The deposition of Ni from aqueous sulphate and sulphamate solutions

Electroanalytical Chemistry and Interfacial Electrochemistry, 44 (1973) 431-443

431

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

THE DEPOSITION OF Ni FROM AQUEOUS SULPHATE AND SULPHAMATE SOLUTIONS

W. DAVISON and J. A. HARRISON

Electrochemical Laboratories, School of Chemtstry, The University, Newcastle-upon-Tyne NE1 7RU (England) (Received 7th September 1972)

INTRODUCTION

Much work has been conducted upon the study of the reduction of Ni 2÷ from an aqueous system of simple salts. The work upon solid electrodes ~-4, produces conflicting results as regards the value of the Tafel slope, the order of the reaction with respect to Ni 2+, and the pH dependency. The foundations for the study of the deposition of Ni 2÷ at mercury electrodes were laid by Vlcek 5 and Dandoy and Gierst 6 who first suggested that the kinetic current which was evident was due to the dehydration of Ni 2+. Work by Hush and Scarrott 7 and most recently by some French workers 8 has concentrated upon further amplification of this basic idea. Based on the work of this present article an entirely different explanation will be presented. To give the study some industrial relevance sulphate and sulphamate are chosen as the major anion species because these form the basis of the two most important plating baths. EXPERIMENTAL

The potential of the working electrode was controlled by use of a potentiostat. Perturbation of the system was effected by superimposing either single step potential pulses or a triangular potential wave. The equipment and cell used have been described elsewhere9. A hanging mercury drop working electrode could be renewed for each measurement. The secondary electrode was a platinum spiral and all measurements were performed against a calomel/saturated KC1 reference electrode. The ohmic drop was found to be negligible. From a knowledge of the geometry of the cell the maximum resistance was estimated as 0.5 fL Sulphate solutions were used with 1 or 10 -1 M Na2SO 4. Nickel in the form of NiC12 6 H 20 , was either 10 .2 or 10 .2 M, but always at most ~o of the total ionic concentration. This excess of swamping electrolyte was sufficient to eliminate migration effects because the observed currents were considerably less than the theoretical diffusion limiting value. All solutions were buffered. Usually boric acid was used to keep the pH in the region 5q5, but phosphate was employed as an alternative. Analogous sulphamate solutions were prepared by replacing Na2SO4 with KNH2SO 3. In all experiments base electrolyte measurements were performed and the resulting current subtracted from the actual so-

432

W. DAVISON, J. A. HARRISON

lution value. This was especially important for short time pulse measurements. All solutions were deoxygenated with N2 before taking measurements. To reduce convection nitrogen was flushed over the surface of the solution while a slow sweep or long time pulse was in process. All reagents used were B.D.H. AnalaR grade except K N H / S O 3. This was prepared by neutralizing B.D.H. "Lab Reagent" sulphamic acid with AnalaR K O H . As a test of purity a comparative base electrolyte was prepared from B.D.H.O.A.S. sulphamic acid. The electrochemical techniques employed could not detect any difference between the two solutions. RESULTS

Generally the data obtained from sulphate and sulphamate solutions were alike. Therefore a detailed discussion of only one of these sets of solutions will be given, but unless otherwise stated it will apply to both. Potential sweep measurements

Figures 1 and 2 show that depending on the ratio of Ni 2+ to anion e~ther a peak or a limiting current was obtained. The limiting current was found to be independent of sweep rate and up to 100 V s- 1 the equipment proved incapable of resolving a further peak. A first order relation for Ni 2 + concentration was found. In Fig. 2 comparison of the currents obtained from sweep measurements at 0.1 V s- a on solutions of 10- 2 M Ni 2 + concentrations and different anion concentrations clearly show that for 1 M anion concentrations the currents were well below the

(A)

(B) 2

macro-

S

5do

lo'oo -ElmV

1~o

46o

66o

86o

l&o

~2oo

-ElmV

Fig. 1. Dependence of potential sweeps on sweep rate. (A) 0.1 M Na/SO4, 0.1 M H3BO3, 10 -z M N1" 2 + ;(a) 0.3, (b) 0.1, ( C ) 0.03, (d) 0.01 V s -1. (B) 1 M KNH2SO3, 0.1 M H3BO 3 10 -2 M Ni2+; (a) 0.3, (b) 0.1, (c) 0.03, (d) 0.01 V s -i.

433

Ni D E P O S I T I O N F R O M S U L P H A T E A N D S U L P H A M A T E SOLUTIONS

t~x

!

4

E u

3

E 2

o

u./o cJ/*

oX~"° (A) A A'°

~A"~'g*

I 1

01

I

9O0

1000

i

i

1100

1200

-ElrnV

g60

950

1dO0

lOgO

- E l rnV

Fig. 2. Dependence of potential sweep on anion species and concentration. All with 10 .2 M Ni 2+, 0.1 M H3BO 3 and at 0.1 V s - t . (a) 0.1 M K N H 2 S O a , ( b ) 0.1 M Na2SO¢ , (c) 1 M KNH2SO3, (d) 1 M

Na2SO4. Fig. 3. Log i vs. E plot showing linear Tafel region, independent of sweep rate. (A) 1 M KNH2SO3, 0.1 M HaBO3, 0.1 M NiZ+; (©) 0.01, (×) 0.03, (D) 0.1 V s -1. (B) 1 M Na2SO¢, 0.1 M H3BO3, 0.1 M NiZ+; (x) 0.01, (@) 0.03, (IS])0.3 V s 1.

diffusion limiting value, while for 0.1 M the currents approach the diffusion limiting value. The low currents make Tafel slope measurements impracticable for some of the curves, but for the higher currents (10 -1 M Ni 2+ solutions) log i vs. E plots were consistent for different sweep rates. Figure 3 demonstrates that a reasonable estimation of the Tafel slope could be made from the central linear portion which will be free of both diffusion and extraneous low currents. Values obtained were for a sulphate solution, 114 mV; for sulphamate, 118 InV. Since we believe the curves to be controlled by a preceding chemical reaction, reaction layer theory should apply, and therefore a plot of l o g ( i l - i ) / i vs. E is applicable. This was prevented because ix was difficult to estimate with any accuracy. Moreover, one end of the plot would be distorted by the presence of small impurity currents. P o t e n t i o s t a t i c pulse m e a s u r e m e n t s

Figure 4 shows that the transients obtained were flat, slowly moving over to the more normal diffusion controlled transient at very high values of potential. F r o m Fig. 5 it is apparent that solutions with 1 M anion concentration produced completely flat curves for the time scale investigated--up to 3 ms--while those of 0.1 M anion concentration yielded transients which approached diffusion control. Figure 6 shows a plot of i vs. t - ~ and clearly demonstrates that the curve is approaching the theoretical diffusion line which is drawn in with a diffusion coefficient, D = 1 0 -5 c m 2 s - 1 . The log i vs. E curves of Figs. 7 and 8, with the current measured at a fixed time of 0.5 ms, show that like the sweep data a fixed anion concentration and variation of Ni 2 + concentration, produced a first order relationship with respect to nickel Low current densities prevented any determination of TaM slopes from the pulse data.

11~

434

W. DAVISON, J. A. HARRISON

(A)

(B)

30

15

6c

T

t

E 20

[::: 4 0 E

L...____


b

E 10

2O

C

-

~

d

d

2 tlms

t/ms

~

~

c t l s ~

Fig. 4. Potential dependence of pulse transients. All from a starting potential of 700 inV. (A) 0.1 M KNH2SO3, 0.1 M H3BO3, 10 -2 M Ni 2+, E = ( a ) -1600, (b) -1400, (c) -1200, (d) - 1 0 0 0 inV. (B) 1 M Na2SO4, 0.1 M H3BO 3, 0.1 M Nia+; E = ( a ) -1600, (b) -1400, (c) -1200, (d) -1000 inV. Fig. 5. Observed transients with base electrolyte current subtracted compared with the base electrolyte current. All taken from a pulse of 700 t600 inV. (a) 0.1 M Na2SO4, 0.1 M H3BO3, 10 -2 M Ni; (b) 1 M Na2SO4, 0.1 M H3BO3, 0.1 M Ni; (c) 1 M Na2SO 4, 0.1 M H3BO 3.

/

30

/ ! 11

i

/ 0

/

t

2O /

E u

i

i

i

0

¢

0 0

/ o

o

0

o

/ c¢9 /i

E

i

0

/

l

lC

/

/

iI

iI

/

/ /

i

lo

1~

2b

Fag. 6. i vs. t - { for 0.1 M Na2SO4, 0.1 M HaBO3, 10 -2 M Ni 2+. The theoretical line for D = t 0 -5 cm 2 s -1 is shown.

Effect of pH and other Iigands present Apart from Ni z +, SO42 - and N H 2 S O j- other species present in solution which may possibly have an effect on the system under study are H +, C1- and boric acid. To test if boric acid took part in the reaction a phosphate buffer was subsiituted. Figure 7 shows that this had no effect on the pulse data. Any possible pH deoendence was tested for by performing sweep measurements

Ni DEPOSITION FROM SULPHATE AND SULPHAMATE SOLUTIONS

1

•

10

435

x

•

x X

EI.)

X

X

x

E

x

x

x

×

×

Et +

x

O

x

1

•

Q

0 [] +

o

x + 0

12~O0

1 4~0 0

1 6 ~0 0

18'00

--

-EImV

Fig. 7. Log i vs. E plots showing change in current with concentrations used. All points taken after 0.5 ms, with SO~ as the major anion species. ( x ) 1 M Na2SO4, 0.1 M H3BO3, 0.1 M Ni2+; pH 5.5 (©) 1 M Na2SO4, 0.087 M KH2PO4, 0.0304 M Na2HPO4, 10 .2 M Ni2+; pH 6.8. ([]) 1 M Na2SO4, 0.025 M KH2PO4, 0.025 M Na2HPO4, 10 .2 M Ni2+; pH 6.2. ( + ) 1 M NazSO4, 0.1 M'H3BO3, 10 .2 M Ni2+; pH 5.5. ( i ) 0.1 M Na2SO4, 0.1 M H3BO3, 10 .2 M Ni2+; pH 5.5.

100 [] [] X

[]

+ ~ O Q

T1C E {J

El

[]

[3

[]

0 0 0 0

E

0

0

X 0

I(?OQ

1200

1400

1600

18~00

-ElmV

Fig. 8. Log i vs. E plots showing change in current with concentrations used. All points taken after 0.5 ms with NH2SO ~ as the major anion species. ( x ) 0.1 M KNH2SO3, 0.1 M H3BO3, 10 -z M Ni 2+. (O) 1 M KNH2SO 3, 0.1 M H3BO3, 10 -2 M Ni 2+. (D) 2.5 M KNH2SO3, 0.1 M H3BO3, 10 :t M Ni 2+. ( + ) 1 M KNH2SO3, 0.1 M H3BO3, 10 1 M Ni 2"-.

on a solution which was progressively made more alkaline with 0.1 M N a O H . The i-E curves recorded were found to be identical up to pH 8.0. Similarly HC1 was added to produce a more acidic solution; there was no resulting change in the i-E curves. Additions of KC1 up to 0.5 M had no effect. In the present w o r k - t h e ionic strength varies between 1 and 0.1. This could affect the activity coefficients of the species involved in the reduction. Data 1° for NiSO4, a more drastic situation than the system being used, gives the mean ionic activity coefficient as 0.15 for 0.1 M N i S O 4 and 0.0425 for 1 M N i S O 4. These

436

w. DAVISON, J. A. HARRISON

measurements do not take into account complexation but merely reflect the fact that complexation is occurring. The basic problem is that activity coefficients are a thermodynamic quantity whereas in investigating the kinetics of the present system it is necessary to consider the effect of the individual species present. We maintain that the results presented here could not be explained on the basis of a change in activity coefficient as this is unlikely to have a ten-fold effect in the experimentally used ionic strength range.

Interpretation of results It is our contention that all the results may be interpreted on the basis of a preceding chemical reaction of the type: ze

~k

~---MXk_ 1

Here X represents the sulphate or sulphamate species. Evidence of complexes of this nature does exist s 1.1:, but no quantitative values for equilibrium constants have yet been established. It is easy to visualize that possible complex equilibria may be: M 2 + ~ - M S O 4 ~ [ M (SO4)2] 2(I) and M 2 + ~- [MNHzS03 ] + ~ - M (NH2SO3)2 -~[M (NH2SO3)a] - ~_ ~-[M(NHzSO3)4] z(II) whilst the possibility of other (e.g. dimeric) species must not be excluded. Depending on the relative values of the stability constants involved it is possible that either an increase or decrease in the ligand species concentration may lead to an increase in any given species in solution. This may be the reducible species, and so the current observed is dependent upon the stability constants and the rate of exchange of the complexes present. Examination of the electrochemically controlled part of the wave should provide information on the species being reduced, while in the limiting current region information should be furnished regarding the total concentration of all species which are in sufficiently rapid equilibrium with the reducible species. The fact that fiat potentiostatic transients and limiting sweep curves independent of v are produced implies that the conditions for the reaction layer theory to hold are in operation and therefore that we are dealing with a preceding chemical reaction. Let us consider: kf

X+B ~ A ze ~ kb

put K = [A]/[B] [X] = kf/k b then reaction layer theory 13 gives:

I1 = zFACRkb(D/kf[X]) ~

(1)

IX] is the concentration of the ligand species and therefore if it is in sufficient excess, it may be taken to be the actual initial concentration supplied to the solution. cR is the concentration of reducible species as used in the reaction layer theory. If preceding equilibria exist which are faster than, and feed into, the chemical reaction

Ni D E P O S I T I O N F R O M S U L P H A T E AND S U L P H A M A T E SOLUTIONS

437

that is experimentally detected then cR may be greater than the equilibrium concentration of A which is the concentration employed in the simple theory. By employing different ligand concentrations in the pulse measurements the CR values and thus the limiting current density, i~, may be effectively changed. Thus for two solutions (1) and (2): i1(1) =

ZFCR(1)kb (D/ke [X(1)])~

i1(2) ---=ZFCR(2) kb (n/kr

(2)

[X(2)])~

(3)

Therefore

fix( 2)]/[x(1)] )* = ( i,(1)/il( :3 ( cR( 2)/~( ~)) i1(1), i1(2) are measured experimentally, [X1] , [X2] are known, thus CR(1)/CR(2)may be calculated. Further information is available from sweep data. Firstly, for the case where a limiting current, independent of sweep rate, is obtained we may use it to check the pulse data as reaction layer theory is applicable in both cases. Secondly, where peaks are obtained we may use the treatment of Nicholson and Shain .4. Their equation:

ik/id = 1/(1.02 + 0.531 (b') ~ [X] ~ k~/kb) is used to determine

0 = (b')~ [x]~

k~/kb

where

b' = an, Fv/RT = 2.3 v/b and b is the experimental Tafel slope. ia was calculated first from the slowest sweep performed. If this did not produce a constant value of O/v ~ then it was apparent that the slowest sweep was not purely diffusion controlled, ia was then raised until a value was achieved which produced the best constant value of O/v ~. This involved a correction ca. 10~. Using the above equations:

O/v* = (2.3/b) ~ [X] ~*kf*~/kb i.e.

kb/(~ IX] )~ = (Z3/b)~/(O/~ ~) TABLE, 1

Predominant ligand

Potential

(Ni 2+)

il(mAcm 2)

NH2SO 3 0.1 M NH2SO ~- 0.1 M NH2SO ~ 1 M NH2SO ~- 1 M SO420.1M SO~1M

1400 1600 1400 1600 1600 1600

10 -2 M 10 2 M 10 2 M 10 -2 M 10 2 M l0 -2 M

35 52 3.3 7 32 0.56

438

W DAVISON,J. A. HARRISON

o

oOO

T i

o

o

2

o

E u .< E

T

0

e~ "7

X XX

0

0

xX X

X

x

v1121V1/2 s-1/2

~

v1121 V1/2 s-1/2

Fig. 9. Plot of iv vs. v ~ to demonstrate limiting value of tp. (©) 1 M KNH2SO3, 0.1 M H3BO3, 10 -2 M Ni2+; ( x ) 1 M Na2SO 3, 0.1 M HaBO 3, 10 -2 M Ni 2+ . Fig. 10. Graph showing constant value of O / v ~ with respect to v+. (©) 0.1 M Na2SO4, 0.1 M H~BO3, 10 -2 M NiZ+; ( x ) 0.1 M KNH2SO 3, 0.1 M H3BO3, 10 -2 M Ni 2+.

using the pulse data from the same solution by combining with eqn. (2)

i1(1) = zFcrt( 1)(2.3D/b)~/(O/v ~) It is clear that sufficient information is available to calculate the various values of CR present in solution. Table 1 shows the limiting current data obtained from pulses after 0.5 ms. These values can be compared with the graph of current vs. sweep rate, Fig. 9, which demonstrates the limiting currents obtained from sweep measurements. Figure 10 gives the values of O/v ~ with respect to sweep rate and shows that a reasonable constant is obtained in both cases. After the appropriate calculations the following reaction layer concentrations are obtained, all for solutions with 10 .2 M Ni 2+. 1 M SO3NH2, cR=8.9x 10-3; 0.1 M NH2SO3, CR=3xl0-2; 1 M SO 2-, CR= 2.3 X 10-3; 0.1 M SO]-, cg=4.2 x 10 -2. It is apparent that these values tend to err on the high side. This is quite explicable if the various inaccuracies of the determination are considered. Firstly solutions with an anion concentration of 0.1 M exhibit results of an almost diffusion controlled type and therefore it is not to be expected that the reaction layer theory will hold in these cases. Moreover, the ligand concentrations, IX], used have been taken as the actual concentrations of ligand supplied to the solutions. Since some complexation is assumed, this cannot be true. If values for all the stability constants of the system were known then the actual equilibrium concentration of ligand species could be evaluated, so improving the above calculation, and the cR values obtained could then be used to determine the identity of the reducible species. The cR values calculated are not entirely redundant, however, because they serve to illustrate that the larger reaction layer concentrations coincide with the solutions with 0.1 M anion concentration and were closest to a diffusion controlled system. Whereas, the smaller concentrations obtained correspond to the solution which

Ni D E P O S I T I O N F R O M S U L P H A T E AND S U L P H A M A T E SOLUTIONS

• 439

produced limiting sweep currents and flat transients, indicative of control by a preceding chemical reaction. A further demonstration that the results may be explained by reaction layer theory is obtained by comparing the experimental sweep results with theoretical sweep curves. In order to facilitate this the curves have been reduced to a dimensionless form by normalizing with respect to the peak current and potential. The resulting plot is of i/i v vs. ( E p - E ) / E p / 2 . Figure 11 shows that for a solution containing 0.1 M sulphamate the experimental curve is very close to the theoretical shape for a reduction controlled jointly by electrochemical step and diffusion. For 1 M sulphamate, Fig. 12 shows the appropriate shape for a process controlled by a combination of electrochemical step, preceding chemical reaction, and diffusion, results. The correct change in shape with respect to sweep rate is obtained and also it may be seen that the part of the wave prior to and including the peak is unaffected by the presence of the slight bump evident after the peak. Thus further justification for treating the data as a system controlled by preceding chemical reaction is provided. Moreover, our action in not taking into account the part of the

o o-..o

1

°

05

.

o~ / o

--o - - - - - - ~ ° - j ° ~

o,

2

1

b

J

(Ep-E)/Epl2--~

Fig. 11. Plot of i/ip vs. (Ep-E)/E~/2 to compare an experimental sweep curve with the calculated curve for a reaction controlled both by electrochemical step and diffusion. (O) 0.1 M KNH2SO3, 0.1 M H3BO3, 10 -2 M Ni2+; 0.03 V s ~, ( - - ) theoretical shape.

1

O~X_~.~C~,..Z

1

o

/

X

-X

X

X

o/r

x

2

1 (Ep

-E)/Epl 2

6

J

-

Fig. 12. Hot of i/i v vs. ( E p - E ) / E p / 2 to compare experimental sweep curves with calculated curves for a reaction controlledjointly by electrochen~ca] step, preceding chemical reaction, and diffusion. ( x ) i M KNH2SO3, 0.1 M H3BO3, 10 -2 M Ni 2+, 0.1 V s - i ; ( O ) 1 M KNH2SO3, 0.1 M H3BO3, 10-2 M Ni 2 +, 0.01 V s-1. Theoretical shapes: ( ) 0 = 1, ( ...... ) 0 = 3.

440

W. DA¥ISON, J. A. HARRISON

1300

1

cl

1200

> IF

o

u.i

o

x

b

x

11od

Q01

0.1

1

10

log ( v / V s - 1 ) ~

Fig. 13. Ep vs. logv. (O) 0.1 M Na2SO4, 10 .2 M Ni 2+, 0.1 M HaBO3; ( x ) 0.1 M KNH2SO3, 10 -2 M Ni 2+, 0.1 M I-I3BO 3.

wave after the peak seems vindicated. The data for producing the theoretical curves was taken from Nicholson and Shain ~4 who also give a graph showing the theoretical dependency of Ep with respect to log v ~-. Figure 13 is an experimentally determined plot for the present work. The curvature is in the predicted direction for a chemical reaction preceding an irreversible charge transfer. Comparison with reduction on solid nickel

An experiment was performed using a nickel wire instead of a mercury drop. High purity nickel wire was fused into a glass rod and sealed from the inside

oj _ 002] 600

800

1000 -ElmV

1200 -

Fig. 14. Comparison ol sweep on Ni electrode with that on Hg electrode. 1 M NazSO4, 0.1 M H3BO3, 10 _2 M Ni 2 at 0.03 V s -1. (a) Ni wire electrode; surface area 0.398 cm 2, (b) base electrolyte on Ni wire electrode. (c) Hg drop electrode; surface area 0.0645 c m 2.

Ni DEPOSITION FROM SULPHATE AND SULPHAMATE SOLUTIONS

441

with araldite. It was so arranged that it fitted into the same cell as used for the hanging mercury drop. The wire was carefully polished with successive grades of emery paper and finally 7-alumina. Linear potential sweeps were then performed upon this working electrode using a solution with and without nickel present. The hydrogen evolution current, as observed from base electrolyte currents is substantial and it proved difficult to obtain reproducible results for the deposition current. Because of this it proved impossible to study the sweep rate dependency. However, Fig. 14 shows an i-E curve produced by comparable sweeps on both nickel and mercury electrodes. The correlation is quite reasonable, especially when it is realised that errors in measurements of surface areas of solid electrodes can be substantial. It is apparent that the reduction current is of a similar magnitude and is occurring at the same potential, so showing that it is probable that the same reaction could be occurring at both the nickel and mercury surfaces. GENERAL DISCUSSION

Kinetic type behaviour for aqueous nickel solutions has been previously observed in polarographic and gatvanostatic measurements 6-8. Dandoy and Gierst and subsequent workers, who have undertaken the experimentation in perchlorate solutions, have interpreted these results on the basis of the slow step involving the dehydration of the nickel cation •

2+

Ni (H20)~ + ~-~ N1 (H2 O), _ x +.x H 2 0 Ni in which the rate of this preceding dehydration reaction is retarded by the potential of the outer Helmholtz plane of the double layer. This argument is dependent upon the ClOg not complexing with the Ni 2+ ions. However, evidence is certainly available for perchlorate forming complexes with other cations 11"15 and some data for complexation with Ni 2 + has been published 12' ~6. Therefore in the concentrated perchlorate solutions used in the experiments performed it is conceivable that complexation may be present. Recent work by Bennes 8 has produced evidence which further disrupts the arguments of Gierst and colleagues. Using a similar system he has obtained a value for the reaction layer thickness from galvanostatic measurements. Bennes calculated this to be much greater than the diffuse l@er thickness. This would suggest that a solution process is prevalent and therefore the outer Helmholtz plane potential can have no significant effect on the measured rate. Bennes suggests, in view of this fact, that the results in perchlorate can be explained by considering the rate determining step as a heterogeneous dehydration occurring at the mercury surface. Although the results for perchlorate might satisfy this argument, for the solutions used in this work large changes in rate produced by variation of sulphate and sulphamate concentrations can not be explained on this basis. Control of the slow step by preceding chemical equilibria of complexes in solution would therefore seem to be the most likely mechanism for the cathodic reduction of nickel from aqueous sulphate and sulphamate solutions. The reaction scheme may be expressed as:

442

W. DAVISON, J. A. H A R R I S O N

[N~/f~n] 2-nz~ - ~ ... [NiXj] 2.j~o ~ ... [NiXk] 2-kz~ ~

[N]LX~_1] 3.k~" ~ . . . . [Ni] 2+ 4,

Ni where NiX k_ 1 is an unknown species occupying any position between the aqueous Ni 2 + cation and the cation of maximum coordination. The proposed equilibria, I and II, show the probable species present, but it must be emphasized fhat until stability constants are determined these can only be suggestions. It is noticeable that the Tafel slope obtained of 120 mV demonstrates a low dependence of the process with respect to potential when compared to the 60 mV value obtained for the aqueous N i - N H 3 system 9. A possible reason for this could be inhibition by one of the species in solution. More work will be necessary on this particular aspect of the subject. LIST O F SYMBOLS

A b b' co

CR D E

E. Ep/2

F I i

f~ ik il

k~ k~ k-1 K n

n~ T t I) z

z~ o~

area of electrode, cm 2 Tafel slope, mV/decade =2.3 v/b bulk concentration of reducible species, tool cm-3 concentration applicable to reaction layer theory, tool cm-3 diffusion coefficient, cm 2 s-1 electrode potential, V peak potential, V potential at half the peak current, V Faraday current, A current density, A cm-2 diffusion controlled current, A cm-2 peak current density, A c m - 2 kinetically controlled current, A cm-2 limiting current density, A cm -2 forward rate constant backward rate constant co-ordination number of complex species being reduced stability constant maximum co-ordination number number of electrons in rate determining step absolute temperature time, s rate of potential change, V s- 1 overall number of electrons charge of anion charge transfer coefficient

Ni DEPOSITION FROM SULPHATE AND SULPHAMATE SOLUTIONS

443

REFERENCES B. Le Gorrec and J. Guitton, C.R. Acad. Sei., Set. C., 272 (1971) 2031. F. Ovari and A. L. Rotinyan, Elektrokhimiya, 6 (1970) 528. I. Epelboin and E. Wiart, J. Electrochem. Soc., 118 (1971) 1577. A. T. Vagramyan, M. A. Zhamagortsyan, L. A. Uvarov and A. A. Yavich, EIektrokhimiya, 6 (1970) 755. 5 A. Vlcek, Chem, Listy, 50 (1956) 828. 6 J. Dandoy and L. Gierst, J. Electroanal. Chem., 2 (1961) 116. 7 N. S. Hush and J. W. Sca~rott, J. Electroanal. Chem., 7 (1964) 26. 8 R. Bennes, C.R. Acad. Sci., Set. C., 273 (1971) 206. 9 W. Davison and J. A. Harrison, J. Electroanal. Chem., 36 (1972) 399. 10 R. A. Robinson and R. H. Stokes, Electrolyte Solutions, Butterworths, London, 2nd ed., 1968. 11 G. Sillen, Stability Constants of Metal-Ion Complexes, Spec. Publ. No. 17, The Chemical Society, Burlington House, 1964. 12 G. Sillen, Stability Constants of Metal-Ion Complexes, Spec. Publ. No. 25, The Chemical Society, Burlington House, 1971. 13 S. G. Mairanovskii, Catalytic and Kinetic Waves in Polarography, Plenum Press, New York, 1968. 14 R. S. Nicholson and I. Shain, Anal. Chem., 36 (1964) 706. 15 G. Archdale and J. A. Harrison, J. Electroanal. Chem., 39 (1972) 357. 16 O. H. Brown, R. H. Nuttall, J. Mcavoy and P. W. A. Sharp, J. Chem. Soc., (A) (1966), 892. 1 2 3 4