[CONTRIBUTION FROM
THE
CHEMICAL LABORATORY OF THE UNIVERSITY OF BERKELEY, CALIFORNIA]
CALIFORNIA,
THE DISSOCIATION OF HYDROGEN IONS FROM THE SULFATES OF AMINOPHENYLBORIC ACIDS CHARLES GERALD CLEAR*
AND
GERALD E. K. BRAXCH
Received November 18, 193'7; revised March 8, 1938 INTRODUCTION
The nomenclature of some derivatives of boric acid is anomalous. Phenylboric acid is CaH5B(OH)z,not C6H5OB(OH)z. After this anomaly normal nomenclature is used, and aminophenylboric acid is NHzCeHaB(OH)z. We shall call B(OH)z the boric acid group. The relationships between the dissociation constants of the aminophenylboric acids have sufficient theoretical interest to warrant a comparison of their values amongst themselves and with the dissociation constants of phenylboric acid and aniline. Since all the factors that govern acidic and basic strength are not known, these comparisons must be made with considerable caution. We shall restrict our comparisons to those between acids and bases of similar types, in the same solvent, and at one temperature, and our discussion to what we believe to be the more important factors. We measured the dissociation constants by a method equivalent to electrometric titration of the sulfate, which we shall treat as a dibasic acid. The pH numbers of solutions containing weighed quantities of the sulfates, neutralized to varying degrees with sodium hydroxide, were measured. The concentrations of the solutions varied between 0.02 N and 0.005 N . From these measurements the first and second dissociation constants of the positive ions were calculated from the equations,
* From a thesis in partial fulfillment of requirements for the degree of doctor of philosophy. 522
DISSOCIATION OF HYDROGEN IONS
523
and
M is the molality of aminophenylboric acid in all its forms. It was obtained from the weight concentration and equivalent weight of the sulfate, the latter being determined from the titration curve of the first acidic hydrogen atom. As the first and second dissociation constants of aminophenylboric acid sulfates are widely separated the above equations and method of determining the equivalent weights are justifiable. When necessary the equation for K1was corrected for the concentration of bifor the second dissociation constant of sulsulfate ion, using 1.2 X furic acid. Only measurements near the half neutralization points of the first and second dissociations were used for calculating K1and Kz. The K's obtained in this way involve the activity of hydrogen ion and the concentrations of the other substances. They are, therefore, neither classical nor thermodynamic dissociation constants. But when determined for identical conditions they measure the relative acid strengths. I n the range of concentrations and degrees of neutralization used, the values for aniline, phenylboric acid, and aminophenylboric acids show no trend great enough to be apparent above the erratic errors, so the values of K measure the relative strengths of the acids under practically identical conditions, viz. in aqueous solution about 0.01 N with respect to sodium sulfate. To check accuracy, a determination of the dissociation constant of anilinium ion was made, starting with aniline sulfate. The value of 2.1 )< was obtained, which agrees well with the accepted value k,/kb = 2.2 x 10-5. As acid solutions of the aminophenylboric acids are reduced by hydrogen in the presence of platinum, a hydrogen electrode could not be used in the acid range. Measurements in this range were made with a glass electrode, the hydrogen electrode being used for alkaline solutions. Over the short alkaline range in which both electrodes could be used, the two gave concordant results. The benzoylaminophenylboric acids are too insoluble in water for aqueous solutions to be used for determinations of acidic strength. The dissociation constants for these compounds in an alcohol-water mixture containing 25 per cent. alcohol by volume were determined, with a hydrogen electrode. This solvent was used for the determination of the
524
C.
a. CLEAR
AND G. E. K. BRANCH
dissociation constants of many boric acid derivatives by Branch, Yabroff and Bettman.l The equivalent weights were taken as those determined by titration to the phenolphthalein end-point in the presence of large amounts of mannitol. The equivalent weights of the sulfates of ortho and meta aminophenylboric acids agreed well with the formula, [CaH4B(OH)zNH3]ft SO4--, that of the para isomer with [CaHd3(OH),NHl]+zS04-2H20, those of ortho and meta benzoylaminophenylboric acids with [CeH6CONHGH4B(OH)z]z - HzO, and that of the para with CaH&ONHCsH4B(OH)2.*
+
RESULTS AND DISCUSSION
The results are summarized in Tables I and 11. For the amine sulfates, the products of the two dissociation constants are listed in column 4. APk = (logloKO - logl&) is a useful quantity for comparing t,he effect of groups on acid strength. An acid strengthening or negative group gives a negative APk value. For the APk of a first dissociation constant, KO is taken as that of aniline sulfate, for a second dissociation constant KO = K , of phenylboric acid in water,’ and KOfor the products is the product of the dissociation constants of aniline sulfate and phenylboric acid. For the benzoyl derivatives KO is taken as the dissociation constant of phenylboric acid in 25 per cent. alcohol.1 These APk’s are included in the tables. Zwitter ions. For an amino acid the relationships between the various equilibria are given by the series of equations:
in which K E is the equilibrium constant of the reaction, true amino acid to zwitter ion, Kl and Kz are the measured dissociation constants, KIA and K1z are the two dissociation constants of the positive ion forming the true amino acid and the zwitter ion respectively, K ~ is A the dissociation constant of the true amino acid and K2z that of the zwitter ion. These constants can be obtained by assuming KIA to be equal to the hydrolysis constant of the ester of the positive ion3 or K1z to the hydrolysis constant of the positive ion of the 1 BRANCH, YABROFF, AND BETTMAN, J . Chem. SOC.,66,937 (1934). *The formulas for the ortho and meta compounds agree with those found by SEAMAN AND JOHNSON,J . Am. Chem. SOC.,63, 711 (1931), who did not prepare the para isomers. a EBERT, 2.physik. Chem., 121,385 (1926). 4 MIYAMOTO AND SCHMIDT, Univ. California Pub. Physiol., 8, 1 (1932).
525
DISSOCIATION OF HYDROGEN IONS
as Log & . ~ [ ~ - N H ~ ~ E H ~ B ( approximately OE)~ equals 1% &CIHIB(OH):I the introduction of a meta amino group does not greatly change the strength of an acid.
+
+
+
Log KZ[~-NH,C~H,B(OTH)~ =P log K~(c,N~NH,),
P being the effect of the borate ion group. As the negativity of the boric acid group is practically zero, P is approximately the effect of the work of TABLE I DI~SOCIATION CONSTANTS OF THE SULFATES OF AYINOPHENYLBORIC ACIDIN WATER AT 25" KY
BUBBTANCB
x
KIKSX
1010
1014
-Aniline sulfate . . . . . . . . . . . . . . . . . . . . . 2 . 1
]
2.9 Phenylboric acid... . . . . . . . . . . . . . . . . . -' 13.7 Sulfate of o-aminophenylboricacid.. . 2 . 7 5.9 1.6 Sulfate of m-aminophenylboric acid. . 3 . 5 15.4 5 . 3 Sulfate of p-aminophenylboric acid.. , 19.5 6 . 7 13.1
-0.22 -0.97
1
-0.05 0.31
I
-0.27 -0.66
TABLE I1 DISSOCIATION CONSTANTS OF THE BENZOYLAMINOPHENYLBORIC ACIDSIN 25 PERCENT. ALCOHOL AT 25"
Phenylboric acid.. ...................................... o-Bensoylaminophenylboric acid. ........................ m-Benzoylaminophenylboricacid. . . . . . . . . . . . . . . . . . . . . . . . p-Benaoylaminophenylboric acid.. .......................
1.97 4.47 5.84 2.9
0
-0.36 -0.47 -0.17
completely separating opposite charges situated with respect to each other
+
-
as in m - NHsCsH4B(OzH). Hence and itpproximately equals
526
C. 0. CLEAR AND 0. E. K. BRANCH
P' being the effect on log KA(m-&H,CsH4CO&H,)of changing the group C02CH3to GOz-. P' is approximately the effect of the work of completely separating opposite charges situated with respect to each other as in
+
- NH3CeH&02-. Obviously, P and P' are approximately equal, and within at least a factor of ten
m
can be calculated from the dissociation constants of m-aminobenzoic acid and methyl m-amin~benzoate,~assuming KIA = Ka(m-6HLkHCOzCHa). Then
KE(m--"zCsHdCOOH)
KE(m-NHzCeH4COOH) =
(9
x io-*) - (2.8 x 2.8 x 104
= 2.2,
and KE (m -NHz C 6 Hd B (OH)2 =
2.2 X 1.4 X lo-' X 4.6 X lo-'' X 3.6 X lo-" 6.6 X
=
Similarly for o- and p-aminophenylboric acids KE must be much smaller than the small values of KE for o- and p-aminobenzoic acids. The true amino acids must be in great preponderance over the zwitter ions in aqueous solutions of the three aminophenylboric acids. The measured values K1 and K2 may be taken as, within experimental error, equal to KIA and K ~ A respectively. In an amino acid the product K1.K2 = KIA.K~A = Kl~eKzz. The difference between K1 KZ and the product of the hydrolysis constant of the parent amine and the dissociation constant of the parent acid (APklk,) is independent of zwitter-ion formation, and of any interaction of the groups in the neutral molecule. It measures the combined effects of the dipole of the acidic group on the charge in the positive ion, of the dipole of the amino group on the charge of the negative ion, of the interaction of the acidic group with the aminium-ion group in the positive ion, and of the interaction of the amino group with the ionized acidic group in the negative ion. The APktk2 of an aminocarboxylic acid should have a much more negative value than that of the corresponding aminoboric acid, as the carboxyl group is much more negative than the boric acid group (-B(OH)z). Thus for m-aminobenzoic acid, if C ~ m m i n g s 'values ~ for its dissociation 0
5
CUMMINGS, Proc. Roy. SOC., 79,82 (1906).
DISSOCIATION O F HYDROGEN IONS
527
constants and the generally accepted values for those of benzoic acid and aniline are used, APklk, = -1.01, while for maminophenylboric acid, APklkn = -0.27. The APk,kr values for the aminophenylboric acids are peculiar, in that the sign for the ortho compound is opposite to that for the other two isomers, and the numerical value of APklk2for the para compound is greater than that for the meta isomer in spite of the greater separation of the groups. The significance of these facts is discussed later. Resonance.-Branch, Yabroff, and Bettman' found that ortho and para ethoxyphenylboric acids are weaker than phenylboric acid while the meta isomer is stronger. They attributed these results to the opposing actions of the negativity of the ethoxy group and its resonance interaction with .the boric acid group. The resonance involved is indicated by the formulas :
From the separation of charges in the quinoidal form it is obvious that this resonance is stronger in the undissociated acid than in its ion, and therefore it tends to reduce the strengths of the acids. This type of resonance is weak in the meta acid for the quinoidal form
6CSI6 I/
has a, high energy level, one of the bonds being between somewhat widely separated atoms. The meta acid is therefore stronger than either its ortho or para isomer. Owing to the weakness of the resonance and the negativity of the ethoxy group the meta acid is stronger than phenylboric acid. This type of resonance is also possible in aminophenylboric acid, and similarly must be greatly reduced in the negative ions and the meta isomer. It cannot occur in the positive ions, the four bonds of the nitrogen atom being already satisfied. It therefore tends to make the first dissociation constants greater than that of anilinium ion, the second dissociation constants less than that of phenylboric acid, the first dissociation constant of the sulfates of 0- and p-aminophenylboric acids greater than, and their second dissociation constants less than, the corresponding constants of the meta isomer. These expected resonance effects are very apparent in both dissociation constants of the sulfate of paminophenylboric acid. Its first dissocia-
52s
C. G. CLEAR AND G. E. IC. BRANCH
tion constant is much greater than that of its meta isomer, and than the Ka of anilinium ion, while its second dissociation is much smaller than that of its meta isomer and than the K , of phenylboric acid. Similarly KZof the sulfate of o-aminophenylboric acid is less than Kz of the meta isomer and than the Ka of phenylboric acid. On the other hand, hand, Kl of the sulfate of o-aminophenylboric acid is anomalous in being smaller than that of its meta isomer. Since this type of resonance is reduced rather than absent in the negative ion, but is impossible in the positive ion, it favors the negative ion over the positive ion. Hence it tends to make KlK2 for the ortho and para compounds greater than KlK2 for the meta compound and than Ka(anilinium, X Ka(phenylbori0 acid). This effect is observable in the para compound, but again the ortho isomer is anomalous, its APklkz being positive. Chelalion.-The low values of K I and KIKz for the sulfate of o-aminophenylboric acid show that there is some factor that stabilizes the positive ion with respect to the neutral molecule and the negative ion. Many anomalies have been observed in the properties of ortho derivatives of benzene. Some of these have been attributed to the formation of chelate rings by an atom bridging the narrow gap between ortho groups. For example lack of hydroxyl infra-red bands in o-hydroxycarbonylbenzenes;'j the anomalously high dissociation of salicylic acid;' and the anomalously low dissociation of o-nitrophenylboric acid.s Chelation in the positive ion of o-aminophenylboric acid is a probable reason for its stability. I n the only chelation possible the proton of the nminium ion binds the nitrogen to an oxygen atom. This implies the
+
+
resonance, (H~NCOH~BO~HZ, H2NC6H4B02H3). To both this theory and that of chelation of o-hydroxycarbonylbenzenes, the objection may be raised that the groups accepting the proton, -B(OH)2 and -COR, respectively, are not sufficiently basic. However, the ability of the boric acid group to function in this way is greatly enhanced by the resonance
+
-+
that arises from the existence of a quinoidal form H2N = C6H4 = BO2H3,
+
as the corresponding form is impossible in H3NC&B02H2. Similarly the ability of the carbonyl group to accept the proton from an o-hydroxyl group is enhanced by the resonance arising from the quinoidal form 0 = C6H4= C(OH)R, as this is much more stable than the corresponding
+
-
form HO = CaH4 = C(O)R, in which the proton is on the hydroxyl group. 8
HILBERT, WULF,HENDRICKS, AND LIDDEL, Nature, 136, 147 (1935);J . Am. Chem.
Soc., 68,548 (1936). 7 8
BRANCH AND YABROFF, J. Am. Chem. SOC.,66,2568 (1934). BITTYAN,BRANCH, AND YABROFF, ibid., 66,1865 (1934).
DISSOCIATION OF HYDROGEN IONS
529
If this chelation is the correct explanation of the low value of K 1 in the ortho compound, it must be more energetic than the chelation of the neutral molecule which would result from the resonance between amino acid and zwitter ion forms. In an o-amino acid, such as this, the distance between the groups is such that one would expect the equilibrium between amino acid and zwitter ion to degenerate to resonance, and a chelate ring to form. This type of chelation should be stronger in o-aminobenzoic acid than in o-aminophenylboric acid, as benzoic acid is stronger than phenylboric acid. Consequently it is not surprising to find that K , for o-aminobenzoic acid is not anomalously low like that for o-aminophenylboric acid, the chelation of the positive ion being compensated by that of the neutral molecule. Negativity.-We shall define the negativity of a substituent group as the property of the group whereby it affects the dissociation of a proton from mother part of the molecule, when the arrangement of the charges in the group are not displaced by resonance interactions with other parts of the molecule. It is evident that this proviso is an idealization that is never rigorously achieved, and hence evidence concerning the negativity is always to some extent incomplete. The negativity of a group arises from the charges and electric dipoles inherent to the group. Since there is always some interaction between a group and the solvent, the negativity of a group strictly refers to the group in a particular solvent. The hydrogen atom is a convenient arbitrary zero of negativity. The effect of a group on the strength of an aromatic acid most nearly measures negativity of the group, when it is in the meta position, for this position is the least favorable to resonance. For the sulfate of m-aminophenylboric acid, APk, = -0.22. The boric acid group is acting as if it were a definitely negative group. But Bettman, Branch and Yabroff found APk for the boric acid group in m-carboxyphenylboric acid = $0.01. This discrepancy is too great for experimental error. As the resonance interaction of a boric acid with a carboxyl group is much less than with an amino group, the true negativity of the boric acid group is more accurately obtained from the dissociation of the m-carboxyphenylboric acid than from that of the sulfate of m-aminophenylboric acid, and is probably very nearly zero. The apparent negativity shown by the group in m-aminophenylboric acid is presumably due to the resonance, "2
&HI /I
which makes the boric acid group appear to be negative.
530
C. G . CLEAR A N D G . E. K , BRANCH
The APk, of maminophenylboric acid is -0.05. The amino group is acting as a negative one. The resonance described above tends to make it act as a posit,ive group. The conjugation of the amino group with the benzene ring charges the ring negatively, and so gives a pseudopositivity to the group. Zwitterion formation has a similar effect. The result therefore shows a definite inherent negativity for the amino group. This agrees with the fact that hydrazine is a weaker base than ammonia, and the expected order of negativities, CHs, NH2, OH, F. However, Ebert' A glycine, and Miyamoto and Schmidt' obtained 5.6 X 10" for K ~ of Both of these results make the amino group positive, and 4.4 X more so than the methyl and other alkyl groups. Miyamoto and Schmidt used 4.54 X 2.5 X 10-lo, and 1.85 X lo-* for Kl, K2, and the hydrolysis constant of the methyl ester ion, respectively. Had they not introduced the basic constant of betaine, but had been content with the assumption that K 2 A = the hydrolysis constant of the methyl ester ion, they A , would have obtained K ~ A = and the equation KI.K2 = K ~ A - K ~ 6.1 X This value is greater than K , of acetic acid, and makes the for sevamino group negative. Edsallg has calculated the values of eral amino fatty acids, and finds that they are greater than the dissociation constants of the corresponding unsubstituted acids, and hence that the amino group acts as a negat,ive group in amino fatty acids. His value for glycine is 5 X The aminobenzoic acids are like the aminophenylboric acids in that they have very similar resonances. Cummings6 has measured the constants necessary to determine for m-aminobenzoic acid. It is doubtful whether he fully corrected for the nearness of the values of K 1 and Ke. We redetermined the necessary constants, finding significant variation only in K2. Our method was the electrometric titration of the sulfates of m-aminobenzoic acid and methyl m-aminobenzoate using a hydrogen electrode. For K1 and K 2 values of the amino acid we solved by means of successive approximations the simultaneous equations ( H + ) 2 [ ( H + ) (Na') (HSO7)l K1 = M[(H+) 2K2J - [(H+) Kz][(H+> (Na+) (HSOhl) and
+
+
+
+
The bisulfate ion concentration equals 0.5M(Ht) 1.2 X H+'
+
EDBALL,
ibid., 66,2337 (1933).
+
+
DISSOCIATION O F HYDROGEN IONS
53 1
being the second dissociation constant of sulfuric acid. The 1.2 >< equations were set up from points near half neutralization of the first and !second acidic hydrogen atoms. Cummings’ values are K1 = 9 X lo4, Kz = 1.63 X loW5and the hy~ 2.8 X lo4. Our drolysis constant of the ester, assumed equal to K 1 = values are K1 = 8.85 X lo4, K , = 2.5 X and K I A = 2.8 X 10“‘. Cummings’ values give K2, = 5.2 X and APkz = +0.10, and ours KPA= 7.9 X and APkz = -0.08. The positivity obtained from Cummings’ values could easily be accounted for by the resonance interactions of the amino with the carboxyl group and with the benzene ring, both of which produce a pseudo-positivity. Our values show a negativity more than sufficient to counterbalance the effects of these resonances. Our conclusion is that the amino group is inherently a negative one. Benzoylaminophenylboric acids.-The AP, vakes for the three benzoylaminophenylboric acids are negative. The benzoylamino group is negative, as might be expected from the high negativities of acyl groups. That the order of strengths is meta > ortho > para > unsubstituted acid, while negative groups produce greater effects in the ortho than in the meta position, indicates that the conjugation between the boric acid and amino groups is not entirely removed by the resonance of the type
0
/I
(CsHoNHC-CsH5,
+
0-
I
CaHbNH=C-CsHrj),
which is introduced with benzoylation. Hydroxyphenylboric acids.-The hydroxyphenylboric acids have possibilities of resonance and chelation very similar to those of the aminophenylboric acids. The resonance enhances dissociation from the phenolic group, and hinders that from the boric acid group. Chelation has the reverse effect. Unfortunately the yield of the ortho isomer was so small that it was impossible to make sure that it was the required compound and to measure its dissociation constant. The amount of p-aminophenylboric acid was so small that it seemed hopeless to attempt to get sufficient of the phenol. However, the meta compound was prepared by the method of Bean and Johnson,lo and both of its dissociation constants were and Kz = 1.4 X lo-” (unmeasured. The results were K1 = 2.8 X corrected for salt effect, which should be appreciable for the divalent ion). The value of K1 indicates that the proton dissociates from the boric acid group, and the negativity of the hydroxyl group is the predominant factor. The value of Kz shows that the group B02H- is positive, a fact 10
BEANAND JOHNSON, ibid., 54, 4415 (1932).
532
C. G. CLEAR AND G. E . IC. BRANCH
that in view of the negative charge hardly needs experimental corroboration. EXPERIMENTAL
The aminophenylboric acids were prepared by the catalytic reduction of the nitro compounds with hydrogen, described by Bean and Johnson,lo except that an equivalent of sulfuric acid was added before reduction in each case. The sulfates were crystallized from the reaction mixture by vacuum concentration. The reduction in acid solution is just as efficient as i t is in neutral solution. if the hydrogenation is stopped as soon as the theoretical amount of hydrogen has been absorbed. If this is not done, the reduction in acid solution proceeds past the amine stage, and ammonia and an unstable boric acid derivative are formed. The ortho and meta benzoylaminophenylboric acids were obtained from the mother liquors from the crystallizations of the aminophenylboric acid sulfates by shaking with benzoyl chloride and sodium carbonate. They crystallize as semianhydrides as reported by Seaman and Johnson.2 It was found necessary to use the pure sulfate of p-aminophenylboric acid, and an exact equivalent of benzoyl chloride to obtain the p-benzoylaminophenylboric acid. The equivalent weight of this compound agreed with the formula, C6H6CO;\iHCsH,B(OH)z. The m-hydroxyphenylboric acid was prepared from the corresponding amine sulfate according to the method of Bean and Johnson.10 The hydrogen electrode was used in junction with a saturated calomel cell, connection being made through a saturated solution of potassium chloride. The electromotive forces were measured by balancing with a Leeds and Northrup potentiometer using a lamp-scale type of galvanometer. The cells were enclosed in an air thermostat a t 25". The solutions were brought to this temperature in a water thermostat. Fresh liquid junctions were provided for all measurements. The glass electrodes were tubes sealed with a thin diaphram of 0.15 Corning glass. They were partially filled with a saturated solution of quinhydrone, into which platinum leads dipped. They are similar to those described by MacInnes and Dole." The glass and hydrogen electrodes were used in conjunction with the same calomel cell and saturated solution of potassium chloride. The current from the glass electrodes was amplified by a De Eds12 amplifier. By means of a multi-pole high resistance switch the same potentiometer and galvanometer could be used for both glass and hydrogen electrodes. The glass electrodes were checked before and after each measurement by observing the pH value of a standard buffer solution. Measurements with the glass electrodes were reproducible within 0.5 millivolts, and those with the hydrogen electrodes within 0.1 millivolts. The two types of electrodes agreed within these limits in solutions varying from p H = 2 to pH = 7. The glass electrodes were not used in more alkaline solutions. Table I11 includes the solvent, the type of electrode, and the molality or range of molalities used; the number of approximately half-neutralized solutions for which pH numbers were determined; the range of dissociation constants calculated from each of these solutions; and the mean of these constants. 11 11
MACINNES AND DOLE,ibid., 62,29 (1930). DEEDS,Science, 78,556 (1933).
ISSOCIATION OF HYDROGEN IONS
3
00
00
0
0
533
534
C. G. CLEAR AND G. E . K . BRANCH SUMMARY
The acid strengths of the positive ions of the three aminophenylboric acids have been compared with each other and with that of anilinium ion. The acid strengths of the three aminophenylboric acids have been compared with each other and with that of phenylboric acid. It has been shown that zwitter ion formation has little to do with these relative strengths. The relative strengths have been explained on the assumption that there is a strong resonance in 0- and p-aminophenylboric acid, that this resonance, though reduced, persists in the negative ions and the meta amino acid, that the positive ion of the ortho acid has a hydrogen bridge connecting the substituent groups, that the boric acid group has practically zero negativity, but that the amino group is negative. It has been shown that the amino group is negative in the amino fatty acids, and in maminobenzoic acid. The dissociation constants of the benzoylaminophenylboric acids and both constants of m-hydroxyphenylboric acid have been measured. The first hydrogen lost by the last-named substance appears to be from the boric acid group.
