxide and converts it usingdinitrogen tetroxide and water to amanganese(II) nitrate solution. Evaporation of the water leaves the crystalline nitrate salt. At temperatures of 400 °C, the salt decomposes, releasingN
2O
4 and leaving a residue of purified manganese dioxide.[9] These two steps can be summarized as:
In another process, manganese dioxide iscarbothermically reduced tomanganese(II) oxide which is dissolved insulfuric acid. The filtered solution is treated withammonium carbonate to precipitateMnCO
3. The carbonate iscalcined 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 withsodium chlorate.Chloric acid, which forms in situ, converts any Mn(III) and Mn(II) oxides to the dioxide, releasing chlorine as a by-product.[9]
Lastly, the action ofpotassium permanganate overmanganese sulfate crystals produces the desired oxide.[10]
The above reaction is an example of potassium permanganate reacting to make manganese dioxide.
Most reactions with potassium permanganate are known to make brown manganese dioxide as a byproduct, where potassium permanganate undergoes aRedox reaction where it reduces and oxidizes a compound with manganese dioxide byproduct.
Electrolytic manganese dioxide (EMD) is used inzinc–carbon batteries together withzinc chloride andammonium chloride. EMD is commonly used in zinc manganese dioxiderechargeable alkaline (Zn RAM) cells also. For these applications, purity is extremely important. EMD is produced in a similar fashion aselectrolytic tough pitch (ETP) copper: The manganese dioxide is dissolved insulfuric acid (sometimes mixed withmanganese sulfate) and subjected to a current between two electrodes. The MnO2 dissolves, enters solution as the sulfate, and is deposited on theanode.[11]
The important reactions ofMnO
2 are associated with itsredox, both oxidation and reduction.
MnO
2 is the principalprecursor toferromanganese and related alloys, which are widely used in the steel industry. The conversions involvecarbothermal reduction usingcoke:[12]
The key redox reactions ofMnO
2 in batteries is the one-electron reduction:
MnO
2catalyses several reactions that formO
2. In a classical laboratory demonstration, heating a mixture ofpotassium chlorate and manganese dioxide produces oxygen gas. Manganese dioxide also catalyses the decomposition ofhydrogen peroxide to oxygen andwater:
Manganese dioxide decomposes above about 530 °C tomanganese(III) oxide and oxygen. At temperatures close to 1000 °C, themixed-valence compoundMn
3O
4 forms. Higher temperatures give MnO, which is reduced only with difficulty.[12]
Hot concentratedsulfuric acid reducesMnO
2 tomanganese(II) sulfate:[4]
The reaction ofhydrogen chloride withMnO
2 was used byCarl Wilhelm Scheele in the original isolation ofchlorine gas in 1774. Scheele treatedsodium chloride with concentrated sulfuric acid:[4]
Standard electrode potentials suggest that the reaction would not proceed...
...but it is favoured by the extremely highacidity and the evolution (and removal) of gaseous chlorine.
This reaction is also a convenient way to remove the manganese dioxideprecipitate from theground glass joints after running a reaction (for example, an oxidation withpotassium permanganate).
Heating a mixture ofKOH andMnO
2 in air gives greenpotassium manganate:
Potassium manganate is the precursor topotassium permanganate, a common oxidant.
Excavations at thePech-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 toNeanderthals . 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 350°C (662°F) to 250°C (482°F), making fire making much easier and this is likely to be the purpose of the blocks.[13]
The predominant application ofMnO
2 is as a component ofdry cell batteries: alkaline batteries and so calledLeclanché cell, orzinc–carbon batteries. Approximately 500,000tonnes are consumed for this application annually.[14]
δ-MnO
2 has also been researched as the primary cathode material for aqueous zinc-ion battery systems. Such cathodes often contain additives to address structural,kinetic, andconductivity-based issues. These carbon additives can include reduced graphene oxide (rGO) andcarbon nanotubes, among others.[15]
A specialized use of manganese dioxide is as oxidant inorganic 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.[16] The mineralpyrolusite makes a poor reagent. Usually, however, the reagent is generated in situ by treatment of an aqueous solutionKMnO
4 with a Mn(II) salt, typically the sulfate.MnO
2 oxidizesallylic alcohols to the correspondingaldehydes orketones:[17]
The configuration of thedouble bond is conserved in the reaction. The correspondingacetylenic alcohols are also suitable substrates, although the resultingpropargylic aldehydes can be quite reactive.Benzylic and even unactivated alcohols are also good substrates. 1,2-Diols are cleaved byMnO
2 todialdehydes ordiketones. Otherwise, the applications ofMnO
2 are numerous, being applicable to many kinds of reactions includingamine oxidation, aromatization,oxidative coupling, andthiol oxidation.
InGeobacteraceae sp., MnO2 functions as an electron acceptor coupled to the oxidation of organic compounds. This theme has possible implications forbioremediation within the field of microbiology.[18]
MnO
2 is used as an inorganicpigment inceramics and inglassmaking.
{{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)