flakestoday
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Metal Complex Questions
Hello all, I am a newcomer to chemistry and this forum, and I had a few questions about metal coordination complexes in general, and specifically
K3(Fe(C2O4)3)*3H2O.
So last week I had a gen-chem lab practical in which we prepared the above iron coordination complex, and the point of this specific practical was to
carefully follow a procedure. They never said anything about the complex we made, though, and I was curious. I know that there is huge variety in the
properties of complexes like this, but would that specific complex be of any actual use in a lab? Also, since this is an ionic compound, does it
dissociate in water? If it dissociates, couldn't it upon coming out of solution rearrange in a different format, for example with a different ratio of
oxalate to iron? I also have the procedure for preparation of the above iron complex if anyone is interested. Thanks!
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Steam
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Excellent question Flakes, To answer your questions, what you made is a potassium ferrooxalate type complex. While I don't know of any specific uses
of the top of my end, you can always just use the cop-out answer of potential uses in pigments and catalysts. In addition, I bet it thermally
decomposes to form finely divided particles of iron oxides which could always be used for something... maybe ferrofluid if you stabilize them.
On your next question you mention this is an "ionic compound" and that it might dissociate in water. While it is true that it probably would be
soluble in water, it would be wrong to classify this compound to that of a salt like sodium chloride or potassium nitrate. What we have here is what
is called a n organo-metallic (as you mentioned) and thus there is a "dative" bond between the metal center and the oxalic acid. One can really get
into the weeds with the exact nature of the bond and there are several theories which predict bonding such as MO-theory. If you are interested, I
highly encourage you to pick up the following textbook which is kind of a rite of passage for all upcoming chemists (especially in the western US). I
would start at the chapter where it teaches you about high-spin and low-spin complexes and how to fill d-orbitals.
https://www.amazon.com/gp/product/0321811054/ref=ppx_yo_dt_b...
Now to your question, because there is actually a bond between the metal and organic, there will not be dissociation in water. The potassium is
another story- it is attracted to the complex through coulombic forces and thus will likely dissociate in water. The waters in the formula are waters
of hydration which are actually bonded to the structure. In fact, they will remain on the structure even at temperatures passed the boiling point of
water. Hard to say the exact temperature which they well be liberated from the structure but my guess would be around 200C. Solubility science is
EXTREMELY complicated and most professors just try to give the hand-wavy explanation of "like-dissolves-like" and pray to their chosen god/gods that
the student doesn't ask for a further explanation. If you want to learn more about solubility than your Lab-TA or even some professors I suggest
reading this short pdf by Abbott:
https://www.stevenabbott.co.uk/practical-solubility/
When it does dissolve in water, it will likely change its shape a bit... but it might be negligible amounts. What is the important take away is that
the Bonds will not be broken so it will keep its overall structure. If you do indeed have a true complex, the washing with water should not remove
oxalate to any significant degree. Often, washing is actually used to purify organometallics to remove excess salt. Washing with your typical lab
water might exchange out some of the potassium ions for sodium something similar. This is to do with HSAB theory (which If you learn anything from
this response learn that HSAB theory actually predicts a great deal of chemical interaction and actually makes a ton of chemistry make sense; I have
no idea why they don't teach it earlier in ones academic career since all you really need to know to use it qualitatively is a basic understanding of
molecular orbitals) , but essentially because the oxalate is a "hard" base it will prefer to complex with hard ions . Sodium is more "hard" than
potassium and thus it will preferentially interact with sodium.
I hope that was what you were looking for in a reply, I'll happily answer any additional questions you have
Cheers!
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DraconicAcid
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This isn't an organometallic compound; it's a coordination compound. An organometallic compound would have a direct metal-carbon bond such as that
found in a carbonyl compound or an alkene complex.
It is ionic, but when it dissolves in water, it will ionize to give potassium ions and tris(oxalato)ferrate(III) ions- the oxalate ligands will not
detach from the iron. Oxalates are ligands; the potassium is just a counterion.
Hard and soft acid/base theory really isn't relevant to this compound- whether you have sodium ions or potassium ions as the counterions (or which
compound is more stable or less soluble) has nothing to do with the hardness of the cation, but the size of it relative to the anion. Washing the
compound with cold water isn't going to exchange out your ions- you'd have to dissolve it in a concentrated solution of a soluble sodium salt and see
which preferentially precipitates.
You can read a bit about coordiantion compounds here: https://openstax.org/books/chemistry-2e/pages/19-2-coordinat... (if your textbook doesn't have a chapter on it).
[Edited on 22-2-2020 by DraconicAcid]
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Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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flakestoday
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Thank you all for the replies! I suspected that this stuff could be very interesting.
| Quote: |
In addition, I bet it thermally decomposes to form finely divided particles of iron oxides which could always be used for something... maybe
ferrofluid if you stabilize them.
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This is very interesting, I watched NileRed's videos on ferrofluids a while back and it seemed to be harder than one would think to get finely enough
divided magnetite nanoparticles. I wonder if there exists a coordination complex which has both iron(II) and iron(III); maybe upon decomposing in air
this could yield magnetite nanoparticles? I bet this could also be useful for thermites!
It seems strange to me that this wouldn't be considered be considered an organometallic, just because the oxalate ion does definitely have carbon, I
suppose I should do some more reading on this. Would something like NaCN be considered an organometallic?
Steam, I am very excited to give the Abbott pdf a read, solubility has always been kind of a mystery to me and until now I have just contented myself
with memorizing the solubility rules -- hopefully this will give some clarity! Do you know if the third edition of the Miessler text is as good as the
fifth? The third edition was the only one on amazon for a price that a broke college student can afford haha!
Thanks everyone for the info and for the reading suggestions!
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DraconicAcid
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Quote: Originally posted by flakestoday  |
It seems strange to me that this wouldn't be considered be considered an organometallic, just because the oxalate ion does definitely have carbon, I
suppose I should do some more reading on this. Would something like NaCN be considered an organometallic? |
It contains carbon, but there's no carbon-iron bond, and therefore not organometallic. We classify stuff as organometallic because they have chemical
properties that sets them apart from organic compounds and traditional inorganic compounds. Metal oxalates and carbonates (and acetates, and
formates, etc) don't share these properties- they're just ordinary inorganic compounds (even though the anions contain carbon).
No textbook on organometallic chemistry is going to waste chapters describing the properties of metal oxalates, carbonates, and acetates, since their
properties aren't significantly different from sulphates, sulphites, or phosphates. Just like a textbook on organic chemistry won't devote chapters
to metal carbonates or oxalates.
Sodium cyanide isn't organometallic, as it's completely ionic- there is no bond between the sodium and the carbon.
If you have a coordination compound with cyanide, such as K3[Fe(CN)6], then you have a covalent bond between the carbon of the cyanide and the iron.
But even that isn't generally considered a true organometallic compound, as cyanide is considered a pseudohalide, and not an organic bit.
You might think, "Well, that seems a bit arbitrary!" It is. But all of our definitions of what is or is not a particular class of compound will have
some arbitrariness to it- compounds exist on their own, not to fit into our definitions, and we have to draw our lines somewhere. It's just like
defining what is and is not an organic compound- no organic chemist will consider sodium carbonate to be an organic compound, even though it contains
carbon.
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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DraconicAcid
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| Quote: Originally posted by flakestoday |
| Quote: |
In addition, I bet it thermally decomposes to form finely divided particles of iron oxides which could always be used for something... maybe
ferrofluid if you stabilize them.
|
This is very interesting, I watched NileRed's videos on ferrofluids a while back and it seemed to be harder than one would think to get finely enough
divided magnetite nanoparticles. I wonder if there exists a coordination complex which has both iron(II) and iron(III); maybe upon decomposing in air
this could yield magnetite nanoparticles? I bet this could also be useful for thermites! |
If you want to get iron oxide particles, you don't want to thermally decompose the coordination compound (with all of its potassium in there), but
iron(II) oxalate. It's easily made, and easily decomposed to give a fine powder which contains iron(II) oxide and iron metal (this powder ignites on
exposure to air, with sparkles if you disperse it). Expose this powder to oxygen at the right temperature, and you can probably get the iron oxide of
your choice.
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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Bedlasky
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Hi.
Ferrioxalate can dissociate in water. There are equilibriums depending on concentration and pH.
About properties - ferrioxalates are photosensitive. Under UV light they decomposes in to ferrous ions, oxalates and carbon dioxide. This property is
used for making blueprints. NileRed have about it two videos:
https://www.youtube.com/watch?v=0e8CMbHfLxM
https://www.youtube.com/watch?v=hYqn8CO2P3E&t=122s
He used ammonium ferric citrate, but potassium ferrioxalate can be used as well.
This reaction is also used for determination of oxalate (CO2 which is form by this reaction is measured by CO2 electrode).
There are more ferric complexes which are also photosensitive - for example citrate and tartrate complexes. I wrote about this complexes article on my
website.
https://colourchem.wordpress.com/2019/06/25/photochemical-ox...
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Boffis
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Another use for trioxalatometallates of this type is as a selective precipitant for certain amino acid. At one time the iron salt described above or
the equivalent Al3+ and Cr3+ complexes were proposed as reagents for the seperation of amino acid formed by the hydrolysis of proteins. I can't
remember which amino acids are precipitated by this group but I stumbled across several papers in the past in journals such as the Journal of
Biological Chemistry and the Biochemistry Journal both are available freely online, atleast for the older stuff that we are interested in here.
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DraconicAcid
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I seem to remember reading something like that- I think it was glycine. And it formed a precipitate with a weird stoichiometry- something like 11
ammonium ions, one glycinium ion, and four tris(oxalato)ferrates.
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flakestoday
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DraconicAcid that is very interesting about the classification of organometallics. Although a bit arbitrary, it does make sense when you consider the
ionic nature of oxalates or carbonates, as you said. It is hard to remember that this stuff doesn't exist to fit our rules sometimes, I guess it's
human nature to think like that though.
Another good point about potassium contamination - if I remember correctly NurdRage had a video about pyrophoric iron and iron oxalate, probably from
the method you described. I'll have to give that a watch at some point soon.
Bedlasky, that photo-sensitivity is very interesting - I bet there is an interesting rabbit hole to go down with ferric complex photo-sensitivity and
blueprint making. I will definitely check out your article and those NileRed videos, thank you for the info.
Boffis, this is interesting as well. Is this related to/the same thing as "salting out" with amino acids? I remember reading something about this
somewhere.
I knew that there was bound to be some cool chemistry behind these complexes, thank you all for the information! I now have some exciting reading in
front of me about this stuff.
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