The electrochemical oxidation of hydroquinone ester derivatives in aqueous solutions

ELECTROANALYTICALCHEMISTRYAND INTERFACIALELECTROCHEMISTRY Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

T H E E L E C T R O C H E M I C A L O X I D A T I O N OF H Y D R O Q U I N O N E DERIVATIVES IN AQUEOUS SOLUTIONS*

409

ESTER

E. P. MEIER, J. Q. CHAMBERS**, C. A. CHAMBERS, B. R. EGGINS ANDC.-S. LIAO

Department of Chemistry, University of Tennessee, Knoxville, Tennessee 37916 (U.S.A.)

(Received 8th March 1971; in revised form 10th June 1971)

INTRODUCTION Hydroquinone esters have received a great deal of attention in recent years as model compounds for the conservation of oxidative energy in biological systems. Mechanistic details of the phosphorylation processes which f o r m A T P in the biological respiratory chain or in photosynthesis are lacking and efforts to identify hypothetical energy-rich intermediates have not met with success. Hydroquinone esters and their oxidized forms have been proposed as intermediates in these processes and they present several attractive features as model compounds. Biological quinones with methyl groups in the 2-position, i.e. the vitamin K's, the ubiquinones, and the tocopherol quinones, are known to be implicated in oxidative phosphorylation. Ester derivatives of model compounds are oxidized at potentials compatible with the biological electron transport chain. Finally, oxidation by a variety of methods produces reactive intermediates which readily form condensed phosphate species. Research in these areas has been reviewed recently 1. These species are also attractive potential intermediates in these processes if the driving force for their oxidation or formation is an electrochemical membrane potential as in the chemiosmosis theory of Mitchell 2 or a conformational change in the hypothesis due to Green 3. In this paper the results of an extensive voltammetric study of several hydroquinone esters are summarized. Previous reports from this laboratory have described the electrochemical oxidation of two compounds in the series in some detaiP -6 This work has been extended to the five hydroquinone phosphate esters and one sulfate ester given below. The qualitative features of the electrochemical oxidation of these compounds are similar and results will only be summarized here. Details are given elsewhere 7,s. EXPERIMENTAL The structures of the compounds studied in this work are given below. The phosphate esters were a generous gift from Hoffman LaRoche, Inc. The diphosphate

* Taken in part from the Ph.D. Thesis of E. P. Meier, University of Colorado, 1969 and the M.S. Thesis of C.-S. Liao, University of Colorado, 1969. ** Address correspondenceto this author. J. Electroanal. Chem., 33 (1971) 409-418

410

E. P. MEIER, J. Q. CHAMBERS, C. A. CHAMBERS, B. R. EGGINS, C.-S. LIAO

ester (6) was obtained as the dihydrate by drying in vacuum 9. One sample of the 2,3dimethylnaphthohydroquinone derivative contained 17~o ethanol of recrystallization. The compounds were analyzed by titration with standard base prior to use. All compounds analyzed for better than 99 ~ purity, except for the phosphate ester of benzohydroquinone (4). Compound (4) had a 5 ~ electroinactive impurity, probably water of hydration. The potassium salt of the sulfate ester of benzohydroquinone (5) was prepared by the method of Yamaguchi 1°. It was recrystallized from methanol prior to use. The pK values of the compounds are given in Table 1. TABLE 1 pK VALUES Compound

pK Values

(1) (2) (3)r (4) (5) (6)

1.9, 6.69, 2.0, 6.73, 1.8, 7.00, 1.5, 6.16, 2.9a 2.6, 3.2,

10.0 10.4 10.8 10.4 6.35, 6.95

a pK 2 was not measured for (5).

Experimental details have been given elsewhere 4- 6. The techniques employed were cyclic and single-sweep voltammetry, chronocoulometry and chronoamperometry. Electron spin resonance spectroscopy was used to detect semiquinone species in the electrolysis solutions of compounds (1), (4), (5) and (6). The semiquinone is formed from the product of the electrolysis (quinone). In all the voltammetric work, the working electrodes were newly prepared carbon paste (Nujol) disk surfaces. Geometrical electrode areas were used in the calculations. Unless noted otherwise, all solutions contained 0.500 M NaC10 4 as a supporting electrolyte and all potentials are referred to an aqueous saturated calomel electrode (SCE). The values of the surface concentrations reported here were obtained by singlestep chronocoulometry xl. Inherent in this technique is the assumption that the amount of electricity required to charge the double layer capacity is independent of the presence of specifically adsorbed molecules. Previous experiments with the diphosphate ester of the series (compound (6)) 6 indicated that this assumption permits approximate surface concentrations to be determined with accuracy within the precision of the data. Determinant errors which are present in the data will be similar for each compound and will not obscure trends in the series. All values given are the average of three separate experiments and the precision was + 0.1 x 10-10 mol cm-2 at the 95 ~ confidence level. o~

~ (I)

/oH

o~

..1oH P r~ ~o~{

CH3

~cH 3 OH

(2)

J. Electroanal. Chem., 33 (1971) 409-418

OH

o~

/OH

o/V~o~ H3C - - ~ C H

3

H3C ~ C H Oil (3)

3

411

OXIDATION OF HYDROQUINONE ESTER DERIVATIVES RESULTS AND DISCUSSION

The structures of the compounds included in these studies follow:

o~ /6H

o ~ /oH

o

o / " P'" oH

o/S~o

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OH

O~l

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(5)

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o/

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In solutions from 1 M sulfuric acid to alkaline buffer solutions (pH > 8.8) the overall electrode process for each compound is the expected irreversible two-electron oxidation to the corresponding quinone. The oxidation process is pH dependent and can be correlated, in acidic and neutral solutions, with the ecec mechanism proposed 6 for (6). In this scheme a chemical step (either a proton dissociation or a phosphate ester hydrolysis in the case of (6)) of the one-electron intermediate occurs prior to loss of the second electron. Also, as was found 6 for compound (6), the oxidation at low pH is complicated by adsorption of the reactant at the carbon paste (Nujol)/solution interface.

Reactant adsorption Each of the phosphate esters is adsorbed to some degree"in aqueous solution of low pH. The values of the surface concentration in moles per cm 2 of geometrical a r e a (Fads) are given in Table 2. The magnitude of the Fads values and the evidence TABLE 2 ADSORPTIONIN 1 M H2SO 4 Compound

1010 F, ajmo I cm- 2

pKal ~+ 0.05)

(6) (2) (1) (3) (4)

1.90 at 1.08 at 0.96 at 0.93 at 0.49 at

2.88 2.00 1.88 1.83 1.54

a

0.000 V" 0.207 V 0.000 V 0.400 V 0.000 V

Equilibration potential, V vs. SCE.

that the uncharged, fully protonated acid species is adsorbed, support the view that these species are specifically adsorbed at the electrode/solution interface. Thus, in these media, ionic adsorption in the diffuse layer makes a negligible contribution to Fads.

This specific adsorption of these esters is independent of the electrode potential before the foot of the oxidation wave. This was verified for compound (3) and is identical6 to the behavior of compound (6). The F~a~ values were also independent of the J. Electroanal. Chem., 33 (1971) 409-418

412

E. P. MEIER,J. Q. CHAMBERS,C. A. CHAMBERS,B. R. EGGINS, C.-S. LIAO

magnitude of the potential step as long as the background region was avoided. The adsorption is strongly pH dependent. The values of Fad~ decrease to zero as the pH is increased. A typical pH profile for compound (3) is shown in Fig. 1. A ,

,

,

,

,

0.Tf i Uo'5c

, ~

6

10 oH

Fig. 1. Variation of Fads with pH for compound (3) in 0.500 M NaC104. limiting value of/"ads at low pH was not observed for any of the compounds studied, although 1 M sulfuric acid was the most acidic solution studied. The values of/"ads correspond to less than monolayer coverage based on geometrical area. The surface roughness factor is not known, but carbon paste electrodes appear to exhibit relatively rough surfaces under magnification, although crevices will be filled with Nujol. This behavior strongly suggests that the neutral molecule is extracted into the Nujol layer of the carbon paste electrode and is adsorbed on the Nujol side of the electrode/solution interface. Extraction of organic species into the Nujol phase of carbon paste electrodes is well documented in the literature 12-14. Further support for this description is afforded by the dependence of/'ads o n the pKa values for the first ionization constant of the reactant species (Table 2). The values of/"ads correlate directly with the p K a values and not with the size of the aromatic n-cloud. This is the expected behavior if the adsorption process is coupled to an extraction equilibrium into the Nujol phase.

,

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J 0.~

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r

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E/V vs. SCE

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Fig. 2. Cyclic voltammograms of compound (l) (0.37 mM) in acidic solns. (pH 1.90) at different sweep rates: (A) 1.56, (B) 15.6, (C) 155 Vs 1. Fig::--3.4g:yclicvoltammograms of compound (1) in acidic solns. (pH 1.90)at differentconcns, of (1) : (A) 0.15, (B) 0.37, (C) 0.74, (D) 1.47 raM. Sweep rate 15.6 Vs-1 in each case. J. Electroanal. Chem., 33 (1971)409-418

OXIDATION OF HYDROQUINONE ESTER DERIVATIVES

413

The surface concentrations were not strongly dependent on the ionic strength of the electrolyte solutions. Experiments at pH 1.90 with compound (6) gave Fads values of 1.5o, 1.57, and 1.67 x 10- lO mol cm -2 in 0.100, 0.500, and 1.000 M NaC10,, respectively. These data suggest a slight salting-out effect, although the variation is within the precision of the measurements. It is significant that the single sweep peak voltammetry parameters are completely consistent with the presence of adsorbed reactant. Adsorption postwaves which overlap the main diffusion wave were observed for compounds (1), (2) and (6). For all the phosphate esters, the experimental values of the current functions (ip/Acx/v, where ip is the peak current (A), A the electrode area (cm2), c the concentration of reactant (tool 1- '), and v the sweep rate (Vs- 1) increased markedly at fast sweep rates and low concentrations. Typical cyclic voltammograms for (1) are shown in Figs. 2 and 3. This behavior is characteristic of classical Brdicka postwaves ~5 and reactant adsorption ~6. The cyclic voltammograms clearly show that the contribution of the adsdrption postwave increases as the sweep rate is increased and as the concentration is decreased. At 1.47 m M (Fig. 3D) the distortion of the peak voltammogram is no longer evident at 15.6 Vs-i. Rough integration by hand of the area under postwaves gave values of Fads which were consistent with the chronocoulometric results. This behavior insures that the Fads values measured by chronocoulometry are not gross artifacts due to improper charging current corrections.

Oxidation mechanisms The voltammetric data fit the general scheme given below. In this scheme R represents either a phosphate group (POaH2, PO3H-, or PO 2-) or a sulfate group (compound (5)). The scheme is similar to those proposed earlier for compounds (1) and (6). It is also similar to the schemes proposed by Vetter and coworkers 17'18 for the oxidation of hydroquinones in aqueous solutions at gold and platinum electrodes. In their terminology the scheme can be described as an eHeR mechanism which becomes a HeeR mechanism at high pH values. In these terms e represents an electron o/R

o/R

\-e\ o XR /R 0

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\

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OH

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o

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0 k4~ofaSt

o J. Electroanal. Chem., 33 (1971) 409-418

414

E . P . MEIER, J. Q. CHAMBERS, C. A. CHAMBERS, B. R. EGGINS; C.-S. LIAO

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415

OXIDATION OF HYDROQUINONE ESTER DERIVATIVES

transfer, H a proton transfer, and R an ester hydrolysis with either P-O or C-O bond fission. Hydrolysis of the phosphate ester of the one-electron intermediate does occur at high pH values, but is a slow process for pH < 9. A mechanism similar to the above has been suggested for the oxidation of derivatives of durohydroquinone in acetonitrile x9. Salient features of the results will be briefly mentioned here. In strongly acid solutions the electron transfer steps and the proton transfer steps are reversible On the time scale of the electrochemical experiment. The overall process is irreversible because kl and k 4 a r e large and the hydrolysis of the two-electron intermediate is irreversible. The narrow peak widths and large current functions (Table 3) support this conclusion. These data were obtained under conditions where the adsorption postwave made a negligible contribution to the total current, i.e. high concentrations and slow sweep rates. One would expect the radical cation species in the scheme to be a fairly strong acid by analogy with hydroquinone. Radical cations of hydroquinone derivatives are readily prepared in concentrated sulfuric acid 2°'21. If the oxidation were carried TABLE 4 CURRENT FUNCTIONSOF DIMETHYLNAPHTHOHYDROQUINONEPHOSPHATE

v/Vs 1

0.012 0.058 0.116 0.588 1.18 11.8 58.8 118

Current functiona,b p H = 1.90

4.45

6.25

8.85

1.86 1.86 1.97 1.88 1.89 2.40 4.13 5.79

1.45 1.41 1.50 1.43 1.49 1.65 1.98 3.32

1.31 1.22 1.30 1.19 1.12 1.02 1.29 1.41

1.06 0.79 0.76 0.73 -0.62 0.63 0.77

a iv/ACx/V ' A cm -2 M -1 (Vs- a) -1/2. b c = 1.15, 1.06, 1.05, and 1.11 m M for pH = 1.90, 4.45, 6.25 and 8.85. TABLE 5 CURRENT FUNCTIONSOF DUROHYDROQUINONEPHOSPHATE

v/Vs- 1

0.012 0.020 0.115 0.198 1.13 11.3 113

Current functiona.b p H = 1.90b

3.90

5.40

7.13

8.30

10.50

1.98 1.94 2.07 2.04 2.42 3.19 5.69

1.50 1.44 1.48 1.39 1.53 1.69 1.93

1.26 1.22 1.21 1.18 1.29 1.21 1.21

1.21 1.18 1.14 1.15 0.96 0.89 0.77

1.02 1.06 0.99 1.03 1.02 0.90 0.83

0.86 1.20 0.85 0.76 0.90 0.85 0.78

ip/Ac~/v, A c m - 2 M 1 (Vs 1)-1/2. b c=0.93, 0.96, 0.95, 0.95, 0.98, 0.98 m M for p H = 1.90, 3.90, 5.40, 7.13, 8.30 and 10.50.

a

J. Electroanal. Chem., 33 (1971) 409~418

416

E. P. MEIER, J. Q. CHAMBERS, C. A. CHAMBERS, B. R. EGGINS, C.-S. LIAO

out in extremely strong acidic media, this scheme would predict that a reversible oneelectron oxidation would be observed. However, under the conditions of this study the one-electron intermediate is rapidly deprotonated. In 1 M H z S O 4 this proton transfer step is reversible, but as the pH increases, the rate of the backward step, k_ I[H+], decreases until this process does not occur to an appreciable extent on the time scale of the experiment. This is manifested by a decrease in the current function and an increase in the peak width. Typical values are given in Table 3 at pH 5-7. The process is diffusion limited, but the current function has decreased because the first irreversible chemical step now follows the first electron transfer and not the second. In addition, the diffusion coefficients are pH dependent, decreasing with increasing pH, which also decreases the current functions. These trends are also seen in the data presented in Tables 4 and 5 where values of ip/Ac~/v are given for compounds (2) and (3) at different sweep rates. These data emphasize the difficulty of making conclusive deductions based on variations of the current function in complex systems. Although these data were taken over a time domain of four orders of magnitude, there are no marked changes in the current functions. In neutral solutions the values tend to decrease with increasing sweep rate, but the decrease is only 20-30~o over A log v = 4. At low pH values (i.e. 1.90) the current functions increase at fast sweep rates due to reactant adsorption. At intermediate pH values a shallow minimum is observed in the variation of the current functions with sweep rate. The solution pH, on the other hand, strongly influences the current function values. For example, the current function of (3) decreases to ca. 60 ~o of its value at pH 1.90 in the neutral region (at 0.020 Vs- 1). The overall process is still a two-electron oxidation to the quinone, however, and the diffusion coefficient of (3) has decreased by only 11 ~o. On the basis of the above mechanism, the current function at pH 7.13 would be 66 ~ of its value at pH 1.90. CF(pH 7.13) 2 (D7,13) ~ CF (pH 1.90) - 2~ \ --D~9~,9 ° = 0.66 The observed value is 0.61. Similar trends are shown for each of the compounds studied; the data are given elsewhere v. The variation of the half-peak potential with pH is ca. 60 mV per pH unit at low pH and becomes independent of pH at high pH values. This is consistent with

50

'/~o - 100

o18

o'.6

0'.4 o12 o'.o -o'.2

ElY vs.. SC E

Fig. 4. Cyclic voltammogram of compound (5) (0.95 mM) in basic solns. (pH 9.50) at 0.060 Vs- 1. J. Electroanal. Chem., 33 (1971) 40941_8

OXIDATION OF HYDROQUINONE ESTER DERIVATIVES

417

the oxidation scheme. In the alkaline region, the behavior becomes more complex and is not fully understood. For each compound studied, the oxidation process is split into two irreversible waves in basic solutions. A typical cyclic voltammogram for compound (5) is shown in Fig. 4. The initial wave can be assigned to oxidation of the completely ionized phenolate analog (pK ca. 10-11) which is supplied to the electrode surface by dissociation of the phenolic hydrogen. The half-peak potential of this wave becomes independent of pH above pH 8. This transition from an eHeR to an HeeR mechanism occurs between pH 6 and 8. Detailed studies at high pH (pH > 9) were carried out only on compounds (1) and (6). The substrates were not stable at high pH, but decomposed via reactions not completely understood. For (1), the radical intermediate has been shown to hydrolyze to the semiquinone in a relatively slow step which is pH dependent 4. In several experiments at pH 12 the initial oxidation of (3) was demonstrated to be electrochemically reversible at slow sweep rates. This process could not be studied quantitatively because (3) was not stable even in solutions thoroughly purged of oxygen. The general oxidation scheme presented here is preferred to one which would invoke pH dependent electron transfer rates to explain the overall irreversible nature of the electrochemical process. In the scheme above the irreversible character observed at intermediate pH values is rationalized by invoking rate processes reaction ((I), (II) and (III) in the Scheme) which become irreversible on the time scale of the voltammetric experiment. ACKNOWLEDGEMENTS

The authors acknowledge the work ofR. J. Bezjian and M. L. Rumpel (National Science Foundation Research Participant, University of Colorado, 1965) who contributed to the early stages of this research. The research was supported by a grant from the U.S. Public Health Service, 5-RO1-GM14815. SUMMARY

The electrochemical oxidation of a series of ester derivatives of hydroquinones has been studied in aqueous solutions at the carbon paste electrode. The series includes simple phosphate and sulfate derivatives of benzohydroquinone and naphthahydroquinone. These compounds have been proposed by others as model intermediates in biological phosphorylation and sulfonation processes. In acidic media the compounds are specifically adsorbed at the carbon paste/solution interface. The extent of adsorption is correlated with pKal of the ester ester derivatives, increasing as the pKa~ value increases. It is proposed that the neutral fully protonated acid species is adsorbed. An ecec oxidation mechanism is suggested to account for the voltammetric data. In acidic solutions an eHeR scheme is proposed (where H represents a proton transfer and R an ester hydrolysis) which transforms into an HeeR scheme between pH 6 and 8. J. Electroanal. Chem., 33 (1971) 409-418

418

E . P . MEIER, J. Q. CHAMBERS, C. A. CHAMBERS, B. R. EGGINS, C.-S. LIAO

REFERENCES 1 V. M. CLARK AND D. W. HUTCHINSON in J. COOK AND W. CARRUTHERS(Eds.), Progress in Organic Chemistry, Vol. VII, Butterworths, London, 1968, p. 100. 2 P. MITCHELL, Nature, 191 (1961) 144. 3 R . A . HARRIS,J. T. PENNISTON, J. ASAI AND D. E. GREEN,Proc. Nat. Acad. Sci. U.S., 59 (1968) 624, 830. 4 C. A. CHAMBERSAND J. Q. CHAMBERS, J. Amer. Chem. Soc., 88 (1966) 2922. 5 E. P. MEIER AND J. Q. CHAMBERS, Anal. Chem., 41 (1969) 914. 6 E. P. MEIER AND J. Q. CHAMBERS, J. Electroanal. Chem., 25 (1970) 435. 7 E. P. MEIER, Ph.D. Thesis, University of Colorado, Boulder, Colorado, 1969. 8 C.-S. LtAO, M. S. THESIS, University of Colorado, Colorado, 1969. 9 L. F. FIESER AND E. M. FRY, J. Amer. Chem. Soc., 62 (1940) 228. 10 S. YAMAGUCHI, Nippon Kagaku Zasshi, 80 (1959) 171. 11 F. C. ANSON, Anal. Chem., 38 (1966) 54. 12 T. KUWANA AND W. G. FRENCH, Anal. Chem., 36 (1964) 241. 13 D. G. DAVIS AND M. E. EVERHART, Anal. Chem., 36 (1964) 38. 14 C. A. H. CHAMBERS AND J. K. LEE, J. Electroanal. Chem., 14 (1967) 309. 15 J. HEVROVSKVAND J. KUTA, Principles of Polarography, Academic Press, New York, 1966, chap. XVI. 16 R. H. WOPSCHALL AND I. SHAIN, Anal. Chem., 39 (1969) 1514. 17 K. J. VETTER, Electrochemical Kinetics, Academic Press, New York, 1967, p. 483. 18 J. K. DOHRMANN AND K. J. VETTER, Ber. Bunsenges. Phys. Chem., 73 (1969) 1068. 19 V. D. PARKER, Chem. Commun., (1969) 610. 20 I. C. P. SMITH AND A. CARRINGTON, Mol. Phys., 12 (1967) 439. 21 A. B. BARABAS,W. F. FORBES AND P. D. SULLIVAN, Can. J. Chem., 45 (1967) 267.

J. Electroanal. Chem., 33 (1971) 409-418