Electroreduction of 3-nitrophthalic acid on mercury electrodes in water hexamethylphosphortriamide (HMPA) mixtures

JOURNAL OF

ELSEVIER

Jou[nai of Eiectroa, alytical Chemistry 423 (1997) 103-108

Electroreduction of 3-nitrophthalic acid on mercury electrodes in water hexamethylphosphortriamide (HMPA) mixtures A. Kalandyk, J. Stroka, Z. Galus Department of Chemistr3". Unit,ersi~,of Warsaw, Pasteura !. PL-02-093 Warsaw. Poland Received 23 January 1996; revised 29 April 1996

Abstract

The electroreduction of 3-nitrophthalic acid anion on the mercury electrode in water + HMPA mixtures was studied with 0.5 M NaCIO4 and 0.05 M NaOH as electrolytes. A detailed kinetic study concerned one-electron reduction of 3-nitroph;" alic acid anion to free radical. The change of the determined rate constants of this process on water + HMPA mixture composition was interpreted in terms of the solvent adsorption on the electrode and reagent solvation. An equation is proposed which describes the observed change of the kinetics. © 1997 Elsevier Science S.A. Keywords: Electroreduction; Aromatic nitrocompounds; Electrode reaction; Mixed solvents

I. Introduction

In recent years we have studied several electrodeposition-type electrode reactions in water + organic solvent mixtures [I-3]. The change of the rate constant with the solvent composition was then described by the relation k, = kw( 1 - 0)" + kR 0"

( I)

where k~, k w and k R denote rate constants measured at the formal potential in mixed solvent, in aqueous solution and organic solvent respectively, 0 stands for the degree of coverage of the electrode surface by adsorbed molecules of organic solvent, and a and b respectively denote the number of water and organic solvent molecules which are removed from the electrode surface in order to make a site for a reactant molecule in the surface layer [4]. Though Eq. (!) was applied successfully to the electrodeposition-type reactions [1-3] in water+ hexamethylphosphortriamide (HMPA) mixtures, the charge-transfer reactions of organic [5] and inorganic reactants [6,7] could not be described properly. Recent studies on electroreduction of nitrobenzene (PhNO 2) [5] and earlier investigations on the Eu3+/Eu 2+ and V 3 + / V 2÷ systems in water+ HMPA mixtures at

* Corresponding author. i Dedicated to Professor Petr Zuman on the occasion of his retirement and in recognition of his contribution to electroanalysis. 0022-0728/97/$17.00 © 1997 Elsevier Science S.A. All rights reserved. Pll S 0 0 2 2 - 0 7 2 8 ( 9 6 ) 0 4 7 8 9 - 4

mercury electrodes [6,7], have confirmed that the influence of this mixed solvent on kinetics depends on the type of reactant. Though the l e-charge transfer process in the PhNO2/PhNO.;- system is practically independent of the increase in the HMPA concentration in water + HMPA solutions [5], a significant decrease of the rate of further 3e-electroreduction of PhNO;- to phenylhydroxylamine (PhNHOH) was observed. The change of the rate constant of PhNO2/PhNO_gand PhNO:,-/PhNHOH systems could be described formally by Eq. (I) with the a parameter equal to 0 and 3.5 [5] respectively. This large difference in the a parameter of both systems, which have reactants of a similar size, may result from two opposite tendencies: (i) the decrease of the rate constant, due to the increase of 0, (as predicted by Eq. (1)) and (ii) the increase of the reactant concentration in the surface phase when, with the increase of the bulk HMPA concentration, the composition of the reactant soivation shell becomes similar to that of the surface layer on the electrode. In order to confirm this assumption it was desirable to analyse the kinetics of the simple, relatively slow, charge transfer process not associated with the chemical transformations of the reactant in the water + HMPA solvent system. Such a process should occur at potentials, where the electrode is significantly covered by HMPA molecules, even at the low content of this solvent in the water +

104

A. Kalandyk et aL ~Journal of Electroanalytical Chemistr3' 423 (1997) 103-108

HMPA mixture. As a model of such reactant, 3-nitrophthalic acid (3-NPA) anion Ph(COO-)2NO 2 was selected. The present work is related partly to the papers of Zuman and Fijatek [8,9] on the mechanism of the electrode reaction of nitrobenzene and its derivatives.

2. Experimental

2.1. Reagents In experiments, spectroscopically pure HMPA (Merck) was used. 3-NPA Ph(COO- )., NO 2 (Reidel de HaSn) and other chemicals used in the present work were p.a. reagents. The sodium salt of Ph(COO-)2NO 2 was prepared from the acid following the procedure given in the literature [10]. All solutions were prepared using triply distilled water. The third distillation was carried out from an allquartz still. In all solutions, 0.5moldm -3 NaCIO 4 was used as a background electrolyte with 0.05 mol dm -3 NaOH to ensure, high pH during experiments. The sample of ~.odium salt of Ph(COO-)2NO ., was dissolved before experiments in the solutions which were deaerated by purging pure argon. Mercury used in experiments was chemically purified and twice distilled.

2.2. Apparatus The experiments were carried out in a three electrode cell. Static mercury electrode (Laboratorni Pristroje) was used as a working electrode, while the platinum foil ( 1 cm:) served as an auxiliary electrode. The potentials were measured vs. the aqueous calomel electrode with a saturated solution of NaCI (SSCE). Cyclic voitammetry (CV) experiments were carded out with the use of an instrument constructed from an Elpan generator type EG-20, potentiostat (Elpan type EP-20) and X - Y recorder manufactured by Yokogawa (type 3077). The electrode kinetic studies were performed using single and reverse pulse voltammetry, with charge recording instead of current sampling. The home-made apparatus was used in these studies, with the current integration times in limits from 1 to lOOms. Usually, in kinetic experiments the current integration started after I ms from the pulse application in order to re,ject a significant part of the charging current. All experiments were carried out at 25 + 0. I°C in a cell with a water jacket.

3. Results The electroreduction of the sodium salt of 3-NPA and the electro-oxidation of the resulting products were carried out in water + HMPA mixtures with HMPA content from

0 to 70 vol%. In order to decrease the error and to improve the precision of kinetic measurements, for each solution the charge-potential dependences were recorded using four different integration times: 9, 36, 64 and 81 ms. Cathodic curves were recorded starting from the initial potential - 0 . 6 0 V up to - 1 . 3 0 V , while in reverse pulse voltammetric experiments the electrode was kept at - 1.3 V and was pulsed step by step up to ---0.1 V. As expected, in all solutions in the cathodic polarization, two charge-potential waves were observed which correspond to the processes Ph(COO- )2NO 2 + ePh(COO-)2NO_;- w a v e ( l )

(2)

and Ph(COO- ),NO_;- + 3e- + 3 H , O ~- P h ( C O O - ) 2 N H O H + 4 O H - wave(II)

(3)

This behaviour is similar to that of nitrobenzene or nitrobenzoic acid under similar conditions [8,9,11-17]. The formal potential of the first wave reladve to the external SSCE moves only slightly to more negative potentials with increase of the HMPA content in the mixture (Fig. l). The half-wave potential of the second cathodic wave shifts at first significantly to more negative potentials, then later moves to less negative values with increase of the HMPA content in the mixture (Fig. 1). This negative [

--i

T ...........

,

1.6 L

l

>

I

U.I i

i

!'

[

i

0.2 0.0 0

.................... 20 40

60

J 80

~HMPA]/vol %

Fig. 1. The dependence of: El~ 2 of the Ph(COO- )2NOC electroreduction ( v ) , Ell,. of the anodic oxidation of Ph(COO-)2NO.;- (O). and El~,- of anodic oxidation of Ph(COO-)2NHOH (O) on the HMPA concentration in water+HMPA mixtures. Integration time equal 9ms. Curves marked with (zx) and ( l ) give the solvent dependence of the formal potential of the Ph(COO- )2 NO2/Ph(COO- )2 NO2- system vs. aqueous SSCE and the ferrocene electrode respectively. (V) and (D) respectively denote the potentials of the anodic and cathodic adsorptiondesorption peaks of HMPA.

105

A. Kaland3k et al. / Journal of Electroanalytical Chemi.~n'r 423 (1997) 103-108

shift of the second wave (II) is so significant that at higher HMPA concentration the reaction in Eq. (3) occurs in the range of the adsorption-desorption potentials of HMPA [18]. This gives rise to the formation of maxima on these waves. The second wave, as expected, was irreversible in all solvent mixtures used. The kinetic parameters of the reactio~ corresponding to this wave were not determined, because the rate of this reaction was influenced to some extent by the desorption process of HMPA. Therefore, only the kinetics of the reaction in Eq. (2) was studied in detail. Using reverse pulse voitammetry with the initial potential equal - 1 . 3 0 V at low HMPA content, two anodic waves were recorded: one at potentials corresponding to the reaction of Eq. (2), and the additional wave due to the process

6r

......

~

,

- f ~

I

18 i

/

5'

i14 12

,~

10

Ph(COO- )2NHOH ~ Ph(COO- )2NO + 2e- + 2H ÷

(4) The reaction in Eq. (4) occurs at the positive side of the adsorption-desorption potentials of HMPA (see Fig. l). Using reverse pulse voltammetry, the height of both anodic waves was, as expected, dependent on the initial potential. When this potential was selected on the plateau of the first cathodic wave, the radical oxidation wave height was at a maximum, while the second wave (Eq. (4)) was not observed. On the contrary, using more negative initial potentials (plateau of the second cathodic wave) the wave corresponding to the reaction in Eq. (4) was the highest. Since the process corresponding to the reaction in Eq. (2) exhibited quasi-reversibility, which is more evident for shorter integration times, using the co, responding chargepotential cathodic and anodic waves (I), we were able to calculate the kinetic parameters of that process using the method of Koutecky [ 19], Randles [20] and Kimmerle and Chevalet [21 ]. The diffusion coefficient of 3-NPA anion in various mixtures of water with HMPA was determined using the polarographic method with recording of the charge instead of current. With increasing HMPA concentration up to 40 vol%, a systematic decrease of the diffusion coefficient is observed (Fig. 2). At still higher HMPA content in the mixed solvent its values increase slightly. The observed decrease, when moving from 0 to 40 vol% HMPA in the mixture may be qualitatively explained both by the increase of tbe viscosity of the medium (Fig. 2) and partly by the re-solvation of reactant, evidenced by the change of the formal potential (Fig. 1). The rate constants' potential dependences were calculated for different solvent mixtures. Under identical conditions, the rate constants calculated from the charge-potential dependences recorded at different integration times were concordant, within the limits of experimental error. Cathodic and anodic Tafel plots constructed from these data were linear in all water + HMPA mixtures used. The

1l,

,

~

-

0

20

40

60

[HMPA]/vol

J2

%

Fig. 2. The dependenceof the diffusion coefficier: of Ph(COO- )2NO2 (O) and the correspondingWalden product (V) on the water+HMPA mixed solventcomposition. cathodic transfer coefficients ce were found to be equal 0.50 _ 0.02 and were virtually independent of the solvent composition. However, the anodic coefficients fl exhibited a change with the concentration of the non-aqueous solvent, from 0.51 (0.1 vol% HMPA), 0.40 _4-0.02 in the range l to 20voi% HMPA to 0.43 (40vo1%) and 0.53 (70volC~ HMPA). From the intersection of the cathodic and anodi~ Tafel dependences, the formal potentials Er and the rate constants at these potentials were determined. These Ef values, when expressed vs. the ferrocene electrode, move more negative with increase in HMPA content in the mixed solvent (see Fig. 1). The rate constants at the formal potentials change in a more complicated fashion. At first, the addition of small quantities of HMPA to aqueous solution leads to a significant decrease of the rate constar amounting to two orders of magnitude. At approximately 5 vol% HMPA the standard rate constant reaches a minimum. Further increase of HMPA concentration induces an increase of k,.

4. Discussion

There is a significant difference in the electrorcduction of PhNO 2 and Ph(COO-)2NO2 anion in alkaline aqueous solutions. Nitrobenzene undergoes 4e- electroreduction to phenylhydroxylamine in one step [ l l - 1 7 ] with a fast transfer of the first electron. Though in aqueous solutions the formation of the free radical PhNO;- is not observed, in non-aqueous solvents such a radical is relatively stable

106

A. Kalandyk et aL / Journal of Electroanalyticai Chemisto' 423 (1997) 103-108

and the standard rate constant of the PhNO2/PhNO.~system is estimated to be of the order of 2cms -~ [5,11,14,17]. In contrast, the electroreduction of Ph(COO-)2NO2 anion occurs in aqueous alkaline solutions in two separated steps, with consumption of one electron and three electrons respectively. The apparent standard rate constant of the system Ph(COO-)2 NO2/Ph(COO- )2 NO2- (first step), is equal 2.82 × 10 -2 cm s -~. The second 3e- electroreduction process is totally irreversible. The dependence of the rate constant of the first l estep on the water + HMPA solvent composition, in a wide composition range, is also different for PhNO 2 and Ph(COO-)2NO2 . While this dependence for the Ph(COO-)2NO~/Ph(COO-)~NO2 - couple is similar to that of the Eu(III)/Eu(ll) system in the same mixed solvent [6], the rate constant of the PhNO2/PhNO;- system was practically independent of the mixed solvent composition [5]. The decrease of the rate constant of Ph(COO-)2NO2 electroreduction, with the rise of the HMPA concentration at very low HMPA content, may be explained by using a previously developed approach. A comparison of the dependence of k, on the solvent composition with the adsorption isotherm of HMPA [1,18] at the formal potential of the Ph(COO-)2 NO2/Ph(COO-)2 NO;- system (Fig. 3(A)) reveals a similar shape for both dependences at low HMPA concentration. Therefore, at low 0 values, when the second term in Eq. (1) can be neglected, we have analysed the dependence of log k~ on log (I - 0) according to the equation [!,2] k~ = k . ( I - 0 ) "

(5)

Such an analysis performed previously [2-7] indicated that, in several systems, a low content of the organic solvent in the mixture with water, may play the role of an inhibitor. The isotherm given in Fig. 3(A) was obtained using the surface coverage 0 calculated from the previously measured differential capacities of the mercury electrode, at various potentials in different water+ HMPA-mixtures [=Q1 t. z ~lh

The dependence of log k~ on log (1 - 0) obtained (Fig. 3(B)) is fairly linear (correlation coefficient equal 0.97) with slope a - 1.2. This a value is different from those obtained for the PhNO2/PhNO ~- and PhNO~-/PhNHOH systems [5], equal to 0 and 3.5 respectively, though all these reactants have similar sizes. The parameter a is defined, according to the inhibition model [4], as the number of water molecules removed from the electrode surface by one molecule of reactant, but its value should be reduced by solvation of the reactant. Some change in solvation of the reactant, in the used solvent mixtures, will be shown further in this discussion. The approximate iinearity of the log k~ vs. log (1 - 0)

A

..~....,~-.-k--

1.0

i

i

11'0 I

A

0.8

("

]1.s

0.6

~o

~z.0 .~

\o

0.4

~

0.2 0.0

o

~

3.0

~

........... ~1.0 1.5 [HMPA] / voi % . . . . . . . .

0.5

0.0

E

35

I

.

2.0

B 1.0[

~

, --

,

-

L _

3.0 35 [ n.0

o , 0.5

, 1.0

°

x""~xx 1.5

2.0

-tog (1-0)

Fig. 3. (A) Isotherm of ad~orplinn of HMPA nn the mercury elet~l,,de ai the t'ormal potential of the Ph(COO- )2 NO2/Ph(COO- )2 NO j- system (&). 0 was evaluated from the capacity measurements [18]. The second curve (C)) gives the dependence of the standard rate constant on the solvent composition. (B) Dependence of - l o g k. on - I o g ( l - 0 ) for the Ph(COO- )2 NO2/Ph(COO- )_~NO:,- system in water + HMPA solutions.

dependence is observed only to l vol% HMPA. In these water-rich mixtures the reactant ions only interact very weakly with HMPA, as shown by the virtual independence of the fo,,'m~.l potentials of the Ph(COO-)2NO2/ P h ( C O O - ) E N O 2 - of the solvent composition (Fig. i). Since the models presented in literature do not describe adequately the dependence of the kinetics of the electrode processes of the charge transfer systems A / A - on mixed solvent composition at larger organic component contents, in this paper the standard rate constant dependence on the formal p,'Jtential was also analysed, ,:~:pressed against the ferrocene electrode. Such an analys,~ was previously used by Broda and Galus [22] to rationalize the kinetics of var~cus electrode processes in pure organic solvents. When the added solvent plays the role of an inhibitor, Alog kJAEf should approach to + ~ in such an analysis. The large value of that parameter obtained (equal to + 66 log units per volt (Fig. 4)) points to a slight change of the reactant solvation in the water-rich region, which in turn may influence the slope of the log k~ vs. l o g ( l - 0) dependence. At higher mole fractions of HMPA in the mixtures, after significant re-solvation, when the composition of the solvation shell of the reactants becomes similar

A. Kahmdyk et al. / Jounlal of Electroanalytical Chemistry 423 (1997) 103-108 1

!

1

i

\

L . _ . _ _ _ _

09

1.0

I

1.1

1.2

-E~/V Fig. 4. The dependence of - l o g k~ on the formal potential of the Ph(COO- )2 NO2/Ph(COO- )2 NO:,- system. The formal potentials are expressed vs. ferrocene electrode.

to that of the electrode surface, the electrode process should be accelerated due to an. increase in the penetration of the surface layer by the reactant. The effect of inhibition should then be diminished and, in consequence, the expenent a in Eq. (1) may be reduced. The second part of log k s - E l relation, with a slope equal - 7 . 8 log units per volt, was observed when the electrode was practically fully covered by HMPA. Now, the change of the kinetics should be primarily dependent o n the change of the reactant solvation with the ~nlvenl composition. In order to describe the increase of the rate constant at higher HMPA content, we propose the following equation: k, = k , ( l - 0)" + kminObexp[-anF( E l - Ef.min)//RT] (6) Eq. (1) may be considered to be the simplified form of Eq. (6), when the formal potential Ef of the system studied, at some solvent composition, is equal to the potential corresponding to the minimum value of standard rate constant kmi,. While the first term of the fight hand side of Eq. (6) describes the change of the kinetics when the formal potential is practically independent of the solvent composition (water-rich region), the second term of Eq. (6) is valid when Ef changes and the first term is negligible, since (1 - 0) a is near to zero. One has then

k s = kminexp[-t~nF( E l - Ef.min)/RT]

(7)

Now, the change of the rate constant with the solvent composition is controlled by the re-solvation of the reactants. Both potentials Ef and Ef.min should be measured with respect to the potential of the solvent-independent electrode. Eq. (7) reflects in fact the BrCnsted relation of the linear dependence of the Gibbs energy of activation AG ~ on the Gibbs energy of reaction ,.s,.," ,-,o~p which may be given ..~,~tr,,~,..,. form AG* = AGo*+ otAG e

(8)

The linear plot of log k s on Ef is shown in Fig. 4. It

107

gives in fact the dependence of the standard rate constant of the redo~ system on the difference in the Gibbs energy of solvation of the oxidized and reduced forms at different solvent compositions. The transfer (BrCnsted) coefficient calculated from the slope of that plot (in Fig. 4) o~= 0.47, is equal to that obtained from the Tafel plots. Its value points to the intermediate solvation of the activated state between those of the final and initial states. In order to learn more about the mechanism of the studied reaction we have also analysed separately the dependence of the cathodic kfh and anodic kbh rate constant on the solvent composition. These two dependences are different (see Fig. 5), though the rates of both processes were measured at the same constant potential, -0.925 V vs. the ferrocene electrode. For the HMPA concentration range lower than I vol%, when re-solvation of reactant occurs to a small extent, the cathodic and anodic processes were similarly inhibited. However, at higher concentration (exceeding 20voi%), when the surface coverage of the electrode by HMPA is practically complete, kfh slightly decreases while kbh significantly increases with the increase in HMPA concentration. This difference in change of kfh and kbh with HMPA composition reflects the difference in solvation of the oxidized and reduced forms, which participate directly in the electrode reactions. Both substrate and product are anions and should be better solvated by water than HMPA molecules; however, the interaction of the reduced form with water should be stronger. This explanation is based on quite different acceptor numbers of both solvents, 54.8 for water and only 10.6 for HMPA [23]. These different acceptor numbers of both solvents, and consequently their different interaction with the oxidized and reduced states, also explains the change of the formal potential with the solvent composition. The plots of log kfh and log kbh o n log CHMPA for CHMPA>_ 5VO1%, at a constant potential, have a slope

Alogkfh/AlogcnMpa = 0 and Alogkbh/AlOgCnMPA = I. These approximate values, which have the significance of

5 ~ z 3 !

V

kbh

0

kfh

]

',.~..-v -~ ).~0.~0....---'0

!

20

;

~ 0 ~

_

1

40

__

. . . . .

I

L ,.J-

60

80

[HMPA] t v01%

Fig. 5. The dependence of anodic kbh and cathodic k~ rate constants of the Ph(COO- ), NO.,/Ph(COO- )2 NO;- system on the HMPA concentration in water~-HMPA mixtures. The data were calculated at -0.925 V vs. ferrocene electrode.

108

A. Kalandyk et al. / Journal of Electroanalytical Chemistr),' 423 (1997) 103-108

the electrochemical reaction orders with respect to HMPA, should be considered with care, because the solvent properties were changed with HMPA concentration and activities rather than concentrations should be used. These data, combined with earlier results, suggest that after the minimum in the cathodic reaction, when k, starts to increase with HMPA concentration, the partly solvated HMPA species present in the bulk of the solution participate. After l e - transfer, the reduced form interacts more strongly with water, which has a higher acceptor number, and loses one molecule of HMPA. In the oxidation process the more strongly hydrated anionic radical of Ph(COO-)2NOginteracts with the one molecule of the solvent and then the electron is transferred to the electrode. Acknowledgements

This work was supported by the State Committee for Scientific Research (KBN) through the University of Warsaw within the BST project. References [I] K, Maksymiuk, J. Stroka and Z. Galus, J. Electroar~al. Chem., 167 (1984) 211.

[2] K. Maksymiuk, J. Stroka and Z. Galus, J. Electroanal. Chem., 181 (1984) 51. [3] A. Mital, K. Maksymiuk and J. Stroka, J. Electroanal. Chem., 272 (1989) 145. [4] R. Parsons, J. Electroanal. Chem., 2 i (1969) 21 I. [5] A. Kalandyk and J. Stroka, J. Electroanal. Chem., 246 (1993"., 323. [6] M. Cetnarska and L Stroka, J. Electroanal. Chem., 234 (1984) 263. [7] W. G6rski and Z. Galus, Electrochim. Acta, 33 (1988) 1397. [8] P, Zuman and Z. Fijalek, J. Electroanal. Chem., 296 (1990) 583; 589. [9] P. Zuman and Z. Fijalek, J. Org. Chem, 296 (1991) 5486. [10] A.I. Vogel, A Text Book of Practical Organic Chemistry, Green, London, 1956. [I 1] W. Kemula and T. Krygowski, in A.J. Bard (Ed.), Encyclopedia of the Electrochemistry of the Elements, Voi. 13, Marcel Dekker, New York, 1979. [12] W. Kemula and Z. Kublik, Roczniki Chem., 32 (1958) 941. !'13] W. Kemula and Z. Kublik, Nature, 182 (1958) 793. [14] M.E. Peover and I.S. Powell, J. Electroanal. Chem., 20 (1969) 427. [15] B. Kastening and L. Holleck, Z. Elektrochem., 63 (1959), 163. [16] L. Holleck and B. Kastening, Rev. Polarograph. Jpn., I 1 (1973) 129. [17] Yu.M. Kargin, L.Z. Manapova, S.V. Kuzovenko, L.M. Vorontsova and T.Ya. Zaretskaya, Electrokhimiya, 19 (1983) 988. [18] K. Maksymiuk, M.Sc. Thesis, Warsaw University, 1983. [19] J. Koutecky, Chem. List., 47 (1953) 323. [20] J.E.B. Randles, Can. J. Chem., 37 (1959) 238. [21] F.M. Kimmerle and J. Chevalet, J. Electroanal. Chem., 21 (1969) 237, [22] J. Broda and Z. Galus, J. Electroanal. Chem., 199 (1986) 233. [23] U. Mayer and V. Gutmann, Monatsh. Chem., 101 (1970) 912,