Subject: Re: Titration of Nitric Acid with Sodium Hydroxide
From: dwilkins@parkrapids.cfa.org (Don Wilkins) 
Date: Aug 09 1995
Newsgroups: sci.chem

>Some friends at the Army Research Laboratory are having problems using
>an auto titrator to titrate this strong acid with caustic for as you
>would guess, the inflection point is poor and so a derivative curve
>can't locate the end point without introducing too much error from one
>lab to the next. Does anyone have an idea as to whether or not an
>additive to alter the ionization of the acid would get around this.
>Much like changing the dielectric constant of a solution by adding an
>organic to control or force precipitation of ionic salts as we
>sometimes employ in the lab.

I have puzzled over this and have a suggestion which should have
occured to me at the onset. Apparently alzheimer's is setting in.

I now wonder if their problem is perhaps caused by the preparation of
the standard NaOH solution. NaOH is notorious for adsorbing carbon
dioxide to form sodium carbonate. An old bottle of NaOH could easily
have over 5% Na2CO3. This will indeed screw up the endpoint because
one is titrating not with a strong base but with a mixture of a strong
base (NaOH) & a weak base (Na2CO3). You can check this quickly by
doing a titration the old fashioned way using phenolphthalein as an
indicator. Titrate to the endpoint and then boil the solution. If it
then requires more NaOH you have a carbonate problem. Let's go back to
some freshman chemistry.

Prepare about 100 cc of a 50% solution of NaOH in water in a
polyethylene bottle. Set this aside for about two weeks (capped). You
may notice a silky white precipitate collecting on the bottom. This is
Na2CO3 which is insoluble in a 50% solution. Do NOT  shake, rattle, or
roll this bottle because the precipitate must settle and you will
prolong the process. After the precipitate has settled PIPETTE about
10.5 cc of this concentrated carbonate free NaOH into a 1 liter
volumetric flask and dilute to volume with water which has been
previously boiled (to remove the CO2) and cooled to RT (for a .1 M
solution). Some auto titraters suggest stronger solutions so adjust
accordingly.

The reason the problem is worse for KOH incidentally is that the same
problem exists with KOH but K2CO3 is quite soluble in conc. KOH so the
above procedure does not work to clean up KOH.. Carbonate in the
titrant not only screws up the endpoint but you will also get the
wrong answer. I will leave the reason as an exercise for the student.
If you can't figure it out e-mail me and I will give the answer but
check out some 75 year old text books first.

The freshmen  would try to sneak into the stock room to get an
unopened bottle of NaOH which of course had far less carbonate. The
chem professor had taught this course before and all of the unopened
bottles had been removed from the stock room.


--
dwilkins@parkrapids.polaristel.net

From: REMOVE_THISdwilkins@means.net (Don Wilkins)
Newsgroups: sci.chem
Subject: Re: Purification of KOH Solution: CO3 Removal
Date: Tue, 07 Sep 1999 02:57:08 GMT

On Mon, 06 Sep 1999 00:58:21 GMT, Uncle Al <UncleAl0@hate.spam.net>
wrote:

>
>
>Ken Westra wrote:
>>
>> I am looking for ways to remove the CO3 ion from a 1 litre solution of
>> 45 % by weight. I need to drop the CO3 concentration from 5 to 15 % to
>> below 1%. I have been looking at precipitation reactions, but I'm not
>> keen on the metal ions that would left over from this sort of reaction.
>
>What is the solubility of K2CO3 in 45% KOH, and vs. temperature?  The
>incredible [OH-] will send commonplace dilute solution ppt chemistry
>out the window.

Too high at room temperature. This works to prepare NaOH but does not
work for KOH.

The old master recommends using lime.

I.M. Kolthoff, Z. anal.Chem. 61,48, (1922)
>
>The only reasonable route might be using fired KOH in degassed
>distilled water and keeping it away from air.

I wouldn't bet much on this one. K2CO3 melts over 800.

Potassium metal dissolved in water (very carefully). Might better be
done by passing moist air (through a KOH or NaOH absorbent) over the
metal in a drybox.
>
>What about silicate?  You are dissolving Pyrex as we sit here and
>chat.

If you are making extremely high purity solutions teflon bottles are
better and even then don't stare at the solution too long. ;-)