Dissociation Constants of Sulfamic Acid, Salicylic Acid, Thymol Blue, and Bromocresol Green in Anhydrous N,N-Dimethylformamide Arthur J. Libbey, Jr. and John T. Stock Department of Chemistry, University of Connecticut, Storrs, Conn. 06268 N,N-DIMETHYLFORMAMIJJE (DMF) is a useful solvent for nonaqueous acid-base titrimetry ( I ) . However, commercial or partially purified D M F is commonly used. Reported pK, values obtained in purified D M F are summarized in Tables I and 11. Excluding the various nitrophenols, constants for only two acid-base indicators are available. This is probably due to the limited use made of spectrophotometric methods. Particularly in the case of 2,4,6-trinitrophenol (picric acid) and hydrochloric, acetic, and benzoic acids, there is considerable disagreement concerning acid strength. The acid dissociation constants in highly purified DMF, of sulfamic acid, salicyclic acid, thymol blue (acid range), and bromocresol green have now been determined spectrophotometrically. All values are referred to p-toluenesulfonic acid, for which pK, = 2.92 ==I 0.08 (standard deviation) is the mean of values reported by Paul et al. (2) or calculated subsequently (3) from their data. These data cover the 50-fold concentration range -0.005M to -0.23M. Because of the absence of data for activity coefficient calculations in DMF, concentration dissociation constants were determined. Dilute solutions of constant total buffer concentration 0.04M were used to minimize interionic effects and to permit reasonable intercomparison of these pK, values. The (~K,)I.values of thymol blue and of bromocresol green were calculated from
where An and AB are the absorbances when a fixed concentration of the indicator is completely in the acid form and in the base form, respectively, while A is the absorbance when both forms are present (4). The term pcH is given by PCH = -log [H+]
=
where ( p K a ) * refers ~ to the buffer acid HA. CHAand CAare the analytical concentrations of HA and of conjugate base A , respectively, (charges are ignored), while AB - A [HIn] = ____ * C I n AB - AA
10.2
o-Chlorobenzoic 12.27,c11.17d Salicylic 8.23,7.94 o-Nitrobenzoic 10.05 p-Nitrobenzoic 10.62 3,5-Dinitrobenzoic 8.48 P 2.95 p-Toluenesulfonic C a C = Conductance, P = Potentiometric, S metric. * DMF contained 10% H20. Potassium salt buffer. Lithium salt buffer.
(13) (2) =
Spectrophoto-
The pK, of p-toluenesulfonic acid was used to obtain the pK, of thymol blue. The latter pK. and absorption measurements of thymol blue in sulfamate buffers led to the pK, of sulfamic acid. Similar stepwise procedure led to the constants for bromocresol green and for salicyclic acid. The reported pK,’s were obtained at wavelengths X giving maximum absorption. Standard deviations refer to the individual steps of the procedure. Determinations made at X +5 nm gave pK,’s that were within one standard deviation from the values at A. The pcH ranges used were kept small to minimize changes in ionic strength (3).
(3)
where CI,is the analytical concentration of the indicator. [H+] was determined by successive approximations. (1) R. S. Kittila, “Dimethylformamide Chemical Uses,” E. I. DuPont de Nemours & Co., Wilmington, Del., 1967, Chap. 32. (2) R. C. Paul, P. S . Guraya, and B. R. Sreenathan,Indian J. Chem., 1, 335 (1963). (3) A. J. Libbey, Jr., Ph.D. Thesis, University of Connecticut,
Storrs, Conn., 1969. (4) A. Albert and E. P. Serjeant, “Ionization Constants of Acids and Bases,” John Wiley & Sons, Inc., New York, 1962. 526
Table I. Dissociation Constants of Acids in DMF Acid PKa Perchloric Strong Hydrochloric Strong 3.55 3.4 Nitric Strong* Hydrobromic 1.78 Sulfuric 4.3 10.5 Hydrogen sulfide 2.13 Fluorosulfuric 3.85 Pyrosulfuric Hydrazoic 8.5 Formic 11.65 Acetic 13.3 11.1 Chloroacetic 9.0 Dichloroacetic 7.2 Benzoic 11.70 12.20
ANALYTICAL CHEMISTRY, VOL. 42, NO. 4, APRIL 1970
( 5 ) J. Badoz-Lambling and G. Demange-Guerin, Bull. SOC.Chim.
Fr., 1964, 1354. (6) J. Juillard, J . Chim. Phys., 64, 1190 (1966). (7) M. T6z6 and R. Schaal, Bull. SOC.Chim. Fr., 1962, 1372. (8) A. B. Thomas and E. G. Rochow, J . Amer. Chem. Soc., 79, 1843 (1957). (9) B. W. Clare, D. Cook, E. C . F. KO, Y . C. Mac, and A. J. Parker, J . Amer. Chem. SOC.,88, 1911 (1966). (10) P. G. Sears, R. K. Wolford, and L. R. Dawson, J. Electrochem. SOC.,103, 633 (1956). (11) S. M . Petrov and Yu. I. Umanskii, Zh. Fiz. Khim., 41, 1374 (1967). (12) J. Juillard and B. Loubinoux, Compt. Rend., 264, 1680 (1967). (13) J. Juillard and A. Mallet, ibid., 264, 2098 (1967).
Table 11. Dissociation Constants of Phenols, Thiophenols, and Acid-Base Indicators in DMF Compound PK, Methoda Reference Phenol >15 S (9) 12.19, 12.14 P (11,13) o-Nitrophenol m-Nitrophenol 13.85 P (11) 12.18, 11.84 P (11, 13) p-Nitrophenol 10.9 S (9) ea. 12.2 S (14) 2,4-Dinitrophenol 6.33, 6.36 P (11,13) 6.00 S (9) 2,5-Dinitrophenol 8.61, 8.78 P (11,13) 2,dDinitrophenol 6.17,6.07, P (11-13) 6.05 2,4,6-Trinitrophenol 3.65, 2.20 P (11,12) 1.20 C (10) 2,6-Dichlorophenol 12.55 P (13) 2,4,5-Trichlorophenol 12.46 P (11) 2,6-Dibromo-45.70 P (13) nitrophenol 2,6-Dichloro-45.71 P (13) nitrophenol Thiophenol 10.7 S (9) 4-Nitrothiophenol 6.3 S (9) Bromocresol green 8.3 S (9) p - Aminoazobenzene 1.05 S (5) a C = Conductance, P = Potentiometric, S = Spectrophotometric. Table 111. pK, of Thymol Blue in DMF pToluenesulfonate buffers X = 400nm. A A = 0.033. A B = 0.760 CTB= 7.0 X 10-'M CEA
x
IO~M 6.81 6.66 6.73 5.04 4.04 3.36 3.33 2.88 a
CA
x
103~
33.7
A
0.070a 0.069b 0.073 0.091 0.105 0.123 0.112b 0.132
33.3 33.3
34.9 36.0 36.7 36.7 37.1
PCH PK~ 3.62 4.86 3.64 4.86 3.64 4.87 3.78 4.84 3.90 4.86 4.00 4.85 4.00 4.86 4.05 4.85 Av 4.86 f 0.01
A A = 0.031; A B = 0.745. A A = 0.032; A B = 0.689.
EXPERIMENTAL
Apparatus. Absorbance measurements with a Beckman Model B spectrophotometer were made at room temperature (22 to 27 "C). The 1-cm Pyrex and Corex cells had tightly-fitted polyethylene caps. Reagents. D M F (Matheson, Coleman and Bell, reagent grade) was twice azeotropically distilled with dry benzene, dried over barium oxide, then vacuum distilled under nitrogen (10-11 torr). Purified D M F was collected in a blackened receiver. The tip of this was inserted into a glove bag, so that D M F could be dispensed under dry nitrogen. Batch yields were approximately 50 %, with specific conductance 1 X 10-7 to 4 X lo-' ohm-' cm-I (25 "C), water content (Karl Fischer) less than 0.0003 Z wt/vol, and base impurity approximately 5 X 10-BM (3). Anhydrous p-toluenesulfonic acid (TSA) was prepared from the monohydrate (Eastman Kodak White Label; J. T. Baker Analyzed Reagent). This was recrystallized from (14) A. J. Parker, Department of Chemistry, Australian National University, Canberra, Australia, 2600, private communication, September 11, 1969.
I
00
400
350
450 WAVtltHGIH
500
550
600
nm
Figure 1. Absorption spectra of 7.74 X 10-5M thymol blue in DMF solutions -O-O-,
0.1094M TSA
-A-A-,
0.1035M sodium p-toluenesulfonate
acetic acid, dried, then twice azeotropically distilled with dry benzene and vacuum dried over phosphorus pentoxide (3). Acidimetric assay of the approximately 20 % batch yields gave 99.6 % to 99.9 % purity. Sodium p-toluenesulfonate (Eastman Kodak White Label) was twice recrystallized from 95% ethanol, dried at 110 "C, and stored over barium oxide. Sulfamic acid (SA) was recrystallized, dried, and stored over barium oxide. Acidimetric assay gave 99.79 purity. Lithium sulfamate was obtained by neutralizing aqueous SA with lithium carbonate, then evaporating to a slush. This was treated with hot methanol, filtered, then acetone was added to precipitate the very deliquescent lithium sulfamate. This was dried and stored over phosphorus pentoxide. The assay (as Li2S04)was 99.1 %. Indicators (sodium salts) and other chemicals were reagent grade and were dried and stored over Drierite or barium oxide. Procedure. All glassware was stored over barium oxide. Recrystallization, solution preparation, etc., were performed under slight positive pressure of nitrogen in a glove bag that contained barium oxide. Cells were filled in the glove bag, which contained microburets for dispensing D M F and stock solutions in DMF. Measurements were corrected for any small absorbances of the buffer solutions. RESULTS AND DISCUSSION
THYMOL BLUE(TB). Figure 1 shows the absorption spectra of the acid ( A ) and base (B) forms of TB in TSA systems. Molar absorptivities are 419, eB 9960 at 400 nm ( B maximum) and 983 eB 49.9 at 545 nm ( A maximum). Complete conversion of TB to A was verified by absorbance measurements in 0.083 to 0.134M TSA solutions. Measurements in 0.051 to 0.104M sodium p-toluenesulfonate solutions showed complete conversion of TB to B. Linear Beer's law plots were obtained for both A and B. Absorbances were unchanged after 2 hr. Table I11 lists pK, values obtained from measurements in TSA buffers. The linear plot of pcH against log [(AB-A)/(A-AA)] has least-squares slope and pcH intercept of -0.98 and 4.83, respectively. Interference from variations in activity or from ion-pair formation thus appears to be minimal (15). SULFAMIC ACID. Low solubility in D M F prevented the use of sodium sulfamate, but the lithium salt was satisfactory (15) I. M. Kolthoff, S . Bruckenstein, and M. K. Chantooni, Jr., J . Amer. Chem. SOC.,83,3927 (1961). ANALYTICAL CHEMISTRY, VOL. 42, NO. 4, APRIL 1970
527
Table V. pK, of Bromocresol Green in DMF Sulfamate Buffers
I
625 nm. A A = 0.001. A B = 0.725 CBCG= 1.325 X 10-6M CHAX 10aM CA X 103M A PCH =
I
Y
E 04
6.60 4.96 3.97 3.31 6.20
rn
02
(1
00 400
450
500 550 WAVtltNGIH nm
600
~~
1.324 X 10-5MBCG in 0.05855M tris
X = 400nm. A A = 0.039. A B = 0.730.
PK8 8.09 8.01 7.94 7.88 7.84 8.08 8.03 7.99 7.98 8.00
as a buffer component. Measurements were made in buffers of total concentration 0.04 and 0.085M. SA proved to be much weaker than expected, so that the mixtures used had low buffer capacity. Kolthoff and Reddy (16) encountered an analogous situation in their work on salicyclic acid in dimethylsulfoxide. The results listed in Table IV give average pK, values of 7.95 f. 0.10 (O.OSSM), 8.02 f 0.04 (0.04M), and 7.99 f 0.07 (overall). Despite the massive correction applied to CA,the agreement of the p& values is good. The results obtained in 0.04M buffers involved smaller changes in ionic strength and are, therefore, the more reliable. The linear plot of pcH against log [ ( A - A A ) / ( A B - A ) ]for the 0.04M results has least-squares slope and pcH intercept of 0.97 and 4.87, respectively. The decrease in pK, with increase in pcH is obvious in the case of 0.085M buffers. A possible cause is a trace of base impurity in the solvent (4). Zaugg and Schaefer (17) reported basic impurity concentrations of from 4.79 X to 11.4 X 10-FMin 17-day old DMF that had been purified in a manner similar to that used in the preient work. Allowance for the presence of 5 X 10-6M of such impurities changes the average pK, values given above to 7.94 f 0.10 (O.O85M), 7.98 f 0.03 (0.04M), and 7.96 f 0.07 (overall). The low acid strength of SA in D M F may be partially due (16) I. M. Kolthoff and T. B. Reddy, Znorg. Chem., 1, 189 (1962). (17) H. E. Zaugg and A. D. Schaefer, ANAL.CHEM., 37,2121 (1964). 528
8.72 8.87 8.98 9.06 8.72
ANALYTICAL CHEMISTRY, VOL. 42, NO. 4, APRIL 1970
~
Table VI. pK, of Salicylic Acid in DMF X = 625 nm. cBCG
CHAX lO’M 2.005 1.344 1.003 0.802 2.014
Table IV. pK, of Sulfamic Acid in DMF CTB= 7.47 X 10-6M CHAX 104~4 C A x 1 0 4 ~ A PCH 847 0.394 0.461 5.06 0.490 5.13 847 0.789 0.508 5.18 847 1.183 847 1,577 0.520 5.22 847 1.972 0.533 5.26 399 0.197 0.5250 5.19 399 0.395 0.550“ 5.27 399 0.592 0.568a 5.31 399 0.789 0,590” 5.38 407 0.590 0. 560b 5.32 a A A = 0.029. A B = 0.758. CTB = 7.70 X 10-’M. * A A = 0.031. A B = 0.745. CTB= 7.74 X 10-6M.
A A = 0.001. AB =
PK8 9.47 9.49 9.50 9.54 9.53 Av 9.51 f 0.03 0.730. CBCG= 1.246 X 10-‘M. 0.110 0.140 0.167 0.183 0.100“
650
Figure 2. Absorption spectra of bromocresol green in DMF solutions -e-&, 1.302 X lW5MBCG in 0.03212M SA -A-A-,
33.40 35.07 36.07 36.71 31.57
a
A A = 0.002. A B = 0.730 1.324 x 10-5~4
=
C A X 102M 2.000 2.660 3.000 3.200 2.017
A A = 0.001. A B =
PCH PK, 8.84 8.84 9.23 8.93 9.41 8.93 8.95 9.56 8.81 8.81 Av 8.89 k 0.06 0.738. CBCG= 1.318 X lO-’M. A
0.131 0.251 0.325 0.385 0. 123a
to the weak solvation of the small sulfamate ion. Zaugg and Schaefer (17) and Parker (18) have discussed the effects of anion size and solvation upon polarizability. BROMOCRESOL GREEN(BCG). Figure 2 shows the absorption spectra of the acid ( A ) and base ( B ) forms of BCG in DMF solutions. These spectra were determined in the presence of SA and of tris-[2-amino-2(hydroxymethyl)-I &propanediol] (tris) respectively. Complete conversion to the acid and base forms was verified and linear Beer’s law plots were obtained. The average BCG molar absorbtivities in solutions 95, of SA concentration from 0.03283M to 0.07897M are eB 55000 at 625 nm (B maximum). In solutions of tris concentration from 0.03269M to 0.05855M, the average values are eA 17500, eB 5020 at 415 mm ( A maximum). There is an isosbestic point at 504 nm. In 0.05885M tris, eB at 625 nm is 3.5 greater than in the 0.03269M solution. Use of this higher instead of the average value of eB at 625 nm raises the pK, values given in Table V by only 0.03 unit. The linear plot of pcH against log [(A,-A)/(A-AA)]has least-squares slope and pcH intercept of - 1.07 and 9.55, respectively. The small spread of pK, values is probably due to the narrow range of pcH, and hence of ionic strength. The marked salt effect in aqueous systems for a HIn-/W2 indicator such as BCG (19) should be enhanced in a solvent of dielectric constant lower than that of water (20). The absorbance of 6.6 X 10-6M BCG in DMF (no buffer added) indicates the base form:acid form ratio of 4.6. For pK. = 9.5, this leads to the conclusion that the DMF used in the present work was approximately 5 X 10-6M in base impurity. (18) A. J. Parker, Quart. Reu. (London), 16, 163 (1962). (19) I. M. Kolthoff and S. Bruckenstein, in “Treatise on Analytical Chemistry,” I. M. Kolthoff and P. J. Elving, Ed., John Wiley & Sons, Inc., New York, 1959, Part I, Vol. 1, Chap. 10. (20) R. G. Bates, “Determination of pH. Theory and Practice,” John Wiley & Sons, Inc., New York, 1964, Chap. 6.
Parker et a/. (9) have also determined the pK, of BCG in DMF, and report a value that is approximately 1.2 units lower than obtained in the present studies. These workers ( 9 ) give pK, values for acetic acid, benzoic acid (Table I), and p-nitrophenol (Table 11) that are 1 to 2 units below the values obtained by other workers. However, Parker (14) has recently revised the value for p-nitrophenol in D M F to ca. 12.2,and has indicated that this revision will change the pK, values reported by Parker et al. (9)for other compounds. SALICYLICACID. Absorption measurements of BCG were made in salicylic acid-sodium salicylate buffers. The isosbestic point remained at 504 nm. Results are listed in Table VI. The linear plot of pcH against log [(A-A,)/(A,-A)]
has least-squares slope and pcH intercept of 1.1 and 8.86, respectively. Juillard and Mallet (13) give 8.23 as the value of the pK, for salicylic acid in DMF. These workers used measurements based on picric acid, for which widely-differing pK, values have been reported (Table 11). The concentration of picric acid in the salicylate buffers is not stated. Petrov and Umanskii (11) have shown that the pK, of picric acid in D M F is extremely concentration dependent. RECEIVFD for review November 28, 1969. Accepted January 27, 1970. Work was supported in part by grants from the U. S. Atomic Energy Commission (Contract AT (30-1)1977)and the University of Connecticut Research Foundation.
Polarography of Tin(lV) in Presence of Chloranilic Acid Clifford C. Boyd and Jeffrey G . Wardeska Department of Chemistry, East Tennessee State University, Johnson City, Tenn. 37601
THE POLAROGRAPHIC REDUCHON Of tin(IV), showing two distinct cathodic waves, occurs only when certain complexing agents are present in the supporting electrolyte. Previous reports describe the use of chloride ( I ) , pyrogallol (2), and 3-mercaptopropionic acid ( 3 )as ligand sources. The observation that tin(1V) complexes with chloranilic acid (2,5-dichloro-3,6-dihydroxy-p-benzoquinone) in solutions sufficiently acidic to prevent hydrolysis prompted the investigation of chloranilic acid as a complexing agent for the polarography of tin(1V). Excellent polarograms with two waves corresponding to the reduction of tin(1V) to tin(I1) and tin(I1) to tin amalgam are produced in citric acid and hydrochloric acid media at low pH values in the presence of chloranilic acid. I
EXPERIMENTAL Apparatus. A Metrohm Polarecord E 261 in conjunction with a Metrohm Polarographie Stand E 354, complete with drop controller for rapid polarography and a silver-silver chloride reference electrode, was used for all the polarographic work. The capillary had a drop time of 2.0sec with a rate of mercury flow of 3.5 mg per sec at an applied potential of -0.30 V and an effective pressure of mercury of 30.1 cm. Rapid polarography was conducted with the drop controller set at 3 impulses per second. A Cary 14 recording spectrophotometer was used for all the spectrophotometric measurements. All polarographic determinations were made at 25.0 f 0.1 “C. Reagents. A stock solution of 0.01M tin(1V) in 1M hydrochloric acid was prepared by the method of Bard (2). Chloranilic acid, C6C12H204, (Fisher certified reagent) was used without further purification. Samples of the complexing agent weighing 0.209 g were made up to 100 rnl with 0.1M citric acid to give a solution which was 0.01M in chloranilic acid. The citric acid used was reagent grade. (1) I. M. Kolthoff and J. J. Lingane, “Polarography,” Interscience
Publishers, Inc., New York, 1952, p 526. (2) A. J. Bard, ANAL.CHEM.,34, 266 (1962). (3) S. L. Phillips and R. A. Toorney, ibid., 37, 607 (1965).
-08
-07
-06
-05
VOLTS
-04 VI
-03
-02
-01
-
00
ApIAgCl
Figure 1. Polarogram of tin(1V) in the presence of chloranilic acid A . 0.1M citric acid, 2.0mM chloranilic acid, pH = 0.9 (blank) B . 0.1M citric acid, 2.2mM chloranilic acid, 0.73mM tin(IV), pH = 0.9, no gelatin
Procedure. Aliquot portions of the 0.01M tin(1V) solution were added to the citric-chloranilic acid solution in the polarographic cell, and the contents were stirred with nitrogen for at least 5 min. After adjusting the pH to the desired value with 2 M hydrochloric acid or 2 M sodium hydroxide, the deaeration with nitrogen was continued for an additional 10 min. The polarograms were scanned over the 0.0 to - 1 .O voltage range. RESULTS AND DISCUSSION
A typical polarogram for the tin(1V)-chloranilic acid complex is shown in Figure 1. The first wave is apparently produced by the reduction of the stannic complex to tin(I1); the second wave is attributed to the reduction of tin(I1) to the amalgam. The diffusion currents, id, and i d z , of the waves are not significantly affected by pH variations of ANALYTICAL CHEMISTRY, VOL. 42, NO. 4, APRIL 1970
529
