11ss
the addition of Cu(hn)+?,curve B, the titration is E. Potentiometric Titrations.-Titrations in nearly identical with the one for copper(I1) ion ex- which base was added to acidified solutions of the cept that no precipitate is formed. The curvature amine, Cu(hn)+2and Cu(hn)z+2 were used to calin the vicinity of the end-point of curve B indicates culate the acid dissociation constants of the amine incomplete reaction a t the stoichiometric end-point. salt, the stability constants of the complexes and In this reaction the complex [Cu(hen-0-) (OH)]O the acid dissociation constants of the complexes. is evidently formed in one step. The reaction be- These are being reported in detail in a further pubtween Cu(hn)?tZ and hydroxyl ion, curve C, shows lication.q For the formation o f Cu(hn)+?, the no definite end-point and a comparison of this value of log K1 is 10.11. For the formation of curve with curves A and B indicates a total reac- Cu(hn)?l2from Cu(hn)+?and hn, the value of log tion of only one mole of hydroxyl ion per mole of K S is 7.31. For the dissociation of the proton Cu(hn)z+z. The initial slope, however, indicates from Cu(hn)+?,pK, is ‘7.21. These results are all that in the large excess of base two equivalents of in 0..5 I1P potassium nitrate. hydroxyl ion per mole of Cu(hn)z12are used. This Summary.--The chart shown in Fig. 6 gives a is additional confirmation that in great excess of scheme which is consistent with all of the data prestrong base, one mole of amine is displaced from sented here. The formulas designated V and VI the 1 : 2 complex. The curve E for addition of are drawn to include the assumption of five-coordiCu(en)z+2 indicates negligible reaction of this nate copper(I1) ion but are not intended to be incomplex with the hydroxyl ion a t this concentra- dicative of the structure of this ion. tion. The curve D for the addition of the mixed Acknowledgment.-The spectra reported here complexes is evidently a n average of curves C and were determined by Linda S. Gallo.
E.
MORGAUTOWU, I T 17.4
O F BIOCHEMISTRL , CURSELL v S I V E R S I T Y RESEARCH, EQUITABLE LIFE .aSSURANCE SOCIETV O F
[ ~ O Y T R I B U T I O YFROM THE DEPARTMEST hlFDlCAL
hfEDICAL COLLEGE AND THE B U R E A U O F THE [TNITED STATES]
Coordination Complexes and Catalytic Properties of Proteins and Related Substances. I. Effect of Cupric and Zinc Ions on the Hydrolysis of p-Nitrophenyl Acetate by Imidazole BY WALTERL. KOLTLY,RICHARD N. DEXTER,’? K I C I L ~ KE.D C I , . ~ R K . ~:N ~ DFRANK R. K. GURD RECEIVED FEBRUARY 8, 1958 The advantages of combining different methods, particularly kinetic and equilibrium techniques, for assessing the rcactivity of polar groups in proteins are described. The model system containing cupric and zinc chlorides, imidazole and imidazolium chloride has been studied simultaneously by two independent techniques that measure the concentration of free basic imidazole quantitatively. The first technique depends on the catalysis by imidazole of the hydrolysis of p-nitrophenyl acetate (NPA4);the second technique involves measurement of the pH of the solution. Over a wide range of conditions the two techniques of measurement do not interfere with one another. The pnitrophenolate ion released on hydrolysis of XPA combines with zinc and cupric ions about 100 times less strongly than does imidazole. Previous observations on the rate of decomposition of N-acetylimidazole are confirmed, and zinc ions are shown to have no detectable effect on the rate of decomposition. The method of Scatchard is applied to the computation of the successive association constants in the Cu(I1)- a n d Zn(1Ibimidazole qystems, with results that differ only slightly from those reported previously.
Detailed knowledge of the reactivity of the polar side-chain and terminal groups of proteins is necessary for an understanding of their structure and function. At least three general approaches may be followed for assessing the reactivity of a group or class of groups : (1) the preparation of stable derivatives under defined conditions; ( 2 ) the measurement of equilibria in which the group in question takes part, for example, equilibria with cations; (3) the measurement of the catalytic activity of the group in question. The third approach has been pursued least vigorously, except for studies specifically dealing with the catalytic properties of the highly specialized “active centers” in enzymes. We have undertaken an extensive investigation of the applica(1) This investigation was supported by research g r a n t KO. H 27739 f r o m t h e National H e a r t Institute, U. S. Public H e a l t h Service ( 2 ) Lederle Medical Student Research Fellow, Cornell University Medical College, 1957. ( 3 ) Holder of a Summer Research Fellowship under a g r a n t from t h e U. S. Public H e a l t h Service t o Cornell University Medical College, 1957.
bility of the catalytic method €or the general assessment o f the reactivity of groups in proteins Knowledge of the catalytic potentialities of a series of model compounds containing these groups may lead also to a better understanding of enzymatic activity. The catalytic activities of proteins are not necessarily limited to well-recognized enzymic functions. We may expect that the side-chain and terminal groups in proteins may be able to catalyze reactions according t o the characteristics of comparable groups in small molecules. Presumably these catalytic effects may be assignable to the action of individual groups or of clusters of groups acting together. The catalytic approach cannot be expected to circumvent the problem of the overlapping reactivities of the various classes of groups in proteins; for example, many of these types of groups are potentially active in general basic catalysis. Therefore, to be exploited most effectively the catalytic approach should be used in conjunction with the
Aug. 20, 19%
COORDINATION COMPLEXES A N D CAT.\LYTIC PROPERTIES OF ISIID.\ZOLE
other measures of reactivity that have been mentioned. I n principle the simultaneous application of kinetic and equilibrium measurements should have particular advantages. First, these two techniques are the best suited for avoiding permanent changes in the properties of a p r ~ t e i n . ~ Second, the two separate types of measurement may depend on somewhat different facets of the reactivities of the groups in a protein. Consequently, if both kinetic and equilibrium measurements are made on the same series of mixtures of protein and metal ion, then the resulting information about metal binding, p H and catalytic activity should be sufficient t o define the reactivities of the several classes of groups with much greater certainty than could be expected from either type of measurement alone. Clearly, any such interpretation of combinations of measurements depends on the lack of interference of one technique with the other. Before attempting to test this combined approach on proteins we have studied a model system containing imidazole, p-nitrophenyl acetate (NPA) and zinc chloride or cupric chloride. The catalysis by imidazole of the hydrolysis of NPX has been studied in detailj6-7 and the formation of complexes of imidazole with Zn(I1) and Cu(1I) has been studied by the measurement of p H values in mixtures a t equilibrium.8 The test of compatibility of the two approaches is to show that (1) the presence of these metal ions does not upset the linear relation between catalytic activity and concentration of the free basic imidazole, .and ( 2 ) the presence of NPA or its reaction products does not upset the measurement of complex formation by the pH method. It will be shown below that these two tests of compatibility are met successfully in this system. Materials and Methods Materials.-Imidazole (m.p. 89.5-90.5'; lit. 89.5-90O)g obtained from Edcan Laboratories, South Norwalk, Connecticut, was used without further purification, after drying for several days over PtOj. Imidazole hydrochloride was prepared by dissolving 1.382 g. of imidazole in 100 ml. of 0.203 N HC1. The p H was adjusted to 3.86-3.90 by the addition of approximately 0.2 ml. of 1 N HCl. The theoretical pH of a 0.203 A4 solution of imidazole hydrochloride (PK' = 7.08) is 3.89. p-Xtrophenyl acetate (XPA) was prepared by the method of Spasovia as described by Bruice and Schmir.6 The NPA was dissolved in 95% ethanol and diluted with water to give a final concentration of 1.0 X 10-3 A4 in 1.90% ethanol-water (v./v.). A fresh solution was prepared dnily. Cupric chloride solutions were standardized iodometrially.^^ Zinc chloride solutions were standardized by backtitration according to Biedermann and Schwarzenbach,i* using ethylenediamine tetraacetate (EDTA). -~
(4) F. R. K. Gurd a n d P. E. Wilcox, Advances i n Pyolein Chem., 11, 311 (1956). ( 6 ) &.I.L. Bender and B. W. Turnquest. THISJ O U R N A L . 79, 1662 ( 1 857). ( 6 ) T.C. Bruice a n d G. L. Schmir, ibid., 79, 1663 (1957). (7) T. C. Bruice a n d G. L. Schmir, ibid., 80, 148 (1958). (8) J. T. Edsall, G. Felsenfeld, D. S. Goodman and F . R. N. Gurd, rhid., '76, 3054 (1954). (9) K. Hofmann, "Imidazole a n d its Derivatives," P a r t I, Interscience Publishers, Inc., N e w York, S. Y.,1963, p. 6. (10) A. Spasov, A n n . Uniu. SoJra I I . Faculte P h y s . Math., Livre 11, 36, 289 (1938-1939); C. A , , 3 4 , 2343 (1940). (11) F. P . Treadwell and W T. Hall, "Analytical Chemistry," 9 t h E d . , John Tl'iley and Sons. Inc. S e w York, PI' Y . . 1942, p. 610. (1%) IV. Biedermann a n d C . Schxvarzenbach. C:72mia, a. .iR il848).
1180
Kinetic Method.-The following general procedure was used in the kinetic determinations. Two tubes were prepared containing identical mixtmes of imidazole, imidazolium chloride, metal salt (as required) and sufficient NaCl to maintain a constant ionic strength of 0.16 during the reaction. The two tubes, along with a solution of NPA we;e equilibrated in a constant temperature bath a t 25.1 f 0.1 . To one of the tubes was added water (1 part per 9 parts by volume), then the solution was mixed and transferred to a cuvette which was placed in the spectrophotometer as a blank solution. To the second tube was then added the same volume of the NPA solution. This addition was made rapidly from a calibrated syringe and the contents of the reaction tube were thoroughly mixed, transferred to a cuvette and quickly placed in the spectrophotometer. The first reading usually was obtained within 1 minute; during the faster reactions a reading was taken after 30 seconds. Unless specified otherwise, the temperature of the thermostated cell compartment of the Beckman Model DU Spectrophotometer was maintained a t 25.5 10.5" by circulating water from a constant temperature bath passing through dual thermospacers. The pH values of both the blank solution and reaction mixture were measured at the conclusion of the kinetic run, after the lapse of at least 8 half-lives. The rate of hydrolysis was determined by following either the rate of appearance of p-nitrophenolate ion (NP-) a t 400 mp, or the rate of disappearance of NPA a t 273 mp. In following the rate of appearance of NP- it is necessary to maintain a constant concentration of hydrogen ion since the absorption at 400 mfi measures only the N P - in an equilibrium mixture with p-nitrophenol. The imidazole catalyst itself was always in sufficient concentration to maintain a constant pH during the reaction. NPA and NPwere shown to follow Beer's law over the ranges of concentration used. The first-order rate constants kobs were obtained from the rate of appearance of NP- using the relation
and from the disappearance of NPA using the relation
where t = reaction time in minutes; 0.D.o is the optical density a t t = 0, which was obtained by extrapolation; O.D. 0: is the optical density a t infinite time, determined after a reaction time of a t least 8 half-lives had elapsed; and 0.D.t is the optical density a t any time after the start of the reaction. Normally the reactions were followed continuously to 50% of completion. For reactions having half-lives greater than 6 hours the 0 . D . a was taken from standard absorption curves of p-nitrophenol, and the reaction was followed to 20% of completion. Measurement of pH.-A Radiometer Titrator Type T T T l a obtained from Welwyn International, 3355 Edgecliff Terrace, Cleveland 11, Ohio, was used. Some of the earlier measurements were carried out on solutions a t room temperature (27-30") ; however most solutions were thermostated in a water-jacketed cell a t 25.1 & 0.1". Unless stated otherwise all measurements reported in the tables were made a t 25.1 10.1'. The PH measurements were regularly reproducible to within 1 0 . 0 1 unit. The instrument was calibrated before and after each measurement by using standard Beckman buffer solutions of p H 4, 7 and 10. All PH readings were corrected by linear interpolation.
Results Catalysis by Imidazole.-The observed firstorder rate constants were taken to be equal to kl k,, where kl is the first-order rate constant for catalysis by basic imidazole of the hydrolysis of NPA and k , the rate constant for hydrolysis in a solution a t the same pH, temperature and ionic strength in the absence of imidazole. Values of k , were computed as follows. Rates of hydrolysis were measured in the presence of a series of acetate, phosphate and borate buffers containing sufficient NaC1 to maintain the ionic strength a t 0.16. For
+
4100
U-. L. KOLTUN, R. N. DEXTER,R. E. CLARKAND F. R. N. GT-RD
Vol. 80
the more alkaline buffers the rates were measured by following the appearance of p-nitrophenolate (NP-) a t 400 mp. The observed rates were plotted as a function of p H for each ionic strength. Then for each of a series of p H values the interpolated rates were plotted vs. ionic strength due to buffer and k, was obtained by extrapolation to zero ionic strength. Near PH 8.0 the measurements were tnade in both phosphate and borate buffers and led to practically identical values of k,. The results for the pH range 4.5 to 9.0 are shown in Fig. 1 in which k, is shown as a function of pH. The inset in Fig. 1 shows the linear relation between k, and the concentration of hydroxyl ion over the range of the latter between 1 X and 60 X AI. In none of the cases reported in this study did k, exceed 5% of the value of kl. 2.0
4.0 IMIDAZOLE
I2'O
I -
6.0
CONGENTHATION,
8.0
lC.0
M x IO3
Fig. 2.-Dependence of kl on concentration of imidazole. The rectangles indicate thr probable errors of the measurements.
the second-order rate constant, k 2 . I 3 The value of kr obtained here a t 25.5 0.3" and ionic strength 0.16 in 0.2% ethanol is 31.3 l./mole min. and may be compared with the values obtained by other workers: 12.7 l./mole min. a t 25" and ionic strength 20 40 60 0.016 in 28.5% eth.ano1(v./v.)~; and approximately [OH-] IO' M 28.2-29.0 l./mole min. a t 26.2" and ionic strength 0.002-0.08 in 5y0 dioxane-water.h Presumably these differences may be ascribed to the differences in solvents used. The relation between k1 and (Im) shown in Fig. 2 is used below to estimate values of (Im) from rate measurements in the presence of Cu(I1) and Zn(I1). During reactions with the longest half-lives, corresponding to (Im) below 0.0004 AI, the temperature showed a tendency to rise. Temperatures ranged between 26.0 and 26.5' instead of 25.2 and 25.8'. This effect introduced a very slight curvature (corresponding to an increase in kz of 8-10%) Fig. 1.-Dependence of k , on PH, 0 . The open ciicle into the lowest extreme of the standard curve, not shows the rate constant measured in the presence of 0.01 M shown in Fig. 2. This dependence of kz on temperaCuC12 or ZnClt. Inset: dependence of k , on (OH-). ture is in keeping with the observations of Bruice and Schmir.6 Values of kl obtained over a wide range of conMeasurement of Free Basic Imidazole in the centrations of imidazole and of imidazolium chlo- Presence of Cu(I1) and Zn(I1): (A) Kinetic ride are shown in Fig. 2 plotted against the con- Measurements.-The finding of a constant value centration of basic imidazole, (Im) , The measure- of kz a t constant temperature and ionic strength ments were made by following the disappearance of in the face of wide variations in the concentration NPA a t 273 mp. The concentration of imidazole of the imidazolium ion confirms that the latter does ranged from 0.000084 to 0.0090 41,and of imidazo- not act as a catalyst in this s y ~ t e m . ~Accordingly, ?~ lium chloride from 0.0180 to 0.0900 M . By con- it is to be expected that imidazole combined in comtrast the initial concentration of NPA was 1.0 X plexes with other cations such as Cu(I1) and Zn(I1) lo-* M , so that the quantity of acid released during will not act catalytically, and that kinetic measurethe hydrolysis was almost always much less than ments on mixtures containing such metal ions may the concentration of basic imidazole, and could not he correlated solely with the concentration of free be expected to affect the p H of the imidazole buff- basic imidazole, (Im). Table IA shows the results of ers by as much as 0.01 unit. This prediction was kinetic measurements performed on Cu(I1)-imiverified experimentally. One exception was ob- dazole mixtures in the manner already described. served a t the extreme condition of lowest (Im), Besides imidazole, imidazolium chloride and NaC1, 0.000084 &I. Here the p H of the control was 4.84 the solutions contained Cu(I1) chloride in the whereas a t the end of the experiment i t was 4.75. quantities listed. Corresponding results for mixIn a case of this sort the initial rate and initial p H (13) Over t h e g H range covered in this s t u d y the proportion of must be used for the computations. imidazole in t h e anionic form is too small t o make a detectable conT h e dope of the line in Fig. 2 may be defined as tribution t o t h e catalysis.? X
PH
Aug. 20, 19%
COORDINATION COMPLEXES AND CATALYTIC PROPERTIES OF IMIDAZOLE 4191 TABLE I CATALYSIS OF HYDROLYSIS OF NPA BY IMIDAZOLE IN THE PRESENCE OF Cu(I1) PB.
--Composition------Initial total molar concn. of CUCI, HIm+CIIm
0.01793 .00896 .00896 .00896 .00896 .00896 .00896 .00896 .00896 .00896 .000944 .000944 .000896 .000896 .000896 .000896 .000896 .000944 .000893
0.0180 .OB0 .0180 .OM0 .0180
0.00360 .00360 ,00540 .00720 .00900 .01080 .0 1260 .01440 .01620 .01800 .000946 .00237 .00360 .00405 .00450 .00540 ,00720 .00142 .01125
.O l s o .OB0 .0180 .OM0 .0180 ,01923 ,01923 .01827 .01827 .OM27 .01827 .OM27 .01923 .0180
-A. kl
x
1 0 :
mln. -
0.103 ,162 ,212 ,346 .497 .610 .880 1.26 1.76 2.20 0.426 2.14 4.34 5.83 6.84 9.35 14.5 0.895 25.2
Kinetic results-log Um)
4.65 4.39 4.25 4.02 3.85 3.76 3.60 3.44 3.29 3.16 3.93 3.18 2.86 2.75 2.66 2.53 2.34 3.59 2.09
Equilibrium resultp13 of blank -log soln. Um)
Y
fiH
0.200 ,397 .596 .793 .989 1.19 1.38 1.52 1.75 1.93 0.878 1.80 2.46 2.51 2.60 2.72 2.90 1.28 3.57
3.99 4.33 4.61 4.81 4.96 5.12 5.24 5.37 5.53 5.67 4.85 5.57 5.98 6.05 6.14 6.26 6.47 5.14 6.73
3.99 4.33 4.61 4.81 4.96 5.12 5.26 5.37 5.53 5.67
4.84 4.50 4.22 4.00 3.87 3.71 3.59 3.45 3.29 3.15 3.95 3.23 2.84 2.77 2.68 2.56 2.35 3.66 2.10
..
5.58 6.00 6.07 6.15 6.27 6.47 5.16 6.73
Y
0.202 .398 ,595 ,792 ,99 1.18 1.38 1.57 1.75 1.93 0.884 1.88 2.40 2.62 2.68 2.94 3.03 1.27 3.60
TABLE I1 CATALYSIS OF HYDROLYSIS OF NPA BY IMIDAZOLE IN THE PRESENCE OF Zn(I1) -Initial
ZnClt
a
Composition total molar concn. ofHIm T I
-
0.0531 0.00203 .01015 .0496 .0496 .01015 .00294 .0203 .00196 .0203 .01015 .0496 .01015 ,0248 .00294 .0203 .00196 .Os03 .00098 .0203 .00294 ,0203 ,00294 .0203 .00245 .0203 .00196 .0203 .0203 .00196 ,00147 .0203 .00098 .0203 .00098 ,0203 .00049 .0203 Measured at 27'; pK' 7.04.
-A.
Im
AI X I!', min.
Kinetic results-log (Im)
4.34 0.184 2.00 3.24 4.14 2.90 3.92 2.93 2.85 4.72 5.21 2.84 6.31 2.72 7.62 2.63 9.06 2.54 11.9 2.43 14.8 2.33 15.0 2.32 15.9 2.29 18.6 2.23 .OlOO 18.4 2.23 .0100 20.9 2.18 .OlOO 24.0 2.11 23.4 .OlOO 2.13 .OlOO 26.9 2.07 Measured at 30°; PK' 6.98.
0.00100 .OlOO .0200 .00250 .00250 .0300 .0200 .00500 .00500 ,00500 .0100 .OlOO .0100 .OlOO
*
tures containing Zn(I1) chloride are shown in Table M NPA was initially IIA. As before, 1 X present during the measurements. The kinetics of the disappearance of NPA were followed a t the wave length of 273 mp. The values of k1 listed were obtained after taking appropriate values of k , from Fig. 1, using the PH values determined as described below. That k, is unaffected b y the presence of Cu(I1) or Zn(I1) was shown by measurements in acetate buffers similar Zinc to those used for the determination of k,. chloride or cupric chloride (0.01 M ) was substituted for part of the NaCl in the buffers containing 0.015 M sodium acetate, total ionic strength 0.16. The observed rates of hydrolysis of NPA were indistinguishable from those found in the absence of di-
-B.
-Y
BH
0.018 * 190 .378 .449 ,551 .575 .729 .898 1.06 1.33 1.80 1.77 2.00 2.07 2.10 2.26 2.36 2.56 2.84
5.45" 5.75" 6.01b 5.87 5.93 6.17" 6.27" 6.15 6.22 6.32 6.43 6.43 6.48 6.53 6.52 6.58 6.63 6.64 6.TO
Equilibrium results-log (I=)
4.28 3.28 2.96 2.90 2.84 2.86 2.76 2.62 2.55 2.45 2.34 2.34 2.29 2.24 2.25 2.20 2.14 2.13 2.07
0.018
.191 .381 .a5 .641 ,576 .737 .a91 1.12 1.49 1.85 1.85 2.00 2.18 2.24 2.51 2.86 2.68 3.14
valent metal ion. The open circle in Fig. 1 shows rates obtained in the presence of Cu(I1) or Zn(I1). (B) Equilibrium Measurements.-After the kinetic measurements had been followed to completion, the solutions were transferred to thermostated vessels for measurement of the equilibrium p H values a t 25.1 f 0.1". The results are listed in Tables I B and IIB. These p H values represent a second measure of the concentration of free basic imidazole: taken together with the concentrations of imidazolium chloride used and the measured value of pK' for imidazole under these conditions (7.08), they yield values for (Im) in each solution studied.8 The accuracy of the values for the concentrations of free imidazole derived from the p H measure-
4192
IV. I,.
I.;OI.TUN. K.
N.DEXTER. I
cVI
z W
n -I
4
0 t-
n
0 IO
50
30
70
REACTIONTIME, MINUTES, Fig. 3.-Course of formation and decomposition of Nacetylimidazole at p H 6.09, 6.77 and 7.62 as showri by a plot of optical density a t 245 mp vs. reaction time in minutcq.
L
I
Fig. 4.-Dependence of rate constant for the decomposition of AT-acetylimidazoleon pH. The open circle' represent measurements made in the presence of 0.025 ,V ZnCI2.
computations of the values of the successive association constants in the Cu(1I)- and Zn(I1)-imidazole systems by the method of Scatchard used prev i o ~ s l y . ~ In ~ ~Tables * I and I1 are listed the computed values of D, the average number of imidazole molecules bound per mole of metal ion present. Values derived from each method of measurement are listed in the appropriate parts of the tables. The data in Tables I and I1 have been analyzed by the method of Scatchard from plots of log Q os. D where Q = ;/(4
- C)(Im)
These plots are shown in Fig. 5. Points derived from the equilibrium measurements are shown in full circles, those from kinetic measurements in open circles. In Fig. 5a are shown the points for the system containing Cu(II), and in Fig. 5b those for the system containing Zn(I1). The curves were constructed using the log K values shown in Table 111. These values were chosen to fit the points derived from PH measurements because the latter were made under conditions of better control of temperature than proved possible with the kinetic measurements. For example, the kinetic meas-
2
I
PH
-
3
4
L'
Fig. 5.- -Plots of log Q 2's. D; the curves tire computed using the values of K listed in Table I11 and are intended to fit most closely the points derived from p H measurements at 25.1' and ionic strength 0.16, shown in closed circles, 0 . Points derived from kinetic measurements are shown in open circles, 0: (a) CliiII)-imidazole, (b) ZnlI1)imidazole.
urements on the Zn(I1)-system were made a t 26.0-26.5' instead of a t 25.2-25.8", and the slightly low values of log Q in Fig. 5b are the expected consequence. The effect of temperature on TABLE I11 COMPUTED INTRISSIC ASSOCIATION CONSTANTS FOR Cu( I1 ) AND Zn(I1) WITH IMIDAZOLE Log ti values for ionic strength 0.16 and 25'. log K1 log ti2 log Ka log K4
Cu(I1) complcxes
Zn(I1) complexes
3.60 3.24 8.06 2.65
1 92 2.14 2.50 2 65
the kinetic measurements per se was taken into account by using the standard curve made under comparable conditions, but of course the association constants for the Zn(I1)-imidazole complexes are slightly lower a t the higher temperature.* Using
4104
Yol.
P A S C H O A L S E N I S E A N D L X A D E L E I N E PERRIER
the usual statistical relations4,* k, = 4x1; k2 = 3~2/2; k, = 2~s/3; k4
~ q / i
the log K values have been converted to the log k values listed in Table IV. Included for comparison are log k values interpolated from those computed previously.l 4 T A B L E 1V SUCCESSIVE ASSOCIATIONCONSTAXTS FOR Cu(I1) Zn(I1) WITH IMIDAZOLE Log k values for ionic strength 0.16 and 25'.
Cu(I1) complexes FormerPresent
4.33 log ki 4.20 3.54 log ks 3.42 2.82 2.88 log ka 2.03 2.05 log kr " Y . Nozdki, F. R. N. Gurd, R. F. sall, THISJOURNAL, 79, 2123 (1957).
AND
Zn(IT) complexes Former" Present
2.57 2.52 2.32 2.36 2.22 2.32 2.01 2.05 Chen and J. T. Ed-
The method of Scatchard for estimating successive constants was chosen in preference to that of Bjerrum because i t yields a value of kl that is probably more precise. If this type of study is to be useful as a model for the reactivity of imidazole groups in proteins it is usually the value of kl that must be estimated most a ~ c u r a t e l y . ~The present values for klare probably more accurate than those previously reported, but the differences are not marked. Comment has been made previously on the errors involved in arriving a t the other successive constant^.^*'^ I n the present study the higher values of v have not been explored because the technique that we have used does not allow accurate measurements of the kinetics when the half-life approaches one or two minutes. The value of is therefore rather uncertain. Nevertheless, the logarithm of the product of the constants for the Cu(I1)-system agrees well with the value of 12.6 determined polarographically by Li, White and Doody.lS (18) X. C. Li, J. M. White and E. Doody, THISJOURNAL, '76, 6219 (1954).
[CONTRIBUTION FROM
so
Discussion The foregoing results show that the kinetic and equilibrium methods of measuring the concentration of basic iniidazole are mutually compatible. The presence of Cu(I1) or Zn(I1) ions does not affect the kinetic measurements] and the presence of KPA or its hydrolysis products does not affect the equilibrium measurements detectably. Studies to be reported later have shown a similar compatibility of the measurements when the imidazole radical is incorporated in the histidyl residue of a peptide. In terms of their potential role in assessing the reactivity of imidazole groups in peptides and proteins, the two methods show an interesting contrast in characteristics. The kinetic method depends for its usefulness on the fact that it does not alter a condition of equilibrium, whereas the equilibrium method depends on the alteration of such a condition of equilibrium and on our ability to describe the new multiple equilibrium quantitatively in terms of the mass law. The nucleophilic properties of the basic imidazole group resemble those of classes of groups in proteins, such as amino and phenoxy1 group^,^^^^^ and probably others. It is to be anticipated, therefore, that other classes of groups in proteins in addition to imidazole groups may react with NPA. IVhether or not the products of the reaction are stable or unstable]the combined approach described here should be directly applicable whenever relatively minute quantities of NPA are employed. Acknowledgments.-The authors are grateful to Dr. Thomas C. Bruice for a gift of 9-nitrophenyl acetate and for making available unpublished results. Dr. George Scatchard was kind enough to offer suggestions concerning the manuscript. The technical assistance of Miss Reta Roth is gratefully acknowledged. (19) F. W. P u t n a m , in "The Proteins," (H. Neurath and K. Bailey, eds.). Vol. IB, Academic Press, New York. N. Y., 1953, P. 893. (20) W. L. Koltun, THISJOURNAL, 79, 5681 (1957).
NEW YORK,N. Y .
DEPARTAMENTO DE QUfMICA, FACULDADE DE FILOSOFIA, CIeNCIAS S i 0 PAULO]
E
LETRASDA
UNIVERSIDADE DE
Spectrophotometric Investigation in the Near Ultraviolet of the Cobalt(I1) Monothiocyanato Complex B Y P.4SCHOAL S E N I S E A N D
MADELEINE PERRIER
RECEIVED MARCH19, 1958 Use has been made of the absorption band in the near ultraviolet of the CoSCN+ ion to determine the formation constant of this ion in acid aqueous solutions of unit ionic strength, by two independent methods. At pH 3.0 and 25 f 1 the average value of that constant obtained by the method of McConnell and Davidson was found to be 10.09 f 0.12 and by the method of corresponding solutions of Bjerrum 10.38 =k 0.07.
Introduction The nature of aqueous solutions containing eobaltous and thiocyanate ions has been investigated spectrophotometrically by several authors.'-4 (1) A. v. Kiss and P. CsokZln z. physik. Chcm., ~ 1 8 6 , 2 3 9(1940). (2) M. LehnC, BuZl. SOC. chem. France, 76 (1951). (3) A. K. Babko and 0.F. Drako, Zhur. Obshchci Khim., SO, 228 (1950);C . A . , 44, 5G84 (1950). (4) P. W. West and G. F. Vries, Anal. Chrm., 43, 334 (1951).
Consecutive formation of complexes was found to occur, the first species formed being the cation ~ J S C N ' . The value of the formation constant of this ion reported by LehnP is very uncertainas stated by the author-since its calculation was based on the rather small difference of the molar extinction coefficients of the complex and cobalt(I1) ions in the visible range of the spectrum. A dif-
