Purification of a ferric oxalate solution for the Kallitype process

An ideal solution of ferric oxalate (say 20%) for kallitypie should :

- be free of Fe2+ ferrous ions (which would cause fogging in the print highlights)
- be free of free oxalic acid (which would cause the formation of a silver oxalate precipitate when mixed with the silver nitrate solution)
- and therefore be yellow-brown (amber) in color

Ferrous ions appear when the solution ages.

Mainly I imagine by a reduction:

Ferric oxalate → Ferrous oxalate (precipitate?) + carbon dioxide (g)

Fe2(C2O4)3 → 2Fe(C2O4) (precipitate?) + 2CO2 (g)

And possibly under the action of UV light (although very weak in aqueous phase) :

UV light + Oxalic acid → Ferrous oxalate (precipitated?) + Carbon dioxide (g)

hv + Fe2(C2O4)3 (aq) → 2Fe(C2O4) (precipitate) + 2CO2 (g)

As for the free oxalic acid, it is always present in more or less quantity depending on the origin of the ferric oxalate powder used in solution.
If it does not pose a problem for the Pt/Pd process (and is even desirable for clarification), it is on the other hand "deadly" for us kallitypists: Put together, free oxalic acid and silver nitrate form a silver oxalate precipitate which ruins the sensitizer to be coated:

Oxalic acid + Silver nitrate → Silver oxalate (precipitate) + Nitric acid

C2H2O4 + 2AgNO3 → Ag2C2O4 (precipitate) + 2HNO3

Two methods are circulating on the net to neutralize ferrous ions and/or free oxalic acid. (And these 2 methods often tend to be opposed)

The first method neutralizes only oxalic acid and involves the use of ferric nitrate for the following redox (?):

Oxalic acid + Ferric nitrate → Ferric oxalate + Nitric acid

3H2C2O4 + 2Fe(NO3)3 → Fe2(C2O4)3 + 6HNO3

Note the presence of nitric acid as a product of this reaction (not dangerous but can apparently cause an increase in contrast due to its oxidizing power?)

The second method involves the use of concentrated hydrogen peroxide. The addition of hydrogen peroxide actually has two effects (correct me if I'm wrong):

First reaction:

Oxalic acid + Hydrogen peroxide → Carbon dioxide (g) + Water

C2H2O4 + H2O2 → 2CO2 (g) + 2H2O

As product of this reaction CO2 gas is released

Second reaction (oxidation of ferrous oxalate to ferric oxalate in acid medium)

Ferrous oxalate + oxalic acid + hydrogen peroxide → ferric oxalate + water

2FeC2O4 + H2C2O4 + H2O2 → Fe2(C2O4)3 + 2H2O

So this is where things get complicated for me and where I reach my limits in chemistry: I don't know which reaction takes precedence over the other... what is their balance?

In other words, does all the hydrogen peroxide react directly with the free oxalic acid (in which case the second reaction cannot take place and the ferrous ions cannot be removed)?

Also: (unless I was mistaken, but it seems logical): The oxidation of ferrous oxalate to ferric oxalate requires the presence of oxalic acid. So an excess of oxalic acid seems to be necessary for the complete removal of ferrous ions.

If my equations are OK, we can understand the dilemma that tears the kallitypist apart at this point: let's imagine that he has an old amber colored solution containing a lot of ferrous oxalate but finally little free oxalic acid... Trying to eliminate the ferrous ions with hydrogen peroxide, he will finally kill all the oxalic acid necessary for the oxidation of ferrous oxalate into ferric oxalate!

I would therefore conclude that the best procedure for purifying a ferric oxalate solution for kallitypes would be the following (which uses only the hydrogen peroxide method, but after the - counterintuitive - addition of oxalic acid):

1. Test for the presence of ferrous oxalate with potassium ferricyanide. If Prussian blue is formed, go to step 3. If not, go to the next step 2.
2. Test for the presence of free oxalic acid with silver nitrate. If a precipitate is formed, we go to step 4. If not, we are done!
3. Add oxalic acid to the ferric oxalate solution, which turns a beautiful clear emerald green. (Counter-intuitive since we try to eliminate the oxalic acid )
4 (Neutralization of ferrous ions and then of the excess of oxalic acid) Addition of very concentrated hydrogen peroxide (to avoid too much dilution) to the ferric oxalate solution + stirring
5. Repeat step 4. as long as gassing (CO2) is observed. The ferric oxalate solution turns yellow-brown (amber). If no more gassing, go back to step 1.

I am waiting for your feedback! I have no degree in chemistry, just some leftovers from my years in high school and my interest in photochemistry. So please be forgiving with possible/probable errors of reasoning

PS: I know I can throw away a ferric oxalate solution that is too old. But this is not the point of the discussion. What I want to do is to keep my ferric oxalate solution as long as possible. Avoid throwing away as much as possible (for me and maybe for the planet too?) With a small 25ml bottle of ferric oxalate solution, I can coat exactly 1m2 of paper. But I never prepare only 25 ml of ferrix oxalate solution... I rather prepare 100ml and bottle it in 4 small 25ml bottles. That's enough to make 4m2 of print... 3 bottles go directly to the freezer. And I use the remaining bottle the next day. But it's rare that I coat 1m2 of paper the day after, so usually this same bottle goes in the fridge itself. But (I repeat myself) even in these conditions, the ferric oxalate solution contains free oxalic acid...

Not being a chemist I suggest checking with Bostick and Sullivan. I buy my supplies for Pt/Pd printing. I do keep the mixed ferric oxalate #1 and #2 refrigerated and have not had problems. I buy the powder and mix fresh after a while if not doing printing for some time.

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First and foremost... a disclaimer: I have never made a kallitype. However, I am a retired chemistry professor. Thus keep these two things in mind as you read my reply.

Let me start by commenting on the very beginning of your post (shown in italics)...

An ideal solution of ferric oxalate (say 20%) for kallitypie should :

- be free of Fe2+ ferrous ions (which would cause fogging in the print highlights)

The way to insure this is simply to use a high quality ferric oxalate.

- be free of free oxalic acid (which would cause the formation of a silver oxalate precipitate when mixed with the silver nitrate solution)

This is a chemical impossibility. Oxalic acid is a weak acid (its second pKa is about 4.5). Thus the acid form is always in equilibrium with its conjugate base (i.e. oxalate). The higher the pH of the solution the smaller the concentration of the acid, but the solution will never be free of oxalic acid. The presence of iron complicates the exact calculations but very approximately by pH 7 or so the acid form with be less than one or two percent of the total.

- and therefore be yellow-brown (amber) in color


Ferrous ions appear when the solution ages.

This statement is almost assuredly not true. I don't see how iron(II) would form spontaneously in this solution if it is kept in the dark. Further more even if there is a slow evolution of iron (II), in an aqueous solution equilibrated with air any iron (II) formed would be oxidized back to iron (III). Thus the concentration of iron(II) would be expected to be very low.

As for the rest of your post, I'll simply say that 'purifying' a solution that has gone bad for what ever reason is unlikely to be worth the effort and expense involved.

Rather, make smaller quantities that you can use up before the solution goes bad. Or even better... make more prints more often so that the solution doesn't have a chance to go bad! Or, continue as you have been; storing solution frozen which should definitely make it last longer.

fgorga said:

First and foremost... a disclaimer: I have never made a kallitype. However, I am a retired chemistry professor. Thus keep these two things in mind as you read my reply.

Let me start by commenting on the very beginning of your post (shown in italics)...

An ideal solution of ferric oxalate (say 20%) for kallitypie should :

- be free of Fe2+ ferrous ions (which would cause fogging in the print highlights)

The way to insure this is simply to use a high quality ferric oxalate.

- be free of free oxalic acid (which would cause the formation of a silver oxalate precipitate when mixed with the silver nitrate solution)

This is a chemical impossibility. Oxalic acid is a weak acid (its second pKa is about 4.5). Thus the acid form is always in equilibrium with its conjugate base (i.e. oxalate). The higher the pH of the solution the smaller the concentration of the acid, but the solution will never be free of oxalic acid. The presence of iron complicates the exact calculations but very approximately by pH 7 or so the acid form with be less than one or two percent of the total.

- and therefore be yellow-brown (amber) in color


Ferrous ions appear when the solution ages.

This statement is almost assuredly not true. I don't see how iron(II) would form spontaneously in this solution if it is kept in the dark. Further more even if there is a slow evolution of iron (II), in an aqueous solution equilibrated with air any iron (II) formed would be oxidized back to iron (III). Thus the concentration of iron(II) would be expected to be very low.

As for the rest of your post, I'll simply say that 'purifying' a solution that has gone bad for what ever reason is unlikely to be worth the effort and expense involved.

Rather, make smaller quantities that you can use up before the solution goes bad. Or even better... make more prints more often so that the solution doesn't have a chance to go bad! Or, continue as you have been; storing solution frozen which should definitely make it last longer.

Click to expand...

Thanks for these very interesting replies Frank !

I have made a small test with 2 drops of 20% ferric oxalate, and added a couple of crystals of ferric nitrate to one : the color shift to amber is obvious, which should indicate that the oxalic acid is indeed getting eliminated from the ferric oxalate solution :

largo said:

Thanks for these very interesting replies Frank !

I have made a small test with 2 drops of 20% ferric oxalate, and added a couple of crystals of ferric nitrate to one : the color shift to amber is obvious, which should indicate that the oxalic acid is indeed getting eliminated from the ferric oxalate solution :

Click to expand...

I am not an authority on the subject by any stretch of the imagination, but the brown color would also mean you have an excess of ferric now. Ferric oxalate is not a simple compound with 2 Fe(3)'s and 3 (C2O4)'s but kind of bunch of them associating to form a structure that gives that green color. If you add more of iron and the structure gets broken up and the color changes.

Rather than doing an experiment on drops, I would do one using ferric nitrate solution and adding it to the sensitizer solution in incremental amounts (not worry about whether there is "excess" oxalic acid or not) and see if precipitates redissolve. Once you have found a minimum amount, then study its effect on the speed and contrast and hope they do not become impractical.

:Niranjan.

This is a fascinating thread to me. All the time I've spent trying to get my head wrapped around ferric oxalate really quickly goes over my pay grade since it's apparently a very complicated substance. At this point my pedestrian understanding is that if you buy good, pure stuff in powder form (I've had best luck with Bostick & Sullivan), it's more yellow than green.

Then mixed up, the solution is always going to be a little greenish, and should be fine for 2-3 months (if stored out of heat and UV light). At some point after that, it seems to precipitate out and form blue crystals, leaving behind water and ???. That's what happened to the batch I mixed up of 20% aqueous solution a little over a year ago. One big blue crystal and some clear liquid.

The interesting thing is that before I got my "good" powder, I bought some from the Formulary. They were very up-front telling me that it was old back stock that had been on the shelf for 20+ years and might not be good anymore, and IIRC they gave it to me for free when I said I still wanted it. I never could get it to stay in solution for more than a couple of hours at a time, and every Kallitype I ever tried to make with it had horribly fogged highlights. The powder itself seemed pretty caked together into smallish rocks, and was a pronounced greenish color.

I've succeeded in mixing up very small quantities of 20% ferric oxalate solution that is working really well for my Kallitype printing with no fogging, using the B&S sourced powder. And I have enough of that powder to last me many years. But I still have that Formulary powder too, quite a lot of it. If there was some home-chemist-accessible way of getting it back into a usable state like the B&S powder, I'd be very interested. I gather that the process of synthesizing ferric oxalate is pretty complicated and tricky to do correctly in terms of producing a good quality result.