Re: WHAT is the pH of a weak-weak salt?
From: Mohammed Farooq (farooq_w_at_hotmail.com)
Date: 06/02/04
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Date: 1 Jun 2004 20:08:17 -0700
Dr. F wrote
>Thank you for the answers, Mohammed. I still have a few questions.
You
>mentioned the formula for the pH of a weak-weak salt:
>> pH = 7 - 1/2 log Ka + 1/2 log Kb
>Would that formula be equivalent to saying pH = 7 + pKa/2 - pKb/2 ?
If
Infact this is the standard way of writing it. Unfortunately you would
not find this realtionship mentioned in many modern books and I
noticed that they delibrately avoid this case you would easliy find
NaCN examples but not of NH4CN, and perhaps your teacher did not
consult "the books of the yore" to answer your question when she said
" I don't know". I found this relation mentioned and derived in a very
old but a very classic book by Kolthoff's" Volumetric Analysis"
perhaps translated from the German, of around 1960's, which was a part
of my intorductory course. He changed the way analytical chemistry was
taught in the world. He was the father of analytical chemistry in USA.
>so, that is interesting because I know that the pH of a 1M weak acid
>is pKa/2 . Also, if the salt were, for example, NH4CN, would Ka in
the
>formula be the Ka of HCN (the acid) or of NH4(+) (the conjugate
acid)?
We are not concerned with conjugate acid-base pair rather with the
acid and the base from which the salt was made ie HCN and NH3 (rather
hypothethical NH4OH).
>I am guessing that Ka would be the Ka of HCN, because that assumption
>gives a basic answer (which I would expect), but I am not sure.
Correct.
>Also, how could pH not depend on concentration? What if the
>concentration was only 0.00001M, how could that tiny amount of salt
>change the pH from 7?
The concentration factor cancels when this relation is derived. There
are many assumptions in this derivation, you know, almost all
equations are usually derived on certain assumptions and
approximations.
>Is this nondependence on concentration somehow
>related to the fact that the pH of a buffer solution is also
>independent of concentration?
Salt a of weak-acid-weak-base is not a buffer. pH of a buffer is given
by Handerson-Hasselbach relation which does involve concentration
terms.
>> Transition metals do form covalent bonds and usually their halides
>> hydrolyze in water generally giving [M(H2O)6](n+) species eg in
>> [Cu(H2O)6] (2+), water molecules are covalently bonded.
>Alright, I know that solvated metal ions can form complexes. I don't
>see how the complexes you mentioned there would make the solution
>acidic. They include whole water molecules, so I don't see how they
>could create H+ ions in solution. How do CuCl2 and ZnCl2 create H+
>ions and make solutions acidic?
Zn salts also form [Zn(H2O)5 OH](2+) in water , thus increaing H+
ions, so does iron and aluminum, solutions of zinc, iron aluminum,
copper are normally acidic.
>> As an excercise can you find out the pH of 1 x 10^-7 M HCl?
>> (Should it be less than 7 or greater?)
>Is it pH = 6.7? That would be the pH if [H+] = 2*10^-7 M.
This was deceptively simple. This is a common misconception that water
_always_ produces 1 x 10^-7 M H+ and 1 x 10^-7 M OH-. This is true
only in pure water when there is NO acid or base added. Hint: Kw
=[OH][H], where [OH] is x, and [H] = (1 x 10^-7 + x). Solve the
quadratic equation, and find out the [H] and finally pH. Kw however
does not change (why? recall for example : 3 x 2 = 6 x 1 = (1+2) x 2).
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