Botanic88
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How can I get simple ferric ions from a ferrioxalate complex?
I am trying to improve an alternative photographic process, which I invented in the 1970's.
This process uses a sensitiser containing ferric ions and ferrioxalate ions.
The ferrioxalate ions are necessary to obtain reasonably fast exposure times.
However the process requires simple ferric ions in order to fix a coat of pigmented gum.
The gum is fixed in a separate part of the process, after the exposure has been made.
If there is some way to convert the unexposed ferrioxalate ions into simple ferric ions then it would greatly improve my process.
Can anyone offer any suggestions?
There are certain restrictions. E.g. it is a photographic process, so heating the chemicals to a high temperature on the paper is not an option!
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blogfast25
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Ah, I get it better now. In your previous thread you were trying to convert a mixture of Fe<sup>3+</sup> and ferrioxalate to
Fe<sup>3+</sup> only, by destroying the complex through addition of Ca<sup>2+</sup> which is supposed to form CaOx
precipitate. Except: in the presence of ‘free’ (a better term than ‘simple’) Fe<sup>3+</sup>, the
ferrioxalate/Ca<sup>2+</sup>*Fe<sup>3+</sup mixture yielded no precipitate, suggesting the complex remained intact
(‘shielded’ by the free Fe<sup>3+</sup> . Let me first see if I can
reproduce your results.
Can you post some data on the solution strengths? You must have some idea of how many grams of this and that you dissolve per litre or 100
ml, as no process could ever be reproducible without pinning down these numbers…
One possible way might be to try and use another cation instead of Ca<sup>2+</sup> because although Ca oxalate is techically insoluble, it
isn't that insoluble: Wki lists the solibility at RT at 0.00067 g/100 ml, that's fairly soluble compared to some really insoluble stuff out
there! This mild solubility may be the reason CaOx isn't very good at 'breaking' the ferrioxalate complex, less so in the presence of free
Fe<sup>3+</sup>.
Using a cation that forms a far more insoluble oxalate (with a much lower Ks compared to CaOx) should be more than worth considering.
Your homework: look up various insoluble oxalates. Neodymium oxalate is one of them. Compare solubilities to CaOx.
[Edited on 22-10-2011 by blogfast25]
[Edited on 22-10-2011 by blogfast25]
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Botanic88
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Blogfast25:
Yes, my main concern is with my photographic process.
However I also found the phenomenon, that a little ferric chloride seems to inhibit the precipitation of calcium oxalate, both surprising and
interesting in itself.
So I hope you will try to repeat that experiment yourself.
The possibility of using other cations had occurred to me. For example the solubility of zinc oxalate is 1.4 E -9
Also I happen to have some of this chemical, which is always an advantage!
However I found that this chemical behaves in the same way as calcium chloride, in respect of the precipitation experiment
and trying it in various ways with my photographic process also showed no signs of improvement.
This makes me think that obtaining free ferric ions by precipitation will not work
unless I can be sure that the precipitate really is calcium oxalate (or some such) and not something like calcium ferrioxalate.
As to the amounts of chemicals I use with my photographic process:
Typically I use about 1 part strong ferric chloride soln with 3 parts water.
With 30 year old paper it is unnecessary to use any oxalate. The ferric chloride alone is slightly light-sensitive; probably due to a reaction with
the paper and/or the size.
With modern papers the ferric chloride can be made light-sensitive with the addition of only 1% w/v potassium oxalate.
However I often use as much as 15% w/v potassium oxalate, in order to shorten the exposure times.
If you would like to try my photographic process for yourself you can find details at
http://www.alternativephotography.com/wp/processes/gum-bichr...
This article also contains references to my earlier on-line articles.
Alternatively you can find all the articles by doing a search on "ferric gum"
I might even tempt you (or other chemists) to participate in my search for an improved process, or to help with the negative-working version of Ferric
Gum, which is chemically another kettle of fish!
Meanwhile, has anyone got any other ideas about getting free ferric ions from ferrioxalate ions?
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anonymous201
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http://firstyear.chem.usyd.edu.au/LabManual/E02.pdf
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blogfast25
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Well, I can confirm the added FeCl3 prevents calcium oxalate from forming with ferrioxalate complex, like Botanic88 claimed. Here’s what I did.
Firstly I made stock solutions of 1 M CaCl2, 1 M acidified FeCl3, 0.8 M oxalic acid and 0.33 M potassium ferrioxalate. The latter was made from
homemade K3Fe(C2O4)3. This salt doesn’t dissolve to well in cold water and I had to apply some gentle heat. Below’s 25.0 ml of a 0.33 M solution.
On cooling small amounts of crystals started to deposit, so 0.33 M appears to be close to the solubility limit for this salt at RT.
Then some tests were carried out in test tubes:
1. 2 ml CaCl2 + 2 ml oxalic acid: precipitation forms immediately.
2. 2 ml CaCl2 + 2 ml ferrioxalate: precipitation is delayed, is less voluminous and the suspension remains green; indicating the complex had not been
entirely ‘broken’.
3. 2 ml CaCl2 + 1 ml FeCl3 + 2 ml ferrioxalate: no precipitation or turbidity occurs.
4. 2 ml CaCl2 + 2 ml FeCl3 + 1 ml oxalic acid: briefly turbidity occurs but this disappeared, indicating the oxalate prefers to form the ferrioxalate
complex, rather than precipitate as CaC2O4.
Tubes:
A second series was as follows:
5. 2 ml CaCl2 + 0.8 ml FeCl3 + 2 ml ferrioxalate
6. 2 ml CaCl2 + 0.6 ml FeCl3 + 2 ml ferrioxalate
7. 2 ml CaCl2 + 0.4 ml FeCl3 + 2 ml ferrioxalate
8. 2 ml CaCl2 + 0.2 ml FeCl3 + 2 ml ferrioxalate
No turbidity or precipitation occurred in any of these four.
Tubes:
Finally Test 8. was repeated on a slightly larger scale in order to measure pH, found to be about 1.0. Such a low value does of course suppress
oxalate ions a bit too. To the solution was then added 2 ml of oxalic acid (nothing happened), then another 2 ml which finally did cause turbidity and
precipitation.
Everything points to the notion that CaC2O4 is essentially too soluble (Ks too high) to effectively snatch the oxalate from ferrioxalate and
break this complex, especially when free Fe3+ is around.
Regards Zinc, for CaC2O4 the solubility product Ks is about 2.7 x 10<sup>-9</sup> (based on a solubility of 0.0067 g/L, Wiki), so ZnC2O4
wouldn’t really fair much better (assuming the value of 1.4 x 10<sup>-9</sup> is indeed the solubility product for ZnC2O4). We probably
need something in the order of 10<sup>-15</sup> or lower.
Edit:
This *.pdf gives solubility products for RE oxalates (page 6, Table 2):
http://www.cheric.org/PDF/JIEC/IE04/IE04-4-0277.pdf
… as ranging from 6 x 10<sup>-30</sup> to 4 x 10<sup>-32</sup>. Cerium and neodymium salts must by now be relatively easy to
obtain. I have some neodymium trichloride hydrate, so I might put that to the test…
[Edited on 22-10-2011 by blogfast25]
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Botanic88
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Anonymous201:
The link you have given is about some experiments with the 'blueprint' mechanism.
These experiments don't produce any of the free ferric ions, which are necessary in my photographic process. First they reduce some ferrioxalate when
exposed to light. Then the resulting chemicals are precipitated as prussian blue.
Blogfast25:
Thank you for your impressive contribution! And thanks for confirming my very sloppy results.
The solubility table in Wikipedia gives the solubility of calcium oxalate as 0.00067 at 20 deg (this is g/100ml)
The same table gives the solubility of zinc oxalate as 1.38 E-9
If this is correct it makes zinc oxalate about 100,000 times less soluble.
Incidentally I still cannot find any data for the solubility of calcium ferrioxalate.
Some comments on your experiments:
| Quote: | | 2. 2 ml CaCl2 + 2 ml ferrioxalate: precipitation is delayed, is less voluminous and the suspension remains green; indicating the complex had not been
entirely ‘broken’. |
If the complex had broken significantly then I would expect to see the colour of the
soln change to a more olive green, as you got when you added ferric chloride in the other test tubes.
The lack of colour change also supports the idea that calcium ferrioxalate is being precipitated, rather than calcium oxalate, because the former
would not result in any (coloured) free ferric ions.
| Quote: | | 4. 2 ml CaCl2 + 2 ml FeCl3 + 1 ml oxalic acid: briefly turbidity occurs but this disappeared, indicating the oxalate prefers to form the ferrioxalate
complex, rather than precipitate as CaC2O4. |
I think your interpretation may be correct. On the other hand I
have noticed that when CaOx is precipitated from KOx it is easy to redissolve the precipitate in an excess KOx. Something like this may also be
happening here. (but perhaps there is no excess Ox in this case)
| Quote: | | 8. 2 ml CaCl2 + 0.2 ml FeCl3 + 2 ml ferrioxalate ... No turbidity or precipitation occurred |
It would be
interesting to know how little ferric chloride could be added to achieve the effect.
| Quote: | | Finally Test 8. was repeated on a slightly larger scale in order to measure pH, found to be about 1.0. Such a low value does of course suppress
oxalate ions a bit too. To the solution was then added 2 ml of oxalic acid (nothing happened), then another 2 ml which finally did cause turbidity and
precipitation. |
This might be explained by assuming that the first additional amount of Ox acid complexed
with the ferric chloride, then the second amount caused precipitation because there were no more free ferric ions to complex with.
To anyone else out there!
I am still hoping to get other ideas to improve my photographic process. E.g. it might be possible to destroy the oxalate instead of precipitating it
out!
[Edited on 23-10-2011 by Botanic88]
[Edited on 23-10-2011 by Botanic88]
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blogfast25
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Botanic88:
The value of 1.4 x 10<sup>-9</sup> for ZnOx is a solubility product (Ks, aka Ksp), not an actual solubility (note the important
difference), see the references below that both state this value. If correct, the molar solubilities of ZnOx and CaOx are very similar and would
explain your results with both:
http://www.ktf-split.hr/periodni/en/abc/kpt.html
http://www.thelabrat.com/protocols/solubilityproductconstant...
Re. your comment on Test 4., if anything, there a shortage of oxalate in that solution (the quantities were chosen very deliberately - hence the use
of known stock solutions). Also we’re in strongly acidic conditions and:
H2Ox(aq) + 2H2O(l) < === > 2 H3O+(aq) + Ox2-(aq); high [H3O+] pushes it to the right, suppressing free oxalate further. Oxalic acid is a
medium-weak acid.
The temporary turbidity (it lasted no more than a minute) was likely due to initial poor mixing, giving perhaps rise to localised high [Ox2-].
$%$%$%$%$%
I also carried out some tests with neodymium trichloride (NdCl3). My small stockpile was unfortunately not in great shape, so some work-up was needed
to obtain a clear (slightly lavender), strong solution of Nd<sup>3+</sup> but of unknown concentration.
With it, 3 more tests were conducted:
9) 2 ml NdCl3 + 2 ml oxalic acid: immediate and permanent precipitation occurred: Nd oxalate. The amount was comparable to 1)
10) 2 ml NdCl3 + 2 ml ferrioxalate: immediate and permanent precipitation, about the same amount as 9). Supernatant liquid remains green, so there’s
excess ferrioxalate. This result is consistent with Nd oxalate being much more insoluble than Ca oxalate: compare it to 2)
11) 2 ml NdCl3 + 1 ml FeCl3 + 2 ml ferrioxalate: no precipitate formed but much later some turbidity did develop. Solution remained olive green.
Disappointing result, in short.
I’m now checking whether the precipitate from 10) will dissolve by adding about 1 ml of FeCl3. It didn’t immediately do so. But with such
insolubility and at RT, dissolution may take some time. Wait and see…
I’m now also pretty convinced that finding an oxalate complex that is significantly stronger than
FeOx<sub>3</sub><sup>3-</sup> will be nigh impossible.
If you do want to experiment with RE oxalates in real photographic conditions, you could prepare so-called ‘lanthanide chloride’, obtained from
dissolving ‘mischmetal’ (German: mixed metal) which is a mixed RE alloy. Mischmetal is what makes up lighter flints, so it’s really OTC.
Dissolve a known quantity of lighter flints in HCl (the alloy is very reactive), filter and you have a RE chlorides mixture of approx. known strength.
Buffering all solutions to pH 5 - 6 may also help precipitating Re oxalates from ferrioxalate solution in the presence of free ferric ions.
[Edited on 23-10-2011 by blogfast25]
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Botanic88
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Blogfast25:
Thanks for your further results.
Also the solubility table in wikipedia must have put the Ks value for zinc chloride, instead of the g/100ml value, by mistake.
I look forward to your final results.
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blogfast25
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Botanic 88:
You meant zinc oxalate (not zinc chloride), right? It seems where Wiki’s solubility table lists only one value at 20 C and in exponential format,
that value is a solubility product but nowhere does it state that clearly…
Counter intuitively, adding FeCl3 to tube 10) didn’t result in the oxalate precipitate to redissolve.
I also added some NaOH (1 M) to some K3FeOx3 solution and it caused Fe(OH)3 to precipitate (Fe(OH)3 is of course highly insoluble - Ksp = 2.5 x
10<sup>-39</sup>. Potassium hexacyanoferrate ((K3Fe(CN)6) doesn’t do
that, so the ferrioxalate complex must be less strong than the hexacyanoferrate complex, as confirmed by:
http://bilbo.chm.uri.edu/CHM112/tables/Kftable.htm
… which gives:
ferrioxalate: Kf = 2 x 10<sup>20</sup>
ferri hexacyanate: Kf = 1 x 10<sup>42</sup>
aluminotrioxalate (AlOx<sub>3</sub><sup>3-</sup>: Kf = 2 x
10<sup>16</sup>
It’s a shame the aluminotrioxalate complex doesn’t have a higher Kf value: it could have been a solution to your problem…
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Botanic88
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Blogfast25:
| Quote: | | 10) 2 ml NdCl3 + 2 ml ferrioxalate: immediate and permanent precipitation, about the same amount as 9). Supernatant liquid remains green, so there’s
excess ferrioxalate. This result is consistent with Nd oxalate being much more insoluble than Ca oxalate: compare it to 2) ......... Counter
intuitively, adding FeCl3 to tube 10) didn’t result in the oxalate precipitate to redissolve. |
An
alternative explanation for this might be that the Nd is precipitating Nd ferrioxalate rather than Nd oxalate. Then the supernatant liquid would
remain green instead of being turned olive green by any free ferric ions being released. Also there would be no reason why adding more free ferric
ions would redissolve the Nd ferrioxalate precipitate.
But in any case I agree with you that there is unlikely to be a practical method to release free ferric ions from the ferrioxalate complex.
Thank you also for the table showing the 'strengths' of various complexes. I wouldn't have known where to find this table. It confirms my impression
that free ferric ions are extremely good at forming complexes and I believe that something like this occurs when ferric chloride fixes gum arabic.
I suspect that the ferric ions form bridges between the -OH bits that stick out the sides of this polysaccharide chain (excuse my lack of correct
terminology). Other chemicals such as borax or free aluminium ions do the same kind of thing, I think. However ferric ions do it much better. For
example the gum doesn't dissolve away in water when it has been fixed with ferric ions, whilst it does so with the other chemicals.
Thanks again for the stuff you have done. The results have confirmed some of my intuitions and you have also lead me to some stuff I didn't know
about.
There are other ways that my photographic process might be improved and I will probably return with more questions in the future!
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blogfast25
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Nd ferrioxalate would in all likelihood be green (or at least somewhat coloured) because the anion is green, so it's unlikely. Nd cations vary in
colour depending on environment.
Ok, 'so long' for now!
[Edited on 26-10-2011 by blogfast25]
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