# Why doesn't EDTA complex with alkali metal ions?

EDTA complexes with all the other metal ions in the periodic table except those from the group 1. Why is this so ? What is the coordination chemistry behind this ?

• That's funny, the label on my bottle says Ethylenediaminetetraacetic Acid, Disodium Salt.
– LDC3
May 28 '15 at 4:43
• @LDC3: That is $\ce{Na2EDTA.H2O}$. In that compound, $\ce{Na+}$ ions are not complexed to EDTA. May 28 '15 at 7:12
• @chemkatku Can you explain how you differentiate between "complexed?" I guess your argument against Na is that it's ionic and not a coordination bond? May 28 '15 at 13:31
• It does. I’m willing to bet my life it does. Only that the energy gained from complexation is so little and the process so easily reversible that there is no inert/favoured complex form that could be observed easily.
– Jan
May 28 '15 at 13:57
• @GeoffHutchison Yes that is what I think. I don't know about any way of complex formation without formation of coordination bonds. Jun 1 '15 at 7:32

If you're going to be picky that a "complex" doesn't form without covalent coordination bonds, then that's your answer. Alkali metals will tend to form cationic salts with EDTA.

Personally, I agree with the other comments. EDTA does form complexes with alkali metals.

I'd classify the disodium salt of EDTA as a complex. It has a metal and an organic ligand, and there's no direct metal-carbon bond (which would make it organometallic). As discussed in another question, it's a pretty arbitrary distinction between an ionic bond and a covalent one:

What happens if the electronegativity difference is exactly 2.1?

This distinction is particularly cloudy in the case of metal complexes, since most metals readily gain positive charge.

• I think crown ethers and cryptands provide unquestionable examples of alkali metal ion coordination? If those do, then that establishes a scale of coordination strength which never truly hits zero. Jun 2 '15 at 1:27
• @NicolauSakerNeto This is my opinion as well. My feeling is that there's a metal, a ligand, and they're close enough for some measure of covalent bonding. All bets are off if the ions are far apart, but in EDTA, it's at least partially covalent. Jun 2 '15 at 2:31
• @NicolauSakerNeto After looking closely at some $\ce{Cu(II)}$ aquo complexes, I have to agree with you, Geoff. Copper indisputably forms complexes, but its interactions with water ligands appear almost purely ionic. In my computation of $\ce{[Cu(H2O)5]^{2+}}$, the line critical point joining the axial $\ce O$ with the $\ce{Cu}$ has $\rho=0.05$ and $\nabla^2\rho=+0.26$, and all of the equatorial LCPs have $\rho\approx 0.07-0.08$ and $\nabla^2\rho \approx +0.4$. Thus, ionic EDTA-$\ce{Na+}$ interactions could also be called coordination. May 6 '16 at 2:03

Alkali metal cations are very "hard", in the HSAB theory of acid-base interactions (roughly corresponding to 'poorly polarizable'), and thus only coordinate in the conventional sense of the term to "hard" ligands such as ethers. Crown ethers, in particular, are known for their alkali metal ligating/chelating capability.

The amine nitrogens and carboxyl oxygens of EDTA are sufficiently 'soft' that, as Jan noted in a comment to the question, there is sufficiently little energetic benefit derived from EDTA chelation of a 'very hard' alkali cation such that the complexes are not practically observed. Instead, the chemistry is dominated by electrostatic interactions of the positively-charged alkali cations with the negatively-charged carboxylate groups (when deprotonated, anyways).

$\ce{Na2EDTA.H2O}$:
First of all the EDTA sodium salt is an ionic compound with complexed water. A sodium ion forms an ionic salt with the carboxylate group in two of the four acetic acid molecules bound to the amine groups. The salt has one water of crystallisation.

Solution Chemistry:
I'd say the high ionic radius and hardness of group I metal ions against the soft carboxylate make chelation less favourable with respect to softer metal ions, such as group III metals and transition metals.

Affinity constants:
\begin{array}{lr}\hline \ce{Al (III)} & 16.13\\ \ce{Ca (II)} & 10.96\\ \ce{Mg (II)} & 8.69\\ \ce{Pb (II)} & 18.04\\\hline \ce{Li+} & 2.79\\ \ce{Na+} & 1.66\\ \ce{Cs+} & 0.15\\\hline \ce{Fe (II)} & 14.33\\ \ce{Fe (III)} & 25.10\\ \ce{Mn (II)} & 14.04\\ \ce{Mn (III)} & 24.80\\\hline \end{array}