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Electrochemical studies of molten alkali sulphates
Etcctrochimica Acta. 1970, Vol. 15, pp. 445 to 458.
Pergamon Prays.Printedin NorthernIreland
ELECTROCHEMICAL ALKALI
STUDIES OF MOLTEN SULPHATES*
B. W. BURROWS~ and G. J. HILLS Chemistry Department, University of Southampton, Southampton, England Abstract-Electrochemical studies in the Li,SOI-K,SOI eutectic melt at 625°C are reported. Experimental evidence suggests that the sulphate ion is not directly reduced at an inert metal cathode. The corrosion of metals such as Fe and Ni in molten sulphates appears to result from reaction with SOS derived from dissociation of SOI*- anions. This dissociation is enhanced by the relative acidity of the metal-ion species. The limiting anodic process was studied by steady state and transient galvanostatic techniques in melts of different basicities. Discharge mechanisms are proposed to explain the observed behaviour. R&sum&-On pr&scnte des etudes electrochimique dun entectique fondu de Li,SO,-K,SOI g 625°C. Avec l’evidence experimentale on propose que l’ion de sulphate n’est pas directement I-eduit & la cathode metallique inerte. La corrosion des m&aux comme le Fe et le Ni darts les sulphates fondus peut r&ulter dune reaction avec le SOa qui vient de la dissociation des anions SO,s-. Cette dissociation est encouraged par l’acidite relative aux esp&cesd’ions m&alliques. On a etudie le processus anodique limitant par des techniques galvanostatiques stationaires et transitoires pour des sels fondus de bacicities differents. On propose des mechanisms de d&charge avec lesquelles on explique les fawns de se cornporter que l’on observe. Zusammenfassung-Elektrochemische Untersuchungen an der eutektischen Schmelze von LirSOaI&SO, bei 625°C werden beschrieben. Die experimentellen Ergebnisse weisen darauf hin, dass das Sulfation an inerten Metallkathoden nicht direkt reduziert wird. Die ICorrosion von Metallen wie Fe und Ni in Sulfatschmelzen scheint auf eine Reaktion mit SO, zuriickzufiihren sein, welches durch die Dissoziation von SO&-* entsteht. Diese Dissoziation wird durch die relative Aziditat des metallischen Ions begiinstigt. Der begrenzende anodische Vorgang wurde mittles Gleichgewichtsund nicht-stationf+.rergalvanostatischer Methoden in Schmelzen verschiedener Basizitiit untersucht. Zur ErklHrung des beobachteten Verhaltens werden Entladungsmechanismen vorgeschlagen. INTRODUCTION
studies in oxyanion melts have been largely confined to nitrate and carbonate systems. l The former are low-melting and thus convenient media to work with, while the latter have considerable technological importance as electrolytes for high temperature fuel cells. Sulphate melts, on the other hand, are relatively high-melting, like the carbonates, but have little practical importance apart from their deleterious effects on boiler tubes in power stations burning fossil fuels.* Consequently, there have been few systematic electrochemical investigations in molten sulphates3 This paper seeks to augment these studies and is especially concerned with the mechanism of the cathodic reduction of sulphate systems. This reduction process is, as suggested above, of particular interest in relation to the corrosive oxidation of fabricating metals by molten alkah sulphates in the temperature range 5oo-800°c.2 Evidence will be presented that the sulphate ion is not directly reduced at an inert metal cathode, and attention has therefore been directed to other routes, which may, for example, involve the equilibrium ELECTROCHEMICAL
so,=zz
so,
+ 02-.
This dissociation reaction is the basis of acid-base * Manuscript received 9 September 1968.
(1) equilibria
in sulphate
t Present address: Tyco Laboratories, Inc., Bear Hill, Waltham, Mass., U.S.A. 445
melts and
446
B. W. BURROWS and G. J. HILLS
can in principle be studied by determining
electrochemically
either the SOS or the
0% concentrations. Equilibrium (1) is evidently related to other equilibria, eg that of metal oxide solubility, of oxygen partial pressure, and the relative acidity of metal ions in solution. Factors such as the Ag/Ag(I) reference electrode, metal/metal-ion potentials, oxygen-electrode behaviour, the general applicability of voltammetric (including rapid scan) techniques, and the limiting cathodic process at ‘inert electrodes, have been investigated with the aim of elucidating the factors controlling the stability of metals in contact with molten sulphates. For completeness the limiting anodic processes in sulphate melts of different basicities have been studied, using both steadystate and transient galvanostatic techniques. EXPERIMENTAL
TECHNIQUE
A vertical, wire-wound furnace was used in all experiments. The furnace, its temperature control, and the cell assembly were all similar to those previously described for use in chloride melt$p6 except that the cell envelope was made of silica. To smooth out temperature gradients within the hot zone of the furnace, a grounded, thickwalled steel cylinder was inserted between the furnace tube and the cell assembly. This cylinder also served to reduce ac noise. The temperature was controlled to 625 & 1°C over a vertical length of 10 cm inside the cell assembly. A calibrated chromel/aIumel thermocouple, sheathed in Pyrex, was used to monitor the temperature in the melt during an experiment. Melt containers of platinum and of recrystallized ahunina were used. Their capacity was between 30 and 40 ml, and they normally contained about 20 ml of melt. Into these were inserted separate compartments consisting of quartz or Pyrex tubes (10 cm long x 1 cm i.d.) sealed at the lower ends by a fritted quartz or Pyrex disk (porosity 2-3). Each compartment contained about 4 ml of melt. The porous disks effectively prevented any significant intermixing of the contents of each compartment with that of the main cell, without introducing any electrical barrier or any appreciable liquid-junction potential. The various meta electrodes, of foil or wire, were of high purity metal from Johnson Matthey Ltd. The iron, cobalt and copper wire electrodes were l-mm diameter and dipped directly into the melt. The palladium and gold wire electrodes were 0.25-mm diameter and were coiled into spirals before being dipped into the melt. The silver wire electrode was of 0.5-mm diameter and was also in spiral form. The platinum counter-electrode was thin foil l-cm square, spot-welded to 0.25-mm diameter platinum wire. All the electrodes were loosely supported by Pyrex tubes through the head of the cell, and all electrical contacts between the electrodes and the potentiometer leads were made isothermally just above the cell assembly to avoid thermoelectric effects. The microelectrodes were made by sealing O.Ol-cm diameter Pt wire into Pyrex. l-cm needle-type microelectrodes were found to give more reproducible results than planar electrodes. They were submitted to a standard pre-treatment : they were heated in the oxidizing flame of a Bunsen, dipped in concentrated HCI, and reheated in the oxidizing flame. This treatment normally gave reproducible and well defined i/V curves and was used throughout the work. Tungsten was also tried as a micro-electrode material but was found to be reactive
Electrochemical
studies of molten alkali sulphates
447
in the sulphate melt, a black sludge being formed on the metal surface after brief immersion. ElectricuZmeasurements The chronopotentiometric measurements were derived from a simple constant current source (O-5-20 mA). The voltammetric measurements were recorded using either a pen-recording polarograph (scan rate 300 mV/min) or a laboratory-made, cathode-ray polarograph (scan rate continuously variable from 1 to 20 V/s). Preparation of the sulphace melt The L&$0,-J&SO, eutectic (80:20 mole-%; mp 535°C) was prepared by mixing Analar Li,SO, . H,O and Analar &SO4 and heating to 200°C. After a few hours at this temperature the mixture was assumed to be dry, since Li,SO, . HZ0 readily loses water at 130°C. The dehydrated mixture was then placed in a translucent silica beaker and fused in a muBe furnace at about 700°C. After fusion, the eutectic meIt was cooled, ground, and stored in a desiccator for further use. A sample of the eutectic melt was prepared for an experiment by placing about 30 g of the ground eutectic in a platinum or an alumina container. This was placed on the bottom of the cell enveIope, and fritted compartments were then placed on top of the solid eutectic. After the cell head had been attached, the assembly was evacuated, silicone high-vacuum grease being used to seal the flanges. The furnace was then raised slowly to enclose the cell in the hot zone. As the eutectic melted under vacuum, it gradually filtered into the f&ted compartments. The melt was then contained under a blanket of argon which prevented the oxidation of exposed metals. However, the solubility of oxygen is so low that for most purposes the melt can be exposed to air without any significant effect. The initial purity of the melt was checked voltammetricahy and was considered satisfactory if the residual current at the needle electrode was < 10 ,uA at - 1 V (Ag/Ag+). All experiments were carried out at 625 1. 1°C. Reference electrodes In accord with earlier studies, 6p7the silver/silver-sulphate electrode system was found to be a satisfactory reference electrode. It was isolated in a separate compartment by a fritted silica disk or by a glass (Jena Supremax) or silica membrane_ The use of Pyrex glass introduced difficulty: it rapidly developed a brown colouration and corresponding reference electrode potentials drifted with time. This phenomenon has been ascribed to exchange of silver ions with alkali ion of the glass, followed by subsequent reduction of silver atoms. ’ Experiments showed that evaporated silver gave the same colouration. Dissolved silver sulphate was generated by anodic dissolution of a silver wire. Concentrations of 5 x 10-s to 10-l molal were prepared and by comparison with another such electrode, the appropriate Nernst response i?AE/i? In m = RT/F was observed. The performance of reference electrodes in moIten salt systems is often relatively poor and limited to reproducibilities of f1 mV. One cause of failure may be discerned in the local recrystalhzations of the silver wire arising from small temperature and concentration gradients. Gentle continuous stirring of the reference eIectrode
3. W. BURROWS and G. J. H~LS
448
system by argon bubbling minimizes these inhomogeneities and enhances the reproducibility. Pairs of electrodes so maintained gave emfs reproducible to O-2 mV over long periods and there is no doubt that, wherever possible, reference electrode systems should be continuously equilibrated. The use of membrane electrodes was also satisfactory although persistent asymmetry potentials of up to 10 mV were often noted. These were sticiently stable not to introduce additional uncertainties, although it was noted that over a period of days the resistivity of the silica or Supremax membrane decreased steadily. RESULTS
AND
DISCUSSION
The results will be presented and discussed in three sections, viz emfmeasurements, cathodic processes and anodic processes.
Metal T;-tmetal-ion systems.
A series of studies were carried
out on cells of the
type where the electrolyte was the Li,SO&,SO, eutectic at 625°C. In the case of the Ag(I), Co(II), Cu(I), Pd(I1) and Fe(I1) systems, the metal ions were formed in situ by anodizing (at 5 mA/cm2) the metal electrode in a fritted compartment. The potentials of the metal electrodes after successive periods of anodizing and stirring were measured against an Ag/Ag(I) reference electrode. For each measured potential an experimental standard potential was calculated by use of the Nernst equation. The average values of these standard potentials together with the mean deviations are listed in Table 1. For comparison the values obtained by other workers are also shown; the agreement is good. It can therefore be concluded that, as with other molten salt systems, the activity coefficients of dilute solutions of metal ions are unity with respect to the ideal state of infinite dilution. The oxygen electrode. The oxygen electrode was investigated by observing the emf of both bright and platinized platinum electrodes (2 cm2) over which dry oxygen was bubbled at 2-5 bubbles/s. The oxide-ion concentration was varied by adding small quantities of CaO which is known to dissolve.s Attempts to reduce oxygen cathodically were not successful. TABLE
1. STANDARD
Couple
PdOWPd(O)
Rh(III)-Rh(0) Cu(IItCu(1) CufB-Cu(0) Co(ITt-Co(O) Fe(IJ)--Fe(O) * Li,SO~-KtSOI eutectic melt t Li,SO,-Na,SO&C,SO, eutectic melt
POTENTIALS
ON
THE
Present work* 625°C mV +530+4
-206 -714 -930
f 2 + 2 * 40
MOLAL
SCALE
IN SULPHATE
Liu6 625°C mV +541 f +387 f +51+3 -202 f
MELTS
Johnson and Laitinen’ 3 3 3
+518 +370 +55 -220 -689
mV =t 4 (580°C) & 4 (575°C) f 4 (575°C) rrt4 (58OOC) f 17 (550°C)
Electrochemical studies of molten alkali sulphates
449
The potential of both electrodes responded only sluggishly to changes in oxygen partial pressure. The eventual change in potential of the platinized platinum electrode on switching from air to pure oxygen was +31 mV (ie from +15 to f46 mV with respect of Ag/Ag(I)) in accord with the redox equilibrium
*O, + 2eT).02
(2)
and the Nernst relationship RT pg” E = E&-j2- + F In 2
D2-1
.
On the other hand, the bright platinum electrode had acorrespondingrest of -40 mV in air and +25 mV in pure oxygen. This suggests an alternative determining reaction, eg Pt + 0, + 2esPtO + 03-,
(3) potential potential (4)
which would predict a potential shift of +63 mV. It would also explain the initial difference in rest potentials as resulting from the cell reaction, Pt + 40, = PtO, the derived standard free energy of formation of PtO being -2.5 Kcal/mol. There is no independent value against which to check this figure, but it may be noted that the corresponding values for PdO and Rho at 625°C are -4.8 and -6-O Kcal/mol respectively.g The response of the electrodes to changes in oxide-ion concentration was immediate, but quite rapidly the negative potentials began to drift back to more positive values, until after an hour or so the original value was restored. It appeared that this reversal of potential resulted from the progressive removal of oxide ion by reaction with the silica of the cell compartment, SiO, + 02- +
SiOa2-.
(5)
The compartment was always badly etched at the end of the experiment and the buffering nature of the containers was further evidenced by addition of an acid stronger than silica (ie NaPO,), whereupon a more positive rest potential was obtained. In spite of the drift in potential, a reasonable estimate could be made of the observed Nernst slope, aE/a In mar-. The results of duplicate experiments illustrated in Fig. 1, suggest a Nernst slope of RT/F and not RT/2F as required by either of the This result was noted for both bright and equilibria (2) or (4) invoked above. platinized platinum electrodes and a similar response has been noted in both nitrater and chloridell melts. Attempts to account for this anomaly have been made in terms of hydrolysis reactions with residual water or of peroxide formation,l* but there is no evidence for either explanation here. Cathodic processes Metal/metal-ion systems. The reproducibility and apparent reversibility of the metal/metal-ion electrode systems was evidence of fast charge-transfer reactions. This was confirmed using cathode-ray voltammetry. Using fast voltage sweep rates (1 to 20 V/s), reproducible peak/shaped i/V curves were observed for the reduction of both Co(H) and Ag(I). Furthermore, linear plots of ii., versus z+‘~, C,,= and A were obtained in accordance with12 - = %
ffi3/2,$D1/2D
ox
,112
3
(6)
450
B. W. BURROWS and G. J. HILLS 210
log
Pm. 1. Potential
[0*-l
of Pt/Op electrode versus Iog [0*-l. bright Pt substrate.
Replicate
experiments
on
where iP is the peak current in A, n, A, D and C,,, have their usual meanings, D is the potential scan rate in V/s and Kis a constant, dependent on temperature, on the rate of the electrode process, and on the state of the product of the electrode reaction. Molal concentration was converted to molar concentration using the density data of James and Liu.ls Assuming the metals to be deposited at unit activity, the constant is given by l-08 P’2 K== ,$fZR”2T”a (7) ’ ie at 625”C, K = 2-l x lo6 C/mol/v112. This enabled the only remaining unknown in (6), the diffusion coefficient, to be determined. For Ag(I), the value of D was 3 x lo-6 cm2/s (42 x Iod for Cu(I), as determined chronopotentiometrically by Liu14). A corresponding evaluation of Dc,(,,, was not possible because the peak voltage was observed to change with sweep rate which, if it is evidence for charge-transfer control, requires a more complex formuIation of iP. Reduction of ml&hate species. A pure sulphate eutectic gave the current/voltage curve shown in Fig. 2, which apart from the anodic peak is similar to that recorded by Johnson and Laitinen .’ The anodic peak was found to be due to oxidation of 02- ions. The products of prolonged cathodic electrolysis were elementary sulphur, oxide ion, and sulphur dioxide. The sulphur was observed, and the SO, and 02were detected voltammetrically (see below). These products differ to some extent from those previously reported.6s7 No cathodic reduction of dissolved oxygen could be detected and no evidence was obtained of oxide ion resulting from reduced oxygen (small concentrations of added oxide were readily detected by anodic voltammetry).
Electrochemical studies of molten alkali sulphates
451
-20
.60 1.C B
0.8
0.4
0
-0.4 E*
--98
-1.2
-1.6
-2-o
-
2.4
V(Ag/Ag,‘Im)
FIG. 2. (a) Background i/r/curve at Pt needle electrode (area 3 x lo-* cml) in L&SO, &SO, at 625°C. (b) i/ Vcurve after addition of Na&O,. (c) PIot of limiting current in sdphate melt at 625°C vs temperature of SOI source (Na&O,). SuIphur trioxide was added to the melt as sodium pyrosulphate. thermally decomposes, Na,S,O, + Na,SO, + SO,,
This material (8)
to give an acid melt, Dissolved SO, could also be produced by addition of powdered silica or sodium metaphosphate. Oxides such as Co0 which were otherwise quite insoluble, dissolved immediately in acid melts. The presence of SO, gave rise to a reduction wave at -0.7 V (Ag/Ag(I)) (cf Fig. 2) and to an anodic dip at -0.3 V on the reverse potential scan, corresponding to the re-oxidation of the SO, so formed. Wrench and InmanX have recently reported the chronopotentiometric reduction of SO, in LiCI-KCI melts subsequent to the addition of metaphosphate. A quantitative study of the reduction of dissolved SO, proved difficult, since it rapidly disappeared from the system and its solubility is evidently negligible at ordinary partial pressures. An attempt was made to control the partial pressure of SO, by passage of argon over Na&O, (mp 401aC) in a separate furnace. The temperature was raised to 420°C and the transpired SOa led into the sulphate eutectic. After 30 min, a voltammogram was taken and the temperature of the source raised to give a new, known partial pressure of SO,. In fact, this experiment did not give very reliable results. Condensation of SO, (bp 44*5”C) occurred, and only at low partial pressures was it possible to record a sensible relation between the voltammetric limiting current and the source temperature (see Fig. 2). There is a further complication resulting from the dissociative equilibrium so, -;-t so, + HO,, (9) which is significant at 625°C. Sulphur dioxide was added in the form of anhydrous sodium sulphite. It decomposed rapidly to give some dissolved SO,, the voltammogram of which is shown in Fig. 3. The reduction peak at ca -1.5 V is considered to be characteristic of SO,,
8. W. BURROWS and G. J.
452
HILLS
although once again it was difficult to make a quantitative and reproducible study of this volatile and sparingly soluble component. The absence of any such reduction waves in the pure sulphate eutectic suggests that the dissociation of sulphate ions into sulphur trioxide and oxide is negligible in extent. That sulphate systems are ultimately reduced (eg by metals) may result from the promotion of this dissociation by associated acid species, such as metal ions (see An attempt to demonstrate the ultimate reducibility of sulphate ions at below). extreme cathodic potentials was made by studying solutions of sodium sulphate in the LiCl-KC1 eutectic, prepared and purified in the usual way.6*16 No cathodic wave attributable to sulphate was ever observed.
so2
SO-
60-
-2-o E.
FIG. 3. i/V curve
v
(Ag/Arrt
in Li,SO,K,SO,
-
‘4
Im)
melt at 625°C after bubbling SO*.
This observation is in agreement with that of Inman and Wrench16 but disagrees with those of SenderofF’ and of Woodall,ls who both reported that SOd2- could be directly reduced at a cathode in molten LiCl-KCl. Bukun and Ukshe,lQ on the other hand, have also reported evidence to suggest that SOd2- is not directly reducible. The observation that the S042- anion is not directly reducible suggests that the oxidizing character of sulphate melts must result from the promotion or acceleration of the dissociation of the sulphate ion. Evidence that this is so can be derived from the emf measurements. There is a “cut-off” value of -0.7 V (ie close to the reduction potential of SO, beyond which stable emfs were not observed. Instead, metals such as iron and nickel gave erratic potentials and other evidence of corrosion by SO,. The equilibria involved in any metal-sulphate system are s0*2-
*
M + SO,1-
so,
+ 02-,
(10)
SO, + MO,
(11)
MO z? M2+ + 02-,
(12)
Electrochemical studies of molten alkali sulphates
453
from which it is evident that at any particular state of oxidation, the acidity will be determined by the solubility of the metal oxide. Insoluble oxides necessarily relate to acidic metal ions. The corrosive oxidation of metaIs such as iron and nickel is stimulated not simply because their standard potentials are negative to that of SO,, but because they promote finite concentrations of SO, through equilibria of the type M2+ + SOaz- T, MO + SO,.
(131
Evidence in support of this hypothesi$ can be found in the thermal instabilities of temperatures of several such solid anhydrous metal sulphates. 2o The decomposition sulphates are Iisted in Table 2 and the relative acidity of the metal ions may be assessed TABLE
2.
Sulphate FeSO,
cuso, znso,
NISO,
COSO, M&O, PbSO, CdSO, MgSO, C=% SrSO,
DECOMPOSITION
TEMPERATURES
Decomposition temperature 537 598 646 675
708 755 803 816
895 1149 1374
OF SOME
DIVALENT
Observed intermediate product 2cuo. so, 3ZnO . 2SOII
METAL
SULPHATES**
Final non-volatile product Fe@, cue
ZnO
NiO Co,O. MU& PbO
Cd0 MgO CaO SrO
in terms of their readiness to form the metal oxide, ie to decompose thermally and displace SO,. Thus FeS04 has a decomposition temperature of 537”C, indicating that Fe2+ ions are quite acidic. The reaction between SO, and Fe is thermodynamically favorable,s Hence Fe is readily oxidized by SO,. Additional evidence of this mechanism can be adduced from the reported corrosiveness of aluminum and ferric iron sulphates22 and of zinc sulphateP towards Fe at are known to be thermally unstable at this 600°C. Both Al,(SO& and F+(SOJ, temperaturezl and ZnSO, is close to its decomposition temperature (see Table 2). Thus these sulphates are sources of SO, by virtue of an acid-base reaction. The evidence cited above strongly suggests that SO, is the oxidizing entity in molten suIphates and that their corrosiveness can be controlled onIy by repressing their acidity. Anodic processes
A series of steady-state overpotential measurements were made at bright platinumfoil eIectrodes in the pure sulphate eutectic and in solutions containing added oxide ion. CompIementary to these measurements were chronopotentiograms recorded at platinum-needle microelectrodes. Care was taken to minimize Ii? drop by using a three-electrode system and ensuring that the gap between the working electrode and the reference electrode was small. Oxygen was bubbled over the foil electrode during steady-state galvanostatic measurements to minimize concentration gradients. The aoodic oxidation of the pure sulphate eutectic is illustrated in Fig. 4. It can be seen from the experiments with decreasing current that there are two Tafel regions At cds less than ca O-6 mA/cm2, the slope corresponds to ca of overpotential. 4
B. W. BURROWSand G. J. HU
454
RT/2P and at higher cds to RT/J? Experiments with increasing currents gave poorly reproducible hysteresis effects, particularly at the higher cds. The over-potentials recorded during an experiment in a sulphate melt containing 9 x 10S2 m of oxide ion are plotted vs log I in Fig. 4. At low cds a Tafel region with a slope corresponding to RT/4F can be drawn through the experimental points. As the cd approached 1 mA/cm2, a limiting current was observed. A calculation based on the well known equation i, =
nFDC,s6
(14)
showed that this limiting current was not likely to be due to mass transport effects. By taking D = 1W6 cm2/s for the oxide ion, 6 = 5 x 1W2 cm for a well stirred melt, (b)
Slope
I
I
-0
-1-o
2.3RT 7
0
log ia (&
-2.0
I
-1.0
,
in mA/cms2)
FIG. 4. Anodic overpotentialat Pt/O, electrode in molten Li,SOcK,SOI eutectic at 625°C L)Slog i. The reference was an unpolarized Pt/O, electrode in the anolyte. (a) Neutral melt. (b) Basic melt.
Co,- = 1-S x l(r moIe/cm3 and n = 4 i, was calculated to be ca 14 mA/cm2, which is significantly higher than that observed, viz 1 mA/cm2. No potential/time dependence was noted in the basic melt. At each cd, a constant over-potential was quickly established, and open-circuit decay of 9 to near zero was rapid. Furthermore, no hysteresis was observed but, as in the pure meIt, the pIatinum electrode became tarnished by a film the thickness of which increased with time of electrolysis and with oxide concentration. In appearance it varied from speckIed dark grey to a more coherent dark metallic grey. It was assumed to be platinum oxide24*26and readily disappeared when heated. Evidence suggesting the presence of an oxide film was obtained from currentreversal chronopotentiograms at low cds. From these, both the anodic formation and cathodic removal of an oxide layer are clear (Fig. 5). Anodic formation was observed only on precathodized electrodes, suggesting that a film of oxide (most likely PtO) is normally present on a bright Pt electrode in a sulphate melt. The emf behaviour of
Electrochemical bright
and platinized
electrodes
studies of molten alkali sulphates
(see above)
also points to the existence
455
of a film of
PtO. The chronopotentiometric observation can be quantified by comparing the charge consumed during either formation or removal of the film (it is the same in both cases) with that calculated for a monolayer of PtO. From the lattice constants of crystalline platinum, it has been estimated 26 that there are l-5 x 10L6 atom/cm2 on the preUsing this value it can be calculated that approxidominant (111) crystal face. mately O-5 x 10-s C/cm2 are required to provide each surface atom of platinum From several of the origina galvanostatic transients it was with an oxygen atom. determined that O-3-08 x 103 C/cm2 were consumed during the oxidation, in reasonable agreement with the calculated value. o-4
o-t-
-o.3o
I I
I 4
I 3
1 2 Tinu.
I 5
i
I
FIG. 5. Chronopotentiogram showing formation and removal of a monolayer of PtO in LipSO,-KISOI eutectic melt at 625°C. 1 mA/cm*.
At high cds in the pure melt (> I mA/cma), a further transition time was observed (Fig. 6), which increased in proportion to the time of cathodic pretreatment. The potential/time trace for this process was sufficiently well defined for the potential E to be plotted against (rlla - +‘2/@z), which plot according to the relation2’
E = Ei +sln
(Zmt”l’=‘2)
(15)
should be linear. As is seen in Fig. 6, this is so and an = 2-9 for the charge-transfer reaction involved. The nature of the process involved in this second transition time is evidently the partially irreversible oxidation of dissolved oxide ions present as a residual impurity. Using the Sand equation, the concentration of residual oxide in the pure sulphate melt was estimated to be 4 x lo-8 m. Thus, transition between the TafeI regions in
B. W. BURROWS
456
0
I
I
0.5
I.0
and G. J. IHILLS
I
Time.
I
I
I.5
CO
2-5
s
FIG. 6. (a) Chronopotentiogram in l&SO,-K,SO, eutectic melt at 625°C showing wave due to oxidation of OS-. 1 mA/cm”. (b) An a lysis of the chronopotentiometric wave for the oxidation of Oz-.
presumabiy corresponds to a change in mechanism from 0% discharge to If this is the case, the transition cd should correspond to a massSO,” discharge. transport-controhed limiting cd. This supposition was checked by calculating i, from (14) taking Co,- = 8 x 10” mol/cm3. A value of ii ca O-7 mA/cm” was obtained, in cIose agreement with the observed value of ca 0.6 mA/cm2. The degree of oxidation of the substrate electrode and its degree of coverage with oxide or adsorbed oxygen are important factors that influence the detailed mechanism of anodic processes in soIutions containing oxygenated species. It is theoretically possible to relate the observed Tafel parameters to the mechanism. Bockris,2s for example, has extended the general theory of consecutive reactions developed by and applied it specifically to the oxygen-evolution reaction. Christiansen2s The treatment has been applied in a similar manner to anodic processes in molten carand cryoliteahuninas2 systems. It is somewhat simpler to apply bonate,30 nitraW in ionic melts, since the number of possible reaction steps can be reduced, due to the absence of aqueous solvent. In a sulphate melt the observation that a surface layer of PtO is probably present on the electrode (see above) further allows the number of possible reaction schemes to be reduced until only the folIowing appear feasible,
Fig.
4
(a) (b)
PtO + 02-* 2Pt0,&2Pto
PtO,
+ 2e, + 0,.
(16)
The Tafel slopes, predicted when either (a) or (b) is the slow step, are summarized in TabIe 3. Langmuir behaviour is assumed for the adsorption of the reaction intermediate, PtO,.
Electrochemical TABLET.
ANALYSIS
OF THE PROPOSED
LiBS06KpS0,
Rate-determining (a) PtO + O*- * (b)
studies of molten alkali sulphates ANODIC REACTION AT 625”C*
457
MECHANISM
step
Predicted tafel slope
PtO, + 2e
RT 2(1 - a)F
2ptoa *2PtO
+ 02
g
for e,,o, -
il for L&o~ -
IN
0 1
In Table 3, ir represents a limiting cd caused by the rate controlling activation reaction having reached a constant maximum at &oa = 1. Considering the q/log i behaviour in a basic sulphate melt, the Tafel slope of RT/4F and the limiting cd previously shown not to be controlled by diffusion in solution can now be understood in terms of mechanism (16) where (b) is the ratelimiting cd has previously been determining step. A similar activation-controlled reported for the hydrogen-evolution reaction on palladium electrodes in 5 N aqueous W-X= In a pure sulphate melt it has already been suggested that the two observed Tafel regions (Fig. 4) correspond to the successive discharge of Oa- ions and SOg2- ions at low and high cds respectively, with a transition at CLI0.6 mA/cmZ corresponding to a diffusion-controlled limiting cd. Arkhipov, Trunov and Stepanov3Q came to a similar conclusion regarding the nature of the anodic process taking place in carbonate melts. The discharge of SOq2- possibly proceeds through a mechanism similar to (16) with intermediate formation of PtU,. A theoretical analysis for this case also leads to the same predicted Tafel slopes. Therefore, with the reaction PtO. -j- SOq2- *
PtO,
+ SO3 + 2 e
(17)
as the rate-controlling step and with (1 - 0~) = 0.5, a TafeI slope of RT/p is predicted, in accordance with the observed value. Furthermore, it can be predicted that if measurements were made at sufficiently high cds than an activation-controlled limiting cd would be observed, corresponding to complete coverage of the electrode by PtO,. It is likely that the discharge of oxide ion in a “pure” sulphate melt also proceeds through a PtO, intermediate, but in contrast to more basic melts, step (a) of reaction (16) must be rate-determining. The only previous measurements of anodic polarization behaviour in sulphate melts are those of Flood and Fsrland. 35 These workers investigated the overpotential associated with the anodic evolution of O2 in sulphate and carbonate melts and in fluoride melts containing added sulphate, carbonate, phosphate and silicate, over the temperature range 6561000°C. The investigation was designed to study the possibility of using the O,/Pt electrode to measure oxide-ion activities; consequently, the systems were buffered with respect to these ions. Steady-state galvanostatic measurements showed that the various systems had some common characteristics. At low cds and high temperatures, Tafel slopes of RT/3F were obtained, whereas at high cds and low temperatures, the slopes were RT/F. Because the temperature coefficient of overpotential was lower by a factor of three in the former case (-0-4 mV/Y), it was concluded that the overpotential was mass-transfer controlled. For Tafel slopes
B. W.
458
corresponding order
to RT/F,
BURROWS and G. J. HI-
the activation so42-
> PO,*
overpotential
in the
> COs2- > SiO,&,
parallel with the strength of the X-O Strict comparison of these results with although it appears that the Tafel slope of present investigation in sulphate melts not
ie in
was found to decrease
bond. those of the present case is not possible, RT/F corresponds to that obtained in the containing added oxide.
authors wish to express their thanks to the Central Electricity Generating Board for Financial assistance, including a Junior Fellowship for one of them (B. W. B.).
Acknowledgements-The
REFERENCES 1. A. C. GRAVES, G. J. HIand D. INMAN, in Advances in Electrochemistry and Electrochemical Engineering, Vol. 4, ed. P. Delahay and C. W. Tobias. Interscience, New York (1966). 2. B. W. BURROWSand G. J. HILLS, J. Inst. Fuel, 168 (1966). 3. B. W. BURROWS, Thesis, University of Southampton (1965). 4. G. J. HILLS and K. E. JOHNSON,J. electrochem. Sot. 108, 1013 (1961). 5. H. A. LATTINEN,W. S. FERGU~ONand R. A. OSTERYOUT*‘G, J. electrochem. Sot. 104,516 (1957). 6. C. H. Lru, J. Am. them. Sot. 66,164 (1962). 7. K. E. JOHNSON and H. A. LAITINEN,J. electruchem. Sot. 110, 314 (1963). 8. D. G. HILL, B. PORTERand A. S. GILLESPIE,J. electrochem. Sot. 105,408 (1958). 9, J. P. cm.JGHLIN, Bull. U.S. Bur. Mines, 542 (1954). 10. A. M. SHAMSELDIN and A. GERGES,Proc. First Aust. Conf. Electrochemistry, p. 562. Pergamon Press, Oxford (1965). 11. R. LI-I-IUWOOD and E. J. ARGENT, Electrochim. Actu. 4, 114 (1961). 12. T. BERZINSand P. DELAHAY, J. Am. ckm. Sot. 75, 555 (1953). 13. D. W. JAMESand C. H. Lru, J. them. Eng. Data 8,469 (1963). 14. C. H. LIU, Analyt. Chem. 33, 1477 (1961). 15. D. L. MARICLE and D. N. HUME, J. electrochem. Sot. 107, 354 (1960). 16. D. M. WRENCH and D. IAN, Electrochim. Acta 12, 1601 (1967). 17. S. SENDEROFP,E. M. KLOPP and M. L. KRONENBERG,Eighth and Ninth Quarterly Technical Progress Reports, Contract NORD-18240 (1960). 18. B. WOODALL, quoted by H. A. LAITINEN, Tafunta 12, 1237 (1965). 19. N. G. BUKUN and E. A. UKSHE, Chem. Abst. 60,1028g, (1964). 20. A. G. OSTROFFand R. T. SANDERSON,J. inorg. nucl. Chem. 9,45 (1959). 21. N. A. WARNER and T. R. INGRAWAMS,Can. J. Chem. 38,2196 (1960). 22. B. K. TORPAN and KH. KH. KHYDREYRV, Chem. Abstr. 54,4313 (1960). 23. V. P. KOCHERGINet al, C&m. A&r. 58, 297a (1963). 24. R. L. EVERY and R. L. GRIMSLEY,J. electroana!. Chem. 9, 165 (1965). 25. G. K. STEPANOV and A. M. TRUNOV, Dokl. Akad. Nauk SSSR 142,866 (1962). 26. H. A. LAITINENand C. G. ENKE, J. electrochem. Sot. 107, 773 (1960). 27. P. DELAHAY, New Instrumental Methods in Electrochemistry, p_ 187. Interscience, New York (1954). 28. J. O’M. B~CKRIS, J. phys. Chem. 24, 817 (1956). 29. A. CHRISTIANSEN. 2. phys. Chem. B33, 145 (1936). 30. G. J. JANZ and F. SAEGUSA, J. electrochem. Sot. 108, 663 (1961). 31. W. E. TRIACA and A. J. ARV~A, Hectrochim. Acta. 10,409 (1965). 32. N. E. RICHARDSend B. J. WELCH, Pruc. First Aust. Co@ Electrochemistry, p. 901. Pergamon Press, Oxford (1965). 33. J. O’M. BOCKRIS, and A. M. AZZAM, Trans. Faraday Sot. 48,145 (1952). 34. G. G. ARKHIPOV, A. M. TRUNOV and G. K. STEPANOV, Trans. Insr. Electrochim. Akud. Nauk SSSR, Uraf’sk. Filial 4,41 (1963); Chem. Abstr. 60, 14124g (1964). 35. H. FLOOD and T. FIZIRLAND,Discuss. Faraday Sot, 1, 302 (1947).
Pergamon Prays.Printedin NorthernIreland
ELECTROCHEMICAL ALKALI
STUDIES OF MOLTEN SULPHATES*
B. W. BURROWS~ and G. J. HILLS Chemistry Department, University of Southampton, Southampton, England Abstract-Electrochemical studies in the Li,SOI-K,SOI eutectic melt at 625°C are reported. Experimental evidence suggests that the sulphate ion is not directly reduced at an inert metal cathode. The corrosion of metals such as Fe and Ni in molten sulphates appears to result from reaction with SOS derived from dissociation of SOI*- anions. This dissociation is enhanced by the relative acidity of the metal-ion species. The limiting anodic process was studied by steady state and transient galvanostatic techniques in melts of different basicities. Discharge mechanisms are proposed to explain the observed behaviour. R&sum&-On pr&scnte des etudes electrochimique dun entectique fondu de Li,SO,-K,SOI g 625°C. Avec l’evidence experimentale on propose que l’ion de sulphate n’est pas directement I-eduit & la cathode metallique inerte. La corrosion des m&aux comme le Fe et le Ni darts les sulphates fondus peut r&ulter dune reaction avec le SOa qui vient de la dissociation des anions SO,s-. Cette dissociation est encouraged par l’acidite relative aux esp&cesd’ions m&alliques. On a etudie le processus anodique limitant par des techniques galvanostatiques stationaires et transitoires pour des sels fondus de bacicities differents. On propose des mechanisms de d&charge avec lesquelles on explique les fawns de se cornporter que l’on observe. Zusammenfassung-Elektrochemische Untersuchungen an der eutektischen Schmelze von LirSOaI&SO, bei 625°C werden beschrieben. Die experimentellen Ergebnisse weisen darauf hin, dass das Sulfation an inerten Metallkathoden nicht direkt reduziert wird. Die ICorrosion von Metallen wie Fe und Ni in Sulfatschmelzen scheint auf eine Reaktion mit SO, zuriickzufiihren sein, welches durch die Dissoziation von SO&-* entsteht. Diese Dissoziation wird durch die relative Aziditat des metallischen Ions begiinstigt. Der begrenzende anodische Vorgang wurde mittles Gleichgewichtsund nicht-stationf+.rergalvanostatischer Methoden in Schmelzen verschiedener Basizitiit untersucht. Zur ErklHrung des beobachteten Verhaltens werden Entladungsmechanismen vorgeschlagen. INTRODUCTION
studies in oxyanion melts have been largely confined to nitrate and carbonate systems. l The former are low-melting and thus convenient media to work with, while the latter have considerable technological importance as electrolytes for high temperature fuel cells. Sulphate melts, on the other hand, are relatively high-melting, like the carbonates, but have little practical importance apart from their deleterious effects on boiler tubes in power stations burning fossil fuels.* Consequently, there have been few systematic electrochemical investigations in molten sulphates3 This paper seeks to augment these studies and is especially concerned with the mechanism of the cathodic reduction of sulphate systems. This reduction process is, as suggested above, of particular interest in relation to the corrosive oxidation of fabricating metals by molten alkah sulphates in the temperature range 5oo-800°c.2 Evidence will be presented that the sulphate ion is not directly reduced at an inert metal cathode, and attention has therefore been directed to other routes, which may, for example, involve the equilibrium ELECTROCHEMICAL
so,=zz
so,
+ 02-.
This dissociation reaction is the basis of acid-base * Manuscript received 9 September 1968.
(1) equilibria
in sulphate
t Present address: Tyco Laboratories, Inc., Bear Hill, Waltham, Mass., U.S.A. 445
melts and
446
B. W. BURROWS and G. J. HILLS
can in principle be studied by determining
electrochemically
either the SOS or the
0% concentrations. Equilibrium (1) is evidently related to other equilibria, eg that of metal oxide solubility, of oxygen partial pressure, and the relative acidity of metal ions in solution. Factors such as the Ag/Ag(I) reference electrode, metal/metal-ion potentials, oxygen-electrode behaviour, the general applicability of voltammetric (including rapid scan) techniques, and the limiting cathodic process at ‘inert electrodes, have been investigated with the aim of elucidating the factors controlling the stability of metals in contact with molten sulphates. For completeness the limiting anodic processes in sulphate melts of different basicities have been studied, using both steadystate and transient galvanostatic techniques. EXPERIMENTAL
TECHNIQUE
A vertical, wire-wound furnace was used in all experiments. The furnace, its temperature control, and the cell assembly were all similar to those previously described for use in chloride melt$p6 except that the cell envelope was made of silica. To smooth out temperature gradients within the hot zone of the furnace, a grounded, thickwalled steel cylinder was inserted between the furnace tube and the cell assembly. This cylinder also served to reduce ac noise. The temperature was controlled to 625 & 1°C over a vertical length of 10 cm inside the cell assembly. A calibrated chromel/aIumel thermocouple, sheathed in Pyrex, was used to monitor the temperature in the melt during an experiment. Melt containers of platinum and of recrystallized ahunina were used. Their capacity was between 30 and 40 ml, and they normally contained about 20 ml of melt. Into these were inserted separate compartments consisting of quartz or Pyrex tubes (10 cm long x 1 cm i.d.) sealed at the lower ends by a fritted quartz or Pyrex disk (porosity 2-3). Each compartment contained about 4 ml of melt. The porous disks effectively prevented any significant intermixing of the contents of each compartment with that of the main cell, without introducing any electrical barrier or any appreciable liquid-junction potential. The various meta electrodes, of foil or wire, were of high purity metal from Johnson Matthey Ltd. The iron, cobalt and copper wire electrodes were l-mm diameter and dipped directly into the melt. The palladium and gold wire electrodes were 0.25-mm diameter and were coiled into spirals before being dipped into the melt. The silver wire electrode was of 0.5-mm diameter and was also in spiral form. The platinum counter-electrode was thin foil l-cm square, spot-welded to 0.25-mm diameter platinum wire. All the electrodes were loosely supported by Pyrex tubes through the head of the cell, and all electrical contacts between the electrodes and the potentiometer leads were made isothermally just above the cell assembly to avoid thermoelectric effects. The microelectrodes were made by sealing O.Ol-cm diameter Pt wire into Pyrex. l-cm needle-type microelectrodes were found to give more reproducible results than planar electrodes. They were submitted to a standard pre-treatment : they were heated in the oxidizing flame of a Bunsen, dipped in concentrated HCI, and reheated in the oxidizing flame. This treatment normally gave reproducible and well defined i/V curves and was used throughout the work. Tungsten was also tried as a micro-electrode material but was found to be reactive
Electrochemical
studies of molten alkali sulphates
447
in the sulphate melt, a black sludge being formed on the metal surface after brief immersion. ElectricuZmeasurements The chronopotentiometric measurements were derived from a simple constant current source (O-5-20 mA). The voltammetric measurements were recorded using either a pen-recording polarograph (scan rate 300 mV/min) or a laboratory-made, cathode-ray polarograph (scan rate continuously variable from 1 to 20 V/s). Preparation of the sulphace melt The L&$0,-J&SO, eutectic (80:20 mole-%; mp 535°C) was prepared by mixing Analar Li,SO, . H,O and Analar &SO4 and heating to 200°C. After a few hours at this temperature the mixture was assumed to be dry, since Li,SO, . HZ0 readily loses water at 130°C. The dehydrated mixture was then placed in a translucent silica beaker and fused in a muBe furnace at about 700°C. After fusion, the eutectic meIt was cooled, ground, and stored in a desiccator for further use. A sample of the eutectic melt was prepared for an experiment by placing about 30 g of the ground eutectic in a platinum or an alumina container. This was placed on the bottom of the cell enveIope, and fritted compartments were then placed on top of the solid eutectic. After the cell head had been attached, the assembly was evacuated, silicone high-vacuum grease being used to seal the flanges. The furnace was then raised slowly to enclose the cell in the hot zone. As the eutectic melted under vacuum, it gradually filtered into the f&ted compartments. The melt was then contained under a blanket of argon which prevented the oxidation of exposed metals. However, the solubility of oxygen is so low that for most purposes the melt can be exposed to air without any significant effect. The initial purity of the melt was checked voltammetricahy and was considered satisfactory if the residual current at the needle electrode was < 10 ,uA at - 1 V (Ag/Ag+). All experiments were carried out at 625 1. 1°C. Reference electrodes In accord with earlier studies, 6p7the silver/silver-sulphate electrode system was found to be a satisfactory reference electrode. It was isolated in a separate compartment by a fritted silica disk or by a glass (Jena Supremax) or silica membrane_ The use of Pyrex glass introduced difficulty: it rapidly developed a brown colouration and corresponding reference electrode potentials drifted with time. This phenomenon has been ascribed to exchange of silver ions with alkali ion of the glass, followed by subsequent reduction of silver atoms. ’ Experiments showed that evaporated silver gave the same colouration. Dissolved silver sulphate was generated by anodic dissolution of a silver wire. Concentrations of 5 x 10-s to 10-l molal were prepared and by comparison with another such electrode, the appropriate Nernst response i?AE/i? In m = RT/F was observed. The performance of reference electrodes in moIten salt systems is often relatively poor and limited to reproducibilities of f1 mV. One cause of failure may be discerned in the local recrystalhzations of the silver wire arising from small temperature and concentration gradients. Gentle continuous stirring of the reference eIectrode
3. W. BURROWS and G. J. H~LS
448
system by argon bubbling minimizes these inhomogeneities and enhances the reproducibility. Pairs of electrodes so maintained gave emfs reproducible to O-2 mV over long periods and there is no doubt that, wherever possible, reference electrode systems should be continuously equilibrated. The use of membrane electrodes was also satisfactory although persistent asymmetry potentials of up to 10 mV were often noted. These were sticiently stable not to introduce additional uncertainties, although it was noted that over a period of days the resistivity of the silica or Supremax membrane decreased steadily. RESULTS
AND
DISCUSSION
The results will be presented and discussed in three sections, viz emfmeasurements, cathodic processes and anodic processes.
Metal T;-tmetal-ion systems.
A series of studies were carried
out on cells of the
type where the electrolyte was the Li,SO&,SO, eutectic at 625°C. In the case of the Ag(I), Co(II), Cu(I), Pd(I1) and Fe(I1) systems, the metal ions were formed in situ by anodizing (at 5 mA/cm2) the metal electrode in a fritted compartment. The potentials of the metal electrodes after successive periods of anodizing and stirring were measured against an Ag/Ag(I) reference electrode. For each measured potential an experimental standard potential was calculated by use of the Nernst equation. The average values of these standard potentials together with the mean deviations are listed in Table 1. For comparison the values obtained by other workers are also shown; the agreement is good. It can therefore be concluded that, as with other molten salt systems, the activity coefficients of dilute solutions of metal ions are unity with respect to the ideal state of infinite dilution. The oxygen electrode. The oxygen electrode was investigated by observing the emf of both bright and platinized platinum electrodes (2 cm2) over which dry oxygen was bubbled at 2-5 bubbles/s. The oxide-ion concentration was varied by adding small quantities of CaO which is known to dissolve.s Attempts to reduce oxygen cathodically were not successful. TABLE
1. STANDARD
Couple
PdOWPd(O)
Rh(III)-Rh(0) Cu(IItCu(1) CufB-Cu(0) Co(ITt-Co(O) Fe(IJ)--Fe(O) * Li,SO~-KtSOI eutectic melt t Li,SO,-Na,SO&C,SO, eutectic melt
POTENTIALS
ON
THE
Present work* 625°C mV +530+4
-206 -714 -930
f 2 + 2 * 40
MOLAL
SCALE
IN SULPHATE
Liu6 625°C mV +541 f +387 f +51+3 -202 f
MELTS
Johnson and Laitinen’ 3 3 3
+518 +370 +55 -220 -689
mV =t 4 (580°C) & 4 (575°C) f 4 (575°C) rrt4 (58OOC) f 17 (550°C)
Electrochemical studies of molten alkali sulphates
449
The potential of both electrodes responded only sluggishly to changes in oxygen partial pressure. The eventual change in potential of the platinized platinum electrode on switching from air to pure oxygen was +31 mV (ie from +15 to f46 mV with respect of Ag/Ag(I)) in accord with the redox equilibrium
*O, + 2eT).02
(2)
and the Nernst relationship RT pg” E = E&-j2- + F In 2
D2-1
.
On the other hand, the bright platinum electrode had acorrespondingrest of -40 mV in air and +25 mV in pure oxygen. This suggests an alternative determining reaction, eg Pt + 0, + 2esPtO + 03-,
(3) potential potential (4)
which would predict a potential shift of +63 mV. It would also explain the initial difference in rest potentials as resulting from the cell reaction, Pt + 40, = PtO, the derived standard free energy of formation of PtO being -2.5 Kcal/mol. There is no independent value against which to check this figure, but it may be noted that the corresponding values for PdO and Rho at 625°C are -4.8 and -6-O Kcal/mol respectively.g The response of the electrodes to changes in oxide-ion concentration was immediate, but quite rapidly the negative potentials began to drift back to more positive values, until after an hour or so the original value was restored. It appeared that this reversal of potential resulted from the progressive removal of oxide ion by reaction with the silica of the cell compartment, SiO, + 02- +
SiOa2-.
(5)
The compartment was always badly etched at the end of the experiment and the buffering nature of the containers was further evidenced by addition of an acid stronger than silica (ie NaPO,), whereupon a more positive rest potential was obtained. In spite of the drift in potential, a reasonable estimate could be made of the observed Nernst slope, aE/a In mar-. The results of duplicate experiments illustrated in Fig. 1, suggest a Nernst slope of RT/F and not RT/2F as required by either of the This result was noted for both bright and equilibria (2) or (4) invoked above. platinized platinum electrodes and a similar response has been noted in both nitrater and chloridell melts. Attempts to account for this anomaly have been made in terms of hydrolysis reactions with residual water or of peroxide formation,l* but there is no evidence for either explanation here. Cathodic processes Metal/metal-ion systems. The reproducibility and apparent reversibility of the metal/metal-ion electrode systems was evidence of fast charge-transfer reactions. This was confirmed using cathode-ray voltammetry. Using fast voltage sweep rates (1 to 20 V/s), reproducible peak/shaped i/V curves were observed for the reduction of both Co(H) and Ag(I). Furthermore, linear plots of ii., versus z+‘~, C,,= and A were obtained in accordance with12 - = %
ffi3/2,$D1/2D
ox
,112
3
(6)
450
B. W. BURROWS and G. J. HILLS 210
log
Pm. 1. Potential
[0*-l
of Pt/Op electrode versus Iog [0*-l. bright Pt substrate.
Replicate
experiments
on
where iP is the peak current in A, n, A, D and C,,, have their usual meanings, D is the potential scan rate in V/s and Kis a constant, dependent on temperature, on the rate of the electrode process, and on the state of the product of the electrode reaction. Molal concentration was converted to molar concentration using the density data of James and Liu.ls Assuming the metals to be deposited at unit activity, the constant is given by l-08 P’2 K== ,$fZR”2T”a (7) ’ ie at 625”C, K = 2-l x lo6 C/mol/v112. This enabled the only remaining unknown in (6), the diffusion coefficient, to be determined. For Ag(I), the value of D was 3 x lo-6 cm2/s (42 x Iod for Cu(I), as determined chronopotentiometrically by Liu14). A corresponding evaluation of Dc,(,,, was not possible because the peak voltage was observed to change with sweep rate which, if it is evidence for charge-transfer control, requires a more complex formuIation of iP. Reduction of ml&hate species. A pure sulphate eutectic gave the current/voltage curve shown in Fig. 2, which apart from the anodic peak is similar to that recorded by Johnson and Laitinen .’ The anodic peak was found to be due to oxidation of 02- ions. The products of prolonged cathodic electrolysis were elementary sulphur, oxide ion, and sulphur dioxide. The sulphur was observed, and the SO, and 02were detected voltammetrically (see below). These products differ to some extent from those previously reported.6s7 No cathodic reduction of dissolved oxygen could be detected and no evidence was obtained of oxide ion resulting from reduced oxygen (small concentrations of added oxide were readily detected by anodic voltammetry).
Electrochemical studies of molten alkali sulphates
451
-20
.60 1.C B
0.8
0.4
0
-0.4 E*
--98
-1.2
-1.6
-2-o
-
2.4
V(Ag/Ag,‘Im)
FIG. 2. (a) Background i/r/curve at Pt needle electrode (area 3 x lo-* cml) in L&SO, &SO, at 625°C. (b) i/ Vcurve after addition of Na&O,. (c) PIot of limiting current in sdphate melt at 625°C vs temperature of SOI source (Na&O,). SuIphur trioxide was added to the melt as sodium pyrosulphate. thermally decomposes, Na,S,O, + Na,SO, + SO,,
This material (8)
to give an acid melt, Dissolved SO, could also be produced by addition of powdered silica or sodium metaphosphate. Oxides such as Co0 which were otherwise quite insoluble, dissolved immediately in acid melts. The presence of SO, gave rise to a reduction wave at -0.7 V (Ag/Ag(I)) (cf Fig. 2) and to an anodic dip at -0.3 V on the reverse potential scan, corresponding to the re-oxidation of the SO, so formed. Wrench and InmanX have recently reported the chronopotentiometric reduction of SO, in LiCI-KCI melts subsequent to the addition of metaphosphate. A quantitative study of the reduction of dissolved SO, proved difficult, since it rapidly disappeared from the system and its solubility is evidently negligible at ordinary partial pressures. An attempt was made to control the partial pressure of SO, by passage of argon over Na&O, (mp 401aC) in a separate furnace. The temperature was raised to 420°C and the transpired SOa led into the sulphate eutectic. After 30 min, a voltammogram was taken and the temperature of the source raised to give a new, known partial pressure of SO,. In fact, this experiment did not give very reliable results. Condensation of SO, (bp 44*5”C) occurred, and only at low partial pressures was it possible to record a sensible relation between the voltammetric limiting current and the source temperature (see Fig. 2). There is a further complication resulting from the dissociative equilibrium so, -;-t so, + HO,, (9) which is significant at 625°C. Sulphur dioxide was added in the form of anhydrous sodium sulphite. It decomposed rapidly to give some dissolved SO,, the voltammogram of which is shown in Fig. 3. The reduction peak at ca -1.5 V is considered to be characteristic of SO,,
8. W. BURROWS and G. J.
452
HILLS
although once again it was difficult to make a quantitative and reproducible study of this volatile and sparingly soluble component. The absence of any such reduction waves in the pure sulphate eutectic suggests that the dissociation of sulphate ions into sulphur trioxide and oxide is negligible in extent. That sulphate systems are ultimately reduced (eg by metals) may result from the promotion of this dissociation by associated acid species, such as metal ions (see An attempt to demonstrate the ultimate reducibility of sulphate ions at below). extreme cathodic potentials was made by studying solutions of sodium sulphate in the LiCl-KC1 eutectic, prepared and purified in the usual way.6*16 No cathodic wave attributable to sulphate was ever observed.
so2
SO-
60-
-2-o E.
FIG. 3. i/V curve
v
(Ag/Arrt
in Li,SO,K,SO,
-
‘4
Im)
melt at 625°C after bubbling SO*.
This observation is in agreement with that of Inman and Wrench16 but disagrees with those of SenderofF’ and of Woodall,ls who both reported that SOd2- could be directly reduced at a cathode in molten LiCl-KCl. Bukun and Ukshe,lQ on the other hand, have also reported evidence to suggest that SOd2- is not directly reducible. The observation that the S042- anion is not directly reducible suggests that the oxidizing character of sulphate melts must result from the promotion or acceleration of the dissociation of the sulphate ion. Evidence that this is so can be derived from the emf measurements. There is a “cut-off” value of -0.7 V (ie close to the reduction potential of SO, beyond which stable emfs were not observed. Instead, metals such as iron and nickel gave erratic potentials and other evidence of corrosion by SO,. The equilibria involved in any metal-sulphate system are s0*2-
*
M + SO,1-
so,
+ 02-,
(10)
SO, + MO,
(11)
MO z? M2+ + 02-,
(12)
Electrochemical studies of molten alkali sulphates
453
from which it is evident that at any particular state of oxidation, the acidity will be determined by the solubility of the metal oxide. Insoluble oxides necessarily relate to acidic metal ions. The corrosive oxidation of metaIs such as iron and nickel is stimulated not simply because their standard potentials are negative to that of SO,, but because they promote finite concentrations of SO, through equilibria of the type M2+ + SOaz- T, MO + SO,.
(131
Evidence in support of this hypothesi$ can be found in the thermal instabilities of temperatures of several such solid anhydrous metal sulphates. 2o The decomposition sulphates are Iisted in Table 2 and the relative acidity of the metal ions may be assessed TABLE
2.
Sulphate FeSO,
cuso, znso,
NISO,
COSO, M&O, PbSO, CdSO, MgSO, C=% SrSO,
DECOMPOSITION
TEMPERATURES
Decomposition temperature 537 598 646 675
708 755 803 816
895 1149 1374
OF SOME
DIVALENT
Observed intermediate product 2cuo. so, 3ZnO . 2SOII
METAL
SULPHATES**
Final non-volatile product Fe@, cue
ZnO
NiO Co,O. MU& PbO
Cd0 MgO CaO SrO
in terms of their readiness to form the metal oxide, ie to decompose thermally and displace SO,. Thus FeS04 has a decomposition temperature of 537”C, indicating that Fe2+ ions are quite acidic. The reaction between SO, and Fe is thermodynamically favorable,s Hence Fe is readily oxidized by SO,. Additional evidence of this mechanism can be adduced from the reported corrosiveness of aluminum and ferric iron sulphates22 and of zinc sulphateP towards Fe at are known to be thermally unstable at this 600°C. Both Al,(SO& and F+(SOJ, temperaturezl and ZnSO, is close to its decomposition temperature (see Table 2). Thus these sulphates are sources of SO, by virtue of an acid-base reaction. The evidence cited above strongly suggests that SO, is the oxidizing entity in molten suIphates and that their corrosiveness can be controlled onIy by repressing their acidity. Anodic processes
A series of steady-state overpotential measurements were made at bright platinumfoil eIectrodes in the pure sulphate eutectic and in solutions containing added oxide ion. CompIementary to these measurements were chronopotentiograms recorded at platinum-needle microelectrodes. Care was taken to minimize Ii? drop by using a three-electrode system and ensuring that the gap between the working electrode and the reference electrode was small. Oxygen was bubbled over the foil electrode during steady-state galvanostatic measurements to minimize concentration gradients. The aoodic oxidation of the pure sulphate eutectic is illustrated in Fig. 4. It can be seen from the experiments with decreasing current that there are two Tafel regions At cds less than ca O-6 mA/cm2, the slope corresponds to ca of overpotential. 4
B. W. BURROWSand G. J. HU
454
RT/2P and at higher cds to RT/J? Experiments with increasing currents gave poorly reproducible hysteresis effects, particularly at the higher cds. The over-potentials recorded during an experiment in a sulphate melt containing 9 x 10S2 m of oxide ion are plotted vs log I in Fig. 4. At low cds a Tafel region with a slope corresponding to RT/4F can be drawn through the experimental points. As the cd approached 1 mA/cm2, a limiting current was observed. A calculation based on the well known equation i, =
nFDC,s6
(14)
showed that this limiting current was not likely to be due to mass transport effects. By taking D = 1W6 cm2/s for the oxide ion, 6 = 5 x 1W2 cm for a well stirred melt, (b)
Slope
I
I
-0
-1-o
2.3RT 7
0
log ia (&
-2.0
I
-1.0
,
in mA/cms2)
FIG. 4. Anodic overpotentialat Pt/O, electrode in molten Li,SOcK,SOI eutectic at 625°C L)Slog i. The reference was an unpolarized Pt/O, electrode in the anolyte. (a) Neutral melt. (b) Basic melt.
Co,- = 1-S x l(r moIe/cm3 and n = 4 i, was calculated to be ca 14 mA/cm2, which is significantly higher than that observed, viz 1 mA/cm2. No potential/time dependence was noted in the basic melt. At each cd, a constant over-potential was quickly established, and open-circuit decay of 9 to near zero was rapid. Furthermore, no hysteresis was observed but, as in the pure meIt, the pIatinum electrode became tarnished by a film the thickness of which increased with time of electrolysis and with oxide concentration. In appearance it varied from speckIed dark grey to a more coherent dark metallic grey. It was assumed to be platinum oxide24*26and readily disappeared when heated. Evidence suggesting the presence of an oxide film was obtained from currentreversal chronopotentiograms at low cds. From these, both the anodic formation and cathodic removal of an oxide layer are clear (Fig. 5). Anodic formation was observed only on precathodized electrodes, suggesting that a film of oxide (most likely PtO) is normally present on a bright Pt electrode in a sulphate melt. The emf behaviour of
Electrochemical bright
and platinized
electrodes
studies of molten alkali sulphates
(see above)
also points to the existence
455
of a film of
PtO. The chronopotentiometric observation can be quantified by comparing the charge consumed during either formation or removal of the film (it is the same in both cases) with that calculated for a monolayer of PtO. From the lattice constants of crystalline platinum, it has been estimated 26 that there are l-5 x 10L6 atom/cm2 on the preUsing this value it can be calculated that approxidominant (111) crystal face. mately O-5 x 10-s C/cm2 are required to provide each surface atom of platinum From several of the origina galvanostatic transients it was with an oxygen atom. determined that O-3-08 x 103 C/cm2 were consumed during the oxidation, in reasonable agreement with the calculated value. o-4
o-t-
-o.3o
I I
I 4
I 3
1 2 Tinu.
I 5
i
I
FIG. 5. Chronopotentiogram showing formation and removal of a monolayer of PtO in LipSO,-KISOI eutectic melt at 625°C. 1 mA/cm*.
At high cds in the pure melt (> I mA/cma), a further transition time was observed (Fig. 6), which increased in proportion to the time of cathodic pretreatment. The potential/time trace for this process was sufficiently well defined for the potential E to be plotted against (rlla - +‘2/@z), which plot according to the relation2’
E = Ei +sln
(Zmt”l’=‘2)
(15)
should be linear. As is seen in Fig. 6, this is so and an = 2-9 for the charge-transfer reaction involved. The nature of the process involved in this second transition time is evidently the partially irreversible oxidation of dissolved oxide ions present as a residual impurity. Using the Sand equation, the concentration of residual oxide in the pure sulphate melt was estimated to be 4 x lo-8 m. Thus, transition between the TafeI regions in
B. W. BURROWS
456
0
I
I
0.5
I.0
and G. J. IHILLS
I
Time.
I
I
I.5
CO
2-5
s
FIG. 6. (a) Chronopotentiogram in l&SO,-K,SO, eutectic melt at 625°C showing wave due to oxidation of OS-. 1 mA/cm”. (b) An a lysis of the chronopotentiometric wave for the oxidation of Oz-.
presumabiy corresponds to a change in mechanism from 0% discharge to If this is the case, the transition cd should correspond to a massSO,” discharge. transport-controhed limiting cd. This supposition was checked by calculating i, from (14) taking Co,- = 8 x 10” mol/cm3. A value of ii ca O-7 mA/cm” was obtained, in cIose agreement with the observed value of ca 0.6 mA/cm2. The degree of oxidation of the substrate electrode and its degree of coverage with oxide or adsorbed oxygen are important factors that influence the detailed mechanism of anodic processes in soIutions containing oxygenated species. It is theoretically possible to relate the observed Tafel parameters to the mechanism. Bockris,2s for example, has extended the general theory of consecutive reactions developed by and applied it specifically to the oxygen-evolution reaction. Christiansen2s The treatment has been applied in a similar manner to anodic processes in molten carand cryoliteahuninas2 systems. It is somewhat simpler to apply bonate,30 nitraW in ionic melts, since the number of possible reaction steps can be reduced, due to the absence of aqueous solvent. In a sulphate melt the observation that a surface layer of PtO is probably present on the electrode (see above) further allows the number of possible reaction schemes to be reduced until only the folIowing appear feasible,
Fig.
4
(a) (b)
PtO + 02-* 2Pt0,&2Pto
PtO,
+ 2e, + 0,.
(16)
The Tafel slopes, predicted when either (a) or (b) is the slow step, are summarized in TabIe 3. Langmuir behaviour is assumed for the adsorption of the reaction intermediate, PtO,.
Electrochemical TABLET.
ANALYSIS
OF THE PROPOSED
LiBS06KpS0,
Rate-determining (a) PtO + O*- * (b)
studies of molten alkali sulphates ANODIC REACTION AT 625”C*
457
MECHANISM
step
Predicted tafel slope
PtO, + 2e
RT 2(1 - a)F
2ptoa *2PtO
+ 02
g
for e,,o, -
il for L&o~ -
IN
0 1
In Table 3, ir represents a limiting cd caused by the rate controlling activation reaction having reached a constant maximum at &oa = 1. Considering the q/log i behaviour in a basic sulphate melt, the Tafel slope of RT/4F and the limiting cd previously shown not to be controlled by diffusion in solution can now be understood in terms of mechanism (16) where (b) is the ratelimiting cd has previously been determining step. A similar activation-controlled reported for the hydrogen-evolution reaction on palladium electrodes in 5 N aqueous W-X= In a pure sulphate melt it has already been suggested that the two observed Tafel regions (Fig. 4) correspond to the successive discharge of Oa- ions and SOg2- ions at low and high cds respectively, with a transition at CLI0.6 mA/cmZ corresponding to a diffusion-controlled limiting cd. Arkhipov, Trunov and Stepanov3Q came to a similar conclusion regarding the nature of the anodic process taking place in carbonate melts. The discharge of SOq2- possibly proceeds through a mechanism similar to (16) with intermediate formation of PtU,. A theoretical analysis for this case also leads to the same predicted Tafel slopes. Therefore, with the reaction PtO. -j- SOq2- *
PtO,
+ SO3 + 2 e
(17)
as the rate-controlling step and with (1 - 0~) = 0.5, a TafeI slope of RT/p is predicted, in accordance with the observed value. Furthermore, it can be predicted that if measurements were made at sufficiently high cds than an activation-controlled limiting cd would be observed, corresponding to complete coverage of the electrode by PtO,. It is likely that the discharge of oxide ion in a “pure” sulphate melt also proceeds through a PtO, intermediate, but in contrast to more basic melts, step (a) of reaction (16) must be rate-determining. The only previous measurements of anodic polarization behaviour in sulphate melts are those of Flood and Fsrland. 35 These workers investigated the overpotential associated with the anodic evolution of O2 in sulphate and carbonate melts and in fluoride melts containing added sulphate, carbonate, phosphate and silicate, over the temperature range 6561000°C. The investigation was designed to study the possibility of using the O,/Pt electrode to measure oxide-ion activities; consequently, the systems were buffered with respect to these ions. Steady-state galvanostatic measurements showed that the various systems had some common characteristics. At low cds and high temperatures, Tafel slopes of RT/3F were obtained, whereas at high cds and low temperatures, the slopes were RT/F. Because the temperature coefficient of overpotential was lower by a factor of three in the former case (-0-4 mV/Y), it was concluded that the overpotential was mass-transfer controlled. For Tafel slopes
B. W.
458
corresponding order
to RT/F,
BURROWS and G. J. HI-
the activation so42-
> PO,*
overpotential
in the
> COs2- > SiO,&,
parallel with the strength of the X-O Strict comparison of these results with although it appears that the Tafel slope of present investigation in sulphate melts not
ie in
was found to decrease
bond. those of the present case is not possible, RT/F corresponds to that obtained in the containing added oxide.
authors wish to express their thanks to the Central Electricity Generating Board for Financial assistance, including a Junior Fellowship for one of them (B. W. B.).
Acknowledgements-The
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