jan1234
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determination of dissociation constant
hey!
i've got this task to develop a quick test to check for free metal-ions in an aqueous solution of some metal-amio acid-complexes.
so far, i have no experimental data on these complexes. i don't know how fast the equilibrium is reached again after precipitating parts of the free
floating metal ions. if i'm able to recipitate all free metal ions while the equilibrium is not reached again, i can go on further than until now.
my favorite one of these complexes is one of copper(ii), so it's quite easy to follow the reactions the color changes. until now i tried to
precipitate Cu(OH)2 via NH3, but since the Cu(NH3)4(H2O)2SO4 complex is soluble in NH3-solutions, i dont think this is the route to go.
i have another indicator here, bis(triethylamino)chloroanilic acid. i couldn't find literature about this one, except for the preparation in a rather
hard to find article in a book (?) i can't even find anymore. it works so far, there are colorchanges and precipitate and so on, but i'm rather
sceptical...
does anyone have input on this one? should i determine the dissociation constants (via potentiometric titration f.ex.?) to check for other ways to
detect free metal ions; or should i go on with rather trial-and-error experiments to check wether the future "test" on this base is suitable or not.
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Sulaiman
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Bubbling H2S precipitates most metallic ions from solution as insoluble sulphides that could be weighed ?
or
A cheap TDS meter may suffice ?
CAUTION : Hobby Chemist, not Professional or even Amateur
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Sir_Gawain
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Ammonia does work to make copper hydroxide, you just need to make sure there is an excess of copper sulfate. I actually prefer using ammonia because
there’s no sodium contamination, plus sodium hydroxide tends to decompose copper hydroxide if it’s too concentrated.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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jan1234
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Quote: Originally posted by Sulaiman  | Bubbling H2S precipitates most metallic ions from solution as insoluble sulphides that could be weighed ?
or
A cheap TDS meter may suffice ? |
hi, thanks for your answer!
the point is not to exactly determine the Text amount of free ions, but from the metal in the complex.
for example: our Cu-complex is made from copper sulphate, and since it's an industrial process there will be cu(ii) from the sulphate in the product.
our goal is now to determine the amount of 'free inorganic copper' or at least a semi-quantitative way of saying 'this is TOO much free cu(ii)' vs.
'that's ok".
that's the point of determining the dissociation constant. if the complex doesn't dissiciate a noticable amount after let's say max. 30 minutes, it's
way easier dermining the amount of free copper. just precipitate it and weigh it.
i'm almost at the point of trying iodometric titration. but since i don't know how stable this complex is (should be stable in 'common conditions',
i.e. pH > 2.5 and < 10.5, no harsh reductants/oxidants, air-exposure), i can't rely on those results i guess.
one of our competitor
[Edited on 7-11-2024 by jan1234]
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jan1234
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| Quote: Originally posted by Sir_Gawain | | Ammonia does work to make copper hydroxide, you just need to make sure there is an excess of copper sulfate. I actually prefer using ammonia because
there’s no sodium contamination, plus sodium hydroxide tends to decompose copper hydroxide if it’s too concentrated. |
Hey!
Well excess of CuSO4 isn't possible, since i want to filter the precipitate and weigh it to determine the amount of Cu2+.
I don't know yet if this is even possible. There is no data on stable the complex is against various chemicals, so i don't know how to precipitate the
cu2+ ions that come from impurities (precursor is a copper(ii) salt) and from the dissociation process of the complex. Thats why i want to determine
the KD, to know if i even have a chance to precipitate all cu2+, filter, titrate part of it, let stand in solution x tomes for peroid y and titrate
every time and so on...
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DraconicAcid
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You can't determine the presence or quantity of free copper ions in solution by adding something that precipitates the copper, as that will probably
also react with your complex.
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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RU_KLO
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Maybe (some ideas)
1) by masking/protecting the complex? so it will not be precipitated when adding the precipitant (dont know if this is a word)
2) adding a solvent, where the complex is soluble, but the salt/metal not.
3) using destilation/boiling - if you dont want to recover the complex- to dryness to remove the amino-etc-complex, whats left is your excess salt.
4) using EDTA (maybe not with copper): "EDTA is mainly used to sequester (bind or confine) metal ions in aqueous solution. In the textile industry, it
prevents metal ion impurities from modifying colours of dyed products."
[Edited on 8-11-2024 by RU_KLO]
Go SAFE, because stupidity and bad Luck exist.
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bnull
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| Quote: Originally posted by RU_KLO | | 1) by masking/protecting the complex? so it will not be precipitated when adding the precipitant (dont know if this is a word) |
That's precisely the word.
Try sodium or ammonium oxalate. Copper (ii) oxalate is practically insoluble in water.
Quod scripsi, scripsi.
B. N. Ull
P.S.: Did you know that we have a Library?
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DraconicAcid
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Depending on the Kd, oxalate can react directly with the complex to kick off the ligands and precipitate copper(II) oxalate. Unless you add too much
oxalate and get [Cu(C2O4)2]2-.
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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bnull
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Ok, I forgot the stability of the complex toward oxalate. Crap.
Norman C. Li and Edward Doody wrote some papers about zinc and copper complexes of amino acids. Polarographic and potentiometric studies etc. I
recommend you take a look. They may give you an idea of what to do.
If there was a way to precipitate the amino acid complex...
Edit: More info on stability of amino acid complexes:
G. Berthon, The Stability Constants of Metal Complexes of Amino Acids with Polar Side Chains, http://publications.iupac.org/pac/pdf/1995/pdf/6707x1117.pdf
O. Yamauchi, A. Odani, Stability Constants of Metal Complexes of Amino Acids with Charged Side Chains--Part I: Positively Charged side chains,
https://www.degruyter.com/document/doi/10.1351/pac1996680204...
[Edited on 8-11-2024 by bnull]
Quod scripsi, scripsi.
B. N. Ull
P.S.: Did you know that we have a Library?
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jan1234
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| Quote: Originally posted by DraconicAcid | | You can't determine the presence or quantity of free copper ions in solution by adding something that precipitates the copper, as that will probably
also react with your complex. |
That's exactly what i am looking for. A reagent that won't react with my complexes, but precipitate the free ions. If the equilibrium shifts towards
the metal ions after precipitating them, due to dissociation of the complex - thats fine if it isnt too fast. Measuring the exact value of m+-ions vs.
Complexed m+ isnt possible by this way, but it doesn't need to be that exact for starters.
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DraconicAcid
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| Quote: Originally posted by jan1234 | | Quote: Originally posted by DraconicAcid | | You can't determine the presence or quantity of free copper ions in solution by adding something that precipitates the copper, as that will probably
also react with your complex. |
That's exactly what i am looking for. A reagent that won't react with my complexes, but precipitate the free ions. |
You won't find one. Copper complexes are labile.
I use sodium anthranilate for gravimetric determination of copper, but it doesn't work in the presence of oxalates or amino acids- you get completely
different precipitates.
The only way you're going to find the concentration of uncoordinated copper is potentiometry, assuming there aren't any other oxidizing agents in
solution. The Nernst Eq'n in your friend.
Spectrophotometry is unlikely to work, because copper(II) complexes tend to all be relatively the same colours.
[Edited on 11-11-2024 by DraconicAcid]
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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jan1234
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| Quote: Originally posted by DraconicAcid | | Quote: Originally posted by jan1234 | | Quote: Originally posted by DraconicAcid | | You can't determine the presence or quantity of free copper ions in solution by adding something that precipitates the copper, as that will probably
also react with your complex. |
That's exactly what i am looking for. A reagent that won't react with my complexes, but precipitate the free ions. |
You won't find one. Copper complexes are labile.
I use sodium anthranilate for gravimetric determination of copper, but it doesn't work in the presence of oxalates or amino acids- you get completely
different precipitates.
The only way you're going to find the concentration of uncoordinated copper is potentiometry, assuming there aren't any other oxidizing agents in
solution. The Nernst Eq'n in your friend.
Spectrophotometry is unlikely to work, because copper(II) complexes tend to all be relatively the same colours.
[Edited on 11-11-2024 by DraconicAcid] |
thanks! that's valuable input 
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