July 20, 1960
3575
IONIZATION CONSTANTS OF XBIINOACIDS
[COSTRIBUTIOSFROM
THE
DEPARTMENT OF CHEMISTRY, BARXARD COLLEGE,COLUMBIA USIVERSITT,XEW YORK,NEW YORK]
The Thermodynamics of Ionization of Amino Acids. V. The Ionization Constants of 3-Methoxy-DL-alanine (0-Methylserine) and Methoxyacetic Acid1 BY EDWARD J. KING RECEIVEDDECEMBER 28, 1959 The thermodynamic functions for the ionization of methoxyacetic acid and 3-methoxyl-~L-alanine (0-methylserine) were obtained a t ten temperatures from measurements on cells without liquid junction. The ionization constants were fitted to the equation -log K = ( A / T ) - B f CT where A , B and C are, respectively, 974.26, 3.6701 and 0.013325 for methoxpacetic acid; 1290.20, 6.0134 and 0.012490 for the ionization of the carboxyl group of the amino acid; and 2788.60, 2.0924 and 0.0064234 for the ionization of the ammonium group. The entropies of ionization of these acids lie between those of the analogous unsubstituted and hydroxyl substituted acids, for the methoxyl group is polar yet is less capable than hydroxyl of forming hydrogen bonds with solvent molecules. The heats of ionization are smaller (less positive or more negative) than those of the reference acids and this is reflected in the greater strength of the methoxy acids. The relative importance of polar effects and solvation in determining the heat of ionization is discussed.
The study of the effect of polar substituents on the ionization of weak acids has been a favorite field of work since the early days of the theory of ionization. Discussion of polar effects generally has been based on the electrostatic interaction between polar group and the ionizing proton or an inductive effect transmitted through the chain attached to the ionizing group. Few such discussions have been based on accurate thermodynamic ionization constants a t several temperatures. As such data have gradually accumulated, attention has been drawn to other effects, particularly which may be revealed not so much by variations in free energy of ionization or ionization constant as by large fluctuations in heat and entropy of ionization from one polar acid to another. Two acids containing the methoxyl group were selected for investigation. The electron withdrawing power of this group is generally considered to be less than that of h y d r ~ x y lyet , ~ methoxyacetic acid, for example, is a stronger acid than glycolic acid. It has been suggested that the unexpected strength of the methoxy acid is a solvation effect.* Measurements were made on three types of cells (Pt)H2 CH39CH?COOH(nzl),CH30CHzCOONa(m), SaCl(nt3) I AgC1-Ag (I) (Pt)Hs !CH30CH?CH(SHa+)COO-(nti), CHBOCH2CH(SHs")COOH, CI-(m,) 1 AgC1-Ag (11) (Pt)H, I CH~OCH~CH(NH~+)COO-(W~), CH~OCH~CH(NHZ)COO-, Na+(mz), NaCl(nz8) I AgC1-Ag (111)
in which the ioiiization reactions were, respectively CHSOCHzCOOH
+ H20 J_ CHaOCH2COO-
f HgO'
(1) This investigation was supported by a research grant, H-1651, from the National Heart Institute of the National Institutes of Health, U.S. Public Health Service. A preliminary report was given a t the American Chemical Society Meeting, New York, N. Y.,in September, 1957. (2) D. J. G. Ives and J. H. Pryor, J. Chem. Soc., 2104 (1955). (3) E. J. King and G. W. King, THISJOURNAL, 78, 1089 (1956). (4) E. 1. King, i b i d . , 78, 6020 (1956). (5) F. S. Feates and D. J. G. Ives, J . Chem. Soc., 2798 (1956). (6) L. P. Fernandez and L. G. Hepler, THIS JOURNAL, 81, 1783 (1959). (7) See, for example, R. W. Taft, Jr.,in "Steric Effects in Organic Chemistry," ed. by M. S. Newman, John Wiley and Sons, Inc., New York, X. Y., 1956, p. 619, and T. C. Bissot, R. W.Parry and D. H Campbell, THIS J O U R N A L , 79, 796 (1957). ( 8 ) L. 4 . Wiles, Chevt. R ~ v s 56, . , 329 (1956).
CH30CHzCH(NHa+)COOH f H?O CHaOCHzCH(NH3+)COO-
+
CH30CH*CH(NHa+)COOH20 CH3OCHzCH( SH?)COO-
+ H30' + H30'
Measurements were made a t ten equally spaced temperatures from 5 to 50" in order that the entropy and enthalpy changes associated with the ionization reactions in the standard state could be calculated. It was hoped that some deductions could be made from these thermodynamic properties about solvation effects on the ionization process. Experimental Two samples of methoxyacetic acid mere prepared by repurifying a commercial product (Eastman Kodak Co.). One portion of crude material was distilled a t a pressure of 13-14 mm. through a vacuum jacketed S'igreux column 20 cm. long. The middle fraction, b.p. 97", was retained. d second sample was prepared by fractionally freezing the crude material four successive times and then fractionally distilling it. Stock solutions of the acid were made by dilution of the samples and were standardized against carbonate-free sodium hydroxide solution. Two samples of 0-methyl-DL-serine, A.p. grade, were obtained from the H. M . Chemical Company, Santa Monica, California and were used without further purification. They were subjected to a' series of tests: assay by formol titration with standard sodium hydroxide solution, water content by the Karl Fischer method, and semiquantitative tests for chloride, iron, ammonia, phosphate and heavy metals 9 For Lot No. 5.8.22 the results were: assay, lOO.07,; water, 0.257,; and less than 0 . 0 0 1 ~ ,of each of the five minor impurities. For Lot No. 114.65.37 they were: assay, 99.92%; ; others, < 0.004% of each. water, 0.087,; C1< 0 . 0 0 5 ~ 0 the Other chemicals were prepared and standardized as in earlier work .lo,ll The apparatus and techniques have likewise been described before.3J0J1 One standard cell was recalibrated by the h'ational Bureau of Standards three months before measurements began. Buffer solutions were prepared by weight from appropriate combinations of acid, sodium chloride, hydrochloric acid and sodium hydroxide. For the acid buffers of Cells I and I1 a separate solution was prepared for each cell whereas for the alkaline buffers of Cell I11 the 18 cell solutions were made by dilution of six stock buffers. The results on the earliest six solutions in Cell I11 have been discarded because of possible contamination with carbon dioxide from the air. The values of pK derived from these six buffers were between 0.005 and 0.007 unit lower than those from the other twelve buffers from which carbon dioxide was excluded. In determining the second ionization constants of amino acids a frequent source of difficulty is the solubility of silver chloride in the alkaline solutions. As a result the electromotive force a t 25" measured a t intervals during two days (9) M. P. Stoddard and M. S. Dunn, J . B i d . Chem., 142, 329 (1942). (10)E. J. King, THISJOURNAL, 73, 155 (1951). i l l ) E. J . King and G. W. King, ibid., '74, 1 2 1 2 (1952).
EDWARD J. KIXG
3576
usually decreases. This was not observed for Cell 111. Only two cells out of the 18 containing alkaline buffers had electromotive forces a t 25' a t the end of the series of measurements that differed by more than 0.1 mv. from the original values. The average difference between initial and final electromotive forces of the other 16 cells was only 0.017 mv . It is necessary to report smoothed values of electromotive forces to save space, but the original values have been used in the calculations that follow. The electromotive forces can be represented as functions of the temperature by Et = El6 ~ ( -t 25) b(l - 25)' (1) The parameters cf this equation when the electromotive forces are in absolute volts, concentrations are in moles per kilogram of water, and the hydrogen pressure in one atmosphere are given in Table I. The standard deviations between observed and calculated values of the electromotive forces are It 0.035 mv. for Cell I, 5 0.035 mv. for Cell I1 and zt 0.066 mv. for Cell 111.
+
+
n?,
v!I
-
Eaa
10%
-108b
0.54958 .53491 ,52179 .50903 ,50245 ,49321 ,52665 .50714 ,49593 .49825 ,49086 ,48206 .53202 ,51696 .48273
590 541 495 450 426 392 506 443 406 415 389 358 526 478 361
-33 -25 -12 0 0 +lo 38 25 2 ,i 10 10 15 20 10 1'
vu
0.010785 , 0 1787 ,03152 ,05128 .06fi12 .09443 .009629 ,01954 ,02967 ,03964 ,06260 .07374 ,01120j .0195lj ,04876
Calculations and Results Since the method of calculating ionization constants from these data is well known,'* discussion will be confined to the only unusual feature. For Cells I and 11, which contain acidic buffers, the apparent hydrogen ion concentration is estimated with'? -log
117"
ml =
ma
0.010013" ,020045 ,03000 .04012 ,05001 ,07010
Ea6
106a
- lO'b
0.48807 .46455 .45197 .44329 .437ll .42800
196 107 60 '28 6 -28
132 115 108 92 95 90
0.49645 .47558 .46928 .46487 .45888 .45386
208 128 105 89 66 4s
95 88 82 82 76 72
ml = 2mt
0 . 02007 ,04008 .OX15 ,05922 ,07405 ,09039 mp
nt 2
EX
0.02195 0.009955 0.010635 0.86194 .84419 .02029 ,02168 .04474 .83940 .02472 .02640 ,05450 ,83462 ,03114 .03284 ,06737 ,03665 .83187 .03476 ,07519 .82865 ,03982 .04199 ,08614 ,009086 .008908 .86791 ,01836 .012760 .85928 .012510 .02577 .85818 ,01461 .01426 .02816 ,84909 ,01878 ,01916 .03870 ,83994 ,02730 .02784 .05623 .82913 ,04728 .01613 ,09112 ' m2 = 0.010029.
10'a
+30 -31 -51 -61 -76 -89 $48 17 14 -17 -50 -90
Temp
WCI
-
d a ) (2)
+ CT
I
(3)
Methoxyacetic 3-Methoxylalanine 3-Methox lalanine acid DK *K1
*d
5 10 15 20 25 30 35 40 45 50
3.5380 3.5440 3.5505 3.5591 3.5704 3,5834 3,5970 3,6127 3.6314 3.6509
Std dev.
k0.00057
*0.00136
974.26 3.6704 0.013325
1290.20 6.0134 0.012490
c 150 165 162 210 165 168 150 142 155 160 150 160
+A
TABLE I1 THE SEGArIVE LOGARITHMS O F THE IONIZATION CONSTAXTS A N D THE PARAMETERS OF EQUATION 3
B
-10'b
+ log
-
pK = ( A / T ) - B
A
3-Methoxylalanine, basic buffers in1
2.3['26RT)(E - Ewe) 2Sd&/(l
The last term on the right is the Debye-Hiickel approximation for the activity coefficient product of hydrochloric acid in the buffer solutiogs. The ion size parameter B was taken to be 4.00 A., an extrathermodynamic choice which has no significant effect on the ionization constant of an acid as weak 3.5) but does influas niethoxyacetic acid (pK ence the value of the first ionization constant of the amino acid (pKl 2)."yi4 The reason for choosing the value 4.00 A. has been presented before.I5 Values of the negative logarithms of the ionization constants, as obtained by extrapolation of the results of the calculations to zero ionic strength, are collected in Table 11. The standard deviations given in the table indicate the errors in each set of intercepts (pK values) as obtained by the method of least squares. The constants A , B and C in Table I1 are the parameters of the Harned and Robinson equation16
O C .
3-h.lethosylalanine, acid buffers
= (5
-
TABLE I OF EQUATION 1 PARAMETERS Methoxyacetic acid 0.011791 ,01954 .03446 ,05474 ,07059 .lo081 .03050 .0618S ,09398 .OS393 .11138 ,15615 ,02382 ,04149 ,15396
Vol. 82
2.0975 2 . os14 2.0647 2.0474 2.0373 2.0298 2.0223 2.0172 2.0150 2.0160
9,7182 9.5756 9,4363 9.3032 9.175G 9.0538 8.9356 8.8220 8.7167 8.6131 10.00138
2788.60 2,0924 0.0064234
The thermodynamic functions associated with the ionization of the methoxy acids are presented in Tables I11 and I V together with those of related acids. They include the absolute temperature 0, a t which K attains its maximum value, and the changes in free energy (AFO), enthalpy (AHo),entropy (AS") and heat capacity (ACpo) associated with the ionization reaction a t 25" in the standard state. Free energy and enthalpy changes are given (12) H . S. Harned and B. B . Owen, "The Physical Chemistry of Electrolytic Solutions," 3rd Ed., Reinhold Publ. Corp., New York, N. Y., 1958, pp. 650-652, 658-660. (13) See ref. 3 and 4 for the meaning of Ewe. (14) R . G.Bates, J . Rmearch N u l l . Bur. Standards, 47, 127 (1951). (16) H.S. Harned and R. A. Robinson, Trans. Faraday Soc., S 6 , 973 (1940).
July 20, 19GO
IONIZATiON CONSTANTS O F
in calories per mole, entropy and heat capacity changes in calories per degree-mole, where one calorie is defined as 4.1840 joules. TABLE I11 IOYIZ.4TION FUNCTIONS O F METHOXYACETIC ACIDA N D RELATED ACIDS Acid
O K 2 8 3 16
AFOzas
IS
Aceticacid'' 4 757 6486 hlethoxq acetic acid 3 570 4871 Glycolic acid1* 3 538 5223
A H 0 2 e ~18
A S o z 16 ~~
A c o p z a 8 18
0
-100
-22
11
-37
295
-960 $180
-19 57 -16 94
-36 -39
270 303
3577
A N I N 0 ACIDS
tion that acids as diverse in structure as formic acid, glycolic acid, bromoacetic acid and N-acetylglycine all have entropies of ionization a t 25' in the range - 16.9 to - 17.3 cal. deg.-' mole-'. Setting aside entropy effects associated with translational and vibrational motions and with the creation of electric fields about the ions as either being intrinsically small or not varying greatly from one acid to another, one is left with entropy effects associated with restricted internal rotations and with specific interactions of the ions and molecules of the acid with molecules of the s o l ~ e n t . ~
TABLE I\' IOSIZATION FUNCTIONS O F 3-hfETHOXYALANINE AND RELATED ACIDS Acid
Ionizing group
PKlB.ls
AF0m ib
AH%B.M
A S k . 18
Acop2SS 13
Glycine10 3-Methoxyalanine Serinelg Glycinexo 3-Methoxyalanine Serine19
-COOH -COOH -COOH -NHs+ -NHI + -xHd +
2.349 2.037 2.187 9.779 9.176 9.209
3208 2780 2984 13340 12520 12560
980 820 1320 10550 10150 10390
-7.56 -6.56
-34 -34 -31 - 12 - 18 - 10
Discussion The data of Tables I11 and I V show that the niethoxy acids are stronger acids than the comparable unsubstituted and hydroxyl substituted compounds. The effect of a substituent can enter as a potential energy contribution to AFO, ;.e., as a polar or inductive effect, and as a kinetic energy contribution dependent on the motions of solute and solvent molecules and ions. The polar effect can scarcely be all important, for the group moment of -0CH3 is less than that of -OH, viz., 1.3 as compared with 1.7 D.*O Furthermore, the methoxyl group is generally considered to have less electron withdrawing power, a t least in aromatic compounds.7 Thus on the basis of a polar or inductive effect alone one would expect methoxy acids to be weaker than their hydroxy counterparts. Two possible interpretations of the data in Tables I11 and IV must be considered: either the methoxy acids are stronger than polar effects would lead one to expect or the hydroxy acids are weaker. For further clues i t is necessary to examine the heats and entropies of ionization. It is convenient to consider the entropy change first, for it is independent of potential energy contributions. Interaction of a substituent dipole with the ionizing proton occurs partly within the cavity of the acid molecule and partly through the solvent. Insofar as i t depends on the dielectric constant of the solvent this interaction will contribute to the entropy change, but theoretical estimates indicate that such a contribution must be small.21 This is also supported by the observa(17) H. S. Harned and R. W. Ehlers, THIS JOURNAL, 68, 652 (1933). (18) L. F. Nims, i b i d . , 18, 987 (1936). ( 1 9 ) P. K. Smith, A. T . Gorham and E. R. B. Smith, J . Biol. C l e m . , 144, 737 (1942); recalculated from the original electromotive force data using the constants recommended by F . D. Rossini, P. Gucker, Jr., H. Johnston, L. Pauling and G. Vinal, THISJOURNAL, 74, 2699 (1952). These are the constants on which the calculations of this paper are based. (20) Landolt-Bbmstein, "Zahlenwerte und Funktionen," ed. A. Eucken and K. H. Hellwege. Sechste Auflage. I. Band, 3. Teil, Springer Verlag. Berlin, 1951, pp. 507-508.
-5.57 -9.05 -7.95 -7.41
e 321 32 1 338 792 659 833
It will be seen from Tables 111 and IV that the entropies of ionization of the methoxy acids are about half-way between those of the reference acids. This is most simply interpreted as a solvation effect. Water molecules about the methyl group in acetic acid are within 5 8.of the carboxyl group and are thus subject to orientation by the negative charge produced by ionization with loss of entropy. 2 2 Water molecules about the hydroxyl group of glycolic acid molecules have their entropy lowered even before ionization by hydrogen bonding with the hydroxyl group. They are then less subject to perturbation by the field created by ionization than those around a methyl group. As a result, less entropy is lost in ionizing glycolic acid than in ionizing acetic acid. In methoxyacetic acid the methoxy1 group contains an oxygen atom with two lone pairs of electrons and is still capable of forming hydrogen bonds* though to a much smaller degree than the hydroxyl group. Moreover, the terminal methyl group in methoxyacetic acid should have a structure-strengthening effect on water similar to that manifested in the behavior of acetic and propionic acids3 The entropy loss on ionization of methoxyacetic acid should therefore be larger than that for glycolic acid but smaller than that for acetic acid. In methoxyalanine the methoxyl group is further from the carboxyl group than it is in methoxyacetic acid, yet it is still within 5 A., the limit of influence of the carboxylate ion on water stru.cture.22 The order of the entropies of ionization of the acids in Table 1V can thus be accounted for by the same reasoning as was used for those in Table 111. The greater strength of methoxy acids as compared with hydroxy acids cannot be an entropy effect, for if this were predominant the methoxy compounds would be the weaker acids. The entropy effect is more than offset by an enthalpy ef(21) F. Westheimer and J. Kirkwood, Trans. Faraday Soc., 43, 77 (1947). (22) D.H . Everett, D. A . Landsman and B . R. W. Pinsent, Proc. Roy. SOC.( L o n d o n ) , 2lSA, 403 (1952).
3578
EDWARD J. KING
1-01, 82
fect. For methoxyacetic acid the heat of ioniza- COOH, which has a terminal methyl group, is tion is, a t first glance, surprisingly large and nega- -150 cal./mole whereas that of hydantoic acid, tive. A second possibility must nevertheless be NHzCONHCH~COOH, is +a90 cal./mole. Groups borne in mind: the anomalous heat of ionization that hydrogen bond to water and break its strucmay be that of the hydroxyl acid. X basis for de- ture increase A I P ; those that repel water with the ciding between the two possible anomalies can be result that i t draws away from them like a conobtained by examining the known heats of ioniza- tractile skin and has its structure strengthened, tion of a wide variety of acids. The comparison decrease AHo. will be made a t 25" where AHo is known with the On the basis of this survey of heats of ionization greatest precision. At higher temperatures all i t is now possible to come to some conclusion about A H o values for carboxylic acids are less positive the behavior of glycolic and methoxyacetic acids. or more negative; because all are shifted to about It is to be expected that in passing from HOCHZthe same extent, the conclusions of this discussion COOH to CH30CH2COOH a decrease in AHo are not altered by the choice of a different tempera- should occur, for the hydrogen bonding power of ture for the comparison. the molecule decreases and a terminal methyl group Because i t includes a potential energy contribu- is introduced. The large difference in heats of tion, AHo is a more complex quantity to interpret ionization of methoxyacetic acid (- 960 cal. 'mole) than ASO. The large difference in the two heats of and acetic acid (-100 cal./mole) is a reflection of ionization of glycine (Table IV) is an illustration the interaction between the polar methoxy group of the importance of the potential energy contribu- and the ionizing proton. Among polar acids iodotion, for i t is a reflection of the difference in binding acetic acid, the least solvated of the monohaloenergy of the proton to the carboxylate and ammo- genoacetic acids,2 is most nearly comparable with nium groups. methoxyacetic acid. These two acids have about Uncharged carboxylic acids can be divided into the same entropy of ionization : - 19.1 and - 19.6 several groups on the basis of their heats of ioniza- cal. deg.-' mole-'. Iodoacetic acid has the more tion. The short chain fatty acids have small nega- negative heat of ionization : - 1370 cal./mole as tive values of AHo a t 25". Acids which have large compared with -960 for methoxyacetic acid, but negative heats of ionization include the monohalo- this is reasonable in view of the larger group gen derivatives of acetic acid, cyanoacetic acid and moment of iodine, 1.6 as against 1.3 D for methoxyl methoxyacetic acid. These are polar acids and the For polar acids in general there is no correlation interaction of the polar substituent with the ioniz- between dipole moment and heat of ionization,2 ing proton leads to a large negative potential en- because the heat of ionization includes a negative ergy contribution to AHo. Glycolic, succinic and contribution from the polar effect and a positive one hydantoic acids, which contain good hydrogen from solvation. Only when solvation is weak, bonding groups, namely: -OH, -COOH and as for iodoacetic and methoxyacetic acids, does -NHCONH2, have positive heats of ionization. an increase in dipole moment correspond to a more Hydrogen bonding between the molecules of these negative heat of ionization. For heavily solvated acids and water lowers the enthalpy of the un-ion- cyanoacetic acid AHo is only -S40 cal./mole, ized state and less enthalpy is lost on ionization. much less negative than that for iodoacetic acid, Other enthalpy effects are associated with changes despite the large polar effect to be expected of in the alkyl chain, the position of the polar group, cyanide with a group moment of 3.4 D.5 The if any, being kept fixed. As the alkyl chain is polar effect for glycolic acid is smaller and is lengthened or branched AHo becomes more nega- swamped by the strong solvation of the moletive, This effect is frequently coupled with a de- cules; the heat of ionization is positive. I t may crease in A s 0 with the result that AFO and pK remain almost unchanged, 3, l5 Extension of the chain takes thus be concluded that the enthalpy of solvation place in passing from glycolic to lactic acid, for ex- effect is responsible for the fact that hydroxy ample, and this is associated with a decrease in AHo acids are weaker than methoxy acids. Grateful acknowledgment is made to Dr. from +IS0 to - 120 cal./mole. Likewise, the heat of ionization of S-acetylglycine, CH,CO?;HCH!- J E Prue for helpful discussions.
