Electrochemical formation of LiAl alloy in molten dimethylsulfone at 150°C

283

1. Electroanal. Chem., 209 (1986) 283-296

Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

E~~OC~CAL FORMATION OF LiAi ALLOY IN NOLAN DI~E~~FO~ AT 150°C

JEAN-PIERRE

PEREIRA-RAMOS,

RICHARD MESSINA and JACQUES PERICHON

~~r~t~~~g ~~~~tr~h~~ie, Cata&se et SynthPse organza

~LE.C.S.0.) C.N. R.S.,

2, rue Hewi-i&rant, 94320 lliiais {France) (Received 2nd January 1986; in revised form 28th April 1986)

ABSTRACT The electrochemical behaviour of the Li+/Li couple is examined briefly in molten ~methylsulfone (DMSO,) at 15OY. It has been found that the Li+/Li system is reversible. However, electrodeposited lithium exhibits some instability which severely limits the use of the lithium anode in molten DMSO, based secondary cells. Therefore we have undertaken the study of the electroformation of the LiAl alloy in 1 mol kg-’ LiClO,/DMSO, at 150°C by investigating the electrochemical incorporation of lithium into an alum~nium electrode by potentiostatic and galvanostatic techniques. The results obtained from electrochemical rne~u~rnent~ and X-ray diffraction experiments have proved that in~o~ration of Li in AI in molten DMSO, leads to the formation of the fi LiAl alloy. Analysis of the chronocoulometric curves has allowed the different processes limiting the rate of incorporation of Li in Al to be specified. Moreover, the diffusion coefficient of lithium into aluminium has been determined from chronoamperometric me~urements: 0.7 X 10-'O Q DLicnpd 1.4X 10-‘” cm* s -‘. Finally, the galvanostatic study has shown that the & Liil alloy can be considered a promising anodic material for molten dimethylsulfone-based secondary batteries.

INTRODUCTION

It is well known that the dendritic nature of electrodeposited lithium [l-3] and its high reactivity towards organic electrolytes [3J are the major factors limiting the practical development of secondary lithium organic cells. Also, the solubility of liquid lithium and its extreme corrosivity often cause difficulties in molten salts secondary batteries [4-61. One way of overcoming these problems is to use lithium intermetallics as negative electrodes. The electroformation of a large number of lithium alloy intermetallics in organic electrolytes was demonstrated by Dey [7]. It appears that aluminium is the most promising alloying metal with lithium [S], since LiAl deposited on aluminium does not grow dendrites and exhibits high rechargeability in many aprotic solvents [8-lo] and in molten electrolytes [5]. A large number of data on LiAl alloys have been reported, such as, for example, emf ~22~728/86/~3.50

0 1986 Elsevier Sequoia S.A.

284

measurements [5,6,11,12] and thermodynamic properties [4f. Furthermore, the mechanism of the electroformation of LiAl alloys at an aluminium electrode has been investigated in various organic electrolytes at ambient temperature [9,11,13-171, in molten salts at 175°C [1X] and over the temperature range from 415OC to 600°C [4,19-231. In addition, recent work has shown that LiAl must also be considered as a promising electrode material for solid state rechargeable lithium cells [24]. We have shown previously that sulfone-based electrolytes, and especially molten dimethylsulfone Li salts, could be used in lithium intercalation batteries operating between 80 and 200°C [25,26]. Though the Li+/Li system was suitable for use as a reference electrode in these media [25], neither the stability of Li, nor the electrochemical behaviour of the Li+/Li couple have been investigated. The present paper is concerned first with the el~tr~eposition of lithium in molten dimethylsulfone (DMSO,) at 150°C (1 mol kg-’ LiClO,/DMSO,). Then, in order to improve the stability of the deposit, we have investigated especially the electrochemical formation of the LiAl alloy in this medium. EXPERIMENTAL

Dimethylsulfone (DMSO,) was obtained from Janssen. DMSq (m.p.: 10S°C) was first recrystallized in water and then twice from absolute methanol, air dried at 90 “C for 48 h and dried again in vacua at 30°C for 12 h. It was then conserved in an argon glove box. Under these conditions, the water concentration did not exceed 5 x 10T3 mol kg-’ [25]. Anhydrous lithium perchlorate (Fluka) was dried under vacuum at 200°C for 12 h and tetrabutylammonium perchlorate was dried under vacuum at 80 “C. The electrolytes were prepared under a purified argon atmosphere. The lithium working electrode and the counter electrodes were made with lithium metal pressed into glass tubes of 0.3 cm and 1 cm diameter, respectively. The reference electrode was formed by a lithium wire in contact with a 1 mol kg-’ lithium perchlorate solution in molten DMSO, and separated by a frit. The aluminium working electrode was an aluminium wire of high purity (99.998%), 2 mm in diameter. The aluminium wire was sealed into a glass tube. The surface of the working electrode was cleaned before each experiment by rubbing it with extra fine emery paper and then with diamond paste using lubricant until its potential was about E 2.2/2.3 V vs. Li/Li+. A 0.25 mm thick ~~~~ foil (Alfa 99.999%) with a geometric area of 1 cm2 was used as the working electrode for galvanostatic measurements. The pyrometallurgical LiAl powder was supplied by La SocietC SAFT. In all cases, the working electrode was held in a parallel plate configuration with the lithium counter electrode. The three electrode cell used for electrochemical measurements has already been described 1251.The cell assembly was placed in an argon filled glove box. During experiments, the atmosphere above the electrolytic solution inside the celI was flushed continuously with argon. The temperature of the cell was controlled to within &-l°C by a Huber T 200 thermostat.

285 The samples for X-ray diffraction measurements were washed with dried and degassed distilled tetrahydrofuran (THF) and then dried in the dry box at ambient temperature. The X-ray analysis was carried out with a C.G.R. (Theta 60) X-ray system using the Co Ka, radiation (X = 0.1789 nm). RESULTS (1)

AND

DISCUSSION

Electrochemical

stua) of the Li ‘/ Li system

A typical cyclic voltammogram for a 4.25 X lo-* mol kg-’ LiClO, solution in molten DMSO, recorded under linear diffusion conditions at a lithium electrode is shown in Fig. la and is concerned with the following electrochemical reaction: Li+ + e- * Li. A linear dependence of the cathodic peak current (I,) versus the square root of the sweep rate (u”*) is obtained for sweep rates from 0.01 to 0.25 V s-l (Fig. lb). The diffusion coefficient of Li+ ions in molten DMSO, for the diffusion controlled process: Li+ + e- --, Li is evaluated from the slope of the straight line plotted in Fig. lb and from chronoamperometric measurements. The current-time transients recorded for two different overpotentials follow the single relationship i = kt- I’*, showing that electrodeposition of lithium ions is diffusion-controlled for

RI.

0.6

Fig. 1. (a) Typical cyclic voltammogram of 4.25 x 10e2 mol kg-’ LiClO, on a lithium electrode in 1 mol kg-’ NBu4C104/DMS02 (sweep rate = 0.17 V s-l, A = 0.07 cm2, starting potential = - 120 mV vs. Li/Li+ ). (b) Dependence of the reduction peak current on the square root of the sweep rate.

286

Fig. 2. Current-time dependence for potentiostatic reduction of a 8.66 X 10m3 mol kg-’ LiClO, solution on a lithium electrode in 1 mol kg-’ NBu,CIO.,/DMSOz at 150°C (A = 0.07 cm2).

about 10 to 25 s depending on the overpotential (Fig. 2). Taking the density of molten DMSO, calculated at 150°C from Smith’s relation [25] into account, the concentration of the Li+ ions can be expressed in mol cmd3. The diffusion coefficient of lithium ions, I&+, in molten DMSO, at 150°C is found to be: 0.7 x 10m5 < D‘i+ Q 3 x lo-* cm2 s-l. These determinations are consistent with the values provided by other authors who found a diffusion coefficient, Q+, of the order of 1O-6 cm2 s-’ in conventional organic electrolytes at ambient temperature 11,271. The voltammetric curve obtained for a 1 mol kg-’ LiClOJDMSO, electrolyte at a lithium electrode shows that the Li+/Li system is reversible (Fig. 3). From the chronocoulometric curves reported in Fig. 4, it can be seen that whatever the overpotential applied to the working electrode, there is a linear dependence between the charge consumed for the formation of metallic lithium and time. Such a result proves that the rate of lithium electrodeposition is not limited by the diffusion process of the Li+ ion. A fundamental criterion with regard to the suitability of lithium electrodes as anodes operating in secondary lithium cells is the anodic stripping efficiency

287

11

.

i . l.!

1r

.

O.!5.

2

1I-

A *

-

0.5.

1.

5. - 1.

-,2-

L

-0.sb

-0.2

0 E/V

0.2

0.4

6

vs. Li/Li’

Fig. 3. Cyclic voltammogram of a 1 mol kg -’ (A = 0.07 cm*, sweep rate = 5 mV s-l).

LiClO,

solution

in DMS02

on a lithium

electrode

obtained after electrodeposition. The electrodeposition of lithium at a platinum electrode gives rise to a stripping efficiency of about 62% (current density: 6.2 mA cm-‘; charge density: 2.8 C cm-*) whereas an immediate anodic redissolution yields a stripping efficiency of 60-808 in propylene carbonate [3], of 65% in sulfolane [28] and more in some mixed sulfolane-solvent solutions [29]. Although the Li+/Li couple behaves as a reversible, Nemstian [25] system in molten DMS02, as in propylene carbonate [1,30], the results obtained above clearly show that the electrodeposited lithium is found to be unstable and probably reacts with the electrolyte. This intrinsic lithium/electrolyte reactivity is usually encountered in any polar organic solvent-based electrolyte [1,3]. As the high reactivity of the deposited lithium towards the electrolyte limits the utilization of the Li anode in molten DMSO,-based secondary cells, we have

288

Fig. 4. Chronoeoulometic curves for reduction at various potentials of a 1 mol kg-’ LiClO, solution at a lithium dectrode (A = 0.07 cm2).

undertaken the study of the e~~trofo~ation of the lithium + mourn alloy which is well known to be better behaved and to be more stable than pure lithium in conventional organic solvents.

The ehxtrochemical fo~ation of the LiA1 ahoy was investigated by studying the electrochemical incorporation of lithium into an aluminium electrode. Figure 5 shows a typical voltammetric curve of the reduction-oxidation reaction of lithium on an aluminium electrode recorded at a scan rate of 5 mV s-’ in the potential range below the ~~~b~urn Li+/Li potential. The anodic dissolution of lithium produced a well defined peak whose area depends on the number of Li+ ions

289

1

0.b-

0.3 -

0.2 .

0.1.

-Of-

-0.3

.

-0.b

.

-OS-

0

.

’ 0.

I lib E/V

Fig. 5. Typical LiCIO.,/DMSO,

P6

0.6

1

IS. Li/Li*

voltammogram for lithium deposition/stripping at 150°C (sweep rate = 5 mV s-l, A = 0.031 cm*).

on

aluminium

in

1 mol

kg-’

reduced during the cathodic process. Although the coulombic charge recovered in the anodic process is lower than that involved in reduction, this anodic redissolution is more important than that obtained for lithium deposition on a platinum electrode. We first studied the electrodeposition of Li+ ions onto an aluminium electrode by chronocoulometric measurements. This method appears to be particularly suitable for such a study and has been applied in other media in order to investigate the rate-limiting process in cathodic deposition of Li onto Al. In this method, the potential is stepped to a desired value and the charge passed during the controlledpotential electrolysis is measured as a function of time by integration of the observed current. The chronocoulometric curves obtained for the reduction of Li+ ions at Al are reported in Figs. 6 and 7.

290

O.4

0.

O., s El

0.

Fig. 6. Chronocoulometric curves for lithium incorporation into aluminium at low overpotentials of the litbium+al&nium alloy formation (1 mol kg-’ LiClO.,, A = 0.031 cm*). The incorporation rate is limited (a) by lithium diffusion in Al to form the o phase and (b) by the reaction of the /3 LiAl alloy formation.

I

0 ,

L

0’

.

Fig. 7. Chronocoulometric curves recorded at various potentials for lithium incorporation into aluminium to form the /3 phase Liil alloy (1 mol kg-’ LiClO,, A = 0.031 cm2). The incorporation rate is limited (a) by the reaction of the B LiAl alloy formation and (b) by diffusion of lithium into the /3 LiAl alloy.

291

For low overpotentials (Fig. 6), for a short time after the onset of the cathodic polarization, the charge passed, Q, is proportional to t’12 (Fig. 6a) indicating that, in line with Astakhov’s theory [31], the incorporation rate of Li into Al is limited by the diffusion of Li in Al. Such results have been observed in other media at low overpotentials and at low current densities [18,22]. However, because of the possible corrosion of LiAl, parasitic currents related to electrolyte decomposition might be included in the measured currents. For longer times, the incorporation rate then changed to a linear dependence on f, as shown in Fig. 6b. It can be seen that Li initially penetrated at a diffusion-limited rate into aluminium to form the a-phase which is a solid solution of Li in Al. The duration of the diffusion-limited incorporation regime diminished with increased overpotential. For longer times, the rate of incorporation appears to be controlled by the formation of the intermetallic compound, /? LiAl. At overpotentials where the rate of Li deposition on the Al surface exceeds the diffusion flux into the electrode, an accumulation of Li in the surface layers most probably occurs and gives rise to the formation of the j3 LiAl alloy. As can be seen from the plots in Fig. 7a, a linear relation exists initially between Q and t but a deviation from linearity occurs later. When the data obtained for t > 150 s are replotted as Q vs. t”‘, a linear relation is observed (Fig. 7b). In accordance with Astakhov’s theory, the rate of incorporation of Li in Al is controlled initially by the P-phase formation (Q = f(t)) and then by a diffusion controlled process. Indeed, the linear dependence, Q = f(t’/*), found for longer times may be ascribed either to the diffusion process of Li into LiAl or to concentration changes (Li+ depletion) in the solution in the vicinity of the electrode [32], i.e. a liquid state diffusion. For low overpotentials, the diffusion coefficient of lithium in Al during the formation of the a phase is determined from the chronoamperometric measurements. As seen in Fig. 8, the Cottrell law is obeyed well according to: i = nFA( cL, - ci’,) D~{,fa,/vr’/2t’/2

where cLi is the saturation concentration of lithium in the a phase at 150°C (= 2.5 at.% Li [33]) divided by the molar volume of pure aluminium (10 cm3 mol-‘) and the initial lithium concentration czi was assumed to be zero. The other symbols have their usual meanings. An estimation of the diffusion coefficient, DLituj, of Li in the a phase of the LiAl alloy is obtained: 0.7 x 10-r’

.
We can find no other studies of lithium diffusion into Al conducted in the temperature region of our work. For comparison, different values have been found at ambient temperature, 1.4 x lo-l2 cm2 s-’ in tetrahydrofuran [14] and 2.4 x 10-l’ cm* se2 in propylene carbonate [16]; DLicaj has also been found in the range 0 7 x lo-” < D L,taj G 5 X lOPro cm2 s -’ for a temperature range of 417-450°C [;9,22,23]. ’ In order to investigate the electrochemical behaviour of the j3 LiAl alloy in Li secondary cells, we have performed galvanostatic measurements at an aluminium

292

36

0 0

0.1

0.2

04

0.5

6.6

Fig. 8. Current vs. (time)- ‘/’ for the incorporation of lithium into aluminium to form the a phase Ltil alloy (A = 0.031 cm2).

Fig. 9. Cathodic (a) and anodic (b) chronopotentiogram in 1 mol kg-’ LiC10,/DMS02 (j = 6.4 mA cme2).

for lithium deposition/stripping

on aluminium

293 TABLE

1

Dependence J/d

%_I/~

cm-’

of the cycling efficiency, 0.5 83

qL,, on the current density

1 88/89

6.4 92/93

12.8 92

20 94

40 92/93

foil. Figure 9 shows a typical cathodic chronopotentiogram corresponding to Li+ reduction (a), followed by an anodic chronopotentiogram (b) corresponding to the oxidation of Li from the previously formed alloy. The polarization observed at the onset of the cathodic chronopotentiogram is attributed to the nucleation of the intermetalhc compound. The cycling efficiency of the process is determined from the ratio of the transition time for the oxidation to the time for which reduction is carried out. The cycling efficiency with respect to lithium is denoted uLi. A stripping potential limit of 2 V was employed. In Table 1, the experimental values of rlLi obtained as a function of the current density (j) when the electrical charge involved in the reduction process is kept constant at 1 C cm-*, are reported. From the data summarized in Table 1, it can be seen that the lithium efficiency is of the order of 90-94% whereas lithium efficiencies of about 91-96% are obtained in common aprotic media [14,34]. The higher values of nL, found for current densities > 1 mA cm-* can be explained by the fact that the formation of the intermetallic compound, LiAl, is enhanced whereas the (Y phase is mainly formed at low current densities. Taking the high discrepancy between the values of diffusion coefficients of Li into Al and into p LiAl into account (D,,,, -+x &(a)), it appears that Li is more accessible from the /I phase for anodic dissolution. Moreover, the time reduction being higher at low current densities, a chemical reaction between the deposit and the electrolyte can occur. The stability test of the /3 LiAl in molten DMS02 at 15O’C was conducted by deposition of Li on an aluminium foil followed by anodic stripping after a time interval. The experiment was performed with a charge involved in the cathodic process equal to 10.8 C cm-* (current density = 20 mA cm- * ) and the anodic redissolution performed immediately led to a lithium efficiency equal to 91%. When the anodic stripping occurred after a time interval of about 3 h, an efficiency of only 85% was obtained. Although this stability problem has not been explored further, we suggest that a chemical reaction between the electroformed LiAl alloy and the electrolyte or the diffusion of Li from the alloy into Al to produce the (Y solid solution phase, or both, are probably involved in the stability of the alloy. We have compared the nature of the electrochemically formed LiAl alloy with that of the pyrometallurgically formed LiAl with a Li concentration equal to 50 at.% Li. The equilibrium potential of an electroformed LiAl alloy corresponding to a concentration of 48 at.% Li (381 mV) is found to be similar to that obtained for the pyrometallurgical LiAl powder pressed on a nickel grid (378 mV). Such a result is in

294 TABLE 2 Comparison of the experimental X-ray diffraction data obtained for the electroformed LiAl alloy and that observed for the pyrometallurgically formed LiAl Electroformed Liil alloy

LiAl intermetallic compound (pyrometallurgical formation) lOd/nm 3.67 2.342 2.25 2.04 a 1.918 1.77 a

Relative intensity 100

hkl

lOd/nm

Relative intensity

hkl

(111)

3.67 3.35 2.342

45
(111)

(220)

2.257

100

(220)



2.026

90

(311)

1.921

25

(311) (400) (331)

1.589 1.458 1.432

15 15 10

(400) (331)

1.591 1.459 1.434

20 30 (10

1.298 1.221

20 60

(422) (333X511)

1.298 1.22

20 10

(422) (333)(511)

1.123 1.075 1.064

25 10 30

(440) (531)

1.123 1.074

20 10

(440) (531)

1.012

85

(620)

a The thermally prepared LiAl powder was pressed on a nickel grid (d/nm

0.204, 0.177).

good agreement with emf measurements performed in other media and is consistent with the electrochemical formation of the /3 LiAl alloy [9,11,13]. An X-ray diffraction experiment on the electroformed alloy was carried out. Lithium was electrochemically deposited on an aluminium foil at a current density of 20 mA cm- 2. The aluminium foil was charged with Li to 65% of capacity (based on conversion of the foil to /3 LiAl 50 at.% Li). The available X-ray diffraction data on the specific intermetallic pyrometallurgically formed LiAl alloy (50 at.% Li) are reported and compared with that obtained for the electrochemically formed LiAl alloy (see Table 2). The X-ray diffraction lines were identified by using the literature data given in ref. 35. The X-ray diffraction data of the electrochemically formed alloy show this to be identical to the thermally prepared intermetallic compound LiAl, i.e., to the /3 LiAl ahoy. The experimental X-ray data listed in Table 2 also indicate the presence of unalloyed aluminium or (Yphase (d/nm 0.2342, 0.2026, 0.1434) and perhaps of lithium (d/nm 0.1434) in addition to the electroformed /3 LiAl alloy. CONCLUSION

The Li+/Li couple has been found to behave as a Nemstian [24], reversible system in molten DMSO, at 150°C but, as observed in many aprotic media, its

295

relative stability severely limits its use as the anode in molten DMSO,-based secondary cells. -Electrochemical measurements and X-ray experiments have proved that incorporation of Li in Al in molten DMSq at 150°C leads to the formation of the p LiAl alloy. The rate of incorporation of Li into Al is first controlled by the j3 phase formation and subsequently by the diffusion of Li in the /3 LiAl alloy. The formation of the (Y phase is only observed at the onset of the incorporation of Li into Al at low overpotentials of the lithium + aluminium alloy formation. In this case, the diffusion coefficient of Li in Al was measured in the range 0.7-1.4 x lo-” cm* s-r at 150°C. -The preliminary galvanostatic measurements reported in this work have demonstrated that the LiAl alloy must be considered a promising anodic material for molten dimethylsulfone-based rechargeable batteries. ACKNOWLEDGEMENT

The authors with to thank Dr N. Baffier from C.N.R.S./L.A. performing X-ray experiments.

302 for his help in

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