J 'ITL!
4608
t"
1990, 55, 4608-4614
Redox Potentials and Acid-Base Equilibria of NADH/NAD+ Analogues in Acetonitrile AgnBs .4nne' and Jacques Moiroux* Facult8 de Phcrmacie, Uninersit4 d e Piccirdis, 3 plnre Louis Deuailly, 80037 Amiens Cedex, France
Kece:ued January 2, 1990
Redox potentials Eo,of seven NADH/NAD+ analogues (acridine,phenanthridine, quinoline, and pyridine derivatives) have been determined in acetonitrile. The pK,'s of the protonated forms of three reduced species AiH have also been determined together with the formation constants pKB, and pKoHof pseudobases resulting from the respective additions of amines and hydroxide to the seven oxidized forms A,+. Structural assignments and quantitative conclusions are based on 'H NMR and UV-vis spectrometries. For derivatives of quinolinium or pyridinium ring systems, there exists a linear correlation between Eo,and pKioHwith a slope of -29.6 mV. Presently, there is considerable interest in acquiring quantitative information about NADH/NAD+ analogues. In particular, the choice of one compound as a reagent for a given nonenzymatic redox transformation,2 e.g., the reduction of ketone^,^^^^^^^^ can be made rationally if the potentials of the redox couples are known. Since the neutral reduced forms AH are generally not water soluble in contrast to the cationic oxidized forms A+, the reactions are usually carried out in dipolar organic solvents, most often a ~ e t o n i t r i l e . ~ ~ The - ~ , ~potentials -~ of the AH/ A+ redox couples are also pH dependent and reactions, such as the reductions of ketones by AH, are general-acid cata l y ~ e d . ~ V *However, * controls of the redox potentials of the AH/A+ couples by means of pH adjustments are possible only in limited pH ranges since the h+'sreact with bases while the AH'S react with acids, and both transformations annihilate the interesting redox abilities of the AH or A+ species.2cvd Fifteen independent equilibrium constants K,, for redox reactions of the type shown in eq 1, where the oxidants k
A,+
+ A,H $ A,H + A,+
(1)
A,(,,r,)+ are a variety of substituted pyridinium, quinolinium, acridinium, and phenanthridinium cations, have been reported in the literature5 for a 4:l mixture of 2propanol and water. As a result, reduction potentials have been estimated for the corresponding A,+ in aqueous solution by assuming that the K 's would be the same and accepting -361 mV as the stanaard reduction potential of the 3-carbamoyl-1-benzylpyridinium cation against the standard hydrogen electrode, i.e.- -503 mV vs SCE (1) UA CNRS 484. 'Synthise et electrochimie de compos& d'intket pharmacologique". Faculti de Pharmacie, Universite Paris V , 4 Avenue de l'observatoire. 75230 Paris Cedex 06, France. Present address: Laboratoire d'Electrochimie MolBculaire, UniversiG Paris VII, 2 P1. Jussieu. 75231 Paris Cedex 05, France. ( 2 ) Reviews: (a) Yasui, S.; Ohno, A. Bioorg. Chem. 1986, 14, 70. ib) Zehani, S.;Gelbard, G. N o w . J . Chim. 1986, IO, 511. (c) Sliwa, MI. Heterocycles 1986,24,181. (d) Sausins, A,: Dubmi, G. HeterocyclPs 1988,
27, 291. (3) (a) Ohno, A.; Nakai, J.; Nakamura, K.; Goto, T.; Oka, S. Bull. Chem. SOC.J p n . 1981,54, 3482. (b) Ohno, A.; Nakai, J.; Nakamura, K.; Goto, T.; Oka, S. Bull. Chem. SOC.Jpn. 1981,54, 3486. (c) Tanner, D. D.; Stein, A. R. J . Org. Chem. 1988,53,1642. (d) Tanner, D. D.; Kharrat, A. J . Org. Chem. 1988,53, 1646. (e) Shinkai, S.; Hamada, H.; Kusano, Y.; Manabe, 0. J . Chem. SOC.,Perkin Trans. 2 1979,699. (fiFukuzumi, S.; Mochizuki, S.; Tanaka, T. J . Am. Chem. SOC.1989, I l l , 1497. (g) Mochizuki, S.; Fukuzumi, S.; Tanaka, T. Bull. Chem. Sot. Jpn. 1989,62, 3049. (h) Fukuzumi, S.; Kitano, T.; Ishikawa, K.; Tanaka, T. Chem. L e f t . 1989, 1599. (4) Zehani, S.; Lin, J.; Gelbard, G. Tetrahedron 1989, 45,733. (5) Ostovic, D.; Lee, I. S. H.; Roberts, R. M. G.: Kreevoy, M. M.J . Org. Chem. 1985,50, 4206.
These reduction potentials span 430 mV. In the present report, we present six independent K,,'s measured in pure acetonitrile and the standard potentials of seven AH,/A,+ redox couples (see Figure l),spanning 280 mV, deduced from the values of the six K,,'s and the standard potential of the 10-methylacridan/ 10-methylacridinium (AIH/A1+) redox couple previously determined in the same ~ o l v e n t . ~ We also present the pK,'s of the AH2+/AH acid-base couples when the reversibility of the protonation of AH has been ascertained and the equilibrium constants for pseudobase formation when the A,+'s react with various types of bases yielding A,OH species or other kinds of adducts depending on the nucleophilicity of the base. For N-benzylquinolinium cations, the analyses of the 'H NMR spectra of the two isomeric forms of the resulting adduct have been completed.
Results and Discussion Redox Potentials. Values found for the standard potentials E", of the redox couples A,H/A,+ (eq 2) are gathered in Table I. These data were obtained according to A,+ + 2e + H+-" A,H (21 or I--
-
procedures described in the Experimental Section. The dependence upon pH is also mentioned. Charge-transfer-type complex formation between A,+ and A,H could cause some error in the reported second-order rate constants k , and k,, if the equilibrium constant K m favoring complex formation is large, irrespective of whetger or not the complex lies on the reaction coordinate8 In agreement with the reported small values of KCT,in neat solvent^,^ it appears that the complex formation is inefficient in the concentration range we used since halving the concentrations did not make significant changes in the k,,and k,,. Qualitatively, it is worth noticing that the order of decrease in oxidizing power established previously in a mixed (6) SCE aqueous KC1 saturated calomel electrode to which all the potentials given in the present paper are referred. (7) Hapiot, P.; Moiroux, J.; Saveant, J. M. J . Am. Chem. Soc. 1990. 112, 133'7. (8) i a ) Rappoport, 2.;
Horowitz, A. J . Chem. Soc. 1964, 1348. (b) Powell, M. F.; Bruice. T. C. J. Am. Chem. SOC.1983, 105,7139. (9) (a) Cilento, G.; Tedeshi, P. J. Riol. Chem. 1961, 236, 907. (b) Cilento, G.; Schreiber, S. Arch. Biochem. Biophys. 1964, 107, 102. (c) Shinkai, S.; Tamaki, K.; Kunitake, T. Bull. Chem. SOC. Jpn. 1975, 48, 1918. (d) Murakami, Y.; Aoyama, Y.; Kikuchi, J.; Nishida, K.; Nakano, A. J . Am. Chem. SOC. 1982, 104,2937. (e) Martens, F. M.; Verhoeven, J. W. J . Phys. Chem. 1981,85, 1773. (0 Itoh, Y.; Abe, K.; Senoh, S. J . Polym. Sci. Polym. Chem. 1987,25, 2871. ( 9 ) Murakami, Y.; Kikuchi, J.; Tenna, H. Chem. Lett. 1985,103. (h) Kalyanasundaran, K.; Colassis, T.; Humphry-Baker, R.; Savarino, P.; Barni, E.; Pelizzetti, E.; Gratzel, M. J A m . Chem. Soc. 1989, 111, 3300.
0022-3263/90/ 1955-4608$02.50/0 0 1990 American Chemical Society
J. Org. Chem., Vol. 55, No. 15, 1990 4609
NADH/NAD+ Analogues in Acetonitrile
Table I. Eauilibrium Constants K;;and Standard Potentials Eaiof the A;H/A;+Redox CouJes in Acetonitrile at 25 "C measured" oxidant reductant measured" k;;lM-' min-l k;;lM-' min-' E0;'ImV vs SCE ., . . (93 f 1) x 10-2 End = (257 - 29.6 PH) f 26 A4+ AlH (52 i 2) x 10-3 K,, = 18 f 1 (22) Eo1= (220 - 29.6 pH) f 25d 34.1 f 0.5 (94 f 3) x 10-4 K13 = 3630 f 170 (490) E"3 = (115 - 29.6 PH) f 25 (54 f 1) x 10-2 (30 f 1) x 10-3 E"7 = (79 - 29.6 PH) f 26 K37 = 18 f 1 (120) (44 f 1) x 10-3 (40 f 1) x 10-3 KTz = 1.10 f 0.05 (1.2) E O 2 = (77 - 29.6 PH) f 26' 2.4 & 0.1 (115 f 5) x 10-4 K25 = 210 f 20 (110) E"5 = (8 - 29.6 PH) f 28 1.10 f 0.02 (110f 5) x 10-3 K,, = 10.0 f 0.5 (540) E O 6 = (-22 - 29.6 pH) f 29 Y
,*e
I
. I
nMean and standard deviation (three separate experiments). *Values deduced from ref 4, which were determined in a mixture of 2propanol and water in the ratio 4:l by volume. = E n j + 29.61 log Kip dFrom ref 7. "educed from the determination of K32 = 19.5 f 1.
Table 11. pK.)s of the Conjugated Acids AiH2+of the Reduced Forms AH in Acetonitrile at 25 "C Ai H A,H AsH pK, 3.3 f 0.1 8.6 f 0.1 8.2 f 0.1
Eoi listed in ref 5 are implicitely given a t pH 7 in water and no straightforward relationship can be established with a corresponding pH value in acetonitrile. Protonation of the Reduced Form AH. The thermodynamic pK,'s of three AiH2+species determined in the present work are given in Table 11. AiH2+ + AiH
+ H+
(3)
Ka = (AiH)(H+)/ (AiHZ+)
I
8r
&
,
A,'
, x
A :
,
X : COCH,
I
04-1
81
A5H
:CONH,
X:CN
A7H
Figure 1. Skeletons of the AiH/Ai+ redox couples. Position numbers are mentioned on the rings of the reduced AiH. Bz = CHPCBHS(benzyle).
solvent5is respected. Quantitatively, the few discrepancies can reasonably be ascribed to changes in the solvatation energies of Ai+ and/or AiH, the greatest amounting to 1 order of magnitude in the value of K3,, i.e., a ca. 30 mV difference on the potential scale. The potential span Eo4/Eo6 observed in acetonitrile is ca. 40 mV smaller than was estimated in water.5 The comparisons of the individual values of Eoi given in Table I with those estimated in ref 5 cannot be meaningful for the following reasons: (i) All the equilibrium constants mentioned in the present paper, including the one leading to the evaluation of Eo1(see ref 7), were determined directly in acetonitrile, whereas the estimations reported in ref 5 result from the combination of equilibrium constants determined directly in 2-propanollwater (4/1 by volume) with a reference potential estimated in water according to the cyanide affinity method.1° (ii) The
To our knowledge, only a tentative evaluation of the pK, of A1H2+in acetonitrile has been reported in the literature." However there was an important flaw in that case since the authors, starting with an unbuffered solution of AIH, mistook the concentration of added HCIOl (not in excess) for the proton concentration (H+)a t equilibrium. For these rather strong acids, the pK, values are water concentration independent, up to 100 mM. Reactions of A+ with Amines and Hydroxide: Spectrophotometric and 'H NMR Identifications of the Adducts at Equilibrium. For all the cations A+, pH-dependent reversible UV-vis spectral changes are observable in basic acetonitrile, suggesting that pseudobase formation occurs (see Experimental Section). A closer investigation of the equilibrated solutions spectra over a wide pH range reveals that primary and secondary amine bases B,H can successfully compete with hydroxide, as nucleophiles, to give amino-pseudobases according to equation 4. Ai+
+ B,H
AiB,H+ + AiB,
+ H+
(4)
UV-vis spectral characteristics of hydroxy and some amino adducts have been determined for cations Ai+, when accessible, and are given in Table 111. Corresponding isomeric structures were, in each case, assigned by reference to spectral data of closely related adducts resulting from nucleophilic additions a t the same sites12-14 and
(lo) (a) Wallenfels, K.; Diekmann, H. Justus Liebigs Ann. Chem. 1959,621, 166. (b) Taylor, K. E.; Jones, J. B. J. Am. Chem. SOC.1976, 98,5689. (c) Kellogg, R. M.; Piepen, 0. J . Chem. Soc., Chem. Commun. 1982,402. (11) Fukuzumi, S.; Ishikawa, M.; Tanaka, T. Tetrahedron 1986,42, 1021. (12) Bunting, J. W.; Meathrel, W. G. Can. J. Chem. 1974, 52, 981.
4610 J . Org. Chem., Vol. 55, No. 15, 1990
Anne and Moiroux
Table 111. Spectrophotometric Characteristics Used for the Determination of the KO, a n d Kg, Equilibrium Constantsa _____ AOH adducts AB, adducts , ,x A, and (r$s)* (e,)c x, (cw's) ,A, ~_ (%a) ._-_.____ 280 nm 356 nm 284 nm (15000) ( t A = 16000) (
