l -ascorbic acid system

JOURNAL OF

ELSEVIER

Journal of Electroanalytical Chemistry 380 (1995) 273-277

Short communication

Formal redox potentials of the dehydro-L-ascorbic acid/L-ascorbic acid system Y a k o v I. T u r ' y a n *, R o n K o h e n Department of Pharmacy, The Hebrew University of Jerusalem, POB 12065, Jerusalem 91120, Israel Received 26 October 1993; in revised form 22 April 1994

Keywords:Formal redox potentials; Dehydro-L-ascorbic acid/e-ascorbic acid system

I. Introduction

2. Potentiometry

The dehydro-L-ascorbic acid/L-ascorbic acid ( D H A A / A A ) redox system is widely used in analytical and electroanalytical chemistry as an analyte, and is also a model for the development of the theory and methods of an analysis [1-10]. This system is of great interest in electrochemistry, biochemistry, pharmacology and other fields [11-13]. Therefore the formal redox potentials (FRPs) of the D H A A / A A system are very important. The FRPs have been investigated insufficiently, especially in relation to the intermediates of the redox process. The aim of the present work was to establish and to analyze the FRPs of the D H A A / A A system including the FRPs for the intermediates of the redox process. Literature potentiometric, polarographic and cyclic voltammetric data were used for this purpose. This is the first determination of the FRPs from polarographic and cyclic voltammetric data. It should be noted that all FRPs presented here correspond to unit concentration of the components of the redox couple and to unit activity of the hydrogen ions. The AA dissociation constants used in the calculations also correspond to the concentration of the acid/base components and to the activity of the hydrogen ions at a given ionic strength I.

For the potentiometric determination of the FRPs the electrochemical equilibria Hy.

+2e-,+2H+(+2e ,+H+;+2e -)

" H 2 A ( H A - , A 2-)

(1)

were established by the addition of a special mediator [14,15]. In Scheme (1) HzA, H A - and A 2- are the undissociated and anionic forms of AA which are in equilibrium and Hy is the hydrated bicyclic form of D H A A [12,16]: H

O

Ho \c

c\/ H-~/~ /Cl H

O H0

~ c/ ~ o c/\ HO

(2) OH

Although other products can also be formed as the result of the D H A A conversion [17], Hy predominates in the equilibrium conditions of the potentiometric investigations. As has been noted earlier [8,19], the most reliable results from the large amount of literature data on the potentiometric determination of FRPs were obtained by Ball [14]: Hy + 2e + 2H + .

• HzA

o,f EHy/H2A =

+0.390V

(3)

* Corresponding author. Present address: The National Physical Laboratory of Israel, Danciger A Bldg., Givat Ram, Jerusalem, 91904, Israel. Elsevier Science S.A. SSDI 0 0 2 2 - 0 7 2 8 ( 9 4 ) 0 3 5 2 4 - 7

at COHAA= [Hy], ([HzA]/[Hy])a2÷ = 1 and •=0.2, where c is the total concentration, [H2A ] and [Hy] are equilibrium concentrations, and all+ is the activity of the hydrogen ions. All potentials here and below are quoted versus the standard hydrogen electrode (SHE).

El. Tur'yan, R. Kohen /Journal of Electroanalytical Chemistry 380 (1995) 273-277

274

The value of E °'f H y / H 2 A [14] agrees with recent results obtained by Oganski [15]: Hy + 2 e -

.

, A2

o ,f 2- = - 0.058 V EHy/A

(4)

at CDHAA~-- [Hy], [A2 ]/[Hy]) = 1 and I = 0.2. Indeed, our calculations of FPRs for different forms of A A (H2A, H A - , A 2 - ) have shown that the FRPs from Refs. 14 and 15 are similar (Table 1). This determination has been carried out on the basis of general methodology [19] using the A A dissociation constants PKH2 A = 4.00 and PKHA = 11.27 at I = 0.2 [15]. In the next step the determination of the FRPs for the subsequent stages of electron transfer in the reversible electro-oxidation of A A are considered. For the analysis of the reversible process (1) different consecutive stages of electron transfer can be chosen. We considered the irreversible stages of the electron transfer based on the data of Deakin et al. [5] assuming that these stages are transformed into reversible stages under the influence of, for example, the mediator [14,15] or activation of the electrode surface [3,5,6]. Thus, we considered the following stages of reversible electron transfer: A 2-

On the basis of general methodology [18,19] and the o,f values of EHy/H2A [14] EAA/A~ "f [20] the apparent o,f F R P EHy/AA- = - 0 . 1 8 2 V at p H 7, CDHAA~-[Hy] and CAA/[Hy]= 1 for stage II in Scheme (5) has been determined by Deakin et al. [5]. Since CAA-= [ A ' - ] a t o,f p H 7, and the value of Eny/AA [5] is close to the true o ,f F R P EHy/A at CDHAA "~ [Hy] and ( [ A ' - ] / [ H y ] ) = 1. We o,|" repeated this calculation (Table 1) using the EHy/A2 o ,f value [15] together with the EHy/H2Avalue [14] and the A A dissociation constants at I = 0.2 (see above), unlike the thermodynamic dissociation constants used in Ref. o,f 5. The FRPs EHy/Aobtained (Table 1) differs slightly from E °Hy/AA 'f = --0.182 V [5] which is nearer the standard redox potential.

3. Polarography In polarography the equilibrium condition for the polarographic A A oxidation process with reversible stages of electron transfer on the dropping mercury electrode (DME) [16,21] given by Scheme (5) should be changed to the following Scheme proposed by Ruiz et al. [16]: A2

HA-

-e

• HA'.

-H +

-e

" A'-

' Hy

(5)

+Jl I apparent F R P

o ,f Eaa./aa = +0.30 V at p H 7,

stage I in Scheme (5) has been determined by Steenken o ,f and Neta [20]. We have recalculated E~fv/aa for the o,f true F R P EHA./rIA- at [ H A - ] / [ H A ' ] = 1 (Table 1). In this calculation the acid dissociation constant PKHA .= --0.4 [5,20] of the radical HA" has been used.

Table 1 The formal redox p o t e n t i a l s of the D H A A / A A

Hy Hy FC FC a b c d e

a a d e

' HA" .

-H +

-e-

" A'-

" FC

H20

~ Hy

-+iF

II

H2A

I

(CAA/CAA') = i (CAA ~ [ H A - I , CAX= [HA'] + [ A ' - I for

D

-e

HA-

H2A

The

+lr

II

III

It should be noted that reversible electron transfer in polarographic A_A oxidation is achieved in the absence of a mediator [16,21]. In Scheme (6) stage III (EC mechanism [16]) is completely irreversible and FC is

system ( I = 0.2)

t~ ° C

E~)'~HzA/V ( N H E )

E D° /'fH A / V ( N H E )

Eb'~A a - / V ( N H E )

E~I'fA./HA / V ( N H E )

E~)'~A- / V ( N H E )

30 25 25 23-25

+0.390 b +0.394 +0.593 +0.584

+0.272 +0.276 +0.476 +0.467

--0.062 -0.058 c +0.141 +0.132

+0.74 +0.74 +0.74 +0.74

--0.17 -0.16 +0.24 +0.22

Potentiometry. F r o m Ref. [14]. F r o m Ref. [15]. Polarography. Cyclic v o l t a m m e t r y

(6)

Y.I. Tur'yan, R. Kohen /Journal of Electroanalytical Chemistry 380 (1995) 273-277 Table 2 The formal redox potential of the F C / H A -

275

redox couple from the polarographic data ( I = 0.2, 25° C)

Buffer a

pH

(El/2) a b/V(SHE)

Reference

o,f EFC/H A-/V

A B A B

7.2 7.2 9.0 9.0

+ 0.209 + 0.203 +0.158 +0.155

[27] [28] [27] [28]

+ 0.477 +0.481 + 0.467 + 0.477 Av. +(0.476 + 0.005)

c d c ~

(SHE)

a A, H3Cit + Na2HPO4; B, HAc + H3PO 4 + N H 4 N O 3 + NaOH. b From the curve of (E1/2) a versus pH. c Average currents; the value of t 1 is not indicated, and we assumed t 1 = 3 s. a Instantaneous currents, t = 4.32 s.

the initial form of D H A A with three free carbonyl groups: H.COH ~f HCOH

/\

/

H C--C

//

(7)

%

O O Scheme (6) is described by Koutecky's equation [22,23]. Taking into account the effect of pH (KH2 A >> art+>> KHA-), Koutecky's equation for the potential at the average polarographic current ( i = I J 2 ) can be written in the following form:

RT (El/2)

a =

o ,f A- + ~ - l n EFC/H

RT all+-- F i n

0.886(kft1) °5

(8) (E1/2) a is the half-wave potential of the kinetic anodic wave of AA oxidation at the reversible electron o,f transfer, E F C / H A - is the FRP for the partial reaction FC + 2 e - + H + . " HA(9)

where

at ([HA-]/[FC])aH+= 1, t I is the drop time and kf is the rate constant of stage III (Scheme (6)). For the instantaneous polarographic current I = 11/2 at KH2A >> all+>> K H A w e changed Koutecky's equation [22,23] slightly on the basis of Refs. 24-26:

RT (El/2)

a = E Fo C,f/ H A

RT

- --In 2F

+

~-~ln

all+

1.386( kft)°"545

(10)

Table 3 The formal redox potentials of the F C / H A -

where t is the time from the beginning of mercury drop formation until the instantaneous current measurement. o,f The EFC/HA- values have been derived using Eqs. (8) and (10). In these calculations we used the (E1/2) ~ values for the average polarographic current from the work of Vavrin [27] (Table 2), t h e ( E l / z ) a values for the instantaneous polarographic current from the work of Ruiz et al. [28] (Table 2), k~= (1.35 + 0.10)× 103 s -~ (pH 7.2, 23° C) from the data of Perone and Kretlov [1] and kf = 3.4 × 102 S 1 (pH 9.0, 25° C) from the data of Jaenicke and Hoffmann [29]. o ,f The EVC/HA7 values (Table 2) were close to each other despite the different pH values and the use of different methods to determine kf and El~ 2 [1,29]. Hence, the difference in the kf values in Refs. 1 and 29 is not caused by the inaccuracy of the kf value in Ref. 29 as was suggested in Ref. 1. This difference can be explained by the influence of pH on the rate constant of the hydration and cyclization reaction of FC (Scheme (6)). This conclusion is confirmed by calculations of the kf values over a wide range of pH (pH 2-7). The value of kf changed with the change of pH. These calculations were based on Eqns. (8) and (10), o ,f the EFC/HA value (Table 2) and the (E1/2) a values from Refs. 27 and 28. The FRPs for FC and other forms of AA ( H 2 A , o ,f A 2-) (Table 1) were determined from EFC/H A similarly to the calculations of the FRPs for Hy and different forms of AA (see section 2). Stage I in Scheme (6) is the same as stage I in Scheme (5), and the value of E Ho a,f. H A - is also the same (Table 1). For stage II in Scheme (6) the value of

redox couple from the cyclic voltammetric data

t~ ° C

Electrode a

Buffer b

pH

Ep/mV (SHE)

AE/mV

u / V s- 1

Reference

E ~ / H A / V (SHE)

25 23 23

HMDE AHGCE HMDE

0.3M Ph 0.1M Ph 1.0M Ph

7.4 7.0 7.2

+ (229 _+ 5) +(244 + 5) + (269 + 5)

25 _+ 5 42 +_ 5 35 + 5

0.1 0.1 200

[31] [5] [1]

+0.471 + 0.472 +0.457 Av. + (0.467 +_ 0.006)

a Ph, phosphate buffer.

Y.I. Tur'yan, R. Kohen / Journal of Electroanalytical Chemistry 380 (1995) 273-277

276 o,f

EFC/A. (Table 1) has been calculated in the same way o,f as was used to calculate EHy/A-- (see section 2).

4. Cyclic voltammetry Scheme (6) has also been used [1] to describe the AA cyclic voltammetric oxidation under conditions of the reversible electron transfer. We have calculated E~/H a for the cyclic voltammetric oxidation of AA (Scheme (6)) using the Nicholson-Shain equation (Eqs. (10) and (83) in Ref. 30). Taking into account the effect of pH (KH2 A >> aN+>> KHA-), the Nicholson-Shain equation for the anodic process can be written in the following form:

Ep=EFc/HA-+0.782F

4Fln

~

- 4F

k u

RT + 2~lnaH+

(11)

where Ep is the anodic peak potential and v is the potential scan rate. In order to use Eq. (11) the electron transfer should be reversible (Scheme (6)). The reversibility of the electron transfer has been verified on the basis of Eqs. (36) and (37) of Ref. [30]. In the case of the anodic process, Eqs. (36) and (37) in Ref. [30] can be written in the following form:

AE

= Ep -

RT Ep/2 = 1.10~--

(12)

where E p / 2 is the anodic half-peak potential. From much literature data on the cyclic voltammetric oxidation of AA [1-7,31], we have chosen to use data from Perone and Kretlow [1] (pH 7.2, Fig. 1 of Ref. [1], Deakin et al. [5] (pH 7.0, Fig. 5 of Ref. [5]) and Karabinas et al. [31] (pH 7.4, Fig. 2 of Ref. [31]) (Table 3) because the value of kf indicated above has been determined at pH 7.2 [1]. The change of kf in the pH interval 7.0-7.4 was neglected. For these data (Table 3) the electron transfer stages in the absence of a mediator are reversible or near reversible as follows from A E values (eqn. (12) and Table 3). The reversibility is observed better with the hanging mercury drop electrode ( H M D E ) than with a glassy carbon electron activated by heating (AHGCE) (Table 3). However, the difference in AE (Eq. (12)) is o ,f not significant for the determination of EFC/H A- (Table 3). In the case of the mercury film ultramicroelectrode [4] the stages of electron transfer are probably irreversible: A E = 50-75 mV. These data [4] have not o ,f been used for the calculation of EFC/HA-. The results reported in Ref. [3] have also not been used by us because a non-standard reference electrode was employed.

o,f

As can be seen from Table 3 the EFC/H A- values obtained from cyclic voltammetric data are almost independent of the potential scan rate u over a wide range of values and are also independent of the type of electrode ( H M D E or AHGCE). This FRP (Table 3) is o,f close to the EFC/HA- derived from the polarographic data (Table 2). This confirms the reliability of the o ,f o ,f o ,f o ,f values of EFC/HA- , EFC/H~A, EFC/A2- and EFC/A. obtained (Table 1).

5. Discussion A new approach, using the polarographic and cyclic voltammetric methods together with the traditional potentiometric method, has been used for the determination of FRPs of the D H A A , / A A system. This method makes it possible to increase information about FRPs including the FRPs of the formation of the e-ascorbic acid radicals and the initial form of the dehydro-Lascorbic acid with three free carbonyl groups. The latter cannot be determined by the potentiometric method because the concentration of dehydro-Lascorbic acid is not known. In the case of the polarographic and cyclic voltammetric methods for the determination of the FRPs of D H A A / A A system it is not necessary to add a mediator to ensure reversibility of electron transfer as in the case of the potentiometric measurements. It is interesting to note that, in contrast with potentiometry, the polarographic determination of the FRPs of metal ions with variable valence [32,33] required the addition of a catalyst (a ligand). It seems that in the case of the D H A A / A A system the mediator influenced not only the rate of the electron transfer but also the rate of the chemical electrode reactions. It should be mentioned that FRPs for Hy and FC shift in the negative direction in the sequence H 2 A H A ~ A 2-. Hence the properties of A A as a reducing agent increased in the same direction [15]. It follows from Table 1 that the FRP values for FC are more positive then the corresponding FRP for Hy. This difference can be explained by the large value of the equilibrium constant K r = [Hy]/[FC]

(13)

of the FC hydration and cyclization (Scheme (6)): FC + H 2 0 .

" Hy

(14)

For example, using the FRPs for Hy and FC at the o ,f same form of AA, EHy/HA = + (0.274 ___0.002) V (Tao,f ble 1) and EFC/HA- = +(0.472 _+ 0.006) V (Tables 2 and 3) we determined the K r value: In K r =

o ,f EFC/H

o ,f

A-

--

EHy/H A-

RT/2F

(15)

Y.L Tur'yan, R. Kohen /Journal of Electroanalytical Chemistry 380 (1995) 273-277 H e n c e , K r = (6 _+ 3) × 106. T h e K r value f o u n d is very large which confirms that the form of Hy is n e e d e d as d o m i n a n t in the solution as was a s s u m e d in the p o t e n tiometric d e t e r m i n a t i o n of the F R P s . I n conclusion, we should note that Scheme (6) describes the stage of the second e l e c t r o n t r a n s f e r more exactly t h a n Scheme (5) since the initial p r o d u c t of the second e l e c t r o n t r a n s f e r is F C (not Hy). T h u s the F R P o ,f of the first EHA./H A = + ( 0 . 7 4 __+0.01) V ( T a b l e 1) a n d the F R P of the second electron transfer is, in fact, EFC/A. _ ,fo = + (0.23 _+ 0.01) V (Table 1).

6. Conclusions (1) T h e formal redox p o t e n t i a l s of different partial redox reactions involving L-ascorbic acid, its a n i o n s H A a n d A 2-, its radicals HA" a n d A ' - , the L-dehydro form with t h r e e free carbonyl groups (FC) a n d the h y d r a t e d bicyclic form of F C (Hy) have b e e n determ i n e d a n d analyzed. (2) It has b e e n shown that the second electron transfer stage in the electro-oxidation of L-ascorbic acid is described m o r e exactly by the formal redox p o t e n t i a l c o r r e s p o n d i n g to the redox couple F C / A t h a n that c o r r e s p o n d i n g to the redox couple H y / A ' - . (3) T h e e q u i l i b r i u m c o n s t a n t of the h y d r a t i o n a n d cyclization r e a c t i o n of F C ( F C + H 2 0 ~ Hy) has b e e n established o n the basis of the formal redox p o t e n t i a l s c o r r e s p o n d i n g to the redox couples H y / H A and F C / H A - . T h e p r e d o m i n a n c e of the Hy form in this e q u i l i b r i u m has b e e n shown.

Acknowledgement T h e a u t h o r s would like to express their g r a t i t u d e to Professor Z. G a l u s for his valuable advice.

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