IONIZATION CONSTANT OF ORTHANILIC ACID
March, 1067
333
TABLEI CONCENTRATIONS OF THE THREE SATURATED PHASESWT. '% --Isoortsnr Temp.
Iso-
OC.
octane
25.0 31.5 33.7
72.2 64.2 57.5
ph8EFluorocarbon
11.9 13.7 8.9
-PerRuorotri-n-butylatninc 180Fluorocarbon
Nitroethane
octane
15.9 22.1 33.6
7.5 12.4 8.7
respectively. The differences in 6 values are shown in column 2 of Table I1 for comparison with those calculated from the data for the binary and ternary systems. The values shown in column 3 were calculated from the critical solution temperatures of the hydrocarbon-fluorocarbon and hydrocarbon-nitroethane binary systems using equation 5 2 zlz*Vlv2>
RT, (z,V, -+ SZV# which has been derived from equation TABLEI1
(5) 1.22
SOLUBILITY PARAMETER DIFFERENCES 611
- 6F
25.0" 31.5' 33.7" 8N
- 8F
25.0' 31.5" 33.7'
To
Phase boundaries aii aN
(1F
2.20 0.95
- 6B
25.0' 31.5' 33.7" IN
(@)'I:
2.30 2.31 2.40
2.34 2.24 2.45
5.48
3.50 3.44 3.48 3.48 3.44 3.33
r-----hitroolliane phase-lsoF~UOCONitrooctane carbon ethane
4.2 2.8 3.5
'26.9 29.5 39.6
1.9 2.9 4.9
71.2 67.6 56.5
agreement with those calculated from the physical properties of the pure components. Similar discrepancies were found in the case of solutions of hydrocarbons and fluorocarbon amines by Simons and L i n e v ~ k y ,Rotariu, ~~ Hanrahan and Fruin' and McLaughlin and Scott.8 Reedz4 has been able to explain a t least part of the aiscrepancy by considering the deviation from the geometric mean approximation due to the differences of the ionization potentials of the molecules. Discrepancies for solutions of hydrocarbons in solvents with a higher internal pressure have been observed and discussed by HildebrandPzs The parameter differences calculated from the phase boundaries in the three-component system agree with those calculated from the critical solution temperatures and compositions of the perfluorotri-n-butylamine-iso6ctane and nitroethaneb~ isooctane systems. The low values for 6~ obtained from the ternary data were unexpected. Acknowledgment.-We wish to thank the Minnesota Mining and Manufacturing Company for the perfluorotri-n-butylamine used in this investigatiori and The National Science Foundation for supporting the program of which this work was a part*
-
3.80 4.53
88.3 84.8 87.8
pharreNitroethane
3.91 3.48 3.50
3.78 5.65 4.20
The differences in 8 values calculated from the data for the ternary and binary systems are not in (22) See ref. 1, p. 253.
(23) J. H. Simona and M. J. Linevaky, J . Am. Chem. SOC., 74,4750
('y: ;'
19, 425 (1955), T. M. Reed, 111, THIa (25) J. H. Hildebrand, J . Cham. Phya.. 18, 1337 (1950).
THE IONIZATION CONSTANT OF ORTHANILIC ACID FROM 0 TO 50' BY MEANS OF E.M.F. MEASUREMENTS' BY It. NORMAN DIEBEL ANI) D. F. SWINEHART Department of Chemistry, University of Oregon, Eugene, Oregon Recebed September 81, 1968
The ionization constant of orthanilic acid has been determined from 0 to 50' by the use of cells without liquid junction. The equation -log K = 1160.68/T 0.0066339T - 3.2314expresses the ex erimental data as a function of temperature in the above temperature range wit,h a standard deviation of 0.0008 in -log for 11 experimental points. The standard entropy of ionization, ASo, was found to be -3.3 e.u., a value much larger than that for sulfanilic and metanilic acids.
+
Introduction The ionization constants and related thermodynamic quantities for sulfanilic acid2 and metanilic acida ( p - and m-aminohenzenesulfonic acids, respectively) have been reported previously from this Laboratory. Both acids are zwitterions with large charge separations and both show a remarkably Taken in part from the maater's thesis of R. Norman Diebel. (2) R. 0.MacLaren sild D. F. Swinehart, J . A m . Chcm. SOC.,7 8 , 1822 (1051). (3) R. D. McCoy and D. 8 . Swinchart, i b i d . , '76, 4708 (1054). (4) D. F. Swinohxrt. papor prrscntcd at the Northwcst Rcgional Meeting of thu A.C.S., Pullnian. Wash.. June 12-13. 1953. (1)
small entropy of ionization. It has been suggestedas4 that the small entropy of ionization of these two acids can be understood on the basis of a qualitative picture of the interaction of water dipoles and these zwitterions. On this basis a prediction was made regarding the entropy of ionization for orthanilic acid. The present paper reports experimental data for the determination of the ionization constant of orthanilic acid (o-aminohenzene.sulfonio acid) as a function of ternperatura and confirms the prediction regarding entropy of ionization.
334
R. NORMAN DIEBELAND D. F. SWINEHART
Previously reported values for the ionization conKtant of orthanilic acid were obtained from conductivity mcasurements and calculations were made without the benefit of the Debye-Huckel interionic attraction theory. These values are Investigator
K
1 ("C.)
Oswald' Boylee
3 . 3 x 10-8 4 . 2 9 X 10-8
25 25
It has been shown by Carr and Shutt' from measurements of the change in dielectric constant with pH for sulfanilic acid solutions that sulfanilic acid behaves as a switterion. Even though orthanilic acid is somewhat stronger than sulfanilic acid, it is not nearly as strong as the unsubstituted aromatic sulfonic acidss and it is reasonable to assume that orthanilic acid also exists in solution largely as the zwitterion. Hence the ionization constant is written for the reaction 0-
'HsNCsHiSOa- = o-HgNCeHiSOa-
+ H+
The general method of investigation is that developed by Harned and co-workers.9 The cells were of the type Pt, Hz/HOr(ml), NaOr(mz), NaCl(ma)/hgC1-Ag where HOr and NaOr are orthanilic acid and its sodium salt, respectively, and ml, m2 and m3 are molalities. By elimination of ~ H + Y H +from the cell potential equation 2.3025912T E = Eo
-
log mH+rncI-yH+yCl- (1)
and t,he thermodynamic ionization constant exp: ession K - ma+nlor-yH + Y O ~ mHOrYliOr
(2)
there results the relation (3)
The ionization constant was determined from eq. 3 by extrapolation of the left-hand side to zero ionic strength using a method of calculation similar to that developed by Hamer'O and essentially identical with that used in the previous papers.208 However, due to the fact that orthanilic acid is the strongest, acid of the three isomers, the calculations were long and tedious. The full calculation required four complete successive approximations at each temperat,ure in order to be sure that the extrapolations of the left-hand side of eq. 3 had converged to give constant values of -log K a t zero ionic strength. This full calculation was carried out at 0, 25 and 50" only. The final hydrogen ion molalities obtained at these temperatures were used to interpolate initial hydrogen ion molalities a t the remaining temperatures and successive approximation started with the resulting values. This procedure saved two complete successive approxi(5) W. Oswald, 2. p h p i k . Chem., 3, 406 (1889). (6) M. Boyle. .I. Ckem. Soc., 116, 1505 (1919). (7) W.Carr and W.J. Shutt, Trans. Faiaday SOC..86, 579 (1939). (8) W. J. Hanior, G . D. Pinching and S. F. Acreo, J . Research Natl. Bur. Slandards. 31, 291 (10483. (9) H. 8. liarned and B. B. Owcn, "The Physical Cheuniatry of Electrolytic Solutions," 2nd Ed., Reinhold Publ. Corp., Now York, N. Y., 1950. (10) W.J. Hanier. J . Am. Cham. Soc.. 66, 800 (1934).
Vol. 61
mations a t the remaining eight temperatures. Howevcr, the effect was to have made four complete approximations a t each temperature. The values of the molal electrode potentials, EO, of the silver-silver chloride electrode have been determined by Harned and Ehlers" and recalculated, along with 2.30259RT/F, in absolute volts by Swinehart'* using the constants of Bearden and Watts.la Experimental The materials and reagents other than orthanilic acid were purified in a manner similar to that described by MacLaren and Swinehart.' Eastman Kodak white label orthanilic acid, when first dissolved in hot water, gave a noticcable odor of sulfur dioxide which was eliminated by addition of strong acid. Therefore the first crystallization was made from a solution containing a proximately 20 ml. of concd. HCl per liter of solution. TIIS was followed by two or three recr stallizations from pure water. The white cr stals were cked in a vacuum dcsiccator over 90-95% H z S d A solution of this product gave no visible test for chloride ion with silver nitrate. Solutions of the im ure acid developed a pink color fairly rapidly when exposei to light. This sensitivity decreased markedly on further purification. Sam les of purified acid in solution in contact with air were l e 8 in the dark for 6 months and did not become colored but turned pink when left standing in daylight for several days. Solutions used in the measurements took a t least one week to become visibly tinted unless exposed to bright sunlight. The various purification operations were carried out in a semi-darkened room. Analysis for purity was made by titration with NaOH which had been standardized us. National Rurenu of Standards potnssium acid hthalatc. Weight buret techniques were ap licd and a p% meter was used to detect the endpoint wkch occurred at pH 7.5. Tho resulting values of purity varied from 99.72 to 99.94%, the value increasing with the drying time. A Sam le, oven dried at llOo, analyzed 100.01% pure. Each flesiccator-dried batch of recrystallized acid was analyzed separately; the difference between the analysis and 100% was assumed to be water. The most likely impurity neglecting moisture is sulfanilic acid. The problem of detecting this impurity in our product and estimating its amount has not been solved satisfactorily. In the ultraviolet absorption spectra of these acids, the molar extinction coefficient of sulfanilic acid is approximately tcn times that of orthanilic acid at 2550 A. We estimate that such measuremcnt8 would permit us to detect 1 or 2% of sulfanilic acid as an impurity but not lesser nmounts. We can only express confidence in the absence of sulfanilic acid to this degree. The cells, electrodes, measuring instruments and procedure of the measurements were the same as those described by MacLaren and Swinehart.*
Results The results are shown in Table I. Each potential difference is the average of 16 potential measurements obtained from two cells, run in duplicate, except where noted in the table by the footnote, "one cell only." These latter values are the averages of eight measurements from one cell only. The average deviations of the potentials from any one cell from the mean potential were about *0.02 mv. The mean potentiaIs of duplicate cells showed an average difference of 0.062 mv. and no difference was greater than 0.16 mv. The data were treated as indicated in the introduction. Representative extrapolations of the last approximation of the left-hand side of eq. 3 S. Iiarned nnd R . W. Ehlera. ibid., 66, 2179 (1933). (12) D. F. Swinehart, kbid., 74, 1100 (1952). (13) J. A. Bewderi and H. M . Watts, Phya. Reu., 81,73 (1951). (11) H.
March, 1957
IONIZATION CONSTANT OF ORTHANILIC ACID
and D , respectively. The calculated values of -log K using these constants together with the experimental values of -log K and the values
K . This quantity is calculated from the equation Standard deviation =
(n2fc) l"
~
E
-
p:
&g
18
2.63
335
R. NORMAN DIEBICL AND D. 17. SWINEHART
336
Vol. 61
rated charges of the zwitterions. Thus, when such
TABLE IT1
ENTROPIES AND HEAT CAPACITY CHANQES FOR THE IONIXA- a zwitterion separates into two independent ions, the change in amount of water bound is much TION OF SOME WEAKACIDSAT 25”
smaller than for neutral acids, resulting in a markedly smaller entropy of ionization for the zwitterion. Water -18.7 -46.6 Sulfanilio and metanilic acids are switterions Formic -17.6 -41.7 with such large charge separations that essentially Acetic -22.1 -36.5 as much water is bound by the zwitterions as for Propionic -22.8 -38.3 the separated ions, yielding a very small entropy efButyric -24.4 -36.4 fect on ionization. Chloroacetic -17.0 -40.0 For the case of orthanilic acid, the charge sepLactic -18.0 -40.7 aration must be markedly less than for its isomers. Glycolic -16.9 -38.8 Space considerations then limit the amount of water Phosphoric (1st H+) - 15.7 -50.7 bound by the zwitterions and results in a much Glycine - 6.9 -32.2 larger entropy effect than for the isomeric acids, dl-Alanine - 8.3 -34.2 ASo being, for this case, approximately ten times dl-Amino-n-butyric - 9.5 -34.1 larger than for the isomers, although not quite as dl-Leucine 9.4 -36.8 large as predicted.a This indicates that the charge dl-Valine - 9.0 -32.7 separation, i.e., the dipole moment, of orthanilic, 3.0 -110 Sulfumic although smaller than that of its isomers, is still 7.9 - 8.4 Taurine considerably larger than that of the aliphatic ami- 0.37 - 7.1 Sulfanilic no acids. Direct measurements of the dipole mo0.57 - 7.7 Metanilic ments of metanilic and orthanilic acids have not - 3.32 -18.6 Orthanilk been reported. It has been claimed that sulfamic acid exists in It is remarkable that the standard entropies of ionization of many uncharged acids are approxi- solution as the neutral molecule, i.e., not as the mately constant at about -20 e.u. and that the zwitterion, ” ~ 1 even though other evidence indiheat capacity changes for ionization hover about cates that it does exist in the crystalline state as -40 cal./degree/mole, However, when the cor- the zwitterion.2a The entropy of ionization of responding values for the classical zwitterions, sulfamic acid” shown in Table 111supports strongly i.e., the aliphatic amino acids, are considered, the the idea that it also exists in solution as the zwitterentropies of ionization decrease to about half the ion. The entropy and heat capacity change for values for the neutral acids. The heat capacit.y taurine agree with other evidencea8J4that this acid changes cleem to be lowered only slightly in magni- exists in solution as a zwitterion. In this picture, tude. The entropies of ionization of sulfanilic and only the heat capacity change for sulfamic acid apmetanilic acids are very small indeed as are the pears to be anomalously high. changes in heat capacity. King has reported difficulty in the interpretation These observations are interpreted following the of e.m.f. data for the estimation of the ionization ideas of Frank and Evans’g and of Powell and Lati- constant for sulfamic acid’’ due to its being a fairly mer.2o An entropy decrease occurs when neutral strong acid and suggests that if an acid with an acmoleciiles are dissolved in water. Part of this en- curately known ionization constant of about 1 X tropy loss is due to rest,riction of the molecule to a fiere available, a method involving a mixture small volume or “cage” of solvent molecules and of a salt of sulfamic acid and the second weak acid, part is diie to thc immobilizat,ionof water molecules as suggested by might be used to check his adjacent to the solute molecule. Now, if the mole- e.m.f. data. It is suggested that sulfanilic or orcule separates into two charged ions, many more thanilic acid might serve this purpose. water molecules are “frozen” and a large entropy Acknowledgment.-This research was supported loss occurs. This effect is so large that, to a first by a fellowship and grant from the Graduate Counapproximation, the entropy is the same for all neu- cil of the University of Oregon. tral, Le., uncharged, acids. (21) P. Baumgsrteo, Ber., 61B,820 (1929). A zwitterion, on the other hand, such as an amino (22) F. A. Kanda and A. J. King, J . A m . Chem. Soc.. 18, 2316 acid, binds many more water molecules due to in- (1961). t,eraction of the water dipoles with the widely sepaAcid
A SO
(cal./dee.)
ACop
(oal./dee.)
-
(28) 0.DeVoto, f3aw chim. ital., 61,897 (1931).
(ID) H. S.Frank and M. W. Evans, J . Chem. Phys., 13,607 (1946). (20).R. E. Powell and W. M. Latimer. Cbid.. 19, 1139 (1961).
(24) M. Freymann and P. Rumpf, Compl. rend.. Pol, 608 (19351. (26) R. a. Bates, J . Am. Cham. Boc., 73, 2259 (1861).
