Redox properties and bond dissociations energies of phenoxyl radicals

Electrochimica Acta 49 (2004) 537–544

Redox properties and bond dissociations energies of phenoxyl radicals Günter Grampp∗ , Stephan Landgraf, Claudia Mure¸sanu1 Institute of Physical and Theoretical Chemistry, Graz University of Technology, Technikerstraße 4/I, A 8010 Graz, Austria Received 14 May 2003; received in revised form 8 August 2003; accepted 6 September 2003

Abstract Reduction half-wave potentials of several phenoxyl radicals have been measured by photomodulated voltammetry in acetonitrile. Two of the investigated radicals exhibit a reversible behavior, but for the others a rather quasi-reversible nature of the heterogeneous electron transfer has to be considered. The measured reduction half-wave potentials are within 70 mV of reported values for formal one-electron potentials of the phenolate–phenoxyl couples. That is why one can assume that the values of half-wave reduction potentials are close to the standard potentials of the electrochemical reactions. Differences of the bond dissociation energy (BDE)(O–H), for the substituted phenols relative to the BDE(O–H) of the phenol were calculated and compared with corresponding data in the gas-phase. © 2003 Elsevier Ltd. All rights reserved. Keywords: Photomodulated voltammetry; Half-wave potentials; Phenoxyl radicals; Relative bond dissociation energies

1. Introduction Phenols were intensively studied because they are involved in many chemical, electrochemical, photochemical or biochemical electron-transfer reactions [1–5] and their acid–base equilibria are widely used as reference data in establishing linear free energy relationships [6–8]. Beside their antioxidant properties [9], they also are important environmental pollutants [10]. Less is known about the redox properties of the phenolate–phenoxyl radical couples. Informations about the electrochemical potentials of several phenolate–phenoxyl radical couples in water, were obtained by pulse radiolysis [11,12] and by cyclic voltammetry [13], but only a few data are available in organic solvents. Direct measurements of the one-electron oxidation potential of some phenolate–phenoxyl couples in acetonitrile, by fast linear sweep voltammetry on ultra-microelectrodes [14] and the determination of the oxidation potentials of several phenoxide and napthoxide ions in DMSO by cyclic voltammetry [15] are reported in the literature. The formal potential of ∗ Corresponding author. Tel.: +43-316-873-8220; fax: +43-316-873-8225. E-mail addresses: [email protected] (G. Grampp), [email protected] (C. Mure¸sanu). 1 Co-corresponding author. Present address: Dept. of Chemistry, Babes Bolyai University, Arany Janos 11, RO 3400 Cluj-Napoca, Romania. Tel.: +40-264-593833; fax: +40-264-590818.

0013-4686/$ – see front matter © 2003 Elsevier Ltd. All rights reserved. doi:10.1016/j.electacta.2003.09.007

phenoxyl radicals derived from sterically hindered phenols could be determined, in acetonitrile, by cyclic voltammetry [16,17]. In this paper, we present the determination of redox potential of various phenoxyl radicals in acetonitrile by photomodulated voltammetry (PMV), a technique successfully employed for the determination of redox potentials of arylmethyl [18,19], aminyl [20] and phenylthiyl radicals [21]. 2. Experimental 2.1. Materials 2.1.1. Photomodulated voltammetry Phenol (Merck) (1), 4-methylphenol (Merck) (2), and 2,6-di-t-butylphenol (Fluka) (5) were distilled under vacuum before use, 4-nitrophenol (Aldrich) (3) and 2,4,6-tri-t-butylphenol (Aldrich) (7) were purified by sublimation, while 3,4-dimethylphenol (Merck) (4), 2,4,6-trimethylphenol (Aldrich) (6), 4-hydroxyl-2,5-di-tbutylphenol (Fluka) (8) and 2-hydroxy-3,5-di-t-butylphenol (Aldrich) (9) were used without further purification. Acetonitrile came from Merck and was passed, together with di-t-butylperoxide (DTBP, Aldrich), over a column of activated aluminum oxide before each experiment. The supporting electrolyte, tetra-butylammonium tetrafluoroborate (Bu4 NBF4 ; Fluka), was purified by the following procedure:

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the electrolyte was dissolved in methylene chloride, the solution was dried overnight with CaCl2 and filtered. The solvent was completely removed by distillation and the residue dissolved in a small amount of ethyl acetate. After that the solid Bu4 NBF4 was precipitated by addition of ether and filtered from solution. Finally, the electrolyte was dried under vacuum (60 ◦ C) and kept under argon. 2.1.2. Determination of pKa Phenols came from the same sources and were purified like above. Picric acid (Merck) was recrystallized from water. Acetonitrile (Merck) was dried over molecular sieve, 3Å, distilled and kept under argon. Ethyl acetate (Merck), methanol (Merck) and tetra-n-butylammonium hydroxide (Fluka) were used without further purification. Tetra-n-butylammonium phenolates were prepared by mixing appropriate amounts of the corresponding phenol and a methanolic solution of tetra-n-butylammonium hydroxide, removing the solvent under vacuum and recrystallizing twice from ethyl acetate. It was not possible to prepare phenolates from 2,4,6-tri-t-butylphenol and 2-hydroxy-3,5-di-t-butylphenol with this method, probably because of steric hindrance effects. 2.2. Instrumentation 2.2.1. Photomodulated voltammetry Instrumentation and method are similar to those described by Wayner et al. [18]. The Teflon cell was designed and built at the University of Aarhus [21] and has a cylindrical chamber (φ = 9 mm, L = 6 mm). The working electrode consists of a gold net (minigrid, 1000 mesh, Goodfellow) sealed in plastic foil, with a hole (φ = 7 mm) in order to ensure contact between net and solution. Electrical contact to the working electrode was realized with a isolated platinum wire. The light beam, from the lamp reached the working electrode through a quartz window, which was one of the walls of the chamber. A steel counter electrode was placed in the outlet of the cell, in order to avoid interferences from products formed at the electrode. As reference electrode we used a Flex-Ref electrode (World Precision Instruments), placed behind the working electrode and close to it, to achieve a low Ohmic drop. Light source (500 W Osram Hg lamp), power supply, IR-water filter, light chopper (home made), a quartz lens and the electrochemical cell were placed in line on a optical bench. The instrumentation included also a voltage scan generator (VSG72, Wenking), a potentiostat (68FR0.5, Wenking), a lock-in amplifier (186 Synchro-Het, PAR) for the detection of the ac component of the current and a X–Y recorder (7045A, Hewlett Packard). The instrumentation did not include any device for the compensation of Ohmic drop. 2.2.2. Determination of pKa Potentiometric measurements were performed with a WTW pH-meter/voltmeter, a glass electrode (Ingold) and a Flex-Ref electrode (World Precision Instruments) as the

reference electrode. The glass electrode was kept in a 3 M aqueous solution of KCl, when it was not in use. Before each experiment it was soaked for 20 min in acetonitrile. The reference electrode was introduced in a 0.1 M solution of tetra-n-butylammonium perchlorate in acetonitrile separated from the phenolate solution by a G4 fritted glass disk. The phenolate solutions in acetonitrile were degased for 30 min prior to the measurements. Experiments were undertaken under argon in a cell provided with a water jacket connected to a recirculatory bath, which kept the temperature at 25 ± 0.1 ◦ C. 2.3. Procedure 2.3.1. Photomodulated voltammetry To generate the phenoxyl radicals, an indirect method was used, by photolyzing di-t-butylperoxide to tert-butoxyl radicals, which abstract a hydrogen atom from the phenol in a subsequent fast reaction: hv

t-BuOOBu-t − → 2t-BuO• k

→t-BuOH + Ph–O• t-BuO• + Ph–OH−

(1) (2)

The frequency of the light chopper was 120 Hz and the sweep rate was varied from 2.5 up to 50 mV/s. Solutions were flowed slowly through the cell, after carefully removing oxygen by bubbling argon through. Potentials were recalculated relative to the aqueous standard calomel electrode, by means of the ferrocene/ferrocenium couple (0.41 V/SCE in CH3 CN [22]). For every compound, three–five voltammograms were recorded at each scan rate. 2.3.2. Determination of pKa The glass electrode was calibrated in four picric acid–tetra-n-butylammonium picrate buffers, by adding appropriate volumes of a 0.01 M solution of picric acid in acetonitrile to 25 ml of picrate solution in acetonitrile (0.001, 0.004, 0.008 and, respectively, 0.01 M). The paH values were determined based on the emf of the cell containing equimolar solutions of the tetra-n-butylammonium phenolate and phenol. Five determinations of the paH were carried out for each phenolate in the range of 0.001–0.01 M. 3. Results and discussion 3.1. Half-wave potentials Photomodulated voltammograms of the reduction of the following phenoxyl radicals were recorded: phenoxyl (1• ), 4-methylphenoxyl (2• ), 4-nitrophenoxyl (3• ), 3,4-dimethylphenoxyl (4• ), 2,6-di-t-butylphenoxyl (5• ), 2,4,6-trimethylphenoxyl (6• ), 2,4,6-tri-t-butylphenoxyl (7• ), 4-hydroxy-2,5-di-t-butylphenoxyl (8• ), 2-hydroxy-3,5-di-tbutylphenoxyl (9• ). The voltammograms of 2• , 5• and 7• are represented in Figs. 1–3.

G. Grampp et al. / Electrochimica Acta 49 (2004) 537–544

with the in-phase setting (φ = 0◦ ), for 2 and 6 there was no more than 10 mV difference between the potentials determined with in-phase and the quadrature (φ = 90◦ ) setting, while for the rest this difference could reach up to 100 mV. The scan rate dependent shift of the half-wave potential is difficult to assign because the effects of Ohmic drop are similar to those due to a slow electron transfer reaction:

0

red

E 1/2

i / nA

-200

539

+e−

-400

PhO• −→ PhO−

-600

or coupled chemical reactions, like the preceding reaction (Eqs. (1) and (2)), or the reversible dimerization of the phenoxyl radical to a p-quinol ether, coupled with its photochemical depletion:

-0.2

0.0

0.2

0.4

0.6

0.8

E / V vs.SCE Fig. 1. Photomodulated voltammogram of 4-methylphenoxyl radical in acetonitrile with 0.08 M Bu4 NBF4 added. Sweep rate: 5 mV/s.

Although all the voltammograms were S-shaped, the red − Ered |, were different as exwidth of the waves, |E3/4 1/4 red −Ered | = pected for a one-electron Nernstian process (|E3/4 1/4 56.4 mV), indicating a rather sluggish electron transfer at the electrode. A common feature of all the radicals is the influence of the scan rate on the position of the half-wave potential. As the scan rate increased the potentials shifted either towards positive direction for an anodic scan, or towards negative direction for a cathodic scan. With increasing scan rate the waves lost their original S-shape and in some cases they disappeared completely at scan rates higher than 50 mV/s. The influence exerted by the phase settings of lock in amplifier, on E1/2 depends on the nature of the radical. For 1 and 3 the wave became visible only

k

2PhO•
dimer hv

(3)

(4)

According to de Tacconi et al. [23], the dependence of the peak potential on the uncompensated resistance RU , for the simple electron transfer: X−
X + e−

(5)

is given by:


RT RT βF k Ep = + ln ln v + IRU βF RT k1 βF

(6)

where R, T, β and F have their usual meaning, k1 is the rate constant of the forward electron transfer, k corresponds to the charge required to obtain a monolayer of adsorbed X, v the sweep rate and I the peak current. The current is also dependent upon the sweep rate, that is why a Ep versus ln v plot will be a curve changing to a straight line only if RU is very small or equal to zero. A similar shape of Ep versus ln v plot is obtained when follow up reactions, like a

Fig. 2. Photomodulated voltammogram of 2,6-di-t-butylphenoxyl radical in acetonitrile with 0.08 M Bu4 NBF4 . Sweep rate: 10 mV/s.

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G. Grampp et al. / Electrochimica Acta 49 (2004) 537–544

0

-200

-400

-600

-800

-1000 -0.7

-0.6

-0.5

-0.4

-0.3

-0.2

-0.1

0.0

0.1

0.2

Fig. 3. Photomodulated voltammogram of 2,4,6-tri-t-butylphenoxyl radical in acetonitrile with 0.08 M Bu4 NBF4 . Sweep rate: 10 mV/s.

dimerization, are involved. The horizontal part of the plot corresponds to the reversible behavior, while the part corresponding to pure kinetic conditions is linearly increasing with ln v, according to [24]:
 RT RT v bF + E0 + ln (7) Ep = a F bF kC0 cRT where the values of the constants a, b and c depends on the dimerization mechanism, k is the rate constant of the forward reaction and C0 the initial concentration of the reactant. We considered that it is possible to extrapolate the considerations mentioned above to explain the E1/2 = f (ln v) dependence from photomodulated voltammograms. A typical half-wave scan rate dependence is illustrated in Fig. 4. As it can be seen, the potential tends to a constant value at

low scan rates. This behavior was typical for all the investigated radicals, which of course seems to put the involvement of dimerization under question, because hindered phenols, like 2,4,6-tri-t-butylphenol do not dimerize [25]. The conred stant value of the half-wave potential of the radicals, ∗ E1/2 was considered from the plateau of the E1/2 = f (ln v) dependence, obtained for in phase setting, when the potential was scanned in cathodic direction. The potentials located on the plateau should be less influenced by the combined effects of uncompensated resistance and slow electrode kinetics or coupled chemical reactions. The E1/2 = f (ln v) dependence was fitted with an asymptotic equation to emphasize the shape of the curve and to identify the position of the plateau. The constant values of the half-wave potentials, red −Ered | and available thus obtained, the corresponding |E3/4 1/4

-0.05

(a)

-0.10

-0.15

(b)

-0.15 -0.20 -0.25

-0.20

-0.30 -0.35

-0.25

Cathodic scan Anodic scan

-0.40 -0.45 -0.50

-0.30

-2.6

-2.4

-2.2

-2.0

-1.8

-1.6

-1.4

log v

-0.35 -0.40 -2.6

-2.4

-2.2

-2.0

-1.8

-1.6

-1.4

Fig. 4. Influence of the scan rate, v, on the half-wave potential of the 2,6-di-t-butylphenoxyl radical in acetonitrile: (a) in phase setting of in lock in amplifier and cathodic scan; (b) quadrature setting of the lock in amplifier.

G. Grampp et al. / Electrochimica Acta 49 (2004) 537–544

literature values of E0 in acetonitrile and in water are summarized in Table 1. In phase setting was preferred in order to avoid possible influences from unwanted radicals as it is possible when an indirect method is used. The scan direction has less influence for the determination of E1/2 as it can be seen from the inset of Fig. 4. Only for two of the investigated radicals one can consider a reversible behavior, namely for the 2,4,6-tri-t-butylphenoxyl, which is known to be reversible, and for the 4-hydroxy red − Ered | values of the 2,5-di-t-butylphenoxyl. The |E3/4 1/4 others indicate rather a sluggish heterogeneous electron transfer. Which means it is not possible to assume from the beginning, that E1/2 is equal to thermodynamically PhOH(sol) → PhOH− (sol) + H+ (sol) •

PhOH− (sol)
PhO (sol) + e− +

−

H (aq) + e
1 2 H2 (g)

negative value of E0 for 8• , comparative to that of the other radicals, may be attributed to the strong electron-donating effect of the OH group in the para position. 3.2. Bond dissociation energies A great number of homolytic bond dissociation energies (BDE) were estimated by means of the electrochemical (EC) method developed by the group of Bordwell et al. [17,26,27]. This method is based on a thermochemical cycle approach originally suggested by Eberson [28], which combines pKa values and electrode potentials to determine gas phase bond dissociation energies: G1 = 2.303RTpKa − FE0

H+ (sol) → H+ (aq)

G2

1 2 H2 (g)

− FE

•

541

(9) (10) (11)

00

(12)

→ H (g)

G3

(13)

PhOH(sol) → PhOH(g)

G4

(14)

G5

(15)

•

•

PhO (sol) → PhO (g) •

•

PhOH(g) → PhO (g) + H (g)

G = 2.303RTpKa + FE0 + G2 − FE00 + G3 + (G5 − G4 )

significant value E0 , for all the couples. Fortunately, it is possible to compare some of our results with those obtained by Hapiot et al. [14] by means of fast scan voltammetry. As it can be seen from Table 1, there is no larger difference than red determined by photomodulated 70 mV between the ∗ E1/2 voltammetry and the potentials reported in the literature, despite of peak widths larger than expected for a Nernstian electron transfer. That is why we assumed that our values are close to the standard potential of the electrochemical reaction within 70 mV. red values in acetoniComparing the E0 , respectively, ∗ E1/2 trile with E0 values in water one can notice the same trend. As the E0 correlates with the free solvation energies of the phenolate ion and the phenoxyl radical according to: G0sol (PhO− ) − G0sol (PhO• ) (8) F where IP is the gas-phase ionization potential of the anion and C a constant that depends on the reference electrode, one can assume that the more positive values of E0 in water than in acetonitrile originates in the different solvation capacity of the anion and the radical in this two solvents. Substituents with strong electron-withdrawing, effect like NO2 , shift the E0 to more positive values comparative to the unsubstituted phenol, while electron-donating groups exert an opposite effect. The latter effect is of course enhanced by the number of the electron-donating groups as one can see by comparing the E0 of 4• with that of 6• and that of 5• with 7• . The large,

IP = C + E0 +

(16)

where E00 means the potential of the normal hydrogen electrode. The bond dissociation energy in the gas phase may be expressed from Eq. (16) as: BDE(O–H)gas = 2.303RTpKa + FE0 + C

(17)

where the constant is: C = G2 + G3 + (G5 − G4 ) − FE00 + TS

(18)

Values of the constant C may be estimated by means of thermochemical data from the literature and the values thus found were successfully employed to determine BDE(O–H) of phenols [12] and BDE(N–H) of anilines [29] in water, or BDE(O–H) of phenols [15], benzylic BDE(C–H) [30] and BDE(S–H) of thiophenols [27] in DMSO. An usual approach to obtain the constant C, is to consider that the solvation free energies of the radical and the neutral molecule cancel each other. In our case, this approximation would lead to a less precise estimation of C, because it is very likely that the phenol is hydrogen bonding with acetonitrile and it can no longer be considered that G5 ≈ G4 . To overcome this situation we preferred to calculate the difference of the BDE(O–H) for the substituted phenols relative to the BDE of the unsubstituted phenol: BDE(O–H) = 2.303RTpKa + FE0

(19)

Literature data for pKa values in acetonitrile are known only for phenol (pKa = 26.6 [31]) and 4-nitrophenol

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G. Grampp et al. / Electrochimica Acta 49 (2004) 537–544

Table 1 Half-wave reduction potentials for some phenoxyl radicals, measured by photomodulated voltammetry at a gold electrode in acetonitrile with 0.08 M Bu4 NBF4 Radical

∗ Ered 1/2

1• 2• 3• 4• 5• 6• 7• 8• 9•

+0.159 +0.052 +0.668 +0.032 −0.151 −0.205 −0.295 −0.374 −0.061

(V vs. SCE)

red red |E3/4 − E1/4 | (mVc )

E0 (V vs. SCEa )

E0 (V vs. SCE in water)

139 174 100 171 133 160 89 86 134

+0.220b +0.038b +0.74

+0.548d ; 0.618e +0.438d ; +0.338e +0.978d ; +1.038e

−0.193 −0.200 −0.289

a Formal one-electron potential of phenolate–phenoxyl couples, from [14], derived from fast scan voltammetry in acetonitrile, as the midpoint between the anodic and the cathodic peak potentials. b From [14], estimated from E0 vs. σ + Hammett plots. c Scan rate 10 mV/s. d From [12], determined by pulse radiolysis. e From [13] estimated from cyclic voltammetry measurements.

(pKa = 20.7 [31]), for the other phenols only the pKa values in water are available (see Table 2). The paH values in acetonitrile were determined by means of the calibration lines obtained from the emf measurements described in the experimental part, at different ionic strength, µ, and assuming that the pKaAN of picric acid is 11.0 [34]: E = −59.675paH + 1033.7 E = −59.913paH + 1055.3 E = −60.878paH + 1066.5 E = −61.447paH + 1069.2

µ ≈ 0.001 M µ ≈ 0.004 M µ ≈ 0.008 M µ ≈ 0.01 M

values of the phenols under study were calcuThe lated from the equation: (20)

where the logarithm of the activity coefficient is given by the Debye–Hückel equation: log fi = −355z2i D−3/2 µ1/2 (the dielectric constant D = 36.0)

(21)

From these measurements we obtained for the phenol the pKa = 26.55±0.15 and for the 4-nitrophenol pKa = 20.5± 0.2 close to the literature values [31]. Table 2 The pKa values of the investigated phenols in water H O

Phenol

pKa 2

Ref.

1 2 3 4 5 6 7 8 9

9.89 10.17 7.15 10.36 11.7 10.86 12.19 – 10.29

[32] [32] [32] [32] [33] [32] [33] –

a

Estimated by Eq. (24).

pKaAN = 1.9 pKaH2 O + 8.3

a

(22)

Instead of pKa s we have calculated the differences relative to the phenol, so that relation (22) became: pKaAN = 1.9 pKaH2 O

pKaAN

pKaAN = paH − log fi

Magonski et al. [35] determined the pKa values of several phenols in organic solvents and they found that the pKa in acetonitrile and in water are linearly correlated:

(23)

BDEs determined by Eq. (19), the pKa s from Eq. (23) and experimental values of pKa obtained by us along with literature data for the gas phase are summarized in Table 3. For 4-hydroxy-2,5-di-t-butylphenol and 2-hydroxy-3,5di-t-butylphenol not even the pKa values in water could be found, so we had to estimated first this values by means of a Hammett type relation [38]:  pKa = 9.94 − 2.23 σ+ (24) As it can be seen, the linear correlation between pKa values in acetonitrile and water is satisfactory for many of the investigated phenols with the exception of hindered phenols like 4-hydroxy-2,5-di-t-butylphenol or 2,6-di-t-butylphenol. There is a difference between the BDEs obtained in solution by the EC approach and data from the gas phase, as it can be seen in Table 3. This may be due to solvent effects, especially to different hydrogen bonding between phenols and solvent. This supposition is very likely, because differences between BDEs of phenols, determined in DMSO or sulfolane, by EC method and BDEs in the gas phase were also noticed [36]. The problem of solvent effects is complicated especially because some of the investigated phenols have large substituents placed in the vicinity of the phenolic OH group, causing steric hindrance, phenols 8 and 9 possess to OH groups and large substitutents. That is why we would like to make the influence of solvent effects to the subject of another paper.

G. Grampp et al. / Electrochimica Acta 49 (2004) 537–544

543

Table 3 Homolytic dissociation bond energies of the investigated phenols Phenol

pKaAN a

pKaAN b

Ered c (V)

BDE(O–H)d (kJ/mol)

BDE(O–H)gas e (kJ/mol)

2 3 4 5 6 7 8 9

0.72 −6.15 1.04 −1.7g 1.8

0.34 −5.95 0.76 3.25 1.65 4.18 – 0.77

−0.107 0.509 −0.127 −0.310 −0.364 −0.477 −0.533 −0.220

−6.22 14.05 −6.32 −39.59 −24.85 −21.11h −72.79 −16.89h

−8.00 ± 25.00 ± −14.71f −26.00 ± −23.00 ± −37.00 ± −35.71f −23.69f

−3.75

4 8 8 4 4

a

Experimental values. Calculated from Eq. (23). c Obtained from data presented in Table 1, by subtraction of the Ered of phenol from the Ered of the substituted phenol. d Calculated by Eq. (19) with experimental values of pK . a e From [36].  + f Estimated by means of a Hammett type relation: BDE(O–H) = (28.31 ± 0.91) σ − (3.11 ± 0.59), from [36]. The σ0+ for substituents in the + + ortho position was estimated as σ0 = 0.66σp , from [37]. g This value is not so reliable because it was not possible to purify this phenolate as many times as the other. h Calculated by Eq. (19) with pK values from column 3. a b

formal potentials [14] reported from fast scan voltammetry measurements. The dependence of the half-wave potentials from the scan rate may be attributed to the quasi-reversible character of the PhO• /PhO− electron transfer or to the involvement of preceding reactions. Values of the O–H bond dissociation energies relative to phenol were calculated by the EC method and compared with similar data from the gas-phase.

4. Conclusions The reduction potentials of phenoxyl and several substituted phenoxyl radicals were determined by photomodulated voltammetry in acetonitrile. Although only two of the investigated radicals showed a reversible behavior according to red −Ered | criterion, the half-wave reduction potentials the |E3/4 1/4 of four of the radicals are within 70 mV of the one-electron

OH

1

OH

OH

OH

CH3

NO2

2

3

OH

OH

CH3

H3C CH3 CH3

CH3 4

5

OH

OH

6 OH OH

OH 7

8 Scheme 1. .

9

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G. Grampp et al. / Electrochimica Acta 49 (2004) 537–544

Acknowledgements We want to thank the Austrian Federal Ministry for Education, Science and Culture and the VW-Foundation (Germany) for financial support. We gratefully acknowledge Professor Steen Uttrup Pedersen from the Department of Chemistry, University of Aarhus (Denmark), for the electrochemical cell and for very helpful discussions.

[14] [15] [16] [17] [18] [19] [20] [21]

Appendix A

[22]

Scheme 1 shows the structure of the investigated phenols. [23]

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[24]

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