Measuring pH in 95% DMSO solutions?

EK khatipovNO-SPAM at
Thu Nov 28 18:26:19 EST 2002

DMSO does not dissociate to form H+ or OH-, thus it does not contribute to
pH (after all, a measure of concentration of H+)
I would suggest just making a 10% CH2ClCOOH in water (4.5 parts H2O, 0.5
parts  CH2ClCOOH ), adjust pH to 5 and mix it with 95 parts of DMSO. Same
concept applies to, e.g., preparation of Tris-buffered phenol (even though
phenol is not miscible with water): you always assume that the inorganic
solvent does not contribute to pH value. Thus, mixing say 100 parts of
phenol with 1 part of Tris buffer pH 8.0 is considered as preparation of
phenol pH 8.0. Same for e.g. 80% acetone in phosphate buffer pH 7.0, etc.,

"David F. Spencer" <DSpencer at Dal.CA> wrote in message
news:DSpencer-6C9D4C.13483528112002 at News.Dal.CA...
> In article <20021030163433.3406.qmail at>, jsp at
> wrote:
> > Hi there,
> >
> >
> >
> > I'm trying to make a 95% DMSO, 4.5% H2O, 0.5% CH2ClCOOH solution at pH
> > by adjusting pH with HCl.
> >
> >
> >
> > Im using a standard pH-meter (Hamilton minitrode), but I can't seem to
> > get stable readings due to the low water activity. If I clean the
> > electrode and re-introduce it into the DMSO solution I get a completly
> > different reading. Does anyone have a good suggestion how to accurately
> > measure the pH?
> >
> There are many layers to the problem presented here.
> First, chloroacetic acid is a fairly strong acid (pKa 2.85) so the fact
> that you have been adding HCl to lower the "pH" to 5 should be the first
> sign that this is not a simple problem. The normal strategy here would
> be to make up a 20X stock of chloroacetic acid (because you have room
> for only 5ml per 100ml) with the pH adjusted to 5 (with say NaOH) and
> add that to the DMSO. However, that would not give a pH of 5 when added
> to the DMSO, assuming that you could measure the pH of this solution
> with a conventional pH meter, and also assuming that the whole concept
> of pH is applicable to what is essentially a non aqueous solution. Note
> that even in aqueous solution adjusting chloroacetic acid to a pH of 5
> would be tricky because that is more than 2 pH units above the pKa so
> less than 1% of the chloroacetic would be free acid.
> The pK of acids (and bases) is a function of several factors the simple
> Henderson-Hasselbach equation does not account for, such as ionic
> strength and the dielectric constant of the solution. The dielectric
> constant for DMSO (45) is considerably less than that of water (78.5 at
> 25C) and that factor shifts the observed pKa of a carboxylic acid up
> (i.e., towards 7). There are equations to calculate this more properly
> but they are not trivial to work with. The ionic strength factor is
> another can of worms.
> You would predict that in 95% DMSO choroacetic acid would be poorly
> ionized and so trying to adjust the pH down to 5 with acid would not be
> useful. The analogy would be trying to adjust the pH of say a NaCl
> solution to 5 or 7, whatever. You can only attain a stable pH if one of
> the solutes in a solution has a pK in the general vicinity of the
> desired end point. With the chloroacetic acid in DMSO example, if the
> observed pKa for the acid has risen to say 6 to 7, there is little
> buffering capacity at pH 5 (if that even means anything in this
> situation). If the chloroacetic acid is essentially unionized in DMSO
> the problem is even more fundamental and so buffering is not possible.
> You did not explain what you were trying to do or why but I assume this
> can't be an established protocol. You might get better feedback here if
> you elaborated more on the specific goals. I assume that you know that
> it is advisable (before dispensing) to purge the DMSO stock bottle (pure
> liquid) with a gas such as nitrogen to clear out the dimethyl sulphide,
> a breakdown product of DMSO.
> Dave
> --
> David F. Spencer, PhD
> Dept. of Biochemistry and Molecular Biology
> Dalhousie University
> Halifax, Nova Scotia, Canada
> DSpencer at Dal.CA

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