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Author: Subject: Drying Manganese Nitrate
Photonic
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[*] posted on 24-9-2018 at 18:27
Drying Manganese Nitrate


Does anyone have a procedure for preparing Mn(NO3)2 crystals?

I prepared a small quantity of it ~<5g and am having trouble getting it to crystallize. "dry out". I believe there may also be a small amount of residual nitric acid as it certainly fumes lightly when air is blown on it at temperature.

I believe the wikipedia stated boiling point is also incorrect. I have been unable to observe it boiling at temperatures to 160C.

The nitrate was observed decomposing on metal surface at temperatures in the 220C-250C range.

Heating the solution to approximately 140-160C on a hot plate produces a viscious light pink liquid. I do not have a particular use for this at this point, so I would like to make crystalline Mn(NO3)2 for my collection.

My other remaining idea is heating with a stream of air, or getting a desiccator to try out, however, I haven't found any procedures in books, or by googling to reference for bp, mp, etc.

Thanks kindly,
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Deathunter88
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[*] posted on 24-9-2018 at 19:24


As a general rule, don't heat nitrate salts of transition metals to dry them, since they will just dissolve in their own water of crystallization. Instead, put the beaker into a Ziploc bag with another beaker of sodium hydroxide, and let it sit until the solution is dry. In your case, you need to add some more water to the syrup and use the above procedure to get nice crystals.
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JScott
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[*] posted on 26-9-2018 at 04:18


I have run into similar difficulties with Nickel salts.

You may find some of the advice I was given in this thread to be of use. It helped me a great deal.

Good luck!
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mayko
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[*] posted on 26-9-2018 at 08:51


Did you observe the formation of a dark brown solid in your Mn(NO3)2 solution? I've had some sitting around, but every time I try heating it down, a dark, fine precipitate appears, even after multiple filtrations. My best guess was MnO2, though I couldn't imagine how it would be forming, and a test with H2O2 didn't form bubbles and did dissolve the solid, so I think that possibility can be excluded.



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[*] posted on 26-9-2018 at 16:09


Quote: Originally posted by mayko  
Did you observe the formation of a dark brown solid in your Mn(NO3)2 solution? I've had some sitting around, but every time I try heating it down, a dark, fine precipitate appears, even after multiple filtrations. My best guess was MnO2, though I couldn't imagine how it would be forming, and a test with H2O2 didn't form bubbles and did dissolve the solid, so I think that possibility can be excluded.


Negative. You do have to be cautious with the temperature, though. Decomposition was not noticed until approximately 220C on a metal surface. Glass was heated to approximately 160C for ~6 hours with no decomposition visible.

Manganese nitrate was prepared with all high grade chemicals obtained from a chemical supplier. 99.98% MnO2, Technical grade oxalic acid, and reagent grade 68% HNO3.
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AJKOER
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[*] posted on 28-9-2018 at 17:30


Generally, a transition metal in one of its lower valence states in the presence of oxygen and H+ lends itself to an electrochemical reaction. For example, with Fe(ll) (see https://wwwbrr.cr.usgs.gov/projects/GWC_coupled/phreeqc/html... and relatedly, comment at https://pubs.acs.org/doi/10.1021/es0501058 ):

Fe(ll) + 1/4 O2 + H+ = Fe(lll)+ 1/2 H2O

although, I prefer the format below as the reaction can apparently result in the formation of a basic salt (like with cuprous forming a basic cupric salt, see Eq (7) at https://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide):

Fe(ll) + 1/4 O2 + 1/2 H+ = Fe(lll)+ 1/2 OH-

Or, upon rescaling:

4 Fe(ll) + O2 + 2 H+ = 4 Fe(lll)+ 2 OH-

Now, in the case of Mn with oxidation states of 2+ to 7+ (see Table 3 at https://chem.libretexts.org/Textbook_Maps/Inorganic_Chemistry/Supplemental_Modules_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_B lock/3_d-Block_Elements/1b_Properties_of_Transition_Metals/Oxidation_States_of_Transition_Metals ), yes you can get MnO2 and more...Note, heating Mn(NO3).xH2O forms an acidic solution in contact with air/oxygen.

I known the equation above is valid for Fe and Cu metals, and per my recollection Co also, but now apparently add Mn.
-----------------------------------------------------------

I have previously derived an underlying radical based chemistry to this reaction based on the supplement, "Impacts of aerosols on the chemistry of atmospheric trace gases: a case study of peroxides radicals"', by H. Liang1, Z. M. Chen1, D. Huang1, Y. Zhao1 and Z. Y. Li, link: https://www.google.com/url?sa=t&source=web&rct=j&... :

R24 O2(aq) + Cu+ → Cu2+ + •O2− ( k = 4.6xE05 )
R27 •O2− + Cu+ + 2 H+ → Cu2+ + H2O2 ( k = 9.4xE09 )
R25 H2O2 + Cu+ → Cu2+ + •OH + OH− ( k= 7.0 xE03 )
R23 •OH + Cu+ → Cu2+ + OH− ( k = 3.0×E09 )

Net reaction: O2 + 4 Cu+ + 2 H+ → 4 Cu2+ + 2 OH-

Interestingly, the slowest reaction is R25, which is also a fenton-like reaction with a known pH dependence.

Also, I suspect that the above radical reaction sequence would be valid for Cr, Ru and Ce metal salts also (see file of 'Review of iron-free fenton-like systems...' at http://www.sciencemadness.org/talk/viewthread.php?tid=91815#... ), so don't try heating any of these transition metal nitrate hydrates either.

[Edited on 29-9-2018 by AJKOER]
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