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The reduction mechanism of the >CO group —III. The electrochemical reduction of isonicotinic acid in an aqueous medium
Electrochimica
Pergamon PII: !soo13-4686(96)00189-2
Acfo, Vol. 42, No. 2, pp. 331-336,
1997
Copyright 0 1996 Elswier Science Ltd. Printed in Great Britain. All rights resewed OOlW686/97
$17.00 + 0.00
The reduction mechanism of the >C=O group -111. The electrochemical reduction of isonicotinic acid in an aqueous medium E. Mathieu, R. Meunier-Prest Laboratoire
and E. Laviron
de Synthbse et d’Electrosynthtse Organometalliques, 21000 Dijon, France
CNRS-URA
1685, 6 bd. Gabriel,
(Received 10 April 1996) Abstract-The electrochemical reduction of isonicotinic acid NRCOOH and its protonated from H+NRCOOH has been examined in an aqueous medium between Hc, = 0 and pH 6.7. As shown by cyclic voltammetry and polarography, a global 2e-, 2H+ reversible transfer is followed by two successive first-order chemical reactions and a 2e-, 2H+ reduction (ECiCzE process). A mechanism is proposed, beginning with the 2e-, 2H+ reduction of YRCOOH (where Y is N or H+N) to the formal diionized forms YRC-(OH)(OH: ), where the first chemical reaction Ci is an internal proton transfer which yields the hydrate YRCH(OH)z; this reaction, which is analogous to the internal proton transfer taking place in the case of nitronic acids in acidic medium has been shown previously to occur in the case of methyl isonicotinate. The rate constant of this process is 2 x 102s-i in neutral medium and 2 x 104s-i in acidic medium. The second chemical reaction Ca (dehydration), which is much slower (rate constant kd < 1.6 x 10m3s-l), involves the loss of a Hz0 molecule to give 4_formylpyridine, which is itself easier to reduce than isonicotinic acid. The variations of the total current and the dehydration constant with pH are analyzed in detail, and compared to those observed for Cformylpyridine. Copyright 0 1996 Elsevier Science Ltd Key words: isonicotinic
acid, Cformylpyridine,
reduction mechanism, hydration
PC.ZHf
INTRODUCTION
YRCOZ -
Previous electrochemical studies of isonicotinic acid NRCOOH [l-8] showed that, globally, a 2ereduction of the c=O bond yields the hydrate of Cformylpyridine, which loses more or less rapidly a molecule of water, thus giving the aldehyde which is immediately reduced (ECrE reaction; cJ equation (a), in which Y represents H+N or N, and Z = OH). YRCOZ 2
2H+
YRCHZOH.
- ZWkd,
(4)
P7
YRCH(OH)Z
c2
2.-.2H+
YRCHO -
E
YRCHzOH.
(b)
We report in this paper the results of our study of isonicotinic acid. EXPERIMENTAL Isonicotinic acid was a commercial product. The experiments were carried out in water at a temperature of 20°C. For pH > 1.80, BrittonRobinson buffers were used with 0.8 M NaCl as supporting electrolyte, and the pH values were measured. For pH < 1.80, HrS04 was employed and the I-b values were taken from literature data (lo]. The potentials are referred to the saturated calomel electrode (see). For pH < 1, the error due to the junction potential and the viscosity were corrected as described previously [ 111.
YRCH(OH)Z G YRCHO 5
E
(a)
A detailed analysis of the reduction of the analog of isonicotinic acid, methyl isonicotinate [9] (equation (a), with Z = OCH,) led to the conclusion that the reaction actually proceeds through a more complex sequence ECiCzE; Ci is also first-order and much faster than Cz (equation (b)) 331
E. Mathieu et al.
332
-87mV
/
-800
G zi
-loo0
-1 200
Fig. 3. Variations of Epa and Ew with log u for isonicotinic acid at HO= 0.7; Cr = 1O-4 M. Voltammograms (k upwards, ia downwards): (a) 0.8 V s-r, (b) 60 V s-r.
-1 400 Om\
-4
I
I
I
I
I
I
I
-2
0
2
4
6
8
10
Cyclic voltammetry
Hoor pH Fig. 1. Variations of E o (O), E O’(0) and Et/z (m) as a function of pH; r = 0.5 s.
The equipment used in earlier work [l I] was employed to obtain the voltammograms and the polarograms. In polarography, the drop time r was 0.5 s and the rate of capillary flow 1.89 mg s-l.
RESULTS Polarography
The variations of the half-wave potentials are shown in Fig. 1; a second wave appears at certain pH values; we have studied only the first. The limiting currents of this wave vary with pH as shown in Fig. 2. We have also indicated the variations of the limiting current of 4-formylpyridine (cJ Ref. [12]), obtained under the same conditions. 5
Examples of variations of the cathodic and anodic peak potentials Ew and Epa with log v are shown in Figs 3 and 4. Below pH 2.2, the shape of the peaks indicate a diffusion-controlled process when log v is smaller than about 0.5 (Fig. 3); for larger values the reaction is of a surface nature. Above pH 2.2, the peaks have a “diffusional” shape whatever the sweep rate (Fig. 4). The changes of the slopes will be analyzed in the discussion, but it must be noticed that the equilibrium potential for the surface (P’) or heterogeneous (P) reaction can easily be obtained in the region where both the cathodic and anodic peaks are seen (Figs 3 and 4). This occurs when v > 1000 V s-i below pH 2.2, or when v > 20 to 100 V s-l above it. DISCUSSION The first 2e- electrochemical reaction
The first 2e- electrochemical reaction can be described by using [13] a tridimensional scheme (Fig. 5) in which the molecules protonated on the pyridine nitrogen lie in the bottom plane. The notations are those used previously [13]; if the protonations are fast (at equilibrium), the kinetics of the system can be treated mathematically [13-151; -850 ‘I
i0
0
,
I5
Ho or pH Fig. 2. Variations of the limiting current (plateau of the polarographic wave) for isonicotinic acid b (+) and 4-formylpyridine it. (m); CT = 5 x lO-4 M, r = 0.5 s. (-) curves simulated according to Nicholson et al. [20] and according to Koutecky [16,24], with log kd = - Ho - 0.39 forHa< -O.landlogkd= -0.29forHo> -0.1.
Fig. 4. Variations of EpI and EF with log u for isonicotinic acid at pH 2.9; Cr = 10e4 M. Voltammograms (ic upwards, i. downwards): (a) 1 V s-l, (b) 100 V s-r.
Reduction mechanism of the > C=O group-111 +mR&g
+“NRC&
333
The chemical reactions
Generally speaking, the presence of a chemical reaction with a rate constant k inserted between two electrochemical reactions (El CE2 mechanism, with EL reversible) causes two types of effect (we shall consider the case of a reduction). (a) The half-wave potential El,1 or the peak potential Ew of the cathodic LSV peak is shifted towards positive potentials. If the rate of the chemical reaction is large enough (purely kinetic case), the shift is given by
log 0.886 + slogkr
E,,2 = I?’ + =$
Q OH +HNRK / 0”;
(1)
E,=&O.78$++logg
(2)
P&t / Q OH NRC: OH;
III k
A-
l8PP .
E rl
A- +
k
for a heterogeneous El reaction followed by a volume C reaction [16, 171, and
. r2
E,,z = I?” - y
A=
log kz (3)
log 1.104 + F
III E,=E”+
Fig. 5. The tridimensional scheme for the first 2 e- reduction of isonicotinic acid.
the solutions obtained are valid for a heterogeneous reaction, or for a surface reaction when a Langmuir isotherm is obeyed. The scheme is kinetically equivalent to two simple successive reactions, with apparent rate constants klapp and kbpp and apparent standard potentials & and Er2(Fig. 5); when the two stages overlap (Erl > Er2), it is thus equivalent to a simple 2e- reaction, with an apparent equilibrium potential Ep = 0.5 (E,, + Er2). In the present case, k,t and k3 are so large that the electrochemical reaction is always at equilibrium except at high sweep rates. The equilibrium diagram, deduced from the values of E? and F’, is given in Fig. 6.
(4)
when both El and C are surface reactions [18, 191. Equations (1) and (3) are valid for mean currents. In LSV, the anodic peak then disappears. In the present case, the slopes of the asymptotes at slow sweep rates do indicate that the first 2e- reaction is followed by a first order chemical reaction. The slope is 14.5 mV/log u when the electrochemical reaction is of a heterogeneous nature or 29 mV when it is of a surface nature (cf. equations (2) and (4) and Figs 3 and 4). The rate constant, which we shall designate as kl, can be determined in certain zones of pH (Fig. 7) when the appropriate asymptote can be associated with E” or I?“, using equations (2) and (4). (b) The second effect is an increase in the total
-800-
The nature of the intermediate P The fact that the first 2e- reduction is reversible when u is large enough demonstrates the impossibility of a reduction to the dihydrate YRCH(OH)z (cf. equation (a)), since this compound is not oxidizable in the potential range available on mercury [12]. Therefore, the reaction sequence is that represented by equation (b). Formally, as shown in [9], the intermediate P can be written as a molecule in which the carbon of the substituent bears a minus charge, protonation occurring on the pyridine nitrogen or on the oxygen of the group (cf. Fig 5). Another possibility is the existence of quinoidal structures [9].
-!900$ -lOOOw -I loo-
-IZoO-
-2
-1
0
1
L
,
4
3
b
I
8
H,or pH Fig. 6. Equilibrium diagram for isonicotinic acid: (e), EO; (0). EO’.
E. Mathieu et al.
334
4
K;
II (2)
k,
+HNC5H4CH0
K;
=
(3)
(‘)kdl
II NC5H4CH0
5 Fig. 9. The protonation and hydration equilibria. I
0-l
0
fl
6
2
Fig. 7. Variations of the logarithm of the rearrangement rate constant k~ as a function of pH: (0) from 81~ - E O,(m) from log Vifor a volume reaction (cf. Fig. 4); (0) from log u, for a surface reaction (cJ Fig. 3).
current in over that id due to the first process. For mean currents, the ratio has been calculated numerically as a function of kr for the E(heterogeneous) C(volume) E process [20] and for the E(surface) C(surface) E process [18]. Using these theories, we have calculated a constant which we shall designate as kd (Fig. 8). The constant kl is always much larger than the constant kd. Since the dehydration rate constants determined previously [7, 12,21-231 are of the order of magnitude of kd, this means that the dehydration (rate constant kd) is preceded by the first-order reaction whose rate is kl, as in the case of methyl isonicotinate (equation (b)). The dehydration reaction Before discussing this reaction, we shall recall the behavior of Cformylpyridine. The limiting current of the 2e- reduction wave (Fig. 2) is determined [12] by the equilibrium between the free and hydrated forms (Fig. 9), and by the rate of the dehydration. Four zones can be defined [12]. In zone II, the aldehyde exists mostly in its hydrated form, and the current is controlled by the
0
-3-i. 4
I.. -2
. 0
I. 2
I. 4
I. 6
1 8
Ho or pH Fig. 8. Variations of the logarithm of the dehydration rate constant kd with pH. (W) our results, (0) [12], (0) [I. The equations of the straight lines are log kd = - Ho - 0.39 for HO< - 0.1 and iogkd = - 0.29 for HO> - 0.1.
dehydration rate. In zone III, the dehydration rate is small, so that it is not the controlling factor determining the limiting current; the part of the current due to the fraction of the free aldehyde existing at equilibrium in solution is large. In zones I and IV, the current increases, because the dehydration is acid- and base-catalyzed, respectively. In zone V, a decrease is observed, owing to the deprotonation of the hydrate (Fig. 9). In the case of isonicotinic acid, the situation is not the same as for the aldehyde, since the hydrate is created near the electrode instead of being in equilibrium with the free form in solution, so that the current (in, Fig. 2) in excess of that (id) of 2e- wave is entirely limited by the dehydration rate. This is expressed by the fact that the kinetic current & for the ECE reaction depends only on kr , [18,201 whereas & for the CE process for the aldehyde is a function of K;kdr 116, 18,241, K; being K[ or Ki. If follows from these considerations that below pH 3.8, where the concentration of the protonated form of the hydrate is large and constant at equilibrium (cJ Fig. 9), ik is larger than r& (Fig. 2). Below pH 0, ik increases because of the acid catalysis. Between pH 0 and pH = pKA = 5.35, ik is small and constant because the protonated hydrated species H+NRCH(OH)r, (rate constant k& is formed directly at the electrode. Between pH 6 and 7 ik = 0, because the species created at the electrode is now the unprotonated form NRCH(OH)r, whose dehydration (rate constant kdl) is too slow [ 121to yield the free aldehyde in sufficient amount during the life of the drop. By contrast, in the case of the aldehyde, a large increase of current is observed above pH 3, because a large amount of the aldehyde at equilibrium with the hydrate is present in the solution. We have calculated an upper limit of the dehydration constant between pH 6 and 7 by studying the product ilr- ‘I6(Fig. 2) as a function of log r; the back-pressure, depletion and spherical diffusion effects were corrected for as in [9], paragraph 4.2. The production ilf-“6 was found to be invariant up to r = 20 s, which allows us to establish that kdl < 1.6 x lOA s-l [9]. Using then the general theory for CE currents in polarography [25], it is easily found that the kinetic fraction in the limiting current of the aldehyde (ii., Fig. 2) is smaller than 1%, ie the current is practically entirely controlled by diffusion of the free aldehyde.
Reduction mechanism of the > C==O group-111
335
Table 1 Values of Ki, Ki and kdl and kd2, and percentages of the different forms Percentage pH < 3 Reference This work t71
1121 WI 1231
Ki
Ki
kd>S-’
ka s-’
Percentage pH > 7
+HNCsHXH(OH)z +HNCsHXHO NC5H&H(OH)z NCsHKHO
3.6 lO-3
1
0.51
< 5 10-1
99.6
0.4
50
SO
22 lo--’ 18 IO--)
1.17 1.23 1.28
0.26 0.08
0.14
97.8 98.2
2.2 1.8
46 45 56
54 55 44
0.11
0.6
0.01
0.04
98.7
1.3
62
38
Above pH 6, the current il decreases rapidly, because the species NRCO; is not reducible in the accessible potential range. Calculation of the dehydration rate constant kdr between pH 0 and pH = pKad = 5.35 from the ratio (k + id)/& =f(/u), using the theory of Nicholson et al. for an ECE process [20] gave kz = 0.51 s-’ (cf. Fig. 8). The equilibrium constant KS can then be deduced from Koutecky’s theory from 4, = f(K$kd) [24, 251; a value of KS = 3.6 x 10e3 is found (Table 1). It can then be calculated that the percentage of free aldehyde is 0.3%, ie that of the hydrate is 99.7%. From Fig. 2, we can deduce that ika = (id/z); all the current is due to the reduction of the fraction of aldehyde at equilibrium. The percentage of aldehyde and hydrate is thus 50%, ie K[ = 1. From the relationship K;Ki = K$K;, we can deduce that the value of pKA3is 2.9 with the above values pK& = 5.35, K; = 1 and K; = 3.6 x 10A3.The value of pKL3is in reasonable accord with that found by one of us [12], ie pKi;3 = 3.36. The complete curve for the percentages of the diverse species can easily be calculated as a function of pH (Fig. 10). We have reported in Table 1 the values which we have determined and those found in the literature for the diverse constants. The values of kdi found in the literature are quite inaccurate, since the upper limit which we have found (5 x lo-’ s-l) was derived from the invariance of the product i1@ (c$ above).
100
Our values of kd2 and KS are much more precise than those of the literature, since KS need not be known to calculate kd2 (ECE process for the acid; cf. above), and is subsequently used to determine K$ uiu the product Kikd2T (CE process for the aldehyde; cf. above). By contrast, previous measurements using the current for the aldehyde between pH 0 and 4 [7, 12, 231 (Fig. 2) give the product K;kdz, so that an estimation of K; must be made; it is necessarily very inaccurate, because the equilibrium is much shifted toward the hydrate. Our measurements do not give K; with a better precision than previously, because its value is in every case obtained directly from the ratio of the limiting current of the aldehyde to the global diffusion current between pH 6 and 8 (Fig. 2).
CONCLUSION The electrochemical reduction of isonicotinic acid is more complex than thought hitherto. Reversible reduction of the acid group is followed by two successive first order reactions, viz. a rearrangement and a dehydration. It has been shown that in the case of an EC reaction with adsorption of the reactants [26], the meaning of the chemical rate constant which is measured by using electrochemical methods is not straightforward; it is not possible a priori to know if it is a surface or “volume” rate constant. For an ECCE reaction, the situation is even more complex; in particular, we cannot establish the nature of kl. However, the slope - 1 of log kd =f(Ho) below Ho = 0 indicates c1priori that the dehydration in this region is a “volume” reaction [26]; it is therefore probable that the same is true also for kd2 when Ho > 0.
REFERENCES
Fig. 10. Percentages a function of pH.
of the diverse forms of the aldehyde
as
1. Y. Nagata and Tachi, Bull. Chem. Sot. Jpn 21, 290 (1954). 2. J. Volke and V. Volkova, Coil. Czechoslov. Chem. Commun. 20, 1333 (1955). 3. H. H. G. Jellinek and J. R. Urwin, J. Phys. Chem. 58, 168 (1954).
336
E. Mathieu et al.
4. H. Lund, Acta Chem. &and. 17, 972 (1963). 5. Wang Zheng-Hao and Hu Zhi-Bin, Electrochim. Acta 30, 779 (1985). 6. R. W. Green and H. Tong, J. Am. Chem. Sot. 78, 4896 (1956). 7. M. U. D. Bhatti and 0. R. Brown, J. Electroanal. Chem. 68, 85 (1976). 8. R. Rodriguez-Amaro, R. Perez, L. Camacho and J. J. Ruiz, J. Electroanal. Chem. 324, 270 (1992). 9. E. Laviron, R. Meunier-Prest and E. Mathieu, J. Electroanal. Chem. 371, 251 (1994). 10. C. H. Rochester, Acidity Functions (Edited by A. T. Blomquist), p. 17. Academic Press, New York (1970). 1I. I?. L&iron, R. Meunier-Prest, A. Vallat, L. Roullier and R. Lacasse, J. Electroanal. Chem. 341. 227 (1992). 12. E. Laviron, Buil. Sot. chim. Fr. 2325 (1961). . ’ 13. R. Meunier-Prest and E. Laviron, J. Electroanal. Chem. 328, 33 (1992). 14. E. Laviron, J. Electroanal. Chem. 109, 57 (1980); 124, 1 (1981). 15. E. Laviron, J. Electroanal. Chem. 146, 15 (1983).
16. J. Koutecky, Collect. Czech. Chem. Commun. 20, 116 (1955). 17. L. Nadjo and J. M. Saveant, J. Electroanal. Chem. 48, 113 (1973). 18. R. Meunier-Prest and E. Laviron, J. Electroanal. Chem. in press. 19. E. Laviron, J. Electroanal. Chem. 35, 333 (1972). 20. R. S. Nicholson, J. M. Wilson and L. M. Olmstead, Anal. Chem. 38, 542 (1966). 21. J. F. Rusling, J. P. Segretario and P. Zuman, J. Electroanal. Chem. 143, 291 (1983). 22. E. G. Sander and W. Jencks, J. Am. Chem. Sot. 90, 6154 (1968). 23. A. Vallat, M. Person and E. Laviron, Electrochim. Acta 24, 1125 (1979). 24. P. Delahay, New Instrumental Methoak in Electrochemistry. Interscience, New York (1954). 25. J. Koutecky and R. Brdicka, CON.Cr. Chem. Commun. 12, 337 (1947); E. Laviron and M. Vincent, Polarography 1964 (Edited by G. J. Hills), vol. 1, p. 397. Macmillan, London (1966). 26. E. Laviron, J. Electroanal. Chem. 391, 187 (1995).
Pergamon PII: !soo13-4686(96)00189-2
Acfo, Vol. 42, No. 2, pp. 331-336,
1997
Copyright 0 1996 Elswier Science Ltd. Printed in Great Britain. All rights resewed OOlW686/97
$17.00 + 0.00
The reduction mechanism of the >C=O group -111. The electrochemical reduction of isonicotinic acid in an aqueous medium E. Mathieu, R. Meunier-Prest Laboratoire
and E. Laviron
de Synthbse et d’Electrosynthtse Organometalliques, 21000 Dijon, France
CNRS-URA
1685, 6 bd. Gabriel,
(Received 10 April 1996) Abstract-The electrochemical reduction of isonicotinic acid NRCOOH and its protonated from H+NRCOOH has been examined in an aqueous medium between Hc, = 0 and pH 6.7. As shown by cyclic voltammetry and polarography, a global 2e-, 2H+ reversible transfer is followed by two successive first-order chemical reactions and a 2e-, 2H+ reduction (ECiCzE process). A mechanism is proposed, beginning with the 2e-, 2H+ reduction of YRCOOH (where Y is N or H+N) to the formal diionized forms YRC-(OH)(OH: ), where the first chemical reaction Ci is an internal proton transfer which yields the hydrate YRCH(OH)z; this reaction, which is analogous to the internal proton transfer taking place in the case of nitronic acids in acidic medium has been shown previously to occur in the case of methyl isonicotinate. The rate constant of this process is 2 x 102s-i in neutral medium and 2 x 104s-i in acidic medium. The second chemical reaction Ca (dehydration), which is much slower (rate constant kd < 1.6 x 10m3s-l), involves the loss of a Hz0 molecule to give 4_formylpyridine, which is itself easier to reduce than isonicotinic acid. The variations of the total current and the dehydration constant with pH are analyzed in detail, and compared to those observed for Cformylpyridine. Copyright 0 1996 Elsevier Science Ltd Key words: isonicotinic
acid, Cformylpyridine,
reduction mechanism, hydration
PC.ZHf
INTRODUCTION
YRCOZ -
Previous electrochemical studies of isonicotinic acid NRCOOH [l-8] showed that, globally, a 2ereduction of the c=O bond yields the hydrate of Cformylpyridine, which loses more or less rapidly a molecule of water, thus giving the aldehyde which is immediately reduced (ECrE reaction; cJ equation (a), in which Y represents H+N or N, and Z = OH). YRCOZ 2
2H+
YRCHZOH.
- ZWkd,
(4)
P7
YRCH(OH)Z
c2
2.-.2H+
YRCHO -
E
YRCHzOH.
(b)
We report in this paper the results of our study of isonicotinic acid. EXPERIMENTAL Isonicotinic acid was a commercial product. The experiments were carried out in water at a temperature of 20°C. For pH > 1.80, BrittonRobinson buffers were used with 0.8 M NaCl as supporting electrolyte, and the pH values were measured. For pH < 1.80, HrS04 was employed and the I-b values were taken from literature data (lo]. The potentials are referred to the saturated calomel electrode (see). For pH < 1, the error due to the junction potential and the viscosity were corrected as described previously [ 111.
YRCH(OH)Z G YRCHO 5
E
(a)
A detailed analysis of the reduction of the analog of isonicotinic acid, methyl isonicotinate [9] (equation (a), with Z = OCH,) led to the conclusion that the reaction actually proceeds through a more complex sequence ECiCzE; Ci is also first-order and much faster than Cz (equation (b)) 331
E. Mathieu et al.
332
-87mV
/
-800
G zi
-loo0
-1 200
Fig. 3. Variations of Epa and Ew with log u for isonicotinic acid at HO= 0.7; Cr = 1O-4 M. Voltammograms (k upwards, ia downwards): (a) 0.8 V s-r, (b) 60 V s-r.
-1 400 Om\
-4
I
I
I
I
I
I
I
-2
0
2
4
6
8
10
Cyclic voltammetry
Hoor pH Fig. 1. Variations of E o (O), E O’(0) and Et/z (m) as a function of pH; r = 0.5 s.
The equipment used in earlier work [l I] was employed to obtain the voltammograms and the polarograms. In polarography, the drop time r was 0.5 s and the rate of capillary flow 1.89 mg s-l.
RESULTS Polarography
The variations of the half-wave potentials are shown in Fig. 1; a second wave appears at certain pH values; we have studied only the first. The limiting currents of this wave vary with pH as shown in Fig. 2. We have also indicated the variations of the limiting current of 4-formylpyridine (cJ Ref. [12]), obtained under the same conditions. 5
Examples of variations of the cathodic and anodic peak potentials Ew and Epa with log v are shown in Figs 3 and 4. Below pH 2.2, the shape of the peaks indicate a diffusion-controlled process when log v is smaller than about 0.5 (Fig. 3); for larger values the reaction is of a surface nature. Above pH 2.2, the peaks have a “diffusional” shape whatever the sweep rate (Fig. 4). The changes of the slopes will be analyzed in the discussion, but it must be noticed that the equilibrium potential for the surface (P’) or heterogeneous (P) reaction can easily be obtained in the region where both the cathodic and anodic peaks are seen (Figs 3 and 4). This occurs when v > 1000 V s-i below pH 2.2, or when v > 20 to 100 V s-l above it. DISCUSSION The first 2e- electrochemical reaction
The first 2e- electrochemical reaction can be described by using [13] a tridimensional scheme (Fig. 5) in which the molecules protonated on the pyridine nitrogen lie in the bottom plane. The notations are those used previously [13]; if the protonations are fast (at equilibrium), the kinetics of the system can be treated mathematically [13-151; -850 ‘I
i0
0
,
I5
Ho or pH Fig. 2. Variations of the limiting current (plateau of the polarographic wave) for isonicotinic acid b (+) and 4-formylpyridine it. (m); CT = 5 x lO-4 M, r = 0.5 s. (-) curves simulated according to Nicholson et al. [20] and according to Koutecky [16,24], with log kd = - Ho - 0.39 forHa< -O.landlogkd= -0.29forHo> -0.1.
Fig. 4. Variations of EpI and EF with log u for isonicotinic acid at pH 2.9; Cr = 10e4 M. Voltammograms (ic upwards, i. downwards): (a) 1 V s-l, (b) 100 V s-r.
Reduction mechanism of the > C=O group-111 +mR&g
+“NRC&
333
The chemical reactions
Generally speaking, the presence of a chemical reaction with a rate constant k inserted between two electrochemical reactions (El CE2 mechanism, with EL reversible) causes two types of effect (we shall consider the case of a reduction). (a) The half-wave potential El,1 or the peak potential Ew of the cathodic LSV peak is shifted towards positive potentials. If the rate of the chemical reaction is large enough (purely kinetic case), the shift is given by
log 0.886 + slogkr
E,,2 = I?’ + =$
Q OH +HNRK / 0”;
(1)
E,=&O.78$++logg
(2)
P&t / Q OH NRC: OH;
III k
A-
l8PP .
E rl
A- +
k
for a heterogeneous El reaction followed by a volume C reaction [16, 171, and
. r2
E,,z = I?” - y
A=
log kz (3)
log 1.104 + F
III E,=E”+
Fig. 5. The tridimensional scheme for the first 2 e- reduction of isonicotinic acid.
the solutions obtained are valid for a heterogeneous reaction, or for a surface reaction when a Langmuir isotherm is obeyed. The scheme is kinetically equivalent to two simple successive reactions, with apparent rate constants klapp and kbpp and apparent standard potentials & and Er2(Fig. 5); when the two stages overlap (Erl > Er2), it is thus equivalent to a simple 2e- reaction, with an apparent equilibrium potential Ep = 0.5 (E,, + Er2). In the present case, k,t and k3 are so large that the electrochemical reaction is always at equilibrium except at high sweep rates. The equilibrium diagram, deduced from the values of E? and F’, is given in Fig. 6.
(4)
when both El and C are surface reactions [18, 191. Equations (1) and (3) are valid for mean currents. In LSV, the anodic peak then disappears. In the present case, the slopes of the asymptotes at slow sweep rates do indicate that the first 2e- reaction is followed by a first order chemical reaction. The slope is 14.5 mV/log u when the electrochemical reaction is of a heterogeneous nature or 29 mV when it is of a surface nature (cf. equations (2) and (4) and Figs 3 and 4). The rate constant, which we shall designate as kl, can be determined in certain zones of pH (Fig. 7) when the appropriate asymptote can be associated with E” or I?“, using equations (2) and (4). (b) The second effect is an increase in the total
-800-
The nature of the intermediate P The fact that the first 2e- reduction is reversible when u is large enough demonstrates the impossibility of a reduction to the dihydrate YRCH(OH)z (cf. equation (a)), since this compound is not oxidizable in the potential range available on mercury [12]. Therefore, the reaction sequence is that represented by equation (b). Formally, as shown in [9], the intermediate P can be written as a molecule in which the carbon of the substituent bears a minus charge, protonation occurring on the pyridine nitrogen or on the oxygen of the group (cf. Fig 5). Another possibility is the existence of quinoidal structures [9].
-!900$ -lOOOw -I loo-
-IZoO-
-2
-1
0
1
L
,
4
3
b
I
8
H,or pH Fig. 6. Equilibrium diagram for isonicotinic acid: (e), EO; (0). EO’.
E. Mathieu et al.
334
4
K;
II (2)
k,
+HNC5H4CH0
K;
=
(3)
(‘)kdl
II NC5H4CH0
5 Fig. 9. The protonation and hydration equilibria. I
0-l
0
fl
6
2
Fig. 7. Variations of the logarithm of the rearrangement rate constant k~ as a function of pH: (0) from 81~ - E O,(m) from log Vifor a volume reaction (cf. Fig. 4); (0) from log u, for a surface reaction (cJ Fig. 3).
current in over that id due to the first process. For mean currents, the ratio has been calculated numerically as a function of kr for the E(heterogeneous) C(volume) E process [20] and for the E(surface) C(surface) E process [18]. Using these theories, we have calculated a constant which we shall designate as kd (Fig. 8). The constant kl is always much larger than the constant kd. Since the dehydration rate constants determined previously [7, 12,21-231 are of the order of magnitude of kd, this means that the dehydration (rate constant kd) is preceded by the first-order reaction whose rate is kl, as in the case of methyl isonicotinate (equation (b)). The dehydration reaction Before discussing this reaction, we shall recall the behavior of Cformylpyridine. The limiting current of the 2e- reduction wave (Fig. 2) is determined [12] by the equilibrium between the free and hydrated forms (Fig. 9), and by the rate of the dehydration. Four zones can be defined [12]. In zone II, the aldehyde exists mostly in its hydrated form, and the current is controlled by the
0
-3-i. 4
I.. -2
. 0
I. 2
I. 4
I. 6
1 8
Ho or pH Fig. 8. Variations of the logarithm of the dehydration rate constant kd with pH. (W) our results, (0) [12], (0) [I. The equations of the straight lines are log kd = - Ho - 0.39 for HO< - 0.1 and iogkd = - 0.29 for HO> - 0.1.
dehydration rate. In zone III, the dehydration rate is small, so that it is not the controlling factor determining the limiting current; the part of the current due to the fraction of the free aldehyde existing at equilibrium in solution is large. In zones I and IV, the current increases, because the dehydration is acid- and base-catalyzed, respectively. In zone V, a decrease is observed, owing to the deprotonation of the hydrate (Fig. 9). In the case of isonicotinic acid, the situation is not the same as for the aldehyde, since the hydrate is created near the electrode instead of being in equilibrium with the free form in solution, so that the current (in, Fig. 2) in excess of that (id) of 2e- wave is entirely limited by the dehydration rate. This is expressed by the fact that the kinetic current & for the ECE reaction depends only on kr , [18,201 whereas & for the CE process for the aldehyde is a function of K;kdr 116, 18,241, K; being K[ or Ki. If follows from these considerations that below pH 3.8, where the concentration of the protonated form of the hydrate is large and constant at equilibrium (cJ Fig. 9), ik is larger than r& (Fig. 2). Below pH 0, ik increases because of the acid catalysis. Between pH 0 and pH = pKA = 5.35, ik is small and constant because the protonated hydrated species H+NRCH(OH)r, (rate constant k& is formed directly at the electrode. Between pH 6 and 7 ik = 0, because the species created at the electrode is now the unprotonated form NRCH(OH)r, whose dehydration (rate constant kdl) is too slow [ 121to yield the free aldehyde in sufficient amount during the life of the drop. By contrast, in the case of the aldehyde, a large increase of current is observed above pH 3, because a large amount of the aldehyde at equilibrium with the hydrate is present in the solution. We have calculated an upper limit of the dehydration constant between pH 6 and 7 by studying the product ilr- ‘I6(Fig. 2) as a function of log r; the back-pressure, depletion and spherical diffusion effects were corrected for as in [9], paragraph 4.2. The production ilf-“6 was found to be invariant up to r = 20 s, which allows us to establish that kdl < 1.6 x lOA s-l [9]. Using then the general theory for CE currents in polarography [25], it is easily found that the kinetic fraction in the limiting current of the aldehyde (ii., Fig. 2) is smaller than 1%, ie the current is practically entirely controlled by diffusion of the free aldehyde.
Reduction mechanism of the > C==O group-111
335
Table 1 Values of Ki, Ki and kdl and kd2, and percentages of the different forms Percentage pH < 3 Reference This work t71
1121 WI 1231
Ki
Ki
kd>S-’
ka s-’
Percentage pH > 7
+HNCsHXH(OH)z +HNCsHXHO NC5H&H(OH)z NCsHKHO
3.6 lO-3
1
0.51
< 5 10-1
99.6
0.4
50
SO
22 lo--’ 18 IO--)
1.17 1.23 1.28
0.26 0.08
0.14
97.8 98.2
2.2 1.8
46 45 56
54 55 44
0.11
0.6
0.01
0.04
98.7
1.3
62
38
Above pH 6, the current il decreases rapidly, because the species NRCO; is not reducible in the accessible potential range. Calculation of the dehydration rate constant kdr between pH 0 and pH = pKad = 5.35 from the ratio (k + id)/& =f(/u), using the theory of Nicholson et al. for an ECE process [20] gave kz = 0.51 s-’ (cf. Fig. 8). The equilibrium constant KS can then be deduced from Koutecky’s theory from 4, = f(K$kd) [24, 251; a value of KS = 3.6 x 10e3 is found (Table 1). It can then be calculated that the percentage of free aldehyde is 0.3%, ie that of the hydrate is 99.7%. From Fig. 2, we can deduce that ika = (id/z); all the current is due to the reduction of the fraction of aldehyde at equilibrium. The percentage of aldehyde and hydrate is thus 50%, ie K[ = 1. From the relationship K;Ki = K$K;, we can deduce that the value of pKA3is 2.9 with the above values pK& = 5.35, K; = 1 and K; = 3.6 x 10A3.The value of pKL3is in reasonable accord with that found by one of us [12], ie pKi;3 = 3.36. The complete curve for the percentages of the diverse species can easily be calculated as a function of pH (Fig. 10). We have reported in Table 1 the values which we have determined and those found in the literature for the diverse constants. The values of kdi found in the literature are quite inaccurate, since the upper limit which we have found (5 x lo-’ s-l) was derived from the invariance of the product i1@ (c$ above).
100
Our values of kd2 and KS are much more precise than those of the literature, since KS need not be known to calculate kd2 (ECE process for the acid; cf. above), and is subsequently used to determine K$ uiu the product Kikd2T (CE process for the aldehyde; cf. above). By contrast, previous measurements using the current for the aldehyde between pH 0 and 4 [7, 12, 231 (Fig. 2) give the product K;kdz, so that an estimation of K; must be made; it is necessarily very inaccurate, because the equilibrium is much shifted toward the hydrate. Our measurements do not give K; with a better precision than previously, because its value is in every case obtained directly from the ratio of the limiting current of the aldehyde to the global diffusion current between pH 6 and 8 (Fig. 2).
CONCLUSION The electrochemical reduction of isonicotinic acid is more complex than thought hitherto. Reversible reduction of the acid group is followed by two successive first order reactions, viz. a rearrangement and a dehydration. It has been shown that in the case of an EC reaction with adsorption of the reactants [26], the meaning of the chemical rate constant which is measured by using electrochemical methods is not straightforward; it is not possible a priori to know if it is a surface or “volume” rate constant. For an ECCE reaction, the situation is even more complex; in particular, we cannot establish the nature of kl. However, the slope - 1 of log kd =f(Ho) below Ho = 0 indicates c1priori that the dehydration in this region is a “volume” reaction [26]; it is therefore probable that the same is true also for kd2 when Ho > 0.
REFERENCES
Fig. 10. Percentages a function of pH.
of the diverse forms of the aldehyde
as
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336
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4. H. Lund, Acta Chem. &and. 17, 972 (1963). 5. Wang Zheng-Hao and Hu Zhi-Bin, Electrochim. Acta 30, 779 (1985). 6. R. W. Green and H. Tong, J. Am. Chem. Sot. 78, 4896 (1956). 7. M. U. D. Bhatti and 0. R. Brown, J. Electroanal. Chem. 68, 85 (1976). 8. R. Rodriguez-Amaro, R. Perez, L. Camacho and J. J. Ruiz, J. Electroanal. Chem. 324, 270 (1992). 9. E. Laviron, R. Meunier-Prest and E. Mathieu, J. Electroanal. Chem. 371, 251 (1994). 10. C. H. Rochester, Acidity Functions (Edited by A. T. Blomquist), p. 17. Academic Press, New York (1970). 1I. I?. L&iron, R. Meunier-Prest, A. Vallat, L. Roullier and R. Lacasse, J. Electroanal. Chem. 341. 227 (1992). 12. E. Laviron, Buil. Sot. chim. Fr. 2325 (1961). . ’ 13. R. Meunier-Prest and E. Laviron, J. Electroanal. Chem. 328, 33 (1992). 14. E. Laviron, J. Electroanal. Chem. 109, 57 (1980); 124, 1 (1981). 15. E. Laviron, J. Electroanal. Chem. 146, 15 (1983).
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