Calibration of a Practical Acidity Scale for Sulfolane as Solvent J. F. Coetzee' and R. J. Bertozzi Department of Chemistry, Uniuersity of Pittsburgh, Pittsburgh, Pa. 15213
SULFOLANE (TETRAMETHYLENESULFONE) possesses major advantages as a solvent for the determination of very weak acids and bases, as well as for the resolution of mixtures of relatively strong acids or bases that are leveled in more reactive solvents. These advantages have been illustrated well by the results of Morman and Harlow ( I ) . We have been interested in calibrating an acidity scale for sufolane, and have given a preliminary report on the preparation and analytical potentialities of anhydrous perchloric acid solutions in this solvent (2). We also pointed out that our preliminary attempts to calibrate an acidity scale had been unsuccessful because the concentration dependence of the potential of both the hydrogen and the glass electrode was greater than Nernstian. We have established subsequently that the problem was caused by a combination of factors: incomplete dissociation of anhydrous perchloric acid, and the presence of residual water in the purified solvent coupled with further uptake of water during measurements, which is extremely difficult to avoid. In a recent publication Benoit et al. (3) reported the results of a thorough study of acid strengths in sulfolane by means of proton magnetic resonance spectrometry and conductimetry, from which they concluded that perchloric acid has a pK value of 2.7. Our value, derived in a somewhat different manner, agrees reasonably well with this number. We now also report the results of hydrogen electrode measurements in perchloric acid solutions and of both hydrogen and glass electrode measurements in several buffer solutions, which allow calibration of an approximate practical acidity scale for sulfolane as solvent. EXPERIMENTAL
Apparatus. Conductance measurements were carried out in Fisher Low Conductivity type cells with lightly platinized electrodes, using an Industrial Instruments Inc. Model RC-12B conductance bridge calibrated with precision resistors. Potential measurements were carried out with a Corning Model 12 Research pH meter. Hillebrand type hydrogen electrodes were cleaned and lightly platinized periodically according to the directions of Ives and Janz (4). A single Corning Triple Purpose glass electrode (No. 476022) was used throughout this study. The reference electrode (designated as AgRE) was an Agl(0.100M AgC104 in sulfolane) electrode with an asbestos fiber tip. Several such electrodes were constructed and the potential of the working reference electrode was checked against a control both before and after each experiment. Reagents. Commercial sulfolane (Shell or Phillips) was purified as described elsewhere (5). Predominantly anPlease address all correspondence to this author. (1) D. H. Morman and G. A. Harlow, ANAL.CHEM., 39, 1869 (1967). (2) J. F. Coetzee and R. J. Bertozzi, ibid.. 41,860 (1969). (3) R. L. Benoit, C. Buisson, and G. Choux, Can. J . Chem., 48, 2353 (1970),
(4) D. J. G. Ives and G. J. Janz, "Reference Electrodes," Academic Press, New York, N. Y.,1961, pp 106-107. ( 5 ) J. F. Coetzee, J. M. Simon, and R. J. Bertozzi, ANAL.CHEM., 41, 776 (1969).
hydrous perchloric acid was prepared using the same technique as before (2). CAUTION. During the past 3 years we have prepared and stored for periods of months numerous batches of perchloric acid at concentrations up to 0.3M without experiencing any explosions. However, Benoit et a[. (3) recently reported explosions at concentrations of 1 M and higher. We urge other workers to handle even dilute solutions with all due care. For many (but not all) analytical purposes HSbC16,which is recommended by Benoit, should be preferable. Baker and Adamson picric acid was recrystallized twice from acetone and dried in vacuo at 45 "C. J. T. Baker Analyzed Reagent salicylic acid was used without further purification. K and K Laboratories 2,6-dihydroxybenzoic acid was treated with Fisher Norit A decolorizing carbon, recrystallized from water, and dried in cucuo at 65 "C. J. T. Baker Specialty Gas Co. Ultrapure hydrogen chloride (99.99 %) and Electronic Grade hydrogen (99.999 %) were used without further purification. G. F. Smith Chemical Co. silver perchlorate was dried in vacuo at 80 "C. Tetraethylammonium picrate as well as tetrabutylammonium salicylate and 2,6-dihydroxybenzoate were prepared by potentiometric titration of a methanolic solution of the acid with one of the tetraalkylammonium hydroxide (Eastman White Label). The solvent was removed under vacuum and the residue dissolved in a mixture of ethyl acetate and the minimum amount of ethanol. The solvent volume was reduced under vacuum while cooling with a dry ice-isopropanol slush, and finally the crystals were dried in a vacuum desiccator over phosphorus pentoxide at room temperature. Procedure. Precautions against uptake of water included drying of glassware other than volumetric equipment in vacuo at 100 "C, thorough rinsing of volumetric equipment with portions of the solutions involved or with pure sulfolane, and carrying out all titrations and other transfers in a glove bag filled with nitrogen. Hydrogen for hydrogen electrode measurements was bubbled at a rate of 1 to 2 bubbles per second through pure sulfolane before passing into the potentiometric cell. Both the hydrogen and the glass electrode typically required 30 to 40 minutes to reach a stable potential. The AgRE was exposed to hydrogen for a minimum length of time, and it was checked that it behaved properly by repeating some of the measurements in a 3-compartment cell with a tetraethylammonium perchlorate salt bridge solution in the center compartment. All potentials reported are reduction potentials cs. AgRE. All potentiometric and conductimetric data were obtained at a temperature of 30.0 =t 0.1 "6. RESULTS AND DISCUSSION
We made numerous attempts to calibrate the hydrogen electrode by titrating blanks of different batches of sulfolane with perchloric acid, but, as reported before (2), the slope of the titration curve invariably exceeded 60 mV per decade up to the vicinity of 0.1M acid added, which is much higher than the maximum concentration of basic impurities expected in the purified solvent. This apparent anomaly was traced to a combination of first and second order effects of water. There are three possible sources of water in an experiment carried out under our conditions: water originally present in the purified sulfolane (typically 1 to 3 x 10-3M), water taken up mainly from the glassware during the preparation of perchloric acid, and water similarly taken up during the ANALYTICAL CHEMISTRY, VOL. 43, NO. 7, JUNE 1971
961
1.501 0
1
4
I
8
I
12
I
I
I
I
I
I
16 20 24 28 32 36 Volume of W a t e r A d d e d , pl
I
40
I
44
I
46
Figure 1. Conductimetric titration of anhydrous perchloric acid in 28 ml of a solution containing 5.0 X 10-2Mtotal acid. Solvent: sulfolane. Titrant: water
7L 0
I
I
I
I
I
2
3
4
J
5
10 2 3 0 c l o ,
Figure 2. Concentration dependence of equivalent conductivity of hydronium perchlorate in sulfolaneas solvent
course of the titration. Contamination by water from the latter two sources turned out to be much more serious than originally expected on the basis of the precautions taken. Anhydrous perchloric acid in sulfolane is a most efficient water scavenger. In addition, the interference by low concentrations of water is further compounded by the fact that anhydrous perchloric acid is incompletely dissociated, while hydronium perchlorate is more extensively dissociated, with the result that the dissociation of perchloric acid is further repressed by the common ion effect of hydronium perchlorate. Consequently, we measured the potential of the hydrogen electrode in a series of individual solutions of perchloric acid, and then determined the composition of the solutions after the completion of the potentiometric experiments. Anhydrous perchloric acid could be determined quite accurately by conductimetric titration with water, and total acid by conventional alkalimetric titration after flooding with water. A typical conductimetric titration curve is shown in Figure 1. A variety of other weak bases were titrated in the same way; 962
ANALYTICAL CHEMISTRY, VOL. 43, NO. 7, JUNE 1971
the results will be presented elsewhere (6). Even for added bases of very large molecular size, such as triamylamine, the conductance always increased up to the equivalence point, and this was our first indication that anhydrous perchloric acid is incompletely dissociated and that the conductivity of dilute solutions formerly measured ( 2 ) was mostly due to hydronium perchlorate. Conductimetric titrations as in Figure 1 were carried out for a series of perchloric acid solutions in the concentration xange of 5 X 10-3 to 2 x lO-lM, and the plot given in Figure 2 was constructed from the equivalent conductivity values at the intersections of the extrapolated linear sections of the titration curves. These are the equivalent conductivity values that would apply if the slight reversion of hydronium perchlorate into perchloric acid, which is rekealed by the curvature of the plot in Figure 1, were repressed, as it would be in the presence of excess anhydrous perchloric acid at the beginning of the titration, Turning to Figure 2, the curvature in the plot at higher concentrations indicates that hydronium perchlorate is slightly associated (which was ignored in our calculations). Our results are not very different from those obtained by Benoit et al. (3) with 72 perchloric acid, which corresponds to an equimolar mixture of hydronium perchlorate and water. Figure 1 also shows that moderate excesses of water have little effect on the conductivity of hydronium perchlorate. However, this fact should not be interpreted as meaning that no further reaction occurs between hydronium ion and water. In fact, potentiometric titration of anhydrous perchloric acid with water, using both the hydrogen and the glass electrode, still shows a considerable change in potential well beyond the equivalence point. Typical results are the following: in the titration of 5 X 10-2M perchloric acid, the potential of the glass electrode at 10 moles of water added per mole of perchloric acid is 185 mV less positive than at 2 moles of water added and 285 mV less positive than at the equivalence point. Sample Calculation, The dissociation constant of anhydrous perchloric acid was evaluated as in the following sample calculation for the data represented by Figure 1. ( 6 ) J. F. Coetzee and R. J. Bertozzi, University of Pittsburgh, Pittsburgh, Pa., unpublished data, 1970.
Total concentration of acid = 50 X 10-sM. Concentration of water added at equivalence point = 42.5 X 10-3M. Therefore, concentration of water present initially as hydronium perchlorate = 7.5 X 10-3M. From Figure 2, for 7.5 X 10-3M hydronium perchlorate, A = 12.6, so that the conductivity K = 9.5 X Initial total K = 1.79 X Therefore, for anhydrous perso that A = chloric acid present initially, K = 8.4 X 1.98. Benoit et al. (3) showed that HSbCle is a typical strong electrolyte in sulfolane. Addition of water to relatively dilute solutions of this acid in sulfolane causes little change in conductance. Hence, the limiting equivalent conductivities of SH+ (where S represents sulfolane) and H30+are similar, and we estimate that AZ,(SHC104) h0(H3OC104) 15. Ignoring the electrophoretic effect, which must be small in the relatively viscous sulfolane, and also the relaxation effect, which cannot be unduly significant judging from the conductivity data presented by Benoit et al. (3), a “concentration constant,” Kc,is calculated as follows. If the degree of dissociation of anhydrous perchloric acid in presence of hydronium perchlorate is given by C L ~= &/Ao = 1.98/15 = 0.13, then
-
-
+
Kc = (CYC~CECIO~ CL,CH~OCIO~)/(~ - CY,)= 2.0 X IO-’
(1)
+
The ionic strength of the solution = [SH+] [H30+] = a c ~ H c i o 4 CH~OC~O =~ 5.5 x 10-3 7.5 x 10-3 = 1.3 x 10-2. From the reduced Debye-Huckel equation the activity coefficient is given by jf = 0.73 and the activity constant by K , = Kcfi2 = 1.1 X Our best value for several solutions containing from 4 to 12 x 10-3M water also was 1.1 x 10-3. Benoit et al. (3),who used perchloric acid prepared by distillation, estimated pK = 2.7, which is in satisfactory agreement with our value, considering the complications caused by water. Typical results obtained with the hydrogen electrode are Q determined after given in Table I. Values of C H ~ O Cwere completion of the potential measurements, and were invariably higher than typical initial values, indicating that further water was taken up under the conditions of our potentiometric experiments. Values of paH were calculated from the following equation with K , = 1.1 X
+
+
where
and
where C , represents the total (analytical) concentration of acid. Finally, values of Eo(H2) were obtained from the relationship
potential or errors in the activity coefficient values used. Consequently, our mean value of +75 mV should be regarded as provisional until additional information on all these factors becomes available.
Table I. Hydrogen Electrode Potentials in Perchloric Acid Solutions Analytical Concentration of Perchloric Acid, C , Property 0.200M 0.150M 0.lOOM 0.050M CH30ClO4, M 0,106 0,061 0,047 0.029 fi 0.40 0.49 0.54 0.62 2.57 2.12 2.96 PaH 2.68 - 17 - 73 -89 -117 E(HA mV EAHz),mV +84 +81 +74 +61 Table 11. Hydrogen and Glass Electrode Potentials in Buffer Solutions Buffer
CHA
2,6-Dihydroxy9.9 X lo-* benzoic acid + 9.9 X its tetrabutylam- 9.9 X monium salt 2-Hydroxybenzoic 1 .OO X acid its tetrabutylammonium salt 2,4,6-Trinitrophenol 1 .OO X + its tetraethyl- 1 .OO X ammonium salt 1 .OO X 10-1 1.02 x 10-4 1.02 x 10-2
+
CA1.00 x 10-2 1.00 x 10-2 1.00 x 10-2
- E(Hz), - H G ) mV mV 1135 463 1052 383 915 245
1.00 x 10-2
1340
653
1.oo x 10-2 1.00 x 10-2 1.00 x 10-2
... ... ... ... ...
353
LOO 1.00
x 10-3 x 10-3
293 230
358 244
Using solutions of HSbCls, instead of HC104, would be a more direct approach. However, we found that the response of the hydrogen electrode in such solutions was sluggish and otherwise suspect by analogy to the known poisoning effect of arsenic compounds on the hydrogen electrode. The potential of the glass electrode was less positive in both HSbC16 and HClOd solutions than would be expected on the basis of the foregoing. We have encountered the same problem with the glass electrode in highly acidic solutions in other organic solvents, and we feel that it may be associated with dehydration of the glass membrane. However, in less acidic buffer solutions, the response of the glass electrode parallels that of the hydrogen electrode fairly closely, as indicated in Table 11. For the buffers consisting of equimolar (10-2M) acid and salt, the following paH values are calculated from the value E,(HZ) = +75 mV us. AgRE: 2,6-dihydroxybenzoic acid, 18.8 ; 2-hydroxybenzoic acid (salicylic acid), 23.6; and 2,4,6-trinitrophenol (picric acid), 17.3. We recommend the picric acid buffer for calibration of the glass electrode. The pK, values of the three acids listed in Table 11, calculated from the paH values of the equimolar (10-2M) buffers, are given by pK,
=
paH (lO-*M)
- logfi
= paH (10-2M)
(7) 1. M. Kolthoff,M. K. Chantooni,Jr., and Chern. SOC.,90, 23 (1968).
+ 0.1
s. Bhowmik, J . A
ANALYTICAL CHEMISTRY, VOL. 43, NO. 7, JUNE 1971
(6)
~
963
~
~
.
can be illustrated in the present case by the following overall scheme: HA(so1vated)
+S
SH+(solvated)
+
A- (solvated or homoconjugated) Properties of S involved are its proton acceptor power, its relative ability to stabilize the species HA, SH+, and A-, its ability to sustain electrolytic separation of the species SH+ and A-, and its own concentration. In comparing the two solvents sulfolane and acetonitrile, the main factors to be considered are likely to be the difference in their proton acceptor power and in their ability to stabilize A- either directly by a hydrogen bonding or a dispersion interaction, or indirectly by promoting homoconjugation as a result of lack of stabilization of HA. The data for 2,ddihydroxybenzoic acid in Table I1 are consistent with the existence of a A- with a formation conhomoconjugate complex AH stant of 240, whereas in acetonitrile its value is 400 (7). However, the main factor to be considered undoubtedly is the difference in proton acceptor power of the two solvents. Acidity function studies gave p K s ~ +values of -12.9 for
-
a
sulfolane (8) and -9.5 for acetonitrile (9). This difference of 3.4 units accounts for only about one half of the observed difference in pK, values. Our experimental information is too limited to justify further speculation. We have described elsewhere (5) an extrathermodynamic procedure enabling us to refer potentials in sulfolane to the water scale. Our procedure leads to the conclusion that the potential of the AgRE (S) is f0.91 V us. NHE (W), and, therefore, that the potential of the NHE (S) is +0.99 V us. NHE (W). Although such attempts to refer potentials in one solvent to another as standard state are controversial, this number does show that very high proton activities can be attained in sulfolane. RECEIVED for review January 25, 1971. Accepted February 24, 1971. From the Ph.D. thesis of R.J.B., University of Pittsburgh, 1971. Work supported by the National Science Foundation under Grant Numbers GP-6478X and GP-16342. (8) S. K. Hall and E. A. Robinson, Can. J. Chem., 42, 1113 (1964). (9) N. C. Deno and M. J. Wisotsky, J . Amer. Chem. SOC.,85, 1735 (1963).
Method for Determination of pK Values of Certain Low Solubility Organic Amines Josip dajal and Harry B. Mark, Jr. Department of Chemistry, University of Cincinnati, Cincinnati, Ohio 45221
IN RECENT STUDIES on the effect of inductive and steric influences on the electrocatalysis of Ni(I1) reduction by certain organic aromatic amine compounds (1-3), it was found that an accurate knowledge of the ionization constants of the protonated amine group was necessary in the interpretation of the data (3). Many of the compounds of interest, such as various naphthalenediamines and alkyl substituted pyridines and phenylenediamines (3), were found to have a very low solubility in aqueous media; solubilities were to lO+M or less (3). With this low solubility of relatively weak base compounds (pK, values of the protonated forms were in the range of 3 to 7), the usual potentiometric titration and spectrophotometric methods (4-8) for such compounds are not really practically. To circumvent the low solubility On leave of absence from the “Rudjer Boskovic” Institute, Zagreb, Yugoslavia. 1
in these cases, it is common to employ mixed solvent systems. such as ethanol-water (9) or dioxane-water (IO) to increase the solubilities to a convenient level. However, ionization constants of protonated amine groups measured in mixed solvent systems are generally different than those measured in aqueous media [for example, differences of 0.14 to 0.43 pH units have been reported for 70z ethanol (8, 9, II)]. Also, the fact that in most cases it is desired to correlate the pK values with chemical properties measured in aqueous media, it is desirable to measure the pK values themselves in aqueous media. The electrode reaction mechanism of the Ni(I1) electrocatalysis can be illustrated by the following scheme (1-3, 12-14): Lad
IT
(1) H. C. MacDonald, Jr., and H. B. Mark, Jr., Div. of Anal.
Chem., 158th National Meeting, ACS, New York, N. Y., Sept. 1969. (21 . , E. Kirowa-Eisner and H. B. Mark, Jr., J. Electroanal. Chem.,
16, 379 (1968). (3) , , J. Caia. H. C. MacDonald, Jr.. E. Kirowa-Eisner, L. R. McCoy, and H.“B. Mark, Jr., unpublished results, 1969. (4) J. M. Vanderbelt, C. Henrich, and S . G. Vander Berg, ANAL. CHEM.,26,726 (1954). (5) T. V. Parke and W. W. Davis, ibid., p 642. (61 . , L. A. Flexser, L. P. Hammett, and A. Dingwall, J . Amer. Chem. SOC.,57, 2103’(1935). (7) E. Sawicki and F. E. Ray, J . Org. Chem., 19, 1686 (1954). (8) G. Schwarzenbach, A. Willi, and R. 0. Bach, Helv. Chim. Acta, 30, 1303 (1947). 964
ANALYTICAL CHEMISTRY, VOL. 43, NO. 7, JUNE 1971
(9) P. H. Grantham, E. K. Weisberger, and J. H. Weisburger, J . Org. Chem., 26, 1008 (1961). (10) A. V. Willi, Helv. Chim. Acta, 40, 2019 (1957). (11) J. M. Vanderbelt, C. H. Spurlock, M. Giffels, and M. W. Eash, Science, 121, 646 (1955). (12) L. R. McCoy and H. B. Mark, Jr., ANAL.CHEM.,37, 591 (1965). (13) H. B. Mark, Jr., and C. N. Reilley, ibid., 35, 195 (1963).
(14) L. R. McCoy, H. B. Mark, Jr., and L. Gierst, J . Phys. Chem,. 72,4637 (1968).
