Elucidation of the Biological Redox Chemistry of Purines Using Electrochemical Techniques Glenn Dryhurstl and N. T. Nguyen University of Oklahoma, Norman, OK 73019 Monika Z. Wrona University of Warsaw, Warsaw, Poland R. N. Goyal Roorkee University, Roorkee, India
Modern electrochemical and related analytical methods provide powerful techniques to investigate the complex redox chemistry of naturally occurring organic compounds. Electrochemical studies cannot only provide mechanistic information of fundamental interest hut also can give valuable insights into the chemical aspects of enzymatic and in vivo redox reactions of such compounds (1,2). This may he illustrated by recent work concerned with the electrochemical oxidation of the purine, uric acid. This discussion will he limited to information obtained at pH 2 7. Detailed information over a wider pH range has been presented elsewhere (3). Electrochemical Studies Uric acid (I, Fig. 7) gives a single voltammetric oxidation peak (I,) at a pyrolytic graphite electrode. At a sweep rate of 5 mV s-' the peak potential, E,, shifts toward more negative potentials according to eqn. (1)(4).
Cyclic voltammetry (CV) of uric acid (Fig. 1) shows that having scanned oxidation peak I,, two reduction peaks (I, and 11,) appear on the reverse sweep (5-12). The peak potentials for peaks I, and I, are separated by more than 56 mV so the species responsible for these peaks are said to form a quasireversible couple (13). With increasing sweep rate, the height of peak I, grows relative to peak I, and, correspondingly, the height of peak 11, decreases. Controlled potential coulometry of uric acid at potentials slightly positive of peak I, shows that for each molecule oxidized two electrons are transferred. The quasi-reversihle nature of the peaks I, and I,, the dE,ld(pH) slope of -55 mV per pH unit, and the coulometric n-value of 2 all support the conclusion that the reaction for peak I, is a 2e-2HC process. The primary product of this process is responsihle for reduction peak I, in CV. However, in view of the fact that peak I, is smaller than peak I,, except at very fast sweep rates, indicates that this product must he very unstable, i.e., it undergoes a rapid chemical follow-up reaction. The nature and kinetics of this follow-up reaction may he elucidated using double potential-step chronoamperometry. This technique may he understood by reference to Figure 1.Thus, an initial potential, El, of 0.0 V is applied to a stationary micrographite working electrode dipping into the quiet (i.e., unstirred) uric acid solution. At E,,of course, no faradaic process (i.e., electrooxidation) occurs and only a small capacitive current flows for a very short time to charge the electrode. After a short equilibration time (e.g., 5 s), the potential is pulsed to 0.6V, E2. The value E 2 is >ZOO mV positive of E, for peak I, so that uric acid is oxidized at a rate controlled only by its rate of mass transport to the electrode surface, i.e., the oxidation current is dependent upon its rate of diffusion. After the potential is held at Ez for a predetermined time (TF, sec) it is pulsed hack to El (0.0 V). At E 2 the
Anna Brajter-Toth University of Maine, Orono, ME James L. Owens Central Soya Corp., Fort Wayne, IN Henry A. Marsh, Jr. Texaco, Inc., Houston, TX
primary product responsihle for peak I, is generated some of which undergoes a chemical reaction. When the potential is returned to E l (which is >200 mV negative of peak I,), a reduction current flows which is dependent on the amount of unreacted primary product remaining in the vicinity of the electrode surface. A typical plot of current versus time response in this type of experiment is shown in Figure 2. The current is measured a t times TF (IF) and ~ T F ( I B )If. the product of the initial electrode reaction at E2 is stable, then the ratio 1 ~ 1has 1 ~a numerical value of 0.2928 (14). However, if the product reacts to form a species which cannot he reduced at E l , the ratio IBIIFwill be perturbed. The order and rate of the perturhrng reaction may he found by monitoring IR/IF with respect to TF and the hulk concentration of uric acid.
Figure 1. Cyclic voltammogram at the PGE of 2 mMuric acid in phosphate buffer pH 7.5.Sweep rate: 500 mV sc'.
Volume 60
Number 4
Awril 1983
315
Time, Sec
Figure 3. R, versus T ~ l t , for , ~ uric acid peak l./peak I, processes. The experiwas 12.5 ms. The first order solution rate constant is mental value of t, 0.406/t,,, = 32.5 s-' (14).
Figure 2.Current versus time response observed during adoubie potential step chronoamperometric experiment.
Usually, the IBIIFratio is normalized by dividing by 0.2928 to give the dimensionless parameter RI. When there is no perturbing reaction, RI = 1.00. When a kinetic perturbation does occur, RI is less than 1.00. The Rr versus time behavior expected for various chemical follow-up mechanisms has been computed (14). The points in Figure 3 show values ofR1 versus T d t - m (where t ~ i is " the real time in seconds at which RT= versus T ~ / t ~expected ,2 for a first order or pseudo-first order EC reaction, i.e., an electrochemical reaction (in this case the 2e-2Hf electrooxidation of uric acid to produce peak I,) followed by a (pseudo) first order chemical reaction. The value of the apparent first order solution rate constant may he computed from experimental values of RI versus T Fas shown in Figure 3. At pH 8, this rate constant has a value of 32.5 s-', which corresponds to a half-life of the primary product of peak I, (responsible for peak I,) which is about 21 ms. I t is helieved that these results indicate that at pH 7-8 the anion of uric acid (I, Fig. 7 ) is oxidized in the 2e-2H+ reaction for peak I, giving the quinonoid anion IS (Fig. 71, and that peak I, corresoonds to the reverse reaction. Peak I,- is -zenerallv smaller than peak I, because the quinonoid I1 reacts rapidly with water. in a pseudo first order reaction, and disappears. The immediate product of this hydration reaction is believed to be the tertiary alcohol 111(Fig. 7 ) . A very powerful technique to search for and to characterize intermediates generated in electrode reactions is thin-layer spectroelectrochemistry (15, 16). A useful cell for studying electrooxidation reactions is shown in Figure 4. A thin slice (- 0.2 mm thick) of optically-transparent reticulated vitreous carbon (RVC) serves as the working electrode. This is sandwiched between two optical quality quartz microscope slides. Typical spectra recorded during and after electrooxidation 316
Journal of Chemical Education
of uric acid in such a cell are shown in Figure 5. Curve 1in = 295, Figure 5A is the UV spectrum of uric acid at pH 8 (A, 240 nm). Upon initiation of the electrooxidation at potentials corresponding to peak I,, the UV hands of uric acid decrease and, correspondingly, two new hands grow in at 315 nm and 225 nm. The latter bands reach their maximal absorbance values at about the time that virtually all uric acid has been oxidized (curve 6, Fig. 5). If, at this point, the electrolysis is
terminated by open-circuiting the RVC electrode, the spectral changes shown in Figure 5B take place. Thus, hoth of the new UV hands decrease with time and ultimately disappear. Clearly, therefore, electrooxidation of uric acid leads to formation of one or more UV absorbing intermediates. By monitoring the decay of absorbance versus time ( A versus t ) , it may he shown that two kinetically distinct intermediates are present hoth of which disappear by first order reactions. At pH 7.5 in a phosphate buffer having an ionic strength of 0.5 M and measurine kinetics at 315nm. the first rate constant is 0.046 s-' and th;second is 0.0021 s-' ( 3 ) ;these rate constants corresvond to half-lives of 15 s and 330 s. respectivelv. These half-li;es are clearly f q greater than that measured f i r the quinonoid primary product (11, Fig. 7, 21 ms) and so the UV-absorbing intermediates cannot he the latter species. Insights into the identity of the UV-absorbing intermediates may be obtained by combining thin-layer spectroelec-
diate species are present, i.e., equivalent to curve 6 in Figure 5A. At this point, the solution within the thin-layer electrode cavity (-200 ILL)is rapidly removed and frozen at -78°C (dry ice-acetone). This very low temperature serves to quench the decomnosition of the intermediates. The frozen sample is
. .
verts >N-H and -0-H groups to >N-Si(CH& and -0-Si(CH& groups, respectively. Such derivatives are usually hoth volatile and thermally stable. Thus, following derivatization, the reaction mixture may he separated
CH>---C\
/0-S'(CH3)3 N-L(CHa)3
and analyzed by GC-MS (10).GC shows one major peak. The component eluted under this peak has been analyzed by hoth electron impact (El) and chemical ionization (CI) mass spectrometry and it is found to have a molecular weight of 472. By employing a different silylating reagent (e.g., fully deuterated BSA of N-methyl-N(t-hutyldimethvlsilyl)trifluo. . roacetamide) a silyl derivative havinga greater mass than the
trimethylsilyl derivative can be formed. From the increase in mass it is easy to calculate (17) that the derivatized intermediate has four silylatable sites. Since each Si(CH& moup .adds 72 amu to the mass of the compound derivatized, the molecular weight of the UV-absorbing intermediate is 472 - (4 X 72) = 184. Thus, the intermediate has a molecular weight of 184 and possesses four silylatable sites. This information fits the tertiary alcohol 111 (Fig. 7) and the hicyclic carhoxylic acid IV (Fig. 7), hoth of which as anions have extended chromovhores and would he expected to absorb in the UV reeion at longer wavelength than the anion of uric acid, i.e., conform to the spectral behavior shown in Firure 5A. That one of the ~ v - a b s x h i nintermediates ~ is in fait the tertiary alcohol I11 and that this species is at least partially responsible for reduction peak 11, observed in CV of uric acid (Fig. 1) may he demonstrated by additional thin-layer spectroelectrochemical experiments (12).The growth of the UV-absorbing intermediate species is shown in Figure 5A when uricacid is oxidized a t peak I, potent,ials. If, instead of open-circuiting the RVC electrode with the resulting spectral decay of the intermediate
a
.
~-
of t h i intermediatespecies decreases and, cor*espondingly, the spectrum of uric acid reappears. In other words, electrochemical reduction of an UV-absorbing intermediate does occur at potentials corresponding to peak II,, and the ultimate product of that reduction is uric acid. It may he concluded on the basis of such information that the very reactive quinonoid I1 (Fig. 7) formed as the primary product of the 2e-2H+ reaction corresponding to peak I, is rapidly attacked by water yielding the tertiary alcohol 111(Fig. 7). This can be electrochemically reduced in peak 11, probably giving a dihydro species VII (Fig. 7). The expected facile dehydration of this compound (18) leads to the observed (Fig. 5C) regeneration of uric acid. To account for the two-step kinetics noted for decay of the UV-absorbing intermediate, i t is believed that a rine contraction reaction of I11 occurs giving IV (Fie. 7), (3). , The ratter compound has exactly the same molecular weight (184) and number of silylatable positions, four, as 111. In addition, IV possesses the same basic chromophore as I11 and, hence, both com~oundswould be exvected to have similar UV spectra. The end-product of electrochemical oxidation of uric acid in a thin-layer cell may be identified by allowing the inter-
-
u
Wavelengthlnm Figure 5 . (A) Spectra of a 1 mMsolution of uric acid undergoing electrooxidation at a RVC electrode in a thin-layer cell in phosphate buffer pH 8.Applied potential: o.9V (peak I,). Curve 1 is the initial spectrum of uric acid; curves 2-6 each hadaduration of 19 s with essenlialiy no time interval between spectral sweeps. (6)After scannlng curve 6 in (A] the RVC eletilrode was open circuited and curves 7 3 5 were recorded. Each curve had a duration of 38 s with essentially no time interval between sweeps. (C) After scanning curve 6 in (A) the RVC electrode was potentiostated at -1.2 V. Each curve had a duration of 19 s with essentially no time interval between sweeps.
Volume 60
Number 4
April 1983
317
mediates to decay as shown in Figure 5B. Then, the resulting nroduct can he freeze-dried. silvlated. and analvzed bv kc-MS (10). The GC-MS behavior is identical to aithentic allantoin (VI, Fig. 7). Large scale electrolysisof 5-10 mg of uric acid dissolved in 25-30 mL of phosphate buffer can be com. . pleted in ahout 1 hour. The resulting product solution may be freeze-dried, redissolved in 1-2 mL of water, passed through a column of gel permeation resin (Sephadex G-10) and eluted with water. If the eluent is monitored at about 200 nm. two well-seoarated ~ e a k are s observed. the first correp m d i n y t u int.rg.;mic phwphate, the secund t o nllantoin 1191. If a irtictiun collectur is uied. the h t m t wnt:~ininrthe secrmd component may be combined, freeze-dried, and the allantoin convenientlv identified by its melting. point, IR and mass .. spectra.l Formation of allantoin VI (Fig. 7) as the end-product of electrochemical oxidation of uric acid is easily rationalized by a relatively slow hydration of carboxylic acid IV (Fig. 7) giving V followed by ring opening and decarhoxylation ( 3 ) . Enzymatic Oxidation of Uric Acid The enzymatic oxidation of uric acid may be conveniently studied usine tvne VIII neroxidase from horseradish ~ e r o x i dase ( 7 , l l ).?&&a1 c b k g e s observed at pH 8 are s & m in Figure 6. Upon initiation of the enzymatic oxidation (Fig. 6A) the UV bands of nric acid (A,,,, = 295,240 nm) decrease and, = 315 nm and 225 simultaneously, new hands grow in at A, nm. If the oxidation is terminated after scanning curve 4 in Figure 6A by addition of the enzyme catalase (which very rapidly destroys HzOz), the spectral changes shown in Figure 6B occur. Thus, both UV hands decrease. Accordindv, it may be concluded that the peroxidase-catalyzed oxidaGon of nric acid bv. HgO7 generates intermediate species having the same -. UV spectrum as the intermediate species generated by electrooxidation of uric acid at an RVC electrode in a thin-layer cell (compare Figs. 5A, B). The kinetics of decay of the enzymatically-generated, UV-absorbing intermediate species agree within experimental error with the results observed for the
electrochemically generated intermediate species. That the intermediate species formed in the two different types of oxidation reaction are the same can be proved by CV and GC-MS. Voltammogram A in Figure 8 shows that having scanned oxidation peak I, of uric acid peak I,, due to reduction of quinonoid I1 (Fig. I ) ,and peak 11,, due to reduction of tertiary alcohol 111(Fig. 7 ) and perhaps carboxylic acid IV (Fig. 71, are observed on the reverse sweep. The voltammogram shown in Figure 8B is for a mixture of peroxidase and H202, where two small voltammetric reduction ~ e a k of s unknown origin are present. Figure 8C is a voltammogram taken about 60 s after initiation of the enzymatic oxidation of uric acid. The first sweep toward negative potentials shows peak 11,, i.e., the intermediate species responsible for peak 11, is formed in the peroxidase-cadyzed rea>tiou. whenthe enzymatic oxidation is complete, CV reveals the peak 11, has disappeared. The UV-absorbing intermediate(s) formed in the enzymatic reaction may be trapped by allowing the reaction to proceed
Peak .I
0
'
This entire procedure is used as an experiment in an undergraduate instrumental analysis laboratory at the University of Oklahoma. The experiment illustrates voltammetry, controlled potential electrolysis, coulometry, freeze-drying techniques, open-column liquid chromatography, mass spectrometry, and IR spectrophotometry.
COOH
Figure 8.(A) Spectra of 0.2 mMuricacid undergoing oxidation in the presence of type Vlll peroxidase (0.4 fit4 and H202 (200 fit4 in phosphate buffer pH 8. Curve 1 is the initial spectrum of uric acid. Subsequent spectral traces had a duration of 75 s with essentially no time interval between sweeps. (8)After scanning curve 4 in (A) the enzymatic Oxidation was terminated by addition of catalase. Spectral traces 5-34 had a duration of 75 s with essentially no time interval between sweeps. 318
Journal of Chemical Education
Figure 7. Reaction mechanism proposed for the electrochemical oxidation of uric acid in phosphate buffers at pH >7.
until the maximal amount of intermediate is present (i.e., curve 4 in Fig. 6A) followed by rapid freezing, freeze-drying, silvlation and GC-MS. The results of such analysis (11) indicate that exactly the same intermediate is formed in the enzvme-catalyzed reaction as in the electrochemical oxidation.
Conclusions
I t is quite evident a t pH 7-8 that the oxidation of uric acid by H202catalyzed by peroxidase and the electrochemical oxidation yield intermediates which q e spectrally, kinetically, and analytically identical. These intermediates then decay to the same end-product. Electrochemical studies reveal that uric acid is oxidized in a quasi-reversible 2e-2H+ reaction to a short-lived quinonoid (11,Fig. 7). In view of the factthat the intermediates I11 and probably IV (Fig. 7) are formed in both the electrochemical and enzvmatic oxidations. and the same
.
-
lyzed reaction. I t is important to note that electrochemical investigations do not provide direct information concerning the way in which peroxidase accomp&shesthe oxidation of uric acid. However, they do provide insights into the initial, unstable intermediates of the reaction and a reasonable reaction pathway. Acknowledgment
Most of this work was supported by the National Institutes of Health through Grant No. GM-21034. Literature Cited (1) Dryhursl, C,., '~Electroch~mistr~ oEBiologics1Mnlecu!es: Academic Press. New York, 7 07" A".,.
(2) Dryhurst, C., Kadish, K. M.. Scheller. F., and Kenneberg. R.. "Biologicai Elmctrochemistry: Vol. I., Academic Press. New Yurk. 1982. 131 Goyrl, R. N., Brajter-Tuth. A,. and Dryhurst. G.. J Elecfraanal. Chem., 131, 181
,."".,.
,,qx9/
1" Brajalter-Toth, A.. Goys1.R. N., Wrana, R. N..I.acaua,T.. Nguyen,N. T., andDryhurst, Cr.,Bin~lectrochem. Rioene~~.,H,418 (19811. (5) Diyhurst,G.,J. Electrochrm. Sac.. 119,1559 (1977.1. (6) Owens, J. L., Marsh, H. A.,and Dryhurst, G., J Electiaanal. Chem.91.231 11978). (7) Ma1sh.H. A,. and Dryhurst, G , J Elerlroanal Chem., 95.81 (19791. (81 Wrona, M. Z., andDrvhurst, G.,Bmchim.Biuphy;. A d o . 570,371 (19791. (9) W i p a , M 2.. Owens, J. L.. and Dwhurst, G., J. Elrrfraanoi Chem.. 105, 295 (1979). G., J Elactroonol. Chrin , 122,205 (19811. (10) Brajter~T~th,A..andDryhurst, 111) Goyal, R. N., Bmjier~Tnth,A , an6 Dryhura. G., Riaelectrochem. Riosnarg , 9, 39 (19821. Riaenerg., 9, 273 (12) Guyal, R. N , Nguyen. N. T., and Dryhurst, G., Bianl~ciroch~m. (1982). Ymk, (13) Adams, R. N., "Electmchemirlry st Solid Electrodes," Marcel Dekker. NPW 1969. ,A. ,118. (la) Child%,W.V.,Maloy,J . T . . ~ e ~ z t h e l y i , ~ . P . . a n d B a r dJ.,J.Elrrirorhem.Soc e"" ",-
Figure 8.Cyclic voltammograms at the PGE in phosphate buffer pH 8 of (A) 0.4 ~ M t y p Vlll e peroxidase, 600 &MH202. initial sweep toward negative potentials: (6) 0.5 mMuric acid, initial sweep toward positive potentials. and (C)0.36 m M Vlll peroxidase. 6 0 0 p M H 2 0 2 60 safter initiation of the uric acid. 0.4 ~ M t y p e enzymatic oxidation. initial sweep toward negative potentials. Sweep rate: 200 mV 5 - ' .
,.",.,. ,,0",3
(161 Murray, R. W.. Heineman, W. R., and O'Dum. G. W.. A n d Chem.. 39,1666 (19671. (16) Heineman, W. K..Anul Chpm, 10,290A (1978). 1171 McCloakev. L A,. in "Basic Princiolenin Nudeis Acid Chemi8try:P. 0.P. T'so (Ediiorl. v& I, AcademicP~ess.N& York. 1 9 7 4 , ~209. . (18, Nielren,A.T., and Hnulihan, W. J , i " "o,gsnicReactlons,"Vol.1.Adams.R..Blstt, A H Hnpkplhridp~,V C , a i r n x T . L..Cram.D. J.andHouse.H.0.. 1EditorsI.John Wileyand Sans, New York. I S G R , ~ ~1429. . (19) 0wens.J. i..,Thomnr,H. H.. andD~yhurst.G.,Anol Chim Acto.96,89 (1978) ~~
User Support Group Announced for NIH/EPA - CIS The NIHJEPA Chemical Information System has announced a new user support group a t Computer Science Corporation in Falls Church, Virginia. The members of the support group have amixture of skills in chemistry, biology, toxicology, and information science that will enable them to assist users of the database effectively. CIS is a eomputer~accessibledatabase which contains a collection of scientific and regulatory information as well as data analysis computer programs in the areas of toxicology, environment, speetroscapy, chemical and physical properties, and nucleotide sequencing. It allows for structure, substructure, and full or partial name searching on over 225,000 unique chemical substances. Further information on the database and the subscription and user costs can be obtained from: CIS User Support Group, Computer Sciences corpora^ tionIATD, 6565 Arlington Rlvd., Falls Church, VA 22046.
Volume GO
Number 4
April 1983
319
