J. Phys. Chem. B 2004, 108, 19119-19125
19119
Rate-Determining Step in the Electrogenerated Chemiluminescence from Tertiary Amines with Tris(2,2′-bipyridyl)ruthenium(II) R. Mark Wightman,*,† Samuel P. Forry,† Russell Maus,† Denis Badocco,‡ and Paolo Pastore‡ Department of Chemistry, UniVersity of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599-3290, and Department of Chemical Sciences, UniVersity of Padua, Via Marzolo 1, 35131 PadoVa, Italy ReceiVed: July 14, 2003; In Final Form: September 27, 2004
The rates and mechanism of coreactant electrogenerated chemiluminescence (ECL) from tris(2,2′-bipyridyl)ruthenium(II) (Ru(bpy)32+) and the tertiary amines, tripropylamine (TPrA) and trimethylamine (TMeA), in aqueous solution were investigated. Transient (0.5 ms) potential steps were used with microelectrodes to investigate the emission time course under a variety of solution conditions. With amine concentrations that are low with respect to Ru(bpy)32+, the emission rises continually during the transient potential step and decays slowly after its termination. In contrast, the emission approaches a plateau during the potential step and is rapidly extinguished afterward with concentrations of Ru(bpy)32+ that are much lower than the amine concentration. At intermediate pH values, the emission intensity increases approximately linearly with pH. The emission after the potential step is unaffected by the rest potential. To simulate these temporal characteristics by finite difference methods, a mechanism employing 15 discrete chemical and electrochemical steps was employed, using literature-based thermodynamic values and electron-transfer rate constants evaluated from Marcus theory. The rate-limiting step was found to be the deprotonation of the amine radical cation. In addition, the simulations required a rate constant for the homogeneous oxidation of the tertiary amine by electrogenerated Ru(bpy)33+ value much below its Marcusian-calculated value to match the experimental data.
Introduction Oxidation of tris(2,2′-bipyridyl)ruthenium(II) (Ru(bpy)32+) in the presence of tertiary amines results in emission from the Ru(bpy)32+ excited state.1 This process, referred to as oxidativereductive electrogenerated chemiluminescence (ECL), is of considerable importance for analytical applications. For example, using tripropylamine (TPrA) as the coreactant, Ru(bpy)32+ ECL has been used in immunoassays and DNA probe assays.2-9 It has several desirable features including high efficiency photon emission, and it can occur in aqueous as well as nonaqueous solutions. To optimize the reaction conditions and electrochemical parameters used in analytical applications, an understanding of the rate-determining steps leading to photon production is required. Several recent papers have contributed to the understanding of this ECL system, and a number of competitive pathways involving both heterogeneous and homogeneous reactions have been identified and their thermodynamic and kinetic parameters estimated (Table 1).10-16 In aqueous solution, Ru(bpy)32+ exhibits a reversible one-electron oxidation (reaction 1, Table 1). Ru(bpy)32+ can also be reduced (reaction 2, Table 1), although this process lies beyond the solvent breakdown in aqueous solution. ECL generation requires oxidation of a tertiary amine (R3N), either by the electrode or by electrogenerated Ru(bpy)33+, to form a radical cation (R3N+•, reactions 3 and 6, Table 1). The electrochemical oxidation of amines is inhibited at noble metal electrodes,14 but it occurs at carbon surfaces as * To whom correspondence should be
[email protected]. Phone: 919-962-1472. † University of North Carolina at Chapel Hill. ‡ University of Padua.
addressed.
E-mail:
were used in this work. The oxidation of TPrA does not exhibit a reverse wave in cyclic voltammetry, but its standard potential has been estimated from fluorescence quenching experiments11 and, indirectly, from its electrochemical behavior.12,16 Before oxidation, the tertiary amine must be deprotonated (reaction 5, Table 1). Central to the overall ECL reaction is the deprotonation of the radical cation to form a free radical (R2(R′)N•, reaction 7, Table 1), where R′ represents the alkyl group that has lost a proton. Estimates exist of its reducing power from ECL generation11 and quenching15,16 experiments (reaction 4, Table 1). In aqueous solutions R3N+• for TMeA17 and, more recently, TPrA,16 have been observed by electron paramagnetic resonance (EPR). In both cases it was found to be a N-centered radical that disappeared in neutral, aqueous solutions at rates of 35 and 3500 s-1 for TMeA and TPrA, respectively. One pathway to generate the light-emitting state is reduction of Ru(bpy)33+ by the neutral radical to produce *Ru(bpy)32+ directly (reaction 8, Table 1). An alternate pathway is reduction of Ru(bpy)32+ to Ru(bpy)3+ by R2(R′)N• (reaction 9, Table 1) followed by the annihilation reaction, reaction 10, or reaction 11, both of which form the emissive state.16 The excited state emits a photon relatively quickly18 (reaction 12, Table 1). Competing with these reactions are the quenching of reactive intermediates by homogeneous reactions (reaction 13-15, Table 1) and heterogeneous electron transfer (reaction 3 and 4, Table 1). As demonstrated previously,15 the overall emission process is kinetically limited when examined on a millisecond time scale. In this work, we describe a comprehensive set of experiments conducted on a millisecond time scale over a range of reagent concentrations and pH values that reveal slow steps in the generation of emission. A simulation of emission by using the
10.1021/jp036034l CCC: $27.50 © 2004 American Chemical Society Published on Web 11/19/2004
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TABLE 1: Reaction Steps, Thermodynamic Parameters, and Rate Constants for ECL Arising from the Ru(bpy)32+/TPrA System heterogeneous reactions
E° ((V) vs Ag/AgCl)
k° (cm/s)
reaction no.
Ru(bpy)33+ + e- ) Ru(bpy)32+ Ru(bpy)32+ + e- ) Ru(bpy)+ R3N+• + e- f R3N P + e- f R2(R′)N•
1.0616 -1.4816 0.911,12 e-1.711,15,16
1 1 0.612
1 2 3 4
homogeneous reactions that affect the emission time course:
kf
reaction no.
R3NH+ + HPO4-2 ) R3N + H2PO4Ru(bpy)33+ + R3N ) Ru(bpy)32+ + R3N+• R3N+• f R2(R′)N• + H+ Ru(bpy)33+ + R2(R′)N• ) *Ru(bpy)32+ + P Ru(bpy)32+ + R2(R′)N• ) Ru(bpy)3+ + P Ru(bpy)33+ + Ru(bpy)3+ ) *Ru(bpy)32+ + Ru(bpy)32+ Ru(bpy)3+ + R3N+• ) *Ru(bpy)32+ + R3N *Ru(bpy)32+ f Ru(bpy)32+ + hυ
108.2 (M-1 s-1)32 100 (M-1 s-1)a 102.73 (s-1)a 1010 (M-1 s-1)b ∼1010 (M-1 s-1)b,c ∼1010 (M-1 s-1)b,c ∼106 (M-1 s-1)a,c 107.2 (s-1)18
5 6 7 8 9 10 11 12
homogeneous reactions that affect the emission intensityd
reaction no.
R3N+• + R2(R′)N• f R3N + P Ru(bpy)33+ + P ) Ru(bpy)32+ + Pirr Ru(bpy)3+ + P ) Ru(bpy)32+ + Pirr
13 14 15
a This work, higher values predicted more emission than was experimentally observed. b Estimated from Marcus theory, using eqs 7 and 9-11 in ref 19. c These values are estimates; the emission time course is relatively insensitive to the values employed. d The kinetics of these reactions do not affect the emission shape and therefore the rate constants could not be estimated.
reactions in Table 1 is used to evaluate the rate constants of some of the individual steps. As summarized in Table 1, the thermodynamics of the reactions that lead to oxidativereductive ECL are quite well characterized and many of the rates can be estimated by Marcus theory for electron-transfer reactions.12,19 Our investigation reveals that reactions 6, Table 1, is slower than generally recognized. In addition, our results support the controversial finding16 that while R3N+• is relatively long-lived, i.e., it deprotonates on a millisecond time scale, it is not available during that time for electroreduction. Reaction 7, Table 1, is the rate-determining step. Experimental Section Apparatus. ECL was characterized in a flow injection analysis system.15 The system consisted of a stainless steel syringe pump, loop injector, and a channel-type electrochemical cell with a channel height of 100 µm and a width of 3 mm. The base of the cell was fabricated with epoxy into which was embedded a disk-shaped, carbon-fiber microelectrode (15-µm radius) and a silver-band electrode. The top of the channel contained an optical window. A photon-counting photomultiplier tube (R4632, Hamamatsu, Bridgewater, NJ) operated at -870 V (Fluke 412B, Everett, WA) was placed adjacent to this window. The output of the photomultiplier tube was amplified (EG&G Ortec VT120A, Oak Ridge, TN) and directed to a multichannel scaler (EG&G Ortec, T-914). Procedures. Phosphate buffer (0.1 M, pH 7.40 unless otherwise noted) was delivered from the pump to the electrochemical cell at a flow rate of 1 µL/min, and the solutions for ECL were introduced to the cell via the loop injector. The silver band electrode was connected to a saturated silver/silver chloride reference electrode that was placed downstream from the electrochemical cell. The potential waveform applied to the working electrode was generated with Hewlett-Packard 34811A Benchlink Arbitrary Waveform software. At the beginning of each cycle the potential was pulsed to 1.4 V from a resting potential of 0.0 V versus Ag/AgCl for 0.5 ms and then returned to the rest potential. The pulse potential was selected to achieve diffusion-limited oxidation of Ru(bpy)32+. Cycles were repeated at 50 ms intervals, and the emissions from 2000 potential steps
were summed. Emission counts were collected in 50-µs bins during and after the pulse potential. The time between each cycle allowed clearance of electrogenerated species from the vicinity of the microelectrode surface. The shape and intensity of photon emission remained constant during data collection. Solid-State Switch. A solid-state switch (MAX326, Maxim, Sunnyvale, CA) was employed in some experiments to open the circuit between the potential generator and working electrode. A circuit diagram for the switch circuitry is available in the Supporting Information. The switch can open in ∼150 ns. Indeed, with test systems (Supporting Information), switching was virtually instantaneous and the isolated electrode (switch open) passed no current. Although ECL persisted beyond the opening of the switch, this was due to diffusional mixing of previously generated intermediates. Finite Difference Simulation. The experimental curves were compared to the results of a simulator that could evaluate multiple potential steps as were used experimentally. The approach used to simulate the ECL systems was similar to recent literature reports.20,21 The implicit finite difference method was used22,23 with θ ) 1 so that it takes the form of the fully implicit Euler method. In this method, diffusion, described by Fick’s second law, ∂R/∂t ) DR(∂2R/∂x2), and homogeneous kinetics, described by kinetic terms, such as kR+•R-• + ..., can be expressed in the following discrete forms:
Rj+1 - 2Rj + Rj-1 R′j - Rj = DR(1 - θ) + ∆t ∆x2 R′j+1 - 2R′j + R′j-1 D Rθ ∆x2 k(1 - θ)Rj+•Rj-• + k(θ)R′j+•R′j-• + ...
(diffusion) (I) (kinetics) (II)
where the prime indicates subsequent time and j references the space grid where the concentrations of all chemical species are defined. The temporal and spatial increments and the diffusion coefficient are given by ∆t, ∆x, and DR, respectively. The implicit method, together with the bimolecular nature of the kinetics, requires solution of a nonlinear system of the implicit
Electrogenerated Chemiluminescence from Tertiary Amines
J. Phys. Chem. B, Vol. 108, No. 50, 2004 19121
terms. Following Rudolph,21,24 we linearized the kinetic terms, R′j +•R′j-•, as follows:
R′j+•R′j-• = R′j+•Rj-• + Rj+•R′j-• - Rj+•Rj-•
(III)
The experiments to be simulated consisted of the summation of photonic emission from successive rapid potential step waves. To simulate this quickly, we used expanding space and time grids25-28 appropriate for square wave experiments:
exp[β(j - 0.5)] - 1 exp(β) - 1
(IV)
exp[ϑ(i - 0.5)] - 1 ti ) ∆t exp(ϑ) - 1
(V)
xj ) ∆x
Here, xj is the position at the center of the space-grid element at time ti, and β and φ, whose values must be lower than 0.5,27 represent the expansion coefficients for space and time, respectively. The dimensions of the diffusion and reaction layers are critical for an accurate simulation. For a square wave of n cycles and characterized by anodic and cathodic step times ta and tc, the diffusion layer can be described as
Yd ) 6xnDmax(ta + tc)
(VI)
where Dmax is the diffusion coefficient of the “fastest” chemical species involved in the mechanism. The reaction layer29 can be defined as
1 µ ∝ 1/3 k
(VII)
where k is the greatest rate constant. At least five grid points were used to describe its shape. Ohmic drop was iteratively computed by the bisecant method from
((
Ru(I ′c + I ′f) ) Ru Cdl
)
E′n,i - E′true,i-1 ti - ti-1 FS
∑e
De
(
+
))
C′e,1 - C′e,0 ∆x/2
(VIII)
x)0
where Ru is the uncompensated resistance, E′n,i and E′true,i are the actual and true potentials, I ′c and I ′f are the charging and faradaic currents in implicit form, C′e,0 and C′e,1 are the concentrations at the electrode surface and in the first grid element, respectively, De is the diffusion coefficient, Cdl is the double layer capacitance, S is the electrode area, and F is the Faraday. The iterative procedure was stopped when the difference between two successive values was lower than 0.0001 V. To account for acid-base effects, complete equilibria and kinetic equations were employed. The simulator did not account for convection. This is reasonable since, with the flow rate employed, the linear velocity at the edge of the Nernst layer (D(ta + tc))1/2 was 21 µm/s, quite slow compared to the time scale of the potential steps. Chemicals. Tris(2,2′-bipyridyl)ruthenium(II) chloride hexahydrate (Ru(bpy)32+), trimethylamine (TMeA), and tri-n-propylamine (TPrA, 99%+), were obtained from Sigma Chemical (St. Louis, MO). Sodium phosphate monobasic was obtained from Mallinckrodt (Paris, KY). All solutions were prepared with distilled deionized water (Mega Pure System MP-3A Corning
Figure 1. ECL time course with 10 mM Ru(bpy)32+ in the presence of low concentrations of TPrA (10, 20, and 50 µM) at pH 7.4. The background subtracted ECL (noisy traces) increases with TPrA concentration. At t ) 0, the potential was stepped to 1.4 V vs Ag/ AgCl and then returned to 0.0 V at 0.5 ms (indicated by the vertical line). The pulse sequence was repeated 49.5 ms later, and the emission shown is the sum from 2000 consecutive potential steps. Simulations (smooth traces) based on the parameters in Table 1 and the experimental concentrations are overlaid on the emission traces. Inset: At low TPrA concentrations (given in µM), the cumulative emission arising from the potential steps (including light emitted after the potential pulse) was normalized for different Ru(bpy)32+ concentrations (2, 5, and 10 mM), averaged (black circles), and plotted against [TPrA]. Error bars show the standard deviation. A linear fit to the logarithmic plot of the experimental data yielded a slope of 1.06.
Glass Works, Corning, NY). The pH of the 0.1 M phosphate buffer was adjusted with 12 M NaOH. Solutions were freshly prepared for each experiment. Results ECL Response on the Millisecond Time Scale. With Ru(bpy)32+ in excess, the emission intensity increases with increasing concentrations of TPrA (Figure 1). The emission increased during the 0.5 ms potential pulse, and then decreased following return of the potential to 0.0 V. The rate at which the emission decayed was relatively slow, decaying to baseline over 20 ms, a relatively long time scale compared to the 0.5 ms generation time. Superimposed on the time courses of the ECL in Figure 1 are the simulated responses, using the rate constants and thermodynamic values summarized in Table 1. With [Ru(bpy)32+] in significant excess over [TPrA], the electrogenerated [Ru(bpy)33+] should be little altered by reactions with TPrA species (reactions 6, 8-11). Indeed, at [Ru(bpy)32+] ) 2, 5, or 10 mM, the decay rate of the emission in solutions containing low concentrations of TPrA (0.01 to 0.05 mM) could be fit to an exponential that was quite similar in rate (data not shown), suggesting first-order behavior in a species originating from TPrA. The cumulative ECL counts at a single [TPrA] were similar with different [Ru(bpy)32+] as shown by the small error bars in the inset, Figure 1, indicating first-order behavior in Ru(bpy)32+ as we previously described.16 The inset also shows that a logarithmic plot of the normalized ECL intensity is linear with [TPrA] with a slope of 1.06. This further indicates that the reaction is first order with respect to [TPrA] under these conditions, as had been suggested in previous work.15 Effect of Rest Potential on ECL. The potential applied to the electrode before and after the potential step, 0.0 V, could diminish the ECL emission by causing reduction of R3N+• or Ru(bpy)33+. Therefore, a solid-state switch was used to disconnect the working electrode after the potential step, preventing reversal of reactions 1 and 3 (Table 1). The switch was evaluated with the oxidation and reduction of 9-phenylanthracene, an annihilation-based ECL system in which the electrogenerated reagents simply diffuse together, react, and emit photons. When
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Figure 3. pH dependence of ECL. ECL from solutions of [Ru(bpy)32+] ) 10 mM and [TPrA] ) 0.05 mM were acquired in 0.1 M phosphate buffer solutions of pH 6.4, 6.9, 7.4, 7.9, and 8.4. The main panel shows the observed emission and simulated responses for the lowest three pH values. The inset shows the cumulative, background subtracted light output (black circles) as a function of pH. The line shows the simulated cumulative emission. The potential program was the same as in Figure 1.
Figure 2. Effect of rest potential on emission intensity and time course during and following a potential step (1.4 V vs Ag/AgCl) at pH 7.4. Experimental emission was indistinguishable whether a solid-state switch in series with the microelectrode was opened at 0.45 ms or always closed. Upper panel [Ru(bpy)32+] ) 10 mM, [TPrA] ) 0.05 mM. Middle panel: [Ru(bpy)32+] ) 0.05 mM, [TPrA] ) 50 mM. Lower panel: [Ru(bpy)32+] ) 5 mM, [TPrA] ) 50 mM. Other conditions were as in Figure 1. Simulations for the fifth polarization cycle, overlaid on the emission, were based on the parameters in Table 1 and the experimental concentrations. Simulations of the switch open or closed gave indistinguishable emission time courses except for the conditions shown in the lower panel. In this case, simulated emission decayed more slowly than the experimental data, and was slowest when the potential was returned to the initial potential.
the working electrode was disconnected from the applied potential, the emission was extinguished for that system (Figure SI 2, Supporting Information). In contrast, when the working electrode was disconnected at the end of the potential step, the emission from the Ru(bpy)32+/TPrA system decayed with the same rate as when it was returned to a reducing potential at all of the concentrations examined. Shown in Figure 2 (upper panel) is a result with Ru(bpy)32+ in excess as in Figure 1. Simulated responses with use of the values in Table 1 were similarly insensitive to the different rest potentials. When the concentration of TPrA was greater than that of Ru(bpy)32+, several changes in the ECL time course were observed. The duration of the emission after the potential step was shorter (Figure 2, middle panel) when TPrA was in 1000-fold excess over Ru(bpy)32+, and the ECL response became more closely time-locked to the applied potential. We previously showed that with TPrA in large excess the emission exhibits first order behavior in Ru(bpy)32+.15 In addition, the ECL approached a plateau while the potential was applied. Thus, the emissive time course indicates that production of the emissive state (reaction
8, Table 1) is approaching a steady state at 0.5 ms, indicating that one of the prior reagents necessary to form it is becoming limiting. With both Ru(bpy)32+ and TPrA at relatively high concentrations, intermediate behavior is observed (Figure 2, lower panel). Simulations of this condition with the values in Table 1 give a worse fit than the other experimental conditions. The rise of the emission is predicted well but the simulated decay is slower than the experimental one. The simulated decay with an applied rest potential of 0.0 V is slower than that with the switch open; this effect was not seen in the experimental data. Effect of pH. Both reactions 5 and 7, Table 1, appear to be simple acid-base reactions. Therefore, we examined ECL from solutions of pH 6.4 to 8.9 with a 200-fold excess of Ru(bpy)32+ over TPrA. As shown in Figure 3, the cumulative ECL emission, with the background subtracted, increased significantly with increasing pH, consistent with previous findings.30 A logarithmic plot of experimental emission versus pH (inset, Figure 3) exhibits a slope of 0.76 for low pH values. This indicates there is first-order relationship between [OH-] and light emission; surprisingly, the data point to only one pH-dependent step preceding the rate-determining step. At the higher pH values, the emission increased less with increased pH than predicted by the simulation, presumably due to competition from the inefficient light-producing reaction between OH- and Ru(bpy)33+ that could deplete the [Ru(bpy)33+] before its reaction with TPrA.15,31 Consistent with this, at pH values above pH 7.4, emission was observed during oxidation of Ru(bpy)32+ in the absence of amine. The simulated pH response shown in the inset of Figure 3 shows increased emission with greater pH values. Interestingly, despite the increased intensities at each pH value, the rate of decay of the experimental emission at each pH was almost identical with a half-life of ∼2 ms. The simulated responses with the use of the values in Table 1 showed the same behavior. Simulated Emission Responses. The ECL emission for the various conditions in Figures 1-3 were simulated by using the values in Table 1, and the calculated responses are superimposed on the experimental ones. As can be seen, the simulation captures the temporal features of the experimental emission traces over a broad range of solution conditions including pH varied over 2 units, Ru(bpy)32+ concentrations over a 500-fold
Electrogenerated Chemiluminescence from Tertiary Amines range, TPrA concentrations over a 5000-fold range, and two different conditions for the potential between steps. The worst fit is obtained when both Ru(bpy)32+ and TPrA are at very high concentrations (Figure 2, bottom panel). While the thermodynamic values for the first three heterogeneous reactions are reasonably well-known, and reaction 4 has been estimated by several groups, assigning rate constants to the remaining reactions required several considerations. Of the 11 homogeneous reactions in Table 1, only the rate constant of the emissive step, reaction 12, is well established. However, the rate constant for several of the reactions can be estimated. The reactions that cause the major features of the emission time course under most conditions are reactions 5-8, Table 1. The rate constant for the amine deprotonation reaction 5, Table 1, can be calculated according to Eigen32 as kf ) (kdiff)/(1 + K), where log(K) ) pKX - pKY- for weak acids HX and HY, and kdiff ) 1 × 1011 M-1 s-1. Thus, with K ) 1010-7.2, kf ≈ 108.2 M-1 s-1 for TPrA. At pH 7.4 and [HPO4-2] ) 0.06 M, a pseudofirst-order rate constant for deprotonation by this pathway is k′ ) 106.6 s-1. Competing with this pathway is proton abstraction by OH-, a process that should have a second-order rate constant near kdiff. However, the low concentration of OH- in a solution buffered at pH 7.4 gives a pseudo-first-order rate constant of ∼104.4 s-1. This reaction becomes more competitive with reaction 5 as the pH is increased. The rate constant for reaction 8 was estimated from Marcus theory by using the thermodynamic values given for reactions 1 and 3, Table 1. Reaction 8 is the major reaction that produces the excited state, and the large value of its rate constant predicted by Marcus theory was required to simulate the approximately linear relationship between pH and ECL emission (Figure 3). Marcus theory also predicts that formation of the ground state by the reactants of reaction 8, Table 1, is relatively slow so it was not included in the mechanism. Computation of the rate constant for homogeneous oxidation of TPrA by Ru(bpy)33+, reaction 6, with Marcus theory, using the E° values of reactions 2 and 3, predicts a much larger rate constant10 than was used in the simulation. However, when the rate constant of reaction 6 was increased above the value in Table 1, the simulated background emission increased because reagents accumulated in the vicinity of the electrode during the repeated potential steps employed. This was not observed experimentally unless the steps were repeated at very close intervals.15 The low rate constant for the deprotonation of the amine cation radical, reaction 7, Table 1, was required to capture the slow decay of the data with low TPrA concentrations (Figures 1 and 3). Reactions 9-11 are minor contributors to the overall emission and the simulation was quite insensitive to their rate constants. This is because, as revealed by the concentration profiles generated by the simulation, the concentration of Ru(bpy)31+ is quite low under all of the experimental conditions examined. However, emission from reaction 9 followed by reaction 11 was considered because it has been experimentally demonstrated.16 The rate constants for these reactions were also calculated from Marcus theory. The rate of reaction 10 has been measured directly in acetonitrile,33 and the Marcusian estimate is consistent with the high value reported. The value given for the rate constant of reaction 11, Table 1, was used because higher values led to the prediction of more emission after the potential step than was measured. The simulated responses do not predict the absolute emission amplitudes; the amplitudes were adjusted to achieve the fits shown. This is because, while the shape of the simulated curves is insensitive to the values of reactions 1315, Table 1, these reactions do affect the amplitude. Therefore
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Figure 4. TMeA ECL. Background subtracted ECL from pH 7.4 solutions containing 50 mM TMeA and (A, top) 0.05 mM Ru(bpy)32+ or (B, bottom) 5 mM Ru(bpy)32+ when the potential is stepped to 1.4 V vs Ag/AgCl for 0.5 ms (dark gray line). Other conditions were as in Figure 1. Simulations superimposed on the data used the rate constants for TPrA given in Table 1 except for the following: reaction 7, the deprotonation rate constant was 2.5 times lower, k7 ) 102.33 s-1, and the pKa of TMeA was taken as 9.74; E° ) 1.1 V was used for Reaction 3.
the simulations employed values for these rates that approached the diffusion-controlled rate constant. The simulation does a very good job at fitting the experimental data under most of the experimental conditions. For example, in prior work we demonstrated that the emission intensity is directly proportional to [Ru(bpy)32+], indicating firstorder behavior in this reagent, when R3N is the excess reagent.15 However, as the concentration of Ru(bpy)32+ was raised in the presence of high [R3N], the cumulative emission intensity approached an asymptote. The kinetic and thermodynamic values in Table 1 predict exactly these behaviors. To fit the data it was necessary to assume that reaction 3, Table 1, was electrochemically irreversible. If it was not, the response in the rest interval without the potential applied was much more intense than that with it applied at a reducing value, in direct contrast with the results in Figure 2. The low value of the rate for reaction 7, Table 1, yields emission curves that exhibit the exponential decay after the potential step observed with high Ru(bpy)32+ concentrations (Figure 1) as well as the steady-state emission profile that is approached at low Ru(bpy)32+ concentrations (Figure 2, middle panel). ECL from TMeA and Ru(bpy)32+. Similar reactions and solution conditions were evaluated for reaction of Ru(bpy)33+ with TMeA. Figure 4 shows the results of a potential step to 1.4 V versus Ag/AgCl at two different sets of concentrations. The TMeA/Ru(bpy)32+ system was 85 times less efficient in the generation of emission than the TPrA/Ru(bpy)32+ system. Nevertheless, the emission curves with TMeA exhibited similar shapes as with TPrA. The decay of the ECL following the potential pulse was slower with TMeA than with TPrA, indicating that the rate of reaction 7 for TMeA is slower than that for TPrA. Indeed, simulations with similar rate constants agreed quite well with the experimental data (overlays, Figure 4).
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SCHEME 1
Discussion The results of this study show that the ECL from Ru(bpy)32+/ tertiary amine systems are kinetically limited when examined on a millisecond time scale. To determine the actual ratedetermining steps, we have used the rate constants for the Ru(bpy)32+/TPrA system, summarized in Table 1, with a finite difference simulation to fit the dynamics to the experimental emission results. The simulations, shown overlaying the data in Figures 1-4, capture the data remarkably well over a very broad range of conditions. These fits required the assumption that Reaction 3, Table 1, is an irreversible reaction. The values for the homogeneous rate constants for reaction 6 and 7 and the E° for reaction 3 were the primary parameters adjusted to achieve a good fit. The unusually slow rate of the deprotonation of the amine radical cation was found to be the major ratedetermining step on the millisecond time scale. An identical reaction sequence was found to describe the emission time course for the Ru(bpy)32+/TMeA system with only a slight adjustment of rate constants. The use of a low rate constant for reaction 6 was somewhat surprising since Marcus theory with the E° values of reactions 2 and 3 predicts a much larger rate constant.10 From experiments in which the catalytic current at a microelectrode was monitored, the combined rate constant for reactions 5 and 6 at pH 6 was estimated as ∼102 M-1 s-1.10 Recognizing the pH dependence of reaction 5, the rate constant for the electron-transfer reaction, reaction 6, Table 1, was extrapolated to have a value similar to the Marcusian computed value. However, the pH correction neglected the role of the conjugate base on the reaction, perhaps leading to an underestimate of the deprotonation rate. A slow rate of reaction 6 is consistent with the general observation that oxidation of tertiary amines is slow.34 Furthermore, stop-flow measurements of the chemiluminescent time course following mixing of Ru(bpy)33+ and TPrA at pH 6.0 had a half-life greater than 5 s.35 Since the slow chemiluminescence must have resulted from the sequence of reactions 5, 6, 7, and 8, at least, the much faster time course in our potential step experiments directly indicates that electrogeneration of R3N+• is a far more efficient route to emission. Similarly, the low emission observed at noble metal electrodes, where the heterogeneous oxidation, reaction 3, Table 1, is inhibited, is consistent with a low rate constant for the homogeneous oxidation of tertiary amines.14 The occurrence of reaction 11, Table 1, has previously been pointed out by Bard and co-workers,16 and they incorporated it in their Scheme 6. However, our simulation indicates it is a minor contributor to the emission. The critical experiment that required us to consider reaction 3, Table 1, as electrochemically irreversible is that shown in Figure 2. The emission was found to be identical whether the potential of the electrode was set to 0.0 V or disconnected after the oxidation step. If reaction 3 is reversible, then applying a potential more negative than that required for electrooxidation should remove R3N+• by electrolysis, resulting in a lowered emission because of a loss of this key reagent. This was not
the case, and the simulation required that R3N+• has a lifetime in the millisecond range to yield the observed emission profiles. This leads to the expectation of a reverse wave for TPrA+• when examined with fast scan cyclic voltammetry, although this has not been observed.36 To explain this, Bard and co-workers proposed that the reduction of R3N+• is obscured by the simultaneous oxidation of R2(R′)N•.16 However, the emission is unchanged whether there is a reducing potential applied after the initial potential step or not. The cause of the electrochemical irreversibility is unclear. We propose Scheme 1 to account for events after formation of the radical cation. Other neutral/radical cation systems, such as those for aliphatic hydrazines, that exhibit a large geometric change upon oxidation, have a slow rate of electron selfexchange.37 This has been attributed to inner-sphere effects on electron transfer.38 In the case of the tertiary amines, the evolution from the pyramidal to planar geometry enables orbital overlap so that the radical can be transferred to the R-carbon. While it may be electrochemically inaccessible, this planar structure would yield an EPR spectrum as has been reported for triakylamines.16,17 Our findings indicate that the deprotonation reaction, reaction 7, Table 1, that yields R2(R′)N•, from the electrochemically formed radical cation, is remarkably slow. The simulation for TPrA employed a rate constant for this deprotonation of 540 s-1, a value within an order of magnitude of that reported for the lifetime of TPrA+•.16 In contrast, reaction 5, the formation of the amine free base, is rapid as expected,32 and an increase in pH increases the emission intensity because of the greater availability of free base for oxidation, as well as its increased formation rate. The slow rate of reaction 7, Table 1, may arise because [R′CH2-•N+R2] is a carbon acid, and these are known to be resistant to base-promoted deprotonation.32,39 Indeed, it has been demonstrated that the deprotonation of amine radical cations exhibits a large isotope effect, clearly demonstrating that deprotonation is rate limiting in their decomposition.40 Consistent with our finding that emission with TMeA is slower than that with TPrA, the former radical has a higher charge density than the latter, which should lower its rate of formation. Collectively, the results of this study provide at least four new insights into coreactant ECL. First, the finding that the homogeneous oxidation of tertiary amines by Ru(bpy)33+ is slow explains why electrochemically based methods that enable direct oxidation of the amine result in higher emission than when amines and Ru(bpy)33+ are homogeneously mixed. Second, the pronounced effect of solution pH on the observed emission is due to deprotonation of the parent amine, but not the deprotonation of the radical cation. This explains why the emission rate is unchanged at higher pH values. Third, although reactions 9-11 are thermodynamically viable, they contribute little to the overall emission under the conditions explored in this work because the concentration of Ru(bpy)31+ is normally quite low. Finally, the results predict that emission can be improved by conditions that promote deprotonation of the radical cation. This
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