CHRIS25
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Many ways to Ferric Oxalate?
Question 1:
Ferric oxalate is non-stoichemetric compound and can not be easily defined....a phrase I have often encountered and have tried to understand. It has
been said that one can pull out three different measures of ferric oxalate from solution, made in the same way and still end up with three different
structurally speaking compounds that can not be measured in a strict stoichemetry way? I recognise after some learning that the oxalate is a ligand
and therefore can be unpredictable in how it attaches itself to the metal ion, but that is as far as my learning has taken me. Can one perhaps
simplify this explanation so that I may then build upon something that I can understand and take the study further?
Question 2:
I have encountered time and time again the use of Ferrous ammonium sulphate, potassium ferricyanide and Hydrogen Peroxide as the recommended method
for making the Ferric oxalate. There has not been one single place on the whole of the internet where I have found the Ferric chloride and oxalic
acid route, which appeared to me to be the most logical route to take. This I do not understand. Another method was precipitating ferrous oxalate in
ferrous chloride - but this latter one seems unnecessarily redundant. I am looking for any other methods, but in the meantime I do not see the
problem with the following:
2FeCl3+3H2C2O4 = Fe2(C2O4)3+6HCl
.......providing one does not allow hydrolysis to take place when mixing the Ferric chloride into solution in the first place. I think Vargouille
kindly suggested to dissolve the ferric chloride in HCl as opposed to water.
I would appreciate comments please, if you could be so kind. This is photography related chemistry so while I have learned a huge amount of theory
through reading and experimentation over the last six months I am not a chemist.
‘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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ScienceSquirrel
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Your best course of action is to follow a recipe eg;
http://www.ericneilsenphotography.com/free_papers/Ferric%20O...
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AJKOER
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OK, per Watt's a direct approach to quote:
"Ferric salts. When ferric hydrate is treated with a quantity of aqueous oxalic acid not sufficient to dissolve it, a yellow powder is formed nearly
insoluble in water, and apparently consisting of neutral ferric oxalate. The same salt is precipitated on adding a small quantity of neutral potassic
oxalate to a ferric salt. It dissolves in oxalic acid, forming a solution which, when exposed to sunshine, gradually assumes a greenish-yellow colour,
gives off carbonic anhydride, and deposits crystals of ferrous oxalate till it becomes quite colourless. Ferric hydrate dissolves in the acid oxalates
of the alkali-metals, forming double salts. Ammonio-ferric Oxalate, ...—A hot solution of ferric hydrate in acid oxalate of ammonium deposits this
salt on cooling in small, anhydrous, rhombic octahedrons, having a greenish-white colour and turning yellow when exposed to light"
Source: "A dictionary of chemistry and the allied branches of other sciences", Volume 4, by Henry Watts, page 258.
Link: http://books.google.com/books?pg=PA258&id=lYXPAAAAMAAJ#v...
Also, a note of why the direct reaction with FeCl3, a chloride of an easily reducible metal, does not work:
"Chlorine does not act on dry oxalic acid; but in presence of water, decomposition quickly takes place, thus:
H2C2O4 + Cl2 = 2CO2 + 2HCl
A similar reaction is produced by bromine, hypochlorous acid, and the chlorides of easily reducible metals. Hence oxalic acid precipitates metallic
gold from auric chloride, especially on boiling;"
---------------------------------------
With respect to the synthesis, a problem, I think, would be in forming your Fe(OH)3 that is free of any Na, K, NH4,..residue from neutralizing FeCl3
to produce Fe(OH)3 as the latter could add an Iron double salt contamination.
A final comment on the much tauted 'recipe', I guess it would be OK if you did not know that Iron oxalate double salts also existed, or you don't care
on such contamination (if your a student don't mention this to your teacher as they take it personally when you embarass them with the truth, or
better, play dumb and ask the question: 'Won't double salts formation be a problem?').
[Edited on 18-2-2013 by AJKOER]
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Vargouille
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The simplest meaning to a non-stoichiometric compound is that the compound has unusual structural properties that means that while Fe2(C2O4)3 can be a
useful approximation, the ratio of ferric to oxalate isn't always the same. There could be inclusion of bioxalate or free acid (in the same way that
green cupric chloride crystals contain HCl), or there could be some reduction of ferric to ferrous.
As for reduction, the strict calculations say that it is to be expected, although the apparently low solubility of ferric oxalate in general (although
not as well defined as that of ferrous oxalate) should keep reduction to a minimum. I can't imagine that the acidic medium has an enormous effect on
the reduction potential, considering that the half-reaction for oxalate is "H2C2O4 --> 2CO2 + 2H+ + 2e-".
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CHRIS25
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Sciencesquirrel: Your best course of action is to follow a recipe eg;
http://www.ericneilsenphotography.com/free_papers/Ferric%20O... I have this but it appears overly complicated and fraught with pitfalls.
Vargouille - thanks for the extra explanation. I have read about splitting reactions up into half re`actions to show their electron status etc,
unfortunately it is all theory, I do not have the experience to apply the theory in practise, so sadly I can not understand how this translates into
practise. I guess that what you say is that there is no need to add extra HCl?
AJkoer: <<<<<< "Chlorine does not act on dry oxalic acid; but in presence of water, decomposition quickly takes place, thus:
H2C2O4 + Cl2 = 2CO2 + 2HCl>>>>>>
So if all that happens is that oxalic acid dissappears for all intense and purposes, why then does the balanced equation result in Ferric oxalate? I
check my balancing with a site and usually if I have completely missed the point then the site gives me a big red sign saying 'Impossible reaction'
Ok I know I should not rely on this, but I can not understand why the Iron and the Oxalic do not join together? (forgive my non scientific
phraseology please).
Also: You said "A final comment on the much tauted 'recipe', I guess it would be OK if you did not know that Iron oxalate double salts also existed,
or you don't care on such contamination (if your a student don't mention this to your teacher as they take it personally when you embarass them with
the truth, or better, play dumb and ask the question: 'Won't double salts formation be a problem?')"
Please explain this, this double salt formula is tauted all over the web, both in science references and photographers references, I am confused now.
[Edited on 18-2-2013 by CHRIS25]
I have to add that ferric oxalate is un-obtainable in UK and Ireland. The only place where I can buy it is in the USA via Bostik and Sullivan, but at
almost 65 dollars extra to ship it is too expensive. So if, contrary to this sites understandable policy of not spoon feeding, someone would like to
offer me a method whereby I can make it then I would appreciate that very much. It appears from this site and a whole load of references that it is a
very difficult product to make. I thought it was straightforward, the reaction balanced stoichemetrically was too good to be true. Help? Perhaps
one could elaborate on the double salt downfalls before I consider perhaps the ferrous oxalate which is easy to make, and then perhaps oxidize that
with H2O2? I really don't know.
[Edited on 18-2-2013 by CHRIS25]
[Edited on 18-2-2013 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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Vargouille
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I would not say that there is no need: for preventing hydrolysis of the ferric ion, extra acid is helpful. I was referring to the oxalic acid, in that
the half-reaction for oxalic acid has two unchanged protons (H+) on either side, and I was expressing doubt that the product from acidic solution
would contain more ferrous contaminant than that of the relatively neutral solution because of that. The reaction that AJKOER suggests would occur is
"2Fe+3 + H2C2O4 --> 2Fe+2 + 2CO2 + 2H+" instead of the simpler "2Fe+3 + 3H2C2O4 --> Fe2(C2O4)3 + 6H+". The first is a redox; the second is a
double-replacement reaction. This is one case in which you have to know the properties of the compounds involved; and I must apologize for forgetting
the reducing properties of oxalic acid.
If you want to go the ferrous to ferric oxalate route, it could work, although I can imagine that the same issues regarding reduction may take place,
as well as accidental oxidation of the oxalate by the hydrogen peroxide. As far as double salts, it could be stabilizing, as, if I remember correctly,
ferrous ammonium sulfate is more stable than ferrous sulfate.
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CHRIS25
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Thanks Vargouille, I will have to really work things out here I suppose. But what you said: "........This is one case in which you have to know the
properties of the compounds involved........" you are right. I am working on that.
‘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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AJKOER
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Per Wikipedia (http://en.wikipedia.org/wiki/Double_salt ) to quote:
"Double salts are salts containing more than one cation or anion, obtained by combination of two different salts which were crystallized in the same
regular ionic lattice. Examples of double salts include alums (with the general formula MIMIII[SO4]2•6H2O) or Tutton salts (with the general formula
[MI]2MII[SO4]2•6H2O).[1] Other examples are potassium sodium tartrate and bromlite.
Note that double salts should not be confused with a complex. When dissolved in water, a double salt completely dissociates into simple ions while
complexes do not; the complex ion remains unchanged. For example, KCeF4 is a double salt and gives K+, Ce3+ and F− ions when dissolved in water,
whereas K4[YbI6] is a complex salt and contains the discrete [YbI6]4− ion which remains intact in aqueous solutions.[1] It is therefore
important to indicate the complex ion by adding a square bracket "[ ]" around it.
In general, the properties of the double salt formed will not be the same as the properties of its component single salts"
------------------------
Now on Ferric oxalate, people make this salt ever day when they use Oxalic acid, a rust remover, to remove stains. The important point to remember is
that in excess, many of the insoluble oxalate salts form soluble complexes with excess H2C2O4. Hence, Watt's comment "ferric hydrate is treated with a
quantity of aqueous oxalic acid not sufficient to dissolve it". I believe on reading Watts that one can use Fe2O3 (as is in the case of rust removing)
in place of Fe(OH)3 although the latter as a fresh precipitate, should be more reactive.
[Edited on 19-2-2013 by AJKOER]
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CHRIS25
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Fe2O3 + 3H2C2O4 = Fe2(C2O4)3 + 3H2O
Ok, you said: << The important point to remember is that in excess, many of the insoluble oxalate salts form soluble complexes with excess
H2C2O4.>>
Ferric oxalate is insoluble as we know, the product is yellow, so Watts's comment and yours above means that if I add "Less" than the stoichemetric
amount of Oxalic acid to the ferric oxide, then I would get a fairly good ferric oxalate precipitating. Sorry, but at this stage in my learning I
can not work/think in anions and electrons and charges since I am still reading about these things. Since I need this substance I can not wait for my
understanding to catch up. Hope you understand.
‘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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CHRIS25
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Further searching has turned up the following information: The web reference given below is apparently an alternative process forum.
http://www.usask.ca/lists/alt-photo-process/alt95/subject.ht...
The following information that is relevant to Ferric oxalate has been selectively culled from those messages for my own learning. As it turns out I
feel that I really want to understand this substance even more and make it, even though it is a huge distraction from the alternative photography that
I do; it has my interest and therefore I might ask you guys for your thoughts and comments to all this, I would be very interested........A Big
Thnakyou.
=======================================================================
Just for kicks, I thought I'd mention Phil Davis's method of making
ferric oxalate. He makes a solution of ferrous oxalate and puts the
beaker in an ice salt bath to cool it way down. Then he slowly adds drop
by drop full strength hydrogen peroxide or as Phil calls it "bomb" grade.
The icing down slows down the reaction enough so that human life is
sustained within 100 yards.
==============
The perennial problem of ferric oxalate stems from the fact that it is
described as an "ill-characterised" substance by chemists, that is: it does
not form a unique crystalline solid of definite, repeatable composition;
this is rare behaviour for a metal salt, most of which are
"well-characterised".
The solid that is obtained from an aqueous solution of ferric and oxalate
ions varies in its colour (from yellow to green), composition and
solubility, depending on the parameters of the method of its preparation,
especially the pH, temperature, time and relative concentrations of the
components. Richard Sullivan has evidently struck on a method of obtaining
a 'user-friendly' variety which has escaped the "major chemical houses",
and good luck to him. I once had a research student prepare ferric oxalate
by three different published methods, and he came up with three different
substances.
The variable amount of water in the formula of solid ferric oxalate
reflects the fact that the ferric iron is to some, variable, extent
hydrolysed (linked to other ferric ions via hydroxy bridges) and therefore
polymerised, which affects the readiness of its solubility in water. The
variability and poor crystallinity of solid ferric oxalate mean that its
molecular structure (which is undoubtedly complex) has not yet been
determined, to the best of my knowledge, and chemists still know very
little about it. Nor are they, (with one or two notable exceptions!) very
interested.
=============
For kallitype (and platinum) I found that a suitable ferric oxalate
solution was obtained by dissolving the following in 20ml distilled water
and leaving to stand for 24hrs:
iron(III) nitrate nonahydrate 8.4g
potassium oxalate monohydrate 5.7g
oxalic acid 0.25g
====================
To my knowledge there is no published formulas for making a powdered form of
ferric oxalate in the standard preparatory text, either Bauer or the
Handbook of Preparatory Chemistry.
=========================
1. Dissolve 2 grams of oxalic acid in 60 ml of destilled water.
2. Stir in 26 grams of sodiumoxalate in the solution. It wont dissolve.
3. Stir in 52 grams of Ferric(III)nitrate, after some stirring the
solution will turn dark green.
4. Add water until 110 ml. Pour into two dark dropper bottles 55 ml in each.
5. Label one of the bottles 1.
6. Label the other bottle 2. Add 0.3 grams of potassiumchlorate for
platinotypes and 0.6 grams of potassiumchlorate for palladiotypes.
========================
‘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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AJKOER
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Quote: Originally posted by CHRIS25  | Fe2O3 + 3H2C2O4 = Fe2(C2O4)3 + 3H2O
Ok, you said: << The important point to remember is that in excess, many of the insoluble oxalate salts form soluble complexes with excess
H2C2O4.>>
Ferric oxalate is insoluble as we know, the product is yellow, so Watts's comment and yours above means that if I add "Less" than the stoichemetric
amount of Oxalic acid to the ferric oxide, then I would get a fairly good ferric oxalate precipitating.... |
Actually, my take of the direct action of H2C2O4 on the insolube Fe2O3 is that we are converting it into the yellow Ferric oxalate. Other synthesis
may be faster, but in my opinion, using more ingredients and heating just invites likelihood of contamination or double salts formation.
Note, for example, that the purest HCl is prepared by burning H2 in Cl2 (this is, actually, the basis of a commercial process).
[Edited on 20-2-2013 by AJKOER]
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CHRIS25
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"""Actually, my take of the direct action of H2C2O4 on the insolube Fe2O3 is that we are converting it into the yellow Ferric oxalate. Other synthesis
may be faster, but in my opinion, using more ingredients and heating just invites likelihood of contamination or double salts formation."""
Ok, it all looks nice on paper, and under normal circumstances I would try this, but have to make iron oxide first and I know it is no good making
rust since this is a hydroxide complex as well, (unless I heat the rust to above 1200c which I suppose I could do). Nevertheless what bugs me is that
this formula is too simple; my previous post above clearly demonstrates that, in the public realm at least, chemists have not published any methods;
and if it were that easy then the above reaction would have been popularized on photographers' forums since 1995. Hope you see this not as an
argument but simply questioning, as if I were in a classroom.
[Edited on 20-2-2013 by CHRIS25]
[Edited on 20-2-2013 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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ScienceSquirrel
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Chris25 does not have to make ferric oxalate as a solid or even as a pure solution.
All he has to make is a solution of ferric oxalate that acts in the way that he requires when making his photographs.
I suspect that these solutions are anionic complexes of ferric iron with oxalate ligands with ammonium, potassium or sodium counter ions.
The best course for him may be to start from iron III nitrate. Readily available and the nitrate anion is not strongly complexing or reactive in
phographic mixtures. Add oxalic acid, etc and then try the solution out.
Make a print and see how it goes.
There are lots of factors to consider.
He will use sunlight to do the exposures and he will make the print paper himself.
We can help him with the chemistry but he is making art and there are no calculations for that.
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CHRIS25
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I do appreciate AJKoer's input and thank him. Sciencesquirel's comments are indeed helpful because I have been so focused on the chemistry of these
substances that I was losing sight of the fact that the ferric oxalate is simply there to be reduced to ferrous oxalate in the paper and replaced with
the silver ion (silver nitrate gets thrown in as well) while the ferrous ion gets washed away during the wash process; I would like to understand
the meaning of counter ions though....I have read that ammonium ferrioxalate is sometimes noted, sometimes a sodium oxalate is thrown in. Since
reading about electron transfers and sharing and the redox reactions and ionic bonding etc I find this theoretical learning is, at this moment
unfortunately, not helping me to understand the use of these counter ions. The actual reduction of ferric ion to ferrous by UV light and similar
processes I do understand, but when these extra ingredients are thrown in I just want to know why, chemically speaking.
‘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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AJKOER
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ScienceSquirrel thank for explaining the context. Here is some background on so called platinum print per Wikipedia (http://en.wikipedia.org/wiki/Platinum_print ):
"Willis introduced the "hot bath" method where a mixture of ferric oxalate and potassium chloroplatinate are coated onto paper which is then exposed
through a negative and developed in a warm solution of potassium oxalate. This is the basic platinotype process which is in use today."
My observation is that the ferric oxalate is not dissolved in excess H2C2O4 as Oxalic acid would react with Potassium chloroplatinate forming
Potassium oxalate which is the photographic developer. Now, whether the Ferric oxalate actually dissolves in the potassium chloroplatinate solution is
not material, as I would employ, in any event, a very fine form of Ferric oxalate. This implies to me the need of producing a very fine form of
Fe(OH)3. One recent research review (see http://www.ijest-ng.com/vol2_no8/ijest-ng-vol2-no8-pp127-146... ) notes to quote "Synthesis of iron oxides in the nano range for various
applications has been an active and challenging area of research during the last two decades. The processes include careful choice of pH,
concentration of the reactants, temperature, method of mixing, and rate of oxidation (Domingo et al., 1994). The morphology of the iron oxide
particles depends on the competition between several processes like nucleation, growth, aggregation and adsorption of impurities (Cornell and
Schwertmann, 1996)".
Also, "Iron oxides (FeOOH, Fe3O4 or γ-Fe2O3) are usually prepared by addition of alkali to iron salt solutions and keeping the suspensions for
ageing. The main advantage of the precipitation process is that a large amount of nanoparticles can be synthesized."
And, "The higher the pH and ionic strength, the smaller the particle size and size distribution width will be, because these parameters determine the
chemical composition of the crystal surface and consequently the electrostatic surface charge of the particles (Tartaj et al., 2006)." And finally,
"Pure goethite was synthesized using 1 M ferric nitrate solution and 10 M sodium hydroxide solutions under controlled conditions. Ferric nitrate
solution was vigorously stirred at room temperature with the simultaneous addition of 10 M sodium hydroxide solution until the pH of the solution
reached 12–12.5. In order to obtain Cu, Ni or Co doped goethites, the respective sulphate solutions were mixed with ferric nitrate solution prior to
alkali addition".
Here is another article where conc ammonia and Ferric nitrate formed very fine Fe(OH)3 structures with appropriate processing (see
http://144.206.159.178/ft/748/86678/1467549.pdf ). After forming the fine Ferric hydroxide, treat with Oxalic acid (not in excess).
[Edited on 21-2-2013 by AJKOER]
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ScienceSquirrel
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This might help;
http://www.bostick-sullivan.com/articles/ferriccoalate.html
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CHRIS25
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Yes I have decided to go down this road, I found this weeks ago, but went off in different directions seeking other ways until I figured that ferric
chloride and oxalic should work, then discovered problems then posted and goodness knows the amount of studying I have been doing just to end up back
at the place where I started three weeks ago. Mmm....I will make the ammonium sulphate and the Iron 2 Sulphate and then the ferrous ammonium sulphate
from that. Now that should be easy shouldn't it?
[Edited on 21-2-2013 by CHRIS25]
[Edited on 21-2-2013 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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