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The effect of micellar media on photoinduced electron transfer reactions in methylviologen-naphthalene ground-state charge-transfer complexes
Journal of Luminescence 47 (1991) 137 145 North-Holland
137
The effect of micellar media on photoinduced electron transfer reactions in methylviologen— naphthalene ground- state charge-transfer complexes Stephan M. Hubig Center for Fast Kinetics Research, The University of Texas at Austin, Austin, TX 78712, US4 Received 17 May 1990 Revised 30 July 1990 Accepted 30 August 1990
We have studied photoinduced electron transfer reactions by photoexcitation of ground-state charge-transfer complexes between methylviologen ions and several differently charged naphthalene derivatives. Using laser flash photolysis and picosecond absorption spectroscopy techniques, we have determined ion pair lifetimes and charge separation efficiencies. One-electron oxidation potentials of the naphthalene compounds have been measured by cyclic voltametry and complex formation equilibrium constants have been obtained from Benesi Hildebrand plots. Since most of the biologically relevant electron transfer reactions occur in or close to membranes, we have been interested in the effects of interfaces on the electron transfer reactions studied. Using sodium dodecyl sulfate micellar dispersions as model interfacial environment we have studied the micellai effects on charge recombination and escape rate constants and on the escape yields.
1. Introduction In recent years many studies on photoinduced electron transfer reactions have been carried out in order to explore nature’s photosynthesis mechanisms and to develop artificial systems for efficient solar energy conversion [1 3]. In order to improve solar energy conversion systems one needs to investigate the electron transfer reactions occurring immediately after photoexcitation. There are in principle two mechanisms of photoinduced charge separation to distinguish: (a) photoexcitation of a potential electron donor or acceptor followed by electron transfer quenching (kq) occurring at an encounter with a quencher; or (b) photoexcitation of a ground-state charge-transfer complex leading directly to a charge-separated (ion pair) state (see scheme 1). According to Mataga et al. [4 8] the two mechanisms also result in two differently structured ion pairs, compact or loose ion pairs. Once the charge separation is completed, there are two possible reaction path0022-2313/91/$03.50 © 1991
ways: charge recombination (krec) within the ion pair or escape (k~5~)of the ions out of the solvent cage around the ion pair, leading to “free” solvated radical ions. The efficiency of charge separation, i.e. the yield of escaped ions (~esc), can be determined according to the following equation: ~esc
=
abs
=
k
kesc + k
.
=
kesc
.
T,
(1)
where [W] represents the molar concentration of either solvated radical ion, D~ or A ‘ab, is the molar concentration of photons absorbed, and T is the ion pair lifetime. According to eq. 1 two experiments are necessary to determine k rec and kesc: (1) The determination of ~esc by measuring both the amount of solvated radicals produced (R~)and the dose of light absorbed (‘abs)~ The radicals can be observed directly, by measuring
Elsevier Science Publishers B.V. (North-Holland)
138
S.M Hubig
(a) D*
A or D + A*
Effect of mice//ar media on electron transfer reactions
transfer absorption band [4 8, 10 14]. In addition to the advantage of fast ion pair formation, ground-state charge-transfer complexes provide another convenience in being simple systems where the ion pair state is the only excited state formed upon irradiation, as opposed to dynamic quenching experiments where the ion pair state is formed from other excited (singlet or triplet) states which complicates the spectroscopical observation of the transients. In this paper we present data on charge transfer complexes between methylviologen ions (MV2~) and several naphthalene derivatives.
+
~kq ~
[D
. .
1
A
II
~
÷
D
+
A
D
A
secondary ET reactions (b) [D..
A]*
[D~ [D.
. .
Al
. .
A
.1
Using laser flash photolysis and picosecond sorption spectroscopy techniques, we have abdetermined escape yields and ion pair lifetimes. Redox potentials were measured by cyclic voltametry, and complex formation constants were obtained from Benesi Hildebrand plots [15]. For electron transfer reactions within ion pairs
D”+A E
Scheme 1. Reaction schemes for mechanisms (a) and (b) (see text) of photo-induced charge separation (D electron donor, A electron acceptor).
their transient absorption spectrum, usually on a microsecond timescale, or indirectly using appropriate trapping reagents [9]. The amount of photons absorbed can be determined using chemical actinometers. (2) The determination of the lifetime T (kesc + kr~) of the initial ion pair [D A 1 by fast kinetic techniques. Charge recombination reactions within ion pairs may occur on very short timescales (T < 1 ns). To measure such short ion pair lifetimes, the initial photoinduced charge separation leading to the ion pair must be completed within an even shorter time period. In the case of charge separation by dynamic electron transfer quenching the charge recombination and the ion escape reaction cannot be followed kinetically because the ion pair formation is too slow compared to the decay, and the only way to obtain charge recombination rate constants from escape yields is by taking an estimated escape rate constant of 5 x 108 ~ 1 [9]. A successful way to produce ion pairs very fast is the photoexcitation of ground-state charge-transfer complexes at a wavelength within the charge—
~
. .
Coulombic forces play an important role as driving force since the distance between the charged species is very small. In order to study such Coulombic effects we chose naphthalene derivatives of different charges including naphthalene (nap), 2-naphthalene sulfonic acid sodium salt (napSO 3 ), 1,5-naphthalene disulfonic acid disodium salt (nap(S03 )2)~ and N-(1-naphthyl) ethylenediamine dihydrochloride (nap-NH~C2H4 NH~)[161. Most of the electron transfer studies reported in the literature were carried out in homogeneous solutions. However, most of the biologically relevant electron transfer reactions occur in or around biomembranes (cell membranes or membranes of organelles) and there is a lack of information about the effect of interfaces on electron transfer reaction parameters [17 22]. For instance, Coulombic forces are expected to play an important role not only within the ion pair, as mentioned above, but also between the charged interface and either charged ion pair partner. There is also little known how interfacial sites influence the formation of charge transfer complexes [23,24] and the redox potentials of electron transfer partners [25]. For this study we chose sodium dodecyl sulfate (SDS) micellar dispersions as model environment and compared the results in homogeneous and heterogeneous media.
S. M. Hubig
/ Effect of micellar media on electron
2. Experimental
electrolyte, except for naphthalene where a 75 : 25 water-acetonitrile mixture was used. The absorption spectra Benesi—Hildebrand plots were obtamed from for a Hitachi U3210 spectrophotometer. Methylviologen dichloride was obtained from Sigma and used as received. For experiments with naphthalene in acetonitrile the hexafluorophosphate salt was used which was made by mixing the dichloride salt with ammonium hexafluorophosphate (Aldrich) in water, yielding a white precipitate. Naphthalene (scintillation grade) and acetonitrile (HPLC grade) were purchased from Aldrich, the naphthalene sulfonic acid derivatives (napSO3 and nap(S03 )2) from Eastman Kodak, N-(1 -naphthyl)ethylenediamine dihydrochloride from FLUKA, and sodium dodecyl sulfate (SDS, “especially pure” grade) from Gallard-Schlesinger (BDH). For all aqueous and micellar solutions Millipore “reagent grade” water was used.
3. Results 3.1. Ground-state charge-transfer complex formation
From the four naphthalene compounds studied in this paper, three including napSO3 and nap (SO3 )2 in water and nap in acetonitrile, form ground-state complexes 2~ ions incharge-transfer homogeneous solution [12]. with The MV one (nap-NH~C forth 2H4NH~)did not show 2 + any in evidenceprobably of complex with MV water, due formation to the strong repulsive
139
-
3.60-
The flash photolysis and the picosecond absorption spectroscopy set-up were described in detail elsewhere [13,26]. In a typical picosecond absorption spectroscopy experiment 300 spectra were accumulated for each delay time setting. The redox potentials were measured using a RDE-4 potentiostat/ galvanostat (Pine Instruments Co.). We chose a glassy carbon electrode as working electrode, a platinum wire as counter electrode, and the standard calomel electrode as reference. All measurements were carried out in aqueous solution containing 0.2 M Na2SO4 as supporting
transfer reactions
-
-
2 80-
~
2.00-
~20 -
.
-
__________________________________
.400
~
328.
3~6.
I
I
~
4~t.
I
WAVELENGTh
Fig. 1. Absorption aqueous solutions containing 2~and (1) 0spectra M, (2)of0.061 M, (3) 0.122 M, (4) 0.1830.01 M M MV and (5) 0.244 M napSO 3
Coulombic forces. However, all four compounds formed complexes in SDS micellar environment. As shown for napSO3 in fig. 1, the sulfonic acid derivatives as well as the unsubstituted naphthalene showed charge transfer absorption2~/napbands (shoulders) around 400 nm. The MV NH~C 2H4NH~ charge transfer absorption band peaks at 485 nm. The exact position of the shoulders were determined by calculating the first derivative of the absorption spectra and we obtamed 356, 378 and 380 nm as shoulder maxima for nap(5O3 ) 2’ napSO3 and nap, respectively. For the three complexes formed in homogeneous solution we determined the complex formation equilibrium constants applying the Benesi Hildebrand method [15] (see fig. 2) and the data are given in table 1. 3.2. Redox potentials
In order to understand the charge transfer absorption spectra as well as the electron transfer rate constants we need to know the one-electron redox potentials of electron donor andvoltametry acceptor molecules involved. Using the cyclic method we measured potentials (peak potentials)one-electron between 0.78oxidation and 1.79 V versus SCE, going from nap-NH~C 2H4NH~to
140
S.M. Hubig
/ Effect of micellar media on
electron transfer reactions
tion and reduction potentials we calculated t~G° values between 1.47 and —2.48 eV.
0.07 0.06 0.05 ~.
0.04
3.3. Laser flash photolysis results
0.03
The transient absorption spectra of all four methylviologen/ naphthalene complexes, recorded on a microsecond timescale upon excitation at 355 nm, showed the reduced methylviologen species (MV~)absorbing at 395 and 605 nm [17] and the oxidized naphthalene species absorbing at 685 nm
0.02
0.01 0.00 0
100
I 200
I
300
[napS0~]
I 400
500
1
2~/napSO Fig. 2. Benesi Hildebrand plot for the MV 3 com2~] 5.5 mM, [napSO plex formation [MV 3 ] = 2.1—10.5 mM.
nap(S03 )2 (see table 1). The unexpected small difference between napSO3 and nap can be explained by the solvent change from water to acetonitrile. As one-electron redox potential for methylviologen we took a value of —0.69 V versus SCE from the literature [27]. Using these oxida-
[28] (~see fig. 3). These are the spectra of the escaped solvated radical ions which recombine in a second order reaction on a ~ts timescale. To determine the initial amount of solvated radicals produced ([MV~]0 in eq. (2)), we extrapolated the second order decay traces at 605 nm to the center of the laser pulse (A(605)0 in eq. (2)). The number of photons absorbed (‘ab, in eq. (2)) was obtained by measuring the laser intensity with the benzophenone actinometer [29] and multiplying this
Table I Oxidation potentials, complex formation constants, ion pair lifetimes, escape yields and electron transfer rate constants for charge transfer complexes between methylviologen and four naphthalene derivatives: 1,5-naphthalene disulfonic acid disodium salt (I), 2-naphthalene sulfonic acid sodium salt (II), naphthalene (III) and N-(1-naphthyl)ethylenediamine dihydrochloride (IV). Derivative
I
II
III
IV
Ground state charges Ion pair charges Oxidation potentials [VI a) Absorption maximum of the CT complex [nm] Complex formation constants in water
2 + /2 +/ 1.79
2+ / + /0 1.50
2 +/0 +/ + 1.48 ~
2 + /2 + + /3 + 0.78
356
378
380 ~
54
19
2
‘~
485 0
Homogeneous solution ~es~
r [ps] 95 ~ k~,~ kre~X10 x108 s
0.045 2125.0
0.004 45 22
2.3
0.88
0.057
0.031
0.020 27 36 5.6
h)
0.016 ~ 35 28 5.6
d)
SDS micellar dispersion [psI 189 142 kr~, X109 s 4.4 6.8 kes~ X 108 s 2.7 2.2 a) vs. SCE in water containing 0.2 M Na 2SO4. 2 ± chloride. acetonitrile using MV2~hexafluorophosphate. b) pure water/acetomtrile mixture (75 : 25) using MV d) charge transfer complex not formed in water (see text). e) within the time resolution of the experiment. r
0.004 58
<0.001 <25 e)
17
>40 e)
0.68
<0.4
S. M. Hubig
/ Effect of rn/cellar media on electron
value by (1 10 A(355)) A(355) being the absorbance at the irradiation wavelength 355 nm. The benzophenone and the solvated radical measurements were carried out under exactly the same flash photolysis conditions, i.e. the same laser intensity and same absorbance of the samples at 355 nm. By using eq. (2), we calculated escape yields taking (605) 14 000 M cm as extinclion coefficient for the methylviologen radical —
transfer reactions
141
-
.25~ -
0
i3~
~ ci) ~
I
07E-
MV~ [30]: [MV~”]0 ~esc J
.
.0±E
—
1
A(605)0
(2)
III
.
490.
.430.
(605) I~(1—10
—
I 550.
I
I 6±0.
I
I
670.
.4(355))
WAVELENGTh
The data for homogeneous solutions and SDS micellar dispersions are given in table 1. The time-resolved spectra obtained on a ps timescale (fig. 4) showed the same absorption bands around 395 and 600 nm, attributed to the MV~ radical cation, and at 685 nm (oxidized naphthalene species). The kinetic traces for nap, napSO3 , and nap(S03 ) 2 observed at 605 nm (fig. 5 (b d)) showed a first order decay on a picosecond timescale to a raised baseline, indicating some long-lived transient species which correspond to the species observed on the ~istimescale (see fig.
______________________________________
Fig. 4. Picosecond time-resolved spectra upon photoexcitation 2~/nap(SO of theMV 3 )2 complex at 355 nm at different delay times: 1
30 ps, 2
—
70 ps, 3
and 6
120 ps, 4 500 Ps.
170 ps, 5
250 Ps,
3). In the case of the nap-NH~C2H4NH~no long-lived species was observed on a ~istimescale, neither did we observe a raised baseline in the picosecond experiment (fig. 5(a)). Taking the observed first order decay rate constants kOb, = ‘r = k rec + kesc and the escape yields ~esc determined as described above, krec and k~,~ could be calculated applying eq. (1). The rate constants measured in a homogeneous solution and SDS micellar dispersion are presented in table 1.
0.30’
1.
~ ~
0.50
__________________________________ -0.10. 3.70
4.70
5.70
6.70
7.70
______
ES
).
ti~,]
Fig. 3. Microsecond time-resolved spectra upon photoexcitation of the MV2 k/nap complex at 355 nm at different delay times: A = 0.4 as, B —1.3 ‘as, C = 3.2 as, 0 = 6.4 as and E=13.5 ~s.
Comparing the oxidation potentials of the naphthalene compounds and the charge transfer 4. Discussion absorption bands of the corresponding charge 2~ions, we note that transfer with increasing complexes oxidation with MV potential the maximum of the charge transfer absorption band decreases. This can be explained by the Mulliken theory [31] which suggests direct excitation in an ion pair state. According to this theory, the energy gap between ground-state and ion pair state is approximately the difference between the ionization energy of the electron donor and the electron affinity of the electron acceptor corrected by an
S. M. Hubig
142 XE—S
/ Effect of micellar media on
‘a’
-
electron transfer reactions
XEO (b)
3.tC
N
2.31’
~
~ IN
1
“I
5.51
Ui
-
N
.701
N
\~
-
\N
-
N
-.
N N
—.506
N
N
N
-
N
-±d.0 I
I
‘-...........N N ~
7d.0
±~.
TINE (PIC0—6ECs)
XEO
-
7~.O
-
XEO
~.
~.
TINE (PIC0—SEC~Sl
XEO
XEO
‘
(c)
(d)
.406
.356 N N N N
.506
.236
N
-.
Ui
N
N
Ui
.206 N
-
N
.076
N N
.
N
--N
N
N
—.001 ~
.0L6 ~i
I
2~0.
~
I
5)0.
‘
.0130
3013.
6013.
9013.
i2~20
TINE (PICO—SECS) XEO 2~/nap-NH~ C 2~/nap (c) theradical MV2~/napSO 2~/nap(SO Fig. 5. Picosecond decay tracescomplex, of the MV~ observed at 605 nm for (a) the MV 2H4NH~’complex, (b) the MV 3 complex, and (d) the MV 3 )2 complex. TINE (PIC0—6E0&)
)tZO
energy term for the Coulombic work of the charge separation. For the interpretation of the flash photolysis data we need to emphasize that the two competing reactions measured, the charge recombination and
the ion escape reaction (see scheme 1), are of very different nature. The former reaction is an electron transfer reaction, the latter a solvation reaction, which causes different driving forces to be predominant. The major driving force in electron
S. M. Hubig
/ Effect of mice//ar media on electron
transfer reactions is usually the ~G° change, but solvent properties play also an important role by affecting the reorganization energy A [32 34]. The solvation reaction is clearly controlled by the electrostatic properties of the solvent or the interfacial region, but also by the Coulombic forces between the ion pair partners and, in the case of the micellar milieu, between the radical ions and the charged interfacial region. The data summarized in table 1 led to the following conclusions: In homogeneous solution complex III in a H20/CH3CN (75 : 25) mixture and complex II in 100% H20 showed very similar recombination rate constants. However, it is difficult to compare the two systems because of the different dielectric constants of the solvents (~(H2O]= 80, c[CH3CN] = 39). Nevertheless, we measured similar oxidation potentials for the naphthalene components in the corresponding solvents which leads to the conclusion that the reorganization energies did not change much going from the water/ acetonitrile (75 : 25) mixture to pure water. The escape yields and rate constants for the same two complexes were quite different. This can be explained assuming strong effects of the solvent on these solvation reactions. On the other hand, the escape rate constants of complex III did not change going from a water/ acetonitrile mixture to pure acetonitrile. But, for charged naphthalene components the solvent effects could be quite different from those for uncharged naphthalene species. In addition, Coulombic forces within the ion pairs may affect the escape rate constants: the [MV~~ nap~] ion pair suffers from stronger repulsive forces within the ion pair than the [MV ± napSO3 ] ion pair, i.e. the nap~~ radical cation may escape more easily than the ~napSO3 radical. Interestingly, the escape rate constant measured for complex III (kesc = 5.6 x 108 s 1) is in very good agreement with the value estimated by Gould et al. [9] as an averaged escape rate constant for all ion pairs listed in ref. [9]. However, the ion pairs described in ref. [9] were produced by encounters between an excited species and a quencher molecule. According to Mataga et al. [5] the structure of such ion pairs is quite different from the structure of ion pairs formed by photoexcitation of ground-state charge. . .
...
transfer reactions
143
transfer complexes for which an escape rate constant of about 2 X s 1 [5] was determined. In any case, we would like to emphasize that the ion escape reactions for complexes as described in this paper, occur in two steps, going from the “cornpact” ion pair via the “loose” ion pair to “free”, solvated radical ions. These two steps may be quite differently affected by the solvent for each complex listed in table 1. Comparing complex I and II, both in water, we notice a lowering of the charge recombination rate constant along with an increase of the redox potential difference between donor and acceptor molecule. This may indicate that we are measuring within the “Marcus-inverted” region [32 34] which was expected for electron transfer reactions with free energy changes, z~G°, between —1.47 and 2.48 eV. The differences in the escape yields and rate constants are difficult to understand at this stage of the investigation. Obviously, neither solvent effects nor effects of Coulombic forces alone can explain these results, and we need to know more about the differences in structure between the two complexes. In SDS solution the data obtained were in a better systematic order. The decrease in krec Y~~th increasing values, going from nap-NH~C2H4NH~ to nap(SO3 ) 2’ again demonstrate the Marcus-inverted region. There is also a clear trend for the escape yields and rate constants: both values systematically increased going from positively charged to negatively charged naphthalene compounds. A possible explanation is that the repulsive Coulombic forces between the negatively charged naphthalene species and the negatively charged SDS micelle surface enhanced the escape reaction. Comparing the data in homogeneous and heterogeneous environment, we found the following interesting results: 2 ~/nap(SO (1) The MV 3 )2 complex shows very similar behavior in homogeneous and heterogeneous environment. This leads to the question where the complex is located in2 +SDS solutions. moiety tends Since the positively charged MV charged micelle to associate with the negatively [17,35], while the negatively charged naphthalene sulfonic acid moiety is repelled, we assume that the complex is only slightly attached to the micelle. —
144
S.M. Hubig
/ Effect of mice//ar media on electron transfer reactions
This assumption is supported by the shifted charge transfer absorption band of the complex and slightly higher escape yields and rate constants going from aqueous to micellar solution. (2) Comparing complex II and III in homogeneous and micellar solution, we first notice that the charge recombination rate constants are lower in SDS micellar dispersion than in a homogeneous solution. Although a final explanation of this interesting result cannot be given at the present stage of our study, we suggest that the lowering of krec may be caused by the different structure and polarity of the surfactant environment around the ion pairs in SDS solution, as opposed to the water or acetonitrile solvent cages in homogeneous solutions. The ion escape reactions show an inverse trend: whereas complex II shows a higher escape yield in SDS, complex III shows a higher ion escape yield in homogeneous solution. Although a comparison of the two complexes is difficult due to the different solvents used (see above), we suggest that this different behavior of the two complexes may be caused by the Coulombic forces between the negatively charged micelle surface and the differently charged naphthalene species. In summary, we note that the rnicellar environment strongly influences the escape yields of free ions out of the solvent cages around the ion pairs. We could show that this is due to the fact that both the charge recombination rate constants and the escape rate constants change going from homogeneous to micellar environment. Whereas the escape rate constants seemed to be governed by the solvent environment and the Coulombic forces between the negatively charged micelle surface and the differently charged ion pair partners, the micellar effect on the charge recombination rate constants is not clearly understood at this stage of our study. We are currently investigating both micellar effects by systematically varying the naphthalene component in the charge-transfer complexes to be studied. Acknowledgements All laser flash photolysis experiments were carned out at the Center for fast Kinetics Research
which is supported jointly by the Biotechnology Resources Program of the Division of Research Resources of NIH (grant RR 00886) and by The University of Texas at Austin. I wish to thank Dr. Bhaskar G. Maiya for the cyclic voltametry measurements and Dr. Anthony Harriman for many valuable discussions.
References [1] S.L.
Mattes atid S. Fond, in. Organic Photochemistry,
Vol. 6, ed. A. Padwa (Marcel Dekker, New York, 1983) P. 233. [21 MA. Fox(Elsevier, and M. Amsterdam, Chanon, eds., Photoinduced Electron Transfer 1988). [3] N. Mataga, Pure Appl. Chem. 56 (1984) 1255. [41N. Mataga, T. Okada, Y. Kanda and H. Shioyama, Tetrahedron 42 (1986) 6143. [5] N. Mataga, H. Shioyama and Y. Kanda, J. Phys. Chem. 91(1987) 314. [61 N. Mataga, Y. Kanda, T. Asahi, H. Miyasaka, T. Okada and T. Kakitani, Chem. Phys. 127 (1988) 239. [7] H. Miyasaka. S. Ojima and N. Mataga. J. Phys Chem 93 (1989) 3380. [8] T. Asahi and N. Mataga, J. Phys. Chem. 93 (1989) 6575. [9] I.R. Gould, D. (1987) Ege, S.L. Chem. Soc. 109 3794.Mattes and S. Farid, J. Am. [10] J.M. Masnovi, A. Levine and J.K. Kochi, Chem. Phys Lett. 119 (1985) 351. [11] S. Sakkararaman, W.A. Haney and J.K. Kochi, J. Am. Chem. Soc. 109 (1987) 7824. [12] G. Jones II and V. Malba, Chem. Phys. Lett. 119 (1985) 105. [13] S.J. Atherton, S.M. Hubig, T.J. Callan, J.A. Duncanson, PT. Snowden and M.A.J. Rodgers, J. Phys. Chem. 91 (1987) 3137. [14] T.W. Ebbesen, L.E. Manring and K.S. Peters, J. Am. Chem. Soc. 106 (1984) 7400. [15] H.A. Benesi and J.h. Hildebrand, J. Am. Chem. Soc. 71 (1949) 2703. [16] The data in this paper indicate that under our experimental conditions the diammonium salt is formed. [17] SM. Hubig, B.C. Dionne and M.A.J. Rodgers, J. Phys. Chem. 90 (1986) 5873. [18] M. Almgren, F. Grieser and J.K. Thomas, J. Phys. Chem. 83 (1979) 3232. [19] S.M. Hubig and M.A.J. Rodgers, J. Phys. Chem. 94 (1990) 1933. [20] S.M. Hubig and M.A.J. Rodgers, Chem. Phys. Lett. 146 (1988) 539. [21] M.A.J. Rodgers, Radiat. Phys. Chem. 23 (1984) 245. [22] M.A.J. Rodgers and S.M. Hubig, in: Radiation Research, eds. EM. Fielden, J.F. Fowler, J.B. Hendry and D. Scott (Taylor & Francis, London, 1987) P. 102.
S. M. Hubig
/ Effect of micellar media on
[23] M. Hamity and R.H. Lema, Can. J. Chem. 66 (1988) 1552. [24] H. Masuhara, H. Tanabe and N. Mataga, Chem. Phys. Lett. 63 (1979) 273. [25] A.E. Kaifer and A.J. Bard, J. Phys. Chem. 89 (1985) 4876. [26] D.C. Foyt, Comp. Chem. 5 (1981) 49. [27] M. Ito, T.J. Kuwana, J. Electroanal. Chem. 32 (1971) 415. [28] R. Gschwind and E. Haselbach, Helv. Chim. Acta 62 (1979) 941. [29] J.K. Hurley, N. Sinai and B. Linschitz, Photochem. Photobiol. 38 (1983) 9
electron transfer reactions
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[30] T. Watanabe and K. Honda, J. Phys. Chem. 86 (1982) 2617. [31] (a) R.S. Mulliken, J. Am. Chem. Soc. 74 (1952) 811. (b) R.S. Mulliken, J. Phys. Chem. 56 (1952) 801. [32] R.A. Marcus, J. Chem. Phys. 24 (1956) 966. [33] R.A. Marcus, Ann. Rev. Phys. Chem. 15 (1964) 155. [34] R.A. Marcus and N. Sutin, Biochim. Biophys. Acta 811 (1985) 265. [35] M.A.J. Rodgers, D.C. Foyt and Z.A. Zimek, Radiat. Res. 75 (1978) 296.
137
The effect of micellar media on photoinduced electron transfer reactions in methylviologen— naphthalene ground- state charge-transfer complexes Stephan M. Hubig Center for Fast Kinetics Research, The University of Texas at Austin, Austin, TX 78712, US4 Received 17 May 1990 Revised 30 July 1990 Accepted 30 August 1990
We have studied photoinduced electron transfer reactions by photoexcitation of ground-state charge-transfer complexes between methylviologen ions and several differently charged naphthalene derivatives. Using laser flash photolysis and picosecond absorption spectroscopy techniques, we have determined ion pair lifetimes and charge separation efficiencies. One-electron oxidation potentials of the naphthalene compounds have been measured by cyclic voltametry and complex formation equilibrium constants have been obtained from Benesi Hildebrand plots. Since most of the biologically relevant electron transfer reactions occur in or close to membranes, we have been interested in the effects of interfaces on the electron transfer reactions studied. Using sodium dodecyl sulfate micellar dispersions as model interfacial environment we have studied the micellai effects on charge recombination and escape rate constants and on the escape yields.
1. Introduction In recent years many studies on photoinduced electron transfer reactions have been carried out in order to explore nature’s photosynthesis mechanisms and to develop artificial systems for efficient solar energy conversion [1 3]. In order to improve solar energy conversion systems one needs to investigate the electron transfer reactions occurring immediately after photoexcitation. There are in principle two mechanisms of photoinduced charge separation to distinguish: (a) photoexcitation of a potential electron donor or acceptor followed by electron transfer quenching (kq) occurring at an encounter with a quencher; or (b) photoexcitation of a ground-state charge-transfer complex leading directly to a charge-separated (ion pair) state (see scheme 1). According to Mataga et al. [4 8] the two mechanisms also result in two differently structured ion pairs, compact or loose ion pairs. Once the charge separation is completed, there are two possible reaction path0022-2313/91/$03.50 © 1991
ways: charge recombination (krec) within the ion pair or escape (k~5~)of the ions out of the solvent cage around the ion pair, leading to “free” solvated radical ions. The efficiency of charge separation, i.e. the yield of escaped ions (~esc), can be determined according to the following equation: ~esc
=
abs
=
k
kesc + k
.
=
kesc
.
T,
(1)
where [W] represents the molar concentration of either solvated radical ion, D~ or A ‘ab, is the molar concentration of photons absorbed, and T is the ion pair lifetime. According to eq. 1 two experiments are necessary to determine k rec and kesc: (1) The determination of ~esc by measuring both the amount of solvated radicals produced (R~)and the dose of light absorbed (‘abs)~ The radicals can be observed directly, by measuring
Elsevier Science Publishers B.V. (North-Holland)
138
S.M Hubig
(a) D*
A or D + A*
Effect of mice//ar media on electron transfer reactions
transfer absorption band [4 8, 10 14]. In addition to the advantage of fast ion pair formation, ground-state charge-transfer complexes provide another convenience in being simple systems where the ion pair state is the only excited state formed upon irradiation, as opposed to dynamic quenching experiments where the ion pair state is formed from other excited (singlet or triplet) states which complicates the spectroscopical observation of the transients. In this paper we present data on charge transfer complexes between methylviologen ions (MV2~) and several naphthalene derivatives.
+
~kq ~
[D
. .
1
A
II
~
÷
D
+
A
D
A
secondary ET reactions (b) [D..
A]*
[D~ [D.
. .
Al
. .
A
.1
Using laser flash photolysis and picosecond sorption spectroscopy techniques, we have abdetermined escape yields and ion pair lifetimes. Redox potentials were measured by cyclic voltametry, and complex formation constants were obtained from Benesi Hildebrand plots [15]. For electron transfer reactions within ion pairs
D”+A E
Scheme 1. Reaction schemes for mechanisms (a) and (b) (see text) of photo-induced charge separation (D electron donor, A electron acceptor).
their transient absorption spectrum, usually on a microsecond timescale, or indirectly using appropriate trapping reagents [9]. The amount of photons absorbed can be determined using chemical actinometers. (2) The determination of the lifetime T (kesc + kr~) of the initial ion pair [D A 1 by fast kinetic techniques. Charge recombination reactions within ion pairs may occur on very short timescales (T < 1 ns). To measure such short ion pair lifetimes, the initial photoinduced charge separation leading to the ion pair must be completed within an even shorter time period. In the case of charge separation by dynamic electron transfer quenching the charge recombination and the ion escape reaction cannot be followed kinetically because the ion pair formation is too slow compared to the decay, and the only way to obtain charge recombination rate constants from escape yields is by taking an estimated escape rate constant of 5 x 108 ~ 1 [9]. A successful way to produce ion pairs very fast is the photoexcitation of ground-state charge-transfer complexes at a wavelength within the charge—
~
. .
Coulombic forces play an important role as driving force since the distance between the charged species is very small. In order to study such Coulombic effects we chose naphthalene derivatives of different charges including naphthalene (nap), 2-naphthalene sulfonic acid sodium salt (napSO 3 ), 1,5-naphthalene disulfonic acid disodium salt (nap(S03 )2)~ and N-(1-naphthyl) ethylenediamine dihydrochloride (nap-NH~C2H4 NH~)[161. Most of the electron transfer studies reported in the literature were carried out in homogeneous solutions. However, most of the biologically relevant electron transfer reactions occur in or around biomembranes (cell membranes or membranes of organelles) and there is a lack of information about the effect of interfaces on electron transfer reaction parameters [17 22]. For instance, Coulombic forces are expected to play an important role not only within the ion pair, as mentioned above, but also between the charged interface and either charged ion pair partner. There is also little known how interfacial sites influence the formation of charge transfer complexes [23,24] and the redox potentials of electron transfer partners [25]. For this study we chose sodium dodecyl sulfate (SDS) micellar dispersions as model environment and compared the results in homogeneous and heterogeneous media.
S. M. Hubig
/ Effect of micellar media on electron
2. Experimental
electrolyte, except for naphthalene where a 75 : 25 water-acetonitrile mixture was used. The absorption spectra Benesi—Hildebrand plots were obtamed from for a Hitachi U3210 spectrophotometer. Methylviologen dichloride was obtained from Sigma and used as received. For experiments with naphthalene in acetonitrile the hexafluorophosphate salt was used which was made by mixing the dichloride salt with ammonium hexafluorophosphate (Aldrich) in water, yielding a white precipitate. Naphthalene (scintillation grade) and acetonitrile (HPLC grade) were purchased from Aldrich, the naphthalene sulfonic acid derivatives (napSO3 and nap(S03 )2) from Eastman Kodak, N-(1 -naphthyl)ethylenediamine dihydrochloride from FLUKA, and sodium dodecyl sulfate (SDS, “especially pure” grade) from Gallard-Schlesinger (BDH). For all aqueous and micellar solutions Millipore “reagent grade” water was used.
3. Results 3.1. Ground-state charge-transfer complex formation
From the four naphthalene compounds studied in this paper, three including napSO3 and nap (SO3 )2 in water and nap in acetonitrile, form ground-state complexes 2~ ions incharge-transfer homogeneous solution [12]. with The MV one (nap-NH~C forth 2H4NH~)did not show 2 + any in evidenceprobably of complex with MV water, due formation to the strong repulsive
139
-
3.60-
The flash photolysis and the picosecond absorption spectroscopy set-up were described in detail elsewhere [13,26]. In a typical picosecond absorption spectroscopy experiment 300 spectra were accumulated for each delay time setting. The redox potentials were measured using a RDE-4 potentiostat/ galvanostat (Pine Instruments Co.). We chose a glassy carbon electrode as working electrode, a platinum wire as counter electrode, and the standard calomel electrode as reference. All measurements were carried out in aqueous solution containing 0.2 M Na2SO4 as supporting
transfer reactions
-
-
2 80-
~
2.00-
~20 -
.
-
__________________________________
.400
~
328.
3~6.
I
I
~
4~t.
I
WAVELENGTh
Fig. 1. Absorption aqueous solutions containing 2~and (1) 0spectra M, (2)of0.061 M, (3) 0.122 M, (4) 0.1830.01 M M MV and (5) 0.244 M napSO 3
Coulombic forces. However, all four compounds formed complexes in SDS micellar environment. As shown for napSO3 in fig. 1, the sulfonic acid derivatives as well as the unsubstituted naphthalene showed charge transfer absorption2~/napbands (shoulders) around 400 nm. The MV NH~C 2H4NH~ charge transfer absorption band peaks at 485 nm. The exact position of the shoulders were determined by calculating the first derivative of the absorption spectra and we obtamed 356, 378 and 380 nm as shoulder maxima for nap(5O3 ) 2’ napSO3 and nap, respectively. For the three complexes formed in homogeneous solution we determined the complex formation equilibrium constants applying the Benesi Hildebrand method [15] (see fig. 2) and the data are given in table 1. 3.2. Redox potentials
In order to understand the charge transfer absorption spectra as well as the electron transfer rate constants we need to know the one-electron redox potentials of electron donor andvoltametry acceptor molecules involved. Using the cyclic method we measured potentials (peak potentials)one-electron between 0.78oxidation and 1.79 V versus SCE, going from nap-NH~C 2H4NH~to
140
S.M. Hubig
/ Effect of micellar media on
electron transfer reactions
tion and reduction potentials we calculated t~G° values between 1.47 and —2.48 eV.
0.07 0.06 0.05 ~.
0.04
3.3. Laser flash photolysis results
0.03
The transient absorption spectra of all four methylviologen/ naphthalene complexes, recorded on a microsecond timescale upon excitation at 355 nm, showed the reduced methylviologen species (MV~)absorbing at 395 and 605 nm [17] and the oxidized naphthalene species absorbing at 685 nm
0.02
0.01 0.00 0
100
I 200
I
300
[napS0~]
I 400
500
1
2~/napSO Fig. 2. Benesi Hildebrand plot for the MV 3 com2~] 5.5 mM, [napSO plex formation [MV 3 ] = 2.1—10.5 mM.
nap(S03 )2 (see table 1). The unexpected small difference between napSO3 and nap can be explained by the solvent change from water to acetonitrile. As one-electron redox potential for methylviologen we took a value of —0.69 V versus SCE from the literature [27]. Using these oxida-
[28] (~see fig. 3). These are the spectra of the escaped solvated radical ions which recombine in a second order reaction on a ~ts timescale. To determine the initial amount of solvated radicals produced ([MV~]0 in eq. (2)), we extrapolated the second order decay traces at 605 nm to the center of the laser pulse (A(605)0 in eq. (2)). The number of photons absorbed (‘ab, in eq. (2)) was obtained by measuring the laser intensity with the benzophenone actinometer [29] and multiplying this
Table I Oxidation potentials, complex formation constants, ion pair lifetimes, escape yields and electron transfer rate constants for charge transfer complexes between methylviologen and four naphthalene derivatives: 1,5-naphthalene disulfonic acid disodium salt (I), 2-naphthalene sulfonic acid sodium salt (II), naphthalene (III) and N-(1-naphthyl)ethylenediamine dihydrochloride (IV). Derivative
I
II
III
IV
Ground state charges Ion pair charges Oxidation potentials [VI a) Absorption maximum of the CT complex [nm] Complex formation constants in water
2 + /2 +/ 1.79
2+ / + /0 1.50
2 +/0 +/ + 1.48 ~
2 + /2 + + /3 + 0.78
356
378
380 ~
54
19
2
‘~
485 0
Homogeneous solution ~es~
r [ps] 95 ~ k~,~ kre~X10 x108 s
0.045 2125.0
0.004 45 22
2.3
0.88
0.057
0.031
0.020 27 36 5.6
h)
0.016 ~ 35 28 5.6
d)
SDS micellar dispersion [psI 189 142 kr~, X109 s 4.4 6.8 kes~ X 108 s 2.7 2.2 a) vs. SCE in water containing 0.2 M Na 2SO4. 2 ± chloride. acetonitrile using MV2~hexafluorophosphate. b) pure water/acetomtrile mixture (75 : 25) using MV d) charge transfer complex not formed in water (see text). e) within the time resolution of the experiment. r
0.004 58
<0.001 <25 e)
17
>40 e)
0.68
<0.4
S. M. Hubig
/ Effect of rn/cellar media on electron
value by (1 10 A(355)) A(355) being the absorbance at the irradiation wavelength 355 nm. The benzophenone and the solvated radical measurements were carried out under exactly the same flash photolysis conditions, i.e. the same laser intensity and same absorbance of the samples at 355 nm. By using eq. (2), we calculated escape yields taking (605) 14 000 M cm as extinclion coefficient for the methylviologen radical —
transfer reactions
141
-
.25~ -
0
i3~
~ ci) ~
I
07E-
MV~ [30]: [MV~”]0 ~esc J
.
.0±E
—
1
A(605)0
(2)
III
.
490.
.430.
(605) I~(1—10
—
I 550.
I
I 6±0.
I
I
670.
.4(355))
WAVELENGTh
The data for homogeneous solutions and SDS micellar dispersions are given in table 1. The time-resolved spectra obtained on a ps timescale (fig. 4) showed the same absorption bands around 395 and 600 nm, attributed to the MV~ radical cation, and at 685 nm (oxidized naphthalene species). The kinetic traces for nap, napSO3 , and nap(S03 ) 2 observed at 605 nm (fig. 5 (b d)) showed a first order decay on a picosecond timescale to a raised baseline, indicating some long-lived transient species which correspond to the species observed on the ~istimescale (see fig.
______________________________________
Fig. 4. Picosecond time-resolved spectra upon photoexcitation 2~/nap(SO of theMV 3 )2 complex at 355 nm at different delay times: 1
30 ps, 2
—
70 ps, 3
and 6
120 ps, 4 500 Ps.
170 ps, 5
250 Ps,
3). In the case of the nap-NH~C2H4NH~no long-lived species was observed on a ~istimescale, neither did we observe a raised baseline in the picosecond experiment (fig. 5(a)). Taking the observed first order decay rate constants kOb, = ‘r = k rec + kesc and the escape yields ~esc determined as described above, krec and k~,~ could be calculated applying eq. (1). The rate constants measured in a homogeneous solution and SDS micellar dispersion are presented in table 1.
0.30’
1.
~ ~
0.50
__________________________________ -0.10. 3.70
4.70
5.70
6.70
7.70
______
ES
).
ti~,]
Fig. 3. Microsecond time-resolved spectra upon photoexcitation of the MV2 k/nap complex at 355 nm at different delay times: A = 0.4 as, B —1.3 ‘as, C = 3.2 as, 0 = 6.4 as and E=13.5 ~s.
Comparing the oxidation potentials of the naphthalene compounds and the charge transfer 4. Discussion absorption bands of the corresponding charge 2~ions, we note that transfer with increasing complexes oxidation with MV potential the maximum of the charge transfer absorption band decreases. This can be explained by the Mulliken theory [31] which suggests direct excitation in an ion pair state. According to this theory, the energy gap between ground-state and ion pair state is approximately the difference between the ionization energy of the electron donor and the electron affinity of the electron acceptor corrected by an
S. M. Hubig
142 XE—S
/ Effect of micellar media on
‘a’
-
electron transfer reactions
XEO (b)
3.tC
N
2.31’
~
~ IN
1
“I
5.51
Ui
-
N
.701
N
\~
-
\N
-
N
-.
N N
—.506
N
N
N
-
N
-±d.0 I
I
‘-...........N N ~
7d.0
±~.
TINE (PIC0—6ECs)
XEO
-
7~.O
-
XEO
~.
~.
TINE (PIC0—SEC~Sl
XEO
XEO
‘
(c)
(d)
.406
.356 N N N N
.506
.236
N
-.
Ui
N
N
Ui
.206 N
-
N
.076
N N
.
N
--N
N
N
—.001 ~
.0L6 ~i
I
2~0.
~
I
5)0.
‘
.0130
3013.
6013.
9013.
i2~20
TINE (PICO—SECS) XEO 2~/nap-NH~ C 2~/nap (c) theradical MV2~/napSO 2~/nap(SO Fig. 5. Picosecond decay tracescomplex, of the MV~ observed at 605 nm for (a) the MV 2H4NH~’complex, (b) the MV 3 complex, and (d) the MV 3 )2 complex. TINE (PIC0—6E0&)
)tZO
energy term for the Coulombic work of the charge separation. For the interpretation of the flash photolysis data we need to emphasize that the two competing reactions measured, the charge recombination and
the ion escape reaction (see scheme 1), are of very different nature. The former reaction is an electron transfer reaction, the latter a solvation reaction, which causes different driving forces to be predominant. The major driving force in electron
S. M. Hubig
/ Effect of mice//ar media on electron
transfer reactions is usually the ~G° change, but solvent properties play also an important role by affecting the reorganization energy A [32 34]. The solvation reaction is clearly controlled by the electrostatic properties of the solvent or the interfacial region, but also by the Coulombic forces between the ion pair partners and, in the case of the micellar milieu, between the radical ions and the charged interfacial region. The data summarized in table 1 led to the following conclusions: In homogeneous solution complex III in a H20/CH3CN (75 : 25) mixture and complex II in 100% H20 showed very similar recombination rate constants. However, it is difficult to compare the two systems because of the different dielectric constants of the solvents (~(H2O]= 80, c[CH3CN] = 39). Nevertheless, we measured similar oxidation potentials for the naphthalene components in the corresponding solvents which leads to the conclusion that the reorganization energies did not change much going from the water/ acetonitrile (75 : 25) mixture to pure water. The escape yields and rate constants for the same two complexes were quite different. This can be explained assuming strong effects of the solvent on these solvation reactions. On the other hand, the escape rate constants of complex III did not change going from a water/ acetonitrile mixture to pure acetonitrile. But, for charged naphthalene components the solvent effects could be quite different from those for uncharged naphthalene species. In addition, Coulombic forces within the ion pairs may affect the escape rate constants: the [MV~~ nap~] ion pair suffers from stronger repulsive forces within the ion pair than the [MV ± napSO3 ] ion pair, i.e. the nap~~ radical cation may escape more easily than the ~napSO3 radical. Interestingly, the escape rate constant measured for complex III (kesc = 5.6 x 108 s 1) is in very good agreement with the value estimated by Gould et al. [9] as an averaged escape rate constant for all ion pairs listed in ref. [9]. However, the ion pairs described in ref. [9] were produced by encounters between an excited species and a quencher molecule. According to Mataga et al. [5] the structure of such ion pairs is quite different from the structure of ion pairs formed by photoexcitation of ground-state charge. . .
...
transfer reactions
143
transfer complexes for which an escape rate constant of about 2 X s 1 [5] was determined. In any case, we would like to emphasize that the ion escape reactions for complexes as described in this paper, occur in two steps, going from the “cornpact” ion pair via the “loose” ion pair to “free”, solvated radical ions. These two steps may be quite differently affected by the solvent for each complex listed in table 1. Comparing complex I and II, both in water, we notice a lowering of the charge recombination rate constant along with an increase of the redox potential difference between donor and acceptor molecule. This may indicate that we are measuring within the “Marcus-inverted” region [32 34] which was expected for electron transfer reactions with free energy changes, z~G°, between —1.47 and 2.48 eV. The differences in the escape yields and rate constants are difficult to understand at this stage of the investigation. Obviously, neither solvent effects nor effects of Coulombic forces alone can explain these results, and we need to know more about the differences in structure between the two complexes. In SDS solution the data obtained were in a better systematic order. The decrease in krec Y~~th increasing values, going from nap-NH~C2H4NH~ to nap(SO3 ) 2’ again demonstrate the Marcus-inverted region. There is also a clear trend for the escape yields and rate constants: both values systematically increased going from positively charged to negatively charged naphthalene compounds. A possible explanation is that the repulsive Coulombic forces between the negatively charged naphthalene species and the negatively charged SDS micelle surface enhanced the escape reaction. Comparing the data in homogeneous and heterogeneous environment, we found the following interesting results: 2 ~/nap(SO (1) The MV 3 )2 complex shows very similar behavior in homogeneous and heterogeneous environment. This leads to the question where the complex is located in2 +SDS solutions. moiety tends Since the positively charged MV charged micelle to associate with the negatively [17,35], while the negatively charged naphthalene sulfonic acid moiety is repelled, we assume that the complex is only slightly attached to the micelle. —
144
S.M. Hubig
/ Effect of mice//ar media on electron transfer reactions
This assumption is supported by the shifted charge transfer absorption band of the complex and slightly higher escape yields and rate constants going from aqueous to micellar solution. (2) Comparing complex II and III in homogeneous and micellar solution, we first notice that the charge recombination rate constants are lower in SDS micellar dispersion than in a homogeneous solution. Although a final explanation of this interesting result cannot be given at the present stage of our study, we suggest that the lowering of krec may be caused by the different structure and polarity of the surfactant environment around the ion pairs in SDS solution, as opposed to the water or acetonitrile solvent cages in homogeneous solutions. The ion escape reactions show an inverse trend: whereas complex II shows a higher escape yield in SDS, complex III shows a higher ion escape yield in homogeneous solution. Although a comparison of the two complexes is difficult due to the different solvents used (see above), we suggest that this different behavior of the two complexes may be caused by the Coulombic forces between the negatively charged micelle surface and the differently charged naphthalene species. In summary, we note that the rnicellar environment strongly influences the escape yields of free ions out of the solvent cages around the ion pairs. We could show that this is due to the fact that both the charge recombination rate constants and the escape rate constants change going from homogeneous to micellar environment. Whereas the escape rate constants seemed to be governed by the solvent environment and the Coulombic forces between the negatively charged micelle surface and the differently charged ion pair partners, the micellar effect on the charge recombination rate constants is not clearly understood at this stage of our study. We are currently investigating both micellar effects by systematically varying the naphthalene component in the charge-transfer complexes to be studied. Acknowledgements All laser flash photolysis experiments were carned out at the Center for fast Kinetics Research
which is supported jointly by the Biotechnology Resources Program of the Division of Research Resources of NIH (grant RR 00886) and by The University of Texas at Austin. I wish to thank Dr. Bhaskar G. Maiya for the cyclic voltametry measurements and Dr. Anthony Harriman for many valuable discussions.
References [1] S.L.
Mattes atid S. Fond, in. Organic Photochemistry,
Vol. 6, ed. A. Padwa (Marcel Dekker, New York, 1983) P. 233. [21 MA. Fox(Elsevier, and M. Amsterdam, Chanon, eds., Photoinduced Electron Transfer 1988). [3] N. Mataga, Pure Appl. Chem. 56 (1984) 1255. [41N. Mataga, T. Okada, Y. Kanda and H. Shioyama, Tetrahedron 42 (1986) 6143. [5] N. Mataga, H. Shioyama and Y. Kanda, J. Phys. Chem. 91(1987) 314. [61 N. Mataga, Y. Kanda, T. Asahi, H. Miyasaka, T. Okada and T. Kakitani, Chem. Phys. 127 (1988) 239. [7] H. Miyasaka. S. Ojima and N. Mataga. J. Phys Chem 93 (1989) 3380. [8] T. Asahi and N. Mataga, J. Phys. Chem. 93 (1989) 6575. [9] I.R. Gould, D. (1987) Ege, S.L. Chem. Soc. 109 3794.Mattes and S. Farid, J. Am. [10] J.M. Masnovi, A. Levine and J.K. Kochi, Chem. Phys Lett. 119 (1985) 351. [11] S. Sakkararaman, W.A. Haney and J.K. Kochi, J. Am. Chem. Soc. 109 (1987) 7824. [12] G. Jones II and V. Malba, Chem. Phys. Lett. 119 (1985) 105. [13] S.J. Atherton, S.M. Hubig, T.J. Callan, J.A. Duncanson, PT. Snowden and M.A.J. Rodgers, J. Phys. Chem. 91 (1987) 3137. [14] T.W. Ebbesen, L.E. Manring and K.S. Peters, J. Am. Chem. Soc. 106 (1984) 7400. [15] H.A. Benesi and J.h. Hildebrand, J. Am. Chem. Soc. 71 (1949) 2703. [16] The data in this paper indicate that under our experimental conditions the diammonium salt is formed. [17] SM. Hubig, B.C. Dionne and M.A.J. Rodgers, J. Phys. Chem. 90 (1986) 5873. [18] M. Almgren, F. Grieser and J.K. Thomas, J. Phys. Chem. 83 (1979) 3232. [19] S.M. Hubig and M.A.J. Rodgers, J. Phys. Chem. 94 (1990) 1933. [20] S.M. Hubig and M.A.J. Rodgers, Chem. Phys. Lett. 146 (1988) 539. [21] M.A.J. Rodgers, Radiat. Phys. Chem. 23 (1984) 245. [22] M.A.J. Rodgers and S.M. Hubig, in: Radiation Research, eds. EM. Fielden, J.F. Fowler, J.B. Hendry and D. Scott (Taylor & Francis, London, 1987) P. 102.
S. M. Hubig
/ Effect of micellar media on
[23] M. Hamity and R.H. Lema, Can. J. Chem. 66 (1988) 1552. [24] H. Masuhara, H. Tanabe and N. Mataga, Chem. Phys. Lett. 63 (1979) 273. [25] A.E. Kaifer and A.J. Bard, J. Phys. Chem. 89 (1985) 4876. [26] D.C. Foyt, Comp. Chem. 5 (1981) 49. [27] M. Ito, T.J. Kuwana, J. Electroanal. Chem. 32 (1971) 415. [28] R. Gschwind and E. Haselbach, Helv. Chim. Acta 62 (1979) 941. [29] J.K. Hurley, N. Sinai and B. Linschitz, Photochem. Photobiol. 38 (1983) 9
electron transfer reactions
145
[30] T. Watanabe and K. Honda, J. Phys. Chem. 86 (1982) 2617. [31] (a) R.S. Mulliken, J. Am. Chem. Soc. 74 (1952) 811. (b) R.S. Mulliken, J. Phys. Chem. 56 (1952) 801. [32] R.A. Marcus, J. Chem. Phys. 24 (1956) 966. [33] R.A. Marcus, Ann. Rev. Phys. Chem. 15 (1964) 155. [34] R.A. Marcus and N. Sutin, Biochim. Biophys. Acta 811 (1985) 265. [35] M.A.J. Rodgers, D.C. Foyt and Z.A. Zimek, Radiat. Res. 75 (1978) 296.