Leafs
Harmless
Posts: 18
Registered: 30-7-2017
Member Is Offline
Mood: No Mood
|
|
manganese dioxide particle size
So I made some "active" MnO2. I'm wondering how fine should I mill this stuff considering it really isn't soluble in anything so presumably it would
be best to get it to a really fine state? I thought perhaps it would break up into smaller pieces in solution but apparently this doesn't happen as
last oxidation I noticed many chunks and my yield wasn't where it should've been.
Anyone with any experience working with this stuff and how to maximize reactivity/yields of this stuff... would be much appreciated. Thanks.
|
|
|
Boffis
International Hazard
Posts: 1836
Registered: 1-5-2011
Member Is Offline
Mood: No Mood
|
|
By its very nature "active" Mn oxide occurs as clumps of fine particles that have a high surface area to mass ratio and are microporous. You don't
want to mak the particles too fine or they become difficult to remove by filteration. The method you use to make it pretty much determines its
characeristic. The type made by the thermal decomposition of manganese carbonate may appear granular but it has a very high surface area per unit
weight and appears perfectly serviceable as prepared. The material made by the various wet processes is often so fine that it is hard to filter
anyway. The dried filter cake may appear massive and chunky abd does not alway disperse readily but this doesn't appear to affect its activity.
|
|
|
Leafs
Harmless
Posts: 18
Registered: 30-7-2017
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by Boffis  | | By its very nature "active" Mn oxide occurs as clumps of fine particles that have a high surface area to mass ratio and are microporous. You don't
want to mak the particles too fine or they become difficult to remove by filteration. The method you use to make it pretty much determines its
characeristic. The type made by the thermal decomposition of manganese carbonate may appear granular but it has a very high surface area per unit
weight and appears perfectly serviceable as prepared. The material made by the various wet processes is often so fine that it is hard to filter
anyway. The dried filter cake may appear massive and chunky abd does not alway disperse readily but this doesn't appear to affect its activity.
|
Thank you for that information. I think I might have to redo another batch because I already milled this pretty fine. I was thinking a celite pad
would catch it.
|
|
|
Leafs
Harmless
Posts: 18
Registered: 30-7-2017
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Boffis | | By its very nature "active" Mn oxide occurs as clumps of fine particles that have a high surface area to mass ratio and are microporous. You don't
want to mak the particles too fine or they become difficult to remove by filteration. The method you use to make it pretty much determines its
characeristic. The type made by the thermal decomposition of manganese carbonate may appear granular but it has a very high surface area per unit
weight and appears perfectly serviceable as prepared. The material made by the various wet processes is often so fine that it is hard to filter
anyway. The dried filter cake may appear massive and chunky abd does not alway disperse readily but this doesn't appear to affect its activity.
|
Hi I further looked into alternative methods to produce "active" MnO2 and the Manganese Carbonate method appears ideal but calls for stirring of the
resulting black solid with 15% HNO3 in water... I'm not too sure what this step does exactly, and I'm thinking it could be skipped as none of the
other methods call for anything like this. The appeal to me in the first place with this method is the fact it doesn't require filtration (without
that step) etc and wont get everywhere. Would appreciate yours or anyone else's opinion who is more skilled in these matters.
The reference I'm referring to.
| Quote: | Preparation of Active MnO2 by Pyrolysis of MnCO3.
Powdered MnCO3 was spread in a one-inch thick layer in a Pyrex glass and heated at 220-280 °C for about 18 h in an oven in which air circulated by
convection. The initially tan powder turned darker at about 180 °C, and black when maintained at over 220 °C. No attempt was made to determine lower
temperature or time limits, nor the upper limit of temperature. The MnO2 prepared as above was stirred with about 1 L of a solution made up of 15%
HNO3 in H2O. The slurry was filtered with suction, the solid was washed on the Buchner funnel with distilled water until the washes were about pH 5,
and finally was dried at 220-250 °C. The caked, black solid was readily crushed to a powder which retained its oxidizing ability even after having
been stored for several months in a loosely stoppered container. |
|
|
|
JJay
International Hazard
Posts: 3440
Registered: 15-10-2015
Member Is Offline
|
|
Wouldn't that step remove unreacted manganese carbonate?
|
|
|
Leafs
Harmless
Posts: 18
Registered: 30-7-2017
Member Is Offline
Mood: No Mood
|
|
yeah google indicates manganese carbonate is soluble in dilute acid. I think I'm just going to overkill it and let it cook in the oven for 48 hours.
It's a mini-oven so it shouldn't be too inefficient. I will also periodically break and mix up the solid.
edit: also getting mixed information as to the decomposition of manganese carbonate. some sources claim it decomposes to manganese (||) oxide, others
say manganese dioxide. hmmmm
| Quote: | | The carbonate is insoluble in water but, like most carbonates, hydrolyses upon treatment with acids to give water-soluble salts. Manganese carbonate
decomposes with release of carbon dioxide at 200 °C to give manganese(II) oxide: MnCO3 → MnO + CO. |
[Edited on 2-8-2017 by Leafs]
[Edited on 2-8-2017 by Leafs]
|
|
|
Boffis
International Hazard
Posts: 1836
Registered: 1-5-2011
Member Is Offline
Mood: No Mood
|
|
The decomposition of manganese II carbonate is complex. Mn II carbonate breaks down in the absense of oxygen to give MnO and CO2 but in the presence
of oxygen and/or moisture it breaks down at lower temperatures to complex multivalent Mn oxides with variable stoichometery. It is possible that the
activation is due to the leaching of Mn2+ from these complex oxides in a manor similar to the leaching of K from the oxides produced by such methods
as the oxidation of Mn sulphate with KMnO4. I make activated Mn oxide from the potassium bearing (8-10% K) Mn oxide that I have left over from the
used of potassium permanganate and KOH to oxidize various heterocyclic and hydrocarbons to carboxylic acids.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
