bolbol
Hazard to Others
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 167
Registered: 3-1-2015
Member Is Offline
Mood: No Mood
|
|
KMnO4 > MnO2 + HNO3 > Mn(NO3)X
Hi everyone
I have two questions regarding the synthesis of manganese nitrate that I want to do.
I have some potassium permanganate laying around so I thought of making a Manganese salt with it.
I want to decompose the KMnO4 to MnO2 but I am not sure if there is any easy way to do that. I know the H2O2 method but that would require tons of
bottles since I cant get my hands on concentrated stuff and i would probably have gallons of liquid to filter off which is just messy..
I also know that heating it will decompose too but I have never attempted it a larger scale so I don't know if its really safe to do so.
Then theres the reaction of MnO2 with Nitric acid. I believe there is no Mn(NO3)4 compound so if Im correct then Im guessing it gets reduced to
Mn(NO3)2?
|
|
DJF90
International Hazard
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline
Mood: No Mood
|
|
Treat the permanganate with stoichoimetric MnSO4. The MnO2 that forms is particularly active for the oxidation of allylic/benzylic alcohols to the
corresponding aldehydes/ketones.
|
|
bolbol
Hazard to Others
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 167
Registered: 3-1-2015
Member Is Offline
Mood: No Mood
|
|
Unfortunately I don't have MnSO4 to begin with and In fact I don't have that kind of stuff around.
Somewhere else someone said to make an alkaline solution and treat it with a reducing agent like sugar or ethanol, but he didn't really go into detail
on how exactly I could make an alkaline solution with it
|
|
woelen
Super Administrator
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 7990
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Dissolve some KMnO4 in water, make it as concentrated as possible.
Add this solution to a solution of sugar, to which also some NaOH is added. Soon, MnO2 settles at the bottom. Let the liquid stand for one day, I
expect a dark brown precipitate to settle at the bottom and having a clear solution above it.
Use 1 teaspoon of sugar, half a teaspoon of NaOH and add half a teaspoon of KMnO4, dissolved in water.
MnO2 does not form Mn(NO3)2 with HNO3 though. You also need some reductor. I would try oxalic acid, which is completely oxidized to CO2 and then
evaporate the liquid to get Mn(NO3)2. I do not expect an easy separation, Mn(NO3)2 is very hygroscopic and I am afraid you will have a hard time
getting dry crystals.
|
|
bolbol
Hazard to Others
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 167
Registered: 3-1-2015
Member Is Offline
Mood: No Mood
|
|
So are u saying MnO2 wouldn't react with nitric acid at all?
|
|
bolbol
Hazard to Others
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 167
Registered: 3-1-2015
Member Is Offline
Mood: No Mood
|
|
Alright so I did it in two different ways to see the difference.
For experiment #1
-Dissolved 6 grams of KMnO4 in 200 mls of water.
-Then I added 6 grams of NaOH in it.
-After all was dissolved, I added around 50 mls of 91% isopropyl alcohol
the solution immediately turned a nice green color and then after a while turned into a brown goo... resembling the color of.. well u know what
For experiment #2
-Dissolved 1 gram of KMnO4 in around 50 mls of Acetone (Not all of it dissolved)
-Added 100 mls of water to the mixture
-Added 1 gram of NaOH
The solution turned a dark color, probably dark brown couldnt really tell, but after a while I noticed the solution was clear and I could see what I
think is the MnO2 particles all over the place.
I've added H2O2 to permanganate before and the resulting solution resembled what I got out of experiment #2.
Whats wrong with #1? are there any water insoluble impurities or should I let the solutions sit for a while
|
|
woelen
Super Administrator
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 7990
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Particles of MnO2 can be very fine and it may be hard to isolate it. Boiling a solution for a while may help, it makes precipitates coarser and easier
to filter.
MnO2 will not react with dilute HNO3, with more conc. HNO3 I also doubt that there will be any reaction. Keep in mind that for the MnO2 to turn into a
manganese(II) salt, you need a reductor and with HNO3 only you don't have that. Adding oxalic acid to MnO2 suspended in dilute nitric acid certainly
will lead to reduction of the MnO2, but with HNO3 alone you won't get Mn(NO3)2.
|
|
bolbol
Hazard to Others
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 167
Registered: 3-1-2015
Member Is Offline
Mood: No Mood
|
|
Theres a difference in color of the product that occured which im not sure if its due to the particle size or contaminates but with isopropyl method
my product was brown and with Acetone and H2O2 the product was a much darker brown.
And yes I understand, I will try hot conc HNO3 with MnO2 tho just to confirm that there wont be a reaction
|
|
AJKOER
Radically Dubious
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
You may consider looking at dry cell batteries as a source for MnO2. To quote from Wikipedia
http://en.m.wikipedia.org/wiki/Battery_(electricity)#Dry_cell :
"A standard dry cell comprises a zinc anode, usually in the form of a cylindrical pot, with a carbon cathode in the form of a central rod. The
electrolyte is ammonium chloride in the form of a paste next to the zinc anode. The remaining space between the electrolyte and carbon cathode is
taken up by a second paste consisting of ammonium chloride and manganese dioxide, the latter acting as a depolariser. In some designs, the ammonium
chloride is replaced by zinc chloride."
-------------------------------
Note the following reaction path:
MnO2 + 4 HCl → Cl2 (g) + 2 H2O + MnCl2 [edit correction]
MnCl2 + 2 AgNO3 →2 AgCl (s) + Mn(NO3)2
For the dry salt, unsure if two moles of moist NO2 acting on MnO2 works (see http://www.sciencedirect.com/science/article/pii/S0925400511... and this educational reference site https://www.google.com/url?sa=t&source=web&rct=j&... ).
I will let others comment on feasibility.
[Edited on 27-1-2015 by AJKOER]
|
|
blogfast25
International Hazard
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by AJKOER ![](images/xpblue/lastpost.gif) | Note the following reaction path:
MnO2 + 4 HCl → Cl2 (g) + 2 H2O
MnCl2 + 2 AgNO3 →2 AgCl (s) + Mn(NO3)2
|
AJ and the case of the magically disappearing manganese!
Don't waste silver nitrate on this. Instead precipitate the MnCl2 as MnCO3(s) with Na2CO3 solution. Filter, wash and dry carefully.
MnCO3 is a very stable, slightly pink solid. It does not oxidise to MnO2 easily.
Dissolve it in any relevant dilute acid (including HNO3) to obtain the relevant Mn(II) salt.
[Edited on 27-1-2015 by blogfast25]
|
|
AJKOER
Radically Dubious
![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif) ![*](images/xpblue/star.gif)
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Yes, and the accidental deletion has been corrected via the magical edit function .
I also added an unknown educational link which actual cites my direct action by NO2 on MnO2 forming the nitrate.
The accidental formation of dry Al(NO3)3 via NO2 was, per recollection, once reported on SM as well. Interestingly, as Aluminum/Al2O3 does resist
direct attack by HNO3 to form a nitrate, Aluminum and Manganese oxide may behave similarly with NO2.
[Edit] Found the thread here http://www.sciencemadness.org/talk/viewthread.php?tid=22790 shown below:
Quote: Originally posted by Antiswat ![](images/xpblue/lastpost.gif) |
[/rquote]
What I found interesting was that anhydrous aluminum nitrate can be prepared from AlCl3 in N2O4 at 0 °C.
[Edited on 23-12-2012 by AndersHoveland][/rquote]
A bit off topic, but when ive made Pb(NO3)2 by HNO3 + Pb i have covered my flask with 10 layers of aluminium foil, and apparently the NO2 formed can
and WILL react with the aluminium foil to make aluminium nitrate, infact this reaction is probably taking place right now, just thought its weird that
aluminium wont react with HNO3 but NO2 reacts with aluminium |
Here is also a general reference thanks to Wikipedia on NO2, to quote:
"Conversion to nitrates
NO2 is used to generate anhydrous metal nitrates from the oxides:[6]
MO + 3 NO2 → M(NO3)2 + NO"
Because NO2 is toxic upon inhalation of small quantities, I would recommend working with NO and O2 passed through broken glass (in the ratio of 2 to 1
which is the cited method enabling mixing of gases and NO2 formation ) brought into contact with the MnO2.
One preparation for NO, not even requiring HNO3 is:
2KNO3 + 6FeCl2 + 8HCl→ 2KCl + 6FeCl3 + 2NO + 4H2O
---------------------------------------------------------
With respect to aqueous methods, one can replace the expensive AgNO3 with say Pb(NO3)2, assuming the Lead toxicity is not an issue.
[Edited on 27-1-2015 by AJKOER]
|
|