George W. Lotimer, Jr. University of Utah Salt Lake City
The Non-Aqueous Titration of the Salts of Weak Acids
A
method for the HCIOl titration of the salts of earboxylic acids in glacial acetic acid was published by Markunas and Riddick in 1951.' This procedure not only effectively complements the lecture discussion of non-aqueous solvents and their general acidbase relationships but also is readily adaptable to the undergraduate laboratory with quite acceptable results. The students are given commercially available C.P. chemicals. The laboratory assistant prepares a 1 M HC104-in-glacial acetic acid solution and dispenses it in small quantities (25 ml). (Experience has shown that solutions made simply by scaling up the directions from the article vary widely in strength; it is, therefore, better to standardize the concentrated solution (- 1.0 M solution) approximately and then adjust the strength accordingly.) The student adds this 1N HCIOI solution to 225 ml of glacial acetic acid containing 5 g of acetic anhydride, allows the solution to stand overnight, and then standardizes it using potassium acid phthalate with crystal violet indicator. The results of the experiment are shown in Table 1. One salt, magnesium acetate, of those initially dispensed, gave a very gradual visual end point change. For this reason the samples of magnesium acetate were withdrawn and the results were not shown in the table. To obtain a sharp end point, dry glassware must be used. Discussion
I n the present introductory course, the student is required to perform two acid-base titrations. I n the first, a weak acid is titrated with aqueous NaOH and the meq/g H + reported; the base is standardized against votassium acid ~hthalate. I n the second, one or the salts listed in able 1 is titrated in glacial acetic acid with HC104 and the meq/g OAc- reported; the HCIOl acid is also standardized against potassium acid phthalate. The use of potassium acid phthalate to standardize both the acid and the base appears invariably to pique I MARKUNAS, P. C., AND RIDDICK, J. A,, And. Chem., 23 (2), 337 (19.51).
the curiosity of the student and the questions in the laboratory session permit an explanation of the principle at the bench. The experiments are, of course, discussed carefully in lecture, and the behavior of the solvent in acid-base relationships is clearly defined and contrasted. Considerable emphasis is laid upon the precise fornlulation of the equilibrium constant which includes the solvent as one of the factors. The idea of an inexperienced group using HCIO., particularly in an organic solvent, may be disquieting to some instructors. However, the initial preparation of the strong reagent by a graduate student, or by the course instructor himself, minimizes many problems. At no time does the titrant need to he heated. If the students are counseled on the use of HC104 and its properties, the exercise of normal laboratory precautions appears to be all that is necessary. Although a potentiometric end point may be substituted for the visual, we have limited our experiments to those in which a crystal violet indicator proves satisfactory since a discussion of electrodes, potentials, and potentiometric titrations is given later in the lecture sequence. Table 1.
Salt
Sodium Acetate Barium Acetate Sodium Formate Potassium
Data for the Titrotion of the Salts of W e a k Acids in Glacial Acetic Acid
% Purity
Meq per gram Average
Range
n
s
ReFound portedb
12.14 11.70-12.91 16 0.31 99.6 7.75
99.0
7.23- 7.93 16 0.18 99.0
99.2
14.57 14.27-15.04 13 0.12 99.1
99.4
Aoirl
Ktlhitlate 4.89 4.81-5.11 12 0.09 99.9 Sodium Salicvlste 6.24 6.07-6.34 14 0.08 99.9 Ammoiium Acetate 12.61 11.93-12.92 6 0.40 97.2
.. . 99.3 98.3
'The averages reflect the values obtained over two quarters. b Markunas and Riddiek n is the number of students doing the determination (3 determinmtions per student). s is the standard deviation calculated from the range of averages reported.
Volume 43, Number 4, April 1966
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