J. Phys. Chem. 1994, 98, 3993-3998
3993
Effects of Solvent and Ionic Strength on the Electrochemiluminescence of 9,lO-Diphenylanthracene Karolyn M. Maness, Joan E. Bartelt, and R. Mark Wightman' Department of Chemistry, University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599-3290 Received: April 26, 1993; In Final Form: December 28, 1993"
The steady-state electrochemiluminescence (ECL) from the annihilation reaction of the radical anion and cation of 9,lO-diphenylanthracene (DPA) was examined a t a double-band microelectrode in acetonitrile, propylene carbonate, 1,2-dimethoxyethane, acetonitrile-toluene, and propylene carbonate-toluene mixtures, all containing 0.1 M tetrabutylammonium hexafluorophosphate. In each solvent, the ECL efficiency was found to be greatest a t large DPA concentration due to the effect of competing side reactions. The E C L efficiencies extrapolated to infinite DPA concentrations were found to be solution dependent and were correlated to the separation of the half-wave potentials for the generation of the radical anion and cation. In acetonitrile-toluene mixtures, the ECL efficiency was also found to increase as the ionic strength of the solution was decreased. In 50:50 acetonitrile-toluene the ECL efficiency increased from 6.3 to 8.8% as the concentration of supporting electrolyte was lowered from 100 to 1 mM. This increase was found to correlate with changes in the rate constant for dissociation of the reaction encounter complex as the ionic strength of the solution was lowered.
Introduction Electrochemiluminescence (ECL) is the generation of light as a result of electron-transfer reactions at electrode surfaces.' One route to ECL is via an ion-annihilation reaction mechanism: A'
+e
-
A"'
cathode
A'
+X +X
-
-
B
(4)
C
(5)
The efficiency of emission is dictated by the relative kinetics of the associated reactions and can be expressed as
k:; [DPA]
9ccl = 9f
(6)
k z [DPA] + k,[X]
+ k,[X]
i- 1
where
$f
is the efficiency of photon emission from the excited
* To whom correspondence should be addressed. @
(7)
(1)
where Z is the initial charge of the electrode reactant, kd and k~ are the diffusion-controlled rate constants for formation and dissociation of the encounter complex, and kietis the rate constant for the electron exchange reaction. For klet this reaction leads to formation of the light-emitting species where Az* is the excited singlet. The other possible electron-exchange reactions lead to nonradiative regeneration of the starting material through either the formation of nonemitting excited states or direct formation of two ground-state molecules. Side reactions of the electrogenerated reactants with impurities in the solution can also take place as shown in reactions 4 and 5 with rate constants of kl and k2, respectively:
A"'
state and k,,"are apparent electron-transfer rate constants given by
Abstract published in Advance ACS Abstracts, March 15, 1994.
0022-3654/94/2098-3993$04.50/0
Because the reaction scheme involves electron transfer, ECL efficiencies have been successfully predicted for some chemical systems with the use of classical Marcus t h e ~ r y . ~Indeed, J the observation of ECL has been cited as confirmatory evidence for the existence of an inverted kinetic region for electron-transfer reactions. For systems which involve reactions whose rates lie far into the inverted region, however, these predictions have not been su~cessful.3~~ Nevertheless, the concept of competing kinetic pathways (eq 6) provides a framework for evaluating factors which could affect ~ E C L . The investigations reported here concern the ECL of 9,lOdiphenylanthracene (DPA), which follows the reaction scheme described above.1,4,8-12 Energy supplied by the radical-ion annihilation reaction is sufficient to form the first excited singlet state with an emission maximum at 425 nm and & = 1. I 3 The absence of a magnetic field effect on the ECL for DPA is cited as evidence that excited singlets are not formed via a route involving triplets.9 However, competitive formation of triplets via the annihilation scheme leads to a measured ECL efficiency of only a few percent.4.5JoJl We report here two successful attempts to modify experimentally the relative kinetics of the ECL reactions. First, altering the solvent alters the energetics of the electron-exchange reaction and also the diffusional formation and dissociationof theencounter complex. Second, alteration of the ionic strength of the solution, which accelerates the rate of bimolecular reactions involving oppositely charged species,lk18 is shown to increase the observed ECL. Both of these investigations employ double-band microelectrodes19 which allow electrochemical investigations in unconventional media, including ECL under steady-state conditions.20
Experimental Section Chemicals. Acetonitrile (ACN, UV grade, Burdick and Jackson, Muskegan, MI), 1,Zdimethoxyethane (DME) (Aldrich, 0 1994 American Chemical Society
3994 The Journal of Physical Chemistry, Vol. 98,No. IS, 1994
Milwaukee, WI), toluene (TOL), and propylene carbonate (PC) (both Burdickand Jackson) weresubjected to three to five f r e e z e pumpthaw cycles for removal of dissolved oxygen and passed through a column of activated alumina to remove residual water prior to solution preparation. 9,lO-Diphenylanthracene (DPA; Aldrich, Milwaukee, WI), was recrystallized twice from absolute ethanol. Tetrabutylammonium hexafluorophosphate (TBAH) was purchased from Aldrich and recrystallized twice from 95% ethanol. Solid chemicals were dried under reduced pressure a t 60 OC and stored in a desiccator before use. All experiments were performed in a nitrogen-containing glovebox. Electrochemistry. A platinum double-band electrode was constructed as described previously.20.2' The dimensions of the electrodes used for all experiments as measured from the current magnitudes during steady-state collector generator experiments2' as well as by examination with an optical microscope were as follows: bandwidth, 5-8 pm; gap width, 4.5-5.3 pm; length, 0.30.33 cm. The electrode was polished with a 0.25-pm diamond paste and rinsed with water and methanol before use. Electrodes were affixed to glass standard-tapered joints (male), which allowed reproducible placement in the electrochemical cell. The cell was made from a glassjoint (female) in which a platinum wire auxiliary electrode and a silver wire quasireference electrode were sealed. A quartz window was sealed on the end of the cell approximately 1 mm from the end of the electrode. Diffusion coefficients (D) and the potential difference between the half-wave potentials of the oxidation and reduction reactions (AE1/2)were determined using a 5-pm radius Pt disk electrode (50 mV s-I). The half-wave potentials were obtained from the maxima of the differentiated cyclicvoltammograms. For solutions containing less than 10 mM supporting electrolyte, h E 1 / 2 determinations were made while a constant 10:1 ratio of TBAHDPA was maintained. Diffusion coefficients were determined from the limiting currents of the steady-state cyclic voltammograms. The reference Dvalue was 2.2 X cm2s-I for ferrocene in acetonitrile with 0.2 M TBAH.22 A bipotentiostat of conventional design was used for the experiments with the microband arrays. Current and potential measurements were recorded on an X-Y recorder or digital oscilloscope (Nicolet 3 10). Photometric Equipment. Electrochemiluminescence was measured using a Hamamatsu R928 photomultiplier tube (Bridgewater, N J ) operated at -600 V. The electrochemical cell could be reproducibly positioned in front of the PMT which was mounted on the side of a light-tight steel box. The light-tight box and PMT were contained inside a glovebox, and all electrical connections were made through standard BNC and MHV connections in the wall of the drybox. The photomultiplier current was measured using a Keithley Model 427 current amplifier (Cleveland, OH) and recorded on a strip chart recorder or digital oscilloscope. The relative Coulometric efficiency for ECL was taken as the steady-state photomultiplier current, I , divided by the steadystate anode current. In the case of transient signals, the ratio was determined at the time of maximum intensity. To convert to an estimated ECL efficiency, &I, the data were compared to that measured for Ru(bpy),+2 in acetonitrile (emission maximum a t 610 nm). The experimental efficiency of this system is -0.05 at 25 OC in acetonitrile,23a value that is often taken as a standard.2 Intensity measurements were corrected for differences in PMT sensitivity at the wavelength of maximum emission with the use of the manufacturer's specifications. For comparisons between solvents, corrections were made for differences in refractive index.24 These corrections impart an error of approximately 10% to the conversion of measured Coulometric efficiencies into reported &I values. Spectral Measurements. Fluorescence spectra of DPA were obtained with a SPEX Fluorolog 2 spectrofluorimeter. Con-
Maness et al.
'ECL
0.050
-
0.025
-
A
'&L 0
0.000
-
50
0
0
5
[ D P A I (mW
10
0.5
l/[DPA]
1.0
1.5
(mM)-'
Figure 1. (A) Variation of &EL as a function of DPA concentration. (B) Double reciprocal plots of ECL data obtained in solutions of different DPA concentrations containing 0.1 M TBAH. (0) 5050 PC-TOL without water removed; (A) 5050 PC-TOL, dried solution.
centrations of DPA ranged from 0.5to 1.5 pM, where fluorescence intensity was linear with concentration. Solutions were prepared in a glovebox to remove oxygen, which quenches DPA fluorescence. Emission spectra were recorded from 400 to 500 nm with an excitation wavelength of 374 nm. Intensity measurements were taken a t the maximum DPA fluorescence intensity, 425 nm. Spectra during ECL were obtained by positioning the electrochemical cell in front of a monochromator (2-nm bandpass, Model H-20, JY Optical Systems, Metuchen, N J ) which was connected to the above PMT.
Results ECL of DPA in 5050 Propylene CarbonateToulene. To generate ECL, the applied potential for each band was chosen to be 250 mV past the El/2 value for the redox couple of interest. This was to ensure that the radical ions are generated at a rate determined by diffusion-controlled mass transport. When both electrodes were stepped to the electrolysis potentials simultaneously, the electrolytic current rapidly reached a steady-state value, and equal currents were observed at each band.25 The signal from the PMT reached a maximum but showed a slow decay with time. For systems with chemically stable intermediates, such as ECL generated from Ru(bpy)j2+,the ECLcurrent remains a t a steady-state value.20 The ECL efficiency was found to increase with DPA concentration in a 5050 propylene carbonate-toluene solvent system (Figure 1A). Examination of eq 6 reveals that this can occur when the electrogenerated radicals undergo reactions other than the desired annihilation reactions.26 A double-reciprocal plot of &CL versus DPA concentration should yield a straight line. The experimental data follow this behavior, and the reciprocal of the y-intercept of this plot gives the relative ECL intensity at infinite DPA concentration (Figure lB), where side reactions are negligible. These plots for both dried and undried 5050propylene carbonatetoluene solutions yield lines of different slopes but similar intercepts, illustrating thesuccess of this approach. Water can react with both the radical cation and anion of DPA; however, a t infinite concentration these competing reactions become negligible. The effect of impurities on the radical ions can also be ascertained from the collector-generator voltammograms for the cathodic and anodic reactions.2' For example, the collection efficiencies were 0.49 and 0.40, respectively, in undried solutions compared to 0.54 in dried solutions. The latter value shows that the radicals have a half-life of greater than 12 s, the upper limit that can be determined with electrodes of the dimensions employed.2' In the undried solutions, the collection efficiencies correspond to a half-life of 1.2 s. ECL of DPA in Other Solvent Systems. The dependence of ECL on DPA concentration was also investigated in propylene carbonate, acetonitrile, acetonitrile-toluene mixtures, and 1,2dimethoxyethane, all containing 0.1 M TBAH. In all cases the ECL efficiency was a function of concentration (Figure 2A). The ECL efficiencies measured from the extrapolated intercepts are
Electrochemiluminescence of 9,lO-Diphenylanthracene
The Journal of Physical Chemistry, Vol. 98, No. 15, 1994 3995
TABLE 2 Electrochemical and Extrapolated ECL Data for DPA in 5050 Acetonitrile-Toluene at Various Electrolyte Concentrations D (cm2s-I X [TBAH] (mM) 10-6) m1/2(V) 4CCl" 100 10 1.o 0.1
1.0
0.5 l/[OPA]
500
400
Wavelength
(mu)-'
Figure 2. (A) Double reciprocal plots of ECL data obtained in solutions of different DPA concentrations containing 0.1 M TBAPF6. (A)50:50 PC-TOL, dried solution; (v)75:25 ACN-TOL; ( 0 )5050 ACN-TOL; (0) ACN; (U) PC; ( 0 )DME. (B) ECL spectra of DPA containing 0.1 M TBAH: (solid line) DME, (dashed line) 50:50 ACN-TOL.
TABLE 1: Summary of ECL and Electrochemical Results for Different Solvent Systems. D (cm2 s-1 X solvent system 10-6) AJ51/2(V) bed* € 0.039 f 0.001 64.4 PC 1.6 3.04 50:50 PC-TOL 3.7 3.13 0.064 f 0.002 0.061 f 0.003 37.5 ACN 13 3.12 75:25 ACN-TOL 11 3.13 0.065 f 0.001 0.064 0.005 5050 ACN-TOL 8.3 3.13 0.090 f 0.007 7.2 DME 13 3.31 I, All solutions contain 0.1 M TBAH. Values of ~ E C Lare from the interceptsof the plots in Figure 2. b Estimatedvaluesbased on comparison of the measured DPA Coulometric efficiency with that for Ru(bpy)3+2, whose & = 0.05 in ACN. Correctionswere made for changesin refractive index in the different solvents and for differences in PMT response at L is that different wavelengths. The error shown in the ~ W values associated with the y-intercept from the linear regression obtained for plots of 1/ q 5 ~ cvs~ 1/ [DPA].
0.0
c
-0.5
-3.0 I
A/IT-
-3.5
?
I I
I
Figure 3. Steady-state cyclic voltammograms obtained at 5 pm radius Pt disk. Scan rate = 50 mV/s. Solid curve, 0.94 mM DPA, 100.9 mM TBAH in 5050 ACN-TOL. Dashed curve, 1.12 mM DPA, 102.3 mM TBAH in DME. summarized in Table 1. The reported values of the ECL efficiency, &I, are in the range previously reported (0.0150.078) .10*26-2* In addition, the diffusion coefficients and difference in half-wave potentials as determined from steady-state cyclic voltammograms in the different solutions are also shown. The difference in half-wave potentials for DPA in the different solvent systems was determined from steady-state cyclic voltammograms recorded a t microdisk electrodes (Figure 3). Since both radical anion and radical cation couples are present in a single voltammogram, the radical cation was used as the reference voltage. The cyclic voltammograms in Figure 3 show an increase in A E l / 2 of 0.19 V with no evidence of ohmic drop as the solvent is changed from 5050 ACN-TOL to DME. ECL Spectra. ECL spectra of DPA were measured in solutions of 5050 ACN-TOL and DME (Figure 2B). The spectrum in ACN-TOL is in agreement with published ECL spectra for DPA.Z7 A small (2 nm) red shift in peak maximum occurs as
3.13 f 0.01 3.13 f 0.01 3.15 f 0.01 3.16 f 0.01
0.064 t 0.005 0.073 f 0.003 0.088 f 0.002 0.089 f 0.003
'Estimated values based on comparison of the measured DPA Coulometric efficiency with that for Ru(bpy)s+2,whose &I = 0.05 in ACN. Corrections were made for changes in refractive index in the different solvents and for differences in PMT response at different wavelengths. Each value is extrapolated from at least four different DPA concentrations, and the error shown is that associated with the y-intercept from the linear regression. 45
40 35
*15 O
*
0.5
8.3 8.5 8.7 8.7
1 L 0
200
400
600
800
1000
l/[DPAl
Figure 4. Double reciprocal plots of ECL data obtained in solutions of different DPA concentrations with various concentrationsof electrolyte: (0)0.1 M TBAH; (A)10 mM TBAH; (V)1 mM TBAH; ( 0 )0.1 mM TBAH.
the solvent is changed from DME to 5050 ACN-TOL, although no broadening of the spectra is seen. ECL of DPA as a Function of Electrolyte Concentration. The ECL efficiency of DPA was measured as a function of supporting electrolyte concentration in both 5050 and 7 5 2 5 (v/v) acetonitrile-toluene mixtures (Table 2, Figure 5 ) . The addition of toluene to ACN allows the concentration of DPA to be increased above the limit (- 1 mM) imposed by the solubility limitations of a ~ e t o n i t r i l ethus , ~ ~ providing a means to minimize the effects of impurities and side reactions. In the case of the 5050 ACNTOL mixtures, the concentration of DPA was systematically varied for each electrolyte concentration examined. An ECL efficiency was then obtained from a plot of 1/$ECL vs 1/ [DPA] in a manner analogous to the studies of the different solvents describedabove (Figure4). TheDPA&CLvaluesin75:25ACNTOL, on the other hand, were the average of four separate experiments for solutions containing 4 mM DPA, a concentration sufficiently high to avoid significant impurity complications. The &CL of DPA was found to increase with decreasing TBAH concentration in both of these mixtures as the electrolyte concentration was lowered from 100 to 0.1 mM (Figure 5 ) . The value of A E 1 p in 5050 ACN-TOL, however, was insensitive to the electrolyte concentration except at the lowest concentrations tested (Figure 5 , Table 2). Identical results were obtained in pure acetonitrile (data not shown). Thus, in the concentration range where the rise in 4EcL occurs, there is no corresponding increase in AE112. Fluorescence Spectra. Fluorescence spectra of DPA were recorded in all of the solvent systems described above both with and without supporting electrolyte. N o solvent or electrolyte specific differences were found in the fluorescence intensity.
Discussion The data in this paper show that the ECL from DPA is sensitive to the solution environment. The ECL efficiency was observed
Maness et al.
3996 The Journal of Physical Chemistry, Vol. 98, No. 15, 1994
0.09
0.08
1
I
0.07
-
0.06
-
0.05
-
4
1
T
a
I’
4
f
P
9ECL
the ground state far into the inverted r e g i ~ n . ~Thus, an increasingly negative AGO will increase the rate of formation of the excited singlet, while the rate of formation of the ground state is decreased. Similar calculations show that kZet,the rate constant for formation of the triplet, is near its maximum, which exceeds k 4 (vide infra). Thus, from eq 7 k2‘ct becomes equal to kd, Le. diffusion controlled. As such, it should be little affected by changes in AGO. Substitution of eq 7 for the excited singlet with k2‘,t = kd, reduces eq 8 to
3.3s 3.30 3.25 3.20
e
e
*
4 3.10
I -4
-3
-2
4ECL
-1
log[TBAH]
Figure 5. ECL efficiency and LiEll2 of DPA as a functionof log([TBAH]1. ECLefficiencies: (0) 7 5 2 5 acetonitrile-toluenecontaining4.0mMDPA,
(v)5050 ACN-TOL. 7 5 2 5 ACN-TOL results are an average of four experiments. The associated error was taken as the standard deviation of the four separate experiments. 5050 ACN-TOL results are extrapolated values from I/&ECL vs l/[DPA]. The error in these extrapolated &ECL measurements was taken as the error in they-intercept of the linear regression. ( 0 )A E l / 2 of DPA measured in 5050 ACNTOL. to be a function of the difference in half-wave potentials for the formation of the radical anion and the radical cation, and these AE1p values are in turn dependent on the specific solvent employed. In addition, the ECL efficiency has been found to increase with decreasing ionic strength in acetonitriletoluene mixtures. These two trends are observed for data which are extrapolated to infinite DPA concentrations, thereby minimizing impurity effects. In addition, the fluorescence and ECL spectra show these changes are not due to changes in the emission process. The increase in &CL of DPA with increasing AElp can be evaluated in terms of conventional electron-transfer rate theories. In the absence of competing side reactions, the relative rates of the individual electron-transfer steps (eq 3) determine $ECL (eq 6). These rates must be corrected by a statistical factor determined by the spin d i s t r i b ~ t i o n .Thus, ~ ~ eq 6 becomes
+ 3kZ’,, + k3’,,)
t#JECL = kl’,t/(kl’et
where Pet, k2‘ct,and k3’ct are the apparent electron-transfer rate constants given by eq 7 for formation of the excited singlet, triplet, and ground state, respectively. The electron-transfer rate constants for these reactions can be characterized by a form of the Arrhenius-Eyring equation:’O
ket = (z)exP(-wr/kbn exp(-AG*/k,T)
(9)
where Z is the liquid phase collision number, k b is Boltzman’s constant, T i s temperature, and w, is the work required to bring the two reactants together. The free energy of activation, AG*, can be related to the standard free energy change, AGO, and the reorganizational energy required for the reaction, A, through simple Marcus theory as
+ A)2/4A
(10)
+ AGes- W,
(1 1)
AG* = (AGO where AGO is expressed as AGO = AG,,
Here AG,, is the free energy change associated with direct formation of the ground state as experimentally determined from AElp(anodic cathodic), and AG, is the energy requbed to reach theexcitedsingletor triplet state (3.09and 1.8 eV,31 re’spectively). At high ionic strength w, has a negligible effect on AGO. For X = 0.74 eV, a value for acetonitrile which is consistent with experimental results for similar aromatic compounds,32 the energetics of the DPA electron-transfer reactions place the formation of the excited singlet in the Marcus normal region and
-
= k1et/(4k1et+ 3 k - 4 ~
(12)
with the assumption that the rate of formation of the ground state is sufficiently inverted as to be insignificant. Thus, an increase in Petresulting from a change in AE1pwill result in an increase in dECL.33 The data in Table 1 qualitatively agree with eq 12 in that &CL increases with increasing AEl/1values. Values of A E l l z increase with a decrease in the solution dielectric constant with the largest measured difference between dimethoxyethane and propylene carbonate. The increased value of AElp in DME is readily understood in terms of its lower ability to solvate radical i0ns,3~ as predicted by the Born equation.’O Similar correlations between solvent dielectric constant, AE1/2, and &CL have been observed for the ECL emission from r ~ b r e n e . ~ ~ , 3 6 The data shown in Table 1 also provide supporting evidence for the selective solvation of species undergoing electron-transfer reactions in mixed solvent systems.3’ The differences between the ECL efficiencies and AEl/l’s in acetonitrile and acetonitriletoluene mixtures are very small, suggesting that acetonitrile selectively solvates the radical-ion reactants regardless of the percentage of toluene added. However, the data do not support an extension of this concept to propylene carbonate-toluene mixtures (Table 1). Although the changes observed in ~ E C correlate L with changes in AE1p, other solvent-related effects should also play a role in changing ( ~ E C L . For example, $ECL should vary with the reorganization energy, A, which decreases with dielectric con~ t a n t . 3Smaller ~ values of A should accelerate electron-transfer reactions located in the normal Marcus region while decelerating those in the inverted region. The combination of this effect and the increased values of A E l p in solvents of low dielectric constant would predict greater increases in C$ECL than are observed. Other ECL reactions accompanied by competing reactions in the inverted region also have proved difficult to model with a simpIe Marcus theorye3s4 Better predictions are obtained when either the additional production of vibrationally excited inverted region products or changes in the electron-transfer reaction distance are considered.39~~ The increase in ~#JECLseen with decreasing supporting electrolyte concentration is not accompanied by changes in AGO since no significant change in AElpoccurs over the supporting electrolyte concentration range where the rise in ~ E C Lis observed (Figure 5 ) . Thus, free energy changes as a result of ion pairing4sv46do not occur in these solutions. Others have also reported changes in &L of other organic systems with electrolyte concentration, but over a limited range (0.2-0.1 M41-42). While it is known that high &CL’S can be obtained with decreasing electrolyte concentration due to a decrease in impurities associated with the electrolyte,43.44 DPA concentration studies performed with 5050 ACN-TOL solutions of varying electrolyte concentration yield intercept ~ E C L ’which S increase as a function of electrolyte concentration, indicating that the effect is not related to impurity concentration (Figure 4, Table 2). Changes in the efficiency of emission from the excited state could result in changes in the observed ECL efficiency; however, the fluorescence efficiency of DPA is unaffected by the presence of electrolyte in solution. The experimentally observed increase in &cLcan be understood through further examination of eq 12. From this equation it can be seen that the increase in ECL efficiency with decreasing ionic
The Journal of Physical Chemistry, Vol. 98, No. 15, 1994 3997
Electrochemiluminescence of 9,lO-Diphenylanthracene
0.0
0.5
1.0
1.5
2.0
k-d/exp(-wr/k,T)
2.5
(lo-’
3.0 3 . 5 x
4.0
8-’)
Figure 6. Equation 17 plotted for DPA ECL efficiencies obtained in (V) 50:50 ACN-TOL and (0)75:25 ACN-TOL data. Values of k 4 and wereca calculated with eq 13 and 14, respectively, with Dexperimentally determined from the steady-state limiting current, as in Table 2, and a = 1.03 X l e 7 cm,* r = 7.45 X l@ cm.I4
strength can occur with either an increase in the rate constant for excited singlet formation or a decrease in the rate constant for dissociation of the encounter complex, k ~ With . constant energetics as determined through A E 1 / 2 measurements, and with the assumption of a constant reorganizational energy over the range of ionic strengths where the initial change in &CL is seen, kIet (eq 9) should vary only with changes in wI, Le. increasing as wr becomes increasingly negative. Likewise kA is also a function of wr and is calculated through the Eigen equation14 as
where DA- and DA+are the diffusion coefficients of the reacting species and a is the sum of the radii of the reactants.l&l* According to Debye-Huckel theory, for ions of unit and opposite charge, w, can be approximated by14
[ l+@rZ/;: exp(-@a&)
w, = _e2 exp(8r\/;)]
(14)
where
and that obtainedfrom theslope ((1.1 i 0.1) X logs-’) isobtained. This, along with the linearity, supports the interpretation of the ionic strength effect. There is, however, a discrepancy between the experimentally obtained intercept and that expected on the basis of an analysis of eq 17. The intercept’s large value suggests the presence of other competing reaction pathways which derease the ECL efficiency from the 25% maximum expected in their absence. Others have also noted that the behavior of DPA ECL fails to completely conform to the simple reaction scheme presented in eq 3.” Although the of DPA is found to be a function of electrolyte concentration, this is not the case for R ~ ( b p y ) 3 +where ~ , both the values of A E l p and the intensity of emission are independent of supporting electrolyte concentration.20 Equation 14 shows that k 4 should increase with decreasing ionic strength for similarly charged reactants such as Ru(bpy)~+Iand Ru(bpy)~+’. However, changes in ~ E C Lfor systems such as R ~ ( b p y ) 3 +with ~ decreasing ionic strength would be difficult to observe since such systems approach 100% efficiency upon populating the emitting state.2.20
Conclusions The efficiency of ECL for DPA has been shown to be sensitive to the solution environment. The efficiency of formation of the excited singlet via the annihilation route can be affected by the value of AGO of the reactants which is a function of the solvating power of the solvent for charged species. Alternatively, the efficiency can be improved by decreasing the rate at which the encounter complex dissociates by lowering the ionic strength. The extreme sensitivity of the ECL efficiency to these and a variety of other solvent and electrode conditions” may explain the difficulty in fitting experimentally obtained ECL efficiencies to those predicted by electron-transfer rate theories. However, the results presented here clearly show that substantial gains in ~ECL can be achieved through appropriate selection of reaction conditions. Since such reaction conditions involve highly resistive media, ECL under such conditions can only be effectively examined through the use of microelectrodes such as those used in this study. Further investigations should lead to much higher efficiencies for these reactions. Indeed, preliminary data indicate that remarkably high efficiencies can be obtained in dimethoxyethane with low concentrations of electrolyte.
Acknowledgment. This research wassupported by theNational Science Foundation (CHE). and p is the ionic strength, t is the static dielectric constant of the solvent, N is Avogadro’s number, e is the electron charge, and r is the radius of reactant A- (or A+) plus that of the dominant counterion in the ionic atmosphere, assuming rA- = FA+. Examination of eq 14 reveals that a decrease in ionic strength results in a corresponding decrease in the work required for the radical ions to stay together. This leads to a decrease in k, (eq 13), an increase in klct (eq 9), and thus an overall increase in ~ E C L (eq 12) with decreasing ionic strength. The effect of changes in k 4 on ~ E C Lwith ionic strength can be most clearly seen through a double reciprocal analysis of eq 12:
Substitution of eq 9 for klctrassuming constant energetics and reorganizational energy, then yields 1/&CL = 4 + {3/k’ct(p
m)){k4/exp(-wr/kbT))
(17)
When plotted in this way, the experimental data have a slope of (2.6f 0.27) X 10-9 s-1 and a y-intercept of 9.4 k 0.5 ( R = 0.97, Figure 6). Good agreement between kIct ( p a), calculated using eq 9 with X = 0.74 eV and 2 = 1011 M-1 s-1 (0.8 X 109 s-l),
-
References and Notes (1) Faulkner, L. R.; Bard, A. J. In Electroonulyrical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1977; vol. 10, p 1. (2) Wallace, W. L.; Bard, A. J. J . Phys. Chem. 1979, 83, 1350. (3) Kapturkiewiez, A. J. Electroanul. Chem. 1991, 302, 131-144. (4) Van Duyne, R. P.; Fischer, S . F. Chem. Phys. 1974,5, 183-197. (5) Tachikawa, H.; Bard, A. J. Chem. Phys. Lett. 1974, 2, 246-251. (6) Itoh, K.; Honda, K,Sukigara, M. Electrochim. Acta 1979,24,11951198. (7) Mussell, R.D.; Nocera, D. G. J . Am. Chem. SOC.1988,110,27642772. (8) Visco, R. E.; Chandross, E. A. J . Am. Chem. SOC.1964,86, 53505351. (9) Faulkner, L. R.;Tachikawa, H.; Bard, A. J. J . Am. Chem.Soc. 1972, 94(3), 691699. (10) Bard, A. J.; Keszthelyi, C. P.; Tachikawa, H.; Tokel, N. E. In Chemiluminescence and Bioluminescence; Cormier, M. J., Hercules, D. M., Lee, J., Eds.; Plenum Press: New York, 1973; p 193. (1 1) Kim, J.; Faulkner, L. R.J . Electroanul. Chem. 1988,242,107-121. (12) Park, S. M.;Bard, A. J.J. Am. Chem.Soc. 1975,97(1 l), 2978-2985. (13) Berlman, I. B. Handbook of Fluorescence Spectru of Aromutic Molecules; Academic Press: New York, 1965. (14) Chiorboli, C.; Indelli, M. T.; Scandola, M. A. R.; Scandola, F. J . Phys. Chem. 1988.92, 156-163. (15) Meyer, T. J. Prog. Inorg. Chem. 1983, 30, 389-441. (16) Petrucci, S. Ionic Interactions; Academic Press: New York, 1971; vols. I and 11. (17) Sutin, N. J . Phys. Chem. 1986, 90, 3465-3466. (18) Marcus, R. A. J. Chem. Phys. 1965.44, 2654-2657.
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