J. Electroanal. Chem., 158 (1983)341-356 Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

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R E D U C T I O N O F C - - C D O U B L E B O N D S O N A M E R C U R Y E L E C T R O D E

P A R T IV. P O L A R O G R A P H I C AND KINETIC B E H A V I O R O F M O N O E T H Y L F U M A R A T E

LUIS CAMACHO and JUAN JOSI~ RUIZ *

Departamento de Quimica Fisica, Facultad de Ciencias, Universidad de C6rdoba, C6rdoba (Spain)

(Received 28th September 1982; in final form 6th April 1983)

ABSTRACT

A polarographic (dc and dpp) and kinetic studies reduction of monoethyl fumarate (MEF) has been carried out in the 0-11 pH range. Effects of pH, drop time and temperature on the limiting current, half-wave potentials, and E vs. log[i/(i L - i ) ] plots are shown. From the C-E curves obtained, the adsorption of MEF on the mercury electrode in acid medium is shown to be present. The Tafel slopes and reaction orders with respect to MEF and H ÷ ion are obtained at potentials corresponding to the foot of the waves. Reaction mechanisms are thus proposed in this zone of potentials.

INTRODUCTION

Previous studies [1,2] of the reduct ion of fumaric and maleic acids on a mercury electrode showed different behavior for both isomers in acid and basic media. On

the other hand, similar studies [3-6] have been carried out on the polarographic reduct ion of diethyl esters of these acids in aqueous solutions. However, as far as we

know, there exists no electrochemical s tudy of monoe thy l esters. For this reason, the aim of this work is, firstly, to carry out a polarographic s tudy of monoe thy l fumara te

(MEF) using dc and differential pulse po la rography techniques, and, secondly, to

de termine the reduct ion mechanisms of this c o m p o u n d on a mercury electrode in the 0-11 pH range.

EXPERIMENTAL

All reagents used were from Merck p.a. quality. Buffered solutions as suppor t ing

electrolyte were used, in each case, with the following componen t s and concentra- tions: 0 - 2 p H range: sulfuric acid; 2-11 p H range: 0.1 M acetic acid, 0.1 M

* To whom correspondence should be addressed.

0022-0728/83/$03.00 © 1983 Elsevier Sequoia S.A.

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phosphoric acid and 0.1 M boric acid. The pH was adjusted with N a O H and the ionic strength to 0.6 M with N a N O 3.

The i -E relationships were obtained at Hg at 25°C in point-by-point measure- ments using an Amel 551 potentiostat and a Hewlett-Packard X - Y recorder. For the automatically registered dc polarographic curves, a Hewlett-Packard 3310A function generator was also used. The differential pulse polarography measurements were carried out with a Metrohm 626 Polarecord. The C - E curves were automatically obtained using the circuit designed by Roldan et al. [7].

An Ingold 303-NS and Pt-4805 saturated calomel reference and Pt auxiliary electrodes were used respectively. The working electrode was a mercury capillary Radiometer B-400 with the following characteristics: m = 1.997 mg s-1, t = 4.50 s, open circuit, in Bri t ton-Robinson buffer solution at pH = 1.6, and h = 40 cm.

All the other experimental conditions were as those described in ref. 8.

RESULTS

Dc polarography

Monoethyl fumarate shows either one or two reduction waves depending on the pH of the solution. The first wave is visible throughout the 0-9.5 pH range while the second is only visible from pH = 6.5, although above pH = 10 it decreases gradually

4

3

2

I

| . . . . j , L . . . . , _ A .

-600 -700 -800 -900 -1000

Fig. 1. Polarograms of MEF c = 5 × 10 -4 M. (a) pH = 1.87; (b) pH = 4.23.

i . . . . .

-1100 E/mV

" , ~ , m ,

4.0 0 O 0 @

3.0

2.0

1.0

2 5 8 pH

Fig, 2. Variation of i L with pH. (©) First wave; (@) second wave.

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with time due to the basic hydrolysis yielding fumaric acid. The shape of the first wave is strongly affected by the pH of the solution. Two typical polarograms are shown in Fig. 1. The variation of the limiting current iL with pH for both waves is shown in Fig. 2. The 1og(iLl/iL2) VS. pH plot gives a straight line: p K ' = 7.87, where pK' is the pH value at which the current iL~ ---- iL2.

On the other hand, when i L is independent of the pH, its variation with the drop time t a indicates that both processes are governed by diffusion overvoltages at the

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top of the waves. The plot of the instantaneous limiting current vs. time t, for a single drop, yielded (~ log iL/O log t ) = 0.19 (this value remains unchanged up to pH = 6.5 in the first wave). The small temperature coefficients support this conclu- sion.

The El~ 2 shifts to more negative potentials at increasing pH (Fig. 3). As can be seen, four linear segments of slopes - 9 3 , -130 , - 6 0 and - 3 6 m V / p H unit are observed for the first wave in the 0-3.5, 3.5-5.3, 5.3-6.5 and 6.5-7.5 pH ranges

0 . 7 0 0

0 . 9 0 0

1.10C

1 . 3 0 0

,co

o

L I . . . . . . . . . . . x . . . . i

2 4 6 8

Fig. 3. Variation of El~ 2 with pH. (O) First wave; (O) second wave.

OO o oo

pH

,o~I,J,il-i, ]

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1 , 0

1 (1

0

•

-O. t~O0 - 0 . 8 0 0 -1 . 0 0 0 -1 . 2 0 0 E/V

Fig. 4. Effect of pH on the log[i/(iL--i)] VS. E plots. First wave. (a) pH = 1.23; (b) pH = 1,94; (c) p H = 2.34; ( d ) p H = 2.97; ( e ) p H - 3 . 9 9 ; ( f ) p H - 4 . 2 3 ; (g )= 4.48; ( h ) p H - 4 . 7 0 ; ( i ) p H = 4,93; (j) p H = 5.08; (k) p H - 5 . 3 1 ; (1) p H = 5.56; (m) p H = 5,81; (n) pH = 6.05; (O) pH = 6.41; (p) pH =6.68; (q) p H - 7.09; (r) pH = 7.82.

1 .0

logl~l

0 .0

--1.0

• " ~ • , i . | | - -

-800 -900 -1 ooo P./mv

Fig. 5. Effect of t d 0n the l og [ i / ( i L - i)] vs. E. First wave pH = 4.12. (a) t d = 3.39 s; (b) t d = 1.64 s; (c)

td = 0.26 s.

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1oql~l

0.0

-I .0

a b c

.-800 -900 -I 000 E/mY

Fig. 6. Effect of T on the log[i/(i L - i)] vs. E plots. Firs t wave p H = 4.17. (a) T = 7.3°C; (b) T = 29.3°C;

(c) T = 45.6°C.

respectively. On the other hand, only one linear segment of slope - 4 4 m V / p H unit in the 7.5-10 pH range, is observed for the second wave.

The E l l 2 shifts to more negative potentials at decreasing i d. The E l ~ 2 vs . log t d

plots are linear of slopes 16, 22, 48, 24 and 11 mV at pH = 0.66, 1.43, 4.12, 6.08 and 7.07 respectively, for the first wave and 29 mV at pH = 9.96 for the second one.

The log[i/(i L - i)] vs. E plots at various pH values for the first wave are shown in Fig. 4. In the 0-1.4 pH range an invariable slope of - 3 0 mV is obtained for these plots. In this zone the instantaneous current i is independent of the square root of the height of mercury column h 1/2, a t the foot of the wave, but varies linearly for i / i L > 0.2. In the 1.8-3.5 and 5.8-6.4 pH ranges two linear segments are obtained. However, in the 3.5-5.8 pH range two linear segments connected by a curved one can be observed. In these cases, the size of the initial segment decreases with increasing pH. For pH > 6.4 an invariable slope of - 4 0 mV is obtained. For the second wave, these plots show one or two linear segments depending on the pH.

In Figs. 5 and 6 are shown the log[i/(i L - i ) ] vs. E plots at various values of td(dro p time) and T(temperature) respectively. As can be seen, a shift to more negative potentials and a diminution of the initial linear segment is observed when decreasing t d o r increasing T.

(a)

-900

-900

-700

1

I i

- I

-900

-700

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(d) -1100

(±) _ _ . . . ,

-900

(e) _2" \ ( j)

Fig. 7. Differential pulse polarograms. Variation with pH. (a) pH = 3.24; (b) pH = 3.64; (c) pH = 3.89; (d) pH = 4.10; (e) pH = 4.30; (f) pH = 4.54; (g) pH--4.82; (h) pH = 5.05; (i) pH = 5.32" (j) pH = 5.94.

Differential pulse polarography

In Fig. 7 the d i f ferent ia l pu l se p o l a r o g r a m s o b t a i n e d in the 3 . 2 - 6 . 0 p H range are shown. In the 0 - 3 . 6 p H range on ly one peak o f c o n s t a n t he ight is observed , but

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from this pH value one peak and one or two shoulders are found, whose variation with pH is rather anomalous.

Determination of the number of electrons

The number of electrons taking part in the overall reactions was calculated by two procedures; one consisting in analyzing the decrease of the limiting current with the time of the electrolysis [9], and the other in analyzing the decrease of the current with the time of microelectrolysis of one drop, using, in this case, a hanging mercury electrode [10]. In both procedures, the applied constant potential was that corre- sponding to the diffusion one. In all cases, the values obtained were close to 2.

C - E curves

Figures 8 and 9 show the C - E curves of MEF (dotted lines) and those corre- sponding to the supporting electrolyte (continuous lines). As can be seen, in acid

~E

~t

c..)

30

\

%

..... , , J .... J. ,, |

-0.zoo -o.4oo -o.6oo -o.8oo E/V

Fig. 8. Variation of C with E' (a) 0.1 M HCI' (b) 0.1 M HCI+ 10 -3 M MEF. Arrow indicates the potential of initiation of the reduction.

349

c~, E

45

30

15

\

\ "\

o

\ \

o

~ o

I l . . . . . I _ ~ •

- 0 . 2 0 0 - 0 . 4 0 0 - 0 . 6 0 0 - 0 . 8 0 0 - 1 . 0 0 0 Ely

Fig. 9. Variat ion of C with E: (a) 0.1 M KCl; (b) 0.1 M K C l + l0 - 3 M MEF. Arrow indicates the potent ia l of init iat ion of the reduction.

medium, the different capacity between both solutions indicates that the MEF is adsorbed on the electrode at potentials close to the reduction. On the other hand, in neutral medium, both curves overlap at potentials close to the reduction, indicating therefore that either weak or negligible adsorption occurs.

Tafel slopes and reaction orders at the foot of the waves

The i -E curves at the foot of the polarographic waves have been registered point by point. From the E vs. log i plot, the Tafel slopes were obtained, and from the

log (il,~ a b

- 0 . 5

- 1 . 0

-1 .5

350

J . . . . . . . i ,.

- 4 5 0 - 5 5 0 - E / m V

Fig. 10. Representa t ion of Tafel 's law. First wave. (a) p H = 0.45" (b) p H = 0.68; (c) p H = 0.93" (d) p H = 1.18.

log li/

0.0

- 1 . 0

- 2 . 0

0 . 5 - 1 .0 pH

Fig. 11. Varia t ion of log i with p H ' (a) E - - 5 4 0 mV; (b) E = - 5 2 0 mV.

Tafel s l ope /mV

Order with respect

to H + ion 3.0

Order with respect

to M E F 1.0

351

pH

< 2 2 - 4 4.4-5.0 5.6-6.2 6.4-8.2 > 9.5

- 3 1 - 3 1 - 3 9 - 5 9 - 4 0 - 5 6

Variable 2.2 2.1 1.0 1.0

1.0 1.1 1.0 1.0 1.0

~A) o ~o~ (i/

- 0 . 5

- 1 . 0

- t .5

- 2 . 0

TABL E 1

Kinetic parameters at different pH ranges

1 , , 1 ,

-4 .0 -~ .o log (~MEF/M)

Fig. 12. Variation of log i with log eMEF; pH -- 1.80. ( O ) E --- - 6 4 0 mV; (e) E = - 6 2 0 mV.

352

log i vs. log CMEF or log[H ÷] plot, at constant potential, the corresponding reaction orders were calculated as well. Typical plots of log i vs. E and log i vs. pH are shown in Figs. l0 and 11 respectively. A noticeable difference in the results thus obtained at different pH's is observed (Table 1).

Figure 12 shows the log i vs. log CME f plot corresponding to the first reduction wave of MEF at pH = 1.8. A linear plot of slope 1 is obtained for values up to 3 X 10 - 4 M approximately, but starting from these values, the slopes of the corresponding lines approach asymptotically to zero, although the Tafel slopes and the reaction orders with respect to H ÷ ion remain unchanged. For pH > 4, this deviation from linearity does not occur, at least in the 10-5_ 10-3 M range.

DISCUSSION

The variation of i e with pH (Fig. 2) is analogous to that observed in other weak organic acids [11-13] and is due to the reduction of both undissociated (with preceding recombination between the dissociated form and H ÷ ion) and the dissoci- ated forms o%

C--C - - H

/ II / /o H ~ C ~ C

(M-)

I second E 2 wove

H ÷ ~--~

pK =3.40

o ~ __C - - C ~ H /

H --c \oH

(MH)

first El wcive

In the 0-1,4 pH range, slopes of - 30 mV and - 93 m V / p H unit obtained in the log[ i / ( iL- i)] VS. E and El~ 2 vs. pH plots respectively, indicate that the wave is produced by a polarographically reversible bielectronic stage in which three protons are taking place. However, the fact that i is independent of h 1/2 at the foot of the wave and tha t it varies linearly for i / i L > 0.2 confirms that the process under study is of a postkinetic type with a rapid chemical reaction [14]. This conclusion is confirmed by the variation of El~ 2 with log td which is linear with a slope of 16 mV [151.

In the 1.4-3.5 pH range and at potentials corresponding to the foot of the wave, a similar behavior is observed. However, the logarithmic analysis and the slope of the E1/2 vs. log t d plot indicate that for i / i e > 0.2, the electronic transfer is possibly irreversible.

On the other hand, an anomalous behavior is observed in the 3.2-5.3 pH range, although the process continues to be governed by diffusion at potential correspond- ing to the limiting current. So, a noticeable change exists in the shape of the wave (Fig. 1). This change is also made evident by means of the logarithmic analysis (Fig. 4), the considerable high values of the slopes of the E1/2 vs. pH and log t d plots and the anomalous shape of the recording obtained by DPP.

Differential capacity curves (Figs. 8 and 9) at either acidic or neutral pH show that MEF adsorption decreases as the potential becomes more negative, that is, as

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the pH increases. Therefore, a greater influence of adsorption on the electrode process is expected in the vicinity of the maximum adsorption potential, namely at pH values < 3.2. Nevertheless, the rds of the electrode process is a chemical stage (and may be homogeneous one) following the transfer at these pH values. But in the pH range of 3.2-5.3 the rds turns out to be a heterogeneous transfer on which the adsorption phenomena exert a greater influence. Therefore, all the phenomena described above could be attributed to this fact.

There are two facts which are against this hypothesis: (1) At the potential of - 1000 mV, adsorption is negligible (Fig. 9). It is from this

potential value that the effect observed on the peaks recorded for DPP is greatest (Fig. 7).

(2) Diethyl fumarate at acidic pH exhibits a mechanism resembling that of MEF with a chemical rds following the electric transfers. Guidelli and co-workers [6] have shown this rds to have a heterogenous nature, which, according to the similarity of the two products contradicts the reasoning that we have made.

The variation of l o g [ i / ( i e - i)] as a function of E (Fig. 4) and the peak shapes of the DPP suggest an alternative possibility: the occurrence of one or two protona- tions prior to electronic transfers [13,16-20]. However, our experimental results do not allow us to make a clear-cut choice.

In the 5.3-9.5 pH range, the limiting current of the first reduction wave decreases at increasing pH. From pH > 6.5 the slope of log i e vs. log t d plots is > 0.19 and a deviation of the El~ 2 v s . pH plot from a straight line of - 3 6 mV slope is observed. This is due to the kinetic character of the process at potentials corresponding to the limiting current. A solution of the appropriate boundary-value problem has been made for us [21] to explain the anomalous variation of El~ 2 v s . pH in this zone.

From the above-mentioned data and conclusions, and assuming, as the cases of fumaric and maleic acids and their corresponding diethyl esters [1-6], that the reaction product is.monoethyl succinate in all cases, the overall reaction scheme at the foot of the wave may be represented by the following sequence of steps

First wave. p H < 4

[MH]aa + 2 H + + 2 e- KI

[RH']., (])

k (II) [mHt]a d --~ [mH]dis rds

where reaction (II)represents the rate-determining step. In a manner similar to that of diethyl fumarate [6], we think that this reaction is the enol-keto transformation, which is general-acid catalyzed"

HO\ / OH HO c cH C H i C + .A ~ \ C j 6 H _

HO , / ~OEt rds HO~, ----CH~CH ~ C , ~ o E t (RH')

NO\ C~ 6H c//° + HA C -----C H - - C H ~ + A- ~ HON" C---- CH ~ C H 2 ~ HO / ~OEt HO / OE t

RH

354

where HA and A - denote any acid and base present in the solution respectively. At pH < 2, the H30 + ion is the main proton donor and thus, the reaction order with respect to H + ion is 3 and the slope of the El~ z vs. pH plot is - 9 3 m V / p H unit. In the 2 -4 pH range, the reaction order with respect to H + ion decreases gradually from 3.0 to 2.5.

Therefore, assuming reaction (II) is the rds and thus considering as negligible the concentration of those species appearing before the rds, i.e. their 0 > 0, the cathodic current value at the foot of the wave may be expressed as

i = 2 F k K K ' c n2 0 MVFCHA exp( - 2 F E / R T ) (1) where k is the rate constant of the rds K the equilibrium constant of reaction (I), c n and ¢HA the concentrations of H + ion and acid present in solution respectively, OME v the coverage of the MEF and

K ' - - exp(2FA~ar~f/RT )

where A~bre f is the potential of the reference electrode. The deviations of the log i vs. log CME F plots from a straight line with a slope of 1

are due to adsorption of the reactant. At a sufficiently low CME F value, a proportion- ality relationship between i and CME v exists, but at very high CME F values, the natural tendency towards the attainment of full coverage (0ME F ~ 1) causes i to become independent of ¢MEF"

First wave. 4. 4 < p H < 5.0

gl M - + H + ~ MH (I)

g 2 MH + H + + e- ~ [RI" (II)

k [RI-+ e - ~ R - (III)

rds

Therefore, assuming the adsorption of MEF on the mercury electrode surface as negligible, i.e. 6ME F ------0, the i - E relationship may be expressed as

i = 2FkKK'CMEF e2 e x p [ - (1 + fl) F E / R T ] (2)

where the same symbology as above is used and

g t = exp [ ( 1 + fl ) F At~r e f / R T ]

where fl is the symmetry factor.

First wave. 5.6 < p H < 6.2

Where the chemical reaction situated between the one-electron transfers is the rds. Ki

M - + n + ~ MH (I)

K 2

MH + e- ~ [MH]-"

_. k [MH] + H + ~ [ R ] .

rds

[ R ] - + e - ~ R-

The i - E relationship may be expressed as

i = FkKK'CMEF c2 e x p ( - F E / R T )

where

K' = exp( r Aeor e f / R T )

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(II)

(III)

(IV)

(3)

First wave. 6. 4 < p H < 8.2

g l M - + H + ~ MH (I)

M H + e + K2.___. [MH]- (II)

_, k [MH] + e- ~ [MH]= (iII)

rds

K4 [MH]= + HA ~ R - + A - (IV)

Where the second one-electron transfer is again the rds and, therefore, the i - E relationship may be expressed as

i = 2FkKK'CMEFC H exp[-- (1 + f l ) F E / R T ] (4)

where

K' = exp[(1 + fl).54)]

Second wave. p H > 9.5

K l

M - + e- ~ [M]- (I)

.._. k [M] + H + ~ [M]- (II)

rds

K 3

.H.A + [ M ] - + e- ~ R - + A - (III)

The i - E relationship may be expressed as

i = FkKK'CMEFCrI exp( - F E / R T ) (5)

356

where

K ' = exp( F Ackref/RT )

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