Y. POCKEB AND J. E. MEANY
1486
The Reversible Hydration of Pyruvic Acid. 11. Metal Ion and Enzymatic Catalysis' by Y. Pocker2and J. E. Meany Department of Chemistry, University of Washington, Seattle, Washington 98106
(Receiued September 1 6 , 1969)
The hydration of pyruvic acid and the dehydration of 2,2-dihydroxypropionate anion were followed at 0.0' using a spectrophotometric method. The dehydration of 2,Z-dihydroxypropionateanion was investigated with respect to its sensitivity to catalysis by the zinc metalloenzyme carbonic anhydrase from bovine erythrocytes and by various transition metal ions in diethylmalonate buffers at pH 6.7. The magnitude of metal ion catalysis follows the order CuZ+ > Znzf > Nizt > Co2+ > CdZt > Mn2+, an order presumably dictated by the ability of the substrate to chelate to these divalent metal ions. The enzymatically catalyzed dehydration of 2,Z-dihydroxypropionateanion to pyruvate anion has not previously been demonstrated, and the anion hydrate represents the first known substrate analog of HCOJ-. The hydration of pyruvic acid was studied in the pH range 1.2 to 2.1 and these investigations allowed the determination of the spontaneous and the hydronium ion rate constants (ko = 2 min-I, ~ H ~ o=+104 1. mol-1 min-l).
Introduction The physiological importance of pyruvic acid lies in its intermediacy in the metabolism of both proteins and carbohydrate^.^ Our interest in this compound is associated with its reversible hydration. Eigen, et u Z . , ~ and Strehlow5 have studied this process but only at relatively low P H . ~ We have shown in part I of this series7 that the ratio of 2,2-dihydroxypropionic to pyruvic acid at equilibrium is dependent on the third power of the water concentration, eq 1, and took this as an indication that the miniCHsCOCOzH
+ (X + 1) HzO
-
A
(CH,C(OH)2C02H) * ~ H 2 0(1)
mum stoichiometric composition of the hydrate in its ground state is CHJC(OH)z.C02H.2Hz0. The pyruvate anion, however, has little tendency to hydrate since the electron withdrawing jnfluence of the carboxyl group has been rendered ineffective with deprotonation. Consequently, as the pH of the reaction media exceeds the pK, of pyruvic acid (pK, = 2.18 a t 25.0°)8 the dehydration of 2,2-dihydroxypropionic acid becomes dominant (CH3C(OH)2C02H)exHz0 E CHsCOCOzHSOf
+
+ xHzO
(2)
Finally, when the p H of the reaction media exceeds the pK, of 2,2-dihydroxypropionic acid (pK, = 3.6 a t 25.0°8),one is able to follow the dehydration of the 2,2dihydroxypropionate anion (CH,C(OH)&02-). ZHZO
CHsCOCOz(X
+
+ 1)HzO
(3)
The present work delineates the sensitivity of this latter T h e Journal of Physical Chemistry
process, (eq 3) , toward metal ion and enzymatic catalysis. These studies were prompted by the proximity of the carbonyl and carboxyl groups in pyruvic acid. Thus the relative positions of the oxygens associated with the carbonyl and carboxyl groups allow pyruvic acid to act as a bidentate ligand and promote the formation of complexes between pyruvic acid and various metal ions.g-ll With the metalloenzyme carbonic anhydrase, it has become increasingly clear that its efficiency as a hydrating and hydrolyzing catalyst12~1a is not a property of the metal ion alone nor of the protein alone. The dehydration of 2,2-dihydroxypropionate by carbonic anhydrase further exemplifies the vast catalytic versatility of this enzyme.
(1) This work was supported by U.S. Public Health Service Grants from the National Institutes of Health. (2) Author t o whom correspondence should be addressed. (3) D. M. Greenberg, "Metabolic Pathways," Academic Press, London, Vol. 1 (1960) and Vol. 2 (1961). (4) M. Eigen, K. Kustin, and H. Strehlow, Z. Phys. Chem., 31, 140 (1962). (5) H. Strehlow, 2.Elektrochem., 66,3921 (1962). (6). Measurements of the rapid reversible hydration of pyruvic acid's6 indicate that reaction 1 is catalyzed by hydronium ions, pyruvic acid molecules, and pyruvate ions. There is also a spontaneous rate of hydration which by comparison with the results of other hydration reactions seemed to the authors486 to be too large to be interpreted entirely from catalysis by water molecules alone. (7) Y. Pocker, J. E. Meany, B. J. Nist, and C. Zadorojny, J . P h y s . Chem., 73, 2879 (1969). (8) G. Kortfim, W. Vogel, and K. Andrussow, "Dissociation Constants of Acids in Aqueous Solution," Butterworth and Co., Ltd., London, 1961, p 335. (9) E. Ferrell, J. M. Ridgion, and H. L. Riley, J. Chem. Soc., 1442 (1934). (10) W. Franke, Ann. Chem., 475,39 (1929). (11) D. L. Leussing, Talanta, 11, (2), 189 (1964). (12) Y. Pocker and J. E. Meany, Biochemistry, 4,2535 (1965). (13) Y. Pocker and D. R. Storm, ibid., 7, 1202 (1968).
REVERSIBLE HYDRATION OF PYRUVIC ACID
1487
Experimental Section Materials. Reagent grade pyruvic acid was an Eastman Organic Chemicals product which was twice distilled under reduced pressure through a short Vigreux column in an atmosphere of nitrogen gas, bp (18 mm) 60-69'. The substrate was always distilled directly before use in order to minimize polymerization. The second distillation was carried out because small amounts of acetic acid were detected in the distillate from early fractions. Thus, the nmr spectrum of aqueous solutions of these early fractions showed the two expected signals arising from the methyl hydrogens associated with pyruvic acid (2-5 ppm) and its conjugate hydrate (3.5ppm) as well as a third peak in between (3.0 ppm), the intensity of which increased upon addition of small quantities of acetic acid.' The preparation and purification of diethylmalonic acid has been described in an earlier paper.ll All other buffer components were commercially available in analytical or reagent grade. Metal ion solutions were prepared in the form of their nitrates or chlorides as obtained from Baker and Adamson in reagent grade. Bovine carbonic anhydrase was a product of Mann Research Laboratories and was assayed a t 2900 Philpot units/mg. Zinc ion analysis of the enzyme was carried out as in our earlier work.I4 The enzyme was stored dry a t - 20". Apparatus. The reactions were followed on a Gilford high-speed recording spectrophotometer, Model 2000, which was equipped with a specially constructed bath, thermostated to 0.0 f 0.02' by means of the apparatus described earlier. Measurements of pH were carried out on a Beckman 101900 research pH meter. Ultraviolet scanning employed a Cary 14 recording spectrophotometer. All nmr measurements were obtained on a Varian DA-60 IL instrument. Method. Because pyruvate ion remains essentially unhydrated around neutral pH,' it is necessary to follow the dehydration of 2,2-dihydroxypropionate in this region. At very low values of pH, an appreciable amount of the acid hydrate exists in aqueous Thus, an equilibrated solution of pyruvic acid and its hydrate resulted from an aqueous solution of pyruvic acid (32% v/v) and small amounts of concentrated HC1 which were added to obtain a p H around 0.35. A calibrated Hamilton syringe was employed to inject 0.01 ml of this solution into 3 ml of the solution to be tested. The total pyruvate concentration was then 0.015 &I. An instantaneous decrease in p H of about 0.3 unit was observed upon addition of the pyruvic acid to 3 ml of 0.1 M diethylmalonate buffer. The final pH of 6-8 was reached long before any appreciable dehydration was recorded. The dehydration was monitored by following the increase in absorbancy a t either 762 or 340 mp. The ultraviolet scanning of pyruvic acid in T H F reveals two maxima, one a t 276 mp and a more intense peak a t
-0.7
-0.8 -0.9 h
-I;
.o
I
8
4 -1.1
Y
-rn 0
-I .2 -1.3
-1.4
-1.5
-~
I
0.25
I
I
0.75 Time (min.)
0.50
I
1.00
Figure 1. Typical kinetic run for the dehydration of 2,2-dihydroxypropionate with 4 x 10-4 M Zn2+ in 0.10 M diethylmalonate buffer at pH 6.7 and 0.0".
388 mp. In acidic aqueous solutions, only one maximum a t 327 mp was observed which was much less intense than that observed at 338 mp in T H F solvent. The pseudo-first-order rate constants were determined from the slope of the straight line obtained by plotting log (Am - A,) vs. time (Figure 1). The results obtained a t the two wavelengths were identical to within experimental error ( f2%). The hydration of pyruvic acid at low pH (pH Zn2+ > Ni2+> Co2+ > Cd2+> Mn2+ Earlier work in these laboratories involving the catalytic efficiencies of these same metal ions with respect to the hydration of 2-pyridinecarboxaldehyde resulted in a similar order22 Cu2+> Co2+> Ni2+> Zn2+ > Cd2+> Mn2+ In both series, catalysis by copper ion is considerably greater than of all other metal ions in this series while that due to cadmium or manganese is relatively poor. Although the two series differ in the positions of cobalt, nickel, and zinc ions, these three ions have almost the same catalytic efficiency toward the hydration of 2pyridinecarboxaldehyde (Table 111). We may also add that the stability constants of metal complexes of N,N'-di-(2-hydroxybenzyl)-ethylenediarnine-NlNdiacetic acid follow the order2a
I
1
I
0.2
0.4
1
0.6 0.8 1.0 CBCAJ x 1o4(mois I.-')
I
1,2
I
1.4
I
Figure 4. Bovine carbonic anhydrase catalysis of the dehydration of 2,2-dihydroxypropionate in 0.1 M diethylmalonate buffer at pH 6.7 and 0.0".
For the hydration of 2-pyridinecarboxaldehyde1 we have postulated that the special arrangement of the ring nitrogen and aldehydic group is necessary for the powerful catalysis afforded by divalent transition metal ions.z2 On the other hand, the corresponding hydration of 4-pyridinecarbo~aldehyde~where the proximity of the ring nitrogen and aldehydic group is absent, is not appreciably susceptible to metal ion catalysis.22 It has also been observed that the hydrations of aliphatic aldehydes which lack such a special (22) Y. Pocker and J. E. Meany, J. Phys. Chem., 7 2 , 655 (1965). (23) F. L. Eplattenier, I. Murase, and A. E. Martell, J. Amer. Chem. Soc., 74,2036 (1952).
Volume 74, Number 7 April 9, 1070
Y. POCKER AND J. E. MEANY
1490 ~~
~~~~~~~~
Table 11: Catalysis of 2,2-Dihydroxypropionate Dehydration by Bovine Carbonic Anhydrase and Various Metal Ions a t O.OQ
Table I11 : Comparison of Metal Ion Catalysis in the Dehydration of 2,2-Dihydroxypropionate and the Hydration of 2-Pyridinecarboxaldehyde 2 ,P-Dihydroxy-
Species
CUB+ Zne + Zn2+
Zn2+
Nil+ cos +
BCA BCA Cdz+ Mna+
propionate,
Buffer
0.1 M 0.1 M 0.1 M 0 1M 0.1 M 0.1 M 0.1 M 0 1M 0.1 M 0.1 M I
I
DEM DEM DEM Acetate DEM DEM DEM Acetate DEM
DEM
6.7 6.7
10,300 3,020
6.4
3,060
5.2 6.7 6.7 6.7 5.3 6.7 6.7
4,920 1,580 1,050 1,040 3,250 185 150
arrangement of basic groups are not accelerated by metal ions except when in the presence of specific environments, such as imidazole buffers.17 In the present work, we have chosen a substrate for which several metal ion complexes are known.+" Again, their ability to chelate to these cations may parallel the accelerative effect these ions have on the process under investigation. One may envisage two limiting possibilities for metal ion assistance, (a) Participation by the direct transfer of water
(1)
(11)
(111)
(b) (i) Participation by acting as a conveniently located general acid
(W
(V)
(ii) I n a similar scheme, one may also invoke metal ion participation through the bridging of water (or hydroxide)
rvru
k (VIII)
Any Of the above mechanisms is consistent with the ability of the substrate to bind to these metal ions, the The Journal of Physical Chemistry
Species
(bMa '/kH,O)'
CUB + Zna + NP co2+ Cd2+ Mn2+
1.04 x 107 3.05 X lo6
f
1.59
x
106
1.06 X lo6 1.86 X 10' 1.51 X lo6
Z-Pyridinecarboxaldehyde,a ( k M 2 +/kH,O)b
2.44 X 106 6.22 X loe 6.88 X loE 7.33 x 106 1.11 x 108 1.02 x 10s
a See ref 13. b k ~ = ~ H % O[ 5 5 . 5 ] ; ka for 2,2-dihydroxypropionate is taken as 0.05 min-1; that for 2-pyridinecarboxaldehyde was taken as 0.25 rnin-'.
powerful catalysis afforded by these metal ions, and the principle of microscopic reversibility. According to Table 111, it is apparent that most of the ratios ICM2+/IC~,o for metal ion catalysis are greater for the hydration of 2-pyridinecarboxaldehyde than for the dehydration of 2,2-dihydroxypropionate. This may be due in part t o the fact that 2-pyridinecarboxaldehyde (and its hydrate) is more rigid than pyruvic acid (and its hydrate) thus enhancing the very special orientation required in the metal ion-substrate complex. Zinc ion catalysis in 0.1 M acetate buffers at pH 5.2 (Figure 3, Table 11) is appreciably more powerful than in 0.1 M diethylmalonate buffer at pH 6.7. There are two major variables which must be taken into account in this comparison. It is conceivable that zinc ions are less available for catalysis in diethylmalonate buffers due to chelation to diethylmalonate diIt is also possible that zinc ion catalysis is a function of pH, the catalysis of the dehydration process being more powerful a t lower values of pH. We are presently considering metal ion catalysis in a variety of buffers at various values of pH which should lead to a more definitive explanation of these data. The zinc metalloenzyme carbonic anhydrase from bovine erythrocytes was also found to catalyze the dehydration of 2,2-dihydroxypropionate. The catalytic were evaluated for this process from coefficients, IC,, the results of a series of runs in 0.1 M diethylmalonate buffers a t pH 6.7 in which only the concentration of the enzyme was varied. A plot of kobad against enzyme concentration gave a straight line, the slope of which is defined as IC,, (Figure 4, Table 11). This enzymatic acceleration affords the first example of a carbonic anhydrase catalyzed dehydration yielding the anion of a keto acid. It should be noted that the concentrations of diethylmalonate and acetate buffers used in the present work were necessarily ea. ten times higher than those ordi(24) V. 8.K. Nair and G. H. Nancollas, J. Chem. SOC.,4367 (1961);
v. s. K. Nair, ibid., 1450 (1965).
REVERSIBLE HYDRATION OF PYRUVIC ACID
1491
narily employed in these laboratories for such enzymatic experiment^.'^^'^^'^ The monoanions associated with the above buffers act as noncompetitive inhibitors of hydrase and esterase activity but as competitive inhibitors of dehydrase activity.25 Thus, it is probable that the enzymatic activity shown in Figure 4 represents a partially inhibited rate. Nevertheless, since the respective inhibitory potencies of these two buffer anions are very similar,25the IC, values reported in the two buffers (Table 11) probably represent rates inhibited by nearly proportional amounts. We find in the present work that the value of IC,, in the dehydration of 2,2-dihydroxypropionate anion to pyruvate anion at p H 5.3 (k,,, = 3250 1. mol-l min-l) is 3.1 times larger than the corresponding value a t p H 6.7 (ICenB = 1040 1. m01-l min-I). It is of considerable interest that the enzyme retains this much activity a t low pH. In general, the pH-rate profiles associated with the enzymatic hydration of COz26and of aliphatic 27,28 are aldehydes12J6and with the hydrolysis of defined as sigmoid curves having points of inflection around neutrality, showing maximum activity a t high values of p H and decreasing to almost negligible values of kenz at lower values of pH. These pH-rate profiles were found to be consistent with the view that the turnover rate, in all these reactions, varies as though dependent on the ionization of a group in the enzyme with pK, ar?und 7.0, only the basic form of which is active. On the other hand, the enzymatic dehydration of bicarbonate anion is again a sigmoid curve but of opposite ~ h a p e having , ~ ~ a~ similar ~ ~ point of inflection around pH 7 but showing a high plateau around p H 5-6 and a negligible rate at pH 8. Here the enzymatic rate varies as though dependent on the ionization of a group in the enzyme with the same pK value (-7.0)) only the acid form of which is a c t i ~ e . *It~ is~ interesting ~~ to note that the enzymatic dehydration of bicarbonate anion is also faster at pH 5.3 than at pH 6.7 but by a 10
-I
k;
0+2104 I.rnole-f rnin:' f
I
(1x1
(XI
Actually, it may be sufficient to have these anions interact with zinc through a water molecule which would function as a bridge.16 In any case the Zn2+Xor Znz+(OH2)X- ion-pair complexes would be additionally stabilized by interaction with a neighboring imidazolium (BH+) group.16 Significantly, the inhibitory effects of anions and the binding of HC03- both decrease with increasing pH as if the cationic binding site lies close to the zinc and has an apparent pK, around 7.0.25127 The hydration of pyruvic acid was investigated in the pH range 1.2-2.1. The experimental rate constants, kobsd, were converted to the respective forward rate by the constants, Icf, by multiplying the values of fraction of hydration corresponding to the pH at which each kinetic run was carried out (Table IV). The fractions of hydration, x,were deduced as a function of pH by means of nuclear magnetic resonance ~ p e c t r a . ~ The catalytic coefficients, ICfo and ICfH80+, were evaluated from the forward rate constants as a function of hydronium ion concentration (Figure 5 ) - Similar studies have been carried out by Strehlow5 a t 25.0" and our results appear to accord qualitatively with his to the extent that the spontaneous rate coefficient appears to be relatively large in comparison to the catalytic coefficient associated with hydronium ion. The data listed ~ several in Table V31 compare ratios of I C H ~ O + / ~for reversible hydration processes. Strehlows attributes
3
ko= 2 .O mi".-' I
I
Figure 5. Catalysis of the hydration of pyruvic acid by HzO and HaO a t 0.0". +
factor of only 2.20130 Work now in progress attempts to deduce the complete pH-rate profile as well as the kinetic parameters associated with the various elementary steps. Enzyme modification studies in our laboratory strongly indicate that the acidic form of an imidazole group in a histidine residue of the enzyme may play a key role as a proton donor in the catalytic dehydration of both bicarbonate and 2,2-dihydroxypropionate anions. However, this does not diminish the sjgnificance of coordination between anion and the essential zinc atom. Indeed, such an association would facilitate a general acid (e.g., imidazolium) promoted dehydration of HC0,- (e.g., via scheme IX) or of CH3C(OH)2C02(e.g., via scheme X)
(25) Y . Pocker and J. T. Stone, Biochemistry, 7,2936 (1968).
(26) (27) (28) (29) (30) (31)
J. C. Kernohan, Biochim. Biophys. Acta, 81, 346 (1964). Y. Pocker and J. T. Stone, Biochemistry, 6 , 668 (1967). Y . Pocker and J. T. Stone, ibid., 7, 4139 (1960). J. C. Kernohan, Biochim. B w p h y s . Acta, 96, 304 (1965). Y. Pocker, L. Guilbert, and R. Reaugh, unpublished results. Y. Pocker and D. G. Dickerson, J . P h y s . Chem., 73, 4005 (1969). Volume 74, Number 7 A p r i l 3,1970
Y. POCKER AND J. E.MEANY
1492 Table IV : Hydronium Ion Catalysis of the Hydration of Pyruvic Acid [HsO+] X 108,~ mol 1.-1 Xb
min-1 kf, min-1 kobsd,
1 .a9 0.78 5.42 4.22
2.83 0.79 6.42 5.07
2.17 0.79 5.82 4.60
2.83 0.79 6.46 5.10
4.31 0.84 7.98 6.70
5.85 0.84 8.74 7.34
5.48 0.84 9.11 7.65
7.80 0.86 11.67 10 03 I
a The average ionic strength for these runs was 0.06. Correspondingly, hydronium ion concentrations were deduced by dividing the experimentally determined values of aEaO+ by a mean activity coefficient of 0.82. * The fractions of hydration, x, refer here to a pyruvic acid concentration of 0.0479 M . It is important to note that the stoichiometric composition of the hydrate includes 2 additional molecules of water, and the fractions of hydration as determined by nmr vary significantly with the absolute concentration of water and hence of pyruvic acid (ref 7).
~
~
~
~
Table V : A Comparison of Catalysis of kHaOt/ko for Various Reversible Hydrat,ion Reactions Substrate
CHsCHO CHaCHpCHO (CH8)gCHCHO (CHs)aCCHO
k~80+/k0
120,000 130,000 194,000 251,000
Reference
17 31 31 31
Subatrate
2-PA 4-PA Pyruiic acid
kEaO+/ko
Referenoe
44,000 168,000 52
11 11 a
Data from the present study.
the relatively powerful spontaneous rate to intramolecular acid catalysis by the carboxyl group. Interestingly, our studies involving the dehydration of 2,2-dihydroxypropionate show very powerful catalysis by the various metal ions as discussed earlier. One might presume in this reaction, and its reverse, both of which involve anions, that the metal ion merely takes the place of a hydrogen ion which is responsible for the comparatively large spontaneous rate of hydration of
The Journal of Physical Chemistry
pyruvic acid. (See Structure VI, eq 8.) However, it must be remembered that, relative to protons, divalent metal ions are weak Lewis acids, and therefore the added role of the metal ion in facilitating hydration may be found either in its strongly polarizing effect on the carbonyl group of pyruvate or in its capacity to participate in the transfer of water by acting both as a bridging group and activator of a properly positioned water molecule in its hydration shell.
