The anodic dissolution of nickel—II. Bromide and iodide electrolytes

THE ANODIC BROMIDE

DISSOLUTION AND IODIDE

OF NICKEL-II. ELECTROLYTES

G. T. BURSTEIN* and G. A. WRIGHT Department

of Chemistry, University of Auckland, Private Bag. Auckland, New

Zealand

(Rweiaed I1 August 1975) Abstract-The kmetics of the anodic dissolution of nickel were studied in acidic bromide and iodide solutions using a potentiostatic sweep technique. The rate of dissolution is almost unaffected by the halide at low concentrations due to the presence of a prepassive film. At higher halide concentrations the prepassivc film is removed and active dissolution occurs. The reaction is first order in Br-, zero order in I- and independent of nH for nH < 2. The Tale1 slope is 82 mV, in accord with earlier

work in chloride solu~ons.

INTRODUCTION

In the earlier paper[l] it was proposed that nickel dissolves anodically in perchlorate solutions through an intermediate prepassive film and that the kinetics of active dissolution from a film-free surface can be studied by using fast potential sweeps, or by the addition of F-. The present paper deals with the kinetics of nickel dissolution from a film-free surface in moderate concentrations of Br- and I-. Several studies have already been made of the kinetics of nickel dissolution in chloride electrolytes. Hollnagel and Landsberg[Z] ; Kronenberg et n1[3] found Tafel slopes between 70 and 90 mV in chloride solutions. They proposed that Cl- was adsorbed on the metal surface and that this inhibited film formation. Other workers[4-101 have observed that chloride accelerated the dissolution process. Heitz[l I] found the chloride order dlog,O@log,,[C1-] = 1, with Tafel slopes of about 50 mV for dissolution in ethanolic solutions containing HCI, LiCl and H,SO,. He proposed a single step mechanism involving the transfer of 2 electrons: Ni + CI- + NiCl’ f 2e-. Kolotyrkin et ul[12] anodised nickel in acidic sulphate solutions containing up to 10m4M KT. They found that the reaction was inhibited by increasing I- concentration, and that the Tafel slope increased from 33 mV without I- to 83 mV in 10m4M I-. Vagramyan et a/[131 observed similar inhibition by I- up to 4 x IOm3M, but also noticed that at higher concentrations the dissolution process was accelerated.

sweep rate was 10 mV s-’ except of sweep rate was being tested.

where

the effect

Electrolytes were prepared from analytical grade reagents using doubly distilled waler. Hydrobromic and hydriodic acids were distilled over HJPOZ in an atmosphere of nitrogen[l4,15]. The resulting solutions were redistilled under nitrogen, diluted to the desired concentration with boiled out water and stored under nitrogen in the dark. Constant ionic strength was maintained with Nat and ClO,. @ = 1.00 M for bromide electrolytes and 2.00 M for iodide electrolytes.) Electrolytes of higher pH were not buffered. Experiments were carried out at T = (298 + 0S)“K. The polarization curves were measured in two groups for each halide: (a) constant [H+] with variable [X-l and(b) constant [X-l with variable [H’]. Each group consisted of continuous forward and reversed sweeps, with the ion of variable concentration being titrated into the electrolyte from an external solution. In each case the metal surface was allowed to reach a steady state before recorded sweeps were made, indicated by successive sweeps being coincident. This procedure was important in cases where the metal became pitted, in order that a meaningful comparison of the reaction rates at various concentrations of the ions in solution could be made. Corrections for the junction potential E, between the cell and reference electrolytes, and the ohtmc potential drop En (see Fig. 1) were applied[l]. Thus E,,,,, = E - E,,, - E, - En.

EXPERIMENTAL

Rotating disc electrodes were prepared from speo troscopically pure nickel rod (Johnson-Matthey, 0.5 cm dia) sealed with epoxy resin into Perspex or Pyrex. The counter electrode was made from a 1 cm2 sheet of bright platinum. Potentials were measured with respect to a saturated calomel electrode and converted to the normal hydrogen scale. The potentiostatic sweep technique was used as described earlier[l]. The * Present Address: Dept. of Metallurgy and Materials Science, University of Cambridge, Pembroke Street, Cambridge CB2 3QZ, England. 311

RESULTS Bromide

electrolytes

Tafel lines obtained in solutions of [Br-] < 0.02 M showed hysteresis between positive and negative sweeps and were approximately independent of bromide concentration (Fig. 2). Figure 1 shows Tafel lines for Br- concentrations greater than 0.02 M. They are linear over almost three decades of current density with no hysteresis. The average Tafel slope measured over 58 sweeps is b = 84 mV, with a maximum variation of +7 mV, and is independent of [Br-] and E. Figure 2 shows the reaction order with respect

G. T. BURSTEINand G. A. WRIGHT

IO

Fig. 1. Tafel lines for nickel dissolution in 1.00 M Hr. XM BrC. w = 101 rad s-', I, x = 0.020 M. b = 92 mV; II, x = 0.039 M, b, = 79 mV; III, x = 0.057 M, b = 79 mV; IV, x = 0.167 M, h = X6mV; V, x = 0.286 M, b = 79 mV; VI, X= 0.412 M, b =79 mV; VII, x=0.800 M, h= X2 mV. Broken lines are corrcctcd for ohmic potential drop.

I

I

20

,

I

1

Fig. 3. Rate of nickel dissolution in 1.00 M Br- as a function of [H’]_w = 137 rad be’. l, E = +20 mV (nhe); 0, E = + 180 mV (nhr).

to Br-. Over a range of potentials alog,oi/~loglo [Br-] = 1.1 k 0.1. In electrolytes of [I-I+] < 0.005 M a solid was observed to precipitate in the electrolyte during potential sweeps. Although the metal did not passivate there was no linear Tafel hehaviour. For electrolytes in which [H+] > 0.005 M, [BY] = 1.00 M the Tafel lines were again linear. Figure 3 shows the reaction order with respect to H+ at various potentials. Thus in bromide electrolytes ~log,oi/alog,o [H’] = 0.0. There was no evidence of prepassivation of the metal[l] in electrolytes of [Br-] > 0.02 M. It was found in bromide experiments that a reproducible electrode surface could be obtained by using preliminary cyclic potential sweeps in a solution of 0.02 M Br-. 1.00 M H+ over a range which extended 100 mV into the cathodic region (reduction of H*), through the active region and into the passive region. By repeating this process several times before the commencement of recorded sweeps a solid film is formed in the anodic region which is reduced in the

cathodic region, giving a smooth film-free surface with no subsequent pitting during the anodic dissolution process, even at higher bromide concentrations. No further cathodic sweeps were made, and any hydrogen produced durmg cathodisation was removed by anodic oxidation. However, this condition was achieved only in solutions of low bromide concentration and did not occur for the set of sweeps from which Fig. 3 was derived, where [BrC] = 1.00 M throughout. In this case pitting occurred, and although the electrode was allowed to reach a steady state the surface area was not compatible with that of Figs. I and 2.

Fig. 2. Rate of nickel dissolution in 1.00 M Hi as a func-

Fig. 4. Drift of cd for nickel dissolution in 2.00 M I-,

tion of [BrCJ.w= 101 rad SC’. n , E= +I0 mV @he); 0, E = +60 mV @he); 0, E = + 110 mV @he); 0,

O.OI3M Hi during continuous potential sweeping at 10 mV s-‘, 0, E = 0 mV @he); a, E = +60 mV @he); l, E = + 120 mV (&).

Iodide

electrolytes

In all experiments in iodide electrolytes the metal surface became pitted. Once pitting had commenced the surface was allowed to achieve a steady state by continuous potential sweeping before continuation of the experiment. Subsequent sweeps showed only small drift with time as is shown in Fig. 4. The surface

t, min

E = + 160 mV @he).

313

The anodic dissolution of nickel-~11.

I

1

I

I

1

I

10

-1.5

E

mv,

-10

Fig. 5. Tafel lines fc-r nickel dissolution in 2.00 M H+, xM I-.w = 126 rad SC’. I, x = 0.040 M, h = 86 mV; II, x = (I.020 M, h = 93 mV; III, Y = 0.246 M, h = 90 mV; IV,x=O.781 M, h=K7 mV; V, .x=3X7 M, h=X9 mV;

area must therefore be approximately constant. However. cc/s were calculated using the total electrode area since the true active area was unknown. Figure S depicts Tare1 lines at different iodide concentrations. Curvature at high current densities is probably a result of the pitted surface. Although a correction was applied for the ohmic potential drop between the Luggin capillary tip and the metal surface, no calculation could be made for the ohmic rcGlance of the electrolyte within the pits. At sufficiently low cd the ohmic resistance is negligible. By extrapolating the linear regions of Fig. 5 to high i, the difference between the measured potential and the extrapolated potential should give the ohmic resistante when plotted against the current I. Most of these plots were curved, but gave approximate resistances between 2 and 20 Q. Depletion of I- within the pits through formation of NiIf (see discussion) was also considered, but must bc rcgardcd as unreasonable at high concentrations of I-. However. Tafel curvature persists even at [I-l = 1.45 M. For i < 2 mA cm-‘, the lines shown in Fig. 5 are linear over a decade of cd: h = 82 mV with a maxi-

-05

0

log, [H+l mol C-’

“he

Fig. 7. Dependence of the rate of dissolution of nickel in 2.00 M I- on [H+].w = 121 rad sm’. 0. E = 0 mV (&cl; 0. E = + 1XIImV (n/w).

mum variation of + I5 mV over 59 sweeps. A plot of log,,i against log,,, [I-] at constant E (Fig. 6) shows that on the pitted surface the dissolution rate is approximately independent of the iodide concentration: for [I-] > 0.07 M, dlog,&/?log,, [I-] = 0.3 h 0.1. The dissolution rate in 2.00 M I- is independent of [H+] (Fig. 7), [Ni*‘], potential sweep rate and electrode rotation rate (for w > 0). nlWUSSlON

In the presence of halide ions the dissolution of nickel proceeds through two parallel paths: the first involving adsorbed water and the formation of hydroxy complexes and the second through adsorbed halide with the formation of halide complexes. The total cd i, = i, + i_r,where i, and i, arc the respcctivc partial rates of the two processes. For low concentrations of the halide i, is small and the reaction rate is i,. In this case the metal prepassivates as described in the earlier paper[I]. Thus Figs. 2 and 6 show only slight effect of the halide at low concentrations. This is consistent with the work of Ammar et alrl61. Kolotyrkin t’f a1[12], Vagramyan et ul[13]. At higher halide concentrahons where i, exceeds i, prepassivatlon is prevented by insufficient production of NiOH and by adsorption of the halide directly onto the melal surface: Ni.H20,d,

+ X-d

Ni.X& + Hz0

(1)

If 0 is the coverage of X,, then KI = a&Y-] (1 - 0)

(2)

Two possible mechanisms are considered: (a)

Ni.X,,+

NiX+ + 2e-

(3)

(b)

Ni. X2, + Nix,,, + r-

(4)

N&i,

lag. [I-l. mot I-’ Fig 6. Dcpcndcncc of the rate of dissolution of nickel in 2.00 M H’ on [I-].w = 126 rad CL. l . E = t 100 mV (nhu); 0. E = + 140 mV (n/w).

-

Nix+ + Y-

(5).

Both of these mechanisms are followed by a rapId solvation equilibrium of the Nix+ ions. The rate of reaction (3) or (4) is given by: @FE i, = 2FklBexp ~

[ RT

1

(f-4

G. T. BURSTHNand G. A. WRIGHT

314

where kj is the rate constant for reaction f and n = 1 or 2 for’mechanism (b) or (a) respectively. Substituting for 0 from equation (2) gives i, = 2FkiK1U(-I(1

@FE RT

+ K,[X-I)-‘exp

[

1

(7).

For reaction (3) rate determining (n = 2) with fi = 0.36. T = 298 K: h = 82 mV. For reaction 141 rate deteimining (n = 1) with j? = 0.72, T = 298 K:‘ d = 82 mV. Both of these values of fl are reasonable and the experimentally observed Tafel slopes do not distinguish between the two mechanisms. The slopes are

similar to those found in chloride solutions[2,3,17]. Mechanism (h) is comparable with that proposed in the absence of halide[l]. For bromide solutions we regard 0 as small and K, [Br-] 4 1. Thus equation (7) becomes: ig, = 2FkjK,[Br-]exp

L1 @FE RT

,

and ~log10i,,/810g10 [Br] = 1, in agreement with the observed values throughout the potential range. In iodide solutions 0 approaches unity (K,[I-] 4 1) and from equation (6) 8log,&/dlog,, [I-] = 0, in agreement with Fig. 6. Both mechanisms are independent of pH, and this is in accord with Figs. 3 and 7. Acknowledgement-One of us (G.T.B.) wishes to acknowledge the New Zealand Universities Grants Committee for a Post-Graduate Scholarship.

Rl?FERENCES

1. G. T. Burst& and G. A. Wright, Electrochim. Acta 2). 95 / 19751. 2. M. Hoilnagci and R. Landsberg, Z. phys. Chem. 212A. 94, 127 (1959). 3. M. L. Kronenberg, J. C. Banter, E. Yeager and F. Hovorka, J. rlectrochem. Sot. 110. 1007 (1963). 4. R. C. V. Piatti, A. J. Arvia and J. J. Podesta, Electrochim. Actu 14. 541 119701. 5. D. L. Piron, E. P. K&&kos and K. Nohe, Corrosion 25. 151 (1969). 6. T. Tokuda and M. B. Ives, Corros. Sci. 11. 297 (1971). 7, G. Trabanelli, F. Zucchi and L. Felloni, Corrosion Sci. 5. 21 I (1965). 8. I. Garz, H. Worth and W. Schatt, Corrosion Sci. 9, 71 (1969). 9. W. Schatt and H. Worth, Corrosion Sci. 9, 869 (1969). 10. G. Masing and G. Roth, Werksro& Korros. Lpz. 3. 253 119521. 11. El .H‘eitz, &ctroch~m. Acta IO, 49 (1965). 12. Y. M. Kolotyrkin, G, Ci. Lopovok and L. A. Medvedeva, Pmt. Met. 5. 1 (1969). 13. A. T. Vagramyan, M. A. Zhamagortsyants, L. A. Uvarov and A. A. Yavich, Prof. Met. 5. 60 (1969). 14. G. Brauer. Handbook of Preparative Inorganic Chemistry Vol. 1, 2nd edn, p. 287. Academic Press, New York ,5. (1963). J. W. Mellor, A ComprehPnsioe Treatise on Inorganic and Theoretical Chemistry Vol. 2. D. 168, Longmans. London (1922). 16.I. A. Ammar. S. Darwish and S. Riad, Electrochim. Acta 13. 1875 (19681. 17. M. Zamin and M. B. Ives, J. electrochem. Sot. 121, 1141 (1974).