QUANTITATIVE INVESTIGATIONS O F AMINO ACIDS AND PEPTIDES. XI11 EQUILIBRIA BETWEEN HISTIDINE AND FORMALDEHYDE' EDWARD H . FRIEDEN,%MAX S. DUNN,
AND
CHARLES D . CORYELL
Department of Chemistry, University of Calzfornia, Los Angeles, Cdtfornia Received October 12, 1042
The polarimetric investigation of the reaction of formaldehyde with typical amino acids has been described previously (2, 3, 4). Such studies have generally been in excellent qualitative and fair quantitative agreement with the results published by other workers from electrometric studies of these equilibria (6). The polarimetric method has now been extended to include the reactions of the basic amino acids histidine, lysine, and arginine with formaldehyde. This paper contains a summary of an investigation of the histidine-formaldehyde reaction. The other reactions will be discussed in a latter communication. In 1935 Levy (7) described the results of potentiometric studies of the equilibria between formaldehyde and the basic amino arids, including histidine. From the change in pH of a solution of histidine (originally a t pH = p&) as formaldehyde was added, he concluded that the cation, R(IH+)NH:COl, and the zwitter ion, R(I)NH:C02, each reacted with but 1 mole of formaldehyde, the former reaction being far less complete than the latter. Only the amino group was presumed to react, the proton shifting in one case to the carboxyl group and in the other case to the previously uncharged imidazole group (represented by I in the above formulae). When formaldehyde was added to histidine originally a t pH = pK8, the observed changes in pH could not be interpreted as easily. Levy's conclusion that, under these conditions, the zwitter ion and the anion, R(I)NH&OT, each reacted with 1 mole of formaldehyde, appeared to account for the observed facts. The failure of histidine to react with more than 1 mole of formaldehyde under the conditions used by Levy is quite surprising. The polarimetric method seemed to offer a convenient means of verifying the conclusions reached as a result of the studhs quoted. Accordingly, polarimetric titrations of solutions of anionic, zwitkr-ionic, and cationic histidine with formaldehyde were performed, using the apparatus an{ procedure described earlier (2,3). The results of these titrations will be described in the experimental part. 1
For the preceding communication in this series, see Frieded, Dunn, and Coryell:
J. Phys. Chem. 47, 20 (1943). This paper is part of a dissertation submitted by Edward H. Frieden to the Graduate School of the University of California in partial fulfillment of the requirements for the degree of Doctor of Philosophy in Chemistry. The authora were aided in this work by grants from the University of California and from Merok and Company. 8 Present Address: Department of Chemistry, University of Texas, Austin, Tcsas.
85
TABLE 1 Polarimetric titration of 0.0324 M I(-)-histidine a n i o n with formaldehyde FORYALDEBYDE
HISTIDWE
moles 91, lifcr
molcr fief litcr
O.Oo0 0,00346 0.00691 0.01031 0.01372 0,01724 0,02045 0.0238 0.0271 0,0304 0,0337 0.0435 0.0530 0.0655 0.1035 0.1760 0.284 0.389 0.561 0.827 1.381 1.940 2.445 3.325
0.0324 0.0323 0.0322 0.0321 0.0320 0.0319 0.0318 0.0318 0.0317 0.0316 0.0315 0.0312 0.0309 0.0306 0.0304 0.0303 0,0300 0,0298 0,0294 0.0287 0,0273 0.0260 0.0247 0.0226
I 00
02
OBSERVED ROTAIION
dcgrccr
-0.212 -0.474 -1.207 -1.765 -2.462 -2.739 -3.170 -3.650 -3.904 -4.027 -4.020 -3.835 -3.669 -3.457 -3.107 -2.625 -2.157 -1.857 -1.539 -1.230 -0.905 -0.730 -0.624 -0.497
I I I I I I I I 0,4 0.6 0.8 LO d d I? 1.6 1.8 ~ O R M A L D E H Y D E ,MOLES Kf?L l T f R
-16.3 -36.7 -93.5 -136.5 -192.2 -214.2 -248.5 -288.0 -308.0 -319.5 -320.0 -308.0 -297.0 -282.5 -255.0 -215.0 -179.8 -150.3 -131.2 -107.2 -82.8 -70.3 -63.0 -55.0
I
2.0
I
2.2
I
2.4
FIG.1. hIolecular rotation plotted against total formaldehyde concentration for anionic I(-)-histidine. The solid curve represents the rotations calculated from the constants derived in the text; the circles are esperimental points. 86
INVESTIGATIONS OF AMINO ACIDS AND PEPTIDES.
Polarimetric titration of 0.0964 M
TABLE 2 -)-histidine zwitter i o n with 18.40 M formaldehude OBSERVED BOTATION
VOLUME
FOBMAWEEYDE
BISTIDINE
ml.
moles pcr lifer
males per lifer
degrccs
0.000 0.040
0.000 0,0084 0.0168 0.0293 0,0376 0.0460 0.0522 0.0584 0.0666 0.0830 0.1140 0.1546 0.2060 0.3062 0.503 0.966 1.793 3.138 4.295 5.110
0.0354 0.0354 0.0354 0.0353 0.0353 0.0353 0.0353 0,0363 0.0352 0.0352 0.0351 0.0351 0.0348 0.0345 0.0340 0.0327 0.0303 0.0265 0.0233 0.0210
-1.054 -1.707 -2.301 -3.014 -3.314 -3.419 -3.506 -3.499 -4.473 -3.371 -3.339 -2.862 -2.521 -2.081 -1.406 -0.703 -0.173 0.121 0.202 0.244
0.080
0.140 0.180 0.220 0.250 0.280 0.320 0.400 0.550 0.750 1.OOO 1.500 2.500 5.000 10.00 20.00 31.2 41.2
87
XI11
-74.5 -123.0 -162.7 -213.0 -232.1 -242.0 -248.5 -248.0 -246.5 -239.5 -223.5 -203.8 -181.0 -150.8 -103.2 -54.0 -14.3 11.4 21.7 29.1
FIG.2. Molecular rotation plotted against total formaldehyde concentration for zuitterionic l(-)-histidine. The solid curve represents the rotations calculated from the constants derived in the text; the circles are experimental points.
88
E. H. FRIEDEN, hl. S. DUXS, AND C. D. CORYELL
TABLE 3
Polarimetric titration of 0.0404 M I ( - )-histidine monohydrochloride with 12.68 Mformaldehyde VOLUPE
FOBMALDEKYDE
HISTIDINE
OBSERVED BOTATION
ml.
moles par liler
molar per liler
degrees
0.00 0.100 0.200 0.300 0.500 0.750 1.010 1.500
0.ooO
2.ooO
0.485
3.000 5.00 7.50 10.00 15.00 20.00
0.716 1,148 1,647 2.100 2.908 3.600
0.0494 0.0493 0.0492 0.0491 0.0489 0.0486 0.0484 0.0480 0.0475 0,0466 0,0449 0.0430 0.0411 0.0308 0,0353
0.158 0.076 0.032 -0.016 -0.078 -0.146 -0.178 -0.249 -0.290 -0.323 -0.297 -0.186 -0.038 0.179 0.319
0.0252 0.0503 0.0752 0.1249 0,1860 0.2495 0.368
0
04
09
I t
I 1.6
8.00 3.80 1.62 -0.76 -3.99 -7.51 -9.21 -12.98 -15.27 -17.31 -16.62 -10.81 -2.31 11.77 22.60
I
I
I
I
2.4
I
2.0
2.8
$2
3.6
FORPIALDEH YDE, MOLES PER LITER \
FIG. 3. Molecular rotation plotted against total formaldehyde concentration for cationic I(-)-histidine.
INVESTIGATIONS OF AMINO ACIDS AND PEPTIDES.
XI11
89
EXPERIMENTAL
Anionic histidine
I ( -)-Histidine (free base, c.P., Lot No. 3), obtained from Amino Acid Manufactures, was used. To the quantity of amino acid taken for the run (about 0.5 g.) was added sufficient sodium hydroxide to bring the solution to pH 12.2, corresponding to 99.9 per cent conversion to the anionic form. The solution was then diluted to 100 ml. Preliminary experiments had indicated that it was necessary to use dilute formaldehyde for the first part of the run, in order to be able to estimate formaldehyde concentrations accurately. Accordingly, the titration was begun with 1.111 M formalin; after 5 ml. of this solution had been added, the titration was completed with 12.56 M formaldehyde. The results of the experiment are summarized in table 1. The data of table 1 are plotted as circles in figure 1, superimposed upon the curve which represents the predicted relation as calculated by the method described. Zwitter-ionic histidine 1( -)-Histidine (free base) was again used. About 0.5 g. of the amino acid was
dissolved in water and the solution diluted to 100.0 ml. No alkali was added. The histidine solution was titrated with 12.49 M formaldehyde. The results of the experiment are given in table 2. The data of table 2 are plotted in figure 2. As before, the solid line represents a series of calculated values. Cationic histidine In order to study the effect of formaldehyde upon the rotation of cationic I ( - )-histidine, the rotation of a solution of histidine monohydrochloride was observed as formaldehyde was added. It has been shown (1) that in a solution 0.06 ill in this salt at 25'C., more than 98.5 per cent of the amino acid exists in the form R(IH+)NH:CO;. Table 3 lists the experimental results. Figure 3 represents the data of table 3. As no attempt was made to calculate the equilibrium constants of the reactions involved, the solid line represents the best smooth curve drawn through all the points. DISCUSSION
It is clear from a consideration of figures 1, 2, and 3, that Levy's formulation of the reactions between formaldehyde and the various forms of histidine is inadequate. In each case the particular ionic species reacts to form two compounds, and it is reasonable to assume that each curve represents the reaction of histidine with first one, and then another, molecule of formaldehyde. This is completely analogous to the corresponding reactions observed for leucine and glutamic acid. Using the terminology introduced by Levy, the various reactions can be formulated as follows: For anionic histidine:
90
E. H. FRIEDES, 11. S. DUNX, AKD C . D. CORTELL
For zwitterionic histidine:
+ F+ZF* + F + ZF:
Z*
ZF*
Liz = (ZF=)/(Z*)IF)
(3)
LL = (zF:)/(zF*)(F)
(4)
For cationic histidine:
+ F -+ CF' + F -+ CF:
'2%
(cF+)/(c~)(F)
=
(5)
Lh = (CFS)/(CF')(F) (6) The evaluation of the above equilibrium constank can be accomplished by the method of successive approximations d e s d i e d for leucine. For anionic CF'
-2
-1.8
-f.g
%+
-3.2
I
1
-2.0
-08
I -0.6
I
-0.4
I -02
I
0,o
I
0.2
I
0.4
L O G Cf3 FIG.4. Logarithmic treatment of the data relating t o the addition of the second mole of formaldehyde to anionic histidine.
histidine the problem is comparatively simple, because the apparent large valuc of LI3 makes the estimation of a2 quite accurate. The initial estimates of a2 and a3 (the molecular rotations of AF- an? AFT) viere -325" and --ioo, respectively. The first approximation using these values indicated that az was more accurately -326" and ai8 was -44.0'. These corrected values werc then used without further modification to calculate the log (hF;)/(AF-) - log (F) relationship. The procedure used for this calculation has been described in the discussion on leucine. When this procedure waq applied to the data of table l (2F > 0.04 X), the resulting curve was a straight line of slope very accurately unity. This curve is reproduced in figure 4. The value of L:, can be determined from the value of log (F) at log
INVESTIGATIONS O F AMIKO ACIDS AND PEPTIDES.
XI11
91
(AFP)/(AF-) = 0, as explained previously. From figure 4, L:, is calculated to be 4.52 with considerable precision. The data of figure 1 do not permit a calculation of LIB. This constant is extremely large, and the concentration of free formaldehyde a t any point on the curve (ZF < 0.04) is a t least as small as the uncertainty involved in the calculation of total formaldehyde concentration. The calculation of figure 1 was calculated assuming that LIS was infinite, but an indirect method of arriving at the value of this constant will be described later.
FIQ.5. Logarithmic treatment of the data relating t o the addition of the second mole of formaldehyde t o zwitter-ionic histidine.
From an inspection of figure 2 it is apparent that the value of LIZis considerably less than &. For the value of &, 012 was estimated to be -300" and as estimated to be 66'. Two approximations were required to establish azaccurately at -274' and 013 a t 45'. The calculation of log (ZF$)/(ZF*) and log (F)resulted in figure 5, whence L X was calculated to be 2.54. Using the figure -274" for 012, the value of LIZ can be calculated from the first part of figure 2. A preliminary estimate of this constant resulted in the value 4.4 X lo2,which appeared to be quite inaccurate owing to the rapid increase in the concentration of ZF:. Accordingly, the rotations listed in table 2 were corrected for this effect [the method of correction has been described previously (3)], and the log (ZF*)/(Z*) -log (F) terms recalculated. The corrected data are plotted in figure 6 , and from the value of log (F) a t log (ZF*)/(Z*) = 0, LIZ was calculated to be 6.6 X 10'. While this constant is undoubtedly subject to some
92
E. H. FRIEDEN, M. S. DUNN, A S D C. D. CORYELL
uncertainty, it is probable that the true value differs from that given by not more than about 10 per cent. The solid line of figure 2 represents the (AM)- ZF relationship expected if L,, = 6.6 X IO', LLp = 2.54 (LZZ= 1.68 X IO3), cy1 = -74.5", cy2 -- -274O, and cyg = 45'. Because it was impossible to estimate the asymptote of figure 3 with any degree of assurance, no attempt was made to cald~latevalues for Ll1 and Lkl. The former constant can be calculated indirectly, however, and this calculation is described later in this paper.
0 0
-2g-
-
I
I
I
I
I
I
I
l
l
1
I
INVESTI(lhTI0NS OF AMINO ACIDS AND PEPTIDES.
93
XI11
coo-
coo-
@*) (-4-1 The above represent the acidic ionization constants of the various forms of histidine, and all are known. By suitable manipulation of equations 3, 5, and 8, the following identity may be established:
The term on the left-hand side of equation 10 is recognized as the reciprocal of
COOH
/ the dissociation constant of the acid R-NHzF \
Levy (7) gives the reasonable
IH+ value 10-8.16for this constant. LU has been calculated to be 6.60 X lo*, and K2 is given by Greenstein (5) as Lll is thus 0.74, a figure which is consistent with the characteristics of figure 3. TABLE 4 Equilibrium Constants for the reactions between hislidine and formaldehyde
1
Constant ..................
This paper.. .............. Levy ......................
LU 0 74 0:59
LIt 8.6.X I@ 4.3 X 10'
I
LIZ
1.88 X 1Q 0
LIS 4.2 X 10' 13.2 X 10'
1
La: 1.45 X 106 0
Similarly, using equations 1, 3, and 9, the relation of L I to ~ LIPis given by equation 11:
94
JAMES W. MCBAIN AND 0. E. A. BOLDUAN
The best value (from Levy) for the dissociation constant of the acid R(IH+)I\'HzFCOO-, is and Greenstein gives 10-9"0 for Ka. L13 turns out to be 4.16 X lo4,a figure which is consistent with the characteristics of figure 1. Lais L13 X 4.52, or 1.45 X 10'. A comparison of the constants calculated by Levy and those given in this paper is given in table 4. In every case in which comparable values are reported, the agreement of the values calculated by the authors with those given previously is quite satisfactory. SUiMMARY
The polarimetric method has been applied to the study of the equilibria between histidine and formaldehyde. Contrary to previous reports, evidence is presented to show that each of the three ionic forms of histidine reacts with first one, and then another mole of formaldehyde. The calculation of the values of five equilibrium constants which characterize these reactions has been described. The values of these constants are in good agreement with the comparable values which have been reported previously. REFERENCES (1) D U X N hf. , S., FRIEDEN,E. H., STODDARD, M. P., A N D BROWN,H. V.: J. Biol. 144, 487 (1942). (2) FRIEDEN,E. H., D U N N ,hl. S., ASD CORYELL, C. D.: J. Phys. Chem. 46, 215 C. D.: J. Phys. Chem. 47, 10 (3) FRIEDEN,E. H., DUNS, M. S.,AND CORYELL, (4) FRIEDEN, E. H., D U N N M. , S., AND CORYELL, C. D.: J. Phys. Chem. 47, 20 (5) GREENSTEIN, J. P.: J. Biol. Chem. 93, 479 (1931). (6) LEVY,M.:J. Biol. Chem. 99, 767 (1933). (7) LEVY,M.: J. Biol. Chem. 109, 365 (1935).
Chem. (1942). (1943). (1943).
OSMOTIC PROPERTIES OF SOLUTIOKS OF SOME TYPICAL COLLOIDAL ELECTROLYTES' JAMES W. McBAIN
AND
0. E. A. BOLDUAN
Department o j Chemistry, Stanjord Cniversity, California Received December
4,1948
Colloidal electrolytes are characterized by the replacement of an ion by conducting colloidal particles or micelles. Hence, the osmotic activity of colloidal electrolytes is correspondingly diminished (6), while their conductivity may remain comparatively high, owing to the other free ion and the conductivity of the charged colloidal micelles. Although much has been published on the electrical behavior, few data have been supplied during the last twenty years on the thermodynamic or osmotic behavior (3) of colloidal electrolytes. 1 Presented a t the Nineteenth Colloid Symposium, which was held a t the University of Colorado, Boulder, Colorado, June 18-20, 1942.
