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White Yeti
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I leave them open to air to allow oxygen to diffuse into the water and react with the iron bits. Unfortunately, evaporation does happen quite quickly,
it loses about a half a litre per month.
I started making iron oxide this way one and a half years ago (damn time flies) after I got hold of a massive 2Kg mass of galvanised iron packaging
from trees. Instead of throwing those away, I decided to recycle them. I got rid of the zinc by putting all the bits in a tub of diluted vinegar.
I really like this process because it's so passive. you don't have to watch over it, you can forget about it and be greeted with a pleasant surprise
instead of an unpleasant one, and you only have to decant the iron oxide about once every 3 weeks.
I know I can clean the rust off, but since ~90% of the iron still hasn't rusted, I'll hold off on that
"Ja, Kalzium, das ist alles!" -Otto Loewi
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AJKOER
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For those lacking a strong acid or wishing to conserve, I would suggest dissolving Iron into a mixture of Bleach (NaClO) and Vinegar (Acetic acid).
The reaction is a slow to vigorous depending on concentration (but works in even 3% HClO solutions). As to the chemistry, that is a little of a
mystery. I would guess Ferric chloride and/or Ferric/Ferrous acetate, or Ferrous Chloride and an acetate salt.
The case for either the chloride or acetate salt rests on the strong smell (Cl2) and reddish brown color of the solution.
The argument for FeCl3 is that HClO could react with Iron as follows:
Fe + HClO --> FeO + HCl
2 FeO + HClO --> Fe2O3 + HCl
HCl + HClO <==> Cl2 + H2O
Fe2O3 + 6 HClO --> 2 Fe(ClO)3 + 3 H2O
But Iron Hypochlorite immediately decomposes, assuming it exists, being very unstable reacting perhaps with excess iron to form Ferric chloride:
2 Fe + Fe(ClO)3 + 3 H2O --> FeCl3 + 2 Fe(OH)3
This reaction is consistent with the thickening (water loss) of the solution and a visual appearance of Ferric hydroxide.
Alternate routes involve the formation of both FeCl3 and FeCl2, or just FeCl2. The argument is based on the fact that Ferrous chloride formation can
proceed in dilute HCl, and it may (or may not) be oxidized subsequently by any available Chlorine:
FeO (s) + 2 HCl (aq) ==> FeCl2 (aq) + H2O
2 FeCl2 (aq) + Cl2 (aq) ==> 2 FeCl3 (aq)
------------------------
The argument for some Iron Acetate (either Ferrous or Ferric) follows from the presence of NaOAc formed originally upon the treating of the Bleach
with Acetic acid (HOAc). For example:
NaClO + HOAc --> NaOAc + HClO
3 NaOAc + FeCl3 --> 3 NaCl + Fe(OAc)3
Note, there appears to exist some reference in the literature as to sodium Acetate and/or Acetic acid acting in the capacity of catalyst.
Interestingly, if HClO is prepared via a different route, like NaClO + H2CO3 (a much slower path to HClO), the result is a green solution consistent
with Ferrous chloride formation.
-----------------------
The argument for both Ferric (or Ferrous) Chloride and Ferric (or Ferrous) Acetate is that any HCl formed, say from hydrolysis of Ferric (or Ferrous)
chloride:
FeCl3 + 3 H2O ---> Fe(OH)3 + 3 HCl
could react with the Iron acetate formed from the Iron chloride:
3 HCl + Fe(OAc)3 --> FeCl3 + 3 HOAc
to regenerate the FeCl3 (or FeCl2 in the case of Ferrous salt). The argument for this case is consistent with the apparent longevity of the solution
(FeCl3 is know to decompose in time with a notable Ferric oxide precipitation). Also, it may be that only FeCl2 is formed but, in the presence of
reddish brown acetate, it's true identity is hidden.
------------------------------------------------
I found an interesting account of the reaction of Iron and HClO in a dated source concerning various experiments performed employing HClO (created via
washed Chlorine reaction on HgO in a water suspension) to ascertain the presence of blood in a murder investigation. Extracted from page 429 of "The
medical times", Volume 13:
"Experiment 14.—An iron blade, with two dark spots of blood upon it, recently applied, was left for thirty seconds in hypochlorous acid, chlorine
was disengaged, and oxide of iron formed; the spots were of a reddish-brown. At the expiration of an hour, they preserved the same colour, but they
were detached in some places, so that the iron at those spots appeared in all its brilliancy. The blade was then plunged into a fresh bath of
hypochlorous acid. After six hours' contact the spots were still brown in the centre, while their circumference, of a dirty red colour, had a kind of
little ridge, formed of sesquioxide of iron. After fourteen hours' immersion, one of these spots became greyish white, and as if encrusted with oxide
of iron, the other was of a reddish-brown, and peeled off in flakes; by leaving the blade in hypochlorous acid for thirty-eight hours, it became
thickly covered with sesquioxide of iron, and the liquid contained a large quantity of sesquichloride of iron ; when the oxide was removed by means of
a stream of water, it was seen that the spot, which had preserved its reddish-brown colour, had still the same hue, but it was only attached to the
blade by a few points near its centre."
This experiment appears, via a different and purer preparation method for HClO, that the reaction with Fe does indeed result in Fe2O3 and FeCl3.
LINK:
http://books.google.com/books?id=BhkCAAAAYAAJ&pg=PA607&a...
--------------------------------------------
Here another source that confirms Cl2, Fe2O3 and FeCl3 creation but discounts, to some extent, the existence of Ferric hypochlorite:
REFERENCE: "A comprehensive treatise on inorganic and theoretical chemistry, Volume 2 By Joseph William Mellor
EXTRACT
"Iron filings immediately decompose hypochlorous acid with a brisk effervescence produced by the evolution of chlorine; the iron is partly oxidized
and in part dissolved as chloride without the formation of any chlorate. A. J. Balard commented on this: "The greater number of other metallic
substances do not decompose hypochlorous acid, and I am yet entirely ignorant of the cause of the peculiar behaviour of iron." P. Grouvelle passed
chlorine through water with iron hydroxide in suspension and a bleaching liquid along with ferric chloride was produced, and he found the liquid
retained its bleaching properties after boiling for a quarter of an hour; but A. J. Balard failed to confirm this statement; he could not make ferric
hypochlorite either (i) by the action of hypochlorous acid on iron hydroxide, for hypochlorous acid does not dissolve ferric oxide; or (ii) by the
action of calcium hypochlorite on ferric sulphate, for calcium sulphate and ferric oxide are produced. Hence, adds A. J. Balard, "ferric hypochlorite
cannot exist;" the results by P. Grouvelle are due to the formation of ferric chloride and hypochlorous acid in dil. soln.; when the mixture was
heated, a portion of the acid distilled off, and the reaction which occurred with the cold soln. was reversed, for ferric oxide and chlorine were
formed. Ferrous oxide is oxidized to ferric oxide by hypochlorous acid."
A free google book. Link:
http://books.google.com/books?pg=PA275&lpg=PA275&dq=...
[Edited on 31-12-2011 by AJKOER]
[Edited on 31-12-2011 by AJKOER]
[Edited on 31-12-2011 by AJKOER]
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Oggas
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Iron(II)Chloride
When trying to produce magnetite I first had to make an iron(III)salt since I already had iron(II)sulfate. My procedure for this was to produce
iron(II)chloride and then to oxidize it. This I did by adding some steel wool to a solution of somewhat diluted 30% hydrochloric acid. I did not want
a too concentrated solution as I thought it might produce an awfull lot of magnetite making it harder to se the magnetic effect. However, having left
the solution of hydrochloric acid and steel wool over night with a lid on the reaction vessel I found a blue solution and some gunk. The solution was
the colour of a very dilute copper(II)sulfate solution. Upon addition of 3% hydrogen peroxide the solution turned a darker yellow-brown colour and I
used it to succesfully produce magnetite. So my question is, why is my solution blue instead of the anticipated yellow?
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woelen
Super Administrator
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Iron(II) is not yellow, it has a pale mint color, somewhat bluish. Besides that, you used steel wool as source of your iron and that is not the most
pure compound. The blue color also could be due to some impurity. I can well imagine that your observation is due to a combination of both things I
mentioned above.
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blogfast25
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AJ SIMPSON:
More unneeded chemical fantasy from Mr Pompous. A moderator assigned the moniker ‘Giver of bad advice’ to your pal Anders Hoveland, one cannot but
wonder how long it will be before the same fate befalls you?
Yes, a mixture of vinegar and household bleach, assuming you can make it acidic, is a slow boat to ferric hydroxide. But most of what you write by way
of reaction mechanisms might go on between your ears but not in actual solutions.
For one, most of the species you invoke are strongly dissociated and the actual reactions occur between ions.
In short, the slow dissolution of iron to ferric chloride or ferric hydroxide most likely proceeds as follows.
Iron is oxidised by oxonium ions (the solution is acidic) to ferrous ions:
Fe + 2 H<sub>3</sub>O<sup>+</sup> == > Fe<sup>2+</sup> + H<sub>2</sub> + 2
H<sub>2</sub>O
The ferrous ions are quickly oxidised by hypochlorite, to ferric ions, Fe3+. In acid conditions it’s presumably the free chlorine that’s the real
oxidiser.
Whether or not the ferric ions will stay in solution (as a mixture of chloride and acetate, since as the salts are dissociated it matters not one jot)
depends on the pH of the solution as well as the concentration of the ferric ions because Fe(OH)3 is extremely insoluble, thus favouring that stuff to
precipitate out. At a guess I’d say that real mixtures of bleach and vinegar would not be able to keep ferric ions in solution because the pH would
be too high.
In short, iron + bleach + vinegar is probably slightly faster than ordinary rusting, depending on conditions. That wasn't too hard, now was it?
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Poppy
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You will never be able to achieve FeCl3 solutions just with HCl 30%, equilibria between Fe(OH)3 and its trichloride is gonna precipitate the later out
of solution, no matter if you filter it, equilibria will do the same again. As fairly as you can a dilute FeCl3 solution is achiavable in HCl. To
complete the processo you need a HCl gas generator bubbling inside a flask with a certain ammount of FeO3, afterwards concentrations high enough will
be reached thus allowing FeCl3 to precipitate out as a salt, not a hydroxide.
I've just made the poorman's FeCl3 at home, I wont take a pic of it because its too dark to visualize the rubbish inside, but it still works fine on
the purpose you talk about, perphaps faster than ideal.
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blogfast25
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Quote: Originally posted by Poppy  | You will never be able to achieve FeCl3 solutions just with HCl 30%, equilibria between Fe(OH)3 and its trichloride is gonna precipitate the later out
of solution, no matter if you filter it, equilibria will do the same again. As fairly as you can a dilute FeCl3 solution is achiavable in HCl. To
complete the processo you need a HCl gas generator bubbling inside a flask with a certain ammount of FeO3, afterwards concentrations high enough will
be reached thus allowing FeCl3 to precipitate out as a salt, not a hydroxide.
I've just made the poorman's FeCl3 at home, I wont take a pic of it because its too dark to visualize the rubbish inside, but it still works fine on
the purpose you talk about, perphaps faster than ideal. |
Hmm…
I’ve actually made FeCl3.6H2O from scrap iron, 36 % HCl and peroxide. If you keep the acidity high enough you can concentrate the FeCl3 solution by
boiling in without any hydrolysis occurring (I reported it somewhere in this forum).
FeCl3.6H2O doesn’t crystallise easily (it’s very soluble) and it took days for the first solid, yellow FeCl3 hexahydrate to appear. After that,
the crystalline mass grew slowly but steadily until nearly all of the mass was solid product.
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Poppy
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Alrighte ye say, HCl wont evaporate out altogether?
I've got mine for months and the equilibrium is just as said 
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blogfast25
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| Quote: Originally posted by Poppy | Alrighte ye say, HCl wont evaporate out altogether?
I've got mine for months and the equilibrium is just as said
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Remember that HCl and water form an azeotrope around 20 % HCl. From that point on azeotrope starts evaporating and the acidity of the liquid is more
or less constant.
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White Yeti
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I've gotten iron(II) in solution that was completely colourless. Iron(II) acetate sealed in an airtight container is colourless in solution. But add
some bleach or hydrogen peroxide, and presto, it turns red.
"Ja, Kalzium, das ist alles!" -Otto Loewi
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