Deactivation of manganese after its reduction at mercury electrodes

ELECTROANALYTICAL CHEMISTRY AND INTERFACIAL ELECTROCHEMISTRY Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

457

DEACTIVATI ON OF M A N G A N E S E A F T E R ITS R E D U C T I ON AT MERCURY ELECTRODES

A N N A D O W G I R D AND Z B I G N I E W G A L U S

Institute of Fundamental Problems of Chemistry, University of Warsaw, Warsaw (Poland) (Received 28th M a y 1971)

The cathodic reduction of manganese(II) on mercury electrodes is complicated since the electron transfer is followed by a chemical reaction of the first order 1'2. The oxidation of manganese from its amalgam is also complex 3. This electroreduction may be represented in a general way by the following reactions: el

Mn 2+ + 2 e ~ Mn

(1)

kf

Mn ~ Mn x (2) kb where Mn x denotes some less active form of manganese amalgam. The kinetics of reaction (2) have not as yet been investigated. In this paper we describe the results of our investigation on the kinetics and equilibrium of this reaction. EXPERIMENTAL

Reagents The solution of manganese(II) was prepared from MnC12 p.a. In all experiments 0.1 M KC1 solution prepared from p.a. salt was used as background electrolyte. All solutions were prepared with triply distilled water. Mercury was chemically purified by prolonged shaking with an acidified solution of Hg2(NO3)2, and then distilled twice under reduced pressure.

Apparatus In the chronovoltammetric and potentiometric experiments a standard hanging mercury drop electrode (HMDE) with a radius of 0.043 cm was used 4. A modification of this electrode with a wide capillary of 2.5 mm in diameter was also applied. Before the experiment, by turning the screw mercury was pushed close to the end of this wide capillary. It was found that with such an electrode the conditions of semiinfinite linear diffusion were satisfied even for an electrolysis of 10 s. The chronovoltammetric curves were recorded with a Radelkis OH-102 polarograph. Manganese amalgam was prepared by an electrolysis of 0.01 M Mn(II) for 30 s at the classical H M D E with a constant current. In order to obtain manganese amalgams of the desired concentrations, the current intensity was appropriately adjusted by changing the resistance in the circuit containing the dry battery (voltage 120 V) and J. Electroanal. Chem., 34 (1972)

458

A. DOWGIRD, Z. GALUS

a standard resistor with a Radelkis type OP-205 potentiometer connected in parallel. The potentials of these amalgams were measured with a Radelkis OP-205 potentiometer. All potentials in this work are given with respect to the saturated calomel electrode. In the chronovoltammetric experiments at scan rates up to 100 V s -1 a Hewlett-Packard voltage generator model 202 A was used. The resulting cyclic chronovoltammetric curves were observed on the screen of an "Osa" oscilloscope: Oxygen was removed from all solutions investigated by passing electrolytically generated hydrogen purified from traces of oxygen by a palladium catalyst. All measurements were carried out at 2 5 _ 0.2° C. RESULTS AND DISCUSSION

The chronovoltammetric curves recorded with mercury electrodes indicate the presence of a following first order chemical reaction, in agreement with earlier observations. The change of the peak potential with the scan rate follows the theoretical predictions for such first-order processes. Cyclic chronovoltammetric curves indicate even better the presence of the following chemical reaction. Typical curves recorded with the use of a classical electrode are given in Fig. 1. Cyclic experiments performed under constant conditions showed practically the same ratio of anodic to cathodic peak current for 2 x 10 -4 M and 2 x 10-3 M manganese(II) in 0.1 M KC1, confirming first-order kinetics for this process. At low scan rates, cyclic curves are very similar to typical curves for reversible processes without chemical complications. An increase in the scan rate leads to a rela-i b

-

-~2

-1.4

-1.6

-1.2

-1.4

-1.6 c-Iv

Fig. 1. Cyclic chronovoltammetric curves of 10-3 M manganese(II) in 0.1 M KC1 recorded at various scan rates: (a) 0.1, (b) 0.025, (c) 0.0042, (d) 0.0005 V s-1. J. Electroanal. Chem.; 34 (1972)

DEACTIVATION OF REDUCED

Mn

459

tive decrease of the anodic current. Measurements in the scan rate range 5-100 V s- 1 showed that the anodic current practically disappeared at potentials close to the reduction potential of manganese(II). The anodic current at potentials more positive than - 1.0 V was then noted. These results led to the conclusion that the chemical process is reversible, and on the basis of the ratio of anodic to cathodic peak currents and the theory developed by Nicholson and Shain 5, the parameter K/(k~+kb) ~, could be calculated, where K = kf/kb is the equilibrium constant of reaction (2). Since this theory is worked out for an electrode process occurring at a plane electrode with a semi-infinite diffusion field, basic experiments were carried out with the modification of the conventional H M D E described above. The investigations carried out with 10- 3 M MnC12 solution in 0.1 M KC1 led to an anodic to cathodic current peak ratio ia/i ~ considerably lower than 1. These ratios are given in Table 1. T ABLE I KINETIC PARAMETER OF DEACTIVATION REACTION OF MANGANESE AFTER ITS REDUCTION ON MERCURY

v/V s- '

iJik

K/(ke + kb)½

4.2 x 10 -3

0.71 0.69 0.67 0.71

2.6 2.9 3.3 2.6

83 x 10 -3

0,65 0.62 0.65

2.7 2.9 2.7

1.25 x 10-2

0.63 0,60

2.6 3.2

2.5 x 10- 2

0,56 0.56 0.59 0,59

3.0 3.0 2.4 2.4

5.0x 10 2

0,53

2.7

The anodic currents were measured according to the Nicholson and Shain procedure. The switching potential was chosen on the basis of preliminary estimates of K and K/(kf-I-kb) -~ to fulfil the conditions of the theory. A considerable help in the determination of a proper i,/ic ratio was gained from cyclic curves recorded for cadmium at the same scan rate and distance between peak and switching potentials as in the case of manganese. F r o m these experiments we could better determine the slope of the base line for anodic currents. Using the current ratios measured and the Nicholson and Shain theory we calculated the parameter K/(kf + kb) + (Table 1). The mean value of this parameter is equal to 2.8 s½, but the accuracy of this determination is not good. Considerable difficulties were met in the calculation of the rate constants kf and kb from this parameter, since the value of K was unknown. We assumed that the transformation of manganese in the amalgam is due to the formation of an intermetallic J. Electroanal. Chem., 34 (1972)

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A. DOWGIRD, Z. GALUS

compound between mercury and manganese. The formation of such compounds in the solid phase has been described in the literature 6. With this assumption we tried to estimate K assuming a dependence between the standard potentials of the couples Mn2 +/Mn-E°.2+/M, and Mn 2 +/Mn(Hg)-E°,2+/M.(Hg). This dependence may be given in the following approximate form :

EMnzO +]Mn(Hg)= E ° , 2

; /M, + (R T / n F ) In

c~at+ (R T /nF ) In K

(3)

where C~atdenotes the concentration of a saturated manganese amalgam and R, T, F and n have their usual significance. On the basis of the measurements of the equilibrium potentials of manganese amalgam of various concentrations shown in Fig. 2, we determined the solubility of manganese in mercury as 8.7 x 10- 3 M. The manganese amalgam was prepared by the reduction of Mn(II) with constant current at the classical H M D E . After the electrolysis was stopped the potential of this manganese electrode was measured with respect to SCE in a time of 12 min. At higher concentrations of amalgam the potential change was observed practically only in the first 60 s. This change occurred due to the tendency to uniform distribution of manganese atoms in the H D M E and manganese ions in the solution phase. The typical potential change of manganese amalgam with time is shown in Fig. 3. The potentials used further in calculations were obtained by extrapolating these plots to t = 0. The solubility of manganese in mercury found is considerably higher than that reported in the literature (from 1-3 x 10-3 wt. ~2,8), but the method used in this work seems to be satisfactory because in the case of copper we found that the solubility in mercury was in good agreement with the literature data 9. Using the concentration of saturated manganese amalgam and its potential, 0 +/Mn(ng)= --1.457 V. Taking the value the formal potential was determined as EM.2 EMn20 +]Mn given by Latimer 1° we could then estimate K as 9.3. With this K value the rate constants are kf = 10 s- 1 and k b = 1.1 s- 1. The values obtained are in our opinion too low, since in the cyclic chronovoltammetric experl-

Ely -1.46

-1.44

-1.42

-1.4(

-1.3~

-3'8

i

-3'4

'

-Jo

i

-2'~

i

' -z2

i

i

-la

,o~ (c/mo,,-'~

Fig. 2. Dependence of the potential of the electrode M n 2 +/Mn(Hg) on the log of manganese a m a l g a m conch. Soln.: 10 -2 M MnC12 in 0.1 M KC1.

J. Electroanal. Chem., 34 (1972)

DEACTIVATION OF REDUCED M n

461

EImV - 1454

-1452 ~1450

Q

•

--1448 --1446 0

i 2

i

i 4

i

i 6

i

i 8

i

"t/min

Fig. 3. Change of the potential of the electrode Mn 2+/Mn(Hg) in time. Solution: 10-2 M MnC1z in 0.1 M KC1. Time Of formation of manganese amalgam 30 s at a current intensity 12,39 gA, Radius of electrode, 0.043 cm. m e n t s the effect o f the s u b s e q u e n t r e a c t i o n was n o t e d even at a scan rate of 100 V s - 1. T h e e r r o r m a y be d u e to low a c c u r a c y in the e s t i m a t i o n of the e q u i l i b r i u m c o n s t a n t ( a p p r o x i m a t e value o f E°u,2+/Un) o r b y a s s u m i n g t o o simple a m o d e l o f the following c h e m i c a l reaction. A k n o w l e d g e of the kinetics a n d m e c h a n i s m of this r e a c t i o n is i m p o r t a n t in the e v a l u a t i o n of the e l e c t r o d e kinetics of the M n 2 + / M n ( H g ) system. SUMMARY T h e kinetics of the d e a c t i v a t i o n of m a n g a n e s e r e d u c e d o n m e r c u r y electrodes has been investigated using cyclic c h r o n o v o l t a m m e t r y a n d p o t e n t i o m e t r i c m e a s u r e ments. O n the basis of the t h e o r y of N i c h o l s o n a n d Shain, the p a r a m e t e r K / ( k f + kb)~ was f o u n d to be 2.8 s ~. T o e v a l u a t e the rate c o n s t a n t s kf a n d kb, the e q u i l i b r i u m cons t a n t of this reaction, K, was e s t i m a t e d . REFERENCES 1 J. BIERNATAND J. KORYTA, Collect. Czech. Chem. Commun., 25 (1960) 38. 2 W. KEMULAAND Z. GAEUS, Roez. Chem., 36 (1962) 1223. 3 R . BIEBER AND G . TRi~IMPLER, Helv. Chim. Acta, 30 (1947) 971. 4 W. KEMUEAAND Z. KUBLIK,Anal. Chim. Acta, 18 (1958) 104. 5 R. S. NICHOLSONAND I. SHAIN,Anal. Chem., 36 (1964) 706. 6 F. LIHL AND H. KIRNBAUER,Z. Metallk., 48 (1957) 9, 61. 7 J. J. LINGANE, J. Amer. Chem. Soc., 61 (1939) 2099; M. VONSTACKELBERG,Z. Elektrochem., 45 (1939) 446. 8 J. F . DE WET AND R . A . W . HAUL, Z. Anorg. AliA. Chem., 277 (1954) 96; G. JANGG AND H . KIRCHMAYER, Z. Chem., 3 (1963) 47. 9 N. H. Ii~vlrqAND A. S. RUSSELL, J. Chem. Soc. (London), (1932) 891. I0 W. M. LATIMER,The Oxidation States of the Elements and their Potentials in Aqueous Solutions, PrenticeHall, New York, 1952. J. Electroanal. Chem., 34 (1972)