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Although Wikipedia says that EDTA is a very soluble salt, preparing a solution (.5M) is a pain in the neck. It takes a lot of time and still doesn't dissolve. Although addition of a few drops of NaOH makes the process a bit easier (still takes a lot of time)

Now, it did not make sense to me of why NaOH made EDTA more soluble. The salt I'm using is a disodium salt of EDTA, so wouldn't Na+ turn out to be the common ion only to decrease the solubility? Then, it might be the effect of pH(?) But, being a salt of a weak acid and a strong base, theoretically shouldn't such a salt be more soluble in lower pH? The EDTA divalent ion thus formed would be a stron conjugate base and accept H+ increasing the solubility?

I know both my assumptions are wrong since adding NaOH actually helped. What might be the actual reason which is counterbalancing the above mentioned 2 effects (are they even valid in this case?)

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  • $\begingroup$ Adding also an acid would encounter the common ion effect since it is only the disodim salt and 2H+ are still out there? Or is it that the base pull out these 2 H+ and makes it more polar?? $\endgroup$
    – Polisetty
    Jun 7, 2016 at 19:09
  • $\begingroup$ sigh If you neutralise it further it gets more soluble and tetrasodium EDTA is much better soluble, what's your problem with that? $\endgroup$
    – Mithoron
    Jun 7, 2016 at 19:25
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    $\begingroup$ Then why pH 8? If it neutralises completely shouldn't it be pH 7. But in manuals, it clearly says that the end product should be pH 8 @Mithoron $\endgroup$
    – Polisetty
    Jun 7, 2016 at 19:38

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A disodium salt of EDTA also has two remaining carboxylic-acid groups, which can be deprotonated by addition of base, like $\ce{NaOH}$. By the way, higher pH is generally better especially if you plan to use it for chelation, i.e. cation-capture. The labs I've taught in the past use buffers of pH ~10 with EDTA for complexometric titration experiments.

Deprotonating an organic acid in general improves solubility in water because the electrostatic interactions of $\ce{-O^-}$ and the solvent (the hydrogens on water) are stronger than interactions of $\ce{-OH^0}$ and water.

A similar example would be 4-t-butylphenol, which is nearly impossible to dissolve in water without addition of a base, e.g. $\ce{NaOH}$. In experiments I had to reach pH of about 12 to strip the compound of the alcohol-proton, yielding a negative charge which interacts more strongly with water.

Nonetheless you can still make a decently concentrated (~0.1M) solution of EDTA in water by stirring alone, if you allow a stir-plate to do the work for ~30 minutes.

In short, it's not the addition of $\ce{Na+}$ that matters, but rather the addition of $\ce{OH-}$, to increase solubility.

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    $\begingroup$ thanks! I'm actually doing a western blot which has a buffer of pH of 7.4 So, why don't they directly sell the tetrasodium salt? Or why is the dis odium salt more feasible? $\endgroup$
    – Polisetty
    Jun 7, 2016 at 20:59
  • $\begingroup$ welcome! Pure (then fully deprotonated, e.g. with pH ~10) EDTA is ideal for chelation. That is, a non-sodium salt. The sodium/other salts have their uses. But in general it's the base that changes solubility, much more than the sodium. A tetrasodium EDTA salt would be more soluble, I imagine. Never used/seen it though. $\endgroup$
    – khaverim
    Jun 7, 2016 at 21:05
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Also be aware that the so called disodium salt is almost never sold as declared (it's not that easy to get the stoichiometry right on industrial skale), and you usually end up with a mixture of compounds that have 2/3 carboxylic acids protonated per molecule. That is the reason why EDTA is not a primary analytical standard, and needs to be standardized after dissolving it. Adding NaOH also increases the ionic strength of the soln. which shifts equilibrium towards dissolution. Common ion effect doesn't play a big role, as Ksp value is quite large.

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