Sciencemadness Discussion Board - A puzzle about the precipitation of calcium oxalate

A puzzle about the precipitation of calcium oxalate

Botanic88 - 19-10-2011 at 01:22

I am puzzled about something that occurs when I precipitate calcium oxalate.

(1) Adding calcium chloride soln to potassium oxalate soln gives an abundant precipitate instantly, as expected.

(2) Adding calcium chloride soln to ammonium iron (III) oxalate soln also gives an abundant precipitate, but it forms over a number of seconds. This delay is probably because the 'ferrioxalate' complex needs to be pulled apart.

(3) Adding a drop or two of iron (III) chloride soln to the ammonium iron (III) oxalate soln, before adding the calcium chloride, either results in no precipitate or one that takes many minutes (or an hour) to occur and then it is not abundant.

What is going on here? I would have thought that iron (III) chloride would be a by-product of the reaction between calcium chloride and ammonium iron (III) oxalate. In that case why would a little iron (III) chloride added at the start prevent a precipitate, whilst the same chemical produced as a by-product seems to allow the precipitate to occur?

AndersHoveland - 19-10-2011 at 02:32

Perhaps it reacts to form some tempory compound which acts as an inhabitor towards crystal growth. This compound may be either Fe(OH)3 or invissible quantities of precipitated iron oxylate.

IPN - 19-10-2011 at 03:31

I think you are just witnessing Le Chatelier's principle in action. Meaning that having that excess iron chloride present in the solution is pushing the equilibrium to the left causing the observed slowing and/or stopping of the reaction.

For further reading, wikipedia is a good place to visit:
http://en.wikipedia.org/wiki/Le_Chatelier%27s_principle

[Edited on 19.10.2011 by IPN]

blogfast25 - 19-10-2011 at 05:55

Quote: Originally posted by IPN

I think you are just witnessing Le Chatelier's principle in action. Meaning that having that excess iron chloride present in the solution is pushing the equilibrium to the left causing the observed slowing and/or stopping of the reaction.

[Edited on 19.10.2011 by IPN]

Hmm... although exact quantities aren't mentioned. a 'drop' of ferric chloride shouldn't really be enough to affect the equilibrium significantly, although much depends on the equilibrium constant (complexation constant) of the oxalo ferric complex. But it's a possible explanation...

Fe3+ + 3 Ox2- < === > Fe(Ox)3-

Excess Fe3+ would drive that to the right, reducing strongly [Ox2-] available for precipitation of CaOx.

And the solublity product of CaOx also defines whether or not precipitation occurs: Ks = [Ca2+] x [Ox2-]

If you have values for the equilibrium constants, then playing around with the equations (and making some simple assumptions) you could determine in which conditions precipitation is likely to occur and in which not.

And if correct, it's potentially a neat little trick to 'mask' oxalate...

[Edited on 19-10-2011 by blogfast25]

Botanic88 - 19-10-2011 at 06:40

Thanks for your suggestions so far.

I didn't measure the quantities accurately but
(1) The ammonium iron (III) oxalate soln was fairly strong; visually about 4 times as much water as the amount of crystals in the test tube.
(2) I used about 2 mls of this soln.
(3) The calcium chloride soln was also fairly strong.
(4) I added first 1drop of strong iron (III) chloride soln, and on a second run I added 2 drops.

I am not convinced that Le Chatelier's principle would explain the big difference in the abundance of the precipitate and the time delay, when such a small proportion of Iron (III) chloride is added.

The suggestion that ferric hydroxide or oxalate might inhibit the growth of calcium oxalate crystals in some way is interesting; I believe both these chemicals can be 'strange' and they are not always clearly defined substances. So who knows what they might do!

Before doing this experiment I had believed that the 'ferrioxalate' complex was held together very strongly. I never managed to pull it apart. But this was because I have always used mixtures of ferric chloride and oxalates. So it came as a surprise to me that the complex could be pulled apart easily, provided there is no ferric chloride around initially!

blogfast25 - 19-10-2011 at 07:08

@Botanic88:

The only way really is to find the equilibrium constants and play with the equations, in order to see if leverage caused by small amounts of extra Fe3+ can suppress calcium oxalate formation. If conditions exist in which a small amout of free Fe3+ cannot prevent the solubility product of calcium oxalate to be exceeded and yet no precipitate forms, then we can start considering other plausible causes...

Botanic88 - 20-10-2011 at 07:11

Unfortunately I'm not familiar with equilibrium constants. I've looked them up but it would take me a while to understand them and use them properly.
So I am stuck with thinking about the matter in more qualitative terms for the time being!

Take the balance you mention blogfast25: Fe3+ + 3 Ox2- < === > Fe(Ox)3-

I follow what you say, that adding Fe3+ ions initially could push the balance rightwards, thus preventing any free Ox2- taking part in precipitation.
But removing Ox2- ions from the left (by precipitation) would push the balance leftwards.
This would eventually result in the same amount of Fe3+ ions on the left as the amount that I put there when I added a drop of ferric chloride initially.
So the precipitation should be self-limiting. But that's not what appeared to me to be happening!

Incidentally are you sure that the dissociation of 'ferrioxalate' ions occurs in this balanced way? I think I imagined that all the available ferric and oxalate ions would form complexes, as much as possible.

******************************************************************************************
Overnight I had a different idea about what could be happening ......
Suppose the precipitate is not actually calcium oxalate. Is it possible that the calcium and ammonium ions simply swap over to produce calcium ferrioxalate?

I have searched for this substance on the internet without success.
But if it exists then it could be insoluble in water but soluble in weak ferric chloride soln.

With this in mind I repeated my earlier experiment but this time I added a drop of ferric chloride to one sample initially (as I had done previously).
Then I added calcium chloride to both samples. One sample precipitated and the other didn't (as happened previously).
Finally I added a drop of ferric chloride to the second sample, to make both samples 'equal'.

The precipitate didn't redissolve quickly, but there maybe some signs that it is doing so very slowly!

blogfast25 - 20-10-2011 at 08:52

Botanic88:

Applying equilibrium theory when multiple equilibria are at play at once is not easy and somewhat mathematically challenging. It involves setting up a system of simultaneous equations in the concentrations of the various species present and solving this system, usually by simplification/iteration.

Ignoring deliberately that oxalic acid is a fairly weak acid (that just adds another equilibrium to the mess!), the species in solution and of interest are Fe3+, Ca2+, Ox2-, Fe(Ox)3 (3-), NH4+. I’m assuming the solution is buffered, so [H3O+] and [OH-] are constant.

Five species would require five independent and simultaneously valid equations to obtain a solution. The five equations are:

* the complex equilibrium for Fe(Ox)3 (3-)
* the solubility product of CaOx
* the neutrality requirement: 3 [Fe3+] + 2 [Ca2+] + [NH4+] + [H3O+] = 3 [Fe(Ox)3 (3-)] + 2 [Ox2-] + [OH-]
* mass balance for iron: cFe = [Fe3+] + [Fe(Ox)3(3-)] with cFe the overall iron concentration in solution
* mass balance for oxalate: cOx = [Ox2-] + 3 [Fe(Ox)3 (3-)] with cOx the overall oxalate concentration in solution

Writing out all the equations, making some reasonable assumptions and simplifying the resulting algebraical mess somewhat, would then allow to determine whether or not the solubility constant of calcium oxalate was exceeded for given values of cFe, cOx and the actual equilibrium constants. No mean feat! Today such calculations would usually make use of powerful mathematical software to plot solutions… Complex sets of equations sometimes show quite intuitively unexpected solutions. The trick would be to extract from this set of equations an expression for [Ca2+] x [Ox2-] = F(cFe, cOx, KFe(Ox)3, Ks, …) with F a function. This would tell in which ‘zones’ [Ca2+] x [Ox2-] > Ks (precipitation occurs) and in which [Ca2+] x [Ox2-] < Ks (no precipitation occurs).

Ferric oxalate dissociates undoubtedly in three distinct steps, much like a triprotic acid (see H3PO4) does. But for the above purpose the overall equilibrium constant could be used.

You’re referring to calcium ferric oxalate but I believe that that is soluble, otherwise using ammonium ferric oxalate to precipitate calcium oxalate makes no sense at all.

I'd really like to replicate your experiment but I don't have any ammonium ferric oxalate. Bugger. Synthesis isn't complicated but adds time. Maybe one fine day...

[Edited on 20-10-2011 by blogfast25]

Botanic88 - 21-10-2011 at 06:49

Blogfast25:

Thanks for taking the trouble to give me all this detail.
If I decide to find out more about equilibrium constants I will certainly try a much simpler scenario to start with!

I didn't buffer the solutions but I was aware that the acid nature of ferric chloride might affect the experiment.
So I checked at the start that it wasn't just acid that prevented the precipitation.
I substituted the single drop of ferric chloride with 4 drops of HCL and this didn't stop the precipitation so much.
(The HCL was from a 50:50 mix of conc HCL and water)
So I am pretty sure that the small amount of ferric chloride is blocking the precipitation somehow.

****************************************************************************************************
At this point I should explain that I didn't do the experiment for purely chemical reasons.
My main intention is to improve an alternative photographic process that I invented in the 1970's.

This photographic process requires the use of the ferrioxalate complex to reduce the iron at a decent rate by photo-reduction
but after the exposure has been made it requires simple ferric ions to fix some pigmented gum.
I have tried various ways to break up the ferrioxalate complex in the past, without much success!

I used not to have any ammonium ferrioxalate. I just mixed up ferric chloride, potassium oxalate and/or oxalic acid and HCL
and discovered which proportions worked, by trial and error.
But recently I purchased some of the pure chemical for the first time, which is how I came to do the experiment!

Perhaps I need to to post a new question, simply asking if people have any idea how to get simple ferric ions from the ferrioxalate complex, in a photographic context.

Anyway thanks again for taking time to look at my question.

blogfast25 - 21-10-2011 at 12:07

Yes, acid could have influence because it pushes the dissociation of oxalic acid back, thereby reducing the amount of free oxalate. But calcium oxalate forms from oxalic acid plus Ca2+, so it shouldn’t be too much of an issue.

I synth’ed some potassium trisoxaloferrate (III) today and will be running my own tests with FeCl3 etc, to see if small amounts of Fe3+ can hinder the precipitation of CaOx and whether I can reproduce your results...

Botanic88 - 22-10-2011 at 00:03

Blogfast25:

Okay, I look forward to seeing your results.

Meanwhile I will go ahead with my more open question as a new topic.
i.e. has anyone got an idea how to get simple ferric ions from the ferrioxalate complex, in a photographic context.