18
I. M. KLOTZ, I. LUCILLE FALLER, AND J. M. URQUHART
SPECTRA OF COPPER COMPLEXES WITH SOME PROTEINS, AMIKO ACIDS, AKD RELATED SUBSTANCES“ IRVING M . KLOTZ, I. LUCILLE FALLER, A N D JEAN M. URQUHART
Department of Chemistry, Northwestern University, Evanston, Illinois Received August 2.9, IS@ I. ISTRODUCTION
Copper complexes with proteins occur in many tissues in both animal and plant species and have been studied extensively for over a century (8). Of particular interest have been hemocyanin, the respiratory pigment of certain invertebrates, and the copper-containing enzymes, such as tyrosinase. Despite the effort expended on the problem, however, there is as yet no clear-cut indication of the nature of the bonds between the metal and any specific protein. In an attempt to contribute to the solution of this problem, an investigation has been made of the spectra of several “artificial” complexes prepared by the addition of cupric ion to crystallized proteins. The optical properties of these complexes, in turn, can be understood best by comparison with those between cupric ion and simple low-molecular-weight substances. 11. EXPERIMEXTAL
The cupric ion used in the spectra described came from a reagent grade sample of CuC12.2H20. The sodium acetate, acetic acid, and citric acid were also of reagent grade. The betaine hydrochloride, glycolic acid, hippuric acid, and glycine were commercial samples of C.P. or “highest purity” grade. The glycine ethyl ester was prepared by Dr. L. C. King and the glycylglycine by Mr. S. Preis of this laboratory. The glycylglycine was obtained as the hydrochloride monohydrate by following the procedure of Schott, Larkin, Rockland, and Dunn
(18). The peptides-triglycine, tetraglycine, pentaglycine, and hexaglycine-were the gift of Dr. E. F. Mellon of the Eastern Regional Research Laboratory. The samples of @-lactoglobulin and @-casein were also obtained from the Eastern Regional Research Laboratory through the kindness of Dr. T. L. McMeekin. The bovine serum albumin and yglobulin were crystallized samples from Armour and Company. The albumin was methylated by exposure to acidified methyl alcohol for 3 days at room temperature, according to the directions of Fraenkel-Conrat and Olcott (9). Hemocyanin was examined in the form of the serum of the horseshoe crab, 1 Presented at the Tuenty-third Kational Colloid Symposium, which was held under the auspices of the Division of Colloid Chemistry of the American Chemical Society a t Minneapolis, Minnesota, June 6 8 , 1949 * This investigation was supported in part by grants from the Office of S a v a l Research (Project S o N R 135054) and from the Carnation Company.
COPPER-PROTEIN
19
COMPLEXES
Lzmulus polyphemus. The serum was obtained by methods described previously (15). The spectrum was obtained on a sample which had been diluted fivefold.3 Absorption spectra were obtained vith a Beckman spectrophotometer, using cells of 1 em. depth, d . Molecular extinction coefficients, e, were calculated from the usual expression, log,, I , ' I = ecd, ivhere c is expressed in moles per liter. 4c
3c
2(
IC
600
700 WAVELENGTH
I
1
aoo
900
I
A (my)
FIG.1. Spectra of 0.01 .lf cupric ion in acetate solutions of various concentrations. T h e pH's were 4.49 for the three louest concentrations and 4.83 for the 1 111 solution. 111. RESULTS AXD DISCUSSIOS
A . Copper complexes with carboxylic acids dmong the simplest of copper complexes with carboxylate ions are those with acetate ion. These complexes have been studied quantitatively by Pedersen (17), n-ho calculated association constants, ko, for the reactions: Cu++
+ CH3COO- = [Cu(CHaCOO-)]+
[Cu++(CHaCOO-)]+
+ CHICOO-
=
[CU(CH~COO-),]
ky = 146 (1) k," = 11 (2)
The spectra of cupric ion in acetate solutions of various concentrations are illustrated in figure 1. In solutions of 0.02 M acetate ion, the position of the peak 3 Experiments with hemocyanin were carried out a t the Marine Biological Laboratory, Koods Hole, Massachusetts.
20
I. M. KLOTZ, I. LUCILLE FALLER, AKD J. M. URQUHART
is a t 790 mp. At this concentration approximately two-thirds of the cupric ion is bound in the monoacetate complex, and perhaps one-tenth is in the form of the diacetate complex, if one may judge from Pedersen’s results (17). On the other hand, in the 1 M solution, with an absorption maximum a t 740 mp, the cupric ion is found to be about 45 per cent in the form of the diacetate, 45 per cent as triacetate, and about 10 per cent as monoacetate complex.
FIG.2. Calculated absorption curves for three species of cupric-acetate complex near pH 4.5.
From these data on spectra, combined with Pedersen’s (17) equilibrium studies, it is possible to calculate approximately the absorption curves which may be attributed to the individual species, Cu(CH&OO-)+, Cu(CH&OO-)*, and Cu(CH&OO-),-, respectively. The curves obtained are illustrated in figure 2. Those for the higher complexes are less reliable than that for Cu(CH3COO-)+, since they are based on the extrapolated portions of Pedersen’s (17) measurements. Xevertheless, it is evident that cupric ion in complexes with carboxyl groups in acetate ions does not have an absorption peak below about 720 mp even at very high anion concentrations. In preparation for the interpretation of copper-glycine spectra, an effort was made to evaluate the influence of the positive charge on the substituent
21
COPPER-PROTEIS COMPLEXES
attached to the a-carbon atom. For this purpose the spectrum of cupric ion in
+
aqueous betaine solutions was examined. Betaine, (CHd&CH,COO-, was chosen, since it contains a cationic substituent on the a-carbon atom and yet cannot ionize off a hydrogen ion.
t
WAVELENGTH
(my)
FIG.3. Absorption spectpa of 0.01 I!. cupric ion in solutions of various a-substituted carboxylate-containing organic ions: A , no added organic ion, pH 4.0; B, 0.015 M betaine, pH 4.0; C, O.Ol8JI hippurate. pH 4.8; D, 0.017 .l! glycolate, pH 4.8; E , 0.01 M c i t r a t e , p H 6.6.
The spectrum of cupric ion in 0.015 .If betaine is illustrated in figure 3. It is apparent from the practically negligible increase in absorption over that of cupric chloride that very little cupric ion has formed a complex with the betaine. The repulsion of the positively charged substituent for the cupric cation evidently is enough to overcome the intrinsic affinity of the carboxyl group. It is of interest in the case of these betaine complexes of cupric ion to calculate the association constant nhich might be predicted from theoretical considerations of electrostatic effects. In this treatment we shall neglect any effects
22
I. M. KLOTZ, I. LUCILLE FALLER, AND J. M. URQUHART
of the (CH3)32- substituent other than that due to the charge. On this basis the difference in the association constants of cupric ion with betaine and with acetate, respectively, may be attributed to the additional electrical work necessary to bring up the metallic cation to the betaine. Following Bjerrum (3), Xeuberger (le), and others, we may obtain this electrical work, W e ,from the equation
W , = -2e$
(3)
where $ is the potential produced by the positively charged trimethylammonium substituent a t the distance t from the carboxyl group, and e is the electronic unit of charge. The coefficient 2 arises from the divalency of the cupric ion, and the negative sign from the fact that electrical work must be done on the system in the reaction
+ + + (CHI)~NCH~COOe CU[(CHI)~NCH~COO-]++
CU++
(4)
+
if we consider only the electrostatic effect of the (CHS)3S- group. For the present calculation, the potential # will be taken as + = - eDr where D is the dielectric constant of the solvent. In view of the approximate nature of the calculation, no attempt will be made to introduce an “effective dielectric constant,” such as has been suggested by Kirkwood and Westheimer (13), to take into account the cavity of low dielectric constant formed by the molecule itself. The electrical work becomes, therefore,
w - -D2e?-r c -
The differences in free energy of formation of the cupric complex with betaine and acetate, respectively, may be expressed, therefore, by the equation: AFCu-betaine
-
AFCu-acetate =
2e2N
(7)
where the introduction of N , Avogadro’s number, converts the electrical work to a mole basis. If a value of 1.25 A. is used for r , since such a value in acidity calculations (16) compensates for the use of the dielectric constant of the solvent for D , we find that the first association constant for the copper-betaine complex, ky, should be k;
=
1.8
x
10-3
(8)
Such a value for an association constant is exceedingly small and implies that practically no cupric ion would be united in a complex with betaine. The observed behavior thus is qualitatively in agreement with this prediction. More detailed experiments would be necessary before quantitative agreement can be checked.
COPPER-PROTEIK
23
COMPLEXES
The absorption spectrum of cupric ion in solution of 0.018 lli hippurate ion, CsH6COSHCH2COO-,has also been examined. In this case the positive charge
H
of the glycine nitrogen has been eliminated but the -Sgroup has been retained. The absorption curve (figure 3) is below that of cupric ion in an acetate solution of roughly equal concentration (0.02 -11).Evidently the suhstituent on the a-carbon atom contributes no specific affinity for the metal cation. .As a result, only the polarity of the group manifests itself. The p d a r substituent in the cy-position should reduce the affinity of the carboxyl group for cupric ion, if electrostatic considerations (3, 13, 16) alone are applicable. In glycolic acid, in contrast to hippuric, the hydroxyl substituent on the acarbon atom does seem to contribute to the stability of the copper complex. From polar considerations alone one would expect a decrease in affinity. dctually, the absorption of cupric ion in 0.01i Jf glycolate solution (figure 3) is somewhat more intense than that in 0.02 Jf acetate solution. Evidently the hydroxyl group stabilizes the Cu+--carhosyl bond, perhaps by entering itself into a bond with the cation. -An enhancement of cation binding by cy-hydroxy groups in organic carboylic acids has been observed also ivith the alkali metal ions (6). .1 marked increase in the formation of complexes of cupric ion is observed with citrate ion (figure 3), in which a molecular extinction coefficient of over 50 is reached in a 0.01 .11 solution. In this case, the added affinity comes not only from electrostatic at traction of the negatively charged substituents but also from the statistical effect of having three carboxylate groups availahle to the cation. I t is of interest ,to note that the wave length of maximum absorption of copper citrate is at 7400 .I,,just as it is in the case of 1 Ji acetate complexes, hut that the intensity of absorption is higher. Since the cupric ion in the 1 Jf acetate solution is roughly equally divided between the di- and tri-acetate complexes, it seems reasonable to surmise that in the citrate comples all three carboxyl groups are linked to the copper. I t is then apparent that even a triple complex of cupricmionsand carboxyl groups does not show an absorption peak below about 7400 A. B. Copper complexes wilh g l y c i w The spectrum of cupric ion in aqueous glycine solutions is highly sensitive to pH, as has been clearly demonstrated previously by Borsook and Thimann (4). Typical data over the p H region of 3 to 8 are illustrated in figure 4.In the most acid solution investigated, at pH 3.09, the spectrum is quite similar to that of an acetate complex both in the height and in the wave length of the peak. Severtheless, it seems unlikely that the cupric ion linkage is through the car-
+
boxyl group, for the electrostatic effect of the cationic -SH, substituent should be about the same as in betaine,4and hence practically no cupric ion should he in the complex at a glycine concentration of 0.02 M. The approximate equality in electrostatic effect of -XH3 and -S(CH3)3 is seen from a comparison of the change in pK of the carboxyl group of the corresponding substituted acetic acids. I n glycine pK, is 2.31 ( 7 ) and in betaine i t is 1.82 (lo), as compared to 4.758 (11) for the parent acetic acid. ApK thus is 2.42 and 2.94, respectively.
24
I. M. KLOTZ, I. LUCILLE FALLER, AND J. 41. URQUHART
As the pH is raised, the spectra of Cu++-glycine complexes shift to shorter wave lengths and to higher extinction coefficients and approach a limit of 6250 A. for A, and about 47 for emax. This shift parallels the ease of removal of a
+
hydrogen ion from -XH3 as the pH is increased. I t seems likely, therefore, that the cupric ion is bound in a complex a t least in part through the amine group. Once an -XH2 . . . Cu++ bond is formed, complex formation with the carboxyl
WAVELENGTH IN MILLIMICRONS
FIG.4 . Dependence of spectrum of copper-glycine complex on pH. The glyoine:Cu++ ratio is 2. The numbers shown are the pH’s as measured by a glass electrode.
group would be facilitated, since the anionic carboxyl would be attracted to the cationic metal. The importance of the anionic COO- group in glycine in stabilizing the cupric ion complex is seen clearly by comparison of the spectrum in the amino acid solution (figure 4) with that in a solution of glycine ethyl ester a t pH 4.50 (figure 5). In the latter case practically no cupric ion is bound in a complex, despite the
+
fact that the pK of the -XHs group is lower than that in glycine ( 7 ) and hence more favorable for the displacement of a hydrogen ion by a cupric ion. The entirely different behavior in glycine indicates clearly that the carboxyl group provides the added attraction necessary to bind the cupric ion.
COPPER-PROTEIN
25
COMPLEXES
+
The displacement of a hydrogen ion from the -NHS group in the binding of cupric ion by glycine is also indicated by the much less intense spectrum of cupric ions in hippurate solution, where the substituted nitrogen does not have much tendency to be an electron donor.
x
(my) FIG.5.Absorption spectra of cupric ion complex with glycine ethyl ester a t various pH’s. The glycine ester. Cu-+ ratio is 2 and the cupric-ion concentration near 0.01 K.The numbers shown are the pH’s as measured by a glass electrode.
137
AND
250
40
E 20
600
700 VtAVELENGTH
goo
900
IO00
IN MILLIMICRONS
FIG. 6. Dependence of spectrum of copper-glycine complex on concentration of the amino acid. The numbers shown are the ratio of glycine Cu++.The cupric-ion concentration was 0.01 M and the p H 4.0 in all cases
The effect of increasing glycine concentration at an acid pH, illustrated in figure 6, is very similar to that of increasing the pH The absorption maximum approaches 6500 -1.and an e of 47 as the glycine concentration is increased to 2.5 M . I t would seem, therefore, that the nature of the bonds in copperglycine complexes is the same a t pH 4 as a t 8, but that there is a lower ratio of glycine: cupric ion at the lower pH, a t equivalent glycine concentrations.
26
I. M. KLOTZ, I. LUCILLE FALLER, AND J. M. URQUH.4RT
C . Copper complexes with polypeptides The spectra a t pH 7.5 of copper complexes with a series of polyglycines, varying from diglycine to hexaglycine, have been compared with each other and
8
6
E 4
2
I
PO0
600
I
I
700
800
\(my) FIG.7 . Spectra of cupric-ion complexes with glycine and with polyglycines (di- to hexaglycine) a t p H 7.5. The zwitterion: Cut+ ratio was approximately 2.1 in each case, except in the case of pentaglycine where i t was 2.7 and hexaglycine where it was 2.4. T h e cupricion concentration was near 0.01 M in each case.
with glycine in figure 7 . It is of interest to note that all of the polyglycines show more intense absorption and have peaks a t lower wave lengths than does glycine. In the latter complex it is generally assumed that the molecule forms a chelate ring with cupric ion, as illustrated by formula I.
COPPER-PROTEIS
27
COMPLEXES
Such chelate ring format ion is unlikely with the larger molecules, since the large rings formed would be strained and relatively unstable. Kevertheless, it is apparent from the intensity and position of the Cu++-polyglycine spectra that the extent of coordination of the cation is a t least as great as in the monoglycine. I n fact, the spectra of the polyglycine complexes are very close to those of tetracoordinated amine complexes of copper (2). A tetracoordinated copper complex with a polyglycine, involving combined carboxyl and amine interaction, 11-ould be possible with a structure such as 11, 0
-C-O-
0
Hz
S-CH,
'I
CO( SHCHzCO),SHCH2C-O-
cu++
Hz S-
Cu++
I1 suggested to us by Professor F. Basolo (1). S o ring strain would be encountered in this type of structure, yet the extent of coordination of the central cupric ion mould be as great in these peptides as in glycine. With the peptides, furthermore,
+
the smaller pK of the -NHs group would make it easier to displace a hydrogen ion by a cupric ion; hence a greater degree of complex formation would be possible at comparable concentrations. The increase in complex formation would account for the higher optical absorption of the cupric ion in peptide solutions.
D. Copper complexes wzth proteins Preceding investigations (14) have demonstrated that cupric ion combines extensively n ith crystallized bovine serum albumin at pH's between 4 and 5 . Preliminary and incomplete spectra indicated that the carboxyl groups on the protein \yere involved in the linkage with cupric ions. A more detailed spectral examination has seemed desirable, however, particularly in view of the more extensive background data described in this paper. For this purpose, therefore, the absorption spectra of Cu++-albumin complexes were esamined in solutions of varying concentration of protein. The results with a 2 per cent albumin solution a t pH 4.5 are illustrated in figure 8. The shape of the curve is w r y similar to that of copper-acetate complexes (figure 1). I t is of interest, therefore, to compare the observed molecular extinction coefficient with the value which might be expected from the number of free -COO- groups from glutamate and aspartate residues in the protein. A 2 per cent solution of albumin (whose molecular weight is 70,000) corresponds to a 0.0003 Af solution. Since there are 93 free COO- groups per mole of albumin (FI), the carboxyl-ion concentration is approximately 0.03 -If. The observed
28
I. M. ICLOTZ, I . LUCILLE FALLER, AKD J. M . URQUHART
molecular extinction coefficient of 25 is quite close to the value one would expect for a 0.03 M solution of acetate ions, as estimated from the data in figure 1. Further evidence of the participation of carboxyl residues on the protein in the binding of copper below pH 5 is available from studies with chemically modified albumins. Thus the spectrum (figure 8) of cupric ion in a 2 per cent
X (mu) FIQ.8. Spectra of 0.01 >\f cupric-ion complexes with bovine serum albumin, parent and methylated, respectively. The spectra a t p H 4.5 were with 2 per cent albumin solutions, t h a t at p H 7 with a 9 8 per cent protein concentration The lowest w i v e is t h a t for 0.01 M cupric ion alone a t p H 4.5.
solution of an albumin sample whose COO- groups have been partially converted to -COOCH, groups is distinctly lower than that of the parent protein. On the other hand, an albumin in which the amino groups of the lysine residues had been acetylated showed no evidence of loss of ability to interact with cupric ion. In this case a spectrum could not be taken, however, for the acetylaminobovine albumin precipitated immediately upon the addition of cupric ions.
COPPER-PROTEIX
COMPLEXES
29
The behavior was very similar to that of polyglutamic acid in the presence of cupric ions. In both cases, apparently, complex formation between cupric ion and the free COO- groups removes the last polar residues in the macromolecule and hence insolubilizes the complex. I n a solution of p H 7 the effect of bovine albumin is distinctly different from that at pH 4-5. The absorption maximum of cupric ions (figure 8) is moved to 670 mp, i.e., toward the region for amine-type bonding. It Tvould seem, therefore, that at the higher pH, C u . . .SH*- linkages become possible in this protein, just as they are favored in the simpler molecules discussed earlier. Spectra at higher
X(m4) FIG.9. Spectra of 0.01 .%I cupric-ion complexes with &lactoglobulin (0.7per cent) at the pH's indicated. The lowest curve is t h a t for 0.01 .If cupric ion alone a t p H 4.5.
pH's have not been investigated as yet, but they would be expected to be moved toward even shorter wave lengths. h t a pH near 5, p-lactoglobulin (figure 9) and bovine serum y-globulin (figure 10) also produce changes in the spectrum of cupric ion nhich by their appearance are consistent with a copper-carboxyl type of linkage. The 7-globulin has the smallest effect of the three natural proteins examined. I t is of interest to note in this connection that the content of free COO- residues is also lowest, per unit weight, in this protein (5). The spectrum of copper-p-lactoglobulin complexes has been examined also at a higher pH, 9.8. The shift in wave length of the maximum (figure 9) is quite pronounced, and its position at 580 mp and molecular extinction coefficient of nearly 60 are quite characteristic of the results obtained with copper complexes with simple amines and Tvith polypeptides. There is thus very strong evidence
30
I. M. KLOTZ, I. LUCILLE FALLER, .4ND J. M. URQUH.4RT
+
that a t pH’s which are sufficiently high for the -XHs groups to lose their positive charge, cupric ion will form a complex with amine groups on the protein. Similar results have been observed a t pH’s near 10 with @-casein(figure 10). Again the absorption maximum is in a region characteristic of the copperamine type of linkage. The result with casein is particularly interesting, since this protein has a very low cysteine (or cystine) content, and hence practically no available sulfhydryl groups. I t is apparent, therefore, that extensive copper
E
I
500
600
100
800
FIG.10. Spectra of 0.01 11.I cupric-ion complexes with bovine serum 7-globulin (1per cent, pH 4.5) and with 8-casein (0.2 per cent, p H 9.9). The lowest curve is t h a t for 0.01 M cupric ion alone at pH 4.5.
binding can occur with proteins with linkages that do not involve -SH groups. With this backlog of spectroscopic information it is of interest to examine a few of the naturally occurring copper proteins. The spectrum of hemocyanin, in the serum of the horseshoe crab, is illustrated in figure 11. The high molecular weight of this protein leads to much scattering a t shorter wave lengths, and hence a steep rise in absorption is observed near 400 mG. There is no doubt, however, that the cupric ion absorption peak occurs a t 580 mp. Similarly, in the coppercontaining serum protein described by Holmberg and Laurel1 (12), the absorption peak occurs a t 605 mw. In both cases the position of the maximum is in the region
COPPER-PROTEIN
COMPLEXES
31
of tetracoordinated copper complexes such as those illustrated in several of the figures in this paper. I n hemocyanin, indirect evidence has been cited (8) that the copper is in the univalent state. I n view of the similarity in spectrum of this copper protein with others described in this paper, it may be worthwhile to reexamine the question of the valence of copper in oxyhemocyanin. From the data presented it is apparent that the absorption peak of cupric ion in complexes shifts toward shorter wave lengths as the coordination number is
A (my, FIG.11. Spectrum of' serum of Liniirlus polyphenrzcs, diluted 1 : 5 , a t pH 7.7 I\'.
coh~cLusIoss
increased. Tetracoordinated copper, even if two of the ligtnds are carboxyl groups, exhibits a spectrum in the neighborhood of 6000 A. Both naturally occurrin4 and artificially constructed copper proteins also absorb in the region of GOO0 A., if the pH of the solution is above 7 . I t seems likely, therefore, that in hoth types of protein the cupric ion is tetracoordinated, and that a t least two of the bonds are to amine groups in the protein molecule. 1he specificity of copper proteins may then arise from a t least two configurational factors. In the metal enzymes, the type of substrate which is susceptible to the metalloprotein may be determined by the arrangement of residues in the neighborhood of bound cupric ion. The specific arrangement of these groups would determine whether the substrate could be bound by the protein a t a posir 7
32
I. Y. KLOTZ, I. LUCILLE FALLER, 4XD J. M. URQUHART
tion adjacent to the cation. Another configurational factor, which may be particularly important in hemocyanin, may lie in the relative positions of bound copper atoms. It has been well established (8) that in oxyhemocyanin the unit weight of protein which combines with one molecule of oxygen contains two copper atoms. I t seems quite possible, therefore, that the oxygen molecule may act as a bridge between two copper ions, if the distance between the metals is properly fixed. Since cupric ion can have a coordination number of five, the linkage to the oxygen may supply the fifth bond. At present, however, such a structural picture lies largely in the realm of speculation. V.
SCMM.4RY
Optical absorptions in the region of 400-1000 m p have been measured for cupric ion complexes n.ith several simple, substituted carboxylate ions, including acetate, glycolate, citrate, hippurate, and betaine. These have been used to evaluate the effects Lyhich various substituents on the a-carbon atom have on the stability of the copper-carboxyl bond. Spectra of cupric ion complexes n-ith glycine ethyl ester have also been obtained as an example of an amine-type combination. The effects of glycine, diglycine, triglycine, tetraglycine, pentaglycine, and hexaglycine, respectively, on the absorption of cupric ions have also been measured. These have been used as a basis for the interpretation of the observed spectra of copper complexes lvith serum albumin, serum y-globulin, P-lactoglobulin, and @-casein,respectively, and of hemocyanin. All copper proteins a t pH’s above 7 have been found to absorb light in the region of 600 mp, a behavior which indicates that the copper is tetracoordinated and that at least two bonds are of the amine type. REFERENCES (1) BASOLO,F.: Private communication. (2) BJERRUM,J., AND NIELSEN,E . J.: Acta Chem. Scand. 2, 297 (1948). (3) BJERRUJI,K.:2 . physik. Chem. 106, 219 (1923). H . , AND THIMAKN, K . V.: J. Biol. Chem. 98, 671 (1932). (4) BORSOOK, ( 5 ) BRAND,E.: Ann. K.Y. Acad. Sci. 47, 187 (1946). (6) CASNAN,R . K., A N D KIBRICK,A , : J. Am. Cheni. SOC.60, 2314 (1948). J . T . : Proteins,Amino Acids and Peptides, pp. 84, 99. Rein(7) COHS,E. J., AND EDSALL, hold Publishing Corporation, S e w York (1943). (8) DAWSON, C. R . , ASD MALLETTE, M . F . : Advances in Protein Chem. 2, 179 (1945). H . , . ~ N D OLCOTT,H. 6 . : J . Biol. Chem. 161, 259 (1945). (9) FRAENKEL-CONRAT, (10) GCSTAFSOK, C . : Ber. 77B, 66 (1944). (11) HARKED,H . S., AND EHLERS,R . W . : J. Am. Chem. SOC.6 6 , 652 (1933). C. G., AND LAURELL, C. B . : Acta Chem. Scand. 2, 550 (1948). (12) HOLYBERQ, J. G., AND WESTHEIJIER, F. H . : J. Chem. Phys. 6, 506 (1938). (13) KIRKWOOD, (14) KLOTZ,I. M,, .4ND CURME,H. G . : J. Am. Chem. SOC. 70, 939 (1948). (15) KLOTZ,I. M., SCHLESINGER, A . H . , ASD TIETZE,F.: Biol. Bull. 94, 40 (1948). A , : Proc. Roy. Soc. (London) A168, 68 (1937). (16) NEUBERGER, K . J . : Kgl. Danske Videnskab. Selskab, Mat.-fys. Medd. 23, KO.12 (1945). (17) PEDERSEN, (18) SCHOTT, H. F., LARKIN,J. B., ROCKLhND, L. B., AND DUNK,M. S.: J. Org. Chem. 12, 490 (1947).
