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The article "Determination of the Dissociation
Constant for a Monoprotic Acid by Simple pH
Measurements" (J. Chem. Educ. 1996,
74, 792) raises questions that teachers may exploit usefully in the classroom.
1. Because of the wide range of pH values in the
experiments, one presumes that the pH meter was
standardized using NBS buffers or the equivalent. What, then is
the meaning of the initial pH (before the acid is added)? If it
is taken to yield an approximation to the activity of
hydrogen ion, then why do the authors ignore the ionic strength
(and activity coefficient) in deducing their values of
Bo?
2. The authors conclude that "it is only necessary
to know the approximate molecular weight [of the acid] as
well as the exact (emphasis added) concentration of the
alkaline solution." and "It is not necessaryto standardize the
solutions."
Are they not using the pH meter as a (poor) way
to standardize their initial NaOH? Why rely on a meter
(an error of 0.05 pH unit corresponds to an error of about
12% in concentration) when it would be so easy to make up
the solutions very accurately with standard dilute NaOH
and NaCl?
3. What is the value in making a fairly complicated
plot to process the data? Since each data point is merely
a simple monoprotic buffer solution, rapid deduction of
K-values with a calculator is easy. In other words, their
equation K1 =
x/(Ao-x)/B is all one needs for each point, and is
not improved by conversion to a plottable linear form.
The answer is that the plot method uses the ratio
of intercept to slope to find K, and therefore has the
advantage that the value of Ao actually cancels, thereby
justifying their assertion that the molecular weight need not
be accurately known. However, isn't it true that for any
acid for which this approach is suitable it is easy to
determine the molecular weight by titration, and that this is
surely instructive for the students?
4. Further, doesn't such a plot, using reciprocals,
weight the points unequally? With individual point
calculations each point is weighted equally in deducing the average
pK.
Note in the figure that the highest point seems "out
of line". How would the intercept of this plot change when
that one point is omitted? (One gets 9.42 for
pKa instead of 9.36.) With individual calculations isn't it easier to note a
deviating point or two?
5. In short, doesn't it seem better to: use
standardized NaOH; correct for ionic strength; use an accurate value
for the molecular weight; and calculate each point
individually? Such an approach is simple, fast, and more reliable,
and doesn't conceal the deviation that any one point may have.
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