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Iron oxide electrodes in lithium organic electrolyte rechargeable batteries
IRON OXIDE ELECTRODES IN LITHIUM ORGANIC ELECTROLYTE RECHARGEABLE BATTERIES S. MORZILLI and B. SCROSATI Dipartimento
di Chimica, University of Rome “La Sapienza”. Italy
F. SGARLATA Dipartimento
di Science della Terra, University of Rome “La Sapienza”, Italy
(Received 18 December 1984;in revised form 28 February 1985) Abstract-The electrochemical reaction of iron oxide (cc-Fe,O,) with lithium in an organic electrolyte cell has been investigated by cyclic voltammetry, polarization curves, atomic absorption analyses and MLissbauer spectroscopy. This reaction involves a multistep insertion process which appears reversible for a wide range of lithium content in iron oxide. Therefore,lithium-richiron oxide may be regarded as a negative electrode, alternative to pure lithium metal, in organic electrolyte batteries, with possible advantages in terms of cycling life. The oerformance of this iron oxide-based lithium electrode has been tested in cells using pristine intercalation kmpounds as the positive counterelectrodes.
INTRODUCI’ION It is now well known that lithium organic electrolyte
batteries may suffer from limited rechargeability beccause of the poor cyclability of the lithium metal electrode[l-33. Due to passivation phenomena resulting from spontaneous reactions with the electrolyte, the plating-stripping efficiency of lithium falls unavoidably below 100% upon prolonged cycling[4]. As a possible solution to this problem, we have proposed the replacement of the lithium metal with a L&rich compound having little or even no tendency of being passivated in the organic electrolyte. In previous papers[5,6] we have shown that fully intercalated compounds, such as LiWOz or LiNb&, may be efficiently used as negative electrodes in lithium cells. In fact, coupling LiWO, with a pristine intercalation electrode, such as TiS, or V,O,, an electrochemical cell in which lithium can be reversibly and repeatedly transferred from one electrode to the other, is obtained. These double intercalation electrode systems have been termed “rocking chair” cells[7 J. However, even if they are capable of sustaining a large number of cycles[5], LiWOz and similar lithiumintercalated compounds, have too high a weight and too low a lithium content to be effectively proposed for practical uses. In fact, for the realization of highvoltage, high-energy rocking chair batteries, the negative electrode must have the following general properties. (i) High lithium activity and small voltage variation during the cell processes (to ensure high operational This paper was presented at the 35th 1SE Meeting in Berkeley, California, U.S.A., August S-10, 1984 (abst. no. A2- 11).
voltages and low excursion between charge and discharge curves). (ii) Low weight and capability for high Li equivalents per mole (to ensure high specific capacity and high energy density). (iii) High reversibility and fast kinetics (to ensure high coulombic efficiency and high cycling rates). Among the various materials examined, we have focussed our attention on iron oxide (a-Fe,OJ), which seems to fulfill the majority of the above requirements. In previous papers[7,8] we have reported a preliminary study of the electrochemical characteristics of this compound and here we examine in more detail the mechanism of its reaction with lithium as well as its concrete possibilities of use as a negative electrode in rocking chair batteries.
EXPERIMENTAL Propylene carbonate (PC), a reagent grade product, was further rmrified by fractional distillation under reduced pressure. Lithium perchlorate, LiClO& reagent grade product, was vacuum dried at 150°C. Iron oxide, Fe,Oa and vanadium pentoxide. V20s, both reagent grade products, were used as received. A three electrode Teflon cell used for the electrochemical measurements is shown in Fig. 1. The working electrode was obtained by pressing pure Fe,OJ, or a mixture of FezO, with 5 w/o graphite, on a copper substrate. The pellet so obtained was placed at one end of the oell and held in place by a threaded Teflon holder having a stainless steel current collector. The counter electrode, realized in a form of a disk (Li) or of a pellet (V,O,), was placed at the other end of the cell. By means of Teflon spacers the area of the working
1271
1272
S. MORZILLI, B. SCROSATI
AND
F.
Fig. I, The experimental
electrode was fixed as 0.007 cm’ and the distance between the two electrodes was maintained at 0.4 cm. The Li working electrode was connected to the call by a Luggin-type capillar inserted through the top slit visibIe in the figure. The electrochemical stability range of the electrolyte was investigated with the same cell in which platinum or carbon, respectively, were used as the working electrode. All the operations and the preparations of the cells were performed in a controlled atmosphere dry-box (less than 6 ppm of humidity). The cyclic voltammetry curves were carried out with a potentiostat driven by a function generator. The Constant current cycling tests were run automatically with the aid of electronic polarity reversing timers inserted between the galvanostat and the cell. The Li analyses were carried out with an atomic adsorption spectrometer, equipped with a flameless system and using a 670.8 nm wavelength. The Li determinations were obtained with the standard addition method. For this purpose the electrodic pellet was dissolved in concentrated HCI, heated to boiling and then diluted to 1000 cm3 with distilled water, Four 1 cm3 portions of this solution were isolated. One of them was diluted to 10 cm3 and the three others were frrst added with known amounts of lithium (ie 0.3,0.6 and 0.9 pg, respectively) using a standard calibration solution (ie LiNO, 1 gl-’ aqueous solution, acidified with HN03) and then also diluted to 10cm3. The original lithium content in the solution under test, was finally obtained by an extrapolation method, a typical example of which is shown in Fig. 2. The Miissbauer spectroscopy (nuclear gamma resonance analysis) was run at constant acceleration (from -4.5 to + 6.5 mm s-i) with Elscint (Haifa, Israel) electronic equipment[9]. A 10mCi Co57 in Rh was used as gamma ray source (14 keV). The powder with resonant Mossbauer element was mixed with dextrin to fill the 125 mm2 area of sample holder.
cell.
r
* -3
SGARLATA
I
0
I
1
6
3 Known additms
1
9
of Li/pg
Fig. 2. Typical example of the graphic elaboration of the atomic absorption analysis, by the standard addition method.
RESULT 1. The electrochemical oxide
AND DISCUSSION reaction
of lithium with iron
To investigate the electrochemical reaction of Li with FezOa, we have considered the cell Li/LiClO,-PC/FezO,.
(1)
A typical discharge curve of this cell is shown in Fig. 3. The initial part of the curve reveals a series of plateaus covering a discharge range between 0 to approximately 0.6 Faraday per mole of Fe,Os. As already remarked in a previous paper[7$ it is rather difficult to identify the reaction sequence associated with this initial series of short plateaus. Huggins and co-workers[lO] reported a thermodynamic study of the ternary Li-Fe-0 system at 4OO*C, which revealed the presence of a series of intermediate phases, such as LiFe,O,, LiFeOl and Li,FeO,. It is unlikely that the room temperature electrochemical reaction ofcell (2) can be described with the sequential formation of these phases.
Iron oxide electrodes
in lithium
organic
electrolyte
1=025
0
I
I
x
Fig. 3. Discharge
curve
I
I 4
t
6
5
Fen 03
of the Li/LiClO,-PC/Fe,O,
The presence of the initial series of plateaus and the fact that the electrochemical reaction can be eventually extended beyond 6 F/mole (see the discharge curve in Fig. 3), tend to indicate a more complex discharge mechanism than the simple progressive displacement of iron by lithium. Therefore, with the aim of clarifying this mechanism, Mijssbauer spectroscopy has been run on the cathodic pellet at various stages of discharge. In Fig. 4 is illustrated the spectrum of the cathodic mass of a cell first discharged to 5.5 F/mole of Fe,O, and then recharged back to 1 F/mole. In Table 1 the least square fitting of resonance data is also reported. The most striking feature of the spectrum of Fig. 4 is the disappearance of the magnetic hyperfine Zeeman splittings typical of the antiferromagnetic iron oxide. In contrast, a fairly sharp doublet of quadrupolar electric interaction is formed. The MGssbauer parameters of
Fig. 4. Mikbauer
~n LI,
1273
batteries
macm-2
I 3
I 2
I
rechargeable
cell. Current
Table
density:
1. Mathematical
elaborations
PI @,I
x=1 jE = 5.5
0.25 mAcm_‘.
-0.19(0.56) - 0.08 (0.61)
of resonance i.s.
q.s.
0.55 0.61
0.74 0.63
PI(%)
+ 0.55 (0.79) +0.55 (1.01)
data
p = positions of spectral lines; w = half-height width. i.s. = isomer shift with respect to sodium-nitroprusside. qs. = quadrupolar splitting. All in mm s- 1 units. x = lithium equivalents per FelOa mole. quadrupolar Lorentzian lines reveal the presence of Fe’+ ions in a slightly deformed octahedron. Therefore, the electrochemical reaction of cell (1) does not involve the reduction of iron oxide but rather the insertion of lithium ions in the third octahedral interstitial site, which is unoccupied in the hexagonal close-packing of oxygen ions of haematite. The orig-
spectrum of the cathodic mass of the cell Li/LiClO,-PC/Fe,O, 5.5 F/mole of FezOs and subsequent recharge to I F/mole.
after
discharge
to
1274
S. MO~LLI,B.SCROSATIANDF.SGARLATA
inal oxide lattice is basically retained and indeed the Xray powder diffraction pattern of LiFe,OB is similar to iron oxide pattern[7$ Owing to insertion, an interaction seems to take place between the 2s electron of lithium and the valence band of iron oxide with consequent loss of relative orientation of spins. Therefore, the electrochemical reaction (1) may be considered as a multistep insertion process and described according to the scheme: xLi + Fe,O,
= Li,Fe,O,,
(1)
where x indicates the equivalents of Li which have reacted with 1 mole of Fe,O,. Initially (ie to x = 1) the octahedral positions are fillied; the value of x ranges from 1 to 6.6. A similar structure is seen in the NaCl-type lattice of LiFeO,[l 1] where the Li and Fe ions occupy together the cation sites. The crystal-chemical behaviour of LiFez03 may be explained in analogy with that of the LiCB graphite intercalation compound. In fact, an investigation by means of inelastic scattering of synchroton radiations has revealed that both charge transfer to the n band and modification of the valence band occurs in LiCJlZ]. Finally, we may comment on the values of the line widths in the resonance data related to the cathodic mass after discharge to 5.5 F/mole (Table 1). These very high widths reveal the spreading of the electric field gradient at iron positions which indeed may be related to a substantial insertion of lithium ions in the lattice. 2. Electrochemical stability A possible concern during the operation of cell (1) may be in the effectiveness of the electrochemical process, since at the relatively low voltages at which the last part of the discharge takes place (see Fig. 3), a partial reduction of the electrolyte could also take place. Therefore, it has appeared to us of priority importance to clarify this aspect by examining the kinetic characteristics of the electrochemical process and its influence on the stability of the electrolyte. The electrochemical stability of the electrolyte used in cell (l), ie the LiClO,-PC solution, depends on various factors, among which are the nature and the morphology of the testing electrode. For instance, if the stability of the electrolyte is measured with cyclic voltammetry, one obtains a range between 2.5 and 4.0 V (us Li) when using a platinum electrode and a range between 1.0 to 4.5 V (us Li) when using a carbon electrode. The response of the two electrodes is then Table 2. Determinations of Li/LiClO,-PC/Li,Fe,O, Sample Al EO A31
Xcalculated
6.0 6.6 5.6 6.6
quite different confirming that the kinetics of the decomposition process of the electrolyte is indeed greatly influenced by the nature and the characteristic of the electrodic substrate. Since it is impossible to measure the decomposition window of the electrolyte on Fe,O, in the conditions of cell (l), one may only reasonably assume that this window might approach that observed on carbon, ie a substrate morphologically similar on a macroscopic scale to that of iron oxide. However, even under this assumption, a concomitant decomposition of the electrolyte cannot be completely excluded during the last part of the discharge of cell (1). Therefore, to finally exclude any participation of the electrolyte in the discharge process, we have determined by atomic absorption spectroscopy the lithium content in the cathodic mass at the end of constant-current discharges. The lithium contents, determined with this technique, are compared in Table 2 with those calculated coulometrically on the basis of the total charge passed through the cell. The two sets of values agree within the experimental error and this allows to finally exclude any electrolyte decomposition side effects during the discharge of cell (1). Under the working conditions of this cell the electrochemical process predominates, probably since this is kinetically favoured over the reduction of the electrolyte. 3. Reuersibility of the elecrrochemical reaction It may be reasonable to assume that the initial part (x i 1) of the discharge of ccl1 (1) involves the intercalation of Li in low energy sites until filling the octahedral interstitial sites, while the second part (1 < x < 6) involves a more extended intercalation process. Illustrative in this respect is the case of the prelithiated Li, -xV,Os compound which allows a fast and reversible intercalation of lithium up to 2 equivalents per mole[ 133. Similarly, the reversibility of the electrochemical process of the Li/Fe,O, cell here discussed, increases as the discharge proceeds, as indicated by micropolariration and cycling tests reported in a previous work[7]. In Fig. 5 is shown slow rate cyclic voltammetry of the Fe203 electrode in the LiClO,-PC electrolyte. The voltammetric response reveals a large cathodic peak around 1.5 V us Li, related to an extended, reversible reaction. This result confirms the reversibility of the electrochemical process in the Lirich region, but also indicates that the kinetics of this process may be controlled by the diffusion of lithium
Li in the cathodic mass of cell at the end of the discharge
Xdetermincd
6.1 6.5 5.4 6.8
Difference -0.1 0.1 0.2 -0.2
the
Difference y0 -2
1.5 3.5 3
= Li equivalentsper FelOj molecalculated coulometrically. %akulat& On these data, the experimental error is AX, = 17;. x&t_,,,& = Li equivalents per FeZ03 determined by atomic absorption analysis. On these data the experimental error is AX1 = 4 %,
Iron oxide electrodes
in lithium organic electrolyte
throughout the iron oxide-based electrode. It may also be emphasized that no reduction of the electrolyte is revealed by the cyclovoltammogram of Fig. 5, confirming the effectiveness of the electrochemical reaction.
rechargeable
batteries
4. L&rich iron oxide as negative batteries
12?5 QlQCtrdQ
in lithium
The results so far discussed indicate that Li,FelOJ electrodes have the capabitity of reacting with a large amount of lithium with small voltage excursions and offer reversibility especially in the I&rich range. Therefore, these electrodes appear as sutiable candidates for use as alternatives to lithium metal in rechargeable organic electrolyte batteries. To verify this assumption, we have assembled the following test cell:
60.
40.
Li6Fe103/L1C104-PC/VZ05.
0.
20 =l + ‘-t
(2)
This cell was realized by passing charge up to 6 F/mole in a cell of type (1) and then replacing the Li electrode with Vz05, ie a typical intercalation electrode in its pristine state. Using exactly the same geometry, we have also assembled the cell:
20
Li/LiC104-PC/V205. 40.
The cycling behaviour of the two cells was compared under the same conditions and the result is reported in Fig. 6. In the same figure are also shown the single polarization (vs a Li reference) of the Li,Fe,O, electrode [cell (2)] and of the Li electrode [cell (3)]. At the initial cycling stage, the polarization of the former is apparently higher than that of the latter electrode_ However, as the cycling proceeds, the performance of the lithium metal electrode deteriorates much faster than that of the lithium intercalated electrode. This is clearly shown in Fig. 7 which reports the changes in polarization of the two electrodes us the cycle number, expressed as AV (polarization of progressive cycles) over AV” (polaritation of the first cycle). It may be noticed that the AV/AV” ratio remains fairly constant
60
80
IQ0
I
I I 5
I 2 5
I 2.0 WV
I 30
I 3.5
(3)
I 40
YS Ll
Fig. 5. Cyclic voltammetry of the FelOB electrode in the LiClO,-PC electrolyte. Scan rate: 50 pV s- ‘. Electrode surt-ace: 0.07 cmz.
Cycle
CycIe
30
/
I
I
08 x
I 08
50
I
m Li,V205
4-c
+i_Y
Cycle
10
Fig. 6. Charge-discharge cycles of the Li6Fe203/V205 and of the Li/V20s cells with the LiClO,-PC electrolyte. Single polarizations (us Li reference) of the LisFe20B and of the Li anodes, are also shown. Current density: 0.3 mAcm_‘. Room temperature.
I
S.
1276
B,
MOIUILLI,
.%S!ROSATl
AND
L16Fer0,/V,0,
Polor~zot~on of L16Fe,03
F.
?bGARl.ATA
cell
Polarlzotlan
the
0
I 15
0
Number of
cycles LI IV,05
3r
cell
Polarlzotlon of tile LI electrode 2
2
2 I IO
I 5 Number
of
I I5
cycles
the
ceil . l
.
.
l
1
t
I 20
0
I
I 5
1
IO
Number
cycles
Fig. 7. Increase of the polarization
I 60
I 45 of
Polar120t Ion ot
I
i;
of
I 30
Number
3
0
cell
the
electrode
with number of cycles: AV/AV’
for about 60 cycles in the case of cell (2), so that the overall cell polarization observed in Fig. 6 is here mainly associable to the V205 cathode. On the contrary, in the case of cell (3). the AY/A Y”ratio increases to almost double after 15 cycles. This behaviour, which clearly reveals the progressive deterioration of the Li metal electrode, is in agreement with the work of Brummer and co-workersll, 141 who have shown that charge+Gcharge efficiency of a Li electrode in PC-based electrolytes, decays to SO % after 1S cycles. These results confirm the favourable cycling capabilities of the lithium composite electrodes and the feasibility of the lithium rocking chair battery concept. However, they also indicate that Li,Fe20s is by far not the ultimate choice in this respect. Other compounds, having easier preparation procedure, faster kinetics and even higher reversibility, should be characterized for reaching a final evaluation of the effective practical importance of this new type of negative electrodes in the lithium battery technology.
Acknowledgements-The authors wish to thank Dr. E. Cardarelli for the Atomic Adsorption measurements. The financial support of the Consiglio Nazionale delle Ricerche, Progetto Finalizzato Energetica 2, is also acknowledged.
ot
15
/
20
cycles
of the first cycle.
REFERENCES 1. R. D. Rauh and S. B. Brummer, Electrochim. Arta 22,75 (1977). 2. F. W. Dampier and S. B. Brummer, Elecfrochim. Acta 22, i 339 ( 1977). 3. V. R. Koch and S. B. Brummer, Elecfrochim. Acta 23, SS (1978). 4. S. B. B rummer, V. R. Koch and R. D. Rauh, Marerials Jar Advanced Buttevies (Edited by D. W. Murphy, J. Broadhead and B. C. H. Steele), p. 123. Plenum Press, New York (1980). 5. M. lazzari and B. Scrosati, J. electrochem. Sot. 127, 773 (1980).
6. B. Scrosati, B. DiPietro, M. Lazzari and B. Rivolta, Power Sotrrces 9 (Edited by J. Thompson), ~-459. Academic Press, London (19S$. _ ‘. _ 7. B. DiPietro, M. Patriarca and B. Scrosati, J. Power Sources 8. 289 (1982). 8. M. Lazzari and B. Scrosati, U.S. Patent No. 4, 464,447, Aug. 7 (1984). 9. G. De Angelis and F. Sgarlata, Per. Min. 48, 287 (1979). 10. N. A. Ciodshaall, I. D. Raistick and R. A. Huggins, Mar. Res. BulI. 15, 561 (1980). 11. D. E. Cox, Phys. Rev. 132, IS47 (1963). 12. G. Loupias, J. phys. Letz. 45, 301 (1984). 13. G. Pistoia, S. Panero, M. Tocci, R. V. Moshtev and V. Manev, Solid Srate Ionics 13, 311 (1984). 14. K. M. Abraham and S. B. Brummer, Lithium Barter& (Edited by J. P. Gabano), p. 371. Academic Press, London (1983).
di Chimica, University of Rome “La Sapienza”. Italy
F. SGARLATA Dipartimento
di Science della Terra, University of Rome “La Sapienza”, Italy
(Received 18 December 1984;in revised form 28 February 1985) Abstract-The electrochemical reaction of iron oxide (cc-Fe,O,) with lithium in an organic electrolyte cell has been investigated by cyclic voltammetry, polarization curves, atomic absorption analyses and MLissbauer spectroscopy. This reaction involves a multistep insertion process which appears reversible for a wide range of lithium content in iron oxide. Therefore,lithium-richiron oxide may be regarded as a negative electrode, alternative to pure lithium metal, in organic electrolyte batteries, with possible advantages in terms of cycling life. The oerformance of this iron oxide-based lithium electrode has been tested in cells using pristine intercalation kmpounds as the positive counterelectrodes.
INTRODUCI’ION It is now well known that lithium organic electrolyte
batteries may suffer from limited rechargeability beccause of the poor cyclability of the lithium metal electrode[l-33. Due to passivation phenomena resulting from spontaneous reactions with the electrolyte, the plating-stripping efficiency of lithium falls unavoidably below 100% upon prolonged cycling[4]. As a possible solution to this problem, we have proposed the replacement of the lithium metal with a L&rich compound having little or even no tendency of being passivated in the organic electrolyte. In previous papers[5,6] we have shown that fully intercalated compounds, such as LiWOz or LiNb&, may be efficiently used as negative electrodes in lithium cells. In fact, coupling LiWO, with a pristine intercalation electrode, such as TiS, or V,O,, an electrochemical cell in which lithium can be reversibly and repeatedly transferred from one electrode to the other, is obtained. These double intercalation electrode systems have been termed “rocking chair” cells[7 J. However, even if they are capable of sustaining a large number of cycles[5], LiWOz and similar lithiumintercalated compounds, have too high a weight and too low a lithium content to be effectively proposed for practical uses. In fact, for the realization of highvoltage, high-energy rocking chair batteries, the negative electrode must have the following general properties. (i) High lithium activity and small voltage variation during the cell processes (to ensure high operational This paper was presented at the 35th 1SE Meeting in Berkeley, California, U.S.A., August S-10, 1984 (abst. no. A2- 11).
voltages and low excursion between charge and discharge curves). (ii) Low weight and capability for high Li equivalents per mole (to ensure high specific capacity and high energy density). (iii) High reversibility and fast kinetics (to ensure high coulombic efficiency and high cycling rates). Among the various materials examined, we have focussed our attention on iron oxide (a-Fe,OJ), which seems to fulfill the majority of the above requirements. In previous papers[7,8] we have reported a preliminary study of the electrochemical characteristics of this compound and here we examine in more detail the mechanism of its reaction with lithium as well as its concrete possibilities of use as a negative electrode in rocking chair batteries.
EXPERIMENTAL Propylene carbonate (PC), a reagent grade product, was further rmrified by fractional distillation under reduced pressure. Lithium perchlorate, LiClO& reagent grade product, was vacuum dried at 150°C. Iron oxide, Fe,Oa and vanadium pentoxide. V20s, both reagent grade products, were used as received. A three electrode Teflon cell used for the electrochemical measurements is shown in Fig. 1. The working electrode was obtained by pressing pure Fe,OJ, or a mixture of FezO, with 5 w/o graphite, on a copper substrate. The pellet so obtained was placed at one end of the oell and held in place by a threaded Teflon holder having a stainless steel current collector. The counter electrode, realized in a form of a disk (Li) or of a pellet (V,O,), was placed at the other end of the cell. By means of Teflon spacers the area of the working
1271
1272
S. MORZILLI, B. SCROSATI
AND
F.
Fig. I, The experimental
electrode was fixed as 0.007 cm’ and the distance between the two electrodes was maintained at 0.4 cm. The Li working electrode was connected to the call by a Luggin-type capillar inserted through the top slit visibIe in the figure. The electrochemical stability range of the electrolyte was investigated with the same cell in which platinum or carbon, respectively, were used as the working electrode. All the operations and the preparations of the cells were performed in a controlled atmosphere dry-box (less than 6 ppm of humidity). The cyclic voltammetry curves were carried out with a potentiostat driven by a function generator. The Constant current cycling tests were run automatically with the aid of electronic polarity reversing timers inserted between the galvanostat and the cell. The Li analyses were carried out with an atomic adsorption spectrometer, equipped with a flameless system and using a 670.8 nm wavelength. The Li determinations were obtained with the standard addition method. For this purpose the electrodic pellet was dissolved in concentrated HCI, heated to boiling and then diluted to 1000 cm3 with distilled water, Four 1 cm3 portions of this solution were isolated. One of them was diluted to 10 cm3 and the three others were frrst added with known amounts of lithium (ie 0.3,0.6 and 0.9 pg, respectively) using a standard calibration solution (ie LiNO, 1 gl-’ aqueous solution, acidified with HN03) and then also diluted to 10cm3. The original lithium content in the solution under test, was finally obtained by an extrapolation method, a typical example of which is shown in Fig. 2. The Miissbauer spectroscopy (nuclear gamma resonance analysis) was run at constant acceleration (from -4.5 to + 6.5 mm s-i) with Elscint (Haifa, Israel) electronic equipment[9]. A 10mCi Co57 in Rh was used as gamma ray source (14 keV). The powder with resonant Mossbauer element was mixed with dextrin to fill the 125 mm2 area of sample holder.
cell.
r
* -3
SGARLATA
I
0
I
1
6
3 Known additms
1
9
of Li/pg
Fig. 2. Typical example of the graphic elaboration of the atomic absorption analysis, by the standard addition method.
RESULT 1. The electrochemical oxide
AND DISCUSSION reaction
of lithium with iron
To investigate the electrochemical reaction of Li with FezOa, we have considered the cell Li/LiClO,-PC/FezO,.
(1)
A typical discharge curve of this cell is shown in Fig. 3. The initial part of the curve reveals a series of plateaus covering a discharge range between 0 to approximately 0.6 Faraday per mole of Fe,Os. As already remarked in a previous paper[7$ it is rather difficult to identify the reaction sequence associated with this initial series of short plateaus. Huggins and co-workers[lO] reported a thermodynamic study of the ternary Li-Fe-0 system at 4OO*C, which revealed the presence of a series of intermediate phases, such as LiFe,O,, LiFeOl and Li,FeO,. It is unlikely that the room temperature electrochemical reaction ofcell (2) can be described with the sequential formation of these phases.
Iron oxide electrodes
in lithium
organic
electrolyte
1=025
0
I
I
x
Fig. 3. Discharge
curve
I
I 4
t
6
5
Fen 03
of the Li/LiClO,-PC/Fe,O,
The presence of the initial series of plateaus and the fact that the electrochemical reaction can be eventually extended beyond 6 F/mole (see the discharge curve in Fig. 3), tend to indicate a more complex discharge mechanism than the simple progressive displacement of iron by lithium. Therefore, with the aim of clarifying this mechanism, Mijssbauer spectroscopy has been run on the cathodic pellet at various stages of discharge. In Fig. 4 is illustrated the spectrum of the cathodic mass of a cell first discharged to 5.5 F/mole of Fe,O, and then recharged back to 1 F/mole. In Table 1 the least square fitting of resonance data is also reported. The most striking feature of the spectrum of Fig. 4 is the disappearance of the magnetic hyperfine Zeeman splittings typical of the antiferromagnetic iron oxide. In contrast, a fairly sharp doublet of quadrupolar electric interaction is formed. The MGssbauer parameters of
Fig. 4. Mikbauer
~n LI,
1273
batteries
macm-2
I 3
I 2
I
rechargeable
cell. Current
Table
density:
1. Mathematical
elaborations
PI @,I
x=1 jE = 5.5
0.25 mAcm_‘.
-0.19(0.56) - 0.08 (0.61)
of resonance i.s.
q.s.
0.55 0.61
0.74 0.63
PI(%)
+ 0.55 (0.79) +0.55 (1.01)
data
p = positions of spectral lines; w = half-height width. i.s. = isomer shift with respect to sodium-nitroprusside. qs. = quadrupolar splitting. All in mm s- 1 units. x = lithium equivalents per FelOa mole. quadrupolar Lorentzian lines reveal the presence of Fe’+ ions in a slightly deformed octahedron. Therefore, the electrochemical reaction of cell (1) does not involve the reduction of iron oxide but rather the insertion of lithium ions in the third octahedral interstitial site, which is unoccupied in the hexagonal close-packing of oxygen ions of haematite. The orig-
spectrum of the cathodic mass of the cell Li/LiClO,-PC/Fe,O, 5.5 F/mole of FezOs and subsequent recharge to I F/mole.
after
discharge
to
1274
S. MO~LLI,B.SCROSATIANDF.SGARLATA
inal oxide lattice is basically retained and indeed the Xray powder diffraction pattern of LiFe,OB is similar to iron oxide pattern[7$ Owing to insertion, an interaction seems to take place between the 2s electron of lithium and the valence band of iron oxide with consequent loss of relative orientation of spins. Therefore, the electrochemical reaction (1) may be considered as a multistep insertion process and described according to the scheme: xLi + Fe,O,
= Li,Fe,O,,
(1)
where x indicates the equivalents of Li which have reacted with 1 mole of Fe,O,. Initially (ie to x = 1) the octahedral positions are fillied; the value of x ranges from 1 to 6.6. A similar structure is seen in the NaCl-type lattice of LiFeO,[l 1] where the Li and Fe ions occupy together the cation sites. The crystal-chemical behaviour of LiFez03 may be explained in analogy with that of the LiCB graphite intercalation compound. In fact, an investigation by means of inelastic scattering of synchroton radiations has revealed that both charge transfer to the n band and modification of the valence band occurs in LiCJlZ]. Finally, we may comment on the values of the line widths in the resonance data related to the cathodic mass after discharge to 5.5 F/mole (Table 1). These very high widths reveal the spreading of the electric field gradient at iron positions which indeed may be related to a substantial insertion of lithium ions in the lattice. 2. Electrochemical stability A possible concern during the operation of cell (1) may be in the effectiveness of the electrochemical process, since at the relatively low voltages at which the last part of the discharge takes place (see Fig. 3), a partial reduction of the electrolyte could also take place. Therefore, it has appeared to us of priority importance to clarify this aspect by examining the kinetic characteristics of the electrochemical process and its influence on the stability of the electrolyte. The electrochemical stability of the electrolyte used in cell (l), ie the LiClO,-PC solution, depends on various factors, among which are the nature and the morphology of the testing electrode. For instance, if the stability of the electrolyte is measured with cyclic voltammetry, one obtains a range between 2.5 and 4.0 V (us Li) when using a platinum electrode and a range between 1.0 to 4.5 V (us Li) when using a carbon electrode. The response of the two electrodes is then Table 2. Determinations of Li/LiClO,-PC/Li,Fe,O, Sample Al EO A31
Xcalculated
6.0 6.6 5.6 6.6
quite different confirming that the kinetics of the decomposition process of the electrolyte is indeed greatly influenced by the nature and the characteristic of the electrodic substrate. Since it is impossible to measure the decomposition window of the electrolyte on Fe,O, in the conditions of cell (l), one may only reasonably assume that this window might approach that observed on carbon, ie a substrate morphologically similar on a macroscopic scale to that of iron oxide. However, even under this assumption, a concomitant decomposition of the electrolyte cannot be completely excluded during the last part of the discharge of cell (1). Therefore, to finally exclude any participation of the electrolyte in the discharge process, we have determined by atomic absorption spectroscopy the lithium content in the cathodic mass at the end of constant-current discharges. The lithium contents, determined with this technique, are compared in Table 2 with those calculated coulometrically on the basis of the total charge passed through the cell. The two sets of values agree within the experimental error and this allows to finally exclude any electrolyte decomposition side effects during the discharge of cell (1). Under the working conditions of this cell the electrochemical process predominates, probably since this is kinetically favoured over the reduction of the electrolyte. 3. Reuersibility of the elecrrochemical reaction It may be reasonable to assume that the initial part (x i 1) of the discharge of ccl1 (1) involves the intercalation of Li in low energy sites until filling the octahedral interstitial sites, while the second part (1 < x < 6) involves a more extended intercalation process. Illustrative in this respect is the case of the prelithiated Li, -xV,Os compound which allows a fast and reversible intercalation of lithium up to 2 equivalents per mole[ 133. Similarly, the reversibility of the electrochemical process of the Li/Fe,O, cell here discussed, increases as the discharge proceeds, as indicated by micropolariration and cycling tests reported in a previous work[7]. In Fig. 5 is shown slow rate cyclic voltammetry of the Fe203 electrode in the LiClO,-PC electrolyte. The voltammetric response reveals a large cathodic peak around 1.5 V us Li, related to an extended, reversible reaction. This result confirms the reversibility of the electrochemical process in the Lirich region, but also indicates that the kinetics of this process may be controlled by the diffusion of lithium
Li in the cathodic mass of cell at the end of the discharge
Xdetermincd
6.1 6.5 5.4 6.8
Difference -0.1 0.1 0.2 -0.2
the
Difference y0 -2
1.5 3.5 3
= Li equivalentsper FelOj molecalculated coulometrically. %akulat& On these data, the experimental error is AX, = 17;. x&t_,,,& = Li equivalents per FeZ03 determined by atomic absorption analysis. On these data the experimental error is AX1 = 4 %,
Iron oxide electrodes
in lithium organic electrolyte
throughout the iron oxide-based electrode. It may also be emphasized that no reduction of the electrolyte is revealed by the cyclovoltammogram of Fig. 5, confirming the effectiveness of the electrochemical reaction.
rechargeable
batteries
4. L&rich iron oxide as negative batteries
12?5 QlQCtrdQ
in lithium
The results so far discussed indicate that Li,FelOJ electrodes have the capabitity of reacting with a large amount of lithium with small voltage excursions and offer reversibility especially in the I&rich range. Therefore, these electrodes appear as sutiable candidates for use as alternatives to lithium metal in rechargeable organic electrolyte batteries. To verify this assumption, we have assembled the following test cell:
60.
40.
Li6Fe103/L1C104-PC/VZ05.
0.
20 =l + ‘-t
(2)
This cell was realized by passing charge up to 6 F/mole in a cell of type (1) and then replacing the Li electrode with Vz05, ie a typical intercalation electrode in its pristine state. Using exactly the same geometry, we have also assembled the cell:
20
Li/LiC104-PC/V205. 40.
The cycling behaviour of the two cells was compared under the same conditions and the result is reported in Fig. 6. In the same figure are also shown the single polarization (vs a Li reference) of the Li,Fe,O, electrode [cell (2)] and of the Li electrode [cell (3)]. At the initial cycling stage, the polarization of the former is apparently higher than that of the latter electrode_ However, as the cycling proceeds, the performance of the lithium metal electrode deteriorates much faster than that of the lithium intercalated electrode. This is clearly shown in Fig. 7 which reports the changes in polarization of the two electrodes us the cycle number, expressed as AV (polarization of progressive cycles) over AV” (polaritation of the first cycle). It may be noticed that the AV/AV” ratio remains fairly constant
60
80
IQ0
I
I I 5
I 2 5
I 2.0 WV
I 30
I 3.5
(3)
I 40
YS Ll
Fig. 5. Cyclic voltammetry of the FelOB electrode in the LiClO,-PC electrolyte. Scan rate: 50 pV s- ‘. Electrode surt-ace: 0.07 cmz.
Cycle
CycIe
30
/
I
I
08 x
I 08
50
I
m Li,V205
4-c
+i_Y
Cycle
10
Fig. 6. Charge-discharge cycles of the Li6Fe203/V205 and of the Li/V20s cells with the LiClO,-PC electrolyte. Single polarizations (us Li reference) of the LisFe20B and of the Li anodes, are also shown. Current density: 0.3 mAcm_‘. Room temperature.
I
S.
1276
B,
MOIUILLI,
.%S!ROSATl
AND
L16Fer0,/V,0,
Polor~zot~on of L16Fe,03
F.
?bGARl.ATA
cell
Polarlzotlan
the
0
I 15
0
Number of
cycles LI IV,05
3r
cell
Polarlzotlon of tile LI electrode 2
2
2 I IO
I 5 Number
of
I I5
cycles
the
ceil . l
.
.
l
1
t
I 20
0
I
I 5
1
IO
Number
cycles
Fig. 7. Increase of the polarization
I 60
I 45 of
Polar120t Ion ot
I
i;
of
I 30
Number
3
0
cell
the
electrode
with number of cycles: AV/AV’
for about 60 cycles in the case of cell (2), so that the overall cell polarization observed in Fig. 6 is here mainly associable to the V205 cathode. On the contrary, in the case of cell (3). the AY/A Y”ratio increases to almost double after 15 cycles. This behaviour, which clearly reveals the progressive deterioration of the Li metal electrode, is in agreement with the work of Brummer and co-workersll, 141 who have shown that charge+Gcharge efficiency of a Li electrode in PC-based electrolytes, decays to SO % after 1S cycles. These results confirm the favourable cycling capabilities of the lithium composite electrodes and the feasibility of the lithium rocking chair battery concept. However, they also indicate that Li,Fe20s is by far not the ultimate choice in this respect. Other compounds, having easier preparation procedure, faster kinetics and even higher reversibility, should be characterized for reaching a final evaluation of the effective practical importance of this new type of negative electrodes in the lithium battery technology.
Acknowledgements-The authors wish to thank Dr. E. Cardarelli for the Atomic Adsorption measurements. The financial support of the Consiglio Nazionale delle Ricerche, Progetto Finalizzato Energetica 2, is also acknowledged.
ot
15
/
20
cycles
of the first cycle.
REFERENCES 1. R. D. Rauh and S. B. Brummer, Electrochim. Arta 22,75 (1977). 2. F. W. Dampier and S. B. Brummer, Elecfrochim. Acta 22, i 339 ( 1977). 3. V. R. Koch and S. B. Brummer, Elecfrochim. Acta 23, SS (1978). 4. S. B. B rummer, V. R. Koch and R. D. Rauh, Marerials Jar Advanced Buttevies (Edited by D. W. Murphy, J. Broadhead and B. C. H. Steele), p. 123. Plenum Press, New York (1980). 5. M. lazzari and B. Scrosati, J. electrochem. Sot. 127, 773 (1980).
6. B. Scrosati, B. DiPietro, M. Lazzari and B. Rivolta, Power Sotrrces 9 (Edited by J. Thompson), ~-459. Academic Press, London (19S$. _ ‘. _ 7. B. DiPietro, M. Patriarca and B. Scrosati, J. Power Sources 8. 289 (1982). 8. M. Lazzari and B. Scrosati, U.S. Patent No. 4, 464,447, Aug. 7 (1984). 9. G. De Angelis and F. Sgarlata, Per. Min. 48, 287 (1979). 10. N. A. Ciodshaall, I. D. Raistick and R. A. Huggins, Mar. Res. BulI. 15, 561 (1980). 11. D. E. Cox, Phys. Rev. 132, IS47 (1963). 12. G. Loupias, J. phys. Letz. 45, 301 (1984). 13. G. Pistoia, S. Panero, M. Tocci, R. V. Moshtev and V. Manev, Solid Srate Ionics 13, 311 (1984). 14. K. M. Abraham and S. B. Brummer, Lithium Barter& (Edited by J. P. Gabano), p. 371. Academic Press, London (1983).