748
J. Phys. Chem. 1990, 94, 748-752
Oxidation Potential of Luminol. Is the Autoxidation of Singlet Organic Molecules an Outer-Sphere Electron Transfer? G. Merinyi,* J . Lind, X. Shen, and T. E. Eriksen Departments of Physical and Nuclear Chemistry, The RojJal Institute of Technology, S - 10044, Stockholm, Sweden (Receiued: June 7, 1989)
The one-electron reduction potential of the LH'/LH- couple was determined by pulse radiolysis to be 0.87 & 0.02 V vs NHE. The rate constant for the electron-transfer reaction between LH' and 02'-(k+,)was measured to be 3 X lo9 M-' s-l while (kT7)was estimated to be less than lo* M-I s-'. In combination with previously the corresponding rate between Lo- and 02'measured data the above values enabled us to quantify the kinetics and thermodynamics of luminol autoxidation. The data were also analyzed in terms of the Marcus-Hush theory. For comparison, a similar analysis was made on the dihydroflavin mononucleotide system (FMNH-). It was concluded that the autoxidation of both LH- and FMNH- and probably of most In contrast, the autoxidation singlet organic anions is initiated by an outer-sphere electron transfer between the anion and 302. of L'-. FMN'-, and other semiquinone radicals is suggested to involve inner-sphere electron transfer.
Introduction Bond formation between singlet ground-state organic molecules and triplet molecular oxygen is spin-forbidden. Autoxidation of such molecules has therefore been assumed to initiate from a one-electron transfer between the molecule and 02.While theoretically compelling this assumption has rarely been substantiated. Among the few studies addressing this issue we find the early works by Russell et al.' where the autoxidation of certain organic anions in dimethyl sulfoxide (DMSO) was shown to result in radical formation. I n aqueous solutions Bruice et a1.2 have provided evidence for electron transfer between 1,5-dihydroflavins and molecular oxygen. I n the literature on the autoxidation of dihydroflavins there are sufficient thermodynamic and kinetic data2-6 extant to warrant quantitative analysis in terms of electron-transfer theory. A similar analysis of luminol (LH-) requires knowledge of the redox potential of the LH'/LH- couple which will be determined in this work. Additional information needed is the rate of luminol autoxidation. While in DMSO oxygen reacts with the luminol dianion7 (L2-) with a rate of ca. SO M-' s-I autoxidation rates in water are too low to be measurable at any pH. For example, even at pH 14, air-saturated luminol solutions are indefinitely stable in the dark. In the present work we shall measure or estimate the rate of the assumed reverse reaction, Le., electron transfer between the luminol semiquinone radical (LH', Lo-) and 02"-. Previous analysis* of data on electron-transfer reactions involving redox couples (mostly inorganic metal complexes) and 0,/02'have resulted in an enormous spread of the apparent couple. Recently, we self-exchange rate ( k e J of the 02/02'determined directly9 k e x ( 0 2 / 0 2 *to - ) be 450 f 160 M-' SKI. In the present work the aqueous autoxidation will be analyzed of both LH- and FMVH- (1.5-dihydroflavin mononucleotide) as well as ( I ) (a) Russell, G. A.; Jansen, E. G.; Becher, H. D.; Smentowsky, F. J. J . Am. Chem. Sot. 1962, 84, 2652. (b) Russell, G. A,; Moye, A. J.; Nagpal, K . J . Am. Chem. Sot. 1962, 84. 4154. ( 2 ) Eberlein. G.;Bruice, T. C. J . Am. Chem. Sor. 1983. 104, 6685. (3) Favaudon, V . E u r . J . Biochem. 1977, 78, 293. (4) Draper, R. D.; Ingraham. L. L. Arch. Biochem. Biophys. 1968. 125,
802.
( 5 ) Anderson, R. F. Biochim. Biophys. Acta 1983. 722, 158. (6) Lind, J.; Merinyi, G. Photochem. Phofobio!., in press.
(7) (a) White, E. H.; Zafiriou, D.; Kagi, H. Hh.; Hill, J . H . M. J . Am. Chem. SOC.1963,86, 940. (b) Gorsuch, J. D.; Hercules, D. M . Photochem. Photobiol. 1972, I S , 567. (c) Seliger, H. H. In Liquid Scintillation Counting, Volume 2; Peng. C . T., Horrocks, D. L.. Alpen. E. L., Eds.: Academic Press: New York, 1980; p 296. (8) McDowell, M. S.: Espenson. J . H.; Bakac, A . Inorg. Chem. 1984. 23? 2232. in
( 9 ) Lind. J . ; Shen, X.; MerOnSi, G . ; Jonsson. B-0. J . Am. Chem. Sor.. press
0022-3654/90/2094-0748$02.S0/0
l.H
I
0 LO2
u7
P" 2
of the corresponding semiquinones L'- and FMN'-.
Experimental Section Sodium chlorite (Alpha) and luminol (Ega Chemie) were recrystallized prior to use. Phenol, 4-iodophenol (both Aldrich), KSCN, NaN,, and the various buffers (all Merck) were employed as received. Water was triple distilled in quartz. Before irradiation the solutions were purged with N 2 0 gas (Aga). Pulse radiolysis was carried out at room temperature with a 7-MeV microtron accelerator. Details of the setup'O and the computerized optical detection system1' have been described elsewhere. The length of the applied pulses was between IO-' and 2 X lO-'s corresponding to doses of 1-2 Gy. The concentration of radicals generated in such pulses was in the vicinity of mol/dm3. Dosimetry was performed by means of aerated aqueous solutions containing 10 m M KSCN. A Gt value of 2.2 X mZ/J for the (SCN),'radical at 500 nm was employed.12 Results The Redox Potential for the L H ' I L K Couple. Attempts have been made with electrochemical methods to determine the one electron oxidation potential of lumin01.l~ As the authors point ( I O ) Rosander, S . Thesis, The Royal Institute of Technology, Stockholm, TRITAEEP-74-16, 1914; p 28. ( 1 1 ) Eriksen, T. E.; Lind, J.; Reitberger, T. Chem. Scr. 1976, 10, 5 . ( 1 2 ) Fielden, E. M. In The Study of Fast Processes and Transient Species by Electron Pulse Radiolysis; Baxendale, J. H., Busi, F., Eds.; Reidel: Dordrecht, Holland, 1982: Nata Advanced Study Institutes Series; pp 49-62. ( 1 3) Epstein, B.; Kuwana, T. Photochem. Photobiol. 1965, 4, 1 157.
0 1990 American Chemical Society
Oxidation Potential of Luminol
The Journal of Physical Chemistry, Vol. 94, No. 2, 1990 149
TABLE I: Pertinent Acid-Base and Redox Equilibria proton equilibrium
+
LH2 + LHHt LHL2- H+ LH' = L'Ht L02H2 LO2H- + H t L 0 2 H - + L022- + H + H202 + HO2- Ht 022- H+ H02H0; + 02*- Ht
+ +
+ + +
PK.
ref
6.7 15.1 7.7 10.4 -16 11.7 15.9 4.69
14
V vs N H E 0.24 -0.16 0.89 0.936 0.80 f 0.02
redox couple L/ L'0 2 / 0 2 ' - ( 1 M 02) 02*-/H202 (PH 7)
EO,
c102*/c102PhO'/PhO-
15 16 17 18 19 20 ref 21 20 20 22 23
-
1 0
out, the measured potential (ca. 0.67 V at p H 13) is not related
to the thermodynamic potential due to irreversible electrode reactions. In the present work this problem is overcome by having the reactants equilibrate faster than irreversible processes have time to occur. Determination of Eo(LH' f LH-) against the CIO,f C10,Couple. Table I collects data pertinent to the redox measurements. C102 oxidizes LH- with a moderate ratez4 of ca. IO6 M-' s-l, suggesting that Eo(LH'/LH-) is close to ca. 0.9 V. The pulse radiolytic experiments were performed at pH 8.1 in N20-saturated solutions containing I O mM luminol and 0.5 M NaN, with the NaCIOz concentration varying from 0.1 to 1 M. In such solutions the following reaction sequence (rates in M-' S-I) takes place: eaq-
+ CIOz-
H20,CIOi
C102
+ OCI- + 20H-
-
0
1mo
1
1 OD
Figure 1. The ratio [CIO;]/[LH-] as a function of the inverse transient optical density a t 450 nm. Conditions: [LH-] = 5 X 10-3-10-2 M, [C102] = 0.1-1.2 M, [N,-] = 0.5 M, pH = 8.1, N 2 0 saturation. The applied dose was 1.4 Gy/pulse. The measurements were made after equilibration.
2.5 X IO9 (ref 25)
+ N z O H20 OH' + N, + OH- 9.1 X lo9 (ref 25) OH' + C102 C102' + OH6.6 X lo9 (ref 25) OH' + N,N3* + OH1.2 X 1Olo (ref 25) N,' + CIOz- ClO,' + N31.9 X lo9 (this work) N3' + LHLH' + N,5.4 X lo9 (this work) OH' + LHvarious radicals (ref 25) 8.8 X IO9 eaq-
0 500
-
Figure 2. Ratio between the rate of equilibration and [LH-] as a function of the ratio [CIO k15is inferred. I n the above scheme it is assumed that protonation equilibria are attained more rapidly than other equilibria. As the rate measurements were performed with rather high buffer concentrations ( 10-3-10-' M) and also since the measured rates did not display general acid-base catalysis this assumption seems reasonable. Figure 4 is a tentative reaction cartoon at pH 11, the optimum pH for luminol chemiluminescence where AG* values were extracted from the Eyring formula for absolute rates. The cartoon clearly visualizes the L 0 2 H 2intermediate as an irreversible energy sink on the reaction path. It is always formed from L + H 2 0 2 . There is also a kinetic control through the protonation state of pair will LH' which decides whether the semiquinone 02*collapse into the hydroperoxide (as in reaction IO) or revert to LH- 0, (as in reaction -6). Strictly speaking, the latter reasoning holds only for homogeneous kinetics. In particular, the question whether a radical pair formed in a solvent cage can attain protonation equilibrium prior to other reactions is an open one. However, it is safe to assume, that, if the Lo- 0,'- pair forms at high pH, i.e., through reaction 7 , it will collapse into LO2,rather than into L2- + 0,. The fact that the reactivity of luminol toward molecular oxygen in water is undetectably low at any pH is clearly reflected in the high redox potential of luminol. This is quantified in Table I1 and Figure 4. Is the Reaction between LH' and 0,'- Adiabatic? The free energy20 of singlet oxygen to '0, is ca. 22 kcal/mol above that of )02.Therefore, equilibrium 17
+
+
+
LH-
+ '0,
LH'
+ 0,'-
(17)
would be slightly endothermic with K17 = 5 X IO-,. On thermodynamic grounds this suggests that the reaction with the measured rate constant k , (see Scheme I) could produce IO2. However, this is very unlikely, as will transpire from the kinetic arguments to follow. The reaction of IO, with LH- has been reported30 to proceed in D20with a rate constant k = 3 X IO7 M-' s-I. However, the reaction produces an unidentified product distinctly different from 5-aminophthalate (P2-). Consistent with this the chemiluminescence yield measured is at least 6 orders of magnitude lower (30) Matheson, I. B. C.; Lee, J. Photochem. Photobiol. 1976, 24, 605.
L'+ 0;
26
L02H-
Figure 4. Reaction cartoon depicting the energetics of luminol autoxidation at pH 1 I .
than would have been obtained had the reaction produced L02H-. First, this shows that the direct reaction LH- '0, L0,His insignificant. Furthermore, it also suggests that the rate constant k17 must be much lower than the observed rate (3 X lo7 M-l s-I 1. Had LH' formed in significantly amounts, it would have been deprotonated either in the solvent cage or subsequent to diffusion out of the cage. In fact, since the experiments were performed at pH 11.7 in the presence of 0.1 M C03,- the rate of deprotonation of LH' would have been on the order of lo9 s-I and thus at least 1% of it would have yielded Lo- prior to reverting to LHin the back reaction -6. Of course, in the sense of Scheme I L'to yield L 0 2 H - and hence chemwould have coupled with 02*iluminescence. We can thus conclude that the rate constant kI7 is less than lo4 M-' s-' a nd therefore k-17 is below IO6 M-' s-l. As this maximum estimated rate is much lower than the measured k , (3 X IO9 M-I s-l ) the reaction between LH' and 0,'- should produce almost exclusively 302. The Autoxidation of Organic Species. The rate constant of self-exchange of the 02/02'couple was experimentally found9 to be 450 M-' s-l. This yields the reorganization energy Xo = Xi, A, = 45.5 kcal/mol. The Xi, value can be estimated from the harmonic approximationz7 to be ca. 16 kcal/mol. This is probably a maximum value from whence the solvent reorganization energy should lie between 30 and 45 kcal/mol, which corresponds to an effective reaction radius of 2-3 A. The self-exchange rate of a large number of aromatic o r g a n i c ~ , ~ * ,such ~ ' - ~as~ phenols,
+
-
+
(31) Meisel, D. Chem. Phys. Lett. 1975, 34, 263. (32) Meisel, D.; Fessenden, R. W. J . Am. Chem. SOC.1976, 98, 7505. (33) Eberson, L. In Electron Transfer Reactions in Organic Chemistry; Springer-Verlag: New York, 1987. (34) Vaish, S. P.; Tollin, G. Bioenergetics 1971, 2, 61. (35) Baxendale, J . H.; Lewin, S . Trans. Faraday. SOC.1946, 42, 126.
J . Phys. Chem. 1990, 94,152-155
752
hydroquinones, and quinones have been determined and found to be close to lo8 M-' s-I. As this rate is determined predominantly by the solvent reorganization energy,,,,A, the effective reaction radii of these couples would seem to lie between 5 and 7 A. In the below analysis the value lo8 M-I s-I w'111 be assumed to hold unless an experimental value is available. Applying the Marcus cross relationshipz7we can then predict autoxidation rate constants and compare them to experimental values. The results of these calculations, which apart from luminol comprise several other organic anions described in the literature, are presented in Table 111. On inspection of the table, some interesting regularities are observed. It is seen that the experimental autoxidation rates of singlet anions are significantly lower than those predicted from the Marcus cross relationship. The opposite is true for the autoxidation of the radical anions. This is particularly apparent upon pairwise comparison of the species deriving from the same parent as surely the geometric change upon one-electron oxidation is minimal. However, the trend is significant for all the species and does not change if we vary the assumed self-exchange rate of the organic couple by an order of magnitude or so. In trying to understand this trend we recall that the rates of self-exchange of couple and even more so of the organics are mainly the 02/02' Furdetermined by the solvent reorganization energies, .,A, thermore, it is almost certain that the effective radii of the organics are larger (by a factor of 2 or more) than that of the 02/02'couple. Consequently, the ,,A, values are not additive, as has been pointed out in ref 33. Denoting by r, and r2, respectively, the and an organic couple reaction radii (in angstrom units) of 02/02'we obtain approximately (in kcal/mol) real) = 90.9(rl-'
+ r2-I
-
2(r,
+ r2)..')
(36) Marcus, R. A. J . Chem. fhys. 1957, 26, 872. (37) Zahir, K.; Espenson, J . H.; Bakac, A. J . Am. Chem. SOC.1988, 110, 5059.
and
+
Aoul(cr)= 90.9((2r1)-' (2r2)-l) where L,(real) and &,,(cr) signify the real solvent reorganization energy of the reaction and the one calculated from the cross relationship. It is readily seen that h,,,(real) - A,,,(cr) = 45.5(r, - r 2 ) 2 / ( r , r 2 ( r 1 r 2 ) ) is always positive. This means that in outer-sphere electron-transfer reactions the real rate constant is always lower than the calculated one (from the Marcus cross relationship). Furthermore, it also transpires that, with realistic values for rI and r2,the rate ratio should not much exceed 1 order of magnitude. The results in Table I11 are thus fully consistent with the assumption that the autoxidation of singlet organics proceed via an outer-sphere electron transfer. However, they also show that the radical anions presented cannot react with oxygen in the same way. The conclusion that radical autoxidation is not an outer-sphere process is not surprising since bond formation between a doublet and a triplet is allowed in contrast to the case of a singlet-triplet reaction. Looking at Table 111 we find that, on the average, the autoxidation rates of the singlet anions are ca. 10 times slower than the predicted ones. Put in another way, couple is ca. 100 the apparent self-exchange rate of the 02/02' times lower than the experimental keX(O2/O2*-). We note that in ref 37 similar apparent self-exchange rates (i.e., 1-10 M-I s-I) were found when the autoxidation of a number of Cr(I1) complexes was analyzed. The authors suggested that these autoxidations were outer-sphere electron-transfer processes. On the basis of the present work we concur with their interpretation.
+
Acknowledgment. Financial support from the Swedish Natural Research Council is acknowledged. Registry No. LH', 123857-54-5; LH-, 66718-60-3; LH,, 521-31-3; V-,89596-65-6; L02H2, 123857-56-7; L2-, 6671 8-67-0; PH2, 5434-20-8; 1 1062-77-4; 0 2 , 7782-44-7; C102, 10049-04-4; CIO,, 14998-27-7; H z 0 2 , 7722-84-1; 4-iodophenol, 540-38-5; phenoxy1 radical, 2122-46-5.
02*-,
Electron Diffraction Studies of the Composition and Structure of Silver Metal Liquidtike Films D. Yogev? S . Shtutina,t and S . Efrima*,+ Ben Gurion University of the Negev. P.O. Box 653, Beer Sheua, Israel (Received: November 15, 1988; In Final Form: April 18, 1989)
We report electron diffraction studies of dried silver metal liquidlike films (MELLFs). These studies are used for fingerprinting the composition and structure of the films. They show the presence of silver in the films with the interplanar spacing characteristic of metallic FCC silver. In addition, they indicate the presence of the surfactant and the silver salt of the surfactant and that of anisic acid (which are needed to form and stabilize the silver MELLFs). The same was found for all the different samples that were investigated, regardless of the detailed composition. Even granular interfacial films behave similarly. A comparison with the results of Raman spectra previously taken from the silver MELLFs reveals differences that are consistent with the model of a condensed colloid and indicates the distribution of the various species within the film.
Introduction
any amarent damage to the film. Aggressive handling may
miscible liquids (an aqueous phase on top of an organic solvent).1-8 We gave these films the name silver metal liquidlike films (silver MELLFs), which is descriptive of their most striking visual properties. On the one hand, the film has a characteristic metallic luster and reflects light as a mirror, much like continuous films of silver. On the other hand, the film exhibits fluid properties: It can be stirred, spilled, and penetrated by hard objects without
'Department of Chemistry. *Department of Physics.
0022-3654/90/2094-0752$02.50/0
(1) Yogev, D.; Efrima, S.J . fhys. Chem. 1988, 92, 5754. (2) Yogev, D.; Efrima, S. J . fhys. Chem. 1988, 92, 5761. (3) Yogev, D.; Efrima, S.; Kafri, 0. Opt. Lett. 1988, 13, 934. (4) Yogev, D.; Deutsch, M.; Efrima, S. J . Phys. Chem. 1989, 93, 4174. (5) Yogev, D.; Efrima, S. Some Chemical Aspects of Silver Metal Liquidlike Films (MELLFs). Submitted for publication. (6) Yogev, D.; Efrima. S. Silver Metal Liquidlike Films (MELLFs): The Effect of Surfactants. Submitted for publication. (7) Yogev, D.; Kuo, C. H.; Neuman, R. D.; Efrima, S . J . Chem. fhys. 1989, 91, 3222. ( 8 ) Yogev, D.; Efrima, S. Macroemulsionsof Silver Metal Liquidlike Films (MELLFs). Submitted for publication.
0 1990 American Chemical Society
