GUIDE TO TITRANT SELECTION AND USE
On the basis of these studies some general rules can be formulated for selecting and using quaternary ammonium titrants. I n selecting the quaternary ammonium base, ethyl groups, phenyl groups, and to a lesser degree propyl groups, should be avoided, As many methyl groups as possible should be used without sacrificing solubility characteristics. The maximum is three for general purpose titrants where solubility requirements are important. The group added to increase solubility should be large (hexadecyl, for example). Although not verified by the present study a fourth rule, based on published work, can be added. The group added for solubility should have as many alkyl substituents on the beta carbon as possible. The above four rules are based on the influence of structure on solubility and stability. For the influence of cation structure on potentiometric and conductometric titration characteristics a previous publication (6) should be consulted. Titrant stability is not a problem if considerable water can be tolerated.
(4) Cundiff, R. H., Markunas, P. C., ANAL.CHEM.28,792 (1956). ( 5 ) Hanhart, W., Ingold, C. K., J. Chem. SOC.1927, 997. ( 6 ) Harlow, G. A., ANAL.CHEW34, 1482 (1962). ( 7 ) Hadow, G. A., Noble, C. M., Wyld, G. E. A., Ibid., 28, 787 (1956). (8) Harlow, G. A., Wyld, G. E. A., Ibid., 30, 69 (1958). (9) Ibid., p. 73. (10) Ibid., 34, 172 (1962). ( 1 1 ) Heiide. H. B. van der. Dahmen. ‘ E. A. M.’F., Anal. Chim. A’cta 16, 378 (1957). (12) Hummelstedt, L. E. I., Hume, D. X., ANAL.CHEM. 13, 1792 (1960). (13) Ingold, C. K., “Structure and Rlech-
(Water is most troublesome in the determination of very weak acids and in the analysis of acid mixtures containing dibasic acids m-here successful resolution depends on maximum K 1 / K 2 ratios (8). It is preferable to use a more concentrated alcoholic titrant in conjunction with a more precise buret rather than to dilute the titrant with inert or basic solvents. Xonaqueous quaternary ammonium titrants of high basicity should be stored in a refrigerator. (It is convenient in practice to refrigerate a stock bottle of titrant (-15” C.) and to transfer quickly a few days’ supply to a small reagent bottle for daily use a t room temperature.)
anism in Organic Chemistry,” University Press, Ithaca, N. Y., 1953. (14) Ingold, C. K., Vass, C. C. N.,J. Chem. SOC.1948, 2072. (15) Kreshkov, A. P., Bykova, L. N., Mkhitaryan, N. A., J . Anal. Chem. ( U S S R ) 14, 530 (1959). (16) Malmstadt, H. V., Vassallo, D. A., ANAL.CHEM.31,863 (1959). (17) Marple, L. W., Fritz, J. S., 140th Meeting, ACS, Chicago, September
ACKNOWLEDGMENT
The author is indebted to R. L. Johnson for his assistance in the early stages of this investigation.
1961
van, Dahmen, E. A. M. F., Anal. Chim. Acta 19, 64 (1958). (19) Patchornik, A., Rogozinski, S. E., ANAL.CHEM.31, 985 (1959). (20) Scatchard, G., J . Chem. Phys. 9, 34 (1941). (18^jMeurs, N.
LITERATURE CITED
(1) Banthorpe, D. V., Hughes, E. D., Ingold, C. K., J . Chem. SOC.1960, 4054. (2) Beringer, F. M., Gindler, E. M., J. Am. Chem. SOC.77, 3200 (1955). (3) Ibid., p. 3203.
RECEIVED for review July 2, 1962. Accepted August 10, 1962.
Study of the Titration of Selected Substituted Phenylureas as Acids in n-Butylamine M. L. CLUETT lndusfrial and Biochemicals Deparfmenf, Experimenfal Sfafion, E. 1. du Ponf de Nemours &
Co., Wilmingfon,
b Some structurally related substituted phenylureas were studied with respect to their behavior as acids in n-butylamine. Half-neutralization potentials (HNP) were used to establish an order of relative acidity, and the correlation between relative acidity and molecular structure is discussed. For comparison with the titration of weak acids in water, the potentiometric titration data were fitted to the form of the usual titration equations and the HNP values were correlated with the Hammett sigma ( g ) values for the chloro substituents on the phenyl group. The titrations were also followed photometrically and conductometrically. These data, together with the potentiometric titration and equivalent conductance data, were used to draw conclusions regarding the extent of reaction between the ureas and the solvent, the nature and extent of reaction between the ureas and the titrant, and the kinds of ionic species present in solution during
could be titrated as acids in nonaqueous media, and furthermore, if adequate differentiation could be obtained to determine the relative effect of various minor changes in the phenyl and/or alkyl substituents on the intrinsic acidity of these compounds. Preliminary work showed that the compounds behaved as acids in a variety of both neutral and basic solvents but that the optimal differentiation was obtained in nbutylamine. The intriguing ability of such a solvent t o enhance and distinguish the intrinsic acidities of these structurally similar compounds led to the more fundamental study of the physicochemical nature of the titration system.
titration. The titrant used was tetra-n-butylammonium mixed hydroxide methoxide in 9 : 1 benzene-methanol.
N
PUBLICATIONS on the applications of nonaqueous solvents in analytical chemistry-particularly in the field of titrations-continue to appear in the literature. Relatively few of these, however, go beyond the practical aspects of simple analysis. Unfortunately, until more effort is made to study the fundamental behavior of these systems, this technique will remain largely empirical, and the full potentialities of nonaqueous solvents in analytical chemistry in general will materialize slowly. This study had its inception in a n attempt to ascertain if certain substituted phenylureas of the general formula UMEROUS
X Y
= =
hydrogen or chloro hydrogen or alkyl
Del.
EXPERIMENTAL
Reagents and Apparatus.
n-Butylamine (Eastman White Label) was purified by distillation over solid potassium hydroxide, collected, and stored under dry, COz-free nitrogen. A 10% forecut and bottom were discarded. VOL. 34, NO. 1 1 , OCTOBER 1962
1491
Figure 1. Potentiometric curves in n-butylamine 1.
2. 3. 4.
5.
6.
titration
3-Phenyl-1-methyl urea 3-Phenyl-1-methyl-1-ethyl urea 3-(p-Chlorophenyl)-l,l -dimethyl urea 3-(3,4-DichlorophenyI)-1,1 -dimethyl urea 3-(3,4-DichlorophenyI)-l -methyl-1 -hydroxyethyl urea Benzoic acid
Tetra-n-butylammonium mixed hydroxide methoxide titrant, 0.1N was prepared in 9: 1 benzene-methanol by the method of Cundiff and Markunas (3). The composition of this titrant was described by Cluett (2). The benzothiazole-2-carbamic acid methylester and most of the substituted ureas used in this study were research samples prepared in this laboratory. Otherwise, commercial reagent grade chemicals were used. Potentiometric Titrations. Illustrative potentiometric titration curves for several of the substituted ureas and for benzoic acid are shown in Figure 1. All titrations were carried out in 30 ml. of purified n-butylamine in a 50-ml. beaker fitted with a rubber gasket cover. One-fourth millimole of each compound was titrated and a dry, Con-free nitrogen blanket was maintained over the solvent throughout
the titration. The titrant was added incrementally from a 5-ml. semimicro buret, and the titrations were followed with a glass and saturated KC1-methanol modified calomel electrode system. An L&N p H indicator was used to measure potentials. Because each titration required practically the same volume of titrant, conditions were standardized adequately t o permit direct comparison of the titration data. A maximum deviation of =t5 mv. was obtained for several determinations of the HXP of benzoic acid. The potentiometric titration of perchloric acid indicated that the maximum potential range in n-butylamine, using tetra-n-butylammonium mixed base in 9: 1benzene-methanol as titrant, is about 525 mv. This value agrees with that reported by Heijde and Dahmen (6). Conductance Measurements. Conductometric titration curves for several compounds are shown in Figure 2 . These data were obtained by titrating approximately 0.25 millimole of each compound in 30 ml. of purified nbutylamine with the tetra-n-butylammonium mixed base titrant. h dry, COz-free nitrogen blanket w s maintained over the solvent throughout the titrations. A platinized platinum dip-type conductivity electrode with a cell constant of 0.11 was used in combination with an Industrial Instruments Conductivity Bridge, Model RC-16B. to follow the titrations. A blank titration curve is also shown. I n Figure 3 are plotted the equivalent conductances of the titration products of two of the substituted ureas us. concentration in n-butylamine. These data were obtained in the following manner: One-half mmole of the compound was dissolved in 25 ml. of purified n-butylamine and neutralized with the exact amount of tetra-n-butylammonium mixed base titrant. The solution was evaporated a t room temperature
,
36-
I
1
Figure 2. Conductometric curves in n-butylamine 1. 2. 3.
4. 5.
3-(3,4-DichIorophenyI)-l,l -dimethyl urea 3-Phenyl-1-methyl urea Benzothiazale-2-carbamic acid methyl ester
HCl Blank
to dryness under a stream of dry, Confree nitrogen, and the residue of colorless crystals was dissolved in 50 ml. of purified n-butylamine to make a O.OlJ.1 solution. The resistance of the solution mas taken a t numerous dilutions with the conductivity apparatus described above, and the corresponding equivalent conductances were calculated. The solution was kept continually under a blanket of purified nitrogen. The dilution limit %as determined by the range of the conductivity bridge. Spectrophotometric Measurements. The ultraviolet absorption spectra for several of the ureas, including compound 15, were determined in water, n-butylamine, and in n-butylamine with a slight excess of the tetra-n-butylammonium base titrant. I n the latter cases, the titrant used was prepared in a 9: 1 isopropyl alcoholmethanol solution to avoid the strong absorption of benzene in the region below 270 mp, In water and n-butylamine the absorption bands viere esI
I
I
1
34 -
Table I. Comparison of Spectrophotometric and Potentiometric End Points
Photometric Potentioe.p. metric (extrape.p. olated (inflection), intercept), ml . mi. 3-Phenyl-lmethvlureaa 3-Phenil-I, 1dimethylureaa
1.40b
0,60
1. 4 4 b
3( 3,4-Dichloro-
phenyl )-1methyl-l-nbutylurea 3( 3,4-Dichlorophenyl)-1,ldimethylurea
/x
--1t
/x
0
on
8-
z --
/X'
9
22 18
0
x / h
16
1 . G5
1.50
2.00
1.91
3( 3,4-Dichloro-
phenyl)-1methyl-l-hydroxyethylurea 1.15 1.10 a Equivalence point, about. 1.0 ml * Estimated.
1492
//j
t
32
)A
0.58
ANALYTICAL CHEMISTRY
titration
0
1.9
I
I
I
2.3
2.7
3.1
J% meq / liter Figure 3. Equivalent conductance of titration products of substituted ureas vs. concentration 1.
2.
3 [3,4-(Dichlorophenyl)] -1,l -dimethyl urea 3- [p-(Chlorophenyl)]-1,1 -dimethyl ureo
sentially identical, exhibiting a wavelength maximum at about 255 mp. In the presence of excess titrant, however, the bands shifted t o a wavelength maximum of about 300 n~w. In all cases, the absorption bands were symmetrical. In Table I data are given which permit a direct comparison between the potentiometric and spectrophotometric end points for a group of substituted ureas ranging from the least to the most acidic ureas used in this study. These data were obtained simultaneously using the apparatus shown schematically in Figure 4. The apparatus was designed to fit into the cell holder for the Beckman DU spectrophotometer. The compartment wall on one side of the cell holder was removed t o make this possible. The cell holder compartment itself was made lightproof by painting the reservoir and side arms with black paint and by using a fitted cardboard cover. A quartz insert was employed to obtain a solution cell path of 0.5 mm. DISCUSSION
Relative Acidity and Molecular Structure. T h e substituted ureas included in this study are listed in Table I1 (Nos. 1 through 15) together with their H N P . The compounds are arranged in order of increasing acidityi.e., the lower the HNP, the greater the acidity. Because all of the extrinsic factors-volume of solvent, size of titration, etc.-which might influence the H X P were reproduced preciselj- in all.of the titrations, it is assumed that the H N P are indicative of the relative intrinsic acidities of the substituted urea molecules. As discussed later in this paper, the titration of these substituted ureas as acids in n-butylamine results in the removal of the proton on the exo-ring nitrogen to form the tetra n-butylammonium salt of the corresponding conjugate base. Thus it is assumed that the differences in acidity are in effect a measure of the ease of proton removal which, of course, is largely dependent upon the electronic influences of the various constituent groupings in the molecule. Neither urea nor any aliphatic substituted urea can be titrated as an acid in n-butylamine; thus the role of the phenyl substitution on the acid strengths of these compounds is apparent. This increase in acidity caused by the phenyl group is undoubtedly due to the withdrawal of the free electrons on the exoring nitrogen into the ring. Wimer (10) used the same explanation for the lower basicity of N-phenyl substituted amides relative t o N-alkyl or unsubstituted amides. As indicated by the data in Table 11, a further increase in acidity, probably due to the inductive (-1,) effect of the electronegative chlorine, is observed in going from the
acidity established is CH3,H < CHs, n-C4Hs < CH3,C2H5 < CH3, CH3 < CHI, C2H40H. Probably the most perplexing feature of this series is the relative positions of the CH3,H and CH3, CH3 groupings. On the basis of inductive (+I*) effect alone, a decrease in acidity would be expected in going from the monoalkyl, hydrogen compounds t o the dialkyl compounds. Such behavior would have been consistent with the apparent increase in base strength of N-alkyl substituted amides compared with the free amides as observed by \Timer (10). (It is noted, however, that the slopes of curves 1 and 2 in Figures 2 and 3 do not Seem to agree with the maximum u / A V values reported in Table I by Wimer.) On the other hand. the results obtained in this study appear consistent with a somewhat analogous system discussed by Pnuling (9)-i.e., the effect of N-alkyl and N,N-dialkyl substitutions on the base strength of guanidine. In this case, the greater electronegativity of carbon over hydrogen explains a restriction on one of the resonance forms, m hich would account for the observed decrease in base strength. In the case of the substituted ureas, therefore, the greater tendency of the alkyl group t o cause the terminal nitrogen not to assume a positive charge would restrict the contribution of one of the three resonance forms of the urea group:
7 ATTAC t i PROPI P E T " RESERVOIR
-
CELL
I c m O U A R T Z CELL WITH INSERT ~~
Figure 4. Apparatus for simultaneous potentiometric and spectrophotometric titrations
phenyl t o the p-chlorophenyl and then t o the 3,4-dichlorophenyl substitutions. As might be expected, the effect on the acidity of these compounds by the various alkyl substituents on the terminal nitrogen is not so pronounced as the effect of the phenyl and chlorophenyl substituents. Actually, such a consistent differentiation of the small differences between these substituents was surprising. Considering three groups of compounds-namely, the phenyl group (compound XOS.1-4, 7 ) ; the p-chlorophenyl group (Nos. 5, 6, 8, 9, 11); and the 3,4-dichlorophenyl group ( S o s . 10, 12-15)-the general order of
Table II.
Half-Neutralization Potentials vs. Molecular Structure
Compound structure
Designation
x
1 2 3
H H H H p-c1 p-c1 H p-c1 p-c1 3,4-DiC1 p-c1 3,4-DiC1 3.4-DiC1 314-DiCl 3,4-DiCl
4 5
6
7 8 9 10 11 12 13
14 15
Relative acidity" (halfneutralization potentials),
Y
A
I'
mv.
870 847
CH3' CiHs H CiHdOH CdH Q CiHs CHs CzH,OH
745
655 613 500
17 18 19 a
480
HC1 HClO4
420
375
The lower the HSP, the stronger the acid.
VOL. 34, NO. 1 1, OCTOBER 1962
0
1493
\
thus decreasing the over-all electron density near the center of the molecule via this mechanism. This effect would then be observed as an increase in acidity. The acidities of the n-butyl, ethyl, and hydroxyethyl groups relative to the methyl group are in the order expected on the basis of the inductive effect alone. In view of the low dielectric strength of the solvent, field effects might also operate in these instances along with the inductive effect. Conformity to Titration Equations. The potentiometric titration data in n-butylamine followed the usual titration equations for acids in water. I n Figure 5, the data for several acids are plotted against log x/( 1 - 2) and/or log 1/(1 - z), where z is the fraction of acid neutralized on the basis of the amount of titrant added. The curve designations correspond to the compounds in Table 11. In the case of weak acids, the theoretical relationship between electrometrically measured acidity and the logarithm of the ratio of the concentration of the conjugate base of the acid and the free acid-i.e., log x/(l - x)-is linear. As shown in Figure 5, this relationship was found to hold for the more acidic ureas (Nos. 8, 14, and 15). The slopes
y(:
9
650
600
5501
A . l 400
-10 08 06 04 0 2 0 0 2 0 4 06 08 + I O LW
Figure 5. Conformity of potentiometric titration data to titration equations
1494
ANALYTICAL CHEMISTRY
yo] 550
,
,
,; 0
,:5"
,
, ,
,
,
,
,
,
,,
[1
-NICH3UCH3l -NICH311C2H10Hl
500
01
02
03
04
05
06
0
Figure 6. Half-neutralization potentials vs. Hammett values
of these lines are 68, 50, and 52 mv., respectively, compared to the theoretical slope of 59 mv. for weak acids in water. The difference between the ultraviolet absorption spectra for compound 15 in water and n-butylamine and its conjugate base formed in the presence of excess titrant also shows t h a t even the most acidic urea is a weak acid-Le., not appreciably ionized-in n-butylamine. In the case of strong acids in water, the theoretical relationship between electrometrically measured acidity and the logarithm of the reciprocal of the acid concentration-i.e., log 1/(1 - z) is linear. Perchloric acid was titrated in n-butylamine, and the data fitted precisely this relationship for a completely ionized acid. The titration data for hydrogen chloride, however, (Figure 5) do not show a linear relationship in either the log z/(l - z) or log 1/(1 - z) plots. This is precisely the theoretical behavior of appreciably, but not completely, ionized acids in water. The data for compound 15 also are plotted on the log 1/(1 - r ) scale for purposes of comparison with the hydrogen chloride data on the same scale. Although not clearly discernible in Figure 5, the original plot of the benzoic acid data showed a slight curvature on the log z/(1 - z) scale. In any case, it appears t h a t benzoic acid is a t most only slightly ionized in n-butylamine. For a direct comparison i t should be noted that the data for compound 17 show a significant curvature on the log z/(l - 2) scale, indicating that this acid is appreciably ionized in n-butylamine. This conclusion is supported fully by the large shift (to longer wavelengths) and changes in shape and intensity of the ultraviolet spectrum of this compound between methanol and n-butylamine. Only a small shift and very slight differences in shape and intensity were observed between the spectra of the compound alone in n-butylamine and the compound plus excess titrant in nbutylamine.
The curve for compound 1 in Figure 5 shows a steadily decreasing positive slope which can be explained on the basis that this extremely weak acid does not react quantitatively with the titrant added. The higher potentials for the corresponding apparent fractions neutralized are caused by the presence of unreacted base. The potentiometric, conductometric, and spectrophotometric titration data for this compound in Figures 1 and 2 and in Table I support this conclusion. Consequently, this acid is too weakly acidic to be titrated quantitatively in n-butylamine. HNP us. Hammett Sigma (u)Values (5). Although no attempt was made to determine pK, for the substituted ureas, i t is reasonable to assume that in water they would exhibit an extremely low and probably immeasurable acid strength. For instance, spectrophotometric measurements on compound NO. 14 in Table I1 showed no shift in ultraviolet absorption in aqueous sodium hydroxide up to 0.5N. However, in 4.2N base the absorption a t of 248 mp decreased about 30% the A, and a new peak formed a t 275 mp, indicating that a t least partial ionization occurs under these conditions. Therefore, because the pK,-HNP relationship could not be determined for the ureas it was expected that should such a relationship be linear, then a linear relationship would be found between H N P and the Hammett U-values for the chloro substituents on the phenyl group. These data are plotted in Figure 6. The expected linear relationship holds well, particularly for the more acidic compounds. The u value for the p-chloro substituent has been reported (8) whereas the Uvalue for the 3,4dichloro substituent was calculated from the pK, for benzoic acid (8) and the pK. for 3,4dichlorobenzoic acid reported by Davis and Hetzer (4). While it may be obvious, nevertheless i t is not frequently pointed out that the relationship between H N P and uvalues is evidence that the glass electrode actually is measuring the activity of Hf in the n-butylamine solutions; otherwise linearity could not hold. 800
6
Y
BENZOIC A C I D
A
I
1 1
,
,
2
4
, 6 D K ~
I
j
,
,
8
1 0 1 2
Figure 7. HNP in n-butylamine vs. pK. for some weak organic acids
Figure 7 shows that the pK, of the ureas cannot be estimated from H N P in n-butylamine on the basis of the HNP-pK, relationship found for other weak acids. Although a linear relationship was found for the structurally similar phenolic acids, there was no correlation of these data with the data determined for the benzoic acids. That the HNP-pK, relationship is greatly dependent upon structure similarities has been reported for other nonaqueous solvents and is true also for n-butylamine. Ionic Species in Solutions. As a result of the finding t h a t the titration of the ureas could be followed conductometrically, the question arose as t o what ionic species (reaction products) are present in solution during the titrations. Figure 3 shows the manner in which the equivalent conductance of the titration products changes with concentration. The property of a minimum followed by a n increase of conductance toward higher concentrations seems to be a characteristic of solutions of electrolytes in low dielectric media (l)-i.e., below D = 20. The dielectric constant for nbutylamine is 5.3. As noted previously, the shift of the ultraviolet spectrum on titration to a single peak a t considerably longer wavelengths also is evidence of ion formation. Still more evidence was obtained by infrared measurements which were carked out by adding the stoichiometric amount of titrant to an n-butylamine solution of the urea, evaporating to dryness a t room temperature under nitrogen, and scanning the residue promptly in carbon tetrachloride. I n the case of the more acidic ureas, the
spectra showed the absence of absorption a t the N-H and normal C=O stretching frequencies. I n the case of the weakly acidic ureas, N-H, 0-H, and normal C=O stretching bands were observed, indicating that the neutralization was not quantitative. It is concluded, therefore, that the reaction between the ureas and the titrant results in the formation of an electrolyte, the constituent ions of which would be, e.g.
difference between the conductance of the n-butylammonium salts and the tetra-n-butylammonium salts is probably a function of the size and structure of the positive ions. Kraus (7) reported that the considerably lower dissociation of partially substituted ammonium salts containing a hydrogen on the nitrogen compared to the dissociation of fully substituted ammonium salts is due to a strong interaction (other than coulombic) between the proton attached to the nitrogen and the negative ion. ACKNOWLEDGMENT
However, as a result of coulombic interaction, extensive in low dielectric media, these ions apparently exist as ion pairs with decreasing dissociation as the concentration increases up to about O.O0144M, where the equivalent conductance curves show a minimum. At higher concentrations, ions of a more complex type-triple ions, etc.apparently form to cause an equivalent conductance increase due to the greater transport values of the multiple charged complex ions. This phenomenon might also account for the increasing slopes of the conductometric titration curves in Figure 2. In Figure 2, even though the benzothiazole-%carbamic acid methyl ester and HC1 ionize appreciably in nbutylamine, the conductance of the starting solutions was no different from that of the ureas. Therefore it was possible to follow the titrations of these acids conductometrically. The marked
The author thanks R. L. Dalton and J. T. Funkhouser for their helpful suggestions in the preparation of this manuscript. LITERATURE CITED
(1) Ann. N. Y. Acad. Sci. 5 1 , (4), 789
(1949).
(2) Cluett, M. L., ANAL.CHEW 31, 610 (1959). (3j Cundiff, R. H., Markunas, P. C., Ibid., 28, 792 (1956).
(4)Davis, M. M., Hetzer, H. B., J. Phys. Chem. 61, 123 (1957). ( 5 ) Hammett, L. P., “Physical Organic Chemistry,” 1st ed., p. 186 McGrawHill, New York (1940). (6) Heijde, van der, H. B., Dahmen, E. A. M. F., Anal. Chim. Acta 16, 378 (1957). (7) Kraus, C. A., Science 90, 281 (1939). ( 8 ) McDaniel. D. H.. Brown. H. C.. J . Ora. Chem. 23. 420 (1958). ’ (9) Pauling, L.,’“Na&e of the Chemical Bond,” 2nd ed., p. 213. Cornell University Press, Ithaca, N. Y. (1948). (10) Wimer, D. C., ANAL. CHEM.30, 77 (1958). ~
RECEIVEDfor review May 1, 1962. Accepted August 10, 1962.
*
A
I
Relationship between pK,(H,O) of Organic Lompounds and €1/2 Values in Several Nonaqueous Solvents L. G. CHATTEN’ and LOYD E. HARRIS College o f Pharmacy, The Ohio Sfate University, Columbus 70, Ohio
b The relationship between the dissociation constant, pKb (HzO), and the half neutralization potential has been determined for a number of phenothiazines and basic amines in five organic solvents. The presence of electrophilic groups appears to have a marked effect on the behavior of the phenothiazine compounds studied. In addition, each phenothiazine compound produced a distinctly colored solution in nitromethane and a possible explanation for this is offered. The corresponding investigation for the amines indicated that linearity between pKb(H20) and EIlz values depended largely upon similarities of structure.
Steric hindrance does not appear to affect this relationship, but in certain instances hydrogen bonding is worthy of consideration.
A
LTHOUGH NONAQUEOUS TITRIMETRY
is a well established technique, the choice of solvent systems for each compound to be titrated has generally been empirical. Solubility, of course, is a prime factor, but experience has usually been the guide. This situation led Riddick ( S I ) to state that “The information on the strength of acids and bases in organic solvents is meager. Present knowledge indicates that the strength will hrtve to be determined for
each acid and base for each amphiprotic solvent used.” I n 1953, Fritz (10) compared the pKb values of a few common laboratory reagents in water with their half neutralization potentials in acetonitrile. Lemaire and Lucas (IS) determined the pK. of three indicators and then utilized this information to obtain the strength of weak bases in glacial acetic acid. The behavior of bases has been studied more intensely in glacial acetic
* Present Address: Faculty of Pharmacy, University of Alberta, Edmonton, Alberta, Canada. VOL 34, NO. 11, OCTOBER 1962
1495
