veerendra
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Ferrous Oxalate precipitation
How precipitate the ferrous oxalate dissolved in the system ?
I dissolved iron bearing material in oxalic acid solution and got the ferrous oxalte, now I want to recover that by simplest method.
I tried alcohal but it did not precipitaed ? If I use NaOH etc than it might precipitae in the form of hydroxides.
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Brain&Force
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Ferrous oxalate isn't supposed to be soluble at all...it should have precipitated.
What was your Fe source? Different steels have different amounts of iron and different components that may make separation a challenge.
Also, the type of alcohol can make a big difference. Methanol and ethanol can often precipitate salts, but long-chain alcohols just refuse to dissolve
in the solution entirely. Don't use rubbing alcohol for separation, that's isopropanol.
At the end of the day, simulating atoms doesn't beat working with the real things...
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blargish
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Try heating the solution.
I have made ferrous oxalate before by dissolving iron in hydrochloric acid to form ferrous chloride, and then adding oxalic acid solution. I got a
yellow solution, but no precipitate, which is odd since ferrous oxalate is supposed to be practically insoluble; however, after heating the solution
(not too hot to prevent decomp of the oxalate) a yellow precipitate formed.
I'm not sure whether this was because the reaction needed heating to proceed to completion, or whether excess HCl was somehow keeping the ferrous
oxalate in solution and needed to be driven off.
[Edited on 18-1-2014 by blargish]
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elementcollector1
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Ferrous oxalate is actually soluble in excess oxalic acid, forming the triferrooxalato complex.
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Zephyr
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I had this same problem, but on adding an exess of iron sulfate it all precipitated nicely. good luck!
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veerendra
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I tried but not succeded
Thanks to all of you.
Source of iron was iron oxide from rocks/ores. I has a yellow solution - I heated it but no precipitation, I add ethenol but it not precipiated
yet.
what does excess oxalic acid means ?
Thanks
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cyanureeves
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a video on youtube states ferrous iron(pyrophoric iron) and shows it decomposing under heat in a test tube and bursting into sparks as the fine iron
is poured out.now ferrous oxalate is not pyrophoric in itself and can be safely stored cant it?i tried it with homemade iron chloride and got a dark
orange precipitate and i tried it with garden copperas sulfate and got a light yellow precipitate as in the video.wiki shows bright orange crystals
yet zhmapper's video shows a light yellow and in fact does make pyrophoric iron.either or i just dont want anything spontaneously combusting in my
shop like michael jacksons hair.i had left over copper chloride and decided to make the red copper acetilyde and let me tell you when heated the stuff
goes BOOSH!and the residue smoke induces a nasty headache.
[Edited on 1-18-2014 by cyanureeves]
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Hexavalent
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It means a greater than stoichiometric amount of oxalic acid needed for the reaction.
For example, if you are using 1 mole of iron (II) chloride, according to the balanced chemical equation using anything more than 1 mole of oxalic acid
is considered an excess;
FeCl2 + H2C2O4 → FeC2O4 + 2HCl
[Edited on 18-1-2014 by Hexavalent]
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
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cyanureeves
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we are having beautiful weather out here that my iron oxalate is completely dry and my homemade iron chloride worked perfectly with oxalic acid and
just witnessed "iron fireflies".i heated the mixture in a test tube and poured out the powder in my garage with doors shut and in the dark and sparks
trickled down.boy i'm having so much fun today i even tried making ammonium ferrous sulfate with 3% peroxide and 15% household ammonia.now will the
cold iron powder still be pyrophoric or will it be pyrophobic?i will soon find out.update:nope it is pyrophobic when cold!it needs more heat and a
xanax.isnt pyrophobic misleading then?i mean it did come in contact with oxygen and nothing happened.oh well.it's been a good day.
[Edited on 1-18-2014 by cyanureeves]
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AJKOER
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I recall reading in Watt's Dictionary Chemistry that many insoluble oxalates will dissolve in weakly acidic solutions. So heating a solution with
excess HCl could drive off the HCl, neutralizing the solution, allowing the oxalate to precipitate.
Now, in the case of an Iron salt and an excess of H2C2O4, the more accurate explanation relates to the formation of a complex.
[Edited on 18-1-2014 by AJKOER]
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blogfast25
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Quote: Originally posted by AJKOER  | | I recall reading in Watt's Dictionary Chemistry that many insoluble oxalates will dissolve in weakly acidic solutions. |
It depends largely on the solubility product of the oxalate in question.
Several competing equilibria rule:
MOx(s) < === > M<sup>2+</sup>(aq) + Ox<sup>2-</sup>(aq) (for a divalent metal) Ksp = [M<sup>2+</sup>] x
[Ox<sup>2-</sup>]
Ox<sup>2-</sup> + 2 H<sub>3</sub>O<sup>+</sup> < === > H<sub>2</sub>Ox + 2
H<sub>2</sub>O, the combined Ka1 and Ka2.
In the presence of a strong Bronsted acid (stronger than oxalic acid), the [Ox<sup>2-</sup>] is suppressed because the stronger acid
pushes the second equilibrium to the right, assuming it is present in sufficient quantity. But if the remaining [Ox<sup>2-</sup>] is
still high enough to exceed the Ksp, then the oxalate will not dissolve.
From the equibrium constants, mass and neutrality balances can be predicted which oxalate will dissolve in which concentration of a strong acid.
[Edited on 19-1-2014 by blogfast25]
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DraconicAcid
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| Quote: Originally posted by blogfast25 |
It depends largely on the solubility product of the oxalate in question.
Several competing equilibria rule:
MOx(s) < === > M<sup>2+</sup>(aq) + Ox<sup>2-</sup>(aq) (for a divalent metal) Ksp = [M<sup>2+</sup>] x
[Ox<sup>2-</sup>]
Ox<sup>2-</sup> + 2 H<sub>3</sub>O<sup>+</sup> < === > H<sub>2</sub>Ox + 2
H<sub>2</sub>O, the combined Ka1 and Ka2.
In the presence of a strong Bronsted acid (stronger than oxalic acid), the [Ox<sup>2-</sup>] is suppressed because the stronger acid
pushes the second equilibrium to the right, assuming it is present in sufficient quantity. But if the remaining [Ox<sup>2-</sup>] is
still high enough to exceed the Ksp, then the oxalate will not dissolve. |
Actually, the second equilibrium is irrelevant. The relevant equilibrium is:
Ox<sup>2-</sup> + H<sub>3</sub>O<sup>+</sup> < === > HOx- + H<sub>2</sub>O
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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blogfast25
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| Quote: Originally posted by DraconicAcid |
Actually, the second equilibrium is irrelevant. The relevant equilibrium is:
Ox<sup>2-</sup> + H<sub>3</sub>O<sup>+</sup> < === > HOx- + H<sub>2</sub>O
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Assuming the hydrogenoxalates are soluble, yes.
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veerendra
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How to leach out Fe from a mixture of Mn and Fe oxides
I have mixture of MnO2 and Fe2O3. I wish to recover Fe as much as possisble.
I am trying oxalic acid.
Can any body know how to do this ?
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elementcollector1
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1) Dissolve mix in H2SO4 or HCl (warning: chlorine gas).
2) Separating about a third of the solution, add bleach and sodium hydroxide. Filter the precipitated hydroxides, wash free of sodium ions, and then
add back into the mother liquor.
3) Over time, the solution should fade to pink or clear, with a brown/orange precipitate. The brown precipitate is your iron hydroxide, and the
pink/clear solution is manganese ion.
Any questions?
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bismuthate
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Add it to HCl.
Add H2O2
Boil it down.
Dissolve the FeCl3 in acetone.
And to answer your question, yes I can know how to do this  .
How would you use oxalic acid?
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blogfast25
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Oxalic acid is oxidised by MnO2 to CO2 and reduced to Mn(II), in the presence of acid. This reaction proceeds well and is quite exothermic, in my
experience.
The problem is that part of the iron will also solubilise, due to the presence of both oxalic acid and another acid (sulphuric acid is needed to aid
the redox reaction). Just how much iron you will 'lose' will depend mainly on how 'hard' the Fe2O3 is. If it's a commercial grade of ferric oxide it
may only solubilise a tiny bit.
[Edited on 21-1-2014 by blogfast25]
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veerendra
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What if we use only oxalic acid without any other acid
Dear sir
Thanks for reply.
But if I use only oxalic acid and dissolve the mixture than what can happen ?
Which will report in solution and how can we precipate that?
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blogfast25
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You need to write the redox reaction equation for:
Reduction: MnO2 === > Mn<sup>2+</sup>
Oxidation: H2C2O4 === > 2 CO<sub>2</sub>
You'll see that it CANNOT proceed without an additional acid, usually sulphuric in this case.
[Edited on 22-1-2014 by blogfast25]
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Boffis
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What is your source material? I have extracted Mn from natural soft earthy black Mn-Fe oxide called "Wad" in the UK. It's easy, you simply grind it to
a powder (its usually pretty soft) and dissolve in any weak or dilute acid with the aid of hydrogen peroxide. Add sulphate ions either as sulphuric
acid or sodium sulphate to precipitate the barium (common contaminant in natural Mn oxides) and filter. It the resulting solution is clear or pink you
simply add excess NaOH and blow air through it or add more hydrogen peroxide to oxidize the white Mn(OH)2 to dark brown Mn4+ hydrated oxides. If the
initial solution is distinctly pink check it for cobalt, this is a fairly common component of natural "wad" and "psilomelane" and complicates the
process. Even 20% acetic acid works though its slow.
It depends what you want the Mn oxide for and what your source is. Some natural Mn oxides are rich in Mn3O4 and braunites which is rather inert though
hot mineral acid and H2O2 will still work.
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veerendra
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Mn and Fe oxide mixtureleaching
Most of the experts say that Mn will not be leached by only oxaluc acid but when I am leaching a mixture of Mn andFe in oxalic acidthe fliterate is
gray color and conatin more Mn oxalte dihydrate why ?
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blogfast25
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Describe better what you actually did. Quantities and such like.
If you used an excess of oxalic acid there will be no need for sulphuric acid, I guess. But like most oxalates, Mn(II) oxalate is poorly soluble in
water: 0.033 g / 100 g water at 30 C, acc. Wiki.
This thread should be merged with the original one.
[Edited on 23-1-2014 by blogfast25]
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bfesser
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Threads Merged
23-1-2014 at 06:07 |
bfesser
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[merged 3 topics]
veerendra, please don't open new topics to ask the same or closely related questions. Could you be more specific as to which ore you are trying to
extract? Do you know the precise mineralization? Is it primarily carbonate, oxide/hydroxide, etc.? Also, knowing which <a
href="http://en.wikipedia.org/wiki/Iron_oxide" target="_blank">iron oxide</a> <img src="../scipics/_wiki.png" /> would help.
<a href="http://en.wikipedia.org/wiki/Category:Iron_minerals" target="_blank">Fe minerals</a> <img src="../scipics/_ext.png" />
<a href="http://en.wikipedia.org/wiki/Category:Manganese_minerals" target="_blank">Mn minerals</a> <img src="../scipics/_ext.png" />
[Edited on 23.1.14 by bfesser]
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