Kinetics of the electrode processes between copper amalgam and copper(II) glyoxylate and pyruvate solutions

Electroanalytical Chemistry and Interjacial Electrochemistry, 60 (1975) 29 39

29

(~2)Elsevier Sequoia S.A., Lausanne Printedin The,Netherlands

KINETICS OF THE ELECTRODE PROCESSES BETWEEN C O P P E R A M A L G A M AND COPPER(II) GLYOXYLATE AND PYRUVATE SOLUTIONS

S. FRON~EUS and C. L. JOHANSSON Inorganic Chemistry 1, Chemical Center, Unit'ersity of Lund, P.O.B. 740, S-220 07 Lund 7 (Sweden)

(Received 14th November 1974)

INTRODUCTION In two previous studies t ,2 of electrode processes at liquid copper amalgam it has been found that acetate and glycollate ions have a considerable influence on the exchange current density, io, which cannot be explained by a Frumkin correction. It was possible to prove that the overall charge transfer at the electrode proceeds step-wise and that the electron transfer step Cu(II)/Cu(I), which in non-complex solutions is rate-limiting, is speeded up very much by the coordination of an acetate or glycollate ligand to the Cu(II) and Cu(I) ions. Since the acetate ion gives an effect similar to that of the chelating 3 glycollate ion, it was concluded that the coordination of the carboxylate group is of decisive importance. Thus, it was suggested that there is a weak but fast specific adsorption of copper complexes on the amalgam with the carboxylate group functioning as a bridge between the metal ion and the electrode, one oxygen atom being bonded to the amalgam surface and the other coordinated to the metal ion. In addition, it was assumed that the carboxylate bridge is electron conducting. To test the theory further, and especially the latter assumption, we have now extended the investigations with a study of the effect of glyoxylate and pyruvate ions on the same electrode reaction, as Taube and Gould a found that in homogeneous redox reactions between Co(III) and Cr(II) carboxylate ligands with a conjugated bond system pendant to the carboxylate group can be still better bridges for the electron transfer. For pyruvate it is known 5 that Cu(II) promotes dimerization, but at the pH used (2.(~2.5) this process is slow, and with freshly prepared solutions electrode impedance and e.m.f, measurements could be carried out without the influence of side reactions. With glyoxylate the complex solutions show no perceptible change with time. THEORETICAL List o f symbols

io

i0~1; io,2

electrochemically determined exchange current density of the overall reaction Cu(am)~-~Cu(II) + 2e exchange current densities of the steps Cu(I)/Cu(am) and Cu(II)/Cu(I)

30 !

S. F R O N A E U S , °

ki,~ ,

!

kj,2

O~j,l ; Cgj,2

flj; 7i K

C. L. JOHANSSON

rate constants of the reactions CuL) ffGu(am) and C u L ~ , f f C u L ) - i anodic transfer coefficients of the same reactions stability constants of CuL 2 i and CuL 1-i equilibrium constant of Cu 2+ + C u ( a m ) ~ 2 Cu ÷

ki,~ = Fk~,a 7~j'l K ~j''/2 [Cu(am)](1 - ~j,, )/2 ky , 2 = Fb' 2R~j,2 ~"~(1 ~j,2)/2 [ C u ( a m ) ] ( 1 - ~ , , ~ ) / 2 - - ' ~ j , 2 ~,I--~J, ,eI PJ *,.

Equations for the exchange current densities According to measurements below on the complex equilibria, the two carboxylate ions under investigation give mononuclear complexes with Cu 2 + at the low total concentrations of Cu(II) used. Furthermore, from the results of the preceding investigation 2 we conclude that also with the present carboxylate ligands the overall charge transfer most probably proceeds step-wise with participation of the different complexes in the solution according to the following reaction scheme ( L - = glyoxylate or pyruvate ion): Cu(am) + j L - ~ C u L 1 -j + e 1-j~

CuL)

(1)

2

--CuLj -j+e-

(2)

Then, for the exchange current densities i0,1 and io,2 we get as final expressions 1

io,l=~.kj,l[Cu2+]~'l/2[L-] j

(0 < c~j,1/2 < 0.5 )

(3)

(0.5 <(1 +%,z)/2 < 1)

(4)

y

io, 2 = ~. kj,2[Cu2+](1 +~J'2)/2[L-] J J

Copper amalgam is a two-phase amalgam even at very low concentrations 6, and for t h a t reason [Cu(am)] was kept constant throughout the measurements. Then the composite coefficients kj,1 and kj,2 should be constants. The validity of eqns. (3) and (4) presupposes that a supporting electrolyte in high and constant concentration is used, so that the qSz-potential of the outer Helmholtz plane and the activity coefficients are approximately constant. A comparison between the different carboxylate systems is then feasible without applying Frumkin corrections to the rate constants. In the preceding paper 2 it was shown how the electrochemically determined io is composed of io, 1 and io, 2. In the derivation two different possibilities had to be taken into account. If the overall disproportionation reaction Cu(II) + C u ( a m ) ~ 2 Cu(I)

(5)

at the amalgam surface is slow compared with the charge transfer steps, then at a.c. polarization using not too high frequencies, and at the low equilibrium concentrations of Cu(I) in carboxylate solutions, the following relationship should hold 2,7.

1/i o = 1/2io, ~+ 1/2io, 2

(6)

In this case the value of i o obtained electrochemically is equal to the true exchange current density of the isotopic exchange reaction between Cu(II) and amalgam (cf. ref. 8).

ELECTRODE PROCESSES OF COPPER CARBOXYLATES

31

If the disproportionation reaction is much faster than the charge transfer steps, then at small polarizations the Cu(I) concentration does not change. This means that we have two parallel single-electron steps leading to the relationship 2 i o = i0.1/2+ io.2/2

(7)

Eqns. (6) and (7) presuppose that i 0 in both cases is calculated from the transfer resistance, Rt, according to the expression R t = R T / 2 F A i o. Here A is the area of the polarizable electrode. For the elucidation of the mechanism of the charge transfer from /o-data it is important to study the dependence of io on [Cu 2+] within different ligand concentration ranges. For this purpose we use the quantity 1 ~. = (c? log io/O log [CuZ+])tL ~

(8)

It is seen that ~ is an average of the exponents for [Cu 2+] appearing in the complete expression for io. From eqns. (3) and (4) it is evident that quite different values of ~ should be expected for i o ~ io, 1 and for i o ~ i o , 2. S-values of about 0.25 and 0.75, respectively, are plausible in these two cases. EXPERIMENTAL

Chemicals

All chemicals used were of analytical grade, when available, and all solutions were made from doubly-distilled water. Copper(II) perchlorate, copper amalgam (0.1~o by weight-), and sodium perchlorate were prepared as described previously 1'2. Buffer solutions were prepared from sodium glyoxylate or pyruvate and perchloric acid. Both the salts had correct molecular weights, determined by cation exchange and alkalimetric titration. The ionic strength of the solutions was 1 M with sodium perchlorate as a supporting electrolyte. E x p e r i m e n t a l details

The cell in the impedance measurements was of the same kind as before 2. The surface area A of the amalgam drop used as the polarizable electrode was 0.040 cm 2. The a.c. bridge and the procedure for obtaining the faradaic impedance from the total cell impedance were also the same as in the previous investigations. Values of-the faradaic impedance were determined at frequencies v between 20 and 2000 Hz. Fairly stable and reproducible values of the total electrode impedance were obtained within a couple of minutes after the amalgam drop had been renewed. For the determination of the stability constants of the Cu(II) complexes, separate e.m.f, measurements were performed using an amalgam electrode. These measurements yielded values of [Cu 2 +]. From other e.m.f, measurements, with the Cu(am) electrode exchanged for a glass electrode, values of [H +] in all the complex solutions were obtained. All ?neasurements were carried out at 25.0_+ 0.2°C. The complex solutions were always de-aerated by a stream of oxygen-free nitrogen. MEASUREMENTS AND CALCULATIONS T h e 91yoxylate s y s t e m

The exchange current density, i0, increases rapidly on addition of glyoxylate

S. FRONA~US, C. L. JOHANSSON

32

ions, and measurable Rt-values were obtained only at glyoxylate concentrations % < 20 m M and for total copper(II) concentrations in the range 0.25 m M < Cc~ < 2.5 mM. As was also the case with the previously investigated carboxylate systems 1'2, it was found that when the faradaic impedance was represented as a resistance, R~, and a capacitance, C~, in series, R~ and (coC~)-1 plotted versus v -~ gave two fairly parallel straight lines (Fig. 1). This "Randles behaviour" proves that the impedance is composed solely of a charge transfer resistance and a diffusion impedance. A possible specific adsorption of the charge transferring copper species on the amalgam surface must be weak, and furthermore be a very rapid and reversible step (cf. ref. 9). Thus, the intercept of the R~-line on the ordinate axis should give the true value of e t.

30

J

2C /" 1C

~/~///~// I1[/ f/o

//

///

/ / /0

//O///°/~"

.

// I

0.05

[

0.10

1 "I)-~/S

L 0.15 t g

Fig. 1. R~ ( ) and (o)C~)-1 ( - - ) as functions of v-{ at differentvalues of Ccu and the glyoxylateion concentration, [L ]. Cc,, [L ]: (V7 II), 1.00 raM, 2.90 mM; ((3 O), 2.3Oram, 1.09 mM. The i0-values found at a constant copper(II) concentration (Ccu= 0.66 mM) and at varying glyoxylate concentrations are given in Table 1. F r o m the acid constant K a = 1.12 x 10 .3 M of glyoxylic acid and the values of [H+], the free ligand concentration, [L-], was calculated. The stability constants of the Cu(II) complexes were determined from the separate e.m.f, measurements with the amalgam electrode (cf. ref. 10). The following values were found: /31=60+3 M 1, /32= (1.2_+0.1) x 10 a M -2, and/? 3 =(4.6_+0.5) x 103 M -3. The errors given are maximum random errors. With these constants, [Cu 2 +] in the solutions used in the impedance measurements was calculated from the relation Ccu = [Cu 2 +] j2/~[L ]i, where P0 = 1. For the determination of the quantity&, impedance measurements were also carried out at varying values of Ccu and with [L ] as a parameter (Table 2). Values of & at Ccu=0.66 m M were then obtained graphically according to eqn. (8). They are collected in Table 1. At the lowest ligand concentrations ( [ L - ] < 4 mM) used in these latter

33

E L E C T R O D E PROCESSES O F C O P P E R C A R B O X Y L A T E S TABLE 1

T H E E X C H A N G E C U R R E N T D E N S I T Y io A N D T H E A V E R A G E T R A N S F E R C O E F F I C I E N T c~ AT Ccu=0.66 m M A N D D I F F E R E N T G L Y O X Y L A T E I O N C O N C E N T R A T I O N S , [L ] [L ]/ mM

[Cu2+]/ mM

io/ mA cDI 2

~:

i o calc/mA cm 2 Eqn. ( 6 )

0 0.23 0.48 1.09 1.90 2.40 2.90 4.10 4.80 5.50 6.30 9.40 17.5

0.66 0.66 0.65 0.62 0.60 0.57 0.56 0.53 0.50 0.49 0.47 0.40 0.275

1.25 4.0 5.6 11.2 18.0 25 29 40 50 55 60 90 150

0.65 0.66 0.7 l 0.73 0.73 0.70 0.50 0.40 0.30

1.10 3.6 5.9 11.3 18.0 21 25 33 35 40 45

Eqn. ( 7 )

65 75 80 85 100 160

TABLE 2 T H E E X C H A N G E C U R R E N T D E N S I T Y i o AT D I F F E R E N T Ccu A N D G L Y O X Y L A T E I O N CONCENTRATIONS, [L-]

[ L- ] / mM

Cc,,/[ Cu 2+ ]

io/mA cmCcu/mM

0.48 ( V ) 1.09 ( O ) 1.90 ([~) 2.90 ( V ) 4.10 ( O ) 4.80 5.50 6.30

1.03 1.07 1.12 1.19 1.27 1.32 1.37 1.43

0.33

0.66

1.00

1.65

2.30

3.75 7.4 10.5 18.0 25 36 40 45

5.7 11.5 17.5 29 41 50 55 60

7.3 15.0 23 40 55 60 70 70

10.0 23 35 60 80 80

13.0 30 50 70

measurements, the exchange current density can be represented fairly well by the expression: io/mAcm- 2 = 1.3 x 102[Cu2+] °65 + 1.8 x 106[Cu2+]°7°[L -] (9) ([Cu 2+] and [L-] in M) The validity of eqn. (9) is proved by the linear relationship obtained in Fig. 2. On the other hand, the very pronounced decrease in ~ for [L-] > 5 m M indicates that eqn. (9) is not applicable at the higher ligand concentrations. The high values of~ obtained at the lowest ligand concentrations make it very probable that i o ~ io,2 within this range of [L ]. Furthermore, on the basis of earlier results obtained for the glycollate system 2 at low ligand concentrations, we presume

34

S. FRONA~US, C. L. JOHANSSON I

I

I -3.5

I -3.0

I

"7

½ 4.25

< E

%

4.00


%) x OO

~z 3.75 I2 O~ O

I -2.5

log ([Cu2+]/m) Fig. 2. log (i 0-1.3 xl0=[Cu2+]°'6S)[L ] 1 as a function of log [Cu 2+] at different Cc. and low glyoxylate ion concentrations. The symbols refer to the [L ]-values given in Table 2.

that eqn. (6) is valid and that io,1 >~ i0,2- Thus, we have io ~ 2 io,2 at [L-] __<4 raM. Values of the composite coefficients ko, 2 and ka,z and of the transfer coefficients %,2 and ~1,2 are subsequently obtained directly from eqn. (9) (Table 3). The errors given are maximum random errors. At higher ligand concentrations ([L-] > 6 raM), on the other hand, &becomes so low that we must conclude that io ~ io,a for these complex solutions. Here we have two possibilities to take into consideration. (I) Eqn. (6) could still be valid, but now with i0,1 ~ io,z, so that io~2io,~. TABLE3 VALUES OF THE PARAMETERS IN EQNS. (3) AND (4) W H E N io. 1 AND io,2 ARE EXPRESSED IN m a cm 2, AND [Cu 2+] AND [L-] IN M

Ligand

J

kj,1

¢~j,l

kj,2

~j,2

HCOCOO-

0 1 0 1

(5_+2)x 102 (1.5_+0.3) x l0 s

0.60 0.60

65_+5 (9_+1) × 105 65 _+5 (1.1 +0.1) x 105

0.30--0.04 0.40_+0,04 0.30 + 0.04 0.40_+0.04

CH3COCOO

Since i o still increases rapidly with [L-i, this explanation implies that io.2 should increase much more rapidly and give a much higher value of kz, z than that found with the glycollate system. However, this explanation does not seem very plausible, as the k 1,z-values of the two systems are of the same order of magnitude. (II) At the higher glyoxylate concentrations, eqn. (7) could be valid instead of eqn. (6). Then, under the assumption that the relatign io, 1 ~> i0,2 still holds, we should have io~io,1/2. The application of eqn. (7) implies that the rate of the disproportionation reaction (5) is very much enhanced by the formation of glyoxylate complexes, so that the concentration of the intermediate Cu(I) does not

ELECTRODE PROCESSES OF COPPER CARBOXYLATES

35

change appreciably at the a.c. polarization used. The latter explanation of the pronounced decrease in ~ seems to be the most probable one, and has been the basis for the further calculations. From the lowest S-values we have concluded that e0.1 = ctl. 1 ~ 0.60, and the value of ko. 1 has been taken from the earlier paper on the glycollate system. Then, with the highest values of io(= io.1/2) in Table 1 an approximate calculation of k1.1 was carried out by means of eqn. (3). Using the parameters in Table 3, the functions io. 1 and i0.2 w e r e computed (Fig. 3), and then i o (Table 1) from eqn. (6) or eqn. (7). On the whole there is good agreement between measured and calculated values of i o, but in the [L-]range where ~ changes rapidly the measured i o is about half-way between the values of i o calculated from eqn. (6) and eqn. (7). I

./

I

/ ./' /

.i

/

2.0

•

./,/

i.

'E LJ ,4 E 1,2 ._o & .J

///

o° ° °

//

8

o~ 1,o _o

1111 / //

0.5

/

/

/

/

/

/

/

/

/

i -3.0

I -2.0

I ,

-1.o

log ( [ L - I / M )

Fig. 3. log '/o,1 ( - ' - ) , log io. z (--) and log i o ( - - eqn. (6)), ( . . . . eqn. (7)) as functions of log[L ] for the glyoxylate system, ccu=0.66 raM. The curves have been calculated using the parameters in Table 3. • refers to the measured i o.

The pyruvate system On addition ofpyruvate ions to non-complex copper(II) solutions, i o increases but not so rapidly as with glyoxylate ions, and consequently measurements could be made at pyruvate ion concentrations, [L-], up to 160 raM. With freshly prepared solutions reproducible values were obtained in the impedance measurements, and also for this carboxylate system the faradaic impedance displayed "Randles behaviour". The/o-values at a constant Ccu and varying pyruvate ion concentrations are collected in Table 4. The acid constant K~= 6.3 x 10 3 M, obtained in the separate e.m.f., measurements, was used for the calculation of [L-]. The [Cu 2 +]-values were computed by means of the potentiometrically determined stability constants of the

36

S. FRON/EUS, C. L. JOHANSSON

TABLE 4 VALUES OF io AND ~ AT Ccu=0.66 mM AND DIFFERENT PYRUVATE ION CONCENTRATIONS, [L ] [L-f~ mM

[Cu2+]/ mM

io/ mA cm 2

o:

io(calc)/ macro-2

0 0.83 1.45 3.05 5.50 7.00 10.8 15.0 21.0 39.5 56.5 72.5 101 158

0.66 0,65 0.65 0.63 0.60 0.58 0.54 0.50 0.45 0.34 0,265 0.215 0.160 0.095

1.25 1.55 2.15 3.7 7.1 8.0 12.5 17.5 24 33 38 45 50 55

0.65

1.10 2.10 2.9 4.7 7.4 8.9 12.5 16.5 21 32 38 40 50 55

0.63 0.63 0.64 0.68 0.70 0.70 0.70 0.65 0.65

copper(II) complexes. The constants obtained, fil =20-+ 1 M -1 and ~ 2 = 115_ 10 M -2, are of the same order of magnitude as those given by Tallman and Leussing 11 for the zinc(II) system. The results from the impedance measurements at different constant [L-] and varying Ccu are reported in Table 5. The quantity ~, determined for Cc, = 0.66 mM from these measurements, is given in Table 4. At moderate values of [L ] the expression i o / m A c m - 2 = 1.3 x 1 0 2 [ C u 2 + ] ° 6 5 + 2 . 1 x 105 [ C u 2 + ] ° ' V ° [ L - ]

(10)

fits the/o-data fairly well. This is evident from Fig. 4. The values ~ = 0.654).70 for TABLE 5 VALUES OF io AT DIFFERENT Ccu AND PYRUVATE ION CONCENTRATIONS, [L-] [U]/ mM

Cc ,//Ft c u 2 + j•

io/mA am - z Ccu/mM

1.45 5.50 ( 0 ) 7.00 ( I ) 10.8 (V) 21.0 (V) 31.0 (O) 39.5 ([B) 51.0 88.0 150

1.03 1.11 1.15 1.23 1.47 1.74 1.98 2.32 3.65 6.59

0.33

0.66

1.00

1.35

2.00

1.45 4.5 5.5 8.6 15.5 17 20 22 30 40

2.10 7.0 8.0 13.0 24 28 33 37 45 55

2.70 8.8 10.0 16.0 34 38 45 50 60 70

3.4 10.5 13.0 21 40 45 50 60 75 90

17.0 29 50 60 70 80

ELECTRODE PROCESSES,OF COPPER CARBOXYLATES

37

,y

I

I

½ o~

'E u 3.25 g-

E d + U

3.00

% x

-~ 2.75 c~

o

# /q ]

-3.5

L

- 3.0

I

- 2.5

[og([Cu2÷]/M)

Fig. 4. log (i o - 1.3 x 102 [Cu 2 +]o.65)[ L -] - t as a function of log [Cu 2 +] at different values of Ccu and the pyruvate ion concentration. The symbols refer to the [L ]-values given in Table 5.

this system make it very plausible that i 0 ~ io.2. Then we presume that eqn. (6) is applicable with io,1 >>io,2 within the whole range of [L-] investigated. Thus, no parameters in the expression for io.~ can be determined. On the other hand, from eqn. (10) and the relationship i 0 ~ 2 io,2, values of kl.2 and el,2 are obtained immediately (Table 3). Using these parameters, io was calculated (Table 4), and it is found that in general the differences between measured and calculated values are within the limits of experimental random errors. At the lowest values of [L-] the calculated i0-values are somewhat higher than the measured ones. This can be explained by the higher acidity used in these solutions in order to decrease [L-]. DISCUSSION

The results obtained in the present investigation prove that glyoxylate and pyruvate ions have a pronounced rate-enhancing effect on the charge transfer step (2). Such a ligand influence was also found in the previously investigated carboxylate systems 1,z. Furthermore, step (2) is the slowest one of the overall charge transfer at all ligand concentrations used. In this respect, the two present systems resemble the acetate system but are different from the glycollate system, where step (1) becomes rate-limiting at higher ligand concentrations 2. This difference can to some extent be explained by the weaker complex formation in the present copper(II) systems, which makes the ratio [Cu(I)]/[Cu(II)] decrease more slowly than in the glycollate system. According to our interpretation of the strong decrease in ~ already at fairly low glyoxylate ion concentrations, the attainment of the disproportionation equilibrium (5) is accelerated very much in the presence of this ligand. In view of the remarkably strong rate-enhancing effect displayed by the glyoxylate ion as a bridging ligand in certain homogeneous redox reactions 4' ~2, it is very plausible that this ion has

38

S. FRONA~US, C. L. JOHANSSON

a similar catalyzing effect on the homogeneous reverse reaction of (5). Thus, since the disproportionation equilibrium is not displaced very much by the complex formation, the forward reaction of (5) should also be catalyzed. It may be pointed out that the strong decrease in ~ sets in at [ L - ] ~ 5 mM, where the fraction [CuL+]/Ccu has attained about half of its maximum value. As to the influence of the Frumkin effect on the rate of the electrode reaction, it can be presupposed that the specific adsorption of glyoxylate or pyruvate ions on the amalgam surface is as weak as that of acetate ions 13. Then, owing to the specific adsorption of perchlorate ions 1~ from the supporting electrolyte, the q~2-potential of the outer Helmholtz plane at the equilibrium electrode potential should be negative and have approximately the same value for all the copper carboxylate solutions investigated. Thus, the rate enhancement of the charge transfer step Cu(II)/Cu(I) cannot be explained for any of these complex systems on the basis of the Frumkin correction, since the average positive charge of the copper species decreases by the complex formation. For a quantitative comparison between the effects of the different carboxylate ligands, the rate constants k;, 2 and k],2, relating to the charge transfer step (2), have been calculated (Table 6). In these calculations the unknown stability constants ~)1 of the Cu(I) glyoxylate and pyruvate systems had to be replaced by the corresponding stability constants of the Ag(I) systems. The following values were obtained from e.m.f, measurements: 71 = 3.0_+ 0.2 M - 1 for the glyoxylate system a n d ~1 = 2.2+ 0.2 M-1 for the pyruvate system. For K [Cu(am)] the same value was used as in previous work 2. Owing to the approximation with respect to 71, no error limits can be given for k],2, but the ratios of k],2-values should be fairly uninfluenced by the approximation. TABLE 6 RATE CONSTANTS RELATING TO THE CHARGE TRANSFER STEP Cu(II)/Cu(l) WITH WATER AND DIFFERENT CARBOXYLATE IONS AS LIGANDS Ligand

k'o,2/cm s 1

H20 CH3COO HOCH2COOHCOCOO CH3COCOO -

0.10+0.01

k'l.2/cra s 1

ill~M-1

11 25 60 13

47+ 1 220± 10 60- 3 20_+ 1

From Table 6 it is seen that all the carboxylate ions used have a striking rate-enhancing effect on the charge transfer step Cu(II)/Cu(I). Furthermore, the variation in k],2 with the different ligands cannot be correlated to the strength of the Cu(lI) complexes formed or to the tendency of the ligands towards chelation. Since the variation is not very great, it could possibly be explained by differences in the gb2-potential. However, it should be emphasized that the sequence of the carboxylate ions according to their rate-enhancing effect, acetate < glycollate < glyoxylate, is the same as that found if these ligands are arranged according to their accelerating effect on

ELECTRODE PROCESSES O F C O P P E R CARBOXYLATES

39

the rate of certain homogeneous redox reactions 4. In the latter case, it has been proved that the carboxylate group functions as an electron conducting bridge. Although the differences in rate-enhancing effect between the different carboxylate ions are much smaller in the electrode reactions than in the homogeneous reactions, we think that the sequence found provides support for our suggestion that the carboxylate group has a similar bridging function in the two types of reactions. ACKNOWLEDGEMENTS

The present investigation is part of a program supported by a grant from the Swedish Natural Science Research Council. The authors wish to thank Miss K. Anderson for her assistance in the measurements and Dr. Robert Carter for his linguistic revision. SUMMARY

The exchange current density at liquid copper amalgam in complex copper glyoxylate and pyruvate solutions has been determined by electrode impedance measurements at the equilibrium potential. To keep the activity coefficients and the ~bz-potential approximately constant, 1 M sodium perchlorate was used as ionic medium. The results indicate that the charge transfer between Cu(I) and Cu(II) via the amalgam, which is the slowest step in the overall charge transfer, is accelerated very much at increasing ligand concentrations. When different carboxylate ligands are arranged according to their accelerating effect, the following sequence is obtained: acetate < glycollate < glyoxylate. From this result it is concluded that the carboxylate group forms an electron conducting bridge between the central ion and the amalgam. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12 13 14

S. Fronaeus, R. Johansson and C. O. ~)stman, Chem. Scr., 1 (1971) 52. S. Fronaeus and C. L. Johansson, J. Electroanal. Chem., 48 (1973) 195. R. Larsson, Acta Chem. Scand., 19 (1965) 783. H. Taube and S. Gould, Accounts Chem. Res., 2 (1969) 321. K. N. Leong and D. L. Leussing in Proc. of the X I V I.C.C.C., Toronto, 1972, p. 608. G. Tammann and T. Stassfurth, Z. Anorg. Chem., 160 (1927) 246. V. V. Losev in B. E. Conway and J. O'M. Bockris (Eds.), Modern Aspects of Electrochemistry, No. 7, Butterworths, London, 1972, p. 314. S. Fronaeus and C. O. Ostman, Acta Chem. Scand., 10 (1956) 769. A. A. Pilla, J. Electrochem. Soc., 117 (1970) 467. S. Fronaeus in H. B. Jonassen and A. Weissberger (Eds.), Technique of Inorganic Chemistry, Volume I, Interscience, New York, 1963, p. I. D. E. Tallman and D. L. Leussing, J. Amer. Chem. Soc., 91 (1969) 6253. H. J. Price and H. Taube, Inorg. Chem., 7 (1968) 1. D. C. Grahame and B. A. Soderberg, J. Chem. Phys., 22 (1954) 449. R. Payne, J. Phys. Chem., 70 (1966) 204.