1178
J. JANATAAND M. B. WILLIAMS
(Po,t
=
$a[Mo]
expkf[Xo]tXexpX2erfcX (AN)
-
0
Reinserting previously defined values for constants and simplifying, one finally obtains
=
4 adt dD kh[L]ada
=
(A201
The current is given by
i
where
=
nFA%,,
(A21)
Oxidation Pathways of 2,2'-Benzothiazolinone Azines. I. Electrochemistry by Jiii Janata and Michael B. Williams P . & P . LabOTatoTy, I.C.I. Ltd., The Heath, Runcorn, Cheshire, England
(Received October 7, 1971)
Publication costs assisted by Imperial Chemical Industries Ltd.
2,2'-Benzothiazolinone azines are oxidized a t a platinum electrode in two steps to the radical cation and dication, respectively. The value of the apparent semiquinone formation constant K' characterizing the stability of the radical cation depends on the acidity of the solution because of the protonation of both neutral azine and of the dication. The dication has been found to decompose autocatalytically yielding the radical cation, ketone, imine, and nitrogen as final products. The decomposition mechanism which is also acidity dependent has been proposed.
Oxidation of 2,2'-benzothiazolinone azines and related compounds has been studied extensively both by chemical' and electrochemical2 methods. Because of some discrepancies between reported electrochemical parameters of these compounds and stabilities of their higher oxidized states w e have reexamined their behavior in media which would not interfere with oxidation processes at more positive potentials. We have observed that the oxidation pattern is profoundly affected by the acidity of the medium and, to a lesser extent, by the concentration of water. All derivatives were found to be sufficiently soluble in aqueous and nonaqueous acetonitrile which offers a relatively wide potential range. Most of the measurements have been done on 2,2'-(3-methylbenzothiazolinone) azine as a model compound. Intermittent checks on the behavior of other derivatives were made to ensure that they follow the same pattern.
Experimental Section All azines studied in this work were synthesized and purified by methods described by Huenig, et ai.' 3Ilkthylbenzothiazolinone hydrazone was prepared according to Besthorn3 and recrystallized from methanol (mp 142"); the supporting electrolyte used in nonaqueous acetonitrile was 0.1 M tetrabutylammonium tetrafluoroborate (TBATFB). The salt itself was prepared by mixing aqueous solutions of ammonium The Journal of Physical Chemistry, Vol. 76, No. 8,1978
tetrafluoroborate with tetrabutylammonium hydroxide and recrystallizing the precipitate twice from 1 :1 ethanol-water mixture. Spectroscopic grade (BDH) acetonitrile was distilled from Pz05and stored over activated molecular sieves Linde Type 4A. All electrochemical experiments were carried out under nitrogen purified over a Cu catalyst and molecular sieves. The nitrogen was presaturated with acetonitrile vapor. The potentiostat used in this work was noncommercial, based on operational amplifier^.^ A Wavetek function generator Model 110 and Tectronix storage oscilloscope lVIodel 549 were used for fast-sweep cyclic voltammetry. The platinum rotating disk electrode (RDE) had a stationary area of 0.020 cm2 (measured by traveling microscope). Because of a small but significant wobble on the shaft, no absolute current measurements have been made with this electrode. Unless stated otherwise the speed of rotation of the electrode was 1320 rpm. The reference electrode for aqueous acetonitrile was Ag-AgC1, in saturated KCI. Silver wire in 0.01 M AgN03 acetonitrile solution served as (1) S. Huenig, H.Balli, H. Conrad, and A. Schott, Justus Liebigs Ann. Chem., 676, 36 (1964). (2) 5. Huenig, H. Balli, H. Conrad, and A. Schott, ibid., 676, 52 (1964). (3) E. Besthorn, Ber. Bunsenges. Phys, Chem., 43, 1579 (1910). (4) E. R. Brown, T. G. McCord, D. E. Smith, and D. De Ford, Anal. Chem., 38, 1119 (1966).
OXIDATION PATHWAYS OF ~,~'-BENZOTHIAZOLINONE AZINES
1179
the reference electrode for nonaqueous studies. All experiments were carried out a t 25.00 =k 0.05'. A Perkin-Elmer spectrophotometer Model 350 was used for optical measurements.
Results Acidity Dependence of the Main Redox System. Reactions dominating the oxidation-reduction pattern of the azines are as follows
Me
Me I
Iv
0
Me
Figure 1. Depencence of half-wave potentials and apparent semiquinone formation constant K' on concentration of H?S04in 3: 1 acetonitrile-water mixture. The concentration of I is 2.46 X M.
c
1-30
m
V
The dissociation constant of the equilibrium I-IV has been determined spectrophotometrically (in 1 : 1 cellosolve-water mixture) as pK1 = 1.8. The second protonation constant K 2 ,which is expected to be close to K1, has not been determined, but it has been observed that the spectrum of the dication changes with acidity of the solution. Thus, the absorption maximum shifts gradually from 515 nm in 0.19 M Hi304 to 505 nm in 1.87 M H2S04. On the other hand, no change in the visible or uv spectrum of I1 has been observed over the range 0.1 to 3 M HzS04. The dependence of the half-wave potentials of the first and second waves and of the logarithm of the apparent semiquinone formation constant K' on concentration of sulfuric acid is shown in Figure 1. Since the definition of p H scale is not valid for the media used in this work, the dependence of half-wave potentials on the acidity cannot be interpreted in the usual waya6 It should be noted, however, that the curves for the half-wave potentials (Figure 1) are symmetrical, which implies that the satme number of protons are involved in both steps. A summary of half-wave potentials and of apparent semiquinone formation constants of 3-methyl-, 3-ethyl-, and 3-ethyl-6-sulfonate derivatives in unbuffered aqueous and nonaqueous acetonitrile are given in Table I. The second wave of 2,2'-(3-ethyl-6-sulfonate benzothiazolinone) azine is poorly developed in both media because of coating of the RDE with insoluble dication2 IV. Oxidation at RDE and Cyclic Voltammetry. Two one-electron oxidation steps should produce two waves
&
Figure 2. Dependence of the ratio of limiting currents on the square root of speed of rotation (in radians sec-I) for (a) nonaqueous acetonitrile, (b) 3: 1 acetonitrile-water mixture, (c) 3: 1 acetonitrile-water mixture, 1 M HzSOd. M. Concentration of I is 5.0 X
of equal height. I n fact the ratio of the two limiting currents, il/iz,changes with experimental conditions. Plots of limiting current of the first wave against square root of speed of rotation (w'/') are linear under all conditions studied which indicates that this oxidation step is diffusion controlled. On the other hand, the ratio in nonaqueous (a) or neutral (b) medium changes with the speed of rotation (Figure 2). I n acid aqueous acetonitrile (e) the ratio remains constant but is less than unity. The ratio of limiting currents is also affected by the concentration of water. It can be seen from Figure 2 curves a and b that the concentration of water determines whether this ratio increases or decreases with w% . Furthermore, the relative position of these two curves depends on the concentration of the depolarizer. There is no doubt that water plays an important role in ( 5 ) L. Meites, "Polarographic Techniques," 2nd ed, Interscience, New York, N. Y., 1965, p 248.
The Journal of Physical Chemistry, Vol. 76, No. 8, 107.9
1180
J. JANATAAND M. B. WILLIAMS
Table I : Summary of Half-Wave Potentials and Semiquin0n.e Formation Constants in Aqueous and Nonaqueous Acetonitrile 7
E"'I/II~ mV
3-Methyl3-Ethyl3-Ethyl-6sulfo-
465 463 550
3 : 1 aoetonitrile-water--
mV
945 950 985 (?)
10-
E"'I/II~
x K'
1.35 1.78 0.22 (7)
+
The Journal of Physical Chemistry, Vol. 76, N o . 8,1978
10
mV
mV
10-8 x
225 210 232
742 723
5.08 4.78
...
' Against Ag-AgC1
a Concentration of water determined by Karl Fischer method is 0.015 M . Against Ag-0.01 M AgN03 0.1 TBATFB reference electrode.
the second oxidation step. However, the mechanism of its action is not explicitly obvious from measurements at a RDE. Both oxidation steps were found to be reversible under our experimental conditions. The difference between Ea/'and E'/d of both waves was between 55.0 and 57.0 mV as predicted in theory for reversible processes. The heterogeneous rate constant determined by cyclic voltammetry'j was 2.8 X 10-2 em sec-' for both waves in nonaqueous acetonitrile. The diffusion coefficient needed for this calculation was determined chron~amperometrically~ as D = 3.3 X em2 sec-'. The cyclic voltammogram of the system is shown in Figure 3. Apart from the two pairs of peaks corresponding to the two main oxidation steps no other oxidation-reduction processes due to the presence of electroactive intermediates were detected by this method. Constant Potential Electrolysis and Characferization of Products. The plot of electrolytic current against charge during constant potential electrolysis at the first wave was linear and had a slope corresponding to one electron as expected. On the other hand, the electrolysis at the second wave yielded a plot which was not linear, the slope increasing steadily as the electrolysis proceeded. The blue radical cation I1 formed during the electrolysis at the first wave was stable while the red dication I11 (or V) produced during electrolysis at the second wave decayed fairly rapidly. This behavior, together with observations made at the RDE, indicates that the second oxidation step is complicated by chemical reactions involving the dication. The first oxidation seems t o be free of any such associated chemical reactions apart from the equilibrium 1-111. AS our studies were made under conditions where both kinematic viscosity of the medium and the liquid junction potential of the reference electrode were liable to change, relating the parameters of the second wave to the first one seems to be more appropriate than expressing them separately. For this reason data given in Figure 2 are expressed as ratios of the two currents even though the variable of interest is the limiting current of the second wave. The only product of constant potential electrolysis a t the first wave was the radical cation 11. Its solution
Nonaqueous acetonitrileG E"*II/III~
c
E'/'II/III~
K'
...
(satd KC1) reference electrode.
-
a6 42-
a
3 1
0 -
h
3
24-
6I
I
I
100
200
300
I
I
400 500 Potentla1(mV,
I
I
I
600
700
000
1
Figure 3. Cyclic voltammogram of I in nonaqueous acetonitrile. Concentration of I is 9.26 X 10-4 M ; the scan rate is 37 mV sec-l.
gave a strong esr signal, but the spectrum was not sufficiently well resolved for further interpretation. The ultraviolet and visible spectra of the I1 were identical with the ones described in the 1iterature.l The solution after electrolysis at the second wave was initially red and gave a visible spectrum corresponding to the dieation.' This solution gradually turned blue and the visible and uv spectrum showed increasing concentration of the radical cation. Besides that, there were characteristic absorption peaks at 287, 279, and 240260 nm gradually increasing in intensity as the dication spectrum faded away. A number of relevant species, namely 3-methyl-2-benzthiazol~ne~ (VI), the protonated form of 3-methyl-2-benzothiazolinone imine (VII), and the protonated form of 3-methyl-2-benzo(6) R. S. Nicholson, Anal. Chem., 37, 1351 (1965). (7) I. Shain and K. J. Martin, J . Phys. Chem., 65, 254 (1961). (8) G. De Stevens, A. Frutchey, A. Halamandaris, and H. A. Luts, J . Amer. Cham. Soc., 79, 5263 (1957).
OXIDATION PATHWAYS
OF
1181
2,2'-BENZOTHIAZOLINONE AZINES
thiazolinone hydrazone9 (VIII) have this type of spectrum. As the dissociation constants of VI1 and VI11 he
5
MI he
VI
IkH+
he VIII
w
were determined spectrophotometrically as pK = 8.38 and 5.93, respectively, these two species were present in their protonized form in most of our experiments. Another conceivable intermediate or a final product could be 3-methyl-benzothiazolium ion (IX).
as+ 'y+ CH3
Ix There was, however, no evidence in the uv spectrum*0 for the presence of this species in any significant concentration. Neither were there lines present in the mass spectrum of the mixture of final products which would indicate its formation. Attempts to isolate and characterize reaction intermediates failed. The separation after electrolysis by paper or thin-layer chromatography or by electrophoresis always resulted in formation of a diffused zone. The major constituents of this zone have been characterized as the radical cation 11, ketone VI, and imine VII. On a large scale, using Ce(SO& to produce V, VI and VI1 (unprotonized) were extracted as final products and characterized by ir, mass spectroscopy, and nmr spectroscopy. I n addition an appreciable amount of nitrogen was also produced in these experiments. The decay of the dication was found to be dependent on the acidity of the medium, strongly acid solutions being more stable. It is evident that the dication decomposes to intermediates which can, under certain circumstances, cause a further reduction of the dication into the radical cation. The overall stability of the dication would. then depend both on its rate of decomposition and on the oxidation-reduction properties of the intermediates. No oxidation of VI and VI1 has been observed at potentials below +1.200 V. It has been shown by Huenig and Ballill that the oxidation potential of hydrazone VI11 depends on the acidity of the solution and that VI11 can participate12 in the oxidation mechanism of I. We have, therefore, investigated the electrochemical behavior of VIII. Oxidation of 3-Methyl-2-benzothiazolinone Hydrazone at RDE. It has been shown'l that the hydrazone can be oxidized a t a platinum electrode in a single two-electron step to an unstable intermediate XI. Any further oxidation of XI could not be followed in phosphate-'
Me XI
Me
X
citrate buffers used by these authors because of the oxidation of the background electrolyte at potentials above +SO0 mV. The first half-wave potential in aqueous acetonitrile was identical with the one measured by Huenig, et al.," and also shifted to more positive potentials with increasing acidity of the medium (Figure 4,curve a). It was followed by further oxidation barely distinguishable on a poorly developed dc voltammogram (Figure 5, curve a). However, it can be seen quite clearly on the derivative voltammogram (Figure 5, curve b). The dependence of the second peak potential on the acidity is shown in Figure 4, curve b. The overall appearance of the two waves and corresponding cyclic voltammograms indicate that both oxidation steps are completely irreversible. No del
I
I
t"l.1.'
I
I
I
I
H m
Figure 4. Dependence of peak potentials (from derivative voltammograms) of oxidation of 3-methylbeneothiaeolinone hydrazone on concentration of HzSOa. (9) K. Soda, Agr. Biol. Chem., 31, 1054 (1967). (10) J. Metzger, H. Larive, R. Dennilauer, R. Boralle, and C. Gaurat, Bull. SOC.Chim. Fr,, 2857 (1964). (11) 8. Huenig and H. Balli, Justus Liebiga Ann. Chem., 628, 56
(1959). (12) R. A. Bartsch, S. Huenig, and H. Quast, J . Amer. Chem. Soc., 92, 6007 (1970). The Journal of Physical Chemistry, Vol. 76, No. 8,1979
1182 I
.I
J. JANATA AND M. B. WILLIAMB I
I
I
I
I
I
I
I
I
I
(
the protonation equilibria I-IV and III-V. The electrode potential for the first one-electron oxidation step is given by the surface concentrations [I10 and 11110
24
20
The dissociation constant of IV is
([H+l= [H+]o) and the total surface concentration of the reduced form is
c1 = PdmtblbV)
Figure 5. Dc voltammogram (a) and derivative voltammogram (b) of 3-methylbenzothiazolinone hydrazone in 3: 1 awtonitrile-water mixture, 0.5 M HSO4. Concentration of VI11 is 4.1 X lo-' M .
finitive conclusions can be made about the nature of the product of the second oxidation step on the basis of these observations. Nevertheless, it could be an unstable diazonium cation XII, the idea, previously rej e ~ t e d " - ~ *mainly on the grounds of the lack of any electrochemical evidence of further chemical oxidation of XI. Nonetheless, whatever this oxidation product
L
XII
I110
+ [IVIO
(4)
Combination of eq 2-4 and introduction of mass transport equations in the usual waylS yields an expression for the half-wave potential
For H+>> KI the half-wave potential will shift towards more positive values with increasing concentration of hydrogen ions. Similarly the expression for the half-wave potential of the second one-electron step is
For H+>> Kz the half-wave potential will shift towards less positive potentials with increasing concentration of hydrogen ions. Subtraction of eq 6 from eq 5 yields
E'"II,III- E'"I/II= E'II/III- E'I/II
+
is, it is quite important that the oxidation of VI11 does occur in two steps and that these steps are electrochemically irreversible.
AE'"
Discussion
In the absence of any protonation the constant K de-
The effect of acidity on the oxidation-reduction pattern of 2,2'-benzothiazolinone azines falls into two categories. The first one includes protonation equilibria I-IV (K1) and III-V (Kz) and affects the stability of the radical cation I1 through the apparent semiquinone formation constant K'. The second one determines the rate of decomposition of the fully oxidized form I11 (or V) and has an effect on the overall chemical stability of the system. Let us assume that, in the first approximation, all the oxidation steps are stable. The equilibrium concentrations of unprotonized forms are given by the semiquinone formation constant which is defined as
fined by eq l can be expressed as16
=
EO^^,^^^
- EO^^^^
=
F
In K
thus AE'~'=
2.303RT log F
+
2.303RT F
Defining K' as the apparent semiquinone formation constant the final expression for is
a'/'
(13) S. Huenig, Angew. Chem., Id.Ed., 7 , 335 (1968).
Oxidation of I at a platinum R D E then gives two waves, separation of which depends on the value of K and on The Journal of Physieal Chemistry, Vol. 76, No. 8, 1079
(14) 8. Huenig, J . Chem. Educ., 46, 734 (1969). (15) Reference 5, p 209. '(16) R. Brdicka, Z. Elektrochem., 47, 314 (1941).
OXIDATION PATHWAYS OF 2,2’-BENZOTHIAZOLINONE
AZINES
1183
creasing acidity of the solution. This means that the reduction of V to I11 in sufficiently acid solutions should not take place and V would decompose via the first where route only. It has in fact been observed that the solution of V containing more than 1.5 M HzS04 decomposes without formation of I1 although this decomposition takes a rather long time. Reduction of V to I1 It follows from eq 11 that K’ will decrease with independs ultimately on the relative value of the standard creasing acidity of the solution, and in the limiting case potential of the second oxidation step of azine with rewhen H+ >> K1, K2the rate of change of log K’ with the spect to the standard potential of the second oxidation acidity should be twice that for E’/?IIIIor E’/’IIIIII. step of the hydrazone in a given medium. While the In fact, slopes of linear extrapolations of curves in Figfirst potential is represented by the half-wave potential ure 1 are AE”’/A[H2S0~]= 0.9 and AAE”’/A[H2S04] &“”II,III, the standard potential of the second hy= 1.8. I n another limiting case when H+
