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The effect of pH on the interaction of enolase with activating metal ions
ARCHIVES
OF
BIOCHEMISTRY
AND
BIOPHYSICS
61,
186-196
(1956)
The Effect of pH on the Interaction of Enolase with Activating Metal Ions Bo G. Malmstrijm From
the Institute
and Lars E. Westlund
of Biochemistry,
Received
September
Uppsala,
Sweden
19, 1955
INTRODUCTION
In earlier studies (1,2) it has been shown that the activation of enolase by metal ions involves the formation of a complex in which one activating ion is bound per molecule of enzyme. However, no definite information on the nature of the groups involved in the binding of the metal ion is available- [cf. Ref. (3)]. The effect of pH on the rate of an enzymic reaction can sometimes be used in characterizing the active site of the enzyme (4, 5), and the pH dependence of metal ion-binding to proteins has also been used to det,ermine the side chains involved (6). Since the complexity of the enolase molecule makes difficult a direct chemical attack, this approach has been used in the study described here as a first step in attempting to identify the metal ion-binding site. As in previous studies (1, 2)) both kinetic and equilibrium measurements have been employed. Thus, the pH dependence of the strength of interaction between crystallized enolase and its activating ions has been investigated by activity measurements as well as by direct binding determinations. To increase the accuracy of the kinetic experiments, the method used for determining enolase activity has been modified, since the kinetic formulation of Warburg and Christian (7), previously employed (l-3, S), is not strictly valid (3). MATERIALS
AND METHODS
Reagentsand Analytical Methods The preparation of enolase, the synthesis of nL-2-phosphoglyceric acid (PGA), and the purification of other reagents were described earlier (l-3). The binding of Zn++ to enolase was measured by the equilibrium dialysis technique, and the concentration of Zn++ was determined by the dithizone method as described previously (1, 8). 186
ACTIVATION
OF ENOLASE
187
Measurement of h%olaseActivity The new method of activity measurement’s was still based on following the change with time (t) in the optical density (D) of the test solution at 240 rnp (7). However, instead of calculating the activity (a) according to the formula of Warburg and Christian (7) the initial velocity [vu = (dL)/dt),] was estimated. The reaction was started by adding 10~1. of enzyme solution to 3 ml. of substrate solution in a l-cm. cuvette. Since the final concentration of enzyme in the test solution was the same as in previous studies (around 10 pg./ml.) (1, 3, 8), the estimate of 210necessitated shorter time intervals between experiment)al points than the 1-min. periods used earlier (1,7). Ideally, a rapid automatic-recording spect,rophotometer should be employed, but this was not available for the present investigattion. Instead, the optical density was determined every 10 sec. for the first 2 min. of the reaction. With low and moderate activities, enough points were obtained on R straight line to allow an accurate calculation of ~0 , while at higher activities the optical density-time plot curved already after lo-20 sec. (see Results). In the lat ter cases, Q was estimated by fitting the experimental data (usually the lo-, 20., and 30.sec. readings) to an empirical cubic equation as described by Livingston (9). Since this approximation is only applicable when the deviation from linearity is small, measurements of high activities involve a relatively large experimental uncertainty. This is well illustrated in an earlier st,udy of the temperature drpendence of the enolase reaction (lo), in which the present met,hod of activity drterminations was employed. In previous investigations (1, 3,8), a Hilger “Uvispek” spect rophotometer was employed, but to allow a more rigorous temperature control, a Beckman DU spectrophotometer with thermospacers was used here. With this instrument, the temperature variation in the cuvette during the period of measurement could be kept as low as f0.02”. The temperature was 23” in all cases. Since t,he ultraviolet absorption of phospho(enol)pyruvic acid (PPA) changes rapidly with wavelength (3,7), the spectral purity of the radiation used must be high. Serial dilutions of PPA gave a straight line when the optical density was plotted as a function of concentration, showing that the band width obtained with the l3eckman instrument was sufficiently small. The extinct,ion coefficient of PP.4 is a function of pH (7); therefore, (dD/dt)o is only a measure of relative activities at a const,ant pH. If values of ~0 at different pH values should be compared, ~!c/rlt must be calculat,ed, where c is the concentration of PPA. Tris(hydroxymet,hyl)aminomethane (THAM)-HCl or phosphate buffers were used in all experiments. Their ionic strengths were adjusted to 0.050 at all pH v:tlues. RESULTS
AND D~scussro~
The Activity Measure It was noted earlier (3) that a calculated by the formula of Warburg and Christian (7) is not quite constant with time. To illustrate this, the primary data of two experiments with different enzyme concentrations are given in Table I. In general, a tends to increase with time; with
188
BO
G.
MALMSTRijM
AND
LARS
TABLE The Measurement
of vu and a with
10 #l. enzyme
WESTLUND
I
Diflerent Concentrations Buffer, pH 6.46
D 1 min.
E.
25 pl. enzyme
D, 10 pl. enzyme
-
D 25 ill. enzyme
of h’nolase
in Phosphate
10 pl. enzyme
a 25 pl. enzyme
0
0.044
0.072
0.895
0.867
-
-
l/6 216 516 1 2 3 5 7 10 w3
0.113 0.177 0.232 0.284 0.337 0.382 0.610 0.760 0.887 0.924 0.939 0.939
0.229 0.353 0.466 0.552 0.627 0.694 0.888 0.934 0.936 0.939
0.826 0.762 0.707 0.655 0.602 0.557 0.329 0.179
0.710 0.586 0.473 0.387 0.312 0.245
0.480 0.482 0.471 0.468 0.475 0.476 0.500 0.534
1.196 1.172 1.210 1.210 1.223 1.261
vo :
0.420
1.086
0.487
1.212
3/6
416
0.939 Average
a:
lower activities the relative increase is often larger than in the examples given here. However, it should be noted that the ratios of the activities calculated as a0 or a (average value) are very nearly the same, so that even if a is not truly constant, its average value is a good measure of the relative enzyme activity. Still, the initial velocity has been employed throughout this investigation, since its use involves fewer assumptions. Kinetically, the lack of constancy of a indicates that the equilibrium treatment (11) is not strictly applicable to the enolase reaction, as has been assumed in earlier investigations (1, 3, 8). It is possible that the lack of equilibrium is not between enzyme and metal ion or enzyme and substrates but rather between the different catalytic forms of the protein evidenced by the temperature studies (10). Such a situation most likely exists in the case of leucine aminopeptidase whose activation by Mn++ is a slow reaction according to the thorough kinetic studies of Smith and co-workers (12, 13). The slow reaction cannot be the formation of the metal ion-protein complex since the dissociation on dilution is instantaneous while the formation constant is around 104, which means that the association is still more rapid. Thus, it would seem that the complex first formed is slowly converted into a more active complex. Preliminary observations on enolase indicate the possibility of such
ACTIVATION
OF
189
ENOLASE
slow conversions to more active forms; more definite information may emerge from kinetic studies with a more refined technique now being planned. Still, the deviations from the Warburg-Christian formula (‘7) are small and the equilibrium assumption must be considered a fairly good approximation. l’he pH Dependenceof the Mg++ Activation of Enolase Only the activation by Mg++ was studied, since its kinetics is simpler than that of Mn++ and Zn++ activation in which account must be takeu of inhibition at high activator concentrations (I). ,4s was first shown by Warburg and Christian (7), the kinetics of Mg* activation can be satisfactorily described by an equation of the following form (3): V,,,M
(1)
” = K, + M
where AI represents the Mg* concentration, I<, a constant, and v,,,,, the asymptotic limit of vo when M approaches infinity. The experimental 0.200
16.0 x Vj3M
RGOx1f3M
OS50 4.00 x 1Ci3M %o dsoo
200~10'~M
l.M) Y Vj3M 0.050
Q50~16~M
No Mg++
FIG. 1. The change in optical density at 240 rnp (D) with time (1) after of enolase t,o 1.2 X 10-a M PGA4 in phosphate buffer, pH 5.5, p = 0.050, ferent roncent,rations of Mg++.
addition with dif-
190
BO
G.
MALMSTRijM
AND
:.-
LARS
E.
[hf$7*+J’x10-3
WESTJJUND
1.00
FIG. 2. The estimation of K. , Eq. (1). The reciprocal of the initial vebcity (IJO-l), estimated from the data in Fig. 1, is plotted as a function of the reciprocal concentration of Mg++.
evaluation of the constant K, is illustrated by Figs. 1 and 2. Figure 1 gives the spectrophotometric data used in determining the values of v. at different activator concentrations. K, is then calculated from the slope and intercept of a Lineweaver-Burk plot (3,14), as shown in Fig. 2. The results of the determination of K, at different pH values are given in Table II. In several of the experiments, the fit to a straight line in the Lineweaver-Burk plot was not as good as illustrated in Fig. 2; it was estimated that the uncertainty in the graphic determination of K, TABLE II The Constant K, , Eq. (l), at Different pH Values with Mg++ as Activator and p = 0.06 Km&s:p.1 PH
Bllffer
5.13 5.5 5.8 6.13 6.45 6.75 7.03 7.0 7.4 7.8
Phosphate “ ‘I I‘ “ I‘ “ THAM-HCl “ ‘I
[HPOr-]
1.1 2.2 3.6 6.1 9.0 11.8 13.7 -
X 10’
1.2 mY PGA
4.5 4.1 2.6 2.8 1.5 1.4 0.20 0.21
2.4 mM PGA
5.4 6.4 4.3 2.8 2.1 1.8 1.8 0.24 0.29
at .W
ACTIVATION
OF
191
ENOLASE
could be as large as 12 %. K, has been determined at two different concentrations of PGA, since it is also a function of the substrate concentration. In the cases where phosphate buffers have been used, the concentrations of HPO,are also given in Table II; these figures will be used subsequently in making certain corrections in the K, values. The data in Table II indicate that the strength of interaction between enolase and Mg* decreases below pH 6.8, since the increase in K, is well oumide the experimental error. It must, however, be remembered that even if the equilibrium treatment is valid in enolase kinetics, K, cannot be assumed to be identical with the dissociation constant of the enzyme-activator complex because the presence of buffer and substrate displaces its equilibrium with free enzyme protein. It has been shown earlier (1, 15) that increasing phosphate concentrations raises the K, values due to formation of MgHP04 . This is also demonstrated by the fact that considerably lower values are obtained in THAM-HCl buffers. Since the concentration of HPOh-- decreases with pH (see Table II), this partly obscures the pH effect on K, . However, the dissociation constant of MgHP04 is known (2) which allows corrections of the K, values [Eq. (6), Ref. (a)]. The corrected values increase approximately by a factor of 10 in going from pH 7.0 to pH 5.1. The effect of substrate concentration on K, depends on the mechanism of activation. Kinetic equations for several types of activation have been derived (16-18). Perhaps the most useful treatment for distinguishing between the various mechanisms is that given by Segal, Kachmar, and Boyer (16)) but unfortunately real cases are often more complicated than the ones considered, as shown, for example, in a study by Boyer’s own group (19). With enolase, the Lineweaver-Burk plots at different substrate concentrations have neither slope nor intercept in common, which corresponds to ‘(case IV” of Segal et al. (16). The physical meaning of K, is then given by the following relation (16) :
K’I x K, K’I+
K, = KI x KI+
s
K’I x K, s
(8
where K1 and Kfl are the dissociation constants of the protein-metal ion complex without and with substrate attached to the enzyme, respectively; KM is the Michaelis constant, and S the substrate concentration. Thus K, is only identical with the dissociation constant K1 , if K1 equals
192
BO
G.
MALMSTRijM
AND
LARS
E.
WESTLUND
K’l. The fact that the constants for Zn* determined earlier (1) by both kinetic and equilibrium measurements agree well with each other indicates this to be approximately true. But in this case Eq. (2) predicts K, to be independent of S, while the data in Table II show that it increases with S. However, the mechanism of Segal et al. (16) ignores removal of activator by complex formation with substrate, which would have the observed effect on K, . If this is also taken into account, and K1 is assumed to equal K’l, the rate equation becomes: ~IEMS ” = (K,
+ S) (K,
x 7
+ M)
(3)
where Kz is the dissociation constant of the activator complex with PGA, and kl is the rate constant for the rate-limiting step (1,3). Since KM and Kz have only been determined at pH 6.8 (1,2) and since only two values of S have been used, the data do not allow a complete test of Eq. (3) as yet; but, as already mentioned, a more extensive study of enolase kinetics is being planned. At pH 6.8, the increase in K, with S is approximately that which would be predicted by Eq. (3). Kz would be expected to increase with decreasing pH, and corrections according to Eq. (3) would therefore not change the conclusion that the interaction of enolase with metal ions decreases below pH 6.8. However, until the applicability of Eq. (3) has been more thoroughly tested, there still rests a great deal of uncertainty in the calculation of the dissociation constant K1 from the kinetic constant K, , so that this conclusion would be less well justified without the confirmation from the equilibrium experiments described in the next section. It is interesting to note that at high concentrations of the activating ion, Eq. (3) transforms into the equation previously found to describe the substrate kiin the present case the inhibition netics of enolase [(Eq. (12), Ref. (l)]. H owever, at high substrate concentrations is due to removal of the act,ivating ion, while a different mechanism was suggested earlier (1). This gives a further example of the well-known fact of enzyme kinetics, that distinct mechanisms may result in identical rate equations [cf. Refs. (20, 21)l.
The E$ect of pH on the Interaction
of Enolase with Zn*
The relatively high dissociation of Mg?+ enolase makes it difficult to determine the binding of Mg++ by equilibrium dialysis [cf. Ref. (l)]. Therefore, the interaction with the more strongly bound Zn++ was
ACTIVATION
OF
ENOLASE
193
studied by this technique. Figure 3 gives the average number of Zni+ bound per molecule of enzyme as a function of the concentration of free Zn++ at three different pH values. From the precision of the analytical methods [see Ref. (l)], it was estimat.ed that the uncertainty in the values of t is about fO.l. Due to limited availability of the crystallized enzyme, only two points were determined at each pH. However, even these few data are sufficient to demonstrate a marked decrease with pH in the strength of interaction between enzyme and metal ion. As has been shown earlier (I), in the low concentration range st.udied, binding occurs mainly at the active site so that the binding curves can be approximately described by t’he following equation [cf. Eq. (22), Ref. (l)] :
where K1 is the dissociation constant ICI was earlier (1) found to 1~ 5 X lO+ 211 at pH 6.8, and the data in Fig. 2 show that, the value is approximately the same at pH 6.5. However, at pH 5.1 a value of around 5 X lows M is obtained by the application of Eq. (2). ‘These data thus confirm the conclusion from the kinetic experiments and justify the use of K, as an approximate measure of K, . The pH dependence of K1 indicates that a side chain with a pK fat its acid dissociation of around 6 is involved in t)he binding of the activating metal ion to enolase. However, such a, conclusion must be drawn with
[znn+ +]
freeX
FIG. 3. The binding pH 7.8; (0) phosphate,
to5
of Zn++ to enolasc at 4” and p = 0.050; pH G.45; (x) phosph:~te, pII 5.1X
(‘0)
THALI-Tf(‘I,
194
BO
G.
MALMSTRijM
AND
LARS
E. WESTLUND
caution since the change in net negative charge of the protein may be sufficient to influence the affinity considerably on purely nonspecific, electrostatic grounds, as is shown, for example in studies by Klotz and co-workers (6,22). In the case of enolase, however, the electrostatic effect is too small to account for the decrease in binding. The difference in stability between pH 7.0 and pH 5.1 corresponds to a standard freeenergy change of 1400 Cal., which, according to the mathematical treatment given by Klotz (22, 23), would result only if the change in charge were 15 units. But in this pH region the only groups titrated are probably the imidazole side chains of histidine (6, 24), and the histidine content of enolase corresponds to only about a third of the required number of residues per mole (unpublished experiments). The small charge is also indicated by the low electrophoretic mobility of enolase at pH 6.8 (1) ; this does not increase much in going from the isoelectric point at pH 5.5 (25) until a pH of around 8 is reached (unpublished experiments). Thus, the pH dependence of the binding of the activating ion to enolase would implicate the direct involvement of a group with a pK of around 6. The only such groups in proteins are usually considered to be imidazole side chains (6, 24), but it is possible that the special steric conditions existing at the active site of an enzyme may considerably influence the acidic strengths of the side chains. Therefore, the conclusion that histidine is involved in the binding of the activating ion to enolase should be regarded as tentative until confirmatory evidence has been obtained by other techniques. Attempts have been made to utilize optical methods (6, 26), but the low solubility of Cu* enolase prevented the use of this approach. Additional evidence must, therefore, probably be derived by chemical techniques [cf. Ref. (6)]. It has been frequently pointed out (l-3) that the interaction of enolase with Zn* is much stronger than that found between a single imidazole group and Zn++ (27). Thus, if histidine partakes at all in the binding of the activating ion to enolase, either two imidazole groups or one imidazole and some other group, e.g., carboxyl [cf. Ref. (3)], must be involved. ACKNOWLEDGMENTS The authors wish to thank Professor A. Tiselius for his stimulating interest many valuable discussions. The investigation was supported by grants from Wallenberg and Rockefeller Foundations.
and the
ACTIVATION
OF
1 9.5
ENOLARE
1. The enolasc activity measure of Warburg and Christian (7) has bee11shown not to be strictly valid. The measurement of the initial velocity of the enolase reaction has been described and the results compared with those obtained by the Warburg-Christian method. The possible kinetic significance of the deviation has been discussed. 2. The kinetics of the Mg* activation of enolase has been studied at several pH values. The physical meaning of the kinetic constants has been discussed in terms of different mechanisms of activation. It is concluded that t’he binding of Mg++ to enolasedecreasesmarkedly below pH 6.8. 3. The interaction of enolase with ZE++ at different pH values has been st,udied by equilibrium dialysis. The results confirm the conclusion from t’ht kinetic experiments. It has been suggested that an imidaeole group of histidine is involved in the binding of the activating ioll to enolase. REFERENCES 1. 2. 3. 4. 5. 6. 7. 8.
9. 10. 11. 1‘2. 13.
14.
MAr,b%TRijiv,
B. G.,
ilrch.
%xhem.
and Hiophys.
iVI.tr,mmijhr,
B. G., ilrch. Biochem. and Biophys. ~IALMSTR~~X, B. G., Arch. Biochem. and Biophys. FRIEDEN, C., AND ALBERTY, R. A., J. Biol. (‘hem.
I,AIDLER, KLUTZ,
K.
J.,
Trans.
Faraday
sot.
46, 345 (195X). 49, 335 (1954). 58, 381 (1955).
212, 859 (1955).
61, 528 (1055).
I. M., in “The Mechanism of Enzyme Action” (Mcl
15. MALMSTR~M, 16. SEGAI,, H.
B. G.,
Nature
171,
392
(1953).
L., KACHMAR, J. F., AND BOYER, P. D., Enzymologia 16, 187 (195“). 17. HULL, H. B., in “The Mechanism of Enzyme Action” (McElroy, W. D., and Glass, B., eds.), p. 141. The Johns Hopkins Press, Baltimore, 1954. 18. FHIEI)ENWALD, J. S., AKD MAENGWYN-DAVIES, G. I)., in “The Mechanism of
196
BO
(3.
MALMSTRGM
AND
LARS
E.
WESTLUND
Enzyme Action” (McElroy, W. D., and Glass, B., eds.), p. 154. The Johns Hopkins Press, Baltimore, 1954. 19. ROBBINS, E. A., STULBERG, M. P., AND BOYER, P. D., Arch. Biochem. and Biophys. 64, 215 (1954). 20. ALBERTY, R. A., J. Am. Chem. Sot. 76, 2494 (1954). 21. OGSTON, A. G., Discussions Faraday Sot. August 1955. 22. KLOTZ, I. M., AND FIESS, H. A., J. Phys. & Cotloid Chem. 66, 101 (1951). 23. KLOTZ, I. M., in “The Proteins” (N eurath, H., and Bailey, K., eds.), Vol. I, part B, p. 727. Academic Press, New York, 1953. 24. ALBERTY, R. A., in “The Proteins” (Neurath, H., and Bailey, K., eds.), Vol. I, part A, p. 461. Academic Press, New York, 1953. (Sumner, J. B., and Myrback, K., eds.), 25. MEYERHOF, O., in “The Enzymes” Vol. I, part 2, p. 1210. Academic Press, New York, 1951. 26. WARNER, R. C., AND WEBER, I., J. Am. Chem. Sot. 76, 5094 (1953). 27. GURD, F. R. N., AND GOODMAN, D. S., J. Am. Chem. Sot. 74, 670 (1952).
OF
BIOCHEMISTRY
AND
BIOPHYSICS
61,
186-196
(1956)
The Effect of pH on the Interaction of Enolase with Activating Metal Ions Bo G. Malmstrijm From
the Institute
and Lars E. Westlund
of Biochemistry,
Received
September
Uppsala,
Sweden
19, 1955
INTRODUCTION
In earlier studies (1,2) it has been shown that the activation of enolase by metal ions involves the formation of a complex in which one activating ion is bound per molecule of enzyme. However, no definite information on the nature of the groups involved in the binding of the metal ion is available- [cf. Ref. (3)]. The effect of pH on the rate of an enzymic reaction can sometimes be used in characterizing the active site of the enzyme (4, 5), and the pH dependence of metal ion-binding to proteins has also been used to det,ermine the side chains involved (6). Since the complexity of the enolase molecule makes difficult a direct chemical attack, this approach has been used in the study described here as a first step in attempting to identify the metal ion-binding site. As in previous studies (1, 2)) both kinetic and equilibrium measurements have been employed. Thus, the pH dependence of the strength of interaction between crystallized enolase and its activating ions has been investigated by activity measurements as well as by direct binding determinations. To increase the accuracy of the kinetic experiments, the method used for determining enolase activity has been modified, since the kinetic formulation of Warburg and Christian (7), previously employed (l-3, S), is not strictly valid (3). MATERIALS
AND METHODS
Reagentsand Analytical Methods The preparation of enolase, the synthesis of nL-2-phosphoglyceric acid (PGA), and the purification of other reagents were described earlier (l-3). The binding of Zn++ to enolase was measured by the equilibrium dialysis technique, and the concentration of Zn++ was determined by the dithizone method as described previously (1, 8). 186
ACTIVATION
OF ENOLASE
187
Measurement of h%olaseActivity The new method of activity measurement’s was still based on following the change with time (t) in the optical density (D) of the test solution at 240 rnp (7). However, instead of calculating the activity (a) according to the formula of Warburg and Christian (7) the initial velocity [vu = (dL)/dt),] was estimated. The reaction was started by adding 10~1. of enzyme solution to 3 ml. of substrate solution in a l-cm. cuvette. Since the final concentration of enzyme in the test solution was the same as in previous studies (around 10 pg./ml.) (1, 3, 8), the estimate of 210necessitated shorter time intervals between experiment)al points than the 1-min. periods used earlier (1,7). Ideally, a rapid automatic-recording spect,rophotometer should be employed, but this was not available for the present investigattion. Instead, the optical density was determined every 10 sec. for the first 2 min. of the reaction. With low and moderate activities, enough points were obtained on R straight line to allow an accurate calculation of ~0 , while at higher activities the optical density-time plot curved already after lo-20 sec. (see Results). In the lat ter cases, Q was estimated by fitting the experimental data (usually the lo-, 20., and 30.sec. readings) to an empirical cubic equation as described by Livingston (9). Since this approximation is only applicable when the deviation from linearity is small, measurements of high activities involve a relatively large experimental uncertainty. This is well illustrated in an earlier st,udy of the temperature drpendence of the enolase reaction (lo), in which the present met,hod of activity drterminations was employed. In previous investigations (1, 3,8), a Hilger “Uvispek” spect rophotometer was employed, but to allow a more rigorous temperature control, a Beckman DU spectrophotometer with thermospacers was used here. With this instrument, the temperature variation in the cuvette during the period of measurement could be kept as low as f0.02”. The temperature was 23” in all cases. Since t,he ultraviolet absorption of phospho(enol)pyruvic acid (PPA) changes rapidly with wavelength (3,7), the spectral purity of the radiation used must be high. Serial dilutions of PPA gave a straight line when the optical density was plotted as a function of concentration, showing that the band width obtained with the l3eckman instrument was sufficiently small. The extinct,ion coefficient of PP.4 is a function of pH (7); therefore, (dD/dt)o is only a measure of relative activities at a const,ant pH. If values of ~0 at different pH values should be compared, ~!c/rlt must be calculat,ed, where c is the concentration of PPA. Tris(hydroxymet,hyl)aminomethane (THAM)-HCl or phosphate buffers were used in all experiments. Their ionic strengths were adjusted to 0.050 at all pH v:tlues. RESULTS
AND D~scussro~
The Activity Measure It was noted earlier (3) that a calculated by the formula of Warburg and Christian (7) is not quite constant with time. To illustrate this, the primary data of two experiments with different enzyme concentrations are given in Table I. In general, a tends to increase with time; with
188
BO
G.
MALMSTRijM
AND
LARS
TABLE The Measurement
of vu and a with
10 #l. enzyme
WESTLUND
I
Diflerent Concentrations Buffer, pH 6.46
D 1 min.
E.
25 pl. enzyme
D, 10 pl. enzyme
-
D 25 ill. enzyme
of h’nolase
in Phosphate
10 pl. enzyme
a 25 pl. enzyme
0
0.044
0.072
0.895
0.867
-
-
l/6 216 516 1 2 3 5 7 10 w3
0.113 0.177 0.232 0.284 0.337 0.382 0.610 0.760 0.887 0.924 0.939 0.939
0.229 0.353 0.466 0.552 0.627 0.694 0.888 0.934 0.936 0.939
0.826 0.762 0.707 0.655 0.602 0.557 0.329 0.179
0.710 0.586 0.473 0.387 0.312 0.245
0.480 0.482 0.471 0.468 0.475 0.476 0.500 0.534
1.196 1.172 1.210 1.210 1.223 1.261
vo :
0.420
1.086
0.487
1.212
3/6
416
0.939 Average
a:
lower activities the relative increase is often larger than in the examples given here. However, it should be noted that the ratios of the activities calculated as a0 or a (average value) are very nearly the same, so that even if a is not truly constant, its average value is a good measure of the relative enzyme activity. Still, the initial velocity has been employed throughout this investigation, since its use involves fewer assumptions. Kinetically, the lack of constancy of a indicates that the equilibrium treatment (11) is not strictly applicable to the enolase reaction, as has been assumed in earlier investigations (1, 3, 8). It is possible that the lack of equilibrium is not between enzyme and metal ion or enzyme and substrates but rather between the different catalytic forms of the protein evidenced by the temperature studies (10). Such a situation most likely exists in the case of leucine aminopeptidase whose activation by Mn++ is a slow reaction according to the thorough kinetic studies of Smith and co-workers (12, 13). The slow reaction cannot be the formation of the metal ion-protein complex since the dissociation on dilution is instantaneous while the formation constant is around 104, which means that the association is still more rapid. Thus, it would seem that the complex first formed is slowly converted into a more active complex. Preliminary observations on enolase indicate the possibility of such
ACTIVATION
OF
189
ENOLASE
slow conversions to more active forms; more definite information may emerge from kinetic studies with a more refined technique now being planned. Still, the deviations from the Warburg-Christian formula (‘7) are small and the equilibrium assumption must be considered a fairly good approximation. l’he pH Dependenceof the Mg++ Activation of Enolase Only the activation by Mg++ was studied, since its kinetics is simpler than that of Mn++ and Zn++ activation in which account must be takeu of inhibition at high activator concentrations (I). ,4s was first shown by Warburg and Christian (7), the kinetics of Mg* activation can be satisfactorily described by an equation of the following form (3): V,,,M
(1)
” = K, + M
where AI represents the Mg* concentration, I<, a constant, and v,,,,, the asymptotic limit of vo when M approaches infinity. The experimental 0.200
16.0 x Vj3M
RGOx1f3M
OS50 4.00 x 1Ci3M %o dsoo
200~10'~M
l.M) Y Vj3M 0.050
Q50~16~M
No Mg++
FIG. 1. The change in optical density at 240 rnp (D) with time (1) after of enolase t,o 1.2 X 10-a M PGA4 in phosphate buffer, pH 5.5, p = 0.050, ferent roncent,rations of Mg++.
addition with dif-
190
BO
G.
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AND
:.-
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[hf$7*+J’x10-3
WESTJJUND
1.00
FIG. 2. The estimation of K. , Eq. (1). The reciprocal of the initial vebcity (IJO-l), estimated from the data in Fig. 1, is plotted as a function of the reciprocal concentration of Mg++.
evaluation of the constant K, is illustrated by Figs. 1 and 2. Figure 1 gives the spectrophotometric data used in determining the values of v. at different activator concentrations. K, is then calculated from the slope and intercept of a Lineweaver-Burk plot (3,14), as shown in Fig. 2. The results of the determination of K, at different pH values are given in Table II. In several of the experiments, the fit to a straight line in the Lineweaver-Burk plot was not as good as illustrated in Fig. 2; it was estimated that the uncertainty in the graphic determination of K, TABLE II The Constant K, , Eq. (l), at Different pH Values with Mg++ as Activator and p = 0.06 Km&s:p.1 PH
Bllffer
5.13 5.5 5.8 6.13 6.45 6.75 7.03 7.0 7.4 7.8
Phosphate “ ‘I I‘ “ I‘ “ THAM-HCl “ ‘I
[HPOr-]
1.1 2.2 3.6 6.1 9.0 11.8 13.7 -
X 10’
1.2 mY PGA
4.5 4.1 2.6 2.8 1.5 1.4 0.20 0.21
2.4 mM PGA
5.4 6.4 4.3 2.8 2.1 1.8 1.8 0.24 0.29
at .W
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ENOLASE
could be as large as 12 %. K, has been determined at two different concentrations of PGA, since it is also a function of the substrate concentration. In the cases where phosphate buffers have been used, the concentrations of HPO,are also given in Table II; these figures will be used subsequently in making certain corrections in the K, values. The data in Table II indicate that the strength of interaction between enolase and Mg* decreases below pH 6.8, since the increase in K, is well oumide the experimental error. It must, however, be remembered that even if the equilibrium treatment is valid in enolase kinetics, K, cannot be assumed to be identical with the dissociation constant of the enzyme-activator complex because the presence of buffer and substrate displaces its equilibrium with free enzyme protein. It has been shown earlier (1, 15) that increasing phosphate concentrations raises the K, values due to formation of MgHP04 . This is also demonstrated by the fact that considerably lower values are obtained in THAM-HCl buffers. Since the concentration of HPOh-- decreases with pH (see Table II), this partly obscures the pH effect on K, . However, the dissociation constant of MgHP04 is known (2) which allows corrections of the K, values [Eq. (6), Ref. (a)]. The corrected values increase approximately by a factor of 10 in going from pH 7.0 to pH 5.1. The effect of substrate concentration on K, depends on the mechanism of activation. Kinetic equations for several types of activation have been derived (16-18). Perhaps the most useful treatment for distinguishing between the various mechanisms is that given by Segal, Kachmar, and Boyer (16)) but unfortunately real cases are often more complicated than the ones considered, as shown, for example, in a study by Boyer’s own group (19). With enolase, the Lineweaver-Burk plots at different substrate concentrations have neither slope nor intercept in common, which corresponds to ‘(case IV” of Segal et al. (16). The physical meaning of K, is then given by the following relation (16) :
K’I x K, K’I+
K, = KI x KI+
s
K’I x K, s
(8
where K1 and Kfl are the dissociation constants of the protein-metal ion complex without and with substrate attached to the enzyme, respectively; KM is the Michaelis constant, and S the substrate concentration. Thus K, is only identical with the dissociation constant K1 , if K1 equals
192
BO
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K’l. The fact that the constants for Zn* determined earlier (1) by both kinetic and equilibrium measurements agree well with each other indicates this to be approximately true. But in this case Eq. (2) predicts K, to be independent of S, while the data in Table II show that it increases with S. However, the mechanism of Segal et al. (16) ignores removal of activator by complex formation with substrate, which would have the observed effect on K, . If this is also taken into account, and K1 is assumed to equal K’l, the rate equation becomes: ~IEMS ” = (K,
+ S) (K,
x 7
+ M)
(3)
where Kz is the dissociation constant of the activator complex with PGA, and kl is the rate constant for the rate-limiting step (1,3). Since KM and Kz have only been determined at pH 6.8 (1,2) and since only two values of S have been used, the data do not allow a complete test of Eq. (3) as yet; but, as already mentioned, a more extensive study of enolase kinetics is being planned. At pH 6.8, the increase in K, with S is approximately that which would be predicted by Eq. (3). Kz would be expected to increase with decreasing pH, and corrections according to Eq. (3) would therefore not change the conclusion that the interaction of enolase with metal ions decreases below pH 6.8. However, until the applicability of Eq. (3) has been more thoroughly tested, there still rests a great deal of uncertainty in the calculation of the dissociation constant K1 from the kinetic constant K, , so that this conclusion would be less well justified without the confirmation from the equilibrium experiments described in the next section. It is interesting to note that at high concentrations of the activating ion, Eq. (3) transforms into the equation previously found to describe the substrate kiin the present case the inhibition netics of enolase [(Eq. (12), Ref. (l)]. H owever, at high substrate concentrations is due to removal of the act,ivating ion, while a different mechanism was suggested earlier (1). This gives a further example of the well-known fact of enzyme kinetics, that distinct mechanisms may result in identical rate equations [cf. Refs. (20, 21)l.
The E$ect of pH on the Interaction
of Enolase with Zn*
The relatively high dissociation of Mg?+ enolase makes it difficult to determine the binding of Mg++ by equilibrium dialysis [cf. Ref. (l)]. Therefore, the interaction with the more strongly bound Zn++ was
ACTIVATION
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193
studied by this technique. Figure 3 gives the average number of Zni+ bound per molecule of enzyme as a function of the concentration of free Zn++ at three different pH values. From the precision of the analytical methods [see Ref. (l)], it was estimat.ed that the uncertainty in the values of t is about fO.l. Due to limited availability of the crystallized enzyme, only two points were determined at each pH. However, even these few data are sufficient to demonstrate a marked decrease with pH in the strength of interaction between enzyme and metal ion. As has been shown earlier (I), in the low concentration range st.udied, binding occurs mainly at the active site so that the binding curves can be approximately described by t’he following equation [cf. Eq. (22), Ref. (l)] :
where K1 is the dissociation constant ICI was earlier (1) found to 1~ 5 X lO+ 211 at pH 6.8, and the data in Fig. 2 show that, the value is approximately the same at pH 6.5. However, at pH 5.1 a value of around 5 X lows M is obtained by the application of Eq. (2). ‘These data thus confirm the conclusion from the kinetic experiments and justify the use of K, as an approximate measure of K, . The pH dependence of K1 indicates that a side chain with a pK fat its acid dissociation of around 6 is involved in t)he binding of the activating metal ion to enolase. However, such a, conclusion must be drawn with
[znn+ +]
freeX
FIG. 3. The binding pH 7.8; (0) phosphate,
to5
of Zn++ to enolasc at 4” and p = 0.050; pH G.45; (x) phosph:~te, pII 5.1X
(‘0)
THALI-Tf(‘I,
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caution since the change in net negative charge of the protein may be sufficient to influence the affinity considerably on purely nonspecific, electrostatic grounds, as is shown, for example in studies by Klotz and co-workers (6,22). In the case of enolase, however, the electrostatic effect is too small to account for the decrease in binding. The difference in stability between pH 7.0 and pH 5.1 corresponds to a standard freeenergy change of 1400 Cal., which, according to the mathematical treatment given by Klotz (22, 23), would result only if the change in charge were 15 units. But in this pH region the only groups titrated are probably the imidazole side chains of histidine (6, 24), and the histidine content of enolase corresponds to only about a third of the required number of residues per mole (unpublished experiments). The small charge is also indicated by the low electrophoretic mobility of enolase at pH 6.8 (1) ; this does not increase much in going from the isoelectric point at pH 5.5 (25) until a pH of around 8 is reached (unpublished experiments). Thus, the pH dependence of the binding of the activating ion to enolase would implicate the direct involvement of a group with a pK of around 6. The only such groups in proteins are usually considered to be imidazole side chains (6, 24), but it is possible that the special steric conditions existing at the active site of an enzyme may considerably influence the acidic strengths of the side chains. Therefore, the conclusion that histidine is involved in the binding of the activating ion to enolase should be regarded as tentative until confirmatory evidence has been obtained by other techniques. Attempts have been made to utilize optical methods (6, 26), but the low solubility of Cu* enolase prevented the use of this approach. Additional evidence must, therefore, probably be derived by chemical techniques [cf. Ref. (6)]. It has been frequently pointed out (l-3) that the interaction of enolase with Zn* is much stronger than that found between a single imidazole group and Zn++ (27). Thus, if histidine partakes at all in the binding of the activating ion to enolase, either two imidazole groups or one imidazole and some other group, e.g., carboxyl [cf. Ref. (3)], must be involved. ACKNOWLEDGMENTS The authors wish to thank Professor A. Tiselius for his stimulating interest many valuable discussions. The investigation was supported by grants from Wallenberg and Rockefeller Foundations.
and the
ACTIVATION
OF
1 9.5
ENOLARE
1. The enolasc activity measure of Warburg and Christian (7) has bee11shown not to be strictly valid. The measurement of the initial velocity of the enolase reaction has been described and the results compared with those obtained by the Warburg-Christian method. The possible kinetic significance of the deviation has been discussed. 2. The kinetics of the Mg* activation of enolase has been studied at several pH values. The physical meaning of the kinetic constants has been discussed in terms of different mechanisms of activation. It is concluded that t’he binding of Mg++ to enolasedecreasesmarkedly below pH 6.8. 3. The interaction of enolase with ZE++ at different pH values has been st,udied by equilibrium dialysis. The results confirm the conclusion from t’ht kinetic experiments. It has been suggested that an imidaeole group of histidine is involved in the binding of the activating ioll to enolase. REFERENCES 1. 2. 3. 4. 5. 6. 7. 8.
9. 10. 11. 1‘2. 13.
14.
MAr,b%TRijiv,
B. G.,
ilrch.
%xhem.
and Hiophys.
iVI.tr,mmijhr,
B. G., ilrch. Biochem. and Biophys. ~IALMSTR~~X, B. G., Arch. Biochem. and Biophys. FRIEDEN, C., AND ALBERTY, R. A., J. Biol. (‘hem.
I,AIDLER, KLUTZ,
K.
J.,
Trans.
Faraday
sot.
46, 345 (195X). 49, 335 (1954). 58, 381 (1955).
212, 859 (1955).
61, 528 (1055).
I. M., in “The Mechanism of Enzyme Action” (Mcl
15. MALMSTR~M, 16. SEGAI,, H.
B. G.,
Nature
171,
392
(1953).
L., KACHMAR, J. F., AND BOYER, P. D., Enzymologia 16, 187 (195“). 17. HULL, H. B., in “The Mechanism of Enzyme Action” (McElroy, W. D., and Glass, B., eds.), p. 141. The Johns Hopkins Press, Baltimore, 1954. 18. FHIEI)ENWALD, J. S., AKD MAENGWYN-DAVIES, G. I)., in “The Mechanism of
196
BO
(3.
MALMSTRGM
AND
LARS
E.
WESTLUND
Enzyme Action” (McElroy, W. D., and Glass, B., eds.), p. 154. The Johns Hopkins Press, Baltimore, 1954. 19. ROBBINS, E. A., STULBERG, M. P., AND BOYER, P. D., Arch. Biochem. and Biophys. 64, 215 (1954). 20. ALBERTY, R. A., J. Am. Chem. Sot. 76, 2494 (1954). 21. OGSTON, A. G., Discussions Faraday Sot. August 1955. 22. KLOTZ, I. M., AND FIESS, H. A., J. Phys. & Cotloid Chem. 66, 101 (1951). 23. KLOTZ, I. M., in “The Proteins” (N eurath, H., and Bailey, K., eds.), Vol. I, part B, p. 727. Academic Press, New York, 1953. 24. ALBERTY, R. A., in “The Proteins” (Neurath, H., and Bailey, K., eds.), Vol. I, part A, p. 461. Academic Press, New York, 1953. (Sumner, J. B., and Myrback, K., eds.), 25. MEYERHOF, O., in “The Enzymes” Vol. I, part 2, p. 1210. Academic Press, New York, 1951. 26. WARNER, R. C., AND WEBER, I., J. Am. Chem. Sot. 76, 5094 (1953). 27. GURD, F. R. N., AND GOODMAN, D. S., J. Am. Chem. Sot. 74, 670 (1952).