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Enzymatic and electrochemical oxidation of uric acid
J. Electroanal. Chem., 95 (1979) 81--90
81
© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
ENZYMATIC AND ELECTROCHEMICAL OXIDATION OF URIC ACID A MECHANISM FOR THE PEROXIDASE CATALYZED OXIDATION OF URIC ACID
HENRY A. MARSH, Jr. and GLENN D R Y H U R S T *
Department of Chemistry, University of Oklahoma, Norman, Okla. 73019 (U.S.A.) (Received 20th March, 1978; in revised form 19th May, 1978)
ABSTRACT The oxidation of uric acid has been studied by thin-layer spectroelectrochemistry using an optically transparent gold minigrid electrode and enzymatically using peroxidase. Under both enzymatic and electrochemical conditions a u.v.-absorbing intermediate ()krnax = 302-304 nm between pH 7--9.3) is observed. This intermediate is proposed to be an imine-alcohol species formed by rapid hydration of the primary reaction product which is a diimine. Under the conditions employed in both the enzymatic and electrochemical studies the imine-alcohol intermediate is hydrated in a first order reaction with an observed rate constant of 3.5 x 10 - 3 s - 1 between pH 7--9.3. This study establishes the similarity between the electrochemical and enzymatic oxidation of uric acid and indicates that electrochemical investigations of the redox behavior of purines can provide unique insights into their biological redox properties.
INTRODUCTION
A recent report from this laboratory [1] supports the view that the electrochemical oxidation of uric acid (I, eqn. 1) at graphite and gold electrodes proceeds b y an initial 2e--2H + reaction to give a diimine species (II, eqn. 1). This diimine is very unstable in aqueous solution and is rapidly hydrated to an imine-alcohol 0
H O
0
H H
N H
H..N~.r~ + 2 H + + 2e
I
O
H K
0
N
N.. kt
N
O
K2/ H20 O
products
H20
* To whom correspondence should be addressed.
~ OH H N H IE
82 (III, eqn. 1) in a first-order reaction. The rate constant and order for the latter reaction were determined by double potential step chronoamperometry using a highly polished wax-impregnated spectroscopic graphite electrode. This electrode was used to minimize adsorption of the diimine which is particularly severe at conventional pyrolytic graphite electrodes [ 1 ]. The rate constant for hydration of the diimine to the imine-alcohol at pH 8 was found to be 32.5 s-1 . At lower and higher pH values this rate constant was at least as large as that observed at pH 8. The imine-alcohol is also unstable but it may be observed as a u.v.-absorbing intermediate upon electrochemical oxidation of uric acid between pH 7--9 by means of thin-layer spectroelectrochemical studies. The imine-alcohol is hydrated in a first-order reaction to give, probably, uric acid-4,5-diol (IV, eqn. 1). Uric acid-4,5-diol then decomposes to the ultimate products which are, at pH 7--9, allantoin, urea and CO2 [2,3]. An underlying rationale for our studies of the electrochemical oxidation of biologically important purines and other N-heterocyclic molecules is that such studies can lead to information bearing on the reaction routes and mechanisms of biological oxidation reactions [3]. Although there are still some uncertainties regarding the electrochemical oxidation of uric acid (e.g., the actual structures of the diimine and imine-alcohol intermediates) the basic reaction mechanism is now largely understood. Accordingly, an investigation was initiated into the enzymatic oxidation of uric acid in order to develop sufficient data to compare the biological and electrochemical processes. The biological oxidation of uric acid with various peroxidase enzymes has probably been investigated most extensively. Earlier work [2--4] has indicated that the products formed upon both electrochemical and enzymatic (peroxidase) oxidation of uric acid are the same. Thus, at low pH the major electrochemical and enzymatic products are alloxan, urea and CO2, while at intermediate pH allantoin, urea and CO2, are formed. Paul and Avi-Dor [5] found that 1-methyluric acid is oxidized in the presence of horseradish peroxidase and that the oxidation is most rapid at about pH 4.0 when the molecule is uncharged (pKa = 5.75). The latter workers proposed that some u n k n o w n primary oxidation product is formed which then undergoes non-enzymatic, pH-dependent decomposition to alloxan or allantoin. Similarly, Canellakis et al. [6], who studied the oxidation of uric acid with a number o f different peroxidase enzymes, also proposed the existence of several u n k n o w n intermediates (including uric acid-4,5-diol) in the enzyme catalyzed process. The work reported here was undertaken to study the oxidation of unsubstit u t e d uric acid in the presence of peroxidase enzymes and H202 and to compare the observed behavior with t h a t observed upon electrochemical oxidation of uric acid. We have utilized primarily rapid scanning spectrophotometric techniques and ultraviolet kinetic measurements to compare the enzyme catalyzed and electrochemical oxidation processes. EXPERIMENTAL
Chemicals Uric acid was obtained from Eastman and was used w i t h o u t additional purification. Peroxidase, obtained from horseradish (crude, RZ ~ 0.3; type VI,
83 RZ ~ 3 . 0 ; t y p e VII, RZ ~ 3.0; type VIII, RZ ~ 3.0 and type IX, RZ ~ 3.2) was obtained from Sigma. The crude enzyme and the various isoperoxidases were stored at --5°C when n o t in use. Catalase was obtained from Sigma and had an activity of 2000 units mg -1 at 25 ° C. Buffer solutions were prepared from reagent grade chemicals. Al! buffers were prepared from potassium phosphates and had an ionic strength of 1.0 M with the exception of certain phosphate buffers at pH 4.0 which had an ionic strength of 0.1 M.
Apparatus Optical measurements utilized either a Harrick rapid scan spectrophotometer r.s.s.) and signal processing module (Harrick Scientific Co., Ossining, N.Y.) or a Hitachi-Perkin-Elmer Model 124 spectrophotometer. Repetitive spectral sweeps utilizing the r.s.s, were recorded on a Tektronix Model 5031 dual beam storage oscilloscope equipped with a Tektronix Model C-70 camera. Absorbance versus time curves were recorded on either a Hewlett-Packard Model 7001A XY recorder or a Sargent Model SRG recorder. For u.v. studies on the enzymatic oxidation of uric acid 1.0 cm optically matched quartz cells were utilized. Optically transparent thin-layer electrochemical cells were similar to those described by Murray et al. [7] and Heineman et al. [8] and utilized a gold minigrid as the optically transparent electrode. The gold minigrid employed had 1000 wires per inch and a transmittance of ca. 50%. Details of the cell design and construction have been presented elsewhere [7--9]. The thin-layer cells used in this study ranged in thickness from 0.009--0.010 cm. Controlled potentials applied to the thin-layer cell were maintained with a Wenking Model LT 73 potentiostat.
Procedure for enzymatic oxidation o f uric acid The most convenient m e t h o d for enzymatic oxidation of uric acid was to prepare suitable solutions of uric acid, t y p e VIII peroxidase, and H202 in a pH 4.0 phosphate buffer having an ionic strength of 0.1 M. Then, 0.5 ml of the peroxidase and 0.5 ml of the uric acid solutions were transferred to a 1.0 cm quartz spectrophotometer cell and the cell placed into the spectrophotometer. The oxidation of uric acid was initiated by adding 0.5 ml of H202 solution. Mixing was achieved by rapidly drawing the resulting solution in and out of a disposable pipet using a small pipetting bulb. The oxidation of uric acid was followed by monitoring its absorbance at 290 nm. When 95% or more of the uric acid had been oxidized, 1.5 ml of a phosphate buffer of the desired pH having an ionic strength of 1.0 M was added to the solution and the wavelength being monitored was switched to 302 nm. The latter wavelength corresponds to the ~max of the absorbing intermediate species between pH 7--9.3. Mixing was achieved in the same manner described previously. For experiments where the oxidation of uric acid with peroxidase and H202 was terminated with the enzyme catalase the following procedures were employed. A solution was prepared from 0.5 ml each of suitable solutions of uric acid and type VIII peroxidase in phosphate buffer pH 4.0, ionic strength
84
0.1 M. The reaction was initiated with a mixture of 0.5 ml of 1200 t~M H2Oe in pH 4.0 phosphate buffer, ionic strength 0.1 M and 1 ml of pH 7.0 phosphate buffer, ionic strength 1.0 M. The latter volume of pH 7.0 buffer was sufficient to give the entire solution a pH of 7.0. When the oxidation of uric acid was ~>95% complete 0.5 ml of pH 7.0 phosphate buffer, ionic strength 1.0 M, containing 0.5 mg catalase was added and the solution mixed thoroughly. The absorbance of the u.v.-absorbing intermediate at 302 nm was then monitored as a function of time. For the pH-step method, the oxidation of uric acid in the presence of peroxidase and H 2 0 2 was carried o u t at pH 4.0 in a phosphate buffer of ionic strength 0.1 M, and when oxidation of the uric acid was ~>95% complete the pH was adjusted with 1.5 ml of a phosphate buffer pH 7.0 with an ionic strength of 1.0 M which contained 0.5 mg of catalase. The solution was mixed thoroughly and again the reaction of the u.v.-absorbing intermediate was followed spectrophotometrically at 302 nm. Test solutions utilized for thin-layer spectroelectrochemical studies were 5 mM in uric acid. The buffer systems consisted of 50% of a pH 4.0 phosphate buffer having an ionic strength of 0.1 M and 50% of a phosphate buffer of the desired pH value having an ionic strength of 1.0 M. The procedures used in thin-layer spectroelectrochemical studies of uric acid were essentially the same as those previously reported [ 1,9 ] except that in studies involving repetitive spectral sweeps a rather larger spectral range (225--375 nm) was scanned. R.s.s. studies of the enzymatic oxidation of uric acid scanned the same spectral region. RESULTS AND DISCUSSION
Preliminary studies of the enzymatic oxidation of uric acid revealed that crude horseradish peroxidase had insufficient activity (i.e., caused only a slow reaction of uric acid) for our purposes. Accordingly, a number of partially purified isoperoxidases isolated from horseradish peroxidase were studied. Type VI peroxidase contains two basic isoenzymes (the term basic indicates that the enzyme migrates towards the cathode in electrophoresis experiments of the t y p e described b y Davis [10] b u t exhibited very little activity for uric acid. Type VII peroxidase, an acidic isoenzyme exhibited only slight uric acid activity. Type IX peroxidase, a basic isoenzyme exhibited no uric acid activity. However, t y p e VIII peroxidase, another acid isoenzyme, exhibited quite high activity for uric acid and this isoenzyme was therefore used in our investigations. The first studies o f the enzymatic and electrochemical oxidation of uric acid were carried out at pH 7.0 using rapid scan spectrophotometry. Typical spectra obtained at pH 7 during the two methods of oxidation are shown in Figs. 1 and 2. Curve 1 in the latter Figures corresponds to the spectrum of uric acid (~max = 287 nm) before initiation of any oxidation. Upon initiation of the oxidation, either enzymatically or under thin-layer spectroelectrochemical conditions, the u.v. peak of uric acid decreases and, correspondingly, a new peak grows in at longer wavelengths. Curve'2 in Fig. i is the spectrum observed about 5 min after initiation of the enzymatic oxidation. The absorbance between 350--320 nm has clearly increased over that observed for the initial uric acid solution.
85 I
_
{
I
I
1
I
m
I
I
I
T
, fi/i
I
/"1 m
II""r"" m
{
I
I
38O
340
300 ~/nm
I 260
I 220
}
{
{
{
{
380
340
300' ),/r,m
260
220
Fig. 1. Rapid scan spectrophotometry of uric acid (100///I//) in the presence of type VIII peroxidase (0,17 p.M) and H 2 0 2 (200 p_M) in phosphate buffer pH 7.0 having an ionic strength of 1.0 M. Curve 1 is the spectrum before initiation of the oxidation. Curve 2 is the spectrum after 5 min oxidation. Repetitive scans of 19 s. Fig. 2. Thin-layer spectroelectrochemical oxidation of uric acid (5 raM) at a gold minigrid electrode at 0.9 V in phosphate buffer pH 7.0 having an ionic strength of 1.0 M. Curve 1 is the spectrum before initiation of the oxidation. Curve 2 is the spectrum when 90% of the uric acid has been oxidized. Repetitive sweeps of 9.4 s.
In Fig. 2 the effect is much more apparent. Thus, curve 2 in Fig. 2 is the spectrum observed a b o u t 30 s after initiation of the electrooxidation of uric acid. In this case the absorbance between 350--320 nm has increased compared to that observed before electrooxidation. However, the Xma~ of the new absorbing species clearly occurs at 302--304 nm. With increasing time the increased absorbance observed at 350--320 nm increases as the uric acid peak diminishes, reaches a maximal value (corresponding to a b o u t curve 2 in Fig. 2) and then decreases. Clearly, the species responsible for the increased absorbance at 350-320 nm must be an unstable intermediate formed in the enzymatic and electrochemical oxidation of uric acid. After complete enzymatic or electrochemical oxidation of uric acid, i.e., after the intermediate species has completely disappeared, the spectrum of the solution corresponds exactly to that of allantoin, the known electrochemical [1] and enzymatic [5,6] product. The r.s.s, curves shown in Fig. 1 for the peroxidase oxidation of uric acid are not as well defined as those shown for the electrochemical oxidation in Fig. 2. This is so because the peroxidase catalyzed oxidation of uric acid is quite slow at pH 7. For example, at pH 7.0 in phosphate buffer having an ionic strength of 1.0 M the observed first order rate constant for oxidation of uric acid (100 pM) in the presence of t y p e VIII peroxidase (0.17/aM) and H202 (200/aM) at 25°C is 0.0032 s -1. This value is significantly smaller than the observed rate constant at pH 4 under otherwise identical conditions (0.0092 s - l ) . Thus, the rate of oxidation of uric acid at pH 7 is quite slow and, as will be shown subsequently, the rate of disappearance of the u.v.-absorbing intermediate is not greatly different from the rate of oxidation of uric acid. Thus at pH 7, under the conditions outlined above, it is n o t possible to observe the spectrum of the intermediate species without therz also being significant quantities of absorbing uric acid present in the solution. The rate of peroxidase-catalyzed oxidation of
86
uric acid may be increased by increasing the enzyme concentration. However, the cost o f the enzyme precluded extensive studies at very high peroxidase concentrations. Kinetic measurements
In order to characterize the u.v.-absorbing intermediate formed on electrochemical oxidation of uric acid, thin-layer spectroelectrochemical studies were carried out between pH 7 and 9.3. Below pH 7 the molar absorptivity of the intermediate species decreased very significantly such that it could n o t be monitored. Between pH 7 and 9.3 the u.v.-absorbing intermediate had a Xmax of 302--304 nm, i.e. ~kma x w a s essentially independent o f pH. The kinetics of disappearance of the u.v.-absorbing intermediate was studied by applying a constant potential to an optically transparent gold minigrid electrode in a thinlayer spectroelectrochemical cell and monitoring the absorbance o f the intermediate species at a wavelength where uric acid exhibited minimal absorbance (e.g. 320 nm). When the absorbance at the latter wavelength reached its maximal value, the electrolysis was terminated and the A 3 2 0 n m v e r s u s time decay was recorded. Analysis of these A versus time curves indicated that the disappearance of the u.v.-absorbing intermediate followed first order kinetics. Some typical first order rate constants are presented in Table 1. It should be noted that the observed rate constants reported in Table 1 are somewhat smaller than those (0.008 s-1 between pH 7 and 9) reported previously [1]. The latter values were obtained in a different buffer system (phosphate plus K2SO4) having an ionic strength of 0.5 M. This suggests t h a t the rate of disappearance of the u.v.absorbing intermediate is affected by ionic strength and/or the buffer constituents. However, no systematic investigation o f this effect has been carried out. Attempts to study the kinetics of disappearance of the u.v.-absorbing intermediate observed on peroxidase/H202 oxidation of uric acid at pH 7 and above were n o t meaningful because the oxidation of uric acid is so slow that the absorbance versus time curves reflect a combination of the rate of oxidation of uric acid and the reaction of the intermediate species. In order to circumvent this problem two experimental approaches were employed. The first method used involved performing the enzymatic oxidation at pH 4.0 where the peroxidase catalyzed oxidation of uric acid occurs most rapidly [5]. Then, when >195%
TABLE 1 Observed first o r d e r r a t e c o n s t a n t s for r e a c t i o n o f t h e u.v.-absorbing i n t e r m e d i a t e f o r m e d o n e l e c t r o c h e m i c a l o x i d a t i o n o f uric acid a t a gold m i n i g r i d e l e c t r o d e Initial c o n c e n t r a t i o n o f uric a c i d / m M
pH a
kobs/s_l
5 5 5
7.0 8.0 9.3
0.0040 0.0030 0.0035 Mean: 0.0035 + 0.0005
a P h o s p h a t e b u f f e r s , see E x p e r i m e n t a l for details of b u f f e r c o m p o s i t i o n .
87
of the uric acid was oxidized the pH was rapidly adjusted to 7 or higher, where the intermediate species absorbs strongly, and the kinetics of reaction of the intermediate could be easily followed b y a spectrophotometric method. This method will be referred to as the pH-step method. A typical rapid scan spectral response using the pH-step method is shown in Fig. 3. Thus, after rapid peroxidase catalyzed oxidation of uric acid at pH 4.0 then, on adjusting the pH to 7.0, the spectrum of the intermediate is essentially identical to that observed for the intermediate obtained upon electrochemical oxidation of uric acid i.e. ~max ~ 302 nm (see Fig. 2). That the spectrum of the u.v.-absorbing intermediate formed u p o n enzymatic oxidation of uric acid is the same as that formed b y electrochemical oxidation is shown more clearly in Fig. 4. The kinetics of disappearance of the u.v.-absorbing intermediate observed upon peroxidase oxidation of uric acid using the pH-step method described was studied b y monitoring the absorbance of the intermediate at 302 nm as a function of time (Fig. 5A). A typical rate plot for such an experiment is presented in Fig. 5B. Clearly, the plot of log A--A+ versus time is linear indicating that the reaction of the absorbing intermediate species follows first order kinetics. Values of the observed first order solution rate constants for reaction of the u.v.-absorbing intermediate obtaining using the pH-step m e t h o d b e t w e e n pH 7 and 9.3 are presented in Table 2. Clearly, b e t w e e n pH 7 and 9.3 the value of the observed first order rate constant (0.0035 s-1) for reaction o f the u.v.-absorbing inter-
I
I
I
I
+L_I
.,,~t
,
I A=0.4
tw +
, ] ~
x
380
I 340
I 300 ~/nm
I 260
I 220
I
I
I
I
I
I
375 355 335 315 295 275 ;',/n rr,
Fig. 3. Rapid scan spectrophotometry of the absorbing intermediate formed upon oxidation of uric acid (200 p_M) with type VIII peroxidase (0.34 ~tM) and H 2 0 2 (400 p.M) at pH 4.0. Following oxidation of uric acid the pH is adjusted to 7.0. Curve 1 is the initial spectrum, curve 2 that obtained about 11 min later. Repetitive scans of 47 s. Fig. 4. Comparison of the spectrum of the u.v.-absorbing intermediate formed upon electrochemical (x) and enzymatic (.) oxidation of uric acid. The enzymatic data points were obtained when uric acid (100 p.M) was oxidized in the presence of type VIII peroxidase (0.17 p.M) and H 2 0 2 (200 ]2M) in phosphate buffer pH 7.0, ionic strength 1.0 M 5 rain after initiation of the oxidation in a 1.0 cm cell. The electrochemical points were obtained when uric acid (5 raM) was oxidized at 0.9 V in a thin-layer cell (90 pM thick) at 0.9 V at a gold minigrid electrode at a time when/> 90% of the uric acid was oxidized. The absorbance values obtained in the latter experiment were all multiplied by a constant factor of 2.24 to obtain the best fit to the enzymatic data points.
88 0.8 0.7
_E0.6 o~
0.5
,~ 0.4 "~ 0.3
~ 0.2 0.1 I
I
I
200
400
600
800
1000
1200
Time/sec
600 B
500 K = 0,0035 s i
400
•
300
2OO
100
-1.4 -1.3
-1.2 -1.1
-1.0
-0.9
-0.8
-0.7
-0.6
-0.5
-0.4
log A -- A ~
Fig. 5. (A) Variation of the absorbance with time observed for the intermediate formed upon oxidation of uric acid (200 pM) with type VIII peroxidase (0.34/.tM) and H 2 0 2 (400 pM) at pH 4 followed by adjustment of the pH to 7.0. (B) Kinetic plot of time versus log absorbanee at 302 nm.
mediate formed enzymatically is identical to that formed electrochemically (see Table 1). The fact that the observed rate constant is independent of pH, at least between pH 7 and 9.3 indicates that the reaction of the intermediate is probably not base catalyzed over this pH range. The data presented in Table 2 also support the conclusion that the reaction of the intermediate is independent o f the enzyme and uric acid concentration. The rate of oxidation of uric acid is, however, directly dependent on the enzyme concentration. Attempts to measure the reaction kinetics for the u.v.-absorbing intermediate below pH 7 were difficult because the molar absorptivity of the intermediate significantly decreases. In addition, thin-layer spectroelectrochemical studies which, because of the very thin reaction cell used, require rather high uric acid concentrations (/>2 mM) to produce measurable concentrations of the intermediate, become very difficult below pH 7 owing to the low solubility of uric acid. However, using the pH-step enzymatic method, the approximate first order rate constant for reaction of the intermediate at pH 6.0 was ca 0.006 s -1
89 TABLE 2 Observed first o r d e r rate c o n s t a n t s for r e a c t i o n of t h e u . v . - a b s o r b i n g i n t e r m e d i a t e f o r m e d o n o x i d a t i o n of uric acid in t h e p r e s e n c e o f t y p e VIII p e r o x i d a s e a n d H 2 0 2 at pH 4.0 f o l l o w e d b y pH a d j u s t m e n t Initial concentration of uric a c i d / p M
pH
HRP concentration/ pM
H202 concentration/ pM
kobs/s -1
20 100 200 100 100
7.0 7.0 7.0 7.0 7.0
0.17 0.17 0.17 0.34 0.034
200 200 200 200 200
0.0039 0.0039 0.0035 0.0037 0.0036
20 100 200
8.0 8.0 8.0
0.17 0.17 0.17
200 200 200
0.0036 0.0030 0.0032
20 100 200
9.27 9.27 9.27
0.17 0.17 0.17
200 0.0033 200 0.0032 200 0.0032 M e a n : 0 . 0 0 3 5 -+ 0 . 0 0 0 4
and at pH 4.8, 0.0095 s-1. These values are much less reliable than those obtained at pH 7--9.3 but they do suggest t h a t perhaps the reaction o f the inter. mediate might be somewhat acid catalyzed below pH 7. In order that the conditions of the enzymatic oxidation paralleled more closely the electrochemical oxidation of uric acid, it was decided to a t t e m p t to oxidize uric acid at pH 7 with type VIII peroxidase and, when a significant q u a n t i t y of the u.v.-absorbing intermediate had been formed, to terminate the oxidation reaction so that the decay of the intermediate could be observed w i t h o u t interference from the continuing oxidation of uric acid. It was found that the most successful method of doing this was to terminate the oxidation of uric acid by addition of a large excess of the enzyme catalase. The latter enzyme rapidly r e m o v e s H 2 0 2 from the reaction mixture. Using catalase to terminate the reaction of uric acid, H202 and peroxidase at pH 7.0 in this fashion (see Experimental for details) it was f o u n d t h a t the u.v.-absorbing intermediate gave an observed first order rate constant of 0.0035 s- i . If the oxidation o f uric acid with peroxidase and H202 was carried out at pH 4.0 and then, when the uric acid was virtually all oxidized, the pH adjusted to 7.0 with a buffer solution containing catalase, the observed first order rate constant for reaction of the u.v.-absorbing intermediate was 0.0034 s- i . CONCLUSIONS
The studies reported here indicate that when uric acid is oxidized either electrochemically or with type VIII peroxidase an intermediate is formed which exhibits a well
90 less of whether it is generated electrochemically or enzymatically. The evidence, therefore, strongly supports the view that the same u.v.-absorbing intermediate is generated b y both electrochemical and peroxidase catalyzed oxidation of uric acid. In our earlier report we demonstrated that under electrochemical conditions at gold or graphite electrodes uric acid is oxidized in a 2e reaction to a diimine species (II, eqn. 1) which is very short-lived (e.g., it has a half-life at pH 8.0 of a b o u t 21 ms). This diimine primary product is attacked b y water in a (pseudo) first order reaction to give the u.v.-absorbing intermediate described in this paper which has been proposed [1] to be an imine-alcohol species (III, eqn. 1). Clearly, since the imine-alcohol is formed under b o t h electrochemical and enzymatic (peroxidase) conditions it seems quite justifiable to conclude that this species is derived from the same precursor under both oxidative conditions, i.e., from the diimine. There is no simple way to observe the diimine precursor formed u p o n enzymatic oxidation of uric acid because it is t o o unstable. The imine-alcohol formed either electrochemically or enzymatically is hydrated to the diol IV (eqn. 1) which then breaks d o w n to the various observed products. We believe that this study establishes for the first time that investigations of the electrochemical behavior of naturally occurring biological molecules can provide unique insights into the possible biological redox reactions of such molecules, Indeed, in the case of the electrochemical and enzymatic (peroxidase) oxidation of uric acid there is every evidence that the reactions are identical. Further investigations are currently underway on a variety of other purine enzymatic and electrochemical oxidations to further substantiate the similarity b e t w e e n the latter redox mechanisms for these compounds. ACKNOWLEDGMENTS
The authors gratefully acknowledge the help and guidance provided b y Dr. Eddie C. Smith regarding the enzymatic oxidation of uric acid. This w o r k was supported b y the National Institutes of Health through Grant No. GM 21034-04.
REFERENCES 1 2 3 4 5 6 7 8 9 10
J . L . Owens, H.A. Marsh a n d G. D r y h u r s t , J. E l e c t r o a n a l . Chem., 91 ( 1 9 7 8 ) 2 3 1 . G. D r y h u r s t , J. E l e c t r o c h e m . Soc., 1 1 9 ( 1 9 7 2 ) 1 5 5 9 . G. D r y h u r s t , E l e c t r o c h e m i s t r y o f Biological Molecules, A c a d e m i c Press, New Y o r k , 1 9 7 7 , Ch. 3. G. D r y h u r s t , Topics in C u r r e n t C h e m i s t r y , 3 4 ( 1 9 7 2 ) 47. K.G. Paul a n d Y. Avi-Dor, A e t a C h e m . Scand., 8 ( 1 9 5 4 ) 6 3 7 . E.S. CaneUakis, A.L. Turtle a n d P.P. C o h e n , J. Biol. C h e m . , 2 1 3 ( 1 9 5 5 ) 3 9 7 . R.W. M u r r a y , W.R. H e i n e m a n a n d G.W. O ' D o m , Anal. Chem., 39 ( 1 9 6 7 ) 1 6 6 6 . W.R. H e i n e m a n , B.J. Norris a n d I.F. Goelz, Anal. C h e m . , 47 ( 1 9 7 5 ) 79. J.L. O w e n s a n d G. D r y h u r s t , J. E l e c t r o a n a l . C h e m . , 70 ( 1 9 7 6 ) 1 9 9 . B.J. Davis, A n n . N.Y. A c a d . Sci., 1 2 1 ( 1 9 6 4 ) 4 0 4 .
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© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
ENZYMATIC AND ELECTROCHEMICAL OXIDATION OF URIC ACID A MECHANISM FOR THE PEROXIDASE CATALYZED OXIDATION OF URIC ACID
HENRY A. MARSH, Jr. and GLENN D R Y H U R S T *
Department of Chemistry, University of Oklahoma, Norman, Okla. 73019 (U.S.A.) (Received 20th March, 1978; in revised form 19th May, 1978)
ABSTRACT The oxidation of uric acid has been studied by thin-layer spectroelectrochemistry using an optically transparent gold minigrid electrode and enzymatically using peroxidase. Under both enzymatic and electrochemical conditions a u.v.-absorbing intermediate ()krnax = 302-304 nm between pH 7--9.3) is observed. This intermediate is proposed to be an imine-alcohol species formed by rapid hydration of the primary reaction product which is a diimine. Under the conditions employed in both the enzymatic and electrochemical studies the imine-alcohol intermediate is hydrated in a first order reaction with an observed rate constant of 3.5 x 10 - 3 s - 1 between pH 7--9.3. This study establishes the similarity between the electrochemical and enzymatic oxidation of uric acid and indicates that electrochemical investigations of the redox behavior of purines can provide unique insights into their biological redox properties.
INTRODUCTION
A recent report from this laboratory [1] supports the view that the electrochemical oxidation of uric acid (I, eqn. 1) at graphite and gold electrodes proceeds b y an initial 2e--2H + reaction to give a diimine species (II, eqn. 1). This diimine is very unstable in aqueous solution and is rapidly hydrated to an imine-alcohol 0
H O
0
H H
N H
H..N~.r~ + 2 H + + 2e
I
O
H K
0
N
N.. kt
N
O
K2/ H20 O
products
H20
* To whom correspondence should be addressed.
~ OH H N H IE
82 (III, eqn. 1) in a first-order reaction. The rate constant and order for the latter reaction were determined by double potential step chronoamperometry using a highly polished wax-impregnated spectroscopic graphite electrode. This electrode was used to minimize adsorption of the diimine which is particularly severe at conventional pyrolytic graphite electrodes [ 1 ]. The rate constant for hydration of the diimine to the imine-alcohol at pH 8 was found to be 32.5 s-1 . At lower and higher pH values this rate constant was at least as large as that observed at pH 8. The imine-alcohol is also unstable but it may be observed as a u.v.-absorbing intermediate upon electrochemical oxidation of uric acid between pH 7--9 by means of thin-layer spectroelectrochemical studies. The imine-alcohol is hydrated in a first-order reaction to give, probably, uric acid-4,5-diol (IV, eqn. 1). Uric acid-4,5-diol then decomposes to the ultimate products which are, at pH 7--9, allantoin, urea and CO2 [2,3]. An underlying rationale for our studies of the electrochemical oxidation of biologically important purines and other N-heterocyclic molecules is that such studies can lead to information bearing on the reaction routes and mechanisms of biological oxidation reactions [3]. Although there are still some uncertainties regarding the electrochemical oxidation of uric acid (e.g., the actual structures of the diimine and imine-alcohol intermediates) the basic reaction mechanism is now largely understood. Accordingly, an investigation was initiated into the enzymatic oxidation of uric acid in order to develop sufficient data to compare the biological and electrochemical processes. The biological oxidation of uric acid with various peroxidase enzymes has probably been investigated most extensively. Earlier work [2--4] has indicated that the products formed upon both electrochemical and enzymatic (peroxidase) oxidation of uric acid are the same. Thus, at low pH the major electrochemical and enzymatic products are alloxan, urea and CO2, while at intermediate pH allantoin, urea and CO2, are formed. Paul and Avi-Dor [5] found that 1-methyluric acid is oxidized in the presence of horseradish peroxidase and that the oxidation is most rapid at about pH 4.0 when the molecule is uncharged (pKa = 5.75). The latter workers proposed that some u n k n o w n primary oxidation product is formed which then undergoes non-enzymatic, pH-dependent decomposition to alloxan or allantoin. Similarly, Canellakis et al. [6], who studied the oxidation of uric acid with a number o f different peroxidase enzymes, also proposed the existence of several u n k n o w n intermediates (including uric acid-4,5-diol) in the enzyme catalyzed process. The work reported here was undertaken to study the oxidation of unsubstit u t e d uric acid in the presence of peroxidase enzymes and H202 and to compare the observed behavior with t h a t observed upon electrochemical oxidation of uric acid. We have utilized primarily rapid scanning spectrophotometric techniques and ultraviolet kinetic measurements to compare the enzyme catalyzed and electrochemical oxidation processes. EXPERIMENTAL
Chemicals Uric acid was obtained from Eastman and was used w i t h o u t additional purification. Peroxidase, obtained from horseradish (crude, RZ ~ 0.3; type VI,
83 RZ ~ 3 . 0 ; t y p e VII, RZ ~ 3.0; type VIII, RZ ~ 3.0 and type IX, RZ ~ 3.2) was obtained from Sigma. The crude enzyme and the various isoperoxidases were stored at --5°C when n o t in use. Catalase was obtained from Sigma and had an activity of 2000 units mg -1 at 25 ° C. Buffer solutions were prepared from reagent grade chemicals. Al! buffers were prepared from potassium phosphates and had an ionic strength of 1.0 M with the exception of certain phosphate buffers at pH 4.0 which had an ionic strength of 0.1 M.
Apparatus Optical measurements utilized either a Harrick rapid scan spectrophotometer r.s.s.) and signal processing module (Harrick Scientific Co., Ossining, N.Y.) or a Hitachi-Perkin-Elmer Model 124 spectrophotometer. Repetitive spectral sweeps utilizing the r.s.s, were recorded on a Tektronix Model 5031 dual beam storage oscilloscope equipped with a Tektronix Model C-70 camera. Absorbance versus time curves were recorded on either a Hewlett-Packard Model 7001A XY recorder or a Sargent Model SRG recorder. For u.v. studies on the enzymatic oxidation of uric acid 1.0 cm optically matched quartz cells were utilized. Optically transparent thin-layer electrochemical cells were similar to those described by Murray et al. [7] and Heineman et al. [8] and utilized a gold minigrid as the optically transparent electrode. The gold minigrid employed had 1000 wires per inch and a transmittance of ca. 50%. Details of the cell design and construction have been presented elsewhere [7--9]. The thin-layer cells used in this study ranged in thickness from 0.009--0.010 cm. Controlled potentials applied to the thin-layer cell were maintained with a Wenking Model LT 73 potentiostat.
Procedure for enzymatic oxidation o f uric acid The most convenient m e t h o d for enzymatic oxidation of uric acid was to prepare suitable solutions of uric acid, t y p e VIII peroxidase, and H202 in a pH 4.0 phosphate buffer having an ionic strength of 0.1 M. Then, 0.5 ml of the peroxidase and 0.5 ml of the uric acid solutions were transferred to a 1.0 cm quartz spectrophotometer cell and the cell placed into the spectrophotometer. The oxidation of uric acid was initiated by adding 0.5 ml of H202 solution. Mixing was achieved by rapidly drawing the resulting solution in and out of a disposable pipet using a small pipetting bulb. The oxidation of uric acid was followed by monitoring its absorbance at 290 nm. When 95% or more of the uric acid had been oxidized, 1.5 ml of a phosphate buffer of the desired pH having an ionic strength of 1.0 M was added to the solution and the wavelength being monitored was switched to 302 nm. The latter wavelength corresponds to the ~max of the absorbing intermediate species between pH 7--9.3. Mixing was achieved in the same manner described previously. For experiments where the oxidation of uric acid with peroxidase and H202 was terminated with the enzyme catalase the following procedures were employed. A solution was prepared from 0.5 ml each of suitable solutions of uric acid and type VIII peroxidase in phosphate buffer pH 4.0, ionic strength
84
0.1 M. The reaction was initiated with a mixture of 0.5 ml of 1200 t~M H2Oe in pH 4.0 phosphate buffer, ionic strength 0.1 M and 1 ml of pH 7.0 phosphate buffer, ionic strength 1.0 M. The latter volume of pH 7.0 buffer was sufficient to give the entire solution a pH of 7.0. When the oxidation of uric acid was ~>95% complete 0.5 ml of pH 7.0 phosphate buffer, ionic strength 1.0 M, containing 0.5 mg catalase was added and the solution mixed thoroughly. The absorbance of the u.v.-absorbing intermediate at 302 nm was then monitored as a function of time. For the pH-step method, the oxidation of uric acid in the presence of peroxidase and H 2 0 2 was carried o u t at pH 4.0 in a phosphate buffer of ionic strength 0.1 M, and when oxidation of the uric acid was ~>95% complete the pH was adjusted with 1.5 ml of a phosphate buffer pH 7.0 with an ionic strength of 1.0 M which contained 0.5 mg of catalase. The solution was mixed thoroughly and again the reaction of the u.v.-absorbing intermediate was followed spectrophotometrically at 302 nm. Test solutions utilized for thin-layer spectroelectrochemical studies were 5 mM in uric acid. The buffer systems consisted of 50% of a pH 4.0 phosphate buffer having an ionic strength of 0.1 M and 50% of a phosphate buffer of the desired pH value having an ionic strength of 1.0 M. The procedures used in thin-layer spectroelectrochemical studies of uric acid were essentially the same as those previously reported [ 1,9 ] except that in studies involving repetitive spectral sweeps a rather larger spectral range (225--375 nm) was scanned. R.s.s. studies of the enzymatic oxidation of uric acid scanned the same spectral region. RESULTS AND DISCUSSION
Preliminary studies of the enzymatic oxidation of uric acid revealed that crude horseradish peroxidase had insufficient activity (i.e., caused only a slow reaction of uric acid) for our purposes. Accordingly, a number of partially purified isoperoxidases isolated from horseradish peroxidase were studied. Type VI peroxidase contains two basic isoenzymes (the term basic indicates that the enzyme migrates towards the cathode in electrophoresis experiments of the t y p e described b y Davis [10] b u t exhibited very little activity for uric acid. Type VII peroxidase, an acidic isoenzyme exhibited only slight uric acid activity. Type IX peroxidase, a basic isoenzyme exhibited no uric acid activity. However, t y p e VIII peroxidase, another acid isoenzyme, exhibited quite high activity for uric acid and this isoenzyme was therefore used in our investigations. The first studies o f the enzymatic and electrochemical oxidation of uric acid were carried out at pH 7.0 using rapid scan spectrophotometry. Typical spectra obtained at pH 7 during the two methods of oxidation are shown in Figs. 1 and 2. Curve 1 in the latter Figures corresponds to the spectrum of uric acid (~max = 287 nm) before initiation of any oxidation. Upon initiation of the oxidation, either enzymatically or under thin-layer spectroelectrochemical conditions, the u.v. peak of uric acid decreases and, correspondingly, a new peak grows in at longer wavelengths. Curve'2 in Fig. i is the spectrum observed about 5 min after initiation of the enzymatic oxidation. The absorbance between 350--320 nm has clearly increased over that observed for the initial uric acid solution.
85 I
_
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300 ~/nm
I 260
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Fig. 1. Rapid scan spectrophotometry of uric acid (100///I//) in the presence of type VIII peroxidase (0,17 p.M) and H 2 0 2 (200 p_M) in phosphate buffer pH 7.0 having an ionic strength of 1.0 M. Curve 1 is the spectrum before initiation of the oxidation. Curve 2 is the spectrum after 5 min oxidation. Repetitive scans of 19 s. Fig. 2. Thin-layer spectroelectrochemical oxidation of uric acid (5 raM) at a gold minigrid electrode at 0.9 V in phosphate buffer pH 7.0 having an ionic strength of 1.0 M. Curve 1 is the spectrum before initiation of the oxidation. Curve 2 is the spectrum when 90% of the uric acid has been oxidized. Repetitive sweeps of 9.4 s.
In Fig. 2 the effect is much more apparent. Thus, curve 2 in Fig. 2 is the spectrum observed a b o u t 30 s after initiation of the electrooxidation of uric acid. In this case the absorbance between 350--320 nm has increased compared to that observed before electrooxidation. However, the Xma~ of the new absorbing species clearly occurs at 302--304 nm. With increasing time the increased absorbance observed at 350--320 nm increases as the uric acid peak diminishes, reaches a maximal value (corresponding to a b o u t curve 2 in Fig. 2) and then decreases. Clearly, the species responsible for the increased absorbance at 350-320 nm must be an unstable intermediate formed in the enzymatic and electrochemical oxidation of uric acid. After complete enzymatic or electrochemical oxidation of uric acid, i.e., after the intermediate species has completely disappeared, the spectrum of the solution corresponds exactly to that of allantoin, the known electrochemical [1] and enzymatic [5,6] product. The r.s.s, curves shown in Fig. 1 for the peroxidase oxidation of uric acid are not as well defined as those shown for the electrochemical oxidation in Fig. 2. This is so because the peroxidase catalyzed oxidation of uric acid is quite slow at pH 7. For example, at pH 7.0 in phosphate buffer having an ionic strength of 1.0 M the observed first order rate constant for oxidation of uric acid (100 pM) in the presence of t y p e VIII peroxidase (0.17/aM) and H202 (200/aM) at 25°C is 0.0032 s -1. This value is significantly smaller than the observed rate constant at pH 4 under otherwise identical conditions (0.0092 s - l ) . Thus, the rate of oxidation of uric acid at pH 7 is quite slow and, as will be shown subsequently, the rate of disappearance of the u.v.-absorbing intermediate is not greatly different from the rate of oxidation of uric acid. Thus at pH 7, under the conditions outlined above, it is n o t possible to observe the spectrum of the intermediate species without therz also being significant quantities of absorbing uric acid present in the solution. The rate of peroxidase-catalyzed oxidation of
86
uric acid may be increased by increasing the enzyme concentration. However, the cost o f the enzyme precluded extensive studies at very high peroxidase concentrations. Kinetic measurements
In order to characterize the u.v.-absorbing intermediate formed on electrochemical oxidation of uric acid, thin-layer spectroelectrochemical studies were carried out between pH 7 and 9.3. Below pH 7 the molar absorptivity of the intermediate species decreased very significantly such that it could n o t be monitored. Between pH 7 and 9.3 the u.v.-absorbing intermediate had a Xmax of 302--304 nm, i.e. ~kma x w a s essentially independent o f pH. The kinetics of disappearance of the u.v.-absorbing intermediate was studied by applying a constant potential to an optically transparent gold minigrid electrode in a thinlayer spectroelectrochemical cell and monitoring the absorbance o f the intermediate species at a wavelength where uric acid exhibited minimal absorbance (e.g. 320 nm). When the absorbance at the latter wavelength reached its maximal value, the electrolysis was terminated and the A 3 2 0 n m v e r s u s time decay was recorded. Analysis of these A versus time curves indicated that the disappearance of the u.v.-absorbing intermediate followed first order kinetics. Some typical first order rate constants are presented in Table 1. It should be noted that the observed rate constants reported in Table 1 are somewhat smaller than those (0.008 s-1 between pH 7 and 9) reported previously [1]. The latter values were obtained in a different buffer system (phosphate plus K2SO4) having an ionic strength of 0.5 M. This suggests t h a t the rate of disappearance of the u.v.absorbing intermediate is affected by ionic strength and/or the buffer constituents. However, no systematic investigation o f this effect has been carried out. Attempts to study the kinetics of disappearance of the u.v.-absorbing intermediate observed on peroxidase/H202 oxidation of uric acid at pH 7 and above were n o t meaningful because the oxidation of uric acid is so slow that the absorbance versus time curves reflect a combination of the rate of oxidation of uric acid and the reaction of the intermediate species. In order to circumvent this problem two experimental approaches were employed. The first method used involved performing the enzymatic oxidation at pH 4.0 where the peroxidase catalyzed oxidation of uric acid occurs most rapidly [5]. Then, when >195%
TABLE 1 Observed first o r d e r r a t e c o n s t a n t s for r e a c t i o n o f t h e u.v.-absorbing i n t e r m e d i a t e f o r m e d o n e l e c t r o c h e m i c a l o x i d a t i o n o f uric acid a t a gold m i n i g r i d e l e c t r o d e Initial c o n c e n t r a t i o n o f uric a c i d / m M
pH a
kobs/s_l
5 5 5
7.0 8.0 9.3
0.0040 0.0030 0.0035 Mean: 0.0035 + 0.0005
a P h o s p h a t e b u f f e r s , see E x p e r i m e n t a l for details of b u f f e r c o m p o s i t i o n .
87
of the uric acid was oxidized the pH was rapidly adjusted to 7 or higher, where the intermediate species absorbs strongly, and the kinetics of reaction of the intermediate could be easily followed b y a spectrophotometric method. This method will be referred to as the pH-step method. A typical rapid scan spectral response using the pH-step method is shown in Fig. 3. Thus, after rapid peroxidase catalyzed oxidation of uric acid at pH 4.0 then, on adjusting the pH to 7.0, the spectrum of the intermediate is essentially identical to that observed for the intermediate obtained upon electrochemical oxidation of uric acid i.e. ~max ~ 302 nm (see Fig. 2). That the spectrum of the u.v.-absorbing intermediate formed u p o n enzymatic oxidation of uric acid is the same as that formed b y electrochemical oxidation is shown more clearly in Fig. 4. The kinetics of disappearance of the u.v.-absorbing intermediate observed upon peroxidase oxidation of uric acid using the pH-step method described was studied b y monitoring the absorbance of the intermediate at 302 nm as a function of time (Fig. 5A). A typical rate plot for such an experiment is presented in Fig. 5B. Clearly, the plot of log A--A+ versus time is linear indicating that the reaction of the absorbing intermediate species follows first order kinetics. Values of the observed first order solution rate constants for reaction of the u.v.-absorbing intermediate obtaining using the pH-step m e t h o d b e t w e e n pH 7 and 9.3 are presented in Table 2. Clearly, b e t w e e n pH 7 and 9.3 the value of the observed first order rate constant (0.0035 s-1) for reaction o f the u.v.-absorbing inter-
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Fig. 3. Rapid scan spectrophotometry of the absorbing intermediate formed upon oxidation of uric acid (200 p_M) with type VIII peroxidase (0.34 ~tM) and H 2 0 2 (400 p.M) at pH 4.0. Following oxidation of uric acid the pH is adjusted to 7.0. Curve 1 is the initial spectrum, curve 2 that obtained about 11 min later. Repetitive scans of 47 s. Fig. 4. Comparison of the spectrum of the u.v.-absorbing intermediate formed upon electrochemical (x) and enzymatic (.) oxidation of uric acid. The enzymatic data points were obtained when uric acid (100 p.M) was oxidized in the presence of type VIII peroxidase (0.17 p.M) and H 2 0 2 (200 ]2M) in phosphate buffer pH 7.0, ionic strength 1.0 M 5 rain after initiation of the oxidation in a 1.0 cm cell. The electrochemical points were obtained when uric acid (5 raM) was oxidized at 0.9 V in a thin-layer cell (90 pM thick) at 0.9 V at a gold minigrid electrode at a time when/> 90% of the uric acid was oxidized. The absorbance values obtained in the latter experiment were all multiplied by a constant factor of 2.24 to obtain the best fit to the enzymatic data points.
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-0.9
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Fig. 5. (A) Variation of the absorbance with time observed for the intermediate formed upon oxidation of uric acid (200 pM) with type VIII peroxidase (0.34/.tM) and H 2 0 2 (400 pM) at pH 4 followed by adjustment of the pH to 7.0. (B) Kinetic plot of time versus log absorbanee at 302 nm.
mediate formed enzymatically is identical to that formed electrochemically (see Table 1). The fact that the observed rate constant is independent of pH, at least between pH 7 and 9.3 indicates that the reaction of the intermediate is probably not base catalyzed over this pH range. The data presented in Table 2 also support the conclusion that the reaction of the intermediate is independent o f the enzyme and uric acid concentration. The rate of oxidation of uric acid is, however, directly dependent on the enzyme concentration. Attempts to measure the reaction kinetics for the u.v.-absorbing intermediate below pH 7 were difficult because the molar absorptivity of the intermediate significantly decreases. In addition, thin-layer spectroelectrochemical studies which, because of the very thin reaction cell used, require rather high uric acid concentrations (/>2 mM) to produce measurable concentrations of the intermediate, become very difficult below pH 7 owing to the low solubility of uric acid. However, using the pH-step enzymatic method, the approximate first order rate constant for reaction of the intermediate at pH 6.0 was ca 0.006 s -1
89 TABLE 2 Observed first o r d e r rate c o n s t a n t s for r e a c t i o n of t h e u . v . - a b s o r b i n g i n t e r m e d i a t e f o r m e d o n o x i d a t i o n of uric acid in t h e p r e s e n c e o f t y p e VIII p e r o x i d a s e a n d H 2 0 2 at pH 4.0 f o l l o w e d b y pH a d j u s t m e n t Initial concentration of uric a c i d / p M
pH
HRP concentration/ pM
H202 concentration/ pM
kobs/s -1
20 100 200 100 100
7.0 7.0 7.0 7.0 7.0
0.17 0.17 0.17 0.34 0.034
200 200 200 200 200
0.0039 0.0039 0.0035 0.0037 0.0036
20 100 200
8.0 8.0 8.0
0.17 0.17 0.17
200 200 200
0.0036 0.0030 0.0032
20 100 200
9.27 9.27 9.27
0.17 0.17 0.17
200 0.0033 200 0.0032 200 0.0032 M e a n : 0 . 0 0 3 5 -+ 0 . 0 0 0 4
and at pH 4.8, 0.0095 s-1. These values are much less reliable than those obtained at pH 7--9.3 but they do suggest t h a t perhaps the reaction o f the inter. mediate might be somewhat acid catalyzed below pH 7. In order that the conditions of the enzymatic oxidation paralleled more closely the electrochemical oxidation of uric acid, it was decided to a t t e m p t to oxidize uric acid at pH 7 with type VIII peroxidase and, when a significant q u a n t i t y of the u.v.-absorbing intermediate had been formed, to terminate the oxidation reaction so that the decay of the intermediate could be observed w i t h o u t interference from the continuing oxidation of uric acid. It was found that the most successful method of doing this was to terminate the oxidation of uric acid by addition of a large excess of the enzyme catalase. The latter enzyme rapidly r e m o v e s H 2 0 2 from the reaction mixture. Using catalase to terminate the reaction of uric acid, H202 and peroxidase at pH 7.0 in this fashion (see Experimental for details) it was f o u n d t h a t the u.v.-absorbing intermediate gave an observed first order rate constant of 0.0035 s- i . If the oxidation o f uric acid with peroxidase and H202 was carried out at pH 4.0 and then, when the uric acid was virtually all oxidized, the pH adjusted to 7.0 with a buffer solution containing catalase, the observed first order rate constant for reaction of the u.v.-absorbing intermediate was 0.0034 s- i . CONCLUSIONS
The studies reported here indicate that when uric acid is oxidized either electrochemically or with type VIII peroxidase an intermediate is formed which exhibits a well
90 less of whether it is generated electrochemically or enzymatically. The evidence, therefore, strongly supports the view that the same u.v.-absorbing intermediate is generated b y both electrochemical and peroxidase catalyzed oxidation of uric acid. In our earlier report we demonstrated that under electrochemical conditions at gold or graphite electrodes uric acid is oxidized in a 2e reaction to a diimine species (II, eqn. 1) which is very short-lived (e.g., it has a half-life at pH 8.0 of a b o u t 21 ms). This diimine primary product is attacked b y water in a (pseudo) first order reaction to give the u.v.-absorbing intermediate described in this paper which has been proposed [1] to be an imine-alcohol species (III, eqn. 1). Clearly, since the imine-alcohol is formed under b o t h electrochemical and enzymatic (peroxidase) conditions it seems quite justifiable to conclude that this species is derived from the same precursor under both oxidative conditions, i.e., from the diimine. There is no simple way to observe the diimine precursor formed u p o n enzymatic oxidation of uric acid because it is t o o unstable. The imine-alcohol formed either electrochemically or enzymatically is hydrated to the diol IV (eqn. 1) which then breaks d o w n to the various observed products. We believe that this study establishes for the first time that investigations of the electrochemical behavior of naturally occurring biological molecules can provide unique insights into the possible biological redox reactions of such molecules, Indeed, in the case of the electrochemical and enzymatic (peroxidase) oxidation of uric acid there is every evidence that the reactions are identical. Further investigations are currently underway on a variety of other purine enzymatic and electrochemical oxidations to further substantiate the similarity b e t w e e n the latter redox mechanisms for these compounds. ACKNOWLEDGMENTS
The authors gratefully acknowledge the help and guidance provided b y Dr. Eddie C. Smith regarding the enzymatic oxidation of uric acid. This w o r k was supported b y the National Institutes of Health through Grant No. GM 21034-04.
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J . L . Owens, H.A. Marsh a n d G. D r y h u r s t , J. E l e c t r o a n a l . Chem., 91 ( 1 9 7 8 ) 2 3 1 . G. D r y h u r s t , J. E l e c t r o c h e m . Soc., 1 1 9 ( 1 9 7 2 ) 1 5 5 9 . G. D r y h u r s t , E l e c t r o c h e m i s t r y o f Biological Molecules, A c a d e m i c Press, New Y o r k , 1 9 7 7 , Ch. 3. G. D r y h u r s t , Topics in C u r r e n t C h e m i s t r y , 3 4 ( 1 9 7 2 ) 47. K.G. Paul a n d Y. Avi-Dor, A e t a C h e m . Scand., 8 ( 1 9 5 4 ) 6 3 7 . E.S. CaneUakis, A.L. Turtle a n d P.P. C o h e n , J. Biol. C h e m . , 2 1 3 ( 1 9 5 5 ) 3 9 7 . R.W. M u r r a y , W.R. H e i n e m a n a n d G.W. O ' D o m , Anal. Chem., 39 ( 1 9 6 7 ) 1 6 6 6 . W.R. H e i n e m a n , B.J. Norris a n d I.F. Goelz, Anal. C h e m . , 47 ( 1 9 7 5 ) 79. J.L. O w e n s a n d G. D r y h u r s t , J. E l e c t r o a n a l . C h e m . , 70 ( 1 9 7 6 ) 1 9 9 . B.J. Davis, A n n . N.Y. A c a d . Sci., 1 2 1 ( 1 9 6 4 ) 4 0 4 .