I
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Wendell Forst Laval Univelsity Quebec, Canada'
~010rimetricDetermination of the Dissociation constant of Acetic Acid
Two years ago r e found ourselves with the Bausch and Lomb "Spectronic 20" colorimeter in the undergraduate physical chemistry laboratory and nothing more exciting to do with it than to determine the absorption spectrum of some colored compound. Interesting experiments have been described2 but they require an expensive spectrophotometer. A careful search of older literature revealed that a colorimetric investigation by Sidgwick, et al.,' dealing with the determination of the dissociation constant of acetic acid could serve as an interesting example of colorimetric technique, having the added virtue that it could be easily adapted to the requirements of a three-hour laboratory period with the colorimeter mentioned above. Following is a description of the modified experiment as used successfully for the past two years.
A.1 = sbsorhancy of a yellow completely dissociated solution of methyl orange; A.2 = absorhancy of a red undissociated solution: A.3 = ahsorbancv of any other solution, containing both dissociate& and undissociated forms
If the hydrogen-ion concentration of the solution is [H,O+] gram-ion/liter, a fraction y of the indicator will be undissociated, and the absorbancy of such a solution will he given by the simple additive expression
+
A.3 = Y A , ~ ( 1 - y)A,1
(1)
If Kin is the apparent dissociation constant of the indicator acid, then =
[H8Ot1(l
- V)
(2)
V
Eliminating y between (1) and (2)
Principle of Method
The "color," or absorbancy, of a methyl orange solution is used as a measure of hydrogen ion wncentration, from which the dissociation constant of acetic acid may be calculated. Absorbancy measurements are made a t a wavelength greater than the isosbestic point of methyl orange (470 mp), so that the yellow form of methyl orange absorbs less than the red form;' hence the absorbancy of the yellow form may be arbitrarily set equal to The experiment may be conveniently considered in two parts. Part I. First it is necessary to determine the indicator constant Kin. As is well known, a t pH above 4.4, methyl orange is yellow, and changes to red when the pH drops below 3.1. I n a solution where 3.1< pH< 4.4, there will be an equilibrium H.0 + H I n HsO+ + I w acid color (red)
orange solution may be conveniently regarded as proportional to ionization. Let
=
base color (yellow)
Hence the concentration of either the red or yellow forms, i.e., the "color," or absorbancy, of a methyl Now on leave of absence st the D e ~ a r t m e n tof Chemistrv. Carolina. Chanrl ill. ~~niversit," - ~ - nf -~ - North -~ " % ( a )PHILLIPS,J. P., J. CHEM.EDUC.,31, 81 (1954). (b) ISKSHALOM,M., FITZPATRICK, J. D., A N D ORCHIN,M., J. CHEM. S. W., J. CHEM.EDUC.,35,514 EDUC.,34,496 (1957). (e) TOBEY, (1958). "SIQWICK, N. V., WORBOYS, W. J., AND WOODWARD, L. A., P ~ o e Roy. . 8oc. (London), 129.4, 537 (1930). SIDGWICK,N. V., AND WOODWARD, L. A,, Proc. Roy. Soc. (London), 130A, l(1931). FORTUNE, W. B., AND MELLON,M. G., J. Am. Chem. Soc., 60, 2607 (1938). q t might be possible to use the Bauseh and Lomb instrument in another experiment to obtain the spectral trmsmission curves far methyl orange a t different pH values by a suitable modification of a. Iaboratoly experiment described in a very recent paper by Tobey' dealing with methyl red.
.
~
~r~~
~
~
This equation may be used for the experimental determination of K;.-- hv " measurine absorhancv a t known values of [H,O +I.
-
Procedure: A stook solution N/8000 in methyl orange is prepared. Fifty milliliters of this solution is pipetted into each of three 250-ml volumetric flasks marl
