The electrode kinetics of nickel in thiocyanate solutions at mercury electrodes

J. Electroanal. Chem., 100 (1979) 791--800

791

© Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands

THE ELECTRODE KINETICS OF NICKEL IN T H I O C Y A N A T E SOLUTIONS AT M E R C U R Y E L E C T R O D E S

TADEUS Z KROGULEC *, A N D R Z E J BAR,~NSKI and ZBIGNIEW GALUS

Laboratory of Electroanalytical Chemistry, Institute of Fundamental Problems of Chemistry, University of Warsaw, Pasteura 1, 02093 Warsaw (Poland) (Received 8th January 1979)

ABSTRACT Using pulse chronocoulometric polarography the kinetic parameters of the Ni(II)/Ni(Hg) system in thiocyanate solutions have be determined. Under the conditions of the experiments the influence of electroreduction of SCN- ions resulting in formation of NiS adsorbed on the electrode was minimized. The standard rate constants obtained are of the order of 10-a cm s -1 , several times larger than earlier reported in literature. The Tafel slopes indicate a two-step process with slow transfer of the second electron and NiSCN ÷ directly participating in the electrode reaction.

INTRODUCTION

Though the electrode reactions of nickel in thiocyanate containing solutions were studied b y a number of workers [ 1--4], the kinetics and mechanism of these reactions was n o t explained sufficiently as yet. A large part of published papers dealt with strictly analytical problems with determination of the stability constants of the nickel(II) thiocyanate complexes or with the kinetics of the chemical reactions preceding the charge transfer proper. J Turjan and Serova [3] proposed the discharge of the neutral Ni(SCN)2 complex in the electrode reactions. Also the discharge of Ni(SCN)~ was postulated [4], and standard rate constants were given o f the order 2 to 3 × 10 -4 cm s-1. However in earlier studies of the kinetics of this system the side reactions of the parallel reduction [5] of SCN- ions leading to the formation of NiS adsorbed on the electrode surface were n o t taken into account. Recent studies of these side reactions revealed that the adsorbed NiS m a y influence to a large extent b o t h the kinetics and the mechanism of this reaction [6,7]. This work was undertaken in order to study the mechanism and kinetics of nickel(II) electroreduction under the conditions where the influence of NiS adsorption is negligible. This should be possible when (a) the nickel(II) concen* Bialystok Division of Warsaw University.

792 tration is n o t very high and (b) an electroanalytical m e t h o d with short generation o f product such as for instance pulse polarography is applied. Applying pulse polarography one could study also the oxidation o f the homogeneous nickel amalgam formed in the cathodic pulse. EXPERIMENTAL

Reagents Reagents used were pro analysis grade and were additionally purified b y crystallization. Nickel(II) perchlorate was purified by electrolysis at a mercury pool cathode with imposed potential a bit more positive than the potential sufficient to initiate the electroreduction of nickel(II). All solutions were prepared from thrice distilled water. The second distillation was carried out from an alkaline solution of permanganate and the third one from an all-quartz still. In measur~ements the ionic strength was constant, the following equality CNaSCN + CNaClO4= 4.8 M was always satisfied. Before measurements all solutions were deoxygenated b y passing through them a stream of electrolytically generated hydrogen purified catalytically from traces of oxygen. Mercury was twice distilled under vacuum and then was chemically purified b y prolonged shaking with a solution of Hg2(NO3)2 acidified with HNO3.

Apparatus All voltammetric experiments were carried o u t in a three-electrode system. The dropping mercury electrode (DME) with the drop time mechanically regulated and the hanging mercury drop electrode (HMDE) were used as indicator electrodes. The counter electrode constituted a platinum foil separated from the investigated solution b y a sintered glass plug. All potentials were measured versus the saturated calomel electrode placed directly in the solutions studied. D.c. polarographic experiments were carried out with the use o f a Radelkis polarograph t y p e OH-102 with a check o f the potential at the polarograph o u t p u t b y a digital voltmeter t y p e V-530. Pulse polarographic curves were recorded with the use o f a home-made instrument and an XY Yokogawa recorder t y p e 3077. The temperature of the solution in the electrolytic cell was kept constant at 25 + 0.1°C.

The procedure of experiments The potential o f the initial polarization used for producing the homogeneous nickel amalgam was equal to --0.80 V (first set of experiments} and --1.50 V (second set). The application o f so negative potentials in the second set was sufficient to reduce NiS which would form on the electrode surface. One should

793

add that the nickel(II) concentration was kept at low value to inhibit (in the cathodic step when the amalgam was produced) the crystallization of the homogeneous nickel amalgam into intermetallic c o m p o u n d with mercury. When starting from above mentioned potentials the DME was pulsed in the anodic direction, at sufficiently positive potentials ( - 0 . 6 to --0.7 V) the oxidation of the amalgam was observed. In the study of the cathodic reduction of nickel(II) the electrode was pulsed to negative potentials starting from - 0 . 5 0 V, at which no reduction of nickel(II) was observed. Such a procedure makes possible the determination of the kinetic parameters of the studied reactions b y applying the m e t h o d of Kimmerle and Chevalet [8]. The studied reactions of oxidation and reduction m a y be described b y the model D-R presented b y the authors mentioned. The analysis of ratios t

Q'(D-R)T Q'(D-D)J~ t gives parameter × which is connected with the rate constant o f electrode reactions X = kbh(t'/DRed) I/2

(I)

X = k ~ h ( t ' / D o x ) '/2

(2)

for anodic and cathodic processes respectively. RESULTS

Preliminary experiments The polarographic and chronovoltammetric curves of solutions containing nickel(II) in concentrations changing from 10 -s M to 2 X 10 -2 M and constant thiocyanate concentration (4.8 M) were recorded. In case of chronovoltammetric curves for nickel(II) concentrations exceeding 10 -3 M one observes a decrease, of reduction current due to blocking of the electrode b y NiS produced in the electrode reaction. These curves are similar to those described in our earlier paper [6]. Similarly on the polarographic curves for this concentration range one observes a minimum n o t recorded at low concentrations of nickel(II) (Fig. 1).

Determination of diffusion coefficients In order to determine diffusion coefficients of nickel(II) ions, i--t curves were recorded with the used of the HMDE. Sufficiently negative potentials were applied to this electrode in order to have currents limited only b y the diffusion rate of nickel(II). The composition o f solutions was similar to that used in kinetic experiments, only to minimize a possible influence of NiS the concentration of nickel(II) was decreased to 2 X 10 -4 M. In analysis of the curves obtained the Cotrell equation for a spherical diffu-

794

-o5

El V vs. SCE

Fig. 1. P o l a r o g r a p h i c c a t h o d i c curves r e c o r d e d for s o l u t i o n s o f Ni(II) w i t h c o n c e n t r a t i o n s 10 -4 (a), 2 X 10 -4 (b), 5 × 10 -4 (c), 10 -3 (d) a n d 5 × 10 -3 M (e) in 1.0 M N a S C N w i t h 0 . 0 0 2 5 % gelatine. D r o p t i m e 3 s.

sion was used

ilt I n = n F A D o x c ° ~ [ 1/(~rDox)'n + tl/2/ro]

(3)

where i~ is the limiting current dependent on time of electrolysis t, Dox and C~x denote diffusion coefficient and bulk concentration of nickel(II), respectively, r0 and A are the radius and the surface area of HMDE, respectively. In agreement with eqn. (3) the plots of it 1/2 vs. t ~/~ were linear for all thiocyanate concentrations studied and by extrapolation to t 1~2 = 0 one could obtain the value (il~/t)t=o directly connected with the diffusion coefficient Dox = ( il v~)~=olr /n2F2(c°ox)2A 2

(4)

Diffusion coefficients calculated in this way are given in Table 1.

S t u d i e s by using p u l s e p o l a r o g r a p h y The cathodic and anodic curves were recorded for 10 -3 M nickel(II) in solution from 0.094 M to 4.80 M thiocyanates. Total salt concentration was kept constant by addition of NaC104. Similar curves were also recorded for constant 4.80 M NaSCN and different nickel(II) concentrations ranging from 2 X 10 -4 M to 2 X 10 -3 M. Typical recorded curves are shown in Fig. 2. At lower than 2 X 10-4M nickel(II) concentrations pulse polarographic curves were irregular, similar to curves recorded for sufficiently high nickel(II) concentration and

795 TABLE 1 Diffusion coefficients of Ni(II) and kinetic parameters of the electrode reaction occurring in t h e s y s t e m Ni2÷/Ni(Hg) CNi2+ = 10 -3 M; CNaSC N + CNaCIO4 = 4.80 M; DNi(Hg ) = 6.5 X 10 -6 c m 2 s -1 ; t e m p e r a t u r e 2 5 ° C

CSCN-/M

10 6 X DNi2+/cm2 s-I

10 3 X ks/cm s-I

0.094 0.15 0.30 0.60 1.20 2.4O 4,80

2.9 2.7 2.5 2.3 2.2 2.5 2.7

0.77 0.90 1.12 1.23 1.24 1.15 0.99

NaSCN concentration lower than 0.1 M. These curves were n o t used in the kinetic analysis. Measurements were carried o u t for several integration times from 4 to 100 ms. For these integration times good reproducibility of slopes of recorded curves and their potentials was found. Detailed analysis was carried o u t for curves obtained with integration time equal to 49 ms which made possible the use of the tabulated values of X in function of Q'(D-R)Tt' Q'(D-D)AAt given b y Kimmerle and Chevalet [8].

2.O

2.0 L5

LO O.5 0 O.5 tO t5

2.0 2.5

Eo

-/

-i~ E/V vs. SCE

0.5 o

O.5 ~.0 t.5 2.0 2.5

Fig. 2. Pulse c h r o n o c o u l o m e t r i c p o l a r o g r a p h i c curves o f Ni(II) r e d u c t i o n (curve 1) and oxid a t i o n o f t h e h o m o g e n e o u s nickel a m a l g a m (curve 2). 10 -3 M Ni(II) in 0.15 M NaSCN and 4.65 M NaClO4. (1) P o t e n t i a l o f initial electrolysis E 0 = - - 0 . 5 0 V a n d t i m e o f this electrolysis T = 4 s, t i m e o f charge m e a s u r e m e n t t I = 49 ms. (2)E 0 = - - 1 . 5 0 V, ~" = 4 s and t 1 = 49 ms. Charge passing in first m i l l i s e c o n d o f electrolysis was n o t i n t e g r a t e d ; t e m p e r a t u r e 25°C.

796

-

,/0-~

~_xRxN% N

N

'

iO-~

~10-

,

YO-3~

~o-

I"~\ \

-o:o

-'8o

-o.'65 E/V

\

\

x

~o-'

-o)o vs.

SCE

Fig. 3. The dependence of the logarithm of the electroreduction of Ni(II) (o) and oxidation of the homogeneous nickel amalgam (A) on potential. 10 -3 M Ni(II) in 4.80 M solution o f NaSCN + NaC104. Concentration of NaSCN: (1) 4.80 M, (2) 2.4 M, (3) 1.2 M, (4) 0.60 M, (5) 0.30 M and (6) 0.15 M. Arrows point the potentials for which the dependence of log kfh and log kbh on the logarithm of SCN- concentration was studied; temperature 25°C.

The dependence of logarithms of k~h and kbh (determined on basis of eqns. 1 and 2) on potential was linear. Slopes were equal to 39.5 and 119 mV per decade for cathodic and anodic reactions, respectively. Dependences of log kbh on potential calculated from anodic curves recorded starting from potential --0.80 V (generation o f the nickel amalgam) were linear for all concentrations of thiocyanates used. When the potential of amalgam generation was changed to --1.50 V the slopes o f log kbh--potential dependences remained unchanged for lower thiocyanate concentrations. At higher concentrations exceeding 2.4 M the process was a bit slowed down for potentials near to the plateau of the anodic pulse polarographic curves, which appeared as a small change of a slope o f recorded curves. The results obtained for some concentrations of thiocyanates are shown in Fig. 3. The standard rate constants were determined at potentials corresponding to intersections of log kfh and log kbh VS. potential lines. The na values determined from these plots were equal to 1.50 and 0.50 for cathodic and anodic processes, respectively. The kinetic parameters obtained are given in Table 1. The analysis of curves obtained for solutions o f different nickel(II) concentration b u t constant concentration of NaSCN equals to 4.80 M points that kinetic parameters are independent of this concentration.

Determination of formal potentials The knowledge of the formal potentials o f the Ni(II)/Ni(Hg) system in thiocyanate solutions was desirable in order to determine the standard rate constant of the electrode reaction and also to determine the distribution of various nickel(II) thiocyanate complexes in the solutions investigated. It should be said that the application of conventional electrochemical methods to solve this problem is not useful because of very limited stability of the

797 0t

Ct2

a5

¢.0

2.0

'

4•I

5.0

4.0 p-i

2[0

4

32

2.0

2.0

~.0

1.0

a2

a5

io cscN-/ M

Fig. 4. The average coordination numbers (p) of Ni*SCN)p2~p complexes (1 and 2) and numbers of ligands ( p - i ) liberated before the charge transfer proper (3) in dependence on the logarithm of thiocyanate concentration. 10 -3 M Ni(II) in 4.80 M solution of NaSCN + NaC104, temperature 25° C.

h o m o g e n e o u s nickel amalgam and also due to f o r m a t i o n o f NiS which adsorbs on the electrode surface in such a system. In consequence the polarographic m e t h o d was used, however with drop time shorter than 2 s and with nickel(II) c o n c e n t r a t i o n in t he solution equal to 4 × 10 -4 M. Th e polarographic curves o b t a i n e d u n d e r such conditions indicated t h a t t h e process was controlled b o t h by t he rate o f t he charge transfer and transport. To determine the reversible half-wave potentials K oryt a's [9] m e t h o d was applied. The intersections o f a tangent (at the f o o t o f the wave) t o log [ i / ( i l - - i)] vs. E curves with a slope characteristic o f t h a t for a reversible process, with the log [ i / ( i l - - i)] = 0 were taken as E ~ potentials. Their analysis shows t h a t the average co o r di nat i on n u m b e r o f nickel(II) with t h i o c y a n a t e ions changes f r o m 1.0 to 3.8 when SCN- ion c o n c e n t r a t i o n increases f r o m 0.1 to 4.8 M. Curve 2 in Fig. 4 shd~vs this change in detail. Curve 1 in this Figure gives similar d e p e n d e n c e however calculated from stability constants r e p o r t e d in the literature [10] ext rapol at ed t o the ionic strength used in these studies. DISCUSSION

Th e s tu d y o f t he kinetics o f t he Ni(II)/Ni(Hg) system in solutions o f thiocyanates is rather c um ber s om e due t o the parallel e l e c t r o r e d u c t i o n of thiocyanates with f o r m a t i o n of NiS which adsorbs at the electrode surface. Using the m e t h o d with a short time of durat i on of the el ect roreduct i on, t he p r o d u c t i o n o f NiS was very low and n o t sufficient to cover the electrode surface to a significant extent. Also o t h e r conditions o f experiments were such t o

798 diminish as much as possible the formation of sulphides. A p r o o f that the investigated process was n o t influenced, under our experimental conditions, b y such side reactions was supplied b y the independence of the kinetic parameters obtained o f the time of pulse duration in the limits 4 to 100 ms. Also the practically identical anodic rate constants obtained with starting potential --0.8 V (at which NiS is stable on electrode surface) and --1.50 V (NiS is reduced) point on the fact that the blocking effect plays only an insignificant role. However a small decrease of the rate of anodic reaction at large SCN- concentrations and application of more cathodic initial potential m a y be due to presence o f NiS on the electrode surface. Though the formation of NiS on the electrode surface is not possible under such conditions, the S 2- ions produced in cathodic reaction m a y interact at the electrode surface in the anodic step with the nickel(II) appearing due to anodic reaction to form NiSad~. Such a process is stimulated b o t h b y a more negative potential o f the initial electrolysis and larger concentration of thiocyanates. The decrease of the oxidation rate is well observed at thiocyanate concentrations exceeding 2.40 M. The rate constants determined under proper conditions are approximately 3 to 4 times higher than earlier reported [4]. The slopes of dependence of log kfh and log kbh on potential give cathodic and anodic n a values equal to 1.50 and 0.50, respectively. These values suggest a step-wise mechanism with a slow transfer of the second electron. For the detailed description of the mechanism of the studied electrode reactions the composition of the complex directly reacting with the electrode is needed. This was found by analysing the dependence of log kfh on the logarithm of thiocyanate concentration at constant potential. Since the slope of the log k~h vs. potential curve was constant for all thiocyanate concentrations, the determination of the composition of the complex could be done at every potential at which the rate constant was determined. The dependence o f log kfh and log kbh on the logarithm of thiocyanate concentration is shown in Fig. 5. 0.4

02

0.5

KO

2.0

5.O

.,~~iO -2

t0 -2

2

40-'

¢0"~ i

i

(IY

02

,

0.5

i

KO

i

,

2.0 5.0 CscN-/M

Fig. 5. The dependence of the logarithm of the rate constant of Ni(II) electroreduction (1)

and oxidation of the homogeneous nickel amalgam (2)on the logarithm of thiocyanate concentration. 10 -3 M Ni(II) in 4.80 M solution of NaSCN + NaC104; temperature 25°C.

799 Such dependence should have a slope equal to p-i and i' --(~ log kfh/3 log[SCN-])E, ~ = p (a log kbh/O log[SCN-])E,~ = i'

i (5)

where p is the coordination number in respect to thiocyanate ions of the complex existing in the bulk of the solution, p-i gives the coordination number of the complex directly participating in the electrode reaction and i' is the coordination number of the complex formed directly in the anodic reaction. The slope of the dependence for the anodic process is constant in the whole region of thiocyanate concentration and is equal to 0.62 (curve 2 in Fig. 5), while the slope of the cathodic dependence changes from 0.55 to 2.8 as the thiocyanate concentration increases from 0.10 to 4.80 M. This change is mainly due to the change of average bulk coordination number of nickel(II) in respect to thiocyanate. The numbers of p-i are shown b y curve 3 in Fig. 4. Determined from the difference of value p-i (curve 3) and p (curves 1 and 2) the parameter i is constant with a value near to 1 for the whole thiocyanate concentration range. The surprisingly low number of i' obtained for the anodic reaction m a y suggest that the primary complex is formed from t w o nickel(II) ions b o u n d to one SCNion, through sulphur and nitrogen atoms. However such a mechanism can not be operative, because kinetic analysis carried o u t with solution of different concentrations of nickel(II) led to the conclusion that the electrode reaction is of the first order in respect to nickel(II). Using lower concentrations of thiocyanates Kutner [ 11] has found the oxidation of the homogeneous nickel amalgam to proceed to the complex Ni(SCN) ÷ identical with that found in this w o r k for the complex directly reduced at the electrode. One m a y arrive to the conclusion that i' equal to 0.62 found in the present work is t o o low. The deviation from the probably true number equal to 1 m a y be due to some interaction of oxidizing nickel with sulphide ions produced in the cathodic step of the formation of amalgam from nickel(II) thiocyanate complex. The not exactly constant water activity for the solutions studied also play here some role. One may mention that the water activity for 4.80 M NaC104 and 4.80 M NaSCN is equal to 0.794 and 0.748, respectively [12]. Consideration of a slight decrease of the water activity in our experiments (maintaining ionic strength constant) with increase of thiocyanate concentration leads to gradual increase of the anodic rate constant kbh which in turn changes the slope of the log kbh vs. log CscN- dependence to a b o u t 0.7. This still low value m a y be due to adsorption of the thiocyanate complexes of nickel(II) on the mercury electrode. Though the chronocoulometric studies do n o t give support [13] for such an adsorption, none the less the precision of such measurements in case of weakly adsorbed complexes m a y be limited. Also the formation of NiS b y decreasing the rate of the anodic reaction m a y in consequence decrease the i' value. There is some possibility o f direct formation of NiS in the course of the anodic reaction in addition to interaction of nickel(II) with sulphides produced in the cathodic step. However the production o f NiS in that w a y should amount to several per cent in respect to charge

800 passing to oxidize nickel amalgam [4]. It was found that the oxidation of metallic nickel in solution of thiocyanate proceeds to NiS only to the extent of 3%. In conclusion the electrode reaction studied m a y be represented b y the following equations Ni(SCN) 2 - p ~ Ni(SCN) ÷ + (p -- 1)SCN-

(6)

Ni(SCN) * + 2 e ~ Ni(Hg) + SCNkbh

(7)

This mechanism is at variance with those proposed earher [3,4] where complexes with 2 or even 3 thiocyanate ligands were suggested to participate directly in the electrode reaction. These considerations apply to low surface coverage of the electrode b y NiS. At higher coverages as a result of the decrease of the rate of the electrode reaction additional steps in case of voltammetric and minima in case of polarographic curves are observed. These polarographic minima are due to adsorption of NiS and its stability on the electrode surface. The explanation o f their nature as due to for instance electrostatic repulsive interaction of negatively charged surface with anionic complexes would be erroneous. Attempts to study the electrode reaction of these complexes b y using the impedance bridge m e t h o d did n o t give satisfactory results [14], because of adsorption of NiS on the electrode. ACKNOWLEDGEMENT This work was supported by a grant of the research project MR 1-11. REFERENCES 1 M. F l e i s c h m a n n , J . A . H a r r i s o n a n d H . R . T h i r s k , T r a n s . F a r a d a y S o c . , 6 1 ( 1 9 6 5 ) 2 7 4 2 . 2 P. D e l a h a y a n d C.C. M a t t a x , J. A m e r . C h e m . S o c . , 76 ( 1 9 5 4 ) 8 7 4 . 3 J a . I. T u r j a n , Z h . Fiz. K h i m . , 31 ( 1 9 5 7 ) 2 4 2 3 ; J a . I. T u r j a n a n d G . F . S e r o v a , Z h . Fiz. K h i m . , 3 4 (1960) 1009. 4 Z. G a l u s a n d Lj. Jefti~, J. E l e c t r o a n a l . C h e m . , 1 4 ( 1 9 6 7 ) 4 1 5 . 5 E. I t a b a s h i a n d S. I k e d a , J . E l e c t r o a n a l . C h e m . , 27 ( 1 9 7 0 ) 2 4 3 ; E. I t a b a s h i , J . E l e c t r o a n a l . C h e m . , 6 0 (1975) 285. 6 T. K ~ o g u l e c , A . B a r a n s k i a n d Z. Galus, J. E l e c t r o a n a l . C h e m . , 5 7 ( 1 9 7 4 ) 6 3 . 7 T. H u r l e n a n d E. E r i k s r u d , J. E l e c t r o a n a l . C h e m . , 5 4 ( 1 9 7 4 ) 3 3 1 . 8 F . M . K i m m e r l e a n d J . C h e v a l e t , J . E l e c t r o a n a l . C h e m . , 21 ( 1 9 6 9 ) 2 3 7 . 9 J, K o r y t a , E l e c t r o c h i m . A c t a , 6 ( 1 9 6 2 ) 6 7 . 1 0 S. T r i b v l a t a n d J . M . C a l d e r o , C o m p t . R e n d . , 2 5 8 ( 1 9 6 4 ) 2 8 8 8 . 11 W. K u t n e r a n d Z. G a l u s , Bull. A c a d . P o l o n . Sci. Ser. Sci. C h i m . , 2 0 ( 1 9 7 2 ) 1 0 4 5 . 1 2 R . A . R o b i n s o n a n d R . H . Stokes, Electrolyte S o l u t i o n s , B u t t e r w o r t h s ~ L o n d o n , 1 9 5 9 . 1 3 G.W. O ' D o m a n d R.W. M u r r a y , J . E l e c t r o a n a l . C h e m . , 1 6 ( 1 9 6 8 ) 3 2 7 . 1 4 K. W o j d e c k a , M. Sc. Thesis, U n i v e r ~ t y o f W a r s a w , 1 9 7 4 .