[PoetryInMotion]

Hi! I am currently facing the following exercise:

Give the chemical equation for the EDTA-Na₂ titration reaction for
a) magnesium ions at pH 10 and
b) zink ions at pH 5 and

The answer on a) is found in the text book:
[tex]\mathrm{Mg^{2+} + H_2EDTA^{2-} \longrightarrow\: MgEDTA^{2-} + 2H^+}\,.[/tex]
At such a high pH as the one in b), however, the EDTA^4- ion isn't stable at all (α_Y4-≈0), which makes me uncertain of whether it is possible to form any ZnEDTA^2- complex at all :/  A table in my book says that EDTA and zink ions reacts quantitatively at pH 4-5, so I guess some kind of complex must be formed, but as far as I know, only EDTA^4- complexes are possible... Anyone who knows? Thanking you in advance

[Borek]

however, the EDTA^4- ion isn't stable at all (α_Y4-≈0)

Please elaborate - that's the first time I hear about EDTA^4- being unstable. At pH high enough it dominates EDTA solutions. At pH around 4-5 its concentration is very low, but it has nothing to do with the stability (could be it is a matter of terminology).

In both cases the product is MeEDTA^2-, so the reaction equation is just a matter of finding the dominating H_nEDTA^-4+n ion.

[PoetryInMotion]

Thanks for your answer, but still I think I'm a bit confused 

If we take a look at the formation equilibrium of ZnEDTA^2-,
[tex]\mathrm{Zn^{2+} + H_2EDTA^{2-} \rightleftharpoons\: ZnEDTA^{2-} + 2H^+}\,,[/tex]
I can't help wondering why the low concentration of H₂EDTA^2- does shift the equilibrium to the left, and thus gives a rather low value on the formation constant of ZnEDTA^2-? If that's the case, it contrasts with the textbook authors, who claim that zink ions and EDTA reacts "quantitatively" at pH 5.

Please elaborate - that's the first time I hear about EDTA^4- being unstable. At pH high enough it dominates EDTA solutions. At pH around 4-5 its concentration is very low, but it has nothing to do with the stability (could be it is a matter of terminology).

Sorry for my confusing terminology. What I meant was that the molar fraction of EDTA⁴⁻ will be very low at pH 5. Stability is, just like you say, something completely different.

In both cases the product is MeEDTA^2-, so the reaction equation is just a matter of finding the dominating H_nEDTA^-4+n ion.

Hmm. Don't you think it is better to stick with H₂EDTA^2- on the left side, since the titrant consists of Na₂EDTA solution?

[Borek]

giving a rather low value on the formation constant of ZnEDTA^2-.

10^16.5 is not that low.

Hmm. Don't you think it is better to stick with H₂EDTA^2- on the left side, since the titrant consists of Na₂EDTA solution?

Not if you are buffering the solution at some other pH.

[PoetryInMotion]

When I said "formation constant" I meant the so-called "stability constant", β', which should be close to 1, if the reaction should be regarded as quantitative. According to the text book β' depends on the molar fraction of both the EDTA^4- ion and the free metal ion.

What I meant was that a low molar fraction of EDTA^4- shifts the equilibrium Zn + EDTA^4-  ZnEDTA^2- to the left, and gives a lower β'. Something similar to this is discussed at the first page of http://www.udel.edu/chem/beebe/Chem120/Chem120%20LAB2EDTATitration.pdf.

Again, sorry for messing up the terminology

[Borek]

Please write definitions of formation constant and stability constant.

Sounds to me like your β' is so called conditional stability constant. But if so, β' close to 1 doesn't sound good to me - I would expect you need much higher values for the ion to be quantitatively complexed. Perhaps it is a matter of definition again.

Quit and dirty calculations show that formation constant for Zn complex is high enough to ensure high conditional constant in slightly acidic solutions. But it was a rather rough evaluation.

[PoetryInMotion]

Have revised some parts of the EDTA chapter and you are right, the full definitions are are as follows (using Zn²⁺ as an example), meaning that the formation constant should be much more than 1 in order for the reaction to considered quantitivave:

Formation constant: [tex]\normalsize\mathrm{K_f=\frac{[ZnY^{2-}]}{[Zn^{2+}][Y^{4-}]}}[/tex]
Effective formation constant: [tex]\normalsize\mathrm{K_f^{\prime\prime}=\alpha_{Zn^{2+}} \alpha_{Y^{4-}}K_f}\,,[/tex] where alpha is the molar rations.

In this case the molar ratio of free zink ion is high (there are essentially no zink hydroxide formed at this pH) and the molar ratio of EDTA^4- is about 10^-7. That means that the effetive formation constant is approximately 10^9.5. I guess that's enough to consider the reaction quantitative.

Surprisingly enough I also found an handout from some weeks ago where reaction formulas on the form
Me + H₂EDTA^2- MeH₂EDTA
are occasionally used at low pH.

Is that even possible? It would be interesting to compare the effective formation constant forZnH₂EDTA with that for ZnEDTA^2-, but I haven't found any relevant data on protonated EDTA complexes, except from this document http://www.publish.csiro.au/?act=view_file&file_id=EN12103_AC.pdf. What do you think?

[Borek]

http://www.chembuddy.com/?left=pH-calculation&right=pH-polyprotic-acid-base

See equations 9.11-9.13 for formulas to calculate speciation of the multiprotic acid (equations are for diprotic, but they are very easy to generalize). Just plug [H⁺], total concentration of EDTA and known Ka values (note: these are overall dissociation constants, not stepwise ones) and you will known concentration of EDTA^4-.

[PoetryInMotion]

http://www.chembuddy.com/?left=pH-calculation&right=pH-polyprotic-acid-base

See equations 9.11-9.13 for formulas to calculate speciation of the multiprotic acid (equations are for diprotic, but they are very easy to generalize). Just plug [H⁺], total concentration of EDTA and known Ka values (note: these are overall dissociation constants, not stepwise ones) and you will known concentration of EDTA^4-.

Thank you. I have now found out that α(EDTA^4-)=2.9*10^-7 and that α(H₂EDTA^2-)=0.93 at a pH of 5. To determine the two different K_f'' values, I also need the K_f for ZnEDTA^2- and ZnH₂EDTA respectively.

The document I linked to before gives some values on these constants, but according to them Kf would be higher for ZnH₂EDTA than for ZnEDTA^2-. To me, that feels rather unlikely, but then, on the other hand, I'm a total beginner on analytical chemistry :-/

[PoetryInMotion]

Oh, this year's preparatory problems for IChO seem to deal with something very similar to exercise b above. Guess it could be interesting for this thread. The problem is presented on page 57 at http://icho2013.chem.msu.ru/materials/Preparatory_problems_IChO_2013.pdf and the solution on page 35 at http://icho2013.chem.msu.ru/materials/Solutions_15_to_35.pdf

What happens is basically that brass is digested in nitric acid and then titrated with EDTA at a pH of 5.5-6.0. In the worked solutions presented above they actually use ZnH₂EDTA in the reaction formula and motivates it with effective stability constants similar to those mentioned above. What do you guys here think? Can this possibly be correct or is the usage of ZnEDTA^2- more reasonable? Perhaps both complexes have a high enough K'' value to co-exist at pH≈6?