421
Determination of Acidity in 80% DMSO-20% Water
Determination of Acidity in 80% Dimethyl Sulfoxide-20% Water’ Ernest H. Baughman and Maurice M. Kreevoy* University of Minnesota, Chemical Dynamics Laboratory, Minneapolis, Minnesota 55455 (Received June 2 7 , 7973) Pub/ication costs assisted by the Nationai Science Foundation
Dissociation constants have been measured by the “ladder” technique for a series of buffer acids and a series of acid-base indicators in 80% DMSO-20% water. The whole range, between dilute HC1 and dilute KOH, has*been spanned. The latter measurements permit the determination of P K H A for water in this solvent. It has a value of 18.38 which correlates well with values in other DMSO-water mixtures. The indicators provide a convenient method for reasonably accurate pH measurements in this solvent mixture in which reliable electrometric p H determination is inconvenient.
Mixtures of dimethyl sulfoxide (DMSO) with water have recently become increasingly popular and useful as reaction media. I t has been apparent in our own and other work that a rapid and reliable method of determining p H in such solvent mixtures would be desirable. Unfortunately, the conventional glass electrode-calomel electrode p H meter responds slowly and sometimes irreproducibly in such m i x t u r e ~ . ~We J have therefore calibrated a set of indicators for use in 80% DMSO-20% water, by weight. This is a convenient solvent for many purposes because it dissolves many ionic and molecular substances freely. It has a dielectric constant (72) only a little below that of pure water.4 It is not highly hygroscopic and shows solvent acid-base properties quite different from those of pure water. The calibration was carried out b y the “ladder” technique,5 starting with dilute HC1. For the most part colorless and indicator acids were alternated on the “ladder.” The overlap is sufficient so that a t least one indicator, capable of giving quantitative results, is available at any pH. The indicators used all have absorptions characteristic of both the acidic and basic forms so that the p H can be determined from a ratio of absorbances, without an exact knowledge of the indicator concentration. The ladder was continued until the p H of dilute KOH solutions could be determined. These latter measurements permit the determination of the autoprotolysis constant of water in this solvent. The dissociation constants, KHA,of the acids, HA, are represented by eq l..Charges are not shown for HA and A
because various charge types were used. Activity coefficients are represented by y. For each indicator two wavelengths, XI and XZ, were chosen. Extinction coefficients of t ~ t ~~ A l, ,and t ~ j re~ ~ A and HA were designated spectively. The wavelengths were chosen so as to minimize the ratios t H A 2 / t ~ ’ and tA2/tA1 which were designated R1 and Rz. Rs is f H A 1 / t ~ A 2 . The R values were determined from spectra of solutions in which the indicators are completely in the acidic or basic form. When the absorbance at each wavelength, A,, is given by
KHA
is given by eq 3a
When €HA1 is not significantly different from zero, which is the case for many of our indicators, eq 3a simplifies to eq Ba ( 4a)
Equation 3a or 4a was used to determine all the indicator dissociation constants reported in this paper. Ionic strength, u, in these measurements ranged between 2 x 10-3 and 4 x 10-4 M . Equation 5 , a form of the Debye-
log Y, =
,
-A ( Z , ) z ~ ” 2
+
1 Bdu”’ Huckel limiting law,6 was used to estimate the activity coefficients of ionic species. In this solvent A is 0.710 when concentrations are expressed as moles per liter, and Bd was taken as 1.42.? The activity coefficients of neutral species were assumed to be unity. All these activity coefficients refer, of course, to infinite dilution in this soluent mixture, as the standard state. In all these solutions ( H + ) was known, either from the concentration of HC1 or from the composition of a buffer. In the case of a buffer the value of KHAfor the buffer was previously determined. Equations 3a and 4a are readily rearranged to 3b and 4b so as to make them more convenient for the determination of ( H + ) from the, now known, K H A values and measured absorbances. They were used in this form to determine ( H + ) for nonindicator buffers to which a very small amount of indicator was added. Since the buffer ratio was known or separately determined in these solutions, KHA could be calculated. In this way the “ladder” was extended. Experimental Section
DMSO (Aldrich Chemical Co.) was twice distilled under vacuum from calcium hydride, discarding the first and last 15% each time. Such DMSO was shown by Karl The Journal of Physical Chemistry. Vol. 78. No. 4. 1974
422
Ernest H. Baughman and Maurice M. Kreevoy
TABLE I: DKRAfor Buffer Acids Compound _______
~~
ClzCHCOOH ClCHzCOOH CHaCOOH p-CNCsHaOHO p-ClCeHaOH 2,4,6- (CH,) ~ C G H ~ Q H HzO
3.87 5.69
0.01 0.01 8 . 0 0 f 0.02 9.83 f 0.01 11.98 f 0.01 13.83 f 0 . 0 3 18.38 f 0.02h =!z =!z
____
~~
2.61 2.82 3.24 1.88
2.60 2.95 4.38
27.22 0,6155 0.4321
0.0623 0.2866 0.3536
289 312 309
250 284 282
1,265 2. 87b 4 , 76c 7.95d 9.38e 10.88f 14.00
a M. Rondell and C. F. Failey, Chem. Rev., 4, 291 (1927). D.J. G. Ives and J. H. Pryor, J. Chem. Soc., 2104 (1955).' H. S. Harned and R. W. Ehlers, J. Amer. Chem. Soc., 66, 652 (1933). G. W. Wheland, R. M. Brownell, and E. C. Mayo, ibid., 70,2492 (1948). C. M.Judson and M. Kilpatrick, ibid., 71, 3110 (1943). G. R. Sprengling and C. W. Lewis, ibid., 76, 5709 (1953). R3 is 0.0263. *For water K H Ais taken as (H+)(OH-)r+?
'
TABLE 11: PKHAfor Indicators ~~
Compound
PKHA (80% DMSO)
Ri
Re
2,4-Dinitrophenol Bromcresol Green Bromcresol Purple Bromthymol Blue Cresol Red 2,4-Dinitrodiphenylamine 2,6-Dichloro-4-nitroanaline
4.16 i 0 . 0 0 5.73 f 0.02 7.63 rt 0 . 0 1 8 . 9 5 i 0.01 10.68 i 0 . 0 2 11.79 i 0 . 0 1 14.46 f 0 . 0 1
0.6692 0.2528 0.2157 0.192 0.2417 '0.2073 0.3970
0.0524 0,0046 0.00647 0.0107 0.00269 0.01527 0,0198
Ra
1.065
X2
419 615 590 620 581 495 465
A1
303 444
424 454 419 357 362
~~
PKHAWO)
4.09c 4.90* 6.30~ 7.30* 8,20a
a W. R. Brode, J. Amer. Chem. Soc.. 46. 589 (1924). I. M.Kolthoff and T. B. Reddy, Inorg. Chem., 1,189 (1962). H . V. Halban and G. Kortiim, Z . Phys. Chem., A170, 351 (1934).
Fischer titration and by nmr measurements to have a residual water content of about O.l%.s The water used was deionized and then d i ~ t i l l e d .Hydrochloric ~ acid solutions were made up by dilution of constant boiling HC1.1° Reagent grade acetic acid (E. I. du Pont de Nemours & Co., Inc.) was redistilled before use. Both 4-chlorophenol (Aldrich Chemical Co.) and 2,4,64rimethylphenol (Aldrich Chemical Co.) were distilled under vacuum (mp 40-41 and 71-71.5", respectively; lit. mp 4111 and 69"12). 4-Cyanophenol (Aldrich Chemical Co.) was recrystallized from benzene (mp 112.6-113.6'; lit. mp 112.0-112.4°13). Both 2,4-dinitrophenol (Eastman Kodak Co.) and 2,4-dinitrodiphenylamine (Aldrich Chemical Co.) were recrystallized from ethanol (mp 112.6-113.6 and 156.6-157.0"; lit. mp 112-11314 and 154.2-155.5",15 respectively). 2,6-Dichloro4-nitroaniline (Aldrich Chemical Co.) was used without purification (mp 190-191"; lit. mp 189'16). Bromcresol Green, Bromcresol Purple, Bromthymol Blue, Methyl Red, and Cresol Red were dissolved in aqueous 5% NaHC03 solution and precipitated from the hot solution by drop-wise addition of aqueous HCl.I7 The process was repeated until no further increase in the extinction coefficient was noted. No attempt was made to duplicate the literature values of Amax or e's as these are solvent dependent.18 (The commercial products appeared to be about 25% indicator and 75% low molecular weight acid which was not identified.) In the future these indicators will all be made available, as concentrated solutions, by the Ventron Corporation through Alfa Products. The ratios, R1, R2, and R3, required by eq 3 and 4, were determined after taking spectra in successively more basic or more acidic solutions till further changes in pH produced no further changes in the spectra. In these solutions, and also in the solutions used to determine KHA values, total indicator concentrations were around 5 x M. In the buffer solutions the total concentration of buffer constituents was around 10-3 M . The buffer constituents were nearly nonabsorbing in the visible, except for 2,4,6trimethylphenol, in which case its absorbance was canThe Journal of Physical Chemistry. Voi. 78. No. 4 . 1974
celled by using, as reference, solution containing the buffer but not the indicator. Thus they did not interfere with the indicator measurements, which were made in 10-cm cells. The indicator concentrations, on the other hand, were so low that absorption by the indicators was less than 1% that of the phenolic buffer constituents a t their maxima in the uv. For the phenolic buffers this enables us to determine the buffer ratios, also, from eq 3 and 4, without a priori knowledge of the composition of the solution, using 0.1-cm cells. The buffer ratios so determined were always in reasonable agreement with those given by the a priori compositions of the solutions, however. The buffer ratios in the carboxylic acid buffers were determined from the stoichiometric compositions of the solutions. These were less subject to error than the buffer ratios in the phenolic buffers because the carboxylic acid buffers are acidic, and, therefore, less prone to absorb COz from the atmosphere. Spectrophotometric measurements were made with a Beckman DK-2 spectrophotometer, calibrated a t 536.4, 453.6, 385.8, 360.8, 333.8, 287.6, and 279.3 nm with holmium oxide (Beckman Std.). All measurements were made a t ambient temperature, but room temperature was maintained within 2" of 25" by air-conditioning, and temperature fluctuations are not thought to contribute significantly to the inaccuracy of the reported K H Avalues. Results Table I gives the mean pKHa values for the (nonindicator) buffer acids. In the case of the phenols X,I hz, R1, Rz, and, where necessary, R3, are also given. Table I1 gives the same information for the indicators. In each case the value is the mean of a t least five separate determinations. The cited uncertainty is the probable error of the mean value. Table I11 shows a typical set of determinations, The largest deviations from the mean pKHA values were around 0.1 and the average deviations from the mean values around 0.02, so that it ought to be possible to use these indicators to determine pH with about that precision.
423
Determination of Acidity in 80% DMSO-20% Water
20
3
i
1
15 13 L
t
Y
30 v)
< x
Y
a.
5
I
0.0
I
I
.I
.It
I
I
-4
.3
XDMSO Figure 1. P K H ~ Oas a function of XDMSO, the mole fraction of DMSO in the solvent. Values of Wooley and Hepler are indicated by circles, those of Das and Kundu by squares, the present
values by crossed circles. Open figures are for P K H , ~ ; closed figures are for p(KHZo/(H20)). If either function of KH,O is extrapolated to the present XDMSO, a value similar to the present one is obtained.
TABLE 111: K E ADeterminations for 2,6-Dichlor0-4-nitroaniline WCla
A485
2 . 5 2 x 10-14 1 . 7 7 x 10-14 1 . 2 7 x 10-14 5.90 X 4.87 X 4.17 X 2 . 6 2 X 10-16 ~ 1 - l a b0 -10 -1ac
0,123 0.160 0.215 0.371 0,401 0.458 0.595 1.010 0.000
ASs6
0.358 0.331 0.313 0.266 0.249 0.222 0.178 0.020 0.401
KEA X 1016
3.30 3.43 3.52 3.60 3.22 3.56 3.72
Determined from the uv spectra of 2.06 X 10-8 M (CHa)aCeHzOHsohtions. Used to determine the spectrum of HA. ' Used to determine the spectrum of A.
Discussion There are two points at which our results can be subjected to critical, external checks. The PKHAof acetic acid has been previously determined, conductometrically, by M0re1.l~He obtained a value of 8.00. This is in excellent agreement with the value given in Table I. Two recent papers have given pKHA values for water in DMSOwater mixtures, 0 to 60% DMS0.20,21For water KHA was defined as ( H + ) ( O H - ) T * rather-than ~ by eq 1. Plots of PKHA or PKHA log (H2O) are, empirically, linear functions of the mole fraction DMSO. Figure 1 shows that such plots extrapolate to our present value of the pKHA for water quite acceptably. Since the PKHA for water depends on all the others, this supports all the present results. In addition, the reported values of pKHA for chloroacetic acid and dichloroacetic acid are reasonably related to those reported22 at an ionic strength of 0.2 M . These correspondences suggest that the inaccuracy of the pKHA values reported are not much worse than their impreci-
+
Figure 2. The relation between pKHA values in the present solvent and those in water. The slope of the linear relation is i.4.
sion, and p H values obtained with the aid of the indicators will have a corresponding reliability. Internal consistency of the present data is further supported by Figure 2. The linear plot of pKHA(solvent) against pKHA(water) is expected23 and observed. The nonunit slope of this plot has certain implications for acidity functions which will be discussed elsewhere. Further, approximate, use of the indicators can be made in DMSO-water mixtures having 0-80% by weight DMSO, if it is assumed that their PKHA values are linear functions of the mole fraction of DMSO. Such relations have been found to have a wide, approximate, validity.22 References and Notes (1) Supported, in part, by the U. S. National Science Foundation through Grant No. GP-31360X. (2) T. B. Reddy, Ph.D. Thesis, University of Minnesota, 1960. (3) C. D. Ritchie and R. E. Uschold, J. Amer. Chem. SOC.,89, 1721 (1967). (4) E. Tommilaand A. Pajunen, Suom. Kemi. B , 41, 172 (1968). (5) W. Bover and P. Zuman, J. Amer. Chem. SOC.,95, 2531 (1973). (6) H. S. Harned and W. E. Owen, "The Physical Chemistry of Electrolytic Solutions," 3rd ed, Reinhold. New York, N. Y . , 1958, p 66. (7) J. Kielland, J. Amer. Chem. SOC.,59, 1675 (1937). (8) We are indebted to Sister Lavonne Abts for the nmr measurements. (9) J. E. C. Hutchins, Ph.D. Thesis, University of Minnesota, 1969. (10) G. A. Hulett and W. D. Bonner, J. Amer. Chem. SOC., 31, 390 (1909). (11) W. J. Wohlleben, Berichte. 42, 4373 (1909). 43, 1194 (12) C. W. Porter and E. H. Thurber, J. Amer. Chem. SOC., (1971) \ . - - . I .
(13) G. W. Wheland, R. M. Brownell, and E. C. Mayo, J. Amer. Chem. SOC., 70, 2493 (1948). (14) W. E. Bachman, J. M. Chenenda, N. C. Deno, and E. C. Horning, J. Org. Chem., 13, 390 (1948). 41, 1016 (15) E. J. Hoffman and P. A. Dane, J. Amer. Chem. SOC., (1919). (16) k. L.'Datta and H. K. Miller, J. Amer. Chem. SOC., 41, 2036 (1919). 48, 2216 (17) W. R. Orndorff and A. C. Purdy, J. Amer. Chem. SOC., (1926). (18) R. Stewart and J. D. O'Donneil, Can. J. Chem., 42, 1681 (1964). (19) J. P. Morel, Bull. SOC.Chim., Fr., 1406 (1967). (20) A. Dasand K. Kundu, J. Phys. Chem., 69, 730 (1973). (21) E. Woolleyand L. Hepler, Anal. Chem., 44, 1520 (1972). (22) N. M. Ballash, E. E. Robertson, and M. D. Sokolowski, Trans. Faraday Soc., 66, 2622 (1970). (23) R. P. Bell, "The Proton in Chemistry," Corneil University Press, Ithaca, N. Y., 1959, Chapter 1 1 1 . The Journal of Physical Chemistry, Vol. 78, No. 4 . 1974
