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> 	I think I would read the reference pretty carefully.  Those pKa's
> are for stronger acids than carboxylic.

Unfortunately the reference with this listing is just a chart with no 
detail info given, the references from which they got those numbers I do 
not have ready access to and will have to look them up...

>> maybe so weirdo gas phase kinetic values...
> 	pKa's are for aqueous solutions, and are _equilibrium_, not kinetic
> rate constants.

Yeah, I don't know why I wrote kinetic (guess it was still too early in the 
morning :-), I should have said thermodynamic... but more along the lines 
of "calculated" values based on computer based calculations which usually 
look at the stuff in the gas phase (or low water concentration) to make the 
modeling easier... But without looking up the original references this is 
just guessing, I don't know why the cited numbers were so low...

> 	There are two parts to deciding whether a complex is "stable" or not.
> One has to look at how easily it comes apart, and how readily the products
> do something else (so it doesn't reform).  The pKa's of the acids and the
> pKb's of the amines (protonation, not ionization) tell you a bit about
> how readily the EDTA will do something else, and it appears from the pKa's
> that two of the acid groups will be essentially competely ionized, and a 
> third
> will be at least partly so (at aquarium pH).  The amine groups will be
> mostly protonated.  I haven't been able to find the pKb's, but I would
> expect them to be about 8-9.  (Does anyone have the numbers?)
> 	Far stronger bases than this form stable compexes with transition
> metal ions, because the +N-H bond in the protonated amine is in competion
> with a +N-M bond that can form.  The point about a chelating agent, is 
> that once
> it starts to latch on to a metal ion, it is likely to latch on with more
> of its parts, because they can't get very far away.  Once it is completely
> latched on, getting off is difficult, because even when one bit comes off,
> it can't get away in the same way that a separate ion could, so is likely
> to reattach before another bit comes away.
> 	This means that a chelate complex is more stable than one would
> predict from just the bond energies involved.  A complex the flies into
> many pieces is more likely to fall apart and less likely to get together.
> In thermodynamic terms, a chelate is more stable because of the entropy
> term in the equations concerning stability.

So would you say that actually an EDTA complex should be relatively pH 
insensitive because the entropy advantage of the chelate insures that even 
a protonated amine will eventually equilibrate to yield a N->metal 

Something else to consider is that the the pKas listed are for the 
uncomplexed tetraacid; the presence of a suitable Lewis acid such as a 
metal should tend to lower the pKas of each carboxylic acid as the ionized 
state (carboxylate) becomes more stabilized due to the availability of a 
metal to chelate. So basically the metal acts as an ionization catalyst by 
lowering the energy of the ionized form (which is metal chelated) thus 
yielding an "apparent" pKa that is lower than that of the acid alone. These 
"apparent" pKas are also going to depend on the identity of the metal being 
chelated as well since if the EDTA doesn't "like" to form a complex as well 
with Metal A vs Metal B then the apparent pKas are going to be higher with 
Metal A. So I guess the take home message here is that EDTA chelating 
ability is not as dependent on pH as it might appear based solely on the 
pKas of the free acid.

-Greg Morin
Gregory Morin, Ph.D.  ~Research Director~~~~~~~~~~~~~~~~~~~~
Seachem Laboratories, Inc.      www.seachem.com     888-SEACHEM