artigo 24.pdf

Corrosion Science, Vol. 40, No. 2/3, pp. 225-234, 1998

Pergamon :(? 1998 Elsevier Science Ltd. All rights reserved.

Printed in Great Britain. 001lb938x/98 $19.00+0.00

PII: SOOlO-938X(97)00122-4

A LEAD-ACID BATTERY ANALOGUE TO IN SITU ANODE DEGRADATION IN COPPER ELECTROMETALLURGY

G. CIFUENTES.” L. CIFUENTESb and G. CRISOSTOMOb

“Departamento de Ingenieria Metalurgica, Universidad de Santiago, Av. B. O’Higgins 3363, Santiago, Chile

‘Departamento de lngenieria de Minas, Universidad de Chile, Tupper 2069, Santiago, Chile

Abstract-When the applied current is interrupted in a copper electrowinning plant, lead anodes tend to corrode

once the plant is back in operation. This work aims to show, both from theoretical considerations and experimental

evidence, that this phenomenon may arise from the spontaneous coupling of the oxidation of metallic lead to lead

(11) sulfate with the reduction of lead (IV) dioxide to lead (II) sulfate on the unenergized anodes. The process is

similar to the one that takes place in an idle lead-acid battery. It is shown that, in the above conditions, the

protective oxide layer on the anode takes between 5 and IO min to dissolve. Estimates are given for the rate of

lead (IV) dioxide dissolution and lead sulfate formation and conclusions are drawn for plant practice. ‘c 1998

Elsevier Science Ltd. All rights reserved

Ke~cords: anode corrosion, copper electrometallurgy, electrowinning, passivation, lead-acid battery. electro-

chemical kinetics, chronopotentiometry. cyclic voltammetry, weight loss test.

INTRODUCTION

It is widely known that lead anodes used in copper electrowinning can corrode when the current is switched on after an interruption. This phenomenon can be explained in terms of the loss of anode passivity that takes place in the absence of an applied current. The latter favours the integrity of the protective lead (IV) dioxide layer on the lead anode, i.e., during plant operation the anodes are anodically protected. This work presents the view that, in the absence of an applied anodic overpotential, electrochemical reactions take place spontaneously on the anode surface, leading to lead corrosion and dissolution of the lead (IV) dioxide coating, as in an idle lead-acid battery. It is also shown that, in the absence of the applied current, copper deposition takes place spontaneously on the lead anode.

THEORY

Electrochemical reactions in a lead-acid battery

In a lead-acid battery, metallic lead and lead (IV) dioxide electrodes are immersed in sulfuric acid. The lead dioxide electrode is obtained by depositing a porous lead (IV) dioxide layer on a lead substrate. The reactions during operation are as follows.‘.’

At the anode, lead dissolves to give lead (II) ion which immediately reacts to form lead (II) sulfate.

Manuscript received 15 May 1997

225

Pb + SO, + PbSO,( s) + 2e (1)

At the cathode, lead (IV) dioxide reduces to Icad (II) ion which immediately reacts to

form lead (II) sulfate.

PbO,+4H +SO, +2e -+PbS0,(s)+2Hz0

The overall reaction is:

(2)

Pb+PbO~+2H~SO-+2PbSO,(s)+2H~O (3)

During charging, the rcversc reactions take place. When the device is not supplying

energy or charging, both forward reactions couple spontaneously on the lead (IV) dioxide electrode. which causes its transformation to lead (II) sulfate. In this case, the

lead substrate is in contact with the electrolyte through the lead (IV) dioxide pores. Standard electrode potentials for all the relevant half-reactions’ are given in Table I.

In copper electrowinning, copper is deposited at the cathode while oxygen evolves

at the lead anode. The clectrolytc is 21 solution of copper (II) sulfate in sulfuric acid. At the operating current (with densities higher than 2OOA mm’). lead anodes are covered

with a passive layer of lead (IV) dioxide. the formation of which is favoured at potentials higher than the equilibrium potential for the oxygen evolution reaction. Under these conditions, anode dissolution is negligible and the anodic reaction is:

At theanode 2H,O-0,+4H. +4e (4)

When the cell current is switched oil’, the oxide phase stays in direct contact with concentrated sulfuric acid and as the oxide film is porous.’ the metallic lead is also in

contact with acid and dissolves. Solid lead sulfate has been shown to form when lead is

immersed in concentrated sulfuric acid.’ The situation is similar to an idle lead&acid

battery. Reactions (I) and (2) couple spontaneously: ( I) takes place at the pores and (2) at the oxide surface. Figure I pictures this process.

The formation of solid lead (II) sulfate causes the break up of the protective film. as lead sulfate molecules are larger than lead (IV) dioxide ones; their molar volume ratio being sulfate: dioxide = 3.X: I. When the broken oxide film becomes detached from the metal surface. a larger area of fresh metal is exposed. Then. lead corrosion continues

to take place. coupling with any available cathodic reaction, e.g. reaction (2) or the reduction of oxygen gas dissolved in the electrolyte, which is reaction (4) in reverse.

Figure 2 is a qualitative Evans diagram” ’ of the reactions taking place on a lead

electrode in the absence of an applied cell current. Under these conditions, the electrode

Table I. Standard equtltbriutn poknttala I,r relevant half-reacttons’

(I) PbSO,(\)+‘e = Pb+SO, E” = -0.36

(2) PbO, $- 4H ’ + SO, + 3 = Pb kSO,(s)+2H,O E” = 1.6X5

(4) 01+311 ’ 4-4~ : 2H,O E” = I .2? V

(9 Cu’ $ 2e = <‘II F.” = 0 34 v .

(6) 2H t 2e = H. E” = u.0 v

A lead-acid battery analogue to in situ anode degradation in copper electrometallurgy 227

Substrate Coating

Pb02/PbS0,

Anode

Fig. 1, When no current is applied to the cell, both the lead substrate and the lead (IV) dioxide

coating react spontaneously to form lead (11) sulfate.

I

-0.36 0.34 1.23 1.69

E

Fig. 2. Qualitative Evans diagram showing the main reactions for a lead annode in acidic copper

sulfate without an applied current. Standard equilibrium potentials for the relevant reactions are

included as a reference. Notice that copper can deposit on lead.

exhibits a mixed potential. Notice that, when no current flows through a copper electrowinning cell, the presence of cupric ions in solution allows the reaction

On Pb Cu2+ + 2e- +Cu (5)

(i.e. copper deposition on the lead anode). This phenomenon has been observed in plant practice. In this case, copper deposition couples with lead corrosion in a cementation- like process. Other possible reactions are a hydrogen ion reduction to hydrogen gas

OnPb 2H++2ee+H2 (6)

and water decomposition with oxygen evolution [reaction (4)]. Reaction (6) is very slow

on lead. In an Evans diagram, the mixed potential on the lead electrode will be

determined by the intersection of the curves representing the sum of all anodic reaction rates and the sum of all cathodic reaction rates which are taking place.

The current density balance on the lead electrode. assuming that all reactions take

place on the same surface area. is as follows

4, l’i,so, + ~il.o (, = lLC,, PI~zo~/ + IL, rr+l + 14 ,,1 ( LI/

Assuming that low concentration reactants (oxygen, in this case) lead to diffusion (mass

transfer) rate control. their rates will be given by Fick’s law

where it = limiting current density. A m ‘: I I= charge number; F = Faraday’s constant.

C equiv ‘: D = diffusion coefficient, m’s ‘: C’,, = reactant concentration in the bulk

solution. mol m ‘: and ;i = diffusion layer thickness. m.

This equation assumes that no migration effects are present. This is justified, as corrosion takes place in the absence of an applied current.

High concentration reactants (lead. lead (IV) dioxide. cupric ion. vvatcr) favour

activation (charge transfer) rate control. so their reaction rates will be given by the high- held approximations to the Butler ~Volmer equation’

and

where the lirst equation holds for an anodic reaction and the second one holds for a cathodic reaction. The meaning of symbols is: i,, = exchange current density, A m ‘:

x,,x, = charge transfer coefficients (anodic and cathodic); rl = overpotential, V (in this case 11 = E,,,,,- 15,); T = temperature, K; R = gas constant, J mol-’ K ‘; f$,,,, = mixed

potential. V; and EC = equilibrium potential. V. In order to quantify the lead corrosion rate and the lead (IV) dioxide consumption

rate, a current balance is required. As a first approximation, if only the main reactions are taken into account, the current density balance is

/PI> I’hXO, = lh,0, IW~,/

Substituting the high held approximations and solving for E,,,,

where the subindex “A” refers to the anodic reaction [reaction (I) above] and the


subindex “C” refers to the cathodic reaction [reaction (2) above]. The expression for the mixed potential is given here for the sake of completeness.

As a second approximation to the lead (IV) dioxide consumption rate, it may be assumed that the anodic reaction is balanced by both the lead (IV) dioxide and cupric ion reduction reactions, so that the current density balance is

zPb,PbSO, = h’b02,PbS04i + ku2+:Cui

Then, an estimate of the cupric ion reduction rate will give an improved value for the lead (IV) dioxide consumption rate. If necessary, further approximations may be obtained by taking into account the remaining reactions as given above.

EXPERIMENTAL METHOD

Potential-time curve

The experiments were carried out in a 180 gpl sulfuric acid and 40 gpl copper solution at 40°C. The anodes were made of a PbCa (0.02%)-Sn (0.5%) alloy. The working electrode (anode) surface area was 0.86cm’ and a stainless steel coiled wire counter- electrode was used. Its surface area was over ten times greater than the working electrode surface area. A silver/silver chloride reference electrode was used. The electrodes were contained in a 500 ml cell. The electrolyte recirculated at a flowrate of 3 1 per minute by means of a peristaltic pump.

A Model PSG 201 Tacussel potentiostat was used to carry out the following procedure:

(1) A constant 200 A m-2 anodic current density was applied during two hours in order to form a lead (IV) dioxide layer on the working electrode.

(2) Then the c.d. was abruptly interrupted and the electrode potential-time curve was recorded (Fig. 3).

(3) Once the potential became stable, the electrode was removed and subject to microscopic observation.

Cyclic voltammetry

In order to characterize the reactions taking place on lead, a cyclic voltammetry was carried out in a three-electrode cell, using the potentiostat and electrolyte described in the previous section. The working electrode was made of the already described lead alloy, the counter electrode was made of platinum and the reference electrode was saturated calomel (SCE). The electrolyte was 180 gpl sulfuric acid and the temperature was 4OC. Sweep rate was 0.17 mV SK’. The voltammetry is shown in Fig. 4.

Copper deposition on the lead anode

In order to establish the rate of formation of a copper deposit on the lead anode after the cell current is interrupted, the following experiment was carried out: a 25cm* copper-free lead probe was subject to an anodic c.d. of 600Amp2 during 8.5 h in a 180 gpl sulfuric acid solution with addition of 40 gpl copper (II) sulfate. The increased probe area-as compared with the one used in the first experiment-was due to the need of obtaining a non-negligible amount of copper deposit. The increased current density and anodizing time were chosen in order to obtain a good oxide layer on a

5 1000

2 2 > 800

E u

600

400

200 __~~~~ ~~ __

0 50 100 150 200

r (seconds)

Fig. 2. Potential agamst tlmc plot. The ~nfexmn point at about 1400mV (agamst SIIWI- Giver

chloride) corresponds to the transformation oflsad (IV) dioxide into lead (II) sulfate. The potential

hecamc qahle at about 400 mV after 300-600 s.

bigger probe. Then the current was interrupted until the probe reached a stable potential value. This happened after 20min. The probe was withdrawn from the solution and carefully washed with warm distilled water. Then it was immersed in pure 18Ogpl sulfuric

acid and subject to an anodic c.d. of 50 A m ’ in order to dissolve any copper deposit.

Before and after the test. the solution was analyzed for copper and lead via atomic

absorption.

Weight variation experiments were carried out in order to establish and quantify the

spontaneous corrosion of the lead anode during periods of current interruption. Two experiments were carried out. In both cases. a lead electrode of 14cm’ total

surface area was treated in a potassium tartrate (2OOgpl), sodium hydroxide (1OOgpl) and potassium iodide (SOgpl) solution in order to remove all surface oxides and sulfates. washed in distilled water and weighed. Then. the lead anode was immersed in 18Ogpl sulfuric acid at 40 C. together with two 304 stainless steel counterelectrodes of the same surface area placed at 2 cm distance on each side. A constant 250 Am 2 current density was applied during 6.5 h with the lead electrode connected as anode. Then, the experiments followed a different path:

(1)

(2)

Products of reaction on the lead anode were removed according to the previously described procedure and the lead anode was weighed. The lead anode was left in solution. without an applied current, for 24 h in order to ensure that all possible reactions ran to completion, Then, all reaction products were removed as before and the lead anode was weighed.


.

60 -

C

. k h

Fig. 4. Cyclic voltammetry. It shows the main reactions taking place at various electrode potentials

(see text). The leftmost peak represents hydrogen evolution and the rightmost one, oxygen evolution.

At potential above 1500 mV the dominant species is lead (IV) dioxide.

Table 2. Weight variation measurements

Initial weight A. Time* Rest time Final weight Difference

(g) (h) (h) (g) (mg)

Case I 1.8162 6.5 1.7516 -64.6 Case 2 1.7503 6.5 24 1.6814 -68.9

*A. Time = anodizing time (at 250A mm ‘).

The results are given in Table 2.

EXPERIMENTAL RESULTS AND DISCUSSION

Copper deposition on the lead anode Microscopy revealed the presence of deposited copper on the lead anode after its

removal from the electrolyte. This proves that reaction (5) above does take place spontaneously after the current is interrupted. The second experiment showed that, after

131 <i. c1suentes (‘I ill.

anodizing, i.e. after the deposited copper was dissolved, it was present in the solution at a concentration of I .73 ppm in a 300ml beaker. This means that, in the experimental conditions. copper deposits on the unenergized lead anode at an average rate of 2 x 1 OPx g cm~‘s~l or 3x 10 “‘mole cm ‘s ‘. No lead was detected in solution. This was to be expected as any lead ions present tend to form lead (II) sulfate which precipitates.

The potentialMime curve (Fig. 3) shows an inflexion point around 1400 mV (Ag/AgCl)

which corresponds roughly to the transformation of lead (IV) dioxide into lead (11) sulfate, which points to the spontaneous occurrence of this reaction.

The electrode potential became stable at around 400mV after 5 IOmin. This value

falls between the equilibrium potentials for reactions (I) and (2) above. i.e. it may represent a mixed potential for these reactions and other reactions in that range. as for

instance reaction (6) above. From these results. it is clear that both the lead (IV) dioxide/lead (11) sulfate and

the cupric ion/metallic copper reactions take place spontaneously after the current is switched off. Both reactions are cathodic and. in order to occur, they need to couple

with an anodic reaction whose equilibrium potential is less than the equilibrium potential

for cupric ion/metallic copper. In this system, the most likely reaction to play this role is the oxidation of metallic lead to lead (11) sulfate.

Some time after the current is interrupted. the anode becomes chemically unstable,

as the lead (IV) oxide. which protects the anode from corrosion, reacts to form lead (II)

sulfate. The time required to achieve anode activation (i.e.. loss of passivity) is between 5

and IOmin. This very brief period of time is likely to be due to the fact that the lead (IV) dioxide layer is very thin (of the order of 50pm), i.e. the amount of oxide present

on the anode is low and disappears rapidly as a result of reaction (2).

Cyclic voltammetries of lead in sulfuric acid solution have been carried out and

discussed by several authors.” ” so the possible reaction are well understood. The cyclic

voltammetry obtained in this work (Fig. 4) is similar to previously published ones and shows known features: (a) formation of lead (11) sulfate by oxidation of metallic lead

to lead (11); (b) passivation due to formation of a lead (11) sulfate film: (c) formation of lead (IV) dioxide by oxidation of lead (11) to lead (IV). Oxygen gas evolution takes place simultaneously: (d) reduction. to metallic lead, of lead (II) sulfate and non-

stoichiometrical lead oxy-sulfates formed during passivation. The presence of sulfate and non-stoichiometrical oxy-sulfates is shown by the presence of three peaks in zone d. The leftmost peak in this zone represents hydrogen gas evolution; (e) reactivation, i.e. oxidation. to lead (IV) dioxide. of the lead sulfates and non-stoichiometrical oxy- sulfates which did not oxidize during the forward sweep.” This behaviour is typical of semiconductor materials such as non-stoichiometrical oxy-sulfates; and (f) reduction of lead (IV) dioxide to lead (II) sulfate and non-stoichiometrical oxy-sulfates.

The voltammetry shows that the oxidation of various species to lead (IV) dioxide begins in the zone between the e and f peaks, that is, it takes place at potentials equal or higher than 1500mV (SCE). This value is below the anode working potential, which


is about 2000mV (SCE). This shows that, at the anode working potential, the dominant species is lead (IV) dioxide.

Weight variation measurements The following discussion refers to Table 2. In Case 1, the mass diference is -64.6mg.

It represents the lead mass which is lost during 6.5 h of normal operation at 250Am-‘. In Case 2, the mass difference -68.9mg. It represents the lead mass which is lost

during normal operation plus the lead mass lost during the spontaneous transformation of lead (IV) dioxide to lead (II) sulfate coupled with the oxidation of metallic lead to lead (II) sulfate while the current is interrupted.

The difference between the two lead mass losses (4.3 mg or 0.3 1 mg cme2) represents the lead mass lost during current interruption when the spontaneous process runs to completion, i.e. all lead (IV) dioxide reacts to form lead (II) sulfate.

For an industrial cell which contains 71 anodes of 2 m2 total surface area each, the lead mass lost during current interruption, if the interruption time is long enough for all the lead (IV) dioxide to form lead (II) sulfate, is of the order of 440g per cell.

Reaction rates From the previous experiment, if a dissolution time of 10min is assumed, then the

average lead dissolution rate is 5.17 x 10p7g cm-*s-’ or 2.5 x lo-‘mol cm-* ss’ during the period of lead (IV) dioxide to lead (II) sulfate conversion.

From the equations given to quantify the corrosion rate, it follows that the lead mass loss must be balanced by the most likely cathodic reactions, i.e. reactions (2) and (5). The copper deposition rate was calculated as 2 x lop8 g cmm2 s-’ or 3 x 10~‘“mole cm -2ss’, so it is less than the lead dissolution rate by a factor of about ten. It follows that the lead dissolution reaction is mainly balanced by the lead (IV) dioxide reduction to lead (II) sulfate. The reaction rate for the latter reaction is slightly less than 2.5 x 10-9mol cm-*ss’. All mentioned reaction rates are average values.

Implications ,for plant practice

These results show that, in plant practice, a current interruption of less than 5min may be acceptable and, in this case, anode damage after the current is switched on is unlikely as a thin film of protective lead (IV) dioxide is still present. They also suggest that an interruption which lasted more than 1Omin would be unequivocally damaging to the anodes. As soon as the current is switched on, the anodes will dissolve until passivation is achieved. This may lead to severe contamination of the copper cathode with lead, causing financial loss.

Plant practice has also shown that, after current interruptions, anode performance decreases, leading to an increased cell voltage. This suggests that lead (II) sulfate-which would cover a fraction of the anode area after the sulfate-forming reactions have taken place-has a lower catalytic effect on oxygen gas evolution than lead (IV) dioxide. Obviously, an increased cell voltage means increased energy costs, also causing financial loss.

In order to preserve passivity, a protective (“back up”) anodic current should be applied at all times to the lead anodes, including the periods during which the operation is interrupted for cathode removal and cell clean up. Despite the fact that protective c.d.s of less than 10 A rnp2 are often recommended in plant practice, unpublished results

from plant tests indicate that the c.d. required to preserve passivity during a longer period of operation interruption (say. several hours) may be considerably higher. Its

value for a specific plant depends on anode and electrolyte composition and other plant conditions. so it should be determined by experiment.

A higher protective current would. of course. demand the use of a more expensive

rectifier. The convenience of installing it would depend on an economic study. For each plant, the likelihood, duration and frequency of current interruptions should be assessed.

The resulting financial losses should then be balanced against the cost of implementing

an improved anodic protection system.

CONCLUSION

Both theoretical considerations and experimental evidence have been put forward to show that lead anode corrosion~ which occurs after a current interruption in copper electrowinning plants arises from the spontaneous coupling of the oxidation of metallic

lead to lead (11) sulfate with the reduction of lead (IV) dioxide to lead (II) sulfate on the unenergized lead anode. which is immersed in an acidic copper sulfate electrolyte. Under

these conditions, the protective lead (IV) dioxide layer on the anode takes between 5 and

10 min to dissolve. It follows that any current interruption in a copper electrowinning plant which lasted longer than ten minutes would lead to anode corrosion, cathode contamination

and financial loss once the current is switched on.

A~~X,7on~lec~l/c~~l~l~i?l The authors would like to thank the National ~‘ommittee for Scientific and Technological

Research (CONICYT. Chile) for the tinancial support given to this line of work vu FONDECYT Project No.

19.50532.

REFERENC’ES

I. Bockrla. J.O’M. and Reddq. A.K.N.. .Moc/w/ ~‘l~~~,/~r,c/lc,,ll/\/,,~,. Vol. 2. Plenum Press. 1970. pp. XXS and 1414.

2. Kortb. G. L~~hrhuc~h A. /+A /,.oc /w~~i<,. Verlag (‘hemic. 1970. pp. 298 299.

3. Weast. R.C (‘I rrl.. (‘KC’ t_lot~tlhooX 1~1 (‘hnui.~~r~~ trutl P/II,,I~ \. 70th edn. CRC Press Inc.. 1989.

4. Ruestchi. P. ./. ~~w~wc~/tcw. So<,., 1973. 120(3). ii I 3.X 5. Steuart. L. S. and Bennion. D. N. Mathematical lnodel of the nnodic oxidation of Icad. ./. E/et rroclw~. Sw..

1994.141(9). 2416 2421.

6. Fontana, M.G. and Grecnc. N.D.. ( ‘o/.I-o(./o/I /%~giww~rrq. 2nd edn. McGraw-Hill. 197X. Chap. 9 & IO.

7. Cifuentcs. L. Harnessing the dynanmcs (~fclcctroche~~~ic~l processes. C%cw. Eg/q/. 1995. 102(I). 102~~10s.

8. Cifuentes. L.. Evans diagranls in copper electrometaIIt1r@~. In Frock. o/‘Co~ppw ‘VS. 3rd /rz/rrrrcr/iorrd ~‘OIZ,~I.CJJ. Vol. 3. Santiago. Chile. 1995.

9. Bialack>. J. A. and Hatmpson. N A. The eft‘cct ol’allo)~ng wth Sh and Ca;Sn on the electrochetmxal properties

of solid lead. ./. /~/~,[,/r.~,~./~rrlr. .%I . 1983, 130(Y). I797 I799

IO. Sharp-e. T. F. Low rate cathodic Ilnear weep voltammetry (LSV) studies on anodixd lead. ./. E~w/w&r~~.

Sot... 1975. 122(7). X45 X51

I I. Mahato. B. K. The cyclic corrosion of the lead-xxi hatter? poutive. ./. &c~/~oc~/Io~~. Sot,.. 1979, 126(j), 3hS-

;74.

artigo 24.pdf

Documents