[Chelsea]

"H4EDTA, the precursor to that ligand. At very low pH (very acidic conditions) the fully protonated H6EDTA2+ form predominates, whereas at very high pH or very basic condition, the fully deprotonated Y4− form is prevalent." from wiki

Can EDTA be used for chelating metal ions (eg. Fe2+, Fe3+, Mn2+) in acidic medium ? if it become fully protonated and is positively charged,ie. H6EDTA2+, can it still act as ligand to positive metal ions? Or does it work only in akali medium, when it is negatively charged and the "six-toothed" ligand is deprotonated?

However, some literature show that iron is more solubilized under ph 3 to 4, can EDTA be used under such condition? ( or that EDTA do not need to iron to be solubilized in the first place?)

Also, I am not sure about the difference in practical uses of disodium and tetrasodium salt of EDTA. Disodium EDTA is acidic in water while tetrasodium one is alkaline. Which one is better in chelating iron from its oxides?

Thx!

[Borek]

You need to select pH (and buffer the solution) depending on all existing equilibria, as they compete for the ions. For Fe³⁺ high pH would mean precipitation of Fe(OH)₃, so we have to use pH that is low enough to not allow precipitation, but high enough for the presence of EDTA^4-. Generally speaking for ions with higher charge (like Fe³⁺) stability constants for EDTA complexes are high enough to allow titration in relatively low pH.

[Chelsea]

thx again and sorry for posting on the wrong board 

I am trying to find the best pH for my work, and may start with ph 7.0 buffer and then higher ph (7.5 8.0 .) and compare the results. it seems unnecessary to try below ph 7.0 as EDTA becomes positively charged ?

thx 

[Borek]

As explained earlier, in high pH you risk hydroxide precipitation.

If you are looking for a pH where the highest fraction of Fe³⁺ is complexed, you should start with equilibrium calculations. While they will be never exact, you should at least be able to find the best range for your tests. My bet is below 7.

[Chelsea]

I still dont understand, why EDTA below 7 can still chelate if its positively charged?
Thx!!

[Borek]

It is a matter of equilibrium. If the stability constant is high enough, you don't need 'naked' Y to complex the metal:

Me + H_nY MeY + nH⁺

(not balanced, but the idea should be obvious).

[Babcock_Hall]

"Chemical Separations and Measurements"  by Peters, Hayes, and Hieftje suggests using pH 4 and salicylic acid as an indicatir for ferric ion titrations with EDTA.  Salicylic acid makes a nice purple complex with iron, and this is the basis for an old test for overdoses of asprin, BTW.

[danteOne]

According to the textbook "Exploring Chemical Analysis" by Daniel C. Harris the pKa values of EDTA are as follows. For the carboxyl protons the pKa's are pKa1 = 0.0 , pKa2 = 1.5 , pKa3 = 2.0 and PKa4 = 2.69. For the ammonium protons the pKa's are pKa5 = 6.13 and pKa6 = 10.37.

The text states "Below a pH of 10.24 most EDTA is protonated and is not in the form Y^4- that binds to metal ions. (281)"
On the next page the text states that "A metal-EDTA complex becomes unstable at low pH because H+ competes with the metal ion for EDTA. At too high a pH, OH^- competes with EDTA for the metal ion and may precipitate the metal hydroxide or form unreactive hydroxide complexes"(282).

The answer to your question is ETDA is only useful at pH above 10.24. I don't know exactly why the protonated form doesn't work but I bet it has something to do with ETDA's specialized shape.

[Borek]

The answer to your question is ETDA is only useful at pH above 10.24

Not true, it depends on the stability constant of the complex. If the stability constant is high enough, conditional stability constant (which takes into account all side reactions, like protonation of the ligand), can be still high enough to allow practical use of the complexing agent.

Compare with the procedure here: http://www.titrations.info/EDTA-titration-aluminum - Al³⁺ is titrated at pH of 5.5. And it works, even if the kinetics forces us to do the back titration.