sasan
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troubles with MnO2?
Hi every one.i am new here.I have a lot of troubles with MnO2
at first it is undeniable that treating this powder with hydrochloric acid will produce MnCl2.nH2O.but I try to do this.I used lab grade MnO2 that is
shiny black not commercial ones that are brownish and contains iron oxides.and treated it with hydrochloric acid and this was also lab grade and
colorless.at the end the color of solution was brown not pinkish colorless due to the color of Mn2+
MnO2+4HCl(aq)=MnCl2(aq)+Cl2(g)+2H2O
what has happened?
and the other question is what is the products of using HF(aq) instead of HCl(aq)?can this reaction releases flourine gas?
MnO2+4HF(aq)=MnF2(aq)+F2(g)+2H2O http://www.sciencemadness.org/talk/images/smilies/cool.gif
or maybe producing MnOF2??
I am a amateur please give your suggestions
<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: removed
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[Edited on 23.2.14 by bfesser]
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blogfast25
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Hi Sasan,
It helps to use the search facility because this is a subject that has already been clobbered to death here.
If your MnO2 has been hard calcined, it may be much less responsive to HCl (which concentration are you using?) In that case the brown is thus likely
to be simply fine, unreacted MnO2. Try prolonged boiling with excess acid.
Mn<sup>2+</sup> reveals its pink colour only when its quite concentrated and quite pure. E.g. pure crystals of MnCl2 hydrate are pink, but
solutions often appear colourless.
With HF the situation is quite different. Mn(IV) isn't capable of oxidising fluoride anions the way it easily oxidises chloride anions. What you will
obtain will largely depend what kind of HF you use. Aqueous HF, despite it's tremendous toxicity, is quite a weak acid and will probably not attack
your MnO2. Anhydrous HF is so dangerous you shouldn't even contemplate it as a beginner/amateur.
[Edited on 22-2-2014 by blogfast25]
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sasan
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Thank you blogfast. I was using cocentrated HCl(aq) 35-37%
If the brown color of the solution was because of the unreacted MnO2,then the filtration was maybe useful.but if filtration couldnt change the
color,then the brown tint was not the unreacted MnO2
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blogfast25
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Quote: Originally posted by sasan  | Thank you blogfast. I was using cocentrated HCl(aq) 35-37%
If the brown color of the solution was because of the unreacted MnO2,then the filtration was maybe useful.but if filtration couldnt change the
color,then the brown tint was not the unreacted MnO2
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Did you use an excess of acid and tried prolonged boiling?
Also, superfine particles may not be so easy to catch on a filter. Fine particles running through the pores of a filter is often the norm, rather than
the exception and to remedy it the filtrate often has to be run over the same filter several times, until the pores are sufficiently clogged.
It remains difficult to explain the brown tint otherwise as you used a reagent grade and few ions are brown. Unreacted MnO2 is still a likely
candidate.
[Edited on 23-2-2014 by blogfast25]
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bfesser
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Thread Moved
23-2-2014 at 05:49 |
sasan
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yes the acid was in excess and I didnt boil it
Chemistry of manganese is somewhat complicated.i think the brown tint is because of the high concentration of the manganese ion(if you make a
concentrated solution of MnSO4.nH2O it will be brownish pink) or maybe brown tint is due to the color of Mn3+(its color is dark brown and in the low
concentration,would make the solution light brown,ofcourse its too unstable and it is exist just in very acidic environment{the excess acid})
Very low concentration of Mn3+ is more suspecious
sorry for my bad english language,Im persian
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blogfast25
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| Quote: Originally posted by sasan | yes the acid was in excess and I didnt boil it
Chemistry of manganese is somewhat complicated.i think the brown tint is because of the high concentration of the manganese ion(if you make a
concentrated solution of MnSO4.nH2O it will be brownish pink) or maybe brown tint is due to the color of Mn3+(its color is dark brown and in the low
concentration,would make the solution light brown,ofcourse its too unstable and it is exist just in very acidic environment{the excess acid})
Very low concentration of Mn3+ is more suspecious
sorry for my bad english language,Im persian |
Yes, manganese chemistry is very rich and colourful. See also the 'sticky' thread on permanganates in the main chemistry section.
Pure manganese (II) salts in dilute solutions are almost colourless. Mn (III) is much more intense: wine red even at fairly low concentrations. But
with HCl you don't get Mn(III), you do with conc. sulphuric acid, though.
Mn(III) oxidises chloride immediately to chlorine, there's a thread of mine on it somewhere.
I remain convinced that on prolonged simmering your solution will eventually clear up, provided there is reserve HCl.
[Edited on 23-2-2014 by blogfast25]
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