Manganese (IV) Oxide (250gr)
Manganese dioxide 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. 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. MnO2 is 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-MnO2 as a possible cathode for lithium-ion batteries.
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 pyrolusite or β-MnO2), with three-coordinate oxide and octahedral metal centres. 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. The α-polymorph of MnO2 has a very open structure with “channels” which can accommodate metal atoms such as silver or barium. α-MnO2 is often called hollandite, after a closely related mineral.
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”.
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 N2O4 and leaving a residue of purified manganese dioxide. These two steps can be summarized as:
- MnO2 + N2O4 ⇌ Mn(NO3)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.
A third process involves manganese heptoxide and manganese monoxide. The two reagents combine with a 1:3 ratio to form manganese dioxide:
- Mn2O7 + 3 MnO → 5 MnO2
- 2 KMnO4 + 3 MnSO4 + 2 H2O→ 5 MnO2 + K2SO4 + 2 H2SO4
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.
The predominant application of MnO2 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. Other industrial applications include the use of MnO2 as an inorganic pigment in ceramics and in glassmaking. It is also used in water treatment applications.
Excavations at the Pech-de-l’Azé cave site in southwestern France have yielded blocks of manganese dioxide, with lots of scratch marks 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.
A specialized use of manganese dioxide is as oxidant in organic synthesis. 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. 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: 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.