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Anomalous
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Inexpensive synthesis of Ferric Chloride
I'm thinking red iron oxide, and 28% muriatic acid.
6HCl + Fe2O3 -> 2FeCl3 + 3H2O
Given 2268g of Fe2O3 (14.175 mol) I calculate that I should need about 2.89 gallons of the muriatic acid.
Will the yield be anhydrous FeCl3 in solution, or will I get the hexahydrate?
Is there a more cost-effective approach to ferric chloride? It appears to me that the reaction above should cost about 1/4 what it costs to buy ferric
chloride.
Thoughts?
[Edited on 4-11-2013 by Anomalous]
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Crowfjord
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You would have solution of Fe3+ and Cl- in aqueous solution. If you boiled it down, you would get the hexahydrate, which would
decompose back to Fe2O3 if you heated further and drove off the rest of the water IIRC.
[Edited on 4-11-2013 by Crowfjord]
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barley81
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Keep in mind that red iron oxide from suppliers often reacts too slowly with aqueous solutions to be of much synthetic use. Calcined red iron oxide
dissolves very slowly in acid.
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Anomalous
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Quote: Originally posted by Crowfjord  | | You would have solution of Fe3+ and Cl- in aqueous solution. If you boiled it down, you would get the hexahydrate, which would
decompose back to Fe2O3 if you heated further and drove off the rest of the water IIRC. |
That sounds right to me. I know you can't get the anhydrous simply by cooking off the water.
Thanks!
[Edited on 4-11-2013 by Anomalous]
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Anomalous
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| Quote: Originally posted by barley81 | | Keep in mind that red iron oxide from suppliers often reacts too slowly with aqueous solutions to be of much synthetic use. Calcined red iron oxide
dissolves very slowly in acid. |
Hmm... I suppose I'll just have to try it. The supplier I use claims its iron(iii) oxide is from a natural source. Perhaps the reaction will be
catalyzed by other oxide impurities?
We'll see. Thanks. 
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elementcollector1
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Doesn't this form FeCl2, not 3?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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barley81
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No, because iron (III) oxide cannot oxidize chloride ion.
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elementcollector1
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I was thinking of iron metal - my mistake.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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Oscilllator
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The iron oxide I obtained from ebay dissolved extremely slowly. I stained a test tube with some iron oxide, and then left it to soak in nitric acid.
The acid was still there over a week later.
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woelen
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Nearly all sources of Fe2O3 sell calcined oxide. This is very inert and hardly dissolves in acid. It takes forever to dissolve even a small amount of
this oxide in a large excess amount of acid. This method of making a solution of ferric chloride is not useful at all. Well before the acid is used up
(even when still more than half of the acid is still remaining) the reaction comes to a near halt.
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AJKOER
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Try and dissolve the iron in a mixture of HCl and H2O2 and a touch of sea salt (or vinegar and bleach appears to work, but the products now include an
acetate). Even dilute starting solutions may work for you. I suspect some electrochemistry afoot (search for my comments/references on a so-called
'bleach battery' where now Iron metal replaces Aluminum and keep a new 100% plated copper penny in the mix as well), hence the progression with even
dilute reactants involving the corrosive oxidation of a metal in the presence of an oxygen source (HOCl) releasing Cl2 into the solution. Basically,
you are using chlorine water, more specifically, Hypochlorous acid as the fuel. The HOCl apparently enables the slow but steady attack of the Fe.
Beware and take precautions, you will get a strong Chlorine smell and the solution of Ferric chloride, even starting with dilute reactants, which can
become quite corrosive. Note, FeCl3 is commonly employed as an etching agent for circuit boards and the like.
What is amazing here is what a strong acid cannot do, a weak one does in the right conditions!
-----------------------------------------------
A related history involves the action of an oxygen source on Copper in the presence of aqueous ammonia. The debate among pundits proceeded apparently
for many many decades. The latest consensus is that the reaction appears to proceed along the lines of a electrochemical oxidation of a metal in the
presence of O2 coupled with some of the usual copper/ammonia chemistry.
[Edited on 4-11-2013 by AJKOER]
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blogfast25
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Dissolve scrap iron in an excess 37 % HCl, filter off insolubles. Oxidise the FeCl2 solution with 35 % H2O2, slowly and with constant cooling.
Evaporate the FeCl3 solution, while maintaining an acid reserve to prevent hydrolysis. FeCl3 is very soluble and difficult to get to crystallise. It's
all been explained somewhere on this forum. UTSF.
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blargish
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A solution of FeCl2 made from iron and hydrochloric acid can even be oxidized to FeCl3 by bubbling air through it. Just leaving
a solution of FeCl2 exposed to air will oxidize it after a bit.
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blogfast25
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| Quote: Originally posted by blargish | | A solution of FeCl2 made from iron and hydrochloric acid can even be oxidized to FeCl3 by bubbling air through it. Just leaving
a solution of FeCl2 exposed to air will oxidize it after a bit. |
Except... it takes forever! Fe(II) oxidises to Fe(III) much faster in alkaline conditions but even then with air oxygen it's a slow boat to China. The
trouble with air oxidation of Fe(II) is that it's a nuisance when you don't want it to happen and far too slow when you need
it to happen.
[Edited on 4-11-2013 by blogfast25]
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bbartlog
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The other issue with oxidizing a solution of FeCl<sub>2</sub> via air is that unless you have an excess of HCl in the solution, you'll get
some mixture of FeCl<sub>3</sub> and oxides. In fact the formation of Fe<sub>2</sub>O<sub>3</sub> may predominate
over the formation of any FeCl<sub>3</sub> if there is no HCl in the solution, just base on my experience with solutions of
FeCl<sub>2</sub> exposed to air (they don't turn red, but instead gradually develop brown gunk).
The less you bet, the more you lose when you win.
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blargish
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by blargish | | A solution of FeCl2 made from iron and hydrochloric acid can even be oxidized to FeCl3 by bubbling air through it. Just leaving
a solution of FeCl2 exposed to air will oxidize it after a bit. |
Except... it takes forever! Fe(II) oxidises to Fe(III) much faster in alkaline conditions but even then with air oxygen it's a slow boat to China. The
trouble with air oxidation of Fe(II) is that it's a nuisance when you don't want it to happen and far too slow when you need
it to happen.
[Edited on 4-11-2013 by blogfast25] |
Well... this is the inexpensive synthesis of Ferric Chloride  .
I think grabbing a straw and blowing bubbles through the solution is as inexpensive as it gets hehe. You're right though, it would take quite a while
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DraconicAcid
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| Quote: Originally posted by blargish | | Quote: Originally posted by blogfast25 | | Quote: Originally posted by blargish | | A solution of FeCl2 made from iron and hydrochloric acid can even be oxidized to FeCl3 by bubbling air through it. Just leaving
a solution of FeCl2 exposed to air will oxidize it after a bit. |
Except... it takes forever! Fe(II) oxidises to Fe(III) much faster in alkaline conditions but even then with air oxygen it's a slow boat to China. The
trouble with air oxidation of Fe(II) is that it's a nuisance when you don't want it to happen and far too slow when you need
it to happen.
[Edited on 4-11-2013 by blogfast25] |
Well... this is the inexpensive synthesis of Ferric Chloride .
I think grabbing a straw and blowing bubbles through the solution is as inexpensive as it gets hehe. You're right though, it would take quite a while
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Add a small amount of weak base to make the solution barely basic, allow it to oxidize by air, and then acidify to dissolve the sludge.
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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blargish
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| Quote: Originally posted by bbartlog | | The other issue with oxidizing a solution of FeCl<sub>2</sub> via air is that unless you have an excess of HCl in the solution, you'll get
some mixture of FeCl<sub>3</sub> and oxides. In fact the formation of Fe<sub>2</sub>O<sub>3</sub> may predominate
over the formation of any FeCl<sub>3</sub> if there is no HCl in the solution, just base on my experience with solutions of
FeCl<sub>2</sub> exposed to air (they don't turn red, but instead gradually develop brown gunk). |
I believe the equation is 12FeCl2 + 3O2 = 8FeCl3 + 2Fe2O3, so you are right, some iron oxide is
produced. But even so, the iron oxide can easily be filtered off, and not much product is lost, as much more ferric chloride is created that iron
oxide.
I guess excess HCl will react with the iron III oxide produced to make ferric chloride and water via 6HCl + Fe2O3 =
2FeCl3 + 3H2O, thus increasing yield a bit.
And draconic's procedure in the last post is to get around the production of the iron oxide sludge, I assume.
[Edited on 4-11-2013 by blargish]
[Edited on 4-11-2013 by blargish]
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Anomalous
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| Quote: Originally posted by blogfast25 | Dissolve scrap iron in an excess 37 % HCl, filter off insolubles. Oxidise the FeCl2 solution with 35 % H2O2, slowly and with constant cooling.
Evaporate the FeCl3 solution, while maintaining an acid reserve to prevent hydrolysis. FeCl3 is very soluble and difficult to get to crystallise. It's
all been explained somewhere on this forum. UTSF. |
My interest is in keeping the FeCl3 in solution.
I did use the search engine, and did not find anything so helpful to me as the content of this thread. In fact, I've spent days working out the
stoichiometry for various reaction paths, finding my mistakes, doing it over again, watching YouTube videos, reading articles, searching this forum
and others. I reached the point that I felt it was sensible to ask a question and be involved with a community that knows better than I. While I may
be new to this hobby, I don't need anyone to remind me to search. Thanks for your contribution, go on being superior and kindly avoid my threads.
To everyone else, thank you for your thoughtful contributions.
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Anomalous
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| Quote: Originally posted by woelen | | Nearly all sources of Fe2O3 sell calcined oxide. This is very inert and hardly dissolves in acid. It takes forever to dissolve even a small amount of
this oxide in a large excess amount of acid. This method of making a solution of ferric chloride is not useful at all. Well before the acid is used up
(even when still more than half of the acid is still remaining) the reaction comes to a near halt. |
This is the spec from my supplier:
"Red Iron Oxide – Natural
Typical Analysis
Fe 5.5 %
Fe2O3 82 %
MgO 1.1 %
Mn 0.4 %
Al2O3 2.9 %
SiO2 8.0 %
Pb 30 ppm
As 20 ppm
Moisture 1 %"
From this can we make an assumption about calcination?
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blogfast25
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Stop being childish. Stop trying to read minds: there's nothing in what I wrote that tries to convey superiority. They're not your threads: this is a
public forum.
No but most of us here have learned the hard way that commercial oxides are usually hard calcined and far more unresponsive to strong acids than some
textbooks would have you believe.
[Edited on 4-11-2013 by blogfast25]
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blogfast25
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You can't just pretend there isn't any water present and hope for the best.
The oxidation half-reaction is:
Fe2+(aq) ===> Fe3+(aq) + e-
The reduction half-reaction in neutral conditions:
O2(g) + 2 H2O(l) + 4 e- ===> 4 OH-(aq)
Then balance the electrons and add Cl- to make both sides charge neutral. You end up with a mixture of FeCl3(aq) and
Fe(OH)3(s). Freshly precipitated Fe(OH)3 will readily dissolve in strong HCl though, so there needs to be no loss of iron. But
air oxidation remains slow. Slight heating, good stirring and good dispersal (fine bubbles) of the air throughout the mix will of course speed things
up.
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blargish
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| Quote: Originally posted by blogfast25 |
You can't just pretend there isn't any water present and hope for the best.
The oxidation half-reaction is:
Fe2+(aq) ===> Fe3+(aq) + e-
The reduction half-reaction in neutral conditions:
O2(g) + 2 H2O(l) + 4 e- ===> 4 OH-(aq)
Then balance the electrons and add Cl- to make both sides charge neutral. You end up with a mixture of FeCl3(aq) and
Fe(OH)3(s). Freshly precipitated Fe(OH)3 will readily dissolve in strong HCl though, so there needs to be no loss of iron. But
air oxidation remains slow. Slight heating, good stirring and good dispersal (fine bubbles) of the air throughout the mix will of course speed things
up.
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Gah, silly mistake by me. It is best to keep the solution acidic with excess HCl, as we get
4Fe2+(aq) + O2(g) + 4H+(aq) = 4Fe3+(aq) + 2H2O(l)
or
4FeCl2(aq) + O2(g) + 4HCl(aq) = 4FeCl3(aq) + 2H2O(l)
Thanks for clarifying
And Anomalous, are you making this ferric chloride solution for a PCB etchant?
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Anomalous
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Yes, that was the plan.
[Edited on 5-11-2013 by Anomalous]
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blargish
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Ok, then . I found these two videos on youtube, one of which details the
inexpensive production of ferric chloride, whilst the other one goes over other PCB etchants that can be made with relatively cheap chemicals.
However, the ferric chloride in the video is made via use of steel wool and HCl, so if you are bent on going the Fe2O3 route, I
guess these will be of no help.
Make Ferric Chloride
http://www.youtube.com/watch?v=43Xsh9J7Sg
10 Other PCB etchants
http://www.youtube.com/watch?v=Q4tWEse2rDI
Sorry if you have seen these already
[Edited on 5-11-2013 by blargish]
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