J. Phys. Chem. 1986, 90, 5813-5818
5873
Effect of Micellar Media-on the Electron-Transfer Reaction between Benzylviologen and Quinones S. M. Hubig,
B. C. Dionne, and M. A. J. Rodgers*
Center for Fast Kinetics Research, Patterson 131, University of Texas at Austin, Austin, Texas 78712 (Received: March 10, 1986; In Final Form: June 17, 1986)
The electron-transfer reaction between benzylviologen radicals (BV.') and quinones (Q) [duroquinone (DQ), menadione (MD), 2,3-dimethylnaphthoquinone(DMNQ), and vitamin KI (VK)] has been studied in aqueous and micellar [sodium dodecyl sulfate (SDS)] solution by electron pulse radiolysis. In order to explain the micellar effect on rate and equilibrium constants of this reaction the equilibrium constants of the distribution of BV and Q respectively between micelles and bulk water have been determined by kinetic evaluation of the reaction between micellized Q and hydrated electrons and the reaction between micellized benzylviologen and oxygen, respectively. The data show that benzylviologen is completely localized in the Stern layer while the quinones are partially micellized. Electron transfer occurs only between BV in the Stern layer and quinone molecules arriving at the micelle surface from the bulk water. No electron transfer between BV" and micellized Q within the micelle was observed. Therefore, although the measured rate and equilibrium constants for the electron transfer between BV" and quinone change with increasing SDS concentration, this change is an apparent one only, being caused by changes in substrate concentrations in the bulk solution owing to the presence of micelles.
Introduction
Understanding the mechanism leading to cell damage caused by ionizing radiation requires a better understanding of the first reaction steps occurring after radiation deposition. These first reaction steps, involving free radicals, are recognized in many cases as electron-transfer reactions. Electron-transfer processes are also the primary steps in photosynthesis, thereby leading to the storage of energy in a biologically useful form. Quinone-related compounds often function as electron acceptors in biological electron-transfer reactions and the chemical behavior of p-benzoquinone and its derivatives toluquinone, ubiquinone, and duroquinone as well as the anthraquinones have been studied extensively. Rate constants and equilibrium constants of redox reactions involving quinones have been measured and one-electron redox potentials have been determined.'" The one-electron potential of duroquinone has been of interest because of its use as a reference compound to determine redox potentials of other biologically relevant compounds. Wardman et al. determined the redox potential of duroquinone (DQ)by studies of the one-electron transfer between benzylviologen (BV) and DQ using electron pulse radiolysis:'
Whereas the overwhelming fraction of the electron-transfer literature refers to dilute aqueous solution, in biological systems electron-transfer reactions mostly occur in and around membranes. At such sites, physical properties such as dielectric constant, electrical field strength, protonicity, and hydrophobicity change rapidly over very short distances. Thus model studies using aggregated systems such as micelles, vesicles, liposomes, and ghost cells constitute better approaches to biological condition^.^^'^-'^ The electrochemical properties of methylviologen in a micellar environment were reported recently.'* We examined the above reaction 2 in micellar [sodium dodecyl sulfate (SDS)] solution with several quinones (Q). Using pulse radiolysis, we have studied the micellar effect on rate and equilibrium constants in order to obtain electron-transfer parameters where both benzylviologen and quinone are located at the micelle. Experimental Section
The redox potential Eo(DQ/DQ'-) calculated via the equilibrium constant K of reaction 2 with the Nernst equation was in good agreement with results obtained by different methods earlier.4~s~9J0 The benzylviologen radical was chosen as electron donor in our work because its redox potential is close to that of the DQ/DQ'couple."J2
Pulse radiolysis experiments were carried out with a previously described19 4-MeV Van de Graaff electron accelerator and computer-controlled kinetic spectrophotometric equipment for fast kinetic measurements. Millipore "Reagent Grade" water was used for all experiments. 2-Propanol (spectrophotometric grade) was obtained from Mallinckrodt. Benzylviologen was used as received from Sigma. Duroquinone (tetramethyl- 1,4-benzoquinone) (Aldrich) was purified by vacuum sublimation or by recrystallization from ethanol. Vitamin K, (2-methyl-3-phytyl-l,4naphthoquinone) was received from Sigma and used without further purification. Menadione (2-methyl-1 ,Cnaphthoquinone) and sodium dodecyl sulfate (specially pure) were used as received from Gallard-Schlesinger (BDH). 2,3-Dimethyl- 1,Cnaphthoquinone was kindly supplied by Dr. David Creed, University of Southern Mississippi. For all experiments in aqueous solution,
(1) Wardman, P.; Clarke, E. D. J. Chem. SOC.,Faraday Trans. I 1976, 72, 1377. (2) Patel, K. B.; Willson, R. L. J . Chem. Soc., Faraday Trans. I1973,69, 814. (3) Ilan, Y.A.; Czapski, G.; Meisel, D. Biochim. Biophys. Acta 1976,430, 209. (4) Meisel, D.;Czapski, G. J . Phys. Chem. 1975, 79, 1503. ( 5 ) Ilan, Y.A.; Meisel, D.; Czapski, G. fsr. J . Chem. 1974, 12, 891. (6) Almgren, M.; Grieser, F.; Thomas, J. K. J. Phys. Chem. 1979,83, 3232. (7) Meisel, D.;Neta, P. J . Am. Chem. Soc. 1975,97,5198. (8) Afanas'ev, I. B.; Polozova, N. I. J. Org. Chem. USSR (Engl. Transl.) 1979, 15, 1621. (9) Baxendale, J. H.; Hardy, H. R. Trans. Faraday Soc. 1953,1140,1433. (10)Wood, P. M. FEBS Lett. 1974,44, 22.
(11) Michaelis, L.; Hill, E. S.J . Gen. Physiol. 1933,16,859. (12) Steckham, E.; Kuwana, T. Ber. Bunsen-Ges. Phys. Chem. 1974.78, 253. (13) (a) Turro, N. J.; Graetzel, M.; Braun, A. M. Angew. Chem., Int.Ed. Engl. 1980,19,675. (b) Rodgers, M. A. J. Radiat. Phys. Chem. 1984,23, 245. (14) Rodgers, M. A. J.; Foyt, D. C.; Zimek, Z. A. Radiat. Res. 1978,75, 296. (15) Almgren, M.; Grieser, F.; Thomas, J. K. J . Am. Chem. SOC.1979, 101,279. (16) Chevalier, S.;Lerebours, B.; Pileni, M. P. J . Photochem. 1984,27, 301. (17) Carbone, A. I.; Cavasino, F. P.; Sbriziolo, C.; Pelizzetti, E. J . Phys. Chem. 1985,89, 3578. (18) Kaifer, A. E.; Bard, A. J. J . Phys. Chem. 1985,89,4876. (19) Foyt, D. C. Comput. Chem. 1981,5, 49.
BV2+ BV*+
-
+ e-aq
+ DQ
BV'+
BV2+ + DQ''
0022-3654/86/2090-5873$01.50/0 , 0 1986 American Chemical Society I
,
5814
The Journal of Physical Chemistry. Vol. 90, No. 22, 1986
Hubig et al.
:I
[ao1
/
; ; ; i /
10 ps 1
' I
I
I
I
1
TIHE
Figure 1. Decay of the absorption of BV" radicals monitored at X = 600 nm after a 250-11selectron pulse: fast component, first-order approach to equilibrium; slow component, second-order decay of the equilibrium due to redox reaction between BV'+ and Q'- radicals. A. = initial absorbance; A , = absorbance at equilibrium.
+
a phosphate buffer (0.002 M KHzP04 0.003 M NaZHPO4)at pH 7 was used. Both micellar and aqueous solutions contained 0.2 M 2-propanol for removal of hydroxyl radicals and were deaerated by either bubbling or blowing and stirring with nitrogen for ca. 1 h.
Results and Discussion Electron Transfer in Aqueous Solution. Pulse radiolysis of a deaerated aqueous solution produces molecular hydrogen, hydrogen peroxide, hydrated electrons, hydrogen atoms, and hydroxyl radicals. These last two are scavenged by 2-propario1, producing isopropyl radicals. Both the hydrated electrons and the isopropyl M BV*+) to radicals reduce the benzylviologen cations (2 X BV" radicals (see eq 1). With an absorbed dose of near 400 rad, the initial concentration of BV'+ radicals produced by one pulse was about 2 pM. Since the Q concentrations used were in all cases much higher (between 20 and 100 pM),the kinetics of the electron-transfer reaction (see eq 2) observed by measuring the change in the absorbance of the BV*+ radical at 600 nm were characterized by a first-order approach to equilibrium (kLbsd) followed by a slower second-order decay due to the electron transfer from BV" to DQ'- (Figure 1). Under such conditions the bimolecular rate constants for the electron transfer can be calculated from the observed first-order rate constants ( k 6 M ) as a function of the quinone concentration according to the following equation: kbbd =
MQ1 + kb[BVz+j
(3)
kf,b are the bimolecular rate constants for the forward reaction and the back-reaction, respectively. As the quinone concentration is varied, the plot of the observed first-order constants k'obsdvs. the quinone concentration gives the bimolecular rate constants k, and kb (Figure 2). In the case of duroquinone the experiments led to a kf = (3.4 f 0.1) X lo9 L mol-' s-l, from an average of four separate experiments. This value was independent of buffer concentration in the range 2-20 mM, which corresponded to an ionic strength change from 0.017 to 0.12. Under the same conditions we found the k b values varying between 1.2 X lo7 L mol-' s-' (ionic strength 0.12) and 3.5 X lo7 L mol-I s-I (ionic strength 0.017). Even though these two constants show an apparent effect of ionic strength, we are reluctant here to examine this quantitatively since the inaccuracies involved in kb determination by extrapolation of plots such as that in Figure 2 tend to obscure such trends. For the same reason we are reticent about the use of the kinetic approach (Kq,hn = kf/kb) for the determination of the equilibrium constant, as previously noted by Wardman and Clarke.' Therefore we adopted the Kq evaluation method by using
g , , 0.000
0.028
0.040
, 0.060
, r e $ ~ O
, 0.100
; 0.128
Figure 2. Determination of the bimolecular rate constants for duroquinone: observed first-order rate constant k 6 , divided by the BV concentrationvs. durcquinone concentrationdivided by the BV concentration. The slope gives the kf value; the intercept is kb. measurements of the concentrations of BV'+ at equilibrium in the presence and absence of duroquinone:' Kq =
Wz+I PQ.7/([BV'+l
PQ1)
(4)
For absorbance measurements at 600 nm, where only the B Y + radical absorbs, the above equation can be modified:
= ( P v z + l ~ / [ D Q l ~ ) (-AA ~w ) / A m
(5)
A. is the initial absotbance at 600 nm in the absence of duroquinone, A , is the absorbance at equilibrium in the presence of duroquinone, and [hVz+]oand [DQlo are the concentrations of reactants prior to the pulse. Measurements using this approach allowed Kq,abs= 137 f 17 to be computed in close agreement with the earlier determinations.' For menadione we obtained results similar to those for duroquinone (Table 11), while the equilibrium constant for dimethylnaphthoquinone = 8.3) was about 1 order of magnitude lowei. when measured by the absorption technique. It is worth noting that for this molecule kb was an order of magnitude higher than for DQ and MD, leading to a value of Kq&n = 14.5, not far from the KqSb value (see Table 11). Both, however, are very much lower than that expected (K = 115) from the published one-electron redox ~ t e n t i a l .In ~ %is, our data are in line with those of Wilson.20 For vitamin KIowe could not measure the rate and equilibrium constants under the same conditions because of its insolubility in water. Nor a u l d we observe electron transfer from benzylviologen radical to vitamin K either in ethanol or in different ethanolaqueous buffer mixtures. Ledwith reportedz1 a strong change in the one-electron redox potential of methylviologen with increasing ethanol content in water. A similar ethanol effect on the reduction potentials of BY' and vitamin K1 could lead to changes in redox (20)
Wilson, I., personal communication.
(21) Ledwith, A. In Biochemical Mechanism of Paraquat Toxicity; Autor, A. P., Ed.; Academic: New York, 1977; p 21.
The Journal of Physical Chemistry, Vol. 90, No. 22, 1986 5875
Benzylviologen-Quinones Reaction in Micelles
K*q
kf YI I O ' 30
Is0
20
loo
"i
0
A
10
50 0
X
A
X
4
4i
[SDS]
in
Mol/L
z
centration for duroquinone. TABLE I: Rate Constants and Equilibrium Constants for the Electron Transfer from Benzylviologen Radical to Duroquinone in Micellar Solution' lo-* kf, L lo-' kb, L [SDS], [micelle]! mol/L mmol/L mol-' s-l mol-' s-I Kk;K., 0 34f 1 2.9 117 137f 17 0 0.020 0.19 6.3 f 0.1 1.0 63 51 f 2 0.030 0.35 4.4 f 0.3 1.2 31 39f5 0.050 0.68 2.6 f 0.4 2.3 11 12f 5 1.1
1.5
1.9 f 0.1 1.8 f 0.2
1.9 2.6
10 7
a
f
*
Q
Figure 3. Equilibrium constants K (A)and Kcp,b (0)and the forward rate constant kf (X) in L m 8 s-l as function of the SDS con-
0.075 0.100
0
f
0.05
0
10f2 6f1
'The quinone concentrations varied between 20 and 100 pM for experiments in purely aqueous solution and between 0.1 and 1.4 mM in the case of micellar solution. bcmc = 8.1 mmol/L and aggregation number N = 62 (ref 22 and 23). potentials sufficient to invert the thermodynamic tendency. Electron Transfer in Micellar Solution. For the same reaction studied in SDS solutions in the range 20-100 mM the kf values and the measured equilibrium constant (K,) varied drastically with detergent and micelle concentration, while kbwas independent of the SDS concentration (for DQ: Table I and Figure 3). Menadione showed similar results (see Table 11). Vitamin K1 and dimethylnaphthoquinone showed no evidence of electron transfer from the BV'+ radical for SDS concentrations between 20 and 200 mM. The micelle concentration [MI is calculated via the following relationship: [MI = ([SDS]- cmc) / N (6) For the critical micelle concentration (cmc) a value of 8.1 mmol/L is taken. The aggregation number N is assumed to be 62.22,23 In order to explain the micellar effect it has to be considered that (22) Turro, N.J.; Yekta, A. J . Am. Chem. SOC.1978, 100, 5951. ( 2 3 ) Almgren, M.;Griesser, F.;Thomas, J. K.J . Chem. SIX., Faraday Trans. 1 1979, 75, 1674.
Q
I
0.000
4.QC0
8.089
0.120
0.160-
0.200
SDS-CONCENTRRTI 0th 1N MOL/L
0.240
Figure 4. First-order rate constants for the electron decay monitored at X = 700 nm in a lo4 M DQ solution as function of the SDS concentration (curve A). Curve B shows the electron decay in absence of
duroquinone. both quinone and benzylviologen are a t least partially micellized, Le., either solubilized in the hydrophobic regions of the micelle or closely associated with the negatively charged Stem layer. We are concerned with the site of the electron transfer, whether it occurs within a micelle containing both a quinone molecule and a benzylviologen radical or between a radical associated with the micelle surface and a quinone molecule arriving a t the micelle surface from the bulk water. This necessitates knowledge of the way that Q and BV2+partition themselves between micelles and bulk solution. Therefore the equilibrium constants of the micellization of both the quinone and benzylviologen were determined by pulse radiolysis experiments. Micellization Equilibrium of Quinones. The micellization of the quinones was studied by measuring the first-order rate constant for the reduction of quinone by hydrated electrons Q
+ e-aq
-
Q'-
(7)
We monitored the decay of the absorption of the hydrated electrons (e-4) at X = 700 nm in both aqueous and micellar solution for several Q concentrations. The bimolecular rate constants obtained for duroquinone, menadione, and dimethylnaphthoquinone in buffered aqueous solution (kaq)and in micellar (0.2 M SDS) solution (kmic) are given in Table 11. Further meafor the electron surements of the first-order rate constants ),:CI( decay in lo4 M quinone solutions were carried out for several SDS concentrations between 0.01 and 0.1 M. The value of ,:k was shown to strongly decrease with increasing SDS concentration (Figure 4). Thus we conclude that as the micelle concentration is increased more Q is micellized and the rate of loss of hydrated electrons by reaction 7 becomes attenuated because of repulsive Coulombic forces between the electrons and the negatively charged micelle ) for SDS surface. The bimolecular rate constants ( k ~ cobtained concentrations above 0.1 M (plateau region in Figure 4) are
Hubig et al.
5876 The Journal of Physical Chemistry, Vol. 90, No. 22, 1986 TABLE II: E4uilibrium Constants K and Bimolecular Rate Constants k (in L mol-' sd) for Various Quinone Compounds
menadione KM,M-I
12000
kaq kmic
(3.1 f 0.03) X 10" (6.5 f 0.2) x 109
haq kb,aq
(3.9 f 0.5) x 109 (1.0 i 0.5) x 107 143 f 9 (2.2 f 0.2) x lo8 (1.3 f 0.2) x 107 21 f 1
duroquinone dimethylnaphthoquinone Micellization Equilibrium 13 000
26 000
vitamin K a
Reaction with Hydrated Electrons
Kas kf,mic kb,mic
Kmic a
(3.1 f 0.2) X 1O'O (4.2 f 0.3) X lo9
(3.8 f 0.1) X 1O'O (2.8 f 0.4) X lo9
Reaction with Benzylviologen Radicals (3.2 f 0.3) (3.4 f 0.1) x 109 (2.9 f 0.7) X 137 f 17 (1.8 f 0.2) X (2.6 f 0.7) X
lo7
X
(1.0 f 0.08) X1O1Ob
no reaction!
lo9
(2.2 f 0.2) x 108 8.3 0.2
no reaction!
no reaction!
no reaction!
*
lo8 lo7
6 f l
K M not measurable because vitamin K , is insoluble in water. bMeasured in ethanol.
regarded as those for the reaction between e-aq and totally micellized quinone. Vitamin Kl reacts with solvated electrons in ethanol, forming the radical anion. However, we were unable to observe any such reaction in micellar solution suggesting that the vitamin is completely micellized in hydrophobic regions inaccessible to hydrated electrons. The electron-decay rate constants obtained for vitamin concentrations between 10 and 70 pmol and for SDS concentrations between 20 and 200 m M were the same as measured in solutions without quinone (k;b@-J= 3 x io5 s-l). For DQ, MD, and D M N Q we calculated the distribution equilibrium constant between micelles and bulk water as follows: ) at X = 700 nm The first-order electron decay ( k i b @ - Jobserved is caused by three additive decays: the electron decay caused by the medium (k:,), the electron decay caused by free quinone in the bulk water (kaq),and the electron decay caused by micellized quinone (kmic) k ' o ~= k'e~+ kaq[Qaql + kmic[Qmicl
SLOPE 1.2835 X l B q
lNTERCEPT 0.89115 X le
STD. DEVN. 0.8530 X 1 9
.
(8)
When the micellization equilibrium Qaq
+M
* Qmic
(9)
is considered, with KM = [QmicI/([Qaql[MI) and
[Qmicl
= [Qt.J
- [Qaq], the following equation is given:
[Qtota~l/[Qaql = 1 + K M I M l
(10)
With eq 8, the concentration of free quinone [Qaq]can be calculated [Qaql
=
(k'obsd
- k'.~ - kmC[Qtotad)/(kaq - kit)
Combining eq 10 and 11, we obtain S = ([Qtotad(kaq - kmic))/(k;bsd
- k'e~- kmic[Qtota11) = (12)
With [Qmul]= 1 X lo4 mol/L (in our experiment), the data for k , and khc from Table 11, and k ; = 3 X lo5 s-l (first-order decay of electrons measured in micellar solution without quinone), S can be. plotted as a function of the micelle concentration (Figure 5 ) . The slope gives the micellization equilibrium constant KM (see Table 11). The value for duroquinone (KM = 13O00 L mol-') is close to that of KM = 9300 L mol-' obtained by solubilization experimenk6 From this we calculate that at the plateau region in Figure 4 the amount of duroquinone present in the aqueous phase is less than 3% of the total. Micellization Equilibrium of Benzylviologen. As for the quinone situation, it is important to know the location of BVZ+ and BV" in the micellar system. W e were unable to obtain quantitative data for these systems but we make the following observations: 1 . The bimolecular rate constant for the reaction
+ e-aq
-
0.1100
0.800
1.600
M I C E L L E COdC%TRFIT
2.000
2 400
I ON I N M O L / L
103
Figure 5. Determination of the micellization equilibrium constant K M : S value as function of the micelle concentration.
1 + KMIMI
BV2+
a00
(11)
BV'+
(13)
1Olo L mol-' s-') is shown in Figure 6.24 2. The BV2+ species with its double positive charge and hydrophobic benzyl groups will have a high affinity for the micellar pseudophase. 3. Monoreduced viologen species are more hydrophobic than their dicationic parent species.I3 4. The reaction X
BV'+
*
-
BV2+
+ 02'-
(14)
was shown to have a 10 times lower rate constant in all the SDS concentrations (0.01-0.1 M)compared to buffered aqueous media and did not depend neither on the BV2+concentration between 100 and 600 pM nor on the oxygen content. Similar observations have been reported for the oxidation of reduced methylviologen ions by 0 x ~ g e n . l ~ ~ 5 . The decay of BV" with oxygen showed clean exponential behavior. This indicates that the micellization equilibrium for BV'+ is much faster than its reaction with oxygen, otherwise the ~~
is reduced by a factor of ca. 10 when measured in micellar SDS solutions. The measurement in buffered water (kbh = (4.5 0.1)
+ O2
~
(24) See also: Farrington, J. A.; Ebert, M.; Land, E. J. J . Chem. Soc., Faraday Trans. I 1978, 74, 665: k,, = 7.9 X 10'O L mol-' s-'.
The Journal of Physical Chemistry, Vol. 90, No. 22, 1986 5877
Benzylviologen-Quinones Reaction in Micelles
SLOPE
INTERCEPT
S T D . DEVN.
we have used the stochiometric concentrations of quinone, the assumption being made that all Q molecules present in the medium are capable of participating in reaction 2. The above example, however, leads to a suspicion that the fraction of Q that is micelle-associated is reacting only very slowly, if at all, with BV'+ tadicals in the same micelle, and therefore the observed reaction is only arising from the residual Q molecules in bulk solution. This is affirmed as follows: If only free quinone reacts with micellized benzylviologen, then the equilibrium constant of the electron transfer K
( [BVz+l[Q+I ) / ( [BV*+l[Qaql 1
(1 5)
will depend on the micelle concentration through [Qas], which is the only quantity in eq 15 that is not independent of the micelle concentration. Then, following the method of Almgren et a1.6 for the case of electron transfer from anthraquinone sulfonate radical anion to duroquinone, the following equation should be valid: Combining eq 10 and 16, we obtain
0.100
81.300
0.500
0.700
81.900
1'.100
B V - C O N C E N T R R T I D N IN MOL/L
1'.300x 1°'
Figure 6. Observed first-order electron decay monitored at X = 700 nm
as function of the benzylviologen concentration in an aqueous solution containing 0.002 M phosphate buffer: The slope gives the bimolecular rate constant of the reduction of BV by hydrated electrons. The intercept zylviologen. difference in rate constants (see 4 above) would lead to a biexponential decay of the BV'+ radicals. These several observations lead to our assertion that BV2+and BV'+ ions are extensively micellized under our conditions. Electron Transfer between Micellized Benzylviologen Radicals and Free Quinone. Above we noted that in micellar media the measured equilibrium constants for reaction 2 depend on the SDS concentration. This can be understood as follows: We start from the premise that all benzylviologen species are micellized and the quinones are partially associated with the micelles over the concentration range used according to the equilibrium constants reported above. Then we need to examine whether the electron transfer occurs within one micelle or between micellized benzylviologen radicals and free quinone arriving from bulk solution. Before we proceed with this examination it is important to consider the distribution of the quinone molecules within the micellar population. Having been formed by e, capture a BV'+ radical at the micellar periphery can, in principle, undergo electron transfer to Q species either comicellized at the same micelle or arriving in the reaction sphere of the B V " radical from bulk solution as a result of normal diffusive motion. The former can only be observed if there is a high probability of finding a Q species occupying the same micelle. In our experiments the surfactant and quinone concentrations were chosen to ensure this. For example, a t [DQ] = 1.4 m M and [SDS] = 50 mM, the micelle concentration is equivalent to 0.68 m M (Table I) and, according to the equilibrium between DQ and micelles, the concentration of DQ micellized is 1.3 mM. Then with Poisson statistics, with a mean value of 1.3/0.68 = 1.9 quinone molecules per micelle, the probability of zero occupancy is 0.15; i.e.,85% of micelles have at least one associated Q molecule. In this solution (Table I) the measured equilibrium constant for electron transfer according to eq 2 is only 9% of the value in buffered water-a reduction that is governed largely by the fall in the observed k f value (Table I). In extracting the kfvalue and the Keq,abvalues,
The data for duroquinone (Table I), when plotted according to eq 17, give a linear plot. Thus the results are consistent with the initial premise, viz, that micellized viologen is reacting only with DQ molecules arriving from the bulk medium. Further support comes from the extracted value of KM = 12000 M-', in good agreement with KM = 13 000 M-' determined as described above (eq 12). The similarity between the rate parameters for M D and DQ (Table 11) leads us to anticipate that the same is true for MD. The fact that the bimolecular rate constant kb for the back electron transfer is independent of the SDS concentration can be understood on this model: If BV*+only reacts with free quinone arriving a t the Stern layer from the bulk solution and if the produced Q'- anion is subsequently ejected from that region by repulsive Coulombic forces, then all the Q'- species are available for further reactions and are not sequestered into hydrophobic regions of the micelle as Q probably is. Furthermore the data of Table I1 clearly show the following trend for electron-transfer reactions in micellar solution: The micellization equilibrium constant KM increases in the order menadione < duroquinone < dimethylnaphthoquinone < vitamin K,; i.e., from menadione to vitamin K1 the hydrophobic properties increase and the solubility in water decreases. The rate constants for the reactions between micellized quinones and hydrated electrons decrease in the order M D > DQ > D M N Q and reach an immeasurably low value for vitamin K,; Le., the more hydrophobic the quinone molecule, the less accessible it is for electrons, probably due to a deeper penetration of the quinone in the hydrophobic region of the micelle. Our major conclusion is that BV" radicals do not react at a measurable rate with quinone molecules associated with the hydrophobic regions of the micelle. The fact that we did observe reduced rates of electron transfer between B V " radicals and menadione or duroquinone respectively in micellar solution is because of the arrival of free quinone molecules from the bulk aqueous phase. In the cases of dimethylnaphthoquinone and vitamin K, we were unable to observe any electron transfer in micellar solution, presumably because the micellization equilibrium constant ensures that, under our conditions, the equilibrium concentration of these quinones in the bulk phase is very small. The reason BV'+ radicals do not undergo electron transfer to Q molecules within the same micelle is not yet clear. Two possibilities suggest themselves: 1. The two entities are located a t different micellar regions (quinones at hydrophobic sites; BV'+ in the Stern layer) and because of this are unable to encounter each other. 2. The micellar environment imparts large enough changes in the redox properties of the participants that electron transfer is no longer thermodynamically feasible. Currently we are devising experiments to attempt a decision between these possibilities.
5878
J. Phys. Chem. 1986, 90, 5878-5882
In passing we note that Graetzel et al.25 have argued that bimolecular rate constants for the reaction of hydrated electrons and electron acceptors associated with anionic micelles should be diffusion-controlled when the gas-phase electron affinity of such acceptors is between 1.6 and 2 eV. Chen et a1.26have published values of EA = 1.65 and 1.7 1 eV for DQ and MD, respectively, but our findings are that the electron-capture rate constants are significantly lower than the diffusion-controlled ones in micellar (SDS) medium (Table 11). This suggests that such rate parameters in micellar media are not governed by electron-tunnelling (25) Frank, A. J.; Graetzel, M.; Henglein, A.; Janata, E. Eer. Bunsen-Ges. Phys. Chem. 1976.80, 547. (26) Chen, E. C. M.; Wentworth, W. E. J . Chem. Phys. 1975,63, 3183.
reactions as indicated earlier25but by normal collisional encounters moderated by electrostatic and hydrophobic forces.
Acknowledgment. The Center for Fast Kinetics Research is supported jointly by the Biotechnology Resources Program of the Division of Research Resources of NIH R R 00886 and by the University of Texas at Austin. Additional support came from NIH Grant G M 31603. We are grateful to B. K. Naumann for his expert technical assistance in the electron accelerator operation and to Dr. David Creed, University of Southern Mississippi, for his generous gift of dimethylnaphthoquinone. Registry No. Benzylviologen radical cation, 49765-27-7; menadione, 58-27-5; durquinone, 527- 17-3; 2,3-dimethylnaphthoquinone,2197-57-1; vitamin K, 12001-79-5.
Micellar Solutions of Sulfate Surfactants Studied by ESR of Nitroxide Radicals. 3. Effect of Added Electrolytes Piero Baglioni, M. Francesca Ottaviani, and Giacomo Martini* Department of Chemistry, Colloid and Interphase Group, University of Florence, 501 21 Firenze, Italy (Received: March 19, 1986; In Final Form: June 24, 1986)
Micellar and premicellar solutions of sodium octyl, dodecyl, and hexadecyl sulfate as a function of surfactant and sodium chloride concentrations were studied by ESR of nitroxide probes. The study was based mainly on an analysis of the line shape, of the hyperfine coupling constant (AN), and of the correlation time for the motion, with the aim to investigate the polarity and the microviscosity of the micellar double layer. The addition of sodium chloride led to a decrease of the surface polarity and to an increase of the interfacial microviscosity. The trend of the (AN) values as a function of chain length showed that the fluid oillike interior with double layer model was no longer valid for surfactant with short chain lengths (Le. octyl sulfate). Possible structuring effects of water near the micellar surface were also discussed.
Introduction In parts 1 and 2 of this series some properties of premicellar and micellar solutions of anionic surfactants (sodium octyl, dodecyl, and hexadecyl sulfate) were studied through the analysis of the ESR line shape of opportunely chosen nitroxide radicals.’*2 The neutral Tempo1 and positively charged TempTMA’ give information on the “polarity” and on the “fluidity” of the diffuse double layer and of the Stern layer.’ Spin probes with long hydrocarbon chains give more details on the “status” of the micellar interface and on the micellization process.2 Both the above works were carried out in the absence of added electrolytes. Changes in the ionic strength and in the amphiphile concentration are indeed known to induce changes in the micellar shape and ~ i z e . ~These - ~ findings are explained by assuming that the micelle shape and size are dependent on an equilibrium among forces of different natures; in particular, the hydrophobic forces are supposed to increase the aggregation number, whereas the electrostatic forces tend to decrease From a geometrical point of view, an increase in the aggregation number leads to a decrease in the area available for each polar group, thus favoring a decreased number of hydrocarbon-water contact points. As a consequence, a decreased interface polarity results. Although a (1) Ottaviani, M. F.; Baglioni, P.; Martini, G. J . Phys. Chem. 1983, 87, 3146. (2) Baglioni, P.; Ferroni, E.; Martini, G.; Ottaviani, M. F. J. Phys. Chem. 1984.88, 5107. (3) Wennerstrom, H.; Lindman, B. Phys. Rep. 1979, 52 ( l ) , 1-89. (4) Lindman, B.; Wennerstrom, H. Topics Current Chem. 1980,87, 1-80. ( 5 ) Rosen, M. J. Surfactants and Interfacial Phenomena; Wiley: New York, 1978. (6) Tanford, C. The Hydrophobic Effects; Wiley: New York, 1980. (7) Israelachvili, J. N.; Mitchell, D.J.; Ninham, B. W. J . Chem. SOC., Faraday Trans. 2 1976, 72, 1525. (8) Israelachvili, J. N. Intermolecular and Surface Forces; Academic: New York, 1985.
lot of information exists on micellar shape and size and on some properties such as solubilization and micellar catalysis, the polarity and the status of water at the micellar interface have been not thoroughly investigated. The use of suitable probes in micellar systems may also be useful to obtain information on the micelle interface region. For instance, Zachariasse and c o - w ~ r k e r s very ~ ~ ’ ~recently verified, by using the phenolbetaine probe., that variations of the interface polarity are correlated with variations of the aggregation number and of the micellar shape. Although the use of probes has been criticized’ because of the perturbations they may introduce even at very localized levels’ and because their precise localization is very difficult to determine, molecular probes for ESR and fluorescence spectroscopy have been extensively employed in studies of membranes, of membranelike structures, of micelle systems, and in other lipid assemblie~.’~*’~ Details on the fluidity, on the polarity, and on the dynamics of these systems, difficult to get with other experimental techniques, have thus been obtained. The micellar and premicellar solutions of the same sulfate surfactant systems investigated in parts 1 and 2’s2 were analyzed in this paper in the presence of different concentrations of sodium chloride with the aim to study the effect of the added electrolyte on the microviscosity and on the micropolarity of the spin probes as a function of the surfactant chain length and concentration.
’
(9) Zachariasse, K. A.; Phuc, N. V.; Kozankiewicz, B. J. Phys. Chem. 1981. 8.5. 2676. (10) Zachariasse, K. A.; Kozankiewicz, B.; KUhnle, W. In Surfactants in Solutions, Vol. 1, Mittal, K. L., Lindman, B., Eds.; Plenum: New York, 1984; p 565. (1 1) Cadenhead, D. A.; Muller-Landau, F. Adu. Chem. Ser. 1975, No. 144, 294. (12) Marsh, D. In Membrane Spectroscopy, Grell, E., Ed.; SpringerVerlag: West Berlin, 1981; pp 51-142, and references therein. (13) Schreier, S.;Polnaszek, C. F.; Smith, J. C. P. Eiochim. Eiophys. Acta 1978, 515, 395.
0022-3654/86/2090-5878$01.50/0 0 1986 American Chemical Society
