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On the anodic passivity of tin in alkaline solutions
Elcctrochhica
Acta. 1964, Vol. 9. pp. 883 to 8%. Pcrgamon FTCSSLtd. Printed in Northern Irehttd
ON THE ANODIC PASSIVITY OF TIN IN ALKALINE SOLUTIONS* A. Laboratory
M.
SHAMS EL DIN
and F. M. ABD EL WAHAB
of Electrochemistry and Corrosion, National Research Centre, Dokki, Cairo, U.A.R.
Abstract-The anodic oxidation of tin was studied galvanostatically in NaOH solutions of different concentrations. Primary passivity is attained in all solutions when the metal is covered with a film of Sn(OH), or SnO. Permanent passivity sets in when Sn(OH), forms as a continuous layer on the electrode surface. In concentrated alkali solutions larger quantities of electricity are charged to the electrode to compensate for the chemical dissolution of the hydroxides. As a result of the competition between the anodic formation and chemical dissolution of Sn(OI-I),, the potential of the tin electrode oscillates between the Sn/Sn(OH), and Sn(OH),/Sn(OH), values. For any one solution the oscillation phenomenon depends primarily upon the polarizing current. Permanent passivity is governed by the relation (i - io)P = const. where i. is the current below which no passivation takes place, T the time elapsed till oxygen evolution and n is a constant. Interruption of the polarizing current when the electrode is at oxygen-evolution potential gives rise to a pH-independent arrest at potentials more positive than any of the known redox values for tin. This arrest is ascribed to the chemical dissolution of metastannic acid formed at oxygen-evolution potentials. The anodic behaviour of the tin electrode in solutions containing stannite and stannate is similar to that in solution8 devoid of these ions. An increase in the tin content of the solution is equivalent to the elevation of the polarizing current ; passivity sets in more readily in the concentrated solutions. R&suru&L’oxydation anodique de Sn est 6tudiCe galvanostatiquement en NaOH de concentration vari6e. Une pas&it6 primaire apparalt dans toutes les solutions, quand le m&al se recouvre d’un film de Sn(OH), ou SnO. Une passivite permanente est due a une couche continue de Sn(OH),. Dam les solutions alcalines concentrees, la charge de l’&ctrode implique des quantit6s d’&ctricit6 accrues, pour compenser la dissolution chimique des hydrox des. Le potentiel de l%lectrode d’&ain oscille db lors entre les valeurs propres a Sn/Sn(OH), et Sn rOH),/Sn(OH),, ce qui traduit une sorte de comp6tition entre la formation anodique et la dissolution chimique de Sn(OI-I),. Ce ph6nom&ne d’oscillation d6pend surtout du courant polarisant. La pas&it6 pcrmanente est gouvem6e par la relation (i - &,) 71 = cte, i. &ant le courant au dessous duquel il n’y a pas passivation, 7 le temps pr&dant l’bvolution d’oxygene, n une constante. Si l’on interrompt le courant polarisant quand le potentiel d’6volution d’Ot est atteint, le potentiel d’arr&, indtpendant du pH, est plus positif que toute valeur r&lox deja connue pour p∈ parret est attribut B la dissolution de l’acide m&a&annique engendn5 aux potentiels d%volution d’0 *. Le comportement anodique de l’&ctrode Sn dans des solutions de starmites et stannates est analogue. zuplemmenhpsslpg_Die anodische Oxydation von Sn in NaOH-Liisungen wurde galvanostatisch untersucht. Prim&e Pas&it&t setzt ein, wenn sich Sn(OH), oder SnO bildet, permanente Pas&it&t, wenn eine zusammenh&ngende Deckschicht von Sn(0I-I). entsteht. In konzentrierten NaOHLiisungen werden grossere Blektrizititsmengen beniitigt, urn die chemische Aufliisung deer Hydroxyde zu kompensieren. Infolge der konkurrierenden Bildung tmd Au&L%mg von Sn(OH), oszilliert das Potential zwischen den Sn/Sn(OH), und Sn(OH),/Sn(OH),-Werten. Die permanente Passivitit gehorcht der Gleichung (i - io)rn = Konst. * Manuscript
received 9 November
1963.
A. M. SIUMS EL DIN and F. M. ABD
884
EL WAHAB
wobei n eine Konstante ist, 7 die Zeit bis zum Einsetzen der O,-Entwicklung und i. der Strom, unterhalb dem keine Passivierung eintritt. Die Abschaltkurven we&n einen pH-unabhlngigen Haltepunkt auf, der bei einem positiveren Potential liegt als siimtliche bekannten Werte der Redoxsysteme des Zinns. Dieser Haltepunkt wird der chemischen Aulliisung der Metazinn-Slure zugeschrieben, welche sich im Gebiet der Sauerstoffentwickhmg bildet. Das anodische Verhalten der Sn-Elektrode ist in Stannit- und Stannat-Losungen ahnlich wie bei Liisungen, welche diese Ionen nicht enthalten. INTRODUCTION ALTHOUGH it is well established that the passivity of tin in alkaline solutions is associated with film formation,14 yet the exact conditions leading to the development of such a film are not well defined. Different authors seem also not to be in agreement regarding the chemical composition of the passivating film.5*7 The initial stages leading to passivity are ascribed by Goldschmidt and Eckardtl and by Foerster and TABLE1.
FREE ENERGY INCREASE IN THE FORMATION OF TIN OXIDES AND HYDROXIDE.5
AG”, Kcal Reaction (l)Sn++O,=SnO (2) Sn + 0% = SnO, (3) Sn + OI + HI = Sn(OI-0, (4) Sn + 2 0, + 2 HI = Sn(OH),
per mole oxide or hydroxide
per mole Sn
-61.5 - 124.2 -117.6 -227.5
-61.5 -1242 -117.6 -227.5
per mole 0 -61.5 -62.1
Dolcha as the dissolution of the metal in the form of stannite until a definite concentration is attained. The potential of the electrode changes then suddenly to oxygen evolution values and the dissolved tin oxidizes to stannate. On the other hand, El Wakkad et al8 conclude that the initial stages of passivity in weakly alkaline solutions are the formation of a layer l-2 molecules thick of Sn(OI-I), on the electrode surface followed by its oxidation to Sn(OH), before oxygen evolution. Shah and Davies0 are of the opinion that oxidation to SnO and dissolution as stannite occur simultaneously. SnO is oxidized to SnO, before oxygen evolution. The present study was undertaken to establish the exact conditions leading to the passivity of tin anodes in alkaline solutions. This information was needed in order to understand the electrode behaviour of different copper-tin alloys and to determine the most appropriate alloy composition for corrosion resistance.lO A report on the oxidation of copper in alkaline solutions has been recently published.ll A primary orienting idea about the oxidation of tin can be gained from a consideration of the standard free-energy change, AG”, of the different oxidation reactions. Not only does this consideration reveal thepossibility of these transformations, but it helps in determining the most likely reaction among a number ofprobable ones. In Table 1 the decrease in free energy accompanying the formation of SnO, SnO,, Sn(OH), and Sn(OH), from their respective elements are grouped. The same table gives also the decrease in free energy expressed per one mole of either the metal or oxygen. The figures of Table 1 show that the formation of SnO, is accompanied by the highest free-energy decrease, when this is expressed in terms of either the metal or
Anodic passivity of tin in alkaline solutions
885
oxygen. This shows that SnO, is the most stable oxide when in contact with either the metal or molecular oxygen. l2 The small difference in AG for the formation of SnO and SnO,-expressed in terms of one mole oxygen-shows that SnO is similarly stable when in contact with the metal. This conclusion is supported by the results of a number of investigators who showed that, upon heating tin in air or oxygen, SnO was formed, either alone or associated with Sn02.13-le A consideration of the dehydration of both Sn(OH), and Sn(OH), shows that the reactions Sn(OH), = SnO + H,O (5) and Sn(OH), = SnO, + 2 H,O (6) are favoured by free-energy decrease of -0.59 and -10.08 Kcal/mole respectively. The hydration of the oxides is, therefore, thermodynamically improbable. In Table 2 a number of possible electrode reactions for tin in alkaline solutions are formulated. The free-energy changes associated with these reactions and the equations relating the electrode potential to the activity of the reactants and resultants are similarly grouped in Table 2. The necessary free-energy data for all these calculations are taken from Latimer’s mon0graph.l’ For the reasons presented by Pourbaix et al,* reactions involving the formation of Sn022-, Sn2032- and/or HSnO,- are not considered here. Similarly, since Sn2+ and St++ are not likely to be present in alkaline solutions, reactions incorporating these ions fall outside the scope of our interest. pH/E diagrams for tin in aqueous solutions have been given by Charlot,la Pourbaix et u1,18*20 Sillen= and Hoar.22 The primary anodic oxidation of tin to Sn(OH), or SnO (reactions 7 and 8) involves practically the same free-energy decrease. The difference between the potentials of the systems Sn/SnO and Sn/Sn(OH),, 12 mV, is correct within the experimental accuracy limits of potential measurements. A potentiometric differentiation between the two forms does not seem, therefore, possible. Calculations involving Sn(OH), and SnO, are of particular interest. Thus, the direct oxidation of tin to Sn(OH), (reaction 9) is improbable since (7) and (8) are thermodynamically more favoured. On the other hand, the direct oxidation to SnO, (reaction 10) is apparently associated with a larger free-energy decrease than (7) or (8). This does not imply, however, that (10) would occur instead of (7) or (8). The difference in AG” between reactions (9) and (10) is due to the irreversible dehydration of Sn(OH),. Normally dehydration reactions are very slow kinetically. The work of Gutbier et UP on the dehydration of Sn(OH),, obtained through the hydrolysis of either Sn4+ or Sn032- ions, confirms this conclusion. Also, since in alkaline solutions the concentration of oxide ion, 02-, is vanishingly small, there is but very little chance that SnO, would result directly and not via Sn(OH),. One can conclude, therefore, that the direct oxidation of tin to SnO, is similarly improbable. The same argument applies also for the anodic oxidation of Sn(OH), or SnO, reactions (12) and (14). Sn(OH), would oxidize to Sn(OH), and not to SnO,. The direct dissolution of tin as stannite, reaction (15), or stannate, reactions (16) and (17), is a further possibility for the tin anode in alkaline solutions. At unit activity of HSnO,- and OH-, reaction (15) is both thermodynamically and kinetically improbable relative to (7) and (8). On the other hand, reactions involving the formation of Sn(OH),2- are accompanied by a larger free-energy decrease than those
A. M. SW
886
EL DIN and F. M. ABD EL WAHAB
incorporating SnOse ion; hence, the former ion is thermodynamically more stable. The larger decrease in free energy accompanying the formation of Sn(OH),B- ion, as compared to HSnO,, does not imply that the former ion would preferentially form. TABLE 2. ELECTBODEREACTIONSFOR THE TIN ELECIXODEIN ALKALINE SOLUTIONS AND THEIR EQUILIBRIUM POTWTIALS AG”, Kcal
Electrode reaction and equilibrium potential
(7) (8) (9) (10) (11) (12) (13) (14) (15) (16) (17) (18) (19) (20) (21) (22) (23)
Sn+2OH-=Sn(OH),+2e - 0.059 log [OH-] Sn+2OH-=SnO+H,O+2e -0.932 - 0.059 log [OH-] Sn + 4 OH- = Sn(OH), + 4 e -0.845 - 0.059 log [OH-] Sn+4OH-=SnO,+2H,O+4e -0-945 - 0.059 log [OH-] Sn(OH), + 2 OH- = Sn(OH), + 2 e -0.752 - ft.059 log [OH-] Sn(OH), + 2 OH- = SnO, + 2 H1O + 2 e -0.971 - 0.059 log [OH-] SnO + 2 OH- + HI0 = Sn(OH), + 2 e -0.740 - 0.059 log [OH-] SnO+2OH-=SnO,+H,O+2e -0.958 - O-059 log [OH-] Sn+3OH-=HSnO,-+H,O+2e -0+0!4 + 0.029 log [HSnO,-] - 0.089 log [OH-] Sn + 6 OH- = SII(OH)~*- + 4 e -0.921 + 0.015 log [Sn(OH),*-] - 0.089 log [OH-] Sn+6OH-=SnO,*-+3H,O+4e -0.888 + O-015 log [SnO,B-] - 0.089 log [OH-] Sn(OH), + 4 OH- = Sn(OH),*- + 2 e -0.922 + 0.029 log [Sn(OH),*-] - O-118 log [OH-] SnO + HI0 + 4 OH- = Sn(OH),*- + 2 e -0909 + 0.029 log [Sn(OH),*-] - O-118 log [OH-] Sn(OH), + 4 OH- = SnOl*- + 3 H,O + 2 e -0.857 + 0.029 log [SnO,*-] - 0.118 log [OH-] SnO+4OH-=SnO,*-+2H,O+2e -0.844 + O-029 log [SnOl*-] - 0.118 log [OH-] HSnOp- + HI0 + OH- = Sn(OH), + 2 e -0.764 - O-029 log [HSnO,-] - 0.029 log [OH-] HSnO,- + OH- = SnO, + H1O + 2 e -@983 - 0.029 log [HSnO,-] - 0.029 log [OH-] HSnOl- + HtO + 3 OH- = Sn(OH),*- + 2 e -0.935 i- 0.029 log [Sn(OH),*-] - 0.029 log [HSnO,-] - 0.089 log [OH-] HSnO,- + 3 OH- = SnOlp- + 2 HI0 + 2 e -0.868 + 0.029 log [Snot*-] - 0.029 log [HSnO,-] - @089 log [OH-]
E = -0.920 E = E = E = E = E = E = E = E = E = E = E = E= E = E = E = E =
(24) E = (25) E =
* The value for the free energy of formation
of SnO*“-,
-137.42
-42.41 -43.00 -77.12 - 87.20 -34.71 -45.19 -34.12 -44.20 -4190 -84.93 -81.92” -42.50 -41.93 -39*51+ - 38.92* - 35.22 -45.30 -43.03 -40.02+
Kcal, is cited by Pourbaix
et al.18
The formation of Sn(OH),2- is kinetically more difficult and hence its direct formation is improbable. The larger free-energy decrease signifies, however, the instability of stannite solutions. Actually stannite solutions are known to disproportionate into metallic tin and stannate. Sn(OH), and SnO could be oxidized by either Sn(OH),, reactions (11) and (13), or to Sn(OH),2- (Sn0s2-), reactions (18)-(21). The latter set of reactions involves a larger free-energy decrease and are apparently more probable. Inspection of the potential, EBo, at unit activity of Sn(OH),2- (SnOs2-) and OH- reveals, however, that
Anodic passivity of tin in alkaline solutions
881
they are very near to that for the Sn/Sn(OH), system, especially so for reactions involving Sn(OH),,a- ion. Since the difference in potential between the Sn/Sn(OH), and Sn(OH),/Sn(OH), amounts to only 168 mV, and as the pH-dependence of the potentials of reactions (18)-(21) is greater than that for reactions (11) and (13), it is probable that Sn(0I-Q preferentially forms instead of Sn(OH),% at low pH values. The oxidation of HSnO,- to Sn(OH),, reaction (22), involves practically the same free-energy decrease as that for reaction (11). The pH-dependence of the potential of both reactions is, however, different and a decision as to which of the two reactions is taking place could be drawn on this basis (assuming constant HSnO,- activity). The oxidation of HSnO,- to Sn(OH),2- (Sn0a2-), reactions (24) and (25) is favoured by a large decrease in AG”. At unit activity of HSnO,, Sn(OH),s- or Sn0s2the potential of the tin electrode varies by O-089 V/pH unit. How far these thermodynamical predictions are correct was of interest to establish experimentally. EXPERIMENTAL TECHNIQUE The electrolytic cell used was as previously described .I1 The tin electrode was prepared by casting analytically-pure tin granules (Merck, Germany) in the form of a rod 8.5 mm in diameter. The total exposed area was 4.17 cm4. Before use the electrode was polished with the finest grade emery paper and then washed with running distilled water. Experiments were conducted mainly in carbonate-free 2 N @H co 14), 0.1 N @H 12.9) and 0.01 N (pH 12.1) NaOH solutions, as well as in 0.1 M Na,COI solutions (pH 11.1). Each experiment was carried out with a freshly-deaerated solution and with a newly-polished electrode. Polarization curves were recorded on a Leeds and Northrup type-G pen recording potentiometer. The technique for tracing cyclic polarization curves was essentially as described before.” Potentials were measured relative to a saturated calomel electrode and then corrected to the normal hydrogen scale. We have studied also the effect of addition of stannite and stannate on the polarization curves of tin. This type of experiment was carried out mainly in O-1 N NaOH solution. Enough solid Sn(OH), or Sn(OH), was shaken with the hydroxide solution for a period of 48 h to ensure saturation. The solutions were then centrifuged and the clear liquids were diluted with O-1 N NaOH to give solutions of definite starmite or stannate contents. The tin concentration is here expressed in percentage relation to the original saturated solution. All measurements were conducted at 25°C. RESULTS
AND
DISCUSSION
The anodic behaviour of tin in dilute alkaline solutions Figure 1 represents the potential variation of the tin electrode when being subjected to cyclic anodic and cathodic polarization in 0.01 N NaOH solution. The polarizing current in this set of curves was 2 mA/electrode. As can. be seen from the anodic half-cycle curves (curves A), there occurred at first a rapid and almost linear change of potential (part I) due to both the decay of hydrogen overpotential and the subsequent charging of the anodic double layer. Following this process, the potential of the tin electrode changed more slowly, giving rise to two distinct arrests, (part II), before finally passing to values characteristic for oxygen evolution (part III). The starting potentials, E,,, of the two steps of part (II) are -0.78 and -0.62 V respectively. These values compare very satisfactorily with those for the two anodic reactions (7) and (1 l), Table 2, at the corresponding pH value. As will be shown later when discussing the behaviour of the tin electrode in more concentrated alkaline solutions, the potential of these two arrests changed by -59 mV for every pH-unit increase. This strongly suggests that during the anodic oxidation of tin, Sn(OH), (or SnO) and Sn(OH), are formed on the metal surface before oxygen evolution. Dissolution as stannite and stannatelBe is not likely to occur since, for a definite HSnO, or Sn(OH),Z- concentration, the potential should change by -89 mV/pH
A. M. SHAMSEL DIN
888
and F. M.
ALID
EL WAHALI
unit. On the other hand, to assume that the observed variation of -59 mV/pH was due to a corresponding decrease in the concentration of either ion implies that this decrease should be constant for every pH-unit change. Since stannite and stannate are freely soluble in alkaline solutions, such a requisite does not seem reasonable. Inspection of the anodic curves reveals also that the second oxidation step consumed larger quantities of electricity than is simply required for the transformation of Sn(OH), into Sn(OH),. This is apparently due to the higher solubility of Sn(OH), in alkaline solutions. Similarly, upon repeated cathodic reduction and anodic oxidation, the
I.E50 -
I.450 -
>
I.050 -
2
0,650 -
2 0 f 0.250 i ::
0.150-
E Y w-o.550
-
-0.950
-
-1.350
-
I
5 (A)
0 *
15
30
Time,
45
60
75
.90
min
FIG. 1. Cyclic polarization curves for the tin electrode in O-01 N NaOH solution. A, anodic half-cycles, C, cathodic half-cycles. Polarizing current = 2 mA/electrode.
quantity of electricity necessary to bring the electrode to oxygen evolution potential increased with every new anodic half-cycle. The increase was mainly confined to the second step and was probably due to a corresponding augmentation of the true surface area of the electrode as a result of the deposition of the metal during the previous cathodic half-cycle. Cathodic polarization curves (curves C, Fig. 1) were obtained by reversing the polarizing current when the electrode was at oxygen-evolution potentials. The potential first dropped to negative values, but levelled up quickly, giving rise to a sharp peak. The potential then changed more slowly before finally dropping to values characteristic for the evolution of hydrogen. The difference in potential between the apex of the peak and the top of the flat maximum increased with the cycle number. The kink in the cathodic curves was observed only if the previous anodic polarization was allowed to reach the oxygen-evolution potential. If the anodic polarization was interrupted at the Sn(OH),/Sn(OH), potential before the current was reversed, the cathodic curves did not show the kink. This suggests that the occurrence of the kink in the cathodic curves is related to a certain transformation taking place
Anodic passivity of tin
in alkaline solutions
889
at the oxygen-evolution potential. As will be shown in the discussion of the anodicdecay curves, oxygen evolution occurs most probably on a film of metastannic acid (H,SnO, or Sn02.H,0).2*7 The features of the cathodic polarization curves could be explained by assuming that at the instant of reversal of the current, the reduction of the actual passivating film, which is only few monolayers thick,**n takes place and the potential of the electrode tends to hydrogen-evolution values. The metastannic acid film then comes in direct contact with the metal. Since the Sn/Sn(IV) potential in alkaline solutions is more positive than the corresponding Sn/Sn(II), the potential of the electrode rises again. Because of the instability of the Sn/Sn(IV) couple, this process is only transient and the reduction proceeds both chemically and electrochemically. Reduction of some of the dissolved tin is expected to occur simultaneously. The cathodic recovery,ll ie the percentage ratio of the quantities of electricity consumed in the cathodic and anodic half-cycles, was not constant and did not exceed 65 per cent. Apparently some of the tin in solution diffused away from the electrode and escaped subsequent reduction. In fact, some small tin dendrites were always found on the platinum electrode (cathode) at the end of the experiment. The behaviour of the tin electrode in O-1 M Na,CO, solutions was similar to that described above. The potentials of the two oxidation arrests were, however, less defined and the electrode required less quantities of electricity to passivate. In some experiments the tin electrode did not reach oxygen evolution and the potential became irregular. Most probably some basic carbonate formed on the electrode surface. The behaviour of the tin anode in concentrated
alkaline solutions
With currents equal to or less than 5 mA/electrode, the tin anode dissolved continually in 0.1 N NaOH solution at the Sn/Sn(OH), potential (-0.85 V). The time/potential curve run, therefore, parallel to the time axis (curve a, Fig. 2). When the current was raised to 7 mA/electrode, curve b, Fig. 2, dissolution as stannite took place for a comparatively long period and then the potential started to oscillate within a range of 120 mV. The frequency of oscillation was higher at early than at advanced stages. Upon further increase of the current, the dissolution at the Sn/ Sn(OH), potential was confined to a short interval before oscillations started (curve c). At 12 mA/electrode the polarization curve showed that after an initial limited segment at the Sn/Sn(OH), potential the tin electrode dissolved mainly at the Sn(OH),/ Sn(OH), potential. Oscillations in potential were observed only at advanced periods. At this stage the amplitude of oscillation increased to ca 180 mV and the frequency decreased to a minimum. At 14 mA/electrode, the electrode dissolved mainly at the Sn(OH),/Sn(OH), potential and did not-within the time of our experiments-change to oxygen evolution. The above-described phenomena were equally observed with increasing as well as with decreasing polarizing currents, provided that-in the latter case-the electrode did not change to oxygen-evolution potentials. When oxygen was allowed to evolve on the electrode and then the current decreased in steps, the electrode did not return directly to the active state. The lowering of the current affected only the oxygen overpotential on the passive anode. This fact is generally made use of in the electrodeposition of tin from alkaline solutions (see later) where the anode is first filmed at high currents to ensure its subsequent dissolution in the quadrivalent state.2 With currents higher than 14 mA/electrode permanent passivity of the tin electrode in 0.1 N NaOH solution readily occurred.
A. M. SHAMSEL DIN and F. M. fiBD EL
0
I5
30
WAHAB
45
60
min
Time,
FIG.2. Anodic polarization curves for the tin electrode in O-1 N NaOH solution. a, 5 mA; b, 7 mA; c, 10 mA; d, 12 mA; e, 14 mA/electrode.
In 2 N NaOH solution the tin electrode behaved in a similar manner as in O-1 N solution. Here, however, much higher currents were needed to suppress the dissolution of Sn(OH), and Sn(dH),, before the electrode could be made to evolve oxygen. In Fig. 3 a typical experiment for the passivation of tin in 2 N NaOH solution is shown. The current was held constant until either a constant potential (dissolution region) or a constant frequency (oscillation region) was attained. As seen from Fig. 3, oxygen would not evolve on the electrode unless a current equal to or greater than 350 mA was imposed on the electrode.
8 j 4.350 .y 0
15
30 Time.
FIG. 3. Anodic polarization
45
60
75
90
min.
curve for the tin electrod& in 2 N NaOH solution using different currents.
Anodic passivity of tin in alkaline solutions
891
Oscillations in time/potential curves under galvanostatic conditions are frequently reporuzdsssab It is generally agreed upon that they denote the alternating formation and dissolution of a film on the surface of the electrode. The behaviour of the tin electrode in concentrated alkali solutions (>O*l N) seems to support the same conclusion. Below 5 mA/electrode in O-1 N NaOH and 250 mA/electrode in 2 N NaOH solutions, the rate of dissolution of the anodically-formed Sn(OH), surpasses or equals the rate of its formation. The potential remains, therefore, constant at the Sn/Sn(OH), value at the corresponding pH. As the current is increased, the discharge of OH- supersedes the dissolution of Sn(OH),, and Sn(OH), forms at its thermodynamical potential. The electrode is thus in a state ofprimary passivity. Due to its acid character, Sn(OH), dissolves freely in the alkali solution to yield stannate. This would expose the Sn(OH), film and hence the potential drops again to the Sn/Sn(OH), value. The process of formation and dissolution of Sn(OH), would repeat itself so long as a permanent layer of stannic hydroxide cannot persist on the electrode surface. As the current is increased, the ability of the hydroxide to remain intact on the electrode surface increases and the potential remains constant at the Sn(OH),/Sn(OH), value. The rate of formation of Sn(OH), then equals the rate of its dissolution. In O-1 N NaOH solution this seems to occur at 14 mA/electrode. Above this current, permanentpussivation of the tin electrode can take place. The curves of Fig. 2 clearly show that before oxygen evolution, both Sn(OH), and Sn(OH), formed on the surface of the electrode. Although difficulty was experienced in completely reproducing the quantities of electricity necessary to bring the electrode to oxygen evolution potentials, the experimental results seem to fit a relation of the form (i - i&P = const.,
(22)
where i is the applied current and T is the time from the start of polarization till oxygen evolution. i,, represents the current below which permanent passivation could not be achieved, here taken as 14 mA/electrode in O-1 N NaOH. The value of n in the above relation approaches O-5. Relation (22) is common for the anodic passivation of metals.a7*28 Anodic decay curves were recorded by switching off the polarizing current when the electrode was at oxygen-evolution potentials. The curves obtained, Fig. 4, show two definite arrests before the potential reached the equilibrium value of the Sn/ Sn(OH), system. The second of these two arrests was pH-dependent and corresponded to the transformation of Sn(OH), to Sn(OH),. The characteristics of the fist, decay arrest are of particular interest. It is pHindependent and occurs at potentials positive to any of the known redox-systems of tin. From experiments in which the tin electrode was anodically polarized to increasing pre-determined potentials, it was ascertained that the arrest would not appear unless the electrode reached oxygen-evolution potentials. In solutions of low pH, self activation was a very slow process. Cathodic pulses of very short duration (cu 1 s) had to be applied to assist the electrode to reach the Sn/Sn(OH), potential (curve c, Fig. 4). The length of the arrest was not proportional to the time of the preceding oxygen evolution. With too long times at oxygen evolution potentials, or with too high passivating currents, the electrode did not return easily to the Sn/Sn(OH), potential and cathodic pulses had to be applied. The start potential, mid-arrest potential and/or end potential of this arrest were not definite, but depended primarily
892
A. M. SHAMS EL DIN and F. M. ABD EL WAHAB
upon how long the electrode evolved oxygen before the decay experiment was started. The plot of the electrode potential USthe logarithm of time (s) during the first decay arrest was generally non-linear. In those few instances where a straight line relationship was fulfilled, the slope of the lines was exceptionally high, amounting to ca 8RT/F. This rules out the probability of a kinetic process governing the process of self activation. As pointed out above, the first decay arrest is pH-independent, does not correspond to any of the known redox-potentials of tin, and is observed only after the electrode
Time,
min
FIG. 4. Anodic decay curves for the tin electrode
a, in 0.1 N NaOH;
b, in 2 N NaOH;
in alkaline solutions. c, in 0.01 N NaOH with cathodic pulses.
has evolved oxygen. This limits its cause to a chemical reaction not involving electron transfer. According to FSrster2 and Bianchi,’ oxygen evolution on tin anodes in alkaline solutions takes place on a layer of metastannic acid, H,SnO,. This means that Sn(OH), undergoes a process of dehydration, which is thermodynamically favoured. The first decay arrest is most probably the result of the chemical dissolution of H,SnO,. The rate of dissolution of the acid would depend upon the alkalinity of the solution in the manner observed experimentally. The recent work of Shah and Davies9 is of particular interest as it was conducted under apparently similar conditions. Shah and Davies reported a single oxidation arrest for tin in weakly alkaline solutions. This was ascribed to the simultaneous oxidation to SnO and HSn02- in equivalent quantities according to 2Sn+SOH-=SnO+HSn02-+2H20+4e.
(23) From thermodynamical data, the free energy change of reaction (23), -84.9 Kcal, was computed, and correspondingly the equilibrium potential in O-1 M sodium borate solution (pH 9.28) was estimated as -0.85 V (relative to Ag/AgCl reference electrode). Because of the close proximity of this value to that of the recorded step, it was concluded that reaction (23) was governing the electrode behaviour. Reaction (23) is, however, kinetically very improbable as a single step. Elementary electrochemical reactions for oxide formation and stannite dissolution (reactions (8)
893
Anodic passivity of tin in alkaline solutions
and (13, Table 2) must, therefore, be considered. Also, the fact that the calculated potential was close to that of the experimentally observed step does not imply that (23) is actually taking place. Since reactions (8) and (15) have comparable AGo values, their summation would naturally lead to a corresponding doubling of the overall free-energy change and, hence, the EBo value would come again to that for the separate reactions. Further, Shah and Davies, in their calculation of the equilibrium potential of reaction (23) at pH 9.28, assumed that the potential would change by $59 mV for
I
.
0
15
30 Time,
45
60
t-75
min
FIG. 5. Anodic polarization curves for the tin electrode in 0.1 N NaOH containing
20 g/l KCl. 1, 10 mA/electrode; 2, 20 ma/electrode; 3, cathodic polarization curve with 20 mA/electrode. every pH unit decrease. This is, however, not the case as far as reaction (15) is concerned; the change is +89 mV. The absence of a definite oxidation arrest corresponding to the formation of Sn(OH), in Shah and Davies’ polarization curves is also strange. This arrest has been recorded in our previous study as well as in the present one. It is even observed with tin alloys containing copper10*2g and zinc.30 We are unable to give a satisfactory explanation for this. It was thought at first that the absence of this arrest in Shah and Davies’ curves was the result of the presence of Cl- ion in solution (added so that the Ag/AgCl electrode would function). This ion is known to interfere with passivation processes.31*3a We conducted, therefore, some passivation experiments in O-1N NaOH solution containing KC1 (20 g/l). The results are shown in Fig. 5. At low currents a small arrest corresponding to the formation of Sn(OH), was recorded. The potential then changed to the Sn(OH),/Sn(OH), value where it remained constant, curve 1, Fig. 5. When the current was raised, curve 2, the instantaneous formation of Sn(OH), and Sn(OH), was followed by the development of a dense white film (probably tin oxy-chloride) on the electrode surface. This film did not adhere strongly to the metal
894
A. M. SHAMSEL DIN and F. M. ABD EL WAHAB
surface and the potential of the electrode became irregular. Interesting is the fact that upon reversing of the current two kinks were recorded in the cathodic polarization curve (curve 3, Fig. 5). This substantiates the above presented idea that the occurrence of a kink is not due to the presence of a particular oxidation state, but rather to the variation in the type of the reduced species. In the above experiment the first kink is due to the reduction of the actual passivating layer. The potential rises to respond to the reduction of metastannic acid coming into contact .with the free tin surface. When this process comes to an end, the potential tends again towards negative values. However, since the electrode is still covered with a layer of oxy-chloride, the potential rises once more until this film is completely reduced. The behaviour of the tin anode in alkaline solutions containing stannite and stannate As a complement to the work on the passivity of tin in pure alkaline solutions, we extended the study to cover its behaviour in alkaline solutions containing stannite and stannate. This point is of practical interest since the electrodeposition of the metal is carried out mainly from such solutions. 84 The behaviour of the tin anode in tin plating solutions has been previously examined, 4*5but-the conclusions are not in agreement. Our study was the recording of galvanostatic time/potential curves in alkaline solutions containing different concentrations of either sodium stannite or stannate (see Experimental Technique). As shown in the previous section, the polarization curve of the tin anode in 0.1 N NaOH with a current of 10 mA/electrode exhibited a short segment at the Sn/Sn(OH), potential before oscillations started. In the presence of dissolved Sn(OH),, the general polarization behaviour of the electrode varied to a lesser or greater extent depending upon the concentration of the added tin hydroxide. With the same polarizing current, and in a solution 10 per cent saturated with Sn(OH),, the part of the polarization curve representing the dissolution of the anodically formed Sn(OH), was greatly reduced, apparently due to the presence of part of the anodic products at the metal/solution interphase. Following this, free dissolution at the Sn(OH),/Sn(OH), potential took place. Oscillations occurred only at advanced times. These were less frequent and the duration of each cycle was longer than in absence of dissolved stannite, curve a, Fig. 6. The sharp peaks observed at the start of each cycle represent the tendency of the electrode to attain a state of permanent passivity that is quickly levelled down by the dissolution of Sn(OH),. In solutions 20 per cent saturated with Sn(OH),, the general behaviour of the tin electrode was as described above save for the fact that the oscillations were still less frequent. In 30-50 per cent saturated solutions, the electrode continually dissolved at the Sn(OH),/ Sn(OH), potential. Oscillations were seldom observed. Within the time limit of our experiments, the electrode did not change to oxygen-evolution potentials. In 60 per cent solutions, and with a current of 10 mA/electrode, the tin electrode was readily passivated within a relatively short period, curve c, Fig. 6. It will be remembered from the results of the previous section that the same sequence of processes was obtained through the application of increasing currents. The presence in solution of the anode products is equivalent, therefore, to the application of higher currents. This, together with the passivation law given above (relation 22), suggests that the process of passivation is governed by the rate of nucleus formation of the hydroxides on the electrode surface.
895
Anodic passivity of tin in alkaline solutions
The passivation of the tin anode in solutions of higher tin content and with a current of 10 n-A/electrode was almost instantaneous. The polarizing current had to be lowered in order to ascertain the formation of both Sn(OH), and Sn(OH),. The behaviour of the tin anode in solutions containing Sn(OH), resembled that in Sn(OH), solutions. Passivity occurred, however, more readily, apparently due to the high tin content of the solutions. Oscillations in potential were observed only in solutions of 10 per cent (or less) saturation. In solutions of 20-40 per cent saturation, continual dissolution at the Sn(OH),/Sn(OH), potential took place. Passivity with a I
I , I
I I
0.25C
I
A-
I
(b)
______d’
J
c
_(a)
v
I I I I
CI g
s w
-1.350
&
-1.750
A-
:(c) I t : 0
15
45
30 Time,
60
75
min
FIG. 6. Anodic polarization curves for the tin electrode in O-1 N NaOH solution of: a, 10 per cent saturation with Sn(OH), b, 40 per cent saturation with Sn(OH), c, 60 per cent saturation with Sn(OH), Polarizing current, 10 mA/electrode.
current of 10 mA/electrode was easily attained in solutions 50 per cent (or more) saturated with Sn(OH),. With increasing tin content, the quantities of electricity associated with the Sn(OH), and Sn(OH), arrests decreased and lower currents were always applied to prove the formation of both hydroxides. In this respect our results are not in agreement with those of H&A,* who reported that the transition current for the active-passive transformation increased with the increase of the Sn(IV) content of the bath. The results of the present investigation allow the following picture about the passivity of the tin electrode in alkaline solutions to be drawn. The primary oxidation of tin in alkaline solution results in the formation of Sn(OH),. Depending upon the magnitude of the polarizing current, and upon the alkalinity of the solution, the formed hydroxide dissolves to yield stannite. Dissolution continues until the saturation of the metal/solution interphase with stannite leads to the crystallization of Sn(OH), on the electrode surface. The subsequent discharge of OH- ion results in the oxidation of Sn(OH), to Sn(OH),. The electrode is then in a state of primary passivity.
896
A. M. SHAMSEL DIN and F. M. f%FID EL WAHAB
Sn(OH),-by virtue of its higher acidity4issolves as stannate. When the rate of dissolution exceeds that of formation, oscillations in potential/time curves are observed. Once the rate of OH- discharge exceeds that of Sn(OH), dissolution, permanent passivation sets in and oxygen can evolve on the anode. At oxygen-evolution potential Sn(OH), partly dehydrates to metastannic acid. The pre-saturation of the solution with the anodic products facilitates the attainment of both the primary and permanent states of passivity, but does not alter in any way the course of oxidation of the metal. REFERENCES 1. H. GOLDS-T and E. EKARDT,Z. phys. Chem. 56,385 (1906). 2. F. FOERSTER and M. DOLCH,Z. Eiektrochem. 16,599 (1909). 3. E. NE~BERY,J. Chem. Sot. iO66 (1916). . 4. G. -EL and A. GRAVEL,Z. Elektrochem. 41, 314 (1935). 5. R. KERR,J. Chem. Sot. Znd. 57,405 (1938); Chem. Abstr. 1601 (1939). 6. G. T. BACHVAUIW,J. Appl. Chem. Moscow, 14,469 (1941); Chem. Abstr. 2221 (1942); G. T. BACH~AL~Wand P. S. Trrov, Khim. Referat Zh. 4,85 (1941); Chem. Abstr. 5659 (1943). 7. G. BIANCHI,Chim. Znd. 29,295 (1947). 8. S. E. S. EL WAKKAD, A. M. Sm EL DIN and J. A. EL SAYED,J. Chem. Sot. 3103 (1954). 9. S. N. SHAH and D. E. DAVY Electrochim. Acta 8,663,703 (1963). 10. A. M. SW EL DIN and F. M. ABD EL WAHAB, unpublished results. 11. A. M. SHAMSEL DIN and F. M. ABD EL WAHAB, Electrochim. Acta, 9, 113 (1964). 12. A. KUTZELNIGG,Z. anorg. Chem. 202,418 (1931). 13. N. BOUNDand D. A. RICHARDS,Proc. Phys. Sot. SlB, 256 (1939). 14. R. K. HART, Proc. Phys. Sot. 65B, 955 (1952). 15. S. C. BRIITONand K.“BRIGHT,MetaNurgia, Manchr 56, 163 (1957). 16. W. E. Boczs. R. H. KACHIKand G. E. PELWSSER, J. Electrochem. Sot. 108,6,124 (1961). 17. W. M. LA&R, Oxidation Potentials, pp. 39, 136; Prentice Hall, New York (1953): ’ 18. E. DELTOMBE,N. DE ZOUBOV and M. POURBAIX,Proc. 7th Meeting CZTCE, Lindau 1955, p. 216. Butterworths, London (1957). 19. G. CHARUIT, Theorie et Mithodes Nouvelles d’dnalyse Qualitative, 3rd Ed., p. 222. Masson, Paris (1949). 20. P. DELAHAY, M. POURBAMand P. VAN RYSSELBERGHE, Proc. 3rd Meeting CZTCE, Milan 1952, p. 15. Manfredi, Milan (1952). 21. L. G. SILLEN,J. Chem. Educ. 29,600 (1952). 22. T. P. HOAR, Trans. Faraday Sot. 33,1152 (1937). 23. A. GUTBIER,G. F. H-0 and H. D~BLING, Ber. 59B, 1232 (1926). 24. S. FIELDand A. D. WEILL, Electroplating, p. 426. Pitman, London (1961). 25. E. S. -ES, J. Chem. Sot. 1533,2580,2878 (1926); 1028 (1929). 26. R. no-1 and G. SERRAVALLE, Z. Elektrochem. 62,759 (1958). 27. R. LANDS~ERG and M. HOLLNAGEL, Z. Elektrochem. 58,681 (1954), 60,1098 (1956); R. LANDSBERG,Z. phys. Chem. 206, 291 (1957); R. LANDSBERG and H. BARTELT,Z. Elektrochem. 61, 1162 (1957). 28. N. A. HAMPTON,J. Electrochem. Sot. 110,95 (1963). 29. S. E. S. EL WAKKAD, T. M. SALEM,A. M. SW EL DIN and Z. HANAFI, J. Chem. Sot. 2857 (1956). 30. S. E. S. EL WAKKAD, A. M. SHAMSEL DIN and H. KOTB,J. Electrochem. Sot. 105,47 (1958). 31. B. ~ANOV, R. Bmm and A. FRUMKIN,Disc. Faraday Sot. 1,259 (1947). 32. H. K. E~MI and A. A. Mou~~A, J. Chem. Sot. 2027 (1958).
Acta. 1964, Vol. 9. pp. 883 to 8%. Pcrgamon FTCSSLtd. Printed in Northern Irehttd
ON THE ANODIC PASSIVITY OF TIN IN ALKALINE SOLUTIONS* A. Laboratory
M.
SHAMS EL DIN
and F. M. ABD EL WAHAB
of Electrochemistry and Corrosion, National Research Centre, Dokki, Cairo, U.A.R.
Abstract-The anodic oxidation of tin was studied galvanostatically in NaOH solutions of different concentrations. Primary passivity is attained in all solutions when the metal is covered with a film of Sn(OH), or SnO. Permanent passivity sets in when Sn(OH), forms as a continuous layer on the electrode surface. In concentrated alkali solutions larger quantities of electricity are charged to the electrode to compensate for the chemical dissolution of the hydroxides. As a result of the competition between the anodic formation and chemical dissolution of Sn(OI-I),, the potential of the tin electrode oscillates between the Sn/Sn(OH), and Sn(OH),/Sn(OH), values. For any one solution the oscillation phenomenon depends primarily upon the polarizing current. Permanent passivity is governed by the relation (i - io)P = const. where i. is the current below which no passivation takes place, T the time elapsed till oxygen evolution and n is a constant. Interruption of the polarizing current when the electrode is at oxygen-evolution potential gives rise to a pH-independent arrest at potentials more positive than any of the known redox values for tin. This arrest is ascribed to the chemical dissolution of metastannic acid formed at oxygen-evolution potentials. The anodic behaviour of the tin electrode in solutions containing stannite and stannate is similar to that in solution8 devoid of these ions. An increase in the tin content of the solution is equivalent to the elevation of the polarizing current ; passivity sets in more readily in the concentrated solutions. R&suru&L’oxydation anodique de Sn est 6tudiCe galvanostatiquement en NaOH de concentration vari6e. Une pas&it6 primaire apparalt dans toutes les solutions, quand le m&al se recouvre d’un film de Sn(OH), ou SnO. Une passivite permanente est due a une couche continue de Sn(OH),. Dam les solutions alcalines concentrees, la charge de l’&ctrode implique des quantit6s d’&ctricit6 accrues, pour compenser la dissolution chimique des hydrox des. Le potentiel de l%lectrode d’&ain oscille db lors entre les valeurs propres a Sn/Sn(OH), et Sn rOH),/Sn(OH),, ce qui traduit une sorte de comp6tition entre la formation anodique et la dissolution chimique de Sn(OI-I),. Ce ph6nom&ne d’oscillation d6pend surtout du courant polarisant. La pas&it6 pcrmanente est gouvem6e par la relation (i - &,) 71 = cte, i. &ant le courant au dessous duquel il n’y a pas passivation, 7 le temps pr&dant l’bvolution d’oxygene, n une constante. Si l’on interrompt le courant polarisant quand le potentiel d’6volution d’Ot est atteint, le potentiel d’arr&, indtpendant du pH, est plus positif que toute valeur r&lox deja connue pour p∈ parret est attribut B la dissolution de l’acide m&a&annique engendn5 aux potentiels d%volution d’0 *. Le comportement anodique de l’&ctrode Sn dans des solutions de starmites et stannates est analogue. zuplemmenhpsslpg_Die anodische Oxydation von Sn in NaOH-Liisungen wurde galvanostatisch untersucht. Prim&e Pas&it&t setzt ein, wenn sich Sn(OH), oder SnO bildet, permanente Pas&it&t, wenn eine zusammenh&ngende Deckschicht von Sn(0I-I). entsteht. In konzentrierten NaOHLiisungen werden grossere Blektrizititsmengen beniitigt, urn die chemische Aufliisung deer Hydroxyde zu kompensieren. Infolge der konkurrierenden Bildung tmd Au&L%mg von Sn(OH), oszilliert das Potential zwischen den Sn/Sn(OH), und Sn(OH),/Sn(OH),-Werten. Die permanente Passivitit gehorcht der Gleichung (i - io)rn = Konst. * Manuscript
received 9 November
1963.
A. M. SIUMS EL DIN and F. M. ABD
884
EL WAHAB
wobei n eine Konstante ist, 7 die Zeit bis zum Einsetzen der O,-Entwicklung und i. der Strom, unterhalb dem keine Passivierung eintritt. Die Abschaltkurven we&n einen pH-unabhlngigen Haltepunkt auf, der bei einem positiveren Potential liegt als siimtliche bekannten Werte der Redoxsysteme des Zinns. Dieser Haltepunkt wird der chemischen Aulliisung der Metazinn-Slure zugeschrieben, welche sich im Gebiet der Sauerstoffentwickhmg bildet. Das anodische Verhalten der Sn-Elektrode ist in Stannit- und Stannat-Losungen ahnlich wie bei Liisungen, welche diese Ionen nicht enthalten. INTRODUCTION ALTHOUGH it is well established that the passivity of tin in alkaline solutions is associated with film formation,14 yet the exact conditions leading to the development of such a film are not well defined. Different authors seem also not to be in agreement regarding the chemical composition of the passivating film.5*7 The initial stages leading to passivity are ascribed by Goldschmidt and Eckardtl and by Foerster and TABLE1.
FREE ENERGY INCREASE IN THE FORMATION OF TIN OXIDES AND HYDROXIDE.5
AG”, Kcal Reaction (l)Sn++O,=SnO (2) Sn + 0% = SnO, (3) Sn + OI + HI = Sn(OI-0, (4) Sn + 2 0, + 2 HI = Sn(OH),
per mole oxide or hydroxide
per mole Sn
-61.5 - 124.2 -117.6 -227.5
-61.5 -1242 -117.6 -227.5
per mole 0 -61.5 -62.1
Dolcha as the dissolution of the metal in the form of stannite until a definite concentration is attained. The potential of the electrode changes then suddenly to oxygen evolution values and the dissolved tin oxidizes to stannate. On the other hand, El Wakkad et al8 conclude that the initial stages of passivity in weakly alkaline solutions are the formation of a layer l-2 molecules thick of Sn(OI-I), on the electrode surface followed by its oxidation to Sn(OH), before oxygen evolution. Shah and Davies0 are of the opinion that oxidation to SnO and dissolution as stannite occur simultaneously. SnO is oxidized to SnO, before oxygen evolution. The present study was undertaken to establish the exact conditions leading to the passivity of tin anodes in alkaline solutions. This information was needed in order to understand the electrode behaviour of different copper-tin alloys and to determine the most appropriate alloy composition for corrosion resistance.lO A report on the oxidation of copper in alkaline solutions has been recently published.ll A primary orienting idea about the oxidation of tin can be gained from a consideration of the standard free-energy change, AG”, of the different oxidation reactions. Not only does this consideration reveal thepossibility of these transformations, but it helps in determining the most likely reaction among a number ofprobable ones. In Table 1 the decrease in free energy accompanying the formation of SnO, SnO,, Sn(OH), and Sn(OH), from their respective elements are grouped. The same table gives also the decrease in free energy expressed per one mole of either the metal or oxygen. The figures of Table 1 show that the formation of SnO, is accompanied by the highest free-energy decrease, when this is expressed in terms of either the metal or
Anodic passivity of tin in alkaline solutions
885
oxygen. This shows that SnO, is the most stable oxide when in contact with either the metal or molecular oxygen. l2 The small difference in AG for the formation of SnO and SnO,-expressed in terms of one mole oxygen-shows that SnO is similarly stable when in contact with the metal. This conclusion is supported by the results of a number of investigators who showed that, upon heating tin in air or oxygen, SnO was formed, either alone or associated with Sn02.13-le A consideration of the dehydration of both Sn(OH), and Sn(OH), shows that the reactions Sn(OH), = SnO + H,O (5) and Sn(OH), = SnO, + 2 H,O (6) are favoured by free-energy decrease of -0.59 and -10.08 Kcal/mole respectively. The hydration of the oxides is, therefore, thermodynamically improbable. In Table 2 a number of possible electrode reactions for tin in alkaline solutions are formulated. The free-energy changes associated with these reactions and the equations relating the electrode potential to the activity of the reactants and resultants are similarly grouped in Table 2. The necessary free-energy data for all these calculations are taken from Latimer’s mon0graph.l’ For the reasons presented by Pourbaix et al,* reactions involving the formation of Sn022-, Sn2032- and/or HSnO,- are not considered here. Similarly, since Sn2+ and St++ are not likely to be present in alkaline solutions, reactions incorporating these ions fall outside the scope of our interest. pH/E diagrams for tin in aqueous solutions have been given by Charlot,la Pourbaix et u1,18*20 Sillen= and Hoar.22 The primary anodic oxidation of tin to Sn(OH), or SnO (reactions 7 and 8) involves practically the same free-energy decrease. The difference between the potentials of the systems Sn/SnO and Sn/Sn(OH),, 12 mV, is correct within the experimental accuracy limits of potential measurements. A potentiometric differentiation between the two forms does not seem, therefore, possible. Calculations involving Sn(OH), and SnO, are of particular interest. Thus, the direct oxidation of tin to Sn(OH), (reaction 9) is improbable since (7) and (8) are thermodynamically more favoured. On the other hand, the direct oxidation to SnO, (reaction 10) is apparently associated with a larger free-energy decrease than (7) or (8). This does not imply, however, that (10) would occur instead of (7) or (8). The difference in AG” between reactions (9) and (10) is due to the irreversible dehydration of Sn(OH),. Normally dehydration reactions are very slow kinetically. The work of Gutbier et UP on the dehydration of Sn(OH),, obtained through the hydrolysis of either Sn4+ or Sn032- ions, confirms this conclusion. Also, since in alkaline solutions the concentration of oxide ion, 02-, is vanishingly small, there is but very little chance that SnO, would result directly and not via Sn(OH),. One can conclude, therefore, that the direct oxidation of tin to SnO, is similarly improbable. The same argument applies also for the anodic oxidation of Sn(OH), or SnO, reactions (12) and (14). Sn(OH), would oxidize to Sn(OH), and not to SnO,. The direct dissolution of tin as stannite, reaction (15), or stannate, reactions (16) and (17), is a further possibility for the tin anode in alkaline solutions. At unit activity of HSnO,- and OH-, reaction (15) is both thermodynamically and kinetically improbable relative to (7) and (8). On the other hand, reactions involving the formation of Sn(OH),2- are accompanied by a larger free-energy decrease than those
A. M. SW
886
EL DIN and F. M. ABD EL WAHAB
incorporating SnOse ion; hence, the former ion is thermodynamically more stable. The larger decrease in free energy accompanying the formation of Sn(OH),B- ion, as compared to HSnO,, does not imply that the former ion would preferentially form. TABLE 2. ELECTBODEREACTIONSFOR THE TIN ELECIXODEIN ALKALINE SOLUTIONS AND THEIR EQUILIBRIUM POTWTIALS AG”, Kcal
Electrode reaction and equilibrium potential
(7) (8) (9) (10) (11) (12) (13) (14) (15) (16) (17) (18) (19) (20) (21) (22) (23)
Sn+2OH-=Sn(OH),+2e - 0.059 log [OH-] Sn+2OH-=SnO+H,O+2e -0.932 - 0.059 log [OH-] Sn + 4 OH- = Sn(OH), + 4 e -0.845 - 0.059 log [OH-] Sn+4OH-=SnO,+2H,O+4e -0-945 - 0.059 log [OH-] Sn(OH), + 2 OH- = Sn(OH), + 2 e -0.752 - ft.059 log [OH-] Sn(OH), + 2 OH- = SnO, + 2 H1O + 2 e -0.971 - 0.059 log [OH-] SnO + 2 OH- + HI0 = Sn(OH), + 2 e -0.740 - 0.059 log [OH-] SnO+2OH-=SnO,+H,O+2e -0.958 - O-059 log [OH-] Sn+3OH-=HSnO,-+H,O+2e -0+0!4 + 0.029 log [HSnO,-] - 0.089 log [OH-] Sn + 6 OH- = SII(OH)~*- + 4 e -0.921 + 0.015 log [Sn(OH),*-] - 0.089 log [OH-] Sn+6OH-=SnO,*-+3H,O+4e -0.888 + O-015 log [SnO,B-] - 0.089 log [OH-] Sn(OH), + 4 OH- = Sn(OH),*- + 2 e -0.922 + 0.029 log [Sn(OH),*-] - O-118 log [OH-] SnO + HI0 + 4 OH- = Sn(OH),*- + 2 e -0909 + 0.029 log [Sn(OH),*-] - O-118 log [OH-] Sn(OH), + 4 OH- = SnOl*- + 3 H,O + 2 e -0.857 + 0.029 log [SnO,*-] - 0.118 log [OH-] SnO+4OH-=SnO,*-+2H,O+2e -0.844 + O-029 log [SnOl*-] - 0.118 log [OH-] HSnOp- + HI0 + OH- = Sn(OH), + 2 e -0.764 - O-029 log [HSnO,-] - 0.029 log [OH-] HSnO,- + OH- = SnO, + H1O + 2 e -@983 - 0.029 log [HSnO,-] - 0.029 log [OH-] HSnOl- + HtO + 3 OH- = Sn(OH),*- + 2 e -0.935 i- 0.029 log [Sn(OH),*-] - 0.029 log [HSnO,-] - 0.089 log [OH-] HSnO,- + 3 OH- = SnOlp- + 2 HI0 + 2 e -0.868 + 0.029 log [Snot*-] - 0.029 log [HSnO,-] - @089 log [OH-]
E = -0.920 E = E = E = E = E = E = E = E = E = E = E = E= E = E = E = E =
(24) E = (25) E =
* The value for the free energy of formation
of SnO*“-,
-137.42
-42.41 -43.00 -77.12 - 87.20 -34.71 -45.19 -34.12 -44.20 -4190 -84.93 -81.92” -42.50 -41.93 -39*51+ - 38.92* - 35.22 -45.30 -43.03 -40.02+
Kcal, is cited by Pourbaix
et al.18
The formation of Sn(OH),2- is kinetically more difficult and hence its direct formation is improbable. The larger free-energy decrease signifies, however, the instability of stannite solutions. Actually stannite solutions are known to disproportionate into metallic tin and stannate. Sn(OH), and SnO could be oxidized by either Sn(OH),, reactions (11) and (13), or to Sn(OH),2- (Sn0s2-), reactions (18)-(21). The latter set of reactions involves a larger free-energy decrease and are apparently more probable. Inspection of the potential, EBo, at unit activity of Sn(OH),2- (SnOs2-) and OH- reveals, however, that
Anodic passivity of tin in alkaline solutions
881
they are very near to that for the Sn/Sn(OH), system, especially so for reactions involving Sn(OH),,a- ion. Since the difference in potential between the Sn/Sn(OH), and Sn(OH),/Sn(OH), amounts to only 168 mV, and as the pH-dependence of the potentials of reactions (18)-(21) is greater than that for reactions (11) and (13), it is probable that Sn(0I-Q preferentially forms instead of Sn(OH),% at low pH values. The oxidation of HSnO,- to Sn(OH),, reaction (22), involves practically the same free-energy decrease as that for reaction (11). The pH-dependence of the potential of both reactions is, however, different and a decision as to which of the two reactions is taking place could be drawn on this basis (assuming constant HSnO,- activity). The oxidation of HSnO,- to Sn(OH),2- (Sn0a2-), reactions (24) and (25) is favoured by a large decrease in AG”. At unit activity of HSnO,, Sn(OH),s- or Sn0s2the potential of the tin electrode varies by O-089 V/pH unit. How far these thermodynamical predictions are correct was of interest to establish experimentally. EXPERIMENTAL TECHNIQUE The electrolytic cell used was as previously described .I1 The tin electrode was prepared by casting analytically-pure tin granules (Merck, Germany) in the form of a rod 8.5 mm in diameter. The total exposed area was 4.17 cm4. Before use the electrode was polished with the finest grade emery paper and then washed with running distilled water. Experiments were conducted mainly in carbonate-free 2 N @H co 14), 0.1 N @H 12.9) and 0.01 N (pH 12.1) NaOH solutions, as well as in 0.1 M Na,COI solutions (pH 11.1). Each experiment was carried out with a freshly-deaerated solution and with a newly-polished electrode. Polarization curves were recorded on a Leeds and Northrup type-G pen recording potentiometer. The technique for tracing cyclic polarization curves was essentially as described before.” Potentials were measured relative to a saturated calomel electrode and then corrected to the normal hydrogen scale. We have studied also the effect of addition of stannite and stannate on the polarization curves of tin. This type of experiment was carried out mainly in O-1 N NaOH solution. Enough solid Sn(OH), or Sn(OH), was shaken with the hydroxide solution for a period of 48 h to ensure saturation. The solutions were then centrifuged and the clear liquids were diluted with O-1 N NaOH to give solutions of definite starmite or stannate contents. The tin concentration is here expressed in percentage relation to the original saturated solution. All measurements were conducted at 25°C. RESULTS
AND
DISCUSSION
The anodic behaviour of tin in dilute alkaline solutions Figure 1 represents the potential variation of the tin electrode when being subjected to cyclic anodic and cathodic polarization in 0.01 N NaOH solution. The polarizing current in this set of curves was 2 mA/electrode. As can. be seen from the anodic half-cycle curves (curves A), there occurred at first a rapid and almost linear change of potential (part I) due to both the decay of hydrogen overpotential and the subsequent charging of the anodic double layer. Following this process, the potential of the tin electrode changed more slowly, giving rise to two distinct arrests, (part II), before finally passing to values characteristic for oxygen evolution (part III). The starting potentials, E,,, of the two steps of part (II) are -0.78 and -0.62 V respectively. These values compare very satisfactorily with those for the two anodic reactions (7) and (1 l), Table 2, at the corresponding pH value. As will be shown later when discussing the behaviour of the tin electrode in more concentrated alkaline solutions, the potential of these two arrests changed by -59 mV for every pH-unit increase. This strongly suggests that during the anodic oxidation of tin, Sn(OH), (or SnO) and Sn(OH), are formed on the metal surface before oxygen evolution. Dissolution as stannite and stannatelBe is not likely to occur since, for a definite HSnO, or Sn(OH),Z- concentration, the potential should change by -89 mV/pH
A. M. SHAMSEL DIN
888
and F. M.
ALID
EL WAHALI
unit. On the other hand, to assume that the observed variation of -59 mV/pH was due to a corresponding decrease in the concentration of either ion implies that this decrease should be constant for every pH-unit change. Since stannite and stannate are freely soluble in alkaline solutions, such a requisite does not seem reasonable. Inspection of the anodic curves reveals also that the second oxidation step consumed larger quantities of electricity than is simply required for the transformation of Sn(OH), into Sn(OH),. This is apparently due to the higher solubility of Sn(OH), in alkaline solutions. Similarly, upon repeated cathodic reduction and anodic oxidation, the
I.E50 -
I.450 -
>
I.050 -
2
0,650 -
2 0 f 0.250 i ::
0.150-
E Y w-o.550
-
-0.950
-
-1.350
-
I
5 (A)
0 *
15
30
Time,
45
60
75
.90
min
FIG. 1. Cyclic polarization curves for the tin electrode in O-01 N NaOH solution. A, anodic half-cycles, C, cathodic half-cycles. Polarizing current = 2 mA/electrode.
quantity of electricity necessary to bring the electrode to oxygen evolution potential increased with every new anodic half-cycle. The increase was mainly confined to the second step and was probably due to a corresponding augmentation of the true surface area of the electrode as a result of the deposition of the metal during the previous cathodic half-cycle. Cathodic polarization curves (curves C, Fig. 1) were obtained by reversing the polarizing current when the electrode was at oxygen-evolution potentials. The potential first dropped to negative values, but levelled up quickly, giving rise to a sharp peak. The potential then changed more slowly before finally dropping to values characteristic for the evolution of hydrogen. The difference in potential between the apex of the peak and the top of the flat maximum increased with the cycle number. The kink in the cathodic curves was observed only if the previous anodic polarization was allowed to reach the oxygen-evolution potential. If the anodic polarization was interrupted at the Sn(OH),/Sn(OH), potential before the current was reversed, the cathodic curves did not show the kink. This suggests that the occurrence of the kink in the cathodic curves is related to a certain transformation taking place
Anodic passivity of tin
in alkaline solutions
889
at the oxygen-evolution potential. As will be shown in the discussion of the anodicdecay curves, oxygen evolution occurs most probably on a film of metastannic acid (H,SnO, or Sn02.H,0).2*7 The features of the cathodic polarization curves could be explained by assuming that at the instant of reversal of the current, the reduction of the actual passivating film, which is only few monolayers thick,**n takes place and the potential of the electrode tends to hydrogen-evolution values. The metastannic acid film then comes in direct contact with the metal. Since the Sn/Sn(IV) potential in alkaline solutions is more positive than the corresponding Sn/Sn(II), the potential of the electrode rises again. Because of the instability of the Sn/Sn(IV) couple, this process is only transient and the reduction proceeds both chemically and electrochemically. Reduction of some of the dissolved tin is expected to occur simultaneously. The cathodic recovery,ll ie the percentage ratio of the quantities of electricity consumed in the cathodic and anodic half-cycles, was not constant and did not exceed 65 per cent. Apparently some of the tin in solution diffused away from the electrode and escaped subsequent reduction. In fact, some small tin dendrites were always found on the platinum electrode (cathode) at the end of the experiment. The behaviour of the tin electrode in O-1 M Na,CO, solutions was similar to that described above. The potentials of the two oxidation arrests were, however, less defined and the electrode required less quantities of electricity to passivate. In some experiments the tin electrode did not reach oxygen evolution and the potential became irregular. Most probably some basic carbonate formed on the electrode surface. The behaviour of the tin anode in concentrated
alkaline solutions
With currents equal to or less than 5 mA/electrode, the tin anode dissolved continually in 0.1 N NaOH solution at the Sn/Sn(OH), potential (-0.85 V). The time/potential curve run, therefore, parallel to the time axis (curve a, Fig. 2). When the current was raised to 7 mA/electrode, curve b, Fig. 2, dissolution as stannite took place for a comparatively long period and then the potential started to oscillate within a range of 120 mV. The frequency of oscillation was higher at early than at advanced stages. Upon further increase of the current, the dissolution at the Sn/ Sn(OH), potential was confined to a short interval before oscillations started (curve c). At 12 mA/electrode the polarization curve showed that after an initial limited segment at the Sn/Sn(OH), potential the tin electrode dissolved mainly at the Sn(OH),/ Sn(OH), potential. Oscillations in potential were observed only at advanced periods. At this stage the amplitude of oscillation increased to ca 180 mV and the frequency decreased to a minimum. At 14 mA/electrode, the electrode dissolved mainly at the Sn(OH),/Sn(OH), potential and did not-within the time of our experiments-change to oxygen evolution. The above-described phenomena were equally observed with increasing as well as with decreasing polarizing currents, provided that-in the latter case-the electrode did not change to oxygen-evolution potentials. When oxygen was allowed to evolve on the electrode and then the current decreased in steps, the electrode did not return directly to the active state. The lowering of the current affected only the oxygen overpotential on the passive anode. This fact is generally made use of in the electrodeposition of tin from alkaline solutions (see later) where the anode is first filmed at high currents to ensure its subsequent dissolution in the quadrivalent state.2 With currents higher than 14 mA/electrode permanent passivity of the tin electrode in 0.1 N NaOH solution readily occurred.
A. M. SHAMSEL DIN and F. M. fiBD EL
0
I5
30
WAHAB
45
60
min
Time,
FIG.2. Anodic polarization curves for the tin electrode in O-1 N NaOH solution. a, 5 mA; b, 7 mA; c, 10 mA; d, 12 mA; e, 14 mA/electrode.
In 2 N NaOH solution the tin electrode behaved in a similar manner as in O-1 N solution. Here, however, much higher currents were needed to suppress the dissolution of Sn(OH), and Sn(dH),, before the electrode could be made to evolve oxygen. In Fig. 3 a typical experiment for the passivation of tin in 2 N NaOH solution is shown. The current was held constant until either a constant potential (dissolution region) or a constant frequency (oscillation region) was attained. As seen from Fig. 3, oxygen would not evolve on the electrode unless a current equal to or greater than 350 mA was imposed on the electrode.
8 j 4.350 .y 0
15
30 Time.
FIG. 3. Anodic polarization
45
60
75
90
min.
curve for the tin electrod& in 2 N NaOH solution using different currents.
Anodic passivity of tin in alkaline solutions
891
Oscillations in time/potential curves under galvanostatic conditions are frequently reporuzdsssab It is generally agreed upon that they denote the alternating formation and dissolution of a film on the surface of the electrode. The behaviour of the tin electrode in concentrated alkali solutions (>O*l N) seems to support the same conclusion. Below 5 mA/electrode in O-1 N NaOH and 250 mA/electrode in 2 N NaOH solutions, the rate of dissolution of the anodically-formed Sn(OH), surpasses or equals the rate of its formation. The potential remains, therefore, constant at the Sn/Sn(OH), value at the corresponding pH. As the current is increased, the discharge of OH- supersedes the dissolution of Sn(OH),, and Sn(OH), forms at its thermodynamical potential. The electrode is thus in a state ofprimary passivity. Due to its acid character, Sn(OH), dissolves freely in the alkali solution to yield stannate. This would expose the Sn(OH), film and hence the potential drops again to the Sn/Sn(OH), value. The process of formation and dissolution of Sn(OH), would repeat itself so long as a permanent layer of stannic hydroxide cannot persist on the electrode surface. As the current is increased, the ability of the hydroxide to remain intact on the electrode surface increases and the potential remains constant at the Sn(OH),/Sn(OH), value. The rate of formation of Sn(OH), then equals the rate of its dissolution. In O-1 N NaOH solution this seems to occur at 14 mA/electrode. Above this current, permanentpussivation of the tin electrode can take place. The curves of Fig. 2 clearly show that before oxygen evolution, both Sn(OH), and Sn(OH), formed on the surface of the electrode. Although difficulty was experienced in completely reproducing the quantities of electricity necessary to bring the electrode to oxygen evolution potentials, the experimental results seem to fit a relation of the form (i - i&P = const.,
(22)
where i is the applied current and T is the time from the start of polarization till oxygen evolution. i,, represents the current below which permanent passivation could not be achieved, here taken as 14 mA/electrode in O-1 N NaOH. The value of n in the above relation approaches O-5. Relation (22) is common for the anodic passivation of metals.a7*28 Anodic decay curves were recorded by switching off the polarizing current when the electrode was at oxygen-evolution potentials. The curves obtained, Fig. 4, show two definite arrests before the potential reached the equilibrium value of the Sn/ Sn(OH), system. The second of these two arrests was pH-dependent and corresponded to the transformation of Sn(OH), to Sn(OH),. The characteristics of the fist, decay arrest are of particular interest. It is pHindependent and occurs at potentials positive to any of the known redox-systems of tin. From experiments in which the tin electrode was anodically polarized to increasing pre-determined potentials, it was ascertained that the arrest would not appear unless the electrode reached oxygen-evolution potentials. In solutions of low pH, self activation was a very slow process. Cathodic pulses of very short duration (cu 1 s) had to be applied to assist the electrode to reach the Sn/Sn(OH), potential (curve c, Fig. 4). The length of the arrest was not proportional to the time of the preceding oxygen evolution. With too long times at oxygen evolution potentials, or with too high passivating currents, the electrode did not return easily to the Sn/Sn(OH), potential and cathodic pulses had to be applied. The start potential, mid-arrest potential and/or end potential of this arrest were not definite, but depended primarily
892
A. M. SHAMS EL DIN and F. M. ABD EL WAHAB
upon how long the electrode evolved oxygen before the decay experiment was started. The plot of the electrode potential USthe logarithm of time (s) during the first decay arrest was generally non-linear. In those few instances where a straight line relationship was fulfilled, the slope of the lines was exceptionally high, amounting to ca 8RT/F. This rules out the probability of a kinetic process governing the process of self activation. As pointed out above, the first decay arrest is pH-independent, does not correspond to any of the known redox-potentials of tin, and is observed only after the electrode
Time,
min
FIG. 4. Anodic decay curves for the tin electrode
a, in 0.1 N NaOH;
b, in 2 N NaOH;
in alkaline solutions. c, in 0.01 N NaOH with cathodic pulses.
has evolved oxygen. This limits its cause to a chemical reaction not involving electron transfer. According to FSrster2 and Bianchi,’ oxygen evolution on tin anodes in alkaline solutions takes place on a layer of metastannic acid, H,SnO,. This means that Sn(OH), undergoes a process of dehydration, which is thermodynamically favoured. The first decay arrest is most probably the result of the chemical dissolution of H,SnO,. The rate of dissolution of the acid would depend upon the alkalinity of the solution in the manner observed experimentally. The recent work of Shah and Davies9 is of particular interest as it was conducted under apparently similar conditions. Shah and Davies reported a single oxidation arrest for tin in weakly alkaline solutions. This was ascribed to the simultaneous oxidation to SnO and HSn02- in equivalent quantities according to 2Sn+SOH-=SnO+HSn02-+2H20+4e.
(23) From thermodynamical data, the free energy change of reaction (23), -84.9 Kcal, was computed, and correspondingly the equilibrium potential in O-1 M sodium borate solution (pH 9.28) was estimated as -0.85 V (relative to Ag/AgCl reference electrode). Because of the close proximity of this value to that of the recorded step, it was concluded that reaction (23) was governing the electrode behaviour. Reaction (23) is, however, kinetically very improbable as a single step. Elementary electrochemical reactions for oxide formation and stannite dissolution (reactions (8)
893
Anodic passivity of tin in alkaline solutions
and (13, Table 2) must, therefore, be considered. Also, the fact that the calculated potential was close to that of the experimentally observed step does not imply that (23) is actually taking place. Since reactions (8) and (15) have comparable AGo values, their summation would naturally lead to a corresponding doubling of the overall free-energy change and, hence, the EBo value would come again to that for the separate reactions. Further, Shah and Davies, in their calculation of the equilibrium potential of reaction (23) at pH 9.28, assumed that the potential would change by $59 mV for
I
.
0
15
30 Time,
45
60
t-75
min
FIG. 5. Anodic polarization curves for the tin electrode in 0.1 N NaOH containing
20 g/l KCl. 1, 10 mA/electrode; 2, 20 ma/electrode; 3, cathodic polarization curve with 20 mA/electrode. every pH unit decrease. This is, however, not the case as far as reaction (15) is concerned; the change is +89 mV. The absence of a definite oxidation arrest corresponding to the formation of Sn(OH), in Shah and Davies’ polarization curves is also strange. This arrest has been recorded in our previous study as well as in the present one. It is even observed with tin alloys containing copper10*2g and zinc.30 We are unable to give a satisfactory explanation for this. It was thought at first that the absence of this arrest in Shah and Davies’ curves was the result of the presence of Cl- ion in solution (added so that the Ag/AgCl electrode would function). This ion is known to interfere with passivation processes.31*3a We conducted, therefore, some passivation experiments in O-1N NaOH solution containing KC1 (20 g/l). The results are shown in Fig. 5. At low currents a small arrest corresponding to the formation of Sn(OH), was recorded. The potential then changed to the Sn(OH),/Sn(OH), value where it remained constant, curve 1, Fig. 5. When the current was raised, curve 2, the instantaneous formation of Sn(OH), and Sn(OH), was followed by the development of a dense white film (probably tin oxy-chloride) on the electrode surface. This film did not adhere strongly to the metal
894
A. M. SHAMSEL DIN and F. M. ABD EL WAHAB
surface and the potential of the electrode became irregular. Interesting is the fact that upon reversing of the current two kinks were recorded in the cathodic polarization curve (curve 3, Fig. 5). This substantiates the above presented idea that the occurrence of a kink is not due to the presence of a particular oxidation state, but rather to the variation in the type of the reduced species. In the above experiment the first kink is due to the reduction of the actual passivating layer. The potential rises to respond to the reduction of metastannic acid coming into contact .with the free tin surface. When this process comes to an end, the potential tends again towards negative values. However, since the electrode is still covered with a layer of oxy-chloride, the potential rises once more until this film is completely reduced. The behaviour of the tin anode in alkaline solutions containing stannite and stannate As a complement to the work on the passivity of tin in pure alkaline solutions, we extended the study to cover its behaviour in alkaline solutions containing stannite and stannate. This point is of practical interest since the electrodeposition of the metal is carried out mainly from such solutions. 84 The behaviour of the tin anode in tin plating solutions has been previously examined, 4*5but-the conclusions are not in agreement. Our study was the recording of galvanostatic time/potential curves in alkaline solutions containing different concentrations of either sodium stannite or stannate (see Experimental Technique). As shown in the previous section, the polarization curve of the tin anode in 0.1 N NaOH with a current of 10 mA/electrode exhibited a short segment at the Sn/Sn(OH), potential before oscillations started. In the presence of dissolved Sn(OH),, the general polarization behaviour of the electrode varied to a lesser or greater extent depending upon the concentration of the added tin hydroxide. With the same polarizing current, and in a solution 10 per cent saturated with Sn(OH),, the part of the polarization curve representing the dissolution of the anodically formed Sn(OH), was greatly reduced, apparently due to the presence of part of the anodic products at the metal/solution interphase. Following this, free dissolution at the Sn(OH),/Sn(OH), potential took place. Oscillations occurred only at advanced times. These were less frequent and the duration of each cycle was longer than in absence of dissolved stannite, curve a, Fig. 6. The sharp peaks observed at the start of each cycle represent the tendency of the electrode to attain a state of permanent passivity that is quickly levelled down by the dissolution of Sn(OH),. In solutions 20 per cent saturated with Sn(OH),, the general behaviour of the tin electrode was as described above save for the fact that the oscillations were still less frequent. In 30-50 per cent saturated solutions, the electrode continually dissolved at the Sn(OH),/ Sn(OH), potential. Oscillations were seldom observed. Within the time limit of our experiments, the electrode did not change to oxygen-evolution potentials. In 60 per cent solutions, and with a current of 10 mA/electrode, the tin electrode was readily passivated within a relatively short period, curve c, Fig. 6. It will be remembered from the results of the previous section that the same sequence of processes was obtained through the application of increasing currents. The presence in solution of the anode products is equivalent, therefore, to the application of higher currents. This, together with the passivation law given above (relation 22), suggests that the process of passivation is governed by the rate of nucleus formation of the hydroxides on the electrode surface.
895
Anodic passivity of tin in alkaline solutions
The passivation of the tin anode in solutions of higher tin content and with a current of 10 n-A/electrode was almost instantaneous. The polarizing current had to be lowered in order to ascertain the formation of both Sn(OH), and Sn(OH),. The behaviour of the tin anode in solutions containing Sn(OH), resembled that in Sn(OH), solutions. Passivity occurred, however, more readily, apparently due to the high tin content of the solutions. Oscillations in potential were observed only in solutions of 10 per cent (or less) saturation. In solutions of 20-40 per cent saturation, continual dissolution at the Sn(OH),/Sn(OH), potential took place. Passivity with a I
I , I
I I
0.25C
I
A-
I
(b)
______d’
J
c
_(a)
v
I I I I
CI g
s w
-1.350
&
-1.750
A-
:(c) I t : 0
15
45
30 Time,
60
75
min
FIG. 6. Anodic polarization curves for the tin electrode in O-1 N NaOH solution of: a, 10 per cent saturation with Sn(OH), b, 40 per cent saturation with Sn(OH), c, 60 per cent saturation with Sn(OH), Polarizing current, 10 mA/electrode.
current of 10 mA/electrode was easily attained in solutions 50 per cent (or more) saturated with Sn(OH),. With increasing tin content, the quantities of electricity associated with the Sn(OH), and Sn(OH), arrests decreased and lower currents were always applied to prove the formation of both hydroxides. In this respect our results are not in agreement with those of H&A,* who reported that the transition current for the active-passive transformation increased with the increase of the Sn(IV) content of the bath. The results of the present investigation allow the following picture about the passivity of the tin electrode in alkaline solutions to be drawn. The primary oxidation of tin in alkaline solution results in the formation of Sn(OH),. Depending upon the magnitude of the polarizing current, and upon the alkalinity of the solution, the formed hydroxide dissolves to yield stannite. Dissolution continues until the saturation of the metal/solution interphase with stannite leads to the crystallization of Sn(OH), on the electrode surface. The subsequent discharge of OH- ion results in the oxidation of Sn(OH), to Sn(OH),. The electrode is then in a state of primary passivity.
896
A. M. SHAMSEL DIN and F. M. f%FID EL WAHAB
Sn(OH),-by virtue of its higher acidity4issolves as stannate. When the rate of dissolution exceeds that of formation, oscillations in potential/time curves are observed. Once the rate of OH- discharge exceeds that of Sn(OH), dissolution, permanent passivation sets in and oxygen can evolve on the anode. At oxygen-evolution potential Sn(OH), partly dehydrates to metastannic acid. The pre-saturation of the solution with the anodic products facilitates the attainment of both the primary and permanent states of passivity, but does not alter in any way the course of oxidation of the metal. REFERENCES 1. H. GOLDS-T and E. EKARDT,Z. phys. Chem. 56,385 (1906). 2. F. FOERSTER and M. DOLCH,Z. Eiektrochem. 16,599 (1909). 3. E. NE~BERY,J. Chem. Sot. iO66 (1916). . 4. G. -EL and A. GRAVEL,Z. Elektrochem. 41, 314 (1935). 5. R. KERR,J. Chem. Sot. Znd. 57,405 (1938); Chem. Abstr. 1601 (1939). 6. G. T. BACHVAUIW,J. Appl. Chem. Moscow, 14,469 (1941); Chem. Abstr. 2221 (1942); G. T. BACH~AL~Wand P. S. Trrov, Khim. Referat Zh. 4,85 (1941); Chem. Abstr. 5659 (1943). 7. G. BIANCHI,Chim. Znd. 29,295 (1947). 8. S. E. S. EL WAKKAD, A. M. Sm EL DIN and J. A. EL SAYED,J. Chem. Sot. 3103 (1954). 9. S. N. SHAH and D. E. DAVY Electrochim. Acta 8,663,703 (1963). 10. A. M. SW EL DIN and F. M. ABD EL WAHAB, unpublished results. 11. A. M. SHAMSEL DIN and F. M. ABD EL WAHAB, Electrochim. Acta, 9, 113 (1964). 12. A. KUTZELNIGG,Z. anorg. Chem. 202,418 (1931). 13. N. BOUNDand D. A. RICHARDS,Proc. Phys. Sot. SlB, 256 (1939). 14. R. K. HART, Proc. Phys. Sot. 65B, 955 (1952). 15. S. C. BRIITONand K.“BRIGHT,MetaNurgia, Manchr 56, 163 (1957). 16. W. E. Boczs. R. H. KACHIKand G. E. PELWSSER, J. Electrochem. Sot. 108,6,124 (1961). 17. W. M. LA&R, Oxidation Potentials, pp. 39, 136; Prentice Hall, New York (1953): ’ 18. E. DELTOMBE,N. DE ZOUBOV and M. POURBAIX,Proc. 7th Meeting CZTCE, Lindau 1955, p. 216. Butterworths, London (1957). 19. G. CHARUIT, Theorie et Mithodes Nouvelles d’dnalyse Qualitative, 3rd Ed., p. 222. Masson, Paris (1949). 20. P. DELAHAY, M. POURBAMand P. VAN RYSSELBERGHE, Proc. 3rd Meeting CZTCE, Milan 1952, p. 15. Manfredi, Milan (1952). 21. L. G. SILLEN,J. Chem. Educ. 29,600 (1952). 22. T. P. HOAR, Trans. Faraday Sot. 33,1152 (1937). 23. A. GUTBIER,G. F. H-0 and H. D~BLING, Ber. 59B, 1232 (1926). 24. S. FIELDand A. D. WEILL, Electroplating, p. 426. Pitman, London (1961). 25. E. S. -ES, J. Chem. Sot. 1533,2580,2878 (1926); 1028 (1929). 26. R. no-1 and G. SERRAVALLE, Z. Elektrochem. 62,759 (1958). 27. R. LANDS~ERG and M. HOLLNAGEL, Z. Elektrochem. 58,681 (1954), 60,1098 (1956); R. LANDSBERG,Z. phys. Chem. 206, 291 (1957); R. LANDSBERG and H. BARTELT,Z. Elektrochem. 61, 1162 (1957). 28. N. A. HAMPTON,J. Electrochem. Sot. 110,95 (1963). 29. S. E. S. EL WAKKAD, T. M. SALEM,A. M. SW EL DIN and Z. HANAFI, J. Chem. Sot. 2857 (1956). 30. S. E. S. EL WAKKAD, A. M. SHAMSEL DIN and H. KOTB,J. Electrochem. Sot. 105,47 (1958). 31. B. ~ANOV, R. Bmm and A. FRUMKIN,Disc. Faraday Sot. 1,259 (1947). 32. H. K. E~MI and A. A. Mou~~A, J. Chem. Sot. 2027 (1958).