J. Phys. Chem. 1993,97, 6710-6714
6710
Spectroscopic Parameters, Electrode Potentials, Acid Ionization Constants, and Electron Radicals and Ions Exchange Rates of the 2,2~-Azinobis(3-ethylbenzothiazolineine-6-sulfonate) Susannab L,Scott, Wen-Jang Chen, Andreja Jjakac,' and James H. Espenson' Department of Chemistry, Iowa State University, Ames. Iowa 5001 I Received: January 5, 1993; In Final Form: March 9, 1993
The characteristics of the colorless ABTSZ- ion [ABTS = 2,2'-azinobis(3-ethylbenzothiazoline-6-sulfonate)] and of the persistent, intensely-colored radical ABTS- have been examined since some reported values are in disagreement. The standard reduction potentials of ABTS'-/ABTSz- and ABTS'-/HABTS- are 0.68and 0.8 1 V us NHE, respectively. A second wave in the CV is associated with ABTSO/ABTS'-, E l p = 1.09 V. The pH dependence of Eo leads to pKa(HABTS-) = 2.2f 0.3;a spectrophotometric pH titration gives 2.08 & 0.02 at p = 1.O M. The protonated radical NABTS' was not detected (pKa < 0), consistent with the higher acidity of the radical relative to its reduced precursor. The EPR spectrum of ABTS- shows a multiline spectrum centered at g = 2.0036 f 0.0004 from the hyperfine coupling to two sets of two equivalent nitrogens as well as six aromatic hydrogens. Line broadening of the methyl and aromatic resonances of ABTS" in the presence of the radical gives an electron exchange rate constant of (4 & 1) X lo7 L mol-' s-l for ABTSZ-/ABTS'- in neutral aqueous solution. In 1.O M perchloric acid, Fe3+ oxidizes HABTS- (k = 1.30 X lo2 L mol-' s-') and Fez+reduces ABTS'- (k= 7.5 X lo2L mol-' s-I). Although the equilibrium constant K = kf/k, for the reaction agrees with that predicted by the measured potentials, the individual rate constants are much lower than those one calculates from the Marcus cross-relation, which is typical for reactions of Fe(HzO)a3+ and Fe(HzO)6Z+.
Introduction The general usefulness in mechanistic investigationsof the ion ABTS2- (see structure) and its one-electron oxidized radical ion ABTS*- warrants their thorough characterization. These ions are used to detect radicals, to evaluate rates of radical formation, and as so-called kinetic probes for reactions other than those of the parent radicals. For example, the stable and colorless ABTS2dianion is oxidized by inorganic radicals (eq 1, X' = HO', Brz'-, RS', NO2, etc.)lV2and by a l k ~ x y land ~ . ~alkylper~xyl~~~ radicals.
-
""& h-4 ABTS~-
+
X*
y
I
5 \
4
+ x-
ABTS*y
1
1
I
C2H5
C2H5
3
(1) -
,s "
The formation of the persistent radical ABTS'- can be used to affirm that X' is an intermediate, both qualitatively and quantitatively, since ABTS'- is readily detected by its EPR spectrum and intense absorption spectrum. Reaction 1 can also be used as a kinetic probe of the reaction of radical X' with substrate S, as in eq 2. For the reactions of ABTSO- and alkyl radicals the situation is reversed. The alkyl radicals do not react withABTSZ-,whereas theydorapidlyform adducts with ABTP-, eq 3.54
-+
X' + s
XS' (or X-
+ S+etc.)
ABTS2- is protonated below pH 3 and report the acid ionization constant of HABTS-. Also we have measured by NMR line broadening techniques the rate constant for the electron exchange between the partners ABTS'-/ABTS2-.
Experimental Section Diammonium 2,2'-azinobis(3-ethylbemthiamline6-sulfonate) was purchased from Aldrich and Sigma. This salt was dissolved in hot (recently boiling) distilled water togivea saturated solution. White crystals that formed upon slow cooling were filtered and air-dried. Dilute aqueous solutions of ABTS- were oxidized to ABTS*- by a slight stoichiometric deficiency of Br2 or Ce(1V). Our results were not influenced by B r or Ce3+and did not depend on the oxidation method. Hydrated ferric perchlorate was prepared by fuming FeC1y6H20 in concentrated HC104 until the solution gave a negative test for chloridewith Ag+, Hydrated ferric perchlorate precipitated upon cooling. Solutions of Fe3+ were reduced to Fez+ over Zn/Hg and were standardized spectrophotometrically with 1,lO-phenanthroline. UV-vis spectra were obtained with a Shimadzu UV3101-PC scanning spectrophotometer. Kinetics were performed by use of a Durrum stopped-flow spectrophotometer thermostated at 25 OC. Pseudo-first-order rate constants were obtained by fitting the absorbance-time data to eq 4. Cyclic voltammetric deterAbs, = Abs,
+ (Abs,-Abs,)
exp(-k,t)
(4)
(2)
R' ABTS*- R-ABTS(3) The ABTS system has also found use as a probe for highly oxidizing intermediates in metal porphyrin system^,^ as a quantitative test for hydrogen peroxide in glucose determinations,'O and as a chromogen in peroxidase assays.]] To use the ABTS2-/ABTS'- system in these applications, one needs reliable information about its kinetic and thermodynamic properties and about the predominant species in aqueous solutions at different pH values. Since some literature values for molar absorptivities and reduction potentials are incorrect,I2 and have been perpetuated,I3we have redetermined these values. We have found that +
0022-3654/93/2097-67 10$04.00/0
minations were performed with a BAS electrochemical analyzer equipped with a freshly-polished glassy carbon electrode and an Ag/AgCl reference electrode. EPR spectra were obtained with a Bruker ER200 D-SRC spectrometer with a 9.5-GHz klystron. Dynamic NMR measurements were performed with a Varian VXR300spectrometer. The transverse relaxation times (T2) were measured by a CPMG ~ e q u e n c e at ~ ~21 J ~OC. Re5ultS Spectraandply.. ThesptrumofthecolorlessABTS2-dianion in water has been reported by others.IJ1J7 We redetermined the
8 1993 American Chemical Society
The Journal of Physical Chemistry, Vol. 97, No. 25, 1993 6711
ABTS Radicals and Ions 1.8
"
"
"
"
"
"
I8
"
"
"
"
I
'
,
"
1.8
y
j
1.6 135
1.4
b a n
10.9
1.2
0
10
40
30
20 IH+l/mM
Figure 1. UV-visible spectra of ABTSZ-(1, pH 7), HABTS- (2, pH l),
andABTS'(3). ABTS'was prepared byincctmplete(-95%) oxidation of ABTSZ-by Ce(IV) to avoid overoxidation to ABTSO. The inset shows a spectrophotometrictitration of ABTSZ-by Ce4+.
Figure2. Spectral shifts accompany the acid-base equilibriumbetween HABTS- and ABTSZ-. The change of the absorbance at 340 nm variw systematicallywith pH. The fitted line from eq 6 is shown.
TABLE I: Electronic Spectra of ABTS Swcies specie X,/nm c/L mol-' cm-1 AB=% 340 3.66 X 1Pa HABTS310 2.09 x lP* 34oc 9-13x 1 0 3 b ABTS417 3.47 x l e d
l
645 728 810 518
.
I
I
.
l
-0 s o 0
tO.0
~~
1.35 X 10"" 1.50 X lPd 1.27 X 10'" 3.6 X 10'
ABTSO a pH 5-7, ref 1. From measurements at pH 0-1, extrapolated as in eq 6. Not a maximum. At pH 0 after precise oxidation by Ce(1V). extinction coefficient, confirming that cmx = 3.66 X 104 L mol-' cm-l at 340 nm, rather higher than a more recently reported figure.12 Some commercial samples required recrystallization before this value was realized. Indeed, ABTS'- could not coexist with some samples of ABTSZ- unless the ABTS- sample had been purified. The spectra of ABTS- and ABTS*- are shown in Figure 1. Although ABTS2- shows no visible peaks, the intensely blue green ABTS'- has maxima at 417,645, and 728 nm. In these determinations, the solutions of ABTS'- were prepared by oxidation with cerium(1V). Extinction coefficients of ABTS*-, which are presented in Table I, are based on the value cited for ABTS- in the preceding paragraph. The 340-nm peak of ABTS2- shifts to shorter wavelengths as the pH decreases. Ultimately, at 1.O M HC104, the maximum liesat310nm. Thisshiftisaresponsetotheacid-baseequilibrium between HABTS- and ABTS-. To evaluate the acid dissociation constant, spectrophotometric titrations at constant ionic strength (1.0 M) were performed at 340 nm. According to the acid dissociation equilibrium (eq 5 ) , the absorbance is related to [H+] by eq 6.
HABTS- = H+ + ABTS*-
K,
(5)
E(UMT1
Figure 3. Cyclic voltammograms at glassy carbon (Ag/AgCI reference) of 5 mM ABTSZ-. Top (a), in 1.5 M HCIO,, sweep rate = 50 mV/a; middle (b), acetatebuffer at pH 4.7, sweeprate 50 mV/s; lower (c), mme as (b) but sweep rate = lo00 mV/s.
pH was varied to obtain independent evidence about the extent of protonation of ABTSZ and ABW-. Cyclic voltammetry of 5 mM HABTS- at a glassy carbon electrode, with 1.5 M Hc10.1 as the supporting electrolyte, gave two reversible one-electron waves, Figure 3a. The f i t wave is the 1 e oxidation of HABTS-. The second arises from oxidation to the doubly-oxidized species ABTSO. The standard potentials obtained from the average of the anodic and cathodic peak positions are ABTS'-
Here and later EAand €6are the molar absorptivitiesof HABTSand ABTS2-,respectively; Cis the total concentration, [HABTS-] [ABTS2-1, and Abs the absorbance measured at a given pH. The nonlinear fit to this model is depicted in Figure 2, giving pK, = 2.08 f 0.02 at p = 1.0 M and ZA = 9.13 X lo3 L mol-' cm-I at 340 nm with CB fixed at 3.66 X lo4 L mol-' cm-l. Electmckdstry. Experiments were conducted to evaluate standard electrode potentials for the relevant half-reactions. The
+
+ e- + H+ = HABTS-
ABTSO + e- = ABTS'-
E , / 2 = 0.81 V (NHE) (7) E , / z = 1.09 V (NHE) (8)
In acetate buffer, pH 4.7, the cyclic voltammogram at low scan rates showed two anodic waves but only one cathodic wave. The peak currents for the first anodic wave and the cathodic wave were the same, provided the voltage scan was rcvcT8od before the second anodic wave was reached, Figure 3b. This pair of
Scott et al.
6712 The Journal of Physical Chemistry, Vol. 97,No.25, 1993
I
I 9
I
6
0
PPM
Figure 5. Proton NMR spectra of ABTS2-/ABTS'in DzO. Top, 0.12 MABTSZ-withexoc88NaS204; middle,0.1 M ABTSs/9.6pM ABTS'; bottom, 0.06 M ABTSG/5.8 pM ABTS'.
to broaden as [ABTS-] increases. This phenomenon shows that the lineshape does depend on the value of kp However, the spectrum is also broadened by decreasing [ABTS-] The simple exchange treatment is no longer valid for this case because the lineshape depends on both species and paramagnetic ABTS'exhibits no observable proton NMR signal. One has to consider both the exchange influence and the paramagnetic influence on the lineshape analysis.lGZ0 Johnsonzoprovided an equation for this situation,
.
Figure 4. EPR spectrum of le M ABTS' in argon-saturated Hz0 at 23 O C . Also shown is the overmodulated spectrum and its double
integration. waves corresponds to the couple ABTS'-
+ e- = ABTS2-
El,2 = 0.68 V (NHE) ( 9 )
When the voltage scan was allowed to proceed to the second oxidation wave, the single cathodic wave increased in size. This is consistentwith rapid comproportionationof the second oxidation product,
+
ABTS' ABTS" = 2 A B T S (10) At high scan rates, as in Figure 3c, the second cathodic wave became visible. The higher rate of comproportionationat higher pH can be attributed to the greater driving force for the reaction (approximately 0.1 V) due to the deprotonation of HABTS-. Cyclicvoltammetry in 0.1 M NaOH gave results virtually identical with those in the acetate buffer. The potentials for the ABTS-/HABTS- and ABTS'-/ABTSZcouples can be combined to give the pKa for HABTS-: pKa = AE0/0.0592 V = 2.2 f 0.3. This agrees with pK,, = 2.08 f 0.02 determined spectrophotometrically. EPR. The ABTSc radical anion exhibits a complex spectrum of overlapping lines because of hyperfine coupling to two sets of two equivalent nitrogens and to six aromatic hydrogens. The spectrum, centered at g = 2.0036 f O.OOO4, is displayed in Figure 4. The spectrum collapses to a broad Lorentzian singlet when overmodulatedor in the presenceof ARTSz-. Attempts to measure the electron exchange rate by quantifying the broadening of the collapsed singlet were thwarted by the limited solubility, -0.12 M, of (NH4)zABTS in aqueous solution. At this concentration, the fast exchange limit has not been attained, and the simple equation relating line width to kw is not valid." ABTSz/ABTS- Electron Exchange. The dynamic NMR experiments were conducted to determine the electron exchange rate constants for ABTS2-/ABTS'-.
1 T,Q
kL41
-3
[ 1 + (4[BJ2/AH2)k2]
where [A] and [B] are [ABTS-] and [ABTW], respectively. With the ABTS*-/ABTS'- system, in order to get a measurable NMRsigna1,onemustkeep [AB='-] low. Sometimesitbecomes too low to determine the concentration accurately. Fortunately, Johnson's equation can be further simplified if more than one NMR signal can be observed. In such a case, with the two sets of nuclei denoted by the superscripts (1) and (2), k is given by
With this equation one can determine the rate constant from measurementsof two sets of NMR signals, the hyperfinesplitting constants (AH) of corresponding nuclei and the concentration of ABTSz-. The TZ measurements were done by the CPMG sequence. This sequence provides the exponential decay of the transverse magnetic moment of the desired nucleus. For a spinspin coupled system
Mxy = M, exp(
-=&) cos(Jrr)
where J is the spin-spin coupling constant. To simplify the measurement, we applied a decoupling frequencyto the methylene group (3.5 ppm) and to one of the aromatic hydrogens H(5) (7.5 ppm) to eliminate the multiplicitieson the methyl group (1 ppm) and on H(7) (7.7 ppm). The transverse magnetic moment decay then becomes
kz
ABTS'-
+ ABTS2- e ABTS" + ABTS-
(1 1)
rate = k,[ABTS&] [ABTS2-] = k,[ABTS2-] (12) The proton NMR spectrum of ABTS- anion (Figure 5 ) begins
The measured TZ'S are listed in Table 11. The hyperfine splitting constants used for methyl group and H(7) were taken from
The Journal of Physical Chemistry, Vol. 97, No. 25, 1993 6713
ABTS Radicals and Ions
TABLE Ik Dynamic N M R Data for the ABTSs/ABTSSystem. [ABSZ-]/
[ABTW/ PM*
0.12 0.10 0.08 0.06 0.05
15 9.6
M
T~=(CH~)/T~CX(H(~))/ kex/ 1 0-2 sc 10-2 sc 1O7 L mol-' s-1 2.8 6.1 3.8 3.5 8.2
1.1 2.6 1.8 1.3 6.5
1.7
5.8 > [HABTS-10,thereaction followed first-order kinetics. The variation of the pseudo-firstorder rate constants with [Fe3+]is linear, as shown in Figure 6. The second-order rate constant, obtained from the slope of the plot depicted, is kl7 = (1.30 f 0.02) X lo2 L mol-' s-l. Under these conditions the reactions went to completion as written, consistent with the standard reduction potentials determined and the concentrations used. Since [H+] = 1.00 M was the only value used, it is assumed but not proved that the reaction partners are those shown in eq 17. The same series of experiments but with added Fez+resulted in a decrease in the absorbance change with a smaller concentration of ABTSO- formed. The kinetic data from these experiments also fit first-order kinetics, with a rate constant that rises with [Fe2+]. Both of these observationsare expected in a reversible approach to equilibrium. The kinetics of reduction of ABTS'by Fe2+were measured in the initial absence of Fe3+,under which conditions the reaction lies entirely to the left. The pseudo-firstorder rate constant varied linearly with [Fez+],resulting in k-17 = 7.5 (*0.3) X lo2 L mol-' s-I; see Figure 6. As before, [H+] was fixed at 1.00 M.
Discpapion The agreement between the molar absorptivitiesand reduction potentials found here and those reported in 1964,16where values wereduplicated, is fully satisfactory. Thesevaluesdiffer,however, from certain determinations reported in 198712as to the molar absorptivities,acid dissociation constants, and reduction potentials. Certain features are deserving of comment. The basicity of a heterocyclic nitrogen is not surprising. The radical derived from oxidation of ABTS2-, largely 'NR3+ in a localized sense, is much less basic. That is, pKaof HABTS' (the protonated radical anion) is
