Sciencemadness Discussion Board - Ferrous Oxalate precipitation

Ferrous Oxalate precipitation

veerendra - 17-1-2014 at 20:39

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.

Brain&Force - 17-1-2014 at 20:52

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.

blargish - 17-1-2014 at 21:04

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]

elementcollector1 - 17-1-2014 at 21:04

Ferrous oxalate is actually soluble in excess oxalic acid, forming the triferrooxalato complex.

Zephyr - 17-1-2014 at 21:17

Quote: Originally posted by elementcollector1

Ferrous oxalate is actually soluble in excess oxalic acid, forming the triferrooxalato complex.

I had this same problem, but on adding an exess of iron sulfate it all precipitated nicely. good luck!

I tried but not succeded

veerendra - 18-1-2014 at 03:56

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

cyanureeves - 18-1-2014 at 09:44

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]

Hexavalent - 18-1-2014 at 10:15

Quote: Originally posted by veerendra

what does excess oxalic acid means?

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]

cyanureeves - 18-1-2014 at 13:41

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]

AJKOER - 18-1-2014 at 14:51

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]

blogfast25 - 19-1-2014 at 10:55

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) < === > M2+(aq) + Ox2-(aq) (for a divalent metal) Ksp = [M2+] x [Ox2-]

Ox2- + 2 H3O+ < === > H2Ox + 2 H2O, the combined Ka1 and Ka2.

In the presence of a strong Bronsted acid (stronger than oxalic acid), the [Ox2-] is suppressed because the stronger acid pushes the second equilibrium to the right, assuming it is present in sufficient quantity. But if the remaining [Ox2-] 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]

DraconicAcid - 19-1-2014 at 11:01

Quote: Originally posted by blogfast25

Actually, the second equilibrium is irrelevant. The relevant equilibrium is:

Ox2- + H3O+ < === > HOx^- + H2O

blogfast25 - 19-1-2014 at 11:03

Quote: Originally posted by DraconicAcid

Actually, the second equilibrium is irrelevant. The relevant equilibrium is:

Ox2- + H3O+ < === > HOx^- + H2O

Assuming the hydrogenoxalates are soluble, yes.

How to leach out Fe from a mixture of Mn and Fe oxides

veerendra - 21-1-2014 at 09:36

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 ?

elementcollector1 - 21-1-2014 at 11:29

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?

bismuthate - 21-1-2014 at 12:52

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?

blogfast25 - 21-1-2014 at 12:57

Quote: Originally posted by veerendra

Can any body know how to do this ?

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]

What if we use only oxalic acid without any other acid

veerendra - 21-1-2014 at 23:45

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?

blogfast25 - 22-1-2014 at 06:21

Quote: Originally posted by veerendra

Dear sir

Thanks for reply.

But if I use only oxalic acid and dissolve the mixture than what can happen ?

You need to write the redox reaction equation for:

Reduction: MnO2 === > Mn2+

Oxidation: H2C2O4 === > 2 CO2

You'll see that it CANNOT proceed without an additional acid, usually sulphuric in this case.

[Edited on 22-1-2014 by blogfast25]

Boffis - 22-1-2014 at 16:13

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.

Mn and Fe oxide mixtureleaching

veerendra - 23-1-2014 at 02:04

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 ?

blogfast25 - 23-1-2014 at 05:53

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]

[merged 3 topics]

bfesser - 23-1-2014 at 06:09

veerendra, please don't open new topics to ask the same or closely related questions.

Quote: Originally posted by veerendra

Source of iron was iron oxide from rocks/ores.

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]