Sciencemadness Discussion Board

Removing Iron contamination from Sulphuric acid

CHRIS25 - 11-4-2014 at 08:07

Ferrous Sulphate Heptahydrate. As far as I could gather but the only site I found that had information about this was a paid PDF. Has anybody any resources about how I may go about this please.

Zyklon-A - 11-4-2014 at 08:41

I have no references to back this hypothesis, but you might want to add a stoichiometric amount of oxalic acid, which should precipitate Iron(II) oxalate:
FeSO4: 7H2O + H2C2O4 ↔ FeC2O4 + H2SO4 + 7 H2O.
The insolubility of Iron(II) oxalate (0.008 g/100mL @ 20°C) will drive the equilibrium nearly completely to the right.;)
Of course this will only work if you know the exact amount of iron sulfate, and the 7 molecules of water produced will dilute the acid slightly.

[Edited on 11-4-2014 by Zyklonb]

macckone - 11-4-2014 at 09:03

The oxalate method will work.
The another option is distillation.
Distilling sulfuric is best done with a glass vacuum distillation apparatus.
The final option is fractional crystallization of the sulfuric acid.
The crystallization route is probably the easiest without adequate lab gear.

Zyklon-A - 11-4-2014 at 09:07

Assuming the acid is ~98%, I wouldn't try distillation, although possible it would be rather dangerous.

CHRIS25 - 11-4-2014 at 09:26

Well that's a start, I have plenty of oxalic acid. I will do a sample. Distillation is out of the question, I have no suitable apparatus for this. My reactions mimics the sort of situation you get from Mining where there is acid mine drainage pollution. I get Fe(OH)3 precipitated out of solution as it continues to evaporate (in the production of Sodium sulphate, PH is at 7.5 - 8.0). I have to keep filtering every 100 mLs of evaporation at the moment though this will cease because the solution is becoming nicely clear now. Mmm, one thing, why do I need to know the exact amount of Ferrous sulphate? You said use stoichiometric amounts of oxalic and and ferrous sulphate, what did you mean by:...Of course this will only work if you know the exact amount of iron sulfate....

Zyklon-A - 11-4-2014 at 10:59

OK, you don't need to know the exact amount, but if you don't, you'll almost certainly add at least slightly more than the necessary amount, thus contaminating your sulfuric acid with excess oxalic acid. Because you intention is to purify you acid, it will be quite illogical to not measure the iron contaminants.
It would quite easy to find out with reasonable accuracy the amount of iron: Simply add an excess of oxalic acid to a small known quantity of your contaminated acid, then using the method of your choice, to find the amount of iron in the acid, by measuring the precipitated Iron(II) oxalate.

[Edited on 11-4-2014 by Zyklonb]

DraconicAcid - 11-4-2014 at 11:05

Quote: Originally posted by Zyklonb  
I have no references to back this hypothesis, but you might want to add a stoichiometric amount of oxalic acid, which should precipitate Iron(II) oxalate:
FeSO4: 7H2O + H2C2O4 ↔ FeC2O4 + H2SO4 + 7 H2O.
The insolubility of Iron(II) oxalate (0.008 g/100mL @ 20°C) will drive the equilibrium nearly completely to the right.;)
Of course this will only work if you know the exact amount of iron sulfate, and the 7 molecules of water produced will dilute the acid slightly.

[Edited on 11-4-2014 by Zyklonb]


I find this suspicious, because the acidity of the solution will prevent complete ionization of the oxalic acid. I haven't tried that specific reaction, but I do remember doing a lab in which we used sulphuric acid of various concentrations to destroy an iron(III) oxalate complex. The dilute sulphuric acid protonated off the oxalate; the concentrated sulphuric acid decomposed the oxalate to give CO2, CO, and water.

CHRIS25 - 11-4-2014 at 11:07

Mm, the method of my choice - I don't have any. But I suppose the only one I can think of is to use Potassium Permanganate to oxidize iron 2 to iron3. I will have to read up on this. Secondly since this is a strong oxidzing agent it might well oxidize other components as well, or am I making life complicated?

Zyklon-A - 11-4-2014 at 11:08

Ah, well like I said, it's entirely hypothetical. Experiment with small quantity's of acid, and be prepared in case CO is released!
[EDIT]
Quote: Originally posted by CHRIS25  
Mm, the method of my choice - I don't have any. But I suppose the only one I can think of is to use Potassium Permanganate to oxidize iron 2 to iron3. I will have to read up on this. Secondly since this is a strong oxidzing agent it might well oxidize other components as well, or am I making life complicated?

Depending on the temperature, this could lead to manganses heptoxide! I would not try that!

[Edited on 11-4-2014 by Zyklonb]

DraconicAcid - 11-4-2014 at 11:13

Quote: Originally posted by CHRIS25  
Mm, the method of my choice - I don't have any. But I suppose the only one I can think of is to use Potassium Permanganate to oxidize iron 2 to iron3. I will have to read up on this. Secondly since this is a strong oxidzing agent it might well oxidize other components as well, or am I making life complicated?


Then you'd have manganese impurities instead of iron ones.

CHRIS25 - 11-4-2014 at 11:19

I guess I will abandon this. I do not have the knowledge to measure iron as suggested. I have been doing quite a lot of reading already about all this and find nothing except references to industrial and environmental documents and one paid PDF, I would not know how to measure Iron as suggested after precipitating the iron oxalate, and the potassium permanganate was off the beaten track. I will simply buy lab grade and be done with it.

[Edited on 11-4-2014 by CHRIS25]

Zyklon-A - 11-4-2014 at 12:12

Just ask for help, you'll need to learn stoichiometry at some point anyway.
Not knowing isn't an excuse to give up!

[Edited on 11-4-2014 by Zyklonb]

CHRIS25 - 11-4-2014 at 12:32

Well good news is I am very comfortable with stoichiometry and titrating. It's just that I do not understand how to measure the iron oxalate precipitate. If I know how much H2SO4, how much oxalic acid I added I do not grasp how to do the following:
H2SO4 + C2H2O4 = ??
17.5M + ??M
98g/mol + 90g/mol

So in the agony of a lack of education here I wrote the above. But If you have the time then I would like to learn. My only titrations that I do are copper chloride - finding out how much H+ free ions are in a solution, measuring the amount of Sodium Oxide in Sodium silicate, finding exact molarities in acids, very limited experiences I am afraid.

[Edited on 11-4-2014 by CHRIS25]

gdflp - 11-4-2014 at 12:38

What you want to do is accurately measure out some known volume of contaminated sulfuric acid, then dissolve and add a large excess of oxalic acid. Weigh some filter paper, filter the iron oxalate through it, dry it, and accurately weigh the filter paper again. You can then figure out how much iron was present by turning this into moles to determine what concentration the iron is in the solution?

CHRIS25 - 11-4-2014 at 13:04

I am not known for common sense. So thankyou. It was when you said ..turn it into moles... that I kicked myself.

aga - 11-4-2014 at 13:30

When we are born, we can breathe, shit and piss in copious amounts.
Babies are born with a certain skill at Vomiting.
Eventually we seem to learn to walk and talk by some happy accident.

After that, the rest is ALL a learning process.

Even the most able cutting edge scientist once knew nothing at all.

For a male, the two main barriers to learning are Ego and Machismo.
Ego tries to think it already knows it all.
Machismo tries to convince it's mating partners/competitors that it knows it all.
Either state prevents further learning.

You seem to be better than that, and so you will truly learn.

blogfast25 - 11-4-2014 at 13:41

Quote: Originally posted by DraconicAcid  

I find this suspicious, because the acidity of the solution will prevent complete ionization of the oxalic acid.


Precisely. Concentrated sulphuric acid will displace oxalic acid from just about any oxalate. That scheme cannot work.

TBH, removing ferrous sulphate from conc. H2SO4 may be seriously difficult and simply not worth trying. Distillation should work but distilling conc. H2SO4 is no sinecure. Perhaps diluting it slightly and then distilling it back to full strength, leaving behind the ferrous sulphate, could be contemplated.

A company I worked for some 30 years ago tried something similar for years: separating 40 % H2SO4 from 4 - 5 % ferrous sulphate (a waste stream). They gave up in the end.

[Edited on 11-4-2014 by blogfast25]

DraconicAcid - 11-4-2014 at 13:44

Quote: Originally posted by blogfast25  
Distillation should work but distilling conc. H2SO4 is no sinecure. Perhaps diluting it slightly and then distilling it back to full strength, leaving behind the ferrous sulphate, could be contemplated.


If you're distilling sulphuric acid, doesn't the water boil off first, leaving more concentrated, high-boiling acid?

I'd suggest saving that acid for things in which iron won't interfere- use it to distill hydrochloric acid from salt, nitric acid from nitrates, dehydrating organic compounds, etc.

CHRIS25 - 11-4-2014 at 13:57

Well it's too late for experimentation, but some further research turned up the fact that Yes Concentrated Sulphuric acid decomposes oxalic acid, where the stress seemed to be on "concentrated". Secondly just for attempting something I have not really tried before, I did the following based on limited knowledge of course, I tried to balance a theoretical situation out based on the presence of Iron:
I am sure you will tell me otherwise, but it was an attempt.
Fe + H2C2O4 + H2SO4 + 3H2O = 2CO2 + 2H2 + SO4 + 2Fe(OH)3 (precipitate)

[Edited on 11-4-2014 by CHRIS25]

hyfalcon - 12-4-2014 at 03:04

Quote: Originally posted by gdflp  
Weigh some filter paper, filter the iron oxalate through it, dry it, and accurately weigh the filter paper again.


This is one thing the OP isn't going to be able to do with concetrated sulfuric acid. One, it's very viscous. Two, concentrated sulfuric acid will eat right through filter papaer. He needs a glass frit if he's going to be filtering high concentration sulfuric.

[Edited on 12-4-2014 by hyfalcon]

Zyklon-A - 12-4-2014 at 05:28

Not necessarily, concentrated sulfuric acid is not very viscous, that's 100% acid. And besides, he could just dilute it slightly, it wouldn't affect the reaction at all.

blogfast25 - 12-4-2014 at 05:52

DA:

Ooopsie. You are right.

Conc. H2SO4 will chew up filter paper, no matter how 'acid resistance enhanced' it may be advertised as, to a black goo. You need other filtering media for that.
Quote: Originally posted by CHRIS25  

Fe + H2C2O4 + H2SO4 + 3H2O = 2CO2 + 2H2 + SO4 + 2Fe(OH)3 (precipitate)

[Edited on 11-4-2014 by CHRIS25]


Fe(OH)3 would never precipitate in those circumstances. Which leads me to another remedy for your Na2SO4/Fe(OH)3 woes. Depending a bit on concentration, Fe<sup>3+</sup> starts dropping out of solution (as Fe(OH)3) from roughly pH 3 and higher. So if you keep your neutralised solution slightly acidic, the Fe(OH)3 will not form and your Na2SO4 will be slightly contaminated with Fe2(SO4)3. To get rid, redissolve your crude Na2SO4, the Fe(OH)3 will then form. Filter or decant and isolate fairly pure Na2SO4.

macckone - 12-4-2014 at 06:41

Fractional crystallization seems so much easier at this point.
It is one of the recommended ways of cleaning up sulfuric acid.
It requires nothing more than a freezer, a glass container,
and something to filter it through. Pre-cleaned fiberglass
cloth seems appropriate. The fiberglass cloth you can buy
at hardware stores usually contains a binder that can be
removed with various solvents. Obviously the solvent
depends on the binder, but I have had successes with
acetone and toluene.

purification of laboratory chemicals (you may need to scroll a little):

http://books.google.com/books?id=4ViVUQi7Z60C&pg=PA618&a...

CHRIS25 - 12-4-2014 at 07:19

Macckone, It means buying more stuff, not really worth it, best spent on new H2SO4. But thanks anyway. Blogfast, the Iron hydroxide does drop out as you say. The PH level of my solution stays between 7.3 and 8.1. That explains why it successfully drops out already, but only upon heating. I think at this stage just buy another batch from lab supplier instead and keep this one and experiment with what has already been said, then use it for non-essential things, whatever they may be.

deltaH - 12-4-2014 at 08:11

You might want to consider passing your solution through a column filled with beads of the protonated form of a strongly acidic ion exchange resin to remove iron impurities (I think manufactures may even have specific resins for such an application). These resins absorb polyvalent cations very strongly and replace them with protons, raising the acidity of your solution. You will need to take care to watch that the last bit of your resin bed has not yet saturated with iron. The resultant solution will then be one of dilute acid which you can then concentrate up by boiling off the water.

However, simplicity and economics probably dictate you to simply dispose of the material safely and buy more as the acid is so cheap :)

[Edited on 12-4-2014 by deltaH]

CHRIS25 - 12-4-2014 at 09:48

:
Quote: Originally posted by deltaH  
You might want to consider passing your solution through a column filled with beads of the protonated form of a strongly acidic ion exchange resin to remove iron impurities

However, simplicity and economics probably dictate you to simply dispose of the material safely and buy more as the acid is so cheap :)

[Edited on 12-4-2014 by deltaH]

Hallo, yes well after careful deliberation and much planning I decided to opt for your second suggestion:)

[Edited on 12-4-2014 by CHRIS25]

[Edited on 12-4-2014 by CHRIS25]

deltaH - 12-4-2014 at 10:38

yeah... hard choice ;) "money talks, bullshit walks"

AJKOER - 12-4-2014 at 17:51

Quote: Originally posted by DraconicAcid  


I find this suspicious, because the acidity of the solution will prevent complete ionization of the oxalic acid. I haven't tried that specific reaction, but I do remember doing a lab in which we used sulphuric acid of various concentrations to destroy an iron(III) oxalate complex. The dilute sulphuric acid protonated off the oxalate; the concentrated sulphuric acid decomposed the oxalate to give CO2, CO, and water.


Per an old thread on the preparation of H2SO4 from the action of Oxalic acid on a sulfate (see https://www.sciencemadness.org/whisper/viewthread.php?tid=18... ), it was indeed observed in an attempt to prepare very concentrated Sulfuric acid that a sudden violent ejection of the H2SO4 was observed. This confirms that high strength H2SO4 does indeed decompose H2C2O4 in apparently an unsafe manner.

The general recommendation from that thread is that it is safest to prepare mineral acids using Oxalic acid by targetting only dilute to moderate concentrations.

[Edited on 13-4-2014 by AJKOER]

CHRIS25 - 13-4-2014 at 00:31

AJkoer, hallo, yes I read this old thread two days ago, this is where I got the general impression that a dilute solution would work. Have yet to experiment, tied up with two other things at the moment. But i will try The oxalic acid method soon and let people know about the results.

CHRIS25 - 14-4-2014 at 04:43

Added oxalic acid to 10mLs H2SO4 in 50mLs water. No precipitate. Does nothing, as has been suggested.