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CHRIS25
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Removing Iron contamination from Sulphuric acid
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.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Zyklon-A
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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]
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macckone
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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.
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Zyklon-A
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Assuming the acid is ~98%, I wouldn't try distillation, although possible it would be rather dangerous.
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CHRIS25
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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....
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Zyklon-A
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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]
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DraconicAcid
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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.
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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CHRIS25
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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?
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Zyklon-A
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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]
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DraconicAcid
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| 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.
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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CHRIS25
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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]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Zyklon-A
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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]
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CHRIS25
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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]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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gdflp
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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?
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CHRIS25
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I am not known for common sense. So thankyou. It was when you said ..turn it into moles... that I kicked myself.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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aga
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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.
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blogfast25
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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]
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DraconicAcid
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| 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.
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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CHRIS25
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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]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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hyfalcon
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| Quote: Originally posted by gdflp | Weigh some filter paper, filter the iron oxalate through it, dry it, and accurately weigh the filter paper again.
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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]
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Zyklon-A
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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.
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blogfast25
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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.
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.
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macckone
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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...
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CHRIS25
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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.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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deltaH
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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]
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