Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
Iron Oxalate
In my efforts to create pyrophoric iron by decomposing iron oxalate dust to iron micro-particles and carbon dioxide, I found that I needed a source of
this salt.
I firstly prepared a solution of iron (II) chloride by reacting steel wool with HCl and filtering, and then made up a stoichiometric solution of
oxalic acid. I dumped the oxalic solution into the filtered iron chloride solution, expecting iron (II) oxalate to precipitate out due to its extreme
insolubility in water.
But nothing happened.
The iron salt solution was pale green and the oxalic acid solution colourless. When mixed, they formed a urine-coloured mixture which was clear.
My first thought was that there may have been a small amount of residual HCl in my iron chloride solution and somehow it was preventing precipitation
(complexing possibly?) of the iron (II) oxalate, so I stuck the mixture on a hotplate and heated to drive it off until no more white fumes were given
off.
Still no precipitate.
Any ideas?
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
nezza
Hazard to Others
Posts: 324
Registered: 17-4-2011
Location: UK
Member Is Offline
Mood: phosphorescent
|
|
I would suspect that the mixture is too acidic for precipitation. Try Sodium or ammonium oxalate as a source of oxalate ion.
|
|
|
neptunium
National Hazard
Posts: 990
Registered: 12-12-2011
Location: between Uranium and Plutonium
Member Is Offline
|
|
oxalic acid is a weak acid HCl is a strong one...nothing should happen untill ALL the HCl is out of the solution i would rinse thoroughly the iron
solution or maybe neutralise it with sodium carbonate ( never mind the sodium they stay in solution)
|
|
|
Polverone
Now celebrating 21 years of madness
Posts: 3186
Registered: 19-5-2002
Location: The Sunny Pacific Northwest
Member Is Offline
Mood: Waiting for spring
|
|
If you want to salvage your existing solution, I second the suggestions that you raise the pH.
In the future you can use a less expensive, more massive form of iron (steel) like nails. If you like, first strip any galvanization with a
preliminary brief soak in dilute hydrochloric acid. Immerse the pieces of metal in warm oxalic acid solution in a bottle. Add a small amount of sodium
chloride or (better) hydrochloric acid to act as a catalyst. The iron (II) oxalate will form as a crust over the metal and needs to be removed by
vigorous shaking from time to time. The dislodged masses of iron oxalate are later easily ground to fine powder.
In summer the solution can be kept warm just by placing it in direct sun; in winter you may want to use a heating pad in a cardboard or foam box to
keep its temperature up. The process works in the cold too but much more slowly.
There are two small advantages to this approach. One is economy of reagents, using only a catalytic amount of hydrochloric acid and a less expensive
form of steel. Since the reagents are cheap this is only a minor benefit. The greater advantage is that a little hydrochloric acid in an oxalic acid
solution will not fume hydrogen chloride the way the neat acid does. You can keep the mixture indoors without dooming every ferrous object in the room
to visible corrosion.
PGP Key and corresponding e-mail address
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Aqueous Iron Chloride is an acid salts which is clearly evident from its hydrolysis:
FeCl2 + 2 H2O <==> Fe(OH)2 + 2 HCl
Boiling is a tedious way to impact the pH by gaseous evaporation of the HCl.
Note, HCl is a much stronger acid than Oxalic acid, and I would doubt the creation of iron (II) oxalate until the solution is neutralized.
You also may wish to see the thread on Iron oxide, as I discuss a way to create a Iron salt without the use of HCl.
[Edited on 4-1-2012 by AJKOER]
|
|
|
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
Too low pH, obviously.
Just buy some green vitriol and dump its fresh, filtered solution (use degassed, distilled water) to a solution of oxalic acid. Green vitriol is
available as a source of iron for plants and is quite cheap.
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Too low pH is one problem, but you also should be careful not to add too much oxalate. Iron forms very stable complexes with oxalate ion. Iron(II)
forms a very stable brown/yellow complex and iron(III) forms a very stable green complex. The funny thing is that the oxalate complexes have colors,
which are almost the reverse of all other iron(II) and iron(III) compounds (usually the first are green and the latter are yellow/brown, but with the
oxalato complex it is the other way around).
So, in another experiment, try to make a solution of potassium oxalate (not sodium oxalate, as that is only sparingly soluble) from slight excess of
oxalic acid and potassium carbonate and dissolve as much as possible iron in HCl. After filtering off the remaining iron, neutralize this solution
with potassium carbonate, such that the solution only remains marginally acidic. Do not add too much carbonate, otherwise iron is oxidized to
iron(III) by oxygen from the air. After this is done, you can mix both slightly acidic solutions. Use excess amount of the iron-containing solution.
[Edited on 4-1-12 by woelen]
|
|
|
Chemistry Alchemist
Hazard to Others
Posts: 403
Registered: 2-8-2011
Location: Australia
Member Is Offline
Mood: No Mood
|
|
Would Iron(III) Chloride work better?
|
|
|
Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
No, as I want the iron (II) oxalate salt and not the iron (III) oxalate . . .so you'd only have to reduce it back again anyway to what you started
with.
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
Endimion17
International Hazard
Posts: 1468
Registered: 17-7-2011
Location: shores of a solar sea
Member Is Offline
Mood: speeding through time at the rate of 1 second per second
|
|
Why are you sticking to the chlorides so hard? Is your goal to obtain iron powder, or to see how viable it is to take the chloride route?
|
|
|
Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
I'm not sticking so hard to the chlorides, its just that I have a very large amount of steel wool and large amounts of HCl available.
I could also use iron (II) sulfate I guess, but my only source of concentrated sulfuric acid is unfortunately impure and I want high grade iron
oxalate (and no, I am not going to distill concentrated H2SO4). I do have some iron sulfate and nitrate powders available, but they are very pure,
very expensive and don't want to waste them on a synthesis like this - rather save such purity for analyses.
Last night I heated my mixture gently on a hotplate again for three hours until no more white fumes were driven off . . .very shortly later a massive
clump of the compound precipitated out in the beaker. Yay!
After a filtration and washing and drying, I had approx 15.6g of my compound, or 94% of the theoretical. Most of it was lost mechanically as it
extremely insoluble in water (I believe 0.0008g in 100ml of water at 20C).
I didn't want to neutralise the original mixture directly using a base because I feared that iron oxides and hydroxides might form and contaminate the
product.
If you're interested, I'm going to use the salt for making pyrophoric iron - heat a small quantity of the crystals in a test tube until they decompose
and turn black . . .at which point you invert the test tube and watch the sparks fly!
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
